(Hmmll Hmwmig fptog THE GIFT OF T^...C T . ( i...^ai.c5^aL.. A,.A.VCA-kS ^\\{ox arV19359° 0rne " Unlvers ">' Lib ™ry liiiiiiH o. in , an ? 1924 031 263 ,, "95l" (Lr (/ The original of this book is in the Cornell University Library. There are no known copyright restrictions in the United States on the use of the text. http://www.archive.org/details/cu31924031263951 A SHORT TEXT-BOOK OF INORGANIC CHEMISTRY BY D R HERMANN KOLBE PROFESSOR OF CHEMISTRY IN THE UNIVERSITY OF LEIPZIG TRANSLATED AND EDITED BY T. S. HUMPIDGE, Ph.D. B.Sc. (Lond.) PROFESSOR OF CHEMISTRY AND PHYSICS IN THE UNIVERSITY COLLEGE OF WALES, ABERYSTWYTH WITH A COLOURED TABLE OF SPECTRA AND NUMEROUS WOOD ENGRAVINGS NEW YORK JOHN WILEY & SONS, 15 ASTOR PLACE ios- EDITOR'S PREFACE. Although the number of short Textbooks of Inorganic Chemistry is large, it is hoped that this little book will supply a definite want among teachers and students, corresponding to that which the Editor has himself felt. The principles which have guided the Author in writing the book are fully stated in his Preface, and with these the Editor entirely concurs ; but in adapting the book for English students certain alterations and additions were necessary, and to these the Author has given his full consent. The whole book has been carefully revised throughout, and the physical constants brought up to date. Considerable additions have been made to the descriptions of water, atmo- spheric air, coal, iron, &c. Short accounts of Gay-Lussac's law, Avogadro's law, and the manufacture of coal-gas have also been introduced into the text. In the Appendix, which is entirely new, a brief account has been given of the methods used for determining atomic and iv Editor's Preface. molecular weights, of Prout's law, and of the Periodic law. The Editor acknowledges his indebtedness to Lothar Meyer's ' Moderne Chemie ' in writing this Appendix. Finally, a series of tables has also been added, which it was thought would prove useful in the laboratory. Some of these have been taken from Landolt and Bornstein's excellent collection of physical tables. The range which the book covers is rather more than that required for the Intermediate Science and Preliminary Scien- tific (M.B.) Examinations of the London University, and the needs of students working for these examinations have been steadily kept in view, but without following the syllabus in a servile manner. Aberystwyth : March, 1884 AUTHOR'S PREFACE. This short Text-book has been written to recall to the memory of students who have attended a course of lectures on Experi- mental Chemistry what they have seen and heard, and to clear up any points which may not have been properly understood. A necessary condition for using a text-book of chemistry successfully is attentive and continuous attendance at the lectures. Students in arts who have not regularly attended a course of lectures may be able to read up afterwards what they have missed ; but a lecture which has not been attended by a student of chemistry cannot be made up by mere reading — neither the notes of the lecture by other students, which ought to be extremely few, nor a text-book, can serve as a substitute for what has not been heard. The chemist has to learn, not by reading nor by hearing alone, but both by hearing and seeing. A person who has not seen the phenomena produced by the union of oxygen and hydrogen, for example, can have no clear conception of them, nor of the chemical change which accom- panies them. Nothing is more foolish than the opinion, which vi Author's Preface. I have often heard from young medical students, that chemistry can be studied from books alone, like other subjects ; that facts which are learnt by heart can take the place of general principles only partially understood. The opinion of chemists as to how many of the immense number of empirical facts now known in chemistry should be introduced into lectures appears to be very different. I quite agree with Fittig when he says : ' Chemistry can no longer be taught as a descriptive science ; in lectures on chemistry the general nature and principles of chemical phenomena should take the first place, and should be illustrated by experiments, but all the many compounds should not be enumerated and described.' But my opinion differs from that of Fittig when he, rather inconsistently, goes on to add . ' But in teaching chemistry in this manner it is still requisite to offer to the student as easily and completely as possible the material on which these general principles are based.' The study of chemistry is similar to that of a language. What is learnt in the lectures scarcely goes as far as reading and parsing ; at most it only includes the rules by which words are built up into sentences. To use the language with success — to speak it — continued practice is required. The same is also true of chemistry : the science is learnt in the laboratory, not in the lecture theatre. The most that can be done in the lectures is to prepare the student for successful work in the laboratory. Author's Preface. vii And although teachers in law, history, and philosophy give the best they have in their lectures, experimental chemistry, as taught in the lecture theatre, must be elementary. To enume- rate many chemical facts loads the memory with ballast, and tends to learning by rote instead of clear perception and after- thought. The problem of the lecturer on chemistry is therefore to give his hearers an. idea of chemical processes and the most important chemical theories without burdening their memories with a large number of mere facts, and thus to prepare them to acquire an accurate knowledge of chemistry by their own practical work. I have adhered to this general principle in writing this short Text-book. I have also endeavoured not simply to give a series of dry facts, but to blend them together into one continuous narrative. Of the students attending chemical lectures, those making chemistry a special study are nearly always in a minority. A complete description of those parts of the science {e.g. the rare elements, the ammoniacal compounds of cobalt and the platinum bases) which have only interest for these few and not for those studying medicine, pharmacy, agriculture, &c, is therefore out of place both in lectures and in elementary books ; and all the more so as the special study of these subjects is better carried on in the laboratory. Such subjects are therefore briefly treated in this Text-book, while others of general interest, such viii Author's Preface. as water, atmospheric air, carbon, carbonic acid, arsenious acid, the detection of arsenic in cases of poisoning, salts, iron, lime, &c, are referred to more fully. The engravings in the text are limited in number : they illustrate only that apparatus which I consider to be especially adapted to recall to the student what he has seen in the lectures. CONTENTS. Introductory . . . i Oxygen . . 8 Hydrogen . 16 Combustion . . -29 Laws of Chemical Combina- tion . . .37 Law of Multiple Propor- tions . . . 42 Law of Volumes . . . 43 The Atomic Theory and Chemical Valency . . 44 Avogadro's Law . 50 Chemical Affinity . . 5 1 Chemical Nomencla- ture . . 54 Chemical Symbols and For- mulae 58 General Remarks on the ' Chemical Elements . . 63 Physical and Chemical Properties of Bodies . . 67 Crystalline Shape . 70 Chemical Properties of Bodies 77 Water. . . 80 Hydrogen Peroxide . 90 Ozone .... 91 The Halogens . , .96 PAGE Chlorine 97 Hydrochloric Acid . . . 104 Oxides and Oxy-acids of Chlo- rine .... 109 Chloric Acid . . .110 Perchloric Acid . . . 113 Chlorine Peroxide . 114 Chlorous Anhydride and Acid 114 Hypochlorous Anhydride and Acid 116 Bromine it8 Hydrobromic Acid . . 120 Bromic Acid . . 122 Hypobromous Acid . 123 Iodine . . . 123 Hydriodic Acid . . . 127 Iodic Acid .... 129 Periodic Acid . . 131 Chlorides of Iodine . .133 Fluorine .... 133 Hydrofluoric Acid . 134 Elements of the Sulphur Group 137 Sulphur . . . 139 Sulphuretted Hydrogen . . 145 Hydrogen Persulphide . 148 Oxygen Compounds of Sulphur 148 Sulphurous Anhydride . . 150 — Acid . . . 154 Hydrosulphurous Acid 15s Sulphuric Acid 155 — Anhydride . . .16^: X Contents. Sulphur — continued. page Nordhausen Sulphuric Acid . 163 Sulphuryl Chloride . . 164 Chlorsulphonic Acid . 165 Nitrosulphonic Acid . 166 Disulphuric Acid . , 166 Thiosulphuric Acid . 167 Dithionic Acid 170 Poly thionic Acids ". . 171 Compounds of Sulphur and Chlorine . . 173 Selenium . . 174 Compounds of Selenium 175 Seleniuretted Hydrogen . 177 Selenious Anhydride and Acid 177 Selenic Acid . . . 17S . Sulphides of Selenium . . 178 Chlorides of Selenium 179 Tellurium 179 Elements of the Nitrogen Group . 1 So Nitrogen . 181 Compounds of Nitrogen and Hydrogen .... 182 Ammonia . . . 183 Hydroxylamine (Oxyam- monia) 186 Amidogen and Imiclogen 188 Ammonium 188 Compounds of Nitrogen and Oxygen - 189 Nitric Acid . 189 — Anhydride . . . 192 — Oxide 192 Nitrous Anhydride and Nitric Peroxide 195 — Oxide 198 Atmospheric Air 200 . Compounds of Nitrogen with the Halogens 204 Nitrogen Chloride . 204 — Iodide 206 Phosphorus . . . . 206 Compounds of Phosphorus and Hydrogen . 211 Phosphorus — continued. page Gaseous Phosphoretted Hy- drogen (Phosphine) . . 211 Compounds of Phosphorus and Oxygen . . 213' Phosphoric Anhydride . 214 Orthophosphoric Acid . 216 Pyrophosphoric Acid . . 219 Metaphosphoric Acid . 220 Phosphorous Anhydride and Acid . . . 221 Hypophosphorous Acid 222 Compounds of Phosphorus and the Halogens . . 223 Phosphorus Trichloride 224 — Pentachloride . 225 Phosphoric Oxychloride 226 Bromides of Phosphorus 227 Iodides of Phosphorus . 228 Compounds of Phosphorus and Sulphur Arsenic Arseniuretted Hydrogen (Arsine) Compounds of Arsenic and Oxygen .... Arsenious Anhydride and Acid Detection of Arsenic in Cases of Suspected Poisoning Arsenic Anhydride and Acid . Compounds of Arsenic with the Halogens . . . Arsenious Chloride, Bromide, and Iodide — Fluoride Compounds of Arsenic and Sulphur Arsenic Disulphide and Tri- sulphide • — Pentasulphide 228 229 231 233 233 237 244 246 247 248 248 248 250 250 Antimony Antimoniuretted Hydrogen (Stibine) . . 252 Compounds of Antimony and Oxygen . . . 253 Antimony Trioxide . . 253 Contents. XI Antimony — continued PAGE Antimonic Anhydride (Anti- mony Pentaxide). 254 — Acid 254 Metantimonic Acid 255 Antimony Tetroxide 256 Compounds of Antimony and the Halogens 256 Antimony Trichloride 256 — Pentachloride 2 57 — Tribromide . 2 57 Sulphur Compounds of Anti- mony . . ... 258 Antimony Trisulphide 258 — Pentasulphide '259 Boron 260 Boric Anhydride and Acid 262 Boron Trichloride . 264 — Trifluoride . 265 — Sulphide 266 — Nitride 266 Silicon 266 Compounds of Silicon 268 Silica and Silicic Acid . 268 Silicon Hydride 272 — Chloride . 273 — Fluoride 274 Fluosilicic Acid 27S Silico-chloroform - 276 Silico-formic Anhydride ; Silicoxalic Acid . 277 Carbon ... 277 Compounds of Carbon . 286 Car bon — contin ued. PAGE Carbonic Anh\ dride or Acid . 286 Carbamic Acid 2 93 Carbonic Oxide . . . 294 Carbon Oxychloride 297 Oxalic Acid 297 Oxamic Acid and Oxamide 299 Carbon Disulphide . . . 300 — Oxysulphide 302 Compounds of Carbon and Hydrogen . . . 3°3 Methane or Marsh Gas . . 3°4 Ethylene 306 Carbon Tetrachloride 309 Cyanogen 3" Hydrocyanic Acid 3*3 Cyanic Acid . 316 Sulphocyanic Acid . 317 Chlorides of Cyanogen 3i3 Coal Gas 319 Titanium . . 322 Titanic Anhydride and Acid . 323 Titanium Tetrachloride 324 Fluotitanic Acid . . 325 Molybdenum . . . 325 Molybdic Anhydride and Acid 326 Tungsten 327 Tungstic Anhydride and Acid 328 Vanadium . . . 329 Niobium and Tantalum . . 330 THE METALS. The Metals Chemical Constitution of Salts Spectrum Analysis . Classification of the Metals . Metals of the Alkalies . Potassium .... Compounds of Potassium Potassium Oxide and Per- oxide ... 33 1 Potassium — continued. Potassium Hydrate • 346 332 — Sulphate . • 349 337 — Acid Sulphate , . • 349 34° — Nitrate • 35° Gunpowder • 35i 342 Potassium Nitrite . 3S 2 342 — Carbonate • 353 346 — Acid Carbonate • 354 — Chlorate 354 346 — Perchlorate • 354 Xll Contents. Potassium — continued. page Potassium Hypochlorite . 354 — Oxalate . . . .355 — Silicate . . . . 35S — Chloride . . . 356 — Bromide . . . 356 — Iodide . . 356 — Fluoride . . . 358 — Acid Fluoride . . 358 — Fluoborate . . . . 358 — Fluosilicate . 359 — Cjanide . . 359 — Cyanate . . 361 — Sulphocyanate . . . 361 Sulphur Compounds of Po- tassium .... 362 Potassium Monosulphide . 362 — Sulphydrate . 363 — Trisulphide and Penta- sulphide . . . . 363 Liver of Sulphur . 363 Detection of Potassium Com- pounds . . . . 364 Sodium ... . 365 Compounds of Sodium . . 367 Sodium Oxide and Peroxide . 367 — Hydrate . . . . 367 ■ — Sulphate . . . 367 ■ — Acid Sulphate . 368 — Sulphite .... 368 — Acid Sulphite . 368 — Thiosulphate . . 368 — Nitrate . , . 369 — Carbonate . . . 369 — Acid Carbonate . . . 374 — Phosphates . . 374 — Borate ..... 376 - — Chloride .... 376 Detection of Sodium Com- pounds . . ... 378 Lithium 378 Lithium Hydrate . 379 — Sulphate . . , 379 — Phosphate . . 379 — Carbonate . , . 380 — Chloride . . . . 380 Lithium — continued. Detection of Lithium Com- pounds Rubidium and Cesium .. . Ammonium Ammonium Sulphate — Nitrate — Phosphates , — Carbonate — Chloride . . . . ■ — Sulphides . Oxyammonium Compounds . Detection of Ammonium . Compounds Metals of the Alkaline Earths Calcium Calcium Oxide . ■ — Hydrate . — Sulphate — Nitrate — Phosphate — Carbonate Bleaching Powder Calcium Silicates Glass Calcium Oxalate — Chloride ■ — Fluoride .... — Sulphides Detection of Calcium Com- pounds . .... Strontium .... Strontium Oxide — Sulphate . — Nitrate — Carbonate — Chloride -- Sulphide . Detection of Strontium Com- pounds Barium Barium Oxide — Hydrate — Peroxide . — Sulphate 380 380 382 382 383 383 383 385 386 386 387 388 3S8 389 39° 394 396 396 398 400 401 401 404 404 405 4°5 406 406 406 407 4°7 407 407 407 408 408 40S 409 410 411 Contents. xiii B AK IUM — continued. PAGE Aluminium — continued . PAGE Barium Nitrate 4" Ultramarine 427 — Carbonate . 411 Aluminium Chloride 428 — Chloride . 411 — Fluoride 428 — Sulphides . . . 412 — Sulphide . 428 Detection of Barium Com- Detection of Aluminium Com- pounds 412 pounds 429 Magnesium 413 Bfryllium . 