QD \^i~ ^5^ CHARLES EDWARD VANCLEEF MEMORIAL LIBRARY BOUGHT WITH THEJNCOME OF A FUND 'GI^EN FO^' THE USE OF THE ITHACA DIVISION OF THE CORNELL UNIVERSITY MEDICAL | COLLEGE ; BY I MYNDERSE VANCLEEF ! I CLASS OF 1674 ' 1921 Cornell University Library QD 151.D39 An inorganic chemistry, "^ 19.2^ 012 38S 5S1 The original of tliis bool< is in tine Cornell University Library. There are no known copyright restrictions in the United States on the use of the text. http://www.archive.org/details/cu31924012388561 AN INORGANIC CHEMISTRY AN INORGANIC CHEMISTRY BY H. G. DENHAM, M.A., D.Sc, Ph.D. Professor of Chemistry, University of Capetown. LONDON EDWARD ARNOLD & CO. 1922 [All Rights Reserved] 7. y q3> / VKe4, BO^ PREFACE The majority of teachers will agree with the Author that Inorganic Chemistry presents a real difficulty to many of those beginning the study of this subject. In many cases this arises from the inability of the student to correlate and classify his observations. Hence the author has endeavoured to system- atise the subject by giving a more than usual prominence to the Periodic System. Immediately after considering the reactions of the halogen and oxygen families of elements, this great generalisaton has been developed, and in studying the remaining elements frequent recourse has been made to it. To aid the student still further in correlating and classifying chemical reactions and principles, the basic, amphoteric or acidic pro- perties of the oxides have been emphasised, since the probable behaviour of compounds derived from the oxides can thus be predicted. This method of development has necessarily resulted in defer- ring the study of the Ionic Theory until the student has mastered the fundamental laws of the science. In itself this is no dis- advantage because the student passes straight from the study of the ionic theory to the chemistry of the metals, where the theory may be most usefully applied. For the ideas concerning oxidation which have been developed in Chapter XI, the author is indebted to Professor B. D. Steele, D.Sc, F.R.S., who has found this method of dealing with the subject of great help to the student when considering the reactions and properties of the halogens. The author has endeavoured to avoid the introduction of a new law or generalisation until it has been found necessary for the interpretation of facts already described, and when the development of the subject requires a more exhaustive treat- ment, a separate chapter or section has been devoted to it. vi PREj.' ^v/o^ Subsequently references are made to the law in order to empha- sise its applications. The useful principle of Le Chatelier has been introduced at a comparatively early stage and frequent references to it will be found. As far as is possible modern technical processes have been given, and occasionally the older processes have been omitted, but as the treatise is in no way historical, the retention of the latter is not justified. The book will be found to contain sufficient for those taking the intermediate science examination of the Universities. The reduced length of the book has been secured not by the omission of essentials but by the grouping of the elements in " families," which method of treatment also brings out the graduation in properties of the elements in passing through the Periodic Table. In conclusion the Author wishes to thank his loyal colleagues, especially Professor H. Tietz, for healthy criticism, as well as for material help in the revision of the proofs. He is also in- debted to Professor F. Soddy, F.R.S., and Messrs. Longmans, Green & Co., for permission to make a short quotation from "The Chemistry of the Radio-Elements." H. G. D. Cape Town. CONTENTS CHAP. PAGE I Introduction ...... 1 II The Quantitative Aspect of Chemistry . 20 III The Nomenclatueb and Language of Chemistry 4 1 IV Oxygen .... 48 V Physical Pbopebties of Gases . . 60 VI Hydrogen ... . .-78 VII Water .... . 92 VIII Solutions, Crystals, Thermochemistry . 106 IX Moleoulab. and Atomic Weights . . .122 X Chlorine and Hydrogen Chloride . . 141 XI The Halogen Family : Fluorine, Chlorine, Bromine, Iodine and their Hydr- acids . 158 XII The Oxides and Oxy-acids of the Halogens 174 XIII Ozone : Hydrogen Peroxide . . .191 XIV Chemical Equilibrium — Law of Mass Action — Dissociation ..... 200 XV Sulphur : its Hydrides and Chlorides . 214 XVT The Oxides and Oxy-acids of Sulphur . 229 XVII The Classification of the Elements — The Periodic Law ...... 254 XVIII Nitrogen — Atmospheric Air — The Rare Gases 266 XIX The Hydrides and Halides of Nitrogen . 277 XX The Oxides and Oxy-acids of Nitrogen . 287 XXI Phosphorus ....... 305 XXII Aesenic, Antimony, Bismuth, Vanadium, Niobium, Tantalum . . . . .328 XXIII Cabbon and its Oxides. .... 346 XXIV Typical Cabbon Compounds . . . 363 Vlll CONTENTS CHAP. PAGE XXV Flame ... .... 380 XXVI Silicon ... . . . 388 XXVII Osmotic Pressure— JIolecular Weights of Dissolved Substances .... 399 XXVIII Conductivity and its Bearing upon the Ionic Theory .415 XXIX Application of the Ionic Theory to Chemical Reactions ..... .430 XXX The Metals .... . . 453 XXXI Group Ia. The Metals of the Alkalies : Sodium, Potassium, Rubidium, Caesium . 466 XXXII Group 1b. Copper, Silver, Gold . . 497 XXXIII Group 2a. Glucinum, Magnesium, Calcium, Strontium, Baeium . . . 521 XXXIV Group 2b. Zinc, Cadmium, Mercury . . 547 XXXV Group 3. Boron, Aluminium, Galliitm, Indium, Thallium, Scandium, Yttrium, Lanthanum 563 XXXVI Group 4. Germanium, Tin, Lead, Zirconium, Cerium, Thorium . . . . 580 XXXVII Group 6a. Chromium, Molybdenum, Tungsten, Uranium ....... 599 XXXVIII Group 7a. Manganese . . . .613 XXXIX Iron, Cobalt, Nickel . . . . 625 XL Ruthenium, Rhodium, Palladium, Osmium, Iridium, Platinum ..... 646 XLI Radio-Activity — The Constitution of Matter 652 Index . . . . . . . .671 International Atomic Weights . . . 684 CHAPTER I INTRODUCTION Facts and their Observation. — The aim of science is the acquisition of knowledge concerning the multipUcity of natural objects found in this luiiverse, concerning the phenomena which occur when these objects undergo change, and of the principles which govern these changes. The beginning of science dates from the earliest recognition of individual phenomena, from the earliest observation of a fact, and it is no exaggeration to say that every well estabhshed fact, every piece of experimental evidence, records a distinct advance in our scientific knowledge. Science, then, is buUt upon facts : these have been hkened to the bricks of which a stately edifice is composed. The actual observation and truthful record of what actually happens, the precise observation of the fact is often extremely difficult to make faithfully. An unconscious tendency to ignore those results which are Kkely to conflict with the experimenter's preconceived ideas is aU too common with the untrained observer. Especially difficult is the tracing of the fact to its cause. No more excellent example of this is to be found in the annals of chemistry than the task which Lavoisier (1770) set himself — to " determine whether water can be changed into earth, as was thought by the old philosophers, and still is thought by some chemists of the day." Amongst the statements quoted in favour of this contention is the fact that plants grow freely in pots, apparently deriving no nourishment other than from the water suppUed to them, yet they thrive and grow. Does not this argue a conversion of water into woody fibre ? Another fact which was often quoted in support of the conversion of water into earth was the observation that water, no matter how often it had been purified by distillation, always left an earth-like residue on evaporation in 1 B 2 AN INORGANIC CHEMISTRY glass vessels. That is the observed fact. The older chemists attributed the cause of this to a conversion of water into earth. Lavoisier doubted the accuracy of this inference. To test his view, he heated a considerable volume of water for some months in a sealed flask, the total weight, as well as the individual weights of the flask and of the water, being known. At the conclusion of his experiment, the fact which he observed was that the weight of the flask and the water contained therein remained unaltered, although he could see soUd, earthy particles floating about within the flask. The flask was emptied, and the water with the earthy residue carefuUy evaporated to dryness. The facts observed were that (a) the flask had lost weight ; (6) the weight of the earthy residue removed after evapora- tion of the water almost exactly equalled the loss in weight of the flask. This experiment enabled Lavoisier to trace the genesis of the " earth " not to a metamorphosis of the water, but to the glass walls of the vessel in which the water had been heated. Methods of Investigation. — The vast accumulation of facts which patient research has placed at the disposal of the scientist is the result partly of the observation of naturaUy occurring phenomena, e.g. tidal movements ; more often, however, they are the result of the observation of experiments expressly designed for some particular purpose. The diSerence between the observation of a naturally occurring phenomenon over which we are unable to exercise control and of an experiment which we ourselves have planned is weU expressed by W. S. Jevons : " Where we merely note and record the phenomena which occur around us in the ordinary course of nature, we are said to observe ; when we change the course of nature by the intervention of our will and muscular powers, and thus produce unusual combina- tions and conditions of phenomena, we are said to experiment. Our experiment differs from a mere observation in the fact that we more or less influence the character of the events which we observe." It will be noticed that the essential point distinguishing the carryuig out of an experiment from the mere observation of some natural phenomenon is that we are able to control the con- ditions under which our experiment is carried out. Every factor, INTRODUCTION 3 such as temperature, pressure, concentration (amount dissolved per litre), which can conceivably exert an influence upon the experiment which we are observing must be brought under control. In our search for new facts we must keep under abso- lute control all such factors ; then and then only are we hkely to be able to carry out our experiments under reproducible conditions, and to obtain results which are in themselves repro- ducible and therefore worthy of being accepted as facts. A second point to be observed in an extended series of experiments planned with a definite object in view, e.g. to test an hypothesis which has been put forward, is that never more than one factor should be varied at the same time. If, for example, the desired reaction does not take place at the temperature first chosen, it should be repeated at a higher or lower temperature, but no other controllable factor such as concentration should be varied. So much, then, for the manner in which an investigator should work. As far as the actual choice of a suitable subject for experiment is concerned, two methods stand out. The first of these is the " empirical," that used by such an investigator as Joseph Priestley, the discoverer of oxygen. This discovery was the result of his trying the effect of heat upon a large number of substances chosen at random. Such a method is tedious, laborious, and rarely leads to discoveries of fundamental import- ance. The method of the true scientist is di£ferent. He moves onward by steady steps, building his bridges as he goes. Every advance he makes is the result of logical and clear reasoning. No finer example of this type of investigator is to be found in the annals of science than Michael Faraday, most patient of investigators, yet one of the most briUiant in achievement. Midway between these two groups of investigators — the empi- ricist and the true scientist — Hes the group that works by intuition. As examples of such men, Davy and Ramsay stand out. The most striking achievements of these investigators appear to have been the result of a hghtning-like intuition or inspiration. They do not know what has led them to their discovery, the chain of reasoning is missing and must be con- structed after the discovery is made for the benefit of those who come after. The Classification of Facts. — It is the work of the scientist to classify and coordinate the facts, the records of observation. 4 AN INORGANIC CHEMISTRY From the careful study of many apparently diverse observations he is able to unravel uniformities, to pick out the elements of sameness among apparently different phenomena, and in this way he is able to classify the records of observations, and out of the whole there emerges an organised body of knowledge — a science. Science, then, is an organised body of knowledge. But owing to the omniscient range of science, it has become necessary to subdivide science into various branches, and so we have (a) abstract sciences — logic, mathematics, astronomy ; (b) concrete sciences such as zoology, botany, geology ; (c) abstract-concrete sciences — physics, chemistry. The actual Une of demarcation between kindred sciences such as physics and chemistry is very difficult to draw, indeed it is becoming increasingly difficult to determine whether such branches as electro-chemistry are to be assigned to the one science or the other. Chemistry deals with the study of the different kinds of matter in the universe, of the ultimate composition of matter and of the phenomena which occur when the different kinds of matter react one with another. The study of the properties and the ultimate composition of different kinds of matter is dealt with in the branch known as Descriptive Chemistry, and embraces that portion of chemistry which is essentially concrete, whereas the consideration of the mechanism of reaction of one substance with another leads us into the more abstract portions of the science, into the branch known as Theoretical Chemistry. The Formulation of Hypotheses, Theories, Laws. After the facts relating to any particular phenomenon have been made the subject of observation, a guess must be made at the possible cause of the effects which have been observed. Such a guess at the truth constitutes what is known as an hypothesis. However probable an hypothesis may seem, there is necessarily involved in it an element of speculation. It is to determine whether such speculation is justified, that experiments must be devised to test the hypothesis. The formulation of the hypothesis does good work in suggesting further experunents which should be undertaken in order to determine whether the hypothesis is to be retained or rejected. Even if rejected, the INTRODUCTION 5 hypothesis in its death agony has done a useful service to science in suggesting a series of experiments which should be carried through. The path of science throughout the ages is httered with discarded hypotheses which have been trampled underfoot. Occasionally it happens that more than one hypothesis may be put forward to account for the observed facts ; in such a case a crucial experiment must be devised to determine which one of the rivals must be sacrificed. In the earth- water experiment of Lavoisier described above, the plausible hypothesis that the earth had its origin in the flame gases which penetrated the glass walls of the flask and were condensed as earth, was destroyed by the crucial experiment of weighing the flask and its contents before and after the experiment. No change of weight occurred and this hypothesis must therefore be ruled out. An hypothesis can only be considered established when it has been rightly tested by experiment and has survived the ordeal. After an hypothesis has been thoroughly tested by experi- ments devised to determine whether the hypothesis does really explain what it pmports to explain ; when it has taken its established position ; when it has proved its utility and adapt- abihty in correlating the results of experiment ; then, and then only, is the hypothesis entitled to rank as a Theory. A theory is a well-established hypothesis. Out of an hypothesis there may emerge a theory ; the theory may develop into a law. Facts of a like kind are carefully considered to see whether any general statement can be made which will cover some featiu-e common to all the facts under consideration. The formulation of the generalised statement summing up the features common to aU the phenomena is an example of what is known as Indvctive reasoning. We are led from the special to the general. This generalised statement or hypothesis is then tested in the inverse way by being apphed to the explanation of new phenomena. This is deduction, which is based upon the principle that what is true of the general or the many, must be true of other individual cases. Should the hypothesis stand the test appHed by deductive processes, it becomes a theory ; and when the theory becomes firmly estab- lished, it passes into a recognised Law of Nature. The Phlogiston Hypothesis and the Calcination of Metals. —No more striking example of the manner in which a falsQ 6 AN INORGANIC CHEMISTRY hypothesis has been put forward to explain certain experimental results, and finally uprooted by means of a carefully designed crucial experiment, is to be met in the study of chemistry than that afforded by the Phlogiston Hypothesis. As early as the eighth century the Arabian chemist Geber made the discovery that, when metaUic lead was calcined (heated) in air, the resulting calx proved to be heavier than the metal from which it was formed. AU through the Middle Ages it was felt that, when substances bum with the production of flame, the same phenomenon was involved as when metals are converted into their calx. The first clear conception of the phenomenon of combustion was put forward by Becher (1635-1682) and Stahl (1660-1734). Combustible substances were assumed to consist of a certain inflammable ^principle or phlogiston combined with the calx (the product of the combustion). When combustion took place, the phlogiston escaped, the calx remained, and in order to regenerate the combustible matter, the calx must be heated with a substance rich in phlogiston, e.g. carbon, sulphur. Combustible substance — >- phlogiston + calx Calx + phlogiston — >- combustible substance (supplied hy Carbon). This hypothesis brought into line the various experimental results given by the combustion of metals, phosphorus, etc., and explained too the regeneration of the metals by heating the calx with carbon. But one point it did not explain. It failed to make clear why the calx was heavier than the metal. Stahl and his supporters then advanced the view that phlogiston actually possessed the property of levity, which is equivalent to saying that phlogiston, in contradistinction to all other substances, possessed a negative weight, so that, when it escaped dur- ing combustion, the resulting calx became heavier. This hypothesis was accepted widely as a sufficient explanation of the phenomenon of combustion. Rey (1630) showed the weak- ness of the hypothesis in his experiments upon the calcination of various metals in air, as a result of which he was led to the conclusion that this increase in weight came from the air, but it was left to Lavoisier (1774) to cast the phlogiston hypothesis into the scrap heap of science and to replace it by the Theory of Combustion which is held at the present time. INTRODUCTION Lavoisier's Experiments upon the Calcination of Metals. — Lavoisier weighed a retort full of air, and determined the weight of a piece of tin. This was placed in the retort, the tip of which was then sealed off. The retort was heated in a char- coal fire. The calx was formed within the retort, and when no further calx was formed, the apparatus was weighed. No change in weight had occurred in the closed vessel, but when the flask was opened, air rushed in. Prom his various experiments he concluded : — 1. In a given volume of air a certain fixed quantity of tin can be calcined. 2. This quantity is greater in the large retort than in the small one. 3. The sealed retort, weighed before and after the experiment, showed no change in weight, which proves that the increase in the weight of the metal, due to its conversion into calx, does not arise from the penetration of " fire matter " from the flame into the vessel. 4. In all calcinations the increase in weight of the metal is practically equal to the weight of the air absorbed in the calx. StiU further support of his conclusions was given by his experiments upon the calcin- ation of mercury (Fig 1). He placed a weighed quantity of mercury in a retort with an S-shaped neck which dipped under a beUjar, floating upon mercury. Thus there was connection between the air of the retort and of the belljar. The volume of the air in the beUjar was marked and then the retort was steadily heated. After two days, red specks were seen floating on the surface of the mercury. These increased in size and number and at the same time the volume of the air in the belljar diminished. After twelve days, when no further change either in the volume of the air or in the quantity of the " red precipitate " could be detected, the experiment was stopped, and the volume of " air " absorbed by the mercury measured. It was about 8 cu. inches. The air which was left in the retort unabsorbed by the mercury, extinguished a burning Fig. 1. 8 AN INORGANIC CHEMISTRY candle, and suffocated a mouse. This residual air Lavoisier named azote (Gk. without life), known now as nitrogen. He gathered the red precipitate and collected the gas evolved by it when it was heated. In this way he obtained about 8 cu. inches of a gas which had the property of causing a candle to bum brilhantly, and of sustaining hfe in a mouse. This gas he called vital air, now known as oxygen (Gk. acid-producer). In this way Lavoisier established the composite nature of air, that it consists of two gases, oxygen and nitrogen, of which oxygen alone is able to support life and to combine ^Yiih metals during calcination. Moreover by mixing nitrogen and oxygen in the proportions in ^hich he had obtained them from air, he was able to form a gas which possessed all the properties of atmospheric air. In reporting his results, he passed the following scathing condemnation of the Phlogiston hypothesis : " Sometimes this principle has weight, and sometimes it has not ; sometimes it is free fire and sometimes it is fire combined with the earthy ele- ment ; sometimes it passes through the pores of vessels, some- times these are impervious to it ; it exjilains transparency and opacity, colours and their absence ; it is a veritable Proteus changing in form at each instant." Lavoisier then proceeded to put forward his own tentative hypothesis to explain the facts which he had established, from which has, emerged the present day theory of combustion {q.v.) and of oxidation. Specific and Arbitrary Properties. — If one chooses and compares a number of pieces of iron, we find that, although we are unable to alter certain properties such as colour, weight, which are common to aU such samples of iron, other properties are capable of alteration. For instance, the iron may be heated or chilled, i.e. its temperature may be altered, or its shape and volume may be altered. Such pixiperties which are capable of alteration by the hand of man are known as arbitrary properties. The study of these properties l)elongs to the domain of Physics. Properties which are common to all samples of a particular sub- stance and which man is unable to alter (e.g. melting point, boiUng point, colour, odour, solubility, etc.), are known as specific pro- perties or properties common to a species. Chemistry is con- cerned with the study of the specific properties of substances. The specific properties are constant for aU specimens of a par- ticular substance, but the arbitrary properties may vary from INTRODUCTION specimen to specimen. The study of the specific properties of substances is of the ixtmost importance, for substances are recognised as being alike chemically when they possess similar specific properties. It is not necessary to compare more than a few specific properties of the different specimens, for it has been found by experience that, if the specimens have a few specific properties in common, they wiU agree in all such properties. Besides being classified according to their ultimate chemical composition, i.e. according to their specific properties, a rough classification of substances may be made by classing them according to their homogeneity. Separation of Mixtures. — Substances which appear hetero- geneous to the naked eye or to the microscope, can at once be classed as mixtures. Such a mixture, lacking as it is in uni- formity, can be resolved into its component parts by taking advantage of the difference in the specific properties of the substances of which the mixture is made. Occasionally the chemist takes advantage of the difference in density of the con- stituents ; for example, a mixture of finely divided clay and iron filings can be separated by elutriation, i.e. the lighter clay particles are washed away from the denser iron fihngs. Many ores which are used in metallurgical processes are ground and treated with water in order to get rid of a large percentage of the worthless earthy dross. The valuable minerals contained in the ores are thereby concen- trated so that their ultimate separation becomes economically possible. Again one may revert to the difference in the melting point to effect the desired separa- tion ; a mixture of paraffin and brass fihngs may be separated by heating the mixture to a suffi- ciently high temperature to melt the paraffin, but not the fihngs, after which the separation can be completed by the use of a fine sieve. This is essentially the same as the process of filtration, in which one of the constituents, water, is liquid at ordinary temperature, while the sieve is replaced by a filter paper (Fig. 2). Fig. 2. 10 AN INORGANIC CHEMISTRY In other cases, the separation is made by expelling one con- stituent by converting it into a vapour. This is based upon the difference in volatility of the different constituents. Iodine and sand may be completely separated by gently heating the mixture, when the iodine volatilises and may be condensed on a cold surface. The non-volatile sand remains. Again, many substances are magnetic to a greater or less extent, others lack this property entirely. A mixture of iron filings and sand can be at once separated by taking advantage of the mag- netic properties of the iron. During recent years, increasing use has been made of this property for concentrating the valuable minerals present in impure ores. The powdered ore is crushed and some of the valuable constituents are then removed by the use of a suitable electro-magnet. Another specific property which is often called into requisition by the chemist when he is sorting out a mixture is the difference in solubility (i.e. the amount dissolved in a given quantity of the solvent). Water, especially if hot, dissolves sugar freely ; but neither hot nor cold water is able to dissolve sand or iron. The substance potassium permanganate dissolves in water, forming a deep purple solution, but it is quite unable to dissolve in ether. The iodide of mercury dissolves slightly in alcohol, though it does not dissolve in water ; cobalt chloride dissolves in water, forming a pink solution, whUst alcohol gives a deep blue solution. Nearly all substances become more soluble as the temperature rises. From these examples it is seen that the solvent necessary to bring any particular substance into a state of solution must be carefully sought out. Because water dissolves potassium per- manganate and sugar, we are not entitled to conclude that it will also dissolve mercury iodide. The solubility of a substance is therefore conditioned not only by the temperature but also by the solvent chosen. A great deal of use is made in industry of the fact that each substance possesses its own distinct solubility in a particular solvent at a chosen temperature. Gold, for example, can be entirely removed from earthy matter by dissolving it in a dilute aqueous solution of potassium cyanide, from which it is after- wards recovered. A mixture of alum and sand can also be separated into its constituents, because one of them, the alum, is freely soluble in water whilst the other is not. The mixture is treated with hot water, allowed to settle, the clear Kquid INTRODUCTION 11 decanted through a filter paper, and the residue several times extracted with water. In this way the sand may be collected upon the filter paper and the alum can be recovered from the ■filtrate. The separation of a homogeneous solution of alum in water wherein every particle of fluid has the same properties, is generally brought about by expelling the more volatile constituent in the form of vapour. The water is slowly driven off by heating, and the time soon arrives when the remaining solvent can no longer retain in solution the whole of the dissolved substance, and a deposit of fine crystals of alum separates. Crystallisation has set in. If the solution consists of one liquid dissolved in another, the separation is made by the process of distillation. Each of the substances of which the solution is formed has its own character- istic boiling point. If the solution is boiled in the apparatus shown in Fig. 3, the vapour which escapes into the water-cooled Fig. 3. condenser consists in general almost entirely of the substance which boils at the lower temperature. The more volatile liquid is collected in the receiver, the liquid of higher boiling point remaining in the distillation flask. The separation of a homogeneous mixture, whether of liquid in liquid, or of solid in liquid, can only be made through the medium of another phase. A phase is a discrete part of a heterogeneous system, which is capable of mechanical separation. 12 AN INORGANIC CHEMISTRY During the distillation the vapour phase is formed, and this vapour phase has a different composition from the residual liquid. The same remarks apply to the crystallization of the solution of alum. The composition of the gaseous phase expelled during the evaporation is quite different from the composition of the residual solution. The latter slowly changes its composition until yet another phase, the pure solid alum, appears. Chemical Change. — If a piece of magnesium wire is heated in the air, there is a bhncling flash and nothing remains but a fine white powder. The specific properties of this substance are entirely different from the specific properties of the magnesium wire. The change which has occurred is very deep seated, quite distinct from the physical change which a solid undergoes when it melts, for there it is only necessary to cool the liquid slightly in order to convert the hquid back into the soUd. The formation of the white powder from the magnesium wire is a good example of what is known as a chemical change inrolring the formation of a chemical compound. A chemical change leads to the formation of a new compound ichich possesses specific properties different from those of the substance from which it is made. To take another example. If finely divided copper dust is mixed a\ ith floA\'ers of sulphur, the heterogeneity of the mixture can be easOy detected by means of a pocket lens, also by the fact that we can dissolve the sulphur away from the copper by taking advantage of its great solubUity in a hquid known as carbon disulphide. But if a mixture of equal parts by weight of copper and sulphur is heated in a test tube, the mass suddenly becomes glowing hot and there is left a dark, hard mass. If this mass is rapidly washed with carbon disulphide in order to remove any superficial particles of sulphur, dried and then powdered, no sign of heterogeneity is revealed, even under a powerful microscope, nor can sulphur any longer be dissolved out by means of carbon disulphide. A ne^v compound, known as copper sulphide, ^\hich the student ^^111 afterwards find does contain sulphur, has come into existence, endowed with its own specific properties of hardness, colour, solubility, etc. In the classic experiments of Lavoisier upon the calcination of mercury in air we have yet another example of chemical change. On heating mercury in air, he found that it became covered A\ith reddisli scales through the combination of the INTRODUCTION 13 mercury with the oxygen of the &ir. When this reddish powder (mercury oxide) was strongly heated he recovered the oxygen again, and also obtained the mercury in the form of a bright metallic mirror on the cooler parts of his apparatus. If the yellow iodide of mercury is heated there is a sudden change in colour to scarlet. The reverse change takes place on cooling. Is this change analogous to the purely physical change which a piece of gold wire undergoes on being heated in a flame, or is it of a more deep-seated nature, involving chemical change 1 No purely chemical test enables the chemist to detect any difference between these two modifications of mercury iodide ; just as, from the chemical point of view, there is no difference between ice and water, so here there is no distinction between these two varieties of mercury iodide. Energy Changes during Chemical Reaction. — If a strip of zinc is placed in dilute sulphuric acid, a vigorous reaction takes place. There is a brisk evolution of gas, the zinc strip is rapidly eaten away and passes into solution, and if a ther- mometer is placed in the solution, there is found to be a marked Platinum Foil Galvanometer Sulphuric Acid Fig. 4. rise in the temperature of the solution. If the solution is after- wards evaporated, white crystals of zinc sulphate can be obtained. The solution of zinc in sulphuric acid can, however, be effected in a somewhat different manner. In a beaker put some dilute sulphuric acid, and also a strip of platinum and of zinc, both of which are connected to the terminals of a galvanometer by means of a copper wire (Fig. 4). 14 AN INORGANIC CHEMISTRY As soon as the zinc and platinuni strip are immersed, the needle of the galvanometer indicates the passage of an electric current. The zinc is slowly corroded, passing into solution. The bubbles of gas again make their appearance, this time, however, on the surface of the platinum. During the whole time the zinc is passing into solution a steady current flows through the circuit, and there is no appreciable rise in the temperature of the solution. TUthough the gas was hberated at the surface of the platinum, no change whatsoever in this metal can be detected. Crystals of zinc sulphate may again be isolated from the acid solution. By a suitable arrangement we have managed to bring about the formation of zinc sulphate with the production of an electric current iastead of the evolution of heat. Now place the solution of zinc sulphate in the apparatus de- picted in Pig. 5. Two platinum strips, connected to the terminals Fig. 5. of a battery, with a galvanometer in circuit, are placed in the hquid. A current of electricity immediately flows through the solution, bubbles of gas escape from one platinum strip, whilst bright metaUic crystals of zinc begin to grow from the other strip. MetaUic zinc has been formed by the decomposition of zinc sulphate, brought about by the passage of the electric current through the solution. These experiments are very important for they bring home the fact that electricity may be consumed or produced during a chemical reaction. Furthermore, if a little mercury and iodine are placed in a INTRODUCTION 15 mortar, moistened with alcohol, and then vigorously rubbed with a pestle, red crystals of mercuric iodide are formed. In other cases, chemical decomposition is brought about at the expense of mechanical energy. A speck of nitrogen iodide {q.v.) if struck with a hammer will explode with extraordinary violence. Mechanical energy may, therefore, bring about chemical reactions, whilst the well-known effects due to the explosion of various compounds speak sufficiently as to the mechanical energy which is made available in certain chemical decompositions. Chemical Energy and its Relation to other Forms of Energy. — If an electric current is passed through the filament contained in an electric light bulb, the filament immediately be- comes a source of Hght. The passage of the electric current heats the filament to incandescence, and we have as a result a conversion of electricity into hght. In the same way we notice that, if a thin platinum wire is immersed in a beaker of water and a current of electricity passed through it, the temperature of the water rises. Electricity has been converted into heat. The student of Physics is famUiar with the thermo-ceU which converts heat into electricity and with the selenium light cell which brings about a conversion of hght into electricity, whilst the change of mechanical energy into heat is known to aU. In order to correlate our ideas concerning the inter-relations of heat, hght, electricity, etc., the hypothesis has been formulated that heat, light, electricity and mechanical energy are all forms of what is known as energy. No better definition of the concept energy exists than that formulated by Ostwald, " Energy is work, and every other thing which can arise from work or be converted into work." Careful experiments on the transformation of one form of energy uito another led Mayer (1842) and others to conclude that, whenever one form of energy is transformed into another form of energy, e.g. electricity into heat, no loss of energy whatsoever occurs. From a given amount of mechanical work it is possible to obtain certain definite amounts of electrical work, hght, heat, etc., and when these are re-converted into mechanical energy, exactly equivalent amounts of this form of energy are obtained. This quantitative relationship concerning the various forms of energy is often formulated in the Law of Conservation of Energy ; " The amount of energy contained within a sealed system is constant, and unalterable." 16 AN INORGANIC CHEMISTRY At first sight, this law appears to clash with the results of experiments which have already been discussed. Whence did the heat energy arise which became manifest when copper and sulphur were heated together ? What Mas the source of energy that was responsible for the creation of electrical energy when zinc dissolved in sulphuric acid ? In the latter case wo began with zinc and sulphuric acid. During the solution of the zinc an electric current was generated, but during the whole of this time, zinc was passing into solution, forming zinc sulphate and liberating a gas. This may be expressed in the equation. Zinc and sulphuric acid — ^ zinc sulphate and gas + electric energy. In the chemical reaction between copper and sulphur we have. Copper and sulphur — >- copper sulphide + beat energy. If the assumption is made that unit quantities of all forms of ^natter have associated icith them definite amounts of chemical energy, these results at once fall within the province of the Law of Conservation of Energy. Whenever a chemical reaction occurs, a change in the energy content of the system takes place. This is seen in the generation or absorption of heat, light, electricity or other forms of energy. The final product of the reaction contains a different store of energy from what the reacting substances contained. As a rule, chemical energy is converted into heat energy or vice versa. Reactions which are accompanied by the evolution of heat are known as Exothermal reactions ; those in which heat is absorbed are classed as Endothermal reactions. At the same time, although chemical reactions are attended by energy changes, it must not be concluded that all energy changes are necessarily associated Avith chemical reactions. Ice changes its state, passing into water, as soon as the requisite heat is supplied ; iodine can be vaporised by the application of heat. These changes are physical, not chemical. Elements and Compounds. — It was Lavoisier who first expressed in precise language the fundamental distinction between " element " and " compound." His experiments upon the nature of the atmosphere, upon the decomposition of mercury oxide into two substances, mercury and oxygen, which resisted all efforts directed toward breaking them down iNTRODlJCTiON 1? into yet simpler substances, led him to put forward the following views. "If we apply the term elements or principles to bodies to express our idea of the last point which analysis is capable of reaching, we must admit, as elements, all substances into which we are able to reduce bodies by decomposition. Not that we are entitled to affirm that these substances which we consider as simple, may not themselves be compounded of two, or even of a greater number of more simple principles ; but since these principles cannot be separated, or rather, since we have not hitherto discovered the means of separating them, they are, with regard to us, as simple substances, and we ought never to suppose them compounded untU experiment and observation have proved them to be so." His failure to break down lime and magnesia into their elements led him to express the view that " we are certainly entitled to consider them simple substances untU, by new dis- coveries, their constituent elements have been ascertained." Some time afterwards, Davy proved the composite nature of these substances by means of an electric current. It is surprising, in view of the statement above, that Lavoisier maintained that " oxy-muriatic gas," now known as chlorine, was a compound gas containing oxygen, for no experiments ever succeeded in proving the presence of oxygen in this gas. Davy maintained that, until the presence of oxygen in oxy-muriatic gas had been definitely estabUshed, this substance must be looked upon as an element. His view ultimately prevailed. As a working definition of an element the following will serve : " An element is a substance which has never been broken down into other forms of matter by the aid of energy which is under the control of man." Note. — Recent advances in certain branches of chemistry out- lined in the last chapter, necessitate a widening of the definition of an element often quoted : " An element is a substance which has never been shown to contain more than one kind of matter." The actual number of elements is comparatively smaU, approxi- mately ninety, but the number of compounds derived from these elements is estimated to be well over, 50,000. Most of these are formed by the combination of two or three elements together, usually not more than three different elements being found in a chemical compound. 18 AN INORGANIC CHEMISTRY Modes of Chemical Reaction. — Several experiments involving chemical change have been described. Two of them show certain features in common — the combination of copper and sulphur to form copper sulphide, and the combination of mercury and oxygen to form mercury oxide. In both ca^es a compound with its o\mi specific properties, ^^'ith its own allotted store of chemical energy per unit of weight, was formed from elements, these elements likewise possessing certaia specific properties and having associated with them a definite quantity of chemical energy per unit of -weight. These chemical reactions afford an example of direct combination or synthesis. As opposed to this type of chemical reaction, consider the breaking down of the compound mercury oxide into its elements, mercury and oxygen. This is an example of decomposition. In all cases of decomposition it is not necessary for ele- ments to be formed. If copper carbonate (green) is put in a test-tube (Fig. 6) with the exit tube dipping below the surface of some lime water, and the test tube carefully heated, a gas is given off which forms a white precipitate in the lime water, 'while the colour of the soUd in the test tube changes from green to black. The actual change which occurs may be expressed thus : — Copper carbonate — > copper oxide + gas. The black powder Avhich remains, we shall see later, contains oxygen as M-ell as copper ; the gas liberated is also compound (carbon dioxide). The action of zinc in displacing a gas (hydrogen) from sul- phuric acid is an example of chemical displacement or substitu- tion. Zinc and sulphuric acid— > zinc sulphate + hydrogen. A piece of iron, plunged into a solution of copper sulphate, becomes covered with a reddish deposit of copper ; if sufficient Fig. 0. INTRODUCTION 19 iron filings are used, the whole of the copper will be displaced by the iron. This is also an example of substitution — Copper sulphate and iron — >-iron sulphate + copper. A fourth type of chemical reaction is illustrated by the action of a few drops of a solution of silver nitrate on a solution of sodium chloride. A white cloud, consisting of sUver chloride, is precipitated. This is an example of double decomposition or mutual interchange. Silver nitrate + sodium chloride— >- silver chloride + sodium nitrate. This type of reaction is extremely common. The absence of a precipitate is no evidence that such an interchange has not occurred ; it merely proves that no insoluble substance can be found by such an interchange. It will be seen later that mutual interchange when solutions are mixed, is the rule rather than the exception, irrespective of whether a precipitate forms or not. Recapitulation The study of chemistry deals with the properties of the different kinds of matter, of the ultimate composition of matter, and of the phenomena which occur when the different kinds of matter react with one another. Each kind of matter has associated with it not only definite specific properties, but also a definite amount of chemical energy per unit of weight. When chemical action takes place, with the production of a new substance or substances, the amount of chemical energy stored in the system changes. If a decrease in the chemical energy has occurred heat or some other form of energy wQl be set free during the chemical reaction. This change in the energy content of the system is a definite con- comitant of chemical change, but it must not be overlooked Ihat when a physical change occurs, e.g. the passing of a liquid into a vapour, there is also a change in the energy content of the system, hkewise in the specific properties of the substance. The line of demarcation between physical and chemical change is, therefore, not easy to draw. Further consideration of this will be given during the discussion of the quantitative aspect of chemical change. CHAPTER II THE QUANTITATIVE ASPECT OF CHEMISTRY Up to the present the phenomenon of chemical reaction has been considered purely from the quahtative or descriptive standpoint ; before a quantitative investigation can be attempted, certain units of measurement require defining. As unit of weight in chemical work, the gram is chosen. This is approximately the weight of 1 c.c. of water at 4° C* Multiples and sub-multiples of this imit are used. The unit of volume is 1 c.c. or the volume occupied by approximately 1 gm. of water at 4° C* For larger volumes, the Utre (1000 c.c.) is often used as the unit. A litre contains about 61 cu. inches and is about 1-8 pints. The unit of length is the centimetre, the unit of time the second. Measurements of temperature are almost invariably referred to the Centigrade scale (for remarks concerning the Absolute scale, see Chapter 5). The freezing point of water on the Centigrade scale is 0°, the boiHng point under atmospheric pressure 100°. The distance between these two limits is divided into 100 equal parts, each of these repre- senting 1 degree (1° C). The unit of energy is the erg. In absolute units, 981 ergs represent the work derived from 1 gm. falling from a height of 1 cm. On account of the inconvenient smallness of this unit, the Joule ( = 10,000,000 ergs) is often used. The unit of heat is the calorie, which is defined as the amoimt of heat capable of raising one gram of water 1° C. in temperature. The Lawf of Conservation of Mass. — In 1808 Thomas Dalton wrote : " We might as weU attempt to introduce a new planet into the solar system, or to aimihilate one already in existence as to create or destroy one particle of hydrogen." The indestructibility of matter appears, therefore, to have been * As the result of work carried out by the International Bureau of Weights and Measures in Sevres, near Paris, it has been found that the weight of 1 c.c. of water at 4° is not exactly equal to 1 gm. but to 0-999973 gm. 20 QUANTITATIVE ASPECT OP CHEMISTRY 21 firmly rooted in the minds of chemists even so far back as the beginning of the nineteenth century. That this should be so is due to the work of the great French chemist Lavoisier, the apostle of the balance. It was he, above aU others, who made chemistry an exact science by his many investigations in which chemical phenomena were watched not only from the quahtative but also from the quantitative aspect. Although many tacitly accepted the great truth that in a chemical reaction there is neither creation nor destruction of matter, it was not till Lavoisier's classic experiments on com- bustion that a convincing proof was furnished of the great law known as the Law of Conservation of Mass. In summing up his experimental results (1785) Lavoisier wrote : " Nothing is created, either in the operations of art or in those of nature, and it may be considered as a general principle that in every operation there exists an equal quantity of matter before and after the operation ; that the quantity of the constituent is the same, and what happens is only changes, modifications. It is on this principle that is founded all the art of performing chemical experiments." In 1860-65 Stas carried out some extraordinarily accurate experimental syntheses, amongst them the formation of silver iodide from known weights of silver and of iodine. Such was the accuracy with which this masterly analyst worked that the weight of the silver iodide actually collected and weighed differed by no more than 0-002 per cent, from the sum of the weights of the sUver and of the iodine. The matter was still further investigated by Landolt. In an elaborate series of experiments(1893- 1909) this chemist was able to show that the weight of the substance formed did not differ by more than his experimental error from the sum of the weights of the reacting sub- stances. In his work he used vessels p _ shaped as in Fig. 7. Of the two reacting substances one was introduced into one leg of the tube, one into the other. The vessel was then sealed off, the whole carefully weighed, and the reagents brought into reaction by tilting the tube. After the temperature of the contents had readjusted itself the final weighing was made. 22 AN INORGANIC CHEMISTRY Landolt's results show that the average difference found in weighing one of these vessels before and after reaction did not differ measurably from two consecutive weighings carried out without reaction having taken place. The truth of the great generalisation known as the Law of Conservation of Mass stands, therefore, on a very firm experi- mental basis. There is, however, another feature about the manner in which this law is obeyed. The more rigid the experi- mental control imder which the experiments testing this law have been carried out, the more exactly does the law hold. This is our first example of what is known as an exact law of science. The very slight differences between theory and experiment are always irregular, sometimes positive, sometimes negative. These errors become smaller and smaller as the experimental conditions are refined and the errors are purely experimental, and do not arise from any conflict between experiment and the law. The sUght differences observed are apparent, not real. On the other hand, we shall soon meet numerous examples where the law does not exactly explain the phenomena which it purports to explain, e.g. the Gas Laws (q.v.). The methods of experimentation may be altered, they may be made more exact, but the differences are always there, indicating a real conflict between the generalised statement and experiment. In such cases the differences observed between experiment and theory are generally always positive or always negative. In such cases the theory ranks as an approximate law of science (examples, Avogadro's Law, the Gas Laws, the Law of Combination by Volumes, q.v.). A useful and instructive verification of the statement that matter is not destroyed even when a candle burns, is given by the experiment shown in Fig. 8. A candle is fixed upon a perforated base fitted into the bottom of a lamp chimney, so arranged that there is a free passage of air through the chimney. Half way up the chimney is placed a piece of wire gauze, upon which are placed a few sticks of sodium hydroxide. The apparatus is fixed to one arm of the balance and sufficient weights are put into the other scale pan to give an exact counterpoise. The candle is then removed, ht and replaced. Although the candle steadily burns away, it is noticed that the chmmey and its contents gradually become heavier. The gases rising from the flame pass upward and are absorbed by the QUANTITATIVE ASPECT OF CHEMISTRY 23 sodium hydroxide. The increase in weight is the weight of the oxygen which has entered the apparatus in order to aid the combustion of the candle. If a piece of sulphur is burnt in an airtight flask, no change of weight occurs. This can be shown in the apparatus illustrated in Fig. 9. A piece of sulphur is laid in the porcelain boat and in im- mediate contact with the sulphur is a piece of platinum wire, the terminals of which can be connected to a battery. In the bottom of the flask there is placed a piece of sodium hydroxide to absorb Fig. 9. the products of combustion. After the whole has been exactly weighed, a current from the battery is sent through the wire in order to inflame the sulphur. When the oxygen within the flask has been used up, the sulphur ceases to burn, and the apparatus is allowed to cool. No change of weight will be detected if the experiment has been properly carried out. The Laws of Chemical Combination The Law of Constant Coinposition. — In 1799 there begpn one of those bitter controversies which have frequently occurred in the annals of science. Berthollet, one of the great French chemists of his day, advanced the view that the chemical com- 24 AN INORGANIC CHEMISTRY position of a substance could be influenced at will by altering the quantities of the reacting substances. This hypothesis at once called forth the most strenuous opposition from Proust. This chemist prepared a basic carbonate of copper and showed that not only did the properties of this substance agree with the properties of the naturally occurring carbonate of copper, but there was a similar agreement in the chemical composition of the two specimens. He also obtained the same weight of copper oxide, carbon dioxide and water from equal weights of the natural and of the artificial carbonate. Hence, Proust wrote : " We must recognise an invisible hand which holds the balance in the formation of compounds ; we must conclude that nature acts not differently in the depths of the earth than on its surface and through the agency of man. . . . Let us recognise, therefore, that the properties of true compounds are as invariable as is the ratio of their constituents. Between pole and pole they are foimd identical in these two respects ; their appearance may vary owing to the manner of aggregation, but their properties, never. The cinnabar (mercury sulphide) of Japan is constituted according to the same ratio as that of Almaden (Spain)." All along the line Proust proved his point, and the Law of Constant Composition or the Law of Fixed Proportions was firmly established on a secure experimental foundation. Every chemical compound, irrespective of its method of pre- paration, always contains the same elements combined together in the same proportions by weight. The vaUdity of this law was made the subject of a most extensive series of experiments at the hands of Stas (1860-65). He prepared silver chloride, etc., by four independent methods, yet he found that the chemical composition of each of these samples agreed within 0-004 per cent. Since the day of Stas many an exhaustive test has been made of the rigidity with which this law is obeyed, and it may be stated that not the sHghtest deviation from this law has been detected. With every improve- ment in the skill and technique of the experimenter, with every improvement in the apparatus, the agreement between the results of experiment and the demands of this law has become more pronounced. The Law of Constant Composition may therefore he classed as one of the exact laws of nature. The Law of Multiple Proportions.— Although Proust, the QUANTITATIVE ASPECT OF CHEMISTRY 25 champion of the Law of Constant Composition, came across various examples of substances containing the same elements combined together in different ratios, he did not succeed in arriving at the law which governed such combinations. That great discovery was left to Thomas Dalton, who, in 1804, enunciated the law of Multiple Proportions generally associated with his name. If two substances, A and B, unite in more than one proportion, giving rise to two or more compounds, then the various weights of A, which unite with a fixed weight of B, bear a simple ratio to one another. The actual experimental results that enabled Dalton to arrive at this law, were, however, so faulty that subsequent analyses, published by other chemists, will be used to exemplify this law. The first exact verification of Dalton's law of Multiple Proportions was carried out by Berzelius some years later (1811-1812). The following table records his results for the analysis of different compounds : — TABLE 1 Elements. Name of Compound. Lead Oxygen . Oxides of lead. I. II. Yellow. Brown. 100 100 7-8 15-6 Weight of oxygen combined with 100 of lead in I : weight of oxygen com- bined with 100 of lead in 11 ^■^ 1-2 -15-6 -^-^ Oxides of sulphur. Oxygen united with 100 of sulphur in I Sulphur , Oxygen , Iron . Oxygen Copper Oxygen Sulphurous. Sulphuric. 100 100 97-83 146-43 Oxides oi iron. Ferrous. Ferric. 100 100 29-6 44-25 Oxides of copper. Red. Black. 100 100 12-3 25 Oxygen united with 100 of sulphur in II - 9'7-83 _ 2 : 2-993 146-43 =2:3 nearly Oxygen combined with 100 of iron in I Oxygen combined with 100 of iron in II = 1^ =2 : 2-99 = 2 : 3 nearly 44-25 Oxygen united with 100 of copper in I Oxygen united with 100 of copper in II =1^ = 1 : 203 = 1 : 2 nearly 25-0 ■^ In all these examples of chemical compounds formed by the combination of different pairs of elements, the quantities of the 26 AN INORGANIC CHEMISTRY second element combined with a fixed quantity of the first element bear a simple ratio to each other, e.g. 1:2, 2:3, 1:2. Some forty years later (1849) Stas subjected this law to an exhaustive test in his analyses of the two oxides of carbon. Taking a fixed weight of carbon, he found that the ratio of the weight of oxygen combined with this carbon in the compound carbon monoxide is to the weight of oxygen combined with the same weight of carbon in the compound carbon dioxide as 1 to 1 -99990 ; the difference between experiment and the demands of the law, viz., 5 in 200,000, represents no more than his experimental error. We may therefore conclude that, so far as it has been tested, the law of Multiple Proportions is one of the exact laws of science. Another example, chosen from modern analyses. Mill per- haps make this law clearer to the minds of the beginner. Five well-known oxides of nitrogen are known. If a fixed weight of nitrogen is taken, say 14 gms., it will be found that the relative proportions of nitrogen to oxygen found in these five compounds are as follows : TABLE 2 Compound. Weight of Mtrogen. Weight of Oxygen com- bined with 14 of Xitrogen. Nitrous oxide . U Nitric oxide U Nitrogen trioxide 14 Nitrogen tetroxide 14 Nitrogen pentoxide 14 8 = = 8-1 16 = = 8 2 24 = = 8x3 32 = = 8x4 40 = = 8x5 The quantities of oxygen combined M'ith this fixed quantity of nitrogen therefore bear to one another the simple ratios 1:2:3:4:5. The actual method of carrying out these experiments would be to take a known weight of one of these oxides and determine the weight of oxygen and of nitrogen contained therein, and so on for all these compounds. Having found by analysis the actual ratio in which nitrogen and oxygen are combined, we can calculate by proportion what weight of oxygen would combine with 14 of nitrogen to form the various compounds, e.g. nitrous oxide is found by analysis to contain 63-63 per cent', of nitrogen and 36-36 per cent, of oxygen, hence 14 of nitrogen QUANTITATIVE ASPECT OF CHEMISTRY 27 combines with 36-36x14 63-63 =8 of oxygen and so on for all the compounds. The methods adopted in the analysis of chemical compounds vary from compound to compound ; the exact methods used in determining the composition of water will be described later (q.v.) but the general principles can be under- stood even at this stage. A weighed quantity of dry copper oxide is placed in the porcelain boat shown in Fig. 10. Dry Hydrogen F.G. 10. A stream of dry hydrogen is passed through the apparatus, and after all the air has been displaced from the apparatus the tube is carefully heated to a duU red. The water generated by the reaction between the hydrogen and the copper oxide is carried into the receiver B, which contains some substance capable of absorbing water vapour and thus preventing its escape. The passage of the hydrogen is continued in order to effect the reduction of the black oxide to the red metal. The increase in weight of B gives the amount of water formed during the reaction, the loss in the weight of the boat supplies us with the knowledge of the weight of the oxygen contained in the water which has been trapped in B. AU the data are now available for calculating' the composition of water, e.g. weight of copper oxide before the experiment „ after „ „ =17-569 21-991 gm. loss in weight (i.e. weight of oxygen contained in the water) = 4-422 weight of water collected = 4-976 weight of hydrogen contained in the water = -554 hence ratio of hydrogen to oxygen =0-554 ; 4-422 = 2-004 : 16. 28 AN INORGANIC CHEMISTRY This experiment also supplies us with the data necessary for determining the composition of copper oxide, for we see that 17-569 of copper combines with 4-422 gm. of oxygen yielding 21-991 gm. of the oxide, hence copper oxide contains 17-569 of copper : 4-422 of oxygen=63-57 : 16. Law of Reciprocal Proportions. — The two great quan- titative laws of chemical composition, the laws of Fixed and of Multiple Proportions, established the position of chemistry as an exact science. The next great advance in this direction was made by Richter and after him BerzeHus. In 1792-1799 Richter drew attention to the quantitative manner in which acids and bases (q.v.) react with one another, but it was BerzeUus who brought into prominence the results of Richter and appHed his methods to the elucidation of reactions between elements. The following table illustrates the method used by Berzelius : — . in the sulphide with 15-42 sulphur ,„„ , ri n 1- ^ ,. sulphur 15-42 „ ^„ 100 parts 01 lead combine ratio — ^ =2-02 \ oxygen 7-7 N in the oxide with 7-7 oxygen ^'in the sulphide with 25-6 sulphur 100 parts of copper combine — =-^ ^2-08 \ oxygen 12-3 \in the oxide with 12-3 oxygen in the sulphide with 58-73 sulphur 100 parts of iron combine — = =1-98 \ oxygen 29-6 in the oxide with 29-6 oxygen. The ratio is, within the Umits of experimental error, oxygen the same whether compounds of lead, copper or iron are compared. Furthermore, Berzelius found that while 100 parts of lead combine with 15-42 parts of sulphur, the resulting lead sulphide produces 146-33 of lead sulphate. This lead sulphate consists of the oxide of lead (already shown to be, lead 100+7-7 oxygen) and an oxide of sulphur. Consequently the weight of sulphur oxide contained in 146-33 gm. of lead sulphate must be 146-33— 107-7 = 38-63. This sulphur oxide must contain the 15-42 of QUANTITATIVE ASPECT OP CHEMISTRY 29 sulphur originally combined with the lead, hence the amount of oxygen=38-63— 15-42=23-21. ^ ,. , sulphur . , , ., 15-42 7-7x2 , Katio 01 — m sulphur oxide = = nearly, oxygen ^ 23-21 7-7x3 "^ whilst he had already shown that lr>/^ ±. J T 1 ,- /'.n the sulphide with 15-42 sulphur 100 parts 01 lead combme < . , , - j -.i r, r, ^ X m the oxide with 7-7 oxygen ^. sulphur 15-42 7-7x2 ratio — = = nearly. oxygen 7-7 7-7x1 These and other similar results led Berzelius to the general- isation known as the Law of Reciprocal Proportions. If a fixed weight of any (standard) element be taken, the weights — multiple or sub multiple — of the various elements com- bining with the fixed weight of the standard element, will also, react with one another. Law of Combining Weights. — ^Most of the oxides of the metals have been analysed by methods essentially similar to that given above ; in other cases the ratio of metal to oxygen has been determined by aUowing a weighed quantity of metal to combine with the oxygen of the air ; the weight of the oxide formed in this way supplies the requisite data for calculating the ratio metal : oxygen. An ele- ment, such as calcium, or mag- nesium, may have the composition of its oxide determined in this way. For such a purpose the apparatus given in Fig. 11 may be used. A weighed quantity of magnesium or calcium is placed in the crucible and after putting on the Ud, the metal is slowly heated. After some minutes the burner is increased to its full power. The crucible is then cooled. The white powder within the crucible, which also contains a Uttle nitride as well as oxide (see magne- sium nitride, p. 525), is then moistened m order to destroy the Fig. 11. 30 AN INORGANIC CHEMISTRY nitride. If the crucible is again heated to a duU red and cooled, the white powder will be found to consist of magnesium or calcium oxide. This is weighed as such and the composition of the oxide then obtained. In the following table is summarised the percentage com- position of the oxides of a few of the elements. A certain weight of the compound is analysed and the weight of each constituent present in 100 parts calculated. TABLE 3 Copper oxide . Copper 79-9 Hence 8 of oxygen combines with Oxygen 20-1 31-8 of copper Water Hydrogen 11-18 1 -008 of hydrogen Oxygen 88-81 Magnesimn oxide. Magnesium 60-32 12- 16 of magnesium Oxygen 39-68 Ferrous oxide. Iron 77-8 27-9 of iron Oxygen 22-2 Mercui-y oxide Mercury 92-59 100-0 of mercury Oxygen 7-41 By a different method of analysis one may arrive at a similar series of results expressing the composition of compounds containing hydrogen. TABLE 4 Hydrogen oxide Hydrogen 11-18 Hence 1-008 of hydrogen combines (water) Oxygen 88-81 with 8 of oxygen Hydrogen chloride Hydrogen 2-7 Chlorine 97-3 35-5 of chlorine Hydrogen bromide Hydrogen 1-34 Bromine 98-66 79-9 of bromine Hydrogen sulphide Hydrogen 5-92 Sulphur 9407 16 of sulphur From the composition of the chlorides the following table may be constructed : — TABLE 5 Hydrogen chloride Magnesium chlor- ide Silver chloride Copper chloride . Hydrogen 2-7 Chlorine 97-3 Magnesimn 25-53 Chlorine 74-47 Silver 75-26 Chlorine 24-74 Copper 47-3 Chlorine 52-7 Hence 35-5 of chlorine combines with 1 of hydrogen 12-16 of magnesium 107-9 of silver 31-8 of copper QUANTITATIVE ASPECT OF CHEMISTRY 31 For the sulphides we have : — TABLTi () Hydrogen sulphide Hydrogen 5-92 Hence 32-07 of sulphur combines Sulphur 94-07 with 2 of hydrogen Mercury sulphide. Mercury 86-18 200_of mercury Sulphur 13-81 Lead sulphide Lead 86-(i2 207-1 of lead Sulphur 13-38 Silver sulphide Silver 77-12 107-88 of silver Sulphur 22-88 In the oxygen table (for reasons which will soon become apparent) we have calculated in a somewhat arbitrary manner the weights of the various elements which combine with 8 parts by weight of oxygen ; amongst the elements combined with oxygen stands hydrogen which combines with oxygen in the ratio 1-008 : 8. Consequently, in the hydrogen table the results have been recalculated to express the weights of the various elements which combine with 1-008 gm. of hydrogen. Amongst the elements combined with hydrogen is chlorine, which combines with hydrogen in the ratio 1-008 : 35-5. The chlorine table has been recalculated in the last column to show the weights of the various elements combining ^^'ith 35-5 parts of weight of chlorine and so on. The agreement is quite the same, whatever element is chosen to start from, nor does it matter what arbitrary value we assign to that element. In the table before us oxygen has been chosen as the standard and the value 8 chosen for it. Had we assigned the number 100 to oxygen (as indeed was actually done in the time of Berzelius),'the numbers obtained for the various elements would differ by the ratio 100/8, but as the weights of all elements would be equally affected, no real change would be caused. For the present the reasons which have led to the choice of 8 parts of oxygen as our standard must be passed over. They will become clear when Chapter IX is perused. Some of the facts recorded in the above table may be grouped together thus : — - Oxygen — Hydrogen — Chlorine — Silver — Sulphur — Lead 8 parts 1-008 36-5 107-9 32-08 207-1 S2 AN INORGANIC CHEMISTRY These quantities, 8 of oxygen, 1-008 of hydrogen, 355 of chlorine, etc., represent chemically equivalent quantities of these ele- ments, and are known as Combining or Equivalent weights of these elements. The combining or equivalent weight may be defined as that weight of an element which will combine with or displace 8 parts by weight of oxygen (or 1008 parts by weight of hydrogen). In general it may be stated that nearly every element may have its combining weight determined by converting a weighed quantity of the element into the oxide, but in cases where direct com- bination between element and oxygen does not occur, it is always possible to combine the element under consideration with another element, the combining weight of which with reference to oxygen is definitely known. Fluorine is such an element, for no oxide of fluorine has been prepared. But Oxygen — Hydrogen — Chlorine — Silver — Fluorine 8 parts 1-008 35-5 107-88 19-0 The combining weight of silver is 107-88. If a weighed quantity of silver is converted into silver fluoride, the ratio 5 -. — is readily fluorme determinable, thence the combining weight of fluorine itself obtained. As a final example, consider the scheme. Oxygen — Hydrogen — Sulphur — Copper 8 parts 1-008 16-03 31-8 The law of Reciprocal Proportions as appUed to this series of reactions, states that, if copper and oxygen combine together to form an oxide, the ratio of weights of the elements combined in a given weight of copper oxide will be 31-8/8 (or some multiple or submultiple) ; this is borne out by experiment, for the most careful analysis of the black oxide shows that the percentage composition of this compound is copper 79-89, oxygen 20-11, from which one calculates that 31-8 parts of copper are combined with 8 of oxygen. But the conception of combining weight is by no means Umited to simple compounds of the type hitherto considered. Exhaustive study of all known compounds of the elements has failed to reveal a single case in which the composition of a compound cannot be represented in terms of the combining weights of the elements con- cerned. The case of lead sulphate, investigated by Berzelius, QUANTITATIVE ASPECT OF CHEMISTRY 33 is one in point ; the composition of the substance lead sulphate shows that 303-1 parts of this substance contain 207-1 of lead, 32-07 of sulphur, 64 of oxygen, i.e., 2 X 103-55 of lead +2 x 16-03 of sulphur + 8 X 8 of oxygen. Having now chosen a certain standard quantity of a standard element, we have been able to assign to aU elements fixed numbers, known as the combining Aveights of all the elements, which indicate the quantities of these elements which are chemically equivalent, and therefore capable of taking part as a chemical unit in all the many chemical reactions in which that particular element may be concerned. This selection of the combining weight as our unit suggests the following definition to embrace aU three laws of chemical combination : — Elements combine with each other in the ratio of their combining weights or of simple whole multiples of these. — Law of Combining Weights. The importance of the combining weight as a chemical unit is at once seen, for the chemical composition of an element might be represented in a very real way by indicating the number of equivalents of each element present in that compound. The percentage amount of an element present in a compound difiEers from compound to compoimd, e.g. water contains 88-81 per cent, of oxygen, magnesium oxide contains 39-68 per cent, of this element, mercury oxide 7-41 per cent, and so on. These numbers tell us little, but if the composition is expressed in such a way as to indicate the number of combining weights of each element contained in the compound, we shall be able to derive the com- position of the compound from our knowledge of the combining weights of the elements concerned. Each substance has its own distinct chemical unit or equivalent. To the chemist it is of httle importance that a piece of zinc shoidd weigh exactly the same as a certain piece of copper, but the statement that a certain weight of zinc will exactly displace a certain (but different) weight of copper from its compound is of very fundamental importance, for it strikes at the very root of chemical classification. This raises a fundamental distinction between chemical and physical phenomena. In the latter we are concerned purely with the physical unit of weight, which is employed for the determination of the quantity of a substance present ; chemical phenomena, on the contrary, are concerned with quantities D 34 AN INORGANIC CHEMISTRY which are chemically equivalent to each other, and which are, therefore, not physically equal in weight. Chemistry employs a unit of its own. With certain important hmitations to be discussed ui the next section, the combining weight of the different elements ofiers a suitable chemical unit in terms of which chemical union may be expressed. Difficulties with regard to the Combining Weights.— In the section dealing with the Law of Multiple Proportions, attention was drawn to the occurrence of numerous oxides of nitrogen. TABLE 7 Percentage Composition. Weight of Kitrogen combined with 8 parts ol Oxygen. 1 Nitrogen. Oxygen. Nitrous oxide . Nitric oxide Nitrogen trioxide . Nitrogen tetroxide Nitrogen pentoxide . 1 63-63 46-66 . 1 36-84 30-43 25-93 36-37 53-33 63-16 69-57 74-07 14 7 4-66 3-5 2-8 In each of these compounds nitrogen possesses a different combining weight. Which of them is correct or are all correct ? The combining weight has been defined as that weight of the element which enters into combination with or displaces 8 parts by weight of oxygen, and granting this definition there is no doubt that a different combining weight must be ascribed to nitrogen in each of the above compounds. A similar state of affairs exists if one considers the different oxides of iron which are known. TABLE 8 Compound. Percentage Composition. Parts of Iron com- Iron. Oxygen. Oxygen. Ferrous oxide Ferric oxide .... Magnetic iron oxide . 77-71 69-93 72-34 22-29 30-07 27-66 27-9 18-6 20-9 QUANTITATIVE ASPECT OF CHEMISTRY 35 It is because of this variability in the combining weight of many elements that the choice of the combining weight as the fundamental chemical unit has had to be discarded, for it is obvious that no systematic method of sorting up the chemical elements into unit quantities represented by the combining weight of the element is possible if that unit quantity varies from compound to compound. Atomic Weight. — ^We have seen that the combining weight, important though it may appear, lacks a certain characteristic which a true chemical imit should possess, viz. invariability. For this reason, although the conception of combining weight still plays a great part in the science of chemistry, as a suitable unit it is unsatisfactory for our chemical unit of quantity must be invariable. The time is not yet ripe to discuss exactly how the true chemical unit, the unit of chemical reaction, has been selected. That will be dealt with in Chapter IX, but in order that the student may be able to avail himself at once of the results of that discovery, a brief outline of the hne of advance wQl be given. In 1801, Thomas Dalton published a very important paper dealing with the constitution of matter. He put forward the hypothesis that matter is made of small particles or atoms ; the atoms of the same element were always similar to one another, of identical weight and so on, but the atoms of different elements were different. Compounds, he postulated, were formed by the joining together of atoms of different elements, 1:1, 1 : 2, 2 : 1, etc. Since each atom has its own deiinite weight, such a view of chemical reaction and of the constitution of matter, leads at once to the Laws of Fixed and Multiple Proportions. But for the diffi- culty of weighing an isolated atom of any element we should have in the absolute weight of the atom the fundamental chemical vmit. By means which wUl be discussed in detail in Chapter IX it has been possible to compare the weight of an atom of the various elements with the weight of an atom of our standard elemeni — oxygen. In 1898 the value 16 was assigned to oxygen as repre- senting its atomic weight. This is a purely arbitrary number, but it serves as a starting point, as our standard. The atomic weights of aU other elements have been determined, relative to oxygen. The atomic weight possesses one important advantage over the combining weight, its invariability in all the many chemical 36 AN INORGANIC CHEMISTRY reactions in which the atom may take part. There is a definite relationship between the combining -weight and the atomic weight as expressed in the equation A.W.=mxC.W., where n denotes a small integer. The atomic weight of iron is 55-8, the combining weight in the compoimd ferrous oxide is 27-9, hence in 55-8 this compound n=2, whilst in ferrous oxide w=r— -^ =3, etc. The student wQl at once see that, for the consideration of chemical reactions, aU the advantages urged in favour of a suit- able chemical unit, e.g. combining weight, as compared with the physical unit of weight, apply with equal force to the conception of the atomic weight, and added to this is the inestimable advant- age of invariability. The atoms of the various elements may be coupled up to form various compounds, one atom may be displaced by another and so on, yet in all these many-sided phenomena the atom preserves its distinct individuahty. Under such an hypothesis it is possible to represent the composition of a compound by denoting how many atoms of A unite with a Element. Oxygen . Hydrogen Sulphur Chlorine . Phosphorus Nitrogen . Alunainiuin Sodium Potassium . Copper Silver . Iron Lead Tin. . Silicon Zinc TABLE 9 Combining Weight. Atomic Weight. 8-0 16-0 : 1-0 1-0 1 8-01 ■( 5-36 32-07 135-5 ■ 1 15-1 35-5 16-2 "( 10-3 31-0 1 7-0 ■ ■ 1 13-5 14-0 1 9-03 27-1 ' 23 23-0 . 1 39-1 39-1 (31-8 163-6 63-6 107-9 107-9 (27-9 118-6 55-85 ( 103-5 1 51-7 207-1 (59-5 129-7 119-0 7-07 28-3 32-7 65-4 QUANTITATIVE ASPECT OF CHEMISTRY 37 certain number of atoms of B, and since each atom of A and of B has its own specific weight, the laws of conservation of mass and the laws of chemical combination become a necessary corol- lary. In table 9 is a list of the more important atomic weights (approximate) with their combining weights. (When an element possesses more than one combining weight, the two most important are given.) Further Consideration concerning Chemical and Phy- sical Change. — A simple criterion that enables one to class a certain phenomenon as being a definite chemical reaction and not a physical change of state is by no means easy to lay down. During chemical reactions we have seen that there is no change whatsoever in weight ; again, during chemical reactions there is always a definite change in the energy content of the system. Some of the chemical energy of the system becomes converted into other forms of energy, e.g. heat, light, electricity. The chemical union of copper and sulphur or of mercury and oxygen produces a substance of entirely different properties from those of the original element, but what of the change of ice into water, of red mercuric iodide into the yellow form ? Are these changes physical or chemical ? Water and ice possess their own specific properties. One is a crystalline solid, the other a mobile liquid. Each has its own density, its own colour. Furthermore, the conversion of ice into water requires a certain definite amount of heat per gram. The energy content of Igram of ice is different from the energy content of 1 gram of water, just as the energy content of equivalent quantities of sulphur and of copper differs from the energy content of the compound formed by their union. A stOI more difSctdt case is that of the conversion of yellow mercuric iodide into the red form. Each of these sub- stances possesses its own definite colour and solubility. The transformation of one into the other is accompanied by a definite heat change and a definite volume change. Are these changes not as deepseated as those produced by the chemical union of copper and sulphur to form copper sulphide ? The ease with which water may be converted into ice and ice into water, the easy reversibility of the change of one modification of mercuric iodide into another as contrasted with the great difficulty experienced in recovering the copper and the sulphur from copper sulphide, appears a serviceable method of dis- 38 AN INORGANIC CHEMISTRY tinguishing the one class of phenomena from the other ; but on the other hand, we shall meet later many chemical reactions which also show the same tendency towards reversibility under a slight change of conditions. As a rule, however, reversibility is more easily effected for physical than for chemical phenomena. During chemical actions, while a sHght change of conditions often caMses partial reversibihty, the change is less sharp than that shown by the sudden transition of yellow mercuric iodide into the red, and of water into ice. But the only safe criterion that enables one to classify a change as being definitely chemical or physical in nature is based upon the chemical tests given by the two substances. There is no chemical test or reaction which is given by ice and not by water. There is no chemical test or reaction given by red mercuric iodide which is not given by the yeUow variety. On the other hand the chemical tests given by mercury and oxygen before combination are entirely different from the chemical reactions given by the compound, mercuric oxide ; the chemical reactions of copper and sulphur are distinct from the reactions of copper siilphide. Sodium is an element which reacts vigorously ^^'ith water, chlorine is a poisonous gas, but the compound formed of these two elements, common salt, has properties which are utterly different from those of the original elements. Again, when a physical change of state occurs, any arbitrary quantity of the substance may be used, whilst in chemical reactions, interaction must take place between quantities which bear a simple ratio to the atomic weights of the elements under consideration. Mixtures and Chemical Compounds.— The student will have already gleaned from the preceding pages some of the fundamental distinctions between these two classes of sub- stances. These may now be briefly discussed. Mixtures are capable of mechanical separation, compounds are not. The separation of a mixture of sulphur and iron filings, either by use of the magnet, which will remove the u-on, or by means of the solvent, carbon disulphide, which has the property of dissolving sulphur, iUustrates the ease with which a mixture may be separated into its component parts. But if the iron and sulphur be brought into chemical combmation, no such physical method of separation will bring about the separation of the iron and the sulphur. Compounds are aliuays homogeneous ; mixtures QUANTITATIVE ASPECT OF CHEMISTRY 39 are, as a rule, Tieterogenecms. A good microscope is generaUy sufficient to make evident the individual components present in the mixture. No microscope is able to distinguish the slightest sign of heterogeneity in a true chemical compound. However far the physical subdivision of the latter is pushed, it remains truly homogeneous. Mixtures which are homogeneous are really solutions. In this case a separation of the mixture into its component parts can be readUy brought about by the process of freezing out or of evaporation, i.e. the creation of a second phase, the composition of which approximates more or less closely to that of one of the components of the mixture. The properties of mixtures are additive, i.e. they are the mean of the specific properties of the components of the mixture, while the properties of a chemical compound are specific and in no way deducible from a knowledge of the specific properties of its constituents. Hydrogen and nitrogen are non-poisonous gases, carbon a black solid, but the compound formed from these (hydrogen cyanide) is a most deadly poisonous gas. Mixtures are generally formed without any pronounced energy changes, whilst the formation of compounds is always attended by a pronounced energy change. The formation of a solution is one of the few physical processes which is attended by any consider- able evolution or absorption of heat. No better example illustrating this statement and the preceding one relating to the additivlty of the specific properties of a mixture, can be given at this stage than that of air. Nitrogen and oxygen when mixed in the ratio of four to one form a mixture, the specific properties (i.e. density, refractive index, colour, etc.) of which can be accurately calculated from a knowledge of the properties of the component gases. Moreover, there is no gain or loss of energy (heat, light, etc.) when the two gases are mixed. Within certain limits, the composition of a mixture may be varied at will, but the composition of a chemical compound is fixed, unalterable. This is the most important criterion of a chemical compound. Here again, it is noticed that solutions are most ahke to chemical compounds, for there is a definite limit of solubility of one substance in another beyond which we cannot go, but the fact that we can vary the composition of solutions at aU is in itself sufficient to justify the classifying of these substances amongst mixtures. 40 AN INORGANIC CHEMISTRY Recapitulation The quantitative study of chemical reaction has led to the formulation of the following laws : — (1) No change of weight occurs during chemical action — Law of Conservation of Mass. (2) Every chemical compound, irrespective of its method of ]jreparation, always contains the same elements combined together in the same proportion by weight — Law of Fixed Pro- portions. (3) If two substances A and B unite in more than one pro- portion, giving rise to two or more compounds, then the various weights of A which unite with a fixed weight of B, bear a simple ratio to one another — Law of Multiple Proportions. (4) If a fixed weight of any one element be taken, the ■s\eights (multiple and submultiple) of the various elements combining with the fixed weight of one standard element, will also react with one another — Law of Reciprocal Proportions. (5) The combining weight is that weight of an element which win combine with or displace 8 parts by -rteight of oxygen (or 1-008 parts by weight of hydrogen). (6) Elements combine with each other in the ratio of their combining weights or of simple whole multiples of these — Law of Combining Weights. Because many elements possess more than one combining weight this number cannot be chosen as the chemical unit. Dalton's atomic hypothesis has led to the choice of suitable atomic weights for the elements, these atomic weights being simple multiples of their combining weights. CHAPTER III THE NOMENCLATURE AND LANGUAGE OF CHEMISTRY The Naming of the Elements. — Although the actual number of chemical compounds is to be measured in thousands, less than one hundred elements have been isolated. Many of these are named according to one of their specific properties, e.g. chlorine (from its green colour), osmium (from its smell), the names of others are derived from localities, e.g. gallium (named after the country, France (Gallia), where the element was discovered), strontium (from Strontian in Scotland), whilst others have been named in a more or less haphazard way, such as uranium (from the planet, Uranus). The Classification of the Elements. — The elements are divided into two main groups — the metals and the non-metals. The metals, which are by far the more numerous, possess those well-known, easily recognisable properties associated with such substances as iron, silver. As a rule, the metals are dense, malleable and ductUe, and they take a pohsh readily. They are, in general, good conductors of heat and of electricity. The distinctive chemical properties common to the metals will become evident at a later stage. The non-metals show no well marked similarity in properties. In a general way we may say that the distinctive properties of the metals are lacking. Many of them are gaseous, nearly all are bad conductors of heat and electricity. A few of the elements, known as the metalloids, possess properties which are intermediate in character between those of the metals and of the non-metals. Symbols . — In order to make possible the building of chemical equations which wiU enable the chemist to express in a concise way the mechanism of a chemical reaction, it has become the 41 42 AN INORGANIC CHEMISTRY custom to assign to each element a symbol. Where the name of the element is common to several languages, the first letter is chosen as the symbol, but in case of ambiguity a second letter is used, e.g. Ga (gaUium), Br (Eng. bromine ; German, brom), but where the name of the element is different in the various languages, an appeal is made to Latin (English, copper ; Grerman, kupfer ; hence the symbol Cu. from the Latin cuprum is chosen), Hg. for mercury (Latin, hydrargyrum). TABLE 10 CliASSIFTCATION OF THE jMOEE ImpOKTANT ELEMENTS (WITH THEIR .SY:irBoi.s) Mct.als. Aluminium, Al. Barium, Ba. Bismuth, Bi. Calcium, C'a. Chromium, Cr. Cobalt, Co. Copper, Cii. Gold, Au. Iron, Fe. Lead, Pb. Magnesium, ilf;. Manganese, Mn. Mercury, Hg. Nickel, iSTi. Platinum, Pt. Potassium, K. Silver, Ag. Sodium, Na. Tin, Sn. Zinc, Zn. Jletalloids. Non-metals. Antimony, Sb. Argon, Ar. Arsenic, As. Boron, B. Bromine, Br. Carbon, C. Chlorine. Fluorine, F. Helium, He. Hydrogen, H. Iodine, I. Nitrogen, N. Oxygen, 0. Phosphorus, P. Silicon, Si. Sulphur, S. Formulae and their Construction. — Not only has the chemist assigned to every element a definite symbol, but, since the time of Berzehus, each symbol carries with it a perfectly definite quantitative significance. The symbol Au is not a shorthand method of writing "gold," but means the atomic quantity of gold, 197-2 parts by Aveight. Similarly, wherever the symbol is met with, the chemist means the atomic quantity (16 parts) of the element, oxygen, CI indicates 35-5 parts of chlorine, etc. By international agreement the gram has been chosen as the chemical unit of weight, so that, in the concrete sense, the symbol Au imphes the presence of 197 2 gm. of gold. NOMENCLATURE 43 The student is warned against overlooking the qtiantitative sig- nificance of the chemical symbol. The number of atomic quan- tities of each element present in a particular compound is shown by suiifixing small numbers to the symbols of the elements present, the suffix unity being always omitted. The compound formed by the combination of one atomic quantity of iron (55'8 parts) to one of sulphur (32 parts) is represented by writing the symbols side by side, e.g. FeS. The compound formed by combining one atomic quantity of iron to two atomic quantities of sulphur (2 X 32 parts) would be represented by the formula FeSj. Such a combination of chemical symbols is known as a, formula. The construction of a chemical formula from the results of analytical operations involves a knowledge of the atomic weights of the elements concerned. Example. — Determine the chemical formula of the compound which contains 2594 per cent, of nitrogen and 74-06 of oxygen (0 = 16, N = 14) . The first step is to find out what multiple of the atomic weight each of these percentages is. The number of atomic quantities of oxygen = = 4-629 25-94 - ^._ „ „ „ » nitrogen = -j^ = 1-853 But 4-629 : 1-853 = 5:2 (nearly). Hence there are five atomic quantities of oxygen combined with two of nitrogen. The simplest formula is therefore N2O5. The student will note that the formula N4O10 represents the compound equally correctly, and in order to distinguish between the various formulse, other data are necessary. This aspect wiU be discussed in Chapter IX. In the absence of sufficient data to enable one to distinguish between the various possible formulse, it is the custom to choose the simplest possible formula, fractional suffixes being avoided. Conversely, the percentage composition of a chemical com- pound can be readily calculated if the formula is known. As an example take the compound AgNOg. The formula indicates that 107-9 parts of silver, 14-0 „ nitrogen, and 3x16= 48-0 „ oxygen 169-9 44 AN INORGANIC CHEMISTRY are combined together, forming 169-9 parts of the compound, silver nitrate. 107-9 The percentage of silver present=100x ^^_-_ =63-51 14 „ nitrogen „ =100Xjggg= 8-24 48 „ oxygen „ =100x ^^g- =28-25 100-00 The Building and Significance of Equations. — The great value of chemical formulae lies not only in the fact that they convey to the chemist a precise knowledge of the composition of the compound, but also in the ease with which he can repre- sent in a quantitative way the details of the reactions in which the substance has taken part. It is the custom to place on the left hand side the formulae of the reacting substances connected by the sign of summation. The arrow head points to the right, i.e. towards the products of the reaction, e.g. Fe+S — y. The equation can be completed if one knows the composition (i.e. the formulae) of the substances produced in the reaction. Sup- pose, for example, the compound FeS, is produced, then Fe-f S->FeS2. It is obvious that two atomic quantities of sulphur will be needed on the left side of the equation. This is shown by placing the integer before the symbol, hence Fe+2S-^FeS2. Provided the formulae of all the substances taking part in the reaction are known, it is a comparatively easy problem to assign to each the requisite factor to make the chemical equation balance. The roasting of lead sulphide (PbS) in the air leads to the formation of lead sulphate (PbSOj), hence PbS+40->PbS04. Four atomic quantities of oxygen wUl be needed to convert the lead sulphide into the sulphate. The conversion of iron into the oxide (FcoOO is summarised in the equation •2Fe + 30^Fe,Qa NOMENCLATURE 45 To call such a symbolical representation a chemical equation is somewhat of a misnomer, for a novice in chemistry will have already realised that in no sense are iron and oxygen equal to iron oxide. True, they produce iron oxide, but the specific properties of the compound are entirely distinct from those of the elements. In the same way the chemical energy of the oxide differs from the chemical energy of the elements. For this reason, the sign of equality has been replaced by the arrow, as this merely indicates the direction in which the reaction proceeds. The sole equality which holds for the two sides of this chemical equation is the mass. Every chemical symbol embodied in such an equation refers to a definite quantity of matter. Therefore, every chemical equation, however simple or complex, is a quantitative expres- sion of the Law of Conservation of Matter. Furthermore, whenever a chemical equation is written down, the imphcit assumption is made of the truth of the laws of chemical com- bination. The equation 2Fe + 30 ^ Fe^Oa (2 X 55-8) (3 X 16) (2 X 55-8 -f 3 X 18) is therefore a quantitative expression, not only of the Law of Conservation of Matter, but also of the Law of Multiple Pro- portions and of the Law of Combining Weights. The student is strongly advised to look upon a chemical equation from the quantitative aspect, rather than as an easy means of summarising a chemical reaction. Above all, a chemical equation is in no sense an algebraic equation and is therefore not amenable to the manipulation to which an algebraic equation is subject. Simple Chemical Calculations. — The recognition of the quantitative nature of a chemical equation is of great importance to the chemist, for it enables him to calculate with accuracy the quantity of any particular substance produced in a chemical reaction, given a knowledge of the quantities of some of the other reacting substances and provided the equation, which is the quantitative expression of that particular reaction, is also known. Nothing can be done unless the equation symboUsing the reaction is known. As an example calculate what weight of copper sulphide can be obtained from 10 gm. of copper. The equa- tion which governs the combination of copper and sulphur is 46 AN INORGANIC CHEMISTRY Cu + S -^ CuS (63-6) (32-0) (63-6 + 32-0= 95-6) The interpretation of this equation tells us that from 63 '6 parts of copper it is possible to obtain 636+32=95-6 parts of copper 10 x95'6 sulphide ; hence from 10 gm. of copper we obtain ws^ — =15-03 gm. of the sulphide. (The student will note that in this and in all other chemical equations the symbol represents the atomic quantity of that particular element.) As a second example the following wUl serve : — ^What weight of zinc wiU be required to yield 20 gm. of zinc sulphate (ZnSOi ) 1 As already described, zinc sulphate is formed by the action of zinc upon sulphuric acid. The equation is Zn+H^SOi -^ ZnS04 + 2H (65-4) (65-4 + 32 + 4 X 16 = 161-4) The equation tells us that 161-4 parts of zinc sulphate are formed from 654 parts of zinc ; the weight of zinc necessary for the pro- 20 duction of 20 gm. of the sulphate is 65-4 X = 8- 10 gm. In this problem it is to be noted that the figure suffixed to a symbol refers to that element alone, e.g. in the compound ZnSOi the four refers to four atomic quantities (4 X 16=64) of oxygen only. A different point is exemplified in the problem : What weight of mercuric iodide is formed from 50 lbs. of iodine ? Hg + 21 - -> Hgl, (200) {2 X 126-9 (200 + 2 = 235-8) 126-9 = 453-8) The expression 21 means two atomic quantities of iodine (2 X 126'9). The weight of mercuric iodide obtainable from 50 lbs. of iodine is therefore 50 X . =89-4 lbs. Very frequently the integer precedes the formula of a chem- ical compoiind, e.g. 2HgO. In such a case, this means twice the formula weight of the compound, 2(200+ 16) =432. In compUcated formulae, Na2CO3,10H2O (washing soda, or the decahydrate of sodium carbonate), we have the formula weight of NajCOj, combined with ten times the formula weight of NOMENCLATURE 47 water, HjO, hence total formula weight =-(2 X 23+12+3 X 16) + 10(2X1 + 16) =106+ 180=286. Other examples involving chemical calculations are to be found at the conclusion of this and subsequent chapters. The student is strongly advised to study them, for nothing in the study of chemistry is so conducive to clear thinking and logical reasoning as the working out of chemical calculations. Questions 1. Give a concise account of the essential features of a chemical reaction. How does a chemical reaction differ from a, physical change of state ? 2. What do you understand by the " properties " of a substance ? Distinguish between " arbitrary " and " specific " properties. 3. Discuss the following statement : " Elements combine with each other in the ratio of their combining weights, or in simple whole multiples of these." 4. Two oxides of nitrogen have the following composition : (1) Nitrogen, 63-64 per cent, oxygen, 36-36 per cent. (2) Nitrogen, 25-94 per cent, oxygen, 74-06 per cent. Enunciate the Law of Multiple Proportions and use the above data to explain the law. 5. A mineral contains 32-79 per cent, sodium, 1302 per cent, alu- minium, and 54-19 per cent, fluorine. Determine the simplest formula of the compound. 6. What weight of oxygen can be obtained from the decomposition of 24 gm. of mercuric oxide ? 7. What weight of sodium chloride (NaCl) is produced from 1 cwt. of sodium ? Na + Cl^>NaCl. 8. What weight of sulphur would be needed to convert 144 gm. of oxygen into sulphur dioxide ? S + 20->S02 9. A compound is found to have the following composition : — Potassium. . . • 35-6 per cent. Iron . . • • 17-0 Carbon . . . • 21-9 Nitrogen . • • 25-5 „ Determine the simplest formula of the compound. 10. Distinguish between a mechanical mixture and a chemical com- pound. 11. Enumerate and disciiss the laws of chemical combination. 12. What weight of 10 per cent, sulphuric acid wUl be required to dissolve 3 lbs. of iron ? CHAPTER IV OXYGEN Historical. — The discovery of oxygen is frequently attri- buted to Priestley (1774). This experimenter found that, on concentrating the sun's rays upon a httle red oxide of mercury, a small quantity of gas was evolved. The actual method by which his result was obtained is illustrated in Fig. 12. This gas, which he called " de- phlogisticated air," was found to be a vigorous supporter not only of combustion but also of life itself. Somewhat prior to this (1773) Scheele, a Swedish chemist, obtained oxygen from various compounds containing oxygen, including the red oxide of mercury, but owing to a delay in the pubHcation of his discovery, his name is less associated with the discovery of tliis element than is that of his contemporary, Priestley. Sun's Rays -Burning Glass Calx of Mercury Mercury Fig. 12. Occurrence. — About the- same time, Lavoisier, the noted French chemist, extended the work of Rey upon the calcination of metals in air, and his experiments offer a remarkably fine example of the application of the scientific method to the inves- tigation of the unknown, leading him to the conclusion that atmospheric air consists of two gases — oxygen and nitrogen, of which the oxygen alone combines with a metal during combustion. His first experiment was carried out with tin. The metal was strongly heated in a sealed, weighed flask for a considerable 4S OXtGEN 4& time, and on cooling, the total weight was found to be unaltered. However, on opening the flask, an inrush of air occurred, showing that part of the air originally present in the flask had been absorbed during the heating of the metal. On reweighing the flask, an increase in weight was observed, which was found to be equal to the increase which the tin alone had undergone. By varying the quantities of tin and air present, Lavoisier was able to show that some uncombined air was always left over, i.e. only part of the air can combine with the tin during calcination. Air apparently contained at least two constituents, one of which was capable of reacting with a heated metal. This conclusion was tested by heating mercury in a retort as already illustrated in Fig. 1. The flask and bell-jar were filled with air. After heating the retort for a day, a steadily increasing quantity of a reddish substance made its appearance, while the volume of air in the bell-jar slowly diminished. When heat appeared to cause no further change (twelve days), the experiment was stopped. About one-sixth of the air originally present had been absorbed by the mercury, and the residual five-sixths had lost the pro- perty of sustaining Ufe and of supporting the combustion of a candle. This gas he named azote (Gk. a privative, f'oi) life), now known as nitrogen. On heating the red powder prepared as above, Lavoisier obtained about the same volume of gas as had been absorbed in the first part of his experiment. This gas caused a candle to biu'n with unusual brilhancy, in fact it pos- sessed in an enhanced degree all those properties which the air had been deprived of by being heated in the presence of mercury. Lavoisier at first called this gas vital air, afterwards changing the name to oxygen (acid-producer), from the fact that the compounds formed by its vmion with many elements gave acid- tasting solutions when dissolved in water. The name oxygm has survived, although its significance rapidly departed as it was soon recognised that many acid substances do not contain oxygen. Oxygen occurs in the atmosphere mixed with about four times its volume of nitrogen, but it is also found in enormous quan- tities in the earth's crust combined with other elements. By comparuig a large number of rock analyses, F. W. Clarke has estimated the average percentage composition of the earth's crust one-half mile deep, including the ocean and atmosphere. E 50 AN INORGANIC CHEMISTRY TABLE 11 Element. Oxygen. . Silicon . Aluminium Iron Calcium Magnesium Sodium. Potassium . Hydrogen . Titanium . Carbon Chlorine Phosphorus Sulphur Barium . Manganese . Nitrogen Other elements Per cent, in Per cent, in ocean. Per cent earth's crust alone. in atmo- sphere. Average. 47-07 85-79 23-03 49-78 27-06 — — 26-08 7-90 — — 7-34 4-43 — — 4-11 3-44 0-05 — 319 2-40 0-14 — 2-24 2-43 1-14 — 2-33 2-45 0-04 — 2-28 0-22 10-67 — 0-95 0-40 — — 0-37 0-20 0-002 0-19 0-07 2-07 — 0-21 0-U — — 0-U 0-11 {}■{>'.) — 0-11 0-09 — — 009 0-07 — — 0-07 — -- 75-54 0-02 0-55 1-44 0-53 Prepaeation 1. From the Atmosphere. — The air is liquefied (for details see Liquid Air, p. 273), and the more volatile nitrogen is allowed to escape. The residue rapidly approximates to Uquid oxygen, and at this stage the evolved gas is compressed into cylinders for sale. At present this is easUy the most economical method for preparing oxygen on a large scale, and has largely displaced the older Erin's process. The efficiency of this Uquefaction process is considerably aided by the demand for nitrogen for the manu- facture of nitrates {see p. 290). Erin's process is based upon the fact that, when barium oxide is heated in contact with the air at about 500°, oxygen is rapidly absorbed with the formation of the peroxide. On raising the temperature to above 800°, the oxygen is evolved with the production of the normal oxide. 500° BaO+0;=±Ba02. 800° Economically speaking, the method in this form is not efficient as it necessitates the alternate heating and cooling of the furnace, OXYGEN 51 but by taking advantage of the efiect of a change of pressure on the reaction, the double process can be carried out at a constant temperature of 700°. The principle of Le Chateher (p. 58) predicts that barium peroxide will tend to form the system — barium oxide +oxygen, if any attempt is made to lower the pressure ; conversely, if compressed air is pumped over heated barium oxide, the system occupying the smaller volume, viz. barium peroxide, will be formed. In practice, air at a pressure of 2 kilogs. per sq. cm. is pumped over the oxide at a tempera- ture of 700°, thereby forcing the reaction to the right. BaO + 0— ^BaOj. A valve at the extremity of the apparatus enables the nitrogen to escape. When the reaction is complete, the action of the pumps is reversed, and the pressure in the apparatus falls to 0-05 kilogs. per sq. cm. -This disturbs the equilibrium in the direction indicated by the arrow BaOa-^BaO+O. The economy of the process is ia this way considerably increased, but still the process cannot hold its own with that based upon the production of hquid air. The compressed oxygen prepared by this method is about 95 per cent. pure. 2. From Oxides. — The historic method of Priestley affords an interesting example of this type. HgO->Hg + 0. Similarly, manganese dioxide and the oxides of silver and of gold evolve oxygen on heating, though the first-mentioned substance does not form the metal but a lower oxide. 3Mn02->Mn304-f 20. 3. From Peroxides. — ^A few peroxides or higher oxides give up their oxygen readUy. Thus, sodium peroxide, when treated with water, effervesces freely with the evolution of considerable quantities of oxygen. Na^Oa +H20^^2Na0H + 0. The gas can be readily obtained by allowing the water to drip from a tap funnel upon the peroxide (Fig. 13). Hydrogen peroxide, especially in the presence of finely divided platinum, breaks down with the formation of water and oxygen. 52 AN INORGANIC CHEMISTRY The platinum may be replaced by finely divided silver, gold, manganese dioxide, even by dust, and these agents are aU found to be imaltered at the end of the reaction. J^,^ Here Fig. 13. 4. From other Oxy- Compounds. — Potassium chlorate is easily broken down by the action of heat, and oxygen is hberated. At about 340° the chlorate melts, and a few degrees higher begins to evolve oxygen. Soon the mass sohdifies again into a mixture of potassium perchlorate and chloride, and the evolution practi- cally ceases. Above 600° a brisk evolution again sets in and only ceases on complete decomposition. KC103^^KCl + 30. (For further details of this reaction see Potassium Chlorate.) An interesting feature of this reaction is that the presence of manganese dioxide enormously facihtates the Uberation of the oxygen from the chlorate. The dioxide does not give up its own oxygen below 400°, the chlorate at 350°, yet a mixtm^e of the two substances wiU give off a steady stream of the gas at 200°. Moreover, the manganese dioxide is found to be quite unaltered at the conclusion of the experiment. Oxygen may also be obtained by the action of heat upon potassium permanganate, or by the action of sulphuric acid on OXYGEN 53 such rich oxy-compounds as potassium dichromate, potassium permanganate, manganese dioxide. Properties. — Oxygen is a colourless, odourless and tasteless gas. It is slightly heavier than air : thus 1 litre of oxygen (at 0° and 760 mm.) weighs 1-429 gm. as against 1-292 gm. for a similar volume of air. Its solubOity in water is considerable, about 5 volumes dissolving in 100 volumes of water at 0° and 760 mm., while at 20° about 3 volumes are dissolved. Pish are dependent upon the air dissolved in water for their supply of oxygen required for the purpose of respiration. The critical temperature (see p. 66) of oxygen is — 118°, the critical pressure {see p. 66) being 50 atmospheres, so that at — 118° a pressure of 50 atmospheres is necessary to produce condensation. Liquid oxygen has a pale blue colour and boils under atmospheric pressure at — 182-5°. The liquid has a specific gravity of 1-13 and is strongly magnetic. By the use of Uquid hydrogen as refrigerant, Dewar succeeded in freezing the liquid to a pale bluish white solid. Prom the chemical point of view oxygen is extremely reactive, combining with all elements except those of the helium group {v. p. 275), fluorine and possibly bromine to form oxides. With a few elements such as potassium, sodium, iron, phosphorus slow oxidation (combination) will take place even at ordinary temperatures, but, in general, the temperature has to be raised considerably in order to commence the reaction. A glowing spHnter of wood, when thrust into a jar of oxygen, bursts into glowing incandescence, producing a gas, carbon dioxide, as the result of the reaction. C + 0,-^00,. Oxygen itself has no action upon Umewater (a solution of Ume or calcium hydroxide in water), but after the splinter has bmnt for a few moments, a httle clear limewater, if introduced into the jar, will at once become turbid owing to the separation of calcium carbonate. Ca(OH)2 + C02-^CaC03 + H20 Sulphur bums in oxygen with a lavender blue flame of great brilliance, forming gaseous sulphur dioxide, which has the characteristic odour of burning sulphru-. If a little water, rendered blue by the addition of blue htmus, is shaken up with this sulphur dioxide, the colour changes to red. 64 AN INORGANIC CHEMISTRY Phosphorus bvims in oxygen with exceptional vigour, pro- ducing copious clouds of white fumes. These fumes consist of particles of sohd phosphorus pentoxide, which dissolve in water with considerable energy to form phosphoric acid. 2P + 50->PA PA+3H20->2H3P04 (Phosphoric acid). This solution also possesses the property of reddening blue litmus. On the other hand, sodium or potassium bum in moderately dry oxygen to form the corresponding oxides and these oxides dissolve in water, giving solutions of the hydroxides which turn red litmus blue. 2Na + 0->Na20 Na20+H20-^2NaOH (Sodium hydroxide). Another element which bums in oxygen with great briUiance is iron. A tuft of " steel wool " tied to the end of a stout iron wire is strongly heated and plunged into a jar of oxygen, on the bottom of which has been placed a layer of sand. Dazzhng scintillations immediately fiU the jar, and globules of molten iron oxide fall to the bottom of the jar (Pig. 14). Combustion and Oxidation. — ^Towards the end of the sixteenth century it was recognised that air was necessary for the maintenance of the flame of a burning candle. Mayow (1674) clearly demonstrated this by means of the following apparatus (Fig. 15). Fio. 14. Fig. 15. OXYGEN 55 The jar was placed over the burning candle, and by means of the syphon tube, the pressure within and without was equalised. The syphon was then removed. The flame of the candle soon diminished in size and finally went out, while the water rose in the jar. This experiment in more capable hands might have forestalled the epoch-making experiment of Lavoisier (see p. 19) and incidentally have prevented the fantastic Phlogiston Theory of Combustion (q.v.) from holding its own till towards the end of the eighteenth century. However, Lavoisier's classic experiments on the calcination of metals, wherein he showed that the increase of weight during the formation of the calx was exactly equal to the weight of oxygen absorbed from the surrounding atmosphere, ultimately uprooted this hypothesis and in its stead established the present theory of combustion. According to this theory, cximbustion is a process of oxidation or combination which is generally attended by the evolution of light and heat. In more recent times it has become recognised that oxygen itself is not essential for combustion, e.g. a candle may burn quite well in chlorine, iron wiU burn in sulphur vapour. In all cases of combustion heat and light are given out, and oxidation occurs. The burning of such a metal as magnesium in a jar of oxygen and its slow oxidation in an atmosphere of this gas are essentially the same chemical process, the only difference being one of speed. Even the burning and the rusting of iron differ but httle, for each of these processes leads to the formation of an oxide of iron, the former having the composition FeaOj, the latter EeaOs+xHgO. So, too, the slow decay of wood leads to the formation of a mixture of carbon dioxide and water, as does the more vigorous combustion of wood, when burnt in oxygen. The removal of waste products from the human blood by the action of oxygen in the blood is but another example of slow com- bustion, the main product of the oxidation — carbon dioxide — being exhaled with each breath. In an elementary way, oxidation may be looked upon as that process which involves the addition of oxygen to an element or compound, as exemphfied in the following equations : C + O-^^CO (Carbon monoxide). C0 + 0->C02 (Carbon dioxide). 56 AN INORGANIC CHEMISTKY S +20-^802 (Sulphur dioxide). S02 + 0->S03 (Sulphur trioxide). On the other hand, oxygen is very often removed from a compound, i.e. the compound undergoes reduction. CuO+C^Cu + C0. (For further treatment of this subject see p. 166.) Spontaneous Combustion. — The actual amount of heat generated by the combustion of a given weight of coal is inde- pendent of the manner in which the coal is burnt, but the tem- peratvrc reached wiU depend upon the conditions under which the coal is burnt. If the combustion proceeds out in the open the actual rise in temperature may be very slight, owing to the dissipation of the heat into space through radiation. Such is the condition when hay. coal, slack, oily cotton waste, etc., are exposed in small quantities to the action of the atmosphere. Such oxidation M'ill take place without appreciable temperature rise, but if a large heap of such material is allowed to undergo slow oxidation under the action of the atmosphere, the heat generated within the mass cannot escape, and the temperature rises, possibly sufficiently to cause the mass to inflame. Such cases of spontaneous combustion have led to many serious fires. The presence of moisture in hay, sheep's wool, etc., is specially likely to bring about spontaneous combustion owing to its accelerating the slow oxidation. Oxides. — Elements combine with oxygen to form oxides. In many cases an element unites with oxygen in more than one proportion, and in such cases the names of the oxides should indicate the proportion in which the elements unite. A per- fectly regular nomenclature does not exist, but the following rules will be of use to a beginner. In the case of metals the termination -ous is generally used to indicate the lowest oxide, e.g. CujO, FeO, cuprous and ferrous oxide, -ic to denote the highest oxide, CuO, Fe^Os, cupric and ferric oxide. Ferric oxide (FejOa) and other such oxides are often called sesquioxides because they contain one and a half units of oxygen to one of metal. Besides these oxides there are the peroxides or dioxides, as, for example, of lead, barium, etc., PbOj, BaOa, the trioxides OXYGEN 57 as CrOa, chromium trioxide. So far as the non-metallic elements are concerned, the oxide containing one atom of oxygen is generally known as the monoxide, e.g. CO, NjO, CljO ; then come the dioxides, CO2, SOj, SiO,, the trioxides, SO3, N2O3, and so on. A large number of the oxides react with water, and the pro- perties of the solutions which are formed differ in a fundamental way. Attention has already been called to the fact that the oxides of sulphur and of phosphorus form solutions in water which redden litmus. On the other hand, certain oxides such as the oxides of sodium, potassium, calcium, etc., have the property of forming a solution which turns litmus blue. Substances like sulphur dioxide and phosphorus pentoxide are known as acid-forming oxides or acid anhydrides, because their solution in water produces an acid. SOa + HjO^HaSOs (Sulphurous acid) PA + 3H20^2H3P04 (Phosphoric acid) On the other hand, many oxides and their hydroxides possess the property of neutralising or destroying the fundamental properties of an acid. These are known as bases or basic oxides. The product of interaction between an acid and a base is always a saZi+ water. Acid + Base — > Salt + Water HCl + NaOH ^ NaCl -f H2O Hydrochloric ar-id. Sodium hydroxide. Sodium chloride. Water. 2HC1 + Na^O -> Sodium oxide. 2NaCl + H,0 A salt may also be formed by replacing the hydrogen of an acid by a metal. H2SO4 + Zn -> ZnSOi + 2H. The oxides and hydroxides which are freely soluble are fre- quently referred to as alkalies. This term is restricted to the hydroxides of sodium, calcium and their kindred elements. It will thus be seen that the term " alkali " and " base " are not synonymous, for every alkaH is of necessity a base, but all bases are not alkaKes. Catalysis. — The decomposition of hydrogen peroxide in the 58 AN INORGANIC CHEMISTRY presence of finely divided platinum affords an interesting example of what is known as catalysis or contact action. The efiect of the platinum is to accelerate a reaction which otherwise takes place very slowly, and at the end of the reaction the platinum is found to be chemically unchanged. Such a substance which has the property of accelerating or retarding the speed of a chemical reaction without itself undergoing any permanent change is known as a catalyst. In the above reaction the accelera- tion may be brought about by many finely divided substances, such as silver, gold, manganese dioxide and even dust itself. There is no question that the catalyst actually takes part in some way in the chemical reaction, though it is not clearly proved whether the participation is of a chemical or of a physical natiu-e. The action of manganese dioxide in facilitating the evolution of oxygen from potassium chlorate is another example of catalysis. It has long been recognised that many chemical reactions are accelerated by a rise in temperature, an increase of 10° generally doubling the rate of the reaction. The addition of about 25 per cent, of manganese dioxide to the potassium chlorate has at least as great an efiect as raising the temperature 300°, i.e. increasing the velocity 2'° times. Another catalytic agent, the presence of which is often neces- sary for the reaction to begin, is moisture, e.g. dry phosphorus wiil not burn in perfectly dry oxygen. Many years ago Ostwald likened the action of a catalyst to a drop of oil on the wheels of a machine. For elementary purposes no more sug- gestive .■simile is required, especially as research has, up to the present, seldom revealed exactly how a catalytic agent is able to modify the velocity of a chemical reaction. Principle of Le Chatelier. — Suppose that a small cylinder, closed by a piston, is fitted with water and ice at 0° C. (the freezing point of water), the system is in equilibrium, and unless heat is supplied to or taken from the system, or unless the pressure is altered, no change in the system wUl occur. But if the pressure is increased by forcing in the piston, what change will be induced in the system ? Ice is specifically lighter than water, so that one gram of ice occupies a greater volume than does one gram of water. The application of pressure to the system will cause ice to melt, because, in that way, the system OXYGEN 59 secures a release from the stress which is being put upon it ; similarly, a decreased pressure leads to the crystallisation of more ice, for the production of this substance with its greater volume tends to prevent the decrease in pressure which the experimenter is endeavouring to bring about. The melting of ice is attended by the absorption of heat, the separation of ice by the evolution of heat. If the system, ice-water, which is in equilibrium, is warmed, that process, which tends to use up the heat being put into the system, sets in, i.e. ice melts. The cooling of the system has the opposite effect. The general principle, exemplified in these experiments, was first enunciated by Le Chatelier. // an attempt is made to change the temperature, pressure or concentration of a system in equilibrium, this equili- brium will undergo such a change as will tend to undo the effect of the change of conditions to which the system has been subjected. The student has already had his attention directed to a practical application of this law in determining the best conditions under which oxygen may be obtained from the air by means of barium peroxide (Brin's process, p. 51), but numberless such examples are to be found in the study of chemistry. The correct applica- tion of this principle is often of the utmost value in predicting the most suitable conditions for carrying out chemical opera- tions. Questions 1. How do the " acid oxides " differ from the " basic oxides '' ? In what way are " salts " related to these two classes of substances ? 2. Describe fully any two methods for making oxygen in the laboratory. 3. What weight of oxygen is necessary for the combustion of 10 gm. of sulphur ? Discuss what happens when the sulphur dioxide formed is treated with water. 4. Illustrate the law of equivalents from the following table : — Potassium chloride. Potassium iodide. Iodine chloride. Potassium 52'5 per cent. Potassium23-6 per cent. Iodine 78'2 per cent. Chlorine 47-6 „ Iodine 76-4 „ Chlorine21-8 5. Give an account of what is meant by " combustion." 6. What is meant by stating that " manganese dioxide catalyses the decomposition of potassium chlorate " ? How would you prove the truth of your statement 1 CHAPTER V JTorricelli's vacuum THE PHYSICAL PROPERTIES OF GASES Chemistry, being an exact science, demands that quantitative expression shall be given to the various chemical reactions. Obviously this can only be obtained by the use of the balance, and whilst this is comparatively easy when solids and hquids are the subject of investigation, the problem becomes more difficult if the weight of a small volume of gas is required. Gases have Weight. — Even as far back as 117 B.C. Hero of Alexandria endeavoured to show that air is a material substance by pressing an empty bottle mouth downwards into water, but it was about 1,700 years later before Galileo showed by direct weighing that air has weight. Shortly after, in 1644, Torricelli pubhshed his now famous barometer tube ex- periment. A glass tube, nearly 4 feet in length, sealed at one end, was filled with mercury, the open end closed with the thumb, and the tube inverted. The thumb was removed under mer- cury and the tube clamped in the position shown in Fig. 16. No air was allowed to enter the tube, yet, instead of the tube remaining fiUed with mer- cury, the level of the liquid sank so that its height above the surface of the mercury in the dish was about 30 inches (760 mm.). TorrieeUi concluded that the weight of the column of mercury counterbalanced the weight of the column of the air pressing on the surface of the dish. Pascal carried the experiment a step further. Since mercury is about 13-6 times as heavy as water, Pascal argued that a column of water 60 Fig. 10. PHYSICAL PROPERTIES OF GASES 61 approximately 34 feet high should be supported by the pressure of the air. Experiment proved this to be true, and by using other liquids of known density, he concluded that the height of the column of liquid is a measure of the pressure, i.e. the weight, of the atmosphere. From these experiments was evolved the modem barometer. It was but a short step to prove the pressure of the atmosphere varied according to the height above sea level at which the observation was made. Atmospheric pressure at sea level, i.e. the standard pressure of the atmosphere, is equal to the weight of a column of mercury of unit area and 760 mm. high. This corresponds to the weight of 1033-3 gm. per sq. cm., or 14-7 lb. per sq. inch. Another direct proof that the air has weight may be obtained by weighing a flask fiUed with air at atmospheric pressure and then repeating the weighing after the air has been removed by means of a good pump. If the air in the vessel is replaced by another gas such as oxygen, it may be shown that all gases possess weight. Boyle's Law. The Effect oe Pressure ufon the Volume OF A Gas at Constant Tem- perature It is now a matter of everyday knowledge that the volume of a gas depends upon the pressure exerted upon that gas. The law expressing the interdependence of volume and pressure was first enunciated by Robert Boyle (1661), and states that the volume of a gas at constant temperature varies inversely as the pressure. This law can be readily tested in the ap- paratus shown in Fig. 17. Any desired amount of the gas under investigation is intro- FiG. 17. 62 AN INORGANIC CHEMiSTKY duced into the graduated measuring tube A by means of the tap G. The mercury levels in A and B are carefully adjusted and the volume of the gas in A read off. This gives the volume of the gas at atmospheric pressure. By raising or lowering B the pressure may be altered at will, and a series of related values for p and v at some constant temperature may be determined. TABLE 12 Values of pv fok Nitkogbn at 15° C. p. in metres of mercury. pi: 0-76 1-0000 5 0-9981 10 . . . 0-9963 15 . 0-9946 20 . . . . 0-9930 25 . 0-9910 An examination of this table shows that there is not exact proportionahty between pressure and volume, but that the value pv, which, according to Boyle's Law, should remain constant, shows a slight but steady decrease. Investigations have shown that this decrease in the value of pv remains, however carefuUy the measurements are made, and by whatever method they are made. In other words, the variation in pv is not due to experi- mental errors, but arises from the fact that the law as enunciated by Boyle does not exactly correlate the facts which it purports to correlate ; in short, the law is an approximate and not an exact law of Nature. Dalton's Law of Partial Peesstjres John Dalton (1802) showed that when two or more gases which do not interact with each other are mixed, each exerts its own definite pressure and the total pressure exerted is equal to the sum of the individual pressures exerted by each gas. The pressure of each individual gas is the pressure exerted by that gas if it alone fiUed the space. If pi, p^, etc., denote the partial pressures of the various gases, Dalton's Law states that P, the total pressure =pi+;p 2+ etc. This law is approximately true, provided there is no interaction between the gases present in the mixture. In many cases, e.g. mixtures of carbon dioxide and sulphur dioxide, whilst there is no chemical reaction, sUght deviations from the law occur owing to molecular attraction between the different kinds of molecules. PHYSICAL PROPERTIES OF GASES 63 Assuming the air to consist of 20 per cent, of oxygen and 80 per cent, of nitrogen, the partial pressure of oxygen present in air at atmosphere pressure is one-fifth of an atmosphere. Con- sequently the amount of oxygen dissolved in 1 htre of water exposed to the air will be only one-fifth of what it would be were the water exposed to pure oxygen at atmosphere pressure. Chaeles' Law. The Influence of Temperature upon the Volume of a Gas at Constant Pressure Towards the end of the eighteenth century Charles carried out experiments to determine the effect of a change of temperatiu^e upon the volume of a gas, and his results, coupled with the more systematic experiments of Gay Lussac (1802), are embraced in the so-called Charles' Law. The same rise in temperature produces in all gases the same increase in volume, provided the pressure be kept constant. Fig. 18 shows a suitable type of apparatus for verifying this law. The coefficient of ex- pansion, i.e. the expansion of 1 litre of gas, produced by a rise in temperature of from 0° to 1° C. for various gases, is given below : Fig. 1 These values are aU approximately equal to -.j ] j, but shght variations are apparent. Moreover, the coefficient of expansion of any particular gas is seen to vary according to the particular temperature the determination is made at. For nitrogen the following results have been obtained : Air . . . 0-003665 Hydrogen 0003667 Carbon monoxide 0.003667 Carbon dioxide 0-003688 Sulphur dioxide 0-003845 TABLE 13 Temperature. Volume in litres. Cocfflciont of Expansion. 0° 10° 20° 30° 40° 1-00000 1-036778 1-073539 1-110287 1-147044 0-0036778 0-0036770 0-0036762 0-0036761 64 AN INORGANIC CHEMISTRY Chailes' Law may be accurately represented by the equation Vf =V„( 1 + -— j, where V„, Y^ are the volumes at 0° C. and at any arbitrary temperature ^'^ C. respectively, .,\x the coefficient of expansion. By plotting this equation one obtains the follow- ing graph : ^ -^ y' f ^ ^ ^ :S ^ ^ -^ ^ ■^ " --- ^ Ter npe 'ra tut '6. -200' -100° 0' rOO° 200° Fig. 19. — Ghaph of Chahles' Law. An analysis of this graph shows that at — 273° C. the volume becomes zero, and below that temperature the volume becomes negative. A negative volume is beyond our conception, and so the temjoerature — 273° C. or 0° absolute is supposed to be the lowest possible temperature. If one takes —273° C. as the absolute zero, 0° C. becomes 273° abs., 100° C. is 373° abs., i.e. T=273+i, where T denotes the temperature on the absolute scale, t on the Centigrade scale. The equation Vt=V„(l + — jmay be written /273+<\ T V(=V„( )~^o'?r'' '^h®!'® T,, T„, denote the temperatures t" C. and 0° C. reckoned on the absolute scale. Consequently the volumes occupied by a gas at different temperatures, the f res- sure being constant, are directly 'proportional to the absolute temperature. The Combined Influence of Temperature and Pressure upon the Volume of a Gas. — Let p^, v^, T„ denote the pressure, volume and absolute temperature of the gas. It is required to consider what changes occur in the volume of the gas when both pressure and temperature change. Let the pressure of the gas, at constant temperature T„, be PHYSICAL PROPERTIES OF GASES 65 changed from p„ to pi, thereby causing a change in the volume from Vg to V. In accordance with Boyle's Law, it follows that PoK=Pi^ (1) Now, at constant pressure p^, let the temperature change from the initial temperature T„ to Ti, thereby causing a change in volume from v to Vi. Then, from Charles' Law, it follows that ^=-^i (2) If the value of v derived from equation (1) be substituted in (2) we obtaui ^^-=^', whence ^=^i^, i.e. ;pi'=RT, where R is the so-called gas constant. This equation is of very fundamental importance in all calcu- lations involving variations in the volumes of gases brought about by alterations in temperature and pressure. Since gaseous volumes are so dependent upon the particular tempera- ture and pressure at which they are measured, it has become customary to reduce them to a standard set of conditions, i.e. to convert the volume actually measured at any arbitrary temperature and pressure into the volume which would be occupied at 0° C. and 760 mm., i.e. at the so-called normal temperature and pressure. Example. — A certain weight of potassium chlorate liberates 450 c.c. of oxygen, measured at 20° C. and 750 mm. pressure ; what is the volume occupied by the gas at normal temperature 273 750 and pressure (N.T.P.) ? Obviously ?;=450X^^X— -=414 c.c. Deviations from Boyle's and Charles' Law. — In dealing with Charles' Law it has already been pointed out that this law demands that at the absolute zero the volume of the gas should become zero. That a definite quantity of matter, even at this temperature, should occupy no space is inconceivable. The explanation of this fallacy lies in the fact that Charles' Law does not apply to the whole volume of a gas, but rather to the spaces between the particles of the gas. As the temperature falls, these spaces shrink ; and if the gas remained a gas, these spaces would disappear entirely at the absolute zero, so that the final volume would be the volume occupied by the actual particles themselves. However, all gases hitherto discovered liquefy before the absolute zero is reached. Up to the present 66 AN INORGANIC CHEMISTRY the absolute zero itself has not been attained, though it is of interest to note that, by boiling heUum under reduced pressure Kammerlingh Onnes recently succeeded in attaining a tempera- ture of —271° C. or 2° abs. Boyle's Law, as has already been stressed, is also only approxi- mately true. More rigid work of modem times has demonstrated that the volume concerned in Boyle's Law is not the whole volume occupied by the gas, but the residual volume obtained when a deduction is made for the space actually occupied by the particles themselves, i.e. ^[v — 6)=fc, where v denotes the total volume and b the volume occupied by the gaseous particles. At low pressures b is negligible compared wth v, but as the pressure rises and the spaces between the particles become more and more constricted, the neglect of the correction 6 causes a considerable rise in the value of fc. Obviously, since b refers to the actual volume occupied by the particles of the gas, it is to be expected that b has a different value for the various gases. Moreover, the particles appear to exert a certain cohesive action upon each other, i.e. the pressure actually registered is less than the true pressure. This cohesive effect is inversely proportional to the square of the volume, so that the pressure has to be in- creased in the ratio 'p/p-\- ^ ' where a is a constant for the individual gas under consideration. The modified equation of Boyle, wherein account is taken not only of the volume of the molecules themselves but also of their cohesive action upon each other, is (p-| )(«— 6)=RT (van der Waals' Equation of State). In this form the equation appears to represent very closely the facts of nature. Critical Phenomena. — If the equation pv=K const, is graphed, a rectangular hyperbola (Fig. 20) is obtained and this is what is approximately obtained in many cases. The results of Andrews with carbon dioxide show a marked difference from these curves (Fig. 21). So long as the temperature of the experiment is above 31°, a reasonably good hyperbola is obtained. At 31° the curve shows a kink, whilst below that temperature the curve consists of three PHYSICAL PROPEETIES OF GASES 67 5 5 C - t V I \ ^ \^ ^ — — — — — — — /o ^0 soi/oiiimeAO so 60 Fig. 20. A III 1 I I d 1 ^ 1 omo^eneous Gas iV tv^ OS ^1 ^\ /yJU.>^^ ^J- "^^ T V fes ^ b _i ^ ^ _^[. 3/7^ i^ ^^=^ d/(7«/a' ^^5 :^ ::::_ 1 i^ w 1 1 1 1 \. "^~ = = = r= C^ V 0' Volume Fio- 21. 68 AN INORGANIC CHEMISTRY parts — a more or less vertical part, a horizontal portion and a sloping portion which gradually becomes asymptotic. The explanation of this irregularity is as foUows : for temperatures below 31° the steadily increasing pressure causes a diminution in the volume and the curve DC is traced out. At the point G hquid carbon dioxide separates from the system and the pressure becomes constant. Any attempt to increeise the pressure results in the conversion of more gas into Hquid with a corresponding decrease in the value of v. At B the last trace of gas hquefies. After this it requires a large increase in the pressure to produce even a sKght decrease in v — the curve BA has been reached. Above the temperature 31° no hquid separates, but at all temperatures below this the separation of hquid is possible ; 31° is the so-called critical temperature of this gas. By measur- ing the co-ordinates of the point E the critical pressure may also be read off. These experiments make it clear why the attempts to hquefy certain gases had been fruitless — the gases had not been cooled below their critical temperature before the increased pressure was apphed, for the diagram shows that no pressure, however big, can cause the hquid to separate from the system unless the temperature is below the critical temperature. The Kinetic -Molecular Theory. — The only mechanical basis upon which scientists have been able to build a theory that will account for the physical properties of sohds, Hquids and gases is on the assumption of a discontinuous structure of matter. The experimental results that have been obtained under the heading of Boyle's and Charles' Laws in particular, receive their most satisfactory explanation on such an hypothesis. The difference between sohds, hquids and gases is one of degree rather than of kind. The particles of matter, whether sohd, hquid or gaseous, are assumed to be incompressible and to be separated ■ from each other by spaces. The effect of pressure is to bring the incompressible particles closer together by reducing the free space between them. These particles are often called molecules — the imaginary units of which mutter in the aggregate is composed. The essential features of the Kinetic-molecular Theory may be summarised in the following statements : — 1. Matter is discrete, and is composed of a finite number of molecules. In gases the volume actually occupied by the mole- cules is very small compared with the free space ; whilst ia the PHYSICAL PROPERTIES OF GASES 69 case of liquids and solids the free space between the molecules is much more restricted — hence their slight compressibility. 2. The molecules of a gas are always in a state of violent motion in straight Hnes. 3. The molecules are continually colhding not only with the walls of the containing vessel but also with one another. 4. Owing to their perfect elasticity, they rebound after a collision without change of momentum. 5. A slight cohesive or attractive force between the molecules exists. This is specially noticeable at high pressures. Under these assumptions the effect of pressure upon a gas is at once explicable. Take a small volume of gas en- closed in a cylinder by a frictionless piston (Fig. 22). If the pressure on the piston is so increased that the volume is reduced to one-half, it follows that the impacts on the piston will be twice as frequent as before, i.e. the pressure will be doubled, hence the volume is inversely proportional to the pressure. The tendency of two gases to diffuse one into the other, the homogeneity of such a mixture, the non- settling of the heavier particles are all ahke explained on the assumptions of this hypothesis. The for- mula embodying Charles' Law, v, = v^ {\-\-at), the pressure being held constant, may also be written Pj=p„(l -\-at), the volume being constant. If the pres- sure is maintained constant, a rise of 1° C. will increase the volume by 2 i;j , whilst, if the volume is kept constant, this rise of temperature will cause an increase in the .pressm-e by the same amount. Pressure, we have learnt, is to be attributed to the impacts of the molecules. When each molecule strikes the wall, it experiences a change of momentum, and it is the sum of these changes of momentum which constitutes the pressure. The laws of physics state that the total change of momentum experienced when a perfectly elastic sphere strikes a wall is 2mv, where m denotes the mass of the molecule and v its velocity. A rise in temperature increases the pressm-e, i.e. the change of momentum caused on impact. Clearly m, the mass, cannot vary with a slight rise in temperature, so that v, the velocity, must be increased by a rise in temperature. The importance of this theory in governing the behaviour of gases cannot be too strongly stressed. Modern work has done Fig. 22. 70 AK^ lNOri,GANIC CKjiMISxEi much to extend the conclusions reached from a study of gases to the parallel case of Uquids, and though many points are still beyond our ken, the conception of the molecule has been of extreme value to the scientist. Although the molecule was introduced into science as a conception only, the recent work of a school of physicists, foremost among them being Jean Perrin, has led to such discoveries that the molecular theory of matter has taken on a vahdity that was httle dreamed of a decade ago. The actual physical existence of molecules has not yet been proved beyond question, but the cumulative effect of the recent evidence in this realm of science has done much to increase the probabihty of their real existence. This evidence has been mainly obtained by the use of a comparatively new instrument — the ultra-microscope. The Ultra-Microscope. — The hmit of visibiUty of a good microscope is about l/j, (^=0-001 mm.). This is not a question of iUumination, but depends upon the fact that the light merely passes round a particle of such small dimensions and continues r f Fig. 23. on through the microscope to the eye as if the particle did not exist. Tyndall showed many years ago that, if a beam of converging hght is brought to bear upon a solution containing in suspension such fine particles that the solution appears PHYSICAL PROPERTIES OF GASES 71 homogeneous, an opalescence is caused, due to the scattering of the beam of light. The ultra-microscope (Fig. 23) is based upon this discovery. A powerful converging beam of light is focussed by means of a plane mirror (7 and convex lens B upon a solution A, in such a way that it enters at right angles to the direction in which it is viewed by the microscope. With such an instrument it is possible actually to see the individual particles even though they do not exceed about 10"' cm. Solutions of such substances as sodium chloride in water, sulphuric acid in water, etc., show no such scattering, i.e. even under the ultra-microscope the solution appears homogeneous, but not so a solution of starch, albumen, sihcic acid, etc. In such cases the instrument at once proves the heterogeneous nature of the solution, that is, the substance is not in true solu- tion, but forms what is known as a colloidal or quasi-hetero- geneous solution. The dimensions of the particles in many of these colloidal solutions have been accurately measured. A few of the average values for such colloidal particles are here recorded. cms. CoUoidal gold .... 5xl0-« „ silver .... 6x10"" lead .... 4x10"" „ sUver iodide . . . 3x10"' These values vary httle from the estimated diameters of ordinary molecules of matter, e.g. hydrogen 2-1x10"^ cm., chlorine 3-7 x 10"* cm., nitrogen 2-9 X 10"* cm. The starch mole- cule 5 X 10"' appears to be even larger than an average colloidal particle of matter. The assumption, therefore, appears reason- able that such colloidal particles may be viewed as molecular aggregates formed from the aggregation of a relatively small number of individual molecules. The ultra-microscope has thus enabled scientists to see particles relatively httle greater than the average molecule. Even more suggestive are the results obtained from investigating the actual movements of these particles. The Brownian Movement. — In 1827 the botanist Brown drew attention to the pecuhar movements of lycopodium powder suspended in water. When observed under the microscope these appeared to be in a violent state of motion which persisted 72 AN INOjRWAiNiu uiiiiiViiBiKY over a long period, both when exposed to the Ught as well as in the dark. The motion was not due to convection currents of any kind, and must owe its origin to the hquid itself. The particles of a colloidal solution, when observed under the ultra-microscope, show the same Brownian movement, except that the movements are much more hvely. As Zsigmondy, one of the pioneers of this work, has said : "A swarm of gnats in a sunbeam will give an idea of the motion. The particles hop, dance, jump, dash together and fly away from each other. . . ." This motion appears perpetual. It was Ramsay who first suggested that it was due to the impacts of the solvent molecules upon the particles — an explanation now generally accepted. By studying the motion of these colloidal particles of different substances, Perrin and his co-workers have arrived at the follow- ing conclusions — The velocity, the frequency of collision, the distribution of the particles through the solvent, are exactly what the molecular theory assumes to be the case with the particles of a gas. The only difference seems to be that the spaces which separate the molecules of a gas are, in the case of colloidal solutions, filled with the solvent. Finally, by actually counting the number of particles seen under the ultra-microscope, the important conclusion has been reached that the number of particles present in a colloidal solution is, under similar conditions of temperature and pressure, equal to the number of molecules of gas com- puted to be present in a corresponding volume of a gas. Evaporation, Solution, Crystallisation. — The process of evaporation, solution and crystallisation receives a ready ex- planation at the hands of the kinetic-molecular hypothesis. The molecules of a hquid are in a violent state of motion, and some of these molecules succeed in escapmg from the surface of the hquid. Should the hquid be exposed to draught, many of these escaping molecules will be carried away from the neigh- bourhood of the hquid, i.e. the liquid will steadily evaporate. On the other hand, if the liquid is in a confined space, of the molecules which escape many will again strike the surface of the hquid and be trapped. So long as the number escaping exceeds the number caught per second, the hquid will evaporate, but sooner or later the time must come when equihbrium will be attained, i.e. the number escaping per second is exactly equal to PHYSICAL PROPERTIES OF GASES 73 the number re-caught per second. The Uquid will now be exerting its maximum vapour tension or pressure for this par- ticular temperature. If the temperature rises, the velocity of the moving particles of Hquid will increase, more wiU escape per second, and equihbrium will not be attained till the liquid exerts a higher tension than at the lower temperature, i.e. the vapour tension wiU rise with the temperature. So, too, in the case of solution or of the reverse process of crystallisation. Molecules of the sohd are continually leaving the surface of the crystal, and these again are continually return- ing to the surface. Solution represents the difference between the number escaping and the number returning ; crystallisation indicates that the number returning exceeds the number escaping, whilst at saturation point an equilibrium of a dynamic, not static, nature has been established. Density of Gases and its Measurement. — The density of a gas is the weight of 1 c.c. measured at 0° C. and 760 mm. pressure. In order to obtain a rough determination of the density of a gas the apparatus in Fig. 24 may be used. The flask, which should be reasonably light, is thoroughly exhausted, and weighed. It is then filled with the gas the density of which is required, at some known temperature and pressure {w the weight of the gas, t and p are thus known). The volume of the flask is obtained by filUng it with water and then weighing it. The difference in weight between the empty flask and the flask full of water, expressed in grams, represents the volume of the flask in cubic centimetres. Fig. 24. The volume v„ =v X 273 P ^noTi '^here t and p denote the «+273 760 temperature and pressure of the gas during the experiment. The density = - . In many chemical operations, it is inconvenient to carry through an actual weighing of the gas evolved in a reaction. 74 AN INORGANIC CHEMISTRY If, however, the gas is collected and its volume read at some known temperature and pressure, the volume at N.T.P. is readily obtained, and from a prior knowledge of the density of the gas, the total weight of the gas may be computed. Vapour Densities of Liquids and Solids and their Mea- surements. — It is of the utmost importance to the chemist to be able to determine accurately and quickly the vapour density of volatile hquids and sohds. Of the various types of apparatus which have been specially designed for this purpose, the following three stand out. The Dumas Method. — The flask used in this method consists of thin glass, and contains about 150 c.c. (Fig. 25). The exit is a long narrow tube. The bulb is weighed fuU of air, and a few grams of the substance intro- duced into the flask. The bulb is then put into a suitable bath at a constant tempera- ture from 10°-30'' higher than the boiHng point of the substance whose vapour density is being measured. The hquid vaporises and the vapour dis- places the air from the apparatus. When the vapour ceases to escape from the exit, the tube is sealed off with a mouth blowpipe. The temperature of the bath and the atmospheric pressure are then recorded. The bulb is cooled, dried and weighed. The next operation is to determine the volume of the bulb. This is done by breaking the tip under water. The bulb when filled with water is weighed and the difference between the weights of the fuU and empty flasks enables one to obtain the volume of the bulb . In this experiment it is necessary to apply an air correction for the displaced air. [f w is the weight of the bulb filled with air at the temperature t and pressure p, Wi the weight of the bulb filled with the sub- FiG. 25. PHYSICAL PROPERTIES OF GASES 75 stance, H2S04 + 0. The sum total of the reaction is the decomposition of water into hydrogen and oxygen and the accumulation of acid round the positive pole. (For fuUer treatment of this reaction read Chapter XXVIII.) The action of sulphuric acid in this case may be hkened to the action of finely divided platinum in promoting the decomposition of hydrogen peroxide — it f acihtates the reaction which is already taking place, and is itself left undecomposed at the conclusion of the reaction, i.e. the sulphuric acid behaves catalytically. There is unfortunately a tendency to relegate to the domain of catalysis all those mysterious reactions the mechanism of which man has not yet been able to unravel. In the minds of many, directly an " accelerated " reaction is thoroughly understood, it ceases to belong to catalytic phenomena. In the case above, probably every detail whereby the sulphuric acid is able to accelerate the liberation of the hydrogen during the electrolysis of water is clearly understood (see Chapter XXVIII), but does that rob the reaction of its place among catalytic phenomena ? Water may also be broken down chemically by the action of such metals as sodium, potassium, calcium, barium, etc., in the cold, and by magnesium, zinc, iron, nickel, when heated, but in many such cases only half the hydrogen contained in the water is set free. Na + H20->NaOH + H Ca + 2H20-^Ca(OH)2+2H 3Fe +4H20->Fe304 + 8H. Some of these metals are Ughter than water, e.g. sodium, and 80 AN INORGANIC CHEMISTRY in order that the gas may be collected in a pure state, the metal must be held underneath the surface by means of a piece of wire gauze or lead tube (Fig. 29). The metals which require heat to bring about the decomposition of water are generally placed in a tube which can be strongly heated and steam is then passed over them (Fig. 30) . 2. Acids may be readily forced to give up their hydrogen by the action of a suitable metal upon them. The acids most commonly used for this purpose are sulphuric and hydrochloric acids, the metals, zinc and iron. In such cases the reaction is much more vigorous when the acid is moderately 'diluted with water : Zn + H „S04 --> Zn SO, + 2H Zn + 2HC1^ ZnCl^ + 2H. In the laboratory the gas can be readily obtained by either of Fio. 29. _|fi"i)|fi ITiK[;ujjfivTni Fig. 30. these reactions, the generating vessel being a flask provided with a thistle funnel. For a more or less continuous supply the Kipp apparatus is generally called into requisition (Fig. 31). HYDROGEN 81 3. Several metals, notably zinc and aluminium, evolve hydro- gen from a solution of sodium hydroxide. Zn+2NaOH- Al+3NaOH -Na2Zn02 + 2H Sodium zincate. -NajAlOa+SH Sodium aluminate. Commercially, hydrogen is obtained during the electrolytic preparation of sodium hydroxide from sodium chloride {q.v.). Purification . — The hydrogen evolved by most of the laboratory methods is more or less impure. The gas escapes from the generator charged with aqueous vapour and is also contaminated with traces of air and various impurities from the chemicals used. Aqueous vapour is readily eUminated by passing the gases through concentrated sulphuric acid or over granulated calcium chloride, both of which absorb moisture with avidity. Traces of oxygen may be removed by passing the gases over heated platinised or paDadiumised asbestos, when the reaction 2H+0- ■H^O Fig. 31. — Kipp's Apparatus. is promoted. The presence of traces of nitrogen is seldom a disadvan- tage. So far as impurities from the chemicals are concerned, the par- ticular purifying agent chosen will depend upon the impurities Ukely to be present in the acid and the metal, and also upon the purpose for which the gas is required. Most freshly prepared gases carry over a fog of finely divided liquid or solid matter. This can readily be filtered out by passing the gas through several cen- timetres of cotton wool. Pure hydrogen for special purposes can best be prepared by taking advantage of the ready solubility of the gas in the metal 82 AN INORGANIC CHEMISTRY palladium, one part of which will dissolve 860 times its volume of hydrogen at ordinary temperatures. This occluded hydrogen may be removed in a very pure state by the combined action of heat and reduced pressiu-e. This property of occlusion is ex- hibited by other metals such as platinum, gold, iron, etc., but in a less marked degree. In the case of palladium there is stiU doubt whether the hydrogen is present in the palladium in the form of a solution or in a definite state of combination as hydride. Hydrogen, dissolved in palladium, has many properties utterly distinct from the properties of ordinary gaseous hydrogen. Besides causing a considerable change in the volume of the palladium, it has the property of reducing or converting ferric chloride (FeCla) into ferrous chloride (FeCla). FeCla + H -> FeCla + HCl. So-caUed nascent or atomic hydrogen, Uberated in the solution of ferric chloride by the action of zinc upon hydrochloric acid, has the same property. A stream of hydrogen dehvered from a Kipp has no such action. Physical Properties. — Hydrogen is a colourless, tasteless, odourless gas of extreme hghtness (density, 0-0695, air=l, weight of 1 htre N.T.P., 0-08995 gm.). Its critical temperature is —241°, B.P. —252°, M.P. (58 mm.), —260° ; solubihty in water 1-9 vols, in 100 at 14°. The gas is non-poisonous. Its specific heat is for a gas abnormally high (3-4). A wire, heated in oxy- gen or air to incandescence by the electric current, ceases even to maintain a red heat when immersed in a jar of hydrogen, the current being unchanged. Closely associated with its extreme lightness is its property of diffusing through porous partitions. Diffusion If a jar of hydrogen is placed, mouth downwards, on a jar of air and the two glass covers separating them removed, it will be found that in a very short time the two gases form a homogeneous mixture and wiU be evenly distributed between the jars. It is immaterial whether the fighter hydrogen or the heavier air is placed on top, the difiusion of the one gas into the other takes place until in a short time there is complete homogeneity. The phenomenon of diffusion can be neatly demonstrated by means of the apparatus shown in Fig. 32. The essential feature is a pot of unglazed porcelain surrounded HYDROGEN 83 by a beaker. The coloured liquid in the U-tube should stand at the same level in each Hmb. On introducing the exit tube from a carbon dioxide generator into the beaker B and Trom a hydro- gen generator into beaker A, a change in the level of the coloured indicator at once sets in. It is found that the hquid rises towards the carbon dioxide beaker B, but is depressed from the hydrogen beaker A. Presently equihbrium sets in and the coloured hquid takes up a permanent position. In apparatus 1 the air can escape more freely from the porous pot than carbon dioxide can enter, a decrease in pressure within the pot results, and this is Fig. 32. shown by the movement of the indicator. Ultimately just as much carbon dioxide enters per second as air escapes — the liquid assumes a permanent difference of level. In apparatus 2 the reverse happens. Hydrogen forces its way through the pores more freely than the air is able to escape and an increase of pressure results. In 1832 Thomas Graham showed that the speed at which a gas can diffuse through a porous membrane is a function of the density of the gas. His Law of Diffusion reads, The relative speed of diffusion of gases is inversely proportional to the square roots of their densities. The density of oxygen compared with hydrogen is 16, hence the rate of diffusion of 84 AN INORGANIC CHEMISTRY VT 1 oxygen is —^ = ~ that of hydrogen. Obviously, ii one measures the velocity with which equal volumes of two gases diffuse through a fine jet, Vj : Vi=VDi : VT>2, and if the density of one gas is knpwn, it is possible to calculate the density of the other. This method of obtaining the density has been utilised in certain cases where it has not been possible to obtain the pure gas in sufficient quantity for a direct weighing, e.g. ozone and radium emanation (q.v.). So, too, it is possible to separate two gases of different density by taking advantage of the varying rate of diffusion. A mixture of electrolytic gas, that is, a mixture of hydrogen and oxygen obtained by the electrolysis of water, if passed through a long clay pipe, will no longer explode on issuing from the pipe, but will show the characteristic tests of oxygen, so great is the proportion of hydrogen which has escaped. Chemical Properties. — Hydrogen, in itself a non-supporter of combustion, burns in oxygen with the production of intense heat, a temperature sufficient for melting platinum being attained . When the oxy-hj^drogen flame impinges upon a piece of hme, it becomes heated to incandescence and in that state gives rise to the so-called hme-light. Although the reaction between hydrogen and oxygen is so vigorous that at 900° explosion occurs, there is no measurable reaction at room temperature. Finely divided platinum, however, has the same accelerating influence in this case as has already been noted for hydrogen peroxide. Hydrogen combines, directly or indirectly, with nearly all non-metals, as well as with a number of the more reactive metals to form hydrides. In the case of the metaUic hydrides, the compounds are relatively unstable, i.e. easily broken down with the liberation of hydrogen. KH+H2O -^ K0H+2H. Potassmm Potassium hydride. hydroxide. Perhaps the most important property of hydrogen is its tendency to combine with oxygen even when that element is already combined with another. CuO+2H-^Cu+H20 ZnO+2H^Zn+H20 FcaOi + 8H-> 3Fe +4H20. HYDROGEN 85 The last of these reactions is of special interest, for it has already been stated that a common method of preparing hydrogen is by the action of steam upon iron filings. 