429 Magnesium Oxide 414 Beryllium Oxide and Hydrate 430 — Hydrate . 415 — Sulphate . . 430 — Sulphate 415 — Chloride 43° — Phosphate? 416 — Carbonate 430 — Carbonate 417 Gallium . 431 — Borate 417 Indium 432 — Silicate . 417 Indium Oxide 43 2 — Chloride 4 T 7 — Sulphate 43 2 Detection of Magnesium — ■ Chloride 432 Compounds 418 — Sulphide . . . 432 Metals of the Earths 419 Yttrium, Terbium, Erbium, Aluminium 419 and Ytterbium 433 Aluminium Oxide 421 Cerium, Lanthanum, and — Hydrate . 421 Didymium . 433 — Salts . 422 Thorium . 434 — Sulphate 422 Zirconium .... 434 Potash-alum 4=3 Zirconium Oxide and Hydrate 435 Aluminium Phosphate . 425 — Sulphate 435 — Silicate 425 — Chloride 435 -Porcelain and Earthenware 426 — Fluoride 435 HEAVY METALS. Iron . . 436 Ferrous Oxide and Hydrate 445 — Sulphate 445 — Carbonate 446 — Phosphate 447 Ferric Oxide and Hydrate 447 — Sulphate . 448 — Nitrate . 449 — Phosphates . 449 Triferric Tetroxide 449 Ferric Acid 45° Ferrous Chloride 45° Ferric Chloride . . 451 Ferrous Iodide 451 — Sulphide . 452 Ferric Disulphide 452 Compounds of Iron and Cvanogf.n 453 Potassium Ferrocyanide 453 Ammonium Ferrocyanide 456 Sodium Ferrocyanide . . 456 Frussian Blue . 456 Hydroferrocyanic Acid . 457 Potassium Ferricyanide. 457 Turnbull's Blue . . . 458 Hydroferricyanic Acid . 458 Sodium Nitroprusside . 458 Detection of Iron Compounds 459 Manganese . 460 j Manganous Oxide and Hy- drate . 461 XIV Contents. Manganese — continued. Manganous Sulphate — Nitrate . . . . — Carbonate — Chloride — Sulphide . Manganic Oxide and Hydrate — Sulphate — Chloride Trimanganic Tetroxide . Manganese Peroxide Manganic Acid Potassium Manganate Permanganic Acid Potassium Permanganate Detection of Manganese Com pounds . Chromium Chromic Oxide . — Hydrate . — Sulphate — Nitrate — Phosphate — Chloride . Chromous Chloride Chromic Acid Chromates Potassium Dichromate — Chromate Sodium Chromate . Ammonium Chromate and Dichromate Chlorchromic Acid and Chro- myl Chloride Detection of Chromium Com- pounds Uranium Uranic Nitrate — Oxide . — Sulphate — Carbonate — Phosphate Uranyl Chloride — Sulphide . Uranous Oxide . ■ — Chloride . Uranoso-uranic Oxide PAGE 461 462 462 462 463 463 463 464 464 464 463 465 466 466 476 468 469 470 47i 472 472 472 473 473 474 47S 476 477 477 477 479 479 480 480 480 482 482 482 482 Cobalt and Nickel Cobalt ... 484 Cobaltous Oxide and Hydrate 485 Cobaltic Oxide and Hydrate . 485 Cobaltous Sulphate 486 — Nitrate 486 — Nitrite 486 — Phosphate 486 — Carbonate 486 — Oxalate 487 — Chloride . 457 Cobaltic Chloride and Am- monia 487 Cobaltous Cyanide. . . . 488 Potassium Cobalticyanide 488 Cobalt Sulphides 488 Smalt 489 Detection of Cobalt Com- pounds .... 489 Nickel .... 490 Nickelous Oxide and Hydrate 490 Nickelic Oxide and Hydrate . 491 Nickel Sulphate 491 — Nitrate . 491 ■ — Carbonate 491 — Chloride 491 — Cyanide . 491 — Sulphide 492 Alloys of Nickel 492 Detection of Nickel Com- pounds 49= Zinc . Zinc Oxide — Hydrate — Sulphate — Nitrate ■ — Carbonate — Chloride — Sulphide .... Detection of Zinc Compounds Cadmium .... Compounds of Cadmium Detection of Cadmium Com- pounds. . . . Lead 493 494 495 495 495 495 495 496 496 497 497 498 Contents. xv LEAD — continued. PAGE RlSMUTK— continued. PAGE Lead Suboxide . 5°i Detection of Bismuth Com- — Oxide • Soi pounds 515 — Hydrate 502 Tin .... 516 Red Lead 502 Stannous Compounds 517 Lead Peroxide . • 5°2 — Oxide and Hydrate 5*7 Salts of Lead • S°3 — Chloride 5i8 Lead Sulphate . • • 503 — Sulphide .... 5i8 — Nitrate . • S04 Stannic Compounds . 519 — Phosphate S°4 — Oxide and Hydrate 519 — Silicates . • 5°4 - — Chloride 520 — Carbonate • • S°4 — Sulphide 520 White Lead • 5°4 Alloys of Tin . ... 521 Lead Acetate 505 Detection of Tin Compounds 522 — Chromate . • 5°6 Copper 5 2 3 — Chloride — Iodide • • S06 506 Cupric Compounds - — Oxide . 524 524 — Sulphide • 5°7 — Hydrate 5 2 5 Alloys of Lead S07 • — Sulphate 5 2 5 Detection of Lead Com- — Ammonium Sulphates 526 pounds • 508 — Nitrate 5 2 7 Thallium . • ■ 5°9 — Phosphate 5 2 7 Thallous Compounds Sio — Arsenite S27 Thallic Compounds • • S12 — Carbonates 527 Detection of Thalliurr Corn- — Acetates 528 pounds . s' 2 • — Chloride . S 2 8 — Bromide 528 Bismuth . • 513 — Sulphide 528 Bismuth Suboxide 514 Cuprous Compounds S 2 9 — Oxide and Hydrate • 5H - — Oxide and Hydrate . 5 2 9 Bismuthic Acid • • 514 — Chloride 5 2 9 Bismuth Nitrate . SIS — Iodide 53° — Chromate 5'5 — Sulphide 53° — Chloride • SIS Alloys of Copper ' 53° — Sulphide 5*5 Detection of Copper . 53i NOBLE u METALS. Mercury . S33 Mercury — continued. Mercuric Compounds • 53S Mercuric Sulphide . 538 — Oxide • S3S Mercurous Compounds 539 — Sulphate S3S — Oxide S39 — Nitrate . • 536 — Nitrate . . S39 — Chromate'. • 536 — Chromate . 54° — Chloride . : • S3 6 — Chloride . . . 54° — Iodide 538 — Bromide . S4i — Cyanide ■ ■ S38 — Iodide . . . 54i XVI Contents. Mercury — continued. PAGE Silver — continued. PACE Mercurous Sulphide • 54i Silver Sulphide • 55° Amalgams . • • 54i Alloys of Silver . • 551 Detection of Mercury Com- Detection of Silver Com- pounds . • • 542 pounds . • 5Si Silve ; . . . 543 Gold .... 552 Silver Oxide 545 Auric Chloride . -553 — Peroxide . • 54° — Oxide and Hydrate 554 — Sulphate • • 546 Potassium Aurate 554 — Nitrate 546 Auric Cyanide • 554 — Phosphate . 547 — Sulphide 555 — Carbonate • 547 Aurous Chloride 555 — • Chromate - ■ 548 — Oxide 555 — Chloride, Bromide, and — Cyanide . 555 Iodide . • 548 Purple of Cassius 556 — Fluoride • S5° Alloys of Gold 556 — Cyanide • 550 PLATINUM METALS. Platinum 558 Palladium — continued. Platinic Chloride 56i Palladous Nitrate . ■ 567 Platinous Chloride 562 Palladic Chloride . 568 Platinic Oxide and Hydrate 562 Palladous Chloride 568 Platinous Oxide and Hydrate 563 — Iodide . 568 Platinic Sulphide 563 Iridium . 568 5% Platinous Sulphide — Cyanide 563 5 6 3 Compounds of Iridium Ammoniacal Compounds of Rhodium . 570 Platinum 563 Osmium ■57° Palladium 566 Osmic Acid 57° Palladium Hydride 566 Ruthenium 57i APPE NDIX. Atomic and Molecular Weights . Rklations between the Atomic Weights Tables ... INDEX . 573 S82 586 .593 TABLE OF SPECTRA Frontispiece TEXT-BOOK OF INORGANIC CHEMISTRY. INTRODUCTORY. THE science of Chemistry is closely related to that of Physics, and it is the aim of both sciences to solve nearly the same problems. The province of both is to investigate the processes of nature per- ceptible to the senses, to discover the course they take, their con- nection with one another, and their causes, and to investigate the laws according to which the forces of nature act, whatever may be the objects or classes of bodies upon which their action is exerted. Thesetwo sciences — Chemistry and Physics — are the foundation of all other natural sciences — of mineralogy and geology, of zoology and botany, as well as of astronomy, and especially of the whole of medical science, which endeavours to fathom the processes pro- duced in a certain class of bodies by chemical and physical forces. If we consider what goes on around us in nature, we perceive the change from day to night, we notice the moon and apparently the sun revolving round the earth, as well as similar movements in other heavenly bodies ; we feel the movements of the air, perceive the change of temperature ; we notice water freezing, and admire the rainbow, the thunderstorm, formation of clouds, &c. All these phenomena belong to the processes of physics ; but where shall we, then, look for chemical processes ? Simple chemical processes are not nearly so common in nature as simple physical processes, and this is certainly the reason why physical phenomena and physical laws were investigated many B .9 2 Text-Book of Inorganic Chemistry. centuries before any idea of a natural law in chemistry was con ceived. We notice plants growing and animals breathing, both which processes go on by the inter-action of chemical and physical forces. But how plants assimilate the constituents of atmospheric air, and work them up chemically into their substance, still remains a mystery to the chemist and physicist at the present day. When a tree is struck by lightning and burns, this is certainly a chemical process, and one much simpler than the growth of a plant. But although man has known from time immemorial how to produce fire, and has had the chemical process of burning wood daily before his eyes, a correct explanation of a process apparently so simple was only obtained at the end of the preceding century. For this reason chemistry may well be called the younger sister of physics. The intelligent observer also perceives numerous other chemi- cal processes going on in nature ; he notices the weathering of rocks, the putrefaction of animal and vegetable matter, the deposi- tion of limestone in caverns, the rusting of iron, the souring of milk, wine, &c. ; but these and similar changes usually -go on so slowly that their progress is difficult to follow. Innumerable chemical processes were introduced by man thousands of years ago. It was found that lime which had lain in the fire — quick-lime — became hot when brought into contact with water, that it possessed generally other properties than the lime which had not been burnt ; metals had been extracted from their ores by heating with charcoal ; white arsenic had been obtained by roasting other ores ; sulphur had been burnt in the air and the suffocating acid gas so produced had been noticed. But these and a hundred other similar changes could only be correctly understood and explained after the common cause of them all — fire — and the process of combustion had been made clear. For this reason, the discovery of oxygen— of that substance which is as necessary for combustion as it is for respiration — bv Priestley and by Scheele, and Lavoisier's first correct explanation of combustion, are the pillars which carry the noble edifice of modern chemistry, based on the foundation of numerous earlier observations. In what manner, then, do chemical and physical processes differ from one another — which phenomena belong to chemistry and which to physics ? The answer may be easily given in a few words, Introductory. 3 but will not be so easily understood by those to whom chemical phenomena are strange. Chemical processes are those by which bodies undergo a material change. In physical processes, pure and simple, the substances in which we notice a change remain materially unchanged. But what is to be understood by a material change ? This will best be explained by a few examples. A stick of sulphur when rubbed becomes electric and attracts small pieces of paper or other light bodies, and then repels them'. The same sulphur, if heated in a test tube, melts to form a pale yellow, clear liquid, and changes at a higher temperature to a reddish-brown gas, which externally has not the remotest similarity with the solid yellow sulphur. But has the sulphur by the rubbing or by the heating undergone a material change, have chemical pro- cesses taken place in both cases ? The answer to this question is a negative. The substance of the sulphur is the same after the rubbing as before, and sulphur gas, as well as liquid sulphur, have the same relation to solid sulphur as water gas or water vapour, and liquid water to solid water or ice. We have here to do with a change in the state of aggregation pro- duced by heat, which only lasts as long as the cause producing it. Just as water gas by cooling again becomes first liquid water and then solid water, so the sulphur gas changes into liquid sulphur and then into solid sulphur, with exactly the same properties as it possessed at first, as soon as the source of heat is removed. But something quite different happens when we heat sulphur in an open vessel, so strongly that it catches fire and burns with a blue flame. In this case, too, it changes entirely into a gas, the colourless gas with the well-known odour of burning sulphur. If we cool the gas, for example, by leading it through a vessel sur- rounded with cold water, we should in vain expect to obtain liquid or solid sulphur from it. By this process the sulphur has suf- fered a material change — i.e. a change which continues after the cause has been removed. The sulphur has united chemically with a constituent of the air, and a new substance has been produced. If we intimately mix together yellow sulphur and metallic iron, both in the state of a fine powder, we obtain a grey substance, which externally has no similarity with either of the bodies from which it was produced. It might be thought that this grey-powder no longer consists of yellow sulphur and black iron, and that the b 2 4 Text-Book of Inorganic Chemistry. two substances had undergone a material change by the ^act of mixing them. We can, however, readily convince ourselves that this mixture really contains unchanged sulphur and iron. The lighter sulphur may be easily separated by the mechanical process of washing, and the iron may be removed by a magnet — which would not be pos- sible with the product of real chemical action. Let us prepare a mixture of four parts of sulphur and seven parts of iron, both finely powdered, and heat the mixture in a test-tube in the gas lamp. In a short time that portion of the mass which is most strongly heated commences to glow ; the glowing then spreads through the entire mixture, and finally ceases of itself. It seems remarkable that at this high temperature none of the volatile sulphur is given off. Sulphur and iron, which when only mixed remain chemically unchanged, attract one another chemically when heated to a cer- tain high temperature, and then both of them suffer a material change. The dark-grey, solid, hard product may be easily pulver- ized, but this powder has now quite different properties to the mere mixture of sulphur and iron. We might try in vain to separate the sulphur by washing or by any other mechanical operation, or to extract the iron by the magnet. Sulphur and iron experience a material change when heated together, and in such a manner that it might be asked whether the new product really contains sulphur and iron, or whether both have not been lost by the operation. This question cannot be decided a priori, nor by mere philoso- phising, but only by experiment, and it has been so decided. If the sulphur and the iron had been lost on their union, if they were no longer contained in the compound, it would not be possible in any imaginable manner to again separate them with their original properties. But chemistry teaches that, in this and other cases, compounds may be again decomposed into their con- stituents, and that the same quantities of these constituents may be again obtained. The process of building up a chemical compound from its constituents, as here illustrated in the case of sulphur and iron, is called synthesis : the reverse of this is analysis. The union, or, as we may say, the coalescence — of the consti- tuents of a chemical compound is so close that by no means can either of them {e.g. the yellow sulphur in the above compound) be perceived in the dust of a compound, not even with the most powerful microscope. It is thus clear why it is impossible to split Introductory. 