3Fe + 4H2O -^ FeaO^ + 8H , i.e. by passing steam over iron, hydrogen and iron oxide are obtained, whilst, if hydrogen is passed over heated magnetic oxide of iron, iron and steam are reformed. If a mixture of iron and water are heated in a sealed tube, it is found on analysis that four products are present, viz. iron, iron oxide, steam and hydrogen, and for a definite temperature a perfectly definite ratio of steam to hydrogen exists. The explanation of this lies in the fact that the reaction is capable of proceeding in either direction, and if the products of the reaction are not able to escape, the time comes when an equilibrium is reached, i.e. as much steam is used up per second in reacting with the iron as is reformed by the action of hydrogen upon the iron oxide in the same time. If steam is passed over heated iron, the hydrogen produced is carried from the system by the excess of steam and the reverse reaction cannot be set up. Such a reversible reaction is of frequent occurrence in chemical operations. The combination of hydrogen and oxygen is of a reciprocal nature, for not only does hydrogen bum in oxygen, but oxygen burns equally freely in hydrogen. If a large jar of hydrogen, mouth downwards, is ignited at the mouth, and a jet of oxygen introduced, as the jet passes the mouth of the jar the oxygen is ignited and continues to burn in the hydrogen atmosphere, producing a deposit of water upon the walls of the jar (Fig. 33). Hydrogen Equivalent. — The importance of the combining weight of an element has already been stressed (p. 32). The combining weight, or, as it is often called, the hydrogen equivalent, may be readily determined by dissolving metals in suitable acids and measuring the hydrogen evolved. The apparatus in Fig. 34 is very satisfactory for this purpose. ^ is a small flask containing the acid (diluted) together with a weighed amount of metal in a small test tube held in position by a thread, B a measuring burette, D an adjustable mercury holder, C a three-way tap for adjusting the height of the mercury. When the height has been adjusted to zero, the flask is tilted in order to bring the acid and metal in contact and the evolved 86 AN INORGANIC CHEMISTRY gases are coUected. When the apparatus has cooled, the volume is recorded, together with the pressure and the temperature. All the data necessary for the calculation of the hydrogen equivalent are now available. , , j Example.-0-l gm. of zinc Uberated 36-2 c.c. of hydrogen measured at 18° and 755 mm. What is the equivalent of zmc ( 36-2 c.c. at 18° and 755 mm. are equivalent to 34- 1 c.c. at N.T.P. 1 c.c. of hydrogen at N.T.P. weighs 0-0000899 gm.. Oxygen Fig. 33. Fig. 34. hence the weight of hydrogen evolved by 0-1 gm. of zinc = 0-003065 gm. Therefore the weight of zinc required to Uberate 1 gm. of hydrogen is 32-6. Valence The reader wUl have already remarked that the hydroxides of the metals do not necessarily possess the same formulae — that of sodium hydroxide has been given as NaOH, aluminium hydroxide as A1(0H)3, calcium hydroxide as Ca(0H)2. The point at once suggests itself as to why one element can hold HYDROGEN 87 in combination three -OH groups whilst another can hold but one ; similarly, for the compounds NaCl, CaCla, AICI3. The atomic weight of calcium is 40, its combining weight or hydrogen . , . «^ mi ■• atomic weight . „ ,, , equivalent is 20. The ratio r-r-; . , is 2, the number combimng weight which also denotes the number of elements or groups with which atomic weight calcium combines. In the case of aluminium. combimng weight 27-1 o T 1. i. i-i, i.- atomic weight ,., , = — =3. In short, the ratio ^^—. . , appears likely 9-03 combimng weight to be of considerable importance in determining the formula of compounds in which a particular element occurs. The number which represents how many atoms of hydrogen, or other elements equivalent to hydrogen, can combine with one atom of a particular element is known as the valence of that element, or, expressed mathematically, , atomic weight valence = r-^~. . , , ■ combmmg weight The valence of hydrogen is taken as unity. If one is able to determine, directly or indirectly, how many atoms of hydrogen combine with an atom of a particular element, one can auto- matically calculate the valence of the element. The conception of valence is of great assistance to the student in mastering the formula of the many chemical compounds which exist. Given the formula of any oxide, e.g. CuO, since oxygen is divalent (=2H), copper must also be divalent. The action of hydrochloric acid must lead to the formation of a compound in which copper is likewise divalent, as expressed in the equation, CuO + 2HCl->CuCl2 + H2O. Carbon forms an oxide CO 2 ; one would therefore expect that if a hydride (or chloride) of carbon exists, the formula would be CHj, (CCI4). Such is indeed the case. Very often, in the elementary stages, it is customary to indicate the valence thus .• mono-valent H', Na', K', CP, Br', I', F', Cu' Ag". divalent Ba", Ca", Sr", 0", Cd", Cu", Fe", Pb", Mg", Sn", Zn)'. trivalent AU", N"', P'", As'", Sb"', Bi"', Fe"'. tetravalent Pb'^, Si'^, C'^, S'^, Sn'^. pentavalent N^, P'', As", Bi", Bi". hexavalent S"'. 88 AN INORGANIC CHEMiSTKY Monovalent elements are known as monads, divalent as diads, trivalent as triads, and so on. Should two of these elements unite to form a compound the formula may at once be written down with a fair degree of certainty. Arsenic oxide is AsgOj. or AS2O3, the chloride of antimony is either SbClj or SbCls and so on. Maximum and Multiple Valence. — In the above table some of the elements appear to possess more than one valence. This is a common feature, for it must be recognised that valence is a very variable property. Take, for example, the following compounds of sulphur — HjS, SO2, SO3. Few chemists are to be found at the present day who do not consider sulphur to be di-, tetra- and hexa-valent in these compounds. Tin forms the oxides SnO and SnOg, giving rise to the chlorides SnCla and SnCli and so on. None the less, elements appear to exhibit a maximum valence towards certain other elements, e.g. carbon never exhibits a valence of more than four towards oxygen, though it often exhibits a valence of less than four, as in carbon monoxide (CO). Nitrogen has a maximum valence of five towards oxygen (N2O5), though other compounds NA NO, N2O3, NO2 exist. The Effect of External Conditions upon Valence. — The general effect of a rise of temperature is to cause a decrease in the valence — the higher chloride of antimony, SbCls, readily evolves chlorine on heating, forming SbClg. On the other hand, an increase of pressure favours the formation of the compound with the higher valence. Thus, by increasing the pressure of chlorine upon antimony trichloride, the pentachloride is obtained. Should an oxidising agent be present, there is generally a tendency for the formation of the compound -with the higher valence, whilst reducing agents have the reverse effect. 2FeCl2 + 2HC1 + ^ 2FeCl3 + H ^O FeCls + H (nascent) -> FeCl^ + HCl. Graphic or Constitutional Formulae. — Formulae are often written in such a way as to show the valence clearly. H — — H indicates that hydrogen is monovalent, oxygen divalent ; ammonia is represented by the formula : /^ N^H and so on. HYDROGEN 89 The valences are represented by the links or bonds. The student is, however, warned against attributing to these linkages or bonds anything in the nature of a force holding the elements in combination. Valence is not a force, but is a measure of the number of atoms which any particular element may hold in combination. By a general development of this pictorial method of repre- senting valences, chemists have been led to the adoption of graphic or constitutional formulae. The assumption is made that the chemical properties of a compound are determined by the arrangement of the atoms in the particles of the compound, and consequently, if two substances exist having the same composition but with the atoms differently arranged, the pro- perties of those compounds wiU be fundamentally different. Graphic formulse, therefore, are built up in such a way as to represent in a pictorial manner the composition of the compound, but they do not indicate that the atoms are so held within the particles of the compound ; indeed, modern work on crystal structure shows the absurdity of such a contention. The examination of crystals by means of X-rays has shown that in a crystal of sodium chloride the atoms of chlorine are arranged in a perfectly definite and systematic manner, but the evidence is strongly against associating a particular atom of chlorine with any particular atom of sodium. Nevertheless, constitutional formulae are of great assistance in enabling the chemist to follow out the reactions of a given substance. A few typical graphic formulae are given : Ferric oxide. Sulphuric acid. Calcium hydroxide. - H-0 O re=0 \/ 0-H 0< S Ca/ re=0 y\^ \0-H H-0 O Radicles. — The constant presence of the -OH group in the hydroxides of the metals, the similarity in properties of the compounds containing this group leads to the conception of a radicle. A radicle or radical is a group of atoms which can enter into or be displaced from combination without itself undergoing any decomposition. In the reaction between zinc and sulphuric acid the SO 4 group passes as a group from the hydrogen to the zinc ; 90 AN INORGANIC CHEMISTRY this is the sulphate radicle. Other radicles are the groups NO 3 (nitrate), CO3 (carbonate), CIO3 (chlorate), NH4 (ammonium), PO4 (phosphate). It is to be noted that each radicle as a group has Its own characteristic valence. Mg(N03)2 may be written .(NO3) /O-N^o Mg<' or, written in detail, ^S\ r) ^(N03) 0-N NaHSOi + H^O Base. Acid. Salt. Water. The resulting salt is known as sodium acid sulphate or sodium hydrogen sulphate. This is our first example of an acid salt, in contradistinction to the normal salts wherein the whole of the hydrogen of the acid has been replaced by the metal, as in the reactions H2SO4 + 2NaOH^ Na2S04 + 2H,0 H2S04+Zn->ZnS04 + 2H. Acids which contain two replaceable hydrogen atoms (i.e. dibasic acids) or more (e.g. tribasic acids H3PO4 etc.) may have their hydrogen replaced in steps. This leads to the formation of acid salts. H 3PO 4 + NaOH -> NaH 2PO4 + H ,0 Sodium di-hydrogen phosphate (acid phosphate). H3PO4 + 2NaOH-^Na2HP04 + 2HoO Di-sodium hydrogen phosphate (acid phosphate), whilst the addition of three molecules of sodium hydroxide causes the formation of the normal salt H3PO4 + 3NaOH->Na3P04 + SH^O. Monobasic acids, e.g. HCl, cannot give rise to acid salts from the very nature of their constitution. HYDROGEN 91 On the other hand, di-acidic bases, e.g. bases containing two hydroxyl groups, may have these hydroxyl groups replaced one at a time. In this way basic salts may be formed. Pb (OH ) 2 + HCl -> Pb (OH )C1 + H 2O Basic lead chloride. Pb(OH)2 + HN03-^Pb(OH)N03 + H20. Basic lead nitrate. Such basic salts are nearly always less soluble than the normal salts. In actual practice basic salts are generally produced by the action of water upon the normal salt, e.g. Hg(N03), + H,0^Hg(0H)N03 + HN03 Mercuric Basic mercuric nitrate. nitrate. Basic salts occasionally arise from the combination of a normal salt with an hydroxide, as in the case of malachite, CUCO3, Cu(0H)2. The fact that the hydrogen of a dibasic acid may be substituted in steps allows the possibihty of building up a mixed salt, e.g. NaHSOl + KOH-^NaKSO^ + H2O. In a similar way one can prepare sodium potassium ammonium phosphate, NaK(NH4)P04, from phosphoric acid, H3PO4. Questions 1. Describe two methods for making hydrogen from sulphuric acid. 2. Give an account of Graham's Law of Diffusion. 3. What weight of water would be obtained by the passage of dry hydrogen over 10 gm. of lead oxide [formulae PbO, atomic weight of lead =207-2] ? 4. Explain what is meant by the " valence of an element." What are the valences of the acid groups in the following compounds : KCl, H3PO4, NajSiOi, CaS04, A]2(S04)3, PCI5. 5. Show how acid salts, basic salts and mixed salts are related to normal salts. Illustrate your answer by means of equations. 6. Construct equations showing how ferrous sulphate may be made from ferrous oxide, cupric nitrate from cupric oxide, ferric chloride from ferric oxide. 7. A quantity of zinc is treated with an excess of sulphuric acid, and the resulting gas, collected at a temperature of 15° and under a pressure of 765 mm. of mercury, is found to have a volume of 400 cc. Calculate the quantity of zinc used. [Zn = 65-4, H = l, 1 litre of hydrogen at 0° C. and 760 mm. pressure weighs 0-0899 gm.] CHAPTER VII WATER Water is one of the most familiar compounds occurring in nature, for it is found not only in the free state, but also in a state of combination in the bodies of animals, plants, etc. Occurrence. — Natural water varies greatly in purity. Sea water contains on the average about 3-5 per cent, of dissolved material ; rain water, too, contains appreciable quantities of both dissolved and suspended matter, the nature of which is determined by the character of the country through which the water flows. Lake Katrine (Scotland) which receives the drainage of a granite area, contains only 0-032 parts of dissolved matter per 1,000 parts of water, while the Thames, draining a chalk area, averages 0-03 per cent, of dissolved matter, very largely consisting of calcium salts. Spring and mineral water varies exceedingly in its purity. Percolating rain-water attacks the various rocks and carries off dissolved salts. This is especially true of the salts of calcium, magnesium, sodium and iron. This solvent action is due to the action of the dissolved carbon dioxide (see p. 531), and is also greatly faciUtated by a rise in the temperature and pressure of the water as it reaches the lower levels. This percolating water afterwards finds its way again to the surface of the earth. Some of the salts commonly found in such water are ferrous carbonate in the chalybeate springs of Tunbridge, Saratoga, N.Y., etc., sulphur compounds in the Hot Springs of Rotorua, N.Z., Baden, etc., magnesium sulphate and chloride at Cheltenham, Bath, etc., sodium sulphate and carbonate in Marienbad, sodium and other chlorides in Baden-Baden, Homburg, etc., boric acid in Tuscany, silica in the Mammoth Springs of Yellowstone Park, U.S.A., and the Hot Lake District of N.Z. The temperature of such spring 92 WATER 93 waters shows considerable variation. Many of those in New Zealand and California are over 50", while some of the deep artesian wells of Queensland give water much too hot for imme- diate use. Purification. — Matter in suspension may generally be rapidly removed by passing it through a suitable filter. Large filter beds of sand and gravel are a necessary adjunct to any town water-supply system. The small amount of mineral material in solution in such water supply systems is rarely deleterious — not so, however, pathogenic (disease producing) bacteria. For chemical purposes it is frequently necessary to prepare water of a high degree of purity. This is most easily attained by distUlation. Solvent action upon glass is much greater than the uninitiated imagine, so that for careful work the storage bottles must either be well steamed out or made of specially non-soluble glass. Physical Properties. — When coohng, water exhibits some- what anomalous behaviour in its volume changes. As the tern- i / 1 1 1 1 / i 1 / t-OOW f / / f / / / / r- y / / / s y 1-0000 v ■^ _ _ _. 0' 16° ■20° Fig. a° iz° Temperature. 35. — The Volume-Tempeeatube Graph of Water. perature falls from its boiling point, there is a steady contraction until the temperature of 4° is reached (Fig. 35) . At this point the 94 AN INORGANIC CHEMISTRY coefficient of expansion changes its sign and further cooling causes the liquid to expand. Water at 4°, therefore, has its maximum density. One c.c. of water at 4° forms our unit of weight — 1 gm. The fact that water has a greater density at 4° than at 0° is of great natural importance in preventing the freezing of lakes and rivers in times of great cold. At 0° water freezes, forming ice. The amount of heat evolved when 1 gm. of water freezes is 79 calories, i.e. is sufficient to raise 1 gm. of water 79°, conversely, the melting of 1 gm. of ice causes the absorption of 79 calories — the so-called latent heat of fvsion. . The temperature at which water boils under atmospheric pressure is 100°. To convert I gm. of water at this temperature into steam at 100° requires an absorption of 537 calories — the latent heat of evapor- ation. The gradual conversion of water into vapour at temperatures below the boiling point is a matter of everyday experience, and the explanation of this can perhaps be most easily supphed by the considera- tion of the following experiment (Fig. 36). Two barometer tubes, iiUed with mer- cury, stand side by side, and the upper portion of tube B is surrounded by a large tube into which water of any desired temperature may be introduced. If a drop or two of water is introduced into tube B, a depression of the mercury is at once apparent. The actual weight of the water in comparison with the weight of the mercury is neghgible, so that the depression cannot be due to the weight of the water itself, a contention supported by noting that, as the tem- peratiu'e of the water in the jacket rises, the depression of the mercury column increases. The amount of water has not varied, hence the change in the height of the mercury column must be due to the pressure of the aqueous vapour given ofiE by the water. This vapour pressure, together with the pressure due to the residual column of mercury, must equal the baro- metric pressure as registered in tube A. Consequently the differ- ence in the height of the mercury columns in the two tube^ Fig. 36. WATER 95 must be a measure of the vapour pressure of the water. Careful measurements have shown that for every temperature there is a fixed vapour pressure. The vapour pressure curve for water Z40 ZW WO ISO WO m m 80 60 4€ 20 i / / / ^ I 1 5 r, ci 1 i y^ .^ / 1 / 1 1 ^ / 1 / t 1 / / / / / /> y ^ ^ Te vpi va :ur '.. Fig. 37. 10° 20° 30° 4-0° 50° 60° 70° 80° -Temperatubb-Vapoub Peessube Ctxrve roR Water. (Pig.' 37) shows the rapid rise in the vapour pressure as the temperature increases. The curve for liquid water ends at 0°, while the upper hmit is 96 AN INORGANIC CHEMISTRY set by the critical temperature, above which water in the hquid form cannot exist. This curve forms a typical \'.P. curve for all liquids. The subject of vapour pressure may, however, be approached from another standpoint. Suppose a glass vessel B with manometer C attached is so arranged that it may be exhausted by an airpump. ^ is a three-way tap to facihtate the evacuation of the apparatus B, and to enable ^vater or any other hquid to be admitted. The evacuation is carried forward by means of the pump until the manometer C indicates zero pres- sure. A small quantity of water is then introduced through A. The manometer will show an abrupt rise in pressure, which is independent of the amount of Avater added, provided that suffi- cient has been introduced to leave hquid water in the system. By varying the temperature of the apparatus a series of vapour pressure readings may easily be obtained. If water is exposed to the atmosphere, it will tend to exert the vapour pressure characteristic of that temperature. As a rule, it rarely succeeds in attaining a value greater than about two -thirds of this. Air draughts and diffusion prevent the aqueous vapour from exerting its maximum pressure ; the more rapidly the vapour is removed from the neighbourhood of the liquid, the more rapidly will the evaporation proceed. Chemical Properties. — Owing to its great stabihty, water, though so necessary to many chemical operations, is not a. reactive substance. Attention has already been drawn to the reaction between certain oxides and water (see p. 57). Acid oxide -I- Water := SO, Sulphur dioxide. + H,0 Acid = H,S03 Sulphurous acid. WATER 97 Basic oxide + Water — >- Basic hydroxide Na^O + H,0 -> 2NaOH Sodium oxide. Sodium hydroxide. CaO + H,0 --> Ca(OH), Calcium oxide. Calcium hydroxide. The soluble basic oxides, commonly called alkalies, impart a soapy feel to the water, and turn red litmus blue. Other less soluble basic oxides, e.g. ZnO, FeO, react so slowly with water that the hydroxide is generally obtained from their salts by double decomposition. FeCla + 3NaOH-^Fe(OH)3| + 3NaCl. The fundamental difference in the nature of the product formed when water reacts upon an oxide has enabled the chemist to divide the elements with more or less sharpness into groups — metals and non-metals. Metals form basic oxides, non-m,etals give rise to acidic oxides. This classification is simple and agrees very weU with that given by the physical examination of the metals. At the same time, it should be borne in mind that a few oxides lie intermediate between the basic and acidic groups, so that the corresponding element cannot with certainty be placed either with the metals or with the non-metals, e.g. arsenic. A second characteristic of water is its tendency to cause hydrolysis, or a breaking down of the compound. This will be dealt with at a later stage. Water is extremely difficult to decompose by the action of heat. Even in the neighbourhood of 2000° the decomposition is only 1-76 per cent. The higher the temperature, the more is the reaction driven to the right H,0->2H + 0. This is why at high temperatures water acts as an oxidising agent, viz. converting iron into oxide of iron, for under these conditions the oxygen produced by the above dissociation is a more active oxidising agent than the hydrogen is a reducer. Another important property of water is its capacity for com- bining with compounds to form crystalline substances. The solutions of many salts, when evaporated, deposit well-defined crystals which on examination are found to contain water, the so-called water o/ crystallisation. This term has come into use H 98 AN INORGANIC CHEMISTRY owing to the fact that, when the water is expelled, the substance generally crumbles to a powder, i.e. effloresces. There is, however, no evidence that this leads to the formation of an amorphous substance — a term indicating the absence of crystal structure and applied to what are really supercooled Uquids, e.g. glass. These substances containing water of crystaUisation are known as hydrates. A hydrate obeys the laws of chemical com- bination, i.e. the substance unites with water in a definite pro- portion by weight. Examples of this are the decahydrate of sodium carbonate, Na2CO3,10H2O, and the pentahydrate of copper sulphate, CuS04,5H20. Each hydrate has its own specific physical properties. Anhydrous copper sulphate is a fine white powder, which may be obtained in colourless needlelike crystals. The pentahydrate occurs as large, blue, triclinic crystals, the trihydrate, CuSOajSHjO, forms bluish crystals, whilst the monohydrate CuS04,H20, appears as a bluish white, friable mass. The main effect of heating such hydrates is to expel part or all of the water of crystalHsation, and as the aqueous solutions formed from the various hydrates of copper sulphate and from the anhydrous salt are all chemically and physically indistin- guishable, it has become the custom to write the water of crystallisation in an easily detachable way, e.g. Na2CO3,10H2O, and not HjoNaoCOjs. When exposed to the atmosphere many hydrates lose their water of crystalUsation. This is some- what akin to the evaporation of water, and just as the water has been shown to evaporate by virtue of its being able to set up a definite pressure, and a measure of that pressure is afforded by the depression of the barometric column, so too, this method should give information concerning the pressure exerted by a salt containing water of crystallisation. If a crystal of the pentahydrate of copper sulphate is introduced above the mercury column of a barometric tube, the column wiU be depressed by an amount depending upon the temperature ; hence a vapour pressure curve for each hydrate may be constructed. Pig. 39 shows the temperature-vapour pressure curve for the hydrates of copper sulphate. At every temperature ea.ch hydrate has its own particular vapour pressure, provided, of course, that the temperature is such that the compound is capable of existence. We are now in a position to predict what will happen when such a hydrate is exposed to the atmosphere. WATER 99 Temperature. Fig. 39. It will endeavour to set up its own definite aqueous vapour pressure in its immediate neighbourhood. Should the surround- ing atmosphere possess a lower aqueous vapour pressure than does the hydrate, the crystal will continue to lose water vapour, and it will pass either into a lower hydrate or into the anhydrous substance. The aver- age vapour pressure of water in the atmo- sphere at 10° is a Uttle over 5 mm., so that hydrates possessing a higher vapour pressure than 5 mm. wiU gener- ally be found to de- compose spo n t a n e- ously when exposed in an open dish. Such a case is represented by Glauber's salt, NaaSO^.lOH^O. Suppose anhydrous copper sulphate is taken and aqueous vapour is admitted to the apparatus at some temperature maintained constant throughout the experiment — say 50°. Nothing happens till the vapour pressure rises to that at which the monohydrate can exist, viz. 4-5 mm. Then, and only then, does the monohydrate begin to form. Any attempt to raise the pressure above 4-5 mm. wiU fail, until the last trace of anhydrous copper sulphate has been converted into the monohydrate. When the conversion into the monohydrate is complete, the pressure of the water vapour may be further raised, and nothing wiU happen till the vapour pressiu-e of the next hydrate, CuSOi, SHjO, has been reached, viz. 30 mm. At that stage the pressure again becomes constant until every trace of monohydrate has been converted into the trihydrate. After that, the pressure may again be raised without any change in the state of the solid compound until a pressure of 47 mm. — the vapour pressure of the pentahydrate — is formed, and no further change in the vapour pressure of the system can be produced untU the reaction. CuS04,3H20 +2H20^CuS04,5H20 is complete. 100 AN INORGANIC CHEMISTRY The diagram in Fig. 40 shows these changes graphically. Many salts show a tendency to deliquesce. This process is the reverse of efflorescence, for should the aqueous vapour pressxire of water in the air exceed the vapour pressure of any hydrate exposed to the atmosphere, condensation of the water wUl take place upon the hydrate. Calcium chloride and sodium hydroxide 50 4-P 30 ZO 10 1 1 1 1 Form ila: We 'ght5 01 ^ water perm 'lecule oFcopt ^er suf vh=ite. 1 Z 3 V Fig. 40. are examples of substances which deliquesce. It is because the vapour pressure of sulphuric acid and calcium chloride is exceed- ingly low that these compounds are called into requisition for drying gases, i.e. to reduce the aqueous vapour pressure of the gas to a neghgible amount. Composition of Water by Weight.— The composition of this compound has been the subject of many important investi- gations, and among the earher of these measurements those of Dumas stand out (1842). His method is shown in Fig. 41. Hydrogen, most carefully purified from possible impurities WATER 101 by passage through various reagents, was passed over heated copper oxide, and the water collected and weighed. CuO + 2H->Cu + H20. The vessel A containing the dried copper oxide was weighed before and after the experiment. The difference in weight gave the amount of oxygen combined in the water which had been collected and dried in the drying tubes, and the ratio hydrogen/ oxygen could then be calculated. That the hydrogen used in the experiment was actually dry, was shown by weighing the pent- oxide tube B before and after the experiment. If no change of weight occurred, it was assumed that the hydrogen passing Fig. 41. through to the copper oxide was quite dry. The water generated in the reaction was largely collected in the receiver C, the residue being absorbed in the tube D containing sohd potassium hydroxide, and in the tube E filled with phosphorus pentoxide. The tube F containing this reagent served as a guard tube, the exit also dipping below the surface of concentrated sulphuric acid in 0. As the mean of many experiments it was found that in the formation of 236-36 gm. of water, the oxygen given up by the copper oxide was 210-04 gm., hence Water=236-36 gm. Oxygen=210-04 Hydrogen = 26-32 The ratio of hydrogen to oxygen was therefore 2 : 15-96. In 1887 Keiser pubhshed results indicating that the ratio 102 AN INORGANIC CHEMISTRY -A=#- obtained by Dumas was possibly somewhat low, the error in Dumas' experiments arising from the occlusion of hydrogen by the reduced copper. The effect of this is to make the weight of oxygen contained in the collected water lower than it should have been. In 1895 E. W. Morley re-determined this very important ratio by the direct weighing of the oxygen and hydrogen which enter into combin- ation, and such was the care and skill displayed that his experiments stand to-day not only as the most trustworthy on the composition of water, but also as a model of experimental skill. The hydrogen used in his experiments was obtained by heatiag palladium hydride under reduced pressure and the oxygen from potassium chlorate. The gases, after thorough drying, were stored in large globes. From the differ- ence in the weights of these globes before and after the experiment, Morley obtained the weights of the gases used in the experiment (Pig. 42). The gases were brought through the taps A and B, passing through the guard tubes containing phosphorus pentoxide, into the reaction vessel C. The rate at which the gases were led into the apparatus was carefully regulated so that a continuous formation of water could be produced by electric sparks across the terminals D. The water was condensed and collected in the lower part of the chamber, its condensation being acceler- ated by a surrounding jacket of ice-water. At the conclusion of the experiment the apparatus was placed in a freezing mixture in order to freeze the water in C and to remove as much of the aqueous vapour as possible. The residual mixture of oxygen and hydrogen was pumped out and analysed, the necessary correction then being made in the weights of the gases used as determined by the globe weighings. As a mean of eleven experiments Morley found : WATER 103 Hydrogen used . . . 3-7198 gm. Oxygen .... 29-5335 „ Water found . . . 33-2533 The ratio of hydrogen to oxygen is thus 2 : 15-879 or 2-016 : 16, a result agreeing very well with values obtained by Morley by other methods. Independent workers have since obtained exact agreement with Morley's results. Composition of Water by Volume. — This may be deter- mined either analytically or synthetically. The decomposition of water by the electric current has already been described iq.v.), and the results indicate that hydrogen and oxygen unite in the ratio of two parts by volume of hydrogen to one part by volume of oxygen. The composition may be demonstrated synthetically in the following way. Into the tube A (Fig. 43) is introduced, say, 20 c.c. of hydrogen and 15 c.c. of oxygen. By means of a small induction coil a spark may be sent through the gaseous mixture. The pressure in the reaction vessel is first reduced as much as possible by lower- ing the levelhng tube B. On sparking, there is a slight explosion followed by a marked diminution in the volume. When the pressure is adjusted, it is found that the residual volume is 5 c.c. Further examination proves that this consisted of oxygen (the steam condensing to water at the temperature of the room). In other words, 20 c.c. of hydrogen combine with 10 c.c. of oxygen, i.e. two volumes of hydrogen with one volume of oxygen. If the experiment is carried out at a temperature above 100°, the resultant volume of the steam may be determined. If the gaseous mixture consists of hydrogen and oxygen in the ratio of 2 : 1 by volume, after explosion the new volume wiU be two- thirds of the combined volume of oxygen and hydrogen, i.e. from three volumes of the reacting gases will be obtained two volumes of steam. The experiment can be neatly performed in the apparatus shown in Fig. 44. A sufficiently high temperature for the jacket can be obtained by the use of amyl alcohol, which boils at 130°. The fact that the volumes not only of the oxygen and hydrogen used in this reaction, but also of the resultant steam, may be 104 AN INORGANIC CHEMISTRY represented by small integers, led Gay Lussac to formulate his Law of Gaseous Volumes. Whenever gases unite, they do so in volumes which bear a simple ratio not only to each other, but also to the volume of any gaseous product of the reaction. The investigation of the effect of temperature and pressure changes upon the volume of a gas has led to the conclusion that Cold Water Inlet Fig. 43. Fio. 44. Boyle's Law and Charles' Law {q.v.) are approximately true, and do not exactly explain the phenomena which they purport to correlate. If, therefore, in the combination of hydrogen and oxygen to form steam, the gases are taken in the exact ratio of 2 : 1 measiu-ed at 130°, they would not be exactly in this ratio at any other temperature, as the coefficients of expansion of oxygen and hydrogen are not quite the same. It follows, there- WATER 105 fore, that if Gay Lussac's Law holds exactly at any one tem- perature, it will not do so at any other temperature — in short, the Law of Combination by Volumes is an approximate law, and not an exact Law of Nature. QtrESTIONS 1. How would you prove that the composition of water is expressed by the formula H2O ? 2. Indicate the essential differences between a hydrate, a solution and an hydroxide. 3. Which contains the greater store of internal energy — equivalent quantities of ice, water or steam, and why ? 4. Compare and contrast " physical " and " chemical " change. 5. Discuss the effect of pressure upon the boiling point of a liquid. 6. How would you prove that hydrates are true chemical compounds ? In the light of your statement would you consider a solution to be a chemical compoimd or a physical mixture ? 7. What weight of water is it possible to prepare from lib. of potassium chlorate and I oz. of zinc ? CHAPTER VIII SOLUTIONS— CRYSTALS— THERMOCHEMISTRY The formation of solutions has already been shown to depend not only upon the solvent but also upon the solute, e.g. mercuric chloride dissolves in alcohol but not in water, whilst cobalt chloride dissolves in both. But there are other factors which also exert an influence upon the formation of a solution — notably temperature and pressure. If a large crystal of copper sulphate is placed in the bottom of a tail cylinder and a column of water poured over it, there appears at first to be no action. However, after a few days, the layers of liquid in the immediate neighbourhood of the crystal take on a more or less blue colour, the depth of the colour shading off as the distance from the crystal increases. If the cylinder is placed in a position where it is not subjected to vibrations, it will be months before the solution is of a uniform concentration. If the experiment is repeated with hot water instead of cold, the speed at which the blue colour is disseminated throughout the liquid shows a marked increase. A rise of temperature undoubtedly accelerates the rate at which the copper sulphate dissolves in water, but does it increase the absolute amount of salt dissolving in a given amount of water ? The answer to this question is obtained by shaking up powdered copper sulphate in a stoppered flask filled with water until no further solution occurs. On raising the temperature, say 50°, it is at once evident that not only is the speed of solution accelerated, but there is also an increase in the absolute amount dissolved during the course of the experiment. On the other hand, if the copper sulphate is replaced by calcium acetate, less salt dissolves at the higher temperature than at the lower, though the increase in temperature certainly increases the rate of solution. The key that enables one to predict whether a substance will 106 SOLUTIONS 107 become more or less soluble as the temperature rises, is the Law of van't Hoff, which is really a special case of the General Prin- ciple of Le Chatelier, already enunciated (p. 58) . The process of solution is attended either by the absorption or the evolution of heat ; ammonium chloride, e.g., absorbs heat when dissolving, whilst calcium hydroxide evolves it. Suppose that a beaker contains a solution of ammonium chloride in contact with crystals of that substance, and suppose that the solution is saturated, i.e. the amount in solution is the maximum quantity which will dissolve in the solvent at this temperature. If the temperature of the solution is raised, that process must set in which will tend to restore the former equilibrium, i.e. which will absorb heat, consequently more solid ammonium chloride will dissolve. . In short, in order that a permanent rise in the temper- ature of the solution may be effected, solid always being present, the solution must become stronger. Substances which absorb heat on solution in their nearly saturated solution will, therefore, be more soluble with rising temperature ; conversely, substances which dissolve in their nearly saturated solution with the evolu- tion of heat wiU be rendered less soluble by a rise in temperature. Should a substance dissolve without the evolution or the absorp- tion of heat, temperature wUl have no effect upon its solubUity. This condition is nearly fulfilled by sodium chloride (Table 14). TABLE 14 Heat of solution. Solubility in 100 gm. of solvent. at 0° 40° 80° Sodium chloride Calcium acetate Calcium hydroxide Ammonium chloride . Potassium nitrate . -1-22 Cal. -f7-0 + 2-8 „ -4-0 „ -8-5 26-3 37-4 0-13 22-7 13-3 26-65 33-2 010 31-4 63-9 27-5 33-5 007 39-6 169 (For units of measurement of heat see p. 20.) Fig. 45 shows the solubility curves (temperature-concentration graphs) for a number of commonly occurring inorganic substances. Let us now investigate the solubility curve from another aspect. Take Fig. 46, the temperature-concentration curve for 108 AN INORGANIC CHEMISTRY ammonium acetate. A solution which has been saturated at a high temperature and which is in contact with soHd crystals, wiU continue to deposit crystals as the temperature falls. In other words, if we start from the point A, as the temperature 40' 50' 60° Fig. 45. 100° decreases, crystals of ammonium acetate wUl fall out of the solution and we shall pass along the curve ABC. Provided the solution is held at the one temperature sufficiently long for equilibrium to be established, the concentration at any one temperature may be read off from the curve. But let the SOLUTIONS 109 experiment be repeated with a saturated solution from which all crystals have been excluded. Under such conditions, it happens frequently that, as the temperature falls, no change whatever in the concentration of the solution occurs, i.e. we pass from the point A to the point D. If, however, a tiny crystal of ammonium acetate is introduced as a nucleus into a solution, the temperature and concentration of which are repre- sented by the point D, instantaneous crystalhsation with a considerable evolution of heat will at once set in, and the con- ditions of equilibrium as regards temperature and concentration will at once be given by a point on the curve ABG. All solutions whose temperature and concentration are represented in the area above the curve ABC will ex- hibit this phenome- non. Such solutions are known as super- saturated. Super- saturated solutions are in a state of metastability , which is at once destroyed by the introduction of a crystal of the solute. A somewhat similar phenomenon is shown in the supercooling of water below 0° without freezing setting in. This meta-stable state is destroyed by the introduc- tion of a spicule of ice. Careful investigation has shown that many solubility curves show sudden breaks or changes in direction. Such a case is ex- emplified by the deca-hydrate of sodium sulphate (Fig. 47). The direction of the solubility curve is, we have seen, fixed by the heat of solution, and it is little probable that this heat of solution will suddenly change its magnitude or its sign. Rather it is probable that a new phase or substance has come into existence which possesses a different heat of solution. This can easily be proved by analysing the solid phase in contact with the solution at temperatures above and below where the break in the curve occurs. In the above case analysis shows that above the temperature 33° where the break comes, the solid phase in contact Temperature. Fig. 46. no AN INORGANIC CHEMISTRY with the solution is no longer the deca-hydrate of sodium hydrate, but anhydrous sodium sulphate itself, even though the sohd added to the liquid is the deca-hydrate. Thus at 33° the reaction NaaSOi.lOHaO— ^NaaSOi+lOHaO sets in. The investiga- tion of the solubihty curve to see whether any well defined breaks occiu- is a valuable aid to the vapour-pressure method (p. 99) for investigating the existence of hydrates or compounds containing water of crystaUisation. The effect of pressure upon solutions of solids or Uquids is very slight. If the formation of a solution is attended by a volume decrease, the appHcation of Le Chateher's principle 60 50 40 30 10 WO' ^^ ^'-'• --, ^ ■^ t~,i ^ ^ ■u p k. - -— / <^ C ^/ l> ) ^' i .f / ,^ ^ Te mp :pa 'run 3. 10 20 30 -30 50 60 70 SO' 90' Fig. 47. — Soltjbility Curve of Sodium Sulphate. predicts that with increasing pressure the substance -will become more soluble, and conversely. Solubility of Gases. — The effect of pressure upon the solubihty of gases is, however, very marked. Suppose a gas is enclosed with a hquid in a cyHnder having a frictionless piston. Particles of the gas strike the surface of the hquid and some are retained by the hquid. Sooner or later the time must come when the number of such particles retained in this way is just coimterbalanced by the number which shp out again. The Hquid is then saturated with the gas. Now let a weight be placed upon the piston, thereby causing an increase in the pressure of SOLUTIONS 111 the gas. That process at once sets in which will tend to restore the pressure to its original value (Le Chatelier's law), i.e. more gas dissolves because, by so doing, the pressure of the gas is diminished. Thus the solubDity of a gas increases with the pressure. This law, connecting pressure and solubility of a gas, is named after Henry, its discoverer, and may be enunciated : At constant temperature the concentration of a saturated solution of a gas is proportional to the pressure at which the gas is supplied. Since, according to Boyle's Law, the volume of a gas is proportional to the pressure, the law may be stated : At constant temperature, a given volume of a liquid will absorb the same volume of a gas at all pressures. A suitable apparatus for testing the solubiUty of gases is shown in Fig. 48. The gas under investigation is admitted into the measuring tube A through the three-way tap b by first raising and then lowering B. The pressure of the gas in A is made equal to that of the atmo- sphere by levelling the mercury in A and B and the volume of gas in A is then read. The absorption vessel G is fiUed with the desired liquid by means of the three-way tap a and vessels A and C put into communication. By raising B and opening the tap c a volume of hquid is run out into the measuring vessel Z) and this represents the amount of gas transferred from A to C. C is shaken thoroughly, and the gas remaining in G transferred back to A by lowering B and opening the tap c under mercury. The levels of mercury in A and B are equahsed and the new volume of gas in A noted. The diminution in the volume of the gas gives the amount absorbed by the Uquid in G. Gases which react chemically with the solvent, e.g. carbon dioxide, ammonia, apparently do not obey Henry's Law. Fig. 48. 112 AN INORGANIC CHEMISTRY C02+H20->H2C03 (Carbonic acid) NH3+H2O— ^-NHiOH (Ammonium hydroxide) The following table shows the variation in the case of ammonia at 0° : TABLE 15 Partial pressure in mm. Weight of ammonia dissolved in 1 litre of water. Weight/p. 20 40 60 80 100 82 148 199 240 2S0 4-1 3-7 3-3 30 2-8 If the law were obeyed the ratio wt/p would be a constant. In this calculation the weight included the total quantity of dissolved gas (that present as NH4OH as weU as NH3). Experiment has shown that if the free ammonia gas only is included in the calculation the law of Henry is obeyed. Solutions of Liquids in Liquids. — -The laws governing the solubUity of solids and of gases in Kquids have been discussed, and it remains to consider the solubility of Uquids in liquids. In many cases it is possible to obtain complete miscibUity of two liquids, e.g. methyl and ethyl alcohol will mix in aU proportions with water to form homogeneous solutions. Often, however, it is found that, on shaking the Kquids together, two layers separate out, the upper one being a solution ot B in. A, the lower of .4 in B. Up to a hmited extent the one is completely soluble in the other, but when the proportion exceeds a certain value, the two layers separate out. This case is well shown by adding carbon disulphide, chloroform or ether very slowly to water. Partition or Distribution between Two Solvents. — -If, to a double layer of partially miscible liquids, a substance is added soluble in both, it wQl distribute itself between the two solvents in a manner proportional to its solubility in each solvent. Iodine in the presence of carbon disulphide and water or of chloro- form and water is an excellent illustration of this. The bulk of the iodine goes into the non-aqueous layer owing to its greater SOLUTIONS 113 solubility in that solvent. By repeated extractions iodine may be removed from an aqueous solution by shaking with small quantities of one of the above mentioned solvents. This is known as the process of extraction. Theoretically it is impossible to remove the iodine completely from the aqueous layer, but practically speaking, three or four extractions suffice to remove aU but the last traces of iodine from the aqueous layer. This arises from the low solubility of iodine in water and its high solubility in carbon disulphide or chloroform. Colloidal Solutions; — Some of the important kinetic properties of these solutions have already been discussed. These solutions, homogeneous only so long as they are not observed under a strong, converging beam of light, are prepared in various ways. If a Httle albumen is introduced into cold water, or a little powdered starch is treated with hot water and afterwards diluted with cold, the solution pro- duced is colloidal . A very simple method of pre- paring colloidal solutions of ■'^'°- *^- many metals, e.g. platinum, has been described. A small glass vessel (Fig. 49) is filled with distilled water ; two pieces of heavy platinum wire connected with the terminals of a 100 volt battery, are cautiously brought into contact below the surface of the water and the circuit rapidly broken. A small arc is formed and the solution rapidly turns brown. The process has to be repeated many times before the solution is sufficiently strong. Such a solution possesses in a marked degree the catalytic properties generally associated with platinum metal. As a rule, coUoidal solutions are easily coagulated by heat, as well as by various chemical reagents, but perhaps their most interesting feature is their non-homogeneity already discussed. A coUoid will pass through all but specially prepared filters. Definition of Solution. — Solutions are mixtures which appear homogeneous in ordinary daylight, and which cannot be separated I 114 AN INORGANIC CHEMISTRY by ordinary mechanical means into their constituent parts. In order to effect such a separation the physical state of one of the components must be changed. The concentration of a solution may be defined as the amount of the solute which may be dissolved by a given weight of the solvent. The concentration of a solution may be qualitatively referred to as dilute and as concentrated. A quantitative method for expressing the concentration of a solution in the terms of a suitable chemical unit will be dealt with in a subsequent chapter. Crystallisation of Salts from Solutions. — By mere evaporation of an unsaturated or saturated solution, the dissolved sub- stance may be ob- tained in the form of crystals ; so the cooling of solutions of diiferent sub- stances wiU induce crystalhsation t o set in. The un- saturated solution of the temperature and concentration represented by A Temperature. Fig. 50. (Fig. 50), wUl deposit crystals so soon as the temperature has fallen to such an extent that the horizontal AB cuts the curve DBC. Should a small quantity of a soluble impurity be present, it is improbable that the solubility of this impurity wUl be exceeded at any stage of the crystaUisation, so that this substance will be left in the mother hquor. If the crystals are redissolved and several times recrystallised, the amount of impurity, except in certain exceptional cases, left in the crystals will be reduced to a negligible amount. This is the process of purification by re-crystallisation. Separation of Two Salts by Crystallisation. — A mixture of two salts, such as sodium chloride and potassium chloride, may be separated by the process of repeated fractional crystal- lisation. At 100°, 100 gm. of water will dissolve 39-8 gm. CRYSTALS 115 of sodium chloride, and 56-7 gm. of potassium chloride. An examination of the solubility curves of these substances (Fig. 51) reveals the fact that the solubility of sodium chloride is almost entirely independent of the temperature (at 20° 36-(> gm. are dissolved), whilst temperature exerts a con- siderable influence upon the solubility of potassium chloride (at 20° 34 gm. are dissolved). Upon cooling down the saturated solution of these salts from 100° to 20°, about 4 gm. of the sodium salt will crystallise out and 22 gm. of the potassium salt. If the impure crystals are redissolved and recrystaUised many times, it is clear that we shall ultimately obtain pure 50 40 30 y t ^ ^' / ^ ^ y 6^ y JC-J*^ c>i Sx 1 g 3/ ^ 1 o yj. y' ^ y ^ odi um Ch on de _^ - •o . . ■ ^ '^ ^ / ^ Te vpt rat ur tL— 40' SO' Fia. 51. 60' 90° potassium chloride from the crystals, and if the mother Uquor is worked up, pure sodium chloride may also be separated. Crystals All crystalline substances of definite composition possess a definite crystalline structure which is characteristic of that substance. Although the crystals of the same substance appear to differ among themselves, it has been found that in reaHty the angles between similar faces of crystals of the same substance are exactly the same and are characteristic of that substance. The constancy of the interfacial angle is thus of paramount import- ance in comparing two crystals of similar form, for, however alike in shape and in size, the crystals are not of the same substance unless the interfacial angle is of the specific dimensions belonging tQ that compound, llfi AN INORGANIC CHEMISTRY Often, however, a compound has the property of crystallising in more than one crystaUine form, i.e. of exhibiting polymorphism. Sulphur, for example, gives rise to rhombic crystals below 95-5'' and to monooUnic crystals above that temperature. Such a sub- stance is dimorphous. A rarer occurrence is the crystallisation in three forms. Titanium dioxide is a trimorphous substance. Crystal Forms. — Although the variety of geometric forms met with in naturally occurring and in artificially prepared crystals is almost infinite, yet they can be classified in six systems according to their planes of symmetry. A plane of symmetry is an imagiimry plane dividing the crystal into two parts such that one part is the mirrored reflection of the other. The following are the six crystal forms : 1. The regular system. This is the most symmetrical system, possessing, as it does, nine planes of symmetry, three of which intersect at right angles to each other, the remaining six at angles of 60° Examples : sodium chloride, garnet, diamond, alum. Fig. 52. 2. The hexagonal system. Crystals possess seven planes of symmetry, one principal plane and six others at right angles to the principal plane. Examples : quartz, calcite, apatite. Fig. 53. 3. The quadratic or tetragonal system. Crystals possess five planes of symmetry, one principal plane and four other planes at right angles to the principal plane and intersecting each other at angles of 45°. Examples : zircon, tin, potassium ferrocyanide, nickel sulphate. Fig. 54. 4. The rhombic system. The crystals possess three planes of symmetry at right angles to each other. Examples : potassium permanganate, iodine, silver nitrate. Fig. 55. 5. The monocHnic system. The crystals possess one plane of symmetry. Examples : gypsum, sodium carbonate (deca- hydrate), tartaric acid. Fig. 56. 6. The triclinia system. The crystals possess no planes of symmetry. Examples : copper sulphate (pentahydrate), anorthite. Fig. 57. The essential difference between a crystaUine and a non- crystalKne substance is one of internal structure, of molecular arrangement. An amorphous or non-crystalline substance is simply a super-cooled hquid which has not taken on any definite CRYSTALS 117 crystalline form in the process of solidification. The properties of an amorphous substance are the same, whatever direction they are measured in. A crystal of glass cut in the form of a calcite crystal will react with hydrofluoric acid at a uniform rate all over the surface, but a crystal of calcite itself dissolves more rapidly in one direction than in another. The thermal conduc- tivity of quartz differs according as to whether the section is cut perpendicular to, or parallel to, the principal axis, whereas that of glass or of rock salt (regular system) is independent of the direction in which the section has been cut with reference to ,d^ Fig. 53. _.!■ Fro. 54. Fig. 55. FiQ. 56. Fig. 57. the axis. If a slab of glass or of rock salt is cut, covered with a layer of wax, and a heated point brought into contact with the surface, it is found that the wax melts so as to form a circle ; so too in the case of quartz provided the section is cut perpendicular to the principal axis, but if the section is cut parallel to this axis the wax melts in the form of an ellipse. Similarly, light passes through crystals of the regular system as well as through an amorphous substance (glass) at the same rate in aU directions, but a differential speed is observed in the case of calcite. When suspended in a suitable solvent, amor- phous substances and crystals of the regular system have been 118 AN INORGANIC CHEMISTRY found to dissolve regularly in aU directions, but crystals of other systems do not dissolve at a uniform rate all over the surface. Isomorphism. — In 1818 Mitscherlich discovered that the crystals of the arsenates and phosphates of potassium were so much alike as to be almost indistinguishable. He concluded that the same number of atoms combined in the same vnanner produces the same crystalline form ; the crystalline form is indepen- dent of the chemical nature of the atoms, and is determined solely by the number and mode of combination. The phosphorus present in potassium phosphate may be replaced by arsenic without effecting the crystaUine form. Such related compounds are known as isomorphs. Many such are now known, not only in the artificial, but also in the natural world, e.g. aragonite (CaCOa), witherite (BaCOa), strontianite (SrCOs), cerussite (PbCOa), all form isomorphous crystals (rhombic). Careful measurements have revealed that, closely though these sub- stances resemble each other in crystal form, there is a small but measurable difference in their interfacial angles. In order to establish the isomorphism of crystals, it is not sufficient merely to establish the similarity in their crystal form. This is, of course, especially true of the cubic system, where similarity of structure does not by any means go hand in hand with similarity of chemical composition. As further criteria of isomorphism are to be noted : 1. The power to form layer crystals or overgrowths. 2. The power to form a continuous series of mixed crystals which show a regularity of physical properties. As an example of the former may be mentioned the power of chrome alum to induce crystallisation in a saturated solution of potassium alum, of zinc sulphate to form an overgrowth of green nickel sulphate. With respect to the formation of mixed crystals, not only should it be possible to form a continuous series of mixed crystals varying in composition from 100 per cent. A to 100 per cent. B, but the physical properties of these crystals must likewise show a steady change as the composition changes from 100 per cent. A to 100 per cent. B. Theemochbmistey The solution of a substance, we have learnt, may be accom- panied either by an evolution or by an absorption of heat, and THERMOCHEMISTRY 119 so with chemical reactions. A piece of sulphur, when burnt in air, evolves heat, and the product of the oxidation, sulphur dioxide, evolves heat when it combines with water or with sodium hydroxide. When the sulphur was oxidised to sulphur dioxide, there was a degradation of the chemical energy stored in the sulphur and the oxygen, and this was manifested in the hberation of a certain amount of heat. The chemical energy of the product formed differs from that of the original elements. This is a general statement — Every chemical reaction in- volves a change, not only in the form of the matter, but also in the energy content of the system. The unit of heat in physical work is the calorie — the amount of heat required to raise one gram of water one degree. In chemistry, however, we do not compare equal weights of sub- stances but formula weights (molecular weights). The heat oj jormation of sulphur dioxide is the heat evolved, measured in calories, when 32-07 gm. of sulphur combines with 32 gm. of oxygen to produce 64-07 gm. of sulphur dioxide. In practice, the calorie of the physicist is inconveniently small, and the kilogram calorie (Cal.) is often adopted — the amount of heat required to raise the temperature of one kilogram of water one degree. S03 + H20->H2SOi +21320 cal. or 21-32 Cal. Seeing that the conversion of sohd to liquid, of liquid to gas, is always associated with a definite heat transference, it is necessary to indicate the state of aggregation, not only of the reacting substances, but also of the products of the reaction, as in the equation, H2 + O -> H2O+59-4 Cals. Gas. Gas. Gas. H2 + O -> H2O + 68-4 Cals. Gas. Gas. Liq. i.e. the condensation of 18 gm. of steam to hquid sets free 9 Cals. The heats of reaction or of solution are readily determined by measuring the heat evolved when measured quantities of the reacting substances are brought into reaction in a suitably designed calorimeter, in which the rise of temperature produced in a known mass of wat«r by these known weights of the reacting substances is measured with an accurate thermometer. Measure- ments made in this way show that in some reactions heat is 120 AN INORGANIC CHEMISTRY evolved {exothermal reactions), whilst in others heat is absorbed {endothermal reactions). The Law of Conservation of Energy {qA>.) predicts that the heat of formation of every substance is numerically equal to its heat of decomposition, but of opposite sign. Ba + O^BaO + 124-4 Cal. BaO — 124-4 Cal. ^ Ba + O. This does not mean that every reaction is capable of reversal if the necessary quantity of heat is at hand. The heat must be supphed at a sufficiently high temperature in order that it may be absorbed. To take an analogous case — a large amount of water in a tank ; the amount of work available depends not only upon the weight of water stored, but also upon the height above sea-level. Generally speaking, if the requisite heat is supplied electrically, the substance will decompose, and in so doing, wiU absorb an amount of heat equal to its heat of formation. As a further deduction from the Law of Conservation of Energy, it follows that the amount of heat evolved during the formation of a given compound is the same, whether the sub- stance is formed directly or in a series of intermediate stages. The heat of formation, therefore, depends only upon the initial and final stages of the system. This is Hess's Law of Con- stant Heat Summation. The following equations exemplify this : — C + O^^CO +26 Cals. CO + 0->C02 + 68 Cals. 0+20-^00^+94 Cals. that is, the direct combination of carbon and oxygen to form carbon dioxide leads to the evolution of 94 Cals. The same amount of heat is evolved if the carbon is first oxidised into carbon monoxide, and the monoxide then converted into dioxide. Many experiments confirm this law. Questions 1. Construct the solubility curve of zinc sulphate from the following data : — 100 gm. of water dissolve the following weights of this salt at the various temperatures : Temperature .... 0° 25° 50° 70° 80° 90° Weight in grams . . 41-9 58 76-8 88-7 86-6 83-7 THERMOCHEMISTRY 121 2. What reason would you assign for the greater solubility of a gaa with rising pressure ? 3. The solubility of pure oxygen in water at 20° is three volumes per 100 volumes. What volume of oxygen would dissolve in » litre of water from the air at 20° and under a pressure of 10 atmospheres ? 4. Give a brief account of the law of Isomorphism. 5. Ice and water are in equilibrium with each other in a tube. What happens if the pressure upon the system is increased (apply the Le Chate- lier principle). 6. On dissolving 100 gm. of anhydrous copper sulphate in water, 9,900 cals. are evolved, whilst the same weight of crystallised copper sulphate (CuSOj.SHjO) gives an absorption of 1,100 cals.; determine the heat of hydration of the anhydrous salt into the penta-hydrate. [The heat of hydration refers to the heat generated when the molecular weight expressed in grams is hydrated, in this case 63-6 + 32 + 64 gm. of CuSO,.] CHAPTER IX MOLECULAR AND ATOMIC WEIGHTS Avogadro's Hypothesis. — The Law of Gay Lussac govern- ing the combination of gases by volume rapidly led to most important results. It had been recognised for some time fDalton's Atomic Hypothesis ; 1802-1805) that the atoms of the elements unite in simple proportions, one to one, etc., whilst the new law of Gay Lussac emphasised the fact that gaseous elements unite in simple proportions by volume. The atomic structure of matter, viewed in conjunction with the experimental law of volumes discovered by Gay Lussac, leads to the inference that the number of atoms, contained in equal volumes of gases, must be in the ratio of simple whole numbers. The simplest assumption is, of course, that equal volumes of gases contain equal numbers of atoms, but a consideration of the reaction between hydrogen and oxygen renders this view untenable. 2H + -^ H2O 2 vols. 1 vol. 2 vols. Suppose that each volume contains n atoms, then : 2n atoms H + m atoms — >■ In atoms H 2O i.e. two atoms of hydrogen must react with one atom of oxygen to produce two compound atoms of steam. As each atom of steam must contain an oxygen atom, it follows that the oxygen atom must be chemically divisible — a conclusion in direct contradiction of the atomic theory, which postulates the non- divisibility of the chemical atom. Similarly in the reaction H -f a. ->HC1 1 vol. 1 vol. 2 vols. every atom of hydrogen and of chlorine must be capable of 122 MOLECULAR AND ATOMIC WEIGHTS 123 division into two parts in order that 2n atoms of the compound gas may be formed. FuU reconciliation between the demands of Dalton's Atomic Theory and of the Law of Combining Volumes was not brought about till the time of Avogadro (1811), a brilUant ItaHan chemist, who was led to the conclusion that the particles present in a gas, whether simple or compound, are really aggregates of atoms, which are capable of subdivision when reaction takes place. These aggregates he named molecules. The essential postulates of Avogadro 's Hypothesis are : — 1. Molecules of an elementary gas are made up of two or more atoms of the same element. 2. Compound molecules are made up of two or more atoms of different elements. 3. The molecule of an element, when reacting with the mole- cule of another element, splits into parts (atoms). 4. Equal volumes of all gases under the same conditions of temperature and pressure contain the same number of molecules. 5. The relative weights of the gaseous molecules can be determined by comparing the weights of equal volumes of the gases, i.e. by comparing their gaseous densities. This conception of Avogadro that gases consist not of atoms but of molecules which are capable of subdivision at the moment of reaction, and that equal volumes of all gases contain, under the same conditions of temperature and pressure, the same number of molecules, has had most fruitful results in system- atising the science of chemistry. Gaseous reactions such as : H + CI. - ■> 2HC1 1 vol. 1 vol. 2 vols. 2H + - > H,0 2 vols. 1 vol. 2 vols. 3H + N — > NH3 3 vols. 1 vol. 2 vols. become at once easy of interpretation. In the reaction between hydrogen and chlorine n molecules of hydrogen unite with n molecules of chlorine to form 2n molecules of hydrogen chloride. H + CI -^ HCl n mols. n mols. 2n mols. i.e. 1 mol. 1 mol. 2 mols. Consequently each molecule of hydrogen and of chlorine must 124 AN INORGANIC CHEMISTRY subdivide into two atoms at the moment of reaction. An investigation of other reactions in which these gases take part confirms the inference that the molecules of hydrogen and chlorine are di- atomic. The equation may then be written in a molecular way H, + C1,->2HC1 The reaction between hydrogen and oxygen to form steam is capable of similar interpretation. 2H + -> H2O 2 vols. 1 vol. 2 vols. 2n mols. n mol. 2n mols. i.e. 2 mols. + 1 mol. 2 mols. This at once suggests the divisibiUty of the oxygen molecule into two atoms — a conclusion in complete agreement with the results obtained from other reactions in which oxygen partici- pates. The equation, written molecularly, then becomes : 2H2 + 02-^2H20. So, too, in the more difficult reaction between hydrogen and nitrogen to form ammonia. N + 3H ^ > NH, 1 vol. 3 vols. 2 vols. n mols. 3 TO mols. 2 n mols. 1 mol. + 3 mols. 2 mols. Assuming the di-atomicity of nitrogen, the equation becomes N2 + 3H,^2NH3 The Relative Weights of the Molecules — Molecular Weights. — ^This fruitful and suggestive hypothesis of Avogadro has played an invaluable part in the development of theoretical chemistry, but in no branch has its influence been so boundless as in moulding our conception of molecular weights. The actual physical measurement of the weight of a molecule has long been recognised as being beyond the achievement of the scientist, but a consideration of the above hypothesis in all its bearings leads to the conclusion that the relative weights of the molecules of gases may be readily determined by a comparison of the weights of equal volumes of the gases. Suppose that a certain volume V, containing n molecules of oxygen, is taken. It follows that the volume V of any other gas MOLECULAR AND ATOMIC WEIGHTS 125 measured at the same temperature and pressure, will also contain n molecules of the gas. If the weight of this volume V of the different gases is determined, it follows that weight of volume V of oxygen _weight of n molecules of oxygen weight of vol. V of any other gas weight of n molecules of any other gas _ weight of one molecule of oxygen weight of one molecule of any other gas. In other words, if oxygen is chosen as the standard gas, and any arbitrary weight assigned to its molecule, the relative weight of a molecule of any other gas may be determined by comparing the weights of equal volumes of oxygen and the gas in question, i.e. by comparing their densities. Owing to hydrogen being the Ughtest known gas, this element was imtil recent years selected as the standard gas. Since hydrogen is diatomic, its molecular weight was taken as 2. In this way the following table of molecular weights has been constructed. In column 3 the density of the standard gas is recorded as 1, and in the last column its molecular weight as 2. TABLE 16 Weight of 1 litre. Density. Molecular weight. N.T.P. H=l. H=2. Hydrogen .... 0-090 1 2 Oxygen 1-429 15-88 31-76 Chlorine 3-166 35-18 70-36 Hydrogen chloride . 1-639 18-21 36-42 Carbon dioxide . 1-977 21-96 43-92 Water 0-8045 8-939 17-78 But for reasons which wUl soon be discussed more fully, the hydrogen standard has in recent years been replaced by the oxygen standard, the molecular weight of oxygen being assumed to be 32. The following table is thus obtaiued : 126 AN INORGANIC CHEMISTRY TABLE 16a Hydrogen . Oxygen . . . Nitrogen . Hydrogen chloride Carbon dioxide Sulphur dioxide . Hydrogen sulphide Water .... Weight of 1 litre. N.T.P. 0-08987 1-429 1-251 1-6398 1-977 2-9266 1-537 0-8045 Observed molecular weight. = 32. 2-016 32-00 28-00 36-46 44-00 6406 34-07 18-016 The values recorded in the last column are obtained from the second by multiplying by the ratio 32/1-429. If the gram is the unit of weight employed, the molecular weight is generally referred to as the gram -molecular weight, or, in short, the molar weight. Gram Molecular Volume. — The weight of one htre of the standard gas (oxygen) at 0° and 760 mm. is 1-429 gm. The volume of oxygen which weighs 32 gm. is thus 32/1-429, i.e. 22-39 litres, or roughly 22-4 litres. This volume is known as the gram-molecular volume — the volume which contains the gram-molecular weight of any gas at 0° and 760 mm. Conse- quently the following rule for finding the gram-molecular weight of gas may be stated : Find the weight of a known volume of the substances at any temperature and pressure at which it is gaseous ; reduce this volume to standard temperature and pressure (0° and 760 mm.) and calculate by proportion the weight of substance which will fill 22-4 Utres under the same conditions of temperature and pressure. The Adjustment of Molecular Weights. — The above method of obtaining the molecular weight — invaluable as it is — yet leads to a more or less approximate value. First of all, the measurement of the density of gas by any of the usual methods employed for determining this constant (q.v. Chapter V), generally involves an error of the order 0-1-0-5 per cent., and secondly, owing to the deviations shown by aU gases to the gas laws, all results obtained by the application of these laws are bound to be of an approximate nature. Avogadro's Law is in itself an approximation, owing to the deviations in the MOLECULAR AND ATOMIC WEIGHTS 127 behaviom- of gases from that of an ideal or perfect gas to which alone the gas laws rigidly apply, so that all deductions as to the molecular weight based upon the application of Avogadro's Hypothesis must have only an approximate value. In this case, however, this is no great hardship, as there is at hand a method for adjusting the molecular weight, provided we know the approximate value of that molecular weight. This adjust- ment depends upon our knowledge of the combining weights of the various elements. The actual value of these combining weights is capable of very exact measurement at the hands of skilled analysts, and they are known to a degree of accuracy far exceeding the accuracy with which a vapour density (i.e. molecular weight) determination can be made. The valency of an element has been defined as the ratio of its atomic weight to its combining weight (see p. 87), hence A.W.=?iC.W. But the molecule is composed of a definite number of atoms, so that the molecular weight =mA.W. =mnG.W. If the integer mn can be determined, it is obvious that the molecular weight may be accurately obtained from the com- bining weight. An approximate value of the molecular weight as given by the vapour density method, coupled with the knowledge of the combining weight, suffices to fix the value of mn, so that the adjusted molecular weight of the element may then be at once obtained from the combining weight. In the case of compounds that are capable of vaporisation the usual vapour density measurement, e.g. that of Victor Meyer, gives the approximate value of the molecular weight. The compound must then be accurately analysed, and the results expressed on a percentage basis. As an example, the compound phosphorus fluoride will serve. Analysis shows that this compound contains 75-40 per cent, of fluorine and 24-60 per cent, of phosphorus, while the unadjusted molecular weight, as determined by the measurement of its vapour density, was 126-46. The next step is to distribute the gram-molecular weight between the two elements in the ratio of 75-40/24-60, i.e. 7K.40 p=!^X 126-46=95-35 24-fiO P=^L^X 126-46=31-11. The combining weight of fluorine is known to be 19, of phos- 128 AN INORGANIC CHEMISTRY 95-35 phorus 6-2, so that mn for fluorine is 3111 19 =5 nearly, whilst for phosphorus mn — 6-2 5 nearly. Hence the adjusted mole- cular weight is 5x19+5x6-2 = 126. The following table of molecular weights of elements and compounds has been obtained by this method : TABLE 17 Weight o£ 1 litre. N.T.P. Hydrogen 0-08987 Oxygen i 1-429 Nitrogen [ 1-2507 Chlorine ! 3-220 Hydrogen chloride .... 1-6398 Carbon dioxide ' 1-9768 Hydrogen sulphide. ... I 1-537 Ammonia 0-7708 Phosphorus trichloride . . 6- 1138 Phosphine 1-5446 Methane 0-7128 Acetic Acid 2-6849 Mercury 8-870 Mercm-ic chloride .... 12039 Molecular weight obtained. 0=32 2-012 32 28-07 72-01 36-96 44-27 34-43 17-26 136-95 34-6 15-96 60-14 198-3 269-5 Molecular weight adjusted. 2-016 32 28-08 70-90 36-46 44-00 34-07 17-06 137-35 34-06 16-03 60-03 200-0 270-9 It is to be noted that the densities of several substances recorded above are purely imaginary, that is, some of the substances do not exist as gases under the conditions of temperature and pressure specified, e.g. acetic acid and mercuric chloride are solids at 0° and 760 mm. The actual densities of such substances are obtained under such conditions of temperature and pressure that they are gaseous, and by the appUcation of the gas laws the density at 0° and 760 mm. calculated, on the assumption that no physical change in the substance occurs. Other Methods of Determining Molecular Weight.— The above method of obtaining this most important chemical con- stant is limited to those elements and compounds the vapour density of which can be determined under suitable laboratory conditions. Many compounds exist wherein this method fails. Fortunately there lie at hand methods for determining the molecular weight of all dissolved substances. The principle MOLECULAR AND ATOMIC WEIGHTS 129 and application of these methods will be discussed at a later stage (see Chapter xxvii). The Atomic Weight and its Detebmination In the following table, the second column gives the adjusted molecular weight, and in columns 3-8 are recorded the actual amounts of the various elements present in the gram-molecular weight, these weights being calculated from a knowledge of the molecular weight, of the chemical composition of the compounds, and of the combining weights of the elements concerned, in accordance with the method aheady discussed for phosphorus pentafluoride. In column 9 is recorded the sum of the amounts of the various elements present in the compound, i.e. the mole- cular weight. TABLE 18 Substance. Mol. wt. Hydro- gen. Chlor- ine. Car- bon. Sul- phur. Phos- phorus. Oxy- gen. Total. Hydrogen chloride . 36-45f 1-008 35-45 36-45 Phosphorus trichloride . 137-39 — 106-35 — — 31-04 137-39 Phosphorus pentachloride 208-39 — 177-25 — — 31-04 — 208-39 Phosphorus oxychloride . 153-39 — 106-35 — — 31-04 16 153-39 Phosphine ... 34-064 3-024 — — 31-04 34-064 Carbon tetrachloride . 153-8 — 141-8 12 — — — 153-8 Ozone 48 — — — — 48 48 Water 18-016 2-016 — — — — 16 18-016 Acetylene 26-016 2-016 1 — 24 — — — 26-016 Carbon disulphide 76-14 — ; — 12 64-14 — — 76-14 Sulphur dioxide . 64-07 — , — — 32-07 — 32 B4-07 Sulphur trioxide . 80-07 — — — 32-07 — 48 80-07 Hydrogen sulphide 34-086 2-016 i — — 32-07 — — 34-086 CarTjon monoxide 28 1 12 — — 10 28 Carbon dioxide . 44 — 12 — — 32 44 Ethylene 28-032 4-032 — ( 24 — — — 28-032 A cursory glance at this table reveals the fact that all the compounds containing carbon contain 12n parts of this element in the gram-molecular weight, n being an integer. The smallest value of w is unity. If the number of volatile compounds of carbon were indefinitely extended, it would still be found that in no case does less than 12 parts by weight of this element occur in the gram-molecular quantity. Similarly, 35'45 repre- sents the smallest quantity of chlorine occurring in the gram- molecular weight of any of the volatile compounds of chlorine, K 130 AN INORGANIC CHEMISTRY 31 04 for phosphorus, 3207 for sulphur, and 16 for oxygen. These numbers, obtained in this way, are chosen as the funda- mental chemical units, i.e. atomic weights, of these elements. The least amount of an element, present in the mole- cular weight of all its known volatile compounds, is chosen as the atomic weight of that element. The atomic weight is always a multiple of the cornbining weight. Summarising, Atomic Weight = w X combining weight Molecular weight = m X atomic weight where the integers n and m may have any value from 1 onwards. Molecular weights and atomic weights are obtained by calcula- tion from actual experimental data. These quantities are quite independent of the present day theory of matter, these con- ceptions " atomic weight," " molecular weight," stand upon experimental facts, and are not concerned with the existence, real or other^vise, of molecules and atoms. Atomic Weights of the Elements.— Table (19 page 131) contains the accepted values of the atomic weights of the elements, as approved by the International Atomic Weights Committee. Atomic Weights versus Combining Weights. — In the section dealing with combining weights (page 34), attention was drawn to the behaviour of certain elements which form more than one oxide. The combining weight of these elements is therefore not a fixed quantity, as each oxide will give rise to a separate combining weight. Thus the two oxides of copper, CujO and CuO, lead to the values 63-57 and 31-79 respectively for the combining weight of copper. This variability in the combining or equivalent weight of an element has prevented these weights being chosen as the unit of weight for the particular element. No such variation occurs in the atomic weight. For each element this is a definite constant. An alteration in an accepted atomic weight can only be brought about either by the discovery of a new volatile compound in which a smaller amount of the element occurs in the molecular weight than has hitherto been accepted, or by the discovery of a sUght error in the chemical analysis upon which the percentage composition is based ; the former discovery would lead to the choice of an atomic weight which MOLECULAR AND ATOMIC WEIGHTS 131 TABLE 19 Atomic Atomic Element. Symbol. weight. 0=16 Element. Symbol. weight. 0=16 Aluminium . Al 27-1 Molybdenimi . Mo 96-0 Antimony . Sb 120-2 Neodymium Nd 144-3 Argon . A 39-9 Neon Ne 20-2 Arsenic . As 74-96 Nickel . . . Ni 58-68 Barium . Ba 137-37 Niton (radiutti Bismuth Bi 208-0 emanation) . Nt 224-4 Boron . B 10-9 Nitrogen N 14-01 Bromine Br 79-92 Osmium . Os 190-9 Cadmium . Cd 112-40 Oxygen . O 16-00 Caesium . Cs 132-81 Palladium . Pd 106-7 Calcium Ca 40-07 Phosphorus . P 31-04 Carbon . C 12-00 Platinum Pt 195-2 Cerium . Ce 140-25 Potassium . K 39-10 Chlorine . CI 35-46 Praseodymium . Pr 140-9 Chromium . Cr 52-0 Kadium Ra 226-0 Cobalt . . . Co 58-97 Rhodium Rh 102-9 Columbium Cb 93-1 Rubidium . Rb 85-45 Copper . Cu 63-57 Ruthenium . Ru 101-7 Dysprosium Dy 162-5 Samarium . Sa 150-4 Erbium . Er 167-7 Scandium . So 45-1 Europium . Eu 152-0 Selenium Se 79-2 Fluorine F 19-0 Silicon . Si 28-3 Gadolinium . Gd 157-3 Silver . . Ag 107-88 Gallium Ga 70-1 1 Sodium . Na 23-00 Germanium . Ge 72-5 Strontium . Sr 87-63 Glucinum (beryl- Sulphur . . S 32-06 lium) . Gl 9-1 Tantalum . Ta 181-5 Gold. . . . Au 197-2 Tellurium . Te 127-5 Helium . He 4-00 Terbium Tb 159-2 Holmium Ho 163-5 Thallium . . Tl 204-0 Hydrogen . H 1-008 Thorium Th 232-15 Indiimi . In 114-8 Thulium . . Tm 168-5 Iodine . I 126-92 Tin ... Sn 118-7 Iridium . Ir 193-1 Titanium . Ti 48-1 Iron .... Fe 55-84 Tungsten W 184-0 Krypton Kr 82-92 Uranium U 238-2 Lanthanum . La 1390 Vanadium . V 51-0 Lead. Pb 207-2 Xenon . Xe 130-2 Lithium Li 6-94 Ytterbium (neo- Lutecium . Lu 175-0 ytterbium) . Yb 173-5 Magnesium . Mg 24-32 Yttrium. Yt 89-33 Manganese . Mn 54-93 Zinc . . . . Zn 65-37 Mercury. Hg 200-6 Zirconium . Zr 1 90-6 132 AN INORGANIC CHEMISTRY is a submultiple of the previous one, the latter to a sHght correction in the accepted value. Moreover, the great generahsa- tion, known as MendeleefE's Periodic Law, which has done more than all else to systematise the study of chemistry, is based upon our knowledge of atomic weights, not equivalent weights. Note. — See Chapter xli. with reference to recent investigations in radio-activity which lead to a modification of current ideas on the invari- ability of atomic weights. Choice of Oxygen as the Standard. — The necessity for the choice of an element as standard has already been stressed. We cannot weigh an atom or even a molecule, so that atomic and molecular weights are relative, not absolute. The weight of a certain arbitrary volume of a gas containing an unknown but constant number of molecules has been chosen as the molecular weight. Long ago, Stas (1860-1865) pointed out that, in all atomic weight determinations, the element whose atomic weight was required should be combined with the standard element. Hydrogen combines with few elements, oxygen with nearly all. This property of oxygen in itself far outweighs the advantage of choosing as the standard the element of lowest atomic weight, and yet until very recent times hydrogen was the accepted standard. But as the atomic weights of nearly all elements were obtained by the analysis of oxy-compounds, it follows that, if the ratio of hydrogen to oxygen, as determined by Dumas (1842) and accepted by the chemists of the day, were incorrect, an alteration in nearly all the atomic weights would be necessary for every fresh determination of the hydrogen- oxygen ratio. Such a position actually arose when Morley (1895) showed that the Dumas ratio of hydrogen to oxygen was incorrect. In order to prevent such a disturbance in the accepted atomic weights, the majority of chemists thought it advisable to choose oxygen as the standard element (Molecular weight=32, Atomic weight=16). Any alteration in the accepted value for the ratio of hydrogen to oxygen, obtained from the analysis or synthesis of water, automatically brings about an alteration in the atomic weight of one element alone — ^hydrogen. The arbitrary nature of the present standard is revealed by the perusal of the following facts : T. Thomson (1825) used oxygen=l as standard ; W. H. WoUaston (1814) oxygen=lO ; J. J, Berzeljus (1830) oxygen=100. MOLECULAR AND ATOMIC WEIGHTS 133 Other Methods of Obtaining the Atomic Weight.— In 1818 Dulong and Petit made an important discovery in their study of the specific heats of solid elements. They found that the product of the specific heat and the atomic weight of an element is approximately a constant, the average numerical value of this constant being 6-4. TABLE 20 Element. Lithium . Sodium . Gold . . . Copper . Bismuth . Calcium . Bromine (solid) Iron . Mercury . Uraniimi Atomic weight. Specific iieat. 6-94 0-941 23 0-29 197-2 0-0304 63-57 0-0923 208 0-0305 40 0-170 80 0-084 55-8 0-11 200 0-0335 238-5 0-0276 6-5 6-7 6-3 5-9 6-3 6-8 6-7 6-1 6-7 6-6 Although the atomic weight ranges from that of the lightest known metal, lithium, to the heaviest, uranium, the so-called Atomic Heat maintains a nearly constant value. A few con- spicuous exceptions stand out — notably carbon, glucinum, boron and silicon — all elements of low atomic weight. TABLE 21 Element. Atomic weight. Atomic heat. Glucinum . 9 3-7 Boron . 11 2-8 Carbon 12 1-7 Silicon . . 28-4 4-5 But even in these cases it has been shown that the atomic heat tends to rise with increasing temperature, e.g. the specific heat of carbon in the neighbourhood of 1,000 is 0-45, giving the value 5-5 for the atomic heat. The specific heat of an element is the amount of heat required to raise one gram of the element one degree in temperature. It follows, therefore, that oil the average, 6-4 calories will raise the atomic weight in grams of an element one degree. In other 134 AN INORGANIC CHEMISTRY words, the atomic weights of all elements have equal capacities for heat. . Dulong and Petit's Law has occasionally been used to bring about an adjustment of the atomic weight of an element when other methods have failed. Such a case was presented by the element, indium. The equivalent of this element, as determined by the analysis of the chloride was 37-8, i.e. 37-8 parts of the metal combined with 35-5 of chlorine. If the chloride had the formula InClo, the atomic weight would be 75-6 ; but certain analogies with aluminium suggested that the formula might be InClj, in which case the atomic weight would be 3 x37-8=113-4. The specific heat of the metal proved to be 0-057. Assuming 75-6 to be the atomic weight, the atomic heat would be 4-5, whilst it the atomic weight were 113-4, the atomic heat would be 6-4 — a result in agreement with the demands of Dulong and Petit's Law. The atomic weight 113-4 was therefore adopted as the atomic weight of the element. Occasionally MitscherUch's Law of Isomorphism (v. p. 118) has been appealed to, in order to settle the atomic weight of an element. If two or more isomorphous compounds are taken, provided that the atomic weight of one member of the isomor- phous series is known, the others may be determined. The following table represents the composition of the isomorphous sulphate and selenate of potassium : — Potassium sulphate. | Potassium Eelenate. 100 parts contain Potassium . . 44-83 Oxygen . . . 36-78 Sulphur . . 18-39 10000 100 parts contain or 127-01 parts contain Potassium. . 35-29 Oxygen . . 28-96 Selenium . . 35-75 10000 Potassium. . 44-83 Oxygen . . 36-78 Selenium . . 45 40 12701 Assuming that sulphur has an atomic weight of 32, that a molecule of the sulphate contains one atom of sulphur, and that the selenate is similarly constituted, then A.W. of sulphur : A.W. of selenium=18-39 : 45-40 45-40 x32 hence atomic weight of selenium= =79-00 18*39 Molecular Weights and Chemical Formulae.— Although MOLECULAR AND ATOMIC WEIGHTS 135 the empirical formula of a compound may be readily computed from data obtained by accurate analysis and from the atomic weights of the elements concerned, it is impossible to state definitely what the exact formula is, until the molecular weight is known. Example — A certain compound contains 84-24 per cent, of sul- phur and 15-76 per cent, of carbon. Its molecular weight, as deter- mined by the Victor Meyer method is 74-9. What is the formula of the compound (atomic weight of carbon =12, sulphur=32-07) ? The compound must contain =:^^^— sulphur atoms and ,^ ^ 32-07 12 carbon atoms, that is, 2-6 sulphur atoms 1-3 carbon atoms, hence there must be two sulphur atoms to every one carbon atom. The simplest formula the compound can have, would be CSj, but the formulae C2S4, C3S4 . . . (CSa)^ are still possible. The molecular weight, however, is found by the density method to be 74-9. The formula CSg leads to a molecular weight of 12+64-14=76-14, consequently n = l and the compound must have the formula CSj. Should the substance be non-volatile, and should its molecular weight not be determinable by any of the methods described in Chapter xxvii., one cannot determine the value of n. In such cases the simplest possible formula which is generally referred to as empirical is chosen. The Atomic Theory. — The foundation upon which the four laws of chemical combination stand, is of a purely experi- mental nature, and in no way hypothetical. But when the imagination of the chemist is called into play in order to picture how it comes about that chemical compounds are formed by the union of carefully adjusted quantities of the reacting substances, irrespective of the particular temperature of the experiment, when the attempt is made to visuahse the mechanism of chemical reaction, and to explain what predetermines the particular imit weight of each element which shall react in all its combinations, the field of pure speculation is entered, and the Atomic Hypothesis is the result of that speculation. Far back in the time of the ancient philosophers the discrete nature of matter was postulated, and this idea has persisted ever since. But it was left to John Dalton (1801) to give the 136 AN INORGANIC CHEMISTRY Atomic Hypothesis life. In a very striking paper concerning atomic structm-e, he was led to the following postulates : 1. Atoms are real, discrete particles of matter which cannot be subdivided by any known chemical process. 2. Atoms of the elements are indestructible. 3. Atoms of the same element are similar to each other and of equal weight. 4. Atoms of different elements have different properties — weight, etc. 5. Compounds are formed by the union of atoms of different elements in simple numerical proportions, 1:1, 1 : 2, 2 : 1, etc. 6. The combining weights of the elements represent the combining weights of the atoms. The confusion existing in the mind of Dalton concerning atomic weight and combining weight, shown in the last postulate, prevented him from arriving at a system of atomic weights such as at present in vogue, and his assumption concerning the non-divisibility of the atom has required modification in view of the discoveries of atomic disintegration in radio-active phenomena, otherwise the Atomic Theory of Dalton remains as potent a factor in explaining chemical phenomena as in the days of its birth. A chemical reaction, such as the union of hydrogen and oxygen, in the light of this hypothesis is considered to be a reaction between the atoms of the reacting gases. What we actually observe is the result of many millions of such atomic reactions. The ready-made " packets " of the various elements enter into combination or replace each other in such a way that there is no residue, e.g. the hydrogen " packet " contained in such a compound as hydrogen chloride will fit without any alteration into a new hydrogen compound. Dalton's assumption that atoms are permanent, coherent wholes, fits in exactly with the facts concerning chemical reaction which experiment has taught us. The Law of Simple Proportions, etc., is the experimental statement of what one would expect to occur if matter has an atomic structure. These ready-made atoms of the different elements combine with each other in the ratio of 1 : 1, 1 : 2, etc., according to the power of the atom to hold other atoms in combination with it. Each atom has a fixed weight, so that chemical combination must take place in fixed (or multiple) proportions by weight. If one atom is replaced by an atom of MOLECtlLAH AND ATOMIC WEIGHTS 137 another element, there will be a definite ratio between the weights of these atoms, hence the Law of Reciprocal Proportions is a logical deduction from the assumptions of the Atomic Theory. Atomic Structure. — There seems httle doubt at the present day that the properties, both of elements and compounds, are influenced not only by the kind of atoms contained therein, but also by the actual geometrical arrangement of the atoms in space. It is often found that two substances possess the same chemical composition and the same molecular weight, and yet the properties differ fundamentally, e.g. urea and ammonium cyanate. The different structure of the molecules in such cases is pictorially represented by graphic formulae. This is a help in enabUng one to explain the various reactions of the different compounds, but must not be looked upon as conveying any real picture as to the structure of the molecule. NH2 0=C^ 0=N = C— NHi ^NH, Ammonium cyanate. Urea. Substances such as urea and ammonium cyanate, possessing the same chemical composition and the same molecular weight, are known as Isomers. In other cases, although the composition of the substances is the same, the molecular weight differs, viz., C2H2 (acetylene) and CeHe (benzene). Benzene is referred to as a polymer of acetylene. The Significance of a Chemical Equation. — An equation may be viewed as a quantitative expression of a definite chemical reaction, and, as such, it embodies many of the most important laws and hypotheses in chemistry. The equation Na+SHj— ^2NH3 calls attention to the atomic structure of these gases, as well as to their molecular nature, i.e. the atomic and molecular hypotheses are implied in such an equation. Passing on to the quantitative aspect of the reaction, since each atom possesses a definite weight, the equation is an expression of the Law of Fixed Proportions. 138 AN INORGANIC CHEMISTRY Again, the Law of Conservation of Mass is involved, inasmuch as there is no loss of atoms in the reaction. The use of the above molecular equation, instead of the atomic equation N+3H— ^-NHg, enables us to record not only the weight relationship, but also the volume changes. The formula Na is always understood to refer to the nitrogen contained in the gram-molecular volume, so that the equation impUes Gay Lussac's Law of Volumes — one volume of nitrogen combining with three volumes of hydrogen to form two volumes of ammonia. Nor must one overlook the fact that the combination takes place between molecules, i.e. Avogadro's Hypothesis is therein implied. Chemical Calculations. — Nothing brings out more clearly the rigidity of the laws of chemistry than the study of problems bearing upon chemical phenomena. Problem 1. — What weight and what volume of oxygen, measured at 10° C. and 750 mm. pressure, may be obtained from the decomposition of 20 gm. of potassium chlorate ? The first stage is to set down the chemical equation which is a quantitative representation of this reaction : 2KC103->2KCl + 302. This states that from two gram molecules of potassium chlorate three gram molecules of oxygen are produced. Two gram molecules of potassium chlorate weigh 2(39+35-5+3x16) = 2x122-5=245 gm., three gram molecules of oxygen weigh 3x2x16=96 gm. Hence the weight of oxygen produced from 96x20 20 gm. of potassium chlorate is — -- — =7-8 gm. From our definition of gram-molecular weight, we know that 32 gm. of oxygen at 0° and 760 mm. will occupy 22-4 htres. Therefore the volume occupied by 7-8 gm. of oxygen at N.T.P.=22-4x 7-8 — =5-46 htres. Hence, the volume occupied at 10° and 750 „ .„ 283 760 ,.^ nim.=5-46x—X^ htres. Problem 2. — The density of a compound of carbon and sulphur relative to hydrogen is 38. The combining ratio of carbon to sulphur is 3 : 16. What is the molecular weight of the compound, MOLECtlLAR AND ATOMIC WEIGHTS 139 and what is its formula ? Give reasons for the adoption of this method of determining molecular weights. The density relative to hydrogen being 38, it follows that the molecular weight must be 38x2=76. Hence the weight 3 of carbon contained in the molecular weight is— X 76 =12, whilst the weight of sulphur is — -x76=64. The number of carbon 12 64 atoms : the number of sulphur atoms =— : — =1 : 2, where 12 and 32 are the atomic weights of carbon and sulphur respectively. The formula is therefore CSj. Problem 3. — A metaUic oxide contains 48 per cent, of oxygen. What is the exact equivalent of the metal ? If the specific heat of the metal is 0-123, find the probable atomic weight of the element. The ratio of oxygen to metal is 48 : 52, whence the equivalent of the metal (the weight combining with 8 parts of oxygen) 52 must be — =8-66. By the law of Dulong and Petit the Sp. Ht . x At. Wt. =6-4, whence the approximate atomic weight is The exact atomic weight is wxcombming weight, so that «X 8-66 =52-03 approximately. The value of n is obviously 6, so that the true or adjusted atomic weight is 6x8-66=52. Problem 4. — One gram of phosphorus combines with 3-4355 gm. of chlorine to form a volatile chloride. One-tenth of a gram of this chloride occupies 22-6 c.c. at 100° and 750 mm. pressure. Assuming that the combining weight of chlorine is 35-45, find the adjusted molecular weight of the compound. From the above analysis, find the combining weight of phosphorus, then from the density results calculate the approxi- mate molecular weight. The next step in the problem is to distribute this molecular weight between the elements chlorine and phosphorus in the ratio given by the above analysis. This enables us to fix the number of combining weights of each element contained in the molecular quantity, i.e. to fix n and m in the equation. True mol. wt.=reX combining wt. of CI +»iX combining wt. of phosphorus. 140 AN INORGANIC CHEMISTRY QmssTiONS 1. Why has oxygen =1-6 been chosen as the standard atomic weight ? 2. Given a large number of volatile compounds of an element, e.g. chlorine, the percentage composition of each compound being known, how is the atomic weight ascertained ? 3. Give an account of Dalton's Atomic Hypothesis. 4. Show how the Law of Isomorphism has been used to determine the atomic weight of an element. 5l Discuss fully the significance of the equation : — 2Hj + 02->2H20. 6. Show the connection between combining weight, atomic weight and molecular weight. 7. What advantages do atomic weights possess over combining weights ? 8. Why are atomic weights relative, not absolute t 9. What volume of oxygen, measured at 15° C. and 750 mm. pressure, is obtained by the decomposition of 40 gm. of sodium peroxide ? 10. The molecular weight of ammonia is 17-03. What is its density referred to air. Analysis shows that it contains 82-27 per cent, of nitrogen and 17-73 per cent, of hydrogen. What is the formula of ammonia ? 11. What is the relative weight of hydrogen chloride, compared (a) with hydrogen, (6) with air as the unit ? 12. Show how the laws of chemical combination and the Atomic Hypothesis are inter-related ? 13. What weight of zinc will evolve sufiflcient [hydrogen to fill exactly a tank holding 3000 galls.? [1 litre = 1"8 pints; 1 litre of hydrogen weighs 0-0899 gm. Tm -- 65. CHAPTER X CHLORINE AND HYDROGEN CHLORIDE Chlorine was discovered by Scheele in 1774, though it was many years before its elementary nature was recognised. Lavoisier (1789) named the gas oxy-muriatic acid, in the beHef that it was related to muriatic acid (hydrochloric acid) in the same way as sulphurous acid is to sulphuric acid. In 1810 Davy showed that hydrochloric acid is a compound of hydrogen and chlorine, and contains no oxygen. By this discovery he established the elementary nature of this gas which he named chlorine (chloros, green) and incidentally wrecked the oxygen theory of acids propounded by Lavoisier — the theory that aU acids contained oxygen. Occurrence. — Chlorine occurs only in the combined state, almost entirely as chlorides of such metals as sodium, potassium, magnesium, which constitute the main part of the dissolved matter in sea-water. In many parts of the earth, the evaporation of inland seas has led to the formation of thick beds of the salts generally found in sea water. This is especially so at Stassfurt, where the combined thickness of the strata amounts to over 1,000 feet. The chief salts in this bed are rock salt (halite) NaCl, sylvine KCl, carnallite KCl,MgCl2,2H20, kieserite MgS04,H20, schonite K2S04,MgS04,6H20, kainite MgSOi, KaS04,MgCl2,6H20, polyhalite MgS04,K2S04,2CaS04,2H30, gypsum CaS04,2H20, anhydrite CaSOi. Preparation. — The methods of preparation of chlorine fall into three main classes : 1. The chlorides of a few metals, such as gold and platinum decompose on heating with the evolution of chlorine. This method is, however, rarely used owing to the cost of the materials involved. 141 142 AN INORGANIC CHEMISTRY 2. The soluble chlorides of the metals and of hydrogen are readily decomposed by the passage of an electric current through their aqueous solution. In all cases the chlorine is liberated at the positive pole, the hydrogen or metal at the negative pole. The apparatus in Fig. 58 is often used to illustrate the process on a laboratory scale. Owing to the ease with ^^•hich nascent chlorine attacks platinum, the anode or positive pole must be made of carbon. As soon as the ciurent is switched on, an effervescence is noticed round the cathode (negative pole) and to a less extent round the anode (positive pole). The reason of the smaller evolution round Fig. 58. this pole is that chlorine is appreciably soluble in the Uquid. In order to decrease this solubihty loss, it is usual to employ a solution of hydrochloric acid saturated with sodium chloride — a solution in which chlorine is Uttle soluble. The reaction in the case of a metallic chloride is represented in the equations : 2KC1^2K + Cl2 the potassium immediately reacting with the water 2K + 2H20-> 2K0H + Hj with the evolution of hydrogen and the formation of potassium hydroxide. Since the hydroxide reacts with the chlorine, it is essential to keep separate the products formed at the two poles. This is the basis of the most important methods for the pre- paration of chlorine — ^the eleetrglygis of solutions of potassium CHLORINE 143 and sodium chlorides in apparatus so designed that the hydroxide formed at one pole cannot react with the chlorine generated at the other. This result is often achieved by the use of a diaphragm, but discussion of these cells will be postponed till the preparation of sodium hydroxide is dealt with (p. 470). The chlorine of commerce is essentially a by-product in the manufacture of the alkali hydroxide. The chlorine evolved from these cells is either Uquefied by compression into iron cylinders or converted into bleaching powder (q.v.). 3. Hydrogen chloride may be oxidised by the oxygen of the air or by a suitable oxidising agent, forming water and chlorine. 4HC1 + 02-> 2H20 + 2CI2. In the case of oxidation by molecular oxygen (i.e. uncombined oxygen) the reaction will only proceed at all freely at very high temperatures. The reaction between hydrogen chloride and oxygen is, however, attended by the evolution of heat, so that, in accordance with Le Chatelier's Law, the higher the temperature the smaller the yield of chlorine. We have again to deal with an equilibrium reaction, and the position of that equihbrium, as predicted by the application of Le Chateher's Law, is dependent upon the temperature. The conditions that favour an increased velocity of reaction, i.e. high temperature, must inevitably lead to a decreased yield of chlorine. The problem received a satisfactory solution at the hands of Deacon in his so-called Deacon Process. By the use of a suitable catalyst he was able to secure a sufficient velocity of reaction at 400°, a temperature which gave him a good yield of chlorine. The catalyst consists of pumice impregnated with copper chloride. It is supposed that the following cycle of operations occurs : — 4CuCl-f02-^2CuaOCl2 2Cu20Cla + 4HC1^^ 4CuCl2 + 2H2O 4CuCl2->4CuCl-f2Cl2 summmg, 4HC1 -f O2 -> 2H2O -f 2CI2 The chlorine is mixed with nitrogen, oxygen, steam and excess of hydrogen chloride. The most injurious of these impurities is hydrogen chloride, and this is easily removed by " scrubbing " or washing the gas. The chlorine is then sufficiently pure for bleaching or for the manufaotxire of bleaching powder. In general, however, the hydrogen of the hydrogen chloride 144 AN INORGANIC CHEMISTRY is oxidised away by the oxygen contained in some compound rich in oxygen, e.g. manganese dioxide, potassium permanganate, potassium dichromate, lead peroxide. The essential point about all these compounds is that they should contain an element which is capable of passing from a higher to a lower valence, e.g. : Mn'^Oa + 4HC1-^ Mn'^Cli + 2H2O Mn'^Cl4->Mn"Cl2 + Cl2. This reaction is generally carried out in an apparatus such as in Fig. 59. Concentrated hydrochloric acid is run upon the manganese dioxide so as to form a thin paste. The mixture is cau- tiously heated and the evolved gas led through a wash bottle containing water. This re- moves most of the hydrogen chloride f= carried over with the chlorine. Owing to its great solubility in water, the gas is collected by upward displacement of the air. There is still a little doubt as to whether the above equation exactly represents the mechanism of the reaction, but considerable sup- port of the above view is derived from the fact that, if the reaction vessel is cooled by ice and, after satiu-ation with chlorine, the contents are poured into water, hydrated manganese dioxide is thrown down : MnCli -1- xH^O -^ Mn02,(x-2)H20 -I- 4HC1. Others, however, are of the opinion that the trichloride, and not the tetrachloride, is the intermediate compound formed during the reaction. Fig. 59. CHLORINE 145 2Mn02 + 8HCl->. 2MnCl3 + Cl^ + ^H^O 2MnCl3-^2MnCl2 + Cl2 It is highly probable that both the tri- and the tetra-chloride play a part in the reaction. In practice one often generates the hydrochloric acid in the vessel itself. An intimate mixture of salt and manganese dioxide is treated with sulphuric acid and heated. The action of the acid upon the salt (sodium chloride) is to liberate hydro- chloric acid, which then reacts with manganese dioxide. NaCl + H^SOi-^ NaHSOi + HCl Mn'^'Oa +4HCl-^Mn"Cl2 + Cl^ + 211^0 If, instead of the dioxide, one takes the monoxide, the reaction follows the usual course between an acid and a base. Mn"0 + 2HC1 —^ Mn"Cl2 + H^O Acid. Salt. Water. Both in the form of the oxide and the chloride the manganese is divalent, and we find that no chlorine is evolved. Other reactions in which chlorine is evolved from hydrochloric acid are given below : PbOa +4HCl-^PbCl4 +2H20-> PbCl^ + Cl^ + 2HjO. Lead peroxide. The reaction between sodium dichromate and hydrogen chloride can be best understood by looking upon the dichromate as a compound of the basic oxide NaaO with the acidic oxide CrOj, the basic oxide reacting in the normal way to form water and sodium chloride, the acid oxide reacting in a manner analo- gous to manganese dioxide, i.e. showing a change of valence. Na^O +2HCl->2NaCl+H20 2Cr"03 + 12HCl-> 2Cr'"Cl3 + GH^O + 301^ Na^Cr A + 14HC1-^ 2NaCl + 2CrCl3 + IB.^0 + 3Cl7 Sodium dichromate. In this reaction the sodium salt is often replaced by the more expensive potassium compound, but as the evolution of chlorine depends upon the presence of the CrOs group, it is obvious that the cheaper compound possesses exactly the same oxidising power per molecule as does the more expensive potassium dichromate. Not only are the sodium salts much cheaper than the corresponding potassium salts, weight for weight, but it L 146 AN INORGANIC CHEMISTRY must be remembered that the gram-moleoular weight of the sodium salt is considerably less than that of the potassium salt (Atomic weight of sodium =23, potassium 39). Another valuable oxidising agent is potassium or sodium permanganate. The reaction is represented thus : 2KMn04 + 16HCl-> 2KC1 + 2MnCl2 + SCl^ + SH^O The equation may be written in a more suggestive way, 2KMn04=KAMnA K 2O + 2HC1 -^ 2KC1 + H 2O Mn^A + 14HC1-^ 2Mn''Cl2 + VH^O + SCl^ K20,Mn™A + 16HC1->2KC1 + 2Mn"Cl2 + SK^O + SCla The reactions in which the three oxides of manganese function, viz. MnO, MnOa, MnjO,, may be suggestively written in the following way : Mn = + 2HCl->MnCl2 + H^O /0+2HCI MnCIa + H^O i+SHCl^^HjO + Cla the upper half being the normal reaction between acid and base, the lower one an oxidation of the hydrogen of the acid to form water and chlorine. ^4HCl + O2 Dry chlorine reacts very feebly, but in the moist state it is a very reactive element. Copper, phosphorus, antimony, boron and silicon ignite spontaneously, while many other elements, such as sodium and iron, burn vigorously if heated sufficiently to start the reaction. The product of the reaction is a chloride : 2P + 3Cl2->2PCl3 Si +2Cl2-^SiCl4 Hydrogen and chlorine combine with explosion, provided the mixture is exposed to a bright light. In the absence of such hght there appears to be no measurable reaction ; in diffused hght the combination proceeds slowly, while the flash of a magnesium light causes a loud explosion. It appears that a certain measure of Ught energy is necessary to induce the reaction to begin, but when once under way the wave strikes through the mixture with explosive violence. The explosion wave is strongest when the ratio of the volume of the gases is 1 : 1, i.e. the ratio necessary for complete reaction. A jet of hydrogen wUl continue to burn in an atmosphere of chlorine with the production of copious fumes of hydrogen chloride. Similarly, if a jar of hydrogen is inverted and ignited below, and a jet of chlorine passed through the zone of burning hydrogen, the chlorine will ignite and continue to burn in the hydrogen atmosphere. It is an interesting example of reciprocal combustion. Chlorine is also an active agent in replacing other elements already present in compounds. A very good example of this is the usual test for chlorine— the liberation of iodine from potassium iodide. A small quantity of chlorine water, added to HYDROGEN CHLORIDE 149 a solution of a soluble iodide, produces a brownish coloration. This test can be made more delicate by shaking the contents of the test-tube with carbon disulphide. The liberated iodine distributes itseK between the aqueous and the carbon disulphide layers (see section on Partition, p. 112), and the lower layer of disulphide takes on a pronounced violet colour. An even more delicate variation of the test suitable when traces of chlorine are being sought for is the starch iodide test. Filter paper is dipped into starch emulsion containing a small amount of potassium iodide. If a strip of this paper is brought into contact with a mere trace of chlorine, a deep blue colour develops. Bromides, too, are decomposed by chlorine : 2KBr + Cla-^ 2KC1 + Br^ 2KI + Cl2^2KCl+l2 A piece of filter paper, moistened with turpentine, wiU ignite when dropped into a jar of chlorine with the production of dense fumes of free carbon and of hydrogen chloride : CioHie + SCla-^ IOC + 16HC1. Very often the chlorine adds itself on to a compound. Thus, when carbon monoxide (g'.f.) is mixed with chlorine and exposed to the sunlight, drops of a highly poisonous substance, known as phosgene (COClg) separate out (see p. 361). Commercial Applications. — The more important uses to which chlorine is put in the commercial world are based upon the reaction described above. The colouring matter present in many fibres is frequently oxidisable by moist chlorine with the production of colourless oxidation compounds — hence its great use for bleaching purposes. Many bacteria are easily acted upon by chlorine and killed, so that its use as a disinfectant is very wide. The production of bromine, the extraction of gold are other examples of the commercial application of chlorine. Possibly its use as a poison gas may be classed under this head. Hydrogen Chloride Preparation. — Three main methods of preparation of this substance are to be noted : 1. The synthesis from hydrogen and chlorine. 2. The action of water upon the chlorides of the non-metalg 150 AN INORGANIC CHEMISTRY 3. The action of a non- volatile acid upon the chlorides of the metals. The first method has already been mentioned and is only of theoretical importance ; the two latter form the commercial and laboratory methods of preparation. The chlorides of the non-metals in nearly all cases react vigor- ously with water, forming two acids — one of these always being hydrochloric acid (the acid formed by the solution of hydrogen chloride in water) : POs + 3H0H -^ P(OH) 3 + 3HC1 Phosphorous acid. SiClj + 4H0H^ Si(0H)4 + 4HC1 Silicic acid. This type of double decomposition in which one of the reacting substances is water, is called Hydrolysis. Hydrolytic reactions are very common, but in most cases an equihbrium is soon reached,* as in the reaction : Na.COs + HjO^NaOH -f NaHCOg Sodium carbonate. Sodium hydroxide. Sodium bicarbonate. The best example of the formation of hydrogen chloride from a chloride is by the action of sulphuric acid upon sodium chloride. Under the conditions attainable in the laboratory the reaction proceeds thus : NaCl + HaSOi-^NaHSOi +HC1. On account of the extreme solubiUty of hydrogen chloride in water, the gas is collected over mercury or by the upward displacement of the air. If a high temperature is attainable, the reaction can be pushed a stage farther : 2NaCl + H^SOi-^Na^SO^ + 2HC1. Sodium sulphate. This is generally the case in technical work, where the objective is to secure sodium sulphate [q.v.) rather than the by-product, hydrogen chloride. The hydrogen chloride evolved from this reaction, when carried out on the commercial scale, is passed through a tower filled with lumps of coke, over which water trickles. This solution of commercial hydrochloric acid, often known as muriatic acid, is put on the market for use in the dye industry, chlorine bleaching industry, etc. * For full treatment of the subject of hydrolysis, see p. 447. HYDROGEN CHLORIDE 151 It is important to note that all the chlorides react with sulphuric acid in this way, but the rapidity of the reaction varies con- siderably. The rate at which the hydrogen chloride is liberated is largely dependent upon the solubUity of the chloride. Other non-volatile acids, such as phosphoric acid, may be used in place of sulphuric acid. The reason of this is seen in the next section. The Reaction NaCl+H2S04^=±NaHSO,+HCl. Rever- sible Reactions. — ^The action of sulphuric acid upon sodium chloride has already been discussed. It leads to the formation of sodium hydrogen sxilphate and hydrogen chloride, which escapes as a gas. If, into a saturated aqueous solution of sodium hydrogen sulphate, a stream of hydrogen chloride is passed, a copious crop of small, cubical crystals soon separates out. These are crystals of sodium chloride, i.e. the reaction NaHSOi+HCl — >NaCl ^ -(-H2SO4 has set in, and will not cease till practically the whole of the acid sulphate has been converted into the chloride. Here, then, are two opposing reactions, and by a mere alteration of the conditions of the experiment the reaction may be driven at will in either direction (cf . the action of steam upon heated iron filings, p. 85). If, however, the experiment is so arranged that none of the participating substances can escape from the system, it will be found that a definite equihbrium will be set up, and at this point the hydrogen chloride will exert a definite but constant pressure. All investigations of such phenomena lead to the conclusion that at this point the rates of the two opposing reactions are equal, so that there is no apparent change, i.e. just as many hydrogen chloride molecules are hberated per second by the action of the sulphuric acid on the sodium chloride as are removed by the interaction between the acid sulphate and hydrogen chloride. But if the system is not a closed one, the hydrogen chloride, which is but slightly soluble in sulphuric acid, will escape so soon as the liquid has become saturated with the gas. After this, hydrogen chloride will escape as fast as it is formed. So far as the reverse action is concerned, the high concentration of hydrogen chloride molecules present when the gas is led through the solution of sodium hydrogen sulphate, brings about just those conditions which, from a kinetic point of view, favour the reaction with the acid sulphate. Sodium chloride is formed, and as this substance is almost insoluble in a solvent so rich in hydrogen chloride (for discussion 152 AN INORGANIC CHEMISTRY of this, see p. 433), the time soon comes when sufficient of the sodium chloride is produced to saturate the solution with this substance. The constant renewal of the supply of hydrogen chloride consequently leads to the steady precipitation of sodium chloride crystals. The essential features in this complicated reaction are : 1. The volatility of the hydrogen chloride. 2. The insolubility of sodium chloride in solutions saturated with hydrogen chloride. In the past, chemists sought to explain reactions in terms of chemical affinity. The reaction between sodium chloride and sulphuric acid, with the liberation of hydrogen chloride, took place, it was presumed, because the chemical affinity between sodium hydroxide and sulphuric acid exceeded the affinity between sodium hydroxide and hydrogen chloride (hydrochloric acid). The latter was therefore expelled from chlorides by the action of sulphuric acid. But the inadequacy of the explan- ation is at once apparent when one considers that, by a sUght alteration of the conditions, hydrochloric acid may be made to expel sulphuric acid from a sulphate. The reaction NaCl+H2S04^zz±NaHS04-|-HCl forms an excellent example of what is known as a reversible reaction. Provided that water is kept from the reaction vessel — ^for hydrogen chloride is exceedingly soluble in this solvent — other acids, such as phosphoric acid, may be used to set free hydrogen chloride from a chloride. The essential feature to note is that the acid must be less volatile than the hydrogen chloride, in which case the hydrogen chloride will tend to escape from the system. The addition of a volatUe acid hke nitric acid to a solution of sodium chloride causes a partial Uberation of the hydrochloric acid : Naa + HNO3 ^z± NaNOa + HCl, until the necessary equilibrium in the solution has been reached. In general, it may be stated that, when one acid is added to a salt, or when two salts formed from different acids are mixed, a mutual interchange of the radicles wiU occur. Whether pre- cipitation or evolution of gas will take place, depends upon the relative solubility and volatility of the products formed. Physical Properties. — Hydrogen chloride is a colourless gas which attacks the mucous membrane of the throat. It is HYDROGEN CHLORIDE 153 slightly heavier than air and is exceedingly soluble in water, 1 c.c. of water at 0° dissolving 520 c.c. of the gas. Density (air=l) 1-2681. Weight of 1 litre=l-6394 gm. at N.T.P. Critical temperature =51 -8°. Boiling point of liquid = —83-7°. Melting pomt of solid=— 112-5°. If a dilute solution of hydrochloric acid is heated, the gas which escapes consists mainly of water. The residue, therefore, increases in strength and the boiling point of the solution rises. This continues until the concentration of the acid has risen to 20-4 per cent., the boiling point being 110°. On the other hand, if a strong solution of the acid is boiled, the escaping gas consists largely of hydrogen chloride ; the solution, therefore, becomes weaker and the boiling point again rises until the maximum 110° is reached (20-4 per cent, strength). An acid containing 20-4 per cent, of hydrochloric acid will distil unchanged. The following diagram illustrates this : — ■ X \ 50 \ \ \ \ 1 \ \ t :$ \ -KCl+H. With ammonia it unites directly to produce solid ammonium chloride : NH3+HCI->NH4C1. The gas is very stable and shows no sign of dissociation below 1500°. Neither the dry gas nor the dry Hquid has any action upon litmus or upon zinc, i.e. in the absence of water hydrogen chloride lacks the fundamental properties of an acid. On the other hand, hydrogen chloride, it dissolved in water, possesses aU the fundamental properties of a strong acid ; metals, such as iron and zinc, are speedily attacked with the hberation of hydrogen. The hydrochloric acid of commerce {muriatic acid) is an aqueous solution saturated with hydrogen chloride and contaminated by small quantities of impurities, e.g. arsenious chloride, ferric chloride and free chlorine. Chlorides. — The chlorides of the metals are usually obtained by the action of hydrochloric acid upon the oxide, hydroxide or carbonate of the metal : Ca(0H)2 + 2HC1^ CaCla + 2H2O CaCOj + 2HC1^ CaClj + H^O + CO^ Evaporation of the resultant solution generally gives a crop of crystals of the chloride of sufficient purity for most purposes. The student should note that this is the general method for the preparation of a salt — the action of the requisite acid upon the oxide, hydroxide or carbonate of the metal, e.g. : ZnO + 2HN03^ Zn(N03)2 + H^O. The chlorides of the non-metals are obtained by the direct action of chlorine upon the non-metal m the absence of water. 2As + 3Cl2->2AsCl3 2P -f 3Cl2->2PCl3 The chlorides of the non-metals are strongly hydrolysed, i.e. HYDROGEN CHLORIDE 155 broken down, in the presence of water, while the majority of the metallic chlorides dissolve readily in water without appreci- able decomposition (see p. 457). The notably insoluble chlorides are silver chloride AgCl, cuprous chloride CuCl, aurous chloride AuCl, mercurous chloride HgaClj, lead chloride PbCla, the last of which dissolves fairly readily in hot water. Composition. — Various experiments have been devised to determine the composition of hydrogen chloride. Method 1. — One of the simplest is the electrolysis of hydro- chloric acid. A special form of apparatus, designed by Lothar Meyer, gives satisfactory results (Fig. 62). In order to secure reUable results, the -f^gjH ■ acid should be con- centrated, and since chlorine is appre- ciably soluble in hydrochloric acid, the current must be passed for some time before the gases are collected, in order to saturate the hquid round the anode with chloriae. During the experiment the gases are allowed to collect m the receivers A and B. The gases are found to separate at equal rates, i.e. during the decomposition of hydrochloric acid by the electric current, equal volumes of hydrogen and chlorine are evolved at the electrodes. Assuming that all the hydrogen and the chlorine comes from the electrolysis of the dissolved hydrogen chloride, it follows that this gas must contain equal volumes of hydrogen and chlorine— a result in conformity with the law of Gay Lussac. Method 2.— A stoppered glass tube is fiUed with dry hydrogen chloride. A little sodium amalgam is introduced, the stopper replaced and the contents shaken. The tube is then opened under mercury, and after the mercury levels are adjusted, it is found that the residual volume of gas, which proves to be A 156 AN INORGANIC CHEMISTRY ^C hydrogen, is haK of the original volume. Therefore hydrogen chloride contains half its own volume of hydrogen, or half a volume of hydrogen from one volume of hydrogen chloride. The application of Avogadro's Hypothesis leads to the conclusion that half a molecule or one atom of hydrogen must be present in each molecule of hydrogen chloride. The formula must consequently be (HCl)^.. From the vapour density meas- _ urement it is known that the molecular weight is 36-5, and as the weight of hydrogen contained in the mole- cule is known by the above reasoning to be 1, the weight of chlor- ine in the molecule must be 35-5 — t h e atomic weight of chlorine ; x must therefore be 1 and the formula HCl. Method 3. — Equal volumes of hydrogen and chlorine are in- troduced into a stout walled gl a s s tube, closed at each end with a stop-cock, and their union brought about by exposure to a mag- nesium light, or by the passage of an electric spark (Fig. 63). The final volume is found to be the sum of the volumes of the hydrogen and chlorine — an experiment which leads to the conclusion that one molecule of hydrogen combined with one molecule of chlorine forms two mole- cules of hydrogen chloride. On the assumption that the molecule of hydrogen and that of chlorine each consist of two atoms, the molecule of hydrogen chloride is composed of one atom of hydrogen and of chlorine. The formula is consequently HCl. Fia. 63. Fig. 64. HYDROGEN CHLORIDE 157 Method 4. — A rather pretty method for demonstrating the composition by volume is to fill a tube A (Fig. 64) with the mixed gases obtained by the prolonged electrolysis of hydro- chloric acid. After the apparatus has been filled, the taps are turned, and the tube held in brine contained in the vessel B. The lower tap is then opened, the pressure adjusted and the volume read. A small amount of potassium iodide solution is cautiously introduced into the tube through the tap funnel. The reaction 2 KI+Cl2^2KCl+l3 at once sets in, and the evolved iodine dissolves in the excess of iodide. The pressure is again adjusted and the volume read. It is foimd that half the original volume is left, and this gas is hydrogen, i.e. hydrogen and chlorine are formed in equal volumes by the electrolysis of hydrochloric acid. From these experiments we may conclude : — ■ 1. That two volumes of hydrogen chloride are formed from one volume of hydrogen and one of chlorine. 2. On the assumption that the hydrogen and chlorine molecules are diatomic, one molecule of hydrogen chloride must contain one atom of hydrogen and one atom of chlorine. 3. The formula is HCl. This formula is confirmed by a determination of the molecular weight, which is 36-47 (H=1008, 01=35-46). Questions 1. Show how the evolution of chlorine from hydrogen chloride is facili- tated by the presence of a compound which contains an element capable of passing to a lower stage of valence. Give examples. 2. Discuss the reaction NaCl+HjSOi^NaHSOj+HCl. 3. Describe fully two methods which serve to prove the composition of hydrogen chloride. 4. Show how the action of water upon sodiimi chloride and phosphorus trichloride differs. What general class of elements form halides which behave like phosphorus in this respect ? 5. Illustrate what is meant by "catalysis" by reference to the Deacon process. 6. Compare the action of chlorine and of hydrogen chloride upon water. What happens when aqueous solutions of these substances are boiled ? 7. How would you convert chlorine into sodium chloride ? CHAPTER XI THE HALOGEN FAMILY- FLUORINE, CHLORINE, BROMINE, IODINE AND THEIR HYDRACIDS GeneraL — The study of the elements, hydrogen, oxygen, and chlorine has revealed that each of these elements has its own individual properties, and these properties show little similarity. There appears no common ground of relationship between them. On the other hand, the elements fluorine, chlorine, bromine and iodine, show a great similarity of behaviour not only in the elementary state, but equally so in the combined state. The sodium salts of these elements (sodium fluoride, NaF, chloride NaCl, bromide NaBr, iodide Nal) all crystallise in the regular system, and have also many chemical properties in common. From this tendency to form salt-like substances is derived their name — halogen (Gk. aX?, salt ; 'yevvdo), to produce), and their compounds are known as halides. Of these elements, chlorine has already been studied in detail, not only because it is the most important from the chemical point of view, but because it is the most frequently met with. Bromine and iodine exhibit almost every property that chlorine does, though in general to a somewhat less degree. They win therefore be studied first, whUe fluorine, which shows more individuality in its reactions than do the other elements, wiU be dealt with subsequently. Bbomine and Iodine Discovery. — The element bromine was discovered in 1826 by Balard, who was investigating the mother liquor left after the crystallisation of salt from the waters of certain salt springs. Balard not only succeeded in isolating bromine by a method 158 BROMINE AND IODINE 159 similar to that used to-day, but established its elementary nature and its relation to chlorine and iodine. The existence of iodine was first suggested by the experiments of Courtois on the aqueous extract of kelp. The mother hquor left from the kelp after the crystallisation of the less soluble salts was heated with sulphuric acid and gave off beautiful violet vapours, but it was left to Gay Lussac (1813) and Davy to estabhsh that the new substance was an element. Occurrence. — Bromine occurs in sea- water and in certain salt springs in combination with sodium, potassium, magnesium and calcium. The salt deposits of Stassfurt in Germany were, until recently, the chief source of the world's supply of bromine, though the bromide deposits at Ohio, U.S.A., now bid fair to rival them. Iodine occurs in small quantities in kelp, in the thyroid glands of animals, in many sea plants and mineral springs, and occasionally as silver iodide. At the present day, the commercial source of iodine is Chili saltpetre, in which small quantities of sodium iodate and iodide occur. Preparation. — There are two commercial methods of making bromine : (a) The action of chlorine upon solutions containing bromides. (6) The action of manganese dioxide and sulphuric acid upon bromides (cf. the preparation of chlorine). The first method is used at Stassfurt where the hot mother hquor containing magnesium bromide is made to percolate down a tower filled with round stones in order to break up the flow. Chlorine is introduced from below, the liberated bromine escapes at the top, and is condensed by passage through a worm condenser. In America the second method is sometimes used. The mother liquor, freed as far as possible from common salt by crystallisation, is treated with the calculated quantity of sulphuric acid and manganese dioxide. MgBr^ + Mn02-f 2H2SOi->MgS04 +MnS04 H-Br^ + 2H2O. The action is essentially the same as in the preparation of chlorine from sodium chloride, manganese dioxide and sulphuric acid ; the latter acid liberates the haloid acid, which is then oxidised by the manganese dioxide. 160 AN INORGANIC CHEMISTRY Another method, suitable where electrical power is cheap, is to electrolyse a solution of a soluble bromide (of. preparation of chlorine). Iodine is prepared commercially either from kelp or from sodium iodate present in ChUi saltpetre. The kelp is burned in carefully constructed ovens in order to prevent loss through volatUisation. The ash is extracted, and as much as possible of the alkahne carbonates, sulphates and chlorides removed by crystaUisation. The liquor is then treated either with chlorine, or with manganese dioxide and sulphuric acid, as in the case of chlorine and bromine. The quantity of chlorine requires careful regulation, as does the amount of manganese dioxide, so that no excess of this re- agent is left over to react with traces of chlorides thereby causing the Hberation of chlorine. The evolved iodine is purified by sublimation in iron retorts. The main part of the world's iodine is, however, obtained from the mother Hquor of Chili saltpetre left after the sodium nitrate has been crystallised out. This Hquor contains considerable quantities of iodine in a highly oxidised form, viz. sodium iodate, NalOg. In order to obtain the iodine in the elementary form, a reducing agent is necessary, the most satisfactory reducer for this purpose being sulphur dioxide. The chemistry of the process is given in the equation : IA + 5S02-^5S03+l2. Actually, the sulphur dioxide is introduced in the form of sodium bisulphite. The equation then becomes : 2NaI03 + 5NaHS03^ 3NaHS04 + 2Na2S04 + H^O + 1^. Physical Properties. — Bromine is a dark red, heavy hquid which boils at 59° with the production of dense, reddish brown fumes. The specific gravity of the liquid is 3-188. The vapour has a most powerful and disagreeable smell, and even when in a diluted state, it attacks the eyes and the mucous lining of the nose and throat, whilst the Hquid exerts a powerful corrosive action on the skin. Bromine is appreciably soluble in water ; a few drops of the Hquid in water readily impart a reddish brown colour to the solution, forming bromine water. The solubility at room temperature is about 3 gm. per 100 c.c. of water. BROMINE AND IODINE 161 At ordinary temperatures iodine exists as a dark blue crys- talline substance of the rhombic form. These crystals have a sp. gr. of 4-933 at 4°. On heating, iodine vaporises readily and at 114° melts. The smell of iodine is somewhat similar to that of chlorine, but less offensive. The colour of the vapour is at first reddish violet, but as the temperature rises, the colour takes on a deep indigo blue. Iodine is but faintly soluble in water, 1 gm. dissolving in about 5^ litres to form a faint brown solution. Chloroform, carbon disulphide, and benzene dissolve it freely with the formation of violet solutions, whUe alcohol and ether give brown solutions. Iodine is also freely soluble in aqueous solutions of potassium iodide or of any other soluble iodide. Such solutions are brown. The increased solubUity is here due to the formation of a poly iodide. KI + I,^^Kl3. The whole of the iodine is not held in combination in the form of polyiodide, but a definite equiUbrium exists between the three reagents. Such solutions give all the reactions of free iodine, and are frequently made use of, if it is desired to bring into solution a large amount of iodine. As the free iodine is removed from the system, the polyiodide splits off more iodine in order to restore the equUibrium, so that the whole of the dissolved iodine, whether free or bound in the polyiodide, is ultimately available for any chemical reaction. Chemical Properties. — The chemical properties of bromine strongly resemble those of chlorine, except that the energy of reaction is generally less marked, e.g. sunlight is not a sufficiently strong catalyst to bring about the combination between hydrogen and bromine. Even in the presence of finely divided platinum, the reaction is slow and by no means of the explosive nature shown by hydrogen and chlorine. Arsenic and its related elements, phosphorus and antimony, ignite when dropped into bromine and burn with the formation of the corresponding halide. This can be safely shown as follows (Pig. 65) : — A few drops of bromine are put in the inner test-tube, the funnel adjusted to prevent splashing, and a fragment of phosphorus dropped into the tube. The reaction is vigorous and must be carried out in a draught cupboard. Bromine possesses mild bleaching powers and, like chlorine, forms an unstable compound M 162 AN INORGANIC CHEMISTRY with water, Bra.lOHaO, when the saturated aqueous solutioi is cooled to 0°. The atomic weight of chlorine is 35-46, of bromine 79-92 and of iodine 126-92. We have seen that there is a stronj similarity between the first two of these elements, except thai the properties are less clearly marked in the case of the heaviei element. The same general change is to be noted in the case oJ iodine ; the reactions of this element are similar to, but less vigorous than, those of its relatives, chlorine and bromine. As an illustration of this, iodine displays the least tendency of al] the halogens to combine with hydrogen, so that hydroger iodide is generally formed not by direct combination, but by indirect means. Iodine combines directly with many metals, e.g. mercury, and with a few non-metals, notably phosphorus. Both chlorine and bromine displace iodine from iodides, whilst chlorine is also capable of displacing bromine from bromides. We shall soon see that this is equivalent to saying that chlorine is a stronger oxidising agent than bromine, bromine than iodine. In fact, the chemis- try of these related elements may almost be summed up in the statement : The oxi - dising power of the halogens, fluorine, chlorine, bromine and iodine steadily diminishes as the atomic weight of the element increases. The minutest trace of iodine, even -0000001 gm., may be detected by the starch iodide test already described as a test for chlorine (p. 149). This blue coloration is supposed to be due to the formation either of a soUd solution of iodine in the starch or of a loose chemical compound between them. The blue colour is destroyed by raising the temperature above 80°, but returns on cooling. Fig. 65. Dissociation of Bromine and Iodine.— The effect of temperature in decreasing the vapour density of chlorine and hence lowering the molecular weight, has abeady been mentioned. A similar effect has been found in the case of the aUied elements, BROMINE AND IODINE 163 bromine and iodine. Table 22 records the experimental results bearing upon this point. TABLE 22 Bromine : Temperature . . . 102° 228° 1570° Vapour Density (air = 1) 5-73 5-52 3-70 Iodine : Temperatm-e . . . 450° 842° 1027° .Vapour Density . . 8-85 6-76 5-75 1570° 6-70 This steady decrease in the vapour density and molecular weight brought about by an increase of temperature is due to the breaking down of the diatomic molecule into single atoms : Br,;=^ attended by 2Br. 21. heat The dissociation is attended by heat absorption and will therefore be increased by a temperature rise (Le Chatelier's Law). Hydeobeomic and Hydeiodic Acids Preparation. — The methods of preparation are essentially the same as given for the preparation of hydrogen chloride, though minor differences occur. 1. Direct combination of the elements. 2. The hydrolysis of the non-metallic bromides and iodides. 3. The action of a non-volatile acid upon the haHde. Neither bromine nor iodine unite vigorously with hydrogen, except in the presence of a catalyst — platinum. Fig. 66 illustrates a suitable method of preparing hydrogen bromide. TsrjnmmMmr^ ^ Fig. Pure hydrogen is bubbled through a wash-bottle containing liquid bromine, which is maintained at a temperature of about 60°, and the mixed gases then pass through a tube containing a spiral of platinum wire, heated electrically to a dull red heat. -164 AN mORGANIC CHEMISTRY Hydrogen bromide may be obtained from the effluent gases by absorption in water. For efficient working there should be an excess of hydrogen, as shown by the escape of bubbles from the exit. A somewhat similar method may be used for the preparation of hydrogen iodide, but in this case the iodine must be kept sufficiently hot to vaporise and the reaction is by no means complete, i.e. the equilibrium H,+I,^=±2HI is set up and at 448° only 79 per cent, of the constituents unite. Phosphorus tri-bromide or -iodide reacts energetically with water in accordance with the equation : PBr3 + 3H0H^ P(0H)3 + 3HBr. PI3 + 3HOH-^P(OH)3 + 3HI. In practice, the tri-halides are prepared in the course of the experiment, and the requisite amount of water allowed to drip upon the tri-halide to bring 1F3 about the necessary hydrotysis. The apparatus in Fig. 67 is suitable for this purpose. In the case of the bromide, red phosphorus is mixed with a httle sand and 'a small quan- tity of water added. Owing to the extreme solubUity of hydrogen bromide in water, excess of this reagent is to be avoided. Bromine is then allowed to run in from the dropping funnel, drop by drop. The gas, consisting of hydrogen bromide and bromine, escapes through a U-tube packed with glass beads and red phosphorus. This acts as a washer which effectually removes any bromine present in the gas. In the case of hydrogen iodide, red phosphorus and iodine are mixed in a dry flask and water allowed to drop upon the tri-iodide so formed. The method of purification is the same as for hydrogen bromide. Fig. 67. BROMINE AND IODINE 165 The third method of preparation is little used. As a suitable non-volatile acid, phosphoric acid is perhaps the most satisfactory to use in evolving hydrogen bromide or iodide from the halide salt. 3NaBr + H3P04^ NagPO^ + 3HBr. 3NaI + HjPOi-^-NajPOi + SHI. Sulphuric acid, which is so commonly used in the preparation of hydrogen chloride, is not a satisfactory reagent in the above reactions, for the hydrogen bromide (or iodide) reacts vigorously with this acid, reducing it to sulphur dioxide, while it is itself oxidised to the free halogen H2S04 + 2HBr^.2H20+S02 + Br2 (HASO3) so that one obtains a mixture of hydrogen bromide, bromine and sulphur dioxide. This side reaction is even more noticeable in the corresponding preparation of hydrogen iodide. Phosphoric acid, on the other hand, is very difficult to reduce, hence its use in preference to sulphuric acid. If an aqueous solution of the halogen acid is required, advantage is taken of the reaction which proceeds readily in aqueous solution. Although the actual amount of iodine in solution is very slight, in the presence of the solid the equilibrium I2 (solid) ^ T; (dissolved) always tends to be set up, and any reagent which removes the dissolved iodine will automatically lead to the solution of more of the solid. By this means one may obtain a solution (of hydriodic acid) of the sp. gr. 1-56, which merely requires filtering before use. The reaction proceeds equally well in the case of bromine. Physical Properties. — The physical properties of hydrogen bromide and iodide show a marked resemblance to the properties of hydrogen chloride, except in so far as these properties are modified by their increased atomic weight. Both are colourless, fuming gases, readily condensable to liquids (hydrogen bromide condenses at — 73° ; hydrogen iodide at 0°). Both gases are very soluble in water, 100 gm. of water dissolving 210 gm. of hydrogen bromide at 10° and about 240 gm. of hydrogen iodide. Hydrobromic and hydriodic acids behave on evaporation in exactly the same way as hydrochloric acid, i.e. they form a 166 AN INORGAiSrtC CHEMISTRY mixture of maximum boiling point. With hydrobromic acid the maximum boiling mixture contains 48 per cent, of hydrogen bromide (B.P. 126°), with hydriodic acid 57 per cent, of the halide (B.P. 127°). Chemical Properties. — The chemical properties of these acids also show a strong similarity to those of their relative, hydrogen chloride. Both gases have no action in the dry state upon litmus, though, when dissolved in water, they form strong acids. In general, however, the properties are somewhat less clearly marked in the case of the hydr-acids of bromine and iodine, but the same gradation of property is noted here as in the elements themselves, i.e. hydrogen bromide always occupies a position intermediate between hydrogen chloride and hydrogen iodide. This is seen in the action of heat upon these compounds, for while hydrogen chloride is extremely stable towards heat, hydrogen bromide undergoes appreciable decomposition at 800°, and hydrogen iodide suffers 21 per cent, decomposition at 448°. So, too, the action of oxidising agents upon the halogen hydr- acids shows a steady change in this family of elements. To obtain chlorine from hydrogen chloride or from hydrochloric acid requires a strong oxidising agent such as manganese dioxide, but iodine can be liberated from hydriodic acid by the action of atmospheric oxygen. It is owing to this oxidising action of the air that hydriodic acid and aqueous solutions of the iodides become brown on exposure to air, the liberated iodine at once dissolving to form the acid Hlg or the corresponding salt. 4HI + 02-^2H20+2l2 2HI + 2l2->2Hi3 The strong tendency of hydriodic acid to combine with oxygen, i.e. to act as a reducing agent, is striltingly illustrated by the vigour of the reaction when a drop of fuming nitric acid is dropped into a jar of hydrogen iodide. Copious violet fumes of iodine and brown fumes of the oxides of nitrogen are evolved. 2HI +2HN03->2H20 +N2O, + 1^. (HAN A) Oxidation and Reduction Hitherto oxidation has been defined in an elementary way as the addition of oxygen to an element (or compound), or the BROMINE AND IODINE 167 removal of hydrogen from a compound. It is now proposed to widen this definition. Oxidation is the term applied to all reactions wherein the valence of any element towards oxygen is increased, or the valence of any element towards hydrogen is decreased ; conversely, reduction implies that the valence of any element towards hydrogen is increased or the valence towards oxygen is decreased. * This definition will now be applied to a number of typical reactions. (a) 4HCl+MnO.,^>MnCl2 + 2H20+2Cl2. The valence of chlorine towards hydrogen in the compound hydrogen chloride, is one ; the valence of chlorine towards hydrogen in elementary chlorine is nil ; hydrogen chloride has therefore been oxidised to chlorine. The valence of manganese towards oxygen in manganese dioxide is four, towards oxygen in the compound, manganese chloride, two ; manganese has therefore undergone reduction during the reaction. At first sight the statement that the oxygen valence of manganese in the compound MnClj is two, may appear un- warrantable, until one recalls that this chloride is derived from the oxide MnO, wherein the valence of the metal towards oxygen is undoubtedly two. As an aid to this conception of oxidation and reduction recourse must frequently be made to the equivalence of the elements, e.g. hydrogen is capable of replacement by metals, oxygen by such elements as sulphur, chlorine, etc. In applying the above view of oxidation, the student is advised to look upon an oxy-acid as a compound of water and an acid anhydride, e.g. H2S04=H20,S03, whilst the salt of an oxy-acid is looked upon as being compounded of a basic and an acidic oxide, e.g. CuSOi = CuO,S03, KNO3 = JK20,N A. (b) 2KBr + Cl2->2KCl-f Br^. In the compound KBr, the valence of bromine towards hydrogen is one, in Bra it is zero ; the bromide has therefore been oxidised to bromine. An application of the same principle to clilorine shows that its hydrogen valence has been raised from nil in CI 2 to one in KCl ; it has therefore been reduced to form chloride. (c) 14HC1 +K2Cr20,->2KCl +2CrCl3 + 7H2O +3CI2. (K20,2Cr03) * It is assumed as an axiom that the valen2e of all elements in the elementary state towards hydrogen or oxygen is zero. 168 AN INORGANIC CHEMISTRY The oxygen valence of chromium has here been reduced from six in KaCraO, to three in CrCls (C1=|0, cf. HCl, HjO) ; the potassium dichromate has consequently been reduced in order to oxidise the hydrogen chloride to chlorine. (d) 2HBr + H2S04^2H20+SO,+Br2. (H^O.SG,) In this reaction the valence of sulphur has fallen from six in SO 3 to four in SO 2. The oxidation of hydrogen bromide has therefore been brought about by the reduction of sulphuric acid to sulphur dioxide. (e) H2S+Br,^2HBr + S. The decrease of the hydrogen valence of sulphur from two to zero, and the increase of the hydrogen valence of bromine from zero to one shows that the hydrogen sulphide has been oxidised to sulphur by the action of the bromine. (/) 2HI+2HN03->2H20 + I, + N,04. (H,0,N.A) The decrease of the hydrogen valence of iodine in this reaction, coupled with the decrease of the oxygen valence of nitrogen, indicates that the hydrogen iodide has been oxidised by the nitric aciil , which has itself been reduced to nitrogen tetroxide. (fif) 2KC103-^ 2KCI + 3O2. (KAC'IA) In KCIO3 the chlorine has a valence of five towards oxygen, whilst in KCl its valence towards hydrogen has risen to one ; the chlorine present in the chlorate has therefore been reduced to the extent of six units. In the oxide CI2O5, oxygen, as usual, is di-valent, hence chlorine is penta-valent, so that the valence of the oxygen atom towards chlorine (or its equiva- lent, hydrogen) is two; in 0. it is nil, hence the element oxygen has undergone oxidation. In all the above reactions the student will notice that the phenomena of oxidation and reduction are inseparable, they are mutually dependent. In applying this method of treatment to processes involving oxidation and reduction, the postulate is laid down that hydrogen is always mono-valent and that oxygen is always di-valent. Under such an assumption, all such processes as the formation and decomposition of the peroxides and ozone are excluded. FLUORINE 169 This theory of oxidation is as much based upon the mono- valence of hydrogen and the di- valence of oxygen as the chemistry of the carbon compounds is built upon the assumption of the tetra-valence of the carbon atom (see p. 363). Auto -oxidation. — In the reaction G\ + H^O ^=:± HCl + HCIO (iH,0,aO) chlorine has a zero valence towards oxygen and hydrogen in the molecule CIg, in the compound HCl its hydrogen valence has risen to one, whilst in the compound HCIO its oxygen valence has become one. This indicates that one atom of chlorine has been reduced to form hydrochloric acid, and the second atom in the chlorine molecule has been oxidised to form hypochlorous acid. In such cases, where part of a molecule has been oxidised at the expense of the rest of the molecule, the term auto-oxidation is applied. Other examples of auto-oxidation occur in the study of the oxy-compounds of the halogens, nitrogen, phosphorus, etc. FLuoEnsrE Occurrence. — Fluorine does not occur free in nature, but in the form of fluorspar, CaF,, cryolite, a double fluoride of sodium and aluminium, NagAlF^, tourmaKne, fluor-apatite, etc., it is widespread. Discovery and Preparation. — As far back as 1670, evidence of the existence of a new element in fluorspar was put forward, but it was not tUl 1807 that Gay Lussac and Thenard succeeded in isolating hydrogen fluoride. All attempts to make the com- pound yield up its fluorine by methods analogous to those used for the preparation of chlorine, etc., failed, until in 1886 H. Moissan achieved the great feat of isolating this extraordinarily active gas. The electrolysis of solutions of hydrochloric acid leads to the evolution of chlorine at the positive pole, but in the case of hydrofluoric acid oxygen is liberated at this pole, while the anhydrous hydrogen fluoride does not conduct the electric current. Moissan overcame this difficulty by adding potassium fluoride to the anhydrous Uquid. This solution of potassium hydrogen fluoride was found to conduct the electric current and to evolve hydrogen at the cathode or negative pole, fluorine at the anode or positive pole. The electrolysis appears to break down the double fluoride thus : 170 AN INORGANIC CHEMISTRY 2KHF2-^ 2HF + F^ + 2K, the potassium immediately reacting with the solvent, hydrogen fluoride, forming potassium fluoride and hberating hydrogen. Owing to tlie great activity of the fluorine, the electrolysis was first carried out in a platinum-iridium vessel, later in one of copper. The copper vessel rapidly acquired a coating of copper fluoride which protected the tube from further action. The electrodes were made of platinum-iridium alloy. The whole apparatus has to be made thoroughly moisture- and air-proof, and in order to reduce the tendency to chemical reaction, the apparatus was immersed in a bath of methyl chloride ( — 23°). Properties. — Fluorine is a pale yellow gas which condenses to a liquid boiling at — 186° Its density is 37-7 (hydrogen=2). As the atomic \\eight is 19, it follows that the molecule is diatomic. Chemically speaking, fluorine is the most active element known. It combines with every element except oxygen and the inert gases (argon, etc.). With hydrogen it combines explo- sively, and such is the energy of this reaction that these elements react even at the temperature of liquid fluorine ( — 186°). Fluor- ine attacks compounds containing hydrogen with the hberation of hydrogen fluoride. Elements, such as iodine, bromirle, phosphorus, arsenic, combine with incandescence. AU metals are acted upon by fluorine, many of them taking fire spon- taneously. Chlorine is displaced from chlorides by the action of fluorine, i.e. fluorine is a stronger oxidising agent than chlorine, and therefore stands at the head of the halogen family so far as oxidising power is concerned. Hydrogen Fluoride Preparation. — The aqueous solution of this substance is generally prepared commercially by the action of hot sulphuric acid upon powdered fluorspar (CaFo). Leaden retorts are generally used, and the evolved gases are collected in water and stored in lead, rubber or paraffin wax bottles. CaFa + H^SOj -^ CaSO, + H^F, f Anhydrous hydrogen fluoride is, however, best obtained by heating potassium hydrogen fluoride in a platinum retort. 2KHP„ --> 2KF + HoFa t FLUORINE 171 Physical Properties. — Hydrogen fluoride is a colourless liquid which fumes strongly in air. It boils at 19-5°. Its physiological effects are unpleasant, for not only are the vapours poisonous, but contact of the Hquid with the skin produces ulcerous sores. It mixes with water in all proportions, forming hydrofluoric acid, and like the other hydr-acids of the halogens, it forms a maximum boiling point mixture of constant composi- tion (B.P. 120°, 37 per cent. HP). The vapour density of hydrogen at 30° indicates that the molecular weight is 40, from which it is to be concluded that the formula representing the composition at this temperature is H2P2. With rising tem- perature the density falls, until at 88° the molecular weight has fallen to 20 (formula HF). Evidence has also been adduced that the formula of the vapour at the temperature of the boiling hquid (19-5°) is H3F3. Compounds of this type, which possess double and triple molecules, are said to exhibit association. When the difference between the properties of such compounds is sharply defined, polymerisation is said to occur, e.g. acety- lene, C2H2, and benzene, GgHe. Chemical Properties. — Probably owing to the different molecular structure of hydrogen fluoride as compared with the other hydrogen haUdes, the action upon metals is less vigorous. Potassium and sodium dissolve in the anhydrous hquid hberating hydrogen, while, among other metals, silver dissolves in the aqueous solution of this substance. Another result of the molecular structure of this compound is that acid salts may be obtained by the replacement of one hydrogen atom with a monad element, e.g. KHF^. The reaction between sUica or silicates (glass, etc.) and hydro- gen fluoride (either in the vapour or dissolved state) is of considerable importance. SiO^ + 2HJ2 -^ SiF^ + 2H2O CaSiOs + SH^Fa-^ SiF, + CaF^ -f- 31i,0 If a glass object is exposed to the action of hydrogen fluoride, the surface is eaten into or etched, owing to the escape of the gaseous silicon fluoride and the crumbling away of the residual calcium and sodium fluorides. Advantage is taken of this in the process of etching. The surface of the glass object to be etched is covered with paraffin wax, and the figures or marks which it is desired to etch into the glass are made in the wax 172 AN INORGANIC CHEMISTRY with a sharp instrument. On exposing the surface to the action either of hydrofluoric acid or of the vapour of hydrogen fluoride, the glass is eaten into where the design was traced out. Compounds of the Halogens with each other Iodine and chlorine form a monochloride ICl and a trichloride ICI3 — the former is a red crystalline substance readily decom- posed by water, the latter a yeUow compound forming long needle-shaped crystals. On heating, the trichloride dissociates into chlorine and the monochloride. In the presence of re- stricted amounts of chlorine, one obtains the monochloride, whilst excess of chlorine, acting upon iodine or iodine mono- chloride, gives the trichloride. Iodine monobromide, IBr, iodine pentafluoride, IP5, and bromine trifluoride, BrFj, are known. The Halogens as a Family The elements fluorine, chlorine, bromine and iodine have been found to possess so many properties in common that the attempt to study them as a group, a family, has already been justified. The following table brings out the gradual change in most of the physical and chemical properties of these elements as the atomic weight increases. TABLE 23 Pkopbrties of the Halogens Property. Fluorine. Chlorine. Bromine. Iodine. Atomic weight Colour State of aggregation . Boiling point .... Sp- gr Solubility per 100 c.e. water at 20° . . . Dissociation .... Oxidising power . 19 Pale yellow Gas -186" M4 (liq.) Decomposes water Not tested 35-46 Greenish yellow Gas ^33-6° 1-55 (liq-) -0001 Decreas 79-92 Reddish brown Liquid 59" 3-19 (liq.) •032 Increases es 126-92 ^'iolet Solid 183" 4-9 (solid) •00015 > Affinity for oxygen . Increases Affinity for hydrogen Decreases > THE HALOGENS AS A FAMILY 173 For the purpose of comparison, the affinity for oxygen has been included in this table, though this property will not be discussed until the oxy-compounds are dealt with. TABLE 24 Properties of the HAiOGEN Hydr-acids Property. Hydrogen fluoride. Hydrogen chloride. Hydrogen bromide. Hydrogen iodide. Mol. weight at 100°. . B.P Solubility in water . Dissociation measurable at Heat of formation in calories 20 19-4° 35% 38-5 (gas) 36-46 —83-7° 42% 1500° 22-0 (gas) 80-93 -68-7° 49% 800° 12-3 (gas) 126-93 -35-7° 57% 180° — 6-0(solid) Reducing power Increases The solubiUty of the salts of these acids often shows a break in regularity. Thus, the silver salts of hydrochloric, hydrobromic and hydriodic acids are almost insoluble, but show a steady decrease in solubility with increasing atomic weight. Silver fluoride is, however, excessively soluble, 182 gm. dissolving in 100 gm. of water at 20°. On the other hand, calcium fluoride is almost insoluble while the other calcium salts are all freely soluble. Possibly this is connected with the tendency of hydrogen fluoride to associate in solution. Questions. 1. Discuss the phenomena of oxidation and reduction, illustrating your answer by reference to the reactions Cl2+2NaOH -> NaCl+NaaO+HjO US + Cl^-^ 2B.CI+S. 2. Describe the conmieroial preparation of iodine. What impurities is commercial iodine likely to contain, and how would you remove those impurities ? 3. Compare and contrast the reactions between sodium chloride and sulphuric acid (hot) and sodium iodide and sulphuric acid (hot). 4 Compare the halogens with each other, with special reference to their oxidising power. Illustrate by examples. 5. How would you prepare a concentrated solution of hydrobromic acid. Give full details and reasons. 6. Compare the reactions 2HI+a2^'2Ha-fIj 2HI+H2SO, -^ 2H,0 + S02+l2- 7. Describe a suitable method for preparing liquid hydrogen bromide. 8. Account for the difficulty of isolating the element fluorine. CHAPTER XII THE OXIDES AND OXY-ACIDS OF THE HALOGENS Nomenclature of Acids. — The majority of non-metals give rise to more than one oxy-acid, so that a systematic method of naming them is desirable. When two such acids are known, that derived from the lower oxide is referred to as the -ous acid, and from the higher oxide is derived the -io acid. 502, sulphur dioxide forms H2SO3, sulphurous acid. (sulphurous anhydride) 503, sulphur trioxide forms H2SO4, sulphuric acid. (sulphuric anhydride) When more than two acids are known, the prefixes hypo- and per- are called into requisition. H2S2O4, hypo-sulphurous acid ; (HS04)2, per-sulphuric acid. The Oxides and Oxy-Acids of Chlorine Chlorine forms three oxides and four weU-defined acids. Their inter-relations are brought out in the following table : Oxide. Oxy-acid. CI2O, chlorine monoxide (hypochlorous anhydride) — >-HC10, hypochlorous acid. - HCIO4, perchloric acid, (perchloric anhydride) On solution in water, the monoxide and heptoxide give rise to hypochlorous and perchloric acids respectively, whilst chlorine dioxide forms a mixture of chlorous and chloric acids. Chlorine Monoxide and Hypochlorous Acid. Chlorine monoxide is readily formed by passing dry chlorine 174 HALOGEN OXIDES AND OXY-ACIDS 175 over precipitated mercuric oxide, previously heated for about an hour to 400° in order to coarsen the grain and so lessen the speed of the reaction. 2a + 2HgO -^ Hg,OCl, + C1,0 The reddish yellow gas is readily condensed to a liquid. Both gas and hquid are unstable, a drop of the liquid exploding on contact with paper, dust, phosphorus, sulphur, etc. The gas is very soluble in water (1 c.c. of water dissolving 200 c.o. of chlorine monoxide at 0°), and the solution is found to contain hypochlorous acid. H20+Cl20^2HC10. Liquid chlorine monoxide, although so unstable, may be safely distilled under reduced pressure. Owing to the danger involved in the preparation of hypo- chlorous anhydride, the acid is generally prepared indirectly. Chlorine is allowed to act upon water until the equihbrium represented in the equation CI2 + H^O ;=± HCl + HCIO has been reached. If either of the acids is removed, this equih- brium will be disturbed, and more chlorine will react in order to restore the equUibrium (an example of Le ChateUer's Law). The addition of chalk (calcium carbonate) achieves this result, for, while the very weak hypochlorous acid has no action upon this salt, the hydrochloric acid is neutralised and thus removed from the system. CaCO, + 2HCl-> CaCla + H^O + CO^. The equihbrium is disturbed, and in the endeavour to restore this, the chlorine reacts with the water and thus steadily increases the concentration of the hypochlorous acid. Freshly precipitated mercuric oxide may also be used in place of calcium carbonate. The solution containing the hypo- chlorous acid is then distilled and a pure product obtained. Another method that brings about a change in the equili- brium, CI2 -f H^O ^=^ HCl + HCIO, is the addition of a base, e.g. sodium hydroxide, to the system. In this case both the acids are neutraUsed and a solution of sodium chloride and sodium hypochlorite results. For many 176 AN INORGANIC CHEMISTRY purposes the presence of the chloride is in no way objectionahle, while the separation of hypochlorous acid from this mixture is easily effected. The addition of an acid such as nitric or boric acid to the solution brings about a chemical change indi- cated in the equations (see p. 439). NaClO + HNOa-^ NaNOs + HCIO NaCl + HN03-> NaNOs + HCl. Of the two acids HCIO and HCl, the former is exceedingly weak. Under these circumstances the relatively strong nitric acid displaces the weak hypochlorous acid almost entirely from its salt, whilst the amount of the stronger hydrochloric acid set free is so small as to be negligible. In order to attain ideal conditions for the separation, nitric acid is added, equivalent in quantity to the amount of hypochlorous acid originally present as hypo- chlorite. Careful distillation will then give a dilute solution of hypochlorous acid. If the neutralisation is effected by a solution of potassium hydroxide or carbonate, the resulting solution of potassium chloride and hypochlorite is often known as Eau de Javd. This solution is frequently used in households as a bleaching and cleansing agent. Properties of Hypochlorous Acid. — Hypochlorous acid is so unstable that it cannot be obtained except in very dilute solution. It shows a great tendency to break down in various ways. Thus, under the action of sunUght, oxygen is evolved. 2HC10->2HCl + 02. But there is a more important decomposition which hypo- chlorous acid undergoes — its auto-oxidation into chloric and hydrochloric acids. 6HC10 -> 4HC1 -f 2HCIO3. Part of the hypochlorous acid is reduced to hydrochloric acid in order that sufficient oxygen may be available for the oxidation of the remairung hypochlorous acid to chloric acid. This auto- oxidation is also greatly accelerated by heat, and as we shall see, it has an important commercial bearing. The hydrochloric acid formed in the above way, takes part in the reaction HCl + HCIO -> HjO + Cla, and as the concentration of hydrochloric acid steadily increases, HALOGEN OXIDES AND OXY-ACIDS 177 chlorine will accumulate in the system. Sooner or later, the concentration of the dissolved chlorine reaches the saturation point and chlorine will escape. The looseness with which the atom of oxygen is held in hypo- chlorous acid accounts for the great activity of this compound as an oxidising agent. This is especially well illustrated by its action upon bromine and iodine : 2HC10 + 12-> 2HI0 + CI2. The hypoiodous acid immediately passes into the more stable system iodic and hydriodic acids (cf. the auto-oxidation of HCIO). In this reaction the replacement of chlorine by iodine is an apparent contradiction of its behaviour in the reaction 2HI+Cl2->2HCl + l2, but the application of the conception of oxidation and reduction already developed brings these apparently contradictory reac- tions into line. We have learnt that the action of chlorine upon hydrogen iodide involves the oxidation of the hydrogen iodide by the chlorine {cf. 2KBr + Clj, p. 167). In the equation : 2HC10 + 1, -^ 2HI0 + CI2 (H,0,C1,0) (H,OJ,0) the oxygen valence of the iodine has risen from nil to one, whilst the oxygen valence of the chlorine has fallen from one to nil. In short, the iodine has undergone oxidation and the chlorine has been reduced. There is thus complete agreement in the behavioiu^ of these halogens in the above apparently contra- dictory reactions. The oxidising power of hypochlorous acid enables this reagent to play an important part in bleaching. Linen and cotton consist essentially of compounds containing carbon, hydrogen and oxygen, coloured by small quantities of coloured compounds. These coloured compounds are much more easily oxidised than the carbo-hydrates forming the ground mass of the fabric, and the products of oxidation are colourless. Hypochlorous acid is the oxidising agent in general use for bleaching the colour out of such fabrics. It is generally generated by the action of very dilute sulphuric acid upon bleaching powder, which may tem- porarily be considered to be calcium hypochlorite. The fabric is thoroughly washed to remove all traces of oil and grease, N 178 AN INORGANIC CHEMISTRY immersed in a bath of bleaching powder, and finally in very dilute sulphuric acid. Care has to be exercised in keeping the strength of the reagents so low that no reaction with the fabric occurs. Bleaching Powder. — The most important salt of hypo- chlorous acid is bleaching powder. If chlorine is bubbled through a solution of a monovalent base, e.g. potassium hydroxide, the reaction : CI2 + H2O -> HCI + HCIO + + KOH KOH I KCI+KCIO+2H2O occurs, but if a bivalent base like calcium hydroxide is utilised, a mixed salt is formed : OH HCI CI Ca/ + ->Ca<( +2H2O ^OH HCIO ^ClO This is the bleaching powder of commerce. It is manufactured on a large scale by the action of chlorine upon slaked hme. In older plants the hme is spread in layers in a large chamber, and exposed to the action of chlorine. The hme must be occasionally stirred in order to expose a fresh surface. The unabsorbed chlorine left in the chamber is removed by blowing in a spray of fine slaked hme. In many of the modern factories mechanical apparatus for the absorption of chlorine is in use (Fig. 68). Several iron cyhnders, each provided with revolving paddles, are set one above the other. Slaked hme passes slowly through the upper cylinder, then into the second cyUnder and so on. Chlorine enters the bottom cyhnder and the unabsorbed gas escapes from the top cylinder. The usual amount of chlorine absorbed is about 33 to 38 per cent. Analysis shows that bleaching powder prepared in this way always contains more or less free chlorine. For a long time the constitution of this salt was the subject of considerable discussion, many considering it to be a mixture of calcium chloride, CaCl^, and calcium hypochlorite, Ca(C10)2, in molecular proportions. Such a substance should, however, be hygroscopic, owing to the pressnce of calcium chloride ; then, too, calcium chloride is soluble in alcohol, but this solvent has no HALOGEN OXIDES AND OXY-ACIDS 179 action upon bleaching powder, so that the evidence favours the constitution : /^ Ca<( ^ClO. On treatment with cold water, a strongly alJiaUne solution is obtained with the precipitation of insoluble calcium hydroxide. 2CaCl(C10) + 2H0H-> CaCl^ + Ca(0H)2 + 2HC10. Acids in the dilute state Uberate hypochlorous acid ; when concentrated, chlorine is evolved. Ca/ +HC1-^ CaCla + HCIO. ^ClO ,n n n □ n ^n^ n n n n n Bleaching powder drops out. Lime enters Fig. 68. In the presence of an excess of the acid (e.g. hydrochloric), the equihbrium : HCl + HCIO ^=± H^O + CI2 will be forced to the right, with the consequence that the solution rapidly approaches saturation point and chlorine is soon evolved. So, too, if sulphuric or even carbonic acid is used. 180 AN INORGANIC CHEMISTRY' CI Ca/ + H2SO4 -> CaS04 + HCl + HCIO. ^ClO j Cl2+H,0 The decomposition of hypochlorites in the presence of small quantities of cobalt nitrate forms a useful example of catalysis, especially as the mechanism of the reaction is clearly under- stood. Cobalt nitrate, in the presence of h3rpochlorites, Hberates small quantities of cobaltous oxide. The hypochlorite then reacts with the cobaltous oxide, thus : NaClO + 2CoO-^NaCl + C02O3 (cobaltic oxide) 2Co203->4CoO + 02. The cycle of reactions indicated in these equations explains sufficiently the acceleration in the decomposition of the hypo- chlorites brought about by traces of cobalt compounds. Chloric Acid. This acid is most easily prepared by adding a calculated quantity of sulphuric acid to barium chlorate. Ba(C103)2 + H^SOi-^BaSOi i + 2HCIO3. After filtration, the clear solution is concentrated by evaporation in vacuo up to 40 per cent. At concentrations above this, spontaneous decomposition ensues, the products of decomposi- tion being chlorine, oxygen, perchloric acid and water. The aqueous solution is a powerful oxidising agent, acting vigorously upon wood, paper, etc. Similarly, iodine is oxidised to iodic acid. 2HCIO3 + 1„-^ 2HIO3 + CI2 (H^CCIA) (H,0,I,Oa.) In this reaction, as in all other reactions in which chlorine and iodine are involved, chlorine has undergone reduction whilst iodine has been oxidised, that is, chlorine always shows itself to be a stronger oxidising agent than iodine. Chloric acid is a strong bleaching agent, even when dilute. The anhydride of this acid, CI2O5, has not yet been isolated. Properties of Chlorates. — Chloric acid, being monobasic, forms salts of the type M'CClOa), M"(C103)2, etc., where M^, M" denote monovalent and divalent radicals respectively. These salts are all powerful oxidising reagents and are more stable than the acid itself. All chlorates evolve oxygen on HALOGEN OXIDES AND OXY-ACIDS l8l heating. Their vigour aa oxidising agents is illustrated by the following experiments : 1. If a few drops of a solution of phosphorus in carbon disuU phide are poured upon potassium chlorate, a loud explosion will occur so soon as the carbon disulphide has evaporated. 2. A piece of phosphorus is placed in contact with crystals of potassium chlorate under water. On pouring a few drops of sulphuric acid on to the crystals by means of a long funnel, the phosphorus bursts into flame. All the chlorates are soluble, one of the least soluble being potassium chlorate. An important reaction of the chlorates is brought about by the action of heat. If potassium chlorate is heated, the sub- stance melts in the neighbourhood of 350° and bubbles of oxygen are evolved, potassium perchlorate being formed {q.v.). 2KC103->KC10i + KCl + 02^ Shortly after, the melt solidifies. This is due to the fact that the perchorate has a much higher melting point than the chlorate. If the temperature be stUl further raised, the mass again melts at a temperature over 600°, and soon afterwards a further evolution of oxygen sets in. The perchlorate is now decomposing in accordance with the equation KC104-^KCl + 202 Preparation of the Chlorates. — Commercially speaking, potassium chlorate stands out as the most important salt of chloric acid. The most utilised method of preparation of this salt is based upon the reactions summarised in the equations : 3CI2 + 6K0H-^ 3KC10 + 3KC1 + SH^O 3KC10 -> KCIO3 + 2KC1 (fK20,Cl,0) (4K,0,C1 A) If the solution is at all warm the decomposition of the hypo- chlorite proceeds rapidly. An analysis of the last equation shows that two molecules of hypochlorite are reduced to potassium chloride in order that one molecule may be oxidised to chlorate, I.e. auto-oxidation has occurred, a common phenomenon in the decomposition of the oxy-compounds of the halogens. On cooling the above solution, the potassium chlorate crystallises out, but the method in this form has one fatal drawback. Of the original six molecules of valuable potassium hydroxide five 182 AN INORGANIC CHEMISTRY are converted into the less valuable by-product, potassium chloride, and this is economically unsound. The essential feature of the above process is to have present any alkali with which to neutralise the hydrochloric and hypochlorous acids formed by the action of chlorine upon water (see p. 178). As a step towards economy, J. von Liebig suggested the use of the very cheap slaked lime in order to effect the desired neutralisation. 6CI2 + 6Ca(OH)2->5CaCl, + Ca(C103)2 + 6H,0. The resultant solution of soluble calcium salts was then con- centrated and the calculated amount of potassium chloride added. By double interchange potassium chlorate was formed, 2KC1 -f Ca(C103) 2-^ CaCla + 2KCIO3 and as this salt is far less soluble than the other salts present in solution, a very effective separation of potassium chlorate from the mother liquor can be effected by concentration followed by cooUng. The crystals obtained in this way are then purified by fractional crystallisation. This method of manufacturing potassium chlorate has been largely superseded during recent years by the modern electro- lytic method. Diu-ing the electrolysis of aqueous solutions of potassium chloride, chlorine is Uberated at the anode (positive pole), potassium at the negative pole or cathode. The inter- action between the potassium and the water leads to the forma- tion of potassium hydroxide with the liberation of hydrogen. If the conditions are such as to promote the mixing of the potassium hydroxide round one pole and chlorine generated round the other, all the conditions necessary for the formation of hypochlorite are present, whilst, if the solution is kept hot, the so-formed hypochlorite will be instantly converted into chlorate. Owing to the sUght solubihty of potassium chlorate, many prefer to electrolyse solutions of sodium chloride (99 gm. of sodium chlorate dissolve at 20° in 100 c.c. of water, as com- pared with 7-2 gm. of potassium chlorate). The sodium chlorate is afterwards converted into the potassium salt by treatment with potassium chloride, with subsequent fractional crystaUisa- tion of the potassium chlorate (see p. 480 for further details). Chlorine Dioxide ; Chlorous Acid, Chlorites. When potassium chlorate is cautiously heated with sulphuric acid, a reddish yellow gas of extreme instability is liberated. HALOGEN OXIDES AND OXY-ACIDS 183 The first reaction is the formation of chloric acid, with sub- sequent auto-oxidation to form perchloric acid, chlorine dioxide and water, 3HCIO3 -^ HC104 + 2C10,+H20. (fH^O.ClA) (iH,0,CK0,) Chlorine dioxide is a heavy gas, readily liquefied. It 's extremely unstable, detonating if a hot wire is brought into it. Onec.c. of water dissolves 20 c.c. of chlorine dioxide at 4°. The reaction is not one of mere physical solution, for the solution is found to contain a mixture of chlorous and chloric acids, i.e. the usual auto-oxidation has occurred. 2C102 + H20-^HC102 + HCIO3 (iH^O.Cl^O,) (iH^O.ClA) Thus, solution of this oxide in potassium hydroxide leads to the formation of potassium chlorate and chlorite. Chlorous acid is very unstable and has not been prepared in the pure state, though several of its salts have been isolated. These retain the property of instability associated with the parent acid. Owing to the looseness with which the oxygen atom is held within the chlorous radicle, chlorites form strong oxidising and bleaching agents. Perchloric Acid ; Chlorine Heptoxide (Perchloric Anhy- dride). Perchloric acid may be obtained by the auto-oxidation of chloric acid (q.v.), more often, however, by the action of sul- phuric acid upon perchlorates. Potassium per chlorate is first prepared by heating the chlorate up to the point of liquefaction (see p. 181), and after cooling, the chloride and perchlorate are separated by fractional crystallisation. KaO^ + H,S04-> KHSO4 + HCIO4. The perchloric acid hberated in this way is removed by distilla- tion, and concentrated by evaporation in vacuo. It is a volatile, fuming liquid which can be boiled with safety under reduced pressure. The pure acid is distinctly unstable, often exploding spontaneously. A drop of perchloric acid falling upon a piece of charcoal causes a violent explosion. Seventy per cent, solutions are, however, perfectly stable. The perchlorates are not decomposed by hydrochloric acid, nor are they reduced by sulphur dioxide. On heating very strongly, they evolve oxygen 184 AN INORGANIC CHEMISTRY and pass into chlorides. Potassium perchlorate is used in quantitative analysis as a means of estimating potassium, as this salt is almost insoluble in the presence of alcohol. Chlorine heptoxide (perchloric anhydride), CIjO,, is prepared by the cautious addition of phosphoric anhydride (P2O5) to perchloric acid, the reaction being carried out in a vessel immersed in a freezing mixture. After distillation, a liquid boiling at 82° is obtained — chlorine heptoxide. The relationship of this explosive Uquid to perchloric acid is shown by its behaviour on solution in water. C1A+H20->2HC104. Thermochemistry of the Oxy-Acids of Chlorine. The decomposition of hypochlorous acid and of chlorine monoxide results in the liberation of a large store of heat. When these exothermal substances are used as oxidising agents, the energy liberated at the moment of decomposition must be added to the energy available when free oxygen is utilised to effect the same oxidation, hence these substances form much more vigorous oxidising agents than does free oxygen. Thus, it often happens that the chemical energy available when oxygen is the oxidising agent is not sufficiently large to bring about a desired oxidation process, viz. the bleaching of fabrics, whereas the greater available energy of the hypochlorous acid molecule is frequently sufficient to bring about the reaction. The following thermochemical equations represent the heat evolved per atom of oxygen available when the molecule of the oxy-acid of chlorine breaks down into hydrochloric acid and oxygen. HCIO aq. — >HClaq.+ 0+ 9-3 Cals. or 9-3 Cals. per atom of 0. HCIO3 aq. — ^HCl aq. +30+15-3 Cals. or 5-1 Cals. per atom of 0. HCIO4 aq. — ^HCl aq. +40+ 0-7 Cals. or 0-17 Cals. per atom of 0. The greater amount of available energy in hypochlorous acid is in itself sufficient to explain the activity of this compound as an oxidising agent. The old assumption that atomic or nascent oxygen, liberated by the decomposition of hypochlorous acid HC10->HCl + is the active oxidising agent, leads nowhere — what one must consider is the amount of energy associated with the oxygen atom at the moment of liberation. HALOGEN OXIDES AND OXY-ACIDS 185 OXY-COMPOUNDS OF BROMINE. Up to the present, no oxides of bromine have been isolated, but the acids, hypobromous and bromic, are known. Hypobromous Acid has been prepared by the action of bromine water upon precipitated mercuric oxide (cf. hypo- chlorous acid). 2Br 2 + HaO + HgO -> HgBr ^ + 2HBrO. The acid is unstable, but can be concentrated by distillation in vacuo. It is a strong oxidising and bleaching agent. Salts of hypobromous acid are prepared by the method utilised for the related acid, hypochlorous acid. Bra + 2K0H-^ KBr + KBrO + H„0. Bromic Acid is made either by the action of sulphuric acid upon barium bromate (cf. chloric acid), or by the oxidation of bromine by chlorine water Br^ + 5CI2 + GHaO-^ 2HBr03 + lOHCl. Here again, it is to be noted that bromine is oxidised, whilst chlorine undergoes reduction. Bromine, like iodine, consequently displays in its compounds a greater affinity for oxygen than does chlorine. The properties of bromic acid ara quite similar to those of chloric acid. Thus it oxidises iodine to iodic acid, 2HBr03 + 12-> 2HIO3 + Bra (HaO.BrA) (HaO,IA) i.e. the oxygen valence of the bromine has fallen from five to nil, whilst the oxygen valence of the iodine has increased by the same amount. This confirms our previous conclusion that bromine is a stronger oxidising agent than iodine. This experi- ment also illustrates the fact that iodine has a greater affinity for oxygen than has_ bromine. OxY-AciDS AND Oxides of Iodine Oxides. Oxy-acida. ( ) HIO hypoiodous acid I2O5 HIO 3 iodic acid (I2O,) „ -^* \ periodic acids I.O4 (— ) Only two oxides are known, the pentoxide and the tetroxide. 186 AN INORGANIC CHEMISTRY Hypoiodous Acid shows no fundamental difference from the corresponding hypochlorous and hypobromous acids. Both in preparation and properties the kinship is so strongly marked that no further comment is required. Iodic Acid is prepared by methods similar to those used for the preparation of bromic acid : Ba(I03)2 + H2S04^BaS04 + 2HI03 I2 + 5CI2 + 6H2O -> 2HIO3 + lOHCl Another method is to oxidise the iodine with hot nitric acid : I2 + 6HN03-> 2HIO3 + 2N2O3 + 2NO2 + 2H2O (3H,0,N,0,) (H^OJ.O^) The potassium salt of iodic acid is readily prepared by heating iodine and potassium chlorate with an aqueous solution of an acid. The essential reaction is represented in the equation : 2KCIO3 + Ij^-> 2KIO3 + CI2 (K.CCl.Os) (K.OJA) (Note the reduction of chlorine by iodine.) Iodic acid is non-volatile and may be obtained in the above experiments by evaporation, and comes down as a white crystal- line soUd, soluble in water. The acid first reddens litmus and then bleaches it. Iodic acid may be heated up to 170°, when it decomposes, leaving the pentoxide behind. This anhydride, laOs, is a white crystalline soUd of considerable stability, which decomposes into its constituents in the neighbourhood of 300°. Owing to the affinity with which oxygen is held by iodine, it is to be expected that iodic acid is a less powerful oxidising agent than either bromic or chloric acid. Experiment bears this out. Amongst the reducing agents which react with this acid may be mentioned hydrogen iodide, hydrogen sulphide and sulphur dioxide. HIO3 + 5HI-^ SH^O + 31, 2HIO3 + 5H2S -> 5S + 6H2O + 1, 2HI0, + 5SO2 + 4H2O -> I2 + 5H2SO4 (H.O.I^O,) (5H,0,S03) The only naturally occurring iodate is sodium iodate, which is found in Chili saltpetre in small quantities, and forms the HALOGEN OXIDES AND OXY-ACIDS 187 starting point for the manufacture of the greater portion of the world's iodine. Periodic Acid. In all cases where one would expect to obtain the monobasic acid, HIO4, the compound HI04,2H20 or H5IO6, separates out. Periodic acid may be obtained by the action of — (a) Sulphuric acid upon barium periodate ; (b) Iodine upon perchloric acid : 2HCIO4 + 2B.fi + la-^ 2HI04,2H20 + 01^. (c) Chlorine upon an aqueous solution of sodium iodate and sodium hydroxide. In this case, the sodium hypochlorite is the active agent in raising the oxidation of the iodate to the higher stage. The sUghtly soluble salt which separates out, is NajHjIOe, an acid salt derived from HjIOj. The periodic acid of the composition m04,2H20, is a colour- less, deUquescent compound which, on heating, breaks down into oxygen, water and iodine pentoxide. Although it has not been possible to prepare the monobasic acid, HIO4, salts of this acid are well known (KIO4, etc.). Nomenclature of the Periodic Acids. — Both acids, HIOi and HjIOe, are derived from the same anhydride I2O, (not yet isolated) : IA+H20-^2HI04 IA+5H20-^2H5lOe The difference between the properties of the salts derived from these acids is comparatively slight, and of a far less funda- mental nature than exists between the salts derived from such acids as iodic and periodic. This arises from the fact that the two periodic acids differ merely in degree of hydration, and not in the degree of oxidation. In such cases it has become the custom to apply the term ortho-acid to the acid of the highest state of hydration known, while the term meta-acid is reserved for the acid with the lowest degree of hydration. Under such a system of nomenclature, HsIOo becomes ortho-periodic acid, HIO4 meta-periodic acid. (A more systematic scheme, however, 188 AN INORGANIC CHEMISTRY would be to reserve the term ortho-acid for the acid of the highest possible degree of hydration, viz. : lA +VH,0->2H,I0-„ PA + 5H,0->2H5P05. Unfortunately, as the term ortho-acid has been ear-marked for the highest hydrated acid known, and as this differs in many acids, e.g. periodic (HjIOc) and phosphoric (H3PO4), the term ortho-acid has acquired a somewhat loose significance.) Iodine Tetroxide, I2O4, has been prepared in a state of more or less purity by the action of hot sulphuric acid upon iodic acid. Yellow crystals of the composition indicated by the above formula are obtained. On heating the oxide decomposes into iodine and oxygen. Under the action of water or of dilute sulphuric acid auto- oxidation occurs. 5I2O4 + 4H20-^ 8HIO3 + 12. Alkalies yield a mixture of iodate and iodide. The Valence of the Halogens The constitution of the oxy-acids of the halogens is stiU a matter of doubt. Some consider that these acids possess the chain linkmg, e.g. H— 0— G, H— 0—0— 0—0— CI. In the carbon compounds, where chain linking is so common, it has been found that the longer the chain, the less stable is the compound. In the case under discussion perchloric acid is the most stable of the oxy-acids of chlorine, and yet on the chain hypothesis it must presumably possess the longest chain. The chief point in favour of this constitution is that it preserves the mono-valency of chlorine and the other halogens, but the discovery of such compounds as ICl3,IF5, has shattered the view, so tenaciously held, of the constancy of the valence of an element. At present the most logical view of the constitution of these acids is based upon the multi-valence of the halogens, e.g. chlorine, possessing any valence from one up to seven. On this assump- tion the structural formulae of the oxy-acids of chlorine would be: H— CI H— 0— CI H— 0— C1=0 H— 0— C1=0 =^0 HALOGEN OXIDES AND OXY-ACIDS 189 .0 H— 0— Cl^O ^O, whilst the heptoxide would have the constitution represented by the graphic formula : II II 0=C1— 0— C1=0 II II o General. — In concluding the study of the halogens, the following points must be noted : 1. With increasing atomic weight, the physical properties of the elements show a steady change ; 2. With increasing atomic weight, the stabUity of the hydrides diminishes ; 3. With increasing atomic weight, the reducing power of the hydrides increases ; 4. With increasing atomic weight, the oxidising power of the halogens diminishes ; 6. With increasing atomic weight, the stability of correspond- ing oxy-compounds increases, i.e. the iodates are more stable than the chlorates ; 6. The stabUity of the oxy-acids (and of their salts) of each halogen increases with the oxygen content. The last statement is of very general importance. With few reservations, it may be stated that the higher the oxide formed from any element, the more acidic in nature is the oxide ; when more than one acidic oxide is formed from one element, the stronger and more stable acid will be derived from the oxide containing the greater amount of oxygen. Questions 1. Describe the preparation of an aqueous solution of hypochlorous acid. What is the effect of adding sodium hydroxide to such a solution ? How would you distinguish an aqueous solution of hypochlorous acid from one of chlorine ? 2. Describe the modern method of manufacturing bleaching powder. Account for the bleaching action of this substance. 190 AN INORGANIC CHEMISTRY 3. How do you account for the greater stability of perchloric acid, as compared with hypochlorous acid ? 4. Compare and contrast the reactions — 2HI + a2 ^2HCl + l2 2HC103 + l2^ 2HI03+Clj. 5. Describe the following reactions from the point of view of oxidation and reduction ; construct equations. (as) An aqueous solution of ferrous chloride and chlorme. (6) Potassium- chromate and hydrobromic acid. (c) Chlorine and an aqueous solution of sulphur dioxide. (d) Hydrogen iodide and iodic acid. 6. What is auto-oxidation ? Illustrate your answer by reference to the action of heat upon sodium hypochlorite. 7. What is the action between chlorine monoxide and a strong solution of hydrochloric acid. 8. How do you account for the lack of oxy-compounds of fluorine ? 9. Devise a suitable commercial method for converting iodine into potassium iodate. CHAPTER XIII OZONE, HYDROGEN PEROXIDE Ozone Historical. — In 1785, Van Marum drew attention to the fact that, if electrical sparks are passed through oxygen, the gas acquires a peculiar smell, but little more was discovered till Schonbein (1839-1880) took up the subject, and discovered various methods by which this substance, which he named ozone, might be prepared. It was not till 1860 that Andrews and Tait conclusively proved that ozone contained no hydrogen and must consist of a form of matter identical with oxygen. Preparation. — Ozone is often formed in small quantities during electrical reactions. During the catalytic oxidation of ether by means of a hot platinum wire, ozone is generated, so, too, in many processes of slow oxidation, e.g. during the slow oxidation of phosphorus. A mixture containing 14 per cent, of ozone in oxygen is liberated by the action of fluorine upon water. At high temperatures oxygen forms small quantities of ozone, which can be detected by suddenly chilling the gases to a low temperature. This can be effectively shown by blowing oxygen past the hot pencil of a Nernst lamp. The usual method of preparing ozone is to expose pure, dry oxygen to the influence of the silent electric discharge. The instrument in general use for this purpose is known as a Siemens' Ozone tube (Fig. 69). There are two concentric tubes. The inner tube is coated on its inner surface with tinfoil, connected to one terminal of an induction coil. The outer surface of the outer tube, which is similarly coated, is connected to the other terminal of the coil. A slow stream of oxygen is led through the annular space between the tubes during the passage of an 191 192 AN INORGANIC CHEMISTEY electric discharge. The gas which issues from the exit may contain up to 7-8 per cent, of ozone. If the mixtiure from the ozonizer consisting of oxygen and ozone is passed through a spiral immersed in Uquid oxygen, a deep blue solution of ozone in oxygen is obtained, and by cautious evaporation a liquid fairly free from oxygen can be isolated. ien D^^ Fig. 69. Properties. — Ozone is moderately soluble in water, 100 volumes of water dissolving about one volume of ozone at 0°, one-half a volume at 12°. It has a very strong and rather unpleasant smell, and, if inhaled too freely, causes headache. Ozone is a most powerful oxidising agent ; it attacks and destroys organic matter (rubber, etc.). Litmus and indigo sulphate are both decolorised by it. Metals, even mercury and silver, are oxidised by it. The action of ozone upon these elements is often quoted as a distinctive test for ozone. A drop of mercury, when shaken with ozonised' air, loses its lustre and its perfect liquidity, forming " tails " upon the glass. If a piece of hot silver foil is plunged into ozonised air, it becomes tarnished, no doubt owing to the formation of a film of silver peroxide. Lead sulphide is converted into lead sulphate by ozone. PbS +403->PbSOi +40^. Another reaction in which ozone behaves as a strong oxidising agent is the following : 2KI + H2O + 03-^ 2K0H + 12 + O2. OZONE 193 Many other oxidising substances, however, also Hberate iodine from potassium iodide. Ozone is readily converted into oxygen by the catalytic action of platinum black, manganese dioxide, etc., as well as by the action of heat. If the ozonised oxj'gen from the oxoniser is passed through a piece of glass tube heated to about 300°, no trace of ozone is found to escape — the decomposition is complete. The formation of ozone is attended by the absorption of a large amount of heat : + 02-^03 — 32,400 Gals. Composition of Ozone. — In determining the volume rela- tions of ozone and oxygen, advantage is taken of the fact that To induction coil- Fig. 70. Fio. 71. ozone is completely absorbed by certain organic oils, viz. tur- pentine, oil of cinnamon. The original experiments of Andrews and Tait on the forma- tion of ozone threw the first light upon the composition of this substance. They took two tubes, to each of which is attached a small manometer to register changes of pressure occurring within (Fig. 70). The tubes contained platinum points, between which a silent discharge could be sent. One tube was filled with pure, o 194 AN INORGANIC CHEMISTRY dry oxygen, the other with air as a blank experiment to control \^ariations in temperature and pressure occurring outside the tabe. The tubes were placed in a water thermostat and sparks passed. The manometer on the tube containing oxygen soon showed a contraction. On heating this tube above 270°, the original volume of gas was again obtained. If, after sparking, a small bulb containing potassium iodide was broken in the vessel, the gas no longer recovered its original volume on heating to 270°. The investigators concluded that ozone contains nothing but oxygen in a condensed form. Newth's apparatus (Fig. 71) enables the quantitative com- position to be investigated. The fundamental principle of the apparatus is the same as in the Siemens' ozonizer except that there is a manometer attached to the inner tube to record pressure variations. The apparatus is placed in a cyUnder containing water at 0°. The armular space between the two tubes is filled with dry oxygen by lead- ing it through the taps E to D. E is then turned off and the manometer brought into connection with the annular space by means of the three-way tap D. The coU is set in action and the contraction noted. The inner stopper is then twisted, and two glass projections C attached to it, crush a bulb containing turpentine. It is then found that twice the original contraction occurs. The first contraction occurs \\hen the ozone is formed, the second when the ozone is completely absorbed by the tur- pentine. Hence the inference that three volumes of oxygen form two volumes of ozone. 302-^203. This formula of ozone has been confirmed by density measure- ments of ozonised air. Direct density measurements, as well as density determinations with Bunsen's diffusiometer, show that the density of ozone is one and a half times that of oxygen. The work of Andrews and Tait established beyond all doubt that the molecule of ozone is built up of atoms of oxygen, whilst subsequent investigation upon the volume relationships when oxygen is converted into ozone has enabled the formula O3 to be assigned to ozone. In ozone and oxygen we have to do with two distinct substances, each endowed with its own specific properties, yet each buUt from atoms of the same kind. Such substances are known as AUotropes. Whenever elements occur HYDROGEN PEROXIDE 195 in more than one form, but in the same physical state, they are known as allotropic modifications of each other. Gaseous and liquid oxygen are therefore not allotropes, gaseous oxygen and gaseous ozone are. The chemical energy contained in the molecular quantity of two allotropes is always different, e.g. 1 gram of a diamond liberates 7840 calories, 1 gram of charcoal sets free 8040 calories. Hydrogen Peroxide Preparation. — Hydrogen peroxide is formed in small quan- tities when a jet of burning hydrogen plays upon the surface of ice-cold water, as well as by the action of moist oxygen upon many metals. Both zinc and copper, when shaken with oxygen and dilute sulphuric acid, produce small quantities of hydrogen peroxide. The metallic peroxides form the usual source of this compound. Sodium peroxide, treated with cold, dilute hydro- chloric acid, reacts thus : Na^Oa + 2HCl-> 2NaCl + H^Oa. This gives an aqueous solution of hydrogen peroxide, from which the pure substance is not easily obtainable. This solution, however, serves to illustrate the properties of hydrogen per- oxide. More often, however, barium peroxide is used. Ba02+2HCl->BaCl2-|-H202. The barium chloride is removed from the solution by the addition of silver sulphate in equivalent quantities : BaClj + AgaSOi -> BaS04 i + 2AgCl i . Barium peroxide, when slowly added to a cold dilute solution of sulphuric acid, reacts in accordance with the equation : BaOj-f H^SOi-^BaSOii +H20i,. The solution is then filtered, and concentrated by distillation under reduced pressure. Another method is exemplified by the equation : BaOa + 00, + H2O -> BaCOj >!' -^ H^Oa. Pure hydrogen peroxide decomposes violently at 100°. Even in solution it evolves oxygen freely at this temperature. Dilute solutions may, however, be concentrated by evaporation, provided the temperature does not rise above 70°. Properties. — Hydrogen peroxide is a colourless, syrupy liquid (sp. gr. 1-5) with a bitter, metallic taste. When brought 196 AN INORGANIC CHEMISTRY into contact with the skin, it produces a painful blister. Hydro- gen peroxide is miscible with water in aU proportions. The pure substance is unstable and decomposes at temperatures considerably below 0°. In aqueous solutions it is much more stable provided bases and salts are absent. A trace of acid increases the stabihty of the solution, and in this condition it may be preserved for a considerable time. If a solution of hydrogen peroxide is heated, oxygen is evolved : 2H,0,^2H,0+02. This action is facUitated, i.e. catalysed, even at low temperatures, by the presence of platinum black, charcoal, powdered metals, and finely divided manganese dioxide. The action of the powder appears to be at least partly of a physical nature, for a strong solution may be kept in a pohshed platinum dish at 60° without showing any signs of decomposition, but if the surface is scratched, bubbles of oxygen form along the scratch and escape. Hydrogen peroxide appears to react in a three-fold way : 1. As a weak acid. 2. As a strong oxidising agent. 3. As a reducing agent. Its acidic nature is shown by its reaction with a base. On adding hydrogen peroxide to a solution of barium or strontium hydroxide the hydrated metallic peroxide separates out : Sr(0H)2 + H2O2 + 6H2O -> Sr02,8H20. Another reaction of the same type is : Na^COs + HaO^-^Na^O^ -f CO2 -f H^O. Hence it is sometimes called peroxidic acid, and the peroxides, its salts, are known as peroxidates. As an oxidising agent, its reactions are numerous. It liberates iodine from hydrogen iodide or from an acidified solution of potassium iodide : 2HI + H,02->2H30+l2. It oxidises sulphides into sulphates : PbS +4H20,->PbS04 +4H2O, a reaction of some sUght importance. White lead, used in paintmgs, often reacts with traces of hydrogen sulphide to the detriment of the painting. If the brownish black lead sulphide HYDROGEN PEROXIDE 197 is then treated with hydrogen peroxide, the sulphide is thereby- oxidised to the white sulphate. Hydrogen peroxide has no action upon silver or upon mercury (cf. ozone). Solutions of hydrogen peroxide are used for bleaching, i.e. oxidising the colouring matter present in silk, hair, ivory, feathers, etc. Dilute solu- tions (3 per cent.) are largely used for antiseptic purposes. Two very delicate tests for hydrogen peroxide are known, both of them based upon the strong oxidising properties of this compound. A few drops of a solution containing a salt of titanium, when treated with hydrogen peroxide, produces a bright orange yellow coloration. This reaction is supposed to be due to the formation of per-titanic acid : TiOs+HaOa^H^TiOi. The test is extremely sensitive and is used as a test for titanium. By preparing solutions containing known quantities of titanium and by using a colorimeter (an instrument for comparing accurately the colours of two solutions) it is possible to estimate one part of titanium in 180,000 of water. Another characteristic test for hydrogen peroxide is the blue colour formed when it is added to an acidified solution of a chromate. An azure blue solution which rapidly decomposes is produced. If the blue solution is shaken with ether, the upper ethereal layer is found to have extracted practically the whole of the colouring matter, and has itself taken on a very beautiful azure tint. It is stated that one part of hydrogen peroxide in 100,000 of water can be detected by this means. The blue compound is much more stable in the ethereal layer, but even here its instability is so great that its composition has not been determined. Many consider that perchromio acid or a loose compound of perchromic acid and hydrogen peroxide is formed. The action of hydrogen peroxide upon silver oxide is generally classed as one in which hydrogen peroxide functions as a reducing agent. Ag20+H202->2Ag+H20+02. In this reaction both the peroxide and the oxide suffer decom position and appear to be reduced. In all cases of oxidation and reduction hitherto met with, oxidation of one substance has taken place at the expense of another which is reduced. Perhaps in this case the more logical point of view would be to postulate that the hydrogen peroxide oxidises the silver 198 AN INORGANIC CHEMISTRY oxide to an unstable peroxide, which at once breaks down thus : Ag,0,->2Ag + 0,. There is nothing strained about such an hypothesis, for silver peroxide has actually been isolated, and the view that such a decomposition should yield the lowest stage of oxidation — the metal — and not the oxide, AgjO, is in conformity with what generally occurs. As an example of this, it may be stated that if the blue perchromic acid, described above, is heated, one does not obtain the next compound, the chromate, but reduction to a yet lower stage of oxidation occurs {see Per- chromic acid). Another reaction of a similar type is that with ozone : Hfi, + 0,-^-^,0 +20^ Composition. — Owing to the instabihty of pure hydrogen peroxide, it has not been possible to determine the vapoiu" density. However, analysis shows that it is composed of hydrogen and oxygen in the ratio of atom to atom, i.e. the formula must be (HO)^. The molecular weight of hydrogen peroxide in aqueous solution {see Chapter xxvii.) is found to be 32, i.e. n must have a value 2 and the formula be HjOj. Peroxides and their Structure. — Two formulae have been suggested for the constitution of hydrogen peroxide : ^\ H— 0— 0-H and ^O = H^ The first of these formulae has been generally adopted as repre- senting the constitution of this compound. True peroxides are held to possess the linking present in hydrogen peroxide, —0—0—. The breaking down of sodium peroxide, barium peroxide, etc., with hydrochloric acid, would thus lead to the formation of hydrogen peroxide : Na— :0 Na-|0 + Ba/ HCl ->2NaCl + H- -0 1 H;C1 H- -0 HCl ->BaCl2 + H- -0 Hla H- -0 HYDROGEN PEROXIDE 199 Several oxides, which also possess two oxygen atoms, behave somewhat differently when treated with acid. PbOa + 4HC1 -> PbCli + 2H,0 -> PbCl^ + CI, + 2H2O MnOa +4HC1-H^ MnCli +2H20-^MnCl2 +CI2 +2H20 With sulphuric acid, these oxides liberate oxygen. It appears that in these oxides the metals are tetravalent, and the liberation of oxygen or chlorine is the result of a breaking down of the tetravalent compound into a compound of a lower state of oxidation (divalent). The formation of PbCli and MnCl4 is therefore most easily explained on the assumption of the following constitution for the oxide ; Jd .0 PbC MnC ^0 In these cases the valency of the metal in the higher oxide is greater than in the lower oxide. The true peroxide Unking is there- fore assigned to those compounds alone which form hydrogen peroxide with acid, whilst to those compounds which, on treatment with acid, show a degradation in their valence is attributed a constitution in keeping with the maximum valence that such a compound can exert. Questions 1. Discuss how the composition of ozone has been established. 2. Ten cubic centimetres of a solution of hydrogen peroxide are treated with colloidal platinum. The evolved oxygen, measured at N.T.P., occupies 15 c.o. Calculate the percentage of hydrogen peroxide present in the solution. 3. Discuss the constitution of the peroxides. 4. What is the action of hydrogen peroxide upon (a) sodium hydroxide, (5) sulphurous acid, and (c) hydriodic acid ? Write equations. 5. Describe any methods suitable for determining the strength of a solution of hydrogen peroxide. CHAPTER XIV CHEMICAL EQUILIBRIUM— THE LAW OF MASS ACTION— DISSOCIATION NuMEitOTJS examples have been given in which the chemical reaction does not continue to a completion, but reaches a state of equilibrium which may be approached from either side, i.e. the products of the reaction wiU react to give the same state of equilibrium as do the reactants themselves. Notable examples of these balanced or reversible reactions are the following : CI2 + H2O ^=± HCl + HCIO NaCl + H2SO1 ^=± NaHSOi + HCl I, ^=±21 2Ba02— ^2BaO + 02 3Fe + 4H,0 ^z=^ FcaO^ + iH^ We have seen how the equilibrium can be upset by the removal of one of the products of the reactions. The time has now come when the principles underlying these reactions should be clearly grasped. Equilibria considered from the Kinetic Hypothesis. — Suppose that hydrogen and iodiae are enclosed in a vessel. Before interaction can take place between a molecule of hydrogen and a molecule of iodine, the two molecules must meet in collision. The Kinetic Hypothesis assumes that the molecules are in a violent state of motion in straight lines, so that such collisions between molecules must be of frequent occurrence. The vigour of these collisions will at times be so great that chemical reaction will occur, and this will cause a steady decrease in the number of collisions between the molecules 200 CHEMICAL EQUILIBRIUM 201 of hydrogen and iodine, i.e. the velocity of the reaction H2+l2-^2HI must diminish as the reaction progresses. Let us now consider the back reaction 2HI — >-H2+l2. At the beginning of the experiment, owing to the absence of hydrogen iodide from the system, the velocity of the reaction 2HI— ^-Ha+Iji is zero, but as the hydrogen iodide begins to accumulate in the system, the chance of colUsion between molecules of this compound increases, and with it the possibility of chemical reaction. Hence the velocity of decomposition of the hydrogen iodide will steadUy increase as the concentration of the hydrogen iodide rises. We have seen that, as the con- centration of the hydrogen and iodine falls, there is a steady decrease in the velocity of the forward reaction, while there is a consequential increase in the speed of the back, reaction. At some point these two velocities must be equal, i.e. just as many molecules of hydrogen iodide will be formed per unit of time by the combination of hydrogen and iodine molecules as are decomposed per unit of time by the inverse reaction. On this point of view, chemical equilibria are essentially of a dynamic, not static nature (compare cases of phy.sical equilibria already discussed, evaporation, etc.). On the above considerations, the velocity of a chemical reaction depends upon the number of collisions between the reacting molecules per unit of time. The number of these collisions is influenced by two factors : (a) temperature ; (6) concentration. A rise in temperature causes an increase in the velocity with which the particles are moving, and so increases the possibiUty of collisions occurring. So far as con- centration is concerned, suppose that V denotes the velocity of the forward reaction between hydrogen and iodine when one gram molecule of each of these gases is present per litre. It is obvious that, if the concentration of the hydrogen is doubled, i.e. if the number of hydrogen molecules per unit of space is doubled, the velocity of reaction will be twice what it was before, i.e. 2V. And so, if the original concentration of the iodine is also doubled, the velocity of reaction will again be doubled ; in short, the velocity of a chemical reaction is proportional to the concentration of each of the reacting substances. This conclusion holds not only for gaseous reactions, but also for those in which liquids and solids are involved. This is not surprising, inasmuch as the Kinetic Hypothesis, originally 202 AN INORGANIC CHEMISTRY postulated for gases, has long since been extended to embrace phenomena in which liquids and solids are concerned. A very illuminating experiment has been devised in order to illustrate the influence of the concentration upon the velocity of a chemical reaction. It is based upon the following equations : HIO3 + SH^SOa-^ 3H2SO4 + HI 5HI + HI03-> SH^O + 3I2 Prepare a solution of iodic acid by dissolving 17-6 gm. of iodic acid in a litre of water and also an N/10 solution of sulphurous acid. Into fom: large beakers put 15, 10, 5 and 2-5 c.c. of the iodic acid solution, and dilute bj' the addition of 100 c.c. of water. Add also to each beaker 20 c.c. of a starch solution. Then introduce equivalent volumes of the solution of the sulphurous acid, and note the time before the blue colour flashes out in each beaker. The period which elapses before the appear- ance of the blue colour is found to be dependent upon the con- centration of the reacting substances. In the reaction between the substances A and B to form the compounds C and D A + B^=^C+D suppose that at the moment of equilibrium the concentration of the reacting compounds, expressed in gram molecules per htre, is represented by a, h, c and d respectively. The velocity of the forward reaction V — y, in accordance with the above considerations, is proportional to the concentration both of A and B, hence V — y=lci.ah. So, too, the velocity of the back reaction V-< — =k„.cd. At the moment of equUibrium V — >-=V-< — so that ki. ab=k,.cd or — ^ =j^ =K. K is generally referred to as the C(t ICi EquUibrium Constant of the reaction A+B^zr±C+D. In order to obtain a meaning for the proportionality factors k^ and k^, let the concentration of both A and B be unity. Then V— >=A;i, and V-< — =k2, that is, ki and k^ represent the velocity of the reactions when the participating substances have unit concentration, ki and A; 2 are known as the velocity con- stants of the reactions, occasionally as the affinity constants of the two reactions, though it is perhaps difficult for the student to appreciate why the term affinity should be applied to two opposing reactions. CHEMICAL EQUILIBRIUM 203 If the reaction is of a more complex type than that discussed above, e.g. 2A + 3B^z±2C+I>. or A + A+B + B + B^i=±C + C+D, V-^ =k^. aXaxbXbxb=kt a^b^, and similarly, a^b^ a" b'' Y -(r- = k^.c^d, whence K = — -^ ; in general, K= ' ' ' . C (t C (t . t , This is the most general expression of the so-caUed Law of Mass Action, for the discovery of which chemistry is indebted to Guldberg and Waage. The above deduction brings home the fact that K, the equili- brium constant, is the ratio of the velocity constants of the forward and the back reactions ; its value is readily obtained by measuring at the equihbrium point the values of all those substances participating in the reaction. In the reaction H2+l2^=::^2HI, it has been found that at a temperature of 448° approximately 80 per cent, of the mixture consists of hydrogen iodide molecules, 10 per cent. of hydrogen and 10 per cent, of iodine. This value is always obtained, whether pure hydrogen and iodine are heated until the equihbrium was reached, or whether hydrogen iodide is broken down into the equihbrium mixture. It follows that „ axa 0-1 X 0-1 1 ^2 . ^t, i- 1- X XI. K.=-— -=———— =—=--^, I.e. the reaction between the hydrogen molecules and the iodine molecules proceeds 64 times as fast as the back reaction between the molecules of hydrogen iodide. Effect on the Equilibrium of altering the Concentration. — If, in the above reaction H2+l2^=i±2HI, when equilibrium has set in at 448°, excess of hydrogen is introduced into the vessel, will a change in the equihbrium be induced or not ? Before the introduction of the excess of hydrogen, we had K=^~- ; let a-\-c denote the concentration immediately after the addition of hydrogen to the system, then, momentarily, we have K= — j~-. To restore this expression to its original 52 value, the numerator must decrease and the denominator increase, i.e. hydrogen and iodine must combine and so increase the concentration of the hydrogen iodide. 204 AN INORGANIC CHEMISTRY Whenever a system is in equilibrium, any increase in the concentration of one of the reacting substances, provided that temperature and volume remain imchanged, will cause such a change in the equilibrium as wiU tend to the removal of some of the added substance. This generalisation is well illustrated by the reaction : FeClj + 3NH4CNS ^=^ Fe(CNS)3 + SNH^Cl All the substances are practically colourless in dilute solution except ferric thiocyanate, which has an intense red colour. The experiment is best carried out by mixing equivalent quantities of ferric chloride and ammonium thiocyanate in a large volume of water, say 5 c.c. of a half normal solution of each of the reagents. The red solution obtained iii this way is roughly divided into four parts, one of which is reserved for the purposes of comparison. The addition of ferric chloride or of ammonium thiocyanate to different beakers causes a marked increase in the intensity of the coloration, whilst the addition of ammonium chloride almost decolorises the solution. The equilibrium existing in the solution can be moved at will by altering the concentration of the reacting substances. The addition of either of the compounds taking part in the forward reaction drives the reaction forward, whilst the addition of one of the products of the reaction reverses the reaction. In the above experiment the addition of ferric thiocyanate was not tried, as the strong red colour of this salt would mask any change produced in the equilibrium. The Effect of a Change of Volume (or Pressure) upon a System in Equilibrium. — If a system is in equilibrium, one can predict the effect of a change of volume upon the equihb- rium by the application of Le Chatelier's Law. Suppose equi- hbrium has been set up in the system Hj + Ij ^ — ^ 2HI. This equation states that one volume of hydrogen combines with one volume of iodine to form two volumes of hydrogen iodide, i.e. the reaction is not accompanied by a volume change. Le Chatelier's Principle (q.v.) states that an increase in pressure will favour the formation of the system with the smaller volume, whilst a decrease of pressure will induce such a change in the equilibrium as will be accompanied by a volume increase. It follows, therefore, that a change in the pressure (or volume) will exert CHEMICAL EQUILIBRIUM 205 no change on the equilibrium H2+l2;=:±2HI, nor, indeed, on any equilibrium wherein the chemical reaction involves no volume change. On the other hand, the effect of an increase of pressure (or decrease of volume) upon the system 1 vol. 1 vol. I vol. will be to force the equihbrium over to the right, as this change will be accompanied by a volume decrease. The system, by adjusting itself in this way, follows a law of Nature — it is Nature's attempt to avoid the increase in pressure which man is attempting to put upon the system. Other examples in which pressure will force the equihbrium to the right are : N, + 3H2;=±2NH3 2BaO + 02^zz±2Ba02 (The two soUds in the latter case occupy zero volume as com- pared with the gas). 200+02^^200^ In the following cases it is to be expected that an increase wiU force the equilibrium to the left. I ^ 9,T N20.^±2N02. Conclusions similar to the above are reached if the Law of Mass Action is apphed to the problem : H2 + l2^:±2HI a b c Let the concentrations of the tlu-ee gases at the equilibrium point be denoted by a, b, c, respectively. Then, by the Law of Mass Action, -^=K. Suppose that the pressure is increased n times, then the concentration will become nb, na, nc. Hence, after the change of pressure is effected, the equation for equili- brium becomes — — —- =K, from which one may conclude that no change in the value of a, b and c will occur as a result of the increased pressure — the equUibrium wiH not be disturbed, as already predicted by Le Chateher's Law. 206 AN INORGANIC CHEMISTRY If the system PCla+Cl^^iriPCls (whence K=-) is subjected a b c ^ ^' to an increase in pressure n fold, the new concentrations will become rva, vh and nc. The equilibrium equation will then , , . naXnb _ . , become at the moment of the compression =iv, i.e. tne nc values of a and b must diminish and that of c increase, in order that the equilibrium constant K may retain its value. Hence an increase in pressure will cause an increase in the concentration of the pentachloride, in agreement with our general deduction above. Conversely, a decrease in pressure upon the above system in equUibrium will lead to a dissociation of the penta- chloride with a consequential increase in the concentration of the trichloride and chlorine. The Effect of a Rise in Temperature upon Systems in Equilibrium. — Here, again, Le Chateher's Law affords valuable information. All chemical reactions are attended by the evolution or the absorption of heat. The Principle of Le ChateUer predicts that, it a system is in equUibrium and an attempt is made to increase the temperature, that change will be induced in the equilibrium which will negative the change which we are attempting to bring about in the system. In short, an increase in the temperature ivill induce that reaction to set in lohich is accompanied by an absorption of heat, whilst a decrease of tempera- ture will favour the reaction which is accompanied by an evolution of heat. (See the Law of van't Hoff, p. 107, which deals with the heat aspect of the General Principle of Le ChateUer.) Suppose that hydrogen iodide, hydrogen and iodine are in equUibrium at 448°. Seeing that the decomposition of the iodide is attended by an absorption of heat, it follows that, it the temperature is raised, further decomposition of the hydrogen iodide wiU occur. Eqtjilibkium Reactions in Chemistry UntU comparatively recent times, the existence of equihbrium reactions was not dreamt of by the chemist. At the present day the view is rather to consider aU reactions as equihbrium reactions. Even in the case of reactions which appear to run to a completion, it is held that the reaction is only apparently complete, and CHEMICAL EQUILIBRIUM 207 that in reality the equilibrium lies so far over to one side as to be practically complete. As examples of this statement, consider the neutraUsation of an acid by a base. Cu(OH) 2 + 2HC1-^ CuCla + 2H2O. This typical reaction was long looked upon as one that ran to a completion, i.e. the base was completely neutralised by the acid with the production of a salt and water. But it is now known that, if this salt is dissolved in a large volume of water, the reaction does reverse, and quantities of acid, readily and accurately measurable by physical means, are set free : CuCl2 + 2H20-^Cu(OH)2 + 2HCl. In other words, this hydrolysis must result in an equiUbrium. It is true that for certain specified bases and acids (e.g. sodium hydroxide and hydrochloric acid) no free acid can be detected in the solution, but in view of the whole multiplicity of acids and bases which do set up an equilibrium when neutralising each other, is it not more logical to conclude that, even in the few cases where such an equiUbrium has not been detected, such a reversible reaction does occur, and its non-observance is merely due to the imperfection of the instruments which the scientist has at his disposal ? The mixing of a solution of sodium nitrate and potassium chloride appears to result in no chemical change, but such a solution is identical with one obtained from mixing together solutions of sodium chloride and potassium nitrate. What other conclusion can be arrived at than that the equilibrium KNOa+NaCl^riKCl+NaNOa has been set up ? (For further discussion of this reaction, see p. 413.) Factors affecting the Equilibrium.— Our investigations concemmg the reaction NaCl+H2S04^zz±NaHS04+HCl led us to the conclusion that, if the hydrogen chloride cannot escape from the system, its concentration in the solution and its vapour pressure will steadily increase until the speed of the back reaction exactly equals the speed of the forward reaction, and an equili- brium between the four compounds is set up. In the terms of the Law of Mass Action, if the concentrations of the substances in the solution at the point of equilibrium are represented by cd a, b, c, d respectively, then K=-v-. If the concentration of d is CIO 208 AN INORGANIC CHEMISTRY lowered, e.g. by aUowing the hydrogen chloride to escape from the system, the equilibrium will be upset, and in order to restore K to its equUibrium value c must increase, and a and b decrease. In other words, the reaction wiU swing to the right in the endeavour to restore the equihbrium. Such displacements of the equihbrium are brought about both by physical and chemical means. The above method of displacing the equilibrium is essentially physical in nature. Another example of this type is afforded by the use of palladium or platinum vessels for reactions in which hydrogen is one of the participating substances. 2HI^r±H2 + l2 (1) 2H20^zz^2H, + 03 . . . . (2) If the compounds hydrogen iodide and water are heated to a definite temperatiu-e, eqiiilibrium will ultimately be set up ui accordance with the above equations. Denoting concentrations in the two reactions by ai, 6i, Ci, and a^, b^, Cj, we obtain . (1) . (2) But if the reactions are carried out in an apparatus provided with a wall permeable to hydrogen (i.e. of platinum or palladium), this gas wiU escape from the system until the pressure is the same within and without ; the concentration of the hydrogen within the apparatus will thus fall below the equUibrium value, and further dissociation into the elements must take place. The most important physical factors which affect equilibria are undoubtedly volatility and solubility. The reaction NaCl + HjSOi^ NaHSOi + HCl which illustrates sufficiently the manner in which an equilibrium is swung over through the escape of a volatile product of the reaction, has already been discussed (see p. 151). The effect of solubility in preventing the establishment of an equilibrium is well illustrated by the reverse reaction {see p. 151). NaHSOi + HCl -^ NaCl i + H2SO4. Another example of this type is the precipitation of barium sulphate by the addition of sulphuric acid to a solution of barium chloride. 61, d, and tta. K: CHEMICAL EQUILIBRIUM 209 BaClj + H,,S04-> BaSO, i + 2HC1, but before considering this reaction in detail, let us consider, in general, the action of sulphuric acid upon solutions of other soluble chlorides, e.g. copper chloride, sodium chloride. Experiment has shown with all certainty that, when dilute sulphuric acid is added to a dilute solution of a soluble chloride, a perfectly definite equilibrium is set up, e.g. CUCI2 + HjSOi^z^ CUSO4 + 2HC1 NaCl + H2SO4 ^:± NaHSOi + HCl provided the sulphate formed is also soluble. If the solutions are sufficiently dUute, no escape of hydrochloric acid occurs, and the equilibrium is clearly defined. In the first of the above equations, if a molecules of copper chloride are treated with 6 molecules of sulphuric acid, x molecules of copper sulphate and 2z of hydrochloric acid will be formed ; hence K=^^:i^L . . . (3) (a — x)(o~x) This equation tells us that, if the same quantity of any soluble chloride, the sulphate of which is also soluble, is treated with sulphuric acid under comparable conditions of concentration (and of temperature), the same amount x of sulphate must be produced : under such conditions, a, b and K in equation (3) have a constant value, whence x must also be constant. But if the solubility of the sulphate is such that x molecules of the sulphate cannot remain in solution, a deposition of the sulphate must occur. In the reaction, BaCL+H2S04->BaSO^^^ +2HC1 long before the equilibrium concentration of barium sulphate is reached, the amount of this substance present far exceeds the amount which the solvent is able to hold in solution, and precipitation ensues. The reaction therefore swings to the right, and the completeness of the precipitation will be measured by the insolubility of the barium sulphate. This method of looking upon the precipitation of such an insoluble salt may appear somewhat artificial, but he who doubts is advised to explain why, under comparable conditions of concentration and of temperature, sulphuric acid will always displace the same fractional amount of hydrochloric acid from a chloride, the sulphate of which is soluble. 210 AN INORGANIC CHEMISTRY Further light upon the question is shed by the reaction : BaSOi + NajCOa ^r^ BaCO, + Na^SOi abed Solid. Solid. Barium sulphate, if boUed with a solution of sodium carbonate, is partially decomposed with the formation of insoluble barium carbonate, whilst measurable quantities of sodium sulphate are found in solution. If the concentrations are represented by a, b, c, d, where a and c are necessarily very small, the Law of Mass Action states that, when equUibrium has been set up K=- ab' but, seeing that solid barium sulphate and carbonate are present, the concentration of these salts must have a constant value at the temperature of our experiment, and may therefore be included in the constant, i.e. we may write A;=- where k =K.-. c The reaction will cease when the ratio of sodium sulphate to sodium carbonate has reached a certain definite value. This result is in strict accord with experiment. An excellent example of a chemical method of influencing the equilibrium has already been studied. Take the reaction CI2 + H2O ^zi± HCl + HCIO. If an alkali is added to the solution, the acids formed are neutralised, and the reaction swings over to the right. The neutralisation of the two acids by the alkaU prevents the reverse reaction HCl + HCIO --> CL + H^O from setting in, so that the forward reaction Cla + H^O^ HCl + HCIO is enabled to progress far beyond the usual value. A somewhat similar example is afforded by the reaction H,+I,^:±2HI. Suppose a tube is filled with hydrogen iodide, and placed in an oven, maintained at a temperature of 448°, with one end of the tube projecting into a freezing mixture. Within the tube inside CHEMICAL EQUILIBRIUM 211 the oven, decomposition sets in until the usual equilibrium at this temperature (10 per cent, iodine, 10 per cent, hydrogen, 80 per cent, hydrogen iodide) has been set up. Iodine wiU at once diffuse into the projecting end of the tube, where it will deposit itself in the solid state. The equilibrium within the tube will be thereby disturbed and more of the hydrogen iodide must dissociate. The vapour pressure of the solid iodine in the cold portion of the tube is so much less than the vapour pressure of the iodine in the hot part of the tube that steady diffusion of the iodine from the place of high pressure to that of low pressure must continue. Dissociation of the hydrogen iodide will there- fore be promoted. By this simple physical device it is possible to ' decompose the hydrogen iodide almost completely into hydrogen and iodine at a temperature (448°) at which only 10 per cent, dissociation occurs normally. On the other hand, if a piece of potassium hydroxide is introduced into a tube filled with equal quantities of hydrogen and of iodine, and the tube placed within an oven maintained at 448°, so soon as hydrogen iodide is formed in accordance with the reaction H,+I,^2HL the hydrogen iodide will react with the alkali KOH + HI-^KI+H^O and the back reaction 2HI->H,+I, is thereby prevented. The result of this wiU be that hydrogen and iodine will continue to combine, the hydrogen iodide being removed as soon as it is formed. Ultimately, the whole of the hydrogen and iodine will have been brought into combination, though the hydrogen iodide js not immediately available. None the less, the equilibrium has been swung completely to the right by this " chemical method." Dissociation and Heterogeneous Equilibria. — In many chemical phenomena the system is heterogeneous. The usual types of heterogeneous reaction met with are soUd and liquid, soUd and gas. So far as the system — solid, liquid — is concerned, the application of the Law of Mass Action is comparatively 212 AN INORGANIC CHEMISTRY simple, for all solids have a specific solubility in every solvent. In the presence of an excess of solid, the solubility of the solid phase at constant temperature has a constant value (see preceding section). But when the system — solid, gas — is considered, a different method of treatment is in vogue. Every solid, like every liquid, has a fixed, though small, vapour pressure at every temperature. A piece of camphor or of iodine sublimes, a sheet of copper suspended in vacuo above a lump of sulphur becomes coated with a film of sulphide. So long as solid is present, there will be a constant vapour pressure of the solid compound, and interaction between sohd and gas will naturally take place through the medium of the gaseous phase. As an example, the reaction CaCOs ^rzfCaO+COa will serve. If the vapour pres- sures of the three compounds are respectively a, b, c, then since be concentrations are proportional to pressure K=— ; but a, b are a constants, since sohd calcium carbonate and oxide are present, so that the equation becomes kt= c. Hence, the pressure of the carbon dioxide has a fixed value at every temperature. This is known as the dissociation pressure of the calcium carbonate. So, too, in the reaction 2Ba02 :; — ^ 2"Ran-|-0^ we are led to the conclusion that at every temperature there will be a constant oxygen pressure exerted by the peroxide — a fact taken advan- tage of in the Erin's oxygen process. As a final example, the reaction 3Fe-f 4H2O ^=:±Fe304+4H2 will suffice. If steam and iron are heated in a sealed tube to a constant temperature, the c d* equihbrium defined by the equation K=^ — is set up. The two soUds, iron and iron oxide, have at the temperature of the experiment a fixed vapour pressure, i.e. c and a are constants and may be included in the constant, hence A-=pOrF=-^ Equilibrium will be set up when the pressures of hydrogen and steam bear to each other a fixed ratio. This is in strict accord with experiment. Questions 1. Explain what is meant by a " reversible reaction." 2. Show the relation which exists between the equilibrium constant and the velocity constants of the two opposing reactions. 3. How may the equilibrium be displaced ? CHEMICAL EQUILIBRIUM 213 4. A mixture of hydrogen and iodine, in the proportion of one volume of hydrogen to three of iodine, was treated at 445° 0. until equilibrium had been established. The equilibrium constant was found to be 02. Calculate the percentage of iodine converted into hydrogen iodide. 5. Iodine dissociates with rising temperature. If the system I2 ^ 21 be in equilibrium at 700° and the temperature be raised to 900°, would it be possible to restore the equilibrium to that which existed at the lower temperature, and if so, by what means ? 6. Discuss the reaction 3Fe + iB.!,0:=^¥esOi + 4H2 from the point of view of the Law of Mass Action. 7. The influence of temperature and pressure upon the reaction N, + 3H2^2NH3 ia shown in the following table : Temperature .... 700° 901° Ammonia (1 atm.) . . . 0-022 0-007 Ammonia (30 atm.) . . . 0-654 0-207 From these results deduce the most suitable commercial conditions for manufacturing ammonia synthetically. 8. State the Law of Mass Action, and illustrate it by reference to the action of water upon bismuth and antimony chlorides {q.v.). 9. Explain what is meant by the " velocity constant " oi a chemical reaction. CHAPTER XV SULPHUR: ITS HYDRIDES AND CHLORIDES Occurrence. — Sulphur occurs freely in the native state, especially in volcanic districts (Sicily, Italy, Iceland, Japan, New Zealand, United States, etc.). Large quantities of this element are also met with in the combined state. Such sulphides as zinc blende (ZnS), pyrites or iron sulphide (FeSj), galena or lead sulphide (PbS), copper pyrites (CuFeSa), cinnabar or merciuy sulphide (HgS) are of frequent occurrence, whilst such sulphates as barytes or heavy spar (BaSOj) and gypsum (CaS04,2H20) have a fairly wide distribution. Extraction. — Natural sulphur is ahvays more or less mixed with earthy matter, from which it is freed by the process of melting. The ore, containing up to 25 per cent, of sulphur, is stacked upon a sloping floor, and the heat produced by the combustion of part of the sulphur causes the remainder to melt. It is then allowed to drain into wooden moulds, and forms roU- sulphur. Further purification is effected by distilling the crude sulphur from earthenware pots. The vapour passes into a large brick chamber and condenses in the form of a fine powder, known as Flowers of Sulphur. The condensing chamber is not artificially cooled, so that its temperature soon rises sufficiently high to cause the sulphur to melt. This is drawn off and run into moulds. A small quantity of sulphur is also obtained by the decom- position of poly-sulphides (q.v.)as well as from the waste products of the alkali works (Chance's process, q.v.). During recent years large quantities of sulphur have been obtained from the sulphur beds of Louisiana. These beds have an average thickness of 150 feet and lie about 450-500 feet below the surface. The method of raising the sulphur to the surface (the Frascb process) represents one of the triumphs of the 214 SULPHUR 215 chemical engineer. A suitable arrangement of pipes down which superheated water is forced is sunk into the bed. The molten sulphur collects round the bottom of this pipe and is forced to the surface by means of compressed air. In order to prevent solidification setting in, the pipe carrying the sulphur in the form of emulsion (i.e. mixed with air) is jacketed with super- heated water. After the well is exhausted, sandy earth is introduced to prevent subsidence. As much as 73,000 tons is obtained from such a well. The average yield per year is about 250,000 tons, and the estimated amount of the deposit is 40,000,000 tons. Sulphur supplied by the Frasch process dominates the world's markets at the present time. The accompany- ing diagram shows a section through a Frasch sulphur pump. IS . S _. . , _ ^. Sulphur Wf^^ Sulphur Physical Properties. ™ ™ — Sulphur is a pale yel- low, brittle, crystalline sohd, insoluble in water but dissolving freely in carbon disulphide and sulphur dichloride. If a solution of sulphur in carbon disulphide is allowed t o evaporate slowly, well-defined crys- tals of rhombic form separate from the solution. The specific gravity of these crystals is 2-03-2-06 and their melting point 114-5°. Sulphur which occurs native belongs to the rhombic system, as also do the crystals present in roll sulphur. If a large evaporating basin is filled with molten sulphur and the contents allowed to cool slowly, a crystalUne crust forms. If this crust is broken through and the unsoHdified Uquid poured out, the lower surface of the crust is found to be covered with long, needle-Uke crystals. These crystals are nearly trans- parent, melt at 119°, have a specific gravity 1-96 and belong to the monoclinic system. So long as the temperature is main- tained above 96°, the crystals remain clear and transparent, but Hot water Sulphur Mot water Sulphur Sulphur vjv' Sulphur 216 AN INORGANIC CHEMISTRY below that temperature, they slowly lose their transparent appearance, owing to their becoming broken into a number of smaller crystals of the rhombic variety. This slow change is accelerated by contact with a crystal of rhombic sulphur. Con- versely, it a crystal of rhombic sulphur is heated above 96°, a slow conversion into monocUnic sulphur sets in. Contact with a crystal of monochnic sulphur will accelerate the change. The temperature 96° is known as the transition temperature of rhombic into monoclinic sulphur. Below 96° rhombic sul- phur is the stable phase, whilst above that temperature rhombic sulphur becomes unstable and tends to pass into the stable monoclinic form. The change S -^ S rhombic monoclinic is attended by the absorption of heat and by a volume increase. These changes are graphically shown in Fig. 73, where the vapour 1 ^ ^j> 1 1 _^ r ^ tt ^~ ~~ ■ -— ^ , ■^ i- y 'B ^ y' ^ A Temperature. 95-5° //4-° 120' Fig. 73. pressure curves of sulphur (rhombic, A-B ; monochnic, B-C ; and liquid, C-D) are plotted as a function of the temperature. In this diagram B denotes the transition point, C the melting point of monoclinic sulphur and E of the rhombic variety in a state of meta-stability. It is reached by the rapid heating of the rhombic form beyond the transition point. Other but less important crystalline forms of sulphur exist. This element affords a very interesting example of a polymorphic svhstance — an element or compound which crystaUises in more than two crystaUine forms. Substances which crystaUise in two forms are known as di-morphic. The behaviour of the molten sulphur is also of interest. At a SULPHUR 217 temperature immediately above the melting point the liquid is of a clear amber colour, but fairly mobile ; but as the tempera- ture rises, the mass becomes more and more viscous, until at 230° it is almost black, and so viscous that the vessel may be inverted without any of the liquid running out. At 260° the viscosity again diminishes appreciably and at 445° the liquid boils. If sulphur is boiled, and then allowed to cool slowly, the reverse series of changes is noticed, the end product being crystalline sulphur, soluble in carbon disulphide. But if sulphur, when heated above 350°, be poured into water, the substance solidifies to a tough, elastic material similar to india-rubber, known as plastic sulphur. In a few days this plastic sulphur becomes hard and brittle, and is found to consist largely of rhombic sulphur, which may, of course, be dissolved out with carbon disulphide. A residue is left, which entirely lacks crystalline form — the so-called amorphous sulphur. Amorphous bodies are really supercooled Hquids, existing at a temperature at which one would expect them to take on the stable crystalline condition. The coohng has been so rapid that the molecules have not had time to arrange themselves in that definite way required for crystalline structure, and nothing more than a general rigidity has set in. Chemical Properties. — Sulphur in the finely divided form unites slowly with oxygen, even at the room temperature ; at the temperature of 300° sulphur ignites in air. Sulphides are readily formed by the action of finely divided sulphur upon metals. The reactions are generally started by gentle heat, and soon the whole mass becomes incandescent, such is the vigom of the combination. A strip of copper foil biu'ns vigorously in sulphur vapour. The determination of the density of sulphur vapour has led to the conclusion that at relatively low temperatures (500°) the molecule of sulphur is a complex one, Sg (molecular weight = 256), but as the temperature rises, the density steadily falls. At 1,000° it has a density which points to its molecule being diatomic Sj (M.W.=64). There appears httle doubt that between the temperatures 500-1,000°, the variable molecular weight recorded is due to the effect of temperature upon the equilibrium Se-^4S,. 218 AN INORGANIC CHEMISTRY From 1000° to 1,700° the molecular weight remains constant at 64, though it is claimed that above this temperature a still further dissociation sets in. Commercially, sulphur is in considerable demand for the preparation of sulphur dioxide, which is required for bleaching and for the manufacture of sulphuric acid. Gunpowder, fire- works, matches, aU contain appreciable quantities of sulphur. Finely divided sulphur is freely used as a fungicide. Hydrides of Sulphuk There are three compounds of sulphiu- and hydrogen — hydro- gen sulphide (HjS), hydrogen disulphide (HaSj), and hydrogen trisidphide (H2S3). Of these the first is by far the most important. Hydrogen Sulphide. Hydrogen sulphide, occasionally stiU known as sulphuretted hydrogen, is found in some mineral waters and also occurs amongst the gases ejected by volcanoes. It is also formed by the putrefaction of animal and vegetable products containing sulphur, viz. eggs. Preparation. — This compound may be made in the following ways : 1. Synthetically, by the action of sulphur upon hydrogen at a temperature somewhat above 300°. At 310°, however, about seven days are required to secure the maximum yield of hydrogen sulphide. An increase in temperatm^e of 10° reduces the time necessary for reaction by about one half, so that in the neighbour- hood of 400° the reaction is completed in a few minutes. On the other hand, it has been found that the combination of hydrogen and sulphur is attended by an evolution of heat, so that the higher the temperature, the less hydrogen sulphide wiU be formed at the equihbrium point, i.e. a rise of temperature displaces the equilibrium Hj+Si^zr^HaS to the left (Le ChateUer's Law). Direct combination of hydrogen and sulphur as a method of preparing hydrogen sulphide is therefore not satisfactory, inas- much as those conditions which favour a rapid reaction give a relatively small yield of the gas. 2. Hydrogen sulphide may be displaced from its salts by the HYDRIDES OF SULPHUR 219 ^ To Condenser action of a less volatile acid. In general practice hydrochloric acid is allowed to act upon iron sulphide : FeS + 2HC1 -> FeCl^ + H^S. Sulphuric acid is less satisfactory than hydrochloric, as the ferrous sulphate formed in the reaction is very apt to choke the generator unless the concentration of the acid is kept fairly low. A type of generator which is much more economical for class work than the ordinary is shown in Fig. 74. The sulphide is placed in the part of the apparatus marked A and this is kept heated by means of a steam jacket or by the electric cur- rent. Strong hydrochloric acid is allowed to drop upon this column of sulphide and the gas is liberated with almost explosive violence. The acid which escapes into the receiver is completely spent. Iron sulphide gener- ally contains free iron, so that hydrogen sulphide, prepared from this reagent, is usually contaminated by the presence of small quantities of hydrogen. For ordinary testing purposes this is no detriment, but where a pure sample of the gas is required, it is usual to use a sulphide of a more definite composition, viz. antimony sulphide, barium sulphide, etc. Fig. 74. Sb2S3 + 6HCl->2SbCl3 3H2S. 3. Hydrogen sulphide is obtained by the extreme reduction of a sulphur compound. Sulphuric acid is thus reduced by the very active reducer, hydrogen iodide, to hydrogen sul- phide, though, if the sulphxxrio acid is strong, reduction to sulphur and sulphur dioxide only may occur. 220 AN INORGANIC CHEMISTRY SHI + H,S04 -> H^S + 4I2 + 4H2O (H,0.S03) Even sulphur itself is reduced by gaseous hydrogen iodide to hydrogen sulphide. 2HI+S-^H,S+I,. This action is of interest, as we have already seen that hydrogen sulphide will reduce an aqueous solution of iodine, forming hydriodic acid insolution (p. 165). These two reactions are in no way contradictory — they are essentially different reactions. The action of hydrogen iodide gas upon sulphur is attended by an evolution of heat, i.e. the energy content of the system after the reaction is less than that before the reaction. The interaction between hydrogen sulphide and an aqueous solution of iodine con- sists of two steps — ^the reaction 2H2S+I2 — >-2HI+2S, which in the hght of what has been said, must be an endothermic reac- tion, i.e. energy is absorbed during the reaction ; and secondly, the solution of the hydrogen iodide in the water to form an aqueous solution of hydriodic acid. The latter process is exothermal, so that the solution of hydriodic acid contains less energy than does the gaseous hydrogen iodide. The amount of heat involved in the second step considerably exceeds that absorbed in the first step, so that the net result of these coupled reactions is that there is a net decrease in the energy of the system. Physical Properties. — Hydrogen sulphide is a colourless gas having a rather sweetish taste and a characteristic offensive odour. It is a powerful poison in the pme state, and even if breathed in a diluted form for a long time, produces headache and nausea. Its poisonous properties may be computed from the statement that one part of hydrogen sulphide in a thousand is fatal to a bird, and one part in two hundred to mammals. At 10°C., 360 c.c. of hydrogen sulphide dissolve in 100 c.c. of water. The gas can be readily expelled from the solution by boiling. It should therefore be collected over hot water in order to lessen the loss through solubility. Hydrogen sulphide was first hquefied by Earaday by means of his bent tube (see p. 147). Hydrochloric acid and iron sul- phide were brought into reaction, and a colourless liquid soon collected in the cooled tube. At — 61° the hquid boils, while at — 83° it freezes. HYDRIDES OF SULPHUR 221 Chemical Properties. — The aqueous solution of hydrogen sulphide is weakly acid, so that this compound must be con- sidered as a feeble dibasic acid. As such, it gives rise to two types of salts — the normal type, KjS, and the acid salts (or hydrosulphides), NaHS. Acid salts of this type are readily obtained by the action of hydrogen sulphide upon solutions of the alkaUes. NaOH + H^S ^=:± NaHS + H^O Ca(0H)2 + 2H2S ^^ Ca(HS)2 + 2H2O. Insoluble sulphides are precipitated by the action of hydrogen sulphide upon aqueous solutions of the heavy metal salts ; CUSO4 + H^S ^ CuS >^ + H2SO4 soluble normal salts of the alkalies are obtained by the action of an alkaline hydroxide on an equivalent quantity of the acid salt. NaOH + NaHS ^=± Na^S + H^O Ca(HS2) + Ca{0H)2 — ^ Ca^S + 2H2O. On evaporation, the equilibrium is swung to the right owing to the escape of the water from the system, and the somewhat unstable sulphide separates out. A consideration of the equihbrium NaOH+NaHS^^NajS+HjO in the light of the Law of Mass Action would suggest that the addition of water to sodium sulphide' wiU lead to the hydrolysis of this substance, i.e. it wUl be partially decomposed by the action of the water. The precise amount of the various sodium salts in the solution win be determined by the amount of water added in accord- ance with the equation ( NaOH) (NaHS ) _ ' (H,0)(Na2S) the symbols in the brackets denoting concentrations. In order that the Law of Mass Action, as embodied in the above equation, should hold, any increase in the concentration of the water must cause a corresponding decrease in the concentration of the sodium sulphide and an increase in the concentration of the hydroxide and the acid sulphide. Liquid hydrogen sulphide forms a crystaUine hydrate with water, of the probable composition H2S, 6H2O, the behaviour of which closely resembles that of chlorine hydrate {q.v.). It exerts 222 AN INORGANIC CHEMISTRY a definite partial pressure both of hydrogen sulphide and of aqueous vapour, and, unless in a sealed system, the compound wiU slowly decompose. Hydrogen sulphide burns in air, forming steam and sulphur dioxide, 2H2S+3O2— >2H20+2S02. There seems httle doubt that in the interior of the flame a dissociation in accordance with the equation HjS— ^-Hj+S occurs, as the temperature hes far above 310°, the temperature where the dissociation becomes noticeable. This view receives confirmation by noting the result when the flame is allowed to impinge upon a cold dish — a deposit of sulphur is obtained. The Uses of Hydrogen Sulphide in Chemical Analysis. — The solubilities of the sulphides of the heavy metals show extreme differences, and upon this is based the use of hydrogen sulphide as a reagent in analytical chemistry. The sulphides may be divided roughly into three main classes : (a) Sulphides insoluble in dilute acids, e.g. copper, sUver, lead, tin, mercury, bismuth, cadmium, arsenic, antimony, gold, platinum. ( b) Sulphides soluble in dUute mineral acids, but insoluble in alkaline solutions, e.g. iron, manganese, cobalt, nickel, zinc. ( c) Sulphides hydrolysed by water, and therefore not precipi- tated by hydrogen sulphide whether in acid or in alkaHne solu- tion, e.g. chromium, aluminium, magnesium, barium, strontium, calcium, potassium, sodium. (Hydrogen sulphide in alkaHne solution precipitates the hydroxides of chromium and of aluminium but causes no precipitation in the case of the other metals.) In the case of group c a reaction such as is shown, CaClj + HjS ^=± CaS + 2HC1 ensues, but owing to the hydrolytic action of the water, no precipitation of the sulphide occurs. 2CaS + 2H2O ^=± Ca(0H)2 + Ca(HS)o. So far as groups a and b are concerned, in all cases double decom- position occurs with the precipitation of the insoluble sulphide. MeCla + H^S -^ MeS nI' + 2HC1 (Me denoting a divalent metal such as cadmium, zinc, etc.). HYDRroES OF SULPHUR 223 As the reaction progresses, hydrochloric acid accumulates in the system and in aU cases the back reaction MeS +2HCl-^MeCl2 +H2S must be reckoned with, hence, in general, one may write MeCU + HjS ^z:± MeS + 2HC1, whence, adopting the usual nomenclature, we obtain (MeCl2)(H,S) (MeS)(HCl)^ ="°"^*- The concentration of hydrogen sidphide is kept constant by bubbUng the gas through the solution, whilst the concentration of the metalUc sulphide is also constant in the presence of the solid which has separated, so that the equation may be written : (MeCl,) _ (HC1)2 And so one sees that the higher the concentration of hydrochloric acid in the solution, the greater must be the amount of metal salt retained in the solution unprecipitated in order that K may retain its constant value. Experience has shown that for most of the metals of Group IIA the value of K is so smaU that the amount of salt which remains in solution is neghgible, i.e. these sulphides are practically insoluble in hydrochloric acid, but this is not true for such metals as cadmium, lead and antimony. The sulphides of these metals are only precipitated by the action of hydrogen sulphide when the concentration of hydrochloric acid does not reach too high a value. In the case of the metals of Group IIIB, even if a neutral solution is taken, partial precipitation only is effected, as the steadily rising concentration of the acid soon brings the reaction to a standstiU e.g.,. ZnCl2+Hi;S^z±ZnS+2HCl. Complete precipitation may, however, be brought about if the acid formed is removed by the addition of an alkali, in practice ammonium hydroxide being used. A consideration of the above solubilities has led to the adop- tion of the scheme of qualitative analysis at present in vogue. Group II of the analytical tables consists of the metals whose sulphides are insoluble in the presence of dilute hydrochloric acid, and Group III of the metals, the sulphides and hydroxides of which are insoluble in the presence of ammonium hydroxide. Group IV and V of those metals the sulphides and hydroxides of which are soluble in the presence of ammonium hydroxide. 224 AN INORGANIC CHEMISTRY Composition of Hydrogen Sulphide.— If a known volume of pure hydrogen sulphide is decomposed by the passage of electric sparks, no change of volume occurs. It follows from Avogadro's Hjrpothesis that in every molecule of hydrogen sulphide there must be one molecule, i.e. two atoms of hydrogen. The formula is therefore H^Sn- The relative density of hydrogen sulphide (0=32) is 34-204. Hence : Molecular weight of hydrogen sulphide = 34-204 Weight of hydrogen in the gram molecule = 2-016 Weight of sulphur in the gram molecule = 32-188 The atomic weight of sulphur is 32-07, so that there can be but one atom of sulphur in the molecule of hydrogen sulphide. The formula is therefore HjS. Reducing Action of Hydrogen Sulphide. — Aqueous solu- tions of hydrogen sulphide undergo the same slow oxidation as do aqueous solutions of potassium iodide (p. 166). 2H,S+02->2H20+2S This reducing action of hydrogen sulphide is frequently met with. It is supposed that much of the world's native sulphur has arisen from the interaction between hydrogen sulphide and sulphur dioxide. 2H2S+SO„^2H20+3S (The student will note that the sulphur atom present in the molecule of hydrogen sulphide is here oxidised to elementary sul- phur, whilst the sulphur atom in the molecule of sulphur dioxide is reduced to sulphur.) Recently a commercial application of this reaction has been effected for the vulcanisation of rubber. The raw rubber is not sufficiently elastic, is too easily attacked by solvents, and too readily affected by temperature changes. The vulcanisation, which was formerly achieved by heating flowers of sidphur with the caoutchouc dough, removes these deleterious properties. In the modern process the caoutchouc is subjected to the action of the above gases, whereby coUoidal sulphur is precipitated throughout the mass. The great advantage of the process lies in the fact that it is more rapid than the old and can be carried out at ordinary temperatures. HYDRIDES OF SULPHUR 225 When hydrogen sulphide is led through concentrated sul- phuric acid, a similar reduction occurs, H^S + H^SOi^^ 2H,0 + S + SO2, {H,0,S03) the sulphur atom of the hydrogen sulphide being oxidised to sulphur and the sulphur trioxide group reduced to sulphur dioxide. As agent for drying hydrogen sulphide, sulphuric acid is consequently useless, and so is calcium chloride. Hydrogen sul- phide, if sealed up in a tube with calcium chloride, would speedily set up an equihbrium through the weak acid hydrogen sulphide replacing a scarcely measurable trace of hydrogen chloride, CaCla + H2S ;z::± CaS + 2HC1, but if the hydrogen sulphide is led over the calcium chloride, the hydrogen chloride formed during this replacement will be carried away, and the back reaction effectively prevented, so that a dry but impure hydrogen sulphide would result. The only suitable reagent for drying hydrogen sulphide is phosphoric anhydride (P2O5). This acidic oxide is not readily reduced like sulphuric acid, nor does it react with a gas of an acidic nature. Hydrogen sulphide reduces nitric acid energetically, and is itself oxidised either to sulphur or to sulphuric acid, according to the strength of the acid. H2S + 2HNO3 -^ 2H2O + S + 2NO2 (H,0,N,05) HjS + 8HNO3 -> 8NO2 + H2SO4 + 4B.fi (4H,0,N,05) Similarly, the chromates, manganates, and permanganates are readily reduced by the action of hydrogen sulphide : K fir JO, +'4:B.S0i + SHaS-^ Cr 2(804)3 + K^SO^ + VH^O + 3S. (KA2Cr03) (Cr203,3S03) In this reaction the hexavalent chromium is reduced to the trivalent stage. There are thus available of the left hand side of the equation 2x6=12 valences, whilst after the reduction the chromium valence towards oxygen isinaU2x3=6. A total of six valences are therefore available for the purpose of oxidising the hydrogen sulphide. The oxidation of hydrogen sulphide Q 226 AN INORGANIC CHEMISTRY requires two valences per molecule, H2S+O — >-H20+S. Hence the gram molecule of dichromate wiU oxidise - =3 molecules of hydrogen sulphide : 2KMn04 + 3H2SO4 + 5H2S ^-KjsSOi + 2MnS04 + SH^O + 5S. (K^CMn^O,) (2MnO,SO:,) In the reduction of the permanganate there are rendered available for oxidation purposes 14 — 4=10 valences, and as each molecule of hydrogen sulphide needs two valences for its oxidation to sulphur, it is evident that five molecules in all can be oxidised by the action of two molecules of permanganate. In both the above reactions the sulphuric acid may be replaced by hydrochloric acid, but in this case a very slight excess only of the acid should be used, otherwise interaction between the hydrochloric acid and the oxidising agent wUl occur with the liberation of chlorine. Another reducing action of hydrogen sulphide is the reaction with the halogens : H2S+Cl2->2HCl+S. Hydrogen Persulphide or Hydrogen Disulphide. When sulphur is shaken with a solution of the alkali sulphides, appreciable quantities of the sulphur pass into solution, and after evaporation a substance of somewhat variable composition is obtained. The composition appears to range from NajSa to NaaSj. If a solution of this persulphide is treated with dilute hydrochloric acid, a precipitate of milk of sulphur is thrown down, and hydrogen sulphide escapes : Na^Ss + 2HGl-> 2NaCl + H^S + 4S, but if the polysulphide is slowly poured into cold hydrochloric acid, the reaction takes a different course. No gas escapes and a heavy yellow oil falls to the bottom of the vessel. Careful fractional distillation of this substance has led to the isolation of hydrogen disulphide, H2S2, and hydrogen trisulphide, HjSj. Other sulphides such as H2S5 have been described, but their chemical identity is doubtful. The disulphide is an oUy liquid with a pungent smell. It is very unstable, decomposing at ordinary temperatures into sulphur and hydrogen sulphide. Hydrogen persulphide burns with a blue flame, and lik e its analogue, hydrogen peroxide, it possesses feeble bleaching properties. CHLORIDES OF SULPHUR 227 Sulphur dissolves readily in it. The trisulphide has very similar properties to the persulphide. Chloeidbs of Sulphur Three such compounds are known : — Sulphur chloride, S2CI2 ; sulphur diohloride, SCI 2 ; sulphur tetrachloride, SCI4. Sulphur chloride is obtained by passing dry chlorine over heated sulphur in a retort. The compound boils at 138° and is readHy separated by condensation. The Hquid fumes in moist air and reacts with water : 2S2CI2 + 3H2O ->4HC1 + H2SO3 + 3S. It has considerable commercial application in the vulcanising of rubber, owing to the ease with which it dissolves sulphur. The vapour density (0 =32) is 135, showing that the molecule is double. Sulphur dichloride is obtained by passing chlorine into sulphur chloride at the temperature of melting ice : S2Cl2+Cl2^2SCl2. The dichloride is a reddish liquid which parts readily with one atom of chlorine. Sulphur tetrachloride can be obtained by saturating sulphur dichloride with chlorine at — 22°. It dissociates readUy, forming chlorine and sulphur chloride. The compoiuid reacts violently with water forming sulphur dioxide and hydrogen chloride, SCl4+2H20->S02+4HCl. Sulphur also forms a bromide S2Br2 and a well defined hexa- fluoride SFg. This gas is chemically inert and is of considerable chemical importance in establishing the hexa-valency of the sulphur atom : H CI CI .CI Fv /F S< S< \S<( FAS^F ^H ^Cl CV ^Cl W ^F The above compounds furnish strong evidence in favour of variable valence. QtTESTIONS 1. Give an account of the commercial sources of sulphur. 2. Discuss the allotropic modifications of sulphur. 3. Discuss the reaction H2+S 1^ H^S fully. (Show the effect of temperature, pressure and other factors upon this equilibrium.) 228 AN INORGANIC CHEMISTRY 4. Give an account of the fundamental principles underlying the use of hydrogen sulphide as an analytical reagent. 5. Discuss the following reactions fully — HjS + SOj-^ H2S + HNO3 (dilute)-^ HjS + NaOH^- HjS + K2CrOi+H2S04 -^ 6. Describe how you would demonstrate the composition of hydrogen sulphide. 7. Classify the sulphides of the metals according to their behaviour towards (a) water, (6) sodium hydroxide, (c) hydrochloric acid. CHAPTER XVI THE OXIDES AND OXY-ACIDS OF SULPHUR The oxides and oxy-acids of sulphur are summarised in the following table : — Oxide. Acid. Sulphur sesquioxide, SjOj Hyposulphuroua acid, 1128204=1120, SjO Sulphur dioxide, SO2 Sulphurous acid, H2S03=H20,S02 Sulphur trioxide, SO3 Sulphuric acid, H2S04=H20,S03 Pyrosulphuric acid, HaS207=H20,2S03 Sulphur heptoxide S^O, Persulphuric acid HjSjOg = H20,S20, Thiosiiphurio acid, HjSjOs Polythionio acids Dithionic acid, HjSjOg Trithionic acid, 1X28308 Tetrathionic acid, HjS^Og Pentathionic acid, HjSsOy SuLPHUE Dioxide Preparation. — Sulphur dioxide is obtained by the combustion of sulphur in air or oxygen ; small quantities of sulphur trioxide are formed at the same time. Part of the sulphur dioxide utilised in chemical industry is obtained in this way, though perhaps even more is obtained by roasting iron pyrites in air : 4FeS2 + 1102->2Fe203 +8SO2. The most useful laboratory method of obtaining this gas is by the action of acids upon the sulphites or acid sulphites. Strong sulphuric acid is allowed to drop from a tap funnel into a flask containing a strong solution of the sulphite : NaHSOa + HsSOi-^ NaHS04 +80^+ H^O Na^SOa +H2S04-^Na2S04 + SO^ + H2O. Sulphur dioxide may also be obtained by the reduction of hot -«?9 230 AN INORGANIC CHEMISTRY strong sulphuric acid. As reducing agents the elements copper mercury, silver, carbon and sulphur may be employed : C + 2H2S04-> CO2 + 2SO2 + 2H2O. In the case of a metal, e.g. copper, acting upon sulphuric acid, it is highly probable that hydrogen is first liberated CU+H2SO4 — >CuS04+2H and this nascent hydrogen im- mediately reduces the hot concentrated acid : 2H+H2S04^.2H20+S02. At any rate, hot concentrated sulphuric acid is reduced to sulphur dioxide by the passage of molecular hydrogen through the liquid. Physical Properties . — S u 1 p h u r dioxide is a colourless gas posses- sing the well-known suffocating smeU associated with burning sulphur. It is more than twice as heavy as air and is, therefore, generally collected by the upward displacement of the air. The gas is an acute poison. Sulphur dioxide is easily Uquefied by passing it through a condenser cooled by ice and salt, and forms a transparent, colourless liquid, boiling at —8° and solidifying at — 76° to a white soUd. The gas is very soluble in water, from which it may be entirely expelled by boiling : I volume of water at 0° dissolves 79-8 vol. of sulphur dioxide 1 .- .. >. »20° „ 39-37 „ „ „ 1 => " .. ..40° „ 18-76 „ „ „ The extreme solubihty of this gas is strikingly sho-svn in the following experiment (Fig. 75) : The flask is filled with sulphur dioxide, the trough with water coloured blue by litmus. The tap is opened and the flask warmed sufficiently to expel any air lodged in the tube. As the flask cools water slowly rises, and the vacuum produced by the absorption of the gas in the water, soon results in the production of a fountain-like jet of reddish coloured water. Fig. 75 OXIDES AND OXY-ACIDS OF SULPHUR 231 Chemical Properties. — The solution of sulphur dioxide in water has an acid reaction. The acid itself has not been isolated, but forms a weU-defined series of salts — the sulphites. Conse- quently, sulphur dioxide is often referred to as sulphurous anhydride, the acid being known as sulphurous acid. Though sulphur dioxide is not inflammable, and not a sup- porter of combustion in the ordinary sense of the word, many metals such as finely divided iron react with it, producing the sulphide and oxide of the metal. Under the action of light sulphur dioxide undergoes decom- position, forming sulphur trioxide and sulphur (auto-oxidation) : 3SO2 ^=±2803 + 8. This photochemical effect is best illustrated by projecting a beam of light through a long tube filled with sulphur dioxide. The originally transparent gas soon becomes murky and foggy, clearly marking out the track of the beam of light (Fig. 76). Fia. 76. On removing the source of light, the equilibrium slowly shifts to the left, and the gas again becomes clear. Reducing Action of Sulphur Dioxide. — Sulphur dioxide has marked bleaching properties, due to its reducing action ; there is consequently a fundamental difference in the chemistry of bleaching by means of sulphur dioxide and chlorine. Some consider that sulphur dioxide reduces the water, and liberates hydrogen, and this hydrogen reduces the colouring matter : SO2 + 2H20-> H2SO4 -f 2H. This view is supported by the fact that articles bleached by sulphur dioxide have their colour readily restored by mild oxidis- ing agents. Moreover, sulphur dioxide will not bleach in the absence of traces of water. Flannel and sponge, bleached by means of sulphur dioxide and left moist, soon turn yellow again 232 AN INORGAJSriC CHEMISTRY on exposure to the air. Rose leaves are early bleached by sulphur dioxide, but on the addition of a small quantity of sulphuric acid the colour is restored. Considerable use is made of sulphur dioxide as a commercial bleaching agent for decolor- ising such articles as straw, sUk, wool and sponge. An aqueous solution of sulphur dioxide reduces chlorine, bromine and iodine : CI2 + 2H2O +S02->2HC1 + H2SO4, hence its use as an " anti-chlor " to remove traces of chlorine left in articles bleached by chlorine. The action of an aqueous solution of sulphiir dioxide upon iodine appears to be a con- tradiction to the statement already made (p. 165) that hydrogen iodide reduces sulphuric acid in accordance with the equation : 2HI+H,S04-^2H20+S02+l2, provided the reduction is effected in aqueous solution. As a matter of fact, we have to deal with an equiUbrium reaction : 2H20 + SO2 + 12 ^r± 2HI + H,_S04, and the position of the equilibrium can, as usual, be varied at wiU by altering the concentrations of the reacting substances. Advantage is taken of this reaction in estimating the amount of sulphur dioxide present in an aqueous solution. An iodine solution of known strength is allowed to run in so long as decoloris- ation occiu-s, but in order to prevent the back reaction from setting in, the titration must be carried out at such a dilution that the reverse reaction between hydriodic and sulphuric acids does not occur. In this way the equilibrium is driVen practically completely to the right. From a knowledge of the strength of the iodine solution used, one can calculate the equivalent quantity of sulphur dioxide by the application of the above equation. Other oxidising agents such as potassium dichromate, potassium permanganate, iodic acid and lead dioxide, are reduced by sulphur dioxide : K,GrA+3SO,+H,S04->Cr2(S04)3+K,S04 + H20 (K20,2Cr03) (Cr203,3S03) 2KMn04 + 5SO2 + 2H2O -^K^SO^ + 2MnS04 + 2H,S04 (K^CMnaO,) (MnCSOs) 2HIO3 + 5SO2 + 4H2O --> 5H2SO4 + 12 (H,0,I,05) (H,0,S03) PbOa + SO^-^PbSO^ OXIDES AND OXY-ACIDS OE SULPHUR 233 Such is the vigour of the reaction between sulphur dioxide and the lead dioxide that the mass becomes incandescent. Oxidising Properties of Sulphur Dioxide. — The oxidation of finely divided iron, magnesium, etc., by sulphur dioxide, with the formation of the metallic oxide and sulphide, yields an illustration of the oxidising action of sulphur dioxide. So, too, does the reaction 2H2S+SO2— >H20+3S. But perhaps an even better example is afforded by the reaction between sulphur dioxide and stannous chloride dissolved in. hydrochloric acid. At first a fine precipitate of sulphur separates out ; this is soon followed by a golden .yellow precipitate of stannic sulphide, i.e. the sulphur present in the dioxide has been reduced to sulphur and then to hydrogen sulphide in order that the tin may be oxidised to the stannic form SO 2 + 3SnCl 2 + 6HC1 -^ SSnCli + H 2S + 2H ^0 with subsequent reaction between some of the stannic chloride and the hydrogen sulphide : SnCli + 2H2S^ SnS2 i + 4HC1. Composition of Sulphur Dioxide. — This is determined by the combustion of sulphur in a measured volume of oxygen. No change of volume occurs, so that, by the appUcation of Avogadro's Hypothesis, one arrives at the conclusion that every molecule of sulphur dioxide contains a molecule of oxygen ; the formula is SnOa. From the density of sulphur dioxide (64) (0=32) it follows that : One gram molecule weighs 64 gm. „ oxygen „ 32 „ Weight of sulphur in the molecule 32 „ The atomic weight of sulphur being 32-07, it foUows that there can be but one atom of sulphur in the molecule of sulphur dioxide. The formula is therefore SO 2. Sulphurous Acid. There seems little doubt that in aqueous solutions of sulphur dioxide a weak acid exists — sulphurous acid : H20+S02^z^H2S03, 234 AN INORGANIC CHEMISTRY but the instability of this acid has prevented its isolation. It is probable that many of the reactions of sulphur dioxide already discussed are really the reactions of sulphurous acid rather than of the dioxide, e.g. an aqueous solution of sulphur dioxide is much more readily oxidised by oxygen than is sulphur dioxide itself. It appears highly probable that in this reaction the sulphurous acid is oxidised at a greater rate than the sulphur dioxide. The removal of sulphurous acid from the system by the action of the dissolved oxygen wiU, in accordance with the Law of Mass Action, lead to the formation of fresh sulphurous acid at the expense of the dissolved sulphur dioxide, and the oxidation continues to progress. An aqueous solution of sul- phurous acid undergoes auto-oxidation, with the formation of sulphuric acid and the precipitation of free sulphur : 3H2S03-^2H2S04+S^1' +H2O. The constitution of sulphurous acid is made clear by an investigation of the properties of the compound, thionyl chloride. Thionyl chloride, SOCI2, is prepared by the action of phos- phorous pentachloride on dry sulphur dioxide : CI CI CI CI 0=S=0+ /P<-C1 -^ 0=S<^ + 0=P<-C1 Cl'^ ^Cl CI ^Cl It is presumed that in this reaction there is no change in the tetravalency of the sulphur atom ; hence the inference that one of the oxygen atoms is replaced by two chlorine atoms, both of these being directly combined to the central sulphur atom. On treatment with water thionyl chloride forms sulphurous acid : o=s/ CI HOH OH + -^ 0=S< + 2HC1. ^Cl HOH ^OH Assuming that the monovalent hydroxyl radicle replaces the monovalent chlorine atom without any structural change in the grouping of the molecule, it follows that sulphurous acid has the symmetrical formula : OH 0=S<^ OH OXIDES AND OXY-ACIDS OF SULPHUR 235 and is therefore a dibasic acid. This view is supported by the failure to prepare from sulphurous acid any inorganic compound which suggests an unsymmetrical constitution for that body. The above constitutional formula for sulphurous acid enables the chemist to express in a concise way the behaviour of this substance in its many reactions, but the student is again warned against looking upon the above constitutional formula as repre- senting the actual orientation of the atoms within the molecule. It is a mere diagrammatical representation of the chemical behaviour of the molecule and of the inter-relations of the atoms. Sulphites. — Two classes of sulphites exist — the normal and the acid sulphites, e.g. NajSOg and NaHS03. The latter, in common with many other acid salts, lose an acid group on heating and pass into the normal salt : 2NaHS03->Na2S03 + H2S03->Na2S03-|-H20+S02. The sulphites, though more stable than the parent acid, are nevertheless readily decomposed by the action of heat. The sulphite undergoes auto-oxidation : 4Na2S03-^3Na2S04 +Na2S. (Na20,S02) (NaASOs) Owing to the instability of sulphurous acid, acids evolve sulphur dioxide from sulphites : Na^SOa + 2HCl-^2NaCl + H,S03^2NaCl + H^O + SO^. The instability of the sulphites is further emphasised by the ease with which they undergo oxidation both in the solid and in the dissolved state : 2Na2S03+02— >-2Na2SOi. The addition of a small quantity of sugar, glycerine, alcohol, etc., to a solution of a sulphite almost entirely prevents this oxidation. These substances act as negative catalysts or retarders of the reaction. On the other hand, a faint trace of copper sulphate increases the rate of oxidation enormously. The tendency of sulphur dioxide to pass into sulphur trioxide, and still more, the tendency of sulphurous acid and the sulphites to undergo oxidation and pass into sulphuric acid and the sul- phates respectively, illustrates the general rule already stressed in the case of the oxides and oxy-acids of the halogens. Lower oxides are, as a rule, of a weaker acidic nature than the higher oxides, whilst the acids formed from the lower oxides are not only weaker but also less stable than those derived from the 236 AN INORGANIC CHEMISTRY higher oxides. There is a general tendency for the lower acid (salt) either to add on oxygen from the atmosphere or, failing that, to undergo auto-oxidation in order to form a more stable system, e.g. : SKClO-^KClOa + 2KC1 4Na2S03->3Na2S04 +1^0,^. SXTLPHTTR TeIOXIDE Sulphur trioxide may be prepared by the foUowing methods : 1. By heating ferric sulphate, Pe2(S04)3— ^-FeaOg+SSOa. 2. By dehydrating sulphuric acid with phosphorus pentoxide (PA). 3. By the direct oxidation of sulphur dioxide. Of these methods the third far outweighs the others in importance, for it forms the basis of the modern method of manufacturing sulphuric acid. The direct oxidation 2S0..4-0.. ^ — ^ 2SO3 proceeds very slowly. The velocity of the oxidation is much accelerated by a rise in temperature, but another factor comes into play which sets an upper limit to the temperature at which it is economically possible to work. The direct oxidation of sulphur dioxide is attended by a considerable evolution of heat, hence a rise of temperature causes dissociation of sulphur trioxide. The equilibrium 2SO2+O2 ^=^ 2SO3, in accordance with Le ChateUer's Law, wiU be swung to the left by the rise in temperatm-e, i.e. the higher the temperature, the smaller the yield of sulphur trioxide obtained when equilibrium has set in. It is true that high temperatures enable us to reach the equUibrium point more quickly, but this advantage is more than counterbalanced by the decreased yield. It has been found that at 400° nearly 98 per cent, of the gases are in a state of combination when equHLbrium has been reached, but unless a suitable catalyst is used, the time taken to reach the equilibrium concentration of sulphur trioxide is too great for the method to be economically possible. Many catalysts have been tried, among them ferric oxide, copper oxide, pumice, but by far the most efficient in its action is platinum. Although a patent has long been held for the preparation of sulphur trioxide by the catalytic oxidation of sulphur dioxide (Phillips, 1831), it was not till 1901 that Knietsch succeeded in making the process a com- mercial success. The improvements based upon the results of Knietsch were twofold : OXIDES AND OXY-ACIDS OF SULPHUR 237 1. It was found necessary to scrub out of the reacting gases every trace of the compounds of arsenic and other impurities which are frequently present in the sulphur dioxide obtained by the oxidation of pyrites. Arsenic compounds, in particular, have a remarkable poisoning action upon finely divided platinum, in which form the platinum is generally used as a catalyst, so that its accelerating power rapidly disappears. If, however, pure gases are employed, the life of the catalyst is practically unlimited. 2. Owing to the great heat of the reaction when sulphur dioxide is oxidised, the temperature of the reaction tends to rise rapidly above 400° and the efficiency of the process would suffer if this were not prevented by adequate precautions. In the modern processes this is effected by admitting the cold gases in such a way as to avoid a local rise in tempera- ture. An analysis of the equation 2SO2+O2 ^=zi 2SO3 from the point of view of the Law of Mass Action shows that an excess of oxygen wUl cause an increase in the equiUbrium concentration of sulphur tri- oxide. In practice, therefore, three volumes of oxygen to two of sulphur dioxide are used. In the modern develop- ments of this process efforts were first made to remove traces of the injurious arsenic by means of coke filters, and the catalytic oxidation was begun in the presence of ferric oxide (Mannheim process). The oxida- tion was completed in the presence of platinised asbestos. In spite of this there was a steady decrease in the efficiency of the platinum. At the present day the difficulty has been avoided by the use of arsenic free sulphur, the catalyst used being platinised magnesium sulphate. Fig. 77 is a diagrammatic representation of a modern contact plant. Fig. 77. 238 AN INORGANIC CHEMISTRY Physical Properties. — Sulphur trioxide exists in two forms, a liquid and a solid. The Kquid, which boils at 46°, is exceedingly volatile at ordinary temperatures, and fumes strongly in air. On cooling it forms white crystals which melt at 14-8°. A crystal- line variety of the substance, somewhat like asbestos in appear- ance, is obtained by heating the liquid to a temperature of about 16° for some time. These crystals also fume, and at 50° volatilise freely. From molecular weight determinations it appears that the solid form is a polymer of the liquid, having a formula (SO3).. Chemical properties. — Sulphur trioxide combines with water with great energy to form sulphuric acid. On account of the great energy of the combination, sulphur trioxide is able to extract the elements of water from such substances as sugar, paper, skin, producing a charring effect. Sulphur trioxide also reacts with extreme vigour with many metalUc oxides, forming sulphates : BaO+S03-^BaS04. In this case the mass becomes incandescent, owing to the heat of the reaction. Sulphuric Acid The remark has been more than once made that the wealth and importance of a nation may be best estimated by noting the extent to which this most important chemical is manufactured and utilised. In nearly every chemical industry it finds its place — ^it is extensively used in the manufacture of explosives, nitric acid, hydrochloric acid, sodium carbonate, fertOisers, and in countless other industries such as electroplating, dyeing, bleaching, etc. Manufactuee of Sulphuric Acid There are two important methods of manufacturing sulphuric acid — the old " Chamber " process and the modern " Contact " process. Contact process. — ^The theory of the contact process has already been discussed under the heading " sulphur trioxide." The preparation of the trioxide by the catalytic action of OXIDES AND OXY-ACIDS OF SULPHUR 239 platinum is the first and most important step in the manufacture of sulphuric acid by this process. The white mist-like sulphur trioxide which escapes from the contact chamber is rapidly- dissolved in 97-98 per cent, sulphuric acid, and sufficient water is run in to keep the acid in the condensing tanks at this strength. The reason for using strong acid as the condensing liquor is that sulphur trioxide in this fog-like state reacts very slowly both with dilute acid and with water. The great advantage of this method lies in its cheapness and in the purity of the acid obtained. The Chamber Process. — In 1740 Ward prepared sulphuric acid commercially from sulphur, nitre and water, but it was not till 1793 that the process was made continuous, and a method evolved essentially similar to that at present in, vogue. In this process the slow oxidation of sulphur dioxide to sulphuric acid is accelerated by the catalytic activity of the oxides of nitrogen. In the presence of the oxides of nitrogen the oxidation of sulphur dioxide to the trioxide proceeds rapidly, the oxides of nitrogen suffering a reduction to lower oxides. These lower oxides of nitrogen are immediately reconverted into the higher oxides by the oxygen, and the cycle is ready to begin anew. Theoretically the catalyst should be. able to oxidise an unlimited amount of sulphur dioxide ; practically, however, appreciable losses occur through various side-reactions, and this steady loss of a valuable catalyst forms one of the main drawbacks of this method. The chemistry of the process can be effectively reproduced on a small scale in the apparatus figured on p. 240. Into the holder A is first delivered a quantity of oxygen through the drying bottle B, and sufficient nitric oxide is then admitted to form deep red fumes of nitrogen peroxide {q.v.). Simultane- ously, sulphur dioxide is admitted through the drying bottle G. A small quantity of moisture is introduced into the reaction flask by bubbling a httle oxygen through the flask D filled with boiling water. White crystals soon begin to form on the surface of the fiask A and rapidly spread over the whole surface. If the gases stUl present in A are now swept out by means of a stream of oxygen delivered through one of the drying bottles until the contents of the fiask appear quite colourless, and then a httle steam introduced into A from the boiling flask, the chamber 240 AN INORGANIC CHEMISTRY crystak at once liquefy and the gases in A once more become a deep brown. The reaction leading to the formation of the white crystals is represented by the equation : /OH 2S02+3N02+H20^2S02 +N0 \0.N0 The nitro- or nitrosyl- sulphuric acid reacts with water thus : /OH /OH 2SO2 +H2O -> 2SO2 +N0+N02. \0.N0 \0H The oxygen present immediately oxidises the nitric oxide (NO) Fig. 78. to the peroxide (NO 2), and the cycle is ready to begin again so soon as a fresh supply of sulphur dioxide is introduced. Although in practice nitrosyl-sulphuric acid is not allowed to separate out, there seems little doubt that the part played in the above apparatus by the nitrogen oxides is essentially the same as in the actual manufacturing process. Owing to the excess of water present, the actions involved are : OXIDES AND OXY-ACIDS OF SULPHUR 241 SO^+NO^ + H^O-^H^SO.+NO 2NO + 02-^2N02 The Chamber Acid Plant.- — In the plant the sulphur dioxide is made by burning sulphur or hydrogen sulphide or by roasting iron pyrites in a current of air. The oxides of nitrogen are derived from nitric acid vapour, obtained by the action of sulphuric acid upon sodium nitrate : NaNOs +H2SOi^NaHS04 + HNO3. The initial reaction between sulphur dioxide and the nitric acid is represented by the equation, 2HNO3 + 2SO2 +H20->2H2S04 +N0 +NOa and, after that, the oxides of nitrogen take up their role of catalyst. Small quantities of nitric acid vapour are continually introduced into the reaction chambers in order to make good any losses of the catalyst during the operation. This loss arises very largely from the reduction of the nitric oxide to nitrous oxide (NjO), which plays no further part in the reaction, and ultimately escapes from the system. The hot gases, consisting of air and sulphur dioxide, are made to pass up a tower — ^the Glover tower — packed with tiles, over which trickles a diluted acid made from the weak chamber acid and the riitrated acid from the Gay Lussac tower {q.v.). This weak acid with the oxides of nitrogen in solution is deprived of these oxides by the incoming hot gases, and at the same time the percolating acid is considerably concentrated during its passage down the tower. Before the acid reaches the foot of the Glover tower, it is completely denitrated, and is sufficiently strong to be used again in the Gay Lussac tower for the recovery of the oxides of nitrogen. The mixture of air, sulphur dioxide and the oxides of nitrogen is swept out of the Glover tower into a series of leaden chambers into which steam is blown. Here the reactions which lead to the formation of sulphuric acid actually take place. The diluted acid collects on the floor of the chamber, and is drawn off as required. This chamber acid has a strength from 60-70 per cent, and, although ready for use in a few industries such as the manufacture of superphosphates, it requires considerable concentration before it is available for industry in general. The concentration of the chamber acid is effected either in the 242 AN INORGANIC CHEMISTRY Glover tower or in leaden pans. In this way it can be concen- trated up to 78 per cent. (sp. gr. 1-7). Further concentration must be carried out in glass, quartz or platinum vessels. The gases issuing from the lead chambers contain the valuable catalyst — the oxides of nitrogen. These are recovered by making the issuing gases pass up the Oay Lussac tower. This is packed with coke or tiles over which concentrated sulphuric acid trickles. The nitrous fumes dissolve in the acid, forming nitrosyl-sulphuric acid, and this nitrated acid is then mixed with the more dilute chamber acid, and pumped to the top of the Glover tower for use in nitrating the incoming gases. The following diagram gives a representation of a chamber acid plant. Steam Fig. 79. Occasionally, the nitric acid necessary to compensate for the loss during the process is introduced into the top of the Glover tower, instead of being generated in retorts placed in the flues of the pyrite burners. Modern Improvements in the Chamber Process. — During recent years attempts have been made to dispense with the leaden chambers by causing the whole of the oxidation of sulphur dioxide into the trioxide to be effected in towers down which pours sulphuric acid containing dissolved oxides of nitro- gen. The general principle of such a process is that a series of Glover towers is constructed down which the nitrating acid OXIDES AND OXY-ACIDS OF SULPHUR 243 pours. During the passage through these towers the sulphur dioxide undergoes oxidation into sulphuric acid. This oxidation takes place very rapidly owing to the high concentration of the nitrogen oxides dissolved in the acid. The gases escape from the Glover towers into a series of Gay Lussac towers in which the oxides of nitrogen are again removed. By carrying out the oxidation in the liquid phase instead of in the gaseous, as in the Chamber process, a tremendous reduction in ground space and running costs can be effected, while from the purely chemical standpoint the method is a big advance on the old Chamber process. Provided the acid required is not stronger than 95 percent., this new process is probably the most efficient known. Impurities present in Commercial Sulphuric Acid.^ Commercial sulphuric acid, prepared by the Chamber process, contains numerous impurities, amongst the most important of which are carbonaceous matter, oxides of nitrogen, arsenic (from the pyrites), sulphur dioxide, and lead sulphate. The last men- tioned compound remains in solution so long as the acid is concentrated, but on dilution separates out. As a source of pure sulphuric acid, this method of manufacture has been practically superseded by the Contact process. Physical Properties. — ^A considerable evolution of heat occurs when sulphuric acid is mixed with water, and at the same time a very considerable contraction in volume is noticed. Owing to the great heat generated on mixing this acid with water, it is always advisable to pour the acid into the water with constant stirring. Pure sulphuric acid has a sp. gr. 1-85 at 15°. At temperatures above 150° the acid fumes freely, giving off sulphur trioxide. A mixture of 98 per cent, sulphuric acid and 2 per cent, of water distils unchanged at 320°. Acids of greater strength than this undergo partial decomposition until the composition faUs to 98 per cent., at which temperature the residue distils unchanged. Chemical Properties. — The chemical properties of pure sulphuric acid differ very appreciably from those of aqueous solutions. The pure acid is by no means stable, dissociating into sulphur trioxide and water at temperatures far below the boiling point. At 450° the dissociation is complete, as is shown by the determination of the vapour density. The anhydrous acid 244 AN INORGANIC CHEMISTRY (HaO.SOa), owing to its instability and its kigh oxygen content, is a powerful oxidising agent. This is exemplified by its beha- viour when treated with carbon, copper, hydrogen bromide, hydrogen iodide and hydrogen sulphide {q.v.). Dilute aqueous solutions of sulphuric acid do not exhibit oxidising power. When sulphuric acid is added to the salts of other acids, a double interchange occurs, NaCl + H2SO4 ^=± NaHS04 + HCl NaNOj + H2SO4 ^± NaHS04 + HNO3 and an eqxiillbrium is set up, provided that the products of the interchange are not allowed to escape. On the other hand, if the acid set free by the sulphuric acid is of a volatile nature, e.g. nitric or hydrochloric acid, and therefore capable of separation from the non-volatUe sulphuric acid, we have at hand a ready and cheap method for setting free such an acid from its salt. Sulphuric acid has a great affinity for water. For this reason it is frequently made use of as a drying agent. Several hydrates of sulphuric acid have been isolated, of which may be mentioned the monohydrate, HaSOijHaO (melting point 8°), and the tetrahydrate, H2S04,4H20 (melting point — 25°). This great affinity of sulphuric acid for water is shown by its power of abstracting the elements of water from many organic com- pounds. Cane sugar, if moistened with concentrated sulphuric acid, rapidly becomes brown, and finally froths up into a carbon- aceous mass. So, also, paper and wood are charred by the action of sulphuric acid. The use of sulphuric acid in the manufac- ture of nitro- glycerine and guncotton (q.v.) is based upon this dehydrating action of sulphuric acid. When added to a soluble salt of barium, sulphuric acid gives a heavy, white, crystaUine precipitate of barium sulphate, insol- uble in hydrochloric acid (distinction from sulphurous acid). This test is used largely both in quahtative and in quantitative analysis. BaCl^ -f H2S04^ BaS04 i + 2HC1. Being a dibasic acid, sulphuric acid gives rise to two series of salts — the normal sulphates and the acid or bi-sulphates. The acid sulphates are obtained by adding half the quantity of base required for complete neutralisation. H2S04+Na0H->NaHS04-fH20, with subsequent evaporation to expel the water ; also by the OXIDES AND OXY- ACIDS OP SULPHUR 245 action of sulphuric acid upon the salt of another acid, provided the temperature is kept low. NaCl + H2S04->NaHS04 + HCl. Normal sulphates are obtained : 1. By the complete neutralisation of the acid by a base. H2S04+2NaOH->Na2S04+2H20. Ca(OH), + H2SO4 ^CaS04 +2ILfi. 2. By the action of the acid sulphate upon a salt at a high temperature. NaNOg +NaHS04->Na2S04+HN03. 3. By the interaction of a metaUic oxide with sulphur tri- oxide. 4. By the oxidation of a sulphide or sulphite. 2Na2S03 + 02-^2Na2S04 CuS+202-^CuS04. The sulphates of the heavy metals (e.g. PeSOi) decompose on heating, forming metaUic oxides and setting free sulphur trioxide. The sulphates of the alkalies (e.g. sodium) and of the alkahne earth elements (e.g. calcium) are stable towards heat. Constitution of Sulphuric Acid. — The formation of sul- phuric acid from its anhydride, sulphur trioxide, suggests that the linking in the molecules of the trioxide and of sulphuric acid is the same. In sulphur trioxide there is no reason for assigning anything but a symmetrical formula to the molecule, and, arguing from the established hexavalency in sulphur hexa- fluoride, SPj, we assign the formula O to the molecule of sulphur trioxide. Sulphuric acid is derived from sulphur trioxide by the addition of a molecule of water. If we retain the hexavalency of the sulphur atom, this would lead to the formula OH x/ S J\ OH 246 AN INORGANIC CHEMISTRY Strong support to this constitution is furnished by the hydrolysis of sulphuryl chloride, SOaClj. This substance is formed by the direct combination in sunUght of sulphur dioxide and chlorine under the catalytic action of camphor. O .CI \S + CI, -> \s/ 0^ 0^ ^Cl The chlorine atoms may be separately replaced by the cautious addition of water. The first action is to form chloro- sulphuric acid, which acid may also be formed by the direct action of sulphur trioxide upon hydrogen chloride. .CI HOH O OH \S< + -> ^S< + HCl 0^ ^Cl 0^ M Chloro-sulphuric acid. Further hydrolysis leads to the formation of sulphuric acid. O CI HOH O OH \S< + -^ S>S< + 2HC1 0^ ^Cl HOH 0^ ^OH In this reaction the not unlikely assumption is made that the hydroxyl radicles take the place of the chlorine atoms which they displace. Pyeosulphueic or Disulphueic Acid When an acid sulphate is heated, water is ehminated and a pyro-sulphate is formed. 2NaHS0i — H^O-^NaaS A- On solution in water the pyrosulphates reform the acid sul- phates — hence it is clear that the pyrosulphates and the sulphates differ only in the degree of hydration. This is further brought out by a consideration of the properties of disulphuric acid. This is readily formed by passing sulphur trioxide into sulphuric acid, H2SO4+SO3— ^HjSaO?, or by the addition of sulphuric acid to a pyrosulphate. Various mixtures of disulphiu'ic acid and sulphuric acid are on the market. Sulphuric acid, containing 80 per cent, of disulphuric acid, is known as oleum, and is used extensively in chemical industry. If the disulphuric acid content falls below 10-20 per cent., it is known as fuming sulphuric acid. On OXIDES AND OXY-ACIDS OF SULPHUR 247 cooling either oleum or fuming sulphuric acid, crystals of disulphuric acid, H2S2O7, separate out. The behaviour and properties of disulphuric acid leave Kttle doubt that it is derived from the condensation of two molecules of sulphuric acid with the ehmination of one molecule of water. 0, ,0H SO2-OH ><- OH + H2O OH SO,-OH 0^ ^OH It is of the same stage of oxidation as sulphuric acid, and differs merely in the degree of hydration. StTLPHUE HePTOXIDE AND PeRSULPHUEIC AcID When a mixture of sulphur dioxide and oxygen is exposed to a silent electrical discharge, drops of an oily liquid appear. The analysis of the Uquid points to the composition as being S2O, — sulphur heptoxide. The oxide is somewhat unstable, tending to pass into the trioxide and oxygen. It dissolves in water with a hissing noise, forming a solution of persulphuric acid. S^O^+H^O^H^SA Persulphuric acid may also be made by the action of concen- trated sulphuric acid upon a well cooled solution of hydrogen peroxide. The electrolysis of concentrated sulphuric acid also leads to the production of persulphuric acid. The cell should be kept well cooled, and the anode should consist of a small platinum wire. The solution of persulphuric acid is very unstable, and slowly evolves oxygen. It is therefore an active oxidising agent. Persulphates. — The persulphates of the alkaUes are easily obtained by the electrolysis of the acid sulphates, e.g. KHSO4, if the anode is of small area and the cell is kept well cooled, white crystals of potassium persulphate KaSjOg, separating out. The small anode promotes the crowding together of the HSO4 groups which are discharged at that electrode. These HSO4 radicles combine at the moment of hberation to form mole- cules of persulphuric acid, which at once react with potassium 248 AN INORGANIC CHEMISTRY bifiulphate in solution and crystals of the somewhat insoluble potassium persulphate separate from the solution. The soUd persulphates are moderately stable, but decompose on heating into pyrosulphate and oxygen. 2K2SA^2K,SA+02. They oxidise iodides to iodine, iodine to iodic acid, and salts of manganese and cobalt to the higher oxide. Most of the per- sulphates, including barium, are soluble. They find commercial appUcation in photography. Sulphur Sesquioxidb. — Hyposulphubous Acid The direct union of sulphur and sulphur trioxide leads to the formation of sulphur sesquioxide. On solution of the sesqui- oxide hyposulphurous acid is not formed, but decomposition into sulphur, sulphuric acid and sulphurous acid occurs. Zinc hyposulphite is formed by the action of zinc on an aqueous solution of sulphur dioxide. Zn+2H2S03-^ZnS204+2H20. By using zinc and a solution of sodium bisulphite saturated with sulphur dioxide, the sodium salt is obtained. The acid itself is only known in aqueous solution. It absorbs oxygen very freely from the air and is a strong reducing agent — hence the apphcation of sodium hyposulphite in the indigo industry. Insoluble blue indigo is reduced by the hyposulphite to indigo white, which is soluble, and therefore able to find its way com- pletely through a fabric. On exposure to air, the indigo white is re-oxidised and insoluble blue indigo separates out throughout the fibre. Sodium hyposulphite in alkaline solution is also used for absorbing oxygen from gaseous mixtures. So strong is its reducing action that a solution of copper sulphate is reduced to copper hydride (cf. hypophosphite). There is still doubt as to the formula of hyposulphurous acid, but the consensus of opinion points to the formula HjSaOi. Thiosulphueic Acid A solution of sodium sulphite, on digestion with finely divided sulphur, forms sodium thiosulphate, NaaSOj-f-S— ^Na^SaOa. This reaction is very much akin to the oxidation of sodium sulphite to sulphate by oxygen, NaaSOs+O-^NajSO^ OXIDES AND OXY-ACIDS OF SULPHUR 249 — Whence the use of the term thio- or sulpho-sulphate. This resemblance in the behaviour of sulphur and oxygen has already been met with in the compounds HjO, HjS ; HjOj, H2S2. Sodium thiosulphate is also formed by the action of sulphur dioxide upon sodium sulphide. The reaction appears to take place in three steps in accordance with the equations Na^S +SO2 +H20-^Na2S03 +H2S 2H2S + S02->2H20+3S with subsequent interaction between the sulphur and the sul- phite formed in the first stage. The addition of an acid to a solution of sodium thiosulphate sets free thiosulphuric acid, but this unstable acid soon decom- poses, hberating sulphur dioxide and sulphur. Na^S^Os +2HC1-^H2S203 + 2NaCl The sulphiu" does not at once make its appearance, owing, probably, to a state of supersaturation having set in. Even weak carbonic acid of the atmosphere causes slight decomposi- tion of a solution of a thiosulphate. As the action of an acid upon a thiosulphate is reversible, the preUminary addition of a httle sulphite drives the reaction to the left, i.e. the decomposi- tion of the thiosulphate is prevented. Sodium thiosulphate, commonly known as hyposulphite of soda, or hypo, is used in photography for fixing negatives. Its action is to dissolve off the sensitised plate any unreacted silver hahde left there by the developer. Na^SaOs + AgCl-^NaCl +NaAgS203. Solutions of iodine are readily reduced by a thiosulphate, forming sodium tetra-thionate. 2Na2S203 + l2-^2NaI -fNa^S A. The reaction is frequently made use of in estimating the amount of iodine present in a solution. The above equation forms the basis of iodiimetry or titration with solutions of iodine. The reaction between chlorine and thiosulphate follows a different path from the above. 4CIj -f-NaaS^Os + 5H20->Na2S04 +H2SO4 + 8HC1. Owing to the large amount of chlorine which one gram molecule of sodium thiosulphate can remove, this salt is utilised freely in 250 AN INORGANIC CHEMISTRY the bleaching industry for removing traces of chlorine from bleached fabrics — hence the name antichlor. Thiosulphuric acid is a dibasic acid, and owing to its relation H— H— to sulphurous acid \s=0, the formula \S=S is H— O^ H— O^ generally assigned to it. PoLYTHioNic Acids Dithionic Acid is prepared by the action of sulphur dioxide upon manganese dioxide suspended in water. MnOa +2S02-^MnSA- The barium salt is formed by treatment with barium hydroxide, and to this is then added the equivalent quantity of sulphuric acid. BaSA +H2S0i->Hi,S A +BaS04 i ■ It is also made by the action of iodine upon sodium sulphite. 2Na2S03 + 21 -> Na^S A + 2NaI. Trithionic Acid. — The potassium salt is most readily obtained by the action of sulphur dioxide upon potassium thiosulphate. 2K2S2O3 +3S02-^S +2K2S30e, the acid being set free by the aid of hydrofluosUicic acid (HjSiFe) upon the potassium salt. Tetrathionic Acid. — The sodium salt is formed by the action of iodine upon sodium thiosulphate. The acid is then set free by the addition of dilute sulphuric acid to the sodium salt. Pentathionic Acid is obtained by passing hydrogen sulphide into a strong aqueous solution of sulphur dioxide. 5H 2SO3 + 5H 2S ^ H aSsOe + 5S + 9H 2O . All the polythionic acids are unstable in aqueous solution, though their salts appear fairly stable. Tetrathionic acid is the most stable of the series. Selenium and Tellurium Closely associated with sulphur is the element selenium. Native sulphur and pyrites both frequently contain selenium. OXYGEN FAMILY OF ELEMENTS 251 Two allotropic forms of the element are known, the red preci- pitated variety amorphous and soluble in carbon disulphide, and the lead grey, semi-metalhc variety obtained by coohng molten selenium. This is the form in which it is used in the selenium Ught-sensitive cells. Tellurium is generally found associated with gold or silver. Several allotropic modifications of tellurium are known. The chemical properties of the compounds of sulphur (A.W. 32), selenium (A.W. 79-2), and tellurium (A.W. 127-5) show the same gradation as has been stressed in the case of the halogen group of elements. All the elements form hydrides HaS, HaSe, HjTe, and as in the case of the halogens, these hydrides decrease in stabiUty as the atomic weight increases. They are all formed by corresponding methods. FeSe + 2HCl^ HaSe + FeCl^ FeTe + 2HC1-^ H^Te + FeCl^. All the elements form dioxides on burning — SO a, SeOa, TeOj. The dioxides aU dissolve in water, forming the -ous acids, H2SO3, HaSeOj, HaTeOj. These are all strong reducing agents. So unstable a compound is selenious acid that sulphurous acid is able to reduce it to metallic selenium. 2HaS03 +HaSe03->2HaS04 + HaO + Se. Selenium dioxide and tellurium dioxide are also obtained by oxidising the element with nitric acid. In all cases the -ous acid is capable of being oxidised to the -ic form, but selenious acid requires the most powerful oxidising agents at our disposal to oxidise it to the -ic stage. Selenic acid itself is a powerful oxidising agent (cf. H2SO4), and wiU dissolve even gold. H2Se04 + 2HCl-> HaSeOs +B.fi+ Cla. TeUuric acid is formed from teUurous acid by the action of chromic acid, and is readily dehydrated by heat, forming telluric oxide. The sulphates, selenates and teUurates are isomorphous. An aqueous solution of selenic acid is a very weak acid, whilst telluric acid is so weakly acidic that it does not affect htmus. The Oxygen Family of Elements The elements oxygen, sulphur, selenium and tellurium form a group of elements extremely analogous to the halogen group. 252 AN INORGANIC CHEMISTRY In this case, however, oxgyen stands somewhat apart. AU the elements form hydrides, the stability and physical properties of which show a similar change with increasing atomic weight, as do the elements themselves. TABLE 25 PROPEKTrES OF THE OXYGEN-StTLPHUE FaMILV Oxygen. Sulphur. Selenium. Tellurium. Atomic weight Boiling point . Sp. gr. (solid) ... Colour of solid 16 -183° 1-43 Pale blue 32 448° 2 Yellow 79 688° 4-3^-8 Reddish 127 1390° 5-9-62 Black Pbopebties op the Hydrides of the Group. Formula. Boiling point. Sp.gr. Dissociation temperature. H,S H,S H^Te 100° -61-8° —42° 0° 1 1-17 2-81 4-48 1800° 400° 150° 0° ^ AU the elements are divalent in the hydride, but form oxides in which they are either tetravalent or hexavalent. M O The following changes in the properties with increasing atomic weight of the element should be noted, and compared with the corresponding behaviour of the halogen family (p. 172). 1. The increase in the specific gravity of the element. 2. The change in the colour of the element. 3. The increase in the boiling point. 4. The decrease in the stability of the hydrides, 5. The decrease in the affinity for hydrogen. 6. The decreasing affinity for oxygen shown by the elements (cf. the comparative ease with which sulphur dioxide is oxidised with the difficulty with which teUtu^ium dioxide is oxidised). So far as the affinity for oxygen is concerned, the group behaves OXYGEN FAMILY OF ELEMENTS 253 in a manner directly opposite to what has been noted for the halogen group, but in all other cases a striking parallelism is noticeable. QtTESTIONS 1. Give a comparative account of the more important compovmds of selenium and tellurium and compare them with the corresponding com- pounds of sulphur. 2. Give an account of the fundamental principles of the Contact process for manufacturing sulphuric acid. 3. Give examples of reactions in which sulphur dioxide behaves (o) as a reducing agent, (6) as an oxidising agent. What reaction woiild you expect to occur when chlorine is bubbled through an aqueous solution of sulphur dioxide. 4. Show how you would carry out the conversion of sulphur into (a) sodium sulphite, (6) sodium bisulphite, (c) sodium thiosulphate. 6. Explain why sulphuric acid is frequently used for the commercial manufacture of hydrochloric and nitric acids. 6. Discuss the chemistry of the Chamber process. 7. What grounds has the chemist for attributing to sulphurous and sulphuric acids the constitution indicated in the formulae — /OH O^ yOH o=s< •sOH Of \0H respectively ? 8. Give an account of the preparation and properties of the persulj)hates. Why are these substances more vigorous oxidising agents than the sulphates ? 9. Write equations illustrating the oxidising action of sulphuric acid. 10. Compare the reaction between an aqueous solution of sodium thiosulphate and (a) chlorine, (6) bromine, (c) iodine. 1 1 . Give a complete interpretation of the equation — 2H2S + 3O2 -> 2H2O + 2SO2. 12. 2-9875 gm. of a certain substance contain -3000 gm. of carbon, •0250 gm. of hydrogen, and 2-6625 gm. of chlorine. It is also found that •25 gm. occupies 631 c.c. at 80° C. and 750 mm. pressure. Find the true molecular weight. (CI. =35-5, H=l, C = 12.) Give a concise but clear account of the various fundamental laws involved in this problem. CHAPTER XVII THE CLASSIFICATION OF THE ELEMENTS. THE PERIODIC LAW In dealing with the elements in the halogen group as well as in the oxygen-sulphiu- group, attention has been called, to the striking similarity in the chemical and physical properties dis- played by the elements in each of these groups. In nearly every case, as we pass from the element of lowest to the element of highest atomic weight in these groups, a perfectly regular change in each chemical and physical property is brought out. To a student who has inteUigently mastered the chemistry of chlorine the properties of bromine and iodine are almost self-evident. The question at once suggests itself to the thoughtful seeker after truth as to whether other elements can be grouped into families, and if so, is there any guiding principle which enables the chemist to sort out the elements into groups or families of related elements ? This is one of the questions which have exercised the minds of the chemists during the past century. 0. W. Holmes has stated : " One storey intellects, two storey intellects, three storey intellects with skyhghts. All fact collec- tors who have no aim beyond their facts are one storey men. Two storey men compare, reason, generalise, using the labours of the fact collectors as well as their own. Three storey men idealise, imagine, predict, their best illumination is from above, through the skylight." It is to the two-storey and three-storey intellects that chemistry is indebted for such great generahsa- tions as the Periodic Law. Many attempts at a classification of the elements have been made in the past, based sometimes upon valency, sometimes upon the basic or acidic nature of the element and so on, but the first great advance was made by J. W. Dobereiner (1829). This chemist was the first to draw attention to the fact that mogt 254 CLASSIFICATION OF THE ELEMENTS 255 of the chemically related elements exhibited a constant atomic weight (iron, nickel, cobalt), or showed a constant difference between neighbouring pairs of elements. CI 35 45 Ca 40 47 S 32 47 Br 80 47 Sr 87 50 Se 79 48 1 127 Ba 137 Te 127 Although Dobereiner's triads fixed the attention of chemists upon the possibUity that atomic weight and similarity in chemical properties were closely inter-related, the law, of which Dobe- reiner's Triads is but a necessary consequence, stiU eluded discovery. In 1863-1866 Newlands drew attention to a suggestive regularity when the elements are arranged in the order of their atomic weights. Every element was found to bear a surprising resemblance to the eighth element preceding or following it. Li Be B C N F 7 9 11 12 14 16 19 Na Mg Al Si P S CI 23 24 27 28 31 32 35 This pecuhar relationship was called by Newlands the Law of Octaves. Owing to faulty atomic weights at that time in use, numerous important exceptions to the generalisation of New- lands stood out, and, as a result, his law soon ceased to attract attention. Three years later, MendeleefE (1869) and Lothar Meyer quite independently put forward a method of classification which is substantially that in vogue at the present time. Mendeleeff pointed out that, if the elements were arranged in order of their atomic weights, a kind of periodicity in their properties became evident, i.e. the properties of the elements are a periodic function of their atomic weights . Such properties as valence, specific gravity, colour, atomic volume (volume in c.c. occupied by the gram atomic weight), compressibility, conductivity, etc., all proved to be periodic functions of the atomic weights of the elements. ■fiQ.!Aejg oij-ioad^ 256 ■H 00 ■dT £ •_ g 00 to OS b oS J3? p-H o "ab «!' M ?0 a »o o ^ t-- OS © 00 a • 10 osS oi ^ ■-H P' O fe2 to « « '- (M dS S O (O (M 10 o w • > Oo CO cc°; ^^9 Dob CO > o M OS j:^' 04 '<-i CO Oco O) Hoc ^ ^ fc" o St Si EH g _, ^ Ph ^ "? o N Hco i 1-H ^S C39 CO 00 «w3 CM § H * T' §^l ■* 00 H t; 3S (M -o^ e S4. N o2 1-^ «i o lr~ CO ■"* c8 CO g ce9 1^9 00 CO Jo 00 Kg CO 00 tiOoo : CO 00 CO m to GQ 03 .2 (D ^Oi S tj 1 i ' 1 '^ i CD ft 13 _^ 01 'o 13 Q t3 © t3 ,0 ft -ts ** 2 !> -0 > TJ J> ts ft H ■v> w H H y S h ^ 1 bo PLl 1 03 -3 T3 T) ts 1 M.2 econd long perio _ i«-2 S 6D-2 -*3 II a III ^11 0^ l-i* 1 ^ M f!^ 1 E-t 34 1 ^ 2S7 258 AN INORGANIC CHEMISTRY The table on i^. 257 is, perhaps, the most useful method of arranging the elements in order to bring out the periodicity of their chemical properties. Each short period in this table con- tains eight elements, and each long period eighteen elements, i.e. a short period of eight elements connected by a transition group of three elements with a short period of seven. The transition or bridge elements have atomic weights lying close together, and fit in between the strongly electropositive halogen family and the strongly electronegative alkaU family.* Many gaps are found to occur in the long series, though more than one gap has been fiUed since the table was first suggested. Thus, in 1869 three gaps existed in the third series — scandium, gaUium, and germajtiium. The vertical columns in the table are referred to as groups, and it is at once seen that aU elements which are closely related fall into the same group. Periodicity of Chemical Properties. — If the series Ne, Na, Mg, Al, Si, P, S, CI is taken as an example, it is found that both the valence towards oxygen and that towards hydrogen vary in a regular way. The gas hehum does not combine with oxygen, Na forms the oxide NajO, magnesium gives MgO, aluminium AI2O3, silicon SiOa (i.e. SiaOj), phosphorus PjOj, sulphur SO3, chlorine CljO,, i.e. the oxygen valence rises steadily from to 7 to drop suddenly back to zero as we pass from chlorine to argon in the next series. The valence towards hydrogen rises steadily from zero to a maximum of four for silicon, thence falling regularly to one for chlorine. The valance towards chlorine shows a somewhat simi- lar change. In general, we have : 8 MCI4 The hydroxyl valence can be examined from two points of view. * Note. — The modern method of referring to the alkalies as electro- negative, the halogens as electropositive, has been adopted. For reasons for departing from the old convention the student should consult Chapter xxx. 1 2 3 4 5 6 7 — MH MH2 MH3 MH4 MH3 MH2 MH — MCI MCI, MCI3 MCI4 MCI5 MCle — CLASSIFICATION OF THE ELEMENTS 259 012 3 4 56 78 — MOH M(0H)2 M(0H)3 M(0H)4 M(0H)5 M(OH)e M(OH),— M0(0H)3, M02(OH)2M03(OH) In the true ortho-hydroxide (for explanation of the term ortho- see Per-iodic Acids, p. 187), the number of hydroxyl groups is equal to the number of the group to which the element belongs. Practically, however, the tendency to hold hydroxyl in combina- tion steadily diminishes from the maximum value 4 for group 4. True ortho-phosphoric acid P(0II)5 breaks down, passing into the so-caUed or tho -phosphoric acid P0(0H)3, whilst the true ortho-sulphuric acid S(0H)6, passes into S02(0H)2. In the seventh group C1(0H)7 breaks down into C103(0H). The great regularity exhibited by the elements of the various groups in their valence relationships towards oxygen, hydrogen, chlorine and hydroxyl is a weapon of extreme value to the student in mastering the chemical formulae of the many compounds derived from the various elements. The student is warned against assuming that the compounds indicated above are the only ones which an element can form. Rather, they represent only the group compounds formed by the element ; copper, for example, falls into group one. Its chloride as such is GuCl, but it also forms the chloride CuCla with a corresponding series of salts of other acids. The Transitional or Bridge Elements. — The elements in the upper octaves of the long series show a much greater similarity than do the elements of the lower octave, i.e. potassium and rubidium are more closely related in chemical behaviour than are copper and silver, whilst a marked difference is noted in the properties of the corresponding elements in the upper and the lower octaves, e.g. between potassium and copper, calcium and zinc, etc. This is, however, less to be wondered at, when it is remembered that we are dealing with a series of eighteen elements. The properties in Series 3 change steadily along the upper octave to manganese, through the transitional elements, iron, cobalt and nickel, to the lower octave, and so on to iodine. Copper is thus seen to occupy a position midway between the strongly electro-negative metal, potassium, and the strongly electro- positive non-metal, bromine. The following table (for which we are indebted to Bayley, 1882), throws this fact into prominence. 260 CLASSIFICATION OF THE ELEMENTS 261 Not only does this method of tabulating the elements emphasise the gradual change in the basic property of the oxide in passing from the strongly electro-negative alkalies to the strongly electro-positive halogens, but another equally important point is brought out. The short period elements are seen to be con- nected by two lines to elements in the long periods, a heavy liae going from sodium to potassium, a light line to copper. The heavy line indicates that the element of the short period has a stronger kinship with the element so connected than with the element connected with the light line. The groups lying to the left prove to be the sub-groups A, while the sub-groups B are found to the right. Thus the elements, Uthium-sodium, bear a stronger relationship to the element potassium, etc. (sub- group lA) than to the elements of sub-group IB (Cu, etc.). On the other hand, nitrogen-phosphorus are akin to the elements of sub-group 5B (As, etc.), oxygen-sulphur with sub-group 6B (Se, etc.), fluorine-chlorine with sub-group IB (Br, etc.). Carbon- silicon, as is to be expected from its position in the middle of the table, occupies a strictly neutral position and may be classified equally well either with sub-group 44 or 45. The rare gases, hehum, argon, etc., which have no tendency to combine with any other element, are generally placed in Group 0, and act as a break between the strongly electro-positive halogen and the strongly electro-positive alkahes. If one considers the hydroxide series : 12 3 4 5 6 7 M(OH) M(0H)2 M(0H)3 M(0H)4 M0(0H)3 M02(OH)2, M03(0H) a marked change is noticed in the character of the hydroxide. All the hydroxides of group 1 are basic, whilst those in group 7 are equally strongly acidic. Oxides in the neighbourhood of group 4 are generally noted for the absence of any marked acid or basic behaviour (see Amphoteric Oxides, Chapter xxii.). This steady change in the properties of the hydroxide is also found in the oxides from which those hydroxides have been derived. We may therefore summarise by stating that as we pass from left to right in the table, there is a steady de- crease in the basic nature of the hydroxide (oxide) and a corresponding increase in the acidic properties of these compounds. So much for the variation of the basic and acidic behaviour for elements in the same horizontal series. 262 AN INORGAMC CHEMISTRY If attention is now fixed upon a group of elements, e.g. Group 1, it is found that the group readily subdivides into two : Li ; Sub-group A contains the alkalis, lithium, Na : sodium, potassimn, rubidium and caesium ; and K I into sub-group B fall the related elements copper, Cu silver and gold. All the groups are capable of a Rbj similar subdivision, though in some cases the Ag elements in the short series fall into sub-group Cs : B rather than into A, as in Group 6. Au A study of the properties of the elements in any sub-group, e.g. lithium, sodium, potassium, rubidium and caesium, shows that as the atomic weight of the element increases, a steady increase in the basic nature of the oxide (hydroxide) sets in. So far as the groups 4, 5, 6, 7 are concerned, the elements in these groups generally form acid oxides, so that it is perhaps more significant to state the generaUsation thus : with increasing atomic weight, the oxides of the elements in any sub-group show a steady increase in their basic properties or a steady decrease in their acidic properties. In speaking thus of the oxides and hydroxides of the elements it has hitherto been intended that the oxide common to the group is referred to, i.e. in Group 5 the oxides NjOj, P2O5, Aa^O^, SbaOj, BijOj show a decrease in their acidic properties. The above statement is, however, also true for any other group series of corresponding oxides, viz. N2O3, P2O3, AsjOg, SbgOj, BijOg, the first of which is distinctly acidic, the last basic. The following table is of use in summarising the preceding statements : TABLE 28 FT g Group 1 2 3 4 5 6 7 8 *< a S ! 