5 up a chemical compound by pure mechanical means. Substances united by chemical force can be only separated by some force acting chemically. Those substances which we cannot further decompose by chemical analysis we call simple bodies, or elements, without wish- ing to assert that they are really undecomposable. They are elementary substances for us only as long as we have not succeeded in decomposing them. The inquiry after the fundamental substances of inorganic and organic nature is as old as the investigation of nature itself; but what we understand to-day by the word chemical element is some- thing quite different from what was formerly so called. The state- ment of the early philosophers, who neglected experiment and exact investigation, that Fire, Air, Water, and Earth wereelements, rather meant that all natural bodies have their origin in these four things, than that they consist of them. And when this question was more critically examined and experimentally investigated, these four elements of the ancients were abandoned. It is now known that what is called fire is not a substance but a phenomenon ; that water consists of two gaseous bodies, as yet incapable of further decomposition, oxygen and hydrogen ; that air is a mechanical mixture of two gaseous substances, oxygen and nitrogen ; and that earth is no single individual substance, but an agglomerate of thousands of very different chemical elements and compounds. Our present chemical elements are the result of experiment. Their number is somewhat large, and, according to experience, will probably increase. We already know more than sixty. When it is considered how easy it is for us to decompose water into its constituents, and to reproduce it from them, the question arises, how was it possible that up to a hundred years ago water was thought to be an elementary substance, how was it possible that one of the many methods which we now know for its decom- position, had not been discovered much earlier ? The principal agent, which was previously almost exclusively employed in order to produce chemical decomposition, was a high temperature, or fire. Crucibles, heating furnaces, distilling appa- ratus were most indispensable in furnishing the laboratory of an alchemist. If water had possessed the property to be decomposed at a high temperature, its compound nature would have been re- cognized long previously. That it remains unchanged at a red- 6 Text-Book of Inorganic Chemistry. heat could only help to confirm the belief that it was undecompos- able. The methods which we now use to decompose water were partially unknown to the chemist of earlier centuries, and were partially wrongly interpreted. To the former, belong the electrolytic force of the galvanic current, and the action of those metals, only discovered later, which decompose water at the ordinary temperature : potassium, sodium, &c. To those processes by which water is decomposed, and which were long known, but were previously falsely explained, belongs the reaction which occurs when iron is brought into contact with water and a strong acid. Everyone who sees this experiment for the first time believes, as was formerly done, that the gas evolved from the iron is extracted from this metal, whereas it really arises from the decomposition of the water. Let us consider this and other processes by which water is de- composed more exactly. If we dip two platinum plates into water which is made slightly acid with sulphuric acid (pure water is a bad conductor of the cur- rent), and then connect the plates with the poles of a galvanic battery of four or six Bunsen's cells,. we notice at the moment when the circuit is completed that gases are evolved from both platinum plates, and that this evolution continues as long as the galvanic current remains unbroken. If two small tubes of equal size, filled with water, are placed over the platinum plates, so that the gas evolved from one plate is collected in the one tube, and that evolved from the other plate in the other (fig. i), we notice the following results : — i. The gas collecting over the positive electrode x regularly fills a smaller space than that collecting over the negative electrode ; exaet measurements have shown that the volume of the latter is just twice as much as that of the former. ii. If a lighted body is brought near the gas from the negative pole, when it is allowed to flow out by opening the stopcock, it catches fire and burns with a feebly luminous flame. The gas from the positive pole does not possess this property ; it does not burn itself, but it supports the combustion of other bodies in such 1 The two pieces of platinum which bring the current to the water are called poles or electrodes. The one which is connected with the zinc of the battery is the negative pole or electrode, the other which is attached to the copper, carbon, or platinum of the battery is the/ Water = ioo-oo „ This percentage composition of water, calculated from the pro- portions by volume in which hydrogen and oxygen unite with one another and from the specific gravities of the two gases, agrees exactly with innumerable analyses and syntheses of water which 38 Text-Book of Inorganic Chemistry. have been made from time to time. 1 It has been further proved that pure water, whatever may be its source, whether obtained from ice, snow, or water vapour, or whether prepared artificially from its constituents, has always the same composition. A chemical compound containing oxygen and hydrogen, like water, but in some other proportion, is not water. We are acquainted with such a compound, called hydrogen peroxide, to which we shall refer later on. This body contains only 5-9 per cent, of hydrogen, and 94-1 per cent, of oxygen. The composition of all the more accurately known chemical compounds has, like that of water, been carefully determined, and we know not only of what elements they consist, but also in what proportions they contain these elements. From these results of analytical chemistry two laws of extreme importance in discussing the regularities in the composition of chemical compounds have been deduced : — i. On the chemical union of two substances their original weight remains unchanged. ii. Every chemical compound contains the simple substances of which it is composed in one, and only one, proportion by weight. From this, however, it does not follow that two substances which contain the same simple substances in the same proportions are therefore identical. Numerous isomeric compounds exist, which, although they have the same elementary and percentage composition, possess different properties, and are quite different bodies. 2 1 One of the simplest methods for determining the composition of water by Weight depends upon the fact that red-hot copper oxide is reduced in a stream of hydrogen to metallic copper, while its oxygen unites with the hydrogen to form water. If we collect this water by some compound which will absorb it (e.g. calcium chloride) and weigh the tube containing this substance before and after the experiment, we know the weight of water which has been formed. And if we also weigh the copper oxide before and after the experiment we find the weight of oxygen which has combined with the hydrogen to produce the known weight of water. It is then always found that 100 parts by weight of water contain exactly 88 '89 parts by weight of oxygen, the remaining 11*11 parts being hydrogen. — Ed. 2 Isomeric compounds, or those possessing the same percentage composi- tion, may be divided into two classes — viz. those which have the same mole- cular weight, but in which the atoms are differently arranged — e.g. propionic acid jc'o-OH' ant * metn ''- acetate Jco^OCH-' ca,lec - "'etamcric com- pounds ; and those of which the molecular weight of the one is some multiple The Lazvs of Chemical Combination. 39 •^ O N, « 8 O CO *+■ n. co CM JN o o « CO CO M 8 *o ;o 13 ON 8 8 a M "O -a IN CM on 8 8R H O W- lO H CO 8 8S IO Tf CO o IS M O CM 'C H IO CJ p. M o O "5 15 s tn u .2 £ £ g.2 j= -1 3 .3 3.2 3.3 -S.S S 3 a;" S b/)==i bfi==j H 'H-'rt H ^ ^ ^"3 "P — Oh X ^3 OS- OH *j3 Oh 3fi "3£ 3fi 3 j3 oh LO lO 8 ON H 8 ■^vO 8 vO H CO O IN JO pM3 O 03 K 1 JO ^t- ;o to JO !>. w 8 Q CO O ON CO rp fni b si vo en 'O W- b O H IO CO ON O 0) l-t CO o H o CS IN o O ON CO o 1> i-i 01 •a H M O H T3 H N l-l 3 M £ '? 'S -S* O O 3 3 T3 to O Gl T3 •d 3 E 1=1 0) c 1) 8 O 00 \Q M PI IN 8 Q ON O CO LO M W N o co -a eo^o O ON m cn u io ■*■ CO CO "H H h ro ' w H M .*3 H '* a o O *3 2 Cfi u ft p. ft - • a u s u C Jh G tH ft o § S3 1- Js ft o QJ OJ 0) a; p U -H aj •r cj • S QJ M ft boa boa ■s °* •s ^ •S a °& -2 S 5 a >*Oi >»a >*a £>Q* ^a 5-tK X O M o ^ ° 3 o 3 o 3 o J3 J3 o ,£ o OU OU OU crtU WU c^U OU UU UU 00 « o t>. M CO o lO vo ■+, o ON H I O io M CO o « « tN'j CO JO oo h o 8 3 CO M "h- to 8 o>b VO H is ci b ON N io H CO H O ON o M O CO I H So o " H IS d c c -a c G C '0 < (3 c fi ■ ^ 1- V-. OK OX offi 3 •3 um us CjK — r— — r— -T— O. — r^ — ■ r— — ,— -nr— ■"^^ — r- ■3 -C M 3 en - "QJ have still no simple relation to one another. If, however, we compare this proportion with that in which the same four elements unite with equal weights of sulphur (100 parts), we see from line 11.6 that this proportion is the same, for, dividing 6-25,198-12, 646-8, 1275-3 by 6-25, we get as before I : 31-7 : 103-5: 204; and further, the fatioin which the four elements unite with chlorine (lines 111.6, lll.c) is also the same. Here, then, we at once perceive definite regularities which may be formulated in the following statement : — The same proportional weights of hydrogen, copper, lead, and thallium always unite with one definite weight of oxygen, sulphur, or chlorine. These elements do not, however, occupy any exceptional posi- tion, and all that has been said for their combining proportions holds also for all the others. We thus obtain the following general law : — The proportion in which two bodies (A and B) coi?t6ijie with a third (C) is also the proportion in which they unite with all other 6odies as well as with one anotlier. of that of the others— e.g. ethylene (CnH^) and propylene (C 3 H 6 ), c&We&polymcric compounds. Examples of these compounds are more common in organic than in inorganic chemistry. — Ed. The Laws of Chemical Combination. 41 Thus, from the accompanying table, we see that the weight of Chlorine (35-5 parts) which unites chemically with 1 part of hydrogen is the same as that which unites with 317 parts of copper, 103-5 parts of lead, or 204 parts of thallium. And, again, that the same relative quantities of hydrogen, copper, &c, which unite with 35-5 parts of chlorine combine also with 16 parts of sulphur or with 8 parts of oxygen. We have thus obtained the following numbers for these ele- ments, which represent the relative weights in which they combine with one another and with other elements : — Hydrogen .... . ro Oxygen 8-o Sulphur 16-0 Chlorine .... . . 35-5 Copper 317 Lead ....... 103-5 Thallium 204-0 Again, suppose we wish to discover in what proportions chlorine and oxygen unite with one another, it is sufficient to know the proportions in which they combine with a third body. We see from the preceding table that 103-5 parts of lead combine with 8 parts of oxygen or 35-5 parts of chlorine. The proportion 8 of oxygen to 35-5 of chlorine is therefore that in which these two elements unite with one another. Hypochlorous anhydride con- tains exactly 8 parts of oxygen united with 35-5 parts of chlorine. In this manner the number representing the quantity of any other element which combines with 1 part of hydrogen, 8 parts of oxygen, 35-5 parts of chlorine, &c., may be easily obtained. If we wish to find this number for the metal sodium, and know that the percentage composition of its compound with chlorine (common salt) is : Sodium 39-32 Chlorine 6o-68 100 00 we need then only calculate the quantity of sodium which would combine with 35-5 parts of chlorine. This we find to be 23 from the proportion : 6o-68 : 39-32: 135-5 '• n i an d tfl i s number, therefore, expresses the weight of sodium which combines with 1 part of hydrogen, 8 parts of oxygen, 16 of sulphur, &c. The relative weights in which bodies combine with one another, 42 Text-Book of Inorganic Chemistry. and in which they displace one another from their compounds, are called their equivalent weights: 317 parts of copper, or 103-5 parts of lead always combine with 8 parts of oxygen, 16 of sulphur, or 35-5 of chlorine. If in a compound of 35-5 parts of chlorine and 103-5 parts of lead, the lead were displaced by some other substance— £.