3 Oxide- . e F Hydride-, p Chloride-. S" Hydroxide- s' 1. MjO MH MCI MOH Bas MjOj (MO) MH, MClj M(0H)2 c nature MjOj MH3 MCI3 M(0H)3 of oxide MjO. fMOj) MH4 MClj (0MH)4 andhyd M2O5 MH3 MCls M0(0H)3 roxlde stea< M^Oe (MO3) MH2 MCle M02(OH)2 dily diminis M2O, MH M03{0H) lies M,Os MCI4 M(on)4 Acidic nature of oxide and hydroxide steadily increases. Applications of the Periodic Law. — The Periodic Law has been of considerable value to the chemist in the following directions : CLASSIFICATION OF THE ELEMENTS 263 L It lias enabled him to estimate or correct the atomic weight of various elements, the atomic weight of which was not known with sufficient accuracy by the ordinary means. 2. It has led to the prediction of new elements. 3. It affords the only systematic classification of the elements. The gradation which has been found to occur when one passes from left to right along a series, as well as the gradation that occurs in a group of related elements, enables the chemist to correlate the properties of the elements in a manner entirely wanting until Mendeleeff's Periodic Law was put forward. The student has already seen how the chemistry of the halogens and of the sulphur group is simplified by studying as a whole the family to which these elements belong, and by noting the gradation in properties that occurs within a family or group as the atomic weight increases. Moreover, it affords us a means of predicting the properties of any one element from a knowledge of the properties of its immediate neighbours in the table, indeed it enables us to predict the general properties of a whole group of elements from a knowledge of its position in the table. The Estimation and Correction of Atomic Weights. — The two methods that are most frequently called into use for the determination of the atomic weight after the chemical equivalent has been found, are : (a) the determination of the molecular weight of a large number of volatile compounds, (6) the determination of the specific heat of the element. The first method is more frequently used in the case of the non-metals, the second for the metals. The evidence that has led up to the adoption of certain atomic weights has at times been very inconclusive owing to the practical difficulty of applying either of the above methods. The case of indium is just such a one. Its equivalent, as determined by Winkler, was 37-8. On the supposition that the element was divalent, the atomic weight was fixed at 37-8 x 2=75-6. The supposition of the divalence of the element was made upon the knowledge of the properties displayed by the compounds of the element known at that time. The value 75-6 brought the metal indium between the non-metals arsenic and selenium (As=75, Se=79) — an obvious misfit. Mendeleeff suggested that the oxide had the formula In ,03, in which case the metal would 264 AN INORGANIC CHEMISTRY fall into the aluminium group between cadmium and tin. The prediction was subsequently fulfilled by the determination of the specific heat of indium (0-057), whence the approximate 6-4 atomic weight= -— - — =112-3, the correct atomic weight being 0-957 37-8 X 3=113-4. Similar alterations were suggested by Mendeleeff for glucinum, uranium and some of the rare earths, and in all cases subsequent research has substantiated the prediction of Mendeleeff. The Prediction of New Elements. — When the table was originally drawn up by Mendeleeff, in order that elements should fall into groups to which their properties entitled them to belong, it was necessary to insert in the table several gaps which Mendeleeff predicted would soon be filled by the discovery of new elements. Three of these unknown elements he named eka-aluminium, eka-boron and eka-sUicon. From a knowledge of the physical and chemical properties of the elements in the immediate neighbourhood of these gaps, he was enabled to predict the properties of these stiU unknown elements. As an example, the properties of eka-aluminium (predicted in 1871) and galhum (discovered in 1875) are summarised. TABLE! 29 Eka-aluminidm. Gallium. Predicted 1871. Discovered 1875. Atomic weight, 69. Atomic weight, 69-9. Density, 5-9. Density, 5-93. Atomic volume, 1-17. Atomic volmne, 1-18. Low melting point. M.p., 301°. Not readily oxidised by air. Only superficially oxidised in air at a red heat. Soluble in acids and alkalies. Soluble in hot hydrochloric acid and in potassium hydi'oxide. Oxide, EI2O3. Oxide is GajOg. Chloride, El^Cle. Chloride, GaaClo or GaCl3 Would form potassium alum. Forms well-defined alimis. Exceptions to the Periodic Law. — Three cases occur where the recognised atomic weight of the element does not warrant the placing of the element in the position which the chemical properties of the element demand that the element should be placed. They are argon (39-88), nickel (58-68), tellurium (127-5). The atomic weight of argon, as at present known, exceeds that of potassium (39-10), but the placing of argon after potassium CLASSIFICATION OF THE ELEMENTS 265 in the table would involve potassium falling into the helium or zero group, and the unreactive argon under sodium. Nickel has an atomic weight somewhat less than that of cobalt (58-97), yet our knowledge of the chemical properties of iron, cobalt and nickel demands that cobalt should foUow iron, not nickel. In the historic case of tellurium, repeated determinations of the atomic weight of this element have failed to give an atomic weight less than that of iodine, though its great resemblance to selenium and sulphur demands its falling in the sixth group — before iodine. It is possible that finaUty in this matter has not yet been reached. Strong evidence of a physical nature points to the probabUity that the atomic weight of nickel is not yet known with accuracy, and that it must possess an atomic weight greater than cobalt. So far, however, the re-determination of the atomic weight of nickel has only confirmed the earlier values. The bearing of recent radioactive investigations upon this subject, as affording a possible explanation of the anomalous position of argon, nickel and tellurium wiU be discussed in the chapter dealing with radio-active change and the structure of the atom. The Position of Hydrogen. — On this subject the opinion of chemists is divided. Some favour its being placed at the head of the alkaUes, others at the head of the halogen famUy. On the one hand it can replace the alkali elements from their compounds, whilst in other compounds (organic) it can be replaced by the halogens. The evidence on neither side is sufficiently strong to decide the question. Questions 1. "The properties of the elements are a periodic function of their atomic weight." Discuss this statement. 2. Show how the periodicity of the chemical properties of the elements is an aid to their classification. 3. Show how the properties of the oxides of the elements may be inferred from the position of the element in the Periodic Table. 4. Discuss the position of the rare gases in the Periodic Classification. 5. Discuss the justification for including copper and silver in the same group as the alkalies. 6. Give a general account of the chemistry of arsenic from a knowledge of the properties of the elements surrounding it in the Periodic Table. CHAPTER XVIII NITROGEN— ATMOSPHERIC AIR— THE RARE GASES Nitrogen History. — Several scientists contributed towards the dis- covery of nitrogen in air, though it was left to Rutherford (1772) to isolate the new element. He removed oxygen from the air by burning phosphorus, charcoal, etc., the gases produced by the combustion being removed by the action of dilute potassium hydroxide. The gas obtained in this way was at first known as mephitic air, a name afterwards altered to azote by Lavoisier to denote its inabihty to support life. The name nitrogen, or nitre-producer, was suggested by Chaptal owing to the element being a constituent of nitre. Occurrence. — In the free state nitrogen constitutes about four-fifths of the earth's atmosphere. From spectroscopic observations it has been concluded that nitrogen also occurs free in certain nebulae. In combination it is found extensively in the form of nitrates in Peru, Chili and Bengal. Nitrogen is also an essential constituent of vegetable and animal matter. Albuminous compounds, for example, contain about 0-15 per cent, of nitrogen. Preparation.^ — Two methods of preparing nitrogen are available : 1. From atmospheric air. 2. From chemical compounds containing nitrogen. The preparation of nitrogen from atmospheric air is effected either by chemical or physical means. If liquid air (p. 273) is allowed to evaporate, the more volatile nitrogen escapes first. Large quantities of nitrogen are obtained commercially by this method for the ammonia and alhed industries. 266 NITROGEN 267 Nitrogen, containing small quantities of argon, can be obtained from air by causing the admixed oxygen to combine with some reducing agent, e.g. copper, phosphorus, etc. Owing to the fact that phosphorus does not bum in very dilute oxygen, the nitrogen obtained from burning phosphorus in air is never pure. Much purer nitrogen is obtained from passing air over copper heated to bright redness. 2Cu + 02->2CuO In this method air, carefully freed from carbon dioxide by passage through a solution of sodium hydroxide and dried by being p&ssed through sulphuric acid, is drawn through copper filings or turnings, maintained at a bright red heat. The disadvantage of this method lies in the fact that the copper is rapidly oxidised and requires renewal. In order to overcome this, a highly satisfactory alteration in the method was devised by Lupton (1876). Air is drawn through a concentrated solution of ammonium hydroxide, so that a mixture of air and ammonia enters the tube containing the hot copper fiUngs. The ammonia reduces the copper oxide in accordance with the equation : 3CuO + 2NH3-> 3Cu + 3H2O + Na, Air enters CdXE22CX Fio. 81. so that one obtains nitrogen from two sources, the air and the ammonia, while the copper is prevented from oxidising (Fig. 81). 268 AN INORGANIC CHEMISTRY Samples of nitrogen prepared from air are always contaminated by small quantities of argon, etc. If pure nitrogen is required, it is necessary to obtain it by breaking down compounds containing nitrogen. As such, ammonium nitrite is generally taken. On gently heating this substance, it breaks down thus — NH4N02-^N2 + 2H20. Owing to the instability of ammonium nitrite, it is usual to prepare this compound in the reaction vessel by mixing ammonium chloride and sodium nitrite NH4CI + NaNOa ^=± NH4NO2 + NaCl. By the usual interchange an equilibrium is set up, which is disturbed by the action of heat upon the unstable ammonium nitrite. Ammonium dichromate also evolves nitrogen on heating. (NH4),Cr20,->Cr,03 +4H2O +N2. Physical Properties. — Nitrogen is a colourless, odourless and tasteless gas which is slightly soluble in water (1-6 volumes in 100 of water at 20°). It has been condensed to a colourless liquid boiHng at — 195° and soUdifying to a white solid at — 214^. Chemical Properties. — Nitrogen is noted for its comparative inertness as an ele- ment. At ordinary temperatures it is practically iudifferent to all elements. Under the action of heat nitrogen combines directly with a number of elements to form nitrides, e.g. magne- sium, lithium, calcium, boron, titanium, car- bon. In these nitrides nitrogen is trivalent. 6Li+N2->2Li3N. Fig. 82. t" 2^- 3 This reaction can be easily shown by means of the apparatus shown in Fig. 82. A stream of nitrogen is passed through the apparatus to displace ATMOSPHERIC AIR 269 Fig. 83. the oxygen. On heating the combustion tube containing the calcium, magnesium, etc., expansion is at first caused, followed by a rapid rise of the liquid in the gauge as the reaction sets in. Active Nitrogen. During recent years it has been shown that an active form of nitrogen can be obtained by subjecting the gas to an electric discharge. In the apparatus devised by Strutt, the nitrogen was drawn in a slow stream into a tube in which two electrodes had been sealed. The nitrogen is activated by the discharge, and is then sucked out of the activating chamber into another tube in which its chemical properties can be investigated. His apparatus is figured below. Amongst the chemical ac- /C ^ tions attributed to active nitro- gen may be mentioned the Toreaction ^ . ^^^^^^^ '^'Nitrogen action of acety- lene upon it. If this gas was drawn into the activating tube containing the active nitrogen, undoubted evidence of the formation of hydrogen cyanide was obtained. The gases were sucked into a solution of sodium hydroxide, and after acidification, a mixttu-e of ferrous and ferric salts was added. A precipitate of Prussian blue was obtained {q.v.). Nitric oxide is converted in the presence of active nitrogen into the more highly oxygenated compound, nitrogen peroxide. Sulphur reacts with the new modification of nitrogen and gives a greenish deposit. Arsenic, sodium, lead, cadmium, and magnesium appear to form a nitride. Many other carbon compounds, such as methane and hexane, react and form hydrogen cyanide ; others, such as ether, produce cyanogen. The spectrum of active nitrogen is different from that of the usual variety. Traces of oxygen have been found to facilitate the production of the active modification: Atmosphbeic Aie Composition. — ^The work of Lavoisier upon the calcination of metals (p. 7), following as it did upon the qualitative investiga- 270 AN INORGANIC CIHEMISTRY tion of Boyle and others, clearly established that air was not an element, but consisted of at least two gases — oxygen and nitrogen. Since that day, the existence of carbon dioxide, ammonia, aqueous vapour, the inert gases (q.v.), ozone, com- pounds of sulphur (SO 2, HaS), ammonium nitrate and suspended matter (dust) has been estabhshed. Of these constituents it appears that oxygen, nitrogen and the inert gases alone are present in approximately constant quantity. All others are subject to considerable variation. If air is drawn from uncontaminated sources, e.g. the surface of the ocean and mountain tops, the proportion of oxygen present is fairly constant, although variations considerably beyond the experimental error undoubtedly occur. The gravimetric analysis shows that about 23 parts by weight consist of oxygen. The method of carrying out the observation is sho^vn in Fig. 84. Air Solution ^""S^O- " sulphuric acid. Fig. 84. Carefully dried air is allowed to stream over heated copper, and the residual nitrogen is collected in a receiver and weighed. This gives a direct measure of the weight of nitrogen and inert gases, whilst the weight of oxygen which was formerly associated with this nitrogen is obtained from the increase in the weight of the copper. In the volumetric analysis of air the oxygen percentage is determined by the reduction in volume when air is shaken with a suitable reducing agent — alkaUne pyrogaUol, sodium hyposulphite or phosphorus. By these means the oxygen is absorbed from a known volume of air. The residual volume of air is then measured and the ratio of oxygen to nitrogen + inert gases computed. Roughly, 21 per cent, of the air is found to consist of oxygen. S'or such a volumetric analysis a Hempel ATMOSPHERIC AIR 271 pipette (Fig. 85), containing the reagent, is frequently employed. In order to determine the actual amount of nitrogen present in the residue after the removal of the oxygen by means of the heated copper, the residual gas is allowed to stream slowly through a heated tube containing calcium or magnesium. Under these conditions the nitrogen is completely removed in the form of a nitride, SMg+Na-^MgaNj, and a small quantity of gas is left uncombined. This is argon and the other gases belonging to the same family. The proportion of argon present in air amounts to about 0-94 per cent, by volume. Carbon dioxide exists in pure air to- the extent of about 0-03 per cent, by volume, but the proportion of this constituent shows very considerable variation. In crowded rooms the per- centage of carbon dioxide may reach ten or twelve times the normal amount. During fogs the carbon dioxide content is abnor- mally high, reaching as much as '0-2 per cent. The cause of this fluctuation lies in the fact that this gas is exhaled by animals as well as by plants. Volcanoes and decomposing vegetable and animal matter also give rise to appre- ciable quantities of carbon dioxide. However, the balance of nature is roughly maintained by the action of growing plants which absorb carbon dioxide during the daylight and evolve oxygen. The easiest method to determine the percentage of this component present in the air is to cause a measured volume of air to bubble through a solution of barium hydroxide of known concentration. The amount of barium carbonate precipitated in accordance with the equation Ba(OH)2+C02-^BaC03 i-f-HjO is easily determinable either by titration or by direct weighing. Fig. 85. 272 AN INORGANIC CHEMISTRY The amount of aqueous vapour present in the air is likewise subject to very considerable fluctuation. The higher the temperature, the greater the amount of aqueous vapour which the air can carry before saturation is reached. Thus, at 0° 4-87 gm. of water vapour are necessary for the saturation of one cubic metre of air, whilst at 20° 17-16 gm. are required. On cooling air which has been saturated at a higher temperature, a precipitation of moisture will occur, provided the requisite nuclei are present upon which the particles of moisture may condense. The amount of water vapour present in air is obtained by allowing a known volume of air to pass through a weighed amount of calcium chloride, sulphuric acid, phosphoric anhydride or other suitable dehydrating agent. The bulk of the water vapour precipitated in the atmosphere by a fall of temperature, condenses upon nuclei of dust particles. It has been shown that air, charged with aqueous vapour and carefully freed from such nuclei, can be cooled much below the point at which a cloud should form, without the separation of any droplets of moisture. On admission of suitable nuclei condensation at once ensues. The presence of dust particles in air is easily made obvious by the passage of a beam of light, the reflection of the light from the surface of the dust particles rendering very evident the path of the beam of light. Small quantities of ammonia, arising from the decomposition of nitrogenous matter and of nitric acid, formed during thunder- storms, are responsible for the presence of traces of ammonium nitrate in the atmosphere. Is the Air a Mixture ? — The removal of oxygen from the air by the action of burning phosphorus in no way affords any indication as to whether the oxygen in the air is present in a state of combination or as one component in a gaseous mixture. The following reasons, based partly upon chemical and partly upon physical evidence, lead us to the conclusion that air is a homogeneous mixture of gases. 1. If oxygen, nitrogen, etc., are mixed in the ratio in which they occur in air, the resultant product has aU the properties of natural air, but when the mixing takes place there is no energy change (heat effect, etc.), as there is when chemical combination occiirs. 2. The proportion by weight in which oxygen and nitrogen ATMOSPHERIC AIR 273 are found in air has no simple relation to the atomic weights of these elements. The atomic proportion in which they occur is 3-77 : 1, a result difficult to reconcile with Dalton's Law of Multiple Proportions. 3. Slight variations in the composition of the air have been noticed, whereas constancy of chemical composition is a criterion of a chemical compound. 4. The physical properties of air (density, etc.) may be readily calculated from a knowledge of the composition of the air and the physical properties of the gases of which it is composed. The power of the air to refract light, to absorb heat, etc., is directly calculable from such data at our disposal as the refractive index of oxygen, etc. This would not be so, if chemical com- bination had occurred between the oxygen and the nitrogen. 5. When air dissolves in water, the nitrogen and oxygen do not dissolve in the proportion in which they exist in air as they would do if in chemical combination, but they dissolve independently in proportion to their partial pressure and to their solubUities. 6. If hquefied air is allowed to boil, the nitrogen escapes first. 7. The oxygen and the nitrogen can be partially separated by diffusion through a porous wall. This is to be expected for a mixture of gases of different density, but not for a compound of two gases. The Liquefaction of the Air. — The behaviour of carbon dioxide under increasing pressures led Andrews (p. 66) to the conclusion that a gas is not capable of liquefaction however high the pressure, unless the temperature is below a certain limit — the critical temperature of the gas under consideration. In bringing about the liquefaction of the less easily condensable gases, it is therefore necessary to secure a low temperature as well as a high pressure. In the machines in use at the present day for the liquefaction of gases, the requisite low temperature is obtained by taking advantage of the fact that all gases, except the theoretically perfect gas, when expanding into a vacuum, experience a fall in temperature. Owing to the slight cohesion which exists between the molecules of a gas which has to be overcome when expansion occurs, work is done at the moment of expansion, and the energy requisite for the tearing apart of the molecules is supplied by the gas, i.e. the gas cools T 274 AN INORGANIC CHEMISTRY itself. The cumulative effect ^ik^Lj lifmmk: Fig. 86. of the incoming compressed air of liquid air fall from the orifice. The Uquid is collected in Dewar flasks. These are double-walled flasks, in which the annu- lar space between the outer and inner walls has been exhausted by means of a special pump. The vacuum thereby created between the walls of the vessel offers a nearly perfect barrier to the transmission of heat ; its of this cooling is attained by allowing the cooled gas to circulate around the pipe leading the compressed gases to the orifice where the ex- pansion is effected. A refer- ence to the diagram, Fig. 86, makes this clear. Air, purified from such chance impurities as carbon dioxide and aqueous vapour, by passage through a drum containing quicklime and afterwards over solid caus- tic soda, is forced through the inner tube under a pres- sure of about 200 atmospheres. By manipulating the valve A, this compressed air is allowed to expand into the receptacle B. The air, cooled in this ex- pansion, flows back through a tube which forms a jacket round the tube conveying the compressed air to the orifice. A cumulative effect is thus secured and the temperature slowly falls, until at last drops Fig. 87. THE RARE GASES 275 efficiency in this respect is still further improved by silvering the inner walls. Liquid air boUs at about —190°, but owing to the volatility of the nitrogen, its composition and boiling point quickly change, and a liquid much richer in oxygen collects in the flask. From this residue oxygen is obtained and compressed into cylin- ders for the market. This is the main source of commercial oxygen. The Rabe Gases Whilst engaged upon investigations upon the density of nitrogen, Lord Rayleigh (1893) made the interesting discovery that the density of nitrogen appeared to vary according to whether its source was atmospheric or chemical. Atmospheric nitrogen gave a density 14-070, nitrogen prepared from am- monium nitrate, oxides of nitrogen, etc., gave a value 14-005 — a difference far exceeding his experimental error. The suspicion was at once formed that a heavier gas was present in small quantity in nitrogen isolated from the atmosphere. In 1894 Sir William Ramsay succeeded in separating the new gas, which he named argon (Gk. ap769, inert, idle). Argon can be obtained by passing atmospheric nitrogen over heated magnesium or calcium. A nitride is formed, and the unreactive argon is collected. Rayleigh obtained a small quantity of argon by sparking a mixture of nitrogen and oxygen in the presence of a solution of potassium hydroxide, which absorbed the oxides of nitrogen as they were formed. After prolonged sparking the excess of oxygen was removed and a smaU bubble of gas remained. The properties of this gas were identical with those of Ramsay's argon. This ex- periment of Rayleigh's formed an interesting repetition of one carried out nearly a century before by Cavendish. This chemist used the same method to determine whether atmo- spheric nitrogen was a homogeneous, single substance, and his conclusion was that, if any part of the nitrogen of the atmo- sphere differed from the rest, it could not exceed 1/120 of the whole. Rayleigh found 1/84. Argon has a density of 39-9. When liquefied, it boils at —186°. It is about 2| times as soluble as nitrogen. Its most characteristic chemical property is its utter lac^ of reactivity, for it combines with no element. 276 AN INORGANIC CHEMISTRY Ramsay further succeeded in isolating minute quantities of other gases, showing the same lack of reactivity — helium (density=4), neon (density=20), krypton (density=82-9), xenon (density=130-2). The argon obtained from atmospheric nitrogen was liquefied by means of liquid air, and this liquid was then subjected to fractional distillation. The more volatile fraction was cooled with liquid hydrogen, and a white solid, neon, separated out, while hehum was left uncondensed. The less volatile fraction was subjected to further fractional distillation and krypton and xenon were thus isolated. From physical evidence it has been concluded that all the rare gases are monotomic, so that the molecular weights of these gases, adduced from their densities, also represent their atomic weights. The spectrum of Ramsay's heUum proved to be identical with the spectrum of an unknown element discovered ia the sun's prominences by Lockyer in 1868, and named by hiTn helium (Gk. •^Xio?, the sun). It had also been known for some years that certaia minerals, e.g. cleveite, fergusonite, uraninite, evolved, on being heated in vacuo, a gas which had been assumed to be nitrogen. Ramsay iuvestigated this gas and found that it was hehum. This gas has also been found in the waters of certain springs. Further discussion of the source and origin of this hehum will be continued in the chapter dealing with radio- activity (Ch. xli.). Questions 1. Show how the composition of air by vohime may be determined. 2 Calculate the percentage composition of air by volume from the following analyses of Bunsen : — Air employed . . . 850-2 c.c. at 547-5 mm. pressure and 64° C. After addition of hydrogen 1010-5 „ 694-5 „ „ 6-4 After explosion . . 811-2 „ 502-8 „ „ 6-2 3. What volume of .air measured at 20° C. and 750 mm. pressure is required for the oxidation of 100 gm. of iron pyrites ? 4. The capacity of a room is 2,500 cu. ft. Find the weight of water vapour contained in it, if the barometric pressure be 720 mm., the tem- perature 20° and the air saturated -with aqueous vapour. (Tension of aqueous vapour at 20° = 17-39 mm.) 5. Give a brief historical account of the isolation of the gases present in the atmosphere. 6. What justification exists for considering the air to be a mixture of gases rather than a chemical compound ? 7. What weight of copper is required for the removal of the oxygen cpntt^ined in 5 litres of air, measured at 10° C. and 760 mm. pressure 7 CHAPTER XIX THE HYDRIDES AND HALIDES OF NITROGEN Four reduction products of nitrogen are known ; they are ammonia, NH3 ; hydrazine, N2H4 ; hydrazoic acid, N3H ; hydroxylamine, NHjOH. The most important haUdes are : Nitrogen trichloride, NCI 3 ; nitrogen iodide, N2H3l3=(NH3,Nl3). Ammonia History. — The alchemists were familiar with ammonia, chiefly in the form of its salts and its aqueous solution ; but it was not tin the time of Priestley (1774) that the gas was isolated by heating sal-ammoniac with lime and collecting the evolved gas over mercury. Occurrence. — In combination with nitric and carbonic acids ammonia exists in small quantities in the air. It is also formed by the decay of vegetable and animal matter, e.g. the odour of ammonia may generally be detected in the neighbourhood of stables. It is evolved with boric acid from the fumaroles of Tuscany, and small quantities are also deposited in the neigh- bourhood of volcanoes as the chloride and sulphate. Preparation and Manufacture. — 1. When complex organic compounds containing hydrogen and nitro'gen are heated in the absence of air, e.g. feathers, horn, coal and hides, ammonia is evolved. This is particularly noticeable it the substance is heated with soda lime, i.e. quicklime which has been treated with a strong solution of sodium hydroxide. Large quantities of ammonia are set free ia the destructive distUlation of coal . This valuable product is washed out of the gas by treatment with water in which the ammonia is freely soluble. The ammoniaoal liquor is treated with an excess of lime and heated. Ammonia is set free and is passed either into hydrochloric or sulphuric acid for the 277 278 AN INORGAMC CHEMISTRY preparation of the chloride or sulphate respectively, or is passed into water to form the aqua ammonia of the commercial world. 2. Ammonia is formed by the action of nascent hydrogen upon oxy-compounds of nitrogen, e.g. nitrates and nitrites. Thus, zinc, sulphuric acid and sodium nitrate react thus : Zn +H2S04-^ZnS04 +2H NaN03 + H2S04-^NaHS04 +HNO3 2HN0 3 + 16H -> 2NH 3 + 6H ,0 3. Ammonia is prepared by heating any of its salts with the alkali-hydroxides or slaked Hme, i.e. with a less volatile base. The equilibrium 2NH4a+Ca(OH)2 ^=±CaCl2-|-2H20+2NH3 which is attained in a sealed system, is not reached in open vessels owing to the escape of the volatile ammonia, and complete decomposition of the ammonium salt results. 4. All nitrides decompose under the action of water with the liberation of ammonia : CaaNa + 6Hi,0 ^ 3Ca(OH)2 + 2NH3. 5. The action of steam upon calcium cyanamide (p. 545) leads to the evolution of ammonia : CaCN2+3H20-^CaC03 + 2NHj. This reaction forms the basis of one of the most important methods for converting atmospheric nitrogen into ammonia. 6. Under the action of electric sparks nitrogen and hydrogen unite to form small quantities of ammonia. On the other hand, if ammonia is subjected to the action of electric sparks, it decom- poses, so that an equUibrium must be reached. Under such conditions only about 2-3 per cent, of ammonia is left unde- composed. The combination of nitrogen and hydrogen is attended by the evolution of heat, so that with rising temperature the percentage of ammonia in the equilibrium mixture must steadily fall (Le ChateUer's Law). TABLE 30 Pressure in Per cent. NH3 in the equilibrium mixture N.+3H.,^ 2XH3 at atmospheres. different temperatures. 550° 750° 850° 950° 1 0-077 0016 0009 005 100 6-71 1-54 0-87 0-54 200 11-9 2-99 1 1-68 1-07 HYDRIDES OF NITROGEN 279 It is therefore evident that, in order that the direct combination of hydrogen and nitrogen may be used as a commercial method of manufacturing ammonia, the conditions must be such that the combination is effected at the lowest possible temperature. In practice the reaction is catalysed by such catalysts as the compounds of osmium, uranium, etc., so also by a mixture of iron and molybdenum. Advantage is also taken of the effect produced upon the equilibrium mixture by increasing the pressure. The application of Le Chatelier's Law predicts that with rising pressure the yield of ammonia should increase {see Table 30). Large quantities of ammonia are now made by this process. Physical Properties. — Ammonia is a gas possessing a pvmgent odour. The density is a little over halt that of air. ATiimoma is extremely soluble in water, one volume of water at 0° dissolving 1298 volumes (760 mm. pressure), at 20° 710 volumes. The great solubility of this gas can be effectively shown in the apparatus described in connection with sulphur dioxide (p. 230). Liquid ammonia boils at — 34° and the solid melts at — 77°. In order to evaporate one gram molecule of ammonia (17 gm.) a relatively large amount of heat is required, a fact which, together with the case of hqui- faction of the gas, accounts for its extensive use as a refrigerant. The gas is liquefied by compression. The heat Kberated during compres- sion is conducted away by the cold water flowing over the condensing pipes. The liquid ammonia then flows into pipes immersed in brine. Evaporation is facilitated by lower- ing the pressure and the heat ab- sorbed in the evaporation is taken from the brine. This cold brine is now ready for circulation through the chilling rooms. There is rela- tively little loss of ammonia or of brine, and the process is a continuous one. Fig. 88. Chemical Properties. — Ammo- nia is not a supporter of combustion nor does it burn ip air. Oxyge n 280 AN INORGANIC CHEMISTRY However, it burns feebly in an oxygen atmosphere with the production of nitrogen, water and small quantities of the oxides of nitrogen. The catalytic oxidation of ammonia is effectively shown by bubbling a stream of oxygen through a strong solution of ammonia (Fig. 88). If a warm platinum wire is suspended over the surface of the ammonia, it is seen to glow vigorously, and the solution is found to contain ammonium nitrate and nitrite. Ammonia is a good reducing agent, converting oxides of metals into metals (see p. 267). When passed over heated sodium or potassium, dry ammonia forms an amide in which one atom of the hydrogen present in ammonia is replaced by the metal : 2NH3 + 2Na-> 2NH2Na + H^. Sodamide has a metallic appearance and evolves ammonia under the action of \\'ater : NHjNa + HOH-^NaOH+NHs. Chlorine and bromine react with ammonia, hberating nitrogen : 2NH3 + 3Cl2-^ 6HC1 + N,. An adaptation of this reaction is occasionally made use of for generating a continuous supply of nitrogen : 2NH4CI + SCla^ 8HC1 + N2. The aqueous solution of ammonia has an alkaJine reaction owing to the formation of ammonium hydroxide : NH3 + H2O ^=± NH4OH. On boiling such a solution the volatile ammonia escapes, and in order to restore the equilibrium, the hydroxide decomposes. Continued boiling will drive out practically the whole of the ammonia, whether combined or merely dissolved. The Composition of Ammonia. — The dissociation of ammonia by the electric spark can be utilised for showing the composition by volume of this gas. If 50 c.c. of ammonia is subjected to an electric discharge in a eudiometer, the volume is found to be nearly doubled. Analysis of the resultant mixture shows that the gases formed are hydrogen and nitrogen, hence two volumes of ammonia produce approxi- mately four volumes of the mixed gases (Pig. 89). Further evidence on the composition of ammonia is obtained HYDRIDES OF NITROGEN 281 by Hoffmann's volumetric method. A long glass tube, divided into three equal portions, is filled with chlorine at atmospheric pressure (Fig. 90). A few drops of concentrated ammonia are allowed to drop into the tube from the dropping funnel. The reaction between the chlorine and the ammonia is attended by considerable heat evolution. When the reaction is complete, the excess of ammonia is removed by a Uttle sulphuric acid. The lower tap is then opened under water, and by immersion in a high cylinder, the volume is read off at atmospheric pr e s s u r e. The residual gas occupies one-third of the original volume and is found to be nitrogen. The hydro- gen of the ammonia has combined with the chlor- ine to form hydrogen chloride in the ratio of one to one. The hydro- gen chloride then forms ammonium chloride, the volume of which is negh- gible. The sulphuric acid merely removes any ex- cess of ammonia. The three volumes of chlorine combine with three volumes of hydrogen, thereby setting free one volume of nitrogen, hence it is inferred that ammo- nia must contain one volume of nitrogen combined with three volumes of hydrogen. On the assumption of the diatomic nature of the molecules of nitrogen and hydrogen, it follows from Avogadro's Hypothesis that ammonia must contain hydro- gen and nitrogen combined in the atomic ratio of 3 : 1. In order to fix the value of n in the formula (NHj)^, the usual vapour density measurement is necessary to get the molecular weight. This proves to be about 17 (0=32). Hence n is unity and the formula is NH,. Fig. 89. Fig. 90. 282 AN INORGANIC CHEMISTRY The Ammonium Salts. — ^Ammonia combines directly with an acid producing a salt NH3+HCI -^NHjCl NH3 +HN03^NH4NOs. In these salts the univalent radicle NH4 is always present, and has received the name ammonium. A solution of ammonia in water contains the base, ammonium hydroxide NHa + HaO^z^NH^OH, and the general properties of the salts, derived from this hydroxide by the action of acids, are so similar to the properties of the salts of sodium, potassium, etc., that the ammonium salts are generally classed with the salts of the alkalies (see p. 491). The nitrogen • present in the ammonium salts is generally considered to be pentavalent, as opposed to its trivalence in ammonia itself. N^H Nn^CI ^H W \H Hydrazine Hydrazine, NaHj, was first prepared by Curtius (1887). It is a colourless liquid which freezes at 1-4°. Its boiling point is 50° at 71 mm., 113-5° at 761-5 mm. It forms a series of salts very similar to the ammonium salts, hydrazine monohydro- chloride N2H4-HCI, hydrazuie dihydrochloride N2H4-2HC1, hydrazine sulphate N2H4-H2SO4, etc. Oxidising agents are vigorously reduced by hydrazine and its compounds. They reduce cupric salts to red cuprous oxide, silver salts to silver, sulphur to hydrogen sulphide, iodine to hydrogen iodide, etc. On exposure to air, hydrazine fumes freely, owing to the formation of hydrazine hydrate, N2H4 + H20^N2H50H. Hydrazine hydrate is a colourless, fuming liquid, boiling at 118-5°. It attacks glass, corks and india-rubber. Hydrazine hydrate is fairly easily oxidised by atmospheric oxygen N2H4-H20+02^^N2 + 3H20. The preparation of hydrazine and its salts is often achieved by organic methods, but the best method is the following : Mix 200 CO. of a 20 per cent, solution of ammonia, 5 c.c. of a 1 per cent. HYDRIDES OF NITROGEN 283 solution of gelatine and 7| gm. of sodium hypochlorite dissolved in 100 c.c. of water. The mixture is boiled for half an hour. After cooling, the flask is placed in ice cold water, and sulphuric acid is slowly added. This causes the separation of hydrazine sulphate, one of the least soluble of the hydrazine salts. On distiUing hydrazine sulphate and potassium hydroxide from a silver or platinum vessel, hydrazine hydrate is liberated : N2H4-H2S04 + 2KOH-^N2H,-Hs,0 + K.SO^ + H^O. Hydrazine itself is obtained by distilling hydrazine hydrate with barium oxide under reduced pressure in an atmosphere of hydrogen : BaO + N^Hi-H^O -^ N^H, + Ba(OH) ,. Hydrazine and its salts are among the most powerful reducing agents known. Hydeazoic Acid, HN This acid was also discovered by Curtius (1890). The most satisfactory inorganic method of preparation is indicated in the equation : NHjNa + N2O -> NaNa+HjO Sodamide. Nitrous oxide. Sodium hydrazoate. The sodium hydrazoate is dissolved in water acidified with sulphuric acid and distilled. Another method is by the action of nitrous acid upon a cold aqueous solution of hydrazine hydrate : N^Hi-HaO +HN02-^HN3 + 3H2O. The acid is obtained by distilling the mixture. Repeated distillation furnishes an acid of about 90 per cent, strength. Further dehydration is effected by means of calcium chloride. The pure acid boils at 37°, an operation fraught with considerable danger owing to the extreme instabiUty of the acid. Hydrazoic acid is a strongly endothermic body, and its explosive decom- position is attended by the liberation of considerable heat : HNjaq. -^ 3N + H + 62-1 Cal. Hydrazoic acid is a fairly strong monobasic acid which attacks many metals freely, e.g. zinc, aluminium, with the evolution of hydrogen and ammonia. Its salts are known as hydrazoates 284 AN INORGANIC CHEMISTRY or azides. Silver azide, AgNa, is white, and insoluble, not unlike silver chloride in appearance. Hydroxylamine, NH2OII Hydroxylamine is obtained by the action of nascent hydrogen upon nitric oxide, nitric acid, etc. Thus, if a stream of nitric oxide is led through hydrochloric acid in which tin is dissolving, hydroxylamine hydrochloride is formed, NO+SH-^NH^OH. The stannous chloride present in the solution is precipitated by the action of hydrogen sulphide, the liquid filtered and evaporated to dryness. The residue is treated with absolute alcohol, in which hydroxylamine hydrochloride is freely soluble. On eva- porating the alcohol, the white crystalline hydrochloride is obtained. The extraction with alcohol is necessary, as ammonium chloride, which is insoluble in alcohol, is also formed by the more complete reduction of the nitric oxide, NO+SH^^NHa+H^O. Nitric acid can be electrolytically reduced to hydroxylamine by using either a mercury cathode or a cathode consisting of amalgamated lead. The easiest method of obtaining hydroxylamine itself is by the distillation in vacuo of hydroxylamine phosphate, (NH20H)3H3P04-^H3P04+ SNH^OH. Nearly the whole of the hydroxylamine can be obtained by this means. It is a white solid melting at 33° and boiling at 58° under 22 mm. pressure. It is very unstable, and even at temperatures below the melting point begins to decompose. Hydroxylamine resembles ammonia in many respects — it forms a base approaching ammonium hydroxide in strength, and from this base a series of salts has been prepared. These salts, like the azides, are all Hable to explode on heatmg. Hydroxylamine is a strong reducing agent, reducing silver nitrate to metaUic silver ; oxidising agents, hke the halogens and potas- sium permanganate oxidise it with extreme vigour. Copper salts in alkaline solution give cuprous oxide, a reaction whereby even 0-00001 gm. of hydroxylamine in a litre of water can be detected. HALIDES OF NITROGEN 285 The accepted constitution of the hydrides of nitrogen is represented in the following constitutional formulae : -OH H H H N H N^H >N— N<( ll>N— H \N- \h W ^H N^ H^ Ainmonia. Hydrazine. Hydrazoic acid. Hydroxylamine. i.e. the nitrogen is considered to be trivalent in aU these compounds, and therefore in a state of unsaturation. The formation of salts from the compounds derived from ammonia, hydrazine and hydroxylamine, is presumed to be attended by a valence change, i.e. the nitrogen atom becomes penta- valent : X' sH 2 c> NCI3 + 4HC1, and also by the electrolysis of ammonium chloride. In this case the chlorine, liberated at the anode, acts upon the ammonium chloride in the solution, producing drops of nitrogen chloride. It is an oUy liquid which is extremely explosive. Nitrogen iodide is formed by the action of strong aqueous ammonia upon powdered iodine. Nitrogen iodide has also been prepared by the action of ammonia upon an alkaline solution of potassium hypoiodite. When wet, it may be handled with safety, but if dry, it decomposes with violence, even when disturbed by a feather or by a falling dust particle. The composition of nitrogen iodide is NHj-NIg. Other similar compounds of nitrogen iodide ?ind ammonia have been isolated. 286 AN INORGAOTC CHEMISTRY Questions 1. Give an account of the commercial synthesis of ammonia. 2. Give a brief comparative account of the hydrides- of nitrogen. 3. What reasons have we for assigning the formula NH3 to ammonia ? 4. Describe the manufacture of hydroxylamine. AMiat are its more important properties ? 5. What steps should be taken to secure : (a) dry ammonia, (6) liquid ammonia ? 6. A eudiometer contains 15 c.c. of ammonia ; after sparking the volume is found to have increased to 29-95 c.c. ; 50 c.c. of oxygen are then added and the mixture exploded. The volume after the explosion is found to be 46-2 c.c. Assumiiig that the temperature and pressvire remained constant throughout the experiment, what inference concerning the composition of ammonia would you draw from these figures ? 7. What is the object of distilling hydrazine under reduced pressure, whenits boiling point is comparatively low (113°) ? CHAPTER XX THE OXIDES AND OXY-ACIDS OF NITROGEN The following oxides and oxy-acids of nitrogen are known : Oxides. Acids. Nitrous oxide (hyponitrous anhydride) NjO Hyponitrous acid (HN0)3 Nitric Oxide NO Nitrogen trioxide (nitrous anhydride) N2O3 Nitrous acid HNO2 Nitrogen tetroxide (nitrogen peroxide) N2O4 Nitrogen pentoxide (nitric anhydride) NjOj Nitric acid Two of these oxides give rise to no acid— nitrogen tetroxide and nitric oxide. So far as the other oxides are concerned, the guiding principle already emphasised in deaUng with the oxides of chlorine and of sulphur will be found of use in correlating not only the properties of the oxides themselves but also of the acids derived from them. As the oxygen content of the oxide increases, the acid-forming tendency of the oxide becomes more 'pronounced, and, in general, the richer the oxy-acid is in oxygen, the stronger is the acid, and the more stable will the acid and its salts be. NiTEic Acid and Nitrogen Pentoxide (Nitric Anhydride) Occurrence. — Salts of nitric acid occur naturally though the distribution is somewhat localised. Immense beds of sodium nitrate (Chili saltpetre) are found in the desert regions on the west coast of S. America, ChUi and Peru. The deposit is over 400 sq. miles in area, and in some parts attains a thickness of five feet. The nitre rock, known as caHche, contains on the average about 20-30 per cent, of sodium nitrate. Small quantities of 287 288 AN INORGANIC CHEMISTRY potassium nitrate are found in the neighbourhood of insanitary villages of India, Arabia, etc. Under the action of nitrifying bacteria, urine and other nitrogenous matter are oxidised to nitric acid and nitrates of calcium and potassium crystallise on the surface. These are extracted by leaching with water and the extract treated with wood ashes (potassium carbonate). The solution of potassium nitrate is then evaporated and the nitre purified by re-crystallisation. Considerable quantities of Bengal Saltpetre are prepared in this way. Ca(N03)2+K2C03->CaC03i +2KNO3. Preparation and Manufacture. — In the laboratory nitric acid is prepared by the action of sulphuric acid upon a nitrate, generally sodium nitrate. NaNOa +H2S04->NaHS04 + HNO3. The principles governing this reaction are similar to those involved in the preparation of hydrochloric acid from salt (q.v.). This reaction is largely used for the manufacture of nitric acid on the commercial scale, except that, at the higher temperature attained in practice, some of the acid sulphate breaks down — NaN03 +NaHS04^Na2S04 +HNO3, so that the complete reaction is represented by the equation 2NaN03 +H,S04-^Na.,S0i +2HNO3. The retorts employed in the operation are large cast-iron cyhn- ders. The vapours are led through a series of earthenware pots where absorption of the acid takes place. The final traces of nitric acid and of the oxides of nitrogen, produced by the decom- position of the nitric acid, are removed by passing the escaping gases through towers, fiUed with coke, down which water trickles. The Niteogbn Cycle in Nature and the Fixation of Atmospheric Nitrogen. — Few elements play a more important part in nature than does nitrogen. This element is an essential constituent of aU animal tissues, and is always present in veget- able matter in the form of complex nitrogenous products. The animal kingdom secures its supply of nitrogenous matter from plant food, and hence the problem of maintaining the supply of nitrogen in a form capable of absorption by plants is one of national importance. Small quantities of ammonium nitrate are OXIDES AND OXY-ACIDS OF NITROGEN 289 formed naturally (see 272), whilst a few plants belonging to the order Leguminosce are fortunate enough to have a ready made supply of nitrogenous compounds at hand. These plants, including the clover, beans, peas, lupins and lucerne, act as hosts for certain bacteria which form nodules upon the rootlets of the plant. These bacteria possess the property of assimilating the nitrogen of the atmosphere and converting it into a form in which it is available for plant food, a process whereby the growth of the host is effectively stimulated. Under the action of nitrifying and de-nitrifying bacteria, complex nitrogenous matter present in the soil is either oxidised to nitrates or broken down to such an extent that the nitrogen is set free and therefore lost to the vegetable kingdom. The upshot of this is that the available nitrogen in the soil is insufficient if anything in the nature of intensive cultivation is desired. SuppUes of readily available nitrogen must therefore be added to the soil in the form of manures and fertilisers. In nearly all cases fertilisers are rich in ammonium salts or in nitrates, both of which are easily assimilated by plant life. The following diagram, for which we are indebted to v. Braun, illustrates the various pro- cesses of nature in which nitrogen plays a part : Atmospheric nitrogen. Original ly bound nitrogen. Plant decomposition. Animal digestion. Destructive distill- ation. Over and above the small quantities of nitrogen rendered available annually by nature, man has at his disposal steadily diminishing supplies of Chili saltpetre, but so great has been the demand for sodium nitrate as fertiliser and for munition pur- poses in the recent Great War, that the complete exhaustion of these Chili saltpetre beds is not far distant. The amount of V 290 AN INORGANIC CHEMISTRY ammonia derived from the destructive distUIation of coal is certainly of value in eking out our supplies of nitrogen fertilisers, but none the less the discovery of fresh sources of supply of these fertilisers has long been of momentous importance to the economic world. During the past twenty years three important methods of fixing atmospheric nitrogen have sprung into importance : 1. The preparation of ammonia from calcium cyanamide (p. 540). 2. The preparation of ammonia by the direct combination of the elements (278). 3. The direct oxidation of atmospheric nitrogen by the oxygen under the stimulus of the electric arc and the subsequent absorp- tion of the oxides of nitrogen in water or in solutions of the alkalies. -In the oxidation of the ammonia derived from either of the first two processes, the conversion of the ammonia into the nitric acid takes place thus : NH3 + 20a->HN03 + HA though under certain conditions the reaction 4NH3 + 302^-2N2 + 6H2O also occurs. The problem of the chemist is to so control these reactions that the second is reduced to a negligible amount. This is done by the use of a grid of platiniun. Ammonia, mixed with ten times its volume of air, is led into the catalysing cham- ber. After the reaction the nitric acid vapours are led into a series of absorption towers where acid of any desired strength can be drawn off. The oxidation of nitrogen by oxygen is really based upon a modification of Cavendish's old experiment (q.v.). Recent work upon the equilibrium N.-fO- ^ — ^ 2NO has shown that the combination is attended by heat absorption, hence the higher the temperature the greater will be the yield of nitric oxide in the equUibrium mixture. Thus for a mixture of equal volumes of nitrogen and oxygen the following values have been recorded : Temperature . . 2000° 2402° 2927° 3000° Nitric oxide per cent. 1-20 2-23 5-0 5-3 whilst it has been computed that at the temperature of the electric arc nearly 10 per cent, of nitric oxide is formed. But as OXIDES AND OXY-ACIDS OP NITROGEN 291 the gases leave the region of high temperature and cool down, the eqtiiUbrium reverses and seeks to readjust itself to the lower temperature prevailing outside the arc, in other words, dissocia- tion of the nitric oxide formed at the high temperature of the electric arc wiU occur. The problem is to render this loss as low as possible. A drop of 10° in the temperature reduces the velocity of the reaction to about one half. At a temperature of 600-700° the velocity of this decomposition has already sunk so low as to be neghgible. Provided that the gases are swept away from the arc and reduced to this temperature sufficiently quickly, a yield approaching the maximum may be attained. The more rapidly the chiUing is effected, the more nearly will the composition of the resulting mixture approach the value ruling at the high tem- perature of the electric arc. f—vJlQililfiil/ One of the most prominent types of arc for achieving this result is that of Birkeland and ^ _ _ _ Eyde. Their ' '\SS:0^://// idea is not to use a short, thick flame, but a long, thin one with the maximum surface. This effect is produced by using an electro-magnet at right angles to the arc. In the above diagrammatical figure the electrodes consist of copper tubes, through which water circulates. The arc is blown out by the electro-magnet into a wheel-Uke disc of flame com- posed of burning oxygen and nitrogen. Other well-known types of arc are those used in the Pauling and in the Schonherr furnaces. The gases containing the nitric oxide are swept into an oxida- tion chamber where the reaction 9,T\rn-|-0. ^^ — ^ 2NO. occurs. The resulting gases are then passed through the usual series of absorption towers. In these towers the nitric oxide dissolves either in water or in alkaUne liquors. In the first towers nearly pure nitric acid is collected, whilst from those at the end of the train nearly pure nitrous acid is obtained in the form of its I I I t I /iV'^ -^ iVi i > ' ^-.< Fig. 92. 292 AN INORGANIC CHEMISTRY sodium salt. In the Norwegian plant (Birkeland-Eyde) it is the practice to convert much of the nitric acid into calcium nitrate by allowing the acid to act upon beds of limestone. Physical Properties of Nitric Acid. — Nitric acid is a colourless liquid boiling at 86° and freezing to a white soUd at —47°. The acid fumes strongly in air. An aqueous solution containing 68 per cent, of acid forms a mixture of maximum boiling point (120-5°). For solutions more dilute than 68 per cent, the distUlate will consist mainly of water, for solutions stronger than 68 per cent, the distUlate will consist mainly of nitric acid. In both these cases the boiUng point of the solution rises as the distillation progresses, until at 120-5'' an acid of constant composition distils over. The concentration of nitric acid must therefore be effected by means other than distillation. The usual practice to secure concentrated acid is to distU from strong sulphvu-ic acid which retains all the water. Chemical Properties. — The chemical properties of nitric acid, other than those of a purely acidic nature, are dependent upon the high oxygen content of this acid and upon the com- parative ease with which some of this oxygen can be liberated. 1. Nitric acid is easily decomposed by heat, even the tempera- ture of distillation being sufficient to bring about incipient decomposition : 4HN03->2H20 +4NO2 + O2. This decomposition is effectively shown by allowing concen- trated nitric acid to drop into a quartz flask heated to a dvill red heat. Copious fumes of brown peroxide escape from the flask. The brown tint of much of the nitric acid of commerce is due to the presence of dissolved fumes of nitrogen peroxide. They can be removed by bubbling a stream of carbon dioxide through the acid. 2. The aqueous solution of nitric acid is comparatively stable (cf. chloric acid). 3. Definite hydrates of nitric acid, HNOj.SHaO ; HNOj.HjO, have been separated by freezing an aqueous solution of nitric acid. These hydrates are very unstable. 4. Nitric acid, even in aqueous solution, is a powerful oxidising agent, especially fuming nitric acid, formed by distilling nitric acid with a little starch. Coal gas will burn freely under fuming OXIDES AND OXY-ACIDS OF NITROGEN 293 nitric acid, whilst organic matter like sawdust, turpentine, etc., is oxidised rapidly with the evolution of copious fumes of the oxides of nitrogen. A striking illustration of this fact is afiorded by the action of fuming nitric acid upon cane sugar, when a copious evolution of carbon dioxide and of the oxides of nitrogen takes place. Similarly, iodine is oxidised to iodic acid, phosphorus to phosphoric acid, hydrogen sulphide, sulphur and sulphurous acid to sulphuric acid, carbon to carbon dioxide, ferrous salts to ferric, e.g. : HjS + SHNOs^^H^SOi + 8NO2 + 4H2O. {4H20,NA) Such an equation is only a rough representation of such a reac- tion, as appreciable quantities of the oxide, NO, are also formed. The concentration of the oxidising agent and of the reducing agent, coupled with the temperature, wiU determine precisely how far the reaction is represented by the equation above and how far by the equation : SRS + SHNOa^- 3H2SO4 + 8N0 + HLfi. (4H,0,N,05) The great oxidising power of fuming nitric acid arises from the presence in nitric acid of dissolved oxides of nitrogen. One of the strongest oxidising agents, aqua regia, is obtained by mixing hydrochloric and nitric acids in the ratio of three to one. This mixture will dissolve the noble metals, platinum, gold, etc. Aqua regia is supposed to contain nitrosyl chloride, NOCl, and free chlorine ; the solvent action is attributed to the action of this free chlorine, 3HC1 +HN03->2H20 + CI2 +N0C1 8HC1 + Pt + 2HN03-^ HaPtClo + 2N0C1 + 4H,0. 5. Nitric acid reacts vigorously with many carbon compounds, particularly when concentrated sulphuric acid is present to absorb the water generated in the reaction. Picric acid is formed by the action of nitric and sulphuric acids upon phenol. CeH50H + 3HO.N02->CoH20H(N02), + 3H20 Phenol. Trinitro-phenol or picric acid. Under similar conditions alcohols are nitrated with the produc- tion of nitro bodies. Thus, when glycerine is slowly added to 294 AN INORGANIC CHEMISTRY a cooled mixture of nitric and sulphuric acid, nitro-glycerine is formed : CH^OH HO.NO2 CH2O.NO2 CHOH + HO.NO2 -> CHO.NO2 + SH^O I I CH2OH HO.NO2 CH^O.NOj By a similar process gun-cotton is formed from cellulose and tri-nitro-toluol (T.N.T.) from toluol, a derivative of benzene. The extremely explosive nature of such compounds as nitro- glycerine, picric acid and tri-nitro-toluol is undoubtedly due to the fact that the elements required for the chemical changes taking place during the explosion are contained within the same molecule, whereas an explosive Uke gunpowder contains the various ingredients in the form of a mechanical mixture. Fine though the grain of the mixture may be and however thoroughly incorporated, the oxidation and decomposition which take place at the moment of explosion must necessarily be of a slower nature than when the reaction is intra-molecular, as in the case of nitro-glycerine, etc. In the nitration of cellulose above referred to, cotton is treated with a mixture of nitric and sulphuric acids. (CeHi„05), + 6HN03^Ci,H,.0,(NO.,)e + 6H,0. After drying, the product is known as gun-cotton. This substance burns briskly, but does not explode unless subject to a very violent shock, such as is given by the explosion of a percussion cap (see Mercury Fulminate). Gun-cotton is too violent an explosive except for such purposes as demanded by naval work (mines, torpedoes, etc.). In order to sober down its action and make it suitable for general explosive work (the fiUing of cart- ridges, etc.), it is dissolved up in acetone together with 30 parts of nitroglycerine and vaseline. This paste mass is rolled out and cut into strips. After the evaporation of the solvent (acetone), the substance is known as cordite. This is extensively used as a propeUant, and has the great advantage that the explosion is practically smokeless, hence the use of cordite in many smokeless powders. Dynamite is prepared by impregnating infusorial earth (kieselguhr) with nitro-glycerine. The extremely dangerous OXIDES AND OXY-ACIDS 01* NITROGEN 295 explosive, nitro-glycerine, is thereby rendered comparatively safe. The explosive properties of the dynamite are regulated by the addition of such compounds as ammonium nitrate, flour, etc . By a judicious selection of the explosives at hand the miner is able to control the type of explosion, and hence avoid unneces- sary pulverising of the rock or ore, such as would be caused by such an explosive as nitro-glycerine or gun-cotton. Nitric acid stains the skin a bright yellow, owing to the formation of xanthoproteic acid. Action of Nitric Acid upon Metals. — ^The difference in the action of this acid upon the various metals arises very largely from the change in the oxidising power of the acid with con- centration. Copper, hke other metals less active than tin, does not evolve hydrogen from nitric acid, but brings about a reduction of this oxidising agent. In the case of dilute acid the primary reduc- tion product appears to be nitrous acid, Cu-|-3HN03-^Cu(N03)2-fH20+HN02. The unstable nitrous acid undergoes decomposition according to the equation 2HNO2— ^-HjO-l-NO-l-NOa, whilst interaction between the nitrogen peroxide and water then takes place. 3NO2 -f H2O -> 2HNO3 -I- NO. The summation of these three interdependent reactions gives : 3Cu -I- 8HN03-> 3Cu(N03)2 -f 4H2O + 2N0. On passing nitric oxide, NO (q.v.), through concentrated nitric acid, copious fumes of nitrogen peroxide are evolved. These arise from the oxidation of the nitric oxide as well as from the reduction of the nitric acid. NO + 2HN03-> 3NO2 + H2O. HANA In the Mght of this equation the reaction between copper and concentrated nitric acid must lead to the evolution of nitrogen tetroxide and not of nitric oxide. The more active metals, e.g. zinc, magnesium, Uberate hydro- gen from nitric acid, but in the presence of this strong oxidising agent, very little free hydrogen actually escapes. From a consideration of the series NA-^N204->N203^^NO->N20->N2^NH20H->NH3, 296 AN INORGANIC CHEMISTRY it is evident that many reduction products are possible, but the ultimate one is ammonia. Probably the determining factor in the reduction is the pressure at which the hydrogen is liberated. When moderately strong nitric acid acts upon zinc, nitrous oxide is evolved, 4Zn + 10HNO3->4Zn(NO3)2 + 5H2O +N2O but with dilute acid ammonia is formed, with subsequent formation of ammonium nitrate by interaction with some of the nitric acid. 4Zn + 9HN03^4Zn(N03)2 + SH^O +NH3 NH3 + HN03^NH4N03. With tin the reduction is less complete, and appreciable quantities of hydroxylamine (q.v.) are produced. Nitrates. — All the nitrates of the metals are soluble in water. The nitrates of the alkaUes evolve oxygen on heating, forming nitrites, and therefore act as strong oxidising agents, 2KN03-^2KN02 + 02. The nitrates of the heavy metals generally evolve nitrogen peroxide and oxygen on heating, 2Pb(N03)2->2PbO + 4NO2 + Oj. From the commercial aspect, the two most important nitrates are undoubtedly the nitrate of sodium and of potassium, the former as a fertiliser, the latter as a constituent of gunpowder. Nitrogen Pentoxide (Nitric Anhydride). When nitric acid is heated with phosphoric anhydride, de- hydration of the nitric acid ensues, 2HNO3 + P A-^ 2HPO3 + N A, and the distillation of the mixture gives white crystals of nitrogen pentoxide. These crystals melt at 30° and the liquid boils at 45°. Nitrogen pentoxide combines freely with water to form nitric acid ; it cannot be kept in the anhydrous condition owing to its slow decomposition into the peroxide and oxygen. 2NA^4NO,+0,. Nitrogen Peroxide (Nitrogen Tetroxide) Nitrogen peroxide is prepared : 1. By the direct combination of oxygen and nitric oxide, 2NO + 02^=±2N02. OXIDES AND OXY-ACIDS OP NITROGEN 297 2, By the decomposition of the nitrates of the heavy metals 2Cu(N03)2->2CuO +2NO2 + 30,. 3. By the reduction of nitric acid by means of a suitable reducing agent, e.g. arsenious oxide. The reaction is generally somewhat complex, and a mixture of nitric oxide and of the peroxide is evolved. The gases are led through a spiral con- denser immersed in a freezing mixture. A blue hquid soon condenses. This consists of nitrogen peroxide together with more or less nitrogen trioxide. Oxygen is bubbled slowly through this Hquid, which is still kept immersed in the cold vessel. The blue colour changes slowly into a yellowish brown. This is liquid nitrogen peroxide. As a reducing agent one may use nitric oxide, NO + 2HNO3 ^=± 3NO2 + H2O. So long as the acid remains strong, the nitric oxide is immedi- ately oxidised, and the removal of the peroxide prevents a reversal of the reaction. Properties of Nitrogen Peroxide. — The most striking feature of this gas is the change of colour and of density which it undergoes on heating. This is shown in the appended table. TABLE 31 Temperature Density (H=2) . Per cent. NO^ molecules Colour .... 26-7° lOOl" 76-6 48-6 20 79-2 Nearly Brown colourless 135° 140° 46-2 46 99 100 Deep brown Deep brown In the hght of these results there is little doubt that at low tem- peratures the formula of nitrogen peroxide is N2O4, hence the name tetroxide which it often bears. With increasing temperature the equihbrium N^O^ ;^ — ^ 2NO; is forced to the right, and at 140° the dissociation is practically complete. It can also be predicted that, if at some definite temperature a mixture is in equilibrium in accordance with the equation N204^iz±2N02, an increase of pressure will drive the equihbrium to the left, whilst a decrease of pressure wiU have the opposite effect (Le Chatelier's Law). At a dull red heat the brown colour of the peroxide again begins to disappear, owing to its dissociation into the colourless gases, nitric oxide and oxygen. The complete equihbrium is defined in the equation : 298 AN INORGANIC CHEMISTRY N,0, ^^ 2NO2 ;=± 2NO + O2 Colourless Deep brown Colourless The comparative ease with which nitrogen peroxide dissoci- ates into nitric oxide and oxygen, 2NO2 ^^2N0 + O2, accounts for the strong oxidising properties of this substance. Hence phosphorus, which is burning strongly, will continue to bum brilliantly in nitrogen peroxide, the temperatiu'c of the burning phosphorus being sufficiently high to effect the decomposition of the gas. Carbon also burns freely in liquid nitrogen peroxide. Potassium iodide is oxidised to iodine, but on the other hand potassium permanganate is decolorised, i.e. reduced by nitrogen peroxide. 5N2O4 + 2KMn04 + 2H2O -> (KjO.MnA) 2KNO3 +2Mn(N03)2 + 4HNO3 (K2O.NA) (2MnO.NA) (2H2O.NA) When passed into warm water, nitrogen peroxide undergoes auto-oxidation (see p. 169) into nitric acid and nitric oxide, 3NO2 -f H2O ^^ 2HNO3 + NO (H2O.N2O5) but in cold water the auto-oxidation foUows a slightly different course, and a mixture of nitrous and nitric acids is produced, 4N02 + 2H20->2HN03 + 2HNO2 (HjO.NA) (H,O.N203) (cf. chlorine peroxide, ClOj). Nitrous Acid and Niteoqek Tbioxide (Nitrous Anhydride) Although the nitrates of the alkahes lose an atom of oxygen on being heated and form the corresponding nitrite, it is usual to facilitate the reduction by stirring lead into the molten nitrate, Pb + NaNO 3 -^ PbO + NaNO ,. The nitrite is dissolved away from the lead oxide and purified by crystallisation. Large quantities of sodium nitrite are also prepared in the absorption towers used to absorb the oxides of nitrogen during the oxidation of ammonia by oxygen (q.v.). When an acid is cautiously added to a cold, dilute solution of a nitrite, a pale blue solution containing nitrous acid is formed. OXIDES AMD OXY-ACIDS OF NITROGEN 299 The acid itself is very unstable, and breaks down at once on warming, 3HN02^HN03 + 2NO+H20, i.e. auto-oxidation occurs — a type of reaction common to the chlorites, hypochlorites, phosphites and hypophosphites. This instabihty of the lower oxy-acid is thus frequently met with. The instability of nitrous acid explains the oxidising properties of this acid. Thus, a mixture of an iodide and a nitrite when acidified liberates iodine. 2HI + 2HN02-^l2 + 2H2O +2N0 Indigo can also be bleached by means of nitrous acid. On the other hand, nitrous acid being a reduction product of nitric acid can often function as a reducing agent, and strong oxidising agents, such as acidified potassium permanganate, react thvis : 5HNO2 + 2KMn04 +3H2S04-> (SH^O.NaOs) (KjO.MnA) K2SO4 + 2MnS04 + 5HNO3 + 3H2O (MnO.SOa) (fH^O.NA) In this reaction there has been a reduction of the oxygen valence of the manganese from seven to two, whilst the valence of the nitrogen towards oxygen has risen from three to five. As was found to be the case for the nitrates and nitric acid, the nitrites or salts of nitrous acid are much more stable than the acid from which they are derived. Nitrogen Trioxide (Nitrous Anhydride). The blue liquid formed by the condensation of the gases evolved during the action of nitric acid upon arsenious acid (see Nitrogen Peroxide), or by the condensation of a mixture of nitric oxide and nitrogen peroxide is generally supposed to be nitrogen trioxide. As soon as the temperature rises, dissociation sets in, NeOj^^NO+NOj, and the more volatile nitric oxide escapes, leaving the peroxide behind. If electric sparks are sent through liquid air, the trioxide is formed, and separates out as a pale blue powder when the liquid is aEowed to evaporate. The main evidence in favour of regarding nitrogen trioxide as a definite chemical substance is based upon various physical 300 AN INORGANIC CHEMISTRY measurements. It is maintained that in the absence of moisture the dissociation NaOa^^z^NO+NOj does not occur, and that dry nitrogen trioxide has a vapour density lying between the extremes 76 and 152 (0=32), i.e. the molecule is either simple, N2O3, or double, NjOe- Nitric Oxide Preparation. — On the commercial scale, nitric oxide is pre- pared by the direct synthesis of its elements as a prehminary step in the manufacture of nitric acid from atmospheric nitrogen (p. 291). In the laboratory it is nearly always obtained by the reduction of nitric acid. Nitric oxide of moderate purity is hberated by the action of nitric acid (sp. gr. 1-2) upon copper turnings. The gas obtained by this method nearly always con- tains small quantities of nitrous oxide and even free nitrogen, the amount of these impurities tending to accumulate as the reaction progresses. Pure nitric oxide is best obtained by reducing nitric acid with ferrous sulphate. In practice, potassium nitrate and ferrous sulphate, in the proportion of one of nitre to four of ferrous sulphate, are introduced into a flask with a httle water, and con- centrated sulphuric acid is slowly run in. On heating, a steady stream of pure nitric oxide is evolved. 2KN03 + H,S04 ->K2S04-f 2HNO3 6FeS04-f2HN03-|-3H2S04^3Fe2(S04)3+4H20+2NO 6FeS04-f2KN03+4H2S04^^3Fe2(S04)3-fK2S04+4H,0+2NO Properties. — Nitric oxide is a colourless gas which is somewhat difficult to liquefy. The critical temperature is — 93-5°, at which temperature a pressure of 71-2 atmospheres is required to liquefy it. At — 167° the hquid freezes to a snow-white solid. When brought into air, nitric oxide has the distinctive property of forming reddish brown fumes of nitrogen peroxide, 2N0-f 02^=±2N02 which is also produced when it is led through nitric acid. NO + 2HNO3 ^:± 3NO2 -f H^O Under the action of oxidising agents either a nitrite or a nitrate is formed. With lead or manganese dioxide nitric oxide yields a nitrite, but with potassium permanganate it forms a nitrate. An aqueous solution of iodine yields nitric and hydriodic acids. OXIDES AND OXY-ACIDS OF NITROGEN 301 In all these reactions nitric oxide functions as a reducing agent. On the other hand, nitric oxide occasionally behaves as an oxidising agent. Thus, nitric oxide and hydrogen yield ammonia when heated in the presence of platinum black, 2NO+5H2^2NH3+2HA In the same way, hydrogen sulphide and the sulphides are oxidised by nitric oxide with the formation of nitrous oxide. Nitric oxide is the most stable of the oxides of nitrogen. At 3000° about 5 per cent, of the nitric oxide suffers decomposition into nitrogen and oxygen ; and so, although the gas does not support the combustion of feebly burning substances, phos- phorus, if burning sufficiently strongly to cause dissociation of the gas, wiE continue to burn freely. Nitric oxide dissolves in an aqueous solution of ferrous sul- phate, forming a somewhat indefinite compound. At 8° the composition is claimed to be 3FeS04.2NO, but the variabiUty itt the composition of the substance as the temperature rises is such as to cast doubt upon its chemical identity. There has been in the past a tendency to caU such loose chemi- cal compounds as 3PeS04.2NO molecular compounds, indicating thereby that the combination between the molecules under consideration is of a weak nature and therefore easily shattered. A sUght rise in temperature in the above case completely expels the nitric oxide. Other examples of these molecular compounds are given by the salts containing water of crystaUisation, salts formed by the action of ammonia upon silver chloride, AgCl.SNHj and 2AgC1.3NH3, and by the solution of carbon monoxide in an ammoniacal solution of cuprous chloride. An oft quoted test for nitric acid is based upon the above reaction. A solution of the substance, supposed to be a nitrate, is placed in a test tube, and a strong solution of ferrous sulphate added. Concentrated sulphuric acid is then poured slowly down the side of the inclined test tube, and on account of its great density will form a layer at the bottom. If a nitrate is present, at the surface of contact nitric acid will be hberated and a brown ring formed. The brown colour is due to the formation of the above-mentioned molecular compound. The solubility of nitric oxide in water is sUght. 302 AN INORGANIC CHEMISTRY The composition of the gas is most easily determined by passing a measured volmne over red hot copper or iron wire. 2Cu + 2N0 -^ 2CuO + N^. The residual volume of nitrogen is found to be one half of the original volume of the nitric oxide. The density of the gas is found to be 29-96 (0=32). Thus One gram molecule of nitric oxide weighs . 29-96 gm. One half gram molecule of nitrogen weighs . 14-02 „ Hence weight of oxygen contained in one gram molecule of nitric oxide . . 15-94 ,, i.e. in one gram molecule there can be but one atom of oxygen (A.W.=16) and one atom of nitrogen (A.W. =14-01). The formula is NO. Nitrous Oxide and Hyponiteous Acid J. Priestley discovered this oxide in 1772 by reducing nitric oxide by means of moist iron filings. The most convenient method of preparation is by the decomposition of ammonium nitrate. Rapid decomposition sets in at about 200°. NHjNOa^N^O +2H2O. The decomposition should be carefully watched, for if the tem- perature rises too high, nitric oxide is also evolved, and the reaction may become so vigorous as to lead to an explosion. Properties. — Nitrous oxide is a colourless gas with a faint smell and a sweetish taste. It is appreciably soluble in water (100 c.c. of water at 10° dissolves 92 volumes at a pressure of 760mm.). The gas is, therefore, collected either over warm water or over mercury. Nitrous oxide is readily liquefied, a pressure of 30 atmospheres condensing it at 0° Nitrous oxide is a vigorous supporter of combustion. Phos- phorus, sulphur, a sphnter of brightly burning wood burn in it with considerable briUiance. In one respect, however, it differs from oxygen — it produces no brown fumes when mixed with nitric oxide, a valuable distinguishing test. Metals do not rust in it. Nitrous oxide, when inhaled, behaves as a mild anassthetic, and owing to the hysteria often produced during the inhalation, it has earned for itself the title " laughing gas." Similar means are adopted for determining its composition as for nitric oxide. OXIDES AND OXY-ACIDS OF NITROGEN 303 Hyponitrous Acid. Although nitrous oxide does not form hyponitrous acid on solution in water, there is little doubt that the oxide is the anhydride of the acid, for, on warming hyponitrous acid, nitrous oxide is evolved. (HN0),^H20+NA Hyponitrous acid is formed by the action of hydroxylamine on nitrous acid in aqueous solution. HNO2 +NH20H-j^ (HN0)2 +H2O. On treatment with sUver nitrate, the silver salt separates out . When this salt is shaken with an ethereal solution of hydrogen chloride, the reaction ( AgNO ), + 2HC1 -> 2 AgCl -^ + (HNO ) , occurs. After filtration and concentration the unstable hypo- nitrous acid is obtained. Prom molecular weight determina- tions in aqueous solutions there is httle doubt that the molecule is a double one. Recapitulation The study of the oxides and the oxy-acids of nitrogen has illustrated the statement aheady stressed in coimection with the corresponding compounds of chlorine and sulphur — as the propor- tion of oxygen in the acid-forming oxide increases, there is a steady increase in the strength of the corresponding acid and in its stability towards heat. Nitrous oxide. Nitric oxide. Nitrogen trioxide. Nitrogen peroxide. Nitrogen pentoxide. Formula Atomic ratio .... State of aggregation Acid N.O 2:1 Gas (HNO), NO 1 : 1 Gas N.O3 2:3 Gas HNO, NO, 1 :2 Liquid N,05 2:5 Liquid HNO3 Strength and stability of the acid increases. The Nitrogen Family. — Nitrogen is the first element in the fifth group of elements (see Periodic table). The elements in group 6 B, nitrogen, phosphorus, arsenic, antimony, and bismuth, 304 AN INORGANIC CHEMISTRY show as great a similarity to each other as do the elements of the halogen or sulphur groups. In the two succeeding chapters the remaining elements of the fifth group will be discussed, and it will be seen that the compounds of these elements bear a pronounced resemblance to the corresponding compounds of nitrogen, except in so far as these properties are modified by the increase in the atomic weight. Questions 1. The formula of nitric oxide is written NO. On what facts is this formiila baaed ? 2. What changes take place when chlorine is passed into a solution of ammonia ? How does this reaction throw light upon the composition of ammonia ? 3. Correlate the more important properties of the oxy-acids of nitrogen with the properties of their anliydridea. 4. Describe the more important commercial methods for manufacturing nitric acid. 5. What importance has nitrogen in nature ? Discuss the fixation of nitrogen. 6. ^^'hat is the action of nitric acid (dilute and concentrated) on the metals ? 7. What evidence exists for viewing nitrous acid both as an oxidising and as a reducing agent ? 8. Discuss the action of heat upon (a) ammonium nitrite, (6) ammonium nitrate, (c) potassium nitrate, (d) lead nitrate, (e) mercuric nitrate. 9. Construct equations for six reactions in which nitric acid functions as an oxidising agent. 1 0. Given nitric acid and any necessary apparatus and chemicals, show how you would prepare (a) nitrogen peroxide, (6) nitrogen trioxide, (c) nitric oxide, (d) nitrous oxide, (e) nitrogen, (/) ammonia. 11. Compare the action of nitrogen peroxide and chlorine peroxide upon sodium hydroxide. CHAPTER XXI PHOSPHORUS Historical. — Phosphorus was discovered by Brand in 1669 while distUHng a mixture of sand and urine. A century passed, however, before a satisfactory method of preparation was dis- covered. The discovery of phosphorus in bone ash by Gahn in 1771 led Scheele to use this source as a basis for preparing phosphorus in fairly large bulk. The method Scheele devised is practically that which was in use till comparatively recent times. Occurrence. — Phosphorus is found widely distributed in nature, generally in the form of phosphates (phosphorite). These are formed partly from the disintegration of rocks, and partly from animal agencies (bones, etc.). Animal excrement and urine contain appreciable quantities of phosphorus as phosphate. The droppings of sea birds on several of the Oceanic Islands, notably Nauru Island, constitute what is known as guano. This substance, the accumulation of centuries, rich as it is in nitrogenous and phosphatic plant foods, is a most valu- able fertiliser. Some of the more important mineral deposits are apatite, 3Ca3(P04)2.CaF2, a corresponding chlor- apatite, and wavelUte, a basic-aluminium phosphate. Manufacture. — Green bone contains about 58 per cent, of calcium phosphate, together with a large percentage of organic matter. On digestion with water under pressure, the gelatinous matter is dissolved out, and the residue consists partly of calcareous matter and partly of calcium phosphate. These degelatinised bones are then subjected to destructive distillation in order to recover the bone oil ; the residue is calcined in 305 X 306 AN INORGANIC CHEMISTRY order to burn off the carbonaceous matter and a moderately- pure form of calcium phosphate, known as bone-ash, remains behind. This is the chief source for the manufacture of phos- phorus. Two methods are in use for making phosphorus on a large scale : (a) The older retort method (essentially that of Scheele). (6) The modern electrical method. In the former method, the bone ash or a mmeral rich in calcium phosphate, is treated with sulphuric acid and the turbid liquor is then filtered. Ca3{P04)2 + 3H2S04^2H3P04+3CaS04. nJ. The solution of phosphoric acid is concentrated, mixed with sawdust or coke, and dried by heating in cast-iron vessels. This causes dehydration of the phosphoric acid with the forma- tion of meta-phosphorio acid {q.v.) H3PO4 — HaO-^HPOs. The dried mixture containing the meta-phosphoric acid is then transferred to retorts, where it is heated to a white heat. The neck of the retort dips under water in order to prevent the access of air. The phosphorus distils over and collects under water. 2HPO3 + 6C ^ H 2 4- 6C0 + 2P. The electrical method is much simpler than the chemical. A natural or artificial phosphate is mixed with sUica (sand) and coke, and the resulting mixture is heated to a very high temperature in an electric furnace (Mg. 93). The reaction may be looked upon as proceeding in two stages : Ca3(P04)2 + 3Si02-^3CaSiO, + P.,0, 3CaO,PA 3(CaO,Si02) and the phosphorus pentoxide is subsequently reduced by the carbon (coke). PA + 5C^2P-f 5C0. The calcium silicate forms a molten slag which is run off at intervals from the furnace. In such a furnace the part played by the electricity is merely to raise the temperature of the fur- nace. Nothing in the nature of electrolysis occurs. For opera- PHOSPHORUS 307 tions requiring a high temperature such furnaces are now in frequent use (see Calcium Carbide, Carborundum). Physical Properties. — Two well-defined allotropic modifica- tions of phosphorus are known, yellow (or white) phosphorus and red phosphorus. YeUow phosphorus is a transparent, wax-like colourless soUd. Its transparency is, however, of brief duration. This is due to a superficial change into the red variety. Yellow phosphorus melts at 44°. It is very soluble in carbon disulphide, alcohol, oU of cloves, ben- zene, etc., though practically insoluble in. water. The molecular weight of yellow phosphorus varies consider- ably with the temperature ; at 313° it is 128, at 1,700° 91-2. Since the atomic weight of phosphorus is 31, it follows that at compara- tively low temperatures the phosphorus molecule is tetratomic (P4), whilst at high temperatures more or less dissociation takes place, probably into Pj. YeUow phosphorus is very poison- ous. Constant exposure to the vapour of phosphorus causes necrosis, a decay of the bones of the nose and jaw. Red phosphorus is prepared by heating the yellow variety to about 250° in the absence of air. At 359° the transformation is almost explosive in its violence. The addition of a Uttle iodine facilitates the conversion, even at low temperatures. Red phosphorus is insoluble in all the solvents which dissolve the yeUow form. If heated under pressure, it melts at about 610°, forming a yellow liquid. The vapour pressure of red phosphorus is very much lower than that of yeUow phosphorus. The following experiment illustrates this. If a sealed tube contains a little yellow phosphorus at one end, and red phosphorus at the other, and if the end containing the red niodjfication is heated Fig. 93. 308 AN INORGANIC CHEMISTRY to 447°, the other end to 358°, it is found that the yellow phos- phorus distib to the hot end of the tube and condenses upon the red. This is due to the fact that at 358° the vapour pressure of the yellow modification is 1,696 mm., whilst at 447° the vapour pressure of the red is only 1,636 mm. Distillation therefore sets in from the place of higher to that of lower pressure. Chemical Properties. — The difference between the chemical properties of the two forms of phosphorus is quite as marked as between the physical. The yellow variety is characterised by extreme activity. With the halogens it unites spontaneously, with sulphur when shghtly warm, whilst oxygen causes slow oxidation even in the cold. It also reacts freely with sodium hydroxide. On exposing yellow phosphorus in the dark to the action of moist air, it glows with a pale greenish light, and at the same time emits fumes consisting of the oxides of phosphorus. Apparently the slow oxidation is the cause of the glowing. The appearance of this glow is very dependent upon the tempera- ture, and the concentration and purity of the oxygen. Thus, phosphorus does not appear luminous in pure oxygen below 15°. In the air the luminosity does not cease till below 0°. One of the products of slow oxidation is ozone, and it is noteworthy that the presence of traces of any gas which destroys ozone, e.g. hydrogen sulphide, ether, turpentine, completely stops the phosphorescence. Yellow phosphorus, although so reactive in moist air, does not react with perfectly dry oxygen. The strong reducing action of yeUow phosphorus is shown by the reduction of sulphuric acid to sulphur dioxide, whilst nitric acid is reduced to the lower oxides of nitrogen. The method of detecting phos- phorus in toxicological analysis is based upon the property of luminescence. The object containing phosphorus is placed in water, and the whole boiled. The issuing steam, charged with phosphorus vapour, glows brightly. The ignition temperature of yellow phosphorus is so low that it is apt to inflame even with the heat of the body. It must therefore be kept and handled under water. Red Phosphorus. — Red phosphorus is much less active than its allotropic modification. It is but feebly oxidised on exposure to the air, and does not ignite until a temperature above 200° has been reached, nor does it react with sodium hydroxide. The following table throws into prominence the difference in the properties of these two varieties of phosphorus. PHOSPHORUS 309 Property . Crystalline form Smell Action of air TABLE 32 Yellow phosphorus Rhombic-dodecahedron (regular) Garlic-like Oxidises with phosphor- Red phosphorus Rhombohedra (hexagonal ) Odourless No oxidation Melting point . Sp. gr. Physiological action Solubility . Ignition temperature Action upon NaOH Action upon halogens 44° Approximately 600° 1-84 2-10-2-14 Poisonous Non-poisonoua Soluble in carbon disulphide Insoluble in these and other organic solvents solvents 30° 240° Brisk action No action Spontaneous Action only upon heating The decreased activity of red phosphorus as compared with the yellow is no doubt closely connected with the smaller energy content of the red. Thus, when the yeUow variety passes into the red, there is a liberation of 4 Cals. per gram atom. The Transformation of Yellow into Red Phosphorus. In the discussion upon the conversion of one crystalUne form of sulphur into the other, emphasis was laid upon the fact that 10 1 1 60 10 50 / t / 40 ^ 1 ' ■5 30 1 / 10 (0 / ■'C ■t / / / / JO ^ y ^ y ^ — — — r^mf ter<. \tui -e. 500° 600° .wo° Mo° Pig. 94. — VaPOUB PaESStrEB-TEMPEEATUBE CUEVE FOE Phosphoeus. Upper Curve, Yellow Phosphorus ; Lower Curve, Red Phosphorus. 310 AN INORGANIC CHEMISTRY below the temperature of 96° the rhombic form alone is stable, whilst above that temperature the monoclinic variety alone is stable. In the case of phosphorus such a definite transition temperature is not known. A comparison of the sulphur vapour pressure-temperature curve (see p. 216) and the corresponding one for phosphorus (Fig. 94) explains this. At no point do the vapour pressure curves of the two forms of phosphorus cut, i.e. at no temperature are both forms stable. A piece of yellow phosphorus slowly but spontaneously changes into the red form, but the direct change of red into yellow is not known. In other words, the yellow phosphorus is always in a state of instability and tends to pass into the stable red modification. Transformation by Steps. — If red phosphorus is vaporised and the vapour condensed, the yeUow modification separates out first. This is an illustration of the principle first pointed out by Ostwald — ^that in aU reactions the most stable state is not reached at once, but a succession of intermediate and less stable states is first passed through. The condensation of phosphorus produces, first of all, unstable yellow phosphorus which slowly, but of its own will, passes into the stable red modification. Similarly sulphur, thrown out of precipitation, settles out in the form of hquid droplets which will not crystallise for weeks. An excellent example of this law is given by the study of the action of chlorine upon sodium hydroxide, the essential reactions being shown in the equations : 24NaOH + 12Cl2^ 12NaOCl + 12NaCl^^ 4NaC103 + 20NaCl 3NaC104+21NaCl Comniercial Uses of Phosphorus. — The match industry absorbs nearly the whole of the phosphorus manufactured. The paste with which matches are tipped consists of yellow phosphorus, an oxidising agent (manganese dioxide, potassium nitrate, or potassium chlorate), and a Httle glue and gum. The various ingredients are heated together in a closed vessel and then the spills are dipped into this paste. After drying they are coated with a layer of varnish. The poisonous nature of these matches has led to many attempts to replace yellow phosphorus by the non-poisonous variety. However, the non-inflamma- PHOSPHORUS 311 tility, and the increased cost of the red phosphorus, coupled with the fact that the " safety " match wiU only strike upon a prepared surface, have mitigated against the general adoption of this type of match. In the safety match the head consists of a mixture of potassium chlorate and sulphur. They are ignited upon a surface prepared from a mixture of red phosphorus, antimony sulphide and powdered glass, worked up into a paste with a little glue. The Molecular Weight of Phosphorus. — Determinations of the vapour density at high temperatures, coupled with freezing point measurements of the depression produced by the ad- dition of yellow phosphorus to a suitable solvent (see ch. xxvii.) indicate that the molecular weight of this modification of phos- phorus 18 four times the atomic weight, i.e. the phosphorus molecule is represented by P4. Some chemists consider that the: red modification is a polymer of the yeUow, but although the: possibility of this is not easily excluded, the evidence in support' of this view is not conclusive. From measurements of the density of the vapour of the red and the yellow varieties there is. no doubt the molecular state of the vapour is the same, whether generated by the vaporisation of the red or of the yellow, i.e. the formula of phosphorus vapour is represented by P4. The Hydrides os Phosphorus Emphasis has already been laid upon the similarity existing in the properties of corresponding compounds formed from a family of related elements, e.g. the hydrides of the halogens (cf . the hydrides of the halogens, p. 173). Phosphorus belongs to the nitrogen group of elements, and the family relationship is brought out not only by the similarity in chemical composition of the hydrides of nitrogen and phosphorus, but also to a certain extent in their properties. The hydrides of phosphorus are: — PH3 Phosphine (gaseous) cf. NH3 ammonia P2H4 Liquid hydrogen phosphide cf. N2H4 hydrazine PijHe Sohd hydrogen phosphide. Phosphine — Gaseous Hydrogen Phosphide. Phosphine, PH3, is made by boiling yeUow phosphorus with a strong solution of sodium hydroxide. The gas which is evolved 312 AN INORGANIC CHEMISTRY is spontaneously inflammable, so that as it escapes into the air, it catches fire. In order to avoid explosions, it is usual to dis- place the air from the generating flask and apparatus with hydrogen or coal gas before the experiment is begun. As each bubble of phosphine escapes from the outlet, a vortex ring consisting of white particles of phosphoric acid is formed. 3NaOH + 4P + SH^O -^ SNaH^POa + PH3 (iNa^O.SH^O.lP.O) Fig. 95. In this reaction sodium hypophosphite is also found, the phos- phorus being oxidised, whilst for the production of phosphine, reduction has occurred, i.e. auto-oxidation has taken place. In this equation it is seen that in the formation of phosphine there is a reduction to the extent of three units, and in the for- mation of the hypophosphite the oxygen valence has risen from to 1. It follows, therefore, that for every one molecule of phosphine produced by reduction, three molecules of the hypophosphite must be formed. With this ratio fixed the rest of the equation can be written down. A second method of preparing phosphine is by the hydrolysis •of a phosphide (q.v.) cf. CaaPs-f 6H2O- Mg3N, + 6H,0- -2PH3-l-3Ca(OH)2 ^2NH3+3Mg(OH), 'The gas in this case, too, is spontaneously inflammable. PHOSPHORUS 313 3. Phosphonium iodide (q.v.) breaks down under the action of potassium hydroxide PHJ + KOH -^ PH3 + KH- H^O cf . NH.Cl + KOH -> NH 3 + KCl + H ,0. Preparation of Pure Phosphine. — The phosphine prepared by either of the two methods first described possesses the pro- perty of spontaneous inflammability ; not so that obtained from phosphonium iodide. This difference in property has been traced to the presence in the gas of small quantities of Hquid phosphine (P2H1), a spontaneously inflammable substance. With the removal of this impurity departs the property of spontaneous inflammability from the phosphine. The liquid phosphine may be removed by the following means : 1. By passing the impure phosphine through a condenser cooled with liquid ammonia. The readily condensable liquid hydrogen phosphide is thereby completely removed. 2. By using an alcoholic solution of sodium hydroxide in the first method described. Liquid hydrogen phosphide is readily soluble in alcohol, and does not escape from the reaction vessel. Properties of Phosphine. — Gaseous phosphine is a colour- less, poisonous gas with an odour suggestive of decaying fish. It hquefies at about — 85°. The gas ignites at 100° and burns with a brightly luminous flame forming metaphosphoric acid and water. PH3+202-^HP03+H20. Unlike ammonia, it is only slightly soluble in water (11 vols, dissolve in 100 vols, of water at room temperature), and does not produce an alkaUne solution. When phosphine is heated, decomposition into the elements takes place, 4PH3^ii±P4-|-6H2. Phosphine ignites and burns when brought into chlorine. PH3 + 4Cl2^- PCI5 +3 HCl. Phosphine acts as a reducing agent when passed into a solu- tion of the salts of many of the heavy metals, forming phosphides. 3CUSO4 + 2PH3-> CU3P2 + 3H2SO4. Calcium phosphide is, however, obtained by heating metalUo calcium with phosphorus under naphtha. Phosphine shows a certain resemblance to ammonia in its reactions with the halogen hydrides. PHa+HI^zziPHiI. 314 AN mORGANIC CHEMISTRY Of these phosphonium compounds the chloride and bromide can best be prepared by the action of phosphine upon the halogen hydrides. These compounds are only stable at low tempera- tures and under high pressures, whereby the dissociation into phosphine and halogen hydride is prevented. The iodide is much more stable and can be readily obtained in the form of brilliant quadratic prisms. The iodide is easily obtained by the method indicated in the following equation : 13P + 91 + 24H2O — V 7PH4I + 2HI + 6H3PO4. Iodine is slowly added to a solution of phosphorus in carbon disulphide. The carbon disulphide is then distilled off and water allowed to drop upon the mixture. A brisk action at once sets in and the phosphonium iodide is volatilised. The chief interest in the phosphonium compounds hes in their resemblance to the ammonium compounds. Tbey bring out in a striking way the family relationship of the elements nitrogen and phosphorus. Liquid Hydrogen Phosphide, P2H4. This substance shows no resemblance to hydrazine, other than in the similarity of composition. It is most easily obtained by passing through a cooled spiral condenser the impin'e phosphine generated from calcium phosphide. The hquid is spontaneously inflammable (B.P. 57°). It is probably also formed when pure phosphine is led through nitric acid. At any rate the passage of pure phosphine through nitric acid restores the inflammability to the phosphine. Solid Hydrogen Phosphide. This is formed by the action of light or heat upon liquid hydrogen phosphide. 5P2H4 — >-6PH3+2P2H. It possesses no analogue among the nitrogen hydrides. From determinations of the freezing point depression produced by the solution of this substance in molten phosphorus, it has been concluded that the molecule has the formula PioH^. The Halides of Phosphoru.s There are two main classes of these compounds — those derived from trivalent phosphorus and those derived from pentavalent phosphorus. PHOSPHORUS 315 Trivalent halides : PF3 (gas), PCI3 (liquid), PBrj (Kquid), PI3 (solid), P^I^ (solid). Pentavalent halides : PF5 (gas), PCI5 (soKd), PBrj (soM). POF3 POCI3 POBr,. Oxyfluoride, Oxychloride, Oxybromide. In general, the halides of phosphorus are obtained by the direct action of the halogen upon phosphorus. They are com- paratively stable, but the presence of water at once leads to hydrolytic decomposition. Phosphorus trifluoride is produced by the action of arsenic fluoride upon phosphorus trichloride in the absence of moisture AsPa + PCl3-^. PF3 + ASCI3. It is a pungent smelUng gas which is readily oxidised to the penta-compound. PF3+F2^=±PI'6 PF3+Br2^=^PF3Br,. Phosphorus trichloride is easily obtained by the action of dry chlorine upon molten yeUow phosphorus or gently heated red phosphorus. It is a colourless hquid boiling at 75-9°. It fumes strongly in moist air and hydrolyses freely in water with the formation of phosphorous and hydrochloric acids. PCI3 + 3H0H^^ P(0H)3 + 3HC1. Phosphorus tribromide is formed by the action of bromine upon red phosphorus (B.P. 172-9°). With water it reacts thus, PBr3 + 3H0H^^P(0H)3 + 3HBr. Phosphorus triodide is produced by acting upon iodine with a solution of phosphorus in carbon disulphide. Its properties resemble those of the other trihahdes. Phosphorus pentafluoride is formed by the oxidation of the trifluoride by means of fluorine, and by the interaction of arsenic trifluoride with phosphorus pentachloride. SAsFs + SPClg^. 5ASCI3 + 3PF5. It is hydrolysed by water forming phosphoric and hydrofluoric acids. PF5 + 4H0H^- H3PO4 + 5HF. Phosphorus Pentachloride is produced by the action of an excess of chlorine upon phosphorus or by passing chlorine 316 AN INORGANIC CHEMISTRY through the trichloride. It is a solid which sublimes readily without melting. In the presence of small quantities of water it reacts thus : PCI5 + HOH -^ POCI3 + 2HC1, but with an excess of water complete hydrolysis is effected. PCI5 + 4H0H -^ H3PO, + 5HC1. Phosphorus pentachloride (and so, too, the trichloride) reacts freely with compounds containing the hydroxyl group, con- verting them into the corresponding chlor-compound. S0,(0H)2 + 2PCl5^SO,Cl2 + 2POCI3 + 2HC1. Sulphuric acid. Sulphiiryl chloride. The vapour density of phosphorus pentachloride falls with the temperature. Thus Cahours obtained the following results : Temperature . . 182° 200° 250° 300° Vapour density . 146-6 140-0 115-2 104-8 Per cent, dissociated 41-7 48-5 80-0 97-3 Theoretical vapour density of PCl5=208-5(O=32). At temperatures above 300° the density falls to a half of that given by the undissociated molecule. At the same time the characteristic colour of the chlorine appears. This is interpreted to mean that with rising temperature the equihbrium PCls^^PCl^ + CL is forced to the right. At 300° the dissociation is practically complete and the vapour density is half of that demanded for. the undissociated molecule. From the application of the Law of Mass Action to the above equation it is seen that (PcisKcy (PCI,) • Thus, if these three substances are in equilibrium in a sealed vessel, and chlorine is pumped in (i.e. the pressure of the chlorine is increased), the equilibrium will adjust itself by the conversion of some of the trichloride into the pentachloride. Wurtz took advantage of this in his determination of the molecular weight of phosphorus pentachloride ta the presence of phosphorus trichloride. The dissociation of the pentachloride was thus prevented, and he obtained the normal value 206-6 at 160°. PHOSPHORUS 317 Phosphorus pentabromide resembles the pentachloride both in the method of preparation and in properties. Of the oxyhalides of phosphorus phosphoryl chloride (phos- phorus oxychloride) ranks first in importance. It is obtained by the partial hydrolysis of the pentachloride (g'.i'.), as well as by the interaction of phosphorus pentachloride and the pentoxide, 3PCl5 + P205->5POCl3. The action of the pentachloride upon any compound containing the hydroxyl radicle, e.g. ethyl alcohol, also results in the formation of the oxychloride. C2H5OH + PCls-^ C2H5CI + POCI3 + HCl. Phosphorus oxychloride is a colourless liquid boiling at 107-2°. Under the action of water it slowly hydrolyses, forming hydro- chloric and phosphoric acids. POCI3 + 3H0H-> H3PO, + 3HC1. The Oxides and Oxyacids of Phosphorus The following table summarises the oxides and oxyacids of phosphorus : — Oxide. Phosphorus trioxide, P2O3 (Phosphorous anhydride) Phosphorus tetroxide, P2O4 Phosphorus pentoxide, PgOj (Phosphoric anhydride) Acid. Hypophosphorous acid, H3PO2 Phosphorous acid, H3PO3 Hypophosphoric acid, H2PO3 Orthophosphoric acid, H3PO1 Pyrophosphoric acid, H4P2O7 Metaphosphoric acid, HPO3 Of these the anhydrides, phosphorus trioxide and pentoxide are the most important. Phosphorus Pentoxide. This oxide is formed by burning phosphorus in an excess of oxygen or air. When pure, it is white and entirely free from smell, though commercial samples of the pentoxide generally have a faint garhc-like odour, due to the presence of traces of the trioxide. Phosphorus pentoxide absorbs water with avidity, and when brought into cold water forms metaphosphoric acid, P205+H20->2HP03, but in hot water orthophosphoric acid is produced, P205 + 3H20-^2H3P04. 318 AN INORGAMC CHEMISTRY Owing to the affinity "for water displayed by the pentoxide, this substance is the most valued of all reagents at the disposal of the chemist for drying gases. This property is often taken advantage of when it is desired to remove the elements of water from a compound, e.g. HNO3— H^O-^NA C2H5OH — HgO — >-C2H4 (ethyl alcohol) (ethylene) Phosphorus pentoxide has a molecular weight which indicates that its true formula is (P205)2, but for the sake of simplicity in writing equations the simpler formida is generally adopted. Phosphoric Acid. Phosphorus pentoxide resembles iodine heptoxide inasmuch as both these oxides -give rise to more than one acid which differ from each other only in the degree of hydration (see Per-iodic Acids, p. 187). From phosphorus pentoxide are derived three acids : Orthophosphoric acid, P2O5 +3H20->2H3P04 Pyrophosphoric acid, P2O5 +2H2O— >H4P20, Meta-phosphoric acid, P2O5 +H20^2HP03 Pyrophosphoric acid is not prepared by the direct action of water upon phosphorus pentoxide, but all the properties of this acid, as weU as its method of preparation, leave no doubt that it is derived from the oxide P2O5, so that its formula may be written 2H20,P205. The relation of the phosphoric acids to the corresponding acid of nitrogen is of interest. Nitrogen pentoxide. forms nitric acid, N20,+H20^2HN03. As indicated in the discussion of the per-iodic acids {q.v.), the true ortho-nitric acid would be represented thus : N205+5H20^2N(OH)5, but, owing to the instabUity of this acid, immediate decomposition to nitric acid (more correctly termed meta-nitric acid) ensues. N(OH)5^HN03+2H20. In the corresponding case of phosphoric acid we have P205+5H20^2P(OH)5 True ortho-phosphoric acid. PHOSPHORUS 319 This losies one molecule of water and passes into what is commonly termed ortho-phosphoric acid, P(OH)5->H3PO, + HA that is to say, the first member of the fifth (nitrogen) group appears less able to form a highly hydrated acid than does the next member. Later on, in dealing with the remaining elements of this group, it will be found that they offer a strict analogy with phosphorus (cf. arsenic, etc.). Moreover, the first element in other groups is also less able to form the highly hydrated acid than are the lower members of its group, (cf. HCIO4 and H5IO0, and H2CO3, H4Si04). Orthophosphoric Acid. — ^This acid is prepared either by the solution of phosphorus pentoxide (phosphoric anhydride) in warm water, or by the oxidation of phosphorus with nitric acid. Large quantities are prepared commercially by the action of sulphuric acid upon bone ash. Ca3(P04)2 + 3H2S04^3CaS04i+2H3POi. Orthophosphoric acid can be obtained in the form of transparent, six-sided crystals (rhombic) by concentration in vacuo, or by heating the acid to a temperature of 140° and cooling the concentrated liquor. It melts at 38-6°. Ortho-phosphoric acid is tribasic and gives rise to three series of salts. Primary salt, H3PO4 + NaOH = NaHjPOi + H^O Secondary salt, H3PO4 + 2NaOH= Na^HPOi + 2H2O Tertiary salt, H3P04+3NaOH= Na3P04 + 3HaO. The normal salt reacts alkaline to litmus, the secondary salt is neutral, whilst the primary salt is acid to litmus. This apparently conflicting behaviour is due to the hydrolysis of these salts in accordance with the equations — NajPOi + 3H2O ^=± 3NaOH + H3PO4 Na^HPO, + 2H2O ;:z^ 2NaOH + H3PO4 NaHs,P04 + H2O ^=± NaOH + H3PO4 In the first reaction the base and acid are produced in equi- valent quantity, in the second the ratio is 2:3, while in the third reaction the ratio has fallen to 1:3. Owing to the fact that sodium hydroxide is a more powerful base than phos- phoric acid is an acid, solutions of the normal salt react as if they have an excess of base, i.e. alkaline ; solutions of the 320 AN INORGANIC CHEMISTRY secondary salt are practically neutral, since the increased strength of the base just counterbalances its lower concentration as compared with the acid present, while in solutions of the primary salt, the excess of acid (3:1) more than counterbalances the greater strength of the base, and the solution reacts acid. Owing to the ready hydrolysis of the tertiary salts, these are stable only in the solid form, or m the presence of a considerable excess of sodium hydroxide, by which means the hydrolysis of the salt is effectively prevented. Na^POi +3H20;=± 3NaOH +H3PO4 i.e. the equilibrium expressed in this equation will be driven to the left. The insolubility of normal calcium phosphate (bone-ash) renders this substance of little direct value as a source of phosphates for plants. By treatment with sulphuric acid it is converted into the primary salt, Ca3(P04), +2H,SO,^CaH,(POj2 + 2CaS04 — a soluble salt, and therefore a salt capable of absorption by plants. This acid salt, known as " superphosphate," is on the market as a fertiliser. On heating phosphoric acid to about 250°, water is expelled, and pyro-phosphoric acid results. 2H3P04-H,0-^H4PA. Further heating leads to the formation of metaphosphoric acid. H4PA— H20->2HP03. By the hydration of meta- or pyro-phosphoric acid, the ortho-acid is formed. The constitution of orthophosphoric acid has been determined with tolerable certainty by its hydrolytic formation from phosphorus pentachloride and phosphorus oxychloride. CI. .01 01 >P^C1+H,0 ^^ 0=P/C1 + 2HC1. V \ci \ni or \ci \ci ,C1 HOH OH =P^:C1 + HOH -^ 0=P^0H+3HC1 ^iCl HOH \0H Pyrophosphoric Acid.— Pyrophosphoric acid is formed not only by heating orthophosphoric acid, but also from di-sodium PHOSPHORUS 321 hydrogen phosphate, which is first heated to convert it into sodium pyrophosphate, 2Na2HP04^NaiPA +H2O. The sodium pyrophosphate is then treated with lead acetate and the difficultly soluble lead pyrophosphate settles out. This is then decomposed by means of sulphuric acid. Pb^PjO, + 2H2SOi-> H.P^O, + 2PbS04 i The pyrophosphates are easily obtained from the mono- hydrogen phosphate or from tertiary phosphates containing the ammonium group. Thus magnesium ammonium phosphate reacts in the following way on heating, 2MgNH,PO,^2NH3 +Mg2PA +H3O. PjTophosphoric acid is a white crystalline solid, readily soluble in water. The conversion of its aqueous solution into ortho- phosphoric acid is considerably accelerated by boiling. The pyrophosphates are moderately stable salts. Although pyro- phosphoric acid contains four hydrogen atoms, and is therefore tetrabasic, only two series of salts have been prepared — e.g. Na^PA and NaaHaP^O,- Metaphosphoric Acid. — This form of phosphoric acid is obtained by the exposure of phosphorus pentoxide to the action of the aqueous vapour of the atmosphere, and also by the partial dehydration of ortho-phosphoric acid. It can be obtained by strongly heating pyrophosphoric acid. The sodium salt is obtained by heating sodium di-hydrogen phosphate or hydrogen ammonium sodium phosphate. NaH^POi^NaPOa +H2O HNaNHiPOi-^NaPOj -f NH3 + H^O. Metaphosphoric acid is a glassy solid, which fuses at 0°. It is soluble in water, slowly passing into the ortho-acid even in the cold. Although metaphosphoric acid is monobasic, a number of salts have been prepared from it which suggest that the molecule of metaphosphoric acid may exist in a polymerised state. Thus from — Mono-metaphosphoric acid we get . . NaPOs, sodiiun metaphosphate Di-metaphosphorio acid we get. . . NajPaOe, sodium di-metaphosphate 322 AN INORGANIC CHEMISTRY Tri-metaphosphoric acid we get. Tetra-metaphosphoric acid we get . Penta-phosphoric acid we get. Hexa-metaphosphoric acid we get . NajPsOg sodium tri-metaphosphate PbjPiOij. lead tetra-metaphosphate (NHjJsPsOis, ammonium penta-metaphosphate Na^PeOig, sodium hexa-metaphoaphate Tests for the Phosphoric Acids. The ortlio-, pyro-, and meta-phosphoric acids with their salts are easily distinguished from each other by their action upon sUver nitrate and albumen, as shown in the subjoined Table 33. TABLE 3.3 Keagent. Orthophosphoric Acid. Pyrophosphoric Acid, Mefcaphosphoric A' id. Silver nitrate Albumen Canary yellow precipitate of AgsPOj No action White crystalline precipitate of AgiP^O, -No action White gelatinous precipitate of AgP03 Coagulation All the phosphoric acids give a yeUow precipitate with a large excess of a hot solution of ammonium molybdate in nitric acid. Phosphobtjs Tkioxide and Phosphorous Acid (Phosphorous Anhydride) Phosphorus Trioxide (Phosphorous Anhydride) is pre- pared by burning phosphorus in a limited supply of oxygen. In order to obtain the trioxide free from the less volatile pentoxide, which is also formed during the oxidation, the products of the oxidation are drawn through a condenser, kept at 60°. The pentoxide separates out in the condenser, whilst the readily volatUe trioxide is caught in a receiver immersed in a freezing mixture. Phosphorus trioxide is a snow-white crystalline soUd, melting at 22-5°. It has a garUc-Uke odour and is highly poisonous. The oxide dissolves slowly in cold water, forming phosphorous acid, but in hot water decomposition occurs (see action of heat upon phosphorous acid). The trioxide oxidises readily to the pentoxide ; in chlorine the oxidation is so vigorous that the temperature of ignition is reached and the trioxide burns with PHOSPHORUS 323 the formation of the oxychloride. On heating phosphorus trioxide in a sealed tube to a temperature of 440°, auto -oxidation takes place with the production of red phosphorus and phos- phorus-tetroxide, 4P203->3PaOi+2P. Vapour density determinations, as well as determinations of the molecular weight of the trioxide when dissolved in benzene, show that the true formula is P4O6, though for reasons of convenience the simpler formula is still in general use. Phosphorous Acid is prepared either by the action of cold water upon the anhydride P2O3, or by the hydrolysis of phos- phorus trichloride, etc. PCI3 + 3H0H-> H3PO3 + 3HC1 PBrg + 3H0H^ H3PO3 + 3HBr (see Preparation of Hydrogen Chloride, p. 150). The easiest method is to pass chlorine through phosphorus melted under water, when the trichloride is instantly hydrolysed. The solution is then evaporated until its temperature has risen to 180° ; on coohng, the crystalline solid separates (M.P. 70-1°). Phosphorous acid is a vigorous reducing agent. It reduces copper sulphate to metallic copper, and the salts of the noble metals are also reduced to the metallic state. Oxygen converts it into orthophosphoric acid. When the acid itself is heated, it undergoes auto-oxidation, forming orthophosphoric acid and phosphine : 4H3PO3 -^ 3H3PO,+PH3 2(3H,O.P,03) KSH.O.P.Os) The valence of phosphorus towards oxygen in the trioxide is three, in the pentoxide it has risen to five, whilst in phosphine the valence has fallen to three towards hydrogen. Expressed mathematically, we may say that in passing from phosphorous acid to phosphoric acid, there has been an increase of valence of two, whilst in the reduction to phosphine there has been a fall of 3+3=6 units. The auto-oxidation must therefore lead to the formation of three times the quantity of the oxidation product as of the reduction product, i.e. three of H3PO1 to one of PH3. The instabihty of the aqueous solution of phosphorous acid as compared with the great stabihty of phosphoric acid affords 324 AN INORGANIC CHEMISTRY further support to the general statement that the lower oxy-acids, as a class, are less stable than the higher oxy-acids (of. nitrite and nitrate, chlorite and chlorate). The power to form acids of a lower degree of hydration is ■not restricted to phosphoric acid, for similar compounds have been isolated for phosphorous acid and its salts. Pyrophosphorous acid (H4P2O5), and its salts (NaaHaPjOs), and metaphosphorous acid HPO2, have been isolated. The constitution of phosphorous acid is still a matter of doubt. The preparation of this acid from phosphorus trichloride, CI P /ci + 3HOH->P(OH)3+3HCI, ^CI is generally held to support the view that the acid is tribasic OH and symmetrical, i.e. its formula is represented by Pe-OH ; ^OH so, also, the existence of a tri-sodium phosphite, NajPOg, supports this view, but there is not the slightest doubt that in the majority of its reactions phosphorous acid behaves as a dibasic acid. Most salts are of the type M2HPO3, while organic derivatives, such as 0=P^OR, where R denotes a radicle, supports the \0R unsymmetrical constitution 0=P^OH. The hypothesis which ^OH best reconciles these confhcting facts is that phosphorous acid exists in two tautomeric forms in equilibrium with each other : OH H P^OH ^z± =P AOH \0H ^OH Tribasic acid. Dibasic acid. Hypophosphoeous Acid Preparation. — The sodium salt of this acid is produced during the preparation of phosphine by the action of sodium hydroxide upon phosphorus {q.v.). If barium hydroxide is PHOSPHORUS 325 substituted for sodium hydroxide, the reaction proceeds similarly : 3Ba(OH)2 + 8P + GH^O-^ 2PH3 + 3Ba(H2P02)2. The resultant solution is treated with dilute sulphiu'ic acid till precipitation is complete, the barium sulphate removed by filtration and the solution concentrated until white crystals of hypophosphorous acid separate out. Properties. — Hypophosphorous acid is a strong reducing agent, reducing the salts of silver, gold, mercury, etc., to metal, and reducing copper sulphate to copper hydride (distinction from phosphorous acid). Zinc and hydrochloric acid reduce both hypophosphorous acid and phosphorous acid to phosphine. On heating hypophosphorous acid, auto-oxidation occurs with complete decomposition iato phosphine and orthophosphoric acid. 2H3P02^2PH3 + H3PO,. All the salts derived from hypophosphorous acid are mono- basic. If the pentavalence of phosphorus is assumed, the constitutional formula which represents this acid would be =p 4h. ^OH Phosphorus Tetroxide Phosphorus tetroxide is prepared by heating phosphorus trioxide in a sealed tube to a temperature of about 440°. 4P203-^2P+3P204. The tetroxide is very hygroscopic, and when dissolved in water undergoes auto-oxidation, giving rise to a mixture of phosphorous and phosphoric acids (cf. N2O4 and CIO 2). P2O4 + 3H,0^ H3PO3 + H3PO4. Hypophosphoric Acid Hypophosphoric acid, H2PO3, is formed along with phos- phorous and phosphoric acids, when phosphorus is allowed to undergo slow oxidation in a moist atmosphere. The acid is stable at ordinary temperatures and does not possess reducing properties, even hydrogen peroxide being without action upon 826 AN INORGANIC CHEMISTRY it. It is readily converted by acids into a mixture of phosphorous and phosphoric acids. Although phosphorus tetroxide does not give rise to hypo- phosphoric acid on treatment with water, the fact that both the acid and the oxide readily imdergo oxidation in aqueous solution in the same way is held to support the view that the tetroxide is the anhydride of hypophosphoric acid. The acid is dibasic and possesses the simple formula, H2PO3. Sulphides oi" PH0.gpH0Rus The most important sulphide of phosphorus is the penta- sulphide, P2S5. It is prepared by gently heating red phosphorus with the calculated weight of sulphur. The reaction proceeds vigorously, and on cooling the sohd pentasulphide is obtained. The residue is purified by distillation in a stream of carbon disulphide, when the distillate solidifies into a yellow crystalUne mass (m.p. 274°). Phosphorus pentasulphide reacts with water, forming orthophosphoric acid and hydrogen sulphide. P2S5 + SH^O-^ 2H3PO4 + 5H2S. Other sulphides are P4S3 and P4S5 ; while oxysulphides are also known. Comparison of the Phosphorus and Nitrogen Compounds A study of the properties of the oxy-compounds of these two elements leads to the conclusion that, widely though these elements appear to differ in properties, in reality certain funda- mental characteristics are found in all similar compounds, e.g. ammonia and phosphine, the auto-oxidation of nitrogen peroxide or tetroxide and of phosphorus tetroxide on solution in water, the instabihty of the lower acids (nitrous and phosphorous) compared with the greater stabihty of the higher acids. Even the comparatively great reactivity of " active " nitrogen and of yellow phosphorus, as compared with the inactivity of ordinary nitrogen and of red phosphorus, brings out the group analogy. The contrasts which stand out are : — 1. The greater oxidising power of nitric acid as compared with phosphoric acid. 2. The greater power of phosphorus to hold the hydroxyl group in combination as compared with nitrogen. 3. The much greater stabihty of the hydrides of nitrogen. PHOSPHORUS 327 QtTESTIONS 1. Compare and contrast the hydrides of phosphorus with those of nitrogen. 2. Compare the oxy-aoids of phosphorus and of nitrogen. 3. Discuss the effect of temperature and pressure upon the equilibrium PCI5 -!- PCI3 + Cl^. 4. Give an account of the three phosphoric acids. What tests may be used to identify these acids ? 5. Give a comparative account of the general chemistry of the elements, phosphorus and nitrogen. 6. Write equations illustrating auto-oxidation from the chemistry of phosphorus and nitrogen. 7. Describe how phosphorous acid may be prepared. What salts does phosphorous acid give rise to ? What is the action of heat upon such salts ? 8. Describe the preparation of phosphine, and compare its properties with those of ammonia. 9. What is an acid salt, basic salt, mixed salt ? How would you classify the following : — CaH^CPOi)^, MgNHiPOi, KH^POj, KH^POj ? 10. Give an accoont of the preparation and more important properties of the halides of phosphorus. 1 1 . What is the action of heat upon calcium superphosphate, magnesium ammonium phosphate, sodium phosphite, pyrophosphoric acid respectively? CHAPTER XXII ARSENIC, ANTIMONY, BISMUTH, VANADIUM, NIOBIUM, TANTALUM The Nitrogen Family. — The above elements, together with nitrogen and phosphorus, comprise the fifth group of the Periodic Table. The resemblance between arsenic, antimony and bismuth is especially well marked. The same gradation of properties is foimd in this group of elements as has already been stressed in the case of the sulphur group : 1. With increasing atomic weight of the element the hydride becomes less stable. 2. With increasiag atomic weight there is a steady decrease in the acid-forming power of corresponding oxides of the elements of this group, i.e. arsenic trioxide will be found to be a stronger acid-forming oxide than bismuth trioxide. The second of these statements forms the key-note to the study of the oxy-compounds of the elements of this group. Arsenic, Antimony, Bismuth Occurrence. — These metals are all found free in nature, and also in a state of combination. Arsenic occurs widely as sulphide in realgar, As 38 3 ; orpiment, AS2S3 ; arsenical pyrites or mispickel, FeAsS ; cobalt glance, CoAsS ; nickel glance, NiAsS ; as arsenide in kupfernickel, NiAs ; arsenical iron, FeAsj and Fe^Asa ; and as oxide in arsenoHte, Aafia. Antimony is generally found associated with arsenic. It occurs as antimony bloom, SbjOa ; antimony ochre, SbaOi ; stibnite, SbaSj ; and antimony blende or red antimony, Sb,03.2Sb,S3. Bismuth occurs rarely in combination with other elements, 328 ARSENIC, ANTIMONY, BISMUTH 329 though bismuth ochre, BiaOj, and bismuth glance, BigSg, are found. Methods of Preparation. — The oxides of arsenic, antimony and bismuth are all readily reduced to the metallic state by reduction with charcoal, but owing to the comparative scarcity of the oxides of these metals, recourse has generally to be made to the sulphides. In the case of antimony and bismuth, these are roasted in order to convert them into the oxides, and the metals are obtained by reduction with charcoal. Antimony sulphide is also reduced by means of scrap iron : SbaSj + SOa-^Sb^O^-H-SSOa Sb204+4C^2Sb+4CO (Bismuth sulphide forms the trioxide) when it is roasted. Large quantities of arsenic are obtained by heating mispickel. The arsenic sublimes and is recovered by condensation. FeAsS^FeS+As. Physical Properties. — All three elements are distinctly metallic in appearance. Their physical properties are summarised in Table 34. TABLE 34 Element. Colour. Crystal form. Speoiflo gravity. Melting point. Arsenic Antimony- Bismuth Steel grey Bluish white Greyish white with reddish tinge. Hexagonal rhombohedra 5-737 6-7-6-8 9-823 480° (under pressure) 630-6° 269° Both bismuth and antimony have the property of expanding on solidification. This property of antimony accounts for its use in type metal (75 per cent, lead, 5 per cent, tin, 20 per cent, antimony). Bismuth is also used extensively in fusible alloys of low melting point. Wood's Metal contains 1 part of tin (m.p. 232°) ; 1 part of cadmium (m.p. 321°) ; 2 parts of lead (m.p. 327°) ; and 4 parts of bismuth (m.p. 271 °) ; yet its melting point is only 60-5°. Such 330 AN INORGANIC CHEMISTRY alloys are extensively used for safety plugs in boilers, in fire alarms, automatic sprinklers, etc. Chemical Properties. — None of these elements is readily oxidised by exposure to the atmosphere, but the tendency towards oxidation increases with the atomic weight. On heating in oxygen or in air, vigorous combustion ensues, with the formation of the trioxide. Both antimony and arsenic unite with chlorine vigorously, bismuth combines with less energy. In general, sulphuric and hydrochloric acids exert very little solvent action, unless in very concentrated form, when small quantities of sulphate and chloride are formed. Nitric acid, both dilute and concentrated, attacks bismuth vigorously, forming bismuth nitrate, Bi(N03)3 ; concentrated nitric acid oxidises both antimony and arsenic producing oxides of antimony and arsenic acid, but the dilute acid has very Uttle action upon these two metals. Both arsenic and antimony resemble nitrogen and phosphorus in forming allotropic modifications. Thus, arsenic forms grey arsenic (hexagonal system), black arsenic (hexagonal system), and yellow arsenic (cubical system). The two latter are both obtained from grey arsenic. If this be heated in a stream of hydrogen, black gUttering crystals first separate out from the vapour phase, and further on a deposit of yellow crystals is obtained. The yellow variety is best prepared by distilling arsenic in a stream of carbon dioxide and suddenly cooling the vapour by passage through a condenser immersed in liquid air. The yellow variety is soluble in carbon disulphide (cf. yellow phosphorus). The separation of the imstable yellow and black varieties from the vapour of grey arsenic is an illustration of Ostwald's law of successive reactions (p. 310). Antimony not only forms the ordinary grey modification, but also a yellow, black and explosive modification. The yellow variety is obtained by allowing oxygen to act upon anti- mony hydride at —90°, whilst at — 40° these substances lead to the deposition of black antimony. Explosive antimony is obtained by the electrolysis of strong solutions of antimony chloride in hydrochloric acid. An amorphous powder is deposited on the platinum cathode, and on scratching this, a shght explosion occurs. At the moment of the explosion droplets of antimony chloride, which had been held in solution throughout ARSENIC, ANTIMONY, BISMUTH 331 the deposited antimony, are thrown out. The yellow, black and explosive modifications are all metastable and tend to pass into the stable grey form. The Hydrides of Arsenic and Antimony (Arsine, AsHs ; Stibine, SbHg) Just as a marked decrease in the stability of the hydrides was noted as we pass from nitrogen to phosphorus, so there is found a corresponding change in the stability of the hydride in passing from arsenic to antimony. Both the hydrides are obtained : 1. By the action of nascent hydrogen upon soluble salts of arsenic or antimony. 2. By the action of an acid upon a zinc alloy of the metal, e.g. ; ZnjAsa + 3H2SOi-> SZnSO^ + 2ASH3. The zinc alloy is formed by fusing the metals together in the proportions indicated in the above equation. Both these methods give a gas contaminated with a considerable quantity of hydrogen. The hydrogen is readily removed by condensing the hydride in a U-tube immersed in liquid air, when the readily condensable hydride liquefies. Properties. — At ordinary temperatures both hydrides are gaseous. They are intensely poisonous, and fairly unstable, for, when passed through a hot tube complete decomposition into the elements ensues. The gases burn with a blue flame, forming the oxide AsjOs or SbjOg. If a cold porcelain dish is depressed upon the flame, a metallic film is produced. If passed into a solution of silver nitrate, the two hydrides behave difl^erently. Arsine reduces silver nitrate to metallic silver and is itself oxidised to arsenious acid, which remains dissolved in the solution, AsHg + 6AgN03 + SH^O^ 6Ag + 6HNO3 + H3ASO3. On the other hand, stibine forms a precipitate of silver antimonide, the whole of the antimony being removed from the solution. SbHa + SAgNOs-^ AgaSb + 3HNO3. Marsh's Test. — ^The detection of smaU quantities of arsenic (and antimony) is best carried out as follows : The substance suspected of containing arsenic is introduced into a generating 332 AN INORGANIC CHEMISTRY flask with zinc and sulphuric acid (both of which must be free from arsenic). The issuing gases are led through a hard glass tube drawn out to a constriction in the middle. At this point a Bunsen flame is placed. If arsine is evolved, it will be decomposed as it passes through the constriction, and a black mirror is deposited in the cool portion of the tube. Stibine gives a similar deposit, but it is possible to distinguish between the two mirrors by the fact that the less volatile antimony is found nearer the hot portion of the tube. Moreover, the arsenic mirror is soluble in a solution of sodium hypochlorite, while the antimony one is not (Fig. 96). Fig. 96. By forming a series of standard mirrors from loiown amounts of arsenic, it is possible to estimate closely the quantity of arsenic present by comparing the mirror formed with that of the nearest standard mirror. The gradation in the properties of the hydrides of nitrogen, phosphorus, arsenic and antimony is seen not only ia their decreased stability towards heat, but also in the fact that neither arsine nor stibine shows any tendency to form salts corresponding to the ammonium and phosphonium salts ; they have lost the basic property possessed by both ammonia and phosphine. This point is further exemplified by the non-existence of a hydride of bismuth. The Halides of Arsenic, Antimony, and Bismuth The halides of these elements may be classed in two groups : 1. The comparatively stable tri-halides. 2. The unstable penta-hahdes. ARSENIC, ANTIMONY, BISMUTH 333 Tri-halides. Peiita-halidcs. ASF3 ASCI3 AsBra ASI3 AsFj AsCls Liquid Liquid Solid Solid Gas Liquid M.p.-8° M.p.-18° M.p. 31° M.p. 146° SbFj SbClj SbBrj Sblg SbFs SbCls Solid Solid Solid M.p. 73° M.p. 92° Solid M.p. 167° Liquid Liquid BiFj BiCla BiBrj Bilj Solid Solid Solid Solid — M.p. 230° M.p. 215° M.p. 439° The tri-halides may all be prepared by the du'ect combination of the elements, e.g. : 2As+3Brij^2AsBr3. The tri-fluorides and tri-chlorides are also readily obtained by the action of the halogen acid upon the oxide of the metal. The Uquor is concentrated and then subjected to distillation, when the volatile halide passes over in the pure state. Bia03+6HF^2BiF3 Sb203 + 6HCl->2SbCl3 SH^O F3H,0 The above table brings out in a very striking way the change in the physical properties of these hahdes brought about by changing not only the halogen but also the other element. The penta-halides are most readily prepared by passing the halogen through the tri-halide at a low temperature. The object of carrying out the operation at a low temperature is to retard the dissociation of the penta-halide — a dissociation which is strongly affected by a rise of temperature. SbClj^rzi SbCl3+ Cl^. There seems still a little doubt as to whether the substance formed by passing chlorine through arsenic trichloride at — 28° is reaUy AsClj, although the analysis agrees with that formula. It is maintained by some that merely a solution of chlorine in arsenic trichloride is obtained. The penta-fluorides are much more stable than the penta-chlorides, the latter readily dis- sociating into free chlorine and the tri-halide. Under the action of water, the tri-haUdes (excepting the 334 AN INORGANIC CHEMISTRY trifluoride of antimony) at once hydrolyse, with the precipitation of an oxide or basic salt. 2AsCl3+6HOH^2As(OH)3+6HCl-> AS2O3 >!' + SH^O+eHCl SbClj + 2HOH->Sb(OH)2Cl + 2HCl-^SbOCl i + H^O + 2HC1 BiCl, + 2H20^Bi(OH)2Cl + 2HCl^BiOCl ^^ + H^O + 2HC1 The hydrolysis of these haUdes is very similar to that of phosphorus trichloride, PCI3 + SH^O^ 3HC1 + P(0H)3, except that in the case of phosphorus trichloride the hydrolysis effected by an excess of water proceeds practically to a com- pletion, whilst the hydrolysis of the trihalides of arsenic, antimony and bismuth is strictly reversible, i.e. a definite equUibrium depending upon the concentrations of the reacting substances is established. This can readily be observed by noting the effect produced by pouring antimony or bismuth trichloride into water or into concentrated hydrochloric acid. In the first case, a copious precipitate is formed, in the second the presence of the acid throws back the equihbrium so effectively as to prevent the separation of any basic salt. Tiui Oxides and Hydeoxides of Arsenic, Antimony and Bismuth : Their Amphotekic Nature n arsenious oxide is shaken with water, a weak solution of arsenious hydroxide is obtained, As,03 + 3H,0 ^=± 2As(OH)3. If hydrochloric acid is added to this solution, the reaction As(0H)3 + 3HC1^ ASCI3 + SH^O sets in, and the removal of the arsenious hydroxide upsets the equihbrium expressed in the equation : AS2O3 + 3H2O ^=± 2As(OH)3. As a consequence, further arsenious oxide dissolves, in order to restore the equihbrium amount of arsenious hydroxide. In short, there is a steady solution of arsenious oxide with the formation of arsenious chloride (AsClj). The neutralisation of the arsenious hydroxide by the hydrochloric acid, ARSENIC, ANTIMONY, BISMUTH 335 As(0H)3 + BHCl-^ AsCla + 3H2O Base. 4- Acid. Salt. 4-Water. is essentially and fundamentally the same as the neutraHsation of sodium hydroxide by means of hydrochloric acid, NaOH + HCl^ NaCl + H^O Base. Acid. Salt. Water. except that in the one case, the reaction proceeds to completion, whilst in the other the reaction is reversible, and leads to an equihbrium. The solution of arsenious oxide or hydroxide in such an acid as hydrochloric acid, therefore, emphasises the capacity of this substance to behave as a base. On the other hand, if a solution of arsenious oxide (i.e. hydr- oxide) is treated with sodium hydroxide, it is found that the oxide (i.e. hydroxide) passes into solution, and from this solution one can obtain on evaporation a sodium salt — sodium arsenite. The reaction is represented in the equation : As(0H)3 +3NaOH^Na3As03 +3H2O. Acid. Base. Salt. Water. Here, again, no difference can be observed from the normal neutralisation of an acid by a base. HCl + NaOH -^ NaCl + 3^0 Acid. Base. Salt. Water. This capacity of arsenious hydroxide to behave either as a base or as an acid is of extreme importance. A similar behaviour is shown by antimony hydroxide, but not by bismuth hydroxide. Sb(0H)3 + 3HC1-^ SbCls + BH^O Base. Acid. Salt. Water. Sb(0H)3 +3NaOH-^Na3Sb03 +3H;0 Acid. Base. Salt. Water. Whether a hydroxide such as the above will behave as a weak base or as a weak acid, will be determined by the nature of the compound added. In the presence of a strong base the weakly acidic nature of the hydroxide will come to the fore, whilst in the presence of a strong acid the weakly basic nature of the hydroxide wUl cause it to function as a base. Oxides or hydroxides which have the power of behaving either as an acidic or basic oxide {or hydroxide), according to whether they 336 AN INORGANIC CHEMISTRY are reacting with a strong base or with a strong acid, are known as amphoteric oxides or hydroxides. The Position in tlie Periodic Table of the Elements forming Amphoteric Oxides. — In the discussion dealing with the Periodic Law it was emphasised that, as one moved from left to right in the table (i.e. from Group 1 to Group 8), the nature of the oxide undergoes a considerable change — the strongly basic oxides of Group 1 passing gradually into the acidic oxides of Group 8. The elements in the groups towards the centre of the table, Groups 3, 4, 5, are therefore more likely to show the property of amphoterism. The majority of the elements forming this type of oxide fall into these groups, though it may be mentioned that no group except Group is free from examples of amphoterism. Secondly, if one considers an element \^iiich forms several oxides, it may happen that the lowest oxide is basic, the highest acidic, and in such a ca.se the middle oxides are generally ampho- teric. Thus, chromium^ forms the oxides, CrO (basic), CrjOj (amphoteric), CrOj (acidic). Again, if one compares the corresponding oxides of a group of related elements, e.g. N2O3, P2O3, AsjO,, SbjOs, BijOa, owing to the fact that there is a general weakening in the acidic properties of the oxide as we pass down such a group, it is generally found that if the first oxide of the group is acidic, the last is basic, and some at least of the intervening oxides are amphoteric. Hence : N2O3, acidic ; P2O3, acidic ; AsjOj, amphoteric ; SbjOj, amphoteric ; Bi203, basic. On the other hand, owing to the increase in the acid properties of the oxide produced by the increased proportion of oxygen present in the pentoxides, aU these are more or less acidic in their behaviour : • N2O5, strongly acidic ; P2O5, strongly acidic ; AsjOj, weakly acidic ; Sb205, weakly acidic ; BijOj, very weakly acidic. A correct appreciation of the significance of the foregoing remarks with regard to the amphoteric behaviour of the trioxides of arsenic and antimony, coupled with the recognition of the periodic variation in the character: of the oxide produced by the increased atomic weight of the element, supplies the key-note to the chemistry of the fifth group of elements. AESENIC, ANTIMONY, BISMUTH 337 The Oxides and Oxy-Compounds of Arsenic, Antimony AND Bismuth The following table summarises compounds of these elements : Oxide. Suboxide Bismuth suboxide, BijOj Trioxide Arsenic trioxide, As^Os (arseniovis oxide) Antimony trioxide, Sb203 (antimonious oxide) Bismuth trioxide, BijOj Tetroxide Antimony tetroxide, SbjOj Pentoxide Arsenic pentoxide, AsjOs (arsenic oxide) Antimony pentoxide, SbjOg (antimonic oxide) the more important oxy- AOID. Ortho-arsenious acid, H3ASO3 Pyro-arseniouB acid, HjASjOs Meta-arsenious acid, HAsOj r t h o-antimonious acid, HaSbOa Pyro-ant imonious acid, HjSb.Oj Meta-ant imonious acid, HSbOa Hypo-a n t i m o n i c acid, H,Sb,03 Ortho-arsenic acid, H3ASO4 Pyro-arsenic acid, HjASjOj Meta-arsenic acid, HASO3 Ortho-a ntimonic acid, HaSbOi Pyro-a ntimonic acid, HiSbjO^ Meta- antimonic acid, HSbOa Bismuth pentoxide, BijOs Meta-bismuthic acid, HBiOs (bismuthic oxide) Arsenic Trioxide — Arsenious Acid Arsenic Trioxide. — Small quantities of arsenic trioxide occur naturally, but the main supply of this compound is obtained during the roasting of arsenical ores. Direct oxidation of the arsenic to the trioxide ensues, and this volatile oxide is separated from the flue gases by passage through a series of flues, when it separates as a flne powder. It is also obtained by the direct oxidation of arsenic by nitric acid. It is put on the market as " white arsenic " and occasionally as " arsenic." Properties. — Three well-deflned modifications occur: (a) The amorphous variety which first separates out when the vapour of arsenic trioxide is slowly condensed. This form is a colourless, vitreous, transparent mass which has an appreciable solubility in water. It can be preserved in sealed tubes, but on exposure to the atmosphere it changes slowly into weU-defined crystals of octahedral arsenic trioxide. The formation of this unstable variety is another fllustration of the separation of the Unstable phase first (Ostwald's Law of Successive Reactions). z 338 AN INORGANIC CHEMISTRY (&) The octahedral variety is formed as above, or by the rapid cooling of the vapour of arsenic trioxide, or by crystaUisation from an aqueous solution of the oxide. This form is much less soluble than the unstable amorphous variety. (c) Prismatic arsenic trioxide (monocUnic system) formed by crystallisation from a hot saturated solution. of arsenic trioxide in. potassium hydroxide. Arsenious oxide is a powerful poison. It is oxidised to arsenic acid by such oxidising agents as iodine, strong nitric acid, etc. Vapour density measurements show that the molecule is repre- sented by the formula, AsiOe. Arsenious Acid, H3ASO3. — The aqueous solution of arsenic trioxide has a feebly acid reaction, presumably through the formation of arsenious acid. Owing to the ease with which this acid is dehydrated, it has never been isolated in the pure state, 2H3ASO3— ^ AS2O3 + 3H2O, though many weU-defined ortho-arsenites (e.g. NajAsOj), pyro-arsenites (e.g. Ca^AsaOs), and meta-arsenites (e.g. KAsOj) are known. On account of the weak acidic nature of arsenious acid, the arseriites undergo very far-reaching hydrolysis, on solution in water, and consequently show a marked alkaline reaction. NajAsOj -f 3H2O ^^ As(0H)3 + SNaOH. The arsenites of the alkalies are easily soluble in ^later, but the arsenites of the heavy metals are all insoluble. Arsenious acid and its salts are mild reducing agents, and readily pass into arsenates. Na3As03 -f H2O + Iz-^ Na3As04 -f 2HI On heating they behave like the phosphites, i.e. they imdergo auto -oxidation and pass into free arsenic and arsenate. Besides the salts (arsenites) formed from arsenic trioxide when it behaves as an acid, a few salts exist wherein the arsenic trioxide undoubtedly exerts its basic function. Thus arsenic trichloride can be formed by the direct action of hydrochloric acid upon the trioxide. Antimony Trioxide. Antimonious Acid Antimony Trioxide is prepared by burning antimony in air, by heating it in a current of steam, or by the decomposition of AESBNIC, ANTIMONY, BISMUTH 339 antimony trichloride or sulphate by means of an alkaline carbon- ate or hydroxide. After treating the trichloride with hot water, SbCl, + 3H0H-^ Sb(0H)3 + 3HC1, the partially dehydrated hydroxide is thoroughly washed, and gently heated in order to form the oxide. Antimony trioxide is a white powder which subhmes at a high temperature. The vapour which has a density corresponding to the formula SbjOc, condenses in two distinct aUotropic modifications — prismatic needles of the rhombic system and regular octahedra of the regular system. The trioxide is practically insoluble in water, dilute nitric acid or sulphuric acid, but dissolves freely in hydrochloric acid, forming antimony trichloride. It is also soluble in aqueous solutions of acid potassium tartrate, forming potassium antimonyl tartrate (tartar emetic), 2[(SbO)K(C4Hi06)]H20. The amphoteric nature of antimony trioxide accounts for the dual type of salt derived from this oxide. Thus, when treated with concentrated acid, such salts as the nitrate, sulphate, chloride, acetate, are formed. These salts, formed as they are of a weak base and a strong acid, are strongly hydrolysed in aqueous solution, and throw down either a basic salt or the hydroxide itself. On the other hand, antimony trioxide, in its capacity to react as an acid, will also dissolve in strong bases, forming the ortho-antimonites. Sb203+ 6NaOH^:± 2Na3Sb03+ 3H2O. Similarly the meta-antimonites, e.g. NaSb02,3H20, have been prepared. This particular sodium salt is of interest owing to its insolubility. No pyro-antimonites have as yet been isolated. Ortho-antimonious acid is readily prepared by decomposing tartar emetic with sulphuric or nitric acid. The antimony sulphate (nitrate) at first formed immediately hydrolyses, and throws down a precipitate of white Sb(0H)3. By cautious dehydration the meta-antimonious acid can be obtained, HSbOj, Pyroantimonious acid has also been reported. Bismuth Tbioxide and Its Salts Bismuth trioxide is formed when the metal is burnt in air, or when the hydrated oxide, carbonate or nitrate is calcined. It is a creamy coloured powder, unacted upon by water. AU other oxides of bismuth are converted into the trioxidepn heating. 340 AN INORGANIC CHEMISTRY As is to be expected from the position of bismuth in the fifth group of elements, the oxide BiaOs has no acidic properties whatever. It is a weakly basic oxide, and, as such, dissolves in nearly all acids to form salts, e.g. nitrate Bi(N03)3, sulphate 212(804)3, chloride BiCla, etc. Owing to the weakness of bismuth oxide as a base, all the salts of bismuth undergo considerable hydrolysis on treatment with water, and almost complete precipitation of the bismuth in the form of a basic salt occurs. Absenic Pentoxidb. Aesenic Acid Arsenic Pentoxide is formed by the oxidation, of arsenious acid with strong nitric acid, with subsequent heating of the arsenic acid in order to dehydrate it : 2H3As04-^^As205 +3H2O. The pentoxide is freely soluble in water, forming ortho-arsenic acid. On heating the pentoxide, it decomposes thus : Ortho- arsenic Acid is formed either from arsenious acid as above, or by the solution of arsenic pentoxide in water. If the ortho-acid is heated to 140-180°, it pa^sses into the pyro-acid, H4AS2O7, which on heating above 200° forms meta-arsenic acid, HAsOj. The analogy with phosphorus is very complete up to this stage, but whilst meta-phosphoric acid cannot be dehydrated, the heating of meta-arsenic acid to a dull redness causes complete dehydration with the formation of the anhydride, AsoOs. Aqueous solutions of pyro- and meta-arsenic acid are unstable, and at once pass into ortho -arsenic acid. On the other hand, pyro- and meta-phosphoric acids are comparatively stable in aqueous solution, passing but slowly into the ortho-acid. Antimony Pentoxide — Antimonic Acid Antimony Pentoxide is prepared by the action of strong nitric acid upon antiinony, as weU as by the dehydration of antimonic acid at temperatures not exceeding 275°. It is a straw-coloured powder, almost insoluble in water. Although antimony pentoxide does not pass into antimonic acid, the inverse process takes place so freely as to leave little doubt that antimony pentoxide is the anhydride of antimonic acid. Antimonic acid, hke its analogues phosphoric and arsenic ARSENIC, ANTIMONY, BISMUTH 34i 4 > acids, occurs in three forms — ortho-antimonic acid, HaSbO pyro-antimonic acid, H^SbaO, ; and meta-antimonic acid, HSbOj. Ortho-antimonic acid is formed by oxidising antimony trichloride with nitric acid and diluting the solution with water. The white powder, 2H3Sb04,H20, is dehydrated by standing over sulphuric acid in a desiccator. On heating to 200° it passes into pyro-antimonic acid. This is also obtained by hydrolysing antimony pentachloride with hot water. 2SbCl5 + THOH-^H^Sb^O, + lOHCl. Meta-antimonic acid is obtained by heating the pyro-acid above 200°. Solutions of all three antimonic acids in hydrochloric acid liberate iodine from potassium iodide. Nearly all anti- monates are derived from meta-antimonic acid, though a few salts of ortho-antimonic acid (e.g., AlSbOi) and pyro-antimonic acid (e.g., KaHjSbaO,) are known. Bismuth Pbntoxide (Bismuthic Acid), Bismuth Tetboxide By the prolonged action of chlorine upon a solution of potassium hydroxide in which bismuth trioxide is suspended, undoubted oxidation of the bismuth occurs. By some it is claimed that potassium meta-bismuthate, KBiOs, separates out, whilst it has been recently asserted that a mixture of the pentoxide and a tetroxide results. The pentoxide is best prepared by the action of ammonium persulphate upon bismuth trioxide suspended in sodium hydroxide. Treatment with hot nitric acid removes the small quantities of the tetroxide formed, and a brown powder of the composition BijOsjEEjO is left. Although this formula corresponds to that of a possible bismuthic acid, HBiOg, this compound has no acidic properties, and appears to be a definite hydrated oxide. Bismuth tetroxide has also been separated from the product of the reaction between chlorine, bismuth trioxide and potassium hydroxide. It dissolves in hot nitric acid, liberating oxygen, whUe it oxidises rhanganous salts to permanganates in the cold. Other Oxides Antimony Tetroxide, SbaO,, is formed by heating the pentoxide above 200°, or by the prolonged heating of antimony or the trioxide. It is a white, non-volatile powder which is little acted upon by acids, except concentrated hydrochloric 342 AN INORGANIC CHEMISTRY acid, in which it is slightly soluble. By fusion with potassium hydroxide and other alkalis, compomids of the formula MaSbjOj are obtained, hence antimony tetroxide is to be considered the anhydride of hypo-antimonic acid, HaSbgOs. Bismuth suboxide, (BiO)^ where x is probably 2, is obtained by the decomposition of a basic oxalate of bismuth at 250-270" in the absence of oxygen. It is a black powder, which on exposure to air spontaneously passes into the yellow trioxide, and which, when treated with the vapour of methyl iodide at 260°, yields red crystals of bismuth subiodide (Bila). These crystals are readily volatilised, and form long needle-like crystals of the ortho-rhombic system. The existence of bismuth suboxide and subiodide is an interesting example of the tendency shown by the members occurring at the bottom of each group to jrield two or more types of salts of different degree of oxidation. The Sulphides of Aesenic, Antimony and Bismuth The following sulphides are known : Disulphide. As.S, red Trisulphides. AS2S3 yellow Sb2S3 orange BioSn brownish black Pentasulphides. AsjSs yellow SbjSs orange Of these realgar (AsjSa), orpiment (AS2S3), stibnite (SbaSj) and bismuth glance (BijSs) occur naturally. The trisulphides are obtained by passing hydrogen sulphide through a solution of the requisite salt dissolved in hydro- chloric acid : 2ASCI3 H-SHaS-^AsaSa ^ + 6Ha. Antimony pentasulphide is formed by the action of hydrogen sulphide upon a solution of antimony pentachloride acidified with tartaric acid. 2SbCl5 + 5H2S-^Sb2S5 ^1^ +10HC1. Arsenic pentasulphide, however, can be obtained by fusing together the elements in the correct proportions. It can best be prepared by the action of cold hydrochloric acid upon ammonium thioarsenate (q.v.). 2(NH,)3AsSi + 6HCl-> AsjSj ^^ +3H,S + 6NH4CI. ARSENIC, ANTIMONY, BISMUTH 343 The Thio- or Sulpho- Salts of Aesenic and Antimony Thio-arsenites and Thio-antimonites. — ^When antimony trisulphide is fused with potassium hydroxide or boiled with its aqueous solution, potassium thio-antimonite and potassium meta-antimonite are formed. 2Sb2S3 + 4K0H -^ BKSbSa + KSbOa + SHaO. Arsenic trisulphide behaves similarly. On treatment with dilute hydrochloric acid, the thio-antimonite reprecipitates the trisulphide : 2KSbS2 + 2HC1^ SbaSa + H^S + 2KC1. Bismuth sulphide, however, is unacted upon by an aqueous solution of potassium hydroxide or sulphide. The striking analogy between the sulphur and oxy-compounds of arsenic (and antimony) is emphasised in the following table, a complete series of thio-acids corresponding to arsenious and antimonious acids as well as to arsenic and antimonic acids existing. Oxide. Acid. Sulphide. Thio-acid. ASjjOs SbjOj As.Oa Sb^Os HjAaOa HaSbOj H3ASO4 H3Sb04 Aa,S3 SbA Aa.Ss SbjSs HjAaSj H3SbS3 H3ASS4 HaSbSi Besides the ortho-thioacids, there also exist pyro- and meta- acids and compounds of the -ous and -ic type. All such acids undergo decomposition on the addition of acid with the preci- pitation of the penta-sulphides. The reactions discussed in the preceding section are frequently made use of in the separation of the sulphides of arsenic, antimony and tin from other sulphides. When the sulphides of these elements are treated with ammonium sulphide, the soluble thio-salts of arsenic, antimony and tin are formed, and in this way a separation effected. In practice, one generally uses a yellow ammonium sulphide, i.e. ammonium sulphide containing in solution one or more atoms of sulphur in excess of that demanded by the formula (NHi)2S, and this dissolved sulphur 344 AN INORGANIC CHEMISTRY is capable of oxidising the thio-arsenites and antimonites into the thio-arsenates and -antimonates. Sb^Ss + SlNHJaS + 2S-^ 2(NH4)3SbS4. After filtering the solution of ammonium thio-arsenate and antimonate from the other sulphides, the solution of thio-salts is treated with hydrochloric acid, and the pentasulphides thrown down for further treatment. jq- Vanadixjm, Niobium, Tantalum j These elements belong to Group 5 A, and, as P such, are to be expected to show a general similarity to the elements of Group 5B. This is indicated in the group arrangement. j As Nb I Vanadium occurs in a few minerals (vana- I Sb dinite, pucherite, etc.), as weU as in certain iron — I ores. The ores are generally heated with sodium j — carbonate and the soluble vanadate so formed Ta 1 extracted with water. This solution is then treated Bi with ammonium chloride, and ammonium vana- date, NH4VO3, separates out. On heating this salt, the pentoxide is left. Vanadium forms a series of oxides similar to nitrogen : V,0, V,0, or VO, V,03, V,0, or V0„ VA- Of these, the lower oxides are feebly basic, the higher amphoteric. Numerous chlorides are formed — VCI2, VCI3, VCI4, VOCI3. The few salts formed from the oxide VO, e.g. VSO4, are readily oxidised by oxygen, etc. The salts of trivalent vanadium, e.g. V2(S04)3, are also relatively unstable, on account of the weakness of the base. Vanadium gives a well defined series of acids — ortho-vanadic acid, H3VO4, pjrro-vanadic acid, H4V2O7, and meta-vanadic acid, HVO3. During recent years vanadium has sprung into commercial importance, owing to its property of increasing the hardness and tensile strength of steel. Niobium and Tantalum occur in the minerals, tantalite and columbite. Niobium gives a mono-, di-, and pent-oxide, the first two of which are basic, the last is acidic and forms niobates. Tantalum gives a dioxide and a pentoxide from which the tantalates are derived. VANADIUM, NIOBIUM, TANTALUM 345 Questions 1. Describe and discuss the hydrides o£ the nitrogen group of elements. 2. Give a general account of the oxides and hydroxides of phosphorus, arsenic and antimony. 3. Discuss the importance of the thio-salts of arsenic and antimony in qualitative analysis. 4. What is an amphoteric hydroxide ? Show how the properties of the element, arsenic, and its compounds, are correlated with the properties of its hydroxides, 5. Using antimony and bismuth as examples, show how the increasing oxygen content of oxides modifies the properties of the oxide and of the saJts derived therefrom. 6. How do you account for the precipitation of the higher sulphide of arsenic by the action of hydrogen sulphide upon a solution of arsenic acid in hydrochloric acid. » 7. Compare the action of water upon the chlorides of nitrogen, phos- phorus, arsenic, antimony and bismuth. 8. How would you prepare arsine, arsenic acid and arsenious chloride from arsenic ? 9. What happens when arsenious oxide is treated (a) with hydrochloric acid, (6) with a solution of sodium hydroxide, (c) with a solution of ammon- ium hydroxide ? 10. Compare the physical and chemical properties of the elements of the nitrogen family. CHAI'TER XXIII CARBON AND ITS OXIDES General. — Carbon is the first element of the fourth group.. Although often classed with the elements of Group 45 (ger- manium, tin and lead), in many ways it displays an equaUy^ striking likeness to the elements of Group 4:A (titanium, zirconium,, cerium and thorium). The reason of this has been discussed in Chapter XVII on the Periodic Law (p. 261). Carbon and its analogues form oxides of the type MO 2, hydrides of the formula MH4 and chlorides of the formula MCI4. The position of carbon, midway between the basic element lithium, and the acidic fluorine, would lead one to expect that carbon dioxide would not have very pronounced acidic properties, and in point of fact, carbon dioxide is one of the weakest acidic oxides known. Occurrence. — Carbon occurs free in nature as the diamond, graphite and, in a less pure form, as coal. Moreover, considerable quantities of carbon in a state of combination are found in the earth's atmosphere as carbon dioxide, while enormous amounts exist as carbonates in beds of limestone and dolomite. Modifications. — Three well-defined modifications of carbon are known : (1) diamond ; (2) graphite ; (3) amorphous carbon, including charcoal of various sources, coal, coke. These varieties differ very much, especially in their physical properties. Their chemical identity is estabUshed by the fact that equal weights of each variety of carbon in the pure state, when burnt, produce exactly the same weight of carbon dioxide, though the fact that a different amount of heat energy is liberated by the combustion of the atomic quantity of amorphous carbon, diamond and graphite respectively is a conclusive proof that- these represent different allotropic modifications, each of these- 346 CARBON AM) ITS OXIDES 347 being associated •with its own definite quantity of heat per unit of weight : Diamond, 93-24 Gals. ; graphite, 93-36 Cals. ; charcoal, 96-98 Cals. At high temperatures (3000° and above) amor- phous carbon is the least stable of the three, while graphite appears the most stable. On the other hand, at low tempera- tures (200°-500°) it appears probable that the diamond is the most stable, so that at some temperature between 200°-3000° graphite and the diamond must possess a definite transition point. Diamonds are found in India, S. Africa, Brazil, Borneo and AustraUa. They are generally found in a more or less rounded condition, but they belong to the cubic system (regular). The diamond is brittle and very hard, the carbide of boron being the only substance which approaches it in hardness. It is a non-conductor of electricity and has a specific gravity of about 3-51 (compare graphite, 2-5 ; amorphous carbon, 1-45-1-70). Owing to its high refractive index, 3-5, the facets of the diamond sparkle with extreme brilliance. The diamond is insoluble in all liquids, and when burnt in an atmosphere of oxygen, a trace of ash is left and the whole of the carbon is converted into carbon dioxide. The high price of diamonds has been responsible for many attempts to produce them artificially. The most promising results have been achieved by Moissan. He dissolved carbon in molten iron at a temperature of 3000° attained by means of his electric arc furnace. He then cooled the mass to a duU red heat by plunging the molten mass into water, thereby causing a high pressure to be exerted on the stUl liquid interior. After dissolving the iron out of the kernel, there were found amongst the residue a few microscopic crystals which Moissan was able to show possessed the properties and the composition of the diamond. Graphite is widely distributed in nature, occmring in Cumber- land, Ceylon, California and Siberia. Traces have also been found in meteoric iron. Graphite is soft, almost soapy to the touch, hence its use as a lubricant. It is a good conductor both of heat and of electricity. A variety of graphite is used for electric light arcs, batteries, etc. It is made by first grinding coke with coal tar or molasses, moulding the paste to the desired shape and baking the product in the electric furnace. The non- 348 AN INORGANIC CHEMISTRY reactivity of graphite has opened up a big field for the utilisation of this substance as anode in many electrolytic processes. Another technical appUcation is in the lead pencil industry, the lead of such pencils consisting of powdered graphite, mixed with clay and sand. Graphite crystallises in the hexagonal system. Amorphous Carbon embraces all those varieties without a definite crystalline structm-e. The commonest forms are lamp- black, coke and charcoal. None of these are pure, though lamp- black approximates to 100 per cent, carbon in its composition. Lampblack is obtained by burning substances rich in carbon in a limited supply of air. The smoke, heavily charged with particles of carbon, is sent through chambers in which blankets are hung. On these the lampblack settles. The hydrocarbons generally present in lampblack may be got rid of by heatiag to a didl redness in a stream of chlorine. Lampblack is used in the manufacture of printer's ink, black varnish, etc. Coke is obtained by heating coal in a closed vessel. About 90 per cent, of the coke consists of carbon, the balance repre- senting the ash of the coal. Wood Charcoal is formed by the destructive distillation of wood. This is effected either in stacks or in retorts. In the older method, stacks of wood are covered with brushwood and finally with turf. The stack is lighted by lowering a candle down a central shaft. The requisite draught is obtained by adjusting small holes left in the turf. In about fifteen days the carbonisation is complete. This method entails a complete loss of aU the volatile products of the distillation. In the retort method, the volatile products are caught. The easily condensable substances, consisting of wood spirit (methyl alcohol), tar, acetone, and acetic acid are utilised in industry, the non-condensable gases, hydrogen, carbon monoxide, methane, etc., being used for heating purposes. Very pure carbon may be made by the destructive distillation of sugar, followed by the ignition of the carbonaceous product in a stream of chlorine. Properties of Carbon. — The amorphous form of carbon has the power of adsorbing on its surface considerable quantities of gases, dye-stuffs, etc. Thus Favre has estimated that 1 c.c. of wood charcoal will adsorb the following amounts of the various gases at ordinary temperatures : CARBON AND ITS OXIDES 349 Adsorbed gas. Quantity adsorbed. Ammonia ..... 