£v copper or hydrogen— the quantity of this substance required would not be the same as the lead displaced, but 31-7 parts of copper, or 1 part of hydrogen, would be necessary. In these compounds I part of hydrogen has the same effect as 31-7 parts of copper or 103-5 parts of lead, and these elements, therefore, displace one another in their compounds not in equal but in equivalent weights. In order to displace 35-5 parts of chlorine from its compound with 1 part of hydrogen or with 204 parts of thallium by iodine, 127 parts of the latter body are required. On the other hand, to displace the 204 parts of thallium by hydrogen, only 1 part is necessary. Finally, if in the compound called copper sulphide, consisting of 16 parts of sulphur united with 31-7 parts of copper, silver be substituted for copper, it is found that 16 parts of sulphur unite with 108 parts of the former metal. THE LAW OF MULTIPLE PROPORTIONS. The number of chemical compounds would be very limited if the elements could only combine with one another in one single proportion. Experiment has, however, shown that most of the elements can unite with other elements in more than one propor- tion, many even in five different proportions, thus producing com- pounds which are usually very different from one another. In such cases, however, the relative quantities of the two ele- ments combining together are always some multiple of their equivalent weight by small whole numbers. As an example of this may be instanced the five compounds of nitrogen and oxygen, nitrous oxide, nitric oxide, nitrous anhydride, nitric peroxide, and nitric anhydride. These compounds contain their constituents in the following proportions by weight. The Law of Multiple Proportions. 43 Nitrous oxide Nitrogen, parts • H Oxygen j parts 8 Parts Nitric oxide ■ H 16 = 8x2 Nitrous anhydride Nitric peroxide . Nitric anhydride • H • 14 • 14 24 32 40 = 8x3 = 8x4 = 8x5 From which it is at once seen that the quantities of oxygen united with one equivalent of nitrogen (14 parts) are multiples of its equivalent weight by the simple integers 1:2:3:4:5. In the same manner one and the same quantity of iron unites with two different quantities of chlorine which are to one another in the proportion 2 : 3. These and innumerable other results form the experimental basis of the Law of Multiple Proportions, which maybe expressed as follows : — The elements unite either in their equivalent weights or in simple multiples of their equivaletit weights. The law likewise holds for compounds as well as elements ; when the former combine to form more complex compounds they do so either in their equivalents or in simple multiples of their equivalents. The statement of this law is due to Dalton : it is only about seventy years old. Dalton first published his theory in detail in the year 1808, at a time when exact chemical investigation was largely carried on in England, France, and Sweden, but when the leading chemists in Germany, following an erroneous path in natural philosophy without any real foundation of fact, long wasted their powers in worthless speculation. The theories of Dalton and his conclusions derived from the law of multiple proportions were, however, more permanent and fruitful. THE LAW OF VOLUMES. 1 Not only is there a simple relation between the weights in which the elements combine together chemically, but for elements in the gaseous state there is also a simple relation between the 1 The Editor is responsible for this statement of Gay-Lussac's law. 44 Text-Book of Inorganic Chemistry. volumes of the constituents and of the compound, if gaseous, which is produced. It will be remembered that when hydrogen gas and oxygen gas unite with one another chemically, they always do so in the pro- portion of two volumes of the former to one volume of the latter, and the volume of the water vapour produced (above 100°) is always two volumes. Further, one volume of hydrogen gas always unites with exactly one volume of chlorine gas, producing exactly two volumes of the gaseous compound called hydrochloric acid. These may serve as examples of a general important law — the law of volumes — which may be formulated as follows : — The volumes of two gases uniting chemically -with one another are always in some simple proportion, and the volume of the cotnpound produced, when gaseous, is also always in some simple proportion to those of its constituents. This law also holds for compound as well as for elementary gases. We owe its discovery to the celebrated French chemist Gay-Lussac, after whom it is often called. THE ATOMIC THEORY AND CHEMICAL VALENCY OR ATOMICITY. The stability with which two elements remain combined after their chemical union, the impossibility to again recognize in water, for example, any property of either of its components, and again to break up water by any mechanical means into hydrogen and oxygen, led to the idea that on the chemical combination of two bodies a penetration of matter and a similar coalescence occurred as that when a mixture of copper and zinc are melted together to form brass. In such alloys neither of the constituents can be dis- tinguished or separated by mechanical means. It is impossible, even with the most powerful microscope, to perceive the copper and zinc in brass, or the gold in gold amalgam. Still, these alloys are not generally considered as chemical compounds, because the metals of which they are composed generally unite in almost any proportion, while the combination of bodies to form real chemical compounds always takes place in certain definite proportions. If we melt together copper and zinc in those proportions by weight in which they both unite with oxygen, sulphur, chlorine, &c. The Atomic Theory. 45 — i.e. in their equivalent weights, an alloy is produced which might be considered as a chemical compound ; but such an alloy is also produced when we employ a little more or a little less copper or zinc than exactly corresponds to their equivalent weights. In no case does either copper or zinc remain behind unchanged, so that we can separate it from the compound. But it is otherwise with real chemical compounds. As the metals may be fused together in any proportion, so may hydrogen and oxygen be intimately mixed in all proportions. But if we transform this mechanical mixture of the two gases into their chemical compound — viz. water, we know that the smallest quantity of hydrogen or oxygen present in the mixture above the equivalent proportion remains behind uncombined and unchanged. If the chemical union of two substances depended upon an intimate coalescence and penetration of the two bodies, it would be difficult to understand why combination only occurs in certain definite proportions by weight, and it would be incomprehensible that, when it does occur in different proportions, this only happens per saltum, in simple multiple proportions. Dalton's atomic theory affords a simple and satisfactory ex- planation of these results. We believe, as Dalton did, that the substances which compose chemical compounds are not capable of indefinite division, but that they consist of particles, which cannot be further divided, either by physical or by chemical means. These particles are therefore called atoms. We imagine that chemical combination is produced by the attraction and juxtaposition of the atoms. The more firmly the atoms are united together, the stronger must be the force by which the combination is produced and maintained. This chemical force of attraction is quite different from any other physical force. It differs, for example, from the force of gravity, which acts through great distances, by which the earth is attracted to the sun, and a falling stone to the earth ; it differs also from the forces of cohesion and adhesion. Chemical attraction never acts at a distance, like gravity, but requires immediate contact of the two bodies, just as the passage of an electric current requires the direct contact of the conducting bodies. Chemical attraction, as we have seen (p. 5), differs also from cohesion, by producing a material change in the bodies upon which it acts. The chemical force which resides in the atoms, called chemical affinity or attraction, is evidently related to the force of electricity. 46 Text-Book of Inorganic Chemistry. Many chemical processes are produced by electricity — combina- tions as well as decompositions. In the electrolysis of water, the hydrogen atoms are set free at the negative pole, and the oxygen atoms at the positive pole. And since bodies charged with oppo- site kinds of electricity attract one another, the conclusion might be drawn that the hydrogen which passes to the negative pole is charged with positive electricity and the oxygen with negative. In the same manner, on the electrolysis of metallic salts, the metals are liberated at the negative pole, and the acids or haloid constituents at the positive pole. From these and other similar results, it was thought that the atoms were charged in different degrees with positive or negative electricity, and that the combination of two heterogeneous atoms was caused by the attraction of the different electricities : this was the fundamental idea in the electro-chemical theory. But this theory leaves so many phenomena unexplained, and even contradicts many chemical reactions, that we do not at pre- sent employ it, although the existence of near relations between chemical and electrical forces cannot be denied. The fact which particularly contradicts the theory is, that the atoms of one and the same element, which were supposed by the theory to be charged with the same kind of electricity, attract one another and one generally form very stable compounds. We content ourselves with the hypothesis that a force produc- ing chemical phenomena resides in the elementary atoms, and that the atoms of. the elements, according to their nature, offer one or more points of attraction to other atoms. When the atoms of heterogeneous elements attract one another by the forces present in them, and form a chemical compound, the number of the atoms of the one element which can be attracted by one atom of the other element is always limited, and, as we may imagine, depends upon the number of points of attraction which the atoms possess. This number is different with different elements. We know that when hydrogen and chlorine unite with one another chemically they always do so in 07ie definite proportion. From this and other results we conclude that the atom of hydrogen only offers one point of attraction to that of chlorine and to those of other elements. The result that when hydrogen unites with oxygen to form water twice as much of the former gas is required as when it unites with chlorine leads to the conclusion that the atom of oxygen The Atomic Theory. 47 possesses two points of attraction with which it secures the two atoms of hydrogen. When carbon unites with hydrogen to form marsh gas and with oxygen to form carbonic acid, the carbon atom does not fix the same number of atoms of the two substances. The atom of carbon, as proved from the composition of innumerable compounds, possesses four points of attraction for the atoms of other elements, and therefore requires four atoms of hydrogen for saturation and only two of oxygen, which offer, not one, but two points of attraction to other atoms. Two atoms of oxygen thus play the same part, in combination with one atom of carbon and with all other elements, as, four atoms of hydrogen ; or, what is the same thing, one atom of oxygen has the same value or valency as two atoms of hydrogen. Those elements of which the atoms, like those of hydrogen, only possess one point of attraction are called monads (mono- valent or monatomic), those with two points of attraction dyads (divalent or diatomic), those with three triads (trivalent or tri- atomic), those with four tetrads (tetravalent or tetratomic), &c. No elementary atom appears to possess more than seven or eight points of attraction, and, as we have remarked, the number of atoms of an element which can be fixed by one elementary atom is, .therefore, limited. Although there are some organic compounds which contain several dozen atoms united with one another, we know how to resolve these apparently complicated proportions into others more simple, resembling the inorganic compounds. The power of the elements to fix a certain number of the atoms of other elements is called their atomicity or valency, and we say that the atomicity of hydrogen is one, that of oxygen two, and that of carbon two or four. Every element possesses a definite greatest atomicity, or, in other words, a certain maximum number of points of attraction of its atom for other atoms. But the atoms of those elements possessing higher atomicity do not always exert their full power. Sulphur, of which the maximum atomicity is six, and which occurs as a hexad element in sulphuric acid, plays the part of a tetrad in sulphurous acid, and of a dyad in sulphuretted hydrogen. We imagine that when all the points of attraction of an atom are not active they are in a state of rest, or that they are latent, until aroused by some new chemical action. Two of the points of 48 Text-Book of Inorganic Chemistry. attraction of nitrogen, the atom of which is pentad in ammonium chloride, are latent in ammonia, in which this element plays the part of a triad. The addition of hydrochloric acid, or indeed of any acid, to the ammonia suffices to awaken these two latent points of attraction and to convert the compound of triad nitrogen into one of the pentad element. In other cases the action of heat or light is also necessary to make the latent points of attraction active: The four points of attraction of the atom of carbon, of which in the compound with oxygen called carbonic oxide only two are saturated, obtain their full value by combination with oxygen to form carbonic acid at, a high temperature, or by combination with chlorine, producing phosgene gas under the influence of direct sunlight. We say that sulphur plays the part of a dyad element, or, more briefly, that it is a dyad in sulphuretted hydrogen, that it is a tetrad in sulphurous acid, and a hexad in sulphuric acid ; and we speak of dyad carbon in carbonic oxide, of triad nitrogen in ammonia without touching the question of what may be the highest atomicity of these elements. We simply pay regard to the results of experiment, according to which the elements unite with one another in certain multiple proportions. The chemical compounds which are produced by the union of the atoms are mechanically indivisible, but can be decomposed into their constituents by chemical forces. The smallest quan- tity of a chemical compound which can exist in the free state is called a molecule. By a molecule of water we mean the compound of two atoms of hydrogen with one atom of oxygen ; by a molecule of carbonic acid, the compound of one atom of carbon with two atoms of oxygen. The molecules of water, carbonic acid, iron sulphide, &c, are just as little perceptible or visible as the atoms. What we call water is an agglomeration of water molecules, and in the same manner the gaseous or solid carbonic acid collected in a vessel consists of a large number of carbonic acid molecules. The question may here be asked, Are the elements as we know them — e.g. ordinary hydrogen gas or solid sulphur — agglomerates of their respective atoms ? Not long ago this was thought so ; we now know that the atoms of the elements can combine not only with the atoms of other elements, but also with themselves. When atoms of hydrogen are set free from a chemical com- The Atomic Theory. 49 pound, and when no other substance is presented to them for which they possess a stronger affinity, they unite with one another to form molecules of hydrogen. By the chemical union of two atoms of hydrogen a molecule is produced, and in the same manner a molecule of oxygen consists of two atoms of this element, a molecule of nitrogen of two atoms, &C. 1 In the molecules pro- duced by the union of the atoms of the same element, the homo- geneous atoms are not bound together with less, but often indeed with greater, affinity than in the compounds of heterogeneous elements. For example, the molecule of nitrogen is a much more stable compound of the two atoms of the element than the com- pounds of nitrogen with oxygen, or with hydrogen or chlorine. Molecules of the same substance never exercise chemical ac- tion upon one another. Chemical combination or decomposition always occurs between the constituents of the molecules of dif- ferent bodies. Two molecules of hydrogen act upon one another as little as do two molecules of chlorine, but if hydrogen and chlorine are mixed, and exposed to the action of light or heat, a decomposition of the molecules of both elements ensues, with the mutual exchange of the atoms of each. Two molecules of hydro- chloric acid are produced from one molecule of hydrogen and one of chlorine. And although in this case, as in all others, it appears to be the molecules which act upon one another, it is in reality the atoms which unite with one another, and which are separated from molecular combination at the moment when the reaction com- We know from experiment that the elements at the moment of their liberation from a compound act more energetically upon sub- stances with which they come in contact than in their ordinary (molecular) condition. Hydrogen and sulphurous acid in the di y state or in the presence of water have no action at all upon one 1 Different molecules may exist of those elements which have a higher atomicity. It is possible that, besides the molecule of nitrogen, which perhaps consists of the two atoms of triad nitrogen united by their three points of attraction, a second kind of nitrogen molecule may exist in which the two atoms are united together by five points of attraction. Up to the present we only know one kind of nitrogen ; but phosphorus, sulphur, charcoal, and other elements exist in different forms or modifications. It is possible that the ordinary phosphorus, which is so easily ignited, owes its difference from the red amorphous phosphorus, a. body endued with much less chemical affinity, to some such cause. 50 Text-Book of Inorganic Chemistry. another, but if we add sulphurous acid to a mixture of zinc and sulphuric acid, which evolves hydrogen, the hydrogen at the moment when it is set free (in the nascent state) exercises a chemical action upon the sulphurous acid : it separates the mole- cules of this compound, uniting with the sulphur to form sulphu retted hydrogen, and with the oxygen to form water. We explain this and all similar phenomena by the hypothesis that in the nascent state the atoms, not the molecules, exert the chemical action, and that the atoms, with their entire force, natu- rally produce chemical changes which the elements in their mole- cular state — after the elementary atoms have united together — cannot effect. AVOGADRO'S LAW. 1 We have here to refer to a very remarkable law concerning the number of molecules contained in different gaseous substances. This law, which is proved from the physical properties of gases may be thus stated : — Equal volumes of all gases, both elementary and compound, when under the same physical conditions — i.e. at the same temperature and under the same pressure— contain an equal number of molecules. It follows from this that molecular weights of all gases must occupy equal volumes when under the same conditions, or, in other words, that their molecular weights must be in the same pro- portion as their densities. If, then, we know the density of a gas or vapour, whether simple or compound, we can at once find its molecular weight, and conversely, if we know its molecular weight its density may be easily calculated. The molecular weight of hydrogen, which consists of two atoms 'and has therefore the relative weight 2, is taken as the normal volume, and is said to occupy two volumes. The molecular weight of every other gaseous substance will therefore also occupy two volumes under the same physical conditions, and to this law there are very few exceptions. The molecular weight of oxygen is 16 x 2, and this gas should therefore be 16 times as heavy as hydrogen, which agrees with experiment (sp. gr. of hydrogen = 0-0692 ; sp. gr. of oxygen => 1-1056 = 16 x 0-0692). Similarly, the molecular weight 1 Introduced by the Editor. Avogadro's Law. 51 of water, consisting of two atoms of hydrogen and one of oxygen, - Q is 2 x i + 16 = 18, whence water gas should be _ = 9 times as heavy 2 as hydrogen; and, finally, the molecular weight of hydrochloric acid, consisting of one atom of hydrogen and one of chlorine, is 1 +35'5=3°'5; the gas should therefore be ^1 = 18-25 times as 2 heavy as hydrogen. These and innumerable other results also agree closely with experimental determinations. Further, ,the weight of a litre of hydrogen being known (0-0895 gramme at 0° and 760 mm.), we can easily find the weight of a litre of any other gas at the normal temperature and pressure, provided we are acquainted with its molecular weight. Thus the weight of a litre of oxygen under normal conditions is 3 2 18 0-0895 x ^T = r '43 2 gramme; of water-gas, 0-0895 x — =0-806 * 2 gramme; of hydrochloric acid gas, 0-0895 * - — -= 1*633 gramme, 2 The densities of gases and vapours are usually compared with that of air as a standard — i.e. the density of air is taken as unity. In order to find from these densities the various molecular weights we must multiply by 14-44, which is the relative density of air com- pared with hydrogen ( , ), and again by 2, because this is the molecular weight of hydrogen. In other words, we multiply by 28-88. The density of water gas, for example, is 0-622, and its molecular weight 0-622 x 28-88 = 17-96, or very nearly the same as its theoretical molecular weight (18). 1 CHEMICAL AFFINITY AND UPON WHAT IT DEPENDS. The chemical force of attraction (affinity) which resides in the atoms of the elements is of different strength. For example, the affinity of potassium for chlorine is so much greater than it is for iodine, that if chlorine is led into a solution of potassium iodide 1 For a more detailed account of Avogadro's law, the student is referred to the Appendix.— Ed. e 2 >'5 2 Text-Book of Inorganic Chemistry. ■the chlorine unites with the potassium, and all the iodine is set free. In the same manner, the affinity of iron for chlorine is so much greater than that of copper, that if a bright rod of iron is placed in a solution of copper chloride, it becomes at once covered with a coating of metallic copper, and at last all the copper is separated out from the compound in the metallic form, while an .equivalent quantity of iron unites with the chlorine which was pre- viously combined with the copper. But, it may be asked, How much greater is the affinity of potassium for chlorine than for iodine, or of chlorine for iron than for copper ? And it might be thought that the experimental answers to such questions ought not to be more difficult than the measure- ment of the force of gravity, the strength of an electric current, and . other physical forces. But chemistry is in this respect far behind physics, and not only because chemistry is a younger science, but also because chemical affinity is largely influenced by a number of iimportant conditions which often occurtogether — by temperature, light, electricity, state of aggregation, quantity, &c, and which often cause a powerful affinity to be weakened, and a weak affinity • to be strengthened. Some examples of this may serve to show how difficult it is to determine the force of these oft-changing chemical affinities. Influence of Temperature on Chemical .Affinity . At ;the ordinary temperature, mercury and oxygen do not ;,possess sufficient affinity to combine with .one another, nor is their affinity strong enough to produce this change at the boiling-point ■ of water ; but at a temperature of about 300° they unite chemically with one another, and form red oxide of mercury. This compound is, however, again decomposed into its constituents at a slightly higher temperature (under 400°). At 4oo°mercury and oxygen have as little affinity for one another as at the ordinary temperature. The metal potassium, when moderately heated in a stream of car- bonic acid gas, abstracts the oxygen from this compound to form an oxide, and then a carbonate, charcoal being separated. But if potassium carbonate is heated with charcoal up to'bright redness, the affinity of the charcoal for the oxygen becomes greater than that of the metal for the same element, and the potassium is again set free. Charcoal, which at the ordinary temperature has a very weak affinity for oxygen, obtains .the most powerful affinity for this element at a -red heat. Chemical Affinity, and upon what it depends. 53; Influence of Light on Chemical Affinity. A mixture of about equal volumes of hydrogen and chlorine, when quite excluded from the light, remains unchanged, and from this we conclude that the atoms of these elements in their molecules do not possess sufficient affinity to combine with one another. But if the mixture is exposed to light, the combination of the two gases, producing hydrochloric acid, goes on the more quickly the more intense the light. It is, further, principally the violet rays of light which produce this chemical change ; and these rays are therefore called the chemically active (actinic) rays. If the mixture contains pure hydrogen and pure chlorine in exactly equal volumes, diffused daylight produces instantaneous combination. Light is thus able to awaken the affinity of hydrogen and chlorine for one another, and so to cause them to combine. On the other hand, stable com- pounds — e.g. those of silver with chlorine or bromine — are some- times decomposed by light. We thus see that light also can cause both chemical combination and decomposition. Influence of the State of Aggregation and of Quantity {Mass) on Chemical Affijiity. Potash has a stronger affinity for sulphuric acid than lime, and sulphuric acid has a stronger affinity for bases than acetic acid ; but notwithstanding this, if we mix a solution of potassium sulphate with one of calcium acetate, calcium sulphate separates out and potassium acetate remains in solution. In this case it is the insolu- bility of calcium sulphate (which is nearly insoluble in water) that causes the stronger acid (sulphuric acid) to separate itself from the stronger base — potash — and to unite with the lime. In the same manner, the quantities or masses of different bodies acting upon one another also influence chemical changes. Sulphuric acid has a much stronger affinity for metallic oxides than hydro- chloric acid, still the latter can partially displace the former in solutions of the sulphates, and the more completely the more it is in excess. If, for example, strong hydrochloric acid in considerable quantity is added to the blue aqueous solution of copper sulphate, free sulphuric acid is obtained with copper chloride, the latter being easily recognized by the green colour which the blue solution gradually acquires. These examples suffice to show that the force of chemica 54 Text-Book of Inorganic Chemistry. affinity residing in the atoms of the element is difficult to measure, since it is influenced and changed in very various ways by other secondary forces and conditions. Before we are in a position to measure accurately the strength of these secondary forces, even when several act together, we shall not be in a position to predict accurately and to calculate mathematically what would be the re- sult when substances, acting chemically upon one another, are brought into contact under different conditions ; although this is always the ultimate goal of all chemical investigation. But as long as we remain far distant from this end, as at present, chemistry is a science in which mathematical calculations can be little used. 1 CHEMICAL NOMENCLATURE. The impossibility of inventing empirical names for many thousand chemical compounds has been long admitted, and at an early date it was proposed to form the names so that everyone might easily understand the nature of any particular compound, and to what class it belonged, by the name given to it. But besides these rational designations, other empirical names, which come to us with the authority of age, and which also possess the merit of being shorter, are still employed, especially by those unacquainted with the principles of chemistry. The chemist occasionally makes use of the names : soda, potash, Glauber's salts, lunar caustic, &c, although in inorganic chemistry such names are but little employed. In organic chemistry, on the other hand, where we have to refer to whole classes of bodies, of whose chemical composition we know but little, as sugar, starch, the glucosides, &c, we must content ourselves for the present with the empirical names. 