178 c.c. Hydrogen chloride .... 166 c.c. Sulphur dioxide .... 165 c.c. Carbon dioxide . ... 97 c.c. whilst Dewar has shown that the amount adsorbed increases strongly as the temperature falls (15 c.c. of nitrogen are adsorbed by 1 c.c. of wood charcoal at 0° and 760 mm., 155 c.c. at — 185° and 760 mm.). Just as certain gases may readily be eliminated by the adsorb- ing action of charcoal, so also may dissolved or suspended matter be removed by this means. Thus, filtration through charcoal wUl remove fusel oil from whiskey, litmus from its solution, strychnine from its aqueous solutions and so on. Charcoal filters are often used for the purification of drinking water, the charcoal removing from the water deleterious organic matter ; but such filters may become a source of danger unless frequently cleansed by calcination to a red heat, as the charcoal soon acquires its maximum charge of adsorbed material. Another example of the commercial utilisation of the adsorbing powers of charcoal is the decolorising of sugar. For this purpose animal charcoal is generally used, This is obtained by heating bones, etc., in closed retorts (see p. 305). The bone charcoal consists of 10 per cent, carbon, 88 per cent, calcium phosphate, and 2 per cent, other inorganic substance, but whatever carbon is present, is so spread through the cellular structure of the calcium phosphate that it is in a highly reactive state. On boUing such bone charcoal with brown sugar, the colouring matter is completely removed. Another property which is possessed in a marked degree by amorphous charcoal is its reducing power. The metallic oxides, when heated strongly with finely divided charcoal are generally reduced to the metallic state : ZnO+C-^Zn + CO PeO +C-^Fe +C0. Advantage is frequently taken of this in the various metallurgical operations. Graphite, too, has considerable reducing powers, but owing to its greater cost is seldom used for this purpose. Carbon combines with hydrogen under the stimulus of the electric arc, forming acetylene (C2H2). With many metals it 350 AN" INORGANIC CHEMISTRY forms carbides, especially at the high temperature of the electric furnace, e.g. calcium carbide CaCj, iron carbide Ye^G. Coal. — Large suppUes of carbon occur naturally in the form of coal. The consensus of opinion, supported as it is by geo- logical and microscopical evidence, is that the various forms of coal represent different stages in the decomposition of vegetable material in the absence of air ; in a measure, therefore, coal is the result of the destructive distillation of wood, etc., carried out at great depths below the earth's smrface. The result of this decomposition is that there is a steady enrichment in the carbon content of the coal as the age of the coal increases, due to the progressive elimination of oxygen and hydrogen in the form of carbon dioxide and water. The accompanying table illustrates this point : TABLE 35 Percentage excluding Ash and Moisture ^ c. H. 0. N. Wood . . ' 45 6 48 1 Peat . . 60 6 32 o Brown coal . . 70 5 24 ] Bituminous coal. . . 85 5 9 1 Anthracite :^C;2H,0 + 2CO,. A great deal of the carbon dioxide compressed into cylinders and put on the market, is from this source. The actual breaking down of the sugar is effected by enzymes, active organic substances of a coUoidal nature contained in certain organisms and possessing the property of bringing about many chemical reactions. The enzyme known as pepsin, which is present in the human stomach, hydrolyses protein foodstuffs into products which are capable of assimilation. Yeast contains both zymase and sucrase. Of these, sucrase is able to break down cane-sugar into two simpler sugars : CijHjjOu +H2O — ^-CfiHiaOs -4-C6Hi20e Dextrose LseviUose. Invert sugar. Zymase then ferments, i.e. causes the decomposition of these sugars. The dextrose is decomposed much more rapidly than its isomer, laevulose. In the fermentation of wine the ferment- ation of the glucose contained in the grape juice is set up by an enzyme present upon grape-skins. Another enzyme is responsible for the conversion of alcohol into acetic acid {q.v.). It may be mentioned that aU the properties possessed by colloids are possessed by enzymes, for these are essentially colloidal in nature. Physical Properties. — Carbon dioxide is a colourless, odour- less gas possessing a feeble acid taste. It is a heavy gas, its density being about one and a half times that of air (gram molecular weight =44 gm.). By virtue of its great density it can be readily collected by the upward displacement of the air ; and it can be poured from one vessel to the other much as an ordinary liquid can. An instructive experiment illustrating this property is the following : A large bell-jar is filled with carbon dioxide, and in the same way as one draws water from a well, beakers fuU of carbon dioxide are withdrawn from the beU-jar and poured into a beaker counterpoised from the arm of a balance. Liquefied carbon dioxide boils at —79° ; and as its vapour pressure at ordinary temperatures is about 60 atmospheres, it can be kept in strong cylinders. If liquid carbon dioxide is allowed to CARBON AND ITS OXIDES 353 evaporate quickly, as by running it into a small canvas bag fixed over the nozzle of the cylinder, the absorption of heat is so great that the liquid soon freezes to a snow-white solid. A mixture of solid carbon dioxide and ether gives a tempera- ture of about —100°, while stUl lower temperatures may be attained by assisting the evaporation by reducing the pressure. At temperatures such as attained by the use of solid carbon dioxide and ether, rubber freezes to a hard, brittle solid, mercury becomes a solid, whilst many chemical reactions which are even explosive in nature at ordinary temperatures, have their velocity so lowered by the reduction in temperature as to be negligible. Thus sodium scarcely reacts with an acid at the temperature of —80°, though at 20° the reaction is explosive in violence. Carbon dioxide is fairly soluble in water, dissolving in its own volume at 15° and 760 mm. pressure. Under high pressures it shows a slight departiu-e from Henry's Law. Soda water is merely an aqueous solution of carbon dioxide saturated under a pressure of 6-10 atmospheres. Chemical Properties. — Generally speaking, carbon dioxide is a non-supporter of combustion, though very reactive metals such as potassium, magnesium, etc., are able to reduce it to carbon and are themselves converted into the oxide. C02+2Mg-^2MgO+C. Carbon dioxide is comparatively stable towards heat, but with rising temperature increasing quantities of carbon monoxide and oxygen are formed, i.e. the equUibrium 2C02^r3^2C0+02 is driven to the right. Temperature . . . 1027° 1170° 1292° Per cent, dissociation . 0-004 0-025 0-064 Carbon dioxide combines freely with strongly basic oxides to form carbonates. Na^O + CO2 -^ Na^COa Basic oxide. Acidic oxide. Carbonate. When passed into water, a smaU part of the dissolved carbon dioxide combines with the solvent to form the weak, unstable carbonic acid, H2CO3. H,0-fC0,^±H,C03. AA 354 AN INORGANIC CHEMISTRY Prolonged boiling can, however, effect the complete expulsion of the dissolved carbon dioxide from its aqueous solution. Carbonic acid has aU the properties of a weak acid. It turns blue htmus feebly red, conducts an electric current, and com- bines directly with the stronger bases to form carbonates. Carbonic acid, being dibasic, gives rise not only to the normal carbonates, but also to the acid carbonates or bicarbonates. NaOH H- H^COa^nziNaHCOa + H^O NaOH + NaHCOa ^r^Na^COj + H^O. Owing to the weakness of carbonic acid as an acid, solutions of the carbonates are appreciably hydrolysed ; normal carbonates which are soluble react alkahne to htmus, whilst the bicarbonates are neutral. The carbonates, with the exception of those of potassium, sodium and ammonium, are insoluble in water. The appreciable hydrolysis of sodium carbonate is often responsible for the precipitation of a basic carbonate on its addition to solutions containing salts of certain metals, e.g. magnesium. The normal carbonates of many metals such as lead can only be obtained by the use of the bicarbonate as the precipitant. Pb(N03), +2NaHC03->PbC03 i +2NaN03 + H2O +CO2. If carbon dioxide is passed through a saturated solution of sodium carbonate, crystalline sodium bicarbonate (sodium hydrogen carbonate) separates out, Na^COa + Kfi + C02->2NaHC03 i . Carbon dioxide has also the power of converting the carbonates of the alkahne earth metals (calcium, etc.) into soluble bicar- bonates. This is readily shoT\Ti by bubbhng carbon dioxide through a solution of calcium hydroxide, Ca(0H)2-f C02-^CaC03N|' +H2O. The precipitate of calcium carbonate which first separates out slowly passes back into solution owing to the formation of this soluble bicarbonate. CaCOj -f H2O -f CO^^^CaCOj + H2C03;=z±CaH2(C03)2. Definite equilibria, as indicated in the above equations, exist and the position of this equihbrium is much affected by temperature changes. Thus, on warming the solution, carbon dioxide escapes CARBON AND ITS OXIDES 355 and the bicarbonate breaks down. The transportation and deposition of calcium carbonate in nature is brought about by natural processes operating in accordance with the reversible equation : CaC03 + H20+C02^=±CaH2(C03)2. The Assimiiation of Carbon Dioxide. The respiration of carbon dioxide by the animal kingdom, coupled with the ejection of appreciable quantities of this sub- stance from volcanoes, vent-holes, etc., would lead to an accumu- lation of this product in the atmosphere, were it not that plants assimilate carbon dioxide, extracting from it the carbon and giving off the oxygen. This maintains the balance in nature. The absorption of carbon dioxide by plants and its conversion into starch, sugar, cellulose, etc., is attended by a very consider- able absorption of heat. This energy is supplied by the sun. The conversion of the carbon dioxide into plant products is mainly carried out in the green chlorophyll cells under the stimulating action of the red and yellow hght rays. Thus, a plant, prevented by means of a shade made of blue glass from receiving the red and yellow Hght, is quite unable to assimilate carbon dioxide. It is interesting to note that in other chemical reactions blue light, i.e. light of short wave length, has a more marked photo-chemical action than has light of another colour. The photo-activity of the blue rays is especially noticed in the production of a latent picture upon a photographical plate. Besides the process of assimilation which involves the absorp- tion of carbon dioxide and the eUmination of oxygen on the part of the plant, ordinary respiration similar to that in animals occurs. Both by day and by night plants breathe in oxygen and breathe out carbon dioxide. This process is not in any sense a photo-chemical one. The amount of carbon dioxide exhaled by a plant during the period of twenty -four hours is, however, very much less than that assimilated under the action of the sun. Composition. — Oxygen, carefully freed from carbon dioxide and moisture, is passed over a weighed quantity of carbon. The resulting gases are led through hot copper oxide in order to oxidise aU traces of the lower oxide, CO, to the dioxide. This is then absorbed in weighed potash bulbs. From the weight of 356 AN INORGANIC CHEMISTRY carbon used and the weight of carbon dioxide found, the ratio of carbon to oxygen present in the dioxide is obtained. This gives the value 12-005 to 32, i.e. the formula is (COa)^. The density of carbon dioxide is approximately 44 (0=32), whence x=l ; in other words, the formula is CO 2. By burning a piece of carbon in the apparatus of figure 97 it can be shown that, when carbon bums, "^ If the volume of carbon dioxide is equal to ■ — the volume of oxygen required for its pro- duction. Hence ^^ C + 02->C02 1 vol. 1 vol. Carbon Monoxide Preparation. — Carbon monoxide is often prepared by the reduction of carbon dioxide with carbon, zinc, iron, etc. C02 + C^±2CO COj+Zn^ZnO CO. The reduction of glowing carbon can be easily demonstrated by a passing stream of carbon dioxide through a hard glass tube packed with charcoal (Fig. 98). The issu- ing gas is led through a wash-bottle con- taining a solution of potassium hydroxide to remove the unreacted carbon dioxide. As the temperature of the charcoal in- creases, there is a corresponding increase in the amount of carbon monoxide formed. This experiment illustrates the principle underlying the formation of Producer Gas iq.v.). During many metallurgical processes wherein carbon is used as a reducing agent, carbon monoxide is liberated as a by- product, Fe A -f 3C-> 2Fe -|- SCO. In the laboratory this gas is prepared by the dehydrating action of sulphuric acid upon oxahc or formic acid. Fig. 97. CARBON AND ITS OXIDES 357 H2C204-H20->CO+C02 Oxalic acid. H.COOH— H^O^CO. Formic acid. When the gas is prepared from oxahc acid, it is usual to wash the gases through a concentrated solution of sodium hydroxide in order to remove the carbon dioxide. A steady stream of carbon monoxide may be obtained by allowing strong formic acid to drop from a tap funnel upon hot, concentrated sulphuric acid. Another method of obtaining a regular stream of this gas is by gently heating a mixture of ten Dilute solution of sodium carbonate. Fig. 98. parts of sulphuric acid by weight with one part of potassium ferrocyanide. The reaction ceases at once on coohng the flask. K,Fe(CN)o + 6H2SO4 + BH^O-^ 2K2SO4 + FeSO^ + 3(NH J2SO4 + 6C0. Properties. — Carbon monoxide is a colourless, odourless, tasteless gas, slightly soluble in water. It is very difficult to liquefy. (B.P., — 190°.) Physiologically, it is a powerful poison. When inhaled, it unites with the hsemoglobin of the blood, producing a bright red compound, carbonyl hsemoglobin. When converted into this compound, the hsemoglobin is perma- mently removed from participating in the oxygen-carbon dioxide cycle in the human system, and if sufficient of the hsemoglobin is removed in this way, the sufferer dies. The more important properties of carbon monoxide are associated with the unsaturated nature of the carbon atom, i.e. the divalent carbon atom tends to pass into the stable tetravalent state. 358 AN INORGAOTC CHEMISTRY CI C=0 + Cl2 -> C^ \ci Carbonyl chloride. ZnO+CO->Zn + C02. Indeed, at temperatures below 867°, it is a more active reduc- ing agent than hydrogen itself. With a few elements, notably nickel and iron, carbon monoxide combines to form the so-caUed carbonyls. When it is passed over hot finely-divided nickel, direct combination ensues. Ni+4C0^z:±Ni(C0)4. From the issuing gas there can be separated, nickel carbonyl, a mobUe, highly refracting liquid (sp. gr. 1-356), boiling at 43° under a pressure of 751 mm. On heating the vapour to a red heat, decomposition into nickel and carbon monoxide takes place. Advantage is taken of this reaction in the Mond process for purifying nickel. Carbon monoxide can be absorbed by passing the gas into a solution of cuprous chloride dissolved in hydrochloric acid or ammonium chloride, when loose chemical compounds between the carbon monoxide and the cuprous chloride are formed. Carbon monoxide is slowly absorbed by soUd potassium hydroxide at a temperature of 100°, forming potassium formate. This fact, coupled with the ready decomposition of formic acid into carbon monoxide and water, suggests that these two com- pounds are related to each other much in the same way that sulphur trioxide and sulphuric acid are, the true dibasic ortho- acid at once changing into a tautomeric monobasic acid. OH C=0 + H20^C(( -> H-C/' ^OH ^OH Unstable. Formic acid. A yet lower oxide of carbon, known as carbon suboxide, C3O2, has been prepared by distilling malonic acid with phosphorus pentoxide under reduced pressure. Water is thereby abstracted from the compound, and the suboxide distils over as a colourless hquid (b.p. 7°) which yields malonic acid again on addition of water . CHaCCOOH)^— 2H20=C302. The suboxide is combustible, and burns with a smoky flame. CARBON AND ITS OXIDES 359 If kept many hours, it polymerises into a red solid, which is soluble in water. Producer Gas. — The reaction expressed in the equation C0,+C^=±2C0 has long had a commercial application in the so-called Producer Gas. With increasing temperature there is a steady increase in the amount of carbon monoxide in the equihbrium mixture. TABLE 36 Ver cent by Temperature. \'olume. ! CO 1 CO 2 500° 5 95 700° 58 42 900° 96 •.'i :s-5 1000° . . . 99-9 01 Owing to the great amount of heat evolved when one gram molecule of the monoxide is oxidised compared with that generated when the carbon is oxidised to the monoxide, C + 0-> CO +29650 cals. CO + O->CO2 + 68000 cals. it is often considered preferable to gasify the fuel in specially designed plants and to carry the producer gas in pipes to wher- ever it is desired to convert its latent chemical energy into mechanical energy. The final combustion is carried out in suitably designed gas engines. It is true that in the cycle of operations a considerable proportion of the store of energy in the sohd fuel is lost to the manufacturer, but in many cases the advantages more than outweigh the disadvantages. In a well- designed producer plant practically the whole of the carbon is converted into carbon monoxide, provided the temperature is maintained above 1000°. Seeing that the oxygen required for the combustion of the carbon is derived from the air, the final product must contain a large proportion of nitrogen. An ideal producer gas consists of 34-7 per cent, of carbon monoxide and 65-3 per cent, of nitrogen by volume. Water Gas. — ^Another form of power gas in which carbon monoxide plays an important part is the so-called Water Gas. 360 AN INORGANIC CHEMISTRY This is made by injecting steam into incandescent carbon, when the equilibrium C+H^O ;=^C0+H2 is set up. This process, unlike that in the producer gas, is attended by a very considerable heat absorption, so that the temperature of the furnace falls and the equiUbrium 2C0^=±C+C0, which must also exist will shift over to the side of the worthless carbon dioxide (cf. preceding section). Under working con- ditions the fuel is raised to incandescence by a blast of air, and this is followed by a jet of steam. During the steam injection water gas is being generated, at other times a damper deflects the stream of furnace gases. The accompanying table by Bunte shows the effect of temperature upon the composition and there- fore upon the calorific value of the water gas. TABLE 37 Temperature. Per cent. strain decomposed. Per cent, analysis of gas. H. CO. CO.^. 674° 838° 1010° . 1125° . . . . 8-8 41 94 99-4 65-2 61-9 48-8 50-9 4-9 151 49 7 48-5 29-8 22 9 1-5 0-6 In many modern plants it is the custom to inject both steam and air together, so that the heat generated in the pro- ducer reaction, C+COg ;; — ^ 2C'0, may counterbalance the heat absorption in the water gas reaction, HjO+C ^r:±C0-|-H2. With such an arrangement the process becomes a continuous one. The absence of luminosity from the water-gas flame prevents the use of this gas except in burners prtnided with a Welsbach mantle, etc. In order to overcome this defect as an iUuminant, a greater illuminating power is conferred upon the water gas by sending it through a hot tower containing brickwork, and into which a spray of oil is blown. As the water gas and oil find their way down the tower, the oil cracks, i.e. the complex hydro- carbons are decomposed into simpler ones. The enriched water CARBON AND ITS OXIDES 361 gas makes a satisfactory illuminant, except that it is excessively poisonous owing to the high content of carbon monoxide. In the manufacture of Oil Gas, e.g. Pintsch Oil Gas, the oil is brought into hot retorts, where decomposition into rich illum- inants is effected. The great advantage of this gas Ues not only in its great illuminating power, but also in its being much more compressible, hence its value for hghting trains, steamers, etc. Carbonyl Chloride and Urea Carbonyl Chloride, COCI2. — This substance, known as phosgene, is formed by the direct oxidation of carbon monoxide by chlorine under the photochemical action of the sun. The combination is also effected by using animal charcoal as the catalyst. It is a liquid boiUng at 8°. It is soluble in benzene and water, but it hydrolyses rapidly in the latter solvent. COCI2 + 2H2O ^=i± H2CO3 + 2HC1. It is on account of this reaction that the constitution of carbonic acid is generally represented by the formula OH 0=C< ^OH Urea. — When treated with ammonia, carbonyl chloride forms urea. ^Cl +HNH, NH, .. -> o=c/ o =c/ Cl +HNH2 ^NH^ +2HC1 2HC1 + 2NH3-^2NH4C1. It is a white, crystalUne solid soluble in alcohol, and therefore easily separated from the ammonium chloride. Urea affords an interesting example of isomerism, for it can be formed from ammonium cyanate by mere heating. NHj.CNO-^O -C— NH, ^NH, The urea and ammonium cyanate have the same chemical com- position but each has its own specific properties, i.e. these substances are isomeric. Until Wohler (1828) discovered the above method of synthe- sising urea, it was always held that a true " organic " substance such as urea could only be formed by the agency of a Uving 362 AN INORGANIC CHEMISTRY organism. From the date of Wohler's discovery, organic chemistry has become less a study of compounds produced by living organisms than a branch of inorganic chemistry, viz. that section deaUng with the compounds of carbon. Sulphides or Carbon Besides the oxy-sulphide of carbon, COS (carbonyl sulphide) (cf. C0+0->C02 and CO+S->COS), there exists a well- known disulphide, CSj. This is formed by the direct union of carbon and sulphur, carried out in specially designed furnaces. The vaporised carbon disulphide is condensed and purified by redistillation. It is a colourless, highly refractive liquid boiling at 46^. It is very inflammable, forming carbon dioxide and sulphur dioxide. It is often used in industries as a solvent for sulphur, phosphorus, rubber, etc. The pure product has a rather pleasant smeU, but the commercial preparations are very offensive owing to the presence of impurities. A compound, thiocarbonic acid, H2CS3, which bears the same relationship to carbon disulphide as carbonic acid does to carbon dioxide, has been prepared. The salts are much more stable than the acid. CaS+CSa-^CaCSa. Questions 1. What is meant by isomerism and polym/irphiam ? Give examples. 2. By what experiments can you prove that carbon monoxide contains half its own volume of oxygen 1 3. What volume changes occur when carbon disulphide is burnt in oxygen ? 4. How is carbonyl chloride made ? What is its action upon ammonia and water ? How does the latter reaction throw light upon the constitu- tion of carbonic acid ? 5. Ten gm. of carbon are burnt in an excess of oxygen and the resultant gas led through a solution containing 60 gm. of sodium hydroxide. What percentage of the sodium hydroxide will be neutralised 1 6. Given a mixture of moist nitrogen, oxygen, carbon monoxide and carbon dioxide, how would you proceed to determine its percentage composition by volume ? 7. How do you accoimt for the evolution of carbon dioxide when hydro- chloric acid is added to a carbonate ? 8. How much sulphur would be required for the production of 200 kUos of 96 per cent, sulphuric acid ? What volume of air would be required to efiect the combustion ? 9. If 10 litres of carbon monoxide at 15° C. and 750 mm. are burnt to carbon dioxide, what volume of the dioxide, measiu-ed at 0° C. and 760 mm. pressure, will be produced ? 10. Discuss the commercial importance and manufacture of producer gas. CHAPTER XXIV TYPICAL CARBON COMPOUNDS Caebon differs from all other elements in so far as the atoms of this element exhibit an extraordinary tendency to combine with one another. This is exempHfied by the fact that nearly three hundred hydrocarbons (compounds of hydrogen and car- bon) have been isolated. These fall into different series or groups. The first of these groups is known as the paraffin or saturated hydrocarbons. The name " paraffin " indicates the most important property of these compounds, their non-reactivity {parum, little ; .affinis, affinity). The Paraffin Series of Hydrocarbons. — The saturated or " paraffin " series of hydrocarbons are derived from methane, CH4, the parent member of the group. If the assumption is made that the carbon atom is always tetravalent, and that carbon atoms possess the power of combining with each other, forming a chain, e.g. I Q Q Q Q Q I the more complicated members of the group can be built up from the simpler by successive replacements of a hydrogen atom by the CH3 group or methyl radicle. Hence, from methane H H H I ' II H— C-H we derive ethane H— C— C— H, I I I H H H 363 364 AN INORGANIC CHEMISTRY from ethane, we obtain propane and so on. H H H I I i H— C— C— C— H I I I H H H PKOPANE An examination of the formula of propane reveals the fact that the two hydrogen atoms attached to the central carbon atom stand in a different relationship to the molecule from what the other hydrogen atoms do, i.e. they are differently situated so far as the configuration of the compound is concerned. Hence, two differently constituted isomers will be derived, according to whether a hydrogen atom attached to an end carbon atom is substituted by the methyl group, or whether one of the internally situated hydrogen atoms is so substituted. H H H I I I From propane, H — C — C — C — H we obtain normal butane ! 1 I H H H if any one of the end carbon atoms is substituted ; whilst the second method of substitution leads to the formation of the isomer, iso-butane. H H H H H H H I I 1 H— C — C — C— H H— C— C— C— C— H, I j I i 1 I I H H— C— H H H H H H I H NOKMAL BFTANE ISO-BUTANE The number of such isomers increases rapidly with the number of carbon atoms in the compound. It is interesting to note that in no single case have more isomers been discovered than pre- dicted by the theory of the tetravalence of the carbon atom, and TYPICAL CARBON COMPOUNDS 365 of the chain Unking of these atoms. The following is a list of the more important saturated hydrocarbons present in the paraffin series of homologues or related compounds : B.P. B.P. General for- Pentane . . CgHi2 normal 36° mvila . ^oHgD f 2 iso- 28° Methane CH4 -164° neo- 10° Ethane . CA —93° Hexane . CgHi4 71° Propane C3H8 —45° Heptane • C,Hi5 99° Butane . C4H10 normal 1° and so on. iso- —17° The great natural source of the paraf&n hydrocarbon is petro- leum oil. In the crude state this is a thick viscous oil of a greenish brown colour obtained by sinking a bore into the oil-bearing strata. At first the oil rushes forth under considerable pressure, but, generally speaking, it is pumped to the surface, and conveyed in pipes to the reservoirs, where it is stored for distribution or purification. Such petroleum wells are of great economic value, especially since the advent of the motor car and the aeroplane. Petroleum wells occur in the Baku district in Russia (Caucasia), Galicia, Cahfornia, Persian Gulf, Ontario, Ohio, Pennsylvania, Mexico, and elsewhere. The actual origin of petroleum is not known with any degree of certainty. The earlier view that the oil is generated by the action of water upon large underground masses of metallic carbides (compare methane, p. 367) has yielded place to the theory that petroleum is the result of the decomposition of vegetable and animal matter under varying conditions of temperature and pressure, such decomposition being effected deep within the bowels of the earth. Before use, the oil has to be carefully refined. The main rectification is effected by the process of fractional distillation, i.e. the distillate passing over between certain definite temper- atures is collected, and after further purification is put upon the market for various purposes. Table 38 indicates the boihng point and average composition of the various fractions. 366 AN INORGANIC CHEMISTRY TABLE 38 Name. B.P. Components. Use. Petroleum ether Naphtha . Benzine or petrol . Kerosene .... Lubricating oil Vaseline . Solid paraffin 70-90° 90-1 20° 120-150° 150-300° CgHig-CgHjo CgHjo-CigHaj solvent, gas fuel solvent, fuel illuminating oil Occasionally sulphur products are present in the oil, and as the combustion of such compounds leads to the formation of the deleterious sulphur dioxide, the oil has to be heated with copper oxide, whereby the sulphur compounds are removed. Treatment with sulphuric acid and sodium carbonate may also be necessary, the precise method of purification depending not only upon the composition of the oil, but also upon the purpose for which it is intended. Ozokerite is a species of natural parafi&n with a greenish opal- escence. After bleaching and purification, it is used in the manufacture of ceresine, a substitute for beeswax used in the manufacture of candles, ointments, etc. Another naturally occurring mixture of sohd hydrocarbons is Asphalt, found in the Pitch Lake of Trinidad, and used extensively for road-making. Unsaturated Hydrocarbons. — ^Another homologous series, known as the Olefines, is derived from ethylene, C2H4. In these compounds the carbon is unsaturated, having the general formula 0^11 3^. All these hydrocarbons are readily oxidised into saturated hydrocarbons by the action of chlorine, bromine, etc. The unsaturation of ethylene is indicated in the formula H H H— C=C— H where the atoms of carbon are joined by a double bond. The presence of the double bond in a compound does not increase its stability, but proves a source of weakness, for it is found that the addition of chlorine, etc., always breaks the double bond. Yet another series of unsaturated hydrocarbons is represented TYPICAL CARBON COMPOUNDS 367 by the acetylene homologues, derived from the general formula CjjHjy.j. In these compounds combination between the carbon atoms is by means of a triple bond. H — C ;= C — H. The Pkbpaeation and Properties of the Typical Hydrocarbons Methane is the first member of the saturated or paraffin series of hydrocarbons. It is the chief component of many natural gases which escape at some stage during the hfe of an oil well. It is also generated in small quantities in marshy and stagnant pools, hence its old name, marsh gas. Its occurrence in many coal mines as the dreaded " fire-damp " is undoubtedly due to the gradual decay of vegetable matter in the absence of air. The " fire-damp " accumulates in the coal, under more or less pressure, and when the coal face is shattered by an explo- sion, it escapes into the mine, often with disastrous effects. This is due to the explosive nature of the mixture of methane and air. After such an explosion, " choke-damp " (carbon dioxide) is set free. CH4 +202->C02 +2HA The formation of methane by the direct combination of carbon and hydrogen can be effected in the presence of finely divided nickel (q.v.) at a temperature of 250°. More often, it is pre- pared by the action of water upon aluminium carbide, AUC3 + 12H20-> 4A1(0H)3 i + 3CH4 ^ or by heating together a mixture of dry sodium acetate and sodium hydroxide. CH, H COONa „„ , ,^ „„ Q _.j^^-> CH4 -f Na^COa . This reaction is best carried out in a copper flask, and as the evolution occasionally takes place with explosive violence, it is advisable to heat the mixture carefully from the top rather than in the usual way. Chemically speaking, methane is very unreactive. Under the stimulus of sunlight, chlorine slowly replaces the hydrogen atom by atom. This substitution is represented in the equations : 368 AN INORGANIC CHEMISTRY « CH4 +CI2->CH3C1 + HC1 Methyl chloride. CH3CI +Cl2-^CH2Cl2 + HCl Methylene chloride or dichlormethane. CHjCl^+Cla^CHCla+HCl Trichlormethaiie or chloroform. CHCI3 + CI2 ^ CCI4 + HCl Carbon tetrachloride. Of these substitution products, both chloroform and carbon tetrachloride have considerable technical importance. Chloro- form is important as an ansssthetic, carbon tetrachloride as a solvent for fats, tars, etc. The extraction of oils from oil-bearing seeds is being increasingly effected by this chemical, as its non- inflammability makes it a much safer product to handle than benzene, petrol, etc. Carbon tetrachloride is also used in fire- extinguishers. Its action in this respect arises from the heavi- ness and the non-inflammabiUty of its vapour, which displaces the air round the burning objects and so chokes out the fire. Ethylene. — Ethylene, CaHj, olefiant gas, a tjrpical repre- sentative of the unsaturated homologues C„I^^^, is formed by the dehydrating action of sulphuric acid, or better still phos- phoric acid, upon ethyl alcohol. C2H5OH — H^O-^CaH^. Ethylene is also formed along with acetylene when any of the saturated hydrocarbons are strongly heated. 2CH 4 — >■ C all 4 -f- 2H 2- It is a gas which burns with a very luminous flame, owing to the separation of carbon during the preHminary stages of the com- bustion. If ethylene is bubbled slowly through bromine, it is absorbed with the formation of ethylene dibromide, a colourless liquid. C,H4+Br,^C2H4Br,. It is definitely known that in this compound the bromine atoms are symmetrically distributed between the carbon atoms ; H H I I H— C— C— H I I Br Br TYPICAL CARBON COMPOUNDS 369 for this etnd other similar reasons the constitution represented in the formula H H I I H— C=C— H is generally chosen for ethylene, the double bond indicating not that the two carbon atoms are more firmly hnked, but that each carbon atom is capable of entering into combination with another element or radicle. This is what happens when ethylene is treated with chlorine, potassium permanganate, sulphuric acid, etc. Acetylene. — This compound is a typical representative of the homologues possessing a triple bond. It is obtained by the direct union of carbon and hydrogen in the electric arc. The same type of decomposition which methane undergoes when strongly heated {q.v.) also takes place when ethylene is passed through a red-hot tube. C2H4 — ^-CaHj + Hj. More often, however, it is made by the action of water upon calcium carbide (p. 544). CaCa + 2H20->Ca(OH)2 + C2H2 f Acetylene is a colourless gas which possesses in the pure state a faint ethereal smell. It is appreciably soluble in water, and is freely soluble in alcohol and acetone (1 volume of acetone dis- solves 300 volumes of acetylene at 15° under a pressure of 12 atmospheres). Such solutions of acetylene in acetone are largely used as a source of acetylene for the oxy-acetylene blowpipe. The gas itself cannot be compressed beyond about two atmo- spheres, as the compressed gas is likely to explode it subjected to a slight shock, hence the use of the acetone solution as a store- house for the acetylene. The temperature of the flame produced by burning a mixture of acetylene and oxygen in specially de- signed blowpipes is in the neighbourhood of 2,500°, whereas the oxy-hydrogen flame scarcely exceeds 2,000°. The oxy-acetylene flame is hot enough to melt iron and steel with comparative ease (hence its extensive use for cutting steel girders, etc.), yet the reducing action of the carbon monoxide and hydrogen is sufficiently marked to prevent oxidation of the molten metal. Its success in welding is associated with both these factors. BB 370 AN INORGANIC CHEMISTRY The unsaturated nature of acetj'lene, as shown in its consti- tutional formula H— C^C— H, is illustrated by the ease with which the halogens are added on by acetylene, forming deriva- tives first of ethylene and finally of ethane. Although acetylene is comparatively stable at high tem- peratures, it is relatively unstable below 1,000°. When passed through a tube heated to about 750-800° it breaks down into its elements C,H3^2C+H,. This ready separation of carbon from acetylene accounts for the high luminosity of the acetylene flame, as well as for the luminosity of certain coal gas flames {q-v.). If acetylene is passed through an ammoniacal solution of cuprous chloride, a reddish brown precipitate of hydrated copper acetyhde, CuaCj.HaO, separates out. This compound is easily dehydrated, and forms the explosive copper acetyhde Cu,C,. Coal Gas Coal gas is obtained bj' heating coal in closed retorts to a temperature of approximately 1000°. When coal is treated in this way large quantities of gas are evolved, the volume and composition of this gas depending upon the temperature of the retorts. This is shown in Table 39, where the volume, illuminating power and composition of the gas evolved from one ton of an English bituminous coal at different temperatures are shown : — TABLE 39 Dark red. Temperature Gas volume in cbm. . Illuminating power, C.P. Hydrogen per cent. Methane Carbon monoxide . Heavy hydrocarbons . Nitrogen Bright red. 650°-700° 925°-975° 234 275 295 17-8 381 43-8 i 42-7 34-5 8-7 12-5 7-6 5-8 2-9 3-4 Bright orange red. 340 15-6 48 .30-7 140 4-5 2-8 (Small quantities of nitrogen are generaUy present.) Now that incandescent mantles have come into use, candle TYPICAL CARBON COMPOUNDS 371 power has ceased to be the objective of the gas works engineer ; the most satisfactory gas plant is that which produces the greatest volume of gas of a sufficient calorific power. Amongst the by-products in the coal gas industry is cohe, the soUd residue left in the retorts. This represents roughly 70 per cent, of the weight of the original coal, and owing to its high carbon content, it forms an excellent fuel, especially for metal- lurgical work. The absence of volatile matter and of the elements of water in coke accounts for the higher temperature given by coke as compared with coal, for no heat is required for the expulsion of the volatile matter. Where there is no direct demand for coke, it is often the practice to introduce steam into the retorts during the distillation. This leads to the formation of water gas through the reaction C+H.O^n^CO+H^. The resulting gaseous mixture is of poorer quality than the " normal " coal gas, but the greater volume of gas obtained compensates for its lower calorific power. The Calorific Power of a fuel is defined as the amount of heat furnished by the combustion of unit weight of the fuel. Engineers generally use the British Thermal Unit (B.T.U.) as the unit of heat — the amount of heat necessary to raise one pound of water 1° P. The number of B.T. units is determined by the combustion of a weighed amount of fuel in a bomb calorimeter. The principle of the Gas Calorimeter is similar, except that the heat generated by the combustion of a known volume of gas is measured. During the destructive distUlation of the coal a liquid to the extent of about 12-13 per cent, of the total weight of coal distilled, is collected. Part of this distiUate collects under water in the hydraulic main and runs off into the tar-well. Before passing into the gas holder, the gas circulates through a long series of water-cooled pipes (the condensers), where the last traces of the tar are removed. Further purification of the gas is effected by passing it through the scrubber, where all traces of ammonia and a little of the carbon dioxide and hydrogen sulphide are removed by the solvent action of the water. Finally, the purifiers, containing slaked lime and then hydrated oxide of iron, remove the last traces of hydrogen sulphide and most of the carbon dioxide. The tar in the tar-weU separates into two layers, the lower one consisting of tar, the upper of the 372 AN INORGANIC CHEMISTRY TYPICAL CARBON COMPOUNDS 373 ammoniacal liquor, from which ammonia (q.v.) is afterwards recovered. From the tar a large number of by-products is obtained, e.g. benzene, toluene, carbolic and creosote oils. These by-products form the starting point of the great coal-tar dyestuff industry, as well as of many explosives. Alcohols The alcohols are hydroxy-compounds which behave in many ways like inorganic bases. They can be synthesised from the halogen compounds by boiling with potassium or sodium hydroxide. C AI + KOH->C2H60H + KI Ethyl iodide. Ethyl alcohol. CH3I + KOH -> CH3OH + KI Methyl iodide. Methyl alcohol. Large quantities of methyl alcohol are obtained during the destructive distillation of wood, while ethyl alcohol is formed during the fermentation of sugar [q.v.). Both methyl and ethyl alcohol are miscible in all proportions with water. The alcohols react with phosphorus pentachloride to form a halide, R _ OH + PClj^- R — CI + POCI3 + HCl where R denotes the radicle CH3, C2H5, etc. Moreover, one hydrogen atom in an alcohol is replaceable by a metal like sodium — ^facts which point to the conclusion that one hydrogen atom is differently connected to the central carbon atom, viz., H H— C-O-H 1 H In their reactions with acids, the alcohols strongly resemble the inorganic hydroxides, e.g., CH3OH +H.HS04->CH3HS04 +H2O Methyl alcohol. Methyl hydrogen sulphate. Acids The two simplest organic acids are formic acid and acetic acid. 374 AN INORGANIC CHEMISTRY Formic acid, H2CO2, as has already been pointed out, liberates carbon monoxide on dehydration. It can be obtained by distilling red ants, but its structure is best shown by its synthesis from chloroform. Consider the reactions : /CI HOH /yG\ HOH Pf-Cl +H0H->P(0H)5 +5HC1^H3P04 + H20+5HC1 VCl HOH ^Cl HOH and < CI KOH ,0H H-C^-Cl +KOH -^ H - C ^OH + 3KC1 -^ ^Cl KOH \0H Orthoformic acid. OH H— C^ +H2O+3KGI Formic acid. O The latter reaction, amongst others, leads to the conclusion that formic acid has the formula, OH which is in agreement with its monobasic nature. The group, OH -< ^O is called the carhoxyl group, and is present in nearly all organic acids. Formic acid is a liquid boiling at 100-1°. Acetic acid, CH3COOH, is the first homologue of formic acid. It is monobasic, and when treated with sodium hydroxide forms an acetate, CH3COOH + NaOH-> CHjCOONa + H^O. Phosphorus pentachloride causes the replacement of the hydroxyl group by chlorine, CH3COOH + PCI5 -^ CH3COCI + POCI3 + HCl. Acetyl chloride. Acetic acid is prepared commercially by the dry distillation of wood (q.v.), as well as by the oxidation of dilute ethyl alcohol. TYPICAL CARBON COMPOUNDS 375 This alcohol is allowed to flow over wood shavings impregnated with the acetous ferment, when oxidation is brought about. C2H5OH + Oa-^CHjCOOH + H2O. As a typical organic, dibasic acid, no better example offers than oxalic acid, H2C2O4. This can be prepared from pine sawdust by fusion with sodium hydroxide. The sodium oxalate is leached out with water and the oxalate precipitated as the insoluble calcium salt. Na2C204+Ca(OH)2->CaC204i +2NaOH. It is also obtained by the oxidation of sugar with nitric acid. The reactions of this acid point to its possessing two carboxyl groups combined with each other : =C —OH =C -OH Oxahc acid is a moderately strong reducing agent, reducing acidified potassium permanganate on warming, 2KMn04 + 3H,S0, + 5H2C204-> K2SO4 + 2MnS04 + 8H2O + IOCO2 ESTEES Esters are formed by the action of an organic acid upon an alcohol. Their method of preparation and general reactions offer a considerable analogy with the inorganic salts : NaOH + HCl -^ NaCl + H^O Base. Acid. Salt. Water. C2H5OH +HOOC-CH3->CH3COOC2H5 +H2O Alcohol. Acid. Ester. Water. There is, however, an important difference between an inorganic salt and the organic ester ; the former is the result of an in- stantaneous reaction, the latter of a very slow process which does not go to completion. In the reaction between ethyl alcohol and acetic acid equilibrium is reached after two-thirds of the reacting substances have combined. Such reactions are generally accelerated by the addition of a little hydrogen chloride or sulphuric acid. On boiling an ester with water or a dilute acid, hydrolysis ensues with the formation of the free acid and the alcohol. With alkali they form the alcohol and a salt of the acid. 376 AN INORGANIC CHEMISTRY CHaCOOCA+NaOH-^CHaCOONa + CjHsOH Ethyl acetate. Sodium acetate. Ethyl alcohol. The esters form the sweet-smelling oils of many plants. Butter, fats and oils are examples of organic esters. Soaps. — Ordinary fat is an ester, formed by the combination of the tri-hydroxy-alcohol, glycerine with various higher fatty acids of the acetic acid family. If these fats are boiled with sodium hydroxide, they are hydrolysed with the formation of the sodium salt of the fatty acid and free glycerine : CijHasCOOCHa HONa CH^OH I I C^HasCOOCH -fHONa -> CHOH + SC^HssCOONa I I Sodium stearate. CijHasCOOCHj HONa CH^OH; Glycerine. In order to separate the soluble sodium stearate from the glycerine, common salt is added, and the sodium salts separate out in the form of a soUd scum. This is what is known as soap. If the fat is hydrolysed with potassium hydroxide, soft soap is obtained. These soaps, on treatment with water, are strongly hydrolysed with the production of sodium hydroxide, hence their cleansing properties. If fats are hydrolysed by means of superheated steam, the free fatty acids and glycerine are set free. The fatty acids can easily be separated from the sweet liquor, and after impurities have been pressed out, the fatty acid residue is used in the manufacture of candles. Ethees The ethers, as a class, bear the same relationship to the inorganic oxides as the alcohols do to the bases. R_0— H Na — 0— H Alcohol. Base. R— O — R ; Na— 0— Na Ether. Oxide. They are very unreactive substances and are almost insoluble in water. Ethyl ether, which is used largely as an anaesthetic as well as a solvent for fats, oils and resins, is prepared by the action of TYPICAL CARBON COMPOUNDS 377 sulphuric ac'd upon an excess of ethyl alcohol. The first step in the reaction is the formation of ethyl hydrogen sulphate. C2H5OH + H-HSO, -> C2H5HSO4 + H2O. On allowing a further amount of alcohol to enter the flask, ether is formed and can be distilled off. C,H,HS04 + C,H,OH^ (C,H,),0 + H.SO^. Benzene and its Compounds All the carbon compounds hitherto discussed are of the chain type, and belong to the aliphatic group of compounds. But there is another large class known as the ring or aromatic compounds, of which benzene, GJin, is the first member. About half of the carbon compounds belong to this group. Benzene is present in small quantities in coal gas. It is a volatile liquid boiling at 80-5°. From benzene and its homologues there can be derived compounds of the same types as have been discussed for the aliphatic group, e.g. alcohols, acids, halides, etc. The properties of these compounds show a general likeness to the corresponding compounds of the aliphatic class. Alcohol, CeHjOH, commonly called phenol, or carbolic acid. Acid, CeHjCOOH, benzoic acid. Halide, CoHsCl, chlorbenzene. Cyanogen — Cyanides — Cyanatbs — Thiocyanates Cyanogen is formed by sending electric sparks between carbon poles in an atmosphere of nitrogen, more conveniently, however, by the action of a solution of copper sulphate upon potassium cyanide. The unstable cupric cyanide, at first thrown down, quickly decomposes, forming cuprous cyanide and cyanogen. 2KGN + CuS04^^Cu(CN), +K2SO4 2Cu(CN) 2 -^ 2CuCN + (CN) , Cyanogen is a poisonous gas which burns with a characteristic violet colour. It unites directly with the alkali metals, forming cyanides, 2Na +(CN)2— >- 2NaCN, and on being passed into a solution of the hydroxides of the alkalies, it forms a cyanide and a cyanate (auto-oxidation) : 2NaOH + (CN)2-^NaCN + NaCNO + H^O Cf . 2NaOH + CI2 -> NaCl + NaClO + HjO 378 AN INORGANIC CHEMISTRY Hydrocyanic Acid, a dilute solution of which is often known as Prussic Acid, is formed by distilling potassium cyanide and sulphuric acid. It is a colourless, intensely poisonous Mquid boiling at 26-5°. When dissolved in water, it has a feebly acid reaction. The soluble cyanides are therefore strongly hydro- lysed in aqueous solution and react alkaline. KCN + H20^=:± KOH + HCN. There seems considerable doubt about the hnking of the atoms in hydrocyanic acid ; the bulk of the evidence on the organic side indicates that two tautomeric forms exist in equilibrium, H— C=N ;=±H— N = C, whilst the tendency of this compound to add on oxygen, sulphur,, etc., is held to support the formula H— N=C, where the C is unsaturated. Cyanates. — Owing to the strong reducing properties of the cyanides, these compounds often contain small quantities of cyanates. These salts are, however, generally prepared by heating potassium cyanide vith litharge or other suitable oxidising agent, KCN+PbO-^KCNO+Pb. The cyanate is extracted by means of alcohol, in which it is freely soluble. Thiocyanates. — If the oxidation of potassium cyanide is effected with sulphur, instead of oxygen, a thio-cyanate is formed : KCN + S->KCNS. The oxidation can be effected either by fusion, or even by mere boiling. The thiocyanates are important analytically owing to their forming a blood red colour in the presence of ferric iron (p. 204). FeClj + 3KCNS ^=± Fe(CNS)3 + 3KC1 . Questions 1 . Give a succinct account of the saturated hydrocarbons . 2. Compare the action of water upon the nitrides, phosphides and carbides. 3. A sample of coal gas had the following percentage composition by volume : Hydrogen 40, methane 35, carbon monoxide, 18, acetylene 5, nitrogen 2; 100 volumes of it are exploded after admixture with 180 volumes of oxygen. Calculate the volume and composition of the result- ing mixture of gases. TYPICAL CARBON COMPOUNDS 379 4. A gaseous substance is found by experiment to be eight times as heavy as hydrogen (measured under similar conditions of temperature and pressure). Fifty c.c. of the gas are exploded with excess of oxygen, and the volume is found to contract after explosion and cooling by approximately 100 c.c. A further contraction of about 50 c.c. occurs after treatment of the residual gas with a solution of caustic soda. The final residue is found to be pure oxygen. What was the original gas, and what its molecular weight ? 5. CalciJate the heat of formation of methane : C+Oj =COj+96-9 cals. Ha+0=H20+68-4 cals. CH4+40=C02+2HjO + 213-5 cals. 6. Compare and contrast the compounds ethane, ethylene and acety- lene. 7. What is isomerism t Illustrate your answer by reference to the paraffin series of hydrocarbons. CHAPTER XXV FLAME Combustion. — The present day theory of combustion dates back to the time of Lavoisier, the great French chemist, who successfully laid low the old " phlogiston " hypothesis, so vigorously championed by Stahl (see p. 6). In 1777 Lavoisier put forward the view that oxygen is necessary for combustion, and that when combustion occurs, the increase in the weight of the substance burnt is exactly equal to the weight of the oxygen which has disappeared from the surrounding atmosphere, and he was able to substantiate these views by incontrovertible experiments. Combustion, then, is nothing but the process oj oxidation accompanied by the development of light and heat. Tht term oxidation is here understood in its widest sense and is not limited to processes in which oxygen takes part, so that one may quite well speak of the combustion of phosphorus in chlorine. Generally speaking, the atmo- sphere surrounding the burning substance is referred to as the supporter of combustion, whilst that which is actually bm-ning is known as the combustible. When coal gas burns in air, the coal gas is viewed as the combustible, the air as the supporter of combus- tion, but after all this is merely a convention, for with a slight alteration in the conditions of the experiment air, i.e. oxygen, can be made to burn in coal gas. This can be strikingly shown by means of the apparatus 380 Fig. 100. FLAME 381 ■d shown in Fig. 100. The hole at the top of the chimney is at first closed. Coal gas is passed through the chimney from below until the air is completely displaced. The escaping coal gas is then lit at ^. If the hole at the top is now uncovered, the flame recedes up the tube from A to B. This flame is, as it were, the normal coal gas flame turned inside out, and is the result of the vigorous reduction of atmospheric oxygen by the surrounding coal gas. By this means air, which is generally the supporter of combustion, is turned into the combustible. The Stbuctube of Flame The Candle. — An examination of the candle flame leads to the conclusion that there are four distinct zones (Fig. 101). At the bottom of the flame there is a small region c, bright blue in colour, which is non- luminous. Just above this there is a dark space d. This con- sists of unburnt gases formed by the decomposi- tion of the heavy hydro- carbons drawn up the wick by the forces of capillarity. That the flame is indeed hollow, can be readily demonstrated either by inserting a narrow tube into the flame and igniting the issuing inflammable gases, or by depressing a piece of asbestos paper upon the flame, when the centre of the paper is found to be quite unmarked by the flame. Above the dark cone is a brightly luminous portion a and surrounding the whole flame there will be seen a faintly luminous mantle h. In each portion of the flame characteristic chemical processes take place. In the area h there is sufficient oxygen present for the complete combustion of the hydrocarbons derived from the candle. The — a Fig. 102. 382 AN INORGANIC CHEMISTRY dark cone d consists oif unburnt gases, and the luminous mantle a owes its luminosity to the presence of incandescent particles of carbon. It is known that ethylene (q.v.), etc., breaks down when strongly heated, forming acetylene, whUst under similar treatment acetylene yields carbon and hydrogen. Presumably, therefore, the heavy hydrocarbons from the wick undergo a s imil ar decomposition. In the outer mantle the mixture of carbon, hydrogen and undissociated hydrocarbons undergoes complete oxidation to the final products — water and carbon dioxide. The Bunsen Burner. — ^This well-known burner is capable of giving two distinct types of flame — the luminous and the non-luminous, the latter being produced when extra air is admitted along with the gas by adjusting the air-holes at the bottom of the burner. The luminous flame is very similar to that of a candle flame, a consists of unbumt gases, 6 is a bright blue layer, forming a cap over the dark zone, c is the large lumi- nous cone and d the outer, non-luminous mantle (Fig. 102a). In the non-luminous flame only three parts. (a) The inner dark cone of unbumt gases of the flame can be distinguished (Fig. 102b) : (6) The blue sheath covering this, (c) The outer mantle. Little progress concerning the structure of the Bunsen burner was made until SmitheUs introduced the special burner, that enabled the gases present between the blue cone and the outer mantle to be removed without disturbing the equiHbrium, as had always been done hitherto when attempts were made to remove the gases by means of a tube. The SmitheUs' burner consists of two co- axial tubes, the outer of which can be easily slid up and down (Fig. 103). Both the tubes are capped with mica. The smaller tube is fitted over a Bunsen burner, and the burner lit with the tops of the two tubes in the same plane. By adjusting the air holes the flame is made just non-luminous, Fig. 103. FLAME 383 and then the outer tube is carefully slid upwards. It is found that the outer cone ascends with the outer tube, whilst the inner blue cone remains upon the inner tube. In this way the two cones are puUed apart. Samples of the gas present between the two cones can be tapped off at will through a side tube. Soon after the introduction of this burner Haber succeeded in showing that the gases between the blue and the outer cones consisted of the oxides of carbon, hydrogen and water in a state of equihbrium, CO+H,O^^CO,+H,. Therefore, in this part of the flame the water gas equilibrium exists, and whatever oxygen penetrates to this part of the flame distributes itseK in accordance with the above equation. It was shown conclusively that the above equilibrium changed with the temperature of the flame (cf. Table 37, p. 360). Especially interesting was the observation that the observed temperature of the inner cone of a Bunsen burner coincided with that demanded by the water-gas equilibrium mixture as found in this portion of the flame. In the innermost cone of a non-luminous flame, there- fore, the carbon is first at- tacked while the hydrogen is left unoxidised. If these facts are applied to the luminous flame where optical tests reveal the presence of solid incandescent particles, there can be little doubt that in the innermost cone the hydro- carbons are broken into a mixture of carbon and hydro- gen. Ethylene is a consti- tuent of coal gas, and the effect of heating this gas is to break it down first into acetylene, and finally into carbon and hydrogen. The presence of solid carbon particles in the luminous flame, therefore, receives a satisfactory explanation from this theory. Strong support of this view is afforded by the formation of acetylene when air burns To aspirator Gas enters Fig. 104. 384 AN INORGANIC CHEMISTRY in coal gas. The gas containing the products of combustion is led through a wash-bottle containing an ammoniacal solution of cuprous chloride (Pig. 104) ; copper acetyUde is precipitated. In the outer mantle of a Bunsen flame complete oxidation to carbon dioxide and water takes place, at least if the supply of air is properly adjusted. If the gas issues at too high a pressure small quantities of carbon monoxide and of hydrogen may escape combustion, and are found among the products of combustion. The action of Air in a Bunsen Burner. — Seeing that the same amount of gas is burnt by a burner whether the flame be luminous or non-luminous, and that the ultimate products of the combustion are the same in both cases, at first sight one might be tempted to conclude that the same temperature would be reached by either type of flame. Such a conclusion would, however, not be justified, because : (1) The air admitted into the burner has to be raised to the temperature of the flame, and will therefore exert a cooling effect. (2) The larger luminous flame loses by radiation into space more than double the amount of heat lost by the more compact non-luminous flame. The net result of this is that the non-luminous flame is consider- ably hotter than the luminous. According to the determination of Fery (1904), the maximum attainable temperature with a Bunsen non-luminous flame does not exceed 1,870°, while Lewes (1895) claims that a luminous flame gives a maximum temperature of 1,330°. The effect of the air in reducing the luminosity of a Bunsen flame is somewhat complicated and probably arises from three sources : (a) Oxidation. — The presence of solid carbon particles m the luminous flame, and the appearance of carbon monoxide in the inner mantle of the non-lumuious flame suggest that the oxygen of the admitted air prevents the separation of the carbon particles by bringing about their oxidation to carbon monoxide. This is supported by the experiments of Lewes, who tested the effect upon the luminosity of a flame by changing the percentage of oxygen in the admitted air. Thus ; FLAME 385 Mixture ofO:N .1:0 1:1 1:2 1:3 1:5 Vol. of gas admitted to 1 of coal gas to give non-luminous flame 0-5 1 1-5 2-0 2-3 (6) The Cooling of the Flame. — As already indicated, all gases introduced into the burner must be heated to the temperature of the flame. Whatever action the admitted oxygen may have in promoting oxidation, the effect of the nitrogen must be to lower the temperature of the flame. Lewes found that the luminosity of the flame was greatly diminished by other inert Gas ..... Vol. required to produce non- luminous flame Nitrogen 2-3 Carbon dioxide 1-3 Air 2-3 This cooling action plays an important part in the luminosity of a flame as is shown by the following facts : 1. A luminous flame becomes much less luminous if it is made to impinge upon a lump of iron, etc. 2. If the admixed gases of a Bunsen burner which normally pro- duce a non-luminous flame are heated prior to ignition, the flame becomes lumin- ous (Fig. 105). 3. If the dilut- ing gas is heated before entering the burner, the luminosity is considerably in- creased. (c) The stability of the hydrocarbons is considerably increased by the presence of the nitrogen. Possibly this is due to the dilution affecting the speed of the dissociation of the heavy hydrocarbons, so /that the ethylene, etc., escapes the usual dissociation as it liasses into the inner mantle, and when it does react at the higher temperature ruling in the upper portions of the flame, no Reparation of carbon takes place. The Explosion Wave and the Bunsen Burner. — The velocity of an explosion wave is conditioned by the rate at CO Fig. 105. 386 AN INORGANIC CHEMISTRY which the combustion spreads throughout a mixture. The velocity with which this explosion wave is propagated, depends upon the relative concentrations of the reacting substances. When a Bunsen burner is burning quietly, fed by the usual mixture (about 2-5 of air to 1 of coal gas), the combustion at the mouth of the burner is explosive in nature, but, as the speed with which the gases are issuing exceeds the speed of the explosion wave, this wave does not strike down the burner. If, however, the ratio of air to gas is increased by reducing the supply of gas, there is a corresponding increase in the velocity of the explosion wave, and, sooner or later, the explosion wave will be propagated at a rate greater than the rate at which the gases are issuing at the jet, hence the burner strikes back. The effect can be beauti- fully illustrated by clamping a long glass tube 5-6 feet long and about 4 cms. wide, over the mouth of a Bunsen burner, the ingress of air being prevented by plugging the junction with cotton wool. The air holes are first of all closed and the jet of gas ht. It is advisable not to turn on the tap full. If the air holes are now slowly opened, the luminous flame gives way to a well-defined double cone. The luminous flame be- comes more and more unstable as the proportion of air increases by opening the air holes, i.e.' the rate of the explosion wave is only slightly exceeded by the speed at which the gases are rushing through the tube. A further increase in the amount of an- admitted causes the wave to strike down the tube with consider- able violence. The flame continues to burn at the top of the burner itself. This tendency of the Bunsen burner to strike back mitigates against attaining the highest possible temperature for the amount of gas consumed, seeing that for complete combtistion over twice as much air is required as this type of burner will bear without Plan of nickel |nd for Meker burner. Fig. 106. FLAME 387 striking back. This difficulty has been overcome in the Meker burner, wherein a mechanical barrier in the form of a deep grid is interposed at the top of the burner to prevent the striking back. With this arrangement and with large air holes the flame secures sufficient air for complete combustion, and there is no inner cone of unburnt gas (Fig. 106). Questions 1. Give an aocomit of the structure of the flame of a Bunsen burner. 2. What theories have been put forward to explain the luminosity of flames ? 3. Write an account of what is understood by the term combustion. 4. By what means may the luminosity of a flame be increased ? 5. How do you account for the brilliancy of the light given out when " incandescent " mantles are vised with a non-luminous gas flame ? CHAPTER XXVI SILICON General Remarks. — Silicon belongs to the fourth group of the Periodic Table, falling immediately below carbon. In many respects a very close resemblance exists between these elements. This is reflected in the strong similarity between the oxides SiOj and CO 2, as well as in the tendency of silicon to form a short series of homologues SiHj, SioHg, etc. Occurrence. — In combination with other elements sihcon constitutes more than a quarter of the earth's crust, ranking next to oxygen. Silicon dioxide (silica) occurs freely in nature in the form of sand, quartz, etc., but enormous quantities of it also occur in a state of combination with bases, forming sUicates, the chief constituent of rocks. Preparation. — Sihcon, like its analogue carbon, can be obtained in more than one modification. Amorphous Silicon may be made by the action of sodium upon the vapour of silicon tetrachloride, the resulting sodium chloride being washed away from the amorphous powder. 4Na + SiClj -^ 4NaCl + Si. In place of silicon tetrachloride, which is inconvenient to handle, it is customary to use sodium sUico-fluoride when the reaction is represented by the equation Na^SiFe + 4Na -^ 6NaF + Si. The sodium fluoride is dissolved away by careful washing with hydrofluoric acid and water. The best method of preparing amorphous silicon is by the action of magnesium on silicon dioxide. 2Mg + Si02->Si +2MgO. 388 SILICON 389 The reaction is generally vigorous, and often leads to the formation of small quantities of magnesium silicide. The resulting mass is freed from magnesium silicide and oxide by the action of hydro- chloric acid. Crystalline Silicon can be obtained by dissolving the amorphous form in molten zinc. If, after cooling, the zinc is dissolved out by means of an acid, crystaUine silicon remains. Good yields are obtained by passing silicon tetrafluoride over aluminium heated in a hydrogen atmosphere. In this case the aluminium acts both as reducing agent and as solvent. As soon as the silicon is formed, it dissolves in the aluminium, but it is thrown out in shining crystals on cooling. A chemically similar method in which a different reducing agent and solvent are employed is the reaction between sodium siUcofluoride, sodium (reducing agent) and zinc (solvent). Properties. — Amorphous silicon is a dark brown powder which is distinctly reactive. It combines freely with chlorine at 450°, bromine at 500°, sulphur at 600°, nitrogen at 1,060° forming a nitride, SijNi. It reacts readily with hydrogen fluoride, but with hydrogen chloride only at a bright red heat. Si+4HCl->SiCl4 + 2H2. With steam it forms sUica and hydrogen at a bright red heat. It is freely soluble in alkalies, forming a meta-silicate and liberat- ing hydrogen, Si + 2NaOH + HaO-^NaaSiOj + 2H2 this reaction forming the basis of a method of manufacturing hydrogen for airships and balloons. It is scarcely attacked by acids, a mixture of hydrofluoric and nitric acids dissolving it slowly with the formation of silicon fluoride. Crystalline silicon occurs in shining metallic octahedra, belong- ing to the regular system. It is brittle and hard enough to scratch glass. It is a good conductor of electricity, comparing well with graphite. Chemically, it behaves much like the amorphous form, but owing to its decreased surface, is less reactive. Both varieties oxidise superficially on heating. It reacts with chlorine, hydrofluoric acid and steam in a similar way to the amorphous form. Both kinds of silicon dissolve on boiling with sodium or 390 AN INORGANIC CHEMISTRY potassium hydroxide, while, with fused potassium carbonate, carbon is displaced, i.e. silicon is able to displace carbon from an oxy-compound at this temperature. KaCOa + Si^^KaSiOs + C. A similar tendency was remarked in the halogen group, e.g. iodine displaces chlorine from a chlorate. In a group of related elements the reducing power of the element is often found to increase with the atomic weight. During recent years silicon has assumed a certain importance commercially through the discovery that its presence in small quantities in steel has a marked influence upon the hardness and other physical properties. Silicon Hydride. — At least three compounds of silicon and hydrogen are known, silico-methane, SiHj, silico-ethane SigHe, and sihco-acetylene SiaH,. Silico-methane, or sUicane, SiHj, is prepared by acting upon magnesium sUicide with hydrochloric acid. MgaSi + 4HCl-> 2MgCl2 + SiH^. A gas which is spontaneously inflammable escapes, consequently the operation has to be carried out in an atmosphere of hydrogen or coal gas. The inflammability arises from the presence of silico-ethane. In order to isolate sUico-methane, the gaseous products are led through a condenser immersed in liquid air, when a mixture of silico-methane and silico-ethane separates out. This mixture is resolved into its constituents by careful fractionation. Silico-methane is a colourless gas, which decomposes explosively on heating above 400°. On burning the gas, siHcon dioxide and water are formed. It reacts most energetically when passed into potassium hydroxide. 2K0H +H2O +SiH4->K2Si03 +4H2. Potassium silicate. It is a more vigorous reducing agent than its analogue methane, CH4, for on passing into a solution of copper sulphate, a mixture of copper and copper silicide is thrown down (cf. reducing action of HCl and HI). The chief interest attaching to silico-ethane and silico-acetylene lies in the fact that they emphasise the strong relationship between the analogues, silicon and carbon. SILICON 391 Silicon Halides The weU-defined nature of the halides of carbon, CCli, Cfil^, CHClj, etc., ah-eady discussed, and the family relationship exist- ing between silicon and carbon would lead one to expect the existence of a similar series of halides in the case of silicon. Such a series does exist, though the number of such compounds is very much less. However, sufficient have been investigated to show that silicon does possess in. a minor degree the same power of chain linking that has led to such fruitful results in the case of carbon. Silicon forms the following halides : SiCli SiaCle SigClg SiHCla Cf. CCI4 C2CI0 C3CI8 CHCI3 whilst bromine, iodine and fluorine form almost similar com- pounds. Of these compounds silicon tetrachloride and tetra- fluoride are of most importance. Silicon Tetrafluoride has been made by the direct action of fluorine upon amorphous silicon, but it is generally prepared by acting upon silica or a silicate with hydrogen fluoride. SiOj + 4HF-> SiF^ + 2H2O. In practice, the hydrogen fluoride is generated from an intimate mixture of calcium fluoride and sulphuric acid. Silicon tetra- fluoride is a colourless, fuming gas, which is freely hydrolysed by water with the precipitation of gelatinous ortho-silicic acid, SiFi -i-4H20-> Si(OH)i i +4HF. The hydrolyticaUy generated hydrofluoric acid immediately combines with another molecule of sihcon tetrafluoride, forming hydrofluosilicic acid. Hydrofluosilicic Acid cannot be obtained in the pure state, for, on concentrating its solutions, the equilibrium, SiF4 + 2HF ^zzi H^SiFe is disturbed by the escape of the volatile sihcon tetrafluoride and hydrogen fluoride. Many salts of this acid have been prepared, of which those of barium and of potassium are fairly insoluble. 392 AN INORGANIC CHEMISTRY Silicon Tetrachloride is produced by the direct action of chlorine upon silicon, or more often on an intimate mixture of siHca and carbon. In this latter reaction the carbon reduces the silica to sUicon, and reaction between the silicon and chlorine at once takes place (cf . preparation of the chlorides of aluminium, chromium, boron, etc.). SiOa + 2C + 2CI2 -> SiCl4 + 2C0. Silicon tetrachloride is a colourless gas which fumes strongly in the air, and boils at 59-2°. It reacts at once with water. SiCli + 4H0H^ Si(OH), + 4HC1. In this respect there is a marked difference from carbon tetrachloride which hydrolyses very slowly. Silicon tetrachloride is also formed when crystalline silicon is heated with hydrogen chloride. A mixture of silicon tetrachloride and silico-chloroform, SiHCls, is formed and the latter can be separated by fractional distillation. SiHco-chloroform is a liquid boiling at 33°. When treated with water, siUco-formic acid is formed. CI KOH /OH H— C^Cl+KOH -^H— C^0H+3KC1 -^ ^Cl KOH ^OH Chloroform. Ortho-formic acid. OH H— C^ +H2O+3KCI ^0 Formic acid. CI HOH OH H— Si(-Cl+HOH — >H— SiA0H+3HCl \C1 HOH ^OH Rilico-chloroform. Silico-formic acid. Silicon Carbide. — Silicon carbide (carborundum), SiC, is made on the commercial scale by fusing together a mixture of sand, coke and a small quantity of salt and sawdust. The fusion is carried out in an electric furnace at a temperature of 3,500°. The object of the sawdust is to impart porosity to the charge, while the salt removes the oxides of the metals, e.g. iron, in the form of the volatile chloride. The essential chemical reaction is given in the equation. SiO, + 3C-^SiC+2CO. SILICON 393 On cooling the furnace a central core of graphite is found, arising from the dissociation of the silicon carbide first formed there. Surrounding this core are found crystals of carborundum, the best grade being found near the core. After grinding and purifying by means of sulphuric acid, it is graded and put on the market as an abrasive. Carborvmdum crystallises in the hexagonal system, its crystals showing a fine play of colours. It is extremely inactive chemically. Oxygen and sulphur are without action even at 1,000°, while chlorine has only a superficial action at 600°. Acids are also without action. When it is fused with potassium hydroxide, a mixture of silicate and carbonate is obtained. Silicon Dioxide. — Silicon dioxide (silica) occurs widely in nature in a fairly pure state, whilst its distribution in the form of siHcate is even more abundant. As examples of silica the follow- ing wiU serve : fUnt and opal, slightly hydrated amorphous forms of sUica ; kiesel-guhr (diatomaceous earth), smoky quartz, which probably owes its appearance to the presence of carbon- aceous material, milky quartz, amethyst, and lastly, quartz or rock crystal. Quartz crystallises in the hexagonal system, hexagonal prisms terminating in hexagonal pyramids being generally found. Many of these crystals are of very considerable size. The purest varieties have a specific gravity 2-68 at 4°. Distinctive colours are often imparted to the crystals of quartz by traces of extraneous oxides, viz. amethyst, which owes its colour to the presence of manganese. Silica occurs in two other modifications — tridymite and cristohalite ; the former of these crystallises in the hexagonal system with a specific gravity 2-33, the latter in the regular system (sp. gr. 2-34). Tridymite is the stable form at high temperatures, though there is still doubt as to its exact transition temperature. The exact melting point of silica is not known, but at 1,500° it is distinctly plastic. When in the molten state, it can be worked into various forms of chemical apparatus. The resistance of quartz vessels towards anything but alkaline liquids is much greater than that of glass ; while, owing to its small coefficient of expansion, vessels made of quartz can be subjected to great temperature variations without fear of fracture. Quartz also finds appUcation in the manufacture of threads for suspension 394 AN INORGANIC CHEMISTRY work, while its transparency to ultraviolet light is taken advantage of in certain branches of spectroscopic work. SUica is exceedingly unreaotive chemically. Hydrogen sul- phide has a slight action above 1,000°, forming silicon disulphide, and fluorine alone among the halogens attacks it. Acids, with the exception of hydrofluoric acid, are without action. The alkaUes, calcium, barium, and magnesium reduce silica to sUicon. When sUica is boUed with sodium or potassium hydroxide, ready solution takes place with the formation of the ortho-siUcate. 4K0H + SiOa^ KiSiOi + 2H2O. On fusing with the alkali carbonates the metasilicates are formed. Na^COj + SiO,-^ Na^SiOa + CO^. The replacement of the weakly acidic carbon dioxide by the still more weakly acidic silicon dioxide is due to the greater volatility of the carbon dioxide at this temperature. Even sulphates are similarly broken down by silica at very high temperatures, owing to the greater volatiUty of sulphur trioxide. Silicic Acid. — ^When the silicate of an alkali is treated with hydrochloric acid, a gelatinous precipitate is slowly thrown down. Na^SiOa + 2HC1 + H^O-^ Si(0H)4 i + 2NaCl. The precipitate is supposed to be hydrated ortho-silicic acid, but although many attempts have been made to obtain the ortho- and meta-acid by the dehydration of this precipitate, there is no definite evidence that such compounds exist. Thus the vapour pressure curve shows no break when water vapour is steadily removed, as does that for hydrated copper sulphate. And yet many compounds of these acids are known. This, after aU, is no new phenomenon, for many cases have already been met with where the acid has not been prepared in the pure state, but well-defined stable salts exist (of. thiosulphuric acid). In the above reaction for the preparation of sflicic acid, if the solution of the silicate is poured slowly into concentrated hydrochloric acid, the silicic acid does not separate out, but remains in colloidal solution. As such, it forms what Graham termed a Hydrosol. The colloidal solution of silicic acid will pass through an ordinary filter paper, so that in order to effect a separation from the sodium chloride, the process of dialysis is SILICON 396 made use of. The solution is placed in a vessel the bottom of which consists of parchment or animal membrane, and the whole is placed in running water. Dissolved salts or crystalloids have the power to pass unhindered through such a membrane, but the larger colloidal particles are unable to escape from the ceU. The gelatinous form of silicic acid, prepared, as above, by the addition of hydrochloric acid to a solution of a soluble silicate, was named a hydrogel by Graham. Peopeetibs of Colloidal Solutions Colloidal particles, whether prepared by dialysis, by the sparking method (p. 113) or by chemical means, possess certain interesting and important properties. As a rule, colloids, e.g. albumen, are readily coagulated by heat. All colloids carry a definite electrical charge ; for instance, it has been shown that if a eoUoidal solution of a metal, sulphur, or a sulphide is electrolysed in a U-tube, the coUoid wanders to the positive pole. This proves that these colloids are negatively charged. A migration of this type is known as electrophoresis or electrical endosmosis. The phenomenon differs from ionic migrations (p. 420) as there is no migration of charged ions in the opposite direction. When the charged colloid reaches the oppositely charged electrode the charge carried by the colloid is neutralised, and precipitation or coagulation of the colloid ensues. The Brownian movement of colloids under the bombarding action of the solvent molecules has already been referred to {q.v.). One would expect that collision between the rapidly moving colloidal particles, followed by coagulation, would ensue, but the fact that these particles carry a like charge and therefore repel each other, effectively prevents coalescence of the particles into larger agglomerates. Among the coUoids which bear a positive charge are the metallic hydroxides, e.g. Fe(0H)3, and many of the organic dyestufis. Another method of producing the coagulation of a colloidal solution is by the addition of a suitable electrolyte. A negatively charged colloid is precipitated by fairly small quantities of a salt containing a divalent cathion, e.g. CaClj (p. 411) ; more effectively stUl, by the action of small quantities of a salt containing a tri- or tetra-valent cathion, e.g. AICI3 ; the efficiency of the electrolyte increases rapidly with a rise in the valence of the cathion. Similarly, the valence of the anion is the prime factor in deter- 396 AN INORGANIC CHEMISTRY mining the coagulating power of a salt towards a positively charged colloid. Thus ferric hydroxide is coagulated by smaller quantities of a sulphate than of a chloride, whilst a phosphate is even more effective in inducing coagulation. As is to be expected, the mixing of two colloids, one positively charged, the other negatively, causes immediate precipitation of both. Silicates No system acceptable to the chemist and the mineralogist has yet been proposed for the classification of naturally occurring silicates. The frequent occurrence of such amphoteric oxides as AI2O3 in these sihcates raises the question as to whether the aluminium is present as an aluminium silicate or as an aluminate- silicate, a question not yet answered in many cases. Again, a compound of the general formula, RiSiOi may be classed as a neutral ortho-silicate derived from orthosUicic acid, but such a salt may also be classed as a basic meta-silicate RgSiOj-RaO. The diificulty of answering such questions as these has prevented a satisfactory scheme of classification from being adopted. Probably that of Groth is the most satisfactory yet proposed. He refers the silicates to the following acids : H^SiO^ H4Si04— H20=H2Si03 2H4Si04— H20=HeSi20, 2H4Si04— 3H,0=H2Si205 SH^SiOj— 4H20=H4Si308 These are summarised in the follo\ving table : OrthosUicic acid Metasilicic acid DiparasiUcio acid Dimetasilicic acid Tri-orthosUicic acid TABLE 40 Name. Hypothetical acid. Silicate. Mono. j Di. j Tri. Poly. Meta- Ortho- . Para- . H.OSiO., 2H20Si02 .SHjOSiOj RaOSiOj ! RaO^SiOa RjO-SSiO^ iR^OSiO., 2R20-2Si02 l2R20-3SiO. SRaOSiOa 3R30-2Si02 SR^O-SSiOz RjOrcSiOj 2R20a;Si02 3R20-a:Si02 As examples of orthosOicates we have garnet, CajAla {Si04)3; zircon, ZrSi04 ; kaolin, H2Al2(Si04)2-H20 ; mica, KHiAl3(Si04)3. SILICON 397 Metasilicates — beryl, gl3Al2(Si03)6; ensfcatite, MgSiOa. Disilicate — serpentine, Mg jSi 20,2H 2O . Trisilicate — orthoclase, KAlSiaOg (potassium felspar). Glass. — Glass is a complex silicate of varying composition. Slight alterations in the composition of the glass exercise a far-reaching influence upon the properties of the glass. In general, one may say that glass is made by the fusing together of such bases as lime, lead oxide, soda or potash with varying amounts of pure sand. Soda glass, formed in accordance with the equation NaaCOa + CaCOs + 6Si02^^NaaSi03,CaSi03,4SiOa + 200^ is very fusible, hence it is often referred to as soft glass. The fusion is effected in foeclay pots, and when the mass has cooled to a pasty state, a little is collected at the end of an iron tube and then blown into a mould of the desired shape. Glass, which has been rapidly cooled, is liable to splinter, and in order to counter- act this, it must be carefully annealed, i.e. cooled in a specially designed Idln. Potash glass, commonly known as Bohemian glass, is much harder and more difficult to fuse than soda glass. Moreover, the glass is less readUy attacked by water and chemical reagents, hence its extensive use in the manufacture of beakers, etc. It is really a potash-lime sUicate. If the hme is replaced by lead oxide, flint glass is obtained. This type of glass is characterised by its high refractive power and is used extensively in the manufacture of lenses and for cut glass ornaments. The addition of various metallic oxides produces a marked difference in the colour of the glass. Small quantities of cobalt oxide yield a blue glass ; finely divided gold and copper scattered throughout the glass in a fine state of colloidal suspension give a ruby red glass. White glass is produced by the addition of bone-ash or cryolite to the melt. In the manufacture of bottle glass sodiuin sulphate is used instead of the more valuable carbonate. On heating glass it passes slowly through a pasty state and finally into a liquid. The absence of a definite melting point and of a definite crystalline structure shows the amorphous nature of glass. It is a supercooled liquid. Crystallisation may be induced in the glass by maintaining it for some time at a temper- 398 AN INORGANIC CHEMISTRY ature somewhat below the softening point. Such glass is said to devitrify. Questions 1. Give a brief account of the preparation and more important pro- perties of colloidal solutions. 2. Compare and contrast the more important compounds of carbon and silicon. 3. Give an account of the hydrides of silicon and compare them with the corresponding compounds of carbon. 4. By what means may one prepare silicon tetra fluoride ? How does it behave when treated with water ? 5. Explain the terms : dialysis, electrophoresis, coagulation. 6. Compare the allotropic modifications of carbon and silicon. CHAPTER XXVII OSMOTIC PRESSURE— MOLECULAR WEIGHTS OF DISSOLVED SUBSTANCES Semi -Permeable Membranes. — The metal palladium pos- sesses the rather unusual property of being appreciably permeable to hydrogen at temperatures above 200°. This permeability is undoubtedly associated with the ready solubility of the hydrogen in the metal. Quantitative experiments on the permeabihty of hydrogen through palladium have been carried out by means of a palladium tube. This was filled with nitrogen at a pressure Px, connected with a manometer and immersed in a vessel fiUed with hydrogen at a constant pressure p,,- It was found that the manometer, which was attached to the palladium tube, showed a steady increase in pressure until its value nearly equalled the sum jJi+i'z- In short, the hydrogen distributed itself between the two compartments of the apparatus until its pressure within and without the palladium tube was the same, or nearly so. A membrane behaving in this way is known as a semi-permeable membrane. In practice, no such membrane is perfect but many satisfactory semi-permeable membranes are known. Thus it has long been known that animal membrane is permeable to water but not to salts dissolved in water. In 1867 Traube showed that a chemical semi-permeable membrane could be readily made by precipitation, though it was PfefiEer (1877) who successfully developed Traube's idea. Pfefier used a porous cell into which he put a solution of potassium ferrocyanide. This ceU was then placed in a solution of copper sulphate, and when the diffused liquids met, a precipitate of copper ferro- cyanide, strongly supported by the walls of the cell, was formed. Nowadays it is the custom to aid the diffusion by means of the electric current, thereby hastening the formation of the mem- brane. Such a chemical membrane is readily permeable to 399 400 AN INORGANIC CHEMISTRY Manometer V water, but opposes a nearly perfect barrier to the passage of dissolved susbtances. Osmotic Peessuee PfefEer, the pioneer investigator in this branch of ^^•ork, placed a sugar solution of known strength within such a cell, a mano- meter was attached, and the whole placed in a vessel containing pure water (Fig. 107). The manometer indicated a slo^^ but steady increase in pres- sure mthin the cell. Water had evidently passed through the membrane into the cell. When the hydrostatic pressure exerted downwards was exactly equal to the tendency of the water to wander inwards, equiUbrium \\"as reached, and the solution exerted its maximum osmotic pressure. This equihbrium is of a purely dynamic nature, and is reached when the number of molecules of solvent passing through the mem- brane in the one direction per unit of time is exactly equal to the number traversing the membrane in the other direction. The wan- dering of the solvent through the membrane is knovm as Osmosis. Many theories as to what constitutes the osmotic pressure have been advanced — an indication in itself that we do not yet understand the mechanism of the phenomenon. Van't HofE put forward the suggestion that the pressure arises from the bom- bardment of the membrane by the dissolved particles in their endeavour to diffuse into the solvent and thus make a solution of uniform concentration. Undoubtedly both solvent and solute molecules are in a violent state of motion, as a result of which uniformity of concentration would be attained, were it not for the intervening membrane. In the apparatus above described the dilution of the sugar solution can only be effected by the entrance of the solvent through the semi-permeable membrane. —Sugar Solution. -Water. ~Porous Wall. Fig. 107. OSMOTIC PRESSURE 401 It is this striving on the part of the solution to attain a uniform concentration which enables us to measure with the above apparatus the osmotic pressure, i.e. the pressute of the solution in excess of that ruhng in the solvent itself. If a porous cell, prepared as described above, and containing a sugar solution, is placed in a corresponding solution which is more concentrated, it is found that solvent wanders out of the porous cell, i.e. there is migration of the solvent from the more dilute to the more concentrated solution. When the solution within and without the cell shows no such migration of solvent, they are said to be isotonic, i.e. of equal osmotic pressure. One of the earliest methods of comparing the osmotic pressure of solutions lay in the use of vegetable cells. Attached to the outer walls of such cells are vegetable protoplasmic tissues, forming a semi-permeable membrane, through which the various salts present in the sap are unable to pass. If such a vegetable cell is placed in a solution more concentrated than the vegetable sap, the protoplasmic layer is seen to shrink away from the cell wall owing to the extraction of water from the sap. On the other hand, if such a cell is placed in a solution more dilute than the sap liquid, water passes through the vegetable wall into the cell, and the whole cell becomes distended. When placed in an isotonic solution, no movement of the protoplasmic layer occurs . Occasionally blood cells (red corpuscles) have been used in place of the vegetable ceUs. Their action under the play of osmotic forces is similar. In solutions more dilute than the blood fluid the cell bursts, thereby colouring the hquid red, whilst in more concentrated solutions water is extracted from the cell. This causes a rise in the specific gravity of the cell contents and the cell sinks. The pain produced when a badly blistered hand is washed in water is another illustration of osmotic phenomenon — water diffuses through the animal mem- brane and the pressure within rises, with consequent pain to the sufferer. Quantitative Aspect of Osmotic Pressure. A large number of experiments were carried out by Pfeffer wherein he measured the osmotic pressure of solutions of varying concentration and temperature, though it was van't HofI (1887) who first successfully correlated the laws governing osmotic pressure discussed below. His deductions were based partly UD 402 AN INORGANIC CHEMISTRY upon the observations of Pfefier, and partly upon his own experiments. The Dependence of Osmotic Pressure upon the Con- centration. — The effect of concentration upon the osmotic pressure of a solution is clearly brought out in Table 41, which embodies some of the recent results of Morse and Frazer. TABLE 41 Osmotic pressure in atmosplieres of sugar solution. observed. 05 1 -25 010 . 2-44 0-20 4-80 0-30 . 7-23 0-50 . 12-08 These results were obtained at approximately 20°. It is seen that the osmotic pressure is directly proportional to the con- centration, i.e. P=A;.c, but the concentration of a solution is inversely proportional to the volume of that solution ; hence k P= — or PV=ifc. Boyle's Law for gases states that the product ^t; is a constant, so that we may say with van't HofiE that Boyle's Law holds not only for gases but also for solutions, provided that, in the case of solutions, the pressure is understood to be the osmotic pressure. Effect of Temperature upon the Osmotic Pressure. — Charles' Law, correlating the influence of temperature upon the pressure (or volume) of a gas may be expressed in the equation Pi=p^ (l-\-at) {see p. 63), where p^, p„, denote pressures at the temperature i° and 0°, and a is the coefficient of expansion. The adjoined Table 42 shows the variation in the observed value of the osmotic pressure of a 1 per cent, sugar solution when the temperature is changed. TABLE 42 Temperature. Observed pressure. Calculated pressure. 0° 0-649 atmospheres _ 6-8 0-664 0-665 13-7, 0-691 0-681 14-2 0-671 0-682 15-5 0-684 0-686 32 716 0-725 OSMOTIC PRESSURE 403 The values in the last column are calculated from the equation P(=Po (l+ai), where P^, P^ denote the osmotic pressures at t° and 0° respectively and a is the usual coefficient of expansion . The agreement between the found and calculated values of the osmotic pressure is distinctly good, hence Charles' Law may also be applied to solutions. The osmotic pressure of a solution is therefore proportional to the absolute temperature {see Charles' Law, p. 63). The General Gas Equation applied to Solutions. — In Chapter V, dealing with the physical properties of gases, it has aheady been shown that a consideration of Boyle's and Charles' Laws leads to the conclusion that the effect of a combined change of -temperature and pressure upon the volume of a gas is ex- pressed in the equation pv=^T. It is now obvious that a similar method of deduction must lead to the conclusion that PV=RT, where P denotes the osmotic pressure of a solution of volume V, R and T having their usual significations. If a 1 per cent, solution of sugar (1 gm. of sugar dissolved in 99 gm. of water) is prepared, the resultant volume at 0° is 99-7 c.c. The molecular weight of sugar is 342, hence the volume of such a solution which wiU exactly contain 342 gm. of sugar wiU be 99-7x342 c.c. = 34-1 htres. The osmotic pressure of such a solution at 0° has been found to be 0-649 atmospheres. If the question is now investigated as to what pressure the gram molecular quantity of a gas will exert if it fills the volume 34-1 litres at 0°, it follows from the gas equation, ^«;=RT, that RT 0-0821x273 ^^^^ , p= — = sj^j =0-657 atmospheres. The osmotic pres- sure exerted by a dissolved substance is therefore identical in value with the gaseous pressure which that substance would exert, if the same weight of substance were in a state of vapour, under the same conditions of temperature and pressure. To exemplify. — ^A given weight of ethyl alcohol, CaHjOH (M.W. = 46), when dissolved in water to form a solution of volume V, wiU exert an osmotic pressure P. Imagine now that the solvent molecules (water) are removed, and that the molecules of alcohol continue to fill the same volume as heretofore, the temperature remaining unaltered. Then the gaseous pressure which the alcohol will exert as a 404 AN INORGANIC CHEMISTRY vapour is exactly equal to the osmotic pressure which it formerly exerted when dissolved in the water. Osmotic Pressure as a Means of Determining Molecular Weight. — Having established that the gas laws are appUcable to the study of dilute solutions, van't HofE was led to the still more important conclusion that the great generalisation of Avogadro — ^that equal volumes of all gases under the same conditions of temperature and pressure contain the same number of molecules — was equally true for dilute solutions, hence equal volumes of all solutions, which exert equal osmotic pres- sures and are at the same temperature, contain an equal number of dissolved molecules. The importance of Avogadro's Law as apphed to gases has already been stressed (p. 124), for upon it is erected the whole modern conception of molecular weights and their determination. To recapitulate, the molecular weight of the standard gas, oxygen, has been chosen as 32, and the volume (22-4 litres), which exactly holds 32 gm. of oxygen at 0° and 760 mm. pres- sure, has been chosen as the gram molecular volume. In order to determine the gram molecular weight of another gas it is merely necessary to know the weight w occupying the volume v at the temperature t and the pressure p, and we are then in a position to calculate by proportion the weight W of the gas which will occupy the volume 22-4 litres at 0° and 760 mm. pressure. This gives the gram molecular weight of the gas under con- sideration. But since Avogadro's Law has been shown to be rigidly apphcable to dUute solutions, it follows as a necessary conse- quence that, if we know the weight of a substance dissolved in a volume V of a solvent at a temperature t°, and if we measure the osmotic pressure P exerted by such a solution, the gram molecular weight of a substance dissolved in 22-4 litres and exerting an osmotic pressure of 1 atmosphere (760 mm.) at 0° can then be calculated by the same method. Example. — An aqueous solution containing 2-0094 gm. of boric acid per Utre exerts at 10° an osmotic pressure of 0-724 atmospheres. In order to determine the molecular weight from these results it is necessary to find what weight of boric acid, dissolved in 22-4 htres, will exert an osmotic pressure of 1 atmo- sphere at 0°. The volume of 1 litre at a pressure of 0-724 atmosphere and OSMOTIC PRESSURE 405 at a temperature of 10°, when reduced to standard conditions, 273 Since 2-0094 gm. occupy 0'698 litres, it follows that the gram molecular volume, 22-4 becomes 1 X 0-724 X---= 0-698 litre. 283 litres, will be filled by 2-0094 X 22-4 =64-4 gm. The molecular 0-698 weight calculated from the formula HBO3, is 62. There is, however, one handicap with this method of obtaining the molecular weight of dissolved substances — the extreme difficulty of measuring the osmotic pressure accurately, so that the method would Water Barometer To condenser. 1% Sugar solution. appear to be of theo- retical rather than of practical importance, were it not for the fact that there are other properties besides the osmotic pressure which are proportional to the concentration or the number of diss olved molecules, e.g. the lowering of the vapour pressure, the rise of the boiUng point, and the lowering of the freezing point. If a series of baro- meter tubes is filled with mercury at some constant temperature, and into the second is introduced a drop of water, into the third a few drops of a I per cent, solution of sugar, into the fourth a few drops of a 2 per cent, solution of sugar, into a fifth a few drops of a 4 per cent, solution, and so on, it will be found that the vapour pressures recorded in the various tubes (i.e. the difference in height between the mercury Pio. 108. 406 AN INORGANIC CHEMISTRY levels in the standard tube and the various tubes containing fluid) are not the same, but steadily diminish with the concen- tration of the solution. This is illustrated in Fig. 108. The vapour pressure of water exceeds that of the 1 per cent, sugar solution by x cm. of mercttry, the 2 per cent, solution by 2a; cm., the 4 per cent, solution by 4.r, and so on, that is, the lowering of the vapour pressure is proportional to the concen- tration of the dissolved substance. This fact has an important bearing upon boihng point and freezing point determinations. If the vapour pressure curves of a pure solvent (e.g. water) and of solutions I (1 per cent.), II (2 per cent.) are graphed, Fig. 109 is obtained. The horizontal hne ABCD represents the 1 1 1 1 1 1 A — — Id \m s/x Wl e — — — — — — — f /_ — A L -- 1 b ... / / / / / / / / V\ / / / % / / / to / / / / Ci. \p A \ , / / f ^ N y / .§ / Soji^ ^\ y X y y S"' f / ^ iy ^i ^' N ^ ^ / '•^ AS f / Temperature. Fig. 109. pressure of 1 atmosphere, and since the boiling point of a sub- stance is the temperature at which the vapour pressiu:e is equal to the pressure of 1 atmosphere, the points B, C, D wiU repre- sent the boihng points of the pure solvent and of solutions I and 11 respectively. BC and BD denote the rise in the boiling point of the solutions. It has been found experimentally that this rise in the boiling point of the solution is proportional to the concen- tration, i.e. to the number of dissolved molecules. If such phenomena as supercoohng are excluded, the vapour curve of the pure solvent terminates at the point X, the point OSMOTIC PRESSURE 407 where the liquid-vapour curve is cut by the soUd-vapour curve XMR. It is obvious that the vapour curve of solution I will cut the sohd-vapour curve XMR at a temperature below X, that is, the freezing point of solution I is M, and MN denotes the lowering of the freezing point caused by dissolving 1 per cent, of sugar in the solution ; and in the same way RP denotes the lowering of the freezing point for solution II containing 2 per cent . of sugar. It has been experimentally shown that the lowering of the freezing 'point is proportional to the concentration, provided the solutions are dilute. For both the lowering of the freezing point and the rise of the boihng point it follows that A =k.c when A is the lowering of the freezing point or the rise of the boihng point and c is the concentration. The concentration may be defined as the ratio of the number of dissolved molecules n to the number of the molecules of the solvent N, so that i^=^k.^^ ; but n=—- js W where w is the weight of the solute (dissolved substance) and W its molecular weight, while N=— where s is the weight of the S solvent and S its molecular weight. w/W Hence A.=Jc. „ . In order to determine k a substance is s/S chosen, the molecular weight of which is known by an independent method, such as the vapour density method. For a solvent such as water, alcohol would serve to enable the value of k to be determined. In actual practice k denotes the depression pro- duced by the solution of 1 gm. molecular weight in 1000 gms. of solvent, and, as under these conditions S is a constant, it may be included in k ; thus, wt. of solute A=Kx mol. wt. of solute x wt. of solvent' A careful determination of the change in the freezing point or the boihng point of a solvent, produced by the addition of a solute, is therefore sufficient to enable the unknown molecular weight of the solute to be evaluated, provided that K has been fixed by an independent experiment with a standard substance. Fig. 110 illustrates the Beckmann type of freezing point apparatus in general use. The apparatus itself consists of an 408 AN INORGANIC CHEMISTRY outer vessel C for holding the freezing mixture and an inner vessel A carrying a stirrer and a very delicate thermometer D of special design. This inner vessel is protected from direct contact with the freezing mixture by an air jacket B. A weighed quantity of the solvent is introduced into the inner vessel and steadily stirred. The temperature slowly falls until the separation of the solid, i.e. freezing, sets in. The tem- perature then becomes constant. Some liquids, e.g. water, are liable to supercool. In such cases it is usual to add an occa- sional spicule of the soUd, and if super- cooHng has occurred, the temperature will instantly rise as crystaUisation takes place . When a series of concordant results for the freezing point has been obtained, Fig. 110. Fig. 111. a weighed pellet of the substance, the molecular weight of which is required, is introduced and the freezing point redetermined. For further details of manipulation the student must consult a textbook of practical physical chemistry. Several types of apparatus are in use for the determination of the boihng point of solutions ; probably that of Beckmann is still the most important (Fig. 111). OSMOTIC PRESSURE 409 C is an asbestos box provided with two chimneys which serve to convey the hot air from underneath the box. The inner tube A is provided with a piece of heavy platinum wire, and either platinum points or garnets to prevent super- heating of the liquid. The outer jacket B is also fiUed with the solvent, and serves not only to prevent radiation losses but also to heat the inner vessel. The greater difficulty of this method as compared with the freezing point method Ues in the tendency towards temperature fluctuations arising from draughts, and direct heating of the inner vessel, as well as to barometric alterations during the course of the experiment. Molecular Weights of Dissolved Substances. — In general the molecular weights of dissolved substances as determined by the above methods show a satisfactory agreement with the values obtained by other methods, e.g. the vapour density method. This is especially true for the majority of the carbon compounds, e.g. sugar, chloroform, alcohol, acetic acid and so on, but conspicuous exceptions exist, where the molecular weights of the dissolved substances, as obtained either from measiu-e- ments of the freezing point, of the boiUng point or of the osmotic pressure, differ strikingly from the molecular weights which our knowledge of the composition of these substances has led us to expect them to have. Nearly all inorganic acids, bases and salts give an abnormally high value for the boihng point, rise and an abnormally large depression for the freezing point, and conse- quently an abnormally low value for the molecular weight as calculated from these values. One- tenth of a gram molecule of ethyl alcohol (4-6 gm.) dissolved in a litre of water gives a depres- sion of 0'189°, whilst one-tenth of a gram molecule of sodium hydroxide (4 gm.) dissolved in the same volume of water, gives a depression of 0-340° ; the depression is therefore double what one would expect. This is true for all mono-acid bases like KOH, for aU monobasic acids, e.g. HCl, and for salts formed by the interaction of such acids and bases. Furthermore, it was found that dibasic acids and salts formed from them showed an even greater discrepancy in very dilute aqueous solution. In aU these cases, however, the divergence between experiment and theory was less marked in stronger solutions. Many of the above substances which behave abnormally when dissolved in water give perfectly normal values if such solvents as benzene, chloroform, etc., are used. 410 AN INORGANIC CHEMISTRY Dissociation in Solution. — ^Attention has already been called to the fact that the vapour density of nitrogen tetroxide steadily diminishes with rising temperature (p. 297). This has been traced to the dissociation or breaking down of the molecules of the tetroxide in accordance with the equation N,0, ^^ 2NO2. Other such cases of dissociation in gaseous systems are known, PCl5^±PCl3 + CI, 12^=^21 and m all cases the view long held is that the decreased density arises from the breaking down of the complex molecules into a greater number of lighter and simpler molecules. In 1887 Arrhenius put forward the plausible hypothesis that the abnormal values of the molecular weight, etc., shown by acids, bases and salts, when dissolved in water, arose from the dissociation of the molecules of these substances into a greater number of simpler parts. An increase in the number of dis- solved molecules (whether undissociated or the result of disso- ciation) must cause an increased effect upon the osmotic pressure, etc., and the results of experiment receive a sufficiently satis factory explanation. The question now arose as to how such a molecule as NaCl could be broken down in aqueous solution. Arrhenius then observed that it is precisely those substances which show abnor- mal behaviour when dissolved in water that form conducting solutions. Aqueous solutions formed from the great mass of substances which behave normally, e.g. the alcohols, sugar, organic acids and bases, do not form conducting solutions, while an aqueous solution of copper sulphate, hydrochloric acid or sodium hydroxide is an excellent conductor of the electric current. Arrhenius was thus led to put forward the hypo- thesis that inorganic acids, bases and salts dissociated in aqueous solution into charged parts (ions), and the extent of this dissociation or ionisation increased with dilution. The mechanism of this dissociation is indicated in the following equations : NaCl ^r±Na+-fCl- H,S04 ^::±H+-fHSO,- ;z:^2H+-[-S0r AICI3 ==±Al+ + -^ + 3Cl- DISSOCIATION IN SOLUTION 411 Na^COa ;=± 2Na+ +CO3- NaOH ^±Na+ + OH- KNO3 ^z:±K+ + N03- It will be noted that in all cases the sum of the positive charges exactly equals the sum of the negative charges. The metals, in all but certain exceptional cases, carry a positive charge, the number of such charges being equal to the valence of the element under consideration. The negative charge is carried by the acid radicle or element. Nomenclature. — A solution which conducts the electric current and thereby suffers decomposition is spoken of as an electrolyte. The ions, which are the means by which the electric current finds its way through the solution are of two kinds — ^the caihion, which carry the positive charge or charges, and during electrolysis drift towards the cathode or negative pole, and the anion, which carry the negative charge or charges, and which under similar conditions drift towards the anode or positive pole. The most common cathions are the metals and hydrogen, the anions are the hydroxyl ion OH ~, and the acid radicles or groups e.g. S04=, POr, COj", C1-, Br-, etc. Objections to the Ionic Theory. — At first this theory met with considerable opposition, but the objections were soon dis- posed of. To the minds of many of the older school of chemists, the revolutionary idea that in an aqueous solution there should exist, side by side as it were, substances which possess such an extreme affinity for each other as do chlorine and sodium, was beyond behef, until it was clearly emphasised that the theory does not postulate the existence in solution of free chlorine and free sodium, in short, the properties of the ions formed from chlorine and sodium, etc., differed entirely from the properties of the elements themselves. There was no justification whatso- ever for the view that atoms charged with electricity, i.e. ions, should possess the same properties as the free elements. Copper, the metal, has a characteristic reddish appearance, whilst the cupric ion is blue. Chlorine, as a gas, is a yellow, poisonous gas, but when the atom of chlorine has associated with it negative electricity, the chlorine ion so formed is colourless, and solutions containing such ions are, in general, non-poisonous. Whenever 412 AN INORGANIC CHEMISTRY solutions of chlorides are poisonous, e.g. mercuric chloride, the deleterious property is associated with the cathion. Again, it was pointed out in criticism of the ionic theory that the abnormal behaviour of salts in aqueous solution could be well explained by the supposition of a far-reaching hydrolysis, that is decomposition into free acid and base under the agency of the water, e.g. NaCl + HOH ^=± NaOH + HCl. But although such an assumption would, if true, explain in a satisfactory way the abnormal freezing and boiling point values given by solutes such as the inorganic salts, the assumption breaks down entirely when appUed to substances such as the inorganic acids and bases, which likewise behave abnormally so far as the osmotic phenomena are concerned. No such decomposition is possible for such bodies as HCl, HNO3, NaOH. Arguments in Favour of the Theory. — Perhaps no more striking evidence in favour of the ionic theory has been put forward than that afforded by the study of the heat of neutral- isation of acids by bases. Whenever a gram equivalent of an acid is neutralised in dilute solution by the equivalent quantity of a base, the heat of reaction is the same, 13,700 cals. NaOH + HCl ->-NaCl + H^O + 13,700 cals. iH^SOi + NaOH-^ INa^SO^ + H ^0 + 13,700 cals. HNO3 + iCa(0H)2-^ |Ca(N03), + H^O + 13,700 cals. In our study of chemical phenomena (Chapter II) we have seen that the outstanding criterion of a chemical reaction is that, when substances react with each other, a definite energy change occurs, a change which is shown in the evolution or absorption of heat, light, etc. This change of energy represents the difference between the energy of the reactants and of the resultants of the reaction, and for fixed quantities of the reacting substances, is a fixed, unalterable magnitude. The inference to be drawn from a study of the above equations is that in all these apparently diverse reactions the same chemical reaction is in reality taking place. Consider the equihbria present in the dilute solutions before mixing. NaOH ^^ Na+ + OH" HCl =i±H++Cl- DISSOCIATION IN SOLUTION 413 On mixing, we have Na+ + OH- + H+ + CI- -^Na ' + 01- + H^O + 13,700 cals. the salt formed remaining dissociated. On eliminating from each side the common terms, the equation reduces to the fol- lowing : H+ +H0-->H20 + 13,700 cals. And so, too, if nitric acid is neutralised by potassium hydroxide. H+ -f NO3- + K+ + OH- ->K+ + N03- + H20 + 13,700 cals. H+ + OH- ->H20 + 13,700 cals. This constancy of the heat evolution when equivalent quantities of acid and base react, is therefore a necessary corollary of the ionic hypothesis. Another interesting case is the reaction in dilute solution between the reciprocal pairs KNO3, NaCl and NaNOg, KCl. No chemical or physical test enables one to distinguish between the solutions obtained by mixing the first pair of salts from the solution given by the second pair. Here, again, an adequate explanation of the experimental facts is afforded by the ionic hypothesis. In dilute solution we have KN03^K++N0r NaCl^Na+ + Cl- summing, KNO3 +NaCl^K+ + N03- + Na+ + CI- and so also NaNOa— >Na+ + N03- KCl -^K++C1- summing, NaNOj +KCl^^Na+ + N03- + Cl- + K+. Whether we mix dilute solutions of potassium nitrate and sodium chloride, or of potassium chloride and sodium nitrate, it is evident that in both cases we shall obtain a solution containing salts completely dissociated, i.e. the final solution will contain the ions Na + , K+, CI", NO3-, and will be in every way indis- tinguishable. The observation has also been made that the permanganates all give the same coloured solutions. When the absorption spectra (see p. 490) of the solutions of the permanganates, e.g. KMnOi, NaMnOiNHjMnOi, are compared, they are found to be identical. This is not to be expected frsm the assumption that the molecules of the permanganates persist in solution, but is to 414 AN INORGANIC CHEMISTRY be expected on the hypothesis that in aqueous solution all permanganates dissociate into the purphsh coloured perman- ganate ion, which exerts its own specific effect on the absorption spectrum, whilst the colourless cathion exerts no such action. Questions 1. What gaseoua pressure would be exerted by methyl alcohol vapour of the same molecular concentration as a 2 per cent, solution of this substance in a non-volatile solvent at 50° ? 2. Show how the laws which govern the behaviour of gases under vary- ing conditions of temperature and pressure also apply to certain solutions. 3. If 1 gm. of a substance, dissolved in 100 c.c. of water, gives 0-1° depression of the freezing point, what is the molecular weight of the substance ? (the molecular lowering of the freezing point for water = 1 -86°). 4. 3 gm. of a substance dissolved in 50 gm. of water gave a rise in the osmotic pressure of 9 atmospheres at 22°. Find the molecular weight of the substance. 5. What depression of the freezing point of water would be produced by dissolving 5 gm. of sugar (CgHj^Oa) in 150 gm. of water ? 6. 0-394 gm. of a substance was dissolved in 1805 gm. of benzene, thereby producing a depression of the freezing point of the value 0-348''. Find the molecular weight of the substance. The molecular lowering for benzene is 5. 7. 2-01 gm. of a substance are dissolved in 50 gm. of ether. The rise in the boiling point is 0-798. Find the molecular weight of the solute, the molecular rise of the boiling point for ether being 30-3. 8. The depression of the freezing point for sodium chloride in water is much greater than that given by an equivalent quantity of sugar, i.e. by a sugar solution of the same concentration. How do you account for this ? 9. Give a succinct account of Arrhenius' theory of dissociation. CHAPTER XXVIII CONDUCTIVITY AND ITS BEARING UPON THE IONIC THEORY The Electrolysis of Salt Solutions. — Before considering this subject from its quantitative aspect, let us investigate the chemical changes produced when solutions containing various solutes are subjected to electrolysis. If two platinum plates are connected through an ammeter to the terminals of a battery, and then dipped into a solution of sugar, there is no measurable deflection of the ammeter needle ; in short, a solution of sugar does not conduct the electric current. A similar result is obtained if pure water is placed in the beaker in place of the sugar solution ; no current is able to find its way through the water, i.e. the water is a non-conductor, or non- electrolyte. If the electrodes are dipped into concentrated sulphuric acid, there is a small but definite deflection, indicating the passage of a weak current through the circuit. If, how- 416 AN INORGANIC CHEMISTRY ever, the sulphuric acid is poured into the water and the contents thoroughly mixed, the ammeter indicates the passage of a strong current through the solution, and a brisk evolution of gas takes place at the electrodes (Fig. 112). An examination of these gases reveals that they are hydrogen and oxygen, the former being liberated at the cathode, the latter at the anode. The non- conductivity of concentrated sulphuric acid and of water and the ease with which the electric current can find its way through a solution of dilute sulphuric acid leads us to conclude that, whilst the concentrated acid is almost free of ions, the dilute acid is rich in these carriers of electricity. In other words, when the sulphuric acid is dissolved in water, dissociation occurs in accordance with the equation H2S04^i±H+ + HS04- ^=±2H+ + S04-. The positively charged hydrogen ion drifts off towards the cathode (negative pole), there to give up its charge and to become elementary hydrogen, whilst the sulphate ion carries its negative charge to the anode (positive pole). At the cathode the hydrogen atoms immediately pass into the molecular state, and escape. At the anode the SO4" ions have their charge neutralised, but, whilst theSOj" ion in the presence of water is perfectly stable, so soon as its charge is neutralised, reaction takes place thus. Sulphuric acid, therefore, accumulates round the anode — a fact readily capable of proof, and the oxygen escapes in the mole- cular form. The results given by a solution of sulphuric acid indicate the presence of hydrogen and sulphate ions. During the electrolysis of a solution of copper nitrate, Cu(N03)2 between platinum plates, a red deposit of copper separates on the cathode, at the anode the solution becomes increasingly acid and oxygen escapes. Copper ions, therefore, migrate towards the oppositely charged cathode, and when their charge is neu- trahsed, the metalUo copper appears as a red film on the cathode. On the other hand, the negatively charged nitrate ions drift off towards the anode, and after becoming electrically neutral, the reaction 4NO3 + 2H2O -> 4HNO3 + O2 occurs. This is evidence in favour of the presence of cupric and nitrate ions in an aqueous solution of copper nitrate. THE IONIC THEORY 417 An interesting point in such an electrolysis is that the minutest current is able to find its way through an aqueous solution of a dissolved acid, base or salt. If the current were necessary to effect the actual decomposition of the dissolved salt, as was once maintained by Grotthus, it would be expected that each solute would require a definite electromotive force before decomposi- tion sets in, i.e. the current would be unable to find its way through the solution until the E.M.P. of the battery was suffi- ciently high to decompose the dissolved salt. [The quantities E.M.F., current, and resistance (the reciprocal of the conducti- "P TIT "pi vity) are connected by Ohm's Law by the equation C=-^ — ^^; for further details concerning this equation the student must consult a text-book of Physics.] If a solution of sodium nitrate, coloured by a little litmus, is electrolysed, it is found that the solution round the cathode becomes alkahne, and round the anode acid. Moreover, hydro- gen and oxygen escape at these poles respectively. In this case, secondary reactions at the electrodes take place. As soon as the sodium ion passes into atomic sodium, reaction with the solvent takes place, and hydrogen is liberated with the formation of sodium hydroxide. At the anode the reaction 4NO3 -f 2H20^^4HN03 + O2 leads to the accumulation of nitric acid round the pole and the evolution of oxygen. The phenomenon of electrolysis is therefore associated with far-reaching decomposition of the dissolved matter. In general, one may say that the electrolysis of a solution of a dissolved salt causes the migration of the metal ion towards the cathode, and the migration of the acid group towards the anode. Whether the metal will be deposited at the cathode depends upon the chemical nature of the metal, i.e. upon its tendency to react with the water after it has passed from the ionic to the atomic state. Acids appear always to Uberate hydrogen from the cathode during the electrolysis ; the acid group may be hberated and escape as such (see electrolysis of hydrochloric acid), or there may be a secondary reaction at the anode leading to the escape of oxygen, and the re-formation of the acid as in the case of nitric and sulphuric acids. It remains now to consider the quantitative aspect of electrolysis. BE 418 AN INORGANIC CHEMISTRY Quantitative Aspect of Electrolysis. — The conductivity of a Solution is dependent upon three factors : 1. The charge or load carried by each ion. 2. The rate at which these ions move through the solution. 3. The number of such ions present in the solution. It is proposed to consider these factors separately. The Charge caeeied by the Ions — Faeaday's Laws The first investigation concerning the quantitative aspect of electrolytic decomposition was undertaken by Earaday (1834). He showed with extreme exactness that the amount of any one substance liberated at an electrode is strictly propor- tional to the quantity of electricity w^hich has passed through the cell. This is Faraday's first law, governing electrolytic decomposition. The amount of copper or of silver deposited on a platinum cathode must therefore be a direct measure of the quantity of electricity which has passed through the circuit, hence the use of the silver and copper voltameters as measures of current. The actual amount of decomposition effected by the passage of a given quantity of electricity — disturbing secondary reactions excepted — is thus independent of the temperature and concentration of the solution, as well as of the current density (current per sq. cm.). Nothing else matters but the quantity of electricity which has passed through the circuit. But if the same current is allowed to flow through a series of solutions arranged as in Fig. 113 it will be found rfTDnri oTsinri ^firr^ (''^'Trt-^ ii ("iiziii pEzi (?idai fPi AgNOs H25O4 NaCl FeClj CuSO^ Fig. 113. that the amount of decomposition in each cell is strictly proportional to the chemipal equivalent of the element, that THE IONIC THEORY 419 is, the current which would liberate 1-008 gm. of hydrogen either from sulphuric or hydrochloric acid, would also effect the decomposition of 23 gm. of sodium, 56/2=28 gm. of iron in a solution of ferrous sulphate, 66/3=18-6 gm. of iron in a solution of ferric chloride, 63-6/2=31-8 gm. of copper, 107-9 gm. of silver, 27-1/3=9-03 gm. of aluminium. Chemically equivalent quantities of the Elements are liberated by the passage of the same quantity of electricity. This second law of Faraday incidentally affords an excellent method of determining the chemical equivalent of an element, or what may be called its electro-chemical equivalent. The Charge carried by an Ion. — Faraday measured with great care the quantity of electricity required to Uberate 1 gm. equivalent of hydrogen and other elements. For all monovalent elements he found that 96,540 coulombs * were necessary to effect the hberation of 1 gm. equivalent of a mono- valent element, from which fact the conclusion may be drawn that every monovalent gram ion has associated with it 96,540 coulombs. Furthermore, for the hberation of 1 gm. equi- valent of a divalent element (e.g. copper -— - =31-8 gm.) the same amount of electricity is also necessary, so that the setting free of a gram ion of such an element as copper requires 96,540 X2 coulombs, a trivalent element hke aluminium needs 96,540x3 coulombs and so on. This " parcel of electricity ", 96,540 coulombs, which is transported by each gram ion of a mono- valent element is often known as a Faraday and is often denoted by © or 0. The intimate connection between the valence of an element and the number of Faradays associated with the gram ion of that element is brought out in the follow- ing equations : " A coulomb, the unit quantity of electricity, will deposit 1/96540 gm. of hydrogen, 31 -8/96040 gm. of copper. When one coulomb passes through the circuit per second, the current is said to have a strength of one ampere. A current of 10 amperes will therefore liberate x 10 gm. of hydrogen per second. Often the term Current Density is used. This denote§ the iiumber of amperes per sq. cm. passing through the circuit, 420 AN INORGANIC CHEMISTRY H3PO4- NaCl : CuSOi; FeCla SnCl^ :H + + HjP04 ±Na-^+Cl- ^Cn+ + + S04- i±Fe+ + + + 3Cl- i±Sn- + + + + 4Cl- r±2H+ + HP04-; :3H++P04 H^SO. :H- + HSOr ^=±2H + + S04- Migration of the Ions. — The second factor to be taken into consideration in estimating the conductivity of a solution is the speed with which the ions move. The actual migration of coloured ions can be made evident in the following experiments. A moderately dilute solution of such a salt as copper sulphate, cobalt nitrate, potassium chromate, etc., is prepared, containmg also a few per cent, of agar- agar. This is introduced into a U-tube (Fig 114). After the jelly has set, the position of the coloured boun- dary is marked in each limb and a dilute solution of a suit- able colourless salt — in this case potassium sulphate — then introduced. The jelly forms a sponge- like network through which ions may travel with- out hindrance, but diffusion currents are prevented. It is advisable to keep the vessel cool during the electrolysis, otherwise the heat generated may melt the agar-agar. Plati- num electrodes are inserted into each hmb and a cmrent of about 40-80 miUiamps. passed through the cell. In the electrolysis of a solution of copper sulphate the blue boundary is seen to move towards the cathode in one limb, and away from the anode in the other. Consequently there is migration of the cupric ions towards the cathode, thereby estabUshing the fact that these ions carry a positive charge. If potassium permanganate were used in the agar-agar solu- tion, the purple boundary would be found to move towards the Fig. 114. THE IONIC THEORY 421 anode, proving clearly that in this case the coloured ion is associated with a negative charge. Accurate measurements of the velocity of migration of these ions have been made by modifying the above apparatus. Fig. 115 shows an apparatus suitable for a quantitative measurement of the speed of migration. EMr:2i^z2^z2Z2n: ^Je lly ■v////,'///////////////m. Fig. 115. It is by such means that the actual velocities of migration of the ions have been measured. Clearly, since the impelHng force is the attraction between the charged ions and the oppositely charged electrodes, the velocity of migration must depend upon the potential gradient between the electrodes. For unit poten- tial difference,* 1 volt per cm., the absolute velocities of some of the ions, expressed in cms. per hour, are K + Na+ Cu+ + H + OH- ci- NO,- 2-05 1-26 1-6 10-8 5-6 212 1-21 The velocity of migration of the hydrogen ion far exceeds that of the other elements ; next in order comes the hydroxyl ion. Probably the most important factor in determining the speed with which an ion moves is its size — the bigger the ion, the greater the frictional resistance it will encounter in its journey towards the electrodes. Solutions containing the more rapidly moving ions, ceteris paribus, will clearly show the greater conductivity. The Number of Ions. — The third factor controlling the conductivity of a solution is the number of ions present per c.c. * A volt is the unit . of potential difference or electromotive force. When a cvirrent of 1 coulomb per second passes through a resistance of 1 ohm, the potential difference between the ends of tlie resistance is 1 volt. 422 AN INORGANIC CHEMISTRY It remains now to see how one may obtain a measure of the number of such ions present. . Suppose a vessel be constructed, one pair of opposite sides being made of a metal suitable for an electrode. These are connected through a galvanometer to the terminals of a cell (Fig. 116). If pure water is poured into this electrolytic ceU, the galvano- meter shows no sign of a current ; if the pure water is replaced by a dilute solution of sulphuric acid, a strong current wiU pass through the ceU, indicating that the resistance of the electrolytic cell has decreased, i.e. the conductivity has increased. If more water is added to the cell, the current continues to increase ; the greater the dilution, the greater the conducting power of the solution. The conductivity continues to increase until a certain maximum value has been attained, and after this, no further dilution has any efEect upon the con- ductivity. It is clear that throughout the series of ex- periments the actual amount of sulphuric acid present between the electrodes has not altered. The explanation of the increased conductivity must, therefore, be sought in the presence of an increased number of carriers or ions in the dilute solution. In the terms of the ionic hypothesis the dissociation expressed in the equation FiQ. 116. :2H+-fS0i H2S0,^=±H+-f-HS0r has been driven to the right with increasing dilution, i.e. the number of conducting ions steadily increases with dilution up to a maximum. In the solution of maximum conductivity dissociation is complete ; further dilution can create no more ions, and there can be no further increase in the conductivity. Theoretically, dissociation wiU be complete at infinite dilution, but practically the dissociation for many acids, bases and salts is complete at N/1000, that is, in a solution which contains 1 gm. equivalent THE IONIC THEORY 423 of the salt dissolved in 1 ,000 litres. A gram equivalent of sodium chloride weighs 23+35-5=58-5 gm., for the acid (base) from which this salt is formed is monobasic (mono-acidic), but in the case of an acid, base or salt containing ions of more than one valence, due allowance must be made. Hence an equivalent solution of sulphuric acid, designated N, contains ~^ — =49 gm. per Li K, f ^ u <- 63-6+32+64 159-6 ^„ „ htre, 01 copper sulphate ^ =~s~ =79-8 gm. per htre. From a consideration of the results of experiments such as the above, Arrhenius put forward the conclusion that the degree of dissociation at any particular dilution is equal to the conductivity of that solution, divided by the conductivity at infinite solution, i.e. a= -' where a denotes the fraction /« = dissociated, fj.^ the conductivity at any arbitrary dilution v which contains 1 gm. equivalent of the salt (the so-called equi- valent conductivity), /n^ the conductivity at infinite dilution. In order to obtain the conductivity at infinite dilution ju^, for substances which dissociate readily, e.g. KCl, HNO3, H2SO4, etc., it is only necessary to determine the conductivity over a range of dilutions and then by extrapolation the value of may be read off the graph. But there exist a large number of sub- stances, such as acetic acid, ammonium hydroxide and so on, which do not form good conducting solutions, i.e. they do not dissociate freely, and even in the most dilute solutions with which we can work, their dissociation is so far removed from completion that such an extrapolation is impossible. In such cases another method is resorted to. The difference between ju ^ for hydrochloric acid and sodium choride must arise from the difference in the speed of migration of the hydrogen and the sodium ions, for these salts dissociate completely at moderately dilute concentrations. If one now determines the conductivity at infinite dilution of sodium acetate — a substance which dis- sociates freely — it is only necessary to add to /^^ for sodium acetate the difference aheady found between ^ ^ for hydrochloric acid and sodium chloride in order to obtain jlc^ for the hypo- the^ically completely dissociated acetic acid. The conductivity 424 AN INORGANIC CHEMISTRY at infinite dilution of all salts, acids and bases which are but weakly dissociated, has been determined by this means. It is interesting to note that the degree of dissociation derived from the equation, a =1X^1^1^, except for certain substances discussed later, agrees closely with the dissociation calculated from abnormal values of the osmotic pressure, lowering of freezing point, etc. This agreement forms one of the strongest quanti- tative supports of the ionic theory. Ostwald's Dilution Law. — It was Ostwald who first applied the Law of Mass Action to the equihbria existing in aqueous solutions. Let 1 gm. molecule of acetic acid be dissolved in the volume V htres of water, and let a denote the fraction ionised. The concentration of the undissociated acid will therefore be , and - will denote the concentration of the ions. HAcet. : 1— a 'H Acet." a V Applying the Law of Mass Action to this equiUbrium, we get (H + )(Acet.-) (HAcet.) =^-, whence --Ic. {\—a)v a, we have learnt, may be most readily measured by determining the ratio fi^/fi ^ ■ The following table for acetic acid gives the corresponding values of v and ^t^,. In the last column are the values of h, calculated from jx^ and each value of //,. TABLE 43 Acetic Acid at 25° V M„ a = 100 °" I 4-34 (I— a)D *• 8 1-193 0-00000180 16 610 1-673 0-00000179 32 8-65 2-380 0-00000182 64 12 09 3-33 0-00000179 256 23-82 6-56 00000180 1024 46-00 ^^= 364 12-66 00000177 THE IONIC. THEORY 425 The constant is a satisfactory one. Many hundred such dissociation constants have been determined, and the appU- cabiUty of the Law of Mass Action to such equiUbria brilhantly estabUshed. Abnormal Behaviour of Strong Electrolytes. — -Although the Dilution Law of Ostwald is obeyed with great exactness by a large number of compounds, many outstanding exceptions occur. Inorganic salts, strongly dissociated acids and bases do not obey this law, and the cause of this has long been a bone of contention among chemists. Amongst the suggestions ]3ut forward to explain this discrepancy, is the statement that a =CiL (Joes not hold for strongly dissociated electrolytes ; again, it has been suggested that a is not known with sufficient accuracy for such electrolytes. Possibly, when experiment reveals the precise reason of this want of agreement, it wiU be necessary to readjust and widen the ionic theory. It is certain that undis- sociated molecules play a more important part in chemical reactions than the original propounder of the theory was at first prepared to concede. The Dissociation of the Common Acids, Bases and Salts. — The degree of dissociation for weakly dissociated electrolytes is given by the expression a=-— ; and until experiment has shown that this expression faOs to give the degree of dissociation of strong electrolytes, we may take it that the dissociation of such electrolytes is given by a=— with sufficient accuracy. A determination of the conductivity at a particular dilu- tion V, coupled with the knowledge oi /i^, obtained by either of the methods outhned on p. 423, enables one to calculate a, the degree of dissociation of the substance at the chosen dilution v. In practice it has become the custom to work either with the gram equivalent or the gram molecular weight, or with definite fractions of these quantities. As an example, it may be stated that the equivalent conductivity of ammonium hydroxide at v=32 is 6-28, whilst the conductivity at infinite 6'28 N dilution is 237, hence a = ---=-026, i.e. an -^ solution of 426 AN INORGANIC CHEMISTRY ammonium hydroxide is 2-6 per cent, dissociated at 25°. NH4OH ; 97-4<%, :NH, + + OH- 2-6% 2-&%. Owing to the effect of temperature in lowering the viscosity of water, the conductivity of an electrolyte generally increases with rise in temperature, although it must not be overlooked that there may also be a change in the number of conducting ions. As a rule, however, a rise of temperature has compara- tively httle effect upon the degree of dissociation of an electro- lyte, the increase in the conductivity being almost entirely due to the increased speed of the ions. The fractional dissociation of the more commonly occurring acids, bases and salts is given in the subjoined table. Acids. Nitric Hydrochloric Sulphuric Sulphuric cone. Hydriodic Chloric Phosphoric Permanganic Carbonic Hydrogen sul phide . Acetic Oxalic Bases. Potassium hy- droxide Sodium hydroxide Lithium hydrox- ide Calcium hydrox- ide N N N/2 N/2, N/2 N/2 N/io: N/IO' N N/10 18= 18° 18° 18" 25° 25" 18" 18° 18° 25° N N 18° 18° N 18° N/64 TABLE 44 0-820 0-784 0-510 007 0-901 0-880 0-170 0-930 0-0017 0-0007 0-004 0-500 0-77 0-73 0-63 0-90 [contd.]. Ammoniiun hy- droxide .Ammonium hy- droxide Barium hydroxide Salts. Sodium chloride Ammonium chlor- ide Potassium nitrate Silver nitrate Copper sulphate Zinc sulphate Potassium sul- phate . Sodiiun sulphate Mercuric chloride, saturated Mercuric cyanide I N 18°; 0-004 N '10 N/64 X 18° N X X N X N X 1S = 18 = 18 = 18= 18° 18° 18° N/10' 18° 0-014 0-92 0-74 0-75 0-64 0-58 0-22 0-24 0-53 0-45 very small very small A general survey of this table reveals the fact that the mineral acids are strongly dissociated, but the degree of dissociation falls off as the basicity of the acid increases ; i.e. the dissociation of phosphoric acid is less than that of sulphuric acid, sulphuric less than nitric and so on. The same rule appUes to the dissociation of the salts. As a rule, salts of two monovalent ions, like NaCl, are under comparable conditions of concentration, as THE IONIC THEORY 427 strongly dissociated as hydrochloric acid and nitric acid. Salts of mono- and di-valent ions, such as CuCla, K2SO4, show less dissociation than do salts of two monovalent ions, whilst salts of two divalent ions (CUSO4) are stiU less dissociated. With the exception of the chloride and cyanide of mercury — salts which do obey the dilution law — and to a less extent the haUdes of cadmium, the salts are all strongly dissociated. The organic acids, as a class, are weakly dissociated, as is the base, ammonium hydroxide. Finally, the dissociation of carbonic acid and of hydrogen sulphide deserves notice. Both these acids are extremely weakly dissociated, a fact which has an important bearing upon the use of hydrogen sulphide as a reagent in chemical analysis. Abnormal Osmotic Phenomena and Conductivity. — It has already been emphasised that it is just those substances which give an abnormally high osmotic pressure, etc., that form conducting solutions. There is a close quantitative connection between these phenomena. If 1 gm. mole of a substance containing N molecules is dis- solved, and the fraction a is ionised, each molecule giving rise to m ions, there will be N[ma + (1 — a)] active particles in solu- tion, aU of them exerting a definite osmotic pressure. If i denotes the ratio of the molar concentration given by the osmotic method to that calculated by the chemical formula, it follows . N[»ia + (1— a)] , I— 1 T^ . .t. . that I = ^:; •", whence a= -. It is therefore N m — 1 possible to calculate i from a knowledge of the degree of dis- sociation of the different salts as given by conductivity data, and i may also be determined from the relation . abnormal lowering of fr. pt. etc., for a particular salt normal lowering of fr. pt. given by non-electrolyte, the solutions being of the same molecular concentration both for the conductivity and the freezing point measurements. In the following table such a comparison is made. The agreement shown in the last three columns affords strong support to the ionic theory. For salts like NaCl i lies between the limits 1 and 2, whilst for salts like K2SO4 it lies between 1 and 3. 428 AN INORGANIC CHEMISTRY TABLE 45 Values of i. Salts. Molecular concentra- tion. Osmotic Pressure. Freezing point. Conductivity. Potassium chloride . Calcium nitrate Lithium chloride . Magnesiiun sulphate . 014 0-18 013 o:!s 1-81 248 1-92 1 -2.") 1-86 2-47 1-94 1-20 1-86 2-46 1-84 1-35 Recapitulation The study of aqueous solutions has established the followng facts : 1. The osmotic pressure of solutions is proportional to the concentration of the dissolved substance (Boyle's Law). 2. The osmotic pressure of a solution is proportional to the absolute temperature (Charles' Law). -3. Under equal conditions of temperature solutions containing the same number of dissolved molecules per litre exert equal osmotic pressures (Avogadro's Law). Deductions from these facts. — Just as Avogadro's Law, appHed to gases, enables us to compare the molecular weights of gases, so also does the appUcability of this law to dilute solutions enable us to determine the molecular weight of dissolved sub- stances ; but, owing to the difficulty of meastiring the osmotic pressure, this method is of little practical importance. However, other properties such as the lowering of the vapour pressure and of the freezing point, the rise of the boiUng point, are likewise proportional to the concentration of the dissolved substance, and hence methods for the determination of the molecular weight o£ dissolved substances have been worked out, based upon the lowering of the freezing point, etc. Further facts. — 4. Many solutes, notably inorganic acids, bases and salts, give abnormally high results for the osmotic pressure and other allied phenomena, leading to the conclusion that these salts possess in aqueous solution a molecular weight which is smaller than what one would expect them to have from a con- sideration of the composition and atomic weights of the elements concerned. 5. These abnormal solutions are, all of them, capable of con- ducting the electric current, whilst substances like sugar, which THE IONIC THEORY 429 possess a normal molecular weight in solution, do not form conducting solutions. To explain these facts Arrhenius put forward the hypothesis that the molecules of such abnormal solutes are dissociated into ions — electrically charged particles — an equal number of positive and negative ions being formed. These ions act as carriers of the electric current, and it is assumed that when the electric charge upon the ions is neutralised at the electrode, the dis- charged ion assumes again all the properties of the element from which it is derived. As a result of further investigation upon the conductivity of aqueous solutions, quantitative values have been derived for the charge carried by the ions, for the speed at which they move through the solution under a given potential gradient, and for the number of such ions present in solution. The easiest method of measuring the extent to which the dissociation of the dissolved substance has progressed, is by the determination of the ratio conductivity at the dilution v conductivity at infinite dilution* In this way it has become possible to compare with ease the amount of dissociation in different solutions of the same solute and of difierent solutes. Perhaps the most important result of such work from the purely chemical point of view is the dis- covery that the dissociation of solutes increases with dilution. Questions 1. What evidence has been adduced in favour of the view that ionic migration takes place during the electrolysis of an aqueous solution of an electrolyte ? 2. Deduce Ostwald's Dilution Law. To what class of substances does this law apply ? 3. What information does ionic migration throw upon the constitution of pota/Ssium permanganate ? 4. Compare the ionisation of a half normal solution of hydrochloric acid, sulphuric acid, sodium sulphate, aliuninium chloride. 5. By what means may the dissociation of a dissolved salt be measured ? 6. Discuss what happens during the electrolysis of copper sulphate between (a) copper electrodes, (6) platinum electrodes. 7. What weight of silver will be deposited from a solution of silver nitrate by the passage of 5 coulombs ? 8. Enunciate Faraday's laws of electrolysis, and show how you would prove them experimentally. 9. Which solution forms the better conductor, a normal solution of sodium chloride or a normal solution of ammonium hydroxide, and why ? 10. Discuss the conductivity of aqueous solutions of dissolved salts. 11. Show how the following substances ionise in solution : aluminium chloride, bariiun chloride, lead nitrate, sodium sulphate, sodium acid sulphate. CHAPTER XXIX APPLICATIONS OF THE IONIC THEORY TO CHEMICAL REACTIONS Properties of the Ions. — Chemists have long been aware that in dilute solutions all soluble chromates possess a bright yellow colour, all the permanganates a purplish tint, all the salts of copper, derived from an acid which is itself colourless, give a characteristic blue-green solution ; similarly, salts of nickel are green, of cobalt pink and so on. A consideration of the equihbrium CuCl2^z:±Cu+ + + 2Cl- brings to our notice that we have in solution undissociated molecules of copper chloride in equihbrium with cupric and chlorine ions, and experiment has shown that in dilute solution this equihbrium is swung to the right, i.e. the dissociation is far-reaching. Other such salts of copper show a similar behaviour : CuS04^^Cu+ + -f SO4- Cu(N03)2 ;iz± Cu+ + + 2N0^ There is no ground for the assumption that the undissociated molecules are blue ; indeed, anhydrous copper sulphate is white, but the assertion that the cupric ion, whether it arises from the dissociation of the chloride, sulphate or nitrate, is always blue appears well founded. Almost incontrovertible evidence in sup- port of this is given by the results of migration experiments with solutions of copper salts. The blue boundary is found to move in dilute solution at practically the same rate, whatever the copper salt present. That the cupric ion generated by all copper salts should do so is to be expected, that the widely differing molecules of copper sulphate, copper nitrate, etc., should do §g would indeed be a striking coincidence, 439 APPLICATIONS OF THE IONIC THEORY 431 • Each ion, then, has its own specific physical and chemical properties, which are just as characteristic of that particular ion as the colour, density, atomic weight, melting point, are characteristic of any particular element. The chlorine ion, for example, will under all conditions react with the silver ion to produce a precipitate of silver chloride, but the chloride atom does not necessarily react in this way. Thus a solution of silver nitrate does not produce an immediate precipitate with chloroform (CHOI 3) because the latter compound is not ionised — it is a non-electroljrfce. Ionic Equilibria in Solutions. — When an acid, base or salt is dissolved in water, an equilibrium is at once set up between the undissociated molecules and the ions. As an example of this, ammonium hydroxide, normal in strength, will serve : NH4OH ;z:^ NHi ^ +0H- 99-6% 04% 0-4% In this case the dissociation is very slight, the main, part of the ammonium hydroxide remaining in the undissociated state. But should a solution of ammonium chloride be taken of the same strength, the dissociation is far-reaching : NH4Cl^z±NH4+ + Cl- 26% 74% 74% The following question now seems pertinent : What would be the effect of adding this highly dissociated solution of ammonium chloride to a solution of ammonium hydroxide ? The introduction of ammonium chloride into the latter solution wiU cause an immediate increase in the concentration of the ammonium ion. In accordance with the Law of Mass Action this must exert its influence upon the equilibrium : NH40H^± NH4+ + OH- f or ^— =^h, the brackets denoting concentrations ; and (NH4OH) in order to maintain k constant, any increase in the concentration of the NH4+ ion must be attended either by a decrease in the concentration of the 0H~ ion, or an increase in the concentration of undissociated NH4OH, or both. As a matter of fact, the equilibrium will be driven to the left through the removal pf hydroxyl ions, a consequential increase in the concentration 432 AN INORGANIC CHEMISTRY of NHiOH resulting. The ionisation of the ammomum hydroxide is therefore repressed. This is of common occurrence. Given two compounds, etc., possessing a common ion, one of the compounds being more strongly dissociated than the other, then the addition of the more strongly dissociated compound will repress the dissociation of the more weakly dissociated substance. This is strikingly borne out in Table 46, where the quantitative effect upon the dissociation of acetic acid by the addition of sodium acetate is tabulated. TABLE 46 0-25 N Acetic Acid + Sodiuji Acetate Concentration of Concentration of sodium acetate added. H ■*" in tlie solution 00 . . 0021 0012.5 . 00034 002.5 . 000017 005 . . 00008 012.5 . 00003 25 . . 00001 The Solubility Product If a substance, e.g. sUver chloride, is shaken with water, solution ensues, that is to say, the equilibrium AgCl^=±AgCl^zz^Ag+ + Cl- Undissolved. Dissolved. is set up. Molecules from the solid pass into the surface layer and thence into the solution itself, until the number of molecules dissolving per second is exactly equal to the number which returns to the surface of the solid per second. The Law of Mass Action states that in the balanced reaction, A^Cl^Ae-^ + Cl- (Ag^)(Cl-)_ and, as the solution is saturated and in contact with the solid the concentration of the undissociated molecules must be fixed, hence (Ag "•" )(C1~ ) =L, a constant known as the solubility prodtict. It follows, therefore, that in all solutions containing sUver and chlorine ions, there cannot be a stable state if the product of these ions exceeds the solubility product. If L is exceeded, combination of the ions takes place to form undissociated silver chloride, and as the solution is already saturated with this substance, the newly formed silver chloride must separate out. APPLICATIONS OF THE IONIC THEORY 433 If a drop of a solution of potassium chloride is added to a solution saturated with silver chloride, the equilibrium AgCl ^=z^ AgCl ;=z± Ag + + CI- Solid. Dissolved. will be upset in accordance with the conclusion arrived at previously. Momentarily, the product (Ag + )(C1") exceeds L, and in order to re-establish the equilibrium, chlorine atoms enter into combination with, silver ions, and silver chloride is precipitated. Only when the concentrations of the silver and chlorine ions have fallen so low that their product is again equal to the solubility product will the precipitation cease. The addition of a drop of a solution of any chloride or of any silver salt would have a similar effect in producing a separation of silver chloride. Applicability of the Solubility Product to Gases of Double Decomposition. — If solutions of barium chloride and of sodium sulphate are mixed, the same method of treatment enables one to see what will happen. Before mixing, the following equilibria exist : Na^SOi ^=± 2Na- + SO4- BaCl2;=^Ba+ + + 2Cl-. If, after mixing, the solubility product of any salt which can be formed, is not exceeded, e.g. BaS04, NaCl, then no separation of a precipitate will occur, but in the case actually before us the solubility product of one substance, BaSO^, is so extremely low that the product (Ba "*" * )(S04~ ) far exceeds the solubility product, and this condition of instability at once leads to the separation of a precipitate of barium sulphate. This substance will continue to separate out until (Ba+ +)(S04°)=L g.gg^. At the conclusion there will be left in solution a very low concentration of Ba *" '^ and SOj" ions, whilst the whole of the sodium chloride remains in solution, partly unionised, but mainly as Na""" and CI" ions. The separation of crystals of sodium chloride by the passage of hydrogen chloride through a saturated solution of sodium hydrogen sulphate (see p. 151) also receives its explanation from this point of view. Here we have to deal with the equilibria. HC1;=±H+ + C1- NaHS0i^^Na++HS04-. 434 AN INOEOANIC: CHEMISTRY The continued introduction of hydrogen chloride into the solution causes such an increase in the concentration of the chlorine ion that eventually the product (Na + )(C1") exceeds Lj^^^, and a separation of crystals of sodium chloride takes place. The explanation is, perhaps, compUcated in this case by the tendency of hydrogen chloride to combine with the solvent, water, and therefore prevent it from exerting its full solvent action upon the salt in solution, but this in no ^^'ay invalidates the above statement of the case. The whole realm of quahtative analysis bristles «ith examples of ionic reactions. The addition of hydrogen sulphide to a solution containing a copper salt affords a suitable illustration, for the presence of minute traces of the ions of copper and of siilphur leads to the precipitation of the insoluble copper sulphide in the sense of the equation, [Cu+-][S=]>L^„s The delicacy of any precipitation test either in quahtative or quantitative Avork is invariably associated with the question : " What is the magnitude of the solubUity product of the insoluble salt, and by what means can that product be exceeded and the salt precipitated V Application to the Precipitation of the Sulphides. — In dealing ^^ith the uses to A\hich hydrogen sulphide is commonly put in chemical analysis, it was shown how it A\as possible to separate the suljihides of certain elements kno\vn as Group II from the sulphides of Group III by varying the concentration of the acid. The sulphides of Group II are precipitated by the action of hydrogen sulphide in the presence of an excess of acid, whilst the presence of an alkali enables this reagent to throw out of solution the sulphides of Group III. As an example, consider the separation of the sulphides of copper and zinc. Hydrogen sulphide is passed through a solution of the two salts, acidified with hydrochloric acid, and a precipitate of copper sulphide seijarates out. The solution is then rendered alkaline by the addition of ammonium hydroxide, and hydrogen sulphide is then able to throw down the zinc sulphide. Hydrogen sulphide in solution forms a very weak acid, feebly dissociated into its ions : HjS ;=± H+ + HS- ;=^ 2H+ + S- . APPLICATIONS OF THE IONIC THEORY 435 Even in an aqueous solution of this feeble acid, the concentration of the S" ion is extraordinarily low. The addition of such a highly ionised acid as hydrochloric acid to an aqueous solution of hydrogen sulphide will cause such an increase in the concentra- tion of the hydrogen ion that the equilibrium is swung to the left, i.e. the concentration of the S" ion is markedly lessened. In order to bring about the precipitation of copper sulphide, the product (Cu + + )(S "^ )must exceed L^^^g, whilst for the separation of zinc sulphide (Zn++)(S=) must exceed L^^g. Of these two constants L^^g and L^ g, the former is very small, the latter relatively large. Even in the presence of hydrochloric acid the concentration of the S" ion is still sufficiently large for the product (Cu "'""'" )(S") to exceed L^^^g, and this sulphide is preci- pitated quantitatively. But not so with zinc sulphide. The extremely low concentration of the S° ion prevents the product (Zn+''")(S") from exceeding the relatively large solubility product L2ng, and no zinc sulphide separates out. The addition of ammonium hydroxide neutralises the free hydrochloric acid, and the further passage of hydrogen sulphide through the alkaline solution leads to the reaction : 2NH4OH + H^S ;r^ {^'H.,),S + 2H2O. But ammonium sulphide, together with the hydrosulphide (NHiHS) which is also formed, is a salt, and therefore strongly ionised : (NH,)2S;=±2NH,+ + S = . The equihbrium defined in this equation lies well to the right, and the concentration of the S" ion becomes so high that the product (Zn""" ■'")(S") far exceeds L2.,g, and the zinc sulphide is now able to separate from the solution. Acids, Bases, Neutralisation, Salts Acids. — The reader's attention has been frequently directed to the fact that a whole class of substances which we know as acids possesses in aqueous solution many striking properties in common. However different in physical and chemical properties these substances may be in the pure state, the aqueous solutions at least show many common properties. Liquefied hydrogen chloride is a non-conductor of electricity, blue litmus is not reddened by it, nor does it react appreciably with metals or metallic oxides ; pure hydrogen acetate is a colourless liquid, 436 AN INORGANIC CHEMISTRY which neither reddens Utmus nor conducts the electric current. It is also without action upon a piece of zinc. But both these substances in aqueous solution become electrolytes, a gram equivalent of each of them liberating the same quantity of hydrogen when allowed to react with a metal. It is true that the velocity with which the hydrogen is set free differs enormously, but the absolute amount generated is the same. All acids have certain properties in common, but each acid appears to possess certain specific properties of its own. This is explaiaed in the terms of the ionic theory by the assumption that all acids dissociate with the production of more or less hydrogen ions : HC1;=^H+ + C1- CH3-COOH^=^H+ + CH3COO- H2SO4 ^::± H+ + HSO4- ^=± 2H+ + SOr The 'properties common to aqueous solutions of all acids are attributed to the presence of hydrogen ions in such solutions ; the properties which are distinctive of each acid are due- to the presence of the various anions in solution, e.g. CI ~ , Br ~ , SO4 " , NO 3 ~ , P04,=' lOj". Sulphuric acid, for example, gives the same reaction with barium chloride as does any inorganic sulphate, because in the solution of such salts there are present SO4" ions, which react with barium ions, forming insoluble barium sulphate. The properties associated A^ith the hydrogen ion may be briefly summarised : 1. Its sourness. 2. The action upon Utmus and other indicators (see p. 449). 3. The tendency to react with certain metals Avith the liberation of hydrogen, e.g. 2H+ +Zn— >H2+Zn+ + 4. The capacity of solutions containing hydrogen ions to react with bases to form salts. 