1 Other things being equal, the strength of the affinity which causes two substances to unite with one another is proportional to the quantity of heat set free on their combination. The quantity of heat evolved when a molecule of any compound is produced under the same physical conditions is always the same, and if certain changes are possible among a system of bodies, that change will take place which produces the greatest amount of heat. Thus, if two elements, A and B, are mixed with a third, C, with which each can unite directly, combination will take place between those two which liberate the greater quantity of heat on their union. — Ed. Chemical Nomenclature. 55 The compounds produced by the union of the elements, and of these compounds again combined together, mostly belong to three classes — bases, acids, and salts. The compounds of oxygen with the other elements, the most numerous of all, are called oxides ; in the same manner, those of sulphur are sulphides, those of chlorine are chlorides, of bromine bromides, &c. The oxides may be divided into two great classes — basic oxides and acid oxides \ those not belonging to either of these classes are mostly indifferent compounds. Those oxides which have a basic character — i.e. which combine with acids to form salts — may be called the oxides proper, and their name is built up from the name of the element with which the oxygen is combined. Thus we say potassium oxide, zinc oxide, &c. In some cases we are acquainted with two basic oxides of an element. We know, for example, two basic oxides of copper, iron, mercury, and other metals. In order to clearly distinguish these from one another, the termination of the metal is changed ; the termination -ic is added to the metal to indicate the compound con- taining the greater quantity of oxygen, and the termination -ous to indicate the compound containing the smaller quantity of oxygen. Thus we say mercuric oxide and mercurous oxide. In cases where these terminations when added to the name of the metal would be harsh, the Latin name of the metal is employed : we say, for example, ferric oxide and ferrous oxide for the two oxides of iron, cupric oxide and cuprous oxide for those of copper, &c. Those basic oxides, as ferric oxide, chromic oxide, which contain the atoms of the metal, and oxygen in the proportion of two to three, are often called sesqui-oxides, or one-and-a-half oxides. Oxides of the metals, which possess neither acid nor basic character, and which contain more oxygen than the basic oxide, are called peroxides. We are acquainted, for example, with man- ganese peroxide, lead peroxide, &c. They give up the excess of oxygen which they contain when heated with strong acids. Suboxides, of which the number is very small, are those indif- ferent oxides of the metals which contain less oxygen than the basic oxides. On treatment with acids they yield the metal and the basic oxide, which is richer in oxygen. Such an oxide is lead suboxide. The oxygen compounds which are said to have an acid charac- ter, and are sometimes called acids, are those which unite with 56 Text-Book of Inorganic Chemistry. bases to form the ordinary or oxygen salts. They form with water the true oxy-acids, and are called anhydrides. The names are built up in the same way as those of the oxides, except that the termi- nation -ic is always added if there be only one compound of the element with oxygen. Thus, the single oxide of carbon of an acid character is called carbonic anhydride. We further distinguish between those anhydrides of the same element which contain different quantities of oxygen by the same terminations as are used for the oxides {-ic and -mis) ; for example, the two anhydrides of sulphur are sulphuric anhydride and sul- phurous anhydride ; of nitrogen, nitric anhydride and nitrous anhydride. Some elements — e.g. chlorine — are capable of forming a larger number of oxides with acid characteristics. These are distinguished by the prefixes per- and hypo-. The oxide which contains more oxygen than chloric anhydride is called perchloric anhydride, that which contains less than chlorous anhydride is called hypochlorous anhydride. The oxy-salts are named from the bases and acids of which they consist. Salts formed from acids with the termination -ic, have the termination -ate, while those from acids with the termina- tion -ous end in -ite. For example, potassium sulphate is the com- pound of potassium oxide and sulphuric acid ; cupric nitrate, of cupric oxide and nitric acid ; mercurous nitrite, of mercurous oxide and nitrous acid. The name hydrate includes two classes of bodies : the acid hydrates (the oxy- or sulpho-acids) and the basic hydrates (the hydrates proper). The acid hydrates, or as they are better called acids, and which are produced by the union of the anhydrides with water, usually contain in their molecule the same number of hydrogen atoms as the normal salt of a monad metal contains of a metal. 1 The names of the acids are derived from those of the corresponding anhydrides ; thus, from sulphuric anhydride we get sulphuric acid, from phosphoric anhydride, phosphoric acid, &c Basic hydrates are the oxides of the metals in which a portion of the metal is displaced by hydrogen — e.g. potassium hydrate, cupric hydrate. They are often produced by the union of the 1 An acid is therefore a hydrogen salt, and the hydrogen which can be displaced by a metal, not necessarily all which the acid contains, is called its Jhplaceable hyJrogen. Accordingly, as an acid contains one, two, three, &c., atoms of displaceable hydrogen, it is said to be mono-, di-, tribasic, &c. — Ed. CJiemical Nomenclature. 57 basic oxides with water, and may be considered, as was formerly the case, as compounds of these two substances. The chemical nomenclature of the acids, bases, and salts, which contain sulphur instead of oxygen is not so well developed, and probably because we are less accurately acquainted with these compounds. They may be called sulpho-acids, sulpho-bases, and sulpho-salts. On the union of a sulpho-acid {e.g. carbon disul- phide or arsenious sulphide) with a sulpho-base {e.g. potassium sulphide or sodium sulphide) a sulpho-salt is produced. Thus carbon disulphide (sulpho-carbonic acid) with sodium sulphide gives sodium sulpho -carbonate; arsenious sulphide (sulph-arsenious acid) with potassium sulphide gives potassium sulph-arsenite, and in this way all other sulpho-compounds are distinguished. The salts which result from the direct union of the halogen elements (chlorine, bromine, iodine, and fluorine) with the metals, and which contain neither oxygen nor sulphur, are called the haloid-salts, and the corresponding acids, containing hydrogen in the place of the metal, are the haloid-acids. Hydrochloric acid, the compound of hydrogen and chlorine, is not to be confused with chloric acid, which contains oxygen as well as these elements. The metals unite with the halogens usually in the same pro- portion as with oxygen, and the corresponding names are the same. To mercuric oxide corresponds mercuric chloride, con- taining a quantity of chlorine equivalent to the oxygen in the oxide; similarly, mercurous. chloride corresponds to mercurous oxide. In cases when elements unite with the halogens in more than two proportions, or when, as in the case of phosphorus, the com- pounds cannot be included under the true haloid-salts, it is customary to indicate the number of chlorine, bromine, &c, atoms contained in the molecule of the compound by the prefixes mono-, di-, tri-, letra-, fienla-, hexa-, &c. The two compounds of phosphorus and chlorine, which contain respectively five and three atoms of chlorine united to one atom of phosphorus, are called phosphorus pentachloride and phosphorus trichloride ; and, in the same manner, we also say carbon disulphide, referring to the compound containing one atom of carbon and two of sulphur in the molecule. Haloid compounds are also known, in which a portion of the chlorine, bromine, &c, is displaced by oxygen. These are called 58 Text-Book of Inorganic Chemistry. oxychlorides, oxybromides, &c. To this class belongs the well- known phosphorus oxychloride. Those compounds in which sulphur displaces a portion of the halogen instead of oxygen are similarly designated — e.g. the compound of phosphorus, chlorine and sulphur is called phosphorus sulphochloride. Alloys are the intimate mechanical mixtures produced by melting metals together. Those alloys which contain mercury are called amalgams. Alloys and amalgams can only be produced from metals. CHEMICAL SYMBOLS AND FORMULA. When we calculate with figures we do not use the words representing them, but their signs. In the same manner, the need of symbols to express chemical substances, and by a suitable combination of these symbols to reproduce chemical thoughts in a brief but general manner, was early recognized in chemistry. It is customary to represent the elements by letters and to express chemical compounds by the juxtaposition of these letters. The symbol H has been chosen to represent the element hydrogen, and I the element iodine ; by placing these symbols together, without any sign between them, the compound of the two elements — hydriodic acid — is represented (HI). In the same way the compound symbol CaO is a simple expression for the compound of the metal calcium (Ca) with the gas oxygen (O) — i.e. common quick-lime. The symbols of the elements are the first letter of their names and usually of their Latin names : thus K represents potassium (kalium), C = carbon, &c. When the names of several elements begin with the same letter, some other letter is added : thus B means boron, Ba = barium, Br = bromine, Mg = magnesium, Cd = cadmium, Ag = silver (argentum), Hg = mercury (hydrargyrum), Fe = iron (ferrum), Sn = tin (stannum), &c. But the chemical symbols of the elements have another and more important meaning. The symbol HI does not only indicate the presence of hydrogen and iodine, but expresses also the quantities by weight of the constituents of hydriodic acid which are contained in its molecule. The symbol H indicates one atom of hydrogen, or the smallest weight which enters into a chemical Chemical Symbols and Formula. 59 compound (which is taken as unity), and, in the same manner, I means one atom of iodine or 127 parts by weight ; while, finally, the compound symbol H I expresses the compound of hydrogen and iodine, which contains one part of the former united with 127 parts of the latter, and of which the molecular weight is, therefore, 127 + 1 = 128. The accompanying table on p. 60 gives the names of the elements, alphabetically arranged, with their symbols and atomic weights, the non-metallic elements being in italics. Since, however, the elements combine with one another not only in simple but also in multiple proportions, and as this must also be represented by chemical symbols, it is customary to ex- press the number of atoms of the one element which unites with the other to form a molecule by a small figure placed to the right of the symbol of the element. For example, S0 2 means the compound of one atom (32 parts) of sulphur with two atoms (2x16 parts) of oxygen, and the formula H 2 means the compound produced by the union of two atoms of hydrogen (2 parts) with one atom of oxygen (16 parts). The figures placed to the right of any chemical symbol refer only to this one symbol, but another large figure placed on the left of a symbol multiplies all that follows up to a comma, full-stop, or plus sign. The formula 2KCl,PtCl 4 , for example, represents the compound of two atoms of potassium, and two atoms of chlorine (or two molecules of potassium chloride), with one of platinic chloride. Similarly, the formula S0 2 < QNa + I0 ^fi * s tne symbolic expression for a compound of sodium sulphate with 10 molecules of water — i.e. with 20 atoms of hydrogen and 10 atoms of oxygen. Every such formula is full of hidden meaning to the chemist. The composition of the above compound (Glauber's salts) might be more briefly expressed by the formula SNa 2 H 20 O 14 , but this only tells us that one molecule contains one atom of sulphur, two of sodium, twenty of hydrogen, and fourteen of oxygen. If we wished to express the fact that ten molecules of water are contained in each molecule of the salt, we might write its formula S0 4 Na 2 + ioH 2 0. But the above formula, S0 2 j q^ + ioH 2 0, in which the constituents of the compound are still further arranged, expresses the fact that the four oxygen atoms in the sodium sul- 6o Text-Book of Inorganic Chemistry. Table of Atomic Weights. Name s £ "tJb Name E B to < % w = SO,{gg The hydrates may be called hydrogen salts, and there are two classes to be distinguished : (i.) those in which the hydrogen takes the place of a metal — the oxy-acids ; (ii.) those in which the hydro- gen displaces an 'acid-radical — the hydrates proper — as, calcium hydrate, potassium hydrate (KOH), &c. Many of these compounds are easily decomposed, some even (copper hydrate) under ioo°. Others are so unstable that they cannot exist at the ordinary temperature of the air, or under ordinary conditions generally. If a strong base is added to an aqueous solution of a mercuric salt, mercuric oxide and not mercuric hydrate is precipitated ; and from an aqueous solution of a carbon- ate, a strong acid will liberate carbonic anhydride (commonly called carbonic acid), and not its hydrate or the true acid. The crystals of many salts, which crystallize without water, often enclose small quantities of water mechanically. On heating these crystals they break into pieces, often with violence, or, as we say, they decrepitate. Water is deposited on solid bodies from the air, and the quantity depends upon the nature of the substance, temperature, relative moisture of the air, &c. In order to remove this moisture, or generally to abstract loosely combined water from bodies which will not bear heating, the substances are placed under a desiccator — i.e. in a closed space in which the air is kept dry, and the water- vapour which is given off is chemically absorbed. This may be best done by concentrated sulphuric acid. The simplest construc- tion for a desiccator is a bell-jar with ground edge, which is placed on a ground-glass plate over a vessel of sulphuric acid and a separate vessel containing the substance. The air in the bell-jar becomes 88 Text-Book of Inorganic Chemistry. quickly dried by the sulphuric acid, and then the moisture, which is evolved from the water into the dry space, is also absorbed. The process may be much accelerated, if the air in the bell-jar is partially or entirely exhausted. Such an arrangement is shown in fig. 34. The ground-glass plate is cemented to a plate of iron, which is borne by a brass support. A hole is made in the glass Fig- 34- and through the iron plate down the brass support to the side tube, above which is a stop-cock. The side tube is connected by a thick piece of india-rubber tubing with a water-air pump, as shown. If now the pump is set in action, the air under the bell-jar soon be^ Water. 89 comes partially exhausted, and, by closing the stop-cock, may be preserved for several days in this condition. In order to drive off hygroscopic water, or water of crystalliza- tion, from those bodies which can be heated without suffering decomposition, an air-bath is used. Such an apparatus of the simplest construction is shown in fig. 35. It consists of a copper vessel of which the parts are brazed together and of which the front side opens to form a door. The substance to be dried is placed on a piece of wire gauze, about 6 centimetres from the bottom of the vessel. After closing the door, the vessel is heated by a gas-lamp, and the flame regulated by the thermometer inserted into the bath from above. In order to get rid of the water- vapour which is thus produced, an opening is made in the top near the thermometer, and a second in one of the side walls just over the bottom. By this arrangement, a slow and continual stream of air flows through the warm bath, carrying with it the water from the substance which is to be dried. go Text-Book of Inorganic Chemistry. HYDROGEN PEROXIDE. Composition : H 2 0„. Besides water there exists another compound of oxygen and hydrogen which contains twice as much oxygen as water. This compound, called hydrogen peroxide or hydroxyl, was discovered by Thenard in 1818. It is a colourless, transparent, viscid liquid, with a peculiar odour, mixes with water in every proportion, is not yet known in the solid state, and is very easily decomposed. The pure com- pound is decomposed even at 15 or 20° into oxygen and water ; in an aqueous solution, especially when containing a little hydrochloric acid, it is much more stable. Hydrogen peroxide cannot be prepared directly from oxygen and hydrogen, but is formed when the oxygen of peroxides com- bines ■ with hydrogen or water. Some peroxides (for example, those of manganese or lead) cannot be used for its preparation ; it is best to employ barium peroxide, obtained by heating barium oxide in a stream of oxygen. The barium peroxide, when finely powdered and mixed with dilute hydrochloric acid, decomposes, not like manganese peroxide when so treated to form water and chlorine, but into barium chloride and hydrogen peroxide : — Ba0 2 + 2HC1 = BaCl 2 + H.,0 2 . Barium Hydrochloric Barium Hydrogen peroxide acid chloride peroxide The preparation of the pure substance is tedious and difficult. The barium may be precipitated out of the solution by dilute sul- phuric acid, and the hydrochloric acid by silver sulphate, and then the sulphuric acid exactly precipitated with baryta water, by which means a dilute aqueous solution of the substance is obtained. This, on concentrating by evaporation in vacuo, decomposes all the more readily into oxygen and water, the more concentrated it be- comes. The chemical behaviour of hydrogen peroxide is very remark- able. The readiness with which it decomposes into nascent oxygen and water makes it one of the most powerful oxidizing agents for many substances—^, for arsenic, sulphurous acid, and many Hydrogen Peroxide. 91 metallic sulphides. Black lead sulphide is easily converted by it into white lead sulphate. In old oil paintings, in which the light colours produced by white-lead have gradually darkened owing to the formation of lead sulphide, the original colours may be again produced by carefully washing with a dilute solution of hydrogen peroxide. This then converts the lead sulphide into lead sulphate, without damaging the picture. Hydrogen peroxide acts, however, upon many substances as a reducing agent and abstracts oxygen. If an aqueous solution is added to silver oxide (or manganese peroxide), the liquid froths up and evolves considerable quantities of oxygen. The atom of oxygen which the peroxide contains more than water unites with an atom of oxygen from the silver oxide to form a molecule, and the silver oxide is at the same time reduced to metallic silver : — H 2 2 + Ag,0 . = H 2 + o 2 + Ag,. Hydrogen peroxide Silver oxide Water Oxygen Silver In the same manner manganese peroxide is reduced to man- ganous oxide, which, when the liquid contains a free acid, unites with it, forming a salt. This remarkable reaction may be used to determine roughly whether a given solution contains much or little hydrogen peroxide. The solution is poured on a little powdered black oxide of manganese contained in a test-tube when a rapid evolution of oxygen at once begins, sufficient to show its presence by kindling a glowing chip of wood. To detect traces of hydrogen peroxide in an aqueous acid liquid, a little ether is added, and one drop of a solution of potassium bichromate ; after shaking, the ether, as it rises to the surface, is found to be tiriged with a beautiful blue colour, due to some com- pound of chromium probably containing more oxygen than chromic acid. OZONE. Composition : 3 . It was long known that on turning the glass plate of an elec- trical machine a peculiar odour may be noticed, very similar to that produced when phosphorus is allowed to oxidize slowly in moist air. Careful experiments have shown that the cause of this odour is a chemical compound produced by these and other pro- 92 Text-Book of Inorganic Chemistry. cesses from the oxygen of the air. To this body the name ozone has been given. Ozone is an allotropic modification of oxygen, and its relation to ordinary oxygen is somewhat the same as that of ordinary phosphorus to the red amorphous modification of this element. In the same manner as ordinary phosphorus has a much more powerful affinity for oxygen, sulphur, and other elements than the red phosphorus, so the affinity of ozone for most substances is very much greater than that of ordinary oxygen. And, again, as common phosphorus is converted into the red variety on heating, so also is ozone converted by heat into ordinary oxygen. Ozone has never yet been prepared chemically pure, and particu- larly never free from ordinary oxygen. In all attempts to convert oxygen into ozone, a considerable . quantity, always the greater part of the former gas, remains unchanged. Our ideas, therefore, of the properties of ozone refer only to a mixture of it with oxygen. It is a colourless gas, with ,a very intense odour, soluble in water, to which it imparts its smell ; it acts in an irritating manner when inspired, and attacks the bronchial tubes in the same manner as chlorine. Ozone is condensed to the liquid state more easily than oxygen, and then forms a blue liquid of varying tints according to the pressure. Ozone occurs in atmospheric air, although only in very minute quantities. It is produced from the oxygen of the air in places where water evaporates quickly, particularly in the neighbourhood of the sea and of the salt-works in Germany, where the brine evaporates spontaneously. Various organic compounds — e.g. oil of turpentine, oil of bitter almonds — possess the property, when exposed to the light, of converting the oxygen of the air into ozone, absorbing it mechanically, and then yielding it up to other bodies. The partial conversion of oxygen into ozone may be brought about by allowing electric sparks to pass through the gas, or by allowing phosphorus, half covered with water, to oxidize slowly to phosphorous acid in a large ' vessel. In the latter case the ozone produced soon combines chemically with the phosphorous acid. The best method to obtain ozone, in comparatively speaking large quantities is to subject a slow stream of oxygen or atmo- spheric air to the action of high electric potential without allowing sparks to pass through the gas. By this silent discharge of the Ozone. 93 electricity through the oxygen which then takes place several per cent, of the gas is converted into ozone— a very remarkable fact, Fig. 36. 94 Text-Book of Inorganic Chemistry. and one which suggests the question whether other gases might not also be changed in the same way. The figure on the preceding page (fig. 36) shows an apparatus which yields large quantities of ozone, and which may be easily made by a skilful glass-blower. Through the centre of the wooden lid of the cylinder A A passes a wide glass tube, open at the top and closed at the bottom. At the lower end of this tube, near the side, a narrow piece of glass tubing is melted, which is bent up- wards and again at right angles (a) above the lid. By this tube the oxygen enters, and by the tube b it leaves the apparatus when partially converted into ozone. In order to produce this conversion of oxygen into ozone, a second rather smaller tube is placed inside the wide one. This tube is also open at the top and closed at the bottom, and the two tubes are melted together air-tight at the line cc, in such a manner that the gas, entering at a, passes through the narrow annular space between the two tubes, and finds an exit at b. The inner tube and the wide cylinder are filled to about the same height with dilute sul- phuric acid (1 of acid to 10 of water) ; in each of them dips a piece of platinum foil to which is melted a piece of stout platinum wire. Both wires are now attached to a powerful induction coil, and at the same time dry oxygen is allowed to enter at a. The two platinum plates and the sulphuric acid with "which they are in contact become strongly electrified, and discharge themselves silently, without the production of sparks, through the space between the two tubes, by means of which the oxygen is partially converted into ozone. Ozone is one of the most powerful oxidizing agents. It so exceeds ordinary oxygen in this respect that it often oxidizes bodies at the ordinary temperature which only combine with oxygen when heated. The metal mercury, which may be kept for years in the air without any alteration, becomes at once covered with a film of the oxide when exposed to the action of ozone. Silver is also oxidized and converted into the peroxide by ozone. Sulphu- retted hydrogen and sulphurous acid are rapidly converted into sulphuric acid, and black lead sulphide, which remains unchanged for years in ordinary oxygen, is at once changed into white lead sulphate. In the same manner as chlorine, ozone easily expels iodine from its compounds. Iodine, when free but not when combined, possesses the property of colouring starch paste blue, even when Ozone. 95 only minute traces of it are present. If, therefore, dilute starch paste is mixed with a solution of potassium iodide and strips of white paper are moistened with the mixture, the paper will become blue when brought into contact with air containing ozone, and the reaction will be all the more rapid the greater the quantity of ozone present. This very delicate test enables us to find out whether air con- tains ozone. 1 The reaction is expressed by the equation : — 2KI + o 3 + H 2 = = 2KOH + 1, + o,- Potassium iodide Ozone Water Potassium hydrate Iodine Oxygen The powerful oxidizing action of ozone makes it an energetic bleacher, and if it should later happen that ozone Could be pre- pared in such quantities and as easily as chlorine, it would be very generally employed for this purpose. The success of the old process of meadow bleaching depends probably upon the action of ozone, this substance being produced in small quantities by the rapid evaporation of the water contained in the linen by the sun's rays. Even indigo, which is otherwise a very permanent colour, cannot withstand the action of ozone. It is soon decolorized and converted into a new chemical compound by the action of ozone upon it. The difference between the molecular composition of ozone and ordinary oxygen has been determined by experiment. It was found that when electric sparks passed through oxygen, or, better, when the gas was subjected to the silent discharge, its volume diminished. The ozone which was produced must therefore be denser than oxygen. Exact experiments yielded the result that by this process three volumes of oxygen were condensed to two volumes of ozone, so that ozone must be one and a half times as dense as oxygen. And it was further observed that when the mixture of ozone and oxygen was again converted into oxygen by heating it regained its original volume. From these experiments, and from the fact that the molecule of oxygen contains two atoms, it is concluded that the molecule of ozone contains three atoms. If this is correct, the powerful oxidizing properties of ozone are explained on the supposition that its molecule breaks up into a molecule and a free atom of oxygen— 1 This test does not prove conclusively that ozone is present, as other bodies, such as nitrous acid, hydrogen peroxide, also colour the paper blue.— Ed. g6 Text-Book of Inorganic Chemistry. 