6. Its power to carry only one Faraday of positive electricity, i.e. it is a monovalent cathion. 6. The high conductivity of solutions containing hydrogen ions, as compared with the conductivity of solutions of salts of equivalent strength. Of the above criteria the last is the only one which is distinctly characteristic. Thus many substances other than acids are sour to the taste, the solutions of many salts turn blue litmus red (see Hydrolysis, p. 447), whilst some acids scarcely possess this property; many other ions are monovalent, e.g. Na+, NH4+, AP]?LICATIONS OF THE IONIC THEORY 437 while a solution of sodium hydroxide will react with zinc or aluminium and liberate hydrogen. Chloroform reacts with sodium hydroxide to form sodium chloride and sodium formate, but chloroform is not an acid. CHCI3 +4:NaOH-^HCOONa +3NaCl +2H2O. The high conductivity of aqueous solutions of acids is, however, distinctive. There is little difference in the velocity of migration of all anions, so that the increased conductivity possessed by solutions of acids as compared with that of a salt under equal conditions of dissociation, must arise from the much greater mobihty of the hydrogen ion. This reasoning is in agreement with the experimental results obtained from the direct measure- ment of the velocity of migration of cathions {see p. 421). In classifying a substance as an acid, it is obvious that its proper- ties as a whole must be considered, not merely some of them. The Reaction of Acids with Metals. — The reaction of an acid with a metal proceeds more smoothly in the presence of more or less water, i.e. ionisation is a necessary preliminary to reaction with a metal. In dilute solution, where the acid- and the salt formed may be looked upon as being approximately completely dissociated, the reaction is essentially one in which the ionic charges borne by the hydrogen ions are passed on to the metal : 2H+ + 2Cl-+re->Fe+ + + 2Cl- + H2 or, eliminating the ions common to both sides, 2H++Fe->Fe+ + +H2. Should there be present in the aqueous solution undissociated molecules of the acid, the two reactions : 2CH3COOH ;=:± 2H+ + 2CH3COO- 2H+ +Zn^Zn++ +H2 must be considered together. The steady removal of hydrogen ions from the system, as indicated in the second equation, will cause the equilibrium defined in the first equation to swing to the right, i.e. there will be a progressive dissociation of the undissociated molecules of acetic acid, until at last the whole of the acid present has reacted, and we are left with a solution of zinc acetate (dissociated) and a quantity of hydrogen equiva- lent to the amount of zinc dissolved will have been set free. A similar disturbance of the ionic equilibria ruling in an 438 AN INORGANIC CHEMISTRY aqueous solution of an acid is often produced by the action of a base. This is especially the case if the acid is weak and there- fore undissociated. We have already learnt that the action of a base upon an acid is essentially a reaction between the hydrogen and hydroxyl ions, water being thereby formed. The effect of adding sodium hydroxide to a weakly dissociated acid like acetic acid may be represented by the equations : CHj-COOH ^=i H+ + CHj-COO- H++0H-->H20. This removal of the hydrogen ions by the action of the base again causes dissociation of the undissociated molecules of acetic acid to begin afresh, and so the neutralisation, i.e. the action between acid and base, continues until the whole of the acetic acid has been converted into sodium acetate (dissociated) and water. The Strength and Dissociation of Acids. — It is precisely in those solutions of acids which show relatively slight conduc- tivity, i.e. sUght dissociation, that chemical tests indicate that we have to deal with a weak acid. An aqueous solution of hydrogen sulphide, although it fulfils several of the criteria given above for acids, yet has no pronounced action upon litmus, nor does its solution possess a sour taste. This is due to the extreme weakness of this acid, i.e. to the very low concentration of the hydrogen ions produced by the dissociation of the acid : H2S;=±H+ +HS-^r±2H+ +8" Many other weakly dissociated acids are known, notably silicic, carbonic, boric, sulphurous, nitrous and hypochlorous acids : H2S03^=±H+ +HSO3- ^^2H+ +SO3- H3BO3 ^± H+ H- H2BO3- ^=^ 2H+ +HBO3- ^::z± 3H+ +BO3- HNO,^=±H++NOr HC10;=^H+ +C10- H2CO3 ^=i H+ + HCO3- ^13^ 2H+ + CO3- . The organic acids, as a class, are weakly dissociated. Many of the above acids are di- and tri-basic. In such cases the dissociation is mainly Umited to the first stage, viz. : H,C03^±H++HC03-, the extent to which the second stage dissociation HC03-=±H++C03- APPLICATIONS OF THE IONIC THEORY 439 proceeds being very slight. This tendency of the di- and tri- basic acids to show an increasing reluctance to ionise into the second and third stages is common to all acids of this type. Even sulphuric acid shows it, for all but very dilute solutions show a high concentration of HSO4 ions and few SO4 ions. Only on very great dilution does the dissociation HSO4 :H^ -SO. occur freely. Acids, then, which are equally dissociated under similar conditions of concentration, are considered to be equally strong. This is due to the fact that such solutions contain an equal concentration of hydrogen ions — ^the ion which is responsible for the acidic property of the solution. Solutions of acids of equivalent strength, i.e. of equal hydrogen ion concentration, possess very closely the same conductivity. This is due to the predominating influence exerted by the hydrogen ion upon the conductivity. Equally concentrated solutions of hydrochloric, hydrobromic and nitric acids possess a greater conductivity than does a corresponding solution of sulphuric acid or phosphoric acid, hence we conclude that the first mentioned acids are stronger than the latter. Similar conclusions are arrived at if one applied tests other than the conductivity one. As a purely chemical test, the rate at which marble is attacked by the acids which do not form insoluble precipitates would serve. If a solution containing equivalent quantities of nitric and sulphuric acids is treated with an amount of base insufficient to neutrahse both acids, the base has been found to distribute itself between the two acids in proportion to their strength. Physical measurements have shown that a little over 70 per cent, of the base is monopolised by the nitric acid, while only 30 per cent, of the base is left for the sulphuric acid. This experiment is in good agreement with the results obtained by the conductivity method. In the following table are tabulated the relative strengths of the more important acids, hydro- chloric acid being taken as 100,^ the solution being N : TABLE 47 Hydrochloric acid Nitric Sulphuric Oxalic Phosphoric Acetic 100 96 60 18 7 04 440 AN INORGANIC CHEMISTRY Although all methods place sulphuric acid as a much weaker acid than nitric acid, yet it has long been common knowledge that, if sulphuric acid is added to a solution of a nitrate, chloride, etc., the nitric (hydrochloric) acid is expelled from the system. In this case, however, another factor comes into play — the greater volatility of the nitric acid, and it is precisely from the neglect of this important factor that chemists erred so long in assigning to sulphuric acid a greater strength than to nitric and hydrochloric acids. Owing to the non- volatility of the sul- phuric acid, practically none of its molecules pass into the vapour phase, but the more volatile nitric acid, which is partly expelled from the nitrate, distributes itself between the Uquid and vapour phases ; and if the system is not a sealed one, the vapour is carried away and a continuous stream of molecules of nitric acid leaves the liquid to pass into the vapour phase in order to estabUsh an equilibrium between the two phases. In short, the sulphuric acid will completely expel aU the nitric acid if the evaporation is carried out in open vessels, and this in spite of the greater strength of the nitric acid. Bases. — Just as the properties common to aU bases have been found to arise from the presence of the hydrogen ion, the characteristic properties common to all bases are to be attributed to the hydroxyl ion generated when any base is dissolved. The property these substances have of neutraUsing or destroying the acidic properties exhibited by an acid is explained in terms of the ionic hypothesis by the assumption that aU such bases dissociate in solution with the formation of a greater or less amount of hydroxyl ions. NaOH^z=:Na+ +0H- Ba{OH)„^r:±Ba++ +20H- NH,0H^:=iNH4+ +0H- Zn(0H)2^z±Zn++ +20H- Al(0H)3^zi±Al H-30H- AgOH;zz:±Ag^ +0H- The properties which one base possesses as distinct from another base are due purely to the cathion, e.g. the formation of a- precipitate with a given reagent : Ba(0H)2 + H,S04-> BaSO^ i + 2B.fi. APPLICATIONS OP THE IONIC THEORY 441 All the above bases react at once with an acid, forming a salt and water, NH4OH + HCl -> NH^Cl + H^O or, written ionically, NH4OH— =iNH4+ +0H- HCl^z=:H-^ +C1- NH4OH + HCl -^ NH4+ + CI- + H2O ; A1(0H)3 + 3HN03-> A1(N03)3 + 3H2O or A1(0H)3 — ^ A1+ + + + 30H- 3HN03^=±3H+ +3NO3-. A1(0H)3 + 3HN03-^A1+ + + +3NO3- +3H,0. The extent to which the base is actually dissociated does not enter into the question. Ammonium hydroxide, in spite of its slight dissociation, reacts as freely with hydrochloric acid as does the highly ionised sodium hydroxide. This arises from the fact that, directly hydrogen ions are introduced into the solution containing the base, combination between the H"*" and OH" ions occurs, and any undissociated base instantly proceeds to split off fresh hydroxyl ions. This continues until neutrali- sation is complete. A number of methods have been devised for measuring the extent to which solutions of the more soluble bases are dis- sociated. One such method is the distribution of an acid between two bases which are present in excess. This and other methods lead to the conclusion that the hydroxides of the alkalis and of the alkaline earth elements are very strongly dissociated, even in relatively strong solutions. The following table records the relative strengths of some of the more important bases, the solutions always being N/40 : TABLE 48. Lithium hydroxide Sodium hydroxide Potassium hydroxide Thallium hydroxide Ammonium hydroxide 100 9S 98 89 2 The hydroxides of the heavy metals are generally insoluble, and their saturated solutions are so dilute that, although as 442 AN INORGANIC CHEMISTRY bases some of them are weak, they are yet widely dissociated. In considering the neutraUsation of the base, magnesium hydroxide, Mg(0H)2, the following equilibria must be kept in mind : Mg(OH), ^=± Mg(OH), ;=^ Mg+ + + 20H- Solid. Dissolved but undissociated molecules. The introduction of hydrogen ions (acid) into such a solution causes the removal of the 0H~ ions, water being thus produced. In order to restore the equilibrium defined in the equation : [Mg^LOHJ^^^ [Mg(OH),] more of the undissociated molecules must break down into ions, and thereby causes a displacement of the equilibrium between the undissociated and the solid molecules. Hence, the net result is that more and more of the solid must continue to dissolve until the whole of the hydroxide has reacted. Alkali is the term applied to the more active and strong bases, viz. KOH, NaOH. Sodium and potassium hydroxides are also known commercially as caustic soda and caustic potash respec- tively. The alkalies, and to a less extent, certain of the other bases (silver hydroxide, magnesium hydroxide, calcium hydroxide, ammonium hydroxide, etc.), have the property of imparting to water a soapy feel. They also turn red litmus blue. Neutralisation. — The reaction between acids and bases, provided these are strongly dissociated, proceeds to a com- pletion. NaOH + HCl -^ NaCl+ H.O . In chemical work it has become the practice to work with gram equivalent quantities, rather than with the gram molecule. A gram equivalent of hydrochloric acid contains 36-5 gm. of 98 hydrogen chloride per Utre, of sulphuric acid — =49 gm. per 98 htre, of phosphoric acid —=32-6 gm. per litre, of sodium hydr- o oxide 40 gm. per htre, of barium hydroxide — =85-5 gm. per litre, i.e. the molecular weight of the acid (base) -^ valency of the cathion (anion). A solution containing 1 gm. equivalent APPLICATIONS OF THE IONIC THEORY 443 per litre is known as a normal solution ; similarly a solution containing half a gram equivalent per litre is referred to as a half normal solution (N/2), etc. A normal solution of an acid contains 1 gm. of hydrogen ions, a normal solution of a base 17 gm. of hydroxyl ions. When 1 litre of a normal solution of acid reacts with 1 litre of a normal solution of a base, there are formed 18 gm. of water : H+ + 0H-- 1 gm. 17 gm. ^H^O 18 gm. If one has available a solution of a base of known strength, say exactly normal, the exact titre or strength of an acid can be obtained by the operation known as titration. A measured volume of the acid is allowed to drain slowly from a pipette or gradu- ated tube into a beaker, a drop of litmus or other indicator added [see p. 449), and the solution of the base slowly run into the acid from a burette (Fig. 117). The colour change of the indicator shows the point of neutrality. At this moment the tap of the burette is turned. From a knowledge of the volume of the base of known strength used and of the volume of the acid taken, one can at once compute the concentration of the acid. Thus, if 25 c.c. of a normal solution of sodium hydroxide neutral- ises 20 c.c. of a solution of sulphuric 25 acid, the acid must be —^=1-25 N. The liberation of a constant amount of heat, 13,700 cals., when 1 gm. equi- valent of an acid is neutralised by 1 Fio. 117. gm. equivalent of a base has already been dwelt upon, especially as regards the support it gives to the ionic hypothesis. In certain cases, however, it has been observed that during the neutralisation the evolution is less than that mentioned above, e.g. in the neutralisation of ammo- 444 AN INORGANIC CHEMISTRY nium hydroxide by a strong acid. This fact, instead of weak- ening the position of the ionic hypothesis, strengthens it. Consider the equations : NH4OH+ HCl -> NH1CI+H2O NH4++OH- H++a- NH,+ + Cl- in this case the heat of neutralisation per gram equivalent is only 12,200 cals. instead of the usual 13,700 cals. generated by the interaction of a solution of a strong base with a strong acid. 'This arises from the fact that ammonium hydroxide is weakly dis- isociated, but as the removal of OH" ions takes place through the agency of the hydrogen ions of the acid, the undissociated imolecules of ammonium hydroxide dissociate further, NH,OH^=:±NH,+ +0H- and this dissociation is attended by an absorption of heat. What is actually measured, then, is the difference between the heat of neutralisation and the heat of dissociation of the ammonium hydroxide. Convincing proof that this base actually dissociates with the absorption of heat is afforded by the fact that this substance shows a marked rise in conductivity with a rise in temperature (Le Chateher's Law). This increase can only come about from an increase in the number or in the velocity of migration of the ions. It is true that the speed of migration increases with a rise in temperature owing to a diminished viscosity, but this effect is in this case quite insufficient to account for the marked increase in the conductivity. Amphoteric Electrolytes. — In Chapter XXII, deaUng with the elements of Group V, emphasis was laid upon the remarkable behaviour of the lower hydroxides (oxides) of arsenic and anti- mony, in so far as these substances possess the dual property of neutralising either a strong base or a strong acid, viz. : As(OH) 3 + 3HC1 ^=± ASCI3 + 3H2O As(0H)3 + 3NaOH ^:z:± NajAsOa + SH^O. The question may well be asked: How do amphoteric -hydroxides fit into such a scheme as presented by the ionic hypothesis ? Arsenious hydroxide has the property of neutralising an acid. APPLICATIONS OF THE IONIC THEORY 445 This indicates that in an aqueous solution of this difficultly soluble hydroxide the following equilibria exist : As(0H)3 ;zi± As(0H)3 ^=± As+ + + + 30H-. Solid. Dissolved. On the addition of such a strongly ionised acid as hydrochloric acid, combination at once ensues between the hydroxyl and hydrogen ions. The above train of equiUbria is thereby upset, and in order to restore the equilibrium, more hydroxide dissolves. The continued addition of acid will therefore lead to the complete solution of the hydroxide. The introduction of the acid has also increased the concentration of the chlorine ions, so that the reaction As"*" "•" + +3Cl"'^r^AsCl3 setsin. At the conclusion of the experiment we shall have in solution a large number of the undissociated molecules of arsenic trichloride and a few ions of arsenic and of chlorine in equilibrium with the very weakly dissociated arsenic trichloride. On the other hand, the fact that arsenious hydroxide dissolves readily in a strong base such as sodium hydroxide, with the formation of a salt certainly argues in favour of the following equilibria existing in an aqueous solution of this substance : As(0H)3 ^^ As(0H)3 ^i^ 3H+ + AsOa^^- Solid. Dissolved. The introduction of hydroxyl ions into the solution in the form of a strongly dissociated base such as sodium hydroxide causes a reaction between the hydroxyl and hydrogen ions, and the above train of equilibria is again upset. In order to restore the concentration of the hydrogen ions to its former equilibrium value more hydroxide dissolves, and will continue to dissolve so long as sodium hydroxide is introduced into the solution. When solution is complete, we shall be left with a solution of the highly dissociated salt, sodium arsenite. An aqueous solution of an amphoteric electrolyte such as arsenious hydroxide must therefore be in equilibrium not only with hydrogen ions, but also with hydroxyl ions. Thus we have : ASO3- + 3H+ ;=± As(0H)3 ^r± As+ + + + 30H- . It is this capacity to split off both hydrogen and hydroxyl ions that confers upon amphoteric hydroxides their distinctive property of reacting either with an acid or with a base to form a salt. The reactions between arsenious hydroxide and sodium 446 AN INORGANIC CHEMISTRY hydroxide on the one hand, and arsenious hydroxide and hydro- chloric acid on the other can be formulated in the following way : AsO," + 3H+ ^=± As(0H)3 ^=± As+ + + + 30H- 3Na++30H- 3C1-+3H+ NajAsOa + SH^O AsClj + 3H2O 3Na++As03= AS+ + ++3C1- The actual amount of arsenious chloride in the dissociated state is, however, very small. Water as an Electrolyte. — It was Davy (1806) who first showed that the electrolysis of water led to the appearance of an acid round the cathode and an alkaU round the anode. This he rightly attributed to the presence of dissolved salts. Careful distillation in gold vessels enabled him to show that the phenomenon arose from the small quantities of salts dissolved from the glass by the water. Modern work, particularly on the part of Kohlrausch, has shown that, however pure the water may be it stUl retains a certain minimum conductivity which has been attributed to the presence of extremely small quantities of hydrogen and hydroxyl ions. H^O^^H+H-OH-. If such a dissociation of the water molecules occurs, it follows that (H+)(OH-) ~(H,or " ' and since the total quantity of water is practically unaltered by the extremely small amount of ionisation which occurs, we may write (H+)(0H-)=A:(H20)=K that is, whenever hydrogen ions are present in aqueous solution, there must also be a definite concentration of hydroxyl ions. It has been found by various physico-chemical methods, such as the conductivity of water distilled from platinum vessels, that the ionic concentration of the hydrogen and hydroxyl ions in pure water is 10"'', hence (H"'")(OH~)=10""i*. In an N/10 solution of hydrochloric acid (85 per cent, dissociated) the concentration of the hydrogen ions (H'*') is 0-085, hence : APPLICATIONS OF THE IONIC THEORY 447 Similarly, in a tenth normal solution of sodium hydroxide the ionic concentration of the hydrogen ion is about l-2xl0~i*. So long as the ionic concentration of the hydroxyl and hydrogen ions in any solution remains the same, the solution is neutral, but any disturbance which leads to an excess of hydrogen ions wiU cause acidity, whilst an increase in the ionic concentration of the hydroxyl ions produces alkalinity. Hydrolysis . — If a salt, formed from a strong base and a weak acid, e.g. KjS, NajCOg, KCN, etc., is dissolved in aqueous solution, it is found that the solution reacts alkaline. This indicates that there is an excess of hydroxyl ions. Whence do these ions come ? The hydrolytic decomposition of sodium carbonate Na^COs + H^O ;rz± 2NaOH + H^COj leads to the formation of an equivalent quantity of free acid and free base, but, whereas sodium hydroxide, being a strong base, undergoes far-reaching dissociation, 2NaOH ^z^ 2Na+ + 20H- the extremely weak carbonic acid is scarcely ionised at all : H,C03^=±H++HC03- . Consequently, there is produced in solution a large excess of hydroxyl ions ; the concentration of these ions rises above 10"', and in accordance with the equation (H+)(OH-) = 10-i* there must be a corresponding drop in the concentration of the hydrogen ion. The solution must therefore react alkahne. All salts formed from a weak acid and a strong base will undergo hydrolysis in aqueous solution, and as a result their solution will react alkaline. The differential dissociation of the strong base and of the weak acid arises from the difference in the specific properties of these substances, and is further accentuated by the influence exerted by the ions of the original salt, e.g. the dissociation of the very weakly ionised carbonic acid : H,C03^=^H+ +HCOr ^=i2H+ +CO3- will be almost completely prevented by the presence of the 448 AN INORGANIC CHEMISTRY carbonate ions arising from the sodium carbonate. On the other hand, the presence of the sodium ions will exert scarcely any influence upon the dissociation of the strong base, sodium hydroxide. But if a salt, formed from a strong acid and weak base, e.g. aluminium chloride, is dissolved in water, we have AICI3 + 3H0H ;=i± A1(0H)3 + 3HC1. The total quantity of acid and base produced by the hydrolysis is again the same, but, whereas the hydrochloric acid is so strongly ionised that its ionisation is not appreciably affected by the presence of chlorine ions generated from the unhydrolysed aluminium chloride, the dissociation of the aluminium hydroxide is so sHght as to be negUgible. This is due to the weakness of aluminium hydroxide as a base, as well as to the overpowering action of the aluminium ions generated by the dissociation of the unhydrolysed aluminium chloride : 3HCl^:z^3H+4-3Cl-. A1(0H)3^^A1+ + + +30H-. As a result, there is a predominance of hydrogen ions in the solution, i.e. (H+)>10"'. Such a solution will therefore react acid to litmus. Other examples of such salts are the chloride and sulphate of copper, zinc, bismuth, chromium, iron (ferric), ammonium, etc. The extent to which such salts as NajCOa, AICI3, etc., are hydrolysed is slight — a few per cent, at most — but if both acid and base from which the salt is formed are weak, far-reaching hydrolysis occurs. Thus, ammonium acetate, ammonium carbonate, and ammonium sulphide are roughly 10 per cent, hydrolysed. It is because of this extensive hydrolysis that it is impossible to prepare such salts as aluminium carbonate by the wet method. If the salt is formed from a strong acid and an equally strong base, e.g. NaCl, no measurable hydrolysis occurs, for, if we concede for the moment that hydrolytic decomposition does take place, equihbria such as the following must be set up : NaCl + HOH ^:^ NaOH + HCl Na++Cl- Na++OH- H+ +C1 APPLICATIONS OF THE IONIC THEORY 449 Since both acid and base are strong, the extent to which they are dissociated will be practically the same, hence the concentration of the hydroxyl ion must equal that of the hydrogen ion; but (H + )(OH~)=10~", so that the concentration of each of these ions must be 10 ~' as in pure water. This is equivalent to saying that there is no measurable hydrolysis in a solution of such a salt as sodium chloride. The two equations A1(0H)3 + 3HCI-^ AICI3 +3H2O AICI3 + 3H2O — ^ A1(0H)3 + 3HC1 throw into sharp contrast the processes of hydrolysis and neutralisation. Hydrolysis is thus seen to be the reverse of neutral- isation. Exact neutralisation of an acid by a base can only be effected when the acid and base are alike strong, but the error involved in the titration of a weak acid by a strong base is relatively small. It is obviously impossible to neutralise exactly a weak acid (acetic acid) by a weak base (ammonium hydroxide) (cf. next section). Indicators and their Action. — Indicators, such as are used in the titration of acids and bases, change their colour according to whether they are in the presence of an excess of an acid or of a base. Litmus is blue in the presence of an alkali, red with an acid ; methyl orange gives a reddish solution with an acid, and a bright yeUow solution with an alkali. Although the original theory put forward by Ostwald does not take into consideration all the equilibria existing in a solution of an acid- base indicator, yet, in the main, the theory propounded by Ostwald does explain the colour changes observed. On this view indicators in general use are weak acids or bases, possessing a distinctive colour which is different from that of the salt formed from the indicator. As an example, phenol- phthalein will serve. This substance is a very weak acid which forms a colourless solution, but in the presence of an excess of alkah, it is bright pink. On the addition of an acid like hydro- chloric acid to a solution containing phenolphthalein, we have the following equilibria : HC1^±H++C1- HPh^:±H+ + Ph-. Owing to the effect of the high concentration of hydrogen ions o o 450 AN INORGANIC CHEMISTRY generated by the dissociation of the hydrochloric acid, the ionisation of the indicator acid is practically completely pre- vented, and as this undissociated acid is in this case colourless, the solution appears colourless. We have in the solution two acids, one of which is extra- ordinarily weak, the other strong and largely dissociated. If a solution of an alkali is run into the acid solution, it wiU enter into combination with the strong acid, and the amount of it which falls to the share of the extremely weak and undissociated indicator- acid is infinitely small ; but so soon as the last trace of the hydrochloric acid is neutrahsed, the base begins to react with the weak indicator-acid HPh, and the sodium salt of this substance is formed. All sodium salts are strongly dissociated, so that, directly the slightest excess of alkah is introduced into the solution, pink anions of the indicator appear in the solution through the dissociation of the sodium salt, NaPh. Ii l shor t, the colour chanfre is due tn the appearance in the snintjnn gi free__a nions of the indicator acid, these anions possepsing a, di fferent colour from the undissociated indicator a,p.id On suc h a theory, red is the colour of th p nnHias ociated litmus acid , blue the colour of the litmus anion , yeUow the colour of the free anion of methyl orange, and red the colour of this undissociated indicator. If the base which is being titrated is a weak one, e.g. ammonium hydroxide, the salt NHjPh, formed as it is from a very weak acid and a comparatively weak base, will be extensively hydro- lysed, NHjPH -f HOH ^:;± NH4OH + HPh, i.e. we shall not get a sharp colom- change at the point of neutrahty, but a considerable excess of ammonium hydroxide will be required to bring about the pink colour of the ion Ph~ . Ammonium hydroxide should therefore never be titrated with a weak acid indicator such as phenolphthalein, but with one of the relatively strong acid indicators, such as methyl orange. With this indicator hydrolysis does not play such an important part in hindering the colour change at the point of neutrality. But it a weak acid, such as acetic acid, is being titrated, it is inadvisable to use a strong acid indicator such as methyl orange, for competition wiU take place between the acetic acid and the methyl orange for the possession of the base. The APPLICATIONS OF THE IONIC THEORY 451 distribution of a base between two competing acids present in excess is proportional to the strength of the acids, provided the acids are present under comparable conditions of concentra- tion. It follows that some of the base will fall to the share of the methyl orange before the whole of the acetic acid is neutralised, i.e. the colour change' will occur before the neutralisation is complete. For such an acid as acetic acid, the weak acid indicator, phenolphtha'ein, should be called into requisition. In view of these remarks it is obvious that no indicator will give a sharp reading if both acid and base which are being titrated are weak. Furthermore, it is always inadvisable to add more than one or two drops of the indicator, otherwise the competition of the indicator for the base will be favoured in accordance with the Law of Mass Action, and the end point will occur too soon. Indicators are frequently called into use for reactions other than between acid and base, e.g. it a solution of sodium chloride is titrated with a solution of silver nitrate of known strength, potassium chromate is added as an indicator. The underlying principle in this type of indicator differs from that already developed for acid-base neutrahsations. So long as chlorine ions are present, the sHghtest drop of the solution of silver nitrate is sufiScient to enable the product (Ag''")(Cl~) to exceed the solubOity product of this salt, and silver chloride will separate out. But there is another possibUity — the separation of silver chromate. This substance is, however, considerably more soluble than silver chloride, i.e., the product (Ag"'')2(CrO4")=L4^g^0jO4 exceeds L^^j. So long as chlorine ions are present, the con- centration of silver ions is kept so low that the solubility product T-ij^g^^tOi cannot be exceeded ; but as soon as the chlorine ions have been practically completely precipitated as silver chloride, the introduction of a drop of the solution containing silver nitrate will enable the product L^^^^jOj to ^® exceeded, and a precipitate of red silver chromate is thrown down. The titration is completed. Questions 1. Discuss Mly the term acid, base. 2. How may the strength of two acids be compared ? 3. Explain fully what happens when hydrogen sulphide is bubbled through an aqueous solution containing mercuric and zinc chlorides. Wliat happens it the solution also contains hydrochloric acid ? 4. Discuss the ionic reactions which are involved when antimoniouB hydroxide reacts with (a) hydrochloric acid, (6) sodium hydroxide. 452 A]S[ INORGANIC CHEMISTRY 5. Explain fully what happens when sodium carbonate and ammoniuni sulphate are separately dissolved in water. 6. Show how hydrolysis and neutralisation are inter-related. 7. Define" solubility product," and illustrate your answer by referring to the precipitation of barium sulphate from a solution containing barium chloride. 8. A saturated aqueous solution of carbon dioxide is treated with sulphuric acid. Explain what happens. 9. What difference would you expect in the behaviour of a solution formed (a) from equivalent quantities of sodium sulphate and sulphuric acid, (6) sodium acid sulphate, the total amount of sulphat-e ion being the same in both solutions ? Why so ? 10. Discuss fully what happens when hydrogen chloride is passed into a saturated solution of acid sodium sulphate. 11. Draw up a list of coloured cathions and anions, and indicate the charge carried by each ion. 12 INIagnesium hydroxide is precipitated by ammonium hydroxide, but dissolves in the presence of ammonium salts. How do you account for this ? (see page 524). 1 3. Formulate, according to the ionic hypothesis, the reactions between (a) calcium oxide and water, (6) calcium hj-droxide and nitric acid, (c) copper sulphate and barium chloride. 14. \Miat weight of hydrogen chloride is contained in a litre of 0!) normal solution of hydrochloric acid ? How many c.c. of a 0-85 normal solution of sodiiun carbonate would be required to neutralise 20 c.c. of the above acid ? ^Miat indicator would you use ? 15. Give a brief account of the behaviour of indicators. CHAPTER XXX THE METALS In the discussion dealing with the Periodic Method of Classifica- tion, emphasis was laid upon the fact that, whilst the basic property of the oxide is most strongly developed in Groups I and II, there is a steady fall in this property as we pass to the groups to the right. This change is most clearly brought out if the following method of classification is adopted : He Li Gl B C N O F Ne Na Mg Al Si P S CI A K Ca So Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Bb Sr Y Zr Nb Mo — Ru Rh Pd Ag Cd In Sn Sb Te I etc. This table at once throws into prominence the fact that aU the typically metaUic elements are clustered at the beginning of each series, whilst the acid-forming elements appear at the end of each series. This gradual transition from elements producing basic-forming oxides appearing on the left of each series to those yielding acid-forming oxides on the right accounts for the occur- rence of many amphoteric oxides (hydroxides) among the metals grouped towards the centre of each series. As an aid to the study of the metals the following points may be stressed : — 1. In passing down the elements of any particular group, e.g. Li, Na, K, Rb, Cs, there is a steady increase in the basicity of the oxide. 2. The metals which form the strong bases appear in Groups I and II ; metals lying to the right of Group II have much less pronounced basic properties. 3. In many cases one or more of the oxides (hydroxides) of the metals are amphoteric — a symptom of very weak basicity on the part of the metal. Amphoterism is of extreme importance in the study of the metals, for it not only explains the existence 453 454 AN INORGANIC CHEMISTRY of such compounds as the aluminates and zincates (q.v.), but it also accounts for the extreme ease with which such salts as the nitrates, chlorides, sulphates of such elements are hydrolysed in aqueous solution. Oxides of metals which are pronouncedly amphoteric in their nature, e.g. aluminium, possess such weak basic properties that all salts formed from weak acids such as carbonic acid, acetic acid, etc., are completely hydrolysed in aqueous solution. Such salts can consequently only be prepared in the dry way. 4. When metals form more than one oxide, e.g. Pb, the lower oxide is the more strongly basic. Chromous oxide, CrO, is distinctly basic, chromium trioxide is acidic, whilst chromic oxide, CrjOa, is amphoteric. 5. The metals which form several oxides and give rise to more than one series of salts are generally found towards the bottom of the table, e.g. tin, lead, bismuth, mercury. 6. When a metal forms more than one series of salts, these salts often exhibit a greater similaritj' to similarly constituted salts of other metals than they do to one another. Thus, mercurous and cuprous chlorides are insoluble, whilst mercuric and cupric chlorides are soluble. Lead chloride (PbCla) is a white crystalline soUd, dissolving without decomposition in hot ^\ater (cf. CuCla, MgClj, HgClj, etc.), but plumbic chloride, PbClj, is a liquid, completely hydrolysed by water, exhibiting properties very similar to siUcon tetrachloride, SiClj. The Formation of Complex and Double Salts. — The direc- tion ill which an ion migrates has already been discussed, as well as the information which the direction of migration throws upon the constitution of a salt. If a salt is at all coloured, the move- ment of a boimdary when a solution is electrolysed under conditions defined on p. 420, affords visual evidence of the charge associated with the coloured ion. It is a matter of common experience to meet -svith salts wherein a metal forms part of the anion. Potassium permanganate, KMnO^, is such a compound, but the isomorphism between this salt and potassium perchlorate, coupled with the pronounced acid nature of the oxide, Mn^O,, precludes us from attributing a complex nature to this salt. There is, however, a whole series of salts which can be formed by the union of molecules of two different salts, and it is to these compounds that the term METALS 455 complex salt is applied. By the action of potassium cyanide upon an aqueous solution of silver nitrate there is first formed a white precipitate of silver cyanide. KCN + AgNOs^-AgCNNl' +KNO3. This precipitate dissolves in an excess of potassium cyanide, forming potassium silver cyanide. AgCN + KCN ^=± KAg(CN)2. This salt is a typical salt of potassium, giving all the reactions of this metal ; but it gives none of the reactions of silver. Moreover, it can be shown by quantitative migration experiments that the silver migrates to the anode as part of the anion, i.e. the salt is ionised thus, KAg(CN),^=±K+ +Ag(CN)r- This compound is formed by the combination of the molecules of potassium cyanide and silver cyanide. Many so-called double salts (q.v.) exhibit similar complex formation to a greater or less extent, but in all cases where marked complex formation occurs, it is found that one of the salts is of a stronger electro- aifinity than the other. The stronger salt appears to force the weaker salt into the anion. In the case of potassium silver cyanide, potassium is an element of extremely high electro- affinity, silver a metal much lower on the electro-affinity scale. Under these conditions, potassium cyanide forces the weaker silver cyanide into the complex, and a stable complex salt results. Similar behaviour is noticed if a strong solution of potassium chloride is added to a solution of cuprio or cobalt chloride ; the former solution is turned a deep brown, the latter a bright blue. In both cases the relatively weaker salts, CuCla and C0CI2, are forced into the complex anion. 2KC1 -f CuCl, ;=::± K.LCuCU] 2K+-t-2Cl- CU++-1-2C1- 2K'+[CuCl4]- Brown. 2KC1 + C0CI2 K2[CoCl4] 2K++2C1- Co-* 2C1- 2K^ [CoCli]- Blue. 456 AN INORGANIC CHEMISTRY The number of such complex or double cyanides, halides and sulphates is very great. They all show more or less dissociation into their original salts, and the extent of this dissociation will determine to what extent the reactions of the individual salts are given. In the case of potassium silver cyanide, although physical methods show that there is a slight dissociation in accordance with the equation, KAg(CN), ^=^ KCN + AgCN 11 K-^+CN- Ag++CN- as well as the main dissociation, KAg(CN)2-^K+ +Ag(CN)r, the concentration of the silver ions in such a solution is so very low that a solution containing chlorine ions causes no precipitation of silver chloride, while hydrogen sulphide produces only a partial precipitation of silver sulphide. Another type of salt, which has much in common with the complex salts, is that known as a double salt. In the study of the metals many examples of double salts will be met with. As examples, the alums {q.v.) wUl serve. Potassium aluminium sulphate is formed by the combination of molecules of potassium sulphate, aluminium sulphate and water, as represented in the formula K2SOi,Al2(S04)3,24H20. Such a salt, when dissolved in water, gives all the reactions of the potassium, aluminium and sulphate ions, but physical measurements have revealed that a considerable amount of complex formation does take place, though the equiUbrium lies well over to the left. K2SO4 + AlaCSOJ, +xH20^zzi K2[Al2(S04)4]24H,0 + (.r -24)H20. Similarly, the double s'alt, K2S04,ZnS04,6H20, does display shght complex formation in aqueous solution, but such a solution gives all the reactions of iJotassium, zinc and sulphate ions. In short, the difference between double salts and complex salts is one of degree, rather than kind, the former term being reserved for salts in which comparatively little complex formation is found in their aqueous solutions. Hydrolysis of the Metallic Halides.— The action of water upon the metaUic and non-metallic halides affords a ready METALS 457 test for determining whether an element is entitled to rank amongst the metals or non-metals. A typical non-metaUic haUde, such as phosphorus trichloride, reacts in the following way on treatment with water : PCl3 + 3HOH^^P(OH)3+3HCl. This reaction, at any rate in open vessels, is practically complete. Phosphorus trichloride suffers complete hydrolysis into phos- phorous acid and hydrochloric acid, though it is possible that phosphorous acid, treated with hydrogen chloride in a sealed system, might show signs of the opposite reaction. A similar decomposition into a mixture of two acids is afforded by the hydrolysis of sulphur chloride : 2S2CI2 + SH^O^ 4HC1 + H2SO3 + 3Si. Such reactions are not readily reversible in the case of the halides of the non-metals. On the other hand, the halide of a metal not only hydrolyses to a much smaller extent : ZnCla + HOH ^n± Zn(OH)Cl + HCl but the reaction is readily reversible. The hydrolysis of such a salt as zinc chloride can be entirely prevented by the addition of hydrochloric or any other strongly dissociated acid. In other words, the hydrolysis of the halide of a metal is strictly reversible. In the case of the metallic halides, the hydrolysis is considerably increased either by a rise in temperature or by an increase in the dilution. A consideration of the equation NH4CI + HOH ^zz± NH.OH + HCl from the point of view of the Law of Mass Action shows that the effect of dilution is only to be expected. Occasionally the hydrolysis of a metallic halide leads to the precipitation of a basic salt, e.g. : SnCla + HOH->Sn(OH)Cl \ + HCl, but the addition of hydrochloric acid at once forces this pre- cipitate back into solution. A similar phenomenon is noticed in the case of antimony, bismuth and arsenic chlorides : SbCla + 2H0H ^=± Sb(0H)2Cl ^^ + 2HC1 BiCla + 2H0H ;rzi Bi(OH) ^Cl \ + 2HC1 AsCls + 3H0H ^:± As(0H)3 ^^ + 3HC1 458 AN INORGANIC CHEMISTRY but here again the addition of hydrochloric acid to the system drives the precipitate back into solution. On such a view, arsenic and antimony belong to the metals rather than to the non-metals, but the reluctance with which these elements assume positive charges has generally led to their inclusion amongst the non-metals, or at least amongst the transition elements, known as the metalloids. Occurrence of the Metals in Nature. — The manner in \\hich metals occur in nature is conditioned mainly by tlie properties of the element. Most of the metals occur as oxides, sulphides, sulphates, carbonates and chlorides, ^\'hilst a few- occur native. They are also widely distributed as sUicat^s. but in this form they rarely have any commercial value. The more important metals occurring as oxides are iron, zinc, tin, copper, aluminium, manganese, chromium and bismuth. As sulphides v,e have lead, silver, cobalt, nickel, antimony, arsenic, bismuth, cadmium, zinc, copper, molybdenum. Iron sulphide is also widely distributed but difficult to utiUse com- mercially. The most important metals occurring as sulphates are lead and the alkaline earth elements (barium, strontium, calcium), magnesium, potassium. The more important carbonates are iron, zinc, copper, lead, calcium, magnesium, manganese, barium and strontium. The chlorides are not very widely distributed except in the case of sodium, potassium and magnesium, the rehcs of old sea beds. Small quantities of sUver chloride (horn silver) are found in nature. Of the metals occuring native, the most important are the noble metals (silver, gold, the platinum group of elements), mercury, copper and arsenic. Prepaeation of the Compounds of the Metals Oxides and Hydroxides. — The oxides of the metals are either prepared by direct oxidation (e.g. lead), or by breaking down the carbonate, CaC03-^CaO+CO,. In such cases it is usually advisable to heat the carbonate in a draught in order to keep down the partial pressure of the carbon dioxide. Oxides are also prepared by dehydrating the hydroxides, Ca(OH),-H20-^CaO, METALS 459 but this method fails for the oxides of the alkali metals, as does that of decomposing the nitrate by heat. Where a metal forms more than one oxide it is not always possible to form all the oxides by the means indicated above. Occasionally, however, different oxides can be prepared in the direct way by varying the temperature and the pressure of the oxygen. The hydroxides are obtained by double decomposition of a salt with a soluble hydroxide, except in the case of the alkaUes and alkaline earths. CUSO4 +2NaOH-^ Cu(0H)2 i + Na^SOi. The hydroxides of the alkalies and of barium are fairly freely soluble, and, if prepared by the above means, must be obtained from the filtrate by evaporation, Na 2CO 3 + Ca(OH) ^ -^ 2NaOH + CaCO 3 i . The tenacity with which the elements of water are retained by the various hydroxides varies exceedingly. The hydroxides of the alkalies resist even a red heat, whilst at the other extreme we have the hydroxides of silver and mercury, which decompose into the oxide at ordinary temperatures. Between these extremes he the hydroxides of such elements as tin, lead, which can be fractionally dehydrated, e.g. Sn(0H)4, HaSnOj. Sulphides. — A variety of methods are available, either the direct union of the elements under the stimulus of heat, or the action of hydrogen sulphide or other soluble sulphide upon an aqueous solution of a salt, or the reduction of the sulphate with carbon . The only soluble sulphides are those of the alkali metals, though even these are much hydrolysed, Na^S + HOH ^z:^ NaOH + NaHS. The sulphides of the alkahne earths are only slightly soluble, though progressive hydrolysis into the hydrosulphide and hydroxide ultimately leads to their solution, 2CaS + 2H0H ^ Ca(HS) , + Ca(OH) , . Other sulphides are practically insoluble. Chlorides. — The method of preparation varies according to whether one desires to prepare the anhydrous or the hydrated salt. If the former, it is usual to start from the elements themselves, or from the oxide, carbon and chlorine. Many oxides are not attacked by chlorine, but if the oxide is heated with a 460 AN INORGANIC CHEMISTRY reducing agent, the metal at the moment of liberation combines with the chlorine to form a chloride, AI2O3 + 3C + 3Cl2^2AlCl3 +3C0. Chlorides containing water of crystallisation are generally obtained by wet methods, e.g. the solution of the carbonate in hydrochloric acid, followed by evaporation. To obtain the anhydrous salt from the hydrated chloride requires more than heat. Magnesium chloride, MgCla.eHjO, leaves an oxychloride on heating. The formation of such an oxychloride arises from the hydrolysis of the chloride, the hydrochloric acid produced thereby being expelled owing to its volatility. In such cases it is usual to heat the hydrated substance in an atmosphere of hydrogen chloride, whereby the hydrolysis is effectively prevented. It is interesting to note that the hydrolysis of the chlorides increases with the valence of the cathion ; tin and lead tetra- chlorides are practically completely hydrolysed in aqueous solution, thaUic chloride (TICI3), auric chloride (AuClj), aluminium chloride (AICI3) appreciably so, the chlorides of the divalent metals comparatively httle except in the case of zmc and stannous compounds, the alkaU chlorides not at all. Electrode Potentials of the Metals If a piece of zinc is placed in a solution of copper sulphate, the reaction Zn + CuSO, -^ ZnSO, + Cu at once occurs. From the ionic point of view, we may write : Zn+Cu-*-+-^Zn++ +Cu. If this experiment is carried out in the usual way, the reaction is attended by the evolution of heat, but by spacially separating the reactions Zn+2®— >-Zn, and Cu++— >Cu+2©, it is possible to generate an electric current in lieu of this thermal effect. This can be effected in the apparatus shown in Fig. IIS. In A is placed a saturated solution of zinc sulphate, in B a saturated solution of copper sulphate, the two vessels being connected by means of a syphon filled with a solution of sodium sulphate. Zinc and copper electrodes are placed in the vessels. A and B respectively, and connected to each other through an ammeter. When the circuit is completed, a current flows. METALS 461 The net result is that zuic passes into solution in the vessel A with the formation of zinc sulphate, i.e. Zn+2@ — >-Zn'*"^ , whUst in B copper sulphate is decomposed and copper deposited upon the copper electrode, Cu'^+-^Cu+2®. This is the essential chemical reaction of what is known as the DanieU cell. The chemical energy available, i.e. the difference between the heats of reaction of the copper and zinc sulphates, has been converted into electrical energy. Strict proportionality between the electrical and the available chemical energy is only found in exceptional cases like the above. For further treatment of this Zinc Electrode. Zinc, Sulphate Solution. Copper Electrode. Copper Sulphate Solution Fig. ]]8. subject the student must consult a text-book of electro- or of physical- chemistry. The question must now be answered as to why zinc is able to displace copper from its solutions, what other elements are able to do this and under what conditions. In order to answer these questions, one must consider what happens when a metal is placed in a solution of its salt. Nemst was the first to put forward the view that, when a metal, e.g. zinc, is placed even in water, there is a tendency on the part of the metal to send positive ions into solution, thereby itself acquiring a negative 462 AN INORGANIC CHEMISTRY charge. A definite potential difference is thus estabhshed between metal and solution.* This tendency of the metal to ionise is often known as its solution pressure. If the solvent already contains in solution a salt of the metal, there will be a definite tendency on the part of the ions to precipitate themselves upon the metal, i.e. the osmotic pressure of the ions will resist the solution pressure of the metal. It is evident that three cases are possible, for the solution pressure may be greater than, equal to, or less than the osmotic pressiure. If the solution pressure is exactly equal to the osmotic pressure, the potential difference between metal and solution will be zero. If, however, the solution pressure exceeds the osmotic pressure, as in the case of zinc, magnesium, the alkalis, the metal wiU acquire a negative charge, the solution a positive. On the other hand, it the solution pressure is less than the osmotic pressure, the metal wiU become positively charged and the layer of solution in contact with it negatively charged. Should a cell be constructed on the lines indicated above, such that one metal is negatively charged, the other positively charged, it is evident that there is every potentiality for the passage of a steady current if the circuit is once completed. The magnitude of the potential difference at each electrode depends upon two factors, the solution pressure of the metal, and the osmotic pressure of its ions. If a set of standard conditions is defined so far as the solution is concerned, we shall have in the potential difference exerted by each metal in such a solution a measure of the tendency of that particular metal to ionise. For the purpose of comparison, solutions which contain exactly one gram ion of the metal per Utre have been chosen. Such solutions are rather stronger than ordinary equivalent normal solutions. In Table 49 are recorded the potentials given by the more important metals in their normal ionic solutions. In accordance with the general practice amongst physical chemists the sign before each potential indicates the potential borne by the metal when in contact with its solution. In this convention, the alkali metals appear as electronegative elements, the non-metals as electropositive. * The potential diSerence between the electrode and the solution is expressed in volts. The capacity of water, stored at a height above sea level, to do work is measured by the expression wt. X h ; the capacity of electricity to do work is measured by the expression quantity of elec- tricity X the potential difference, i.e. coulombs x volts = electrical energy in joules. Potential diSerence la, in a sense, analogous to the height above sea level at which one '''<■= -"i ' ■ Potasaium Soditun . Barium . Strontium Calcium . Magnesium Zinc . Iron . Cobalt . Nickel METALS TABLE 49 -2-9 Tin . . -2-5 1 Lead . -2-4 I Hydrogen -2-3 J Copper . -1-9 1 Arsenic . -1-5 ! Bismuth . -0-50 Mercury . -0-07 Silver . + 0-045 Platinum . + 0-05 ' Gold . . 463 + 0-1 + 0-12 + 0-277 + 0-60 + 0-6 + 0-7 + 1-02 + 1-04 + 1-2 + 1-4 A similar series has been prepared for the more important anions, though in this case various devices have to be adopted in order to enable the reaction x-^x + e to be measured. In many cases, e.g. oxygen, electrodes of platinum in which the element is soluble, are used. TABLE 50 Fluorine Chlorine Bromine + 2-24 + 1-68 + 1-37 Iodine Oxygen + 0-90 + 0-67 A consideration of Table 49 throws into prominence the electro-affinity of the elements and shows that, of all the metals, the alkalies display the greatest tendency to ionise, the noble elements (gold, platinum) the least. If a cell be constructed of two metals, e.g. zinc and silver, each in contact with a solution of its salt normal ionic in strength, the total electro- motive force of the cell is given by the difference of the electrode potentials of the metals, and the positive current wUl flow through the cell from the zinc electrode to the silver ; its magnitude, expressed in volts, will be — 0-50 — ( + 1-04) = — 1-54 volts. In view of the above considerations, it foUows that, if a metal is placed in a solution of any salt derived from a metal occupying a position lower in the electrode potential series, e.g. magnesium in a solution of copper sulphate, the more noble metal (copper) wiU be displaced from the solution by the more electromotively active element, magnesium. In the same way, hydrogen, dissolved in platinum, wiU reduce silver nitrate to metaUic silver. The question of one metal expelling another from one of its dissolved salts receives a complete explanation from the standpoint of the electrode-potential series. The actual effect ^•xe:vtfv\ nn i!bf> TOR cgjJtude of the potential by a change in the 464 AN INORGANIC CHEMISTRY concentration of the ions is not a great one, for Nernst has shown that for monovalent ions a change in the concentration from N to N/10 causes only 0-068 volt difference, for divalent ions — — volt, etc. Of course, when a metal such as zinc is placed in a solution of a salt of another metal, e.g. copper sulphate, the potential of the zinc against the solution is a maximum, as the opposing osmotic pressure is at first vanishingly small. Electrolytic Decomposition — Its Application to Refining. If a solution of hydrochloric acid is electrolysed between platinum or other inactive electrodes, a galvanometer being in circuit, it is found that if the current is interrupted, an electric current which rapidly becomes weaker passes through the solution in a direction opposite to that of the first current. This is known as the polarisation current, and the E.M.F. of this polarisation current is compounded of the discharging potentials of the cathion and anion. This polarisation current arises from the fact that the gases evolved at the electrodes convert these into hydrogen and chlorine electrodes. When the current is broken, the Hj ! HCl , CI 2 cell sends a current in a direction opposing the original ciurent. The minimum difference in potential necessary for the decomposition of an electrolyte is known as the discharging potential. In order that continuous electrolysis may take place, the current passed through the cell must have an E.M.F. not less than the opposing E.M.F. of the polarising current. Since the E.M.F. of a cell is made up of the potentials of the electrodes, the discharging potentials must also be made up of the same two potentials. Given the electrode potentials of the cathion and anion of the salt concerned, we are thus able to calculate the necessary E.M.F. required to electrolyse the solution, e.g. a normal ionic solution of copper sulphate can just be electrolysed by a current of +0-60— (+2-2)= —1-60 volts; for zinc sulphate, —0-50— (+2-2)= —2-70 volts. If the solution contains salts of more than one metal, e.g. copper and zinc sulphates, seeing that the discharging potential of copper sulphate is — 1-60 volts, of zinc sulphate —2-70 volts, it follows that the metal ions may be fractionally precipitated on the cathode. A current with an E.M.F. slightly in excess of 1-60 volts wUl precipitate the METALS 465 whole of the copper, and an increase in the E.M.F. to a value above — 2-70 volts will deposit the zinc. It may again be mentioned that the exact value of the discharging potential is dependent not only upon the solution pressure of the metal, but also upon the osmotic pressure of the ions ; the latter effect is shght. An increase in dilution from 1/10 to 1/1000000 normal (the limit of analytical determinations) will cause an increase of less than 0-3 volt for a monovalent, and 0-15 for a divalent metal. The whole principle of electrolytic analysis and of refining is based upon the foregoing conclusions. It has been remarked that during the electrolysis of potassium silver cyanide, KAg(CN)2, potassium migrates to the cathode, the ion Ag(CN)2 to the anode, and yet silver is thrown down on the cathode. Consider the equation : KAg(CN)2 ^=± K+ + Ag(CN)r t Ag+ + 2CN- The actual amount of dissociation of the complex anion we have aheady found to be extremely small. The positive current will therefore be carried almost exclusively by the potassium ions, but owing to the electrode potential of the potassium being very much greater than that of silver, the latter element, which has partaken in the electrolysis to a very limited extent, wiU be deposited from the solution. In short, all the ions present in the solution take part in the electrolysis, i.e. in the conduction of the current through the solution, but those ions will be liberated at the electrodes, the separation of which proceeds with the least expenditure of work ; the ions which have the lowest discharging potentials will be discharged first. Questions 1. Compare the behaviour of the aluma and the complex cyanides in aqueous solution. 2. Explain what is meant by complex salt, double salt. Illustrate your answer by examples. 3. To a strong solution of cobalt chloride, potassium chloride is added. Discuss what happens. 4. Define solution pressure, electrode-potential, polarisation. 5. During the electrolysis of solutions containing potassiimi cupro- cyanide, the cathode becomes coated with a layer of copper. How do you account for this ? 6. Explain fully what happens when (a) sodium chloride, (6) aluminium chloride is dissolved in water. 7. Indicate general methods for preparing the metallic chlorides in an anhydrous state. H H CHAPTER XXXI GROUP lA : THE METALS OF THE ALKALIES The elements of this group — ^lithium, sodium, potassium, rubidium and caesium — afford an excellent example of that gradation in physical and chemical properties which Mendeleeff has correlated with the change in the atomic weight of the element. This is shown in Table 51 : TABLE 51 Physical and Chemical Properties of the Alkalies Metal. Atomic weight Sp. gr. . . Atomic volume Melting point Boiling point . Sp. heat . Caesium. Increase in reactivity of the metal and in the basic properties of the oxide and hydroxide The alkaU metals show a steady increase in reactivity with rising atomic weight. Thus the temperature attained during the reaction between lithium and water is not sufficiently high to cause the hydrogen to ignite, whilst caesium reacts with almost explosive violence. The metals all give the group oxide, MjO. and under certain conditions the peroxide M2O2 or M2O4 is also produced. Hydrogen reacts with all the members of this sub-group, forming definite crystalline hydrides. These hydrides decompose in water, producing the hydroxide and hydrogen. The hydroxides are strong bases in aqueous solution, due to their wide dissociation. Lithium hydroxide is the least dissociated and the least active of these bases. The hydroxides show great 466 SODIUM, POTASSIUM 467 resistance to heat, and are the only hydroxides known which are not dehydrated by heat ; and so with their carbonates, that of hthium alone losing appreciable quantities of carbon dioxide at 800°. The carbonates as well as the phosphates are freely soluble in water, with the exception of lithium, which again shows properties more in common with those of the alkaline earth elements (Group 2a). In general the salts of the alkaUes are soluble in water, and with few exceptions (notably fluorides and carbonates) the salts of sodium are more soluble than those of potassium. Sodium, Potassium The salts of sodium are widespread. They are found in the sea as sodium chloride, in many old salt beds, as Chih saltpetre (NaNOs) and as a constituent of many rocks and minerals. Potassium is found in the salt deposits as chloride (sylvine), and as carnallite (potassium magnesium chloride, KCl,MgCl2,6H20), as a double sulphate (schonite, K2S04,MgS04,6H20), also in the potash felspars, in land plants and in sheep's wool (suint, an oily sweat on the skin). Isolation of the Metal. — Although it had long been felt that sodium and potassium hydroxides should yield a metal on decomposition, all efforts to prove this failed until H. Davy (1807) hit upon the idea of passing an electric current through the fused hydroxide. Aqueous solutions of the hydroxide he had already found gave on electrolysis hydrogen and oxygen, the elements of water. Davy's inspiration to electrolyse the molten sodium and potassium hydroxide gave the desired result, small metalhc globules rising to the surface and burning brightly. He exposed a piece of the solid hydroxide to the atmosphere in order that a film of moisture might be collected on the surface. The moist hydroxide was then placed upon a disc of platinum connected to the negative pole of a battery, and a platinum wire, connected to the positive pole, was made to touch the hydroxide. A vivid action took place, especially round the positive pole, while little globules separated round the negative pole. As Davy says, " Some of these globules burnt with explosion and bright flame, as soon as they were formed, and others remained, and were merely tarnished, and finally covered by a white film which formed on their surface. These globules numerous experiments soon showed to be the substance I was 468 AN INORGANIC CHEMISTRY in search of, and a peculiar imflammable principle, the basis of potash." At present a modification of Davy's method forms the com- mercial method of manufacturing sodium. In the Castner Process the molten sodium hydroxide is contained in the pot A (Fig. 119), which can be heated by ring burners, B the cathode, is capped by the vessel E, to which is attached a wire gauze screen. The temperature must be kept about 315-320°. During the electrolysis the molten sodium collects in E, while the hydrogen also liberated at the cathode is allowed to escape. Oxygen escapes from the anode. Attempts to prepare sodium by the direct electrolysis of sodium chloride (m.p. 800°) have hitherto met with little success. The electrolytic decomposition of potassium hydroxide presents several difficulties, some of which have been practically overcome. There is a great tendency for the metal precipitated on the cathode to come out in a fine sus- pension. To prevent this reacting with the oxygen round the anode a porous cell must be used. The cathode is placed within this cell. Potassium is also prepared by means of a suitable reducing agent. K2C03+2C^^2K-f3CO. There is, however, considerable danger in the use of this cheap reducer — quantities of an explosive compound, potassium carbonyl, KeCgOe, are formed unless the potassium vapours are very rapidly cooled. The improved Castner method is based upon the use of iron carbide as reducer. Iron fihngs and pitch are first heated together, and then the carbide (FegC) is used to reduce potassium hydroxide. Properties.— Both elements are silvery white, lustrous metals which are readily oxidised by the air. In order to prevent this oxidation, they are always kept under a layer of petroleum Sodium, the less active of the two, melts at 97°, potassium at 62-5°. With water vigorous action occurs, producing the hydr- oxide and hydrogen. This reactivity is shown by their abihty Fig. 11£ SODIUM, POTASSIUM 469 to reduce carbon dioxide with the formation of carbon and a carbonate. Both elements combine violently with sulphur and the halogens. The salts of these metals impart a character- istic colour to the Bunsen flame, the sodium salts an intense yellow, the potassium salts a violet tint. Potassium has few commercial uses, but sodium is used extensively in the manu- facture of sodium cyanide, sodium peroxide, as well as in the manufacture of many azo-dyes. Both sodium and potassium form characteristic amalgams with mercury. These are definitely crystalline substances and have all the properties of definite chemical compounds, such as fixity of composition. Both metals dissolve in liquid ammonia, but when treated with gaseous ammonia in the neighbourhood of their melting points they forin amides, e.g. : 2NH , + 2Na -^ 2NaNH 2 + H 2. Sodamide. 2NH3 + 2K -^ 2KNH2 + Hj. Potasaamide. These substances decompose in the presence of water, yielding the hydroxide and free ammonia. Hydrides. — The hydrides are formed by passing hydrogen over sodium or potassium, heated to about 350°. A white crystalline substance is obtained, from which any unreacted metal may be removed by treatment with liquid ammonia. These crystals have a perfectly definite composition, KH or NaH ; they are decomposed by water, forming the hydroxide and hydrogen. NaH -f HaO -^ NaOH + H2. Oxides. — The group oxide, KjO and Na^O, may be prepared by heating the nitrate with the metal, 2KNO3 + lOK^^eK^O +N2. The oxides combine with water with avidity, forming the hydroxide. In a more or less impure form the oxide may also be obtained by heating the metals in a limited supply of air, but considerable quantities of the peroxides are also formed. Potassium oxide has been prepared in a pure state by the incom- plete oxidation of the metal in an atmosphere of dry oxygen at a low pressure, followed by the distillation in vacuo of the unreacted metal. Lemon yellow crystals are left. 470 AN INORGANIC CHEMISTRY Peroxides. — Sodium peroxide is prepared by heating sodium to 300-350° in a stream of dry air, freed of carbon dioxide. The compound, prepared in this way, is yellow and is very resistant towards heat. It is prepared on the large scale by passing sodium, heaped upon trays, through tubes maintained at a temperature of 300-500°. The peroxide is very stable in the absence of moisture, but. when it is treated with water, vigorous reaction takes place. 2Na202 +2H20-^4NaOH + 0^. If the temperature is kept low, and especially if- an acid be present, hydrogen peroxide is formed. If anhydrous sodium peroxide is exposed to the action of air, freed from carbon dioxide, a stable hydrate, NaaOa.SHjO, is produced. These white crystals dissolve in water \^ithout the evolution of oxygen, pro^aded the temperature does not exceed 30°. Such solutions can even be used for the recrystallisation of the hydrated peroxide, as well as for the preparation of hydrogen peroxide. Sodium peroxide finds frequent use commercially as a strong oxidising agent, also in analytical separations. Potassium peroxide, K2O4, can be obtained as a yellow mass when potassium is burnt in the air. It decomposes water, forming the hydroxide, hydrogen peroxide and oxygen. It is reported that the peroxides, K2O2, K2O3 and K2O4, have been prepared by the action of oxygen upon a solution of potassium in liquid ammonia. Hydroxides. — Two methods are available for the preparation of the hydroxides, the purely chemical method and the electro- lytic, the latter of which is steadily displacing the former. The chemical method is based upon the reaction : Ca(0H)2 + K2CO3 ;=:± CaCOj i + 2K0H. Calcium hydroxide, suspended in water, is run into a boiling solution of the carbonate of potassium (sodium) contained in large iron vessels. The product (Ca++)(C03") far exceeds the solubility of calcium carbonate, and this salt separates out. The steady increase in the concentration of the OH" ions from the dissociation of the alkali hydroxide formed, exerts an influence upon the equihbrium : Ca(OH) 2 ;=± Ca(OH) 2 ;r:± Ca + + +20H " Solid. Dissolved. SODIUM, POTASSIUM 471 driving it to the left, i.e. as the solution becomes richer in potassium hydroxide the concentration of the Ca++ ions falls off, until it falls so low that the solubility product of CaCOs is no longer exceeded, and the reaction ceases. This is equivalent to saying that the solubility of calcium hydroxide in caustic potash is no greater than the solubility of calcium carbonate, and so the driving force responsible for the conversion of the calcium hydroxide into the calcium carbonate becomes zero and the reaction slows down. In order to overcome this disadvantage the operation is sometimes carried out in dilute solution, but a more economical method is to work with fairly strong solutions of sodium or potassium carbonate and recover the unchanged alkaU carbonate from the filtrate. A great deal of commercial caustic soda and potash is prepared in this way from the sodium carbonate obtained by the Le Blanc method {q.v.). Considerable quantities of pure sodium hydroxide are obtained by heating to redness iron oxide and sodium carbonate prepared by the Solvay soda-ammonia process. The resulting sodium ferrite is raked out and lixiviated. Practically pure sodium hydroxide can be prepared in this way : NaaCO, + Fe203-^2Nare02 + CO2 2NaFe02 + H^O -^ 2NaOH + Fe^Oa i The great progress of recent years in the alkali industry centres round the electrolysis of aqueous solutions of the chloride. If an aqueous solution of sodium chloride is electrolysed between inactive electrodes, hydrogen is liberated at the cathode, chlorine at the anode, while the solution roimd the cathode becomes alkaline owing to the formation of sodium hydroxide. This arises from the migration of Na""" ions towards the negative pole where their positive charge is neutralised, and the liberated sodium then reacts with the water, Na"*"— >-Na+®, where denotes one Faraday of electricity, 96,540 coulombs, Na + HOH -> NaOH + H, while at the anode the reaction Cl~-^Cl+0 takes place. Another reaction is, however, also possible — the chlorine round the anode may diffuse through the liquid and react with the cathodic liquid, sodium hydroxide, producing hypochlorite and chlorate {q.v.). In order, then, to procure sodium hydroxide from the electrolytic decomposition of sodium chloride, it is necessary to prevent the chlorine from diffusing into and reacting 472 AN INORGANIC CHEMISTRY with the sodium hydroxide. In many technical cells this is effected by means of a diaphragm, the great disadvantage of which Hes in the increased resistance offered to the electrolysis. In the Griesheim cell, used for making both potassium and sodium hydroxide, a number of cement boxes (acting as dia- phragms) are placed in a rectangular iron box. The anodes Metres 2 Fig. 120. which consist of iron oxide are placed within the cement boxes, whilst the iron box forms the cathode. The whole cell is steam jacketed. The great advantage of this cell is its simphcity of working (Fig. 120). A different method is adopted in the Solvay cell for preventing = - 7777^7777^ f — ^ = A -_-7 - " - :-L-^ K: "i-^/ >>-. ^ W^- ■--=^= =3o? :5€4 -_- : ; a .^^ L^^^J ^^^^ ^^^A ^^^ ^^^ ^^^^ ^^^ ^^H 1 li ^- \i H Fig. 121. interaction between the products of electrolysis. The cell consists of a large rectangular trough through which flows a steady stream of mercury and of brine. The current enters the cell through carbon or platinum anodes, the mercury forming the cathode. The sodium ions, on their discharge at the cathode, form an amalgam with the mercury. The rate of SODIUM, POTASSIUM 473 flow of the mercury is so adjusted that the amalgam is of the desired strength when it leaves the cell. After removal from the cell the amalgam is decomposed in a separate vessel, Hg^a, + 2/H,0 ^ 2/NaOH +yB. + xKg and the mercury is returned to the cell (Fig. 121). Potassium hydroxide may be prepared by the same means. In the Acker process molten sodium chloride forms the electrolyte, molten lead the cathode. The lead-sodium alloy is decomposed by the action of steam and the lead utilised again. The hydroxides, both of potassium and of sodium, find con- siderable commercial application, especially in the manufacture of soaps. Both hydroxides are characterised by their affinity for water and for carbon dioxide. The aqueous vapour pressure for solid potassium hydroxide is so low that the solid rapidly liquefies, owing to the absorption of aqueous vapour from the surrounding atmosphere. Solutions of these hydroxides are frequently used in the laboratory for the absorption of carbon dioxide. In nearly all cases sodium hydroxide is just as effective a reagent as the potassium compound. It is also much cheaper and less of it is required, e.g. for the precipitation of an insoluble hydroxide. CUSO4 + 2NaOH-^Cu(OH), -HNa^SOj CuSO,+2KOH-^Cu(OH)2+K2S04 Thus 159 gm. of copper sulphate are exactly precipitated by 80 gm. of sodium hydroxide [2x(23 + 16+l)=80], while 112 gm. of potassium hydroxide are required to produce the same effect [2x(39-fl6-fl) = 112]. Chlorides. — Both chlorides occur widely in nature. Sodium chloride (rock salt) is generally in a rather impure state and it must be purified by recrystalhsation. Impure specimens of sodium chloride are distinctly deUquescent, due to the presence of small quantities of magnesium chloride. In order to remove this the solution is treated with sodium carbonate, when insoluble magnesium carbonate is precipitated and removed by filtration. Very pure sodium chloride may be jorepared by the action of hydrogen chloride upon a saturated solution of sodium chloride {see p. 151). Potassium chloride is obtained in large quantities from carnallite, KCl,MgCl2,6H20. This salt is crushed and digested in large 474 AN INORGANIC CHEMISTRY tanks with the mother liquor left from preceding operations, and containing sodium and magnesium chlorides. The magnesium sulphate also present in the camallite (25 per cent.) and most of the sodium chloride (12 per cent.) remain undissolved. From the liquor, saturated with carnaUite, there is deposited on cooling crystals of potassium chloride, followed by a deposit of carnainte itself. The potassium chloride is recrystallised while the camallite is put through the vats again. Both chlorides are anhydrous and belong to the regular (cubic) system. At 25° there is little difference in their solubility in water, NaCl 36 gm., KCl 35-5 gm. per 100 c.c. of water, but at 100° potassium chloride is much more soluble (NaCl 39-1, KCl 56-6 gm.). Potassium chloride is also sUghtly more volatile. The importance of these chlorides from the commercial point of view is inestimable as they form the starting point of nearly every salt of potassium and sodium. Bromides and Iodides. — The bromide and iodide may be prepared by the action of the corresponding acid upon the hydroxide or carbonate. They may also be obtained by the action of bromine or iodine upon the hydroxide, e.g. : 6K0H + 3Br2->5KBr + KBrO, + SH^O. After evaporation, the residue is heated either with or without carbon to convert the oxy-salt into the halide. Large quantities of potassium iodide are also made from iron iodide (Feglg). This is formed by rubbing together iron filings and iodine under water, with subsequent treatment with iodine to convert the ferrous iodide first formed into the ferro-ferric iodide. The ferro-ferric iodide is then treated with a solution of potassium carbonate, when the following reaction takes place : Fe,l, + 4K2C03-^ FcaO, + SKI + 4CO2. Sodium and potassium bromides are used medicinally as well as in the preparation of photographic plates. Polyiodides and Polybromides. — Solutions of potassium iodide are able to dissolve relatively large amounts of iodine, forming deep brown solutions. Such solutions give all the characteristic reactions of free iodine, but the actual isolation ^ of certain polyiodides leaves no doubt that such solutions of SODIUM, POTASSIUM 475 potassium iodide and iodine are represented by the equilibrium equation KI+I,^:±Kl3. This is supported by the fact that, if iodine is added to a solution of potassium iodide, practically no change in the freezing point of the solution is produced ; the interpretation of which experi- ment leads to the postulation of a complex molecule, so that no actual increase in the number of molecules present in the solution is produced. If a solution of sodium thiosulphate or other reagent which reacts with iodine is added to the brown solution of iodine and potassium iodide, immediate reaction takes place and'tLe whole of the iodine which has been added to the solution of the potassium iodide will react with the added reagent. This is due to the fact that as soon as iodine is removed from the system by the action of the thiosulphate, etc., the above equilibrium is upset and more iodide is set free, and so on till the whole has reacted. A heptiodide, KI, has also been isolated. Potassium bromide also shows a tendency towards the formation of such polyhalides. Sodium halides possess this property to a very minor extent, though we shall see that both rubidium and caesium give such compounds. Investigation has shown that the stability of these polyhalides increases with the atomic weight of the alkali. Fluorides. — Sodium and potassium fluoride are obtained by evaporating the hydroxide or carbonate with an excess of hydrofluoric acid in a platinum dish. The sodium salt is also prepared from cryoUte, NagAlFg. This mineral is treated with an excess of sodium hydroxide, when the difficultly soluble sodium fluoride is left behind. Potassium fluoride is much more soluble than the sodium salt. Both fluorides form an acid salt KHFa, NaHFj. These acid salts afford a ready method of preparing hydrogen fluoride, for this gas is evolved on heating strongly. The occurrence of these acid salts, coupled with the abnormal solubility of the fluorides, the solubility of which lies intermediate between that of the bromide and iodide, is held to support the view that the fluorides possess the double formula Carbonates. — Sodium carbonate is obtained almost exclu- sively from sodium chloride, the older method of obtaining it from 476 AN INORGANIC CHEMISTRY the ash of seaweed having ceased to possess commercial importance. Owing to the difSculty of obtaining supplies of this necessary commodity duriag the Napoleonic wars, a prize was offered by the French Academy in 1791 for the discovery of a suitable method for converting sodium chloride into the carbonate. As a result of this, the Le Blanc process, which is stUl in use, was developed. In this method the sodium chloride is first treated with sulphuric acid in cast-iron pans. NaCl + H2SO4 -> NaHS04 + HCl. The hydrogen chloride is led through condensing towers thrqugh which water percolates, thereby forming hydrochloric acid. The solid mass of bisulphate and unchanged chloride is then raked on to the hearth of a reverberatory furnace and heated strongly. Decomposition ensues, and sodium sulphate, known as salt cake, results : NaHSO^ + NaCl-> Na^SOi + HCl. Salt cake contains roughly 95 per cent, of sodium sulphate. The next stage in the operation is to heat the salt cake with limestone and coke. The reaction proceeds really in two distinct stages : Stage 1: Na^SO^ +2C-^Na2S +2C0a Stage 2 : Na^S + CaCOj^ CaS + Na^COa. The black residual mass, consisting of sodium carbonate, calcium sulphide, calcium oxide and other impurities, is known as black ash. The sodium carbonate is removed from black ash by the process of lixiviation. The coarsely broken up black ash is placed in a series of tanks provided with a perforated false bottom. Fresh water flows first upon ash which has already been nearly completely extracted. The hquid from this vat is syphoned upon the ash in the second vat, and so on. The object of the perforated bottom is to allow the concentrated solution of sodium carbonate to collect below the specifically lighter water. It is the underlying concentrated Hquor which is syphoned into the next compartment. The essential feature of the oper- ation is that the partially exhausted black ash is extracted by fresh water, whilst the nearly saturated solution is brought upon the fresh black ash. In this way the minimum quantity SODIUM, POTASSIUM 477 of water is required to dissolve from the black ash practically the whole of the sodium carbonate. A temperature of about 32-37° is maintained as sodium carbonate has its maximum solubility at this temperature. The heat of the furnaces is utilised for the concentration of the solution ; crystals of the monohydrate, NajCOa.HzO, separate out. The monohydrate is either con- verted into the anhydrous salt hy heat or recrystaUised at a temperature below 32° when the deca-hydrate, NajCOsjlOHjO, is obtained. The waste products from the vats contain a considerable amount of calcium hydrosulphide, a substance which evolves hy- drogen sulphide on exposure to air. Several processes have been patented for the recovery of the sulphur from this waste, notably that of Chance. By this method use is made of carbon dioxide, whereby hydrogen sulphide is ultimately Hberated and converted into sulphur by combustion in an insufficient supply of air. Ca{HS) 2 -f CO 2 -f H aO -> CaCO 3 -f 2H ,S ^ 2H2S+02-^2H20+2S The method of Le Blanc is, however, steadily yielding ground to the more modern and more efficient process of Solvay. In this process a saturated solution of ammonia in brine is first prepared by forcing ammonia through a tower down which brine is percolating over a series of baffle plates. After cooling, the saturated brine is then " carbonated " by percolation through a Solvay tower, meeting an upward current of carbon dioxide. The reaction NHiOH + C02->>. ,,' ,-' / G E ^ V C/ / y '' 20-5° Z5-5' 30-5° 35-5° Temperature. Fig. 122. 2NaHC0 3 + 9H ^0 . The bicarbonates, even at the ordinary temperature, exert an appreciable dissociation pressure, and with rising temperature the loss of carbon dioxide increases rapidly. 2NaHC03^Na2C03 + H^O + CO^ ^ Aqueous solutions of these bicarbonates evolve carbon dioxide on warming, a test not given by normal carbonates. Owing to the hydrolytic decomposition of these bicarbonates, NaHCO^ + HOH ^r± NaOH + H2CO3, and the weak ionisation of the carbonic acid, such solutions react faintly alkaline in dUute solution, though concentrated solutions of potassium bicarbonate are practically neutral. Sodium bicarbonate is used in the manufacture of baking powder. When it is mixed with cream of tartar (acid potassium tartrate), an evolution of carbon dioxide takes place. This reaction explains the chemistry of baking powder. Chlorates. — The chlorates of sodium and potassium have long been j)repared by the action of chlorine upon an alkaU but recent developments in electrolysis have almost completely displaced the older chemical process. In the electrolytic decomposition of sodium or potassium chloride the intermixing and interaction of the products of the electrolysis are pro- moted as much as possible. At the working temperature of the cell (70°) interaction between the chlorine and the hydr- oxide proceeds with great velocity. If sodium chlorate is being prepared, the chloride is continually added to the cell untU the solution is nearly saturated with the chlorate (up to 750 gm. per litre). The liquor is then rim out and the chlorate separates on cooling. In the preparation of potassium chlorate, the slight solubility of this substance soon leads to the saturation of the liquid, which is then drawn off and allowed to crystaUise. The mother liquor is re-saturated with potassium chloride and returned to the cell. Owing to the sUght solubUity of potassium chlorate, Nvith its attendant difficulties, some manufacturers SODIUM, POTASSIUM 481 prefer to make sodium chlorate and then convert this into potassium chlorate by treatment with potassium chloride. The applications of potassium chlorate are essentially associated with the strong oxidising action of the chlorate group. Thus, it is employed in the manufacture of matches, as well as in the preparation of fireworks. Sodium chlorate, owing to its some- what hygroscopic nature, finds no application when a solid oxidising agent is required. The action of heat upon potassium chlorate, producing first potassium perchlorate and then the chloride, has already been the subject of discussion (see p. 181). Hypochlorite. — ^Fairly large quantities of sodium hypo- chlorite are now prepared electrolytically as a bleaching agent for textiles, though such liquors are not sufficiently powerful for sewage treatment or other operations where a strong solution is required. The great advantage of the electrolytic bleaching agent is that the fabric is not exposed to the action of acid or alkah. Several processes are at work for the production of this product, but the reactions are essentially the same — the electrolysis of brine at a temperatiu-e of about 21° under such conditions that reduction of the hypochlorite round the anode is prevented. Perchlorates. — The perchlorates are formed by the decom- position of the chlorates by heat, but the electrolytic oxidation of chlorate into the perchlorate is the present commercial method. A strong solution of sodium chlorate is electrolysed between platinum electrodes, and under suitable conditions of current density and of E.M.F. oxidation to the perchlorate takes place. As the sodium salt is deliquescent, it is converted into the potassium salt by the action of potassium chloride. Potassium perchlorate finds use in the fireworks and explosive industries. Both perchlorates form rhombic crystals which are isomorphous with the permanganates. Nitrates. — Sodium nitrate occurs in extensive deposits in ChUi and Peru. The salt is leached out, and the solution evapor- ated in order to crystallise the sodium nitrate from its accompany- ing salts. Commercial saltpetre is used as a fertiliser as well as in the manufacture of potassium nitrate and of nitric acid, whilst the mother liquor obtained during the leaching and crystallising of ChiU saltpetre is used for the preparation of iodine from the sodium iodate contained therein (see p. 160). II 482 AN INORGANIC CHEMISTRY Sodium nitrate is converted into the corresponding potassiiun salt by treatment with potassium chloride. A strong solution of Chili saltpetre is treated ^\'ith a solution of potassium chloride and sodium chloride at once separates out : KCl + NaNOa -^ KNO3 + NaCl i By taking advantage of the slight effect of temperature upon the solubUity of sodium chloride and of the marked influence of temperature upon the solubility of potassium nitrate, it is possible to separate most of the nitrate in a moderately pure state. The mother liquor, still containing more or less of the potassium nitrate, is further concentrated and another crop of crystals of potassium nitrate is obtained. The decay of organic matter, especially in hot countries (Bengal, etc.), also leads to the formation of potassium nitrate. In such countries, nitrifying bacteria, acting upon urea and other organic matter often present in the soil, convert the " organic " nitrogen into nitrates. The soil soon becomes charged with the nitrates of calcium and potassium. These are leached out, potassium carbonate added to decompose any calcium nitrate present, and the filtrate containing potassium nitrate concentrated and crystallised. Potassium nitrate finds considerable use in the manufacture of gunpowder. This substance contains about 75 per cent, of potassium nitrate, 10 per cent, of sulphur, 14 per cent, of charcoal and I per cent, of Mater. The chemical reaction which takes place \\'hen gunpowder is burnt, is supposed to be represented by the equation : 2KN03+S+3C^K,S+N2+3C02. The explosion is occasioned by the sudden liberation of a large volume of gas at a high temperature. When potassium nitrate is heated by itself, oxygen is evolved and potassium nitrite remains, hence its use as an oxidising agent. Sodium nitrate behaves similarly. Potassium nitrate is dimorphous, occurring as rhombic crystals and as rhombohedra (hexagonal) isomorphous with sodium nitrate. Nitrites. — The nitrites of the alkahes are easily obtained by reduction of the nitrate with lead or iron. They can also be formed by the electrolytic reduction of the nitrate. Sodium SODIUM, POTASSIUM 483 nitrite is used extensively in the preparation of organic dye stuffs. Sulphates. — ^The manufacture of sodium sulphate by the Le Blanc process has already been described. It is also a by-product in the manufacture of nitric acid from sodium nitrate. A certain amount of it is also obtained from kieserite (MgSOj.HaO), MgSOi + 2NaCl -^ Na^SOi + MgCl^. On coohng the solution crystals of sodium sulphate are thrown out. Besides the anhydrous form a hepta-hydrate, NaaSOi.THaO, and a deca-hydrate, NaaSOijlOHgO, Glauber's salt, are known (see Fig. 123 and p. 110). 10° 20° 30° Fig. 123. 4-0° 50° --, 50 ' ~~ •~~ ^^ / T ^ %. V n / 1 6^/ ^ S^, i>^ i£e — rs / 1 40 ,^/ % c 37 J / iO '' / 0^/ K4Fc(CN)6 +2CaC03 i takes place ; the insoluble carbonate is removed by filtration and yellow crystals of potassium ferrocyanide are obtained from the filtrate. It is also formed by heating together nitrogenous matter (horn, blood, etc.) with iron filings and potash. The mass is leached and after the evaporation of the solution crude potassium ferrocyanide separates out. SODIUM, POTASSIUM ' 485 This salt is used analytically as a reagent for iron {q.v.). It is of interest because it gives all the reactions of potassium and none of the ordinary chemical reactions of iron. Migration experiments have established that its composition is such that the iron forms part of a complex anion, i.e. in solution the dissociation proceeds thus : K,Fe(CN)e ^=i 4K+ + [Fe(CN))„]^ On treatment with dilute hydrochloric acid, ferrocyanic acid is formed. It separates as a white powder. This acid is compar- able in strength with hydrochloric acid, Uberating oxahc and acetic acids from their salts. The action of moderately strong sulphuric acid upon potassium ferrocyanide with the evolution of carbon monoxide has been discussed (see p. 357). The older method of preparing potassium cyanide was by heating together potassium ferrocyanide and potassium carbonate. K4Fe(CN)e +K2C03->6KCN + FeO +CO2. The product is impure as it always contains small quantities of the cyanate (q.v.). Another method of preparing potassium cyanide is by forcing ammonia through molten potassium carbonate in the presence of carbon. K2CO3 + 2NH3 +4C-^2KCN + SCO +3H2. The main use of potassium cyanide in industry is for dissolving gold and silver from their ores (q.v.), for electroplating and in photography. There is, a priori, no reason for using the potassium salt for these purposes, for it is the cyanide ion which is the effective agent in promoting the solution of gold, etc. Recently manufacturers have prepared a mixed cyanide by acting upon potassium ferrocyanide with sodium. K4Fe(CN)e + 2Na -^ 2NaCN + 4KCN + Fe. This sodium-potassium cyanide has all the solvent properties of the pure potassium cyanide itself, and less of it is required to effect the desired solution. Even this mixed cyanide, however, is being rapidly replaced by sodium cyanide, large quantities of which are prepared by the Castner process. Dry ammonia is passed over molten sodium in the absence of air, yielding sodamide.- 2NH 3 -f 2Na -^ 2NH ^Na + H 3. 486 AN INORGANIC CHEMISTRY The fused sodamide is treated with powdered coal and sodium cyanamide formed. 2NH2Na + C-^Na^CNa +2H2. At a yet higher temperature the excess of carbon reduces the cyanamide to cyanide Na^CNa + C-^BNaCN. Potassium cyanide is very soluble in water, and very deli- quescent. The aqueous solution is alkaline and smells freely of hydrogen cyanide. K+ +CN- +HOH;=iK+ +0H-+ HCN. Hydrocyanic acid is so weak that it is liberated from its com- pounds by the action of the carbon dioxide of the air. Potassium cyanide is a powerful reducing agent, reducing oxides to the metal. SnO 2 + 2KCN ^ 2KCN0 + Sn . Cyanates. — These are formed by heating the cyanide in the air or with a suitable oxidising agent. KCN + PbO ^KCNO + Pb. Potassium cyanate is a white crystalline powder, freely soluble both in water and in alcohol. Ammonium cyanate, NH4CNO, is of interest, for Wohler (1828), on evaporating a solution of this substance, found urea, NH, C0<( ^NHj an isomer of ammonium cyanate. This was the first synthesis of a substance which had hitherto been unknown outside the animal world, for urea had long been known to occur in the liquid excrement of animals as the chief product of decomposition of compounds of nitrogen in the animal body. Thiocyanates are obtained by oxidising the cyanide with sulphur. KCN + S-^KCNS. They are very deliquescent and are used analytically as reagents for ferric iron, with which they give a bright red colouration (see p. 204). SODIUM, POTASSIUM 487 Persulphates and Per carbonates. — The preparation of the persulphates as well as their more important properties have been discussed under the heading of persulphuric acid (q.v.). It is reported that, if a saturated aqueous solution of potassium carbonate is electrolysed at a temperature between — 10° and — 15°, and the anodic liquid evaporated, a bluish white powder separates out of the composition, KgCaOe. The yield is improved by employing a concentrated solution, and a high current density, i.e. a small anode and a high current, to facilitate the crowding together of the KCO"^ ions round the anode. Potassium percarbonate possesses well-marked oxidising properties. With dilute acid it liberates hydrogen peroxide so that it can be used to replace this agent in such reactions as the bleaching of indigo, oxidation of chromic salts to chromates, etc. The corresponding sodium percarbonate is less well known and has probably not yet been isolated in the pure state. Sulphides . — ^The action of hydrogen sulphide upon potassium and sodium hydroxides leads first to the formation of the hydrosulphide. KOH + HjS ^i=± KHS + H^O. Evaporation leads to the separation of crystals of the hydrated hydrosulphide, 2KHS,H20. The normal sulphide is prepared from a solution of the hydrosulphide by the addition of the hydroxide. NaHS + NaOH ^± Na^S + H^O. Owing to the extensive hydrolysis of the normal sulphide an excess of the hydroxide is advisable in order to force the equilibrium to the right before evaporation. Crystals of the pentahydrate separate out. The sulphides are also formed by the reduction of the normal sulphates by means of carbon or hydrogen. Solutions of the sulphides take up an excess of sulphur very readily, forming yellow solutions, from which various polysulphides may be isolated, KjSo, K4S,, KjSf., etc. Sodium sulphide is used technically in the tanning industry to promote the removal of hair from hides. Sodium Thiosulphate. — Some of the methods of preparation of this compound, as well as its more important reactions, have been described under the heading, Thiosulphuric Acid (p. 248). 488 AN INORGANIC CHEMISTRY It is also occasionally made by the oxidation of sodium hydroxide with sulphur : 6NaOH + 12S -> Na^SaOa + 2^(iS: + SH^O. It is also prepared from the Le Blano tank waste containing CaS. This is partially oxidised in solution to the thio-salt which is then treated with sodium carbonate, CaS„03 + Na2C03^-CaC03^^ +Na2S203. The thiosulphate is obtained by the crystallisation of the filtrate. Sodium Silicate. — The meta-silicate of sodium is formed by fusing together silica and sodium carbonate, also by the wet method when freshly precipitated siUcio acid is treated with the calculated quantity of sodium hydroxide. If an excess of silica is fused with sodium carbonate, and the resulting melt boiled for some time, a syrupy liquid, linown as water-glass, is obtained. This form is readily soluble in water. It forms a useful cement and is also used for fire-proofing wood, besides being a valuable egg preservative. Analytical Tests. — With chloroplatinic acid, HaPtClg, the potassium ion gives a yellow precipitate of potassium chloro- platinate, a test none too delicate even in the presence of alcohol. An excellent test is that of Carnot. A test-tube is rinsed with dilute nitric acid and the contents poured out, and two drops of a normal solution of sodium thiosulphate and of bismuth nitrate introduced. Ten c.c. of alcohol are added and then a few drops of a fairly strong solution containing the salt of potas- sium. A copious bright yellow precipitate confirms potassium. This precipitate consists of potassium bismuth thiosulphate, K3Bi(S,03)3. The only salts of sodium which are even relatively insoluble are the fluosilicate, Na.SiFj, and sodium hydrogen pyroanti- monate, NaaHjSboO,. Lithium The chemistry of lithium calls for little comment, for the chemistry of lithium is that of sodium except in so far as the weakened basic nature of lithium oxide and hydroxide modifies LITHIUM 489 its reactions. Thus we find that lithium carbonate parts with carbon dioxide on being heated, as do the carbonates of the calcium group. Next to the phosphate and the iluoride the carbonate is the least soluble salt of Hthium. Lithium gives a series of compounds very similar to those of sodium, e.g. a nitride, oxide, peroxide, sulphide, etc. In nature Uthium is found in small traces in many silicates ; lepidolite (lithia mica) and petalite (Uthium aluminium silicate) form the chief sources of supply of this element and its compounds. The metal is obtained by the electrolysis of the fused chloride, or of a fused mixture of lithium and potassium chlorides. The heat of reaction between lithium and water is not sufficiently high to cause the evolved hydrogen to ignite. This is to be expected from the weakened basic nature of this element as compared with sodium. Otherwise its properties are sufficiently indicated in Table 49. The Spectroscope and Spectrum Analysis. When white hght passes through a glass prism, and the resultant beam is projected upon a screen, the beam of light is sorted out into a band of colours called a spectrum, the colours ranging through red, orange, yellow, green, blue, indigo to violet. This is best shown by allowing the beam to pass first through a narrow slit (Fig. 124). Violet. Indigo. Blue. Green. Yellow. Orange. Red. Slit Fia. 124. In the spectrum it is found that the violet portion is most bent out of its original path. The explanation of this is that white light is composed of vibrations of varying wave length. Light of short wave length (violet) will be most impeded in its passage through the prism, and consequently most swung out of its course ; the red rays of longer wave length will be less deviated from their path. The light of every wave length will give rise 490 AN INORGANIC CHEMISTRY to an image of the slit upon the screen, and as a result, we obtain a broad coloured band, known as a spectrum. If the light comes not from the sun, or an arc light, but from a Bunsen flame coloured by a potassium salt, the continuous spectrum is no longer seen, but merely two bright lines, one in the red and one in the blue portion of the spectrum, accompanied by numerous fainter lines. Each element used to colour the flame, e.g. sodium, Uthium, etc., in whatever compound it may occur, \\'ill give rise to its own characteristic lines and no two elements have ever been found to give the same spectrum Une.s. The instrument whereby it is possible to observe the spectrum directly (i.e. without the aid of a screen) , is known as a spectro- scope, and such is the delicacy of this method of investigation that it is possible to detect 0-00006 milligms. of calcium and even 0-0000005 milligms. of sodium. Careful examination on the part of Frauenhofer (1814-1815) revealed the interesting fact that, whereas the spectrum of the arc light is perfectly continuous, there was no such continuity in- the sun's spectrum, but a large number of dark lines — the Frauenhofer Unes — were seen. The discovery was also made that the position of these lines corresponded in nearly all cases with the position of lines already mapped out for elements known to the chemist in his laboratorj'. The explanation of this is conveyed in the following experiment. If the light from an arc is made to pass through a Bunsen flame coloured yellow with sodium chloride, the spectrum is foxmd to be no longer continuous, but to contain two dark lines in place of the usual two bright Unes in the sodium spectrum, and so on for other elements. It appears, therefore, that the Ught from the arc is robbed during its passage through the yellow flame of those vibrations which normally give rise to the yellow lines in the spectrum. The absorbing power of the yellow flame is thus greater than its emissive power. This experiment supphed the key to the Frauenhofer dark lines . The glowing sim centre gives rise to a perfectly continuous spectrum, as does the arc hght, but during the passage of the sun's light through the glowing, but cooler photosphere, it is robbed of certain vibrations which the photosphere is itself capable of emitting at a lower intensity. The appearance of these dark lines in the sun's spectrum, occurring as they do in exact agreement with the position of known terrestrial elements, is a proof of the existence of these elements in the sun. AMMONIUM 491 When Ramsay discovered helium, it was fomid that its spectrum coincided exactly with certain dark hnes which Lockyer had attributed to the presence of an unknown element in the sun, which he named helium (q.v.). Absorption spectra are obtained by passing the light through a layer of gas or solution before it is examined by the spec- troscope. A series of dark bands or lines is generally to be Rubidium and Caesium The use of the spectroscope as an aid to investigation was soon shown to the world by the discovery of the new elements, rubidium and caesium. Bunsen (1861) and Kirchhoff found several new lines in the blue end of the spectrum given by the salts derived from the Diirkheim mineral waters. To the element giving rise to these lines they assigned the name caesium. Shortly after, in working up the mineral lepidolite, they obtained spectroscopic evidence of yet another element — rubidium. Smee then, other sources of supply have been discovered. Rubidium has been found in the ashes of plants and in many mineral waters. Caesium, besides occurring with rubidium in the above sources of supply, is found in pollux, a rich caesium-aluminium silicate occurring in America. The properties of these two elements and of their compounds are strongly akin to those of potassium. This is shown by their tendency to form polyiodides, KI+I,^=±Kl3 CsI-fI,^=^CsI, RbI + l2^=±Rbl3, as well as a few higher polyiodides, such as Rbl,. Rubidium and caesium also form difficultly soluble salts with chloroplatinic acid, HaPtCle- Ammonium The compounds of the ammonium radicle show sufficient relationship to those of the alkalies to justify their inclusion at this stage. Ammonium Amalgam. — If a little sodium amalgam, prepared by dissolving a small piece of sodium in mercury, is put into a solution of ammonium chloride, it is found to swell 492 AN INORGANIC CHEMISTRY up to several times its original volume. A similar result is obtained if a little mercury is put in ammonium hydroxide, the negative pole of a battery dipped into this and the solution electrolysed. This is supposed to be due to the formation of an ammonium amalgam. The action of this amalgam upon solutions of copper sulphate, zinc sulphate, etc., supports this view, for some of the metal is precipitated, and ammonium sulphate i)asses into solution. A mixture of hydrogen and ammonia exerts no such action, nor will such gases dissolve in mercurj'. Ammonium Halides Ammonium Chloride. — This salt is prepared by the action of ammonia upon hydrochloric acid. Much of it is obtained from the gas liquor (p. 277), the crude material being purified by sublimation. It is heated in large iron or earthenware pots and collected on the dome-shaped cover. The volatUised product forms a tough, fibrous, crystalline solid. When prepared by crystallisation, it occurs in cubes or octahedra. The action of heat upon this .substance is in this case more than mere sublima- tion, for it has been found that the density of the vapour is such as to indicate a breaking down of the molecule mto ammonia and hydrogen chloride. That this is so can be shown by taking advantage of the difference in the rates of diffusion of these gases (Graham's Law, p. 83) (Fig. 125). Fig. 125. A stem of a clay pipe is passed through a glass tube and held in position by means of t«o corks. Between the pipe and the tube is placed a little ammonium chloride. If a piece of Utmus paper is placed in the compartment with the ammonium chloride in a AMMONIUM 493 short time it will be found that the paper has turned red, whilst, if a gentle stream of air is. sent through the pipe, the issuing gas will be found to turn red litmus blue, i.e. molecules of ammonia have diffused through the porous walls more rapidly than the heavier molecules of hydrogen chloride, leaving a predominance of the latter in the outer compartment. Ammonium Bromide and Iodide show properties similar to those of the chloride, except that in the case of the iodide the heat which dissociates the iodide is sufficient to liberate iodine by its action upon the equilibrium 2HI^iziH, + I,. Ammonium Hydroxide has often been referred to as a weak base, its dissociation into the ions NHi"^ and 0H~ being very slight. The effect of adding an ammonium salt to a solution of ammonium hydroxide is discussed on page 524. Ammonium. Nitrate is formed by adding nitric acid in equivalent quantity to a solution of ammonium hydroxide. Four different types of crystals are known, each possessing a distinct transition temperature. Those existing under ordinary conditions of temperature are rhombic in form, isomorphous with potassium nitrate. Ammonium nitrate is most extensively used in the manufacture of fireworks and explosives. Heat causes decomposition in accordance with the equation : NH.NO 3 -^ N 2O + 2H 2O . Ammonium Nitrite is somewhat unstable, though it can be prepared by saturating ammonium hydroxide with nitrous acid. The solution must be concentrated in vacuo at ordinary tem- peratures. Ammonium Carbonate and Bicarbonate. — Ammonium bicarbonate is formed during the manufacture of sodium car- bonate by the Solvay ammonia-soda process (q.v.) It can be readily obtained by saturating a solution of ammonium hydroxide with carbon dioxide and concentrating the solution. The salt smells appreciably of ammonia, owing to the dissociation NH^HCOa-^NHj + HaO+COa. In aqueous solution far reaching hydrolysis of this salt takes 494 AN INORGANIC CHEMISTRY place. This is due to the weakness of ammonium hydroxide as a base and of carbonic acid as an acid : NH,+ +OH- 2H++CO3- 11 11 NH.HCOa+HOH ^=± NH.OH + H2CO3 NH3+H2O CO2+H2O On account of the volatility of the ammonia and carbon dioxide, these gases escape on warming the solution, i.e. ammonium bicarbonate may be completely broken down in solution by the action of heat. The dry salt suffers a similar decomposition. Ammonium carbonate, (NH4)2C03, is generally contaminated by the presence of small quantities of ammanium carbamate. The following formulae bring out the relationship between ammonium carbonate, ammonium carbamate and urea. 0— NH, 0- NH4 NH2 o=c/ 0=0/ o=c<; ^0- NHi ^NHs ^1* Ammonium carbonate. Ammonium carbamate. The carbonate may be freed from the carbamate by means of alcohol, in which the carbamate is freely soluble. Ammonium Sulphate is obtained in considerable quantities by neutralising the ammonia of the gas Uquor with sulphuric acid. It is used extensively as a nitrogen fertiliser. Ammonium Sulphide and Hydrosulphide. — The prepara- tion of these salts offers no striking difference from the method adopted for the corresponding salts of potassium and sodium. The hydrolysis of the sulphide, formed as it is of a very weak acid and base, is unusually large : Urea. NH ++0H- NH/ +HS- 2NH4+ +S- ^zz± (NHJjS + Hi,0 ^iziNH.OH + NH.HS NH3 + H2O NH3+H2S AMMONIUM 495 Moreover, the products of the hydrolysis, the hydroxide and the hydrosulphide, besides being in equilibrium with their ions are also in equilibrium with the dissolved gases, ammonia and hydrogen sulphide. Since these gases are readily expelled by the action of heat, it follows that an aqueous solution of this salt may be entirely decomposed by boiling, at times a great advantage in analytical work. The action of ammonium sulphide in effecting the separation of arsenic, antimony and tin from the rest of the " insoluble " sulphides by the formation of the thio-salts has already been discussed {q.v.). Phosphates. — Of the numerous phosphates of ammonium the most important is the secondary sodium ammonium phos- phate, NaNHjHPOi. This is used in bead tests, and is known as microcosmic salt. During the heating the following decom- position occurs : NaNH4HP04->NaPO.,-|-H20 +NH3 the metaphosphate polymerising into a hexa-metaphosphate. Questions 1. Discuss the use of the spectroscope in qualitative analysis. 2. Describe any electrolytic process which is employed on a large scale for the manufacture of caustic soda. 3. What important salts of sodium and potassium are prepared electro- lytically. Briefly indicate the salient features of each method. 4. Account for the great success of the Solvay ammonia-soda process. Why is this method not satisfactory for the manufacture of potassium carbonate ? 5. Compare the electrolytic and chemical methods of manufacturing caustic soda from the point of view of efficiency, economy and purity of product. 6. An aqueous solution of sodium sulphide is alkaline but the hydro- sulphide gives a neutral solution. Account for this. 7. Describe the modern method of manufacturing sodium cyanide. What is the commercial importance of this substance ? CHAPTER XXXII SUB-GROUP IB : COPPER, SILVER, GOLD General Principles of Metallurgy. — Metallurgy may be defined as the art of extracting metals from their ores on a commercial scale. In many cases there is less difficulty en- countered in producing the crude metal than in freeing it from small quantities of deleterious impurities. It is reported that even 0-05 per cent, of bismuth in gold is sufficient to render that metal valueless for the purpose of coinage, so brittle does it become. Comparatively little difficulty is experienced in separating the crude metal when it exists in the free state, scattered though it may be throughout more or less gangue. Gold is an excellent example of a metal occurring native. It is readily removed from the powdered gangue by treatment ■with cyanide (see p. 516) or by the process of amalgamation. In other cases the metal is separated by fusion with or without the addition of a fiux, and the layers of metal and slag are then easily separated by taking advantage of the difference bet\^een their specific gravities. A slag arises from the fusion of various types of ore, and is generally a silicate, occasionally a phosphate (see Thomas' Basic Slag, p. 629). If the ore contains an excess of silica, as is usually the case, limestone is added as the flux, and combination then ensues between the lime formed in the furnace and the excess of silica. Occasionally borax or fluorite is added as a flux ; with the former there is direct combination between the borate and the oxide (cf. borax bead) to bo removed, the use of the latter being physical rather than chemical. In general, sulphide ores are first broken down by roasting, i.e. heating in an atmosphere of air to eUminate sulphur as much as possible and expel volatile impurities such as arsenic. Other impurities are often converted into oxides, and as such, are brought into combination with the flux and removed as slag. 496 COPPER, SILVER, GOLD 497 Several important methods stand out for the purification or refining of the crude metal. One of these is the process of electrolysis (cf. copper, zinc), wherein the crude metal forms the anode, whilst the solution contains a suitable salt of the metal to act as electrolyte. A high degree of purity can be attained by this means. In a number of cases, impurities are removed by the oxidising action of a blast of air at a high temperature, as in the conversion of cast or pig iron into wrought iron, copper matte into copper. During recent years the commercial preparation of metals by electrolysis has come to the fore. Aluminium is now exclu- sively made by the electrolysis of aluminium oxide dissolved in molten cryolite, sodium from the fused hydroxide, magnesium from fused carnalHte, etc. CoppEE, Silver, Gold The position of these elements in the first group of the Periodic Table appears at first sight somewhat unjustified. The ease with which these metals part with their oxygen, their inactivity towards this element, the comparative weakness of their oxides as bases stand in sharp contrast to the metals of Group 1a. a great deal of this difficulty disappears if the method of writing the Periodic Table, given on p. 260, is adopted, of which the following is a modification : — lA 2A 3A 4A 5A 6A 7A 8 IB 2B SB 4B 6B 6B -B He Li Ge B C — — — — . — — . — — N F Ne Na Mg Al Si — — — ^ * ^ — — . — — P S CJ A K Ca Sc Ti v f'r Mil i'e Co Ki Cu Zn Ga Ge As Se Br Kr Eb Sr Y Zr Nb Mo — Ku Eh Pd Ag t'd In Sn Sb To I etc. The change in basicity in passing from sodium to chlorine is, in the long series, spread over seventeen elements and copper is seen to fall mid -way between the element, potassium, and the element, bromine. In brief, the elements of Group 1b are not to be expected to form compounds closely akin to those of the alkalies, except in so far as they give rise to the group oxide MjO, and a series of salts of the same valence as do the alkalies, e.g. CuCl, AgCl ; with regard to the actual properties of such compounds, they are to be expected to be possessed of a distinctly weaker electro-affinity. Their basicity will, in general, lie intermediate between the elements noted for the strength of their oxides as bases (the alkaHes) and the elements whose oxides are acidic (the halogens). KK 498 AN INORGANIC CHEMISTRY The following statement summarises the more important properties of Group 1b : — 1. The elements give rise to two or more oxides, and to two series of salts, viz. : — Oxides. Salts . CuaO. Cuprous oxide. CuCl, Cul, etc Cuprous salts. CuO. Cupric oxide. CuClj, CuSOi, etc. Cupric salts. Ag40. Silver suboxide. Ag,F. Sub-halide sil- ver salts. Ag^O. Silver oxide. AgCl, AgNOj, etc. Normal silver salts. Ag^O^. Silver peroxide. AujO. Atirous oxide. AuCl, etc. Aurous salts. AuO. Gold monoxide. AuSOj, AuCl 2 etc. Gold monosul- phate, etc. Au,03. Auric oxide. AuClo, etc. Auric salts. The element with the heaviest atomic weight gives rise to the greatest variety of salts ; this is in conformity with the behaviour shown by other elements at the bottom of Groups 2b, 3b, 4b, etc. 2. The carbonates and hydroxides are practically insoluble in water, silver hydroxide alone showing an appreciable solubility (contrast the alkaU hydroxides and carbonates) ; the carbonates evolve carbon dioxide on heating and the hydroxides readily break dovm into the oxides (cf. the carbonates and hydroxides of the alkalies). 3. The metals, excepting copper, show no tendency to combine with oxygen. 4. The ions of silver and copper combine with ammonia to form complex cathions : Cu + + + mNH 3 ^- [Cu. wNH 3] + + 5. The halides and cyanides of these elements are of such weak electro -affinity that, in the presence of the halides and cyanides of the alkalies, they are forced into the complex anion : KCl+AgCl->K[AgCl2] 2KC1 +CuCl2^K2[CuCl,] AgCN + KCN -^ K[Ag(CN) ,] 6. In all cases solutions of salts of the higher stage of oxidation are reduced to the lower by heating with the metal : Ag++Ag^i:±Ag,+ Cu++ -f Cu^rz±2Cu+ Au+ + + + 2Au ^rr± 3Au+ COPPER 499 Copper Occurrence. — Copper occurs native in Chili in the form of a copper sand containing from 60-90 per cent, of metal, also in large quantities near Lake Superior. It is widely distributed as ruby copper, CuaO (U.S.A.), melaconite, CuO (U.S.A.); as basic carbonates — malachite, CuC03,Cu(OH)2, and azurite, 2CuC03,Cu(OH)2 ; as sulphides in chalcocite or copper glance, CujS, and chalcopjrrite, CuFeSa. Small quantities are also found in the feathers of some birds. Physiologically, its occurrence in the hsemocyanin of the blood of the mollusca is interesting ; its role is evidently similar to that of the iron salts in the hsemoglobin of red blood. Metallurgy. Copper ores, free of sulphur, are comparatively easy to smelt. The ore is mixed with coke and any necessary flux, put in a blast furnace, and the result of the operation yields a slag and bUster copper (98 per cent. Cu). Only a single chemical operation is necessary in this step — the reduction of the oxide by the carbon. Such ores are, however, comparatively rare. The metallurgy of the sulphide ores of copper is a much more difficult process. This arises from the great tendency of sulphur to combine with copper ; the affinity between these two elements is so great that copper will expel iron from iron sulphide. The actual method chosen depends upon the grade of ore, the impurities present and so on. Pyrite Matte Smelting. — The charge, consisting of the sulphide ore, containing copper and iron with traces of gold and silver, is brought into the blast furnace without any preliminary roasting. A large part (70-80 per cent.) of the sulphur is burnt off, the heat of combustion of this sulphur maintaining the furnace at the requisite temperature (Fig. 126). The remainder of the 500 AN INORGANIC CHEMISTRY sulphur is left in combination with the copper and the iron, formuig the artificial sulphide known as matte. In this process the slow expensive preliminary roasting is avoided. Within the furnace a large part of the iron is converted into the oxide FeS + 0->FeO + S, which unites vnth the silica to form a slag. The matte produced in the process contains 40 per cent, of copper, whUst as much as 95 per cent, of the iron originally present is converted into the oxide and removed in the slag. The pyrite matte is then converted into bUster copper before refining. Fig. 127. Reverberatory Smelting. — Ores which are at all fine in grain and therefore unsuitable for treatment in the blast furnace, are generally converted into matte by reverberatory smelting. Considerable quantities of the oxides of sulphur escape, the sulphur which remains is found in combination with the copper and the iron. Silica which is often added to the furnace charge removes the iron and other oxides as slag (Fig. 127). Blister Copper. — The molten matte is run into a converter COPPER 501 furnace, lined with a siliceous lining. Air is blown through the molten mass, thereby oxidising the iron and other easily oxidis- able metals. These are removed by the siliceous lining as slag. After the iron has been sufficiently slagged, the furnace is tilted and the slag run off. At this stage the copper content has risen to about 75 per cent., and the iron has been almost eliminated. After the slag has been rxm off, the blast is again turned on, when the reactions, 2Cu2S+302-^2Cu20+2S02 CuaS + 2Gu20^. 6Cu + SO2. take place. As soon as the appearance of the flame indicates the completion of the above interactions, the blast is shut off and the copper poured into moulds. As it cools, occluded sulphur dioxide escapes, giving the copper the well known blistered appearance. Refining of Blister Copper. — Blister copper contains cuprous oxide, sulphur, iron, arsenic, and often silver and gold. If the precious metals are present in sufficient quantity, the blister is refined electrolytically, otherwise the refining is carried out in the copper refining furnace. Compressed air is forced through the molten mass, thereby removing the iron, arsenic and sulphur either by volatOisation or by oxidation. The oxides of copper are reduced by poHng, i.e. forcing poles of wood below the surface. The hydrocarbons of the wood effectively reduce the last traces of copper oxide. Electrolytic Refining. — Blister copper, especially that con- taining gold and silver, is frequently refined by this means. The electrolyte is copper sulphate, acidified with sulphuric acid. Anodes formed from blister copper and cathodes of electrolytic copper are suspended alternately in the bath. During the electrolysis copper dissolves from the anode and is precipitated on the cathode. At the anode the reaction Cu + 2©->Cu+ + takes place. Copper sulphate is therefore formed, while at the cathode Cu+ + ->Cu + 2e occurs, i.e. copper is precipitated. The electrolyte consequently maintains its concentration unaltered. The impurities, gold and 502 AN INORGANIC CHEMISTRY silver, present in the blister copper, coUect in the anode mud which is periodically removed and worked up for the noble metals. Properties. — Copper is one of the best conductors of heat and electricity, but this property is modified materially by the presence of traces of many impurities, e.g. arsenic. The metal is fairly hard, and appears red by reflected light. It can be obtained in regular crystals (octahedra). The melting point of copper is 1,084°. Copper is of great industrial importance. Enormous quantities are used in the electrical industry, as well as in the manufacture of many alloys (brass, 1 of zinc to 2-5 of copper ; nickel and copper coins, bronze, etc.). Oxygen has no action upon copper when dry, but a thin film of cuprous oxide is soon formed if the oxygen is moist. Of the dilute acids, nitric acid alone has an appreciable solvent action upon it, liberating nitric oxide (q.v.). Strong hydrochloric acid in the presence of air (oxygen) dissolves it slowly. The thin fUm of hydrogen produced by the reaction Cu+ 2H+->Cu + + +H5 is oxidised by the dissolved oxygen, and hence solution proceeds slowly. The action of strong sulphuric acid and of nitric acid has already been discussed ; the former evolves sulphur dioxide, the latter nitrogen peroxide. The low electro-potential of copper explains its inability to hberate hydrogen from acids unless aided by some suitable oxidising agent (NO3, O^, etc.). Hydrogen, dissolved or occluded in platinum or in charcoal, will precipitate copper from a solution of copper sulphate. General Stjbvey of the Compounds of Copper. There are two basic oxides of copper — cuprous oxide, CuaO and cupric oxide, CuO. From each of these oxides there is derived a series of salts, the cuprous salts in which copper is monovalent, and the cupric salts in which copper is divalent. As examples of the cuprous salts, we have cuprous chloride, CuCl, cuprous bromide, CuBr, cuprous iodide, Cul, cuprous cyanide, CuCN, and cuprous sulphate, Cu^SOi. The cupric salts are much more numerous, for besides the haUdes, CuCl 2, CuBr 2, many salts from oxy-acids have been isolated, such as the nitrate, Cu(N03)2. COPPER 503 The chemistry of the cupric and cuprous salts is closely bound up with the equation Cu+ + + Cu^=±2Cu + , that is, cuprous salts are prepared by the reduction of cupric salts with metalUc copper. The above equiHbrium is swung to the right by a rise of temperature (i.e. the formation of the cuprous compound is attended by an absorption of heat). It is therefore advisable to use as high a temperature as possible to effect the reduction. In all the cuprous salts the compounds possess properties characteristic of other similarly constituted salts, whUst the cupric salts show a general similarity to other salts derived from a divalent cathion. Hence we find a striking resemblance in the insolubihty of the cuprous halides, the silver hahdes, the aurous halides, whilst the cupric halides are very similar to the salts of lead, zinc, cadmium and mercury. This rule is a general one. When an element forms salts in which it exhibits more than one valence, the properties of the salts of each type may be fairly safely inferred from a knowledge of the properties of the salts of a corresponding type. Lead subchloride, PbCl, is insoluble (cf. CuCl, AgCl), lead chloride, PbCla, is relatively much more soluble (cf. CuCl 2, HgClj), lead tetrachloride, PbClj, is a liquid, strongly hydrolysed in water hke other such compounds (cf. SnCli, SiCli). Cupric Compounds Cupric Oxide and Hydroxide. — Cupric oxide is a black substance, formed by heating copper in a stream of oxygen, or by decomposing certain oxy-salts, e.g. nitrate, carbonate. Its hydroxide is obtained in the wet way by the action of a soluble base upon a soluble compound of copper : Cu++ +20H-->Cu(OH)2i. The hydroxide is imstable and passes readily into the oxide. The hydroxide dissolves freely in ammonia with the formation of a deep blue solution. Migration experiments in a U-tube have shown that the blue ion is positively charged, and consists of the complex (Cu.2NH3)+^. Cupric oxide plays an important role in the combustion of carbon compounds, the hydrogen of which it is desired to oxidise to water, the carbon to carbon dioxide. The carbon compound is burnt in an atmosphere of 504 AN INORGANIC CHEMISTRY oxygen, and the products of combustion are led over heated copper oxide in order to complete the oxidation to carbon dioxide and water, which are then absorbed and weighed. Cupric Chloride. — Cupric chloride is obtained in the anhy- drous form, either by the direct combination of chlorine with copper, or by dehydrating the hydrated chloride, CuCl2-2H20, in a stream of hydrogen chloride. The presence of the hydrogen chloride prevents the hydrolysis of the chloride and the formation of a basic chloride. The crystals are brownish yeUow and very hygroscopic. Hydrated cupric chloride is obtained by dissolving copper in hydrochloric acid to which a little nitric acid has been added to oxidise away the fihn of hydrogen and so promote the solution. The hydrated salt forms bluish green crystals. Dilute solutions of this compound are blue, and during the electrolysis of such solutions the copper migrates to the cathode, i.e. they contain the cupric ion Cu.''"'' But in strong solutions of cupric chloride, or in solution to which considerable quantities of hydrochloric acid, potassium chloride, etc., are added, the colour is a deep green. Electrolysis of such solutions reveals that much of the copper present migrates to the anode, i.e. is part of a complex anion. CuCl2 + 2Cr ^=:^[CuCl4]- Deep green. Increased concentration, either of cupric chloride or of the chlorine ion, or a rise in temperature, promotes the formation of the complex. Another type of complex is produced from a solution of copper chloride by the addition of ammonium hydroxide. At first a basic salt is thrown down, but this dissolves in an excess of ammonium hydroxide, forming a deep blue solution. This phenomenon is given by nearly all the copper salts to which an excess of ammonium hydroxide has been added, and arises from the presence of the deep blue complex cathion (Cu.4HN3)'+. Many salts of the type Cu(NH3)^A 2 have been isolated, where x varies from 2 to 6, A denotes the anion, C1-, etc. Cupric Bromide and Iodide. — Anhydrous cupric bromide forms jet-black crystals, and is best obtained by the cautious dehydration of the hydrated salt in a stream of hydrogen bromide. To obtain the hydrated salt, CuBr24il20, hydro- bromic acid is saturated with copper hydroxide or copper COPPER 505 carbonate and the liquid concentrated at a low temperature. Solutions of this salt show the same colour changes as do the chloride, but to a more marked degree. Dilute solutions are blue, strong solutions are a deep brown. The colour change from blue to brown is intensified by the addition of such chlorides and bromides as KCl,NaBr, etc. Cupric iodide is unknown. Attempts to prepare it, e.g. by the action of potassium iodide upon a solution of copper sulphate, invariably lead to the formation of a mixture of cuprous iodide and iodine : 2CUSO4 + 4KI^ 2CuI ^+12-^+ 2K2SO4. The iodine is easily removed either by the use of a suitable solvent, by reduction with sulphur dioxide, or by solution in an excess of the added iodide. Cupric Cyanide. — Cupric cyanide exhibits properties some- what similar to the iodide. The addition of potassium cyanide to a solution of copper sulphate causes the precipitation of the unstable cupric cyanide : CuS04+2KCN->Cu(CN)2 i +K2SO4. Immediate decomposition of the imstable cyanide takes place 2Cu(CN)2->2CuCN ^^ + (CN)^ ^ In the presence of an excess of potassium cyanide the cuprous cyanide passes into solution, forming potassium cuprocyanide. In the formation of this salt the following equilibria must be considered : CuCN + KCN K Cu^ -CN- K+ Cu(CN)2] CN- K++[Cu(CN),]- Cu+ +2CN- The primary cause of the cuprous cyanide passing into solution is the disturbing of the equilibrium CuCN ^=± CuCN Solid. Dissolved, by the action of the CN" ions in promoting the formation of the complex anion : CuCN + CN-^[Cu(CN)2]- 506 AN INORGANIC CHEMISTRY In such a solution the cuprous ion concentration is so extremely low that, when hydrogen sulphide is bubbled through the solution, no precipitate falls out. The dissolved hydrogen sulphide gives rise to the equUibrium H,S^z^2H++S-, but the cuprous ion is so low in concentration that the solubility product of cuprous sulphide is not reached, or, expressed mathematically, (Cu+)2(S = ) 2H2O + N2O4. The three hydrates and their vapour pressure-temperature graphs have already been dealt -ndth (p. 98). Cupric sulphate combines with the sulphates of the alkalies and of ammonium to form double or complex sulphates, CuSOi,Na2S04,6H20. These occur in large light blue, monocUnic crystals. Copper sulphate (blue vitriol) is used extensively in preserving timber, as a germicide and insecticide, for electroplating, as a mordant in dyeing (q.v.) and in caUco printing. Cupric Carbonate .^Cupric carbonate occurs only in a basic COPPER 507 form. Thus the addition of a solution of sodium carbonate to a solution of copper sulphate produces a precipitate of the com- position, CuC03,Cu(OH)2,a;H20. The non-separation of the normal carbonate is no doubt closely associated with the hydrolysis which such a salt, formed of a weak base and a very weak acid, would undergo. Cupric Acetate. — A basic cupric acetate Cu3(OH)2.(C2H302)4, verdigris, is produced by the action of acetic acid (vinegar) upon copper under the oxidising influence of the atmosphere. Green crystals of the normal acetate, Cu(C2H302)2,H20, can be obtained by crystallising the basic acetate from acetic acid. The basic salt has a certain commercial importance, as it is used in the preparation of Paris Oreen, Cu(C2H302)2>Cu3(As03)2, a double acetate-arsenite of copper which is used as a fungicide. Scheele's green, CuHAs03, an arsenite of copper, has similar properties. Cuprous Compounds Cuprous Oxide. — This oxide can be prepared by the reduc- tion of cupric hydroxide with glucose. Cuprous oxide is a bright red, crystalline powder which, in naturally occurring specimens, crystallise in the regular system (octahedra). It dissolves in hydrochloric acid, forming the soluble hydrogen cupro-chloride Cu 2O -f 2HG1 -> 2CuCl + H 2O CuCl-fHCl->H[CuCl2] It is also soluble in ammonium hydroxide, forming a cupro- ammonia hydroxide, e.g.' Cu(NH3)^0H. Under the action of the oxy-acids, decomposition in the sense of the equation 2Cu+^z±Cu+Cu+ + occurs. Hence sulphuric acid gives cupric sulphate and metallic copper, not cuprous sulphate. Cuprous Sulphate. — Cuprous sulphate is formed by the action of metallic copper upon a hot solution of cupric sulphate, but on cooling the reaction reverses, and copper is precipitated. Cuprous sulphate has been isolated by boiling dry dimethyl- sulphate mth cuprous oxide : Cu^O + (CH3)2S04^Cu2SOi + (CH3)20 Dimethyl ether. It is only stable in the absence of water. 508 AN INORGANIC CHEMISTRY Cuprous Chloride and Bromide. — These salts are both obtained by the reduction of the corresponding cupric salts with metaUic copper. Copper, cupric haUde and hydrochloric acid are boiled together and the insoluble hahde separates out. These salts are insoluble in water, but are appreciably hydrolysed by hot water, giving hydrated cuprous oxide : 2CuCl + HOH ^i^ Cu^O + 2HC1. At the same time cupric chloride is formed, in accordance with the equation : 2CuCl^:r±Cu+CuCl2. These halides dissolve freely in solutions of hydrochloric acid and of the alkah haUdes. The mechanism of the reaction resembles that involved in the solution of cuprous cyanide in potassium cyanide. The solution obtained by dissolving cuprous chloride in hydrochloric acid is used in gas analysis as an absor- bent of carbon monoxide. Cuprous chloride and bromide are also formed by the dissocia- tion of the corresponding cupric salt on heating : 2CuCl2->2CuClH-Cl2. Sulphides. — Corresponding to the two oxides, CuaO andCuO, are the two sulphides, cuprous sulphide, Cu^S and cupric sulphide, CuS. Cuprous sulphide is obtained from the latter by reduction in a hydrogen atmosphere, but it can also be formed in the wet way by the action of a solution of ammonium sulphide upon copper turnings in the absence of air. Cupric sulphide, as prepared in the laboratory by the action of hydrogen sulphide upon a solution of copper sulphate, always contains more or less cuprous sulphide, due to the reaction, 2CuS-^Cu2S-f S SiLVEE Occurrence. — As one of the noble metals, the widespread appearance of this element in the native state is only to be expected. Native copper always contains smaU quantities of silver. In combination with other elements it is found as silver glance, AgaS, in which form it is often associated with galena, PbS. Less important ores are horn silver, AgCl, pyrargyrite, a thioantimonite, AggSbSa, and proustite, a thioarsenite AgsAsSa. SILVER 509 Metallurgy. — Considerable quantities of silver are recovered from the anode mud precipitated during the electrolytic refining of copper. This mud is dried and thrown into molten lead. The silver (and gold) passes into the fused lead and is recovered by cupellation, i.e. the lead is oxidised into Utharge, which acts as a slag in removing the oxides of other metals. The silver (gold) button is left unoxidised. Silver sulphide occurs in a more or less pure state, but is even more extensively found with galena. Silver ores, free of lead, are often treated with lead-bearing ores, the lead of one ore being used to collect the silver of the other. Sulphide ores are first roasted in order to effect oxidation. The charge of roasted ore, together with coke and sufficient lead-bearing ore to raise the lead content to at least 10 per cent., is introduced into the blast furnace with iron ore and Hmestone as a flux. The process within the furnace is represented by the equations : PbS-fFeO+C^Pb+FeS+CO PbSOi -f FeO -f 5C -^ Pb + FeS + SCO. On the passage down through the furnace the lead collects the reduced silver, and is drawn off at the bottom. If the silver ore is a very rich one, the silver content will pay for extraction. This is done by cupellation. Otherwise the silver-lead alloy is concentrated by the Parkes' or the Pattinson Process. (a) The Parkes' Process. — This process is extremely satis- factory and is based upon the greater solubihty of silver in zinc than in lead (cf . the partition of iodine between carbon disulphide and water, p. 112). A lump of zinc is thrown upon the molten lead-silver alloy, and practically the whole of the silver passes into the upper zinc layer. When the temperature has fallen sufficiently for the zinc-silver layer to sohdify, it is removed, the zinc separated by distillation and the silver remains in the retort. (6) The Pattinson Process. — This is less satisfactory than the Parkes' process, except for ores containing bismuth. The process is best understood by reference to Fig. 128. A represents the melting point of water. This melting point is depressed by the addition of a solute, e.g. silver nitrate. AB marks out the freezing points for solutions containing increasing quantities of silver nitrate. At all points along this curve the solid which separates out is ice. C represents the melting point 510 AN INORGANIC CHEMISTRY of pure silver nitrate, CB the freezing points of fused silver nitrate depressed by the addition of increasing quantities of water. The solid which separates along CB is always silver nitrate. At B, both ice and silver nitrate can exist side by side, for at that point the two curves AB and CB intersect. All solutions to the left of B will deposit ice on freezing, to the right of B silver nitrate. The Pattinson process is based upon a similar temperature-concentration diagram (Fig. 129). '^ / / / / / / / V / ^ ? / / 6 / tt / / MP A 1 / 0' / \ / \ / / 'a 100% Water Concentration Fig. 128. lOOX 100% Silvernitrate Pb UP. 1 1 M.P r / / / A ^ / / / / / ) / / A / / / / / / i H _j _ Concentration Fig. 129. MP I00\ Ag A represents the melting point of pure lead, C that of silver, AB the freezing point curve of melts containing increasing quantities of silver. Any melt of a concentration to the left of B will deposit crystals of lead on cooling. C is the melting point of pure silver, CB the freezing point curve of silver to which in- creasing quantities of lead have been added. All melts of a concentration lying to the right of B wiU deposit crystals of silver on cooling. Melts of the composition B soUdify to a heterogeneous mass of lead and silver crystals. The lead obtained from the smelting furnace has a cojicen- tration to the left of B. On cooUng such a melt, crystals of pure lead separate out. These are removed by means of a ladle, and the composition of the fused remainder must consequently change towards B. At B no further separation can be made. The result of the operation is that a silver-lead alloy has been SILVER 511 obtained which is sufficiently rich in silver (2^ per cent. ) to pay for cupeUing. The lead is also recovered from the litharge left in the cupellation. Wet methods are also used for the extraction of silver, such as extraction with cyanide solutions. AgCl + 2NaCN->Na[Ag(CN)2] + NaCl AgaS +4NaCN-^2Na[Ag(CN)2] +Na2S. The silver is precipitated by means of zinc. In Ziervogel's process the ores are roasted to convert the silver sulphide into sulphate. This is then dissolved out, and the sUver recovered by treatment with scrap copper. Other ores are roasted with salt and the silver chloride dissolved by means of sodium thiosulphate or a strong solution of brine. AgCl + NaCl^Na[AgCl2]. In Mexico, owing to the scarcity of fuel, the Patio method is used. The chemical reactions are summarised in the equa- tions : CuCl2 + AgS -> CuS + 2AgCl 2 AgCl + 2Hg -^ Hg2CU + 2Ag. In the actual operation the finely powdered ore is treated with salt. After standing for a day mercury is added together with a mixture of copper and iron salts. After this mixture has been thoroughly incorporated by treading, the mud is washed away, the amalgam filtered and then distilled. The silver remains, the mercury is recovered. Properties of the Metal. — Silver is a white, lustrous metal of great malleabihty and ductihty. It melts at 960°. Molten silver absorbs oxygen freely, but gives it off upon cooling, causing the well-known spitting. It is an excellent conductor of heat and electricity. Owing to its softness, it is generally employed in the form of an alloy. Silver coinage is a copper- silver alloy which contains 90 per cent, of silver in U.S.A., and formerly 92-5 per cent, in Great Britain, but owing to the enhanced value of silver, the percentage of this metal in British coinage has recently been reduced to 50. Frosted silver orna- ments are made by first heating the ornaments and then immers- ing them in sulphuric acid, thereby dissolving out the copper which has been oxidised by the preUminary heating. ' ' Oxidised ' ' 512 AN INORGANIC CHEMISTRY silver is made by dipping silver in a solution of potassium hydro- sulphide, whereby a thin film of silver sulphide is formed. Silver does not combine with oxygen or water, but reacts freely mth their analogues, sulphur and hydrogen sulphide. Dilute acids have httle action, but it dissolves freely in the presence of moderately strong nitric acid, and of hot sulphuric acid (note the presence of the oxidising agent to remove the hydrogen film). The hydroxides of the alkaU metals have no action upon it, hence its use for alkah fusions. Oxides of Silver. — Silver forms three oxides — the suboxide AgiO, the oxide Ag20, and the peroxide Ag202. The peroxide is formed by the action of ozone upon silver foil, and is also deposited at the anode in a rather impure state during the electrolysis of silver nitrate. It gives rise to no salts. The normal oxide, Ag20, forms a large number of salts which have the characteristic properties of this type of salt. From this oxide are derived the hahdes, e.g. AgCl, the nitrate, carbonate, sulphate, etc. The lowest oxide, silver suboxide, Ag^O, is reported to have been made by the action of steam upon silver subfluoride, AgjF, at a temperature of 180°. 2Ag2F + H20->Ag40 +H2F2. It is very unstable, but possesses a certain measure of import- ance as it gives rise to several salts, e.g. subfluoride, Ag2F, and subchloride, AgjCl. Silver Subfluoride is obtained by heating finely divided silver in a saturated solution of silver fluoride in a platinum dish. Bronze-Uke crystals soon appear on the surface, but on treatment with water they break down. Ag2F->AgF+Ag. The existence of other subhaloid salts of silver has been fairly definitely estabUshed by physico-chemical means ; they possibly play a part in photo-chemistry. Silver Oxide and Hydroxide.— The action of a soluble hydroxide upon a dissolved silver salt is to give a pale brown precipitate which is probably the oxide, Ag20. The alkahne reaction given by an aqueous solution of silver oxide argues in favour not only of the existence of the hydroxide in solution, but also of its strength as a base. The oxide decomposes rapidly SILVER 513 at 250°. Silver oxide, like cuprous oxide, dissolves in ammonia, forming a soluble silver-ammonia hydroxide, which is as strongly dissociated as potassium hydroxide. From this solution there separate on evaporation black, shining crystals of an explosive nature, often called fulminating silver, possibly NAgj. Silver Halides.— These salts are precipitated by the action of a soluble halide upon a solution of a soluble salt of silver. Their insolubihty increases with the atomic weight of the halogen. The chloride is white, the bromide a pale yeUow, the iodide yellow. They are all sensitive to Ught, hence their use in photo- graphy (q.v.). The chloride is freely soluble in ammonia, the bromide less readily soluble, but the iodide is almost insoluble. The solubility arises from the equilibria : AgCI — =± AgCl ;=^ Ag ' -f C;i - Solid. Dissolved. NH,+ + OH- ;=±NH,OH ^^ NH3 + H^O Ag+ -f 2NH3 -=±[Ag.2NH3] + The depression of the concentration of the silver ions brought about by the combination with ammonia induces the first reaction to swing to the right. For the same reason these halo- gens dissolve in sodium thiosulphate and in potassium cyanide. The first of these reagents forms a soluble sodium salt, which gives in solution the complex ion [Ag(S203)2]-. Potassium cyanide forms the soluble potassium argento-cyanide KAg(CN)2. Silver Nitrate and Carbonate. — Silver nitrate is formed by the action of nitric acid upon the metal. It is freely soluble in water (100 gm. of water at 0° dissolve 115 gm. of silver nitrate, at 100° 910 gm.). The aqueous solution is easily reduced by organic matter to metalhc silver, hence its use in marking inks. SUver nitrate crystaUises in the rhombic system. It absorbs ammonia, forming AgNOajSNHa. This reaction is similar to that given by silver chloride, which produces two such ammonia- compounds, AgCljSNHs and 2AgCI,3NH3, the formation of each compound being conditioned by the pressure of the ammonia. Owing to the strength of silver hydroxide as a base, an aqueous solution of silver nitrate is neutral to htmus. For the same reason, normal silver carbonate is precipitated by the action of a soluble carbonate upon a solution of silver nitrate. Silver carbonate is soluble in an excess of carbon dioxide, forming a L L 514 AN INORGANIC CHEMISTRY bi-carbonate. On beating, tbe carbonate breaks down into silver, oxygen and carbon dioxide. Other salts are silver sulphate, AgaSOi, which is isomorphous with sodium sulphate, but much less soluble ; silver sulphide, AgS, which is precipitated by the action of hydrogen sulphide upon all solutions of silver salts, complex or otherwise, though the precipitation is not complete in the presence of potassium cyanide. Electroplating. In this process the electrolyte is potassium or sodium argento- cyanide. The anode is made of a plate of sUver, the cathode is the object to be electroplated. In such a solution the foLtbwing equiUbria are established : KCN + AgCN ; K^ K[Ag(CN),] CN- i^g++CN- K^ +[Ag(CN)2]- Ag+ 2CN- The actual concentration of silver ions wiU be extremely low as most of the silver is in the form of the complex anion. As has already been shown (p. 465), the current to the cathode is carried mainly by the potassium ions, but at the electrode itself that process will take place which involves the least expenditure of work. This is the separation of the silver ions, not of the potassium ions. As soon as silver ions are thus discharged to metal, fresh silver ions wiU be created by the dissociation of the complex in order to restore the concentration of the Ag"^ ions to their equihbrium value ; in short, the complex ion acts as a huge reservoir for the supply of silver ions. At the anode silver dissolves to form silver cyanide with the discharging cyanide ions, and this silver cyanide is converted into the complex anion. Ag++CN-->AgCN AgCN + CN-->[Ag(CN)2]- The mechanism of the electrolysis of potassium cuprocyanide is similar. Photography. — Modern photography is based upon the tendency of the haloid salts of silver to undergo incipient decom- SILVER 515 position on exposure to light. The precise nature of this change is not yet definitely established, some maintaining that the darkening of the silver halide under the action of Hght arises from the formation of a subsalt : 4AgCl;=±2Ag2Cl + Cl2, others asserting that the hght causes a reduction to metaUic silver, 2AgCl^=±2Ag + Cl,, this silver being precipitated in the molecular state throughout the unreduced hahde. The action of hght is hkely to be inhibited in either case by the presence of chlorine, so that the observation that silver chloride remains perfectly white in an atmosphere of chlorine even though exposed to the action of hght, does not help towards the elucidation of the problem. In order to take advantage of this photo-chemical activity of the silver salts, it is therefore necessary to have present something which will bind or remove the chlorine. Gelatine is such a substance. A photographic plate consists of a thin film of gelatine, con- taining an emulsion of silver bromide, spread uniformly over a glass plate. If such a sensitised plate is exposed to the action of hght, the silver bromide is affected in such a way that the most intense change of the sensitised material occurs where the light has been strongest, i.e. the amount of incipient reduction into subhahde or metal is proportional to the intensity of the light. The plate is then placed in a developer, a reducing agent, amongst such being pyrogallol and potassium ferrous oxalate. The developer first attacks those portions of the film where the hght has already begun the reduction. Finely divided silver is thereby precipitated on the plate in those parts where the action of the hght was strongest, in fact the density of the deposit is proportional to the hght intensity ; where the action of the hght has been greatest, the silver deposit will be densest. The reducing agent is poured off before it begins to attack those portions of the film where hght has not already set up incipient reduction. The plate is now washed and placed in a solution of hypo — sodium thiosulphate — which dissolves out the unattacked silver bromide left in the film. The negative is now fixed and may be safely exposed to the hght. To obtain a print, the negative is placed over a piece of sensitised paper, i.e. paper prepared similarly to a photographic plate. The picture is 516 AN INORGANIC CHEMISTRY printed by exposure to light, and, as before, the action upon the sensitised paper vdll be greatest where the intensity of the light has been strongest. The negative protects the sensitised film from the action of the light in proportion to the thickness of the silver deposit. When the contrast of the picture is sufficient, it is again fixed by placing it in a solution of sodium thiosulphate to remove any unreacted silver bromide present in the film upon the paper. The last stage in the operation is to tone the print. It is immersed in a solution of gold chloride or of potassium- platino-chloride, KaPtClj. The silver of the print is replaced by the more electro-positive (noble) element, gold or platinum, and a richer tone produced. Gold Occurrence. — Gold is generally found in quartz veins and in river-bed gravels. It is frequently found nearly pure, but it also occurs in many copper ores. A gold-silver teUuride (Ag.Au)Te2, is found in Colorado. Metallurgy. — Rich alluvial gold-bearing gravels can be washed in a cradle, but in general, the gold content is far too small for this rough process to yield returns. The chief process is a combined amalgamation- cyanide one. The gold-bearing quartz is crushed in stamping mills and the finely pulverised sludge r»in over amalgamated copper plates. A large part of the gold is retained on the plates, the remainder passes away in the taiUngs, and is subjected to the cyanide treatment, which is so efficient that less than ^ oz. of gold per ton of ore pays for ex- traction. The solution of the gold in the cyanide (less than 0-1 per cent, in strength) depends upon the following reactions : 2Au + 4NaCN -f 2H2O -f O2 -^ 2NaAu(CN)2 -f 2NaOH + HaO^ 2Au + 4NaCN -f H2O2 -> 2NaAu(CN)2 + 2NaOH Oxygen is therefore necessary to effect solution of the gold in the cyanide. The gold is recovered from the cyanide solution either by electrolysis or by precipitation with zinc. In the chlorination process the ore is first roasted, then treated in a moist state with chlorine. Auric chloride, AuCls, is formed and extracted with water. It is precipitated from the solution by means of a suitable reducing agent, e.g. ferrous sulphate, hydrogen sulphide, etc. 2Aua3 +6FeS04^-2Fe2(SOj3 +2FeCl3 +2Au. GOLD 617 The separation of gold and silver from the less noble metals (e.g. lead) is made by cupelling the sample. This is effected by heating the impure metal with lead and borax in a cupel (bone- ash crucible). Oxidation of the baser metals to oxide takes place, and this oxide is either absorbed by the cupel or blown away. A button containing the gold and silver remains. The silver is removed by the action of nitric acid (the process known as quartation). The solution of silver in nitric acid does not proceed readily from a silver-gold aUoy if the gold content exceeds 2-5 per cent. In such cases a suitable amount of pure silver is fused up with the button before separation by means of acid is attempted. Gold is also refined electrolytically. Properties. — Gold is extremely malleable and ductile, gold leaf being made not exceeding 0-000004 inch in thickness. In order to secure greater hardness gold is generally alloyed with copper, as in the gold coinage of the different countries. The melting point of gold is 1062°. Chemically speaking, gold is very unreactive. It is not attacked by oxygen, nor by hydrogen sulphide (cf. copper, silver), nor does it dissolve in hydrochloric, sulphuric or nitric acid. Aqua regia, however, attacks it {see p. 293), as does selenic acid, no doubt owing to the readiness with which this acid undergoes reduction. Other oxidising agents, such as a hot solution of potassium permanganate and sulphuric acid or hypochlorites and sulphuric acid, attack it. Gold is dissolved in small quantities when a gold anode is used in the electrolysis of a strong acid. It is also attacked by fused alkalies and nitrates. General Propeeties of the Compounds op Gold The outstanding feature of this element is the tendency of its salts to pass into complex anions. This is undoubtedly due to the nobility or weak electro -affinity of the element. The weak- ness of its hydroxides, AuOH and Au(0H)3, as bases accounts for the extreme hydrolysis to which salts formed from these bases are subjected. Thus, such salts as auric sulphate can only be kept in solution when a large excess of acid is present. Oxides. — The oxides of gold are aurous oxide, AujO, gold monoxide, AuO, and auric oxide, AU2O3. In accordance with 518 AN INORGANIC CHEMISTRY the general rule, the highest oxide possesses the weakest basic properties, in fact auric oxide and hydroxide are amphoteric. AU three oxides give rise to salts, of which those derived from the aureus and auric oxides are the best known. The element is an interesting example of the general tendency of the bottom member of each B subgroup to give rise to two or more series of salts (cf. mercury, thallium, lead, bismuth). Aureus oxide, AuaO, is a violet powder formed by the decom- position of its salts by means of a soluble hydroxide. A few derivatives of this oxide are known, aurous sulphide, AuS, aurous thiosulphate, AU2S2O3, aurous chloride, bromide and iodide and a few complex salts, such as potassium aurocyanide KAu(CN)2. As a base, aurous hydroxide is extraordinarily weak — a notable exception to the general rule that the element at the bottom of each group forms the strongest base of the group. Gold Monoxide if difficult to obtain pure, but several salts are known, AuCL, AuS, AuBr, AuSOj. AU these salts tend to decompose into a mixture of aurous and auric salts : 2Au++ ^i:±Au+ + ++Au-^ Auric Oxide and Hydroxide. — Auric oxide, AU2O3, is a brown powder formed by adding potassium hydroxide to a solution of auric chloride. The hydroxide can be obtained pure by precipitating auric chloride with magnesium carbonate, the precipitate being freed from magnesium carbonate by washing with dilute nitric acid. Part of the auric hydroxide is left undissolved in the pure state. Auric hydroxide is amphoteric. With strong bases (KOH, etc.) it gives the well-known meta-aurate. KOH + Au(OH)3^^KAu02 +2H2O. Crystals of the compound KAu02,3H20 are obtained on evaporating the solution. Of the salts formed from the oxide AuoOs functioning as a base, the most important is auric chloride. This is formed during the solution of gold in aqua regia, also by the direct action of chlorine upon gold. Its tendency to complex formation is very marked. With hydrochloric acid it forms chlor-auric acid, H(AuCl4),4H20, whUst the alkali chlorides give rise to comple.^ salts such as NaAuCl4,2H20, 2KAuCl4,H20, etc. The solution obtained by dissolving auric chloride in water has aU the properties of GOLD 519 a complex acid of the formula H2(AuCl30). Red crystals of this composition are obtained on evaporating the solution. - On heating auric chloride breaks down into aureus chloride and chlorine. The addition of an excess of potassium cyanide to a solution of auric chloride leads to the formation of potassium auri-cyanide. 4KCN + AUCI3 -^ K[Au(CN)4] + 3KC1. This solution, as well as that of the corresponding aureus com- poimd, is used in electroplating. Auric sulphide can only be formed with difficulty. It has been prepared by the action of hy- drogen sulphide upon dry Uthium auric chloride, LiAuCli. Hydro- gen chloride is set free and the Uthium chloride is washed out with alcohol. Auric sulphide breaks down at once on treatment with water, yielding a mixture of gold and the lower sulphides. If a solution of stannous chloride is added to a solution of sodium chloraurate, NaAuClj, reduction of the chloraurate to metallic gold takes place, while the stannous chloride is oxidised to stannic chloride. 2NaAuCl4-l-3SnCl2->2NaCl -|-2Aui -f-SSnCl^ SnCli +4HOH^>Sn(OH)4 ^ +4HCI. Immediate hydrolysis of the stannic chloride takes place, with the precipitation of stannic hydroxide. As this substance separates out, it takes with it the gold, which had remained in a colloidal state. The colour of the precipitate ranges from red to violet, according to the conditions of precipitation. This forms the so-called Purple of Gassius, which is used for gilding porcelain. Questions 1. Give an account of the chemical and physical phenomena involved in sUver-eleotroplating. 2. Starting with an alloy of silver and copper, how would you prepare a sample of pure silver nitrate ? What is the action of the following substances upon a solution of silver nitrate : (a) arsine, (6) ammonium hydroxide, (c) sodivmi hydroxide, (d) zinc, (e) dilute hydrochloric acid, (/) concentrated hydrochloric acid ? 3. Give an account of the metallurgy of copper. 4. Discuss what happens when a concentrated solution of potassium chloride is added to a deep-green solution of copper chloride. 5. Discuss what happens when ammonium hydroxide is slowly added to a solution of copper sulphate. 6. Compare the oxides and compounds of copper, silver and gold. 7. Show how the oxides of gold become less basic with the increasing oxygen content of the oxide. Use this fact to compare the more important compounds derived from the oxides of gold. 520 AN INORGANIC CHEMISTRY 8. Silver chloride dissolves in solutions of ammonium hydroxide, potas- sium cyanide, sodium chloride (if strong), and sodium thiosulphate. Briefly account for these facts. 9. Give an account of the chemistry of photography. 10. Give a general method for reducing the cupric compounds to the cuprous state and apply your method to the preparation of cuprous chloride. 11. Discuss the ionic reactions involved when solutions of potassium cyanide and copper chloride are mixed, the former in considerable excess. 12. What is the action of heat upon (a) silver oxide, (6) cuprous oxide, (c) silver nitrate, (d) cupric nitrate, (e) cupric chloride, (/) auric chloride ? 1 3. What examples of double salt and complex salt formation does the study of the compounds of copper reveal ? 14. Compare the elements of Group 1a and 1b, and indicate the justification (or otherwise) for grouping these sub-groups together. CHAPTER XXXIII GROUP 2A: GLUCINUM, MAGNESIUM, CALCIUM, STRONTIUM, BARIUM These elements, of which the last three are often referred to as the alkaHne earth elements, form a family, the properties of which show the same general gradation as has been noted for the alkaU metals. Group 1a. With increasing atomic weight there is a steady increase in the basicity of the hydroxide. Glucinum hydroxide is amphoteric, barium hydroxide is a strong base. The strength of the hydroxides of calcium, strontium and barium is testified to by the strong electrolytic dissociation which these bases undergo when in aqueous solution. Again, the hydroxides become more stable towards heat as one passes from glucinum to barium ; calcium hydroxide is easily dehy- drated, but barium hydroxide does not part with water at a red heat, it merely melts. The solubihty of the hydroxides increases steadily with the atomic weight of the elements of this group. The solubility of the sulphates is, however, in the inverse order, barium sidphate being the least soluble, glucinum sulphate the most soluble (100 grams of water at 18° dissolve approximately 100 grams by weight of glucinum sulphate, but only 0-00023 gm. of barium sulphate). The elements themselves display a similarity to the elements of the alkali group, except that they are somewhat less reactive. This is particularly so with glucinum. Glucinum does not react with water even on boiling ; hot water is slowly decomposed by magnesium, rapidly by calcium, strontium and barium, but in no case is the energy of the reaction sufficient to cause the escaping gas, hydrogen, to inflame. The carbonates are all insoluble in water, but with the possible exception of glucinum, soluble bicarbonates are formed in the presence of an excess of carbon dioxide. The 521 522 AN INORGANIC CHEMISTRY carbonates show an increasing reluctance to part with carbon dioxide as the atomic weight of the metal increases. The action of heat upon the nitrates of the elements of this group is also worthy of note. In all cases the nitrate breaks down into the oxide, giving off oxygen and nitrogen tetroxide. This reaction distinguishes them from the alkahes with which they otherwise have much in common, and brings them more into hne with the metals of the heavy metals : cf. 2Pb(N03)2^2PbO + 2N204 + 02. Glucinum The position of glucinum, also known as beryUium, in the Periodic Table, intermediate between the strongly metaUic element, hthium, and the metalloid, boron, accounts for the weak nature of glucinum hydroxide as a base. The salts derived from glucinum oxide and hydroxide are those given by a typical divalent metal, e.g. GICI2, G1(N03)2, but the extreme weakness of this basic oxide accounts for the ease with which these salts are hydrolysed. The warming of the nitrate to 100° forms the basic salt G1(N03)2,G1(0H)2,2H20. Moreover, the amphoteric nature of the hydroxide leads to the existence of another type of salt, the glucinates ; G1(0H)2 + 2NaOH->Na2G102 + 2H2O. This amphoteric behaviour, coupled with the ease with which the carbonate breaks down into the oxide, led chemists to assign a false position to this element in the Periodic Table. It was at first held to be a member of the third group, aUied to alumin- ium, but subsequent work soon allotted to it its correct position in the vacant place existing at the head of Group 2. The methods of preparation of the metal and its salts resemble strongly those described for its analogue, magnesium (q.v.). The pro- perties of the salts of glucinum differ but little from the pro- perties of the corresponding salts of magnesium, except in so far as the weaker basic nature of glucinum hydroxide causes a modification in the specific properties of the salt. Magnesium Occurrence. — Magnesium is widely distributed in nature as carbonate (magnesite, MgCOg, dolomite, MgC03,CaC03), MAGNESIUM 523 chloride (bischoffite, MgCl^.eHaO, carnaUite, MgCla.KCl.eHaO), sulphate (kieserite, MgSOi.HaO, kainite, MgCl.KaSOj.SHjO, schonite, KaSOi.MgSOi.eHjO), as well as in many silicates, such as enstatite, asbestos, talc, ohvine (MgaSiOi). Isolation of the Metal. — ^Magnesium is prepared by the electrolysis of anhydrous fused carnaUite. The electrolysis is effected in an iron pot which serves as a cathode, the anode being of carbon. The older chemical process, based upon the reduction of the chloride with sodium, is being steadily displaced. MgCl 2 + 2Na --> 2NaCl + Mg. The element is often pressed into wire or ribbon, while hot and put on the market in that form. Properties. — Magnesium is a briUiant white metal when freed from the superficial layer of oxide with which it is generally covered. It remains bright in an atmosphere of dry oxygen, but tarnishes rapidly in the presence of moisture. It attacks hot water, Uberating hydrogen. Dilute acids dissolve it freely. Magnesium burns freely in carbon dioxide, reducing it to carbon. Metallic magnesium combines with nitrogen on heating, forming the nitride : 3Mg+N,->Mg3N2, hence its use in separating nitrogen from argon. As an element, it must be classed among the more reactive ones. It forms a carbide, a silicide, Mg2Si,MgSi, boride, MgsBa, sulphide, MgS and hydrosulphide, Mg(HS)o, selenide, MgSe and possibly a hydride. The vigour with which it combines with oxygen has led to its use in flashlight photography, as well as in fireworks. Oxide and Hydroxide. — Magnesium oxide is prepared by the decomposition of the carbonate, or hydroxide or by the hydrolysis of magnesium chloride, brought about by the actior of steam upon the fused chloride : MgCla + H20->MgO + 2HC1. Again, the mother liquor from the Stassfurt deposits, rich in magnesium chloride, is treated with milk of Ume, and the resultant hydroxide decomposed. Owing to its infusibility, magnesium oxide is used extensively 524 AN INORGANIC CHEMISTRY for crucibles and for lining furnaces, generally in the form of bricks made from magnesite or dolomite. Magnesium hydroxide is feebly alkaline in reaction and only slightly soluble. It can be prepared by adding a solution of a strong base to a dissolved salt of magnesium. If ammonium hydroxide is added to a solution of magnesium chloride, a precipitate of magnesium hydroxide separates : MgCl2 + 2NH40H ^=:± Mg(OH)2+2NH4Cl. This precipitate redissolves if more ammonium chloride is added, from which result one may rightly infer that before the addition of this ammonium chloride precipitation of the magnesium hydroxide was only partial, i.e. an equUibrium was set up. In order to throw magnesium hydroxide out of solution, it is necessary that the solubility product of this substance should be exceeded, i.e. (Mg+ + )(OH-)2>Ljjg(ojj,^. The addition of ammonium hydroxide, weakly ionised though this base is, yet introduces sufficient hydroxyl ions to enable the solubihty product of the magnesium hydroxide to be exceeded, and the precipitate separates out. How, then, does the ammonium chloride drive the precipitated magnesium hydroxide back into solution ? The answer to this question is bound up in the equations : NH40H^z:ziNH4+ +0H- NH4C1;=±NH4+ +01". The dissociation of the ammonium hydroxide is very small (0-4 per cent, for a N. solution), whilst the salt, ammonium chloride, is strongly dissociated (N. solution is 75 per cent, dis- sociated). The high concentration of ammonium ions, brought into the solution by the introduction of the highly ionised ammonium chloride, consequently throws back the dissociation of ammonium hydroxide, i.e. the equUibrium : NH4OH ;zi±NH4+ + OH- is driven over to the left, and such a reduction is thereby brought about in the concentration of the hydroxyl ions that the solu- bility product (Mg+ + )(0H~)2 no longer exceeds Ljjg,ou)2 and the precipitate passes back into solution. The same conclusion can be arrived at by considering the action of ammonium chloride upon a saturated solution of magnesium hydroxide : MAGNESIUM 525 Mg(OH)2;=^Mg(OH)2;^Mg++ +20H-^ Solid. Dissolved. l==±2NHiOH 2NH,Cl^i:^2Cl-+2NH,+ J When ammonium ions are created in the solution by the dis- sociation of the added ammonium chloride, the equilibrium NHiOH ^z± NH4+ + OH- must be established between the ammonium ions from the ammonium chloride and the hydroxyl ions from the dissolved magnesium hydroxide ; and owing to the weakness of the ammonium hydroxide as a base, this equilibrium is not attained untU a marked reduction in the concentration of the hydroxyl ions is effected. In order to restore the equilibria defined in the upper equations, more of the undissociated magnesium hydroxide must dissociate and more of the sohd therefore passes into solution. The greater the amount of ammonium chloride added, the more magnesium hydroxide will be taken into solution. Magnesium Chloride. — Although this salt occurs naturally it is generally obtained from the mother liquor of camaUite, after the potassium chloride has crystaUised out. If carnalUte (MgCl2,KCl,6H20) is treated with a small quantity of water at 25°, decomposition of the double salt occurs, and about 85 per cent, of the potassium chloride is thrown out of solution. Magnesium chloride is worked up from this mother liquor. On heating hydrated magnesium chloride loses appreciable quantities of hydrogen chloride owing to hydrolysis. The usual plan of carrying out the dehydration is in an atmosphere of hydrogen chloride. Many attempts have been made to utihse the waste liquors of the Stassfurt deposits as a source of hydrochloric acid and chlorine. The preparation of hydrochloric acid is based upon the strong hydrolysis of the chloride on heating. 2MgCl2 + H^O^MgCMgCl^ + 2HC1. The evolution of chlorine is brought about by acting upon the heated chloride with air and steam : 4MgCl2 + 2H2O + 0^-^ 4MgO + 4HC1 + 2Cl,. Magnesium Sulphate. — ^The common form of this salt is the heptahydrate (Epsom Salts), but as in the case of the chloride, quite a number of hydrates can be obtained. The sulphate in 526 AN INORGANIC CHEMISTRY combination with 1, 4, 5, 6, 7, 12 molecules of water of crystal- lisation has been reported. Magnesium Carbonate. — The normal carbonate occurs naturally, but can also be obtained in various hydrated forms by crystallising a solution of the carbonate dissolved in an aqueous solution of carbon dioxide. The carbonates precipitated by the action of an alkali carbonate are basic. In this case the precipitate seems to be a mixture of the normal carbonate with the hydroxide, thrown down through the hydrolysis of the alkali carbonate : Na^COs + HOH =± NaHCOa + NaOH 2Na+ +CO3- H+ +0H- -> Na+ +HGO3- Na+ OH- The solubility product of magnesium hydroxide is so low that the concentration of the hydroxyl ions present in an aqueous solution of sodium carbonate is sufficient to cause the separation of appreciable quantities of magnesium hydroxide along with magnesium carbonate. Magnesium Sulphide. — The sulphides of magnesium exhibit no conspicuous difference from those of calcium (q.v.). Magnesium Phosphates. — The insolubility of magnesium ammonium phosphate, MgNH^POj.GHaO is made use of in esti- mating magnesium. It is thrown down on treating a solution of a magnesium salt with a mixture of ammonium hydroxide and sodium phosphate. On ignition the phosphate decomposes into the pyrophosphate, in which form it is weighed. 2MgNH,P04,6H,0 -> Mg^P^O, + 7H,0 + 2NH3 Calcium, Steontium and Barium Occurrence. — These elements aU occur in the form of carbonate, e.g. limestone, CaCOj, strontianite SrCOg, witherite BaCOa. The sulphates are also widely distributed ; gypsum CaS04,2H20, celestine SrSOi, barytes or heavy spar, BaSOi. Calcium is also found in the form of fluorspar CaFj, and phosphate [phosphorite Ca3(POi)2, apatite 3Ca3(P04)2,CaF2]. Preparation of the Metals. — These metals were first isolated by Davy by a method analogous to that used for the CALCIUM, STRONTIUM, BARIUM 527 isolation of the alkali metals. A piece of the hydroxide was placed on a strip of platinum which served as anode, the cathode being a drop of mercury in a depression on the surface of the hydroxide. After the electrolysis the amalgam was distilled in a hydrogen atmosphere and the alkali earth metal remained. In the modern method the difficultly fusible hydroxide is replaced by the more easily fusible chloride. Calcium is the only element of the three which is prepared on the commercial scale, and even that has a very limited application. In the electrolysis of calcium chloride the anode consists of an iron or carbon pole, the cathode is an iron rod which just dips below the surface of the electrolyte, which is maintained at a temperature of 780-800°. As the electrolysis proceeds, the iron cathode is slowly raised, the adhering calcium then serving as cathode. The metal is obtained in a massive, coherent form, free from any impurity except an adhering film of calcium chloride. Properties of the Metals. — The metals are silvery white crystalUne substances which react very vigorously with water as well as with dilute acids. They reduce fuming sulphuric acid to sulphur, concentrated sulphuric acid to hydrogen sulphide. When heated they combine with hydrogen, oxygen, nitrogen, the halogens, sulphur and carbon, while they reduce carbon dioxide, forming the oxide of the metal and also the carbide. Hydride and Nitride. — The hydrides and nitrides of cal- cium, strontium and barium are obtained by the direct com- bination of the elements, or by passing the requisite gas over an amalgam of the metal (general formula of these compounds MHj, M3N2). The temperature required to induce the formation of the hydride increases rapidly in passing from calcium to barium. Calcium hydride, CaHa, is formed at a duU red heat, strontium hydride at a bright red heat, barium hydride only at a temperature above 1,200°. The hydrides decompose on treatment with water or dilute £Lcids C ff. '• CaH^ + 2H20->Ca(OH)2 + 2H2. Calcium hydride is sometimes used commercially for the pre- paration of hydrogen. The nitride of calcium, etc., as do all other nitrides, breaks down on treatment with water, and yields a hydroxide and ammonia : CasNa + 6H20-^3Ca(OH)2 + 2NH3. 528 AN INORGANIC CHEMISTRY If it Mere only possible to make metaUic calcium cheaply, calcium nitride would form an excellent means of making the huge stores of atmospheric nitrogen available for plant Ufe. Carbonates. — Of the three carbonates that of calcium stands out in importance, not only from its natural importance, but owing to its commercial appUcations. It is extensively distri- buted as limestone, an irregularly crystalUne mass, as the definitely crystalline marble, as chalk, a deposit coiLsisting of the calcareous skeletons of shell fish, coral insects, etc., and as the definitely crystaUine calcite and aragonite. Calcite be- longs to the hexagonal system, occurring in a variety of forms. When transparent and colourless, it exhibits double refraction, and is kno-^vn as Iceland Spar. In this form it is used a great deal for polarising Ught. Another crystalUne form is aragonite, which belongs to the rhombic system. This modification, although occurring freely in nature, is really in a state of insta- bility at ordinary temperatures, tending to revert to the stable form, calcite. The transition temperature aragonite ^ ^ calcite Ues somewhere about 50°. If calcium carbonate is pre- cipitated at temperatures above this, aragonite separates out, at lower temperatures calcite appears. The existence of ara- gonite in nature is an example of the extreme reluctance displayed by unstable soHds to revert to the more stable form (cf . red and yeUow phosphorus). Oxides and Hydroxides. The carbonates of calcium, strontium and barium all dissociate on heating into the oxide and carbon dioxide MCOs^^MO + COj, the extent of the dissociation increasing with the temperature. The process is very similar to that of evaporation ; at every temperature the carbonate exerts a definite dissociation pressure or partial pressure of carbon dioxide (see Chapter XIV, p. 212). The following table exempUfies this in the case of calcium carbonate : TABLE 52. Temp. . 547° 610° 625° 740° 810° 812° 865° P. . 27 46 56 255 678 753 1,333 If the partial pressure of carbon dioxide at any temperature exceeds the dissociation pressure for that particular temperature, combination wiU occur in the sense of the equation CaO+C02;=±CaCO„ CALCIUM, STRONTIUM, BARIUM 529 until the pressure of the carbon dioxide has fallen to the equili- brium value for that temperature. At 812° the pressure is almost equal to that of an atmosphere. If the calcium carbonate is heated in open vessels, the carbon dioxide generated by the dissociation wiU escape, the dissociation pressure wiU not be reached, and the carbonate will continue to dissociate. In order to obtain the oxide from the carbonate most easily, it is therefore necessary to carry out the dissociation under such conditions that the carbon dioxide will not be able to reach the equiUbrium pressure for that temperature. This is best done by heating the carbonate in a suitable furnace or Idln through which a steady draught finds its way. As is only to be expected from the stronger basic nature of the oxides of strontium and barium, the carbonates of these elements are much less dis- sociated under similar conditions of temperature than is the carbonate of calcium. The oxides of strontium and of barium are therefore obtained by special means. The kilns in which calcium carbonate is broken down into lime are charged with limestone and a fire started. The resulting rise in temperature, coupled with the draught of air which removes the carbon dioxide from the sphere of action, effects the decomposition into the oxide. Another method which is largely used outside America to reduce the concentration of carbon dioxide and so facilitate the dissociation is to mix the limestone with coal in alternate layers. The action of the coal is twofold ; by its combustion the necessary heat is generated, while the excess of coal present reduces the carbon dioxide to carbon monoxide, thereby lowering its partial pressure and aiding the dissociation. The oxides of strontium and barium can be made on the commercial scale from their carbonates by taking advantage of this method of helping the dissociation. The carbonates are mixed with powdered coal, and fired in specially constructed kilns. Strontium and barium oxides can also be obtained by the reduction of the sulphate with coal and subsequent interaction with steam at a red heat. BaSOi + 4C -^ BaS + 4C0 BaS+HjO-^BaO+H^S The oxides of the alkali earth elements are aU white, amorphous MM 530 AN INORGANIC CHEMISTRY powders which, melt only in the intense heat of the electric oven. When the oxide is treated with water, reaction takes place, the violence of the reaction increasing with the atomic weight of the, metal. If the addition of the water is made carefuUy, the dry hydroxide is obtained. The addition of water to calcium oxide (quicklime) to form slaked hme is exemplified in the equation CaO+H20->Ca(OH)2. Calcium hydroxide (slaked lime) is easily dehydrated by heat, its dissociation pressure at 350° amoimting to 100 mm. of mercury, but at this temperature the tension of the aqueous vapour above barium hydroxide is negUgible. The hydroxides of the alkaline earth elements are strong bases ; this, combined with its cheapness, accounts for the extensive use of calcium hydroxide in industry (see preparation of potassium chlorate, etc.). A suspension of calcium hydroxide in water is often used industrially, and is known as milk of lime. The solubUity of calcium hydroxide decreases with the tempera- ture. (At 0° 0-131 gm. Ca(0H)2 dissolves per 100 gm. water, at 100° 0-060 gm.) The solubihty of barium and strontium hydroxides increases with rising temperature (at 0° 1-5 gm. of BaO dissolve per 100 gm., at 80° 90-77 gm.). If carbon dioxide is passed through a solution in which solid calcium carbonate is suspended, considerable quantities of the carbonate pass into solution. It has been shown that practically the whole of the calcium exists in the form of the soluble bicar- bonate. A solution, saturated with carbon dioxide at atmo- spheric pressure, contains 1-3 gms. at 13-2°. Evaporation of such a solution does not yield the soUd bicarbonate, owing to the reaction : Ca(HC03),,->CaC0 4- 3+C02^ +H2O. Pure water at 16° dissolves only 0-013 gm. of calcium carbonate. It is reported that calcium bicarbonate heis been isolated by the action of a solution of potassium bicarbonate on a solution of calcium chloride at 0°. The equUibrium defined in the above equation plays a great part in the sculpture of the earth's crust. Rain water, charged with carbon dioxide, removes in solution large quantities of calcium bicarbonate from beds of limestone. In this way large underground caves are often formed in such beds. Very often, CALCIUM, STRONTIUM, BARIUM 531 too, as the water with its dissolved salt drops from the roof of the cavern, the equilibrium expressed in the equation CaCO 3 + CO a + H 2O ^=± Ca(HCO 3) a is disturbed by the escape of some of the carbon dioxide. This is due to pressure and temperature changes. As soon as the carbon dioxide escapes a deposit of calcium carbonate is thrown down. In this way stalactites slowly grow from the roof, and, where the drippings strike the floor, a slow growth of stalagmites takes place. Hard Water. — The geological importance of the soluble calcium bicarbonate, formed by the action of water charged with carbon dioxide upon beds of limestone, is enormous. With the erosive action of such water upon beds of limestone we are not here concerned, but the presence of the soluble calcium bicarbonate in natiiral waters has a great influence upon the utility of such water. Water contaiaing dissolved salts of magnesium and of calcium is called hard water, the term being applied owing to the difficidty of securing a lather with soap. Soap is a sodium or potassium salt of an organic fatty acid. When soap is treated with water containing dissolved salts of calcium or magnesium, interaction takes place with the precipitation of the insoluble calcium or magnesium salt of the fatty acid. The insoluble precipitate is the well-known scum given by hard waters when they are used for detergent purposes. The formation of the scum continues until the whole of the dissolved magnesium and calcium salts has been precipitated ; then and only then does the soap begin to lather. In analysis the hardness of water is expressed in degrees of hardness, which may be defined as the number of grains of calcium carbonate (or its equivalent in other salts) present per gallon of water. Hardness of water is often referred to as of a temporary or a permanent nature. Temporary hardness is readily removed by one of two methods : 1. The water is heated, when the reaction Ca(HC03)2-^CaC03i+H20+C02| sets in. The carbon dioxide which is responsible for holding the calcium carbonate in solution and thereby rendering the water hard, is expelled by the boiling. 532 AN INORGANIC CHEMISTRY 2. The water is treated with a suitable base, e.g. slaked lime, to remove the excess of carbonic acid which is rendering the calcium carbonate soluble : Ca(HCO 3) , + Ca(OH) , ^ 2CaC0 ,i +211,0. The precipitated carbonate is allowed to settle and the clear Liquid decanted ofE. Permanent hardness, which arises from the presence of the dissolved sulphates and chlorides of calcium and magnesium, is more difficult to destroy. It can be effected most readily by the addition of sodium carbonate : CaSOj +Na2C03^CaC03>l' +Na2S0,. Temporary hardness is also removed by the addition of sodium carbonate. The fanuliar use of washing soda (NaaCOa.lOHaO) is based upon these reactions. One of the most important of the modern methods of softening water is based upon the use of permutite, an artificial zeoUte or complex sodium silicate. If hard water is allowed to flow over coarse lumps of permutite, the whole of the calcium and mag- nesium is removed from the water in the form of an insoluble permutite. The sodium passes into solution and the exchange is complete. On the other hand, when the permutite ceases to be effective, regeneration can be brought about by passing a weak solution of common salt through the calcium permutite whereby the sodium compound is reformed. Another interesting process is the use of corrugated aluminium plates. Apparently colloidal aluminium is carried into solution, and during its coagulation carries down the salts of calcium and magnesium. The necessity of getting rid of the hardness of water before use in a boiler arises from the fact that the precipitation of the insoluble salts of calcium and magnesium leads to the formation of the boiler scale. The appearance of the scale within a boiler causes a marked decrease in the efficiency, as the boUer scale is a bad conductor of heat. This scale is also apt to crack away from the iron surface, especially if the water in the boiler gets low. The introduction of water at this stage may well lead to the bursting of the boiler. The presence of hard water in boilers is also responsible for corrosion, especially if the chlorides and nitrates of calcium and magnesium are present. This is due to the acid set free by the hydrolysis of these salts. CALCIUM, STRONTIUM, BARIUM 533 Mortar.— Mortar is made by mixing slaked lime and sand to a thick paste with water. The setting of the mortar is due to the reaction : Ca(0H)2 + C02->CaC03 + H^O. The crystals of calcium carbonate form from the surface and, when interlaced with the sand, produce a rigid mass. The sand acts purely in a physical way, firstly in making the mortar more porous to the carbon dioxide, and secondly, in preventing undue shrinking when the carbonate forms. But for the presence of the sand cracks would develop as the setting progresses. Cement. — Cement is made by heating together a mixture of limestone, clay and sand until the mixture nearly sinters. The whole is then ground to a fine powder. When the cement is mixed with water, it sets to a hard stone-Hke mass. This setting appears to be a somewhat complex chemical reaction ; probably the main reaction is the formation of calcium alumin- ium silicates. Carbon dioxide plays no part in the setting of a cement, in fact the hardening of a cement proceeds just as smoothly under water as in the air. Peroxides. — Barium peroxide is the best known of these compounds. It is prepared by the direct oxidation of barium oxide at a high temperature, though, as the oxidation proceeds with the evolution of heat, the equilibrium : 2BaO+02;=^2Ba02 will be shifted to the left with rising temperature. The dissociation of barium peroxide into barium oxide and oxygen is fundamentally the same as the dissociation of calcium carbonate. In both cases the pressure of the evolved gas is constant for any chosen temperature. 2Ba02— ^2BaO+02 CaCOs — =±CaO+C02 The dissociation pressures for barium peroxide are given below : TABLE 5.3 Temp. p. in ram. mercury 525° 20 660° . 65 720° . 210 750° . 340 890° . 670 S34 AN INORGANIC CHEMISTRY Barium peroxide is prepared technically by heating barium carbonate and coal in an oven at 1,200°. The barium oxide so formed is cooled in a stream of air to the neighbourhood of 500° — a temperature at which the dissociation pressure of the per- oxide is very small, hence the monoxide is almost completely oxidised to the peroxide. Other technical methods are the breaking down of barium nitrate, and the heating of barium carbonate with barium carbide. It is essential that the barium oxide first formed: should be in a fine state of division. The ease mth which the barium oxide absorbs oxygen at 600° and the marked influence of temperature and pressure in promoting the dissociation of the peroxide has been responsible for the use of this substance as a means of obtaining oxygen from the air (see Erin's process for manufacturing oxygen, p. 50). If barium peroxide is treated with water, a crystaUine hydrate, Ba02,8H20, is formed. This hydrate can also be prepared by the action of hydrogen peroxide upon barium hydroxide. Hydrated strontium and calcium peroxides can be prepared by a similar method. Cautious heating of the hydrated peroxide of these elements at 100-120° yields the peroxide itseK, CaOa and SrO,. Neither of these peroxides can be prepared by the direct action of oxygen upon the oxide. Both calcium and strontium peroxide break down at high temperatures, yielding oxygen and the ordinary oxide. Halides. — Calcium chloride occurs as a by-product in many important operations, e.g. the manufacture of ammonia from ammonium chloride, the chemical method of preparing potassium chlorate, so that the direct preparation from marble (CaCOs) and hydrochloric acid is rarely resorted to. The interesting feature about this salt is the number of hydrates which it yields. The solubflity curve for calcium chloride illustrates this (Fig. 130). In this graph every point of intersection indicates the condi- tions of temperature and concentration where two solids exist side by side. Each hydrate has perfectly definite limits of temperature and concentration within which it and it alone can exist in a stable state. Moreover, each hydrate has its own definite vapour pressure which is fixed for each temperature. The conditions imder which the various hydrates are formed CALCIUM, STRONTIUM, BARIUM 535 at some constant temperature by an increase in the vapour pressure, are exactly as already discussed for the hydrates of copper sulphate (p. 98). If the hydrates of calcium chloride are heated, hydrolysis -50° 0° 50' WO" 150° / / erripei ati jre J / ^c^ / ■^ ^ R^ s ^ TJ k 100 ^ fe i ^s 1 Ve 150 \ s N \ \ p 200 ^^ > ^? \ \ 9^/1 \ 1 \ \ t« >c^ h Fig. 130. leads to the escape of appreciable quantities of hydrogen chloride, and the fused mass contains considerable quantities of the oxide. In order to avoid this as far as possible, the usual device is resorted to of carrying out the dehydration in an atmosphere of hydrogen chloride. This is generally achieved by adding 536 AN INORGANIC CHEMISTRY concentrated hydrochloric acid before the highly concentrated solution is heated strongly. The evolved hydrogen chloride possibly also aids in imparting a desirable porosity to the sub- stance. The avidity with which combination ensues between water and calcium chloride, coupled with the very low pressure of the hydrates formed from this substance, is responsible for the extensive use of this reagent for the absorption of aqueous vapour. Absolute dryness is, however, not obtained, for every hydrate has its own fixed vapour pressure for every temperature, and it is obviously impossible for the partial pressure of the aqueous vapour to be lowered below that of the hydrates which are present in the drymg tube. Thus, at 36-5° CaCL,2H20 in contact with CaCl2,H20, has a vapour pressure of 4 mm. A more efficient drying agent is concentrated sulphuric acid. This is because the lowest hydrate of sulphuric acid has a smaller vapour pressure than the lowest hydrate of calcium chloride. The vapour pressure of meta-phosphoric acid (HPO3) is immeasurably small so that phosphorus pentoxide forms Om; best drying agent. Strontium chloride is prepared from the carbonate (strontianite) by the action of hydrochloric acid or from the sulphate by the method described for barium chloride. It can be obtained in the anhydrous form, also as a di- and hexa- hydrate. The latter hydrate is isomorphous with the hexahydrate of calcium chloride. Barium chloride is prepared commercially by melting together barytes, mixed with coal, and calcium chloride : BaSOi + CaCla + 4C^ BaCl^ + CaS + 4C0. The mixture is kept molten tUl the blue flame of the carbon monoxide ceases to be seen. The cooled mass is rapidly extrac- ted with water in order to extract the soluble barium chloride. Another method is to pass hydrogen chloride over a glowing mass of barytes and coal. Barium chloride differs both from strontium chloride and from calcium chloride in not being at all hygroscopic. It forms two hydrates, the mono- and the di-hydrate, both of which can be completely dehydrated without the loss of hydrogen chloride (cf. calcium chloride). This is due to the stronger basicity of barium hydroxide as compared with calcium hydroxide. Both calcium and strontium chlorides form definite compounds CALCIUM, STRONTIUM, BARIUM 537 with ammonia, MCl2,a;NH3 ; no evidence concerning a similar compound for barium chloride is known. Bromides, Iodides, Fluorides. — The bromides and iodides of calcium, barium and strontium call for little comment. Like the chlorides, they are all freely soluble in water and also give rise to a number of hydrates. Of the fluorides that of calcium occurs naturally as fluor- spar or fluorite in the form of cubes and octahedra. This is the most widely spread of all minerals containing fluorine, and therefore finds extensive use in the manufacture of hydro- fluoric acid {q.vJ). Natural specimens are often beautifully coloured, others show a well marked fluorescence on heating. Calcium fluoride is difficultly soluble in water and may therefore be prepared by double decomposition. It is used in metallurgy as a flux. Barium and strontium fluorides are also somewhat insoluble, though less so than the calcium salt. Sulphates. — All the alkahne earth elements are found in nature as sulphates, but calcium sulphate is the most widely spread. This salt occurs in large masses as anhydrite, CaSOi (rhombic), and also as gypsum CaS04,2H20. Clean fine- grained masses of CaSOijSHaO are known as alabaster. Besides these modifications there is no doubt that a more soluble form of anhydrite can be prepared by artificial means, whilst the hemihydrate, 2CaS04,H20, is also known. The solubility method for investigating the temperature at which these salts pass one into the other is of httle use, as their solubility is too small. Other methods have shown that at 107° there is a sharp, well-defined transition point above which the hemi-hydrate is stable, and below which the dihydrate is the stable phase : 107° CaS04,2H20 ^==: 2CaS04,H20. The soluble anhydrite is obtained by warming precipitated calcium sulphate in vacuo at 80-90° in the presence of sulphuric acid. It is more soluble than the natural anhydrite, and has the property of combining with water with extreme rapidity and vigour. Plaster of Paris is made by heating gypsum to about 120° until nearly all the water of crystaUisation has been expelled. The resultant mass consists of a mixture of the anhydroua 538 AN INORGANIC CHEMISTRY sulphate, the hemi- and the di- hydrate. On forming a plastic mass of plaster of Paris with water, rapid combination ensues and the change into the dihydrate is accompanied by a sufficient expansion to enable a sharp cast to be made of any mould in which the plaster sets. The explanation of the setting is to be sought in the fact that the three salts present have difierent solubilities. The two phases which are imstable in the presence of water, viz. the anhydrite and the hemihydrate, are more soluble than the stable phase, CaSOi,2H20. The solution which is saturated with respect to the unstable phases is therefore supersaturated with regard to the dihydrate, and crystaUisation sets in, i.e. the crystals of dihydrate grow until a mass of interlocking crystals is produced. The greater solubility of the unstable phase is really another expression of the law that the unstable phase possesses a greater vapour pressure. If plaster of Paris is heated strongly, complete dehydration occurs ; such " dead burnt " plaster sets very slowly when treated with water, probably owing to the absence of nuclei of the dihydrate. Plaster of Paris is frequently used for obtaining casts of ornaments, statues, etc., as well as for giving rigidity in various surgical operations. Calcium sulphate, although soluble in water with difficult j', dissolves more freely in the presence of sulphuric acid and such salts as ammonium sulphate. Advantage is taken of its solu- bDity in the latter reagent to effect its separation from strontium sulphate. As is usual in such cases, the increased solubility is to be attributed to the formation of a soluble double salt. Barium sulphate is the least soluble of the sulphates of the alkaline earth elements, though its solubUity in the presence of strong acids, especially sulphuric, is considerable. It is highly probable that the increased solubihty of the sulphates of barium, strontium and calcium in sulphuric acid is due to the formation of an acid sulphate, e.g. Ba(HS04)2. An iateresting point has been investigated in connection with the solubility of barium sulphate — ^the greater solubility of very small crystals as compared with the solubUity of more massive crystals. It has been shown that, it the size of the crystals is less than 0-0001 mm. in diameter, an increased solubility is to be expected. This is a surface tension effect. If crystals of CALCIUM, STRONTIUM, BARIUM 539 varying size, some of them less than 0-0001 mm., are present in a solution, it has been noticed that the crystals of a diameter less than the critical size (0-0001 mm.) disappear. This arises from the fact that the solution which is saturated with respect to the smaller crystals will be supersaturated with regard to the larger crystals. The latter wiU therefore grow at the expense of the former, until the smaller crystals have disappeared. A somewhat similar phenomenon is observed if a little sulphur is heated in a tube. The vapour condenses on the cool portions of the walls in the form of Uquid, although the temperature is much below the melting point. This is an illustration of the separation of the unstable phase first (cf. Law of Successive Reactions, p. 310). In a few hours it wiU be seen that some of the drops have changed into crystals, and wherever a crystal has formed, it is seen to be surrounded by a halo. This is due to the stable phase (crystal) possessing a lower vapour pressure than the unstable liquid drops. As a result of this there is distUlation from the place of greater pressure to that of lower. Occasionally the halo is seen without a central crystal, but in all such cases it wiU be observed that the central drop in the midst of the halo is a large one. From this we may conclude that large drops, i.e. drops above a certain Umitiag size, possess a fixed vapour pressure, but drops of smaller dimensions than this Umiting size possess a larger vapour pressure, hence the distillation. Chlorates. — Calcium chlorate is formed during the chemical preparation of potassium chlorate {q.v.), but owing to its hygro- scopic and deUquescent ' nature, it is rarely isolated. The corresponding barium salt is not hygroscopic and can be kept in the dry, anhydrous condition. It is used in fireworks. " Greenfire " is made by the combustion of an intimate mixture of barium chlorate, sulphur and carbon. Barium chlorate can be prepared by the action of chlorine upon a solution of barium hydroxide, by neutralisiag a solution of barium hydroxide with chloric acid and by allowing ammonium chlorate to react with barium carbonate in alcoholic solution, barium chlorate being insoluble in this solvent. Sulphides. — Barium sulphide is of considerable technical importance. The main source of barium is the naturally occurring sulphate, but owing to the insolubility of this mineral, 540 AN INORGANIC CHE:\nSTRY it has first to be reduced to sulphide before other salts can be prepared from it. The reduction is effected by heating together barium sulphate and coal : BaSOi +4C->BaS +4C0. The sulphide can easily be converted into oxide or chloride {q.v.). Barium sulphide is soluble in water, but at the same time undergoes considerable hydrolji:ic decomposition : 2BaS + 2H2O -^ Ba(HS) ., + Ba(OH) 3. It is therefore less soluble in an alkaline solution than in pure water. The sulphides of barium and to a less extent of calcium and of strontium possess the property of phosphorescence, maintaining their luminescence in the dark, hence their use in luminous paints. This property of phosphorescence seems to be dependent on the presence of certain impvu"ities, such as bismuth. Calcium sulphide is prepared by a method similar to that described for barium, CaSOi + 4C -^ CaS + 4C0, also by methods indicated in the following equations : 2CaO+3S^-2CaS+S02 2CaO + CS ., -^ 2CaS + CO ^ CaS04 + 4C0 -> CaS + 4C0 ,. Its properties resemble those of barium sulphide very closely. If hydrogen sulphide is passed through a solution of barium or calcium hydroxide, or if barium sulphide is dissolved in water containing hydrogen sulphide, crystals of the hydrosulpMde can be isolated from the solution. The solutions of the hydro- sulphides possess the property of dissolving appreciable quantities of sulphur, forming polysulphides (cf. the sulphides and hydro- sulphides of the alkahes). Phosphates. — Calcium phosphate occurs in several minerals in a more or less pure form as phosphorite, Ca3(P04)2. apatite, 3Ca3(P04)2,CaF2. whilst guano also contains considerable quantities of calcium phosphate. The importance of this compound hes in its use as a fertiliser. Nature has made available for plant life suppUes of phosphates arising from the decomposition of rocks, but the steady depletion of the soil by such crops as wheat necessitates the use of phosphatic CALCIUM, STRONTIUM, BARIUM 541 fertilisers. The phosphatic minerals are not in themselves suitable for this purpose owing to their insolubility. In order to make them readily available to plant Ufe, the mineral phos- phates (also bone ash, q.v.) are converted into the more soluble primary phosphates by the action of chamber acid. CaaCPOi)^ + 2H2SO4 + 5H,0-^ CaH4(P04)2,H20 +2[CaS04,2H,0] This so-called " superphosphate " is fairly soluble in water and by its use a more rapid distribution of the phosphates for plant life is effected. The secondary phosphate CaHPO} and orthophosphate, Ca3(P04)2 are relatively unimportant. The phosphates of cal- cium, strontium and barium are freely soluble in the presence of strongly dissociated acids. The explanation of this is similar to that put forward to account for the solubility of calcium oxalate in a strongly dissociated acid {q.v.). Silicates. — The siUcates of the alkaUne earth elements all occur naturally. The simplest and most important of these naturally occurring silicates is woUastonite, CaSiOj (mono- clinic). It can also be prepared by fusing together calcium oxide or carbonate with silica : CaO-f Si02->CaSi03. Many of the complex silicates, e.g. mica, contain calcium silicate as an integral part of their composition. Calcium Oxalate. — Calcium oxalate is precipitated by the addition of a soluble oxalate to a solution of a calcium salt. CaCla + (NH4),C204^2NH4C1 + CaC^O^ ^^ The separation of the precipitated calcium oxalate takes place because the product (Ca+ + )(C20i") exceeds the extremely small solubility product of calcium oxalate. The addition of a highly dissociated acid, e.g. hydrochloric acid, to a solution in which calcium oxalate is suspended, causes the salt to pass back into solution. The question arises, "Why does the product (Ca+ + )(C204-) no longer exceed the solubility product ^Qg^CiOi ' In a saturated solution of calcium oxalate the following equilibria are set up, 542 AN INORGANIC CHEMISTRY CaC204^=±CaC204^=^Ca++ +C2O4" Solid. Dissolved. Owing to the extreme insolubility of calcium oxalate one may safely assume that, what httle calcium oxalate is present in solution, wiU be almost completely dissociated. If a solution of hydrochloric acid is added to the saturated solution of calcium oxalate, we are, in effect, bringing into the system undissociated molecules of hydrochloric acid in equihbrium with its ions, HC1^=±H++Cl3, and as hydrochloric acid is a strong electrolyte, it will be widely dissociated. After the mixing of the two solutions, there will be present the ions Ca + + , C2O 4 = , H ■*■ , CI" . But wherever hydro- gen ions and oxalate ions are present in a solution, the equih- brium 2H++C204-^=:±H2C204 must be estabhshed. Combination will at once take place between the ions to produce the equUibrium concentration of undissociated oxaUc acid. To what extent will this combination take place 1 Oxalic acid is a very weak acid, so that few free hydrogen ions and oxalate ions can remain in solution, i.e. far- reaching combination wiU occur between the ions H"*" and CaOj", leading to the formation of undissociated oxalic acid. This reaction 2H+-f C204=^=±H2C204 exercises a disturbing influence upon the equihbrium, CaC204 ^=± CaC204 ^± Ca+ + + C2O4- Solid. Dissolved. removing oxalate ions, and in order to re-estabhsh this equili- brium, more of the undissociated calcium oxalate must dissociate. This leads to more of the undissolved calcium oxalate passing into solution. The continued addition of hydrochloric acid will therefore lead to the solution of the precipitate of calcium oxalate. The matter may also be stated in this way. The effect of adding hydrochloric acid to the saturated solution of calcium oxalate is to set up the equihbrium, 2H++C204=;=±H2C204, forcing it to the right, a necessary consequence of the Law of CALCIUM, STRONTIUM, BARIUM 543 Mass Action. The concentration of the oxalate ions is repressed to such an extent that the product (Ca+ + )(C204-) no longer exceeds L^^c^o^, and calcium oxalate passes back into solution. Consider now the system CaC204+ 2HCi ^=±CaCl2 + H2Ca04 It M t I Ca+++C204= 2H++2C1- Ca+++2C1 -2H+ +C2O4- If by any means one can remove the hydrogen ions from the solution, the dissociation H2C204^=^2H++G204- wiU be promoted ; if the removal of the hydrogen ions is continued, the time wiU come, sooner or later, when the concen- tration of the oxalate ions reaches such a value that (Ca+-)(C204-)>L,,^04 and a precipitate of calcium oxalate again separates out. This repression of the concentration of the hydrogen ions may be brought about by the addition of a soluble hydroxide, e.g. NH4OH, or a soluble carbonate Na2C03 to the solution. Such reagents react thus 2NH4OH + H2C2O4 ^=± (NH4)2C204 20H- 2H++C2O4- 2NH4++C2O The oxalate thereby formed is strongly dissociated, so that the concentration of the oxalate ions soon exceeds its critical value and the precipitate falls out. The action of sodium acetate in promoting the precipitation of calcium oxalate from a solution of calcium oxalate in hydro- chloric acid is not so obvious. Before the addition of the solution of sodium acetate we have CaCl2;=±Ca++-|-2Cl- H2C204^I^2H++C204- If a strong solution of sodium acetate, in which the equihbrium Na.C2H302 ^=± Na+ + G2H3O2- already rules, is added, we shall have present in solution the ions H+ and C2H302~, and these must react to estabhsh the equih- brium H+ +C2H3O2- ^^H.C^HaOa. 544 AN INORGANIC CHEMISTRY Acetic acid, H.C2H3O2, is one of the weak acids, and as sucli, will be but slightly dissociated, and the equilibrium will he far over to the right. Combination must therefore occur between the hydrogen and the acetate ions until their concentration has fallen to this equihbrium value. The continued addition of sodium acetate to the solution will force this equihbrium stiU further to the right, i.e. still more hydrogen ions will be removed from the solution in the form of undissociated acetic acid. The removal of hydrogen ions must, however, exert its influence upon the equilibrium H2C204^^2H++C204- and the dissociation of the undissociated oxaUc acid will be promoted. This leads to a steady increase in the concentration of the oxalate ions, and the product (Ca+ +)(C204=) soon exceeds the solubihty product L^j^c^g^, i.e. calcium oxalate will fall out of the solution. From a consideration of the above reactions the foUo^^dng general statement may be put forward : In general, an insoluble salt of an acid A will dissolve, when treated with an acid B, provided that the acid B is more highly dissociated than the acid A. The greater the difference in the strengths of the two acids, the more completely will the insoluble salt pass into solution. Secondly, if an insoluble salt is dissolved by the action of a more highly dissociated acid, the insoluble salt can be throvm out of solution by any agent which will repress the ionisa- tion of the added acid, i.e. repress the concentration of the hydrogen ions in the solution. Such an effect can be brought about by adding a soluble hydroxide or carbonate or by the addition of a strong solution of a salt of a very weak acid, e.g. sodium acetate. Calcium Carbide. — Calcium carbide is prepared by heating a mixture of lime and carbon in an electric furnace. CaO+3C->CaC2 + CO. The raw materials consist of freshly burnt lime and po\\dere(l anthracite, coal and in some cases tar. One of the most effective types of furnace in use is that of the Alby (Fig. 131. One electrode enters the base of the furnace, whilst the other electrode consists of a group of carbon rods. The molten carbide is run off and cast into blocks. In the pure state it is colourless, but is generally grey owing to the presence of impuri- CALCIUM, STRONTIUM, BARIUM 545 ties. It finds technical use in the manufacture of acetylene {q-v.) and also of calcium oyanamide. Calcium Cyanamide.— Calcium cyanamide, CaNjC, is pre- pared by forcing nitrogen at a temperature of about 800° into a retort filled with calcium carbide. CaCa + Na ^=^ CaNaC + C. Fig. 131 The temperature must not be allowed to rise above 1,400°, for the reaction is reversible. After the completion of the reaction the calcium cyanamide is removed, crushed and put on the market as a fertUiser. Its value as a fertUiser is based upon the decom- position it undergoes in the presence of water. CaNjC + 3H2O -> CaCOa + 2NH3 . A mixture of calcium cyanamide and carbon, prepared as above, is on the market as the fertihser " nitrohme." Calcium cyanamide is finding increasing application as a means of manufacturing ammonia from atmospheric nitrogen {see Fixation of Nitrogen, p. 278). Questions 1. A precipitate of magnesium hydroxide is formed on the addition of ammonium hydroxide to a solution of a magnesium salt, but in the presence of ammonimn salts, such precipitation does not take place. Account for this. N N 546 AN INORGANIC CHEMISTRY 2. Discuss the technical methods of softening water. 3. How does calcium carbonate occur in nature ? Discuss the con- ditions of stability of the different crystalline forms. 4. Compare the dissociation of barium peroxide with that of calcium carbonate. 5. Give a comparative survey of the more important properties of the compounds formed from the alkaline earth elements. 6. How do you account for the solubility of calciimi oxalate in hydro- chloric acid ? 7. How do you account for the solubility of calcium sulphate in a strong solution of ammonium sulphate ? 8. The solubility curve of calcium chloride shows several breaks. What inference is to be drawn from this ? 9. Barium sulphate is boiled with a solution of sodium carbonate. Account for the presence of sodium sulphate in the solution (Ch. XIV). 10. Compare the stability of the bicarbonates of the alkali and of the alkaline earth metals. 11. In qualitative analysis, barimn chromate is precipitated in the presence of acetic acid, whilst calcium and strontium chromates are not thrown down. Apply the ionic hypothesis to explain tliis. 12. Compare the action of heat upon the following compounds of calcium and barium : (a) carbonates, (6) nitrates, (c) hydroxides. CHAPTER XXXIV GROUP 2B : ZINC, CADMIUM, MERCURY General Relationships of the Family.— This sub-group again shows a steady gradation in the basic property of the oxide and hydroxide. Zinc hydroxide is amphoteric, a fact of great importance in the study of the chemistry of this element. The other elements have basic oxides. The general properties of the compounds of zinc, cadmium and mercury differ Uttle from those of similarly constituted salts of other metals. The stability of the chlorides and sulphates of the group decreases as the atomic weight increases. Cadmium forms two oxides, a sub-oxide, CdjO, and a normal oxide, CdO, but only one series of salts is known. Mercury, however, gives rise to three oxides, mercurous oxide, Hg20, mercuric oxide, HgO, and mercury peroxide, Hg02. Two well defined series of salts occur, the mercurous salts derived from mercurous oxide and the salts obtained from mercuric oxide. The group affords an interesting example of the tendency of the heaviest element of the B group to form more than one series of compounds. Zinc, Cadmium Occurrence. — ^Zinc occurs as carbonate (zinc spar, calamine, ZnCOs), sulphide (zinc blende, ZnS), oxide (zincite, ZnO), as silicate (willemite, 2ZnO,Si02), and as franklinite, a mixture of the oxides of Zn, Fe, Mn. Traces of cadmium are generally found associated with zinc in the carbonate and sulphide ores. Metallurgy. — The ore, which is in nearly all cases the car- bonate, oxide or sulphide, is first converted into oxide by the operation of calcining or of roasting. 547 548 AN INORGANIC CHEMISTRY Retort with condenser attached. ZnCOa— ^ZnO+COs (for the carbonate) 2ZnS + 30, ^ 2ZnO + 280, \ 2ZnS + 70 -> ZnS04 + ZnO + SO. for the sulphide ZnSOi-^ZnO+SOa i AU these various products are formed in the roasting of a sulphide ore. The crushed ore is then mixed with powdered coke and heated in specially designed retorts. ZnO+C->Zn + CO. Owing to the volatihty of the zinc the metal vapour distils over and coUects in a receiver. The main difference in the various methods used in the reduction and distillation stages is concerned with the shape, size and arrange- ment of the re- torts and furnace. There is a general tendency for the Rhenish type of retort (Fig. 132) to come to the fore. To the mouth of each retort is luted a fireclay conden- ser, in which the zinc coUects, and, attached to the cooler end of the condenser, is a conical iron nozzle to catch the zinc dust. The retorts are arranged in tiers, as many as 144 retorts being placed in a furnace. The loss of zinc during smelting operations is fairly high, as much as 10-15 per cent, being lost in the best plants. The impure spelter is refined by redistillation or by electrolytic separation. The electrolytic method of treating poor zinc ores has made rapid progress of late years. One important method is to convert the zinc into oxide. This is then treated with a solution of calcium chloride and carbon dioxide. ZnO_+ CaClj + CO^-^ CaCOa -|- ZnCl^. After iron and manganese have been removed from the solution, ZINC, CADMIUM 549 it is rendered slightly acid with hydrochloric acid, and electrolysed between an iron cathode and carbon anodes. The zinc is extremely pure, averaging about 99-96 per cent. A still more important electrolytic method of treating poor ores is to roast them, in order to oxidise the sulphide. The roasted ore is then leached out with dilute sulphuric acid ; the electrolyte, containing zinc sulphate in solution, is then electro- lysed between electrodes of lead oxide. The separation of zinc blende from galena (PbS) affords an interesting example of the froth flotation process. By means of an air blast the finely powdered ore is agitated with water to which a little eucalyptus or other oil has been added. The galena is wetted by the water and sinks, but owing to a surface tension effect the blende is buoyed up by the oil droplets, forming a scum. This method of separation is frequently used to effect a concentration of the mineral, e.g. molybdenite, from the accompanying earthy dross. Cadmium is generally obtained as a by-product in the metal- lurgical preparation of zinc. Being more volatile than the zinc, the cadmium collects in the portion which distils over first. This is again reduced with carbon and redistilled. Pure cad- mium is obtained from the electrolysis of cadmium sulphate, the crude metal forming the anode. Properties. — Zinc is a bluish-white, crystalline metal melting at 419°, boiling at 920°. Although brittle at ordinary temper- atures, zinc becomes quite ductUe at 100-150°, while at 200° it is so brittle that it can be powdered with ease. Cadmium also has a crystaUine appearance and is more ductile than zinc. It melts at 320°. Both metals are attacked vigorously by the mineral acids, though pure zinc reacts very slowly with pure sulphuric acid. Contact with a more electropositive (noble) metal, e.g. copper or platinum, causes the zinc to dissolve freely. This action is essentially electrolytic, for the hydrogen separates not from the surface of the zinc but of the more noble metal, i.e. in the cell zinc I acid | noble metal — ^ a current flows through the acid in the direction of the arrow, causing zinc to dissolve and hydrogen to be evolved. Zinc 550 AN INORGANIC CHEMISTRY dissolves freely in hot sodium or potassium hydroxide. This arises from the solubility of zinc hydroxide in these reagents. Zn + 2H0H ;=^ Zn(0H)2 + Ha Zn(0H)2 + 2NaOH ^ Na2Zn02 + 2H2O Sodium zincate. The great reactivity of zinc, arising from its high electrode- potential, has led to its use for protecting iron from corrosion. The iron is cleaned and either dipped into molten zinc or coated electroljrticaUy with that metal. When such galvanised iron is exposed to the atmosphere, the zinc is first corroded as it is the more reactive metal. Zinc is also used in many batteries, and in the manufacture of alloys (brass, bronze, etc.). Chemical Chabacteristics of the Salts The chemistry of the salts of zinc is intimately connected with the behaviour of the oxide towards acids and bases. Owing to its amphoteric nature, zinc oxide (hydroxide) forms two classes of salts, the normal type in which zinc functions as a divalent cathion, ZnClz, ZnSOj, ZnCOa, and a class of salts known as zincates, in which the zinc forms part of the acid group. These are produced by the solution of the hydroxide in an excess of sodium or potassium hydroxide. Zn(0H)2 + 2NaOH ;=i± Na^ZnOj + 2H2O. Acid. Base. Salt. On account of the weakness of zinc hydroxide as an acid, these zincates undergo very considerable hydrolytic decomposi- tion in aqueous solution, i.e. the above reaction will be reversible. The action of acids and bases in dissolving zinc hydroxide is shown in the scheme, /Zn+ + + 20H-^=±Zn(0H), Zn(0H)2 ^=±2H+ +ZnO. = . Dissolved. Dissolved. ^2C1- 1 -2H+ \\i // 20H-+2Na'-' Zn(OH), -1 Solid. I Zn+ + +2C1- + 2H20 2H2O +2Na+ + ZnO^ The addition of hydrochloric acid will remove hydroxyl ions and the disturbance of the equilibrium will lead to the solution of more zinc hydroxide ; in the same way, the addition of sodium ZINC, CADMIUM 551 hydroxide will lead to a reduction in the number of hydrogen ions, the equUibrium will again be disturbed and more zinc hydroxide wiU pass into solution. If ammonia is added to a solution of zinc or cadmium chloride, the precipitate which first separates out passes back into solu- tion. This is due in both cases to the formation of solu,ble ammonia salts, in which the ammonia is associated with the metal to form complex cathions [Zn.KNHs]^ ^ where x varies from 1 to 6, according to the strength of the ammonia. Cadmium forms a very strong basic oxide, CdO, which gives rise to a series of typical salts, such as CdCU, CdSOi. It is also reported that a suboxide, Cd20, has been prepared by breaking down cadmium oxalate in vacuo at a temperature of 300° ; but no salts of this oxide have been described. Oxides. — Zinc oxide is prepared either by burning zinc or by decomposing the carbonate. It is a white powder which turns yellow on heating. The white colour returns on coohng. As zinc white, the oxide finds considerable use as a white paint. Although possessing less covering power than white lead (q.v.) zinc white has the great advantage of not turning black on exposure to the fumes of hydrogen sulphide. Cadmium oxide is brown and is obtained by methods similar to those used for zinc oxide. The Halides. — Zinc chloride is made by dissolving zinc in hydrochloric acid, evaporating the solution to dryness and fusing the residue. The salt is very hygroscopic, and owing to the weakness of zinc hydroxide as a base, and the resultant hydrolysis, the dehydration leads to the loss of hydrogen chloride and a considerable amount of basic oxychloride is found in the fused chloride. ZnCU + HOH ^± Zn(OH)Cl + HCl 2Zn(0H)Cl ^r^ Zn^OClj + H^O If one desires a pure sample of the chloride, the dehydration must be carried out, as usual, in an atmosphere of hydrogen chloride. A solution of zinc chloride is used for injecting into wood as a preservative, i.e. to prevent the growth of organisms. It is also used in soldering, its use in this respect being based upon its solvent action upon the oxides. The hydrochloric acid, set free by the hydrolysis of the chloride, keeps the surface free of oxides so that the solder may take. 552 AN INORGANIC CHEMISTRY Cadmium chloride can be obtained not only in the anhydrous state, but also as a definite mono- and penta- hydrate. It is not hygroscopic, and, as is to be expected from the greater strength of cadmium oxide as a base, it is much less hydrolysed than zmc chloride. AU the hahdes of zinc and cadmium give rise to a large number of double haUde salts, with such haUdes as potassium chloride ; hence the existence of such salts as NajZnClj, K2CdBr4,. etc. In all these salts the salt of stronger electro-affinity (e.g. the alkali hahdes) forces the salt of weaker electro-afBnity (e.g. ZnCL, CdBra, etc.) into a complex anion (see p. 454). Carbonates. — Normal sodium carbonate throws down a basic zinc carbonate (cf. magnesium), but by using the acid carbonate the normal zinc carbonate is precipitated. ZnCls +2NaHC03 -^ ZnCOa i + 2NaCl + H2O + CO2. Both the normal and acid sodium carbonates throw down the normal carbonate of cadmium. Sulphates. — Zinc sulphate is made by roasting zinc blende in the air. The roast is afterwards extracted with water, and crystals of the hepta-hydrate ZnS04,7H20 obtained from the solution. These are isomorphous with the similar salts of magnesium, iron (FeS04,7HoO), cobalt and nickel. The term vitriol is often applied to these sulphates. Cadmium sulphate, 3CdS04,8H20, is obtained by crystallising a solution containing this salt. This salt forms large monochnic crystals, which pass into the monohydrate at 74°. Other hydrates are known. This sulphate is not isomorphous with the vitriols. Zinc sulphate forms a considerable number of double sul- phates, such as K2S04,ZnS04,6H20. These sulphates are monoclinic and are isomorphous with the double sulphates of iron (Fe"), cobalt, copper, etc. These compounds may be repre- sented by the general formula M2'S04,M"S04,6H20, where M' represents a monovalent metal or radicle, such as Na, K, NHj, M" a divalent metal such as Cu, Fe, Co, Hg, Zn, Mg. Cadmium sulphate also forms many double sulphates, e.g. Na2S04,CdS04,2H,0. Sulphides. — The precipitation of zinc sulphide by the action of hydrogen sulphide upon an aqueous solution of a zinc salt is summarised by the equations : ZINC, CADMIUM 553 ZnS04 + H2S H2SO4 + ZnS Z11++ +804= 2H+ +8= ^=±2H+ +804- Zn+ + +S- As the concentration of the acid increases, there is a consequen- tial rise in the concentration of the hydrogen ions : H2S04^=±2H++804-. This, in turn, plays its part in the equilibrium H2S^r±2H + +8", i.e. the ionisation of the hydrogen sulphide is thrown back. In short, through the production of sulphuric acid in the course of the reaction, the time soon comes when the concentration of the S" ion is so low that the product (Zn''" ■^) (8") no longer is able to exceed the solubility product L2ns, and no more zinc sulphide separates out. If sodium acetate is added at this stage, a removal of the hydrogen ions is effected in accordance with the equation H+ + C2H3O2- ^r:± H.C2H3O2 ; (cf. calcium oxalate, p. 543). the dissociation of the hydrogen sulphide is thereby promoted, and a further precipitation of zinc sulphide results. Cadmium sulphide is used as a [yeUow pigment. It is appreciably soluble in] hydrochloric acid, and for effective precipitation the solution should be kept only sufficiently acid to prevent the precipitation of such sulphides as zinc sulphide, manganese sulphide, etc. Separation of Copper from Cadmium. If a solution containing salts of copper and cadmium is treated with potassium cyanide, a precipitate is thrown down which redissolves in an excess of the reagent. The reactions for copper salts under these conditions have already been dis- cussed. In the case of cadmium compounds the reactions are as follows : CdCU + 2KCN -^ Cd{CN)2 + 2KC1 Cd(CN)2 + 2KCN -^K2[Cd(CN)4] II 1 1 I I I Cd+++2CN- 2K++2CN- 2K+ +Cd(CN)4- 11 Cd++ +4CN- 554 AN INORGANIC CHEMISTRY On passing in a stream of hydrogen sulphide, the equilibrium is established. Before a precipitation of cadmium sulphide can be obtained, the product (Cd+ + )(S-) must exceed the solubility product Lpjig. It has already been stressed that, in the presence of potassium cyanide the concentration of copper ions is so extremely low that the solubility product L^ ^^^ is never exceeded, and no precipitate of copper sulphide is thrown down. With cadmium salts, however, the dissociation Cd(CN)7 ^=iCd + + +4CN- yields a sufficiently high concentration of cadmium ions for the product (Cd + "^ ) (S " ) to exceed the value of the solubihty product, L^<,> ^^^ yellow cadmium sulphide separates out. Mercury Metallurgy. — The separation of mercury from its ore, cinna- bar (HgS) is effected by roasting. HgS + 02->Hg+S0,. Occasionally it is heated in retorts with lime. 2HgS + 2CaO -> 2Hg + 2CaS + 0.. The metal is distilled from the retort and recovered by sending the furnace gases through a series of condensation flues. It is then filtered through chamois leather. Moderately pure mercury may be obtained either by allowing fine drops of mercury to fall through a column of dilute nitric acid, or by putting the mercury and nitric acid in a filter flask and keeping the whole thoroughly stirred by means of a brisk current of air. Distillation in vacuo yields the purest mercury. Properties. — Mercury is a silver-white liquid which does not tarnish in aii-. (M.P. -^39-4°, B.P. 357-3°). It is attacked only by ozone, the halogens, hydrogen sulphide and iodide, hence its extensive use in the manipulation of gases. It is a good con- ductor of heat and electricity. The vapour is monatomic, is a non-conductor of electricity and is extremely poisonous. Mer- cury is not attacked by hydrochloric acid or by dilute sulphuric acid, but hot concentrated sulphuric acid soon dissolves it with the formation of sulphur dioxide and even hydrogen sulphide. MERCURY 655 Hg + H2SO4 -^ HgSOi + 2H 2H + H2SO, --> 2H2O + SO2 8H + H2S04^.4H20 + HaS Nitric acid, especially when hot and strong, attacks mercury freely, forming either mercurous or mercuric nitrate, according to the relative quantities of acid and metal. The solution of mercury in an acid is apparently only achieved when the acid is a sufficiently strong oxidising agent to oxidise the film of hydrogen precipitated upon the surface of the mer- cury. Dilute sulphuric acid is not a sufficiently strong oxidiser to do this, hot concentrated sulphuric acid is. Mercury reacts vigorously with certain metals, e.g. sodium, whUst in other cases slow solution of the metal in the mercury takes place. The product is generally referred to in both cases as an amalgam. Physico-chemical research has shown that many of these amalgams contain definite cheniical compounds, dissolved in an excess of mercury ; sodium gives the definite crystaUine compounds NaHgj,, NaHgj. Some of these amalgams find commercial appUcation. Thus a copper-cadmium amalgam is used by dentists for filling teeth. General Properties of the Compounds of Mercury Attention has already been called to the fact that mercury forms three oxides of which two, mercurous oxide HgaO, and mercuric oxide HgO, give rise to a series of salts, the mercurous and the mercuric salts respectively. Until comparatively recent times the mercurous salts were held to be monovalent, the mercuric divalent, but research has definitely established that the mercurous salts are also divalent, the weight of the -ous cathion being twice as great as the weight of the -ic cathion. The formulae of a few typical mercurous salts are Hg2Cl2, HgaSOi, HgslNOa),. These ionise thus : Hg2Cl2^z±Hg,^-+2Cl- whereas mercuric chloride and sulphate ionise in the following manner ; HgCl2^z±Hg' + +2Cl- HgSO,^=±Hg+++SOr. 556 AN INORGANIC CHEMISTRY The chemistry of the compounds of mercury is closely bound up with the equation : This states that the mercuric salts are reduced by mercury to the mercurous state. This equation embodies the general method of preparing the mercurous salts — the reduction of the corre- sponding mercuric salt with mercury. The hydroxides are both unstable, and soon pass into the oxides. As bases they are feeble, so that both the mercurous and the mercuric salts are extensively hydrolysed in aqueous solu- tion. In fact, many of the salts can only be retained in solution when a considerable excess of acid is present e.g. Hg2(N03)2, Hg(N03)2, etc. For a similar reason no mercuric carbonate has been prepared. Just as we found that the cuprous and aurous salts were comparatively insoluble, so we find that all but a few of the oxy-compounds of mercurous mercury are insoluble. Mercuric chloride, HgCl2, is soluble in water (cf. CuCU, CdCU, ZnCl2, PbC'lo, etc.), mercurous chloride, Hg2Cl2, is insoluble (cf. AgCl, AuCl, CuCl, etc.). The haUdes of the mercuric type are freely soluble in the presence of haUdes such as NaCl, KI, etc. This is due to the formation of soluble complexes, viz. : Hgl2+2KI;=z±K2(Hgl4). In many of these complex salts the ionic concentration of the mercury ion is extremely low. Thus, the addition of potaasiimi hydroxide to a solution of potassium mercuric iodide causes no precipitation of mercuric hydroxide or oxide. Intimately linked with this is the fact that mercury sulphide dissolves freely in hydrogen iodide : HgS + 2HI ^izi: H2S + Hgl, Hgl2+2HI^±H2(HgI,). The salts of mercury of both types differ from their analogues, cadmium and zinc, in their behaviour towards ammonia. No evidence concerning the existence of such ions as (Hg.xNHa) has been furnished. On the other hand, a new type of compound is obtained by the action of ammonium chloride upon mercuric chloride, partial replacement of the hydrogen present in the MERCURY 557 ammonium group is effected and compounds are formed of the type ^ \ H An— 01. The halides and cyanides of mercury differ from all other similarly constituted salts (except those of cadmium) in the slight degree to which they are ionised. An aqueous solution of mercuric cyanide scarcely conducts the electric current. An- other specific property of the salts of mercury which distinguishes it from the other members of this sub-group is the ease with which they can be sublimed, and to a less extent their highly poisonous nature. It is owing to the comparative insolubility of the mercurous salts that some of them (calomel) can be used medicinally. Mercuric Salts. Mercuric Oxide. — If a solution of a mercuric salt is treated in the cold with sodium hydroxide, the precipitate of mercuric oxide thrown down is yellow. If one works at a higher tem- perature, the colour is red. Whether there are two distinct modifications or whether the difference in colom: arises merely from a difierence in the granular structure of the particles is doubtful, though the bulk of the evidence favours the latter view. The red oxide is generally precipitated by breaking down the nitrate or by heating the element with oxygen above 300°. When heated strongly, it first turns black and then evolves oxygen. In aU its reactions the yellow form is more reactive than the red, probably owing to its finer state of subdivision. This finer grain of the yellow oxide would also account for its greater solubUity (cf. influence of size of grain upon the solubility of barium sulphate, p. 538). Mercuric Halides. — Mercuric chloride, corrosive sub- limate, HgClz, is made by subliming mercuric sulphate with sodium chloride. HgSOi + 2NaCl-> HgClj + Na^SOi- It crystallises in white rhombic crystals. Its solubility shows a rapid rise with the temperature (at 0° 5-7 gm. dissolve in 100 558 AN INORGANIC CHEMISTRY gm. of water, at 100° 53-9 gm. dissolve). Both mercuric chloride and bromide are freely soluble in alcohol, ether and benzene. The chloride is easily reduced to the mercurous state by such reducing agents as stannous chloride, 2HgCl2 + SnCls -> HgaCU i + SnCli, while in the presence of an excess of stannous chloride the reduction goes still further. HgaCla + SnCU -^ 2Hg >^ + SnCU. A large number of complex salts has been prepared from mer- curic chloride, bromide and iodide. They are generally obtained by crystaUisation from a solution containing the mercuric halide dissolved in the presence of another haUde, e.g. KCl, NaBr, etc. All the alkaU metals give one or more such salts, e.g. KHgCl3,H20 (starry needles soluble in water), K2HgCl4,H20 (rhombic, soluble in water), KHg2Cl5,2H20 (rhombic). A dilute solution of corrosive subUmate is an excellent ger- micide, hence its use in surgery. On the other hand, when taken internally, it is extremely poisonous. Mercuric Bromide differs little from the chloride. Mercuric Iodide is of interest owing to its dimorphous occur- rence. Below 126° the red is the stable form, above that tem- perature the yellow. In precipitating the iodide from a hot solution of potassium iodide, there is first thrown down a yeUow precipitate, the unstable modification, but this soon changes into the stable red form. It is an interesting example of Ost- wald's Law of Transformation by Steps (cf. p. 310). A similar phenomenon is observed it fused mercuric iodide is allowed to cool. The tendency of mercuric iodide to form complex iodides is shown by its great solubihty in an excess of potassium iodide. The great stabiUty of such complexes of mercuric iodide is shown by the non-precipitation of mercury hydroxide (oxide) on the addition of a soluble hydroxide to solutions of such complex salts. Mercuric Cyanide. — Mercuric cyanide is prepared by acting upon the oxide with hydrogen cyanide. It is freely soluble in an excess of potassium cyanide, forming a complex cyanide, K2Hg(CN)4. It is fairly readily soluble in water, yet owing to its shght ionisation neither sodium hydroxide nor potassium iodide gives a precipitate with a solution containing mercuric MERCURY 559 cyanide. Hydrogen sulphide alone will decompose mercuric cyanide. Mercuric Sulphate.— This salt is formed by heating mercury with concentrated sulphuric acid, or by acting upon the oxide with sulphuric acid. It occurs in rhombic crystals. Its aqueous solutions are strongly hydrolysed and throw down insoluble basic salts if an excess of sulphuric acid is not added. When gently warmed, it becomes yellow, then red, and finally decomposes into mercury, oxygen, sulphur dioxide and mercurous sulphate. Mercuric Nitrate. — ^Mercuric nitrate is produced by the action of an excess of nitric acid upon mercury. The aqueous solution throws down a basic nitrate, Hg(N03)2,2HgO,H20. The precipitation of this can be prevented by the addition of nitric acid according to the equation 3Hg(N03)2 +3H20^=^Hg(N03)2,2HgO,H20 +4HNO3. Mercuric Sulphide. — ^Mercury sulphide occurs naturally as cinnabar, but it can also be obtained as a black, amorphous powder by the wet method, as well as by rubbing together mer- cury and sulphiu". The black variety is unstable, and in the pre- sence of a suitable solvent (e.g. sodium sulphide) slowly turns into the stable red form. The solution becomes saturated with regard to the black, and therefore supersaturated with respect to the stable red modification. Mercuric sulphide is exceedingly insoluble and unreactive. Concentrated nitric acid attacks but does not dissolve it. Its solution can be effected by means of hydriodic acid or by aqua regia. The red sulphide is used as a red paint (vermiUion), but has less vogue than red lead owing to the ease with which the less noble (more electronegative) metals, Fe, Zn, Pb, etc., displace it. HgS+Fe^Hg+FeS. Mercury Fulminate. — Mercury fulminate, Hg(0NC)2, is formed when mercury is treated with nitric acid and alcohol added to the mixture. It is a white, explosive powder, used extensively in percussion caps. Mercurous Salts. Mercurous Oxide. — ^Mercurous oxide is thrown down from a solution of a soluble mercurous salt by the addition of a soluble base. 560 AN INORGANIC CHEMISTRY Hg2(N03)2 +2NaOH^Hg20 ^^ +2NaN03 +H2O. The brownish black powder is HgaO, the hydroxide under- going almost instantaneous decomposition. Under the action of light and heat it passes into mercury and mercuric oxide. At 100° it unites with oxygen, forming the oxide HgO. Mercurous Halides. — Mercurous chloride, bromide, etc., are obtained by subliming the mercuric salt with an excess of mercury, HgCl3+Hg->Hg2CU or, more frequently in the commercial \\orld, by heating a mixture of mercuric sulphate, mercury and sodium chloride. The vapour is condensed. Under the action of hght slight decomposition into mercuric chloride occurs, Hg.Cl^^z^HgCU+Hg the white powder darkening shghtly. Attempts were made to determine the vapour density of this compound in order to de- cide whether the formula should be HgCl or HgsClj. Although the experimental results gave a value approximating to 240, the conclusion could not be drawn that the formula was HgCl(M.W. =200+35=235), because it was discovered that, when a strip of gold leaf was placed in the apparatus, it became amalgamated. This proved that free mercury was present, so that the results might equally well be explained on the assumption that the molecule Hg2Cl2 had been broken down at the tem- perature of the experiment into Hg+HgCL, thereby causing a doubling of the number of molecules and a halving of the molecular weight. Other methods have since shown that the formula is Hg2Cl2. The mercurous hahdes are insoluble in water. Mercurous chloride (calomel) is extensively used for medicinal purposes. Mercurous Nitrate. — Mercurous nitrate is formed by dissolving mercury in dilute nitric acid, or by reducing the mercuric nitrate with mercury. It forms monooUnic crystals, Hg2(N03)2,2H20. Owing to the tendency of this salt tohydro- lyse, a clear aqueous solution can only be obtained in the pre- sence of an excess of nitric acid. Many basic mercurous nitrates have been described. Mercurous Sulphate. — ^Mercurous sulphate is prepared by rubbing mercuric sulphate with mercury, or by precipitation MERCURY 561 from a solution of mercurous nitrate by the addition of sodium sulphate. It forms smaU monoclinic crystals of slight solubility. It dissolves readily in hot sulphuric acid. Mercurous Sulphide. — Mercurous sulphide is unstable at ordinary temperatures. Like the chloride, it tends to undergo auto-oxidation, forming mercuric sulphide and mercury. Mercurous Carbonate. — ^Mercurous carbonate is precipi- tated by the addition of sodium carbonate or bicarbonate to a solution of mercurous nitrate. It is a yellow powder which darkens on exposure to Ught. Under the action of heat it forms carbon dioxide, mercury and mercuric oxide. Mercury Amine Compounds. The capacity of mercury to displace hydrogen from the ammonium radicle has already been mentioned. This is especi- ally noted in the case of the mercuric compounds. The addition of ammonium hydroxide to mercuric chloride gives a white pre- cipitate known as infusible white precipitate. H / CI— N=Hg \ H Mercurous chloride gives the same salt, together with an excess of mercury, i.e. the equUibrium Hg, + -^^=ziHg+Hg-' + is swung to the right, the resulting precipitate appearing black. If mercuric iodide is dissolved in an excess of potassium iodide, the soluble potassium mercuric iodide is formed Hgl2+2KI^=z±K2(HgI,). The addition of potassium hydroxide to this solution causes no precipitation of mercuric oxide or hydroxide, on account of the very low concentration of mercury ions present, i.e. the complex is extremely stable. This alkahne solution of potassium mercuric iodide is known as Nessler's Reagent, and it is an exceedingly delicate test for traces of ammonia. A minute trace of ammonia produces a distinct yellow coloration due to the formation of 00 562 AN INORGANIC CHEMISTRY oxy-dimercurio ammonium iodide, OHgjNHal. By preparing a series of standard solutions of known strength or by the use of a colorimeter one can estimate accurately very small traces of this base. Mercury Peroxide. — Mercury peroxide is reported to be formed as a reddish brown powder when mercury is left in contact with hydrogen peroxide. It is extremely unstable, breaking down into oxygen and mercury. Questions 1. Discuss the reaction ZnCla + HjS ^ZnS + 2HCl. ^\^lat effect has the addition of ammonium acetate upon this equilibrium ? 3. A solution of copper and cadmium chlorides is treated with an excess of potassium cyanide, and hydrogen sulphide passed through. Account for the precipitation of cadmium sulphide and the non-precipitation of copper sulphide. 3. Discuss the action of (a) dilute nitric acid, (6) strong nitric acid, upon zinc and mercury respective! 3^. 4. Give an accoimt of the preparation and properties of zinc hydroxide. 5. Give a comparative account of the oxides of the metals, zinc, cadmium and mercury. 6. Discuss the reaction Hg+ + +Hg^Hg2+ + and show how this reaction may be used for the preparation of the mercurous salts. 7. What is the action of (a) ammoniima hydroxide, (6) sodium hydroxide, upon a solution of zinc sulphate ? Construct ionic equations showing the mechanism of the reactions. CHAPTER XXXV GROUP 3 : BORON, ALUMINIUM, GALLIUM, INDIUM, THALLIUM; SCANDIUM, YTTRIUM, LANTHANUM The position of these elements in the middle of the Periodic Table is no doubt responsible for the difficulty of determining whether the two short period elements — boron and aluminium — fall into the A or the B sub-group. The distinctions which have been noted in the properties of the A and B sub-groups of Groups 1 and 2 have almost disappeared. An even better example of this is afforded by Group 4. The evidence in favour of allotting boron and aluminium a definite position is conflicting ; for the purpose before us the existence of the double sulphates (alums) of aluminium, gallium and indium is perhaps sufficient justification for classing together boron, aluminium, gallium, indium and thallium in Group 3b. Group 3A — Scandium, Yttrium, Lanthanum The rare elements, scandium, yttrium and lanthanum, are aU trivalent in their more important compounds, giving the group oxide M2O3, and the chloride MCI3. In all cases the oxides are basic, the basicity increasing in passing from scandium to lanthanum. They form stable carbonates and give well-defined double suJpliates with the alkali sulphates, quite distinct in composition and crystal structure from the alums. Scandium is the eka-boron predicted by Mendeleeff. It occurs in fergusonite and gadolinite. Yttrium and its related elements occur in cerite, monazite, yttrotantaHte, etc. Closely associated with yttrium and lanthanum are the rare earths, amongst which may be mentioned cerium, praseo- 563 564 AN INORGANIC CHEMISTRY dymium, neodymium, samarium, gadolinium, dysprosium, holmium, erbium, thulium, ytterbium. The properties of these elements are so much alike that it is often impossible to effect a separation. Opinion is still divided as to how these elements are to be fitted into the Periodic Table. A rough division of these rare elements into two groups can be effected by saturating the neutral solution of the oxides in nitric acid with potassium sulphate. Double sulphates of the yttrium group (Y, Yb, Er, Ho, Tm, Tb, Sc) remain in solution while the double sulphates of the cerium group (Nd, Pr, La, Ce, Sm, Gd) are rendered insoluble. Group 3B — Boron, Aluminium, Gallium, Indium, Thallium BOEON General Remarks. — Boron is the least metallic of the ele- ments of the group, a fact ^hich is not to be wondered at when one considers its position between glucinum with its amphoteric oxide and carbon with its acidic oxide. Boron, as an element, is to be classed as a metalloid, a transition element lying between the pronouncedly metallic elements and the non-metals. Its chemical nature is shown by its tendency to form borides with metals and a stable hydride (cf. nitrogen), furthermore by the stability of the borates. These facts compel one to class it as a non-metal, but the distinct amphoterism exhibited by its only oxide, B2O3, shows that the element has also certain properties generally associated \\ith the metals. In other words, boron trioxide is a weak acid oxide, but also possesses very weakly defined basic properties. Boron is habitually trivalent in its compounds. Occurrence. — The acidic nature of boron trioxide, coupled with its non-volatility, accounts for the occurrence of this element in the form of borates. Of the more important minerals containing boron may be mentioned borax (tincal, NaaBiO,), boric acid (H3BO3), boracite (MgCL+2Mg3BgOi5), colemani'te (Ca2B50ii,5H.,0), borocalcite (CaBiOj.iHaO), boronatrocalcite (Na2B40„2CaB407,a;Il20). Preparation of the Element.— Amorphous boron can be prepared by the i eduction of the oxide with a strong reducing agent, e.g., sodium, aluminium, magnesium. The fused mass is BORON 565 heated with dilute hydrochloric acid, whereby the excess of magnesium, etc., is dissolved out, leaving a brown mass of amorphous boron. Other methods are indicated in the equations : BF3 + 3Na -^ 3NaF + B 2BCI3 + 3H2-> 2B + 6HC1. Crystalline boron can be obtained by reducing boron trioxide with a considerable excess of aluminium. The reduced boron dissolves in the fused aluminium, from which it separates in a crystalline form as the temperature falls. Properties. — ^Amorphous boron is rather more active than the crystalline variety. It is a powerful reducing agent, decomposing steam and sulphur dioxide at a red heat, whilst at 1,200° carbon monoxide and sUica are reduced. Concentrated nitric and sulphuric acids slowly attack it, forming boric acid. Crystalline boron is much less easily oxidised. Even hot, concentrated nitric acid has little action upon it. Both forms of boron dissolve in fused potassium or sodium hydroxide, forming a borate. 2B + 6KOH^2K3B03 + 3H With many elements direct combination takes place at a high temperature with the formation of borides, e.g. sHicon boride, SiBs ; a few metals reduce boron trioxide to a boride : 6Mg +B203->Mg3B2 +3MgO. Hydride. — By reacting upon crude magnesium boride with hydrochloric acid and fractionating the evolved gases by means of Uquid air, a number of gaseous hydrides of boron have been isolated, e.g. BaH^, BiH^o, B^B.^^, etc., probably somewhat similarly constituted to the many hydrides of the neighbouring elements, carbon and sihcon. Whether boron trihydride BH3 exists is a matter of doubt, though its polymer BjHe has been separated. The existence of the hydride B3H3, reported by Ramsay, is doubtful. Boron Nitride. — ^This compound can be formed by the direct combination of boron and nitrogen at a temperature of 600°- 800° ; also by the reduction of boron trioxide with carbon in a current of nitrogen : B,03+3C+N3->2BN + 3CO. 566 AN INORGANIC CHEMISTRY It has been prepared on a large scale from borocalcite, carbon and nitrogen at a temperature of 1 ,400°-l ,800° : CaB.O, + 8C + 6N->4BN + CaCN^ + 7C0. This method was tried as a commercial means of fixing atmo- spheric nitrogen, but on account of the high temperature required for the reaction, the method proved unsatisfactory. The nitride is a white solid, which is decomposed by the action of water, alkali hydroxides and acids : BN + 3H20-^B(OH)3+NH3. Boron Carbide. — This compound is made by heating a mixture of sugar carbon and amorphous boron in the electric furnace. It forms brilliant black crystals, which are extremely hard, hence their use for polishing diamonds. It is very resistant to the action of chemical reagents, though fused potassium hydroxide decomposes it with the evolution of carbon monoxide. Halides.' — -AU the halogens combine with boron to form halides, BF3 (gas), BCI3 (hq.), BBr, (Uq.), BI3 (sohd). These compounds can be obtained by the direct combination of the elements, or by heating the oxide -nith carbon in an atmosphere of the halogen : B203+3C + 3Cl2-^2BCl3 + 3CO. Other methods are also available, e.g. B2O3 + SCaFa + SH^SOi^ 2BF3 + SCaSO^ + SH^O. The haUdes are generally purified by fractional distillation. Boron trichloride is a colourless, fuming hquid, which is practically completely decomposed by water. BCI3+ 3HOH^B(OH)3 + 3HCl. The bromide behaves similarly. The iodide forms large, leafy, hygroscopic crystals, which also suffer hydrolytic decomposition in water. Boron Trifluoride behaves somewhat differently when treated with water. An interaction occurs between the hy- drolyticaUy formed hydrogen fluoride and the undecomposed boron fluoride, yielding hydrofluoboric acid. BF3 + 3H0H-> 3HF + B(0H)3 BFj + HF^HBF^ BORON 567 (cf . the action of silicon fluoride on water and the formation of hydrofluosilicic acid, HaSiFo). Although hydrofluoboric acid can only be obtained in solution, its salts, the borofluorides, are well known, KBF4, Zn(BFi)2, and Al(BF4)3,a;H20, forming examples of such salts from metals of Groups 1, 2 and 3. Boron Oxide — Boric Acid. — Boron forms only one oxide, the chemical identity of which can be considered to be definitely established. It is most easily prepared by dehydrating boric acid : 2H3B03-3H,0->B203. It is a glassy, hygroscopic mass which is very resistant to heat. Its non-volatility is frequently taken advantage of for getting rid of more volatile oxides, e.g. : K2S04 + B,03-^2KB02 + S03l^ Boron oxide (hydroxide) displays rather weak acidic properties, hence the existence of the borates, while at the same time the capacity of this oxide to function as a very weak base leads to the formation of such compounds as boron phosphate, BPO4 (prepared from boric acid and phosphoric acid). It also accounts for the existence of boryl pjrrosulphate, BO.HS2O7. Boron hydroxide, B(0H)3, is commonly known as boric (boracic) acid. It occurs naturally in. jets of steam, issuing from the around (Tuscany). The steam condenses in lagoons of water surrounding the outlets, and as the liquor concentrates, boric acid crystallises out in characteristic leafy crystals, which are soapy to the touch. It is but sKghtly soluble in cold water, 4-0 gm. dissolving in 100 gm. of water at 20°, 29-1 gm. at 102°. The aqueous solution is a very weak acid, scarcely affecting litmus. It finds extensive use as an eye lotion and mild antiseptic. On heating orthoboric acid to 100°, partial dehydration occurs with the formation of metaboric acid : H3BO3 — H^O^^HBOa. Further heating at 140-160° produces tetraboric acid, The dehydration can be pushed even further, and the non- volatile oxide obtained (cf. the dehydration of orthophosphorie acid). Very few salts of orthoboric acid are known, though 568 AN INORGANIC CHEMISTRY certain tri-alkyl esters have been isolated, B(0R)3 where R denotes a radicle such as C2H5. The metaborates and tetraborates are common. Of the former we have NaBOa, Ca(B02)2, KBO2, Pb(B02)2, and several others. Of the tetraborates by far the most important is borax, Na2B4O„I0H,O. A pentahydrate is also known. Borax is made from tincal by extraction with water. It is also obtained from boronatrocalcite by boiling with a solution of sodium carbonate. The filtrate is allowed to crystallise and crystals of borax separate out : Na2B40,.2CaB40,+2Na2C03^-3Na2B40, +2CaC03^^ Aqueous solutions of sodium tetraborate react strongly alkaline. TWs arises from the weakness of the acidic group (BOH) 3. Na^B.O, + 2H0H ^:± 2NaOH + H2B4O,. The sodium hydroxide being a strong base, will be extensively dissociated, whilst the weak tetraboric acid is practically undissociated. There is thus an excess of hydroxyl ions in the solution and this is shown by the alkaline reaction. If the formula NajBiO, is written 2NaB02.B203, it is seen that borax contains an excess of the acid oxide. Upon this fact is based the use of borax in borax bead testing. If a trace of a metaUic oxide is fused with sodium tetraborate, interaction between the metaUic oxide and this excess acid group takes place, and the complex borates formed have characteristic colours. The use of borax as a flux in soldering is based upon the same principle. Perborates. — The perborates of the alkali metals are prepared by the action of hydrogen peroxide or the alkali peroxides upon an aqueous solution of boric acid, or a borate : NajBiO, + 2NaOH + 4H202^^ 4NaB03 + SH^O. They are also formed in small quantities at the anode during the electrolysis of the borates. Sodium perborate, NaB03,4H20, crystallises in large mono- clinic prisms, and is stable in air, provided no carbon dioxide is present. On warming its aqueous solution, oxygen is evolved. It is an energetic oxidising agent, oxidising ferrous salts to ferric, manganous salts to manganese dioxide, etc. Other perborates are Ba(B03)2, Mg(B03)2, €0(603)2, Cu(B03)2. ALUMINIUM 569 Aluminium Occurrence. — Aluminium occurs extensively in the form of oxide, halide, aluminate and silicate. The ruby, sapphire and corundum are more or less pure oxide, AI2O3, while in the hydrated form aluminium oxide is found as diaspore, AlaOj.HaO ; hydrargyUite, AlaOs.SHaO ; and bauxite, Al203,2H20 ; as aluminate, the spinel group serves as an example, gahnite, ZnAljOi (ZnO,Al203) ; spinel, MgAljOi ; chrysoberyl, BeAljOi. As double fluoride cryolite is of the greatest importance in the metallurgy of the element ; finally, as silicate, its occurrence is widespread, the felspars and clays being good examples of this class of naturally occurring aluminium compounds. Potash felspar or orthoclase is KAlSiaOj. Metallurgy. — The entire world's production of aluminium is obtained by the electrolysis of fused aluminium oxide, fused cryohte, NagAlFo, acting as the sol- vent. The ceU con- sists of a wrought- iron box, the bot- tom of which is lined with carbon plates. These serve as the cathode. The anodes are a set of carbon rods arranged in rows and suspended in the fused electrolyte. The alumina used during the electro- lysis must be quite pure. Nowadays the purification is generally efiected by the Bayer process, which is described later. During the electrolysis carbon dioxide escapes from the anode, whilst aluminium collects at the bottom of the cell and may be run off at mtervals. There is very little loss of cryolite durmg the electrolysis (Fig. 133). Properties. — Aluminium is exceedingly Ught, tough and of great tensile strength (M.P. 657°). It is attacked superficially by oxygen and a protecting film of closely adhering oxide formed. With frequent annealing it can be rolled into sheets or drawn into wire. As a conductor of heat and electricity it is good. Some of its alloys find extensive use in the construction of aeroplanes owing to their toughness and hghtness. Aluminium ■Electrolyte -Aluminium. Ftq. 133. 570 AN INORGANIC CHEMISTRY bronze (95 per cent, copper) is easily fusible, Ught, strong and elastic, and takes a beautiful polish. Aluminium is used in the metallurgical purification of iron. Small quantities of aluminium are added to the molten iron to remove dissolved gases, and hence prevent the formation of air holes in the iron. In a finely divided form aluminium is a most energetic reducing agent. A mixture of aluminium powder and ferric oxide [thermite) is often used for welding in situ : The fuse is fired by means of magnesium wire and a httle barium peroxide. At the high temperature generated during the reaction, the ii'on melts and runs into the crack round which the thermite has been packed. This is the basis of the Goldsehmidt " alumino- thermic " process of ^^•elding. All the oxides, except that of magnesium, are similarly reduced by aluminium. Aluminium reacts freely with hydrochloric acid, but is scarcely attacked by nitric or cold sulphuric acids. Hot sulphuric acid, however, effects solution, sulphur dioxide being evolved. It dissolves easily in a boiling solution of the alkah hydroxides. This is closely associated with the solubUity of aluminium hydroxide in sodium or potassium hydroxide, so that the surface of the metal is kept free of any protecting film. : 2 Al + 6NaOH ^2Na 3AIO 3 + 3H ^ . Sodium Aluminate. It is also dissolved by an aqueous solution of sodium carbonate. With chlorine, the formation of the chloride proceeds vigorously ; combination with such elements as phosphorus, carbon, sulphur, is only brought about at fairlj' high temperatures. Chemical Characteristics of the Compounds of Aluminium The chemistry of aluminium is indissolubly bound up in the properties of its oxide, ALOs, and its hydroxide A1(0H)3. The position of the element in the Periodic Table as the second member of Group 3, standing roughly midway between sodium, which forms a strongly basic oxide, and chlorine with its acid- forming oxide, is accountable for the amphoteric nature of this oxide and hydroxide. In a saturated solution of aluminium hydroxide the following equiUbria are set up : ALUMINIUM 571 A1+ + + +30H-^^A1(0H)3 Al(OH)3^r±3H+ + A103 Dissolved. Dissolved. "^^ A1(0H)3 -^ SoUd. The addition, either of a strong base or of a strong acid will disturb these equilibria, the former by the removal of hydrogen ions, the latter by the removal of hydroxyl ions, so that in either case more of the hydroxide must pass into solution. By the addition of strong acids to aluminium hydroxide there will be formed aqueous solutions of aluminium salts in which aluminium functions as a trivalent base, e.g. : 2A1(0H)3 + 3H,S04 ;=± A1,(S04)3 + 6H3O. The weakness of aluminium hydroxide as a base accounts for the ease with which such salts of aluminium undergo hydrolytic decomposition ; at the same time it explains the non-existence of salts, such as the carbonate and sulphite, and the decomposition of aluminium sulphide in aqueous solutions. On the other hand, the action of such bases as sodium hydroxide upon aluminium hydroxide leads to the formation of a new type of salt in which aluminium functions as part of the anion. These salts are termed ahiwAnates ; their formation and constitution are made clear by the following equations : Al(OH)3^=^3H+ +A103= 3NaOH ^^ 30H- + 3Na + A1(0H)3 + 3NaOH ;^ 3H2O + NasAlOa ^:± 3H20+3Na++A103= The extreme weakness of aluminium hydroxide, considered as an acid, accounts for the far-reaching decomposition when these aluminates pass into aqueous solution. Alutniniutn Oxide and Hydroxide. — ^Aluminium oxide occurs in nature in the crystaUine form (corundum). When it is tinted with a trace of a chromium compound, the ruby is formed, with a trace of cobalt the sapphire. Artificial rubies are prepared by igniting powdered alumina containing a trace of chromium oxide. Emery powder is an impure grade of corundum, containing ferric oxide. Although freshly precipitated aluminium hydroxide is freely soluble in acids, the fused oxide is very unreactive. Acids have practically no action upon it. It can be brought into solution 672 AN INORGANIC CHEMISTRY by fusion with potassium bisulphate or hydroxide, forming in the one case aluminium sulphate, in the other potassium aluminate. Aluminates. — ^A number of naturally occurring aluminates are known (spinel ruby, MgAlaOj ; gahnite, ZnAljOi ; iron spinel, FeAlaOi, etc.). These are salts derived from meta- aluminic acid : H3AIO3 — H20->HA102 Ortho-aluminic acid. Meta-alviminic ticid. The formation of alvuninates by the action of a strong base upon aluminium hydroxide has already been discussed. This reaction has become of ever increasing importance, as it affords the industrial chemist a means of preparing from impure bauxite aluminium hydroxide of sufficient purity for use in the electro- lytic manufacture of aluminium. In the dry process of purifying bauxite the mineral is ground, and roasted with sodium carbonate in a reverberatory furnace to sintering. The sodium aluminate is rapidly leached out with water, the iron oxide and silica being left behind : Al20(OH)i + 3Na2C03->2Na3A103+3CO,+2H20. The hot solution of sodium aluminate is then treated with carbon dioxide, the hj'droxide separating out : 2Na3A103+3C02 + 3H20-^2Al(OH)3i +3Na2C03. An improvement in the method of precipitating the aluminium hydroxide was effected by Bayer, who secured the desired effect merely by stirring in a little aluminium hydroxide. Nearly all the alumina separates out in a very pure form. Apparently, the sodium aluminate is largely hydrolysed into colloidal aluminium hydroxide, and this coagulates when nuclei are stirred into the solution. Bayer has also put forward an economical wet process. The bauxite is digested under pressure with concentrated sodium hydroxide and the meta-aluminate passes into solution : Al^OCOH)! + 2NaOH-> 2NaA102 + SH^O. The filtered solution is then vigorously stirred and aluminium hydroxide separates from the solution. The whole of the sodium hydroxide is then recovered. Aluminium hydroxide is often precipitated upon the fibre ALUMINIUM 573 of certain fabrics before exposTire to the action of the dyestufl. The aluminium hydroxide absorbed upon the surface of the fibre forms a loose chemical compound with the dyestuff. The actual colour taken on by the fibre depends not only upon the dyestuff chosen, but also upon the particular mordant precipitated upon the fibre to biad the dye. Other such mordants are the hydroxides of tin, iron and chromium. Halides. — Hydrated aluminium chloride, AlCl3,6H20 can be prepared by the crystallisation of a solution formed by saturating hydrochloric acid with aluminium hydroxide. Any attempt to dehydrate this hydrate results in a residue of aluminium oxide being left. The hydrolysis may, however, be effectively prevented by carrying out the dehydration in an atmosphere of hydrogen chloride. Other methods of preparing anhydrous aluminium chloride are based upon the equations : AI2O3 + 30 + 3a2-> 2AICI3 + 300 2Al + 30l2-^2AlCl3 4AI3O3 + 3S2OI2 + 9Cl2-> 8AICI3 + 6SO2 The sulphur chloride and chlorine may be replaced by carbon tetrachloride. Anhydrous aluminium chloride is a white, hygroscopic soHd, which sublimes readily. On exposure to air, it fumes freely, owing to the aqueous vapour of the atmosphere causing hydrolysis : AICI3 + 3HOH->Al(OH)3 +3HC1. Aluminium chloride is soluble in alcohol, ether and other organic solvents. Quite a number of double hahde salts have been prepared from aluminium chloride, e.g. AlCls.NHiCl ; A1C13,KC1 ; AlCl3,NaCl ; 2A1C13,3KC1 ; 2AlCl3,3NaCl ; AlCl3,3Na01 ; A1C13,3KC1, etc. Aluminium bromide and iodide can be prepared by means similar to those adopted for the chloride. Their properties resemble those of the chloride. Aluminium fluoride is best prepared by evaporating a solution of aluminium in hydrofluoric acid and subhming the residue in a carbon tube in an atmosphere of hydrogen. It is insoluble in water and forms small cubes. Hydrated fluorides of aluminium can be obtained from the solution of aluminium fluoride in hydrofluoric acid. Cryolite is a double sodium aluminium fluoride, NagAlFe, or AlF3,3NaF, sinylar to the double chlorides enumerated above. 574 AN INORGANIC CHEMISTRY Aluminium Sulphate.— The sulphate is prepared from bauxite or from kaolin (china clay). For pure samples of aluminium sulphate the bauxite is first converted into pure aluminium hydroxide by the methods described above, and the precipitated aluminium hydroxide is then treated with sulphuric acid. Commercial aluminium sulphate is, however, obtained by heating the naturally occurring bauxite with sulphuric acid. The product contains appreciable quantities of iron, and is only fit for such commercial operations as the sizing of paper and for the purification of sewage, etc. In manufacturing aluminium sulphate from kaolin, the calcined clay is treated with sulphuric acid, and after some time a mass known as alum cake separates : H2Al2(SiOi)2 + 3H,S04-> Al2(S04)3 + 2H,Si03 + 2H,0. The alum cake is lixiviated with water and the clear liquid evaporated tiU crystallisation sets in. Pure aluminium sulphate is used extensively as a mordant, the actual mordant being aluminium hydroxide which is precipitated upon the fibre. Basic sulphates of aluminium are known. Potassium Aluminium Sulphate — Alums. — If a strong solution of aluminium sulphate be treated with potassium sulphate, well defined octahedral crystals of potassium aluminium sulphate crystallise from the solution. These have the com- position, K2S04,Al2(S04)3,24H20. Quite a number of similarly constituted salts have been prepared ; these are named the alums. The general formula, M2'S04,M2'"(S04)3,24H20, appHes to them all. M' denotes a monovalent metal such as K, Na, Rb, Cs, NH4, whilst M™ denotes a trivalent metal, Fe, Mn, Al, Cr, Ga, In. These salts, as well as the corresponding selenates, are all isomorphous. In aqueous solutions of the alums, the double salts are largely decomposed into the single salts and their ions, K2S04,Al2(S04)3,24H20 ;=± K2SO4 + Al2(S04)3 + 2m^0, but at the same time it must not be overlooked that these double salts give definite evidence of complex formation, i.e. the equUibrium Y.lAU^O,),M'R.O] ^± K2S04,Al2(S04)3,24H20 ^r±, K2S04+Al2(S04)3 + 24H02 ALUMINIUM 575 exists, but the equilibrium in this case lies far over to the right. As already stated, the difference between double and complex salts is merely one of degree, rather than of land. Potassium aluminium sulphate is largely made from alunite, a basic potash alum. This is heated, when decomposition occurs, 2KA1 3(0H),(S0 4) 2 ^ K 2SO4, Al ,(S0 J 3 + 4A1(0H) 3. The calcined mass is then extracted with water. The con- centrated liquor is allowed to crystallise in wooden vessels. These crystals generally contain very small quantities of iron oxide, hence their faint colour. An improvement on this process is to treat the calcined alunite with sulphuric acid and the solution is then allowed to crystallise. Alum first separates from the solution, and it the mother liquor is further concentrated, the soluble aluminium sulphate crystallises. Potash alum is also obtained by treating bauxite with sulphuric acid, aluminium sulphate being thereby formed. To the hot solution containing this salt, there is added a strong solution of potassium sulphate. On cooling, potassium aluminium sulphate separates as a fine mass. Potash alum is used as a mordant, in the dressing of skins, in sizing paper and in making fire-proof materials. Its solutions react acid to Utmus. Potash alum is very much more soluble in hot water than in cold (at 0° 3-9 gm. dissolve per 100 c.c. ; at 100°, 357-3 grams dissolve). Nitride. — This compound is produced either by direct syn- thesis at 800-1,000°, or by heating together a mixture of alumina, carbon and nitrogen. AI2O3 + 3C +Na-^2A1N + 3C0. It is a grey amorphous powder which is decomposed by water, yielding aluminium hydroxide and ammonia. Aluminium Carbide. — Aluminium carbide is obtained by strongly heating a mixture of aluminium powder and carbon in a furnace to a temperature from 900-1,200°. At temperatures above 1,400° it dissociates into its elements. Aluminium carbide is slowly decomposed by water, forming aluminium hydroxide and methane : AI4C3 + 12H,0 ^4A1(0H)3 + 3CH4. Aluminium Sulphide.— This compound can be prepared by 576 AN INORGANIC CHEMISTRY heating alumina to redness in a stream of carbon disulphide, or in accordance with the following equations : Al^Oj + SH^S-^Al^Ss + SH.O AI2O3 + 3C + 3S -> AI2S3 + SCO^. Owing to the far-reaching hydrolysis which it undergoes, it cannot be prepared in the wet way. AI2S3 + 6HOH^^2Al(OH)3 s^ + SH^S. Aluminium Acetate. — Aluminium acetate, prepared by treating aluminium sulphate with lead acetate, 3Pb(C,H302)2 + Al,(SOi)3-^2Al(C,H302)3 + SPbSO^ i or by the action of acetic acid upon aluminium hydroxide, is used freely as a mordant, and in colour-printing upon cloth. This salt, being constituted from a weak acid and a very weak base, is almost completely hydrolysed by hot water, hence its appUcation as a mordant. AHC^H 30^)3 + 3H0H ;=z± A1(0H)3 + 3CH3COOH. Earthenware, Pottery, Porcelain. — Felspathic rocks (granite, etc.) are slowly attacked under the combined action of water and carbon dioxide. The potash is largely removed in the dissolved state and the residual aluminium sUicate, known as clay, is left. Kaolin is very pure clay, H2Al2(Si04)2,H20. Moistened clay has the property not only of retaining its shape for some time, but, m lien heated carefuUy, sinters into a hard, porous mass without undergoing any change of shape. The use of clay in the manufacture of bricks, tiles, pottery, porcelain, etc., depends upon this. Poorer quaUties of clay, containing iron and other impurities, are used in the manufacture of bricks, tUes and earthenware. The redness of the fired product is due to the oxidation of the iron to the ferric state. Finer varieties of pottery are made from a high grade clay, together with quartz and felspar. The material is moulded into the desired shape and fired in kilns, wherein a temperature of a little over 1,000° can be maintained. An englobe or slip is frequently introduced on to the surface of the pottery, especially where the colour of the latter is somewhat undesirable. The slip consists of a pure clay mixed with some of the constituents of the glaze. The sUp is placed between the pottery and the glaze, its sole purpose being to give a finer GALLIUM, INDIUM, THALLIUM 577 surface to which the glaze may adhere. The glaze itself may be a salt glaze, introduced by throwing the salt into the furnace, or a mixture of fusible Cornish stone and felspar. If it is desired to colour the pottery, suitable metallic oxides, or glasses prepared from them, are either introduced into the glaze or into the sUp. Porcelain, the finest kind of pottery, is translucent and almost glassy. The heating has been carried nearly to the point of fusion; moreover, the constituents employed are of the finest quality. Ultramarine is prepared by fusing together kaolin, sodium carbonate, sulphur and charcoal in the absence of air. The resultant mass, which is of a green colour, is washed with water, dried, mixed with sulphur and heated in the air until the desired shade of colour is acquired. The constitution of ultramarine is uncertain, but it is considered to be a complex sodium aluminium sihcate, with some sodium sulphide. Hydrochloric acid liberates hydrogen sulphide from it. Ultramarine is employed in removing the yellow tint from sugar, paper, linen and cotton goods, as well as for laundry purposes. It is also used as a pigment. Gallium, Indium, Thallium In passing from the short period elements, boron and aluminium, to the long period elements, gallium, indium and thallium, there is a fairly regular change in the basicity of the oxide. Boron trioxide, we have already seen, is amphoteric, though its basic properties are less strongly developed than are its acidic ; aluminium oxide is amphoteric ; the oxides (hydroxides) of gallium and indium are also amphoteric, whUe that of thallium is weakly basic. The gradual strengthening in the basic property of the group oxide in passing from boron to thallium is in agreement with that shown by the oxides of other groups, e.g. the nitrogen group. Gallium forms two types of salts, the divalent from GaO and the trivalent from Ga 2O 3. The latter are strongly hydrolysed . Gallons salts tend to break down in accordance with the equation, 3Ga++;rzi2Ga+ + ++Ga. Gallates are also known. Well defined alums, isomorphous with the alums of aluminium, iron, etc., have been isolated. pp 578 AN INORGANIC CHEMISTRY Indium forms three types of salts, the monovalent from InjO, divalent from InO and the trivalent from InjOj. Salts of the two lower types undergo auto-oxidation in aqueous solution in accordance with the equations : 3InCl^=i2In+InCl3 SlnCla ^=i± In + 2InCl3. The trivalent salts are strongly hydrolysed in aqueous solution. Well defined alums, isomorphous with potassium aluminium sulphate, have been prepared. Indates are also known : In(0H)3 + 3K0H ^=± K3ln03 + SH^O. Both gaUium and indium are found in zinc blende. Thallium is often found in the blendes (sulphides) of iron, zinc, and also associated with tin and copper selenide. It is recovered from the chamber mud and flue dust from vitriol factories where thalliferous pyrites has been used. It was discovered by Crookes in such deposits by the aid of the spectro- scope. Thallium gives two oxides of which the lower, TI2O, is strongly basic. The salts, TlGl, TI2SO4, etc., obtained from this oxide strongly resemble the corresponding salts of silver, even to the insolubility of the haUdes. It is interesting, too, that thaUous chloride darkens under the action of Hght, possibly due to a similar photo-chemical reduction to that which takes place when silver chloride is exposed to the action of light. The sesqui-oxide, TI2O3, gives a series of salts, e.g. TICI3, Tl2(S04)3, in which thaUium is trivalent. The weakly basic nature of TI2O3, compared with TI2O, accounts for the strong hydrolysis which the trivalent salts undergo. ThaUic hydroxide is insoluble in alkali hydroxides. Although thallic sulphate forms double sulphates with the alkali sulphates, these are not isomorphous with the alums. A number of thaUous-thallic compounds are also known, e.g. TlCls.TlCl. Questions 1. Discuss the preparation and commercial importance of sodium aluminate. 2. Give a comparative account of the carbonic acid salts of the alkali and alkaline earth elements. 3. Discuss the amphoteric behaviour of aliuninium hydroxide and show how the properties of the aluminiimi compounds are intimately connected with the properties of the oxide. GALLIUM, INDIUM, THALLIUM 579 4. Give an account of the commercial preparation, the constitution and properties of borax. 5. Why is ammonium hydroxide capable of dissolving zinc hydroxide, but unable to dissolve aluminium hydroxide ? 6. What is an alum ? What elements are capable of forming alums, and what does this capacity to yield alums signify ? 7. Give an account of the modern methods of purifying bauxite. CHAPTER XXXVI GROUP 4 : GERMANIUM, TIN, LEAD, TITANIUM, ZIRCONIUM, CERIUM, THORIUM The fourth group of elements is shown in the following scheme : C Of these, the elements carbon and silicon have already been treated in detail, it remains now to see how far Si, the properties of these elements and their compounds Ti \ persist through the other members of the group, and Ge how far these properties are modified by the increas- Zr I ing atomic weight of the element. Sn The formula of the typical group oxide is MOj, of Ce I the chloride MCI4, of the hydroxide M(OH)i. All I — the elements of the group form compounds of this Th ! type, though the occurrence of these compounds does Pb not justify the conclusion that no other oxides, chlorides, etc., exist. All that the Periodic Law demands is that the elements in each group shall exhibit certain fundamental similarities which justify their inclusion in one group. The appended table throws into prominence the group characteristics of these elements. TABLE 54 Group 4B Element At. wt. . Sp. gr. . Melt. pt. . Oxide Ortho-acid Meta-acid Chloride . Boil. pt. . Carbon. 12 2 -3-3 -5 CO2 weakly acidic [C(OH),] H2CO3 carbonic CCl, 76° Silicon. 28-3 2-35 1,500° (about) SiOa weakly acidic Si(OH)4 H^SiOj silicic SiClj 59° German- Tin. Lead. ium. 725 119 307 1 5-47 5 -8-7 -3 11-4 900° 231° 326° GeOj SnOj PbO^ ampho- ampho- ampho- teric teric teric Ge(OH)4 Sn(0H)4 Pb(OH)j H^GeOs HaSnOj HjPbOa germanic stannic plumbic GeCl. SnCl. Pbaj 87° 114° decom- poses 580 (GtERMANIUM, TIN, LEAD Grotip 4A 581 Element . Titanium. Zirconium. Cerium. Thorium. Oxide . . TiOa ZrOa CeOa ThOj amphoteric amphoteric basic basic Hydroxide . Ti(OH), Zr(OH)i Ce(0H)4 Th(0H-)4 Chloride . TiCli ZrCl^ CeOlj ThCl^ Group 4B Germanium This element forms two oxides, GeO and GeOa, both of which are amphoteric. Both oxides give rise to chlorides, fluorides, etc., whilst the hydroxides are soluble in acid and alkali. Germanium has many properties which are suggestive of carbon and silicon, e.g. it forms a weU-defLned germanium chloroform, GeHClg (cf. CHCI3, SiHCls) ; by the hydrolysis of this compound ger- mano-formic acid is formed. GeHCla +3NaOH cf. CHCI3 +3NaOH - >3NaCI +HGeOOH +H2O ■ SNaCl +HCOOH +H2O Tin, Lead General Remarks. — The facts summarised in contain the key to the chemistry of these elements. TABLE 56 Oxides Oxides Table 55 00 SnO CO 2 Si02 SnO 2 stannous stannites salts SnCla SnSOi NajSnOj stannic stannates salts SnClj Sn(S0i)2 NajSnOa Pb^O PbO- PbgO, Pb^Os Pb02 sub- plumbous salts Pbl PbCl plumbous plumbites salts PbS04 PbCl 2 NaaPbOa no salts plumbic plumbates salts Pb(S04)2 Pb{Acet)4 NaaPbOg — ^-> II ^•§ as Increase in acidic nature of dxide. Decrease in basic nature of oxide. 582 AN INORGANIC CHEMISTRY Of the two oxides SnO, PbO, it follows that PbO is the stronger base, but SnO will be the stronger acid-forming oxide, hence salts derived from PbO functioning as an acid, e.g. PbCla, PbSOi, win be more stable and less hydrolysed than corresponding salts formed from SnO, e.g. SnCla, SnSOj. On the other hand the fact that stannous oxide, SnO, is a stronger acidic oxide thaa plumb- ous oxide, PbO, accounts for the fact that when these oxides function as bases, salts formed from the oxide, SnO, will be more stable and less hydrolysed than salts formed from the oxide PbO, i.e. the stannites will be more stable than the plumbites. So , too, the stannates are more stable than the plumbates. So far as the relative strength of the salts formed from the different oxides of the same element are concerned, in accordance with the general rule that the higher oxide is the more acidic, the stannates are more stable than the stannites, the plumbates more stable than the plumbites. But where the oxides function as a base as in the compounds PbC^, Pb(S04)2, the lower valent salts, being derived from the stronger base, wiU always be found to be the more stable. Tin Occurrence and Extraction. — Tin occurs native in small quantities, but the commercial source of tin is cassiterite or tinstone, SnOa, which consists of tetragonal crystals coloured brown by iron and other impurities. This ore occurs in Corn- wall, Malay Peninsula, Austraha, East Indies, and Germany. The ore is first concentrated by washing away earthy impuri- ties from the crushed ore. It is then roasted to oxidise the sulphides of iron and to remove arsenic. The roast is washed, or treated electro-magnetically to separate it from the magnetic wolfram, a mineral of nearly equal density and with which it is often associated, and the tin oxide so obtained is reduced by heating with carbon (coal) in a reverberatory furnace. The molten tin is run off and cast into ingots. Further purification of this block tin is effected by heating it shghtly above its melting point so that it can run away from its associated dross. The molten metal is further stirred with a billet of wood when the more easily oxidisable impurities are removed in the form of oxide. Properties. — Tin resembles carbon in existing in three distinct allotropic modifications, rhombic, tetragonal and grey tin. The latter variety is prepared by cooling the crystalline TIN 583 modification well below 18°, when the crystals crumble into powder as if attacked with a disease. This phenomenon was first noticed during a very cold winter in Russia, when many tin organ pipes crumbled into powder. Physico-chemical investi- gation showed that this apparent disease arose from the conver- sion of the unstable crystalline modification into the stable powdery form. The equiUbria representing these changes are expressed in the equation : Rhombic tin ^ tetragonal tin ^ grey tin. 161° 18° Tin is a soft malleable metal, so ductile that it can be beaten into tin foU. It is in much demand for coating sheets of iron in order to prevent them from rusting, a process generally effected by dipping thin sheets of iron into molten tin. Copper cooking vessels are treated similarly. Many alloys are used commer- cially, e.g. bronze (tin, zinc, copper and lead), soft solder (tin, lead), pewter (tin, lead), Britannia metal (tin, antimony, copper). Tin is attacked by hydrochloric acid, forming stannous chloride, SnCl2 and hydrogen. Hot sulphuric acid forms stannous sulphate and sulphur dioxide. No doubt, the hydrogen at the moment of Uberation is oxidised by the fairly active oxidising agent H2SO4. Sn +H2S04->SnS04 +2H 2H + H2SO4 -^ SO2 + 2H2O Dilute nitric acid, when cool, dissolves tin with the formation of Stannous nitrate, the nascent hydrogen reducing some of the nitric acid to ammonia. 4Sn + 10HNO3-^4Sn(NO3)2 +3H2O +NH4NO3. Ordinary concentrated nitric acid produces stannic nitrate which is hydrolysed to meta-stannic acid. Sn + 4HN03->H2Sn03 + H2O + 4NO2. Tin dissolves slowly in a boiUng solution of sodium hydroxide. This reaction is most easily explained by the hypothesis that tin reacts with water in accordance with the equation : Sn + 2H20->Sn(OH)2 + H2. Owing to insolubihty of the hydroxide with the consequent for- mation of a protecting surface layer, the reaction in pure water stops at once, but in the presence of sodium hydroxide this 584 AN INORGANIC CHEMISTRY protecting film is instantly removed by the formation of the soluble stannite. Sn(OH), + 2NaOH->Na2Sn02 +2H2O. Oxides and Salts of Tin. Tin forms two oxides — stannous oxide, SnO, and stannic oxide, Sn02- From stannous oxide are derived : 1. Stannous salts in which tin behaves as a divalent cathion, Sn-CU. 2. Stannites in which the tin forms part of the acid group or anion, Na2Sn02. From stannic oxide are derived : 3. Stannic salts in which the tin functions as a quadrivalent cathion, Sn"(S04)2. 4. Stannates in which the tin forms part of the anion, — NajSnOa. Stannous Oxide and its Salts. — Stannous oxide is a black powder formed by the breaking down of stannous oxalate in the absence of air. SnCaOi -> SnO + CO + CO2. When heated in the air, it passes into stannic oxide. Stannous oxide, being an amphoteric oxide, dissolves in the alkaUne hydroxides to form stannites, and in strong acids to form normal metallic salts. SnO + 2HCl-^SnCl2 + H2O. Stannous chloride, SnClj, is the most important of the starmous salts. It is generally obtained by the action of tin upon hydro- chloric acid. Owing to the weakness of the base, SnO, stannous chloride undergoes considerable hydrolytic decomposition on solution in water, with the precipitation of an oxychloride. SnCls + H2O ^ Sn(OH)Cl -f + HCl. In accordance with the Law of Mass Action, this hydrolysis will be effectively prevented by the addition of hydrochloric acid. Stannous chloride is a strong reducing agent, passing into stannic chloride. Thus ferric chloride is reduced to the ferrous state 2FeCl3 + SnCla -^ SnCl^ + 2FeCl2 or, written ionicaUy, 2Fe+ + ++Sn+ + — >Sn'- + + + +2Fe+ + TIN 585 It will reduce sulphur dioxide to hydrogen sulphide (p. 233), mercuric chloride to mercurous chloride, and it in excess, to metallic mercury. 2HgCl2 + SnCl^-^HgaCU + SnCl, ngfiU + SnClj -> 2Hg + SnCU The stannites are most easily formed by the fusion of stannous oxide with the alkaline hydroxides, or by the solution of stannous hydroxide in an excess of the desired hydroxide. Thus, it sodium hydroxide is added to a solution of stannous chloride, a precipi- tate of stannous hydroxide is thrown down. If this is treated with a further supply of sodium hydroxide a solution of sodium stannite is formed. The stannites are unstable, and on heating, undergo auto -oxidation into tin and the more stable stannate (cf. nitrites, phosphites, arsenites). The stannites are strongly hydrolysed in water, forming the hydroxide. Stannic Oxide and its Salts. — Stannic oxide, SnOj, is the chief ore of tin. It can be obtained by igniting tin, or more readily by dehydrating meta-stannic acid (q.v.). When heated, it turns yeUow, but returns to its original colour on coohng. After ignition, stannic oxide is very unreactive, and is only acted upon by fused potassium and sodium hydroxides with the production of stannates. If prepared at a low temperature, stannic oxide will also dissolve freely in nitric and sulphuric acids, with the formation of stannic nitrate and sulphate re- spectively. Stannic oxide (hydroxide) is amphoteric in nature so that it forms two classes of salts, those in which tin functions as a normal quadrivalent cathion, and those in which the oxide forms part of the anion or acid group. SnO, +2NaOH^Na2Sn03 +H2O 2Na+ + SnOa" SnO^ + 2H2S04^Sn{S04)2 +2H2O Sn+ + -'-'-+2S04- Stannic chlwide is readily obtained by the action of chlorine upon tin or upon stannous chloride. It is a fuming hquid which boils at 114°. Various crystalUne hydrates, SnCli.SHaO ; 586 AN INORGANIC CHEMISTRY SnCli.SHsO ; SnCli.SHzO, have been isolated, of which the first has considerable commercial importance as a mordant in dyeiag. The slight conductivity of an aqueous solution of stannic chloride indi- cates that very f ew Sn + + "*■+ ions exist in the solution ; on standing, the conductivity slowly rises, due to the progressive hydrolysis SnCl^ + 4H0H -^ SnCOH), + 4HC1 4H+ +4C1-. The stannic hydroxide (stannic acid) does not separate out but remains in colloidal solution. Stannic chloride combines with the alkaU chlorides to form complex chlorides, 2NaCl, SnClj, or more correctly, Na2SnClG. These complex chlorides, of which many have been prepared, are derived from the acid, HjSnCle, which ionises into 2H"'' and SnCle". Ammonium stannic chloride (NH4)2SnCl6, which is used as a mordant in dyeing cotton, is an example of this type of compound. Other examples of the stannic salts are stannic bromide, SnBrj, stannic sulphate, Sn(S04)2, produced by the solution of stannic hydroxide in sulphuric acid, and the unstable stannic nitrate, Sn(N03)4, which is formed in variable quantities during the solution of tin in not too concentrated nitric acid. When ammonium hydroxide is added to stannic chloride, a ^\hite gelatinous precipitate is thrown down. If dried in the air, this has the composition Sn(0H)4, and represents ortho-stannic acid, but if dried over sulphuric acid, water is given up and meta-stannic acid H2Sn03, is left. In the moist state, stannic hydroxide, in accordance with its amphoteric nature, dissolves freely in nitric, sulphuric and hydrochloric acids as well as in sodium and potassium hydroxides. The stannates of the alkaU-metals are readily formed by fusing together stannic oxide and an alkali hydroxide. These are freely soluble and form stable salts. The other stannates are insoluble (cf. the carbonates and sihcates). Meta-stannic acid is of interest as it has long been recognised that it exists in two isomeric modifications. When produced by the hydrolysis of stannic chloride as above, or by the addition of hydrochloric acid to sodium stannate, NaaSnOa -f 2HC1^2NaCl -f HsSnOj it forms the so-called a-modification. The main characteristic of LEAD 587 a-meta-Btannic acid is its ready solubility in nitric, sulphuric and hydrochloric acid, and in sodium and potassium hydroxides. /3-Meta-starmic acid (HaSnOj)^ is formed by the action of nitric acid upon tin. It dissolves slowly on boiling with sodium hydroxide, forming a soluble sodium stannate NaaSngOn. On treatment with hydrochloric acid, it forms an oxychloride with properties distinct from those of the corresponding compound formed by the solution of the a-meta-stannic acid in hydrochloric acid. These characteristic reactions point to the possibility that the |8-acid, although of the same empirical composition, is yet a polymer of the a-acid. On fusing /S-meta-stannic acid with sodium hydroxide, the same a-stannate is obtained as from a-raeta-stannic acid. Stannous and Stannic Sulphides. — Stannous sulphide is produced either by the direct action of tinfoil upon sulphur vapour, or by the passage of hydrogen sulphide through a solution of stannous chloride. It is a dark brown powder. Stannic sulphide is a yellow compound obtained by the action of hydrogen sulphide upon a solution of stannic chloride, or by heating together tin amalgam, ammonium chloride and sulphur in a retort. Beautiful golden scales are left in the retort, and are used as a pigment under the name of " mosaic gold." Stannic sulphide dissolves freely in potassium hydroxide or in a solution of the sulphides of the alkaUes. SnSs + (NHJ^S-^ (NH4)2SnS3 Ammonium thiostannate. Advantage is taken of this reaction in separating tin sulphide along with the sulphides of arsenic and antimony (q.v.) from the sidphides of other metals. The tin is precipitated in the form of stannous sulphide which is but shghtly soluble in pure ammonium sulphide. In the presence of yellow ammonium sulphide, (NH4)2S+a;S, oxidation of stannous sulphide to stannic sulphide is at once effected by the excess of sulphur present, and the stannic sulphide passes into solution as the soluble thiosalt. Lead General Remarks. — Lead produces five well-defined oxides, of which two are " mixed oxides," and therefore incapable of forming salts, whilst the other three oxides give rise to well- defined salts. The steady change in the properties of the oxide 588 AN INORGANIC CHEMISTRY as the oxygen content changes affords an excellent illustration of the general statement that with increasing oxygen content the oxide of an element becomes more acidic. Thus: PbjO, Lead suboxide . . Distinctly basic. PbO, Lead monoxide . . Amphoteric, but the basic proper- ties predominate. Pb^Oa, Lead sesquioxide | Mixed oxides Pb304, Lead tetroxide . j PbOo, Lead dioxide . . . Amphoteric, but the acidic pro- perties are more strongly devel- oped than the basic. From lead suboxide a few salts such as Pbl, PbCl, Pb2S0j, etc., have been prepared. In these salts the lead is either mono- valent, or similar to mercury in the mercurous salts, i.e. Pb2" On the latter assumption the salts would have a formula such as Pb2l2, etc. From lead monoxide two types of salts are derived, in one of which lead functions as a normal divalent cathion, e.g. : PbCl2 ;=±.Pb •- '+2C1-, Pb(S04);z:±Pb+++S04^. These salts are formed from lead monoxide, acting as a base, by the action of the necessary acid. PbO + 2HCl-^PbCl2 4- HoO. The second type of salt formed from lead monoxide are known as the plumbites ; these are obtamed by the action of an alkaline hydroxide upon PbO acting as an acidic oxide, PbO -f 2NaOH->Na2PbO, -f H^O. These plumbites, owing to the weakness of the acidic oxide are very readily hydrolysed in aqueous solution, i.e. the action is reversible. From lead dioxide, two series of salts are obtained. From the oxide Pb02, in its basic capacity, such salts as PbCl^, Pb{S04)2, Pb AC4, are formed. These salts are unstable and tend to pass into the corresponding salt of divalent lead. Owing to the weak- ness of the base, Pb02.. such salts are completely hydrolysed on treatment with water. From Pb02, functioning as an acidic oxide, are derived the plumbates. PbOa +2NaOH-^Na2Pb03 -fHoO. LEAD 589 These are more stable than the plumbites, but the acidic power of the oxide is still so weak ihsd, appreciable hydrolysis of the plumbates occurs. Occurrence. — A small quantity of lead occurs native, but the most important of the ores of lead is undoubtedly galena, lead sulphide. This ore is generally associated with, silver sulphide. Fairly large deposits of cerussite (PbCOs) and rela- tively small deposits of anglesite (PbS04) occur. Method of Extraction. — Lead sulphide is roasted in a current of air, either in a blast or in a reverberatory furnace, when partial oxidation to lead monoxide and sulphate is effected. 2PbS -f 3O2 ^ 2PbO + 2SO2 PbS+202-^PbS04 The air is then shut off and the temperature of the furnace raised. Interaction between the sulphide, sulphate and oxide occurs, in accordance with the equation : PbS + PbSOi -> 2Pb + 2SO2 PbS + 2PbO -^ 3Pb + SO2. When the ore is impure, reduction is effected by means of metallic iron, or a mixture of iron ore and coke. PbS+Fe^FeS+Pb. The presence of impurities such as antimony, tin, copper, etc., causes the lead to be brittle. These are removed by heating the metal in a shallow reverberatory furnace. Most of the admixed impurities oxidise more readily than does lead, and rise to the surface as a scum. The silver, however, requires special treat- ment, and is removed either by the Pattinson's Process or the Parke's Process {q.v.). Lead is also purified electroljrtically. Properties.— Lead is a soft metal which, when freshly cut, has a bright metalHc lustre, but on exposure the surface rapidly tarnishes, owing to the formation of an oxidation fihn. Its melting point is 326°. Because of its unreactive nature, it is employed largely in industry for Uning reaction chambers, etc., e.g. in the sulphm-ic acid plants. The crystalhne nature of the metal is shown in the electrolysis of a solution of a lead salt, or when a piece of iron is put in a solution of such a salt as lead acetate, when the well known lead tree results. Lead salts have a powerful and deleterious physiological action. 590 AN INOEGANIC CHEMISTRY Pure water is without action upon lead, but in the presence of air oxidation ensues, with the formation of sUghtly soluble lead hydroxide. In the presence of carbon dioxide, a basic carbonate is thrown down. Hard water, containing dissolved carbonates and sulphates, can be safely transported through leaden pipes, as these soon become covered with a layer of insoluble sulphate and carbonate. Water, free from dissolved sulphates and carbonates, should be carried through leaden pipes only after filtration through beds of hmestone, when the water carries into solution sufficient calcium carbonate to form the desired protective layer of lead carbonate inside the pipes. Nitric acid attacks lead freely, forming lead nitrate and the oxides of nitrogen. Dilute sulphxiric and hydrochloric acids have but sUght action, as the surface soon becomes covered with an insoluble layer of sulphate or chloride, and the action ceases. With hot concentrated acid, rapid solution occurs, owing to the solubility of the chloride (sulphate) in the concentrated acid. Sodium hydroxide dissolves lead very slowly on boiling. This reaction is probably twofold : Pb+2H20^Pb(OH)2+H2 Pb(0H)2 + 2]SraOH^-Na2Pb02 + 2H2O. Solution proceeds as rapidly as the lead hydroxide is removed. Lead Suboxide and its Salts. — Lead svhoxide, Pb20, is formed by decomposing lead oxalate under such conditions that the evolved carbon monoxide does not exert any reducing action. 2PbC204^Pb20 +C0 + 3CO2. This is effected either by carrying out the decomposition in a stream of nitrogen, or better still, by removing the gaseous products of the reaction by means of a suitable pump. Lead suboxide is a greyish-black powder, which breaks down into metal and lead monoxide on treatment with acid or alkaU. The lead monoxide then reacts to form a lead salt. The ten- dency of lead suboxide to decompose thus : PbsO-^^Pb-f PbO, has rendered futile all attempts to obtain salts of this oxide by the direct action of an acid. Under the action of methyl iodide vapour at 260° the suboxide reacts to form a bright yellow LEAD 591 powder — lead subiodide, Pbl. On treatment with dilute acetic acid, decomposition at once occurs with the separation of lead. 2PbI->Pb+Pbl2. Other salts such as PbBr, PbCl, PbaSOj, have been prepared by similar means. It is interesting to note that the solubility of these salts is appreciably less than that of the normal divalent salts. So, too, it will be remembered that cuprous chloride, mercurous chloride, and aurous chloride are all less soluble than their corresponding higher chlorides, e.g. cupric, mercuric and auric chlorides. Lead Monoxide and its Salts. — Lead monoxide {Litharge), PbO, is a buff-coloured powder formed when lead or any other oxide of lead is strongly heated in the air. In the laboratory it is readily obtained by heating the nitrate or carbonate. It is known commercially as massicot and as htharge. Considerable quantities are used technically in the manufacture of varnishes, flint glass, etc. Lead monoxide dissolves sUghtly in water, forming lead hydroxide, Pb(0H)2, which gives an alkaUne reaction. This hy- droxide, which is more readily obtained by the action of ammo- nium hydroxide upon a solution of lead nitrate, dissolves freely in acids to form the divalent salts of lead. Pb(0H)2 +2HCl^PbCl2 +2H2O and also in the hydroxides of sodium and potassium, but not of ammonium. This is due to the formation of soluble plumbites. Pb{0H)2 + 2NaOH ^=± Na2Pb02 + 2H2O. Plumbites of the alkahne earths are also known. The plumbites are all strongly hydrolysed, in accordance with the rule that salts formed from a weak base or acid or both will undergo hydrolytic decomposition. In the case of ammonium plumbite, both acid and base are so weak that the salt is completely decomposed by water, i.e. ammonium hydroxide cannot exert any solvent action upon lead hydroxide. Lead Chloride, Bromide a.nd Iodide {Plumbous Chloride, etc.) are all relatively insoluble in cold water, the solubility decreasing as the atomic weight of the halogen increases. The salts are much more poluble in hot water. They are formed by the action of potassium chloride, etc., upon an aqueous solution of lead nitrate or acetate. In the presence of a considerable excess of the 592 AN INORGANIC CHEMISTRY precipitant, re-solution of the lead halide occurs. Thus lead iodide dissolves in excess of potassium iodide to form the soluble potassium lead iodide. Pbl2+2KI->K2Pbl4. Lead nitrate is formed by dissolving lead or lead monoxide in nitric acid. The salt is deposited from this solution in regular octa- hedral crystals. It is one of the most soluble of the lead salts, 50 gm. dissolving in 100 gm. of water at the ordinary tempera- ture. On heating lead nitrate, oxygen, nitrogen peroxide and lead monoxide are produced (see p. 296). Lead carbonate is formed by adding ammonium bi-carbonate to a solution of lead nitrate. Ammonium carbonate precipitates a basic carbonate of lead the exact composition of which varies with the temperature and concentration of the solution. The most important of the basic carbonates of lead is white lead, 2PbC03,Pb(OH)2 — which is in great demand industrially as a pig- ment. Numerous processes have been patented for the manu- facture of this pigment, the best of which is undoubtedly the old Dutch method . Gratings of lead are placed above dilute acetic acid contained in pots , the whole being buried in manure. The decom- position of this furnishes the heat and the carbon dioxide necessary for the reaction. Under the action of the air lead hydroxide is formed, and this reacts with the acid to form basic lead acetate. This is slowly decomposed by the carbon dioxide generated by the decomposing manure with the formation of basic lead carbonate and acetic acid. In this cycle of operations there is theoretically no loss of acetic acid, so that a relatively small amount of acetic acid can convert a very large amount of lead into the basic carbonate. White lead is especially valuable owing to its great covering power. Lead svlphate, which occurs naturally as anglesite, forming rhombic crystals isomorphous with heavy spar, can be preci- pitated as a white powder, insoluble in water and dilute sulphuric acid, but appreciably soluble in nitric and strong sulphiiric acids. It is also soluble in potassium hydroxide, ammonium acetate or tartrate. The addition of potassium hydroxide upsets the equilibrium PbSOi ^=± PbS04 ^zz± Pb + + + SO4 - Solid Dissolved LEAD 593 by converting some of the Pb' + ions into plumbite ions, PbOa" in accordance with the equation Pb++ +40H- ^zz^PbOr +2H2O and in order to restore the equUibrium more lead sulphate must dissolve. Lead sulphide occurs naturally as galena, but it can also be precipitated as a black powder by the action of hydrogen sulphide upon a solution containing a lead salt. PbCla + H2S -> PbS + 2HC1. Lead sulphide is attacked by dilute nitric acid, and converted into lead nitrate and sulphur, whilst the concentrated acid oxidises it into lead sulphate. Lead Dioxide and its Salts. — Lead dioxide is prepared by the action of suitably strong oxidising agents upon an alkahne solution of lead hydroxide ; as oxidising media sodium hypo- chlorite, chlorine and bromine find frequent use. NajPbOa + CaCl(ClO) + HjO-^ 2NaOH + PbO^ + CaCU. It is also deposited at the anode when a solution of a lead salt is electrolyticaUy decomposed. The dioxide is a powerful oxidising agent. Pb02+S02->PbS04 Pb02+4HCl-^PbCl3+Cl2+2H20. In the case of hydrogen sulphide the reaction is so vigorous that when the gas impinges upon lead dioxide, it bursts into flame. Owing to the fact that acids do not liberate hydrogen peroxide from it, lead dioxide is considered to be graphically represented by the formula Pb<|^ and not by the chain formula Pb 2Ca2Pb04 + 4CO2. The process is reversible, i.e. under the action of carbon dioxide the ortho-plumbate decomposes with the evolution of oxygen. This equation summarises the chemistry of the Kastner method of manufacturing oxygen from the air. Lead Dioxide and Accumulators (Storage Battery). — If two lead plates coated with lead monoxide are put into dilute sulphuric acid, a deposit of lead sulphate appears on each plate. PbO+H2S04^PbS04^^ +H2O. If an electric current is now passed through the ceU, the hydrogen hberated at the cathode reduces the lead sulphate to spongy lead, whilst at the anode oxidation of the sulphate to disulphate occurs, and this is at once hydrolysed to lead dioxide and sulphuric acid. 2PbS04 + 2H2SO4 + 02-^2Pb(S04)2 + 2H2O Pb(S04), +2H20^^Pb02 -f 2H2SO4. The lead dioxide forms a dark coherent layer on the surface of the anode. The cell, Pb I H2SO4 aq. 1 PbO, is now charged and current may be obtained from it. During the discharge of the cell the processes at the electrodes are represented by the equations, Pb -f SO4- -^ PbS04 -f 20 Pb02+4H++S04=->PbS04 + 2H20-f 2© i.e. during the discharge the lead plate is maintained at a nega- tive potential, the lead dioxide plate at a positive potential, and the current flows in the direction of the arrow within the cell. Pb I H2SO4 aq. 1 Pb02 -> When the cell is completely discharged, both plates are again LEAD 695 covered with lead sulphate, and on sending an electric current through the cell in the opposite direction, the lead/lead dioxide plates may be again reformed, i.e. the cell is recharged. The whole process is epitomised in the equations : Pb + PbOj + 2H2SO4 = 2PbS0i + 2H2O Pb I 2H2SO4 Pb02=PbS04 2H2O PbSOi +a;H2S04 discharge charge Such a cell is known as an accumulator, and a series of such cells constitutes a storage battery. It evidently serves as a convenient storehouse of electricity. The main problem is to secure a high capacity in the cell so that it wiU be able to give a high current for a lengthy period. This is efiected by preparing the plates of leaden gratings, the openings of which are filled, after charging, with spongy lead and lead dioxide respectively ; whilst the maximum surface is secured by bringing the plates as close together as is possible. Lead Tetrachloride {Plumbic Chloride) is prepared by passing chlorine through lead chloride suspended in hydrochloric acid. On the addition of ammonium chloride, ammonium plumbic chloride, 2NH4Cl,PbCl4, separates out. When this is acted upon by strong sulphuric acid, lead tetrachloride is obtained as a heavy yellow oil. At— 15° it freezes. With a httle hydrochloric acid,' it forms a yellow crystalline substance, probably HaPbClg. On the addition of a httle water in the cold, it forms a hydrate which slowly evolves chlorine ; in the presence of much water it hydrolyses, forming lead dioxide and hydrochloric acid. PbCl4 + 2H2O -> PbO^ + 4HC1. Numerous double or complex salts are formed, e.g. KaPbCU, which are comparatively stable. Other salts in which lead dioxide functions as a base, are lead disulphate, Pb(S04)2, lead tetracetate, Pb{C2H302)4, lead dihydrogen phosphate, Pb(H2P04)4. Of these the phosphate is the most stable, the other salts hydrolysing completely on the addition of water. 596 AN INORGANIC CHEMISTRY Lead Sesquioxide and Lead TetroKide.— Lead sesquioxide, PbjOs, is obtained as an orange- coloured precipitate when sodium plumbite in alkaline solution is treated with oxidising agents, e.g. hypochlorite or bromine. On heating it evolves oxygen, and forms the monoxide. Nitric acid decomposes it with the formation of lead nitrate and lead dioxide, whilst hydrochloric acid forms lead chloride and chlorine. Hence lead sesquioxide behaves as if it were lead meta-plumbate, PbOjPbOa. On this hypothesis, the action of nitric acid is first to decompose this salt into its oxides, and then the monoxide reacts with the formation of lead nitrate. In the case of hydrochloric acid, the lead monoxide forms the chloride, whilst the dioxide produces lead chloride and chlorine. Lead Tetroxide (minium, red lead), PbgO,, is formed by heating lead monoxide in air to a temperature of 400°-500°. The reason of this is evident from a consideration of the dissociation pressures of this oxide. 2Pb30,^6PbO+02 Temperatm-e . . 445° 500° 555' 636" Pressure in mm. .5 60 183 763 The partial pressure of oxygen in air is 150 mm., so that at a temperature somewhat below 555° the dissociation pressure of the tetroxide exceeds the partial pressure of the oxygen in the air, i.e. the tetroxide wiU decompose. It can only be formed in air at temperatures below this. Lead tetroxide behaves very similarly to the sesquioxide on treatment with an acid. Thus dilute nitric acid forms lead nitrate and lead dioxide, which is thrown down as a precipitate. A somewhat similar structure is therefore attributed to it as to the sesquioxide, and it is considered to be lead orthoplumbate, PbsPbO^ or 2Pb0,Pb02. Red lead is used extensively in the manufacture of flint glass and in pottery glasses. It is also used largely as a paint. Group 4A Titanium, Zieconium, Cbeium, Thorium These elements aU form the same group oxide MO 2, and the basicity of this oxide increases from titanium to thorium. Titanium and zirconium dioxides are both amphoteric and the corresponding oxides of cerium and thorium are basic. This TITANIXBi GROUP 597" is in strict accordance with the general behaviour exhibited by Group 4b and by other groups. Titanium is widely distributed in nature as brookite, anatase, and rutUe, trimorphic varieties of TiOa, also as ilmenite, FeTiOa. The metal is beginning to be used in the preparation of certain steels. The oxides TiO, TiaOa, TiOa all exist, and give rise respectively to the divalent, trivalent and quadrivalent salts, whilst the dioxide also forms titanates. On oxidising a solution of a titanium salt with hydrogen peroxide, a yeUow solution is produced, which is supposed to arise from the formation of titanium trioxide, TiOa, in solution. Titanium combines readily with nitrogen to form a nitride, TiaNj. Zirconium occurs as a sihcate, ZrSiOi. It forms only two oxides, Zr203 and Zr02, of which the sesquioxide gives rise to no salts. From the dioxide a well defined series of quadrivalent salts is formed, as well as several zirconates. Cerium is comparatively rare, but is found in cerite and monazite. It forms the oxides, CoaOs and CeOa, both of which give rise to salts. The trivalent salts derived from Ce203 are the more stable. Thorium occurs in thorianite, monazite, orthite, etc. It forms but one series of salts derived from the basic oxide, ThOa- This oxide, commonly known as thoria, is in considerable demand in the incandescent mantle industry. Mantles, made of knitted silk or cotton, are dipped into a solution of thorium nitrate containing 1 per cent, of cerium nitrate, and then ignited. After the mantle has been " burnt," the oxides thoria and ceria are left, forming the framework of the mantle. The ceria seems to be the active agent in promoting the intense hght emission of these mantles. The action seems to be physical rather than chemical. The efficiency of the mantle is seriously affected if the composition of the thoria-ceria mixture departs from that specified above. Questions 1. What happens when hydrochloric acid acta upon (a) lead dioxide, (6) lead sesquioxide, (c) lead tetroxide, and what do these facts enable us to infer concerning the constitution of these oxides ? , , , ., 2. What is the action of sodium hydroxide upon (a) lead hydroxide, (6) lead dioxide, (c) ammonium chloride, (d) stannous chloride ? 598 AN INORGAISriC CHEMISTRY 3. Compare the stability of the stannites and plumbites, the stannates and plmnbates. 4. Tabulate the oxides of lead and indicate any types of compounds to which they give rise. 5. Show how the basicity of the oxides and hydroxides of lead slowly changes as the amount of oxygen in the oxides increases. Give other examples where this has been noted. 6. Compare the stannates and thio-stannates. 7. Compare the properties of the oxides of lead, especially in their behaviour towards acid and alkali. Discuss the action of nitric acid upon these oxides. 8. Compare the properties of the oxides of tin ; compare and contrast the properties of the salts derived from these oxides. 9. Give a comparative account of the tetrachlorides of Group 4. 10. Compare the properties of the various dioxides derived from the elements of Group 4. CHAPTER XXXVII GROUP 6A: CHROMIUM, MOLYBDENUM, TUNGSTEN, URANIUM Group 6 of the Periodic Table comprises the elements : — Of these, the subgroup B— oxygen, sulphur, [ selenium, tellurium — has already been studied. S The elements of the A subgroup show a general ^^ similarity not only to each other, but also to Cr ^ the elements of the B subgroup. They all form I Se trioxides, CrOj, M0O3, WO3, UO3, and these Mo I oxides, hke the trioxides SOa, SeOa, TeOj, are I Te acidic. From these oxides, acids similarly con- W stituted to sulphuric acid are formed, of the I general formula H2MO4. Corresponding salts U of these acids are generally isomorphous, e.g. K^CrOi and K2SO4. Chromium Occurrence and Metallurgy. — Chromium occurs in chrome iron ore or chromite, FeO,Cr203, as crocoisite, PbCrOj, and as chrome ochre, Cr203. Large quantities of chromium are prepared by the Goldschmidt method. CraOa + 2A1 -> 2Cr + AI2O3 . The reduction is carried out in a large refractory clay crucible. An intimate mixture of chrome oxide with slightly less than the theoretical quantity of aluminium is placed in the crucible, and covered with a layer of barium peroxide and aluminium. This is fired by means of a piece of magnesium ribbon. The reduced chromium is found under the slag of fused alumina. 599 600 AN INORGANIC CHEMISTRY Chromium is used in the manufacture of certain steels, such as high speed tool steels, which contain a high percentage of chromium. AUoys of iron containing 50-60 per cent, of chro- mium are produced by the reduction of chromite in an electric furnace. Chrome steel is exceedingly hard and tough, hence its use for armour plating, safes, crushing mills, stainless cutlery, etc. Properties. — Chromium is a hard steel-grey metal with a melting point above that of platinum. It dissolves freely in hydrochloric acid, forming chromous chloride and hydrogen. In the presence of nitric acid it becomes unreactive or passive. Such passive chromium will no longer react with hydrochloric acid unless it is warmed. With concentrated sulphuric acid it forms a deep coloured solution, containing sulphates of chromium, and sulphur dioxide is hberated. Chemical Chakacteeistics of the Compounds of Cheomtum The somewhat complicated chemistry of this element arises from the fact that three oxides exist, all of which give rise to salts. Tabulating these beside the oxides of sulphur, we have : S Cr — CrO basic 502 acidic — S2O3 acidic CrjOa amphoteric 503 acidic CrOa acidic S2O7 acidic — Chromous oxide and hydroxide are distinctly basic, and give rise to a perfectly normal series of divalent salts, the properties of which differ but slightly from the properties of other similarly constituted salts, e.g. CuCla, PeCl2. In one point only do the chromous salts stand apart from many other such salts — ^the ease with which they are oxidised by the air or other oxidising agents. The chromous salts are therefore strong reducing agents. Prom the sesquioxide, functioning as a base, there is derived a series of trivalent salts, CrCls, Cr2{S04)3, etc. The sesquioxide with its higher oxygen content must comport itseK as a distinctly weaker base than the chromous oxide. Such salts, derived from chromium sesquioxide, wiU therefore be more extensively CHROMIUM 601 hydrolysed than the ohromous salts, e.g. CrCls is more strongly hydrolysed than CrCla. The sesquioxide of chromium can, however, also function as an acid. It will dissolve in strong bases to form chromites. Cr{0H)3 + 3NaOH->Na3Cr03 + SH^O sodium orthochromite. Cr(0H)3 +NaOH->.NaCr02 +2H2O sodium metachromite. Chromium trioxide (chromic anhydride) is a strong acid- forming oxide. It gives rise to the chromates. CrOj + 2NaOH^Na2Cr04 + H2O. Chromium trioxide and the chromates are strong oxidising agents. When they are reduced under atmospheric conditions, the reduction is from CrOj to CrjOa, i.e. from hexavalent to trivalent chromium, 2Cr03^Cr203+30. Every oxidation process in which chromium trioxide or a chromate takes part involves this reduction. Ohromous Oxide and Salts Chromous Oxide, CrO, is obtained as a black powder by the oxidation of chromium amalgam in the air. When heated in the air, it inflames and passes into the sesquioxide. Chromous Hydroxide, Cr(0H)2, is obtained as a yellowish precipitate when sodium hydroxide is added to a solution of chromous chloride {q.v.) in the absence of air. It is rapidly oxidised by air to the chromic state. On treatment with acid, chromous hydroxide dissolves, producing a chromous salt. In this way chromous sulphate has been prepared. Chromous Chloride, CrCl2, is prepared by the action of hydrochloric acid upon chromium in the absence of oxygen, or by the reduction of chromic chloride {q.v. ) in a stream of hydro- gen. The anhydrous salt is colourless, but the solution is a bright sky-blue. A solution of chromous chloride can be obtained by reducing a solution of chromic chloride with amal- gamated zinc. The pure chromous chloride (hydrated) can be obtained from the sky-blue solution by the precipitation of chromous acetate which is then redissolved in hydrochloric acid and a stream of hydrogen chloride led through the cooled solu- 602 AN INORGANIC CHEMISTRY tion. A solution of chromous chloride forms a very strong reducing agent. Chromous chloride, dissolved in a solution of hydrochloric acid, wiU evolve hydrogen in the presence of a piece of platinum foU. Chromous Sulphide, CrS, can be precipitated from a solu- tion of a chromous salt by the action of ammonium sulphide. The stabiUty of this sulphide in the presence of water is a striking example of the much stronger basic nature of chromous oxide as compared with chromic oxide, for the sulphide, CrjSa, cannot be prepared in the wet way. Chromous Sulphate, CrSOi, prepared by dissolving chro- mium in dilute sulphuric acid in the absence of air, forms fine blue crystals, isomorphous with ferrous sulphate, FeSO,. Chromous Carbonate, CrCOa, is prepared by precipitating the chloride with potassium carbonate. Double salts with the alkah carbonates have been obtained, these being noted for their strong reducing action. Chromic Oxide and Salts. Chromic Oxide, CrjOa (Sesquioxide) and Hydroxide. — The greenish blue hydroxide is thrown down from a solution of chromic sulphate or chloride by the addition of an excess of ammonium hydroxide. The oxide is best prepared by dehydrating the hydroxide or by the direct decomposition of ammonium dichromate. (NHOaCr^O, -^ Cr^Oa + N2 + 4H,0, and by the reaction KaCr.O, + S ^^ Cr,03 + K0SO4. Ignited chromic oxide is not attacked by acids. It is used in the manufacture of various paints and pigments, e.g. Guignet's green, formed by reducing potassium dichromate with boric acid. Both the oxide and hydroxide dissolve in strong bases, forming chromites which are derived either from ortho- chromous acid HjCrOa, or from the meta- acid, HCr02- Aqueous solutions of the chromites are hydrolysed on boiling, with the precipitation of the hydroxide. The fusion of a chromic salt or oxide with potassium carbonate in the presence of air, or better stiU, with a Httle nitre, produces the characteristic colour of the chromate (q.v.). CHROMIUM 603 Chromic Chloride. — ^Anhydrous chromic chloride is prepared by the action of chlorine upon a heated mixture of chromic oxide and carbon. CraOa + 3C + SCl^ -> 2CrCl3 + SCO. Peach-coloured scales of the composition CrClj volatihse from the reaction tubes. These crystals dissolve in water with extreme slowness, but the solution is rapid in the presence of many reducing agents, e.g. chromous chloride, stannous chloride. Hydrated chromic chloride, CrCl3,6H20, can be prepared in two distinct modifications, a green and a blue. Ordinary solu- tions of chromic chloride appear to contain varying proportions of the two varieties, depending upon the temperature and concentration. At low temperatures the equilibrium lies on the side of the blue variety, and as the temperature rises, the proportion of the green form in the solution increases. It is interesting to note that the change is not in this case a mere change of crystal form (cf . mercuric iodide), but there is a marked change in the chemical nature of the two chlorides. Silver nitrate precipitates the whole of the chlorine from a solution of the blue variety, but this reagent causes the precipitation of only one-third of the chlorine present in the green modification. The velocity with which the equilibrium of the two varieties is attained in aqueous solution is very slow, but if sufficient time is allowed, the same final conditions are reached whether the solution is made from the blue or the green variety. It is con- sidered that the two chlorides are related to each other as in the scheme, [Cr(H,0)e]Cl3 ^=^ [CrCl2(H,0)4]Cl + 2H20 [Cr(H20)e]+ + + + 3C1- [CrCl2(H20)4]+ + Cr Blue. Green. Chromic Sulphate — Chrome Alum. — Chromic sulphate can be prepared by acting upon chromium hydroxide with sulphuric acid. It forms bluish violet crystals of the composi- tion Cr2(S04)3,18H20. On boiling the blue solution of chromic sulphate a similar change occurs as has already been described for chromic chloride. The green modification, formed at high temperatures, differs fundamentally from the normal blue modification as only one 604: AN INORGANIC CHEMISTRY third of the SO4 present in the compound can be precipitated by the action of barium chloride, whilst the whole of the SOj present in the blue solution can be thrown down as barium sulphate. The exact constitution of the green modification is not yet clearly understood. As the temperature of the green solution slowly falls, the colour changes back to blue. Chromic sulphate forms with the sulphates of the alkalies a series of double sulphates which belong to the class known as the alums. These possess the same general formula M'2SOi,Cr2(SOi)3,24H20, as those described in connection with aluminium sulphate, and are also isomorphous with them. Chrome alum is best prepared by the reduction of potassium dichromate with sulphur dioxide or by the addition of alcohol and sulphuric acid. We may write provisionally, KjCtjO, +S0, +H2SO4 +H,0->K,S04,Cr2(S04)3,24HsO. (K,0.2Cr03) (CrsOj.SSOa) As in aU cases of oxidation by means of dichromate, the two molecules of chromium trioxide present in the dichromate are reduced to CrjOs, leaving available three oxygen atoms for the oxidation of the substance to be oxidised, in this case sulphur dioxide, hence three molecules of sulphur dioxide are capable of being oxidised in accordance with the equation : 3S02+30->3S03. In order to form one gram molecule of potassium sulphate and one of chromium sulphate, foxu- molecides of sulphuric acid will be needed, of which three wiU be provided by the oxidation of the sulphur dioxide. The completed equation will therefore read : K2Cr20, + 3S02+H2S04 + 23H20->K2S04,Cr2(S04)3,24H20 If alcohol is used as the reducing agent, the actual oxidation is the conversion of alcohol into aldehyde, e.g. : C2H5OH+ 0->CH3.CH0 + H2O, i.e. three molecules of alcohol can be oxidised by one gram molecule of the dichromate. Moreover, four molecules of sulphuric acid will be needed for the formation of the two sulphates. The final equation will then be : KjCraO, + 3C2H5OH + 4H2SO4 + I7H2O -^ K2S04,Cr2(S04)3,24H20 + 3CH3CHO. CHROMIUM 605 Chrome alum crystallises in large dark, purple-red crystal^ belonging to the regular system (octahedra). It finds com- mercial use in dyeing, calico-printing and tanning. Chromic Nitrate. — Chromic nitrate, Cr(N03)3,9H20, is obtained by the action of nitric acid upon chromic hydroxide. The bluish solution also turns green on boiling and returns to its old colour on cooling. The Chkomates and Dichromatbs Chromic trioxide gives rise to two kinds of salts, the chro- mates and the dichromates. CrO, + 2KOH-^K2Cr04 -\- H^O 2Cr03 + 2KOH^K2Cr20, -f- H2O The chromates and dichromates are therefore of the same degree of oxidation ; they differ merely in the ratio of acid to base. The dichromate contains an extra molecule of chromium trioxide, and as this is the active oxidising agent, it is therefore always more economical to use the dichromate as the oxidising agent instead of the chromate, though the mechanism of the reaction is quite the same. If one considers for a moment the analogous compounds C*r03 -^ K2Cr04 -> KpCr.O, (K20,2Cr03 ) S03->K2S04->K2S20,(K20,2S03) potassium pyrosulphate. the group relationship between the elements, sulphur and chromium, is thrown strongly to the fore. The only important difference between the dichromates and the pyrosulphates is that the latter compounds regenerate the acid sulphates on solution in water. On heating, 2KHSO4 -HoO->K2S20„ but on solution, K2S2O, + H2O -> 2KHSO4. Preparation of the Chromate and Dichromate.— Poto- sium chromate is prepared by roasting finely powdered chrome iron ore with the purest possible hme and potassium carbonate. 4FeO,Cr203 + 8K2CO3 + lO^-^SK^OrO, + 2'Fe,0s + 8CO2 K2O.G1O, The Hme is added to keep the mass thoroughly porous, so that the oxygen of the air can penetrate through the mass, and bring about the necessary oxidation. The reaction is essentially an 606 AN INORGANIC CHEMISTRY oxidation of the sesquioxide, contained in the chrome iron ore, to the higher oxide, CrOa. The resulting mixture of potassium and calcium chromate is then treated with a solution of hot potassium carbonate, when the reaction CaCrOi +K2C03->KaCr04 + CaCOa i takes place. The filtered Uquor is then concentrated and allowed to crystaUise. Sodium chromate is prepared by a similar method. When potassium chromate is dissolved in water, appreciable hydrolysis occurs, for chromic acid is a weaker acid than its analogue, sulphuric acid. KaCrO^ + 2H2O ^± 2K0H + HzCrOj. The addition of an acid (e.g. sulphuric acid) must exert an influence upon the equilibrium, as the concentration of the hydroxide would be lowered. The result of such an operation would be to cause a considerable increase in the concentration of the free chromic acid. So soon as appreciable quantities of chromic acid are formed, the reaction K^CrOi + H^CrOi ^=± K^Cr^O, + HoO sets in, i.e. potassium dichromate is formed. This is shown by the sudden change in the colour of the solution from yellow to orange red. This is the method adopted in industry to convert the yellow chromate into the orange red dichromate. Efforts have been made to reduce the cost of manufacturing potassium dichromate by first preparing sodium dichromate from the relatively cheap sodium carbonate. A very strong solution of sodium dichromate is then allowed to react with potassium chloride. NasCrjO, + 2KGl-^K2Cr20, + 2NaCl. The potassium dichromate, the solubiUty of which is mfluenced greatly by the temperature, is easUy separated from the sodium chloride by crystallisation. (100 gm. of water dissolve 5 gm. of potassium dichromate at 0°, 129 gm. at 100° ; at 9° 62 gm. of the sodium salt dissolve in 100 gm. of water). During recent years sodium dichromate has come increasingly to the fore for commercial purposes, and owing to its greater cheapness has largely displaced the potassium salt. CHROMIUM 607 If one examines the equations : KjCrOi + 2H2O ^=i 2K0H + H^CrOi KjCrOi + H2Cr04 ^=± KaCr^O, + H^O, it is evident that the upper equihbrium will be driven to the left, by the addition of potassium hydroxide, and the decrease in the concentration of the free chromic acid will upset the equihbrium in the second equation, i.e. the potassium dichromate will break down. This is shown by the colour change produced on adding potassium hydroxide to an orange red solution of the dichromate. Polychromates higher than the dichromates have been prepared by the action of chromic acid upon the dichromate, e.g. potas- sium tri-chromate, K2Cr30]o (KaOjSCrOa) ; potassium tetrachro- mate, 'KJQr^Ou (K20,4Cr03). Oxidising Reactions of the Dichromates and Chromates. — In all cases where these salts are used as oxidising agents, a change of valence of chromium towards oxygen occurs. CrViOa^Crm^Oa. With the knowledge that chromic salts are always produced by the reduction, it becomes comparatively easy to write down the equation for any such oxidation-reduction process. In the oxidation of hydxiodic acid by potassium dichromate we may write provisionally : KsCr^O, + HI + H^SO.-^KaSOi + ^2(804)3 + 12 + H2O. Hydrogen iodide is oxidised to iodine and water in accordance with the equation : 2HI + 0->H20+l2. Therefore, one gram molecule of dichromate with its three available oxygen atoms will oxidise six molecules of hydrogen iodide, and four molecules of sulphuric acid will be necessary to form the sulphates of potassium and chromium. The completed equation then becomes : KaCr^O, + 6HI + 4H2S04->K2S04 + Cr2(S04)3 + 3I2 + 7H2O The action of hydrogen sulphide on an aqueous solution of 608 AN INORGANIC CHEMISTRY potassium dichromate containing sulphuric acid can be handled similarly. From a knowledge of the reaction H^S+O-^H^O+S, we can at once set down the completed equation : K^Cr^O, + SHjS + 4H,SO,-^K,S04 + Cr,(S04)3 + 7H,0+3S. If sulphtiric acid is not added before passing in the hydro- gen sulphide the equation would be : K A2O, + SH^S + HjO -> 2K0H + 2Cr (0H)3 + 3S Another common reaction of this type is the oxidation of ferrous sulphate to ferric sulphate. In this case we may write provisionally KAA+H2S04+FeS04-^K2S04+Cr2(S04)3+Fe2(S04)3+H20 (K20.2Cr03) (FeO.SOa) CrA-SSOg) (FeaOa-SSOj) To convert two molecules of ferrous oxide into ferric oxide one atom of oxygen is required. 2FeO + O^Fe203. One gram molecule of potassium dichromate wiU therefore oxidise six molecules of ferrous oxide, i.e. ferrous sulphate. In order to produce ferric sulphate, sulphuric acid must be pro- vided. Hence : 6FeS04 + 30 + 3H2S04-^ 3Fe2(S04)3 + 3B.fi. As before, four molecules of sulphuric acid must also be added for the formation of potassium sulphate and chromium sulphate. The completed equation wiU consequently read : K^Cr^O, + 6FeS04 + 7H2SO4 -^ K,S0„ + 01-3(804)3 + 3Fe2(S04)3 + m,0. AH the above reactions are equally weU carried out by potas- sium chromate except that, for every molecule of the dichromate, two molecules of the chromate must be used in order to provide the same amount of available oxygen, and an extra molecule of sulphuric acid will be necessary to form the extra molecule of potassium sulphate. 2K2Cr04 + 6FeS04 + 8H2S04-> 2K2SO4 + Cr2(S04)3 + 3Fe,(S04)3 + SH^O 2KAO4 + 3H2S+ 5H2S04->2K2S04 + Cr2(S04)3 + 3S + 811,0 CHROMIUM 609 Other Reactions of the Bichromates. — Under the action of heat the dichromate breaks down in accordance with the equation 4X^20, --> 4KAO4 + 2Cr203 + 3O2. (KAOi-CrOs) The excess group of CrOa suffers decomposition (cf. the action of heat upon chromic anhydride, CrOs). Owing to the hydrolysis to which the dichro mates are subject K^CrA +HOH-^K2Cr04 +H2Cr04, or written ionically, Cr^O," +OH--^2Cr04- +H+, there is always present in solution an appreciable concentration of the chromateion, CrO",. This accounts for the fact that the addition of potassium dichromate to a solution which contains an ion of a heavy metal always leads to the precipitation of the insoluble chromate. 2Pb(N03)2 +K2Cr20, +H20-^2PbCr04 4- +2HNO3 +2KNO3 Of the heavy metal chromates, that of lead, PbCr04, and of silver, Ag2Cr04, are often met with in analysis. The bright yellow precipitate of lead chromate forms a useful test for the presence either of lead ions or of chromate ions. In titrating sodium chloride with silver nitrate a drop of potassium chromate is added to the solution of the sodium chloride, and the appear- ance of the red silver chromate shows that the precipitation of the silver chloride is complete (cf. p. 451). Chromic Anhydride — Chromium Trioxide. — This sub- stance is prepared by the addition of concentrated sulphuric acid to a strong solution of a soluble dichromate. Na^Cr^O, + H2S04->Na2S04 +2Cr03 +H2O. Red needles of the anhydride separate out from the solution. It dissolves in water, forming chromic and dichromic acids. HaO+CrOa-^HaCrOi H20+2Cr03^H2Cr20, As an oxidising agent, it has exactly the same properties as the dichromates and chromates. 2Cr03 + 3H2S + 3H2S04^Cr2(S04)3 + 3S + BH^O. EE 610 AN INORGANIC CHEMISTRY On being heated, chromic anhydride breaks down into chromic oxide and oxygen. 4Cr03-^2Cr303+302. Chromyl Chloride. — When potassium dichromate, a chloride and sulphuric acid are heated together, a reddish brown liquid distils over. This is chromyl chloride. KaCraO, + 3H2SO4 + 4NaCl-> K2SO4 + 2Na2S04 + 2Cr02Cl2 + SH^O The reaction is essentially one between the chromic anhydride present in the dichromate, and the hydrogen chloride Uberated by the action of sulphuric acid upon the sodium chloride. CrOa + 2HC1 -^ Cr02Cl2 + H2O Chromyl chloride boils at 118°. It is hydrolysed by water thus : <3r02Cl2 +2H20->H2Cr04 + 2HCl. No corresponding bromide or iodide are known, so that the above reaction forms a satisfactory test for detecting a chloride in the presence of a bromide or an iodide. The distillate is rendered alkaUne ^\'ith ammonium hydroxide, acidified with acetic acid, and if the addition of lead acetate leads to the for- mation of a yellow precipitate, the presence of a chloride in the original mixture may be inferred. (The object of adding a shght excess of ammonium hydroxide is to remove any bromine which may have been Uberated by the action of the sulphuric acid upon any bromide present.) Just as the hydrolysis of the analogue, sulphuryl chloride, leads to the formation of sulphuric acid [q.v.), so we find that chromyl chloride undergoes a similar hydrolysis. O^ CI HOH OH '>Cr<^ + -> >Cr< +2HC1 O ^CI HOH 0^ ^OH OH Although the intermediate chlorochromic acid, /C!r^ , has not been prepared, some of its salts are known. Potassium chlorochromate is formed when the dichromate is treated with hydrochloric acid and the solution is allowed to crystallise. MOLYBDENUM, TUNGSTEN, URANIUM 611 .OK K,Gt,0,+2KCI-^ VCr/ +H,0. O' ^Cl Perchromic Acid.— If hydrogen peroxide is added to an acidified solution of a chromate, a deep blue colour appears. The compound causing this blue colour can be extracted in ether, in which the substance is not only more soluble but also more stable. This forms a deUcate test for hydrogen peroxide as well as for chromates. The colour disappears on boihng, and the greeny blue colour of the trivalent salt of chromium comes out. A number of perchromates have been prepared, but their composition is too vague to allow of their chemistry being clearly understood. Some consider that hydrogen peroxide adds itself on in a way similar to water of crystallisation ; possibly an acid, H2Cr208, corresponding to HjSaOs, is formed. MoLYBDBIfUM, TtJNGSTEN, UbANIUM. These elements all form the group oxide MO3 and the corre- sponding acid H2MO4. Molybdenum. — This element occurs fairly extensively as molybdenite, M0S2, and as wulfenite, lead molybdate. The metal has come into considerable importance during recent years for the manufacture of aUoy steels. These molybdenum-steel alloys, besides being very hard, have the remarkable property of retaining their temper, even when heated to a high tempera- ture. Large quantities of molybdenum steel are used for high speed tools, rifle barrels, etc. The metal is obtained either by the Goldschmidt method or by reduction in an electric furnace. Molybdenum forms numerous oxides, M02O3, M0O2, M0O3 and possibly MoO. The most important oxide is the trioxide. This is obtained by roasting molybdenite in the air. Molybdenic anhydride (M0O3) dissolves in the hydroxides of the alkalies, forming molybdates, e.g. (NH4)2Mo04, which is used largely in analytical work for the precipitation of phosphorus as ammonium phospho-molybdate . Numerous chlorides have been prepared, M0CI2, M0CI3, M0CI4, M0CI5 and MoCle. Tungsten. — This element occurs as scheelite, CaW04, and wolfram, FeW04. In order to obtain the metal the tungstate is first roasted with sodium carbonate. The sodium tungstate is 612 AN INORGANIC CHEMISTRY then extracted and free tungstic acid obtained by the addition of an acid. This acid is dehydrated and the tungstic oxide reduced in the electric furnace. Large quantities of tungsten are used in the manufacture of special tool steels (cf. molybdenum). Tungsten forms two oxides, WO2 and WO3, and the chlorides WCI2, WCI4, WCI5 and WCle. Sodium tungstate is used commercially as a mordant, and also for rendering fabrics less inflammable. Uranium. — ^Uranium is found in pitchblende, camotite, samarskite and in other rare minerals. Numerous oxides have been isolated. Uranous oxide, UOj, a reddish brown basic oxide, which on ignition passes into UjOs , gives rise to the uranous salts in which uranium is tetravalent. The group UO2 is often known as uranyl. Uranic oxide, UOi, or uranyl oxide, UO2.O, a yellow oxide, being of an amphoteric nature, gives rise to the uranic compounds as well as to the uranates. K2U2O, or K2O.2UO3 shows a considerable likeness in structure to the dichromates. Uranic oxide also reacts with acids to form the basic uranyl salts, UO2.SO4, UO,(N03)2.6H.O, etc. Uranium peroxide, UO4, formed by the action of hydrogen peroxide upon a solution of uranyl nitrate, gives rise to the peruranates, e.g. Na2U05. Other oxides are U308(UO,.2U03), UjOslUOa-UOj) and U0O3. The chlorides VCij, UCI4, UCI5 and UO2CI2 are known. The radioactive phenomena of the salts of uranium will be discussed in Chapter XLI. Questions 1. Ciive a concise account of the behaviour of potassium dichromate as an oxidising agent. 2. Tabulate the oxides of chromium, and compare their properties, ^Vhat class of compounds does each oxide give rise to ? 3. Construct six equations in which potassivim dichromate functions as an oxidising agent. 4. Give an account of the chemistry of the dichromates and chromates. 5. How may the cuprous salts be derived from the cupric salts ? Briefly compare the properties of the cuprous salts with the properties of the corresponding silver salts. 6. What are the conditions under which hydi'ogen peroxide is formed ? What evidence is afforded by its reactions as to its constitution ? An aqueous solution of hydrogen peroxide, when warmed with colloidal platinum, gave off 2-5 c.c. of oxygen, and left 37 c.c. of water, both being measured at 4° C. [p = 760 mm.] What was the composition of the solution ? 7. Describe the occurrence and metallurgy of tungsten. What com- mercial application has this element ? CHAPTER XXXVIII GROUP 7A MANGANESE The elements of the seventh group are : F The short period elements, fluorine and chlorine, I have been classed with bromine and iodine, forming CI the halogen group. Manganese is the sole repre- ^\ sentative of Group 7a. Although in its lower Mn \ stages of oxidation it shows a great resemblance to I Br the salts of the metals of corresponding valence, — j there is a fairly well marked similarity between j I manganese and the halogens, when the group oxide — M2O7 and its compounds are taken into consider- ation. From manganese heptoxide is formed permanganic acid, HMnOiCcf. CI2O7 and HCIO4). The permanganates and per- chlorates, for example, are isomorphous, an unfaihng indication of strong chemical analogy. They are also strong oxidising agents. Occurrence. — The ores of manganese are widely distributed, especially as the oxide. Amongst the more important ores are pyrolusite, MnOa ; braunite, MuaOa ; hausmannite, MngOj ; manganite, MnjOajHsO ; manganese spar, MnCOs ; the sulphide, sulphate, silicate, tungstate and phosphate also occur. Metallurgy and Properties. — Manganese can be prepared by the reduction of an oxide with carbon at a white heat. This is the basis of Moissan's method of reduction in an electric fur- nace. Although this method is still used commercially, it has been largely displaced by the Goldschmidt process. 3MnO, +4A1^3Mn + 2Al203. Manganese is a grayish metal with a reddish tinge. As a rule, the metal is hard and brittle, but its properties are largely 613 614 AN INORGANIC CHEMISTRY aSected by impurities. It melts at 1,260° and boils at 1,900°. Manganese is rapidly oxidised in air. It decomposes water with the evolution of hydrogen, even in the cold. Manganese dis- solves freely in acids forming manganous salts. Mn + 2HC1 -> MnCU + H2 . When heated, it is attacked by boron, silicon, phosphorus, forming a boride, etc. Manganese steels have considerable industrial importance, as have alloys of manganese, with copper, bismuth, zinc, aluminium, etc. The Chemistry of the Compounds of Manganese Few elements exist which rival manganese in the number of oxides which it forms and in the variety of salts to which these oxides give rise. The general rule that, with an increasing oxygen content of the oxide, an increase in the acid nature of the oxide and a decrease in the basic nature is to be expected, finds excellent confirmation in the study of the oxides of man- ganese. The oxides and their compounds are : Oxides : Name. Formula. Property. Salts. Manganous oxide MnO Basic Manganous salts MnCl,, MnSOj Manganese tetroxide Mn^Oi Mixed oxide Manganese sesquioxide Mn^Oa Feebly basic Manganic salts Mn,(S0)3 ■O^^MnQi.MnSj — ^ Manganites CaMnOs Manganese dioxide MnO 2 Amphoteric Manganese trioxide MnO 3 Acidic Manganic acid and manganates, HjMnOi.KjMnOi Manganese heptoxide Mn,0, Strongly Permanganic acid and acidic permanganates, HMnOi.KMnO^ Manganous Oxide and the Manganous Salts. — Manganous oxide, MnO, is a greyish green powder, obtained by heating the higher oxides in a stream of hydrogen. It can be prepared by the decomposition of manganous hydroxide and carbonate, also by fusing anhydrous manganese chloride and sodium carbonate with a httle ammonium chloride. The mass is then hxiviated. MANGANESE 615 Manganous hydroxide is thrown down by the addition of a soluble hydroxide to a solution of a manganous salt. It is rapidly oxidised in the air to the higher oxides (MnsOi, MnaOa) so that the pure hydroxide can be prepared only in an atmosphere free of oxygen. The manganous salts may be prepared either from the oxide and hydroxide by solution in the required acid, or, as in the case of the strong mineral acids, one can proceed direct from the metal. The manganous salts are pink in colour ; in the solid form they show no tendency to oxidise (cf . the chromous and ferrous salts). They are also stable in solution provided there is not an excess of alkah. Manganous chloride, MnCU, and sulphate, MnS04, form double salts with the alkali and ammonium chlorides and sulphates respectively. Upon this fact depends the separation of iron from manganese by the addition of ammo- nium chloride to solutions containing salts of iron and manganese. Solutions of manganous chloride combine with ammonia to form stable complex salts of a type similar to those given by copper, zinc, silver, cobalt and nickel, e.g. (Mn,a;NH3)Cl2. A solution of this salt, alkaUne as it is, is rapidly oxidised by oxygen, forming a precipitate of manganic hydroxide, Mn(0H)3. Owing to manganous hydroxide being a relatively strong base the manganous salts formed from the strong acids undergo little hydrolysis (cf. manganic salts). In general, there is considerable kinship between the manganous salts and other similarly constituted salts (cf. MnCls, CuCU, MnSOi, ZnSOi, etc.) Manganese, therefore, affords another example of the rule that, when a metal gives rise to more than one series of salts of different degree of oxidation, a considerable similarity will be found between these salts and other similarly constituted salts of the same degree of oxidation. Manganous chloride is produced by the action of hydrochloric acid upon manganous carbonate. It is, however, also obtained from the waste hquors of the Weldon process of manufactur- ing chlorine from manganese dioxide. After removal of impur- ities such as iron, the solution is brought to crystallisation. Several hydrates are known, e.g. MnCU.eHaO ; MnCl2,4H20 ; MnCl2,2H20. The anhydrous chloride, prepared by the direct action of chlorine upon manganese, forms a pale rose-coloured mass. On heating, this melts to an oily liquid, which begins to 616 AN INORGANIC CHEMISTRY lose hydrochloric acid in the presence of moisture, forming oxides of manganese. Manganous sulphate is prepared from manganese dioxide and sulphuric acid. The paste is heated strongly to break down the manganic sulphate and to convert any iron present into the oxide. The mass is then Hxiviated. Quite a number of crystal- line hydrates have been prepared from this salt, but the general principles governing their formation are precisely the same as described for the hydrates of copper sulphate, etc. Manganese tetroxide — MngOi. — This oxide is formed as a red powder when the other oxides are heated in air. The oxide is very stable towards heat. Its behaviour on treatment with an acid shows that it is a mixed oxide or inner salt (of. PbaOa, PbsOj). On being treated with dilute sulphuric and nitric acids, the following reactions take place : jMn304 + 4HNO3 -^ 2Mn(N03)2 + Mn02,2H20 Mn304 + 2H2S04-^2MnS04 + Mn0.,2H,0. With hydrochloiic aeid, chlorine is evolved. Mn304 + 8HC1^. SMnCla + 4H2O + CL. These reactions, considered alone, would indicate that man- ganese tetroxide is a mixed or compound oxide, 2MnO,Mn02. But, on the other hand, it has been shown that acetic acid leaves behind a precipitate of manganic oxide, MnoOs. MnaOi + 2CH3.COOH-^_Mn(C2H302)2 -|-Mn203 -f H^O, whilst cold concentrated sulphuric acid forms a mixture of manganous and manganic sulphates. These reactions point to the constitution MnO,Mn203. Probably two forms of the oxide exist as in the scheme : MnO,Mn203 Mn304^ ^2MnO,Mn02 Manganic Oxide and the Manganic Salts. — Manganic oxide, Mn203, can be prepared by heating any of the other oxides in a stream of oxygen. It is a brownish black oxide. The behaviour of manganic oxide is somewhat similar to that of the tetroxide. If it is treated with dilute sulphuric acid, a pre- MANGANESE 617 cipitete of manganese dioxide is thijown down, whilst manganous sulphate IS found in the solution. Dilute nitric acid behaves in the same way, These facts are strong evidence in favour of the constitution MnCMnOa. On the other hand, the existence of weU defined salts of trivalent manganese such as MnFa ; MnCl3,2KCl ; Mn2(S04)3 ; MnP04,2H20 ; Mn(C2H302)3 ; and the manganic alums M'2S04,Mn2(S04)3,24H20; bear convincing testimony to the basic nature of the oxide Mn203. The conflicting behaviour of manganic oxide is, perhaps, best explained by the assumption that an equilibrium Mn203^:i:±MnO,Mn02 exists. Of the manganic salts the acetate is especially easy to obtain. The anhydride of acetic acid is allowed to act upon manganous nitrate ; the nitric acid set free oxidises much of the manganese to the trivalent form. Brown crystals of manganic acetate separate out. Manganic fluoride has been obtained by the action of fluorine upon manganous iodide. It forms characteristic double salts, NaaMnFj. Manganic chloride has not been isolated in the pure state, though there is little doubt that it is formed by the action of hydrochloric acid upon manganic oxide. It forms several double salts which have been prepared in the pure state, e.g. K2MnCl5. Manganic suljihate has been obtained as a green powder by the action of concentrated sulphuric acid upon freshly precipitated manganese dioxide at 110°-140°. Solutions of manganic sulphate are rapidly hydrolysed even in the cold. Mn,(S04)3 + 6HOH->2Mn{OH)3 + 3H2SO4. It is extremely unstable. Manganic sulphate forms a series of alums, isomorphous with other alums. These are more stable than the sulphate. It is an interesting example of the greater stabihty conferred upon a salt by the presence of another salt of stronger electro-affinity. Similar behaviour is shown by manganic chloride and fluoi'ide. Thus, manganic fluoride is hydrolysed by water in accordance with the equation MnFa +3HOH-^Mn{OH)3 +3HF, but the double salt, formed by the addition of potassium fluoride, is scarcely hydrolysed. In agreement with these speculations it has been found that the alums formed from the strongly 618 AN INORGANIC CHEMISTRY electro -negative elements, caesium and rubidium, are the most stable. Manganese Dioxide. — ^Manganese dioxide, MnOj, is a black substance, obtainable by heating manganous nitrate. In the hydrated form it may be prepared by the reduction of potassium permanganate by means of manganous sulphate (Volhard method of estimating manganese). 2KMn04+3MnS04+7H,0->2KHS04+H2S04+5Mn02,H20 4- Hydrated manganese dioxide is also thrown down from solutions of manganous salts when such oxidising agents as sodium hypochlorite are used. Naturally occurring manganese dioxide, pyrolusite, is used in the bottle industry, smaU quantities being added to the glass to oxidise the greenish ferrous sUicate to the comparatively colourless ferric compound. Manganese dioxide functions as a very weak base, hence salts formed from it wiU suffer extensive hydrolysis. Of the few salts derived from this oxide, the sulphate, Mn(S04)2, acetate, Mn(C2H302)4 and sulphide, MnSj, are worthy of note. The chloride, MnCli, and fluoride, MnFj, have not been isolated in the pure state, but well defined double salts of these compounds are known, e.g. KaMnClj, KjMnFg, etc. When treated with hydrochloric acid, manganous chloride and chlorine are formed, while sulphuric acid leads to the formation of manganous sulphate and the Kberation of oxygen. These reactions, coupled with the failure to secure evidence of the formation of hydrogen peroxide, have led to the view that manganese dioxide is not a true peroxide ; its formula is therefore not Mn/ O but Mn \ ^O (cf. p. 199). Manganites . — Besides the salts derived from manganese dioxide functioning as a base, a number of salts are known in which the acidic group appears to be derived from this oxide. Manganese dioxide is the anhydride of manganous acid and gives rise to the manganites. As a definite acid, manganous acid itseK is not known, but numerous salts from it have been prepared, MANGANESE 619 Fusion of manganese dioxide in potassium hydroxide does not lead to the simple manganite of the type, M2Mn03(M20,Mn02), but polymanganites are produced. Examples of such polyman- ganites are K20,2Mn02; KaO,5Mn02,a;H20, etc., Manganites of the alkahne earth elements of a more simple type have been prepared by fusion methods. Lime and anhydrous manganous chloride lead to the formation of brownish red crystals of calcium ortho-manganite, Ca2Mii04(2CaO,Mn02). By somewhat similar means black, iridescent crystals of the meta-manganite, CaO.MnOj have been formed. Manganites of calcium are also formed during the Weldon mud process for utihsing the waste pro- duct, manganous chloride, obtained during the manufacture of chlorine (q.v.). The liquor is treated with slaked hme and air is blown through the mass. Oxidation of the manganous hydroxide occurs, leading to the formation of calcium manganites CaMn03(CaO,Mn03) and CaMn205(CaO,2Mn02). When these are treated with hydrochloric acid, they behave like a mixture of manganese dioxide and calcium oxide. CaO.MnOa + 6HC1 -> CaCl2 + MnClj + CI2 + SH^O. Manganates and Manganese Trioxide. — If manganese dioxide is fused with potassium hydroxide or carbonate either in the presence of air or of an oxidising agent such as potassium nitrate, a dark green mass is obtained. 2Mn02 + 4K0H + 0^ -^2K2Mn04 + 2H2O. If the mass is extracted with water, and the liquid evaporated at a low temperature in vacuo, dark green crystals of potassium manganate, KjMnOi, separate out. These crystals are iso- morphous both with potassium sulphate and chromate. The crystals dissolve without decomposition in an aqueous solution of an alkah hydroxide, but decompose in neutral or acid solu- tions. The reason of this is to be sought in the extreme weakness of manganic acid, hence the hydrolysis in the sense of the equation : K2Mn04 +2HOH^H2MnO, +2K0H. The presence of potassium hydroxide throws back this hydro- lysis, hence the greater stability of the manganates in alkahne solutions. Manganic acid itself is somewhat unstable, and, it 620 AN INORGANIC CHEMISTRY present to any extent, undergoes auto -oxidation. In neutral solutions, we have : 3H2Mn04 -> 2H2Mn04 +MnOji +2HjO (HjCMnOs) (HaCMn^O,) The hexavalent manganese present in the manganic acid is reduced to the tetravalent state in MnOj, a decrease of two units, whilst in permanganic acid the valence has risen to seven, an increase of one unit. Consequently, two molecules of perman- ganic acid must be formed at the expense of one molecule of manganese dioxide. Owing to the fact that tetravalent man- ganese always undergoes reduction to the divalent state in acid solution, the auto-oxidation in acid solution proceeds somewhat differently : .-.H JlnOj + H,S04->4HiMn04 + MnSOi + 4H,(3. Potassiuiu and sodium manganates are strong, though unstable oxidising agents. They can, however, only be effectively used for this purpose in alkaUne solution as the acid solutions are too unstable. Under such conditions the manganate is reduced to the hyilrated manganese dioxide. Manganese Trioxide. — This unstable oxide is prepared by allowing the green solution of potassium permanganate in sul- phuric acid to (hup upon anhydrous sodium carbonate. The whole must be kept thoroughly cool. Red vapours are evolved, which are condensed in a suitable U tube, surrounded by a freezing mixture. This red soUd is very dehquescent. At ordinary teiiiperatures it decomposes slowly. Permanganates — Manganese Heptoxide. — The hydrolysis of potassium or sodium manganate leads to the formation of potassium permanganate, 3K,MnO, +6H0H ;;zi± 3H,:\InO^ + 6K0H 2HMn0 , + MnOo + 2H.,0 the permanganic acid ultimately being converted into the potassium salt by the potassium hydroxide. As a commercial method of preparing potassium permanganate this method leaves something to be desired, as only two-thirds of the manganate is raised to the higlier stage of oxidation. A MANGANESE 621 more effective procedure is to pass chlorine or ozone through a solution of the manganate. 2K2Mn04+Cl2-> 2KMn04+2KCl 2K2Mn04 + O3 + H2O -^ 2KMn04 + 2K0H + 0^ The latter equation embodies the chemistry of an important commercial method of carrying out the desired oxidation of manganate to permanganate. Potassium permanganate is also prepared by the electrolytic oxidation of the manganate at an iron or nickel anode. Sodium permanganate is made in the same way, but owing to its great solubility, it is rarely put on the market in the solid form. Potassium permanganate forms dark purple, rhombic crystals which have a greenish metaUic lustre. They are isomorphous with potassium perchlorate. If potassium permanganate is slowly added to well cooled sulphuric acid, a deep oHve liquid is formed. After thoroughly cooUng the liquid, a few drops of water are added, and oily drops of manganese heptoxide, MujO,, separate. When rapidly heated, manganese heptoxide explodes. The substance is volatile, is stable in dry air, but in the presence of moisture it reacts slowly, forming ozone and permanganic acid. Manganese heptoxide is a very powerful and vigorous oxidising agent. Substances such as paper, ether, fat, sulphur, etc., cause instant inflammation and even explosion. Permanganic acid can be prepared from barium permanganate and sulphuric acid. Ba(Mn04)2 + H2S04-»2HMn04 + BaSOi. Dilute solutions can also be formed by dissolving the hept- oxide in water. The acid is unstable in solution, decomposing slowly into hydrated manganese dioxide and oxygen. Per- manganic acid is a strong acid. This is shown by the high conductivity of its aqueous solutions. Potassium permanganate is frequently used as a source of pure oxygen. 2KMn04^^K2Mn04 + MnOa + O2. This decomposition proceeds freely at temperatures above 240°. A mixture of potassium permanganate and sulphur, glycerine or phosphorus is violently explosive. If a soluble hydroxide is added to a solution of potassium 622 AN INORGANIC CHEMISTRY permanganate, the colour changes from a purple to a dark green. This is due to the formation of the manganate. 4KMn04 + 4K0H ^i^ 4K2Mn04 + 2H2O + O^. Even a weak acid, such as carbonic acid, causes the change back to permanganate. SKsMnOi + 2H2CO3 ->2KMn04 + MnO, ^ + 2K2CO3 + 2H2O. The precise manner in which potassium permanganate will behave in the presence of a reducing agent depends upon the acidity or alkahnity of the solution. If the solution is acid, the permanganate is always reduced to the divalent state, salts of manganous oxide, MnO, being formed, i.e. Mn207 — >-2MnO+50. If the solution is alkahne, the reduction proceeds only to manganese dioxide, which separates out as a iiocculent preci- pitate. MnA-^MnOa + SO. Consequently, in acid solution an extra atom of oxygen is avail- able for oxidation from every molecule of the permanganate, containing. ^MuaO,. As a rule, potassium permanganate is employed in acid solu- tion for the purpose of determining the quantity of reducing agent present. The sharpness with which the first trace of the pink colour of the permanganate may be detected, makes this salt a most useful reagent for the estimation of solutions containing reducing agents such as sulphurous acid, ferrous sulphate, etc. In nearly all cases the solution is kept acid by the addition of sulphuric acid, as hydrochloric acid itself would be oxidised by potassium permanganate. The estimation of the quantity of oxalic acid in solution is represented in the equation : 2KMn04 + 5H2C2O4 + 3H2S04-> K,S04 + 2MnS04 + IOCO2 + SH^O In passing from MuaO,— >-2MnO, five oxygen atoms are avail- able for the oxidation of the oxahc acid, H2C204 + 0-^H20+2C02 hence, five molecules of oxaUo acid wiU be equivalent to two molecules of potassium permanganate, KMn04 (the student will note that as the oxide from which this salt is derived is MuaO,, MANGANESE 623 it is usual to work with the double molecule rather than with the gram molecule itself). Moreover, three molecules of sul- phuric acid wiU be needed to ensure the formation of potassium and manganous sulphates. In the reaction between hydrogen peroxide and potassium permanganate, we have as the primary equation : 2KMn04 + HA + H,S04^ K,S04 + MnS04 + H,0 + O2. The reaction H2O2+O— ^-HaO+Oj shows that the double molecule of permanganate (2KMn04) will oxidise five molecules of hydrogen peroxide ; three molecules of sulphuric acid must be added. The completed equation is : 2KMn04 + 5H2O2 + 3H2S04->K2S04 + 2MnS04 + SH^O+SO^. A similar reaction is the following : 2KMn04 + 5HNO2 + 3H2S04-> K2SO4 + 2MnS04 + 5HNO3 + 3H2O Consider the oxidation of ferrous sulphate. The primary equation is : 2KMn04 + FeSOi + H2S04^> (Mn20,) (FeO) K2SO4 + MnS04 + Fe2(S04)3 + H2O. (MnO) (FePa) In this reaction the oxidation is : 2FeO-f O-^FeaOs. Consequently, ten molecules of ferrous sulphate can be oxidised by the double molecule of the permanganate. For this, three molecules of sulphuric acid are required for the formation of potassium and manganous sulphates, while five molecules will be needed in forming the ferric sulphate, in all, eight molecules of acid. The completed equation must be : 2KMn04 + lOFeSOi + SHjSOi^ K2SO4 + 2MnS04 + 5Fe2(S04)3 + SH^O. Questions 1. Tabulate the oxides of manganese, and indicate the types of com- pounds, if any, derived from them. 2. Discuss the action of hydrochloric acid on the various oxides of manganese. 624 AN INORGANIC CHEMISTRY 3. Compare and contrast manganese with the halogen members of Group 7. 4. Construct three equations in which potassium permanganate acts as an oxidising agent (a) in acid solution, (6) in alkaline so ution. 5. What is meant by a mixed oxide, or an inner "alt ? Illustrate your answer by reference to the oxides of manganese. 6. Describe the preparation of potassium permanganate from man- ganese dioxide. 7. Construct equations to illustrate the action of concentrated sulphviric acid upon (a) carbon, (6) manganese dioxide, (c) potassium bromide, (d) potassium bromide and manganese dioxide mixed together. CHAPTER XXXIX IRON, COBALT, NICKEL These elements form the bridge elements in the first long period, their atomic weights lying very close together (Fe=55-85, Co=58-97, Ni=58-68). It is to be noted that, although the atomic weight of cobalt is greater than that of nickel, there is a far greater analogy between compounds of iron and cobalt than between those of iron and nickel, hence the elements will be studied in the above order. The oxides of these elements are : Jron. Cobalt. Nickel. Ferrous oxide, FeO Cobaltous oxide, CoO Nickelous oxide. NiO Ferric oxide, Fe./)., C'obaltic oxide, Nickelic oxide, CO2O3 NiaOj FerrosG-ferric oxide Cobalto-cobaltic ox- Nickelo-nickelic Fe30, ide, C03O4 oxide, NisO^ Iron Occurrence. — Iron occurs abundantly as oxide, carbonate and sulphide. The more important oxides include magnetite (mag- netic iron ore), FcsOi ; specular iron ore (red haematite), FeaOj ; limonite (brown haematite), 2Fe203,3H.20. Spathic iron ore or siderite consists of ferrous carbonate, FeCOa. The sulphide, pyrites, FeS.2, is widespread, but owing to the difficulty of removing the sulphur, it has little commercial value. Metallurgy Cast Iron. The first stage in the reduction of the iron ore is to subject it to roasting in order to expel carbon dioxide, water, sulphur and 625 SS 626 AN INORGANIC CHEMISTRY other volatile matter, as well as to oxidise ferrous to ferric oxide. Magnetite and red ha?matite do not require this prehminary roasting. The calcined ore is introduced into a blast furnace together with fuel and the necessary flux, hmestone. The function of the blast furnace is to reduce the oxides of iron to the metal and at the same time to ehminate as many of the impurities as possible in the form of slag. The blast enters the fur- nace through the tuyeres A (Fig. 134), after being heated in the hot blast stoves. The principal re- ducing agent in the blast furnace is carbon monoxide. The reduc- tion of ferric oxide to metallic iron is expressed in the equations : SFe^Oa +C0^2Fe304 + CO. FejO^ -I- CO ^ 3FeO + CO, FeO+CO->Fe+C02 The ferric oxide is almost entirely reduced in the upper portion of the furnace, and as the reduced iron descends through the furnace, its carbon content slowly increases. The slag formation is greatest near the twyers where the tem- perature is highest ; at this point the oxides of sulphur and phos- "" ■ phorus are reduced, and the ele- ments pass into the molten iron. The source of the heat is twofold, the principal supply being due to the combustion of the fuel in the neighbourhood of the twyers, but in the upper portion of the furnace a considerable amount of heat is evolved from the exothermal process : Fe^Oa + 3C0^2Fe -|- 3C0o. The molten iron, which is tapped ofi at intervals, is known as pig iron or cast iron. Cast iron, prepared by the above process, contains approxi- mately 3'5-4 per cent, of carbon, from 2-5-3-6 per cent, of silicon, 1 per cent, of manganese, varying quantities of sulphur and IRON 627 phosphorus, with a certain amount of carbon combined in the form of carbide. Cast iron has a sharp melting point which varies, however, with the composition of the iron. Wrought Iron. Wrought iron is made from pig iron by what is known as the 'puddling process. This is carried out in a reverberatory furnace, over the bed of which is spread a refractory Uning and a layer of hammer-slag (basic iron siUcate). The slag fuses during the operation, and forms a compact working bed. During the early part of the operation most of the silicon and manganese and part of the phosphorus are oxidised, and after the iron has melted, the oxidation of the impurities con- tinues. In order to facihtate the oxidation of the carbon and the remainder of the phosphorus, the iron is stirred with puddlers and gathered into large balls, known as " blooms." These puddle balls are withdrawn from the furnace and hammered under a steam hammer to expel the slag. In the modern pud- dling process, although some of the oxidation of the impurities is effected by the oxygen of the atmosphere, much of it arises from interaction with magnetic oxide, FcaOj, derived from the lining. The impurities enter into the slag after oxidation in the form of silicate, phosphate, etc. Wrought iron melts at about 1,500-1,550° but it softens suffi- ciently for welding in the neighbourhood of 1,000°. Wrought iron is tough, malleable and fibrous in structure. It can be roUed into plates or drawn into wire. A good wrought iron should contain not more than 2 per cent, of carbon and should be practically free of phosphorus. Steel. Steel is a variety of iron containing from 01-1 '5 per cent, of carbon. Soft steel approximates to wrought iron in composi- tion, averaging from 01-0-2 per cent, of carbon, whilst hard tool steel may contain from 0-9-1 5 per cent, of this element, A medium steel runs about 0-5 per cent, carbon. The properties of steel vary closely with the carbon content. Steel rich in carbon is hard and brittle, and has certain of the characteristics of cast iron, whilst a soft steel is not far removed in properties from wrought iron. Metallurgy. — Five processes are in use for the preparation of steel : 628 AN INORGANIC CHEMISTRY (1) The cementation process; (2) the crucible process; (3) the Bessemer process ; (4) the open hearth process ; (5) the electric furnace process. The Cementation Process. — This method, which is, however, gradually being displaced, is based upon the diffusion of carbon through iron at high temperatures. Bars of wrought iron are packed in wood charcoal and the whole heated for about ten to eleven days at a temperature of 1,000-1,100°. The furnace is then slowly cooled off (fourteen days), so that the complete process takes nearly four weeks. During the cementation the carbon enters the outer layer of the iron and gradually diffuses within. Such steel, known as blister steel, because of its bUs- tered appearance, is never homogeneous. Much of it is used for crucible steel, but some of it is heated and rolled into bars. This finds appUcation as spring steel and as shear steel, a high grade tool steel. The diffusion of carbon into \^Tought iron is similar to the diffusion of one metal into another. Cylinders of gold and lead have been clamped together for a considerable period of time, and it has been shown that there has been a slow, but sure, diffusion of one metal into the other. The diffusion is, of course, excessively slow, as is to be expected in the case of the metals, but in the cementation process a great saving of time is effected by keeping the materials in the neighbourhood of the softening point of iron. Crucible Process. — Steel, made by the cementation process, always contains more or less enclosures of slag, whilst much of it, from the very nature of its preparation, is not homogeneous. These difficulties are overcome by bringing the steel into a state of fusion in large crucibles. The best crucible steel is made from bHster steel, carefully selected so that a steel of the desired composition and properties results. Often scrap steel is added with carbon, whilst in other works pig iron is broken down by fusion with high grade wrought iron. Crucible steel is often of the highest quahty, suitable for the manufacture of razors, surgical instruments, etc. Moreover, liy the addition of requisite quantities of other metals such as molybdenum, tungsten, chromium, etc., the manufacturer is able to impart any desired hardness or toughness to the metal. Special tool steels which have the remarkable property of maintaining their hardness even at a red h^at are now of great technical importance, IRON 629 Bessemer Process. — In this process the impurities present in cast iron are removed by a blast of air forced through the molten mass. The manganese and silicon are oxidised and enter into the slag. The carbon is oxidised to carbon monoxide, in which form it burns at the mouth of the "converter" (Fig. 135). When the carbon is completely oxidised, as is shown by the marked decrease in the size of the carbon monoxide flame at the mouth of the converter, the converter is tilted over, the blast turned off, and there is added the necessary weight of molten spiegeleisen —an iron, manganese aUoy containing a known percentage of carbon — in order to give a steel of the desired composition. If the pig iron is free from phosphorus, the furnace may have an acid lining, but if the phosphorus is present to a greater extent than 005 per cent., a basic hning is added. As such, dolomite (calcium magnesium carbonate) is used. This is thoroughly calcined, mixed with tar, and apphed to the surf ace of the converter. Lime is also added during the oxidation. The oxides of phosphorus are thereby converted into calcium phosphate, which is ground and put upon the market as a fertiUser, Thomas' Bask. Slag. After the re- moval of all impurities by oxidation, the desired quantity of spiegeleisen is introduced into the furnace. The Open Hearth Procass.— The great advantage of this pro- cess is that a wide range of materials may be treated. Like the Bessemer process, it can be used for the treatment of pig iron, which is very low in phosphorus, in which case the lining is acid, and for phosphorus-rich pig iron (basic lining of the open hearth) ; but the open hearth method has a still greater flexibihty inas- much as pig iron, containing a moderate amount of phos- phorus (005-0-5 per cent.) can be successfuUy treated by this Fig. 135. 630 AN INORGANIC CHEMISTRY process. The furnace is of the reverberatory type, the oxidation being effected by means of air blown over the surface of the molten metal. In order to secure the requisite high temperature, both the gas fuel and the air are heated before they are brought into the furnace. This pre-heating is efiected by using up the waste heat of the furnace gases before they are allowed to escape. A typical open-hearth reverberatory furnace is shown in Fig. 136. This method of preheating necessitates the use of two pairs of chambers filled with bricks, each pair being alternately used for coohng the furnace gases and for heating up the gas fuel and the air. In the acid open hearth process, where practically no phos- phorus is present, the charge consists of pig iron and scrap. The manganese and siUcon are oxidised first, but the carbon is not attacked until practically the whole of the other impuri- ties has been re- moved. Haematite is often added to assist the oxidation of the carbon. When the amount of the carbon has fallen to the desired extent, the charge is tapped. In the basic open hearth process, phosphorus is oxidised at the same time as the carbon. The oxidation is facUitated by the addition of haematite ore. After the oxidation of the impurities has been completed, spiegeleisen is added. This process is used largely for the manufacture of mUd steel. The Electric Furnace Process. — This process is coming increas- ingly to the fore, especially in America, for the manufacture of car- bon steels and alloy steels of a high grade. The great advantage appears to lie in the high temperature which can be attained. The highly basic slag is kept thoroughly fluid, and as a result, the per- centage of sulphur and phosphorus falls so low as to be negligible. Chemical Properties. — Iron dissolves freely in dilute acids. If pure, iron hberates hydrogen from dilute sulphuric and hydro- Entrance of hot producer gas. Exit flue $ases to coolinScham- bers. Fig. 136. IRON 631 chloric acids, but the action with nitric acid yields ammonia. With concentrated nitric acid, the iron becomes " passive " (cf . chromium). Such passive iron no longer displaces hydrogen from acids, nor is copper displaced from copper sulphate solu- tions. Such passivity can be removed by scratching or by treatment with a reducing agent. It is, however, not yet con- clusively proved that the passivity is due to a thin film of adhering oxide, as was once thought. Concentrated sulphuric and hydro- chloric acids produce small quantities of hydrocarbons when they act upon steels, This is due to a reaction with the carbide present in the steel (ef. calcium carbide, aluminium carbide). When iron is exposed to a moist atmosphere, it soon becomes covered with a loose film of reddish brown oxide, known as rust. Although the phenomenon of rusting was, until comparatively recent times, considered to be a direct oxidation of the iron by means of the atmosphere, investigation has revealed that the process is by no means so simple. Amongst the facts wiiich have been accumulated upon the subject of corrosion or rusting of iron, the following stand out : (a) Dry air has no action upon iron at ordinary temperatures. (b) Air-free water has also no action at ordinary temperatures. (c) Pure water and pure oxygen together have no action upon pure iron. {d) The presence of an acid is not necessary for corrosion. (e) Ordinary samples of iron and steel rust in the presence of moist air. (/) Iron will not rust m the presence of pure water, pure oxygen and pure carbon dioxide. One of the most promising hypotheses put forward to explain the phenomenon of rusting lay in the assumption that carbon dioxide was the necessary agent in bringing about the corrosion. This is embodied in the equations : 2Pe -f 2H2O +O2 +4C02^^2Fe(HC03)2. 2¥e{B.CO3h + x-RJ0+O^¥e,O„(x + 2)11,0+4:002. A small quantity of carbon dioxide is therefore sufficient to carry on the corrosion as it is constantly regenerated. Unfor- tunately, however, this promising hypothesis does not explain the fact that iron will not rust in the presence of pure oxygen, water and carbon dioxide. Another theory, which carries a great deal of support, attributes the rusting to slight differences 632 AN INORGANIC CHEMISTRY in the composition of the iron whereby its solution pressure {q.v.) is afiected. Some parts of the iron become electro -negative to others, and small local currents are set up. The corrosion is, on this view, an electrolytic phenomenon. Finality does not yet appear to have been reached in the matter. The Tempering of Steel. — If steel is heated to redness, and then suddenly chUled by being plunged into cold water, the steel becomes so hard that it wiR scratch glass. If this hardened steel is again heated and allowed to cool under certain standard con- ditions, the hardness may be reduced to any desired extent. This is known as tempering the steel. The chilling to which the steel has been subjected in the process of hardening, " crystal- hses " the equiUbrium existing at the high temperature when the different modifications and compounds of iron present within the metal are characterised by extreme hardness. The chilling preserves this " false " equUibrium, and owing to the slowness with which solids revert to the stable phase, no notice- able change occurs in the hardness from day to day. If the steel is reheated and allowed to cool more or less slowly, changes take place in the constitution of the steel, and the condition of affairs pertaining to any particular temperature is approached, if not actually reached. Any desired degree of hardness may thus be imparted to the steel. Steel cooled slowly becomes comparatively soft. By a suitable adjustment of temperature and rate of cooUng the annealer can vary the physical properties of the metal within a wide range. The Chemistry of the Compounds of Irox The chemistry of iron centres round the oxides. Ferrous oxide, FeO . . Basic. io i Ferric oxide, FcoOj . . Amphoteric, but mainly basic. gg I Ferroso-ferric oxide, Fe:,0 , Mixed oxide. |o I (FeOa) Xot isolated, but gives com- 1 1 i pounds, ferrates, similar to "^ ^ 4' *he chromates. The lowest oxide is a moderately .strong base, giving rise to the ferrous salts ; ferric oxide is w eakly basic in the ferric salts, but also forms a few ferrites from the acid, HFeO-j. The tri- oxide has not been isolated, but a few derivatives of this acidic oxide are known — the ferrates. IRON 633 Ferrous Oxide and the Ferrous Salts Ferrous Oxide is formed as a black, crystalline powder, when carbon dioxide is reduced by iron, or when iron oxalate is broken down in the absence of air. It is easily reduced by hydrogen to the metal, or oxidised by the air to the sesquioxide, Fe^Oa. Ferrous hydroxide is precipitated as a white precipitate on the addition of a soluble hydroxide to a solution of a ferrous salt in the absence of air. In the presence of oxygen, immediate oxidation to a dirty green and finally brown hydrated oxide occurs. It dissolves in solutions of ammonium salts, being, Uke magnesium hydroxide, sufficiently soluble in water to require a fairly high concentration of 0H~ ions for its precipitation. The ferrous salts can be prepared by the solution of ferrous hydroxide in the requisite acid. These salts are strong reducing agents, and are only stable in the absence of oxygen. In this respect they show a strong resemblance to the compounds of divalent chromium, but not to the manganous salts which are stable under such conditions. The ferrous salts, derivedras they are from a moderately strong base, are only sUghtly hydrolysed. They exhibit properties characteristic of other similarly consti- tuted salts. So far as solubility is concerned no great difference is to be found between the ferrous salts and the corresponding salts of divalent copper, manganese, etc. Ferrous Chloride can be obtained by passing hydrogen chloride over iron filings, as well as by the action of hydrochloric acid upon the hydroxide or carbonate. The chloride can be obtained in colourless, shining scales which are very deliquescent. When heated in the air, it is oxidised to ferric chloride which volatilises, and ferric oxide remains. 12FeCl2 + 3O2 --> SFeClj + SFeaOa. In aqueous solution it is rapidly oxidised by the air. Several hydrates of ferrous chloride are known, e,g. FeCU.xHaO, where X is 1, 2, 4. or 6. Ferrous Sulphate. — This compound can be prepared by the following method : Fe + H2SO4 -> FeSOi + H^ as well as by the oxidation of iron pyrites in the air. 2FeS2,+ 7O2 + 2H,0 -^ 2FeS04 + 2H2SO1. 634 AN INORGANIC CHEMISTRY The liquor draining away contains ferrous sulphate and sulphuric acid. This is treated with scrap iron. This has the double effect of reducing any ferric sulphate as well as converting the sulphuric acid into ferrous sulphate. The neutral solution is concentrated until crystals of green vitriol, FeS04,7H20, separ- ate. This hydrate is isomorphous with the sulphates of zinc and magnesium, forming one of the vitriols which crystaUise in the rhombic system. This heptahydrate is dimorphous, as it also crystaUises in the monocHnic form. The latter appears to be the stable form at ordinary temperatures, for crystals of the rhombic system can only be obtained it the saturated solution is brought to crystallisation by means of a crystal of the sulphate of zinc or magnesium (rhombic). Crystals of ferrous sulphate oxidise readily to a basic ferric sulphate. Solutions of ferrous sulphate, as well as the sohd crystals, combine with nitric oxide to form a loose unstable compound with a deep brown colour, hence the use made of this reagent in testing for nitric acid. Ferrous sulphate forms a well defined series of double salts with the alkah sulphates, amongst them ferrous ammonium sulphate, Mohr's salt, FeS04,(NH4)2S04,6H20. The crystals of this salt are comparatively stable and even in solution the rate of oxida- tion is very shght (cf. stability of the alums, p. 217). Ferrous sulphate is used in the manufacture of iron mordants and Prussian blue [q.v.). The extract of nut-galls contains a considerable amount of tannic acid, which, with ferrous sulphate, forms ferrous tannate, a soluble and almost colourless salt. A solution of this salt, to which a little blue-black dye and some gum-arabic have been added, produces an ink. When this ink is exposed to the atmosphere, the ferrous salt is oxidised, and a fine black precipitate, consisting of ferric tannate, is formed. The presence of the dye is necessary to make the writing visible before the oxidation of the ferrous tannate has had time to be effected. Ferrous Sulphide. — ^This is a black substance of metallic appearance, obtained by the direct action of sulphur upon heated iron turnings. It is used extensively in the preparation of hydrogen sulphide. Ferrous sulphide can be made in the wet way by the addition of ammonium sulphide to the solution of a soluble ferrous salt. Ferrous sulphide is soluble in IRON 635 acids, hence its non-precipitation by the action of hydrogen sulphide. FeCla + H^S ;=± FeS + 2HC1. Ferrous Carbonate. — Besides occurring naturally, this sub- stance can be prepared by treating a cold solution of pure ferrous sulphate with sodium carbonate. A white precipitate is thrown down which soon turns brown on exposure to the air. It is appreciably soluble in the presence of carbon dioxide, owing to the formation of a soluble bicarbonate, hence its presence in many chalybeate spring waters. Ferric Oxide. Ferric Salts Ferric Oxide is a reddish brown powder, formed during the roasting of ferrous sulphate, or by the ignition of the hydroxide or a ferric salt containing a volatile acid. It is also found naturally as red haematite and as specular iron (steel grey in colour). The powdered red oxide is used for poUshing purposes as rouge, and also as a pigment, Venetian red. The corre- sponding hydroxide is thrown down by the addition of ammonium hydroxide to a solution of a ferric salt. As is to be expected from the increased oxygen content of this oxide, its behaviour as a base is characterised by considerable weakness. As a result of this, salts formed by the interaction of this base with an acid, e.g. Fe(0H)3 + 3HCI ^=± FeClj + SHaO, are subjected to considerable hydrolysis in aqueous solution. This hydrolysis can be effectively demonstrated by dialysing the solution of ferric chloride. The solution is put in a bell jar closed with a sheet of parchment or animal membrane. The beU jar is stood in a current of running water. Hydrochloric acid diffuses through the membrane and is carried away. The hydrolytic equilibrium defined in the equation FeCla + 3H0H ^z± Fe(0H3) + 3HC1 is thereby upset, and more hydrochloric acid is generated. In this way the whole of the hydrochloric acid may be spUt off and carried away. The ferric hydroxide does not appear as a brown flocculent precipitate, but remains in colloidal solution (q.v.). Not only does ferric oxide (hydroxide) function as a weak 636 AN INORGANIC CHEMiSTKY base but, if fused with an alkali hydroxide it takes on the properties of a weak acid ; in short, ferrites derived from the acid HFeOa or HaFejO, are formed. Sodium ferrite is produced by heating together ferric oxide and sodium carbonate to a bright redness. Other ferrites are zinc ferrite, ZnO.FcaOs; magnesium ferrite, MgOjFejOa, etc. Ferric Chloride. — Anhydrous ferric chloride, FeClg, is best obtained by heating iron powder or wire in a stream of dry chlorine. It is also produced by the action of hydrogen chloride upon hydrated ferric chloride at a dull red heat. Ferric chloride subhmes Ln dark scales. These crystals are dehquescent and dis- solve in water and in alcohol. From determinations of its vapour density at high temperatures it is concluded that the formula is FeCla. Hydrated ferric chloride may be obtained by the concentra- tion of solutions formed by dissolving iron in hydrochloric acid, the reaction being completed by the oxidation of the ferrous chloride by means of chlorine. Fe+2HCl^FeCU + H.. 2FeCl2 + CU->2FeCl3. Yellow, deliquescent crystals of the composition FeGl3,6H20 are obtained. Other hydrates are also known. Ferric chloride can be reduced by boiling with iron, 2Fe+ + ^+Fe-^3Fe+-^ or by such reducing agents as stannous chloride, hydrogen sulphide, sulphur dioxide, etc., 2FeCl3 + SOa + 2H20-$^2FeCl2 + 2HCI -f- B.SO, 2FeCl3 + SnCl,-^ 2FeCl, + SnClj. or written ionically, 2Fe <■ + + + Sn ' + — ^ 2Fe + ' + Sn + + + + Aqueous solutions of ferric chloride are generally tinted yellowish brown, owing to the presence of colloidal ferric hydroxide. The addition of an excess of hydrochloric acid renders the solution nearly colourless by preventing the hydrolysis. Iron Bisulphide. — Iron disulphide occurs naturally in large quantities as the mineral, FeSa. B is dimorphic, being found not only as glittering yellow, golden cubes, and octahedra IRON 637 (pyrites), but also as the less stable marcasite (rhombic). The compound can also be prepared artificially by heating the lower sulphide with sulphur. It is not attacked by dilute hydro- chloric or sulphuric acids, but is readily dissolved by nitric acid with the precipitation of sulphur. Ferric Sulphide, FcgSg, is formed as a yellow mass either by fusing together the free elements or by oxidising FeS with sulphur. It is obtained as a black precipitate by the action of a soluble sulphide upon a solution of a ferrous salt. Ferric Sulphate. — This compound can be obtained by oxidismg ferrous sulphate with nitric acid in the presence of sulphuric acid. Owing to the weakness of the base Fe(0H)3, ferric sulphate is easily hydrolysed on boiling, with the precipi- tation of basic sulphates. If the requisite amount of potassium sulphate is added to a solution of ferric sulphate, violet octa- hedra of iron alum, K2S04Fe2(S04).-,,24H20, separate out. Ferroso-Ferric Oxide Ferroso-ferric oxide, FcsOj, occurs in large quantities as the mineral, magnetite. As its name implies, it has marked mag- netic properties. It is also produced by the action of steam and of air upon red-hot iron. This oxide is in no way a basic oxide, for it gives rise to no distinctive salts on being treated with an acid. Under the action of hydrochloric acid, it yields a mixture of ferrous and ferric chlorides. This and other similar reactions indicate that ferroso-ferric ox^de is a mixed oxide, FeOjFcaOs, similar to the oxides PbaO,), MuaOj. This oxide is therefore ferrous ferrite and is allied to calcium ferrite, CaO,Fe203. The Cyanides. — The industrial methods of preparing potas- sium ferrocyanide have already been treated (see p. 484). It remains to discuss its constitution and the constitution of its kindred salt, potassium ferricyanide. If a solution of potassium cyanide is added slowly to a solution of ferrous chloride, the first reaction is the precipitation of ferrous cyanide, which, however, immediately dissolves in an excess of the reagent. FeCla -f 2KCN-^Fe(CN)2 + 2KC1 Fe(0N)2 + 4KCN->K,[Fe(CN)e] Potassiwn ferrocyanide, 638 AN INORGANIC CHEMISTRY Potassium ferrocyanide is essentially a ferrous salt, sometimes written 4KCN,re(CN)2, hence the terminology adopted — ferrocyanide. On the other hand, if potassium cyanide is added to a solution of a ferric salt, the reactions will be indicated in the following equations : FeCla +3KCN-^Fe(CN)3 +3KC1 Fe(CN)3 +3KCN-^K3[Fe(CN)fi]. This salt is derived from ferric chloride, and is known as a ferricyanide. The relation of the ferrocyanide to the ferricyanide is shown in the following reaction, 2K,[Fe(CN)„]i^' + Cl2^2K3[Fe(CN)e]™ + 2KC1. The valence of the ferrocyanide group towards hydrogen (or its equivalent, potassium), has been reduced from four to three, i.e. it has been oxidised by the action of the chlorine. The same effect could have been produced by first oxidising ferrous chloride with chlorine and then adding the requisite amount of potassium cyanide. If a solution of potassium ferrocyanide is added to a solution of a ferric salt, a deep blue precipitate, known as Prussian blue, is thrown down. 4Fei"Cl, + 3K4[Fe(CN)e]iV-^Fe4i"[Fe(CN)e]3i^'+12KCl Ferric ferrocyanide. Solutions of potassium ferricyanide, when added to soluble ferric salts, produce only a shght brown coloration. If potassium ferricyanide is added to a ferrous salt, TumbuU's blue is thrown down. It is supposed that the reaction 3Fe"Cl, + 2K3[Fe(CN)e]i"-^Fe3"(Fe(CN)e]2"i + 6KC1 Ferrous ferricyanide. occurs, though it is also claimed that this is immediately con- verted into Prussian blue by the oxygen of the air. Potassium ferricj^anide, as prepared by the oxidation of the ferrocyanide by such oxidising agents as chlorine, nitric acid, etc., forms red monoclinic prisms. Ferricyanio acid is obtained by decomposing lead ferricyanide with dilute sulphuric acid, but it is unstable. Neither potassium ferro- nor ferri- cyanide gives a precipitate with ammonium hydroxide, nor any other reaction of the ions COBALT 639 Fe ' +or Fe + + ' . This shows the great stabiUty of these complex salts and is an indication of the extent to which the equilibrium Fe(GN)2 + 4KCN ^=± K4[Fe(CN)e] is pushed to the right. The absence of all reactions of iron in such complex salts as the ferro- and ferri- cyanides is in striking contrast to the behaviour of the alums, which give all the reactions of the metals contained therein, e.g. potassium alu- minium sulphate gives all the reactions both of potassium and of aluminium. It is because of this differential behaviour that such salts as the alums are generally classed as double salts, though, as already indicated, the difference between double and complex salts is purely one of degree. Iron Carbonyl. — When carbon monoxide is led over finely divided iron at 80", a pale yeUow Hquid is formed (B.P. 102-5° at 760 mm.). Above 180° it decomposes rapidly, forming a mirror of iron. Its molecular weight in benzene, as well as its vapour density, indicates that its formula is Fe(C0)5, iron penta- carbonyl. Iron tetra-carbonyl Fe(C0)4, and a hepta-car- bonyl Fe(CO)„ of similar properties have also been prepared. In aU cases the stabihty is much less than in the case of nickel carbonyl (q.v.). Ferrates. — Although the oxide FeOa has not been isolated, a few salts derived from this acidic oxide have been prepared. If iron is used as an anode in a concentrated solution of potassium or sodium hydroxide, the solution becomes almost black, and on the addition of more hydroxide, a reddish precipitate of potassium ferrate may be obtained. Solutions of potassium ferrate are unstable, and break down on boiHng into the ferrite with the evolution of oxygen. A similar compound has been prepared by passing chlorine through a solution of potassium hydroxide in which ferric hydroxide is suspended. A satis- factory method of obtaining sodium ferrate is by the fusion of iron fihngs or iron oxide with sodium peroxide. The ferrates of barium, strontium and calcium have been prepared by double decomposition. Cobalt Occurrence and Metallurgy. — Cobalt occurs in cobaltite CoAsS, smaltite CoAsa, and glance-cobalt (CoFe)AsS. The 640 AN INORGANIC CHEMISTRY arsenical ores are first worked up into the oxide and this is re- duced by the Goldschmidt method or with carbon. The actual manner of working up the ore into the oxide varies. Amongst the methods used may be mentioned roasting with salt and extracting with hydrochloric acid. The solution of the chloride is then precipitated. The process used for the separation of cobalt from other metals, e.g. bismuth, copper, nickel, varies from works to works. Cobalt is refined electrolytically. It has no commercial apph- cations. Pure cobalt is malleable, and tough ; it is slowly oxidised when heated in the air. It dissolves readily in acids, and at a red heat it decomposes steam. The Chemistry of the Compounds of Cobalt Cobalt forms three oxides, cobaltous oxide CoO, cobaltic oxide C02O3, and cobalto-cobaltic oxide C03O4. Of these the first is fairly strongly basic, comparable with cupric oxide, ferrous oxide and manganous oxide. Salts derived from cobalt- ous oxide are stable in air (contrast ferrous salts), and in other respects show no pronounced- differences from similarly con- stituted salts of other divalent metals. Cobaltic oxide is a very feeble base, its salts being strongly hydrolysed. Moreover, aqueous solutions of cobaltic salts are unstable, and in the presence of platinum they break down into cobaltous salts with the evolution of oxygen. Cobalto-cobaltic oxide is exactly similar in behaviour to its analogues, FejOi, MuaOj, Pb304. It is obtained by the strong heating of the other oxides. Cobaltous Oxide and the Cobaltous Salts. — Cobaltous oxide is prepared by the reduction of the higher oxides in a stream of hydrogen at a temperature not exceeding 350°. The corre- sponding hydroxide is precipitated by the addition of sodium hydroxide to a solution of a cobaltous salt. Cobalt chloride can be obtained in the anhydrous form as blue, crystalline scales by the ignition of finely divided cobalt in a stream of chlorine. Several hydrates of this salt are known. Dilute aqueous solutions of this salt are a pale pink in colour, but concentrated solutions are a deep blue. The same colour is given by dilute solutions if the temperature is raised. So, also, the addition of a strong solution of potassium chloride to a COBALT 641 moderately strong, but pink solution of cobalt chloride, brings out the characteristic deep blue colour. There seems Httle doubt that the addition of potassium chloride to a solution of cobalt chloride causes the formation of a blue complex salt, e.g. C0CI2 + 2KC1 ;:z^ K2[CoCl4] C0+++2CI- K + + C1- 2K++C0CI4- Pink. Blue. In strong solutions of cobalt chloride auto -complex formation occurs with the production of a similar complex salt of cobalt, Co + + + 2C1- + CoCL ;=± Co[CoCl4] Co + Pink. C0CI4" Deep blue. the predominating colour being a deep blue. A rise of tempera- ture drives this equiUbrium to the right and intensifies the blue colour. The ease with which the faint pink of the hydrated cobalt chloride is converted into a deep blue by a rise of temperature is responsible for the use of this material as a " sympathetic " or invisible ink. Cobalt Sulphide. — Cobalt sulphide is of interest as it is not precipitated by hydrogen sulphide in the presence of an acid, but, when once it has been thrown out of solution, dilute acids have but little action upon it. Apparently, there are two modi- fications, one of which, the less stable, is soluble in acids, the other almost insoluble. In conformity with the usual rule, the unstable variety would be formed first on precipitation, afterwards turning into the more stable, but less soluble form. The presence of the acid would inhibit the initial formation of the unstable sulphide. Smalt. — Smalt is prepared by roasting a fairly pure cobalt ore, free from iron and sulphur. The oxide is then fused with sand and potassium nitrate in large earthenware pots. A cobalt silicate of deep blue colour is formed. This is ground and put on the market as a pigment for painting china, etc. TT 642 AN INORGANIC UJli<;iViii5xm Cobaltic Oxide, Cobaltic Salts.— Cobaltic oxide is prepared by the gentle ignition of cobalt nitrate. The hydroxide is obtained by the precipitation of a cobaltous salt by means of sodium hypochlorite. It is a black powder which is decomposed by hydrochloric acid with the evolution of chlorine. The only important cobaltic salt which has been isolated is cobaltic sul- phate. This has been obtained by the electrolysis of cobaltous sulphate in a platinum dish which serves as anode, the cathode consisting of a platinum wire immersed in a porous ceU contain- ing dilute sulphuric acid. The aqueous solution of cobalt sulphate slowly evolves oxygen, especially in the presence of platinum. Cobalt alums, M2S04,Co2(S04)3,24H20, isomorphous with the alums of Fe™, Mn™, Cr"', etc., have been described. These are stable in dry air, but they decompose rapidly in aqueous solution with the reduction of the cobalt to the divalent state. Complex Cyanides. — The constitution of the complex cyanides of cobalt is similar to that of the iron complex cyanides. The addition of potassium cyanide to a solution of cobaltous chloride produces potassium cobaltocyanide, C0CI2 + 6KCN^K4[Co(CN)6] + 2KC1. cf. FeCl2 + 6KCN->K4[Fe(CN)e] -|-2KC1. Chlorine, or even oxygen, oxidises the cobaltocyanide to the cobalticyanide, 2K4Co(CN)e + CI2 -^ 2K3[Co(CN)6] + 2KC1. There is, however, one important difference between the complex cyanides of cobalt and of iron : whereas in the case of iron the -ous cyanide is the more stable, the converse is found to be the case with the cobalt complex cyanides. The addition of an acid to a solution of potassium cobalticyanide precipitates the corre- sponding acid. This acid, as well as its salts, give none of the reactions of cobalt. Other Complex Cobalt Salts. — If ammonium hydroxide be added to a solution of a cobalt salt, the precipitate of cobalt hydroxide, at first thrown down, soon redissolves with the forma- tion of a complex salt, in which the ammonia forms part of the complex cathion (COjKNHs)''' ■*■ (cf. zinc, copper, nickel, cadmium). Potassium cobalti nit rife is obtained as a yellow precipitate NICKEL 643 when a solution of a cobaltous salt, acidified with acetic acid, is treated with a solution of potassium nitrite. The precipitate has the composition K3Co(N02)6. Nickel Occurrence. — Nickel is found in nature as nickel pyrites, NiS ; nickel glance, NiAsS ; nicolite (kupernickel), NiAs ; pentlandite, NiS, 2FeS ; and garnierite, a hydrated silicate of nickel and magnesium. Metallurgy. — Large quantities of nickel are obtained from the sulphide ores round Ontario. These are dressed, roasted and then smelted in blast or reverberatory furnaces in order to produce a matte, consisting largely of the sulphides of nickel, iron and copper. This is then oxidised in a converter similar to that used in the Bessemer steel process. In this way practically the whole of the iron is removed. The resulting mass consists of the sulphides of nickel and copper. This nickel matte is refined, either by repeated fusion with sodium sulphate and coal, or by the Mond process. The latter process is based upon the alternate formation and decomposition of nickel carbonyl, Ni(C0)4 (p. 358). This compound is formed at 50°-80°, and broken down into nickel and carbon monoxide at a higher temperature. Very pure nickel is obtained in this way. Properties. — ^Nickel is a white, hard metal which is capable of taking a high polish. It dissolves freely in dilute nitric acid, but hydrochloric and sulphuric acids have Uttle action upon it. Large quantities of nickel are used in making alloys for coinage purposes, etc., as weU as for nickel-plating. In this case the bath consists of a solution of ammonium nickel sulphate to which an excess of ammonium hydroxide has been added. Finely divided nickel has a marked catalytic action. It facihtates the reduction of ethylene to ethane by means of hydrogen, carbon monoxide to methane, etc. Commercial use is now being made of this property in the hydrogenation (reduc- tion) of oils. Linseed and cottonseed oils are now reduced by having nickel powder suspended in them and hydrogen blown through them. The metal assists the oil to combine with the hydrogen, yielding a fully saturated oil. 644 AN INORGAJVIO UllJ^;iViii5xis,r Chemistry of the Compottnds of Nickel Three oxides of nickel are known — nickelous oxide, NiO, nickelic oxide, ^263, and mokelo-nickeKc oxide, Ni304. Of these the nickelous oxide alone gives rise to a definite series of salts, all of which are divalent. Little need be said of the salts of nickel. They resemble the salts of cobalt to an unusual extent, differing only in colour and in the formation of the double or complex cyanides. If potassium cyanide is added to a solution of a salt of nickel, the following reactions occur, 2KCN + NiCl2->Ni(CN)2 + 2KC1 Ni(CN)2 + 2KCN->K2[Ni(CN)4] The potassium nickelocyanide is therefore differently con- stituted from its analogue, potassium cobaltocyanide, K4Co(CN)6. There is, however, another important difierence ; whereas the reaction 4KCN +Co(CN), ^=± K4[Co(CN)6] 4K + + 4CN -C0 + ++2CN- 4K^ [Co(CN),F is so complete that no measurable amount of cobalt ion is left in the solution, in the corresponding reaction leading to the formation of potassium nickelocyanide the equUibrium does not lie completely to the right, so that an appreciable concentra- tion of nickel ions are left in solution. Consequently, when sodium hypobromite or hypochlorite is added to a solution of potassium nickelocyanide, there are sufficient Ni "^ "^ ions capable of being oxidised to Ni ■*■ "^ '^ to ensure that the solubihty product of Ni(0H)3 will be exceeded, and this hydroxide separates from the solution. The solubihty product of cobaltic hydroxide is not exceeded and a complete separation can be effected. So far as the cobalt is concerned, the only effect is to oxidise the cobaltocyanide to the more stable cobalticyanide. The appear- ance of nickehc hydroxide instead of nickelous hydroxide is, of course, due to the oxidising action of the hypochlorite. Nickel- ous salts give rise to complex ammonia cathions (NijxNHs) + "*" on treatment with an excess of ammonium hydroxide (cf. Zn, Co, Cu). Nickelo-nickelic oxide, Ni304, is formed either by prolonged NICKEL 645 heating of nickel chloride in a stream of moist oxygen, or by fusing metallic nickel with sodium peroxide. On heating to redness decomposition into the monoxide results. Nickelic oxide, ^263, is reputed to be formed by the ignition of the carbonate or nitrate in the air. With hydrochloric acid chlorine is evolved, with sulphuric acid oxygen. Recent work has thrown doubt upon the chemical identity of this substance. It is claimed that NiaOa is in reahty an intimate mixture of the monoxide and a dioxide, Ni02. The oxide NijOi would, on this view, be nickelous nickelite, 2NiO,Ni02. The matter has not been definitely cleared up. Questions ' 1. Discuss the composition and action of tKe gases present in the various portions of an iron-smelting blast-furnace. 2. Compare and contrast the chromates, manganates and ferrates. 3. Discuss the composition and constitution of the complex cyanides of iron, cobalt and nickel. 4. An excess of potassium cyanide is added to a solution containing the chlorides of nickel and cobalt. Sodium hydroxide and bromine are added and the solution warmed. A precipitate of nickelic hydroxide alone is thrown down. Explain this. 5. How would you prepare the following from metallic iron : (a) ferric sulphate, (6) anhydrous ferric chloride, (c) ferrous oxide, {d) ferrous chloride. 6. Compare the Bessemer and the open-hearth processes for manu- facturing steel. 7. Tabulate the oxides of iron, cobalt and nickel, and indicate what types of compounds, if any, are derived from them. CHAPTER XL RUTHENIUM, RHODIUM, PALLADIUM, OSMIUM, IRIDIUM, PLATINUM These elements fall into the last group of the Periodic Table, immediately below iron, cobalt and nickel. The first three elements, ruthenium, rhodium and palladium form the bridge elements in the second of the long periods, the remaining three the bridge elements of the third long period. The chief deposits of these valuable metals are the Ural Mountains, CaHfornia, AustraUa and Brazil. As might be expected from the noble nature of these elements, they are always found in the uncombined state. The elements of the first triad — ruthenium, rhodium and palladium — are more fusible and more easily brought into combination with oxygen than are their analogues. Ruthenium and Osmium. — Both these elements unite fairly readily with oxygen. They form the following oxides : Ruthenium. Osmium. Sesquioxide, RU2O3 Monoxide, OsO Dioxide, RUO2 Sesquioxide, OS2O3 Tetroxide, RuOj Dioxide, OSO2 Tetroxide, OSO4 Most of the salts of ruthenium are derived from the sesquioxide, e.g. RUCI3, RU2S3, a few from the dioxide, e.g. ruthenium sulphate Ru(S04)2. Ruthenates (cf. ferrates) and per-ruthenates are known, the former being derived from the unknown oxide, RuOa, the latter from the unknown RuaO,. Potassium ruthen- ate, K2Ru04,H20, is obtained by igniting a mixture of ruthen- ium, potassium hydroxide and nitrate. Greenish metafile crystals of the ruthenate can be dissolved out. Osmium tetroxide, " osmic acid " as it is often called, is used 646 RUTHENIUM, OSMIUM 647 in microscopic work for staining purposes. This arises from the ease with which it can be reduced to the metallic state. The oxide is formed by direct oxidation or by heating the metal in steam. It subUmes in gUstening, needle-Uke crystals. The oxide dissolves in water, but the solution does not give an acid reaction. The physiological action of this substance is severe, the lungs and the eyes aUke being attacked by its vapour. Osmium forms a few salts in which it is divalent, e.g. OsCU, some in which it is trivalent, OsCla, and a few in which it is tetravalent, OSI4. Osmates are also formed from the unknown oxide OsOa. Potassium osmate, K20s04,2H20, is thrown out by the addition of alcohol to a solution of the tetroxide in potassium hydroxide. Rhodium and Iridium. — These elements are characterised by their extreme reluctance to enter into combination with acids, even aqua regia having no action unless the metal is in a fine state of division. The metals, when in a finely divided state, are attacked by chlorine at a dull red heat. The metals are hard, iridium being often employed to alloy with platinum in order to increase its hardness. The more important oxides are the sesquioxides, Rh203, Ir203, and the dioxides, Rh02 and IrOa. The most important salts are those derived from the sesqui- oxides. From rhodium such salts as RhCls, RhaSs, Rh2 (804)3, etc., have been prepared. Iridium not only forms such salts, but it also gives rise to a considerable number of double or com- plex salts, e.g. K3lrCl6,3H20. The general formula of such salts is 3MCl,IrCl3, where M denotes K, Na, NH4, Tl, Ag, Ru, Cs. Iridium dioxide also forms a few salts, e.g. IrCU. Iridium tetrachloride displays the same tendency to unite with neutral groups of the alkaU halides as does the trichloride. The general formula of these double salts is 2MCl,IrCl4, where M denotes K,Na,NH4,Ru,Cs,Ag,Tl. None of the oxides of these elements displays any acidic tendency. Palladium. — Palladium differs from its analogue, platinum, in the ease with which it can be brought into solution. Concen- trated nitric acid attacks it freely, as does fairly concentrated sulphuric acid. Palladium gives three oxides, the monoxide, PdO, the dioxide, PdOj, and a sesquioxode, PdaOa. Of these the monoxide and dioxide are important, as they give rise to 648 AN INOiUjAiNiu (-/lULmioj-i^i the pallodous and palladic salts, though a few salts from the sesquioxide are also known. The paUadous salts, e.g. PdClj, PdSOi are stable, but on ignition they break down, leaving a residue of the metal. PaUadous chloride dissolves readUy in an excess of potassium chloride, and from the solution crystals of the composition K2PdCl4 can be obtained. Palladic salts, e.g. PdCl,, are unstable, but can be converted into stable compounds by coupling them with alkaU salts. From palladic chloride one can obtain salts of the formula MaPdCle, where M denotes K, Rb, Cs, NH4, etc. The facUity with which palladium absorbs hydrogen has already been commented on (p. 827). Although the view was long held that a definite chemical compound (hydride), PdaH, was formed, physico-chemical evidence upon this subject fails to substantiate such a view. The absorption appears to be nothing but a solubility effect. Palladium, charged with hydrogen, is a very active reducing agent. It will reduce ferric salts to ferrous, potassium ferricyanide to ferrocyanide, chromate to chromous salt, chlorine to hydrogen chloride, etc. Platinum Metallurgy. — The crude platinum is digested with aqua regia, the platinum and its aUied elements being thereby taken into solution. After filtration, the filtrate is evaporated to dryness with the addition of hydrochloric acid. The aqueous extract containing the chlorides is then treated with ammonium chloride and the insoluble ammonium chloroplatinate separates from the solution. This is ignited in a muffle and the platinum remains. (NH4)2PtCl6->Pt + 2CI2 +2NH4CI. Properties. — Platinum is a silver- white metal. Its ductihty and malleability enable it to be drawn out into fine wire or beaten into the foil. On account of its shght chemical activity it is used freely for the purpose of making crucibles and in many electrical instruments. Amongst the few elements which act upon it are nascent chlorine, carbon and phosphorus. A phos- phate under reducing conditions attacks platinum sufficiently to render it dangerously brittle. Platinum is attacked by hot, concentrated sulphuric acid, as well as by fused alkalies. PLATINUM 649 Platinum has the property of catalysing many chemical reactions, especially when in a fine state of division. The catalytic activity of colloidal platinum, prepared by the method of sparking under water (q.v.), is easily affected by traces of poisonous gases, e.g. hydrogen sulphide, hydrogen cyanide, carbon monoxide. Some of the reactions which are catalysed by platinum are, H2+Br2-^2HBr 2H2+02->2H20 2SO2 +02^-2803 A very active form of platinum may be obtained by precipitating the metal from its solution by such reducing agents as an alkaline solution of formaldehyde, or of glucose. There seems little doubt that the catalysis is brought about rather by physical than chemical means, i.e. the platinum does not enter into combination with the reacting substances, but aids the reaction by bringing the reacting substances together in a highly concen- trated form upon its surface (cf . Chap. XIV to see the effect of concentration upon the velocity of a chemical reaction). Compounds of Platinum. — Platinum forms numerous oxides of which the most important are the monoxide, PtO, the dioxide, Pt02, and the trioxide, PtOs. The monoxide and dioxide give rise to a well-defined series of salts. Platinous oxide is obtained by the addition of potassium hydroxide to a solution of a platinous salt. It is a black powder which dissolves readily in hydrochloric and sulphuric acids. Strong heating causes decomposition into the metal. It is capable of reducing hydrogen peroxide and acidified potassium permanganate. PtO + H2O2 -> Pt02 + H2O. Platinous chloride may be prepared by the direct action of chlorine upon the metal at 360°, as well as by heating the tetra- chloride to a temperature of 300°-360°. Platinous chloride is insoluble in water, but dissolves freely in hydrochloric acid, yielding chloroplatinous acid, HaPtClj. Many salts, derived from this acid, have been isolated. Potassium platinochloride or chloroplatinite KsPtCli, is obtained by reducing chloroplatinic acid, HzPtCls, with sulphur dioxide. Potassium chloride is then added and crystals of the platinochloride separate from the 650 AN INORGANIC UHJiiVli»iii.x solution. This salt is used in making platinum prints. A complex cyanide, similar in constitution to the chloride, has been prepared. Potassium platinocyanide, K2Pt(CN)4,3H20, is made by dissolving ammonium platinicyanide in a boihng concentrated solution of potassium cyanide to which potassium hydroxide has been added. After the evolution of ammonia has ceased, the solution is evaporated until crystaUisation sets in. It is also formed when spongy platinum is boiled in a solution of potassium cyanide. It forms long, yellow needles. Both potassium and barium platinocyanide show strong fluorescence in the solid state, especially under the action of the radium rays and X-rays. If a hot solution of this salt be saturated with chlorine, colourless crystals of a chloroplatinicyanide, K2Pt(CN)4Cl2,2H20, may be obtained on evaporation. Hydrated platinum dioxide, Pt02,xH20, is obtained from the decomposition of platinio chloride by means of sodium hydr- oxide. The residue is treated with dilute acetic acid to dissolve out the alkali. The oxide is amphoteric and gives rise not only to such salts as PtCI,, Pt(S0i)2, but also to the platinates. Plalinic chloride is best prepared by heating chloroplatinic acid, HzPtClc, in a current of dry chlorine at 300°-360°, or in a current of hydrogen chloride at 165° It is moderately soluble in water, but easily soluble in the presence of hydrochloric acid, forming chloroplatinic acid. PtCl4 + 2HC1^ HjPtCle. If platinic chloride is dissolved in water, the solution has an acid reaction and will decompose carbonates. This is supposed to be due to the formation of the acid, H2[PtCl4(OH)2], electrolysis reveahng that the platinum is part of a complex anion. Chloroplatinic acid, H2PtCl6,6H20, is obtained by dissolving platinum in aqua regia and evaporating with hydrochloric acid until the whole of the nitric acid is expelled. It forms brownish- red deUquescent prisms. Numerous salts of the type M2PtCl|;, where M is a monovalent metal, have been prepared. Potassium and ammonium chloroplatinates are remarkable for their com- parative insolubihty. Platinic sulphate, Pt(S04)2, is obtained by dissolving platinum sponge in concentrated sulphuric acid at 350°. Hydrated platinic oxide dissolves in an excess of potassium hydroxide, forming potassium platinate. PLATINUM 651 Pt(0H)4 +2KOH^K2Pt(OH)6. Sodium and potassium platinates have been prepared. Their constitution is represented by the formula M2[Pt(0H)(i] rather than by MsPtOajSHaO. These platinates are isomorphous with the stannates. CHAPTER XLI RADIO-ACTIVITY— THE CONSTITUTION OF MATTER Electric Discharge in a Gas at Low Pressure. — If a vacuum tube, containing a gas, be connected to a good pump, and the aluminium electrodes joined to the terminals of an induction coil, some very beautiful and interesting results may be observed. At first, a luminosity is seen to shoot from the elec- trodes, somewhat in the nature of a forked brush-hke discharge, Fig. 137. Fig. 138. and as the pressure of the gas is reduced, the dark space separ- ating the two luminous areas steadily shrinks, until practically the whole tube is filled with the glow. As the pressure falls still further, the cathode glow is seen to consist of two parts — the cathode glow surrounding the cathode with a velvety light and the negative glow, separated from the cathode glow by the so-called Crookes' Dark Space. The positive glow soon begins to break up into flickering striae. As the pressure is still further roduced (below 05 mm.) the negative glow extends, and with it, the Crookes' Dark Space, while the positive glow retreats. Ultimately, as the exhaustion approaches completion, the Crookes' Dark Space extends until 652 KADIO-ACTIVITY 653 it fills the whole of the tube, while the glass walls begin to fluoresce brightly, the precise colour of the fluorescence depending upon the composition of the glass. With soda glass the fluorescence is distinctly bluish. As the gas within the tube becomes stUl further attenuated, a beam of bluish light is seen to leave the cathode, and, if the evacuation has been pushed sufficiently far, these rays strike upon the glass round the anode and cause vivid fluorescence. These are the Cathode Rays. Amongst the more important properties of the cathode rays may be mentioned the following : 1. The rays travel in straight lines normal to the surface of the cathode. An ob- ject interposed in the path of the cathode rays throws a shadow >.^^ ^^^ , v upon the fluoresc- „ ,„. ingwall(Pig. 139). 2. The rays have the property of passing through matter, provided the thickness does not exceed a certain critical thick- ness. Lenard was able to coax the rays through an aluminium window, and experiment with them in the open. Such rays are known as Lenard Rays (q.v.). 3. The cathode rays raise the temperature of bodies upon which they fall. When focussed upon a metal, the rays raise the temperature to such an extent that the metal may melt, e.g. platinum becomes white hot. 4. The cathode rays can be deflected from their course either by the aid of a magnet or by means of an electric field. The deflection which is produced upon a beam of cathode rays by a known electrical field, is such as to leave no doubt that the cathode rays consist of negatively charged particles shot out from the cathode. The view put forward by some of the German school of phy- sicists was that the cathode rays were ether disturbances, but Crookes suggested that they were some form of radiant matter, projected from the electrode under the action of the intense electric forces. This view of the material nature of the cathode rays is now universally accepted. These negatively charged 654 AN INORGANIC CHEMISTRY particles shot out in this way from the cathode are commonly called negative electrons. Whenever a physicist is caUed upon to deal with matter carrying an electric charge, the ratio e/m, where e denotes the electrostatic charge carried by the mass m, becomes at once the subject of investigation. The experiments of J. J. Thomson and his school proved that e/m for the cathode rays had a value approximately 1,800 times as large as the ratio e/m for the hydrogen ion. Either, therefore, the charge carried by the negative electron is 1,800 times as large as the charge on the hydrogen ion, or the mass of the electron is only 1/1800 of the mass of the hydrogen atom. By methods which it would take us into the realms of pure physics to pursue, it has been indubit- ably established that e, the charge carried by the electron, is always the same as that carried by the hydrogen ion, i.e. 96,000 coulombs. The mass of the negative electron must consequently be only 1/1800 of the mass of the hydrogen atom. Moreover, it has been shown that the cathode rays are ejected from the electrode with a velocity from 2 to 10x10* kilom. per sec. (Ught travels at 30x10* kilom. per sec). Although the electrodes used in the vacuum tubes have been made of many different metals, the mass of the electrons gener- ated within the tube remains unaltered. Under the influence of X-rays as well as of intense heat, aU kinds of matter can be made to hurl forth a stream of negatively charged electrons, and we shall soon see that during radioactive disintegration, negative electrons are hurled forth with a velocity approaching that of light, yet, with few reservations, it may be stated that the mass of the electrons is always constant, 1/1800 of the hydrogen atom. This astounding result surely seems to point to the conclusion that under the play of intense electric (and radioactive) forces, these negatively charged electrons are shaken from or shot out of the atoms of the most diverse elements. These results receive their most logical explanation on the assumption that the nega- tive electron is a fundamental part of all matter ; as J. J. Thomson put it, " The explanation which seems to me to account for the facts in the most simple and straightforward way is founded on the view that the atoms of the different chemical elements are different aggregations of particles (electrons) of the same kind." These electrons apparently differ both in number and in arrangement for all the different chemical elements. RADIO-ACTIVITY 655 To Electroscope Lenard carried out some interesting experiments upon the power of the various elements to absorb the so-called Lenard rays. He found that, for all elements except hydrogen, the absorption of the Lenard rays is proportional to the weight of the metal or gas ; nothing else matters but the weight. Lenard deduced from these experiments that equal weights of different elements must contain an equal quantity of some common con- stituent. He also derived support to these conclusions from the observation that when Lenard rays are brought into an attenuated gas, some of them are deflected from their straight-Une path, due, he concluded, to the repulsion exerted upon a negatively cha*ged electron when it came within the sphere of influence of the negative electrons with the atoms of the attenuated gas. Radio-active Phenomena In 1896 Becquerel showed that crystals of uranium potassium sulphate were able to reduce the silver salts of a photographic plate, even though the plate be well wrapped up in black paper. This effect was given by aU the salts of uranium, whether dis- solved or in the sohd states. He also noted that when a salt of Earth uranium is brought into the neighbour- hood of a charged gold- leaf electroscope, the leaf slowly coUapsed. The Becquerel rays, be- sides affecting the photographic plates, must therefore possess the property of making the air in its neighbourhood a conductor, for how else could the insulation of the electroscope have been destroyed ? This property of ionis- ing gases is so strongly developed, and at the same time is so capable of exact and deUcate measurement that it has entirely superseded the old photographic method of investigating the nature of the Becquerel rays. Fig. 140 shows the manner in B To Cells — *- Fig. 140. 656 AN INOROAMC UlimMBTiiiS which such electrical measurements are made. The sub- stance to be investigated is spread over the plate B in the ionisation chamber. B is kept charged to a definite high potential by means of a high voltage battery. The forma- tion of charged ions is shown by the electroscope, connected to A. Shortly after, Mme. Curie (1898) showed that the compounds of thorium also possessed the power of generating Becquerel rays. Soon after this an examination was made of the minerals which contain thorium and uranium and the surprising discovery was made by Mme. Curie that some of the minerals were much more radioactive than either pure uranium or pure thorium ; for example, pitchblendes were four times as active as metallic uranium. This large activity would best be accounted for by the supposition that a hitherto unknown substance was present which possessed the property of radio-activity to an even greater extent than uranium itself. Mme. Curie then began the task of tracking and isolating this radioactive product. In her analytical separation she tested both precipitate and filtrate by means of the electroscope, and was thus able to ascertain whether the active material had been con- centrated in the precipitate or in the filtrate. In this way she was able to isolate two very active substances — one of these which was separated from bismuth she named Polonium, the other Radium. The radium salt proved exceedingly difficult to separate from its analogue, barium, but by repeated crystallisa- tions she was able to separate the less soluble radium bromide in a pure state. The chemical properties of the salts of radium are extremely like those of the alkaline earth elements, especially barium. The metal radium was obtained by the electrolysis of radium chloride with a mercury cathode. The amalgam was afterwards heated in a current of hydrogen in order to volatihse the mercury. The metal proved itself very reactive, dissolving freely in water and acid. The atomic weight of the element was obtained from the analysis of radium chloride. It was found to be 226. But besides the interesting chemical and physiological properties possessed by radium and its salts, such as their power to decom- pose water into hydrogen and oxygen, to incite phosphorescence of the diamond, of zinc sulphide and of the platinocyanides, these substances possess to an extraordinary degree the power 1 '^ 1^ Fig . 141. RADIO-ACTIVITY 657 of fogging a photographic plate and of causing the discharge of a gold-leaf electroscope. If the rays generated by radium bromide be subjected to the action of a magnet or of an electric field, the rays are sorted out into three distinct types (Fig. 141). They are : The Alpha Bays. — Under the action of the iljil electric forces these are iiiii "^-rays bent in such a way as to jijij show that they consist of l';ii; positively charged par- , i;i'i tides. These positively ^, ~~^~^Ju ijiii charged rays have been "«|., iJI '^-rays inve stigated by the methods which have already been discussed in connection with the ca- thode rays, and research has revealed that they consist of positively charged helium atoms, moving with a velocity about one- tenth that of light. These rays have Uttle penetrative power and are stopped by a few layers of paper or a few cm. of air. The Beta Bays. — These proved to be identical with the cathode rays of the vacuum tube. They are, in fact, negatively charged electrons shot out of the radium with a velocity approximating to that of Ught. These rays possess a far greater penetrating power than the alpha rays, for they can produce measur- able effects after passing through a thickness of paper or aluminium foil 100 times as thick as is required to stop the alpha rays. The Gamma Bays. — These are not deflected by the most powerful magnetic field. Their penetrating power is very- intense ; indeed, it has been recorded that they can pass through a sheet of lead 8 cm. thick. All the experimental evidence goes to show that the rays are identical with the X-rays. They are ether pulsations set up by the bombardment of the sohd salts by the stream of negatively charged electrons (Beta rays) produced within itself, in every way comparable with the ether pul- sations (X-rays) propagated through space when the cathode tube is being bombarded by the rapidly moving cathode rays. uu 658 AN INORGANIC CHEMISTRY The Disintegration of the Radio-active Elements. — During the investigation of the properties of uranium, Crookes noticed that if a uranium salt was precipitated by means of ammonium carbonate and the precipitate treated with an excess of the reagent, nearly the whole of the precipitate first thrown down redissolved. This solution, although it contained the uranium, was found to have lost its radio-activity, whilst the smaU residue had acquired the whole of the radio-activity pre- viously possessed by the uranium salt. This residue Crookes named Uranium X. Still more interesting was the result that in the course of a few months the inactive uranium had re- acquired its activity, whilst that of the uranium X had entirely disappeared. A similar phenomenon was observed by Ruther- ford and Soddy in the case of Thorium and Thorium X. The only interpretation of these results is that the activity of uranium is due to some substance which is produced spontaneously from the uranium. The radio-activity of normal uranium represents the equilibrium value wherein just as much of the active product is produced per unit of time as undergoes spontaneous decom- position. The most striking feature about this radio-active decay and regeneration is that it is totally unaffected by temperature changes. It proceeds as freely at the temperature of liquid air as at a white heat. Moreover the radio-activity of a salt of one of the radio-active substances is strictly proportional to the amount of radio-active substance present and is entirely inde- pendent of the particular acid radicle with which the radio-active element is associated. These facts prove that the phenomenon of radio-activity is a property of the atom. Moreover no known molecular decomposition is uninfluenced by a change of temperature from the neighbourhood of the absolute zero to a white heat, so that we are forced to the con- clusion that radio-activity is a result of atomic disintegra- tion. During the disintegration of thorium a whole series of disin- tegration products have been shown to be formed, some of which have but a transient existence, whilst the life of others is measured in geological epochs. RADIO-ACTIVITY 656 Th.*:^Meso-th. I _^Meso-th. Ill^_.Eadio-th*-^ Th X*^ 232 228 228 228 224 ~^ -^- Th. Emanation "'-'^ Th. A*-^ Th. B^-^sTh. C *r^Th D*r^ 220 216 212 212 208 end product. 208 An interesting feature was discovered concerning thorium emanation. It proved to be a gas (atomic weight 220) which has all the properties of the rare gases (argon, etc.). Under no conditions could this gas be brought into combination with another element. A similar disintegration series has been constructed for uranium and radium. The constant association of these two elements suggested the hypothesis, now generally accepted, that radium itself is one of the disintegration products of uranium. Ur. r-^^Vr. Xi'^iT^Ur. Xj^T^Ur. 2^-^^ lonium'^-^^ Ra°^^ 238 234 234 234 230 226 ft,"* Ra. Emanation°''^^Ra. A*i^Ra. B*;:^Ra. C''>l^Ra.D'jf^ 222 218 214 214 210 ^ x-' Ra. E\^Ra. F*^end product. 210 210 206 From the chemical point of view, the most interesting of this series other than the end product is the radium emanation, produced from radium by the expulsion of an alpha ray. Radium emanation, like the emanation of thorium, proved to be a gas, endowed with all the properties of the argon family of gases. The name Niton is now assigned to the radium emanation. 660 AN INORGAJNiU UJUUMlHiKi The following experiment, carried out by Rutherford with niton, is of surpassing interest for the hght which it throws upon atomic disintegration. A very thin-walled glass tube A was sealed into an outer tube B, which was thoroughly evacuated. A was then filled with heUum under pressure, but no trace of the heUum spectrum could be detected in the attached discharge tube C. No leak could therefore exist between A and B. The helium was then carefully removed and radium emanation passed into A. The glass walls of A were sufficiently thin to allow the passage of the alpha rays given o£E during the disintegration of the niton, but, once within the outer vessel B, they were unable to escape through the thicker outer walls into the atmo- sphere. In a few hours the heUum spectrum became visible in C. Atoms of helium, then, are thrown off during the disintegration of niton into another radio-active solid. This affords the first definitely established example of the transmuta- tion of one element into another. B V V Fig. Ml. Are the Radio-active Disintegration Products Elements ? • — The degradation of the atom of radium into niton and helium raises the question whether radium is to be classed as a definite chemical element, and if so, to what extent does the generally accepted definition of an element require amplification. The usual definition of an element stresses its non-composite nature. From such a view point, the radio-active elements can scarcely be recognised as elementary in nature ; but on the other hand, they do possess many of the criteria of a true chemical element. The various radio-active elements and their disintegration pro- ducts have a perfectly definite atomic weight, certain distinct and definite chemical properties are associated with them, each of them possesses a spectrum distinct from the spectrum of all other elements, while their decomposition is utterly and entirely beyond the control of the chemist. The maimer in which radium decomposes into its emanation at a rate which is independent of the temperature of its surroundings is so distinct from the usual decomposition to which a composite molecule is subject that no RADIO-ACTIVITY 661 reasonable doubt can now exist that there is a fundamental distinction between the breaking down of a composite molecule like water, and the disintegration of a radio-active element — the one is molecular, the other atomic. The view which has been put forward to explain radio-active disintegrations — a view which is now generally accepted — is that the phenomenon is purely intra-atomic, and in no way akin to the breaking down of the molecule of a compound into its con- stituent atoms. For this reason a slight modification of the old definition of an element is required (q.v.), and in the definition of an element already quoted (p. 17) emphasis is laid upon the non -divisibility of the atom of the chemical element by any means under the control of man. The Chemistey oe the Radio-active Elements — Theie Position in the Peeiodic System — Isotopes When one takes into consideration the extraordinarily small amounts of the radio-active elements at the disposal of the physico-chemists who have investigated the properties of these substances, it is a matter of wonder that their chemical properties have been tracked down with the certainty with which the trained analyst is wont to undertake the investigation of a sample of a normal inorganic salt. " Separations of two chemically distinct radio-elements," writes Soddy*, " are effected as completely as in ordinary chemical analysis, always provided that if such a mechanical operation as filtration is involved, there must always be sufficient quantity of matter to form a filterable precipitate. Under such conditions it is always found, for example, that polonium is completely precipitated in acid solution by sulphuretted hydrogen. Ionium is not so precipi- tated, but is by ammonia, whilst radium is precipitated by neither reagent, but is completely by sulphuric acid. So that, in presence of precipitable quantities of any of the members of these three analytical groups, the complete separation of these three radio-elements, chosen as examples, from one another is as easy with absolutely unweighable quantities as ordinary analysis. Going further, it is equally possible to show that polonium resembles bismuth more nearly than any other of the metals of Group 2, but does not resemble it absolutely, and is in * The Chemistry of the Radio-Elements (Longmans, Green & Co.) 662 AN INORGANIC CHEMISTRY fact a new type of chemical element, though it has never yet been obtained in quantity sufficient to be worked with, apart from inactive material to serve as a vehicle for its transport. Similarly, radium resembles barium more nearly than it does any other metal of Group 4, but does not resemble it absolutely, and is, in fact, another new type of chemical element.'" "Lastly, ionium not only resembles thorium more nearly than any other member of Group 4, but is found to be chemically identical in properties with it. Ionium, in presence of sufficient thorium to act as a vehicle, may in quantities, however infini- tesimal, be completely and simply separated from any mixture containing any or every known element, by the simple device of separating the thorium therefrom by known methods. Ionium is a new rat^io-element, but is not a new chemical type of element. This is a new conception. In every detail of its radio-active properties, the velocity of range of its a-rays, its period of change, the character of its product and of its parent, ionium is distinct and different from every other radio -element. But in its chem- ical nature it is not new. On account of its identity in this respect, with the element thorium, every detail of its chemical nature is known " by proxy " as completely as though it were obtainable, like thorium, by the ton. Ionium cannot be obtained " radio-chemically pure " on account of its long period and the presence of thorium in all minerals from which it may be obtained." Here, then, are two elements, ionium and thorium, which appear chemically inseparable, and yet possess a different atomic weight. Other such cases are on record. Radium D, or radio-lead (210), is one of the degradation products of radium. It has its own radio-active properties, but from the chemical point of view its properties are indistinguishable from those of ordinary lead. In 1917 Richards and Hall subjected a sample of Austrahan carnotite, containing lead and radium D, to a searching investigation. The lead extracted from this mineral was known to contain radium D ; but although lead made from this source was subjected to over 1,000 re-crystallisations, no measurable change in the atomic weight of the lead resulted. For the explanation of this result we must turn to the Periodic Law. The expulsion of an alpha ray carrying two positive charges has been shown by Soddy and Fajans to shift the position of the I^,±i.u±KJ-Al^ L 1 /IT Y 663 radio-active element two groups towards the left, so that the new radio-active element not only has an atomic weight of four units less than its parent, but its properties will be such as one would expect from an element lying in a group two spaces to the left of the parent element. Furthermore, the expulsion of the beta particle produces a shift of one place to the right. Given the position of the parent element in the Periodic Table and the type of rays thrown out where disintegration occurs, it is possible to assign to every radio-active element its group-position in the Periodic System. This has been done in Table 56, a method of arrangement first put forward by Soddy. With this arrangement we see that different radio-active elements with different atomic weights, derived from different parents and yielding different disintegration products, fall into the same position in the Periodic System. Into each position there is fitted not a single chemical element, but a definite chemical type, aU the elements of the same type occupying the same place in the Periodic classification and exhibiting identical chemical properties. Such elements are known as Isotopes {isos, equal ; topos, place). An examination of Table 66 shows that the end-product from the disintegration both of uranium and thorium falls into Group 4b, under lead. Such end products would be chemically inseparable from ordinary lead. During the disintegration of uranium (atomic weight =238) eight alpha particles (charged helium atoms) are expelled. Neglecting the neghgible loss of weight due to the expulsion of five beta particles, the atomic weight of the end product should be 238— (4x8) =206. In the same way thorium which expels six alpha particles during its disintegration, should give rise to an end product with an atomic weight 232 — (6 x 4) =208. Ordinary lead has an atomic weight of 207-2. The view held at the present time is that lead, such as we know it, is a mixture of two isotopes, the end products arising from the degradation of uranium and thorium. A brilliant confir- mation of this hypothesis was carried out by Soddy. The mineral thorite contains 57 per cent, thorium, 1 per cent, uranium and 0-4 per cent. lead. If all the lead is of radio-active origin, is stable, and is derived from both parents (thorium and uranium) by radio-active disintegration on the Hnes described above, then 95-5 per cent, of the lead found in thorite should consist of the isotope from thorium (atomic weight =208), and »e f" :g i s^ .<= ■j> i:^ S "f^ >i t m <» in e -5? k "U i« — 'a/qei oipouaj py^ ui 33B/d guo /y/uo si sjm — -^-j.y fc 664 RADIO-ACTIVITY 665 4-5 per cent, of the isotope from uranium, from which the atomic weight of lead may be calculated, i This gives the value 207-9. An accurate atomic weight determination of lead gave the value 207-77 as the mean of eight concordant determinations. Again, from a very pure pitchblende which contained very Httle thorium but a great deal of uranium, a sample of lead, in every way indistinguishable from ordinary lead, was separated, which gave an atomic weight 206-046. Strong support to the views of Soddy, should such be needed, is afforded by the work of Aston. This investigator has used a purely physical mode of analysis, known as the positive ray method of gas analysis. By this means he has investigated some nineteen non-radio-active elements, and has found distinct evidence of heterogeneity in the elements boron, neon, sihcon, chlorine, argon, bromine, krypton, xenon and mercury ; in short, his results indicate very clearly that these elements, as the chemist knojvs them, are composite in nature, each of them consisting of a mixture of two or more isotopes. Strenuous efforts have been made to separate the isotopes of chlorine by taking advantage of physical properties, such as diffusion, which are dependent upon density, but so far without success, though such a separation has been claimed for the hydrogen chloride formed from these two isotopic modifications of chlorine. Frac- tional diffusion through pipeclay has also enabled the partial separation of neon into its isotopes to be effected, whilst the method of fractional distillation of mercury has given direct evidence of the presence of isotopes in this element. In the Ught of these results no reasonable doubt, either of the theory of isotopes, or of the radio-active origin of lead, remains. The suggestion has been made that the three outstanding exceptions in the Periodic Law (Ar — K, Co — Ni, Te — I) may arise from the presence of small quantities of isotopes which the various chemical methods of separation have failed to separate. This hypothesis has recently been confirmed by Aston's work on the composite nature of argon (atomic weight of isotopes 40 and 36). Another interesting and vitally important aspect of this scientist's work is the fact that the atomic weights of all the isotopes which he has investigated are represented within the hmits of experimental error by whole numbers. The unique atomic weight possessed by chlorine (35-46) 666 AN INORGANIC CHEMISTRY arises, therefore, from the presence of the isotopes of atomic weight 35 and 37 in such proportions that the resultant atomic weight is represented by the value at present used by the chemist, 35-46. Although the view of Harkins, that the atoms of the chemical elements are built up of atoms of hydrogen and hehum cemented together by electrons, may yet require modification, there seems httle doubt in the minds of the progressive chemists of the day that the old Proutian hypothesis of the common origin of matter has come back to life. Our ideas on the constitution and genesis of matter are at present in a state of flux ; in this connection the writer cannot do better than quote the suggestive words of Prof. E. C. C. Baly (1914), " As did his forefathers of pre-Avogadro days, so also does the chemist of to-day now await that great generahsation which shall co-ordinate and hnk up all the threads to found a new philosophy. Radio-activity . . active nitrogen and oxygen, atomic disintegration, atomic weight variation, all will be unified and embodied in the new philosophy of the twentieth century. Then will a new chemistry in its greater meaning emerge as a phoenix from the glowing parental fires of the many chemistries of to-day." Since the above was penned. Sir Ernest Rutherford, to whom so much of our knowledge of radioactive phenomena is due, has published an account of his experiments, in which he has investigated the effect of bombarding molecules of oxygen and of nitrogen with the heavy a- particles. There seems little reason to doubt that, by means of this intense bombardment, this investigator has actually succeeded in disrupting particles of these gases, the disrupted atom of oxygen yielding atoms of an isotope of helium, whilst the nitrogen atom, under similar treatment, produces atoms of hydrogen and of the helium iso- tope. The transmutation of the elements has begun, the dream of the alchemist appears at hand. The Structure of the Atom During the past decade the old " sohd " view of the atom has yielded place to the modern conception of the electro-magnetic nature of mass. The ease with which the beta particle can find its way through a sheet of metal is in itself sufficient to cast doubt upon the older conception of matter. Further, the mass of the electron, smaU though it be, has been shown to be entirely STRUCTURE OF THE ATOM 667 due to the charge upon it. The modern view is that the atom consists of electrons and positive electricity, nothing else. StartUng though such an hypothesis may be, it is the only one that fits the facts known to us at the present day. Thus no experiments have yet been described which require for their elucidation an assumption involving the presence of anything but positive and negative electricity. In the experiments with alpha rays Rutherford observed that every now and then a ray was swung violently out of its course. He interprets these deflections as being due to a repulsion between the positively charged alpha particle and a highly concentrated positive charge at the centre of the atom. In short, Rutherford postulates the atom as consisting of a highly concentrated nucleus of positive electricity with concentric rings of electrons round it. This view of the atom accounts for the fact that no experiments have yet succeeded in shaking positive electrons out of matter. The number of electrons present in a given atom has been estimated in various ways. The most rehable of the results indicate that the number of electrons present in the atom is approximately equal to half the atomic weight of the element. It has been suggested that the number of electrons present within the atom is equal to the atomic number of the element (i.e. its number in the Periodic System of Classification, H=l, He=2, Li=3, etc.). Should this prove to be so, the atomic number would have a deeper fundamental significance than the atomic weight itself. Should the connec- tion between the atomic number and the number of the electrons present within the electron be definitely estabhshed, the Periodic System of Classification wiU take its place as one of the great laws of Nature. Crystal Structure. — A supersaturated solution of sodium acetate, which is brought to crystaUisation by the introduction of a small crystal of sodium-acetate, presents a very beautiful phenomenon. The rapid but regular manner in which the fairy- Kke net of crystals spreads throughout the mass from the very moment of inception has long been viewed as an expression of some great law of Nature, which is equally responsible for the symmetry displayed by a large perfect crystal of quartz as for the equally perfect symmetry possessed by the minutest crystal of this substance. Many views have been put forward to account 668 AN INORGANIC CHEMISTRY for the orderly arrangement of the molecules within a crystal, but it is only within comparatively recent times that definite light was thrown upon this subject. On p. 489 it has been shown that a ray of white hght can be spUt up into a spectrum by its passage through a prism. An- other method, the theory of which wiU be found discussed in any standard textbook on Optics, is based upon the use of a diffrac- tion grating. This consists of an almost innumerable number of fine parallel Unes, drawn so closely together that as many as 10,000 Unes are found in a centimetre. Such a grating is per- fectly satisfactory for ordinary hght waves. The view has been held for some years that X-rays differ from ordinary hght rays merely in the shortness of their wave length, and that an X-ray spectrum could be obtained if it were possible to obtain a suffi- ciently fine grating. Theory has predicted the impossibUity of ruling such a grating, but Prof. Laue (1912) conceived the idea that the regular grouping of the atoms within a crystal might provide such a grating. His idea was put to the test in 1913 with the following apparatus. F S Fig. 14.3. F represents an X-ray tube, ^4, B and C are screens which are used in order to obtain a fine pencil of X-rays. These impinge upon the crystal X. S is a, sighting screen, P a photographic plate inserted after the adjustment has been made. On devel- oping the plate it was found that, besides the central black patch given by the rays which had passed straight through the crystal, a symmetrical pattern of spots, as predicted by Laue, was found. The subject was carried further by W. L. Bragg, who utihsed this idea, as a means of delving down into the interior mechanism of a crystal. We know that the intensity with STRUCTURE OP THE ATOM 669 which X-rays are reflected from an atom is a function of the number of electrons within it, and therefore of its atomic weight. Bragg obtained an X-ray spectrum for sylvine (KCl) and sodium chloride, and found that the reflection from the potassium (39-0) and chlor- ine (35-5) atoms was practically indistinguishable, but in the case of sodium chloride a distinct difference is seen in his photo- graphs according to whether reflection has taken place at the heavy or the Ught atom. He found that aU the results could be accounted for by assigning to both crystals the structure indicated in Eig. 144, the metalhc atoms being re- presented by dots, the chlorine atoms by circles. In the case of sylvine the dots and circles become indistinguishable. It would thus appear that the regu- larity of structure arises, not from the regular arrangement of the molecules of the salt, but of the constituent atoms. So far as the crystal form is concerned, any particular sodium atom is equally situated with respect to six different chlorine atoms ; the purely chemical co-ordination of atoms within the molecule has given place to a method of arrangement which is governed by the laws of physics. Exactly what happens during the fusion or solution of crystal is as yet unknown. This method of investigation in the hands of Bragg and others has already thrown most valuable light upon the crystal structure of many of the elements and their compounds, among them the diamond zinc blende, copper, calcium, fluoride, etc. Fig. 144. INDEX Acids are indexed under "arid"; salts under the name of the base. ACCUMULATOK, 594 Acetylene, 367, 369 Acid, 57, 435 — anhydride, 57 — oxide, 57 — salts, 90 Acid, acetic, 374 — antimonic, 340 — antimonious, 338 — arsenic, 340 — arsenious, 338 — auric, 518 — bismuthic, 341 — boracic, 564, 567 — boric, 564, 567 — bromic, 185 — carbonic, 353 — chlorauric, 518 — chloric, 180 — chlorochromic, 610 — chloroplatinic, 650 — chlorosulphuric, 246 — chlorous, 183 — chromic, 605 — cyanic, 378 — disulphuric, 246 — dithionic, 250 — ferricyanic, 638 — ferrocyanic, 485 — formic, 374 — hydrazoic, 283 — hydriodic, 163 — hydrobromic, 163 — hydrochloric, 154 — hydrocyanic, 378 — ' hydrofluoboric, 566 — hydrofluoric, 171 — hydrofluosilicic, 391 — hypoantimonic, 342 — hypobromous, 185 — hypochlorous, 147, 175, 210 ; action of, in bleaching, 177 ; as oxidising agent, 177 ; thermochemistry of, 184 — hypoiodous, 186 — hyponitrous, 303 — hypophosphoric, 325 — hypophosphorous, 324 — hyposulphurous, 248 — iodic, 186 — manganic, 619 — meta, 187 ■- — meta-phoaphoric, 321 — muriatic, 154 Acid, nitric, 287 ; action on metals, 295 ; decomposition of, 292 ; fuming, 292 ; manufacture, 288, 290 ; oxidising action of, 292 — nitrosylsulphuric, 242 — nitrous, 298 — ortho, 187 — ortho-phosphoric, 319 — osmic, 646 — oxalic, 375 — pentathionic, 250 — perchloric, 183 — perchromic, 611 — per-iodic, 187 — permanganic, 620 — peroxidic, 196 — persulphuric, 247 — phosphoric, 318 ; meta-, 321 ; ortho-, 319 ; pyro-, 320 ; tests for, 322 — phosphorous, 323 — picric, 293 — plumbic, 588 — plumbous, 588 — polythionic, 250 — prussic, 378 — pyrophosphoric, 320 — pyrosvilphuric, 240 — selenic, 251 — selenious, 251 — silicic, 394 — stannic, 586 ; a, 586 ; ^, 587 — stannous, 584 — stearic, 376 — sulphuric, 238 ; chamber process, 239 ; constitution, 245 ; contact process, 238 ; fuming, 246 ; pro- perties, 243 — sulphurous, 233, 234 — telluric, 251 — tellurous, 251 — tetrathionic, 250 — thiocarbonic, 362 — thiosulphuric, 248 — trithionic, 250 — ■ xanthoproteic, 295 Acids, action on metals, 437 ; on sulphides, 223, 434 — • dissociation of, 425 — ionisation of, 425 — properties of, 436 — strength of, 438 Acidimetry, see titration Adsorption, 348 671 672 INDEX Affinity, chemical, 152 ; constant, 202 Air, a mixture, 39, 272 — composition, 269 — liquefaction, 273 Alabaster, 537 Alcohol, 373 — ethyl, 373 — methyl, 373 Aliphatic hydrocarbons, 377 Alkali, 57, 467 — ammonia-soda process, 477 — - bicarbonates, 479 — carbonates, 475 — electrolytic preparation, 471 — Ije Blanc's process, 476 — Solvay process, 477 Alkalimetry, see titration Alkaline earths, 521 ; family relation- ships, 521 AUotropes, 195 Alpha rays, 657 Alumina tes, 572 Aluminium, 569 — acetate, 576 — bronze, 570 — carbide, 367, 575 — chloride, 573 — halides, 573 — hydroxide, 571 — nitride, 575 — oxide, 571 — sulphate, 574 — sulphide, 575 Aluminothermic process, 570 Alums, 456, 574 Alunite, 575 Amalgam, 655 Amethyst, 393 Amides, 280 Ammonia, 278 -*• composition, 280 — metal compounds, 498, 504, 513, 550, 615, 642, 644 — preparation, manufacture, 277 — silver compounds, 513 — soda process, 477 — synthetic, 279 Ammonium amalgam, 491 — bicarbonate, 477, 493 — bromide, 493 — carbamate, 494 — carbonate, 493 — chloride, 492 — cyanate, 361, 486 — hydrosulphide, -194 — hydroxide, 444. 493 .".i4 — iodide, 492 — molybdate, 611 — nitrate, 493 — nitrite, 493 — phosphates, 495 — salts, 282 — sulphates, 494 — sulphides, 494 Ammonium thiocyanate, 486 — thiosalts 343, 587 Amorphous substances, 117, 217 Amphoteric electrolytes, 444 — oxides and hydroxides, 334 position in Periodic Table, 330 Anatase, 597 Angleaite, 589 Anhydride, 57 Anhydrite, 141 Anions, 411 Anode, 411 Antichlor, 232, 250 Antimonates, 341 Antimonites, 339 Antimony, 328 — allotropic modifications, 330 — blende, 328 — bloom, 328 — halides, 332 — hydride, 331 — ochre, 328 — pentoxide, 336, 340 — sulphides, 342 — tetroxide, 341 — trioxide, 334, 338 Apatite, 305, 526 Aqua regia, 293 Aqueous vapour pressure, 94, 96 Aragonite, 118, 528 Argon, 275 Aromatic hydrocarbons, 377 Arsenates, 340 Arsenic, 328 — allotropic modifications, 330 — halides, 332 — hydrides, 331 — pentoxide, 336, 340 — sulphides, 342 — trioxide, 336, 337 Arsenical iron, 328 ; pyrites, 328- Arsenites, 338 Arsenolite, 328 Arsine, 331 Asbestos, 523 Asphalt, 366 Atmosphere, 269 — composition of, 269 — impurities in, 270, 272 — pressure of, 61 Atom, properties of, 136 — structure of, 666 Atomic disintegration, 655, 060 — heats, 133 — number, 667 — structure, 137, 666 — theory, 135 — weights and their determination, 129, 133 ; correction of, 263 ; list of, 131 ; versus combining weights, 35, 130 Aurates, 518 Auto-oxidation, 169, 183, 231, 234, 235, 298, 312, 323, 325, 338 IKDEX 673 Avogadro's hypothesis, 122 Azides, 284 Azurite, 499 Baking soda, 480 Balanced reactions, 200 Barium, 528 — carbonate, 528 — chlorate, 539 — chloride, 536 — halides, 536 — hydrides, 527 — hydroxide, 521, 530 — nitride, 527 — oxides, 529 — peroxide, 533 — sulphate, 538 — sulphide, 539 Barytes, 214, 526 Bases, 90, 440 — strength and dissociation, 441 Basic oxide, 57 — salts, 9-1 Basicity of acids, 90, 443 Bauxite, 569 — purification of, 572 Becquerel rays, 655 Benzene, 377 Benzine, 366 Beryl, 397 Beryllium, see glueinum Berzelius, 25, 28 Bessemer process, 629 Beta rays, 657 Bicarbonates, 354, 530 Birkeland Eyde process, 291 Bischoffite, 523 Bismuth, 328 — glance, 329, 342 — halides, 332 — ochre, 329 — oxide, 334 — pentoxide, 341 — salts, 339 — subiodide, 342 — suboxide, 342 — tetroxide, 341 — trioxide, 339 Black ash, 476 Black lead, 348 Blast furnace, 626 Bleaching, 177 — powder, 178 Blue vitriol, 506 Boiler scale, 532 Boiling point, 406 elevation of, 406 Bone ash, 305, 349 Boracite, 564 Borax, 564, 568 Borides, 565 Borocalcite, 564 Boron, 564 ■ — • carbide, 566 Boron halides, 56G — hydrides, 565 — nitrides, 565 — oxide, 567 — phosphate, 567 Boronatrocalcite, 564 Boyle's law, 61, 65 Brass, 502 Braunite, 613 Erin's oxygen process, 50 Britannia metal, 583 Bromates, 185 Bromides, 166 Bromine, 158 — dissociation of, 162 — physical properties of, 160 — preparation, 159 Brookite, 597 Bronze, 583 Brownian movement, 7 1 Bunsen burner, 382 — flame, 382 Butane, 364 Cadmium, 547, 549 — ammonia compounds, 551 — carbonate, 552 — halides, 551 — oxide, 551 — separation from copper, 553 — sulphate, 552 — sulphide, 553 Caesium, 491 Calamine, 547 Calcination of metals, 5 Calcite, 528 Calcium, 526 — bicarbonate, 530 — carbide, 369, 544 — carbonate, 528 — chlorate, 182, 539 — chloride, 534 — cyanamide, 278, 545 — halides, 534, 537 — hydride, 527 — hydroxide, 528 — nitride, 527 — oxalate, 541 ■ solution in acids, 542 — oxide, 528 — peroxide, 534 — phosphate, 540 — phosphide, 313 — silicates, 541 — sulphate, 537 — sulphide, 540 Calculations, chemical, 43, 45 Calomel, 560 Calorie, 20 Calorific power, 371 Calx, 6 Carbides, 350 Carbon, 346 — amorphous, 348 X X 674 j.±y±JAii.c\. Carbon compounds, isomerism of, 364 — dioxide, 351 ; assimilation of, 355 ; in air, 271 — disulphide, 362 — modifications, 346 — monoxide, 356 — oxysulphide, 362 — suboxide, 358 — tetrachloride, 368 Carbonates, 354 Carbonyl chloride, 361 — sulphide, 362 Carbonyls, 358 Carborundum, 392 Carnallite, 141, 467, 523 Camotite, 612 Cassiterite, 582 Castner*s process, 468 Catalysis, 58, 79, 180, 236 Cathions, 411 Cathode, 411 — rays, 653 — tube, 653 Cavendish, 275 Celestine, 526 Cell, 461, 463 — Daniell, 461 — Griesheim, 472 — Solvay, 472 Cement, 533 Cerite, 563, 597 Cerium, 563, 597 Cerussite, 118, 589 Chain linking, 188, 198, 363 Chalcocite, 499 Chalcopyrite, 499 Chalk, 528 Chamber acid process, 239 Chance process, 477 Charles' law, 63, 65 Charcoal, 348, 349 — bone, 305, 349 — wood, 348 Chemical change, 12, 37 — combination, 18 — compounds and mixtures, 38 — decomposition, 19 — displacement, 18 — equations, 44 ; their significance, 45, 137 ; equilibria, 200 and concentration, 203 ; pressure, 204 ; temperatiu'e, 206 ; displacement of, 203, 204, 206, 207 ; heterogeneous, 211; reactions, modes of, 18; speed of, 201 — substitution, 18 Chemistry, definition of, 4 — and physics, distinction between, 8 Chili saltpetre, 159 China clay, 574 Chlorates, 181, 480 Chlorides, 154 — preparation and general properties, 155, 459 Chlorine, 17, 141 — chemical properties, 147 — dioxide, 182 — heptoxide, 184 — hydrate, 147 — monoxide, 174 — preparation of, 141 — physical properties, 147 Chloroform, 368, 374 Chromates, 605 Chrome alum, 604 — iron ore, 599 — ochre, 599 — steel, 600 Chromic anhydride, 609 — chloride, 603 — nitrate, 605 — oxide, 602 — sulphate, 603 Chromite, 599 Chromites, 600, 601, 602 Chromium, 599 — oxides, 600 — trioxide, 609 Chromous compounds, 601 Chromyl chloride, 610 Chryeoberyl, 569 Cinnabar, 214, 554 Clay, 574, 576 Cleveite, 276 Coagulation, 113, 395 Coal, 350 — gas, 370 Cobalt, 639 — complex cyanides, 642 — glance, 328, 639 — sulphide, 641 Cobalti -nitrite potassium, 642 Cobaltic oxide and salts, 642 Cobaltite, 639 Cobalto-cobaltic oxide, 640 Cobaltous oxide and salts, 640 Coke, 348, 371 Colemanite, 564 Colloidal solution, 71, 113, 395, 635 coagulation of, 113, 395 Columbite, 345 Combining weight, 32 ; determination of, 29 ; difficulties of, 34 Combustion, 54, 380 — reversed, 380 — spontaneous, 56 Complex salts, 454 Composition of earth's crust, 50 Compounds, 17 — molecular, 3(Q — preparation of, 458 Concentration, 1 14 Conductivity, 415 — equivalent, 423 — of electrolytes, 415, 418 Constant boiling point, acids of, 153, 165, 171, 292 Contact action, see catalysis INDEX 675 Contact process, 236 Copper, 497, 499 — ammonia compounds, 498, 503, 504 — blister, 500 — complex cyanide, 505 — compounds, general survey, 502 — hydride, 248, 325 — matte, 499 — metallurgy, 499 — occurrence, 499 — refining, 501 — ruby, 499 — separation from cadmium, 553 — sulphate, 506 hydration of, 98 Cordite, 294 Corrosive sublimate, 557 Corundum, 569 Coulomb, 419 Cream of tartar, 480 Christobalite, 393 Critical phenomena, 66 Crocoisite, 599 Cryolite, 475 Crystals, mixed, 118 — structure of, 667 — system of classification, 116 Crystallization, fractional, 114 — from solution, 11, 73, 114 Crystallographic systems, 116 Crystalloids, 395 Cupellation, 509, 517 Cupric acetate, 507 — bromide, 504 — carbonate, 506 — chloride, 504 — cyanide, 377, 505 — hydroxide, 503 — iodide, 504 — oxide, 503 — sulphate, 98, 506 — sulphide, 508 Cuprous bromide, 508 — chloride, 508 ■ — cyanide, 505 — iodide, 506 — oxide, 507 — sulphate, 507 — sulphide, 508 Current density, 419 Cyanates, 378 Cyanides, 378 — of iron, 637 Cyanide process, gold, 516 Cyanogen, 377 Dalton, 25, 62, 135 Davy, 3, 159, 446 Deacon process, chlorine, 143 Decomposition, electrolytic, 464 Deliquescence, 100 Density of gases, measurement, 73, 74 — vapour, 74 Deviations from Gas Laws, 65 Dewar's flask, 274 Dialysis, 394, 635 Diamond, 347 Diaspore, 569 Dichromates, 605 Diffusion, 82 — Graham's law of, 83 Dilution law, Ostwald's, 424 Dimorphic substances, 216 Discharging potential, 464 Disintegration of radio-active elements, 658 Dissociation, 200 — in solution, 410 — of acids, bases, salts, 425 hydrates, 98 — pressure, 212 Distillation, 11 — destructive, 348, 371 — ■ fractional, 365 Distribution between two solvents. 112 Dobereiner's triads, 255 Dolomite, 522 Double salts, 455 Dulong and Petit's law, 133 Dyad, 87 Dyeing, 573 Dynamite, 294 Dysprosium, 564 Earthenwabe, 576 Earth's crust, composition 50 Eau de Javel, 176 Efflorescence, 98 Eka-aluminium, 264 — -boron, 264 — -silicon, 264 Electrical terminology, 411 Electric discharge in a gas, 652 — furnace, 306 Electro-a£6nity, 455, 463 Electro-chemical equivalent, 419 Electrode, potential, 460 Electrolysis, 415 — Faraday's laws of, 418 — of water, 79, 446 Electrolytes, 411 — abnormalities of strong, 425 — amphoteric, 444, 550, 571 — conductivity of, 415 Electromotive force, 417 — series, 463 Electronegative elements, 462 Electrons, 654 Electrophoresis, 395 Electroplating, 465 Electropositive elements, 462 Elements, 16, 660 — bridge, 259 — metallic and non-metallic, 41 — transmutation of, 666 — transition, 259 Emanation, 659 676 Emery, 571 Endosmosis, electrical, 395 Energy changes, 13 ■ — chemical, 16 Energy, law of conservation of, 15 Enstatite, 397, 523 Enzymes, 352 Epsom salts, 525 Equations, 44 — significance of, 44, 137 EquiUbria, ionic, 410, 431 — heterogeneous, 211 Equilibrium, chemical, 200 — constant, 202 — effect upon, of concentration changes, 203 ; of pressure changes, 204 ; of temperature changes, 206 — factors aiiecting, 207 Equivalent solutions, 423 — weight, 32 — conductivity, 423 Erbium, 564 Esters, 375 Etching, 171 Ethane, 363 Ethers, 376 Ethylene, 368 Evaporation, 72 — kinetic theory of, 72 — latent heat of, 94 Explosion wave and Bunsen burner, 385 Extraction, 113 Fahaday, 3, 147, 418, 419 Fat, 376 Felspar, 397, 467 Fergusonite, 276, 563 Fermentation, 352 Ferrates, 639 Ferric alum, 637 — chloride, 636 — ferrocyanide, 638 — hydroxide, 635 — oxide, 635 — sulphate, 637 — sulphide, 637 Ferroso -ferric oxide, 637 Ferrous carbonate, 635 — chloride, 633 — ferricyanide, 638 — oxide, 633 — sulphate, 633 — ■ sulphide, 634 Fertilisers, 289, 305, 320 Fire-damp, 367 Fixation of nitrogen, 290 Flame, chemistry of, 383, 384 — Ivmiinosity, 383, 384, 385 — structure of, 381 Flotation, froth, 549 Flint, 393 Fluorine, 169 Fluorspar, 526 Flux, 496 Formulae, 42 — chemical, and molecular weights,. 134 — graphic, 88 Franklinite, 547 Frauenhofer's lines, 490 Freezing point, depression of, 407 Fulminate, mercury 559 Fusion, latent heat of, 94 Gabounite, 563 Gadolinium, 664 Gahnite, 669 Galena, 214, 589 Gallium, 577 Galvanised iron, 549 Gamma rays, 657 Garnet, 396 Gamierite, 643 Gas, coal, 370 — kinetic theory of, 68 — laws, 61, 63 — liquefaction of, 68, 273 — oil, Pintsch, 361 — solubility, 110 — tar, 371 Gay Lussac, 63, 104 Germanium, 581 Glance, cobalt, 639 Glaserite, 483 Glass, 397 Glauber's salt, 483 Glaze, 576 Glucinima, 522 Gold, 497 — compounds, 517 — metallurgy, 516 — oxides, 517 Goldschmidt's process, 570 Graham's law of diffusion, 83 Gram molecular weight, 126 ; volume, 126 Graphite, 347 Green vitriol, 634 Griesheim cell, 472 Guano, 305 Guignet's green, 602 Guncotton, 294 Gunpowder, 482 Gypsum, 141, 214, 526, 537 Haematite, 625, 635 Halides, 158 Halite, 141 Halogen compounds with each other, 172 — family, 158, 172 — hydr-acids, 158 Halogens valence of, 188 Hardness of water, 531 Hausmannite, 613 Heat of formation, 119 Helium, 276 Holmium, 564 INDEX 677 Homologous seriee, 360 Homologues, 365 Horn silver, 508 Hydrargyllite, 569 Hydrates, 98 — dissociation of, 98 — vapour pressure, 98 Hydrazine, 282 — hydrate, 282 — sulphate, 282 Hydrazoates, 283 Hydrides, 84 Hydrocarbons, 363 — saturated, 363 — unsaturated, 366 Hydrogel, 395 Hydrogen, 78 — action upon iron oxide, 85, 212 — bromide, preparation, 163, 164 — chemical properties, 84 — chloride, composition, 155 ; prepar- ation, 149 ; properties, 152, 154 — disulphide, 226 — equivalent, 85 — fluoride, 170 — in Periodic System, 265 — iodide, preparation, 163, 164 — ion, 435, 438 — nascent, 82 — perojdde, composition, 198 ; pre- paration, 195 ; properties and reactions, 196 — persulphate, 247 — persulphide, 226 — phosphides, 311 — preparation of, 78 — properties, physical, 82 — purification, 81 — solubiKty in palladium, 82 — sulphide, preparation, 218 ; re- ducing action of, 186, 224, 225 ; \ise in analysis, 222 — trisulphide, 226 Hydrogenation of oils, 643 Hydrolysis, 150, 207, 315, 319, 334, 338, 339, 341, 447 — of metallic halides, 207, 456 ; non- metallic halides, 164 Hydrosol, 394 Hydrosulphides, 221 Hydroxides, 57, 458 Hydroxylamine, 284 Hydroxyl ion, 440 Hypo, 249 Hypochlorites, 176, 178 Hypophosphites, 324 Hyposulphite, 249 Hypothesis, 4 — phlogiston, 5 Iceland spab, 528 Ideal gas, 127 Ilmenite, 597 Incandescent mantles, 597 Indicators, 449 Indigo, oxidation of, 248 Indium, 578 Ink, 634 lodimetry, 249 Iodine, 158 — properties of, 160, 161 — dissociation of, 162 — monochloride, 172 — pentoxide, 186 — preparation, 160 — tetroxide, 188 — trichloride, 172 Ionic eqmlibria, 431 ; effect of excess of one ion, 432 — theory, 410, 411 lonisation, degree of, 423 — in solution, 410 — modes of, 420 Ions, migration of, 420 — number of, 421 — properties of, 430 — speed of migration of, 421 Iridiiun, 647 Iron, 627, 630 — alum, 637 — carbonyl, 639 — oast or pig, 626 — complex cyanides, 637 — corrosion, 631 — galvanised, 549 — ore, magnetic, 625, 637 ; spathic, 625 ; specular, 625 — passive, 631 — wrought, 627 Isomeric changes, 361 Isomers, 137, 361 Isomorphism, 118 — atomic weights by, 134 Isotonic solutions, 401 Isotopes, 663 Joule, 20 Kainite, 141, 483, 523 Kaolin, 396 Kerosene, 366 Kieselguhr, 393 Kieserite, 141, 523 Kinetic molecular theory, 63 Kipp's apparatus, 80 Krypton, 276 Kupferniokel, 328, 643 Lampblack, 348 Lanthaniun, 563 Laughing gas, 302 Lavoisier, 1, 6, 7, 17, 21, 48 Law, 5 — approximate, 22, 62, 104 — Avogadro's, 22, 122 — Boyle's, 61 — Charles', 63 — of chemical combination, 23 678 Law of combination by volume, 22, 104 — of combining weights, 29, 33 — of conservation of mass, 20 — • — energy, 15 — of constant composition, 24 — of constant ion product, 432 — Dalton's, of partial pressure, 62 — of Dulong-Petit, 133 — exact, 22 — Faraday's, of electrolysis, 418 — Gas, 22, 61 — Gay Lussac's, of gaseous volumes* 22, 104 — Graham's, of diffusion, 83 — Henry's, of solubility of gases. 111 - — Hess', of constant heat summa- tion, 120 — Isomorphism, 118, 134 — Le Chatelier's, 58, 107, 143, 204, 206 — of Mass action, 203 — ilendeleeff's Periodic, 255 — of Multiple Proportions, 25 — of Octaves, Newland's, 255 — Ostwald's Dilution, 424 — of Reciprocal Proportions, 28 — of Transformation by steps, 310 — van 't Hoff's, 107 Lead, 587 — carbonate, 592 — dioxide, 593 — dioxide and accumulators, 594 ; and salts, 593, 695 — halides, 145, 591 — monoxide and salts, 591 — nitrate, 592 — oxides and salts, 588 — sesquioxide, 596 — silver separation, 509 — suboxide and salts, 590 — sulphate, 592 — sulphide, 593 - — tetrachloride, 595 — tetroxide, 596 — white, 592 Le Blanc process, 476 Lenard rays, 653, 655 Lepidolite, 488 Lime, 529 — milk of, 530 — quick, 530 — slaked, 530 Limestone, 526, 528 Limewater, 530 Limonite, 625 Liquefaction of gases, 68, 273 Liquid air, 273 Litharge, 692 Lithiimi, 488 Litmus, 449 Lixiviation, 476 Luminosity of flames, 384, 385 Magnesia, 523 Magnesite, 522 Magnesium, 522 — carbonate, 526 — chloride, 525 — hydroxide, 524 — nitride, 523 — oxide, 523 — phosphate, 526 — sulphate, 625 — sulphide and hydrosulphide, 623,526 Magnetite, 625 Malachite, 499 Manganates, 619 Manganese, 613 — alums, 617 — dioxide, 52, 58, 144, 618 ; salts of. 618 — heptoxide, 620 — oxides, 146, 614 — spar, 613 — tetroxide, 616 — trioxide, 620 Manganic oxide, 616 — salts, 616 Manganite, 613 Manganites, 618 Manganous oxides, 614 — salts, 614 Marcasite, 637 Marsh gas, 367 Marsh's test, 331 Matches, 310 Mayer, 16 ilayow, 54 Meker's burner, 387 Melaconite, 499 MendeUeff's Periodic Table, 256 Mercuric cyanide, 558 — halides, 13, 557 — nitrate, 559 — oxide, 557 • — sulphate, 559 — sulphide, 559i Mercurous carbonate, 561 — chloride, 660 — nitrate, 560 — oxide, 659 — sulphate, 560 — sulphide, 561 Mercury, 647, 554 — amine compounds, 561 — fulminate, 559 — peroxide, 562 — salts, preparation and properties, 554 Metal ammonia cathions, 498, 550 Metalloids, 41 Metallurgy, principles of, 496 Metals, 41, 97, 453 — extraction from ores, 496 — occurrence of, 468 — preparation of compounds of, 468 Methane, 367 Methods of investigation, 2 Methyl orange, 449 INDEX 679 Mica, 396 Microcosniic salt, 494 Migration of ions, 420 Minivim, 596 Mispiokel, 328 Mixtures and chemical compounds, 38 — separation of, 9 Molar weights, 126 Molecular compounds, 301 — theory, 68 — weights, 124, 134 ; adjustment of, 126 ; boiling point method, 406, 408 ; freezing point method, 407, 408 ; of gases, 123 ; of dissolved substances, 399-411 Molybdates, 611 Molybdenite, 611 Molybdenum, 611 Monad, 87 ilonazite, 563, 597 Mond process for nickel, 358, 643 Mordant, 573 Mortar, 533 ^Mosaic gold, 587 Naphtha, 366 Nascent action, 82, 184 Natural gas, 367 Neodymium, 564 Neon, 276 Nessler's reagent, 561 Neutralisation, 90, 412, 442 — ionically considered, 442, 449 — thermal effects, 412 Newlands, 255 Nickel, 643 — carbonyl, 358, 643 — glance, 328, 643 — pjo-ites, 643 — separation from cobalt, 644 Nicolite, 643 Niobium, 345 Niton, 659 Nitrates, 296 Nitre, 266, 287 Nitric anhydride, 296 — oxide, 290, 300 Nitrides, 268, 275 Nitrites, 298 Nitrogen, 8, 266 — active, 269 — cycle in nature, 288 — fixation, 290 — halides, 285 — hydrides, 277 — iodide, 285 — oxides, 287 — pentoxide, 296 — peroxide, 297 — preparation, 267 — properties, 268 — trichloride, 285 — trioxide, 299 Nitroglycerine, 294 Nitrolime, 545 Nitrosyl chloride, 293 — sulphuric acid, 242 Nitrous anhydride, 299 — oxide, 302 Nomenclature of acids, 174 ; chemis- try, 41 ; oxides, 56 ; periodic acids, 187 Non-metals, 41, 97 Occlusion, 82 Oil gas, 361 Oil, hydrogenation of, 643 Olefiant gas, 368 Oleflnes, 366 Oleum, 246 Olivine, 523 Opal, 393 Open hearth process, 629 Orpiment, 328, 342 Orthite, 597 Ortho-acids, 187 Orthoclase, 397 Osmium, 646 Osmosis, 400 Osmotic phenomena, 401 ■ — pressure, 400 ; abnormal, 427 ,- and molecular weights, 404 Ostwald's dilution law, 424 Oxidation, 54, 166 Oxides, acidic, 57 — amphoteric, 334 — basic, 57 — mixed, 616 — nomenclature, 56 — preparation of, 458 Oxygen, 8, 48 — choice of, as standard, 132 — family of elements, 252 — from liquid air, 50 — in air, 48 — preparation, 50, 594 — properties, 53 Oxy-acetylene flame, 369 Oxy-hydrogen flame, 84 Ozokerite, 366 Ozone, composition, 193 — preparation, 191 Palladitjm, 647 — hydride, 82, 648 Paraffins, 363 Paris green, 507 Parke's desilverisation process, 509 Partition, 112 Patio process, silver, 511 Pattinson desilverisation process, 509 Pentlandite, 643 Perborates, 568 Perehlorates, 183 Periodic law, exceptions to, 264 — System, applications of, 262 ; of classification, 254 Permanent hardness, 531 680 Permanganates, 620 Permutite, 532 Peroxidates, 196 Peroxides, linking of, 198 — preparation of oxygen from, 51 Persulphates, 247 Petalite, 488 Petrol, 366 Petroleum, 365 Phase, 11 Phenol, 293, 377 Phenolphthalein, 449 Phlogiston hypothesis, 5 Phosgene, 361 Phosphates, 319 Phosphides, 313 Phosphine, 311 Phosphonium compounds, 314 Phosphoretted hydrogen, see phosphine Phosphorite, 305, 526 Phosphorus, 305 — commercial uses, 310 — hydrides, 311 — manufacture, 305 ■ — oxychloride, 317 — pentabromide, 317 — pentachloride, 315 — pentafluoride, 315 — pentoxide, 317 ~ red, 307, 308 — sulphides, 326 — tetroxide, 325 — tribromide, 315 — trichloride, 315 — trifluoride, 315 — tri-iodide, 315 — tri-oxide, 322 — yeUow. 307. 308 Photochemical action, 355 Photography, 249, 514 Physical change, 12, 37 Pig iron, 626 Pitchblende, 612 Plaster of Paris, 637 Platinised asbestos, 81 Platinum, 648 — as catalyst, 81, 649 — compounds, 649 Plumbates, 588, 693 Plumbic chloride, 595 Plumbites, 588, 591 Polarisation, 464 Pollux, 491 Polonium, 656 Poly bromides, 474 Polychromates, 607 Polyhalite, 141 Polyiodides, 161, 474, 491 Polymer, 137 Polymorphous substance; 116. 216 Polysulphides, 226, 487 Porcelain, 576 Potassamide, 469 Potash, 478 Potassium, 467 — aluminium sulphate, 674 — antimonyl tartrate, 339 — bicarbonate, 479 — bisulphate, 484 — bromide, 474 — carbonate, 478 — carbonyl, 468 — chlorate, 52, 168, 480 — chloride, 473 — chromate, 605 — cyanate, 486 — cyanide, 485 — cobalt i-cyanide, 642 — cobalti-nitrite, 642 — cobalto-cyanide, 642 — dichromate, 145, 605 — ferricyanide, 638 — ferrocyanide, 484, 637 — fluoride, 475 — hydride, 84, 469 — hydrosulphide, 487 — hydroxide, 473 — iodate, 186 — iodide, 474 — isolation of, 467 — nickelo -cyanide, 644 — nitrate, 482 — nitrite, 482 — oxides, 469 — percarbonate, 487 — perchlorate, 481 — permanganate, 146, 620 — peroxide, 470 — persulphate, 487 — polybromides, 474 — polyiodides, 474 — pyxosulphate, 484 — silver cyanide, 465, 514 — sulphate, 483 — sulphide, 487 — thiocyanate, 486 Potential, 421, 462 — difference, 421 — discharging, 464 — electrode, 460 Pottery, 576 Powder, smokeless, 294 Praseodymium, 563 Precipitation, ionic theory of, 433, 434 Pressure of the atmosphere, 61 — osmotic, 399 — partial, law of, 62 — solution, 462 Priestley, 3, 48 Producer gas, 359 Propane, 364 Properties, arbitrary, 8 — physical, of gases, 60 — specific, 8 Prussian blue, 638 Pucherite, 344 Purple of Caasius, 619 Pyrites, 214, 637, 643 INDEX 681 Pyroantimonates, 341 Pyrolusite, 613 Pyrosulphates, 246, 605 QuAiiTATiVE analysis and ionic theory, 433, 434 Quartation, 517 Quartz, 393 Quick lime, 530 Radicles, 89 Radioactive disintegration, 655, 658 Radioactivity, 655 Radium, 656 — emanation, 659 Ramsay, 3, 275 Rare gases, 275 Reaction endothermal, 16 — exothermal, 16 — velocity and temperature, 58 Realgar, 328, 342 Recrystallisation, 114 Red lead, 596 Reduction, 56, 166 Refining, electrolytic, 464, 501 Refrigeration, 279 Resistance, 417 Reversed combustion, 380 Reversible reactions, 85, 152, 200, 206 Rhodium, 647 Richter, 28 Rook salt, 473 Rouge, 635 Rubidium, 491 Ruby, 571 Rusting of iron, 55, 631 Ruthenium, 646 Rutile, 597 Saltpetre, 160, 186, 467, 482 — Bengal, 288 Salt, 57 — acid, 90 — basic, 90 — complex, 454 — double, 455 — inner, 616 — mixed, 91 — normal, 90 Samarium, 564 Samarskite, 612 Sapphire, 569, 571 Saturated carbon compounds, 363 Scandium, 563 Soheele, 48, 141 Soheele's green, 507 Scheelite, 611 Schonite, 141, 467, 483, 523 Science, 4 Selenium, 251 Semi-permeable membrane, 399 Separation, electromagnetic, 10, 582 Serpentine, 397 Siderite, 625 Silica, 393 Silicane, 390 Silicates, 396 Silico-aoetylene, 390 — -chloroform, 392 — -formic acid, 392 — -methane, 390 Silicon, 388 — amorphous, 389 — carbide, 392 — crystalline, 389 — dioxide, 393 — halidea, 391 — hydrides, 390 — nitrides, 389 — tetrachloride, 392 — tetrafluoride, 391 Silver, 497, 508 — ammonia chloride, 513 ; com- pounds, 498 — carbonate, 513 — electroplating, 514 — fulminating, 513 — halides, 513 — hydroxide, 512 — lead alloys, 509 — metallurgy, 509 — nitrate, 513 — oxide, 197 — oxides, 512 — potassium cyanide, 465, 513 — properties, 511 — sodium thiosulphate, 513 — sub-fluoride, 512 — sulphate, 514 — sulphide, 514 Slag, 496 — Thomas' basic, 629 Slaked lime, 530 Slip, 576 Smalt, 641 Smaltite, 639 Smithell's burner, 382 Soap, 376 Sodamide, 280, 469 Soda water, 353 Sodium, 467 — aluminate, 570, 571, 572 — aluminium fluoride, 573 — amalgam, 469 — bicarbonate, 480 — bisulphate, 484 — bromide, 474 — carbonate, 476 — chlorate, 480 — chloride, 473 — cyanamide, 486 — cyanate, 486 — cyanide, 485 — dichromate, 145, 606 — hydride, 469 — hydrosulphide, 487 — hydroxide, 470 ; action on zinc aluminium, 81 682 Sodium hypochlorite, 481 — hypophosphite, 325 — hyposulphite, 248 — fluoride, 475 — iodide, 474 — isolation of, 467 — nitrate, 287, 482 — nitrite, 482 — oxide, 469 — perborate, 568 — percarbonate, 487 — perchlorate, 481 — peroxide, 51, 470 — persulphate. 487 — polybromide, 475 — polyiodide, 475 — properties, 468 — pjTOSulphate, 484 — silicate, 488 — stannate, 581, 584, 586 — stannite, 581, 584, 585 — sulphate, 110, 483 ; and hydrogen chloride, 151, 207 — sulphide, 487 — tetraborate, 568 — thiocyanate, 486 — thiosulphate, 249, 487 Softening of water, 532 Solder, 583 Solubility, 10, 107 ; — curves, 107 ; breaks in, 109, 479 — of gases, 110 — product, 432 ; and precipitation of sulphides, 434 ; and double de- composition, 433 — repression of, 433 — small crystals, 538 Solution, 72, 113 — colloidal, 71, 395 ; preparation of, 113, 395 — definition of, 113 — freezing point of, 407 — heat of, 107 — normal, 442 — of insoluble compounds, 432 — pressure, 462 — saturated, 109 — supersaturated, 109 — vapour pressure of, 405 Solvay ammonia-soda process, 477 Solvay cell, 472 Spectra, absorption, 491 Spectroscope, 489 Spectrum analysis, 489 Speed of reaction and temperature, 58 Spinel, 569, 572 Stalactites, 531 Stalagmites, 631 Stannates, 585, 586 Stannic chloride, 585 — hydroxide, 586 — oxide, 585 — salts, 581, 586 Stannic sulphide, 587 Stannites, 585 Stannous chloride, 584 — hydroxide, 585 — oxide, 584 — sulphide, 587 Starch test for iodine, 149, 162 Stas, 21, 26 Stassfurt salt deposits, 141, 159 Steam, action upon iron, 85, 212 Steel, 627 — Bessemer process, 629 — cementation process, 628 — crucible process, 628 — electric furnace process, 630 — open hearth process, 629 — tempering, 632 — tool, 600, 611, 612 Stibine, 331 Stibnite, 342 Storage battery, 594 Strontium, 527 — carbonate, 528 — halides, 536 — hydrides, 527 * — hydroxide, 521. 530 — oxide, 529 — peroxide, 534 — sulphate, 521, 537 Strontianite, 118, 526 Substitution, 18, 367 Sugar fermentation, 352 Siiint, 467 Sulphates, acid, 244 — normal, 245 Sulphides, 221, 459 — classification of, 222 — precipitation of, 222, 223 Sulphites, 235 Sulphocyanides, see thiocyanates Sulphur, 214 — amorphous, 217 — bromide, 227 — chemical properties of, 217 — chloride, 227 — dichloride, 227 — dioxide, 229 ; oxidising properties, 233 ; reducing properties, 231 — extraction, 214 — heptoxide, 247 — hexafluoride, 227 — hydrides, 218 — monoclinic, 215 — plastic, 217 — rhombic, 215 — s^quioxide, 248 — tetrachloride, 227 — trioxide, 236 Sulphury 1 chloride, 246, 316 Superphosphate, 320 Supersaturation, 109 Sylvine, 141, 467 Symbols, chemical, 41 Synthesis, 18, 149 INDEX 683 Talc, 523 Tantalite, 345 Tantalum, 345 Tar, coal, 371 Tartar emetic, 339 Tellurium, 251 Temperature, absolute, 64 — and reaction velocity, 58 Temporary hardness of water, 530 Thallium, 578 Theory, 5 — ionic, 410 — kinetic, of gases, 68 Thermochemistry, 118 — of oxy-acids of chlorine, 184 Thio-antimonates, 344 — -antimonites, 343 — -arsenates, 344 — -arsenites, 343 Thiocyanates, 378 Thionyl chloride, 234 Thomas' basic slag, 629 Thorianite, 597 Thorium, 597 Thulium, 564 Tin, 582 Tincal, 564 Tin oxides and salts, 584 Tinstone, 582 Titanium, 597 — test for, 197 Titration, 443 T.N.T., 294 Transformation by steps, 310 Transition point, 216 Triad, 88 Tridymite, 393 Tvmgsten, 611 TumbuU's blue, 638 Type-metal, 329 XJltbamabine, 577 XJltramicroscope, 70 Units of measurement, 20, 119 Unsaturated compounds, 366 Uranium, 612 Uraninite, 276 Urea, 137, 361, 494 Valence, 86 Vanadinite, 344 Vanadium, 344 Van der Waals, 66 Vapour density, measurement of, 74 — pressure, 73 ; of hydrates, 98 ; of water, 95 — osmotic phenomena, 405 Vaseline, 366 Velocity of chemical reactions, 201 Verdigris, 507 - Vitriol blue, 506 — green, 634 Vitriols, 552 Volume, gram molecular, 126 — of gas, influence of temperature and pressure, 64 Vulcanisation, 224, 227 Washing soda, 479 Water, 92 — action of metals on, 79 ; on oxides, 57 — as electrolyte, 446 — composition by volume, 79, 103 ; weight, 100 — Dumas' analysis, 100 — electrolysis of, 78, 446 — gas, 359 ^- glass, 488 — hard, 531 — ionisation of, 446 — Morley's synthesis, 102 — occurrence, 92 — of crystallisation, 97 — proofing, 575, 612 — properties, 93, 96 — softening, 532 — ■ vapour pressure of, 95 Wavellite, 305 Weight, atomic, 129 — combining, 32, 34 — equivalent, 32 — gram molecular, 126 — molar, 126 — molecular, 124 Weldon process, 619 White lead, 592 Willemite, 547 Witherite, 118, 526 Wolfram, 611 Wollastonite, 541 Wood distillation, 373, 374 Wood's metal, 329 Work, 15 Wulfenite, 611 X aAYS, 657 Xenon, 276 Ytterbium, 564 Yttrium, 563 Yttrotantalite, 563 Zibbvogel's silver process, 511 Zinc, 547, 549 — ammonia compounds, 551 — blende, 547 — carbonate, 552 — halides, 551 — oxide, 551 — spar, 547 — sulphate, 552 — sulphide, 552 Zincates, 549, 550 Zincite, 547 Zircon, 396 Zirconium, 597 INTERNATIONAL ATOMIC WEIGHTS Oxygen = 16. Aluminiiun . Al 271 Neodymium Nd 144-3 Antimony . . Sb 120-2 Neon Ne 20-2 Argon . A 39-9 Nickel . Ni 58-68 Arsenic As 74-96 Niobium Nb 93-1 Barium Ba 137-37 Niton . . m 222-4 Bismuth . Bi 208-0 Nitrogen N 14-008 Boron . B 10-9 Osmium Os 190-9 Bromine Br 79-92 Oxygen 16-00 Cadmium . Cd 112-40 Palladium . . Pd 106-7 Caesium . Cs 132-81 Phosphoru 3 P 31-04 Caleivrax Ca 40-07 Platinum . Pt 195-2 Carbon C 12-005 Potassium K 39-10 Cerium Ce 140-25 Praseodym ium Pr 140-9 Chlorine CI 35-46 Radium Ra 226-0 Chromium . Cr 52-0 Rhodium Rh 102-9 Cobalt . . Co 58-97 Rubidium Rb 85-45 Copper. . Cu 63-57 Ruthenium . Ku 101-7 Dysprosium . Dy 162-5 Samarium Sa 150-4 Erbium Er 167-7 Scandivun Sc 45-1 Europium Eu 152-0 Selenium Se 79-2 Fluorine . F 19-0 Silicon Si 28-3 Gadolinium Gd 157-3 Silver . Ag 107-88 Gallium Ga 70-1 Sodium Na 23-00 Germanium . Ge 72-5 Strontium Sr 87-63 Glucinum . . Gl 9-1 Sulphur S 32-06 Gold . . . Au 197-2 Tantalum Ta 181-5 Helium . He 4-00 TeUurimn Te 127-5 Holmium . Ho 163-5 Terbium Tb 159-2 Hydrogen . . H 1-008 Thallium Tl 204-0 Indium In 114-8 Thorium . Th 232-15 Iodine . I 126-92 Thulium Tm 168-5 Iridium Ir 193-1 Tin. . Sn 118-7 Iron Fe 55-84 Titanium Ti 48-1 Krypton . Kr 82-92 Tungsten W 184-0 Lanthanum . La 139-0 Uranium XJ 238-2 Lead . Pb 207-20 Vanadium V 51-0 Lithium . Li 6-94 Xenon Xe 130-2 Lutecium . Lu 175-0 Ytterbium Yb 173-5 Magnesium Mg 24-32 Yttrium Yt 89-33 Manganese Mn 54-93 Zinc Zn 65-37 Mercury . Hg 200-6 Zirconium Zr 90-6 Molybdenum . Mo 96-0 Printed in Great Britain by Butler & Tanner, Frame and London / ^4 .' ' ■'■' -.-■> '"\