3 = 3 + O— the latter then exercising the same force of affinity as oxygen in the nascent state — e.g. when set free by the electric current. Ozone, like hydrogen peroxide, is thus a valuable and easily obtained source of nascent oxygen. It may possibly assume very considerable importance in medicine and the arts at no distant date. THE HALOGENS OR SALT-PRODUCERS. The four elements chlorine, bromine, iodine, and fluorine are grouped together under this name because they form salts by direct union with the metals (a\s, sea-salt, and yemaa, I produce). These salts are called haloid-salts to distinguish them from the oxy- and sulpho-salts. The haloid-salts consist of two elements, the halogen element and a metal — for example, sodium chloride, NaCl ; while the oxy- and sulpho-salts contain at least three elements, a metal and a simple or compound radical united to- gether by oxygen or sulphur — for example, potassium nitrate, N0 2 -OK. It was thought at one time that every acid and every salt must contain oxygen, and it was long imagined that this element was present in the haloid-salts and that chlorine, bromine, &c, were compounds of oxygen. As, however, all attempts to detect oxygen in the halogens, and as salts, of the halogens really con- taining oxygen were, discovered later — for example, potassium chlorate (C10 2 -OK) — the halogens are now considered as ele- ments,. as long as we are unable to decompose them. The chemical compounds of the halogens, particularly those of chlorine, bromine, and iodine, are so similar that they are often difficult to distinguish from one another. It is thus all the more remarkable that these three elements should possess such very different physical properties in the free state. Chlorine is a greenish-yellow gas, bromine a dark red-brown liquid, and iodine is a dark grey crystalline solid with metallic lustre. Chlorine. gj CHLORINE. Chemical Symbol : CI. — Atomic Weight: 35-5. This element does not occur free in nature, but only combined ; its chief compound is common salt (sodium chloride), in which it is united with the metal sodium. But although common salt is one of the most widely distributed chemical compounds, chlorine has only been known for about 100 years ; it was discovered in 1774 by Scheele, one of the discoverers of oxygen. Chlorine is a greenish-yellow gas with a powerful and peculiar odour, and owes its name to its colour (xXapos, greenish-yellow). It possesses a density of 2-45 compared with air as unity. Its molecular weight is, therefore, 2-45x28-88 = 71 (p. 51), and its molecule, like that of hydrogen, contains two atoms. At a tem- perature of — 40 , or under a pressure of about ten atmospheres at the ordinary temperature, it is condensed to a dark yellow liquid, with a specific gravity of 1 -3. It is dissolved by water in larger quantities than oxygen. One volume of water at the ordinary temperature absorbs about 2-5 volumes of the gas, but at 0° a much greater quantity. At the latter temperature it forms a crystalline compound with water, having the composition, Cl 2 + ioH 2 0. If chlorine is breathed, even in small quantities, it attacks the organs of respiration, causes inflammation of the bronchial tubes, and so produces a very troublesome cough. In larger quantities it may produce inflammation of the lungs, followed by death. The chemist must therefore exercise caution in experimenting with this substance. In the ordinary sense of the word, chlorine is not, like hydrogen, a combustible body — i.e. if strongly heated in the air it does not burn. But if an atmosphere of hydrogen be substituted for air, chlorine burns freely, and, conversely, a jet of hydrogen burns in chlorine with a pale, livid flame, producing hydrochloric acid. From this we conclude that the affinity of chlorine for hydrogen is much stronger than for oxygen. Chlorine cannot support the combustion of a piece of glowing charcoal or of a chip of wood, as carbon and chlorine never combine directly with one another. For the preparation of chlorine its hydrogen compound (hydro- H 98 Text-Book of Inorganic Chemistry. chloric acid) is almost universally employed. In order to set the chlorine free from this compound, it is only necessary to abstract and convert the hydrogen into some non-volatile chemical com- pound. It will at once be remarked that the strong affinity of hydrogen for oxygen might be employed, and the chlorine obtained according to the equation : — 2HCI + O = Cl 2 + H 2 0. Hydrochloric Water acid And, as a matter of fact, considerable quantities of chlorine may be obtained by passing a mixture of hydrochloric acid gas and air over red-hot pieces of brick (Deacon's process). But by this method the chlorine remains mixed with large quantities of nitrogen, de- rived from the atmospheric air, and the process, though suitable for technical purposes, is quite unfit for the preparation of the pure gas. It is better to employ compounds containing oxygen, and par- ticularly those only loosely combined with the element, such as the peroxides. Most of the metallic oxides are so changed by the action of -hydrochloric acid that the corresponding chloride and water are produced, thus : — MnO + 2HCI = MnCl 2 + H 2 0. Manganous Manganous oxide chloride Some peroxides behave in the same manner — e.g. manganese per- oxide : — Mn0 2 + 4HCI = MnCl 4 .+ 2H a O. Manganese Manganese peroxide tetrachloride But these metallic perchlorides are very unstable bodies, and are easily decomposed into the lower chloride and water ; — MnCl 4 = MnCl 2 + Cl 2 . Manganese Manganous tetrachloride chloride From these two equations it is seen how chlorine may be pre- pared from manganese peroxide and hydrochloric acid, and how only one half of the chlorine is set free. The preparation of chlorine, according to this method, is best carried out by placing a small quantity of the powdered black oxide of manganese in a large glass flask, and then adding commercial hydrochloric acid until the flask is about one-third full. It is not advisable to introduce more than this into the flask, since the mass froths up during the evolution of the chlorine. When the black Chlorine. 99 oxide has been thoroughly mixed with the hydrochloric acid, it is gently heated over a gas flame. The reaction begins at the ordinary temperature, and only requires a very gentle warmth to continue it. The atmospheric air is then gradually expelled by the heavier chlorine, and finally the gas itself issues from the tube, which passes through the cork in the neck of the flask. The chlorine so prepared is always mixed with vapours of hydrochloric acid. In order to separate these vapours, and, at the same time, to dry the gas or abstract the water-vapour which it contains, it is passed through two wash-bottles, of which the first is Fig- 37- half filled with water, and the second with concentrated sulphuric acid. The water absorbs all the hydrochloric acid, together with a little chlorine, and the sulphuric acid removes most of the water- vapour, which is carried over from the first wash-bottle. Chlorine, unlike oxygen and hydrogen, cannot be collected over mercury, as it combines chemically with this metal at the ordinary temperature. If it is to be collected over water, the water is first made luke-warm, as warm water absorbs much less of the gas than cold. If a vessel is required full of dry chlorine, the tube delivering the gas is led to the bottom of the vessel (fig. 37). The chlorine then gradually displaces the lighter atmospheric air, and the u 2 100 Text-Book of Inorganic Chemistry. yellow coloured gas may be readily perceived as it rises in the jar. In order to lessen the communication with the external air, the mouth of the jar is loosely closed with a plug of cotton- wool. The chlorine is led into the jar until the interior appears uniformly coloured, by which time most of the air has been expelled. All such operations with chlorine must be conducted in a good draught of air, on account of the poisonous nature of the gas. In this method of preparing chlorine, sulphuric acid and common salt may be used instead of the hydrochloric acid. An intimate mixture of one molecular weight of common salt with at least two molecular weights of manganese peroxide is prepared, and an excess of dilute sulphuric acid poured over the mixture. In this way all the chlorine is obtained from the common salt, and the end products of the reaction are acid sodium sulphate, man- ganous sulphate, water, and chlorine : — 2NaCl + Mn0 2 + , Qn (OH 3S0 2 J, OH = Sodium chloride Manganese peroxide Sulphuric acid S0- 2 O 3 Mn + ,, n fONa 2SC M0H + Cl 2 + 2 Manganous sulphate Acid sodium sulphate But the preparation of chlorine from commercial hydrochloric acid and black oxide of manganese is always to be preferred ; it is more rapid and more advantageous, and the bye-product, man- ganous chloride, may be easily purified. Chloride may also be obtained by the electrolysis of a saturated aqueous solution of common salt, or of hydrochloric acid, by em- ploying electrodes of gas-carbon, and placing a bell-jar filled with the same liquid over the positive electrode. The chemical affinity of chlorine for most of the metals and non-metals is very powerful, and in many cases it even exceeds that of oxygen. The elements, oxygen, fluorine, nitrogen, and carbon, do not, however, combine directly with it, even at the highest temperatures. But compounds of these elements (except fluorine) with chlorine may be prepared in an indirect manner. The affinity of chlorine for hydrogen is particularly powerful. Equal volumes of the two gases, mixed in darkness or candle-light, unite to form hydrochloric acid when exposed to direct sunlight, or when brought into contact with a burning body or an electric spark. The quantity of heat which is set free on the union of the Chlorine. 1 1 two gases and the consequent expansion is so great that even thick vessels are burst with a loud explosion. A mixture of one volume of hydrogen and one volume of chlorine when ignited pro- duces the same effects, and for the same reason, as a mixture of two volumes of hydrogen with one volume of oxygen. A jet of hydrogen, however, when ignited in the air and plunged into a jar of chlorine, continues to burn quietly with a pale flame, also producing hydrochloric acid. The affinity of chlorine for hydrogen is under some circum- stances stronger than that of oxygen for the same element. An aqueous solution of chlorine, exposed to the light, gradually loses its colour ; the chlorine expels the oxygen from the water and hydrochloric acid remains behind. Sulphur unites with chlorine at the ordinary temperature, and more quickly when heated, to form the red volatile liquid, sulphur chloride. Phosphorus, which at the ordinary temperature neither ignites in air nor in oxygen, catches fire at once in chlorine, forming one of the chlorides of phosphorus. Antimony, when finely powdered and shaken into a jar of chlorine, burns with light and heat to form antimonic chloride. Copper, in a finely divided state, as Dutch metal, behaves in the same manner, the compound produced being cupric chloride (CuCl 2 ), corresponding to that formed when copper burns in oxygen, cupric oxide (CuO). Sodium burns in a stream of chlorine, and becomes incrusted with sodium chloride, which is identical in every respect with the naturally occurring common salt. A lighted taper, the wax or fat of which consists essentially of carbon and hydrogen, when introduced into a jar of chlorine, con- tinues to burn, but only feebly and with a dull red light, large quantities of soot being at the same time liberated. The hydrogen of the wax unites with the chlorine, while the carbon is set free. The same phenomena is shown still more strikingly with volatile hydrocarbons rich in hydrogen — for example, oil of turpentine. Many organic substances are decomposed by chlorine at the ordinary temperature, and in diffused daylight, the decomposition being often accompanied by a change in colour. If chlorine is led into a blue solution of indigo, or into ordinary black ink, or if chlorine water is added to either of these liquids, they both be- come yellow owing to the formation of new chemical compounds. 102 Text-Book of Inorganic Chemistry. Fresh flowers, and calico coloured with organic colours when moistened with chlorine, lose their colours, or, in other words, they are bleached. We are acquainted with the new chemical compounds produced by the action of chlorine on indigo and on ink ; we know that indigo-blue is converted into yellow isatin chloride, which can be separated and obtained in a crystalline form. But in most other cases, particularly with regard to the unknown colouring matter of flowers, we are as little acquainted with the products of decomposi- tion, when bleached by chlorine, as we are with the chemical nature of the colours themselves. The action of chlorine in these processes often consists in the abstraction of one or more atoms of hydrogen, which unite with the chlorine to form hydrochloric acid, while a similar number of atoms of chlorine takes the place of the displaced hydrogen. This process of substitution, as it is called, often takes place in many colourless organic compounds ; the well-known colourless phenol, (carbolic acid), for example, is easily converted by chlorine into monochlorphenol, dichlorphenol, and even trichlorphenol. In other cases, chlorine acts as an oxidizing agent. We have already seen that chlorine can decompose water in sunlight, forming hydrochloric acid and oxygen ; and if at the same time substances capable of easy oxidation are present — and to these belong many organic colouring matters — they become oxidized by the oxygen set free from the water. In most cases, as in the conversion of indigo into isatin chloride, both processes go on together. The bleaching properties of chlorine make this substance d great practical value. Linen which was previously bleached by exposing it to sunlight on green meadows for weeks together, is bleached by chlorine in as few minutes as the old method required weeks. In the arts, free chlorine is, however, not employed, but a compound called bleaching powder or chloride of lijne, which easily yields up its chlorine, and which is much easier to manage and much less dangerous than the free gas. Books or engravings, which, in course of time, have become stained, may be rendered completely white by placing them for a short time in fresh chlorine water. The printing ink, which consists essentially of carbon (lamp-black), is unacted upon by the chlorine. Characters in ordinary writing ink, which is an organic compound, are at once completely bleached. Chlorine. 103 In all cases in which chlorine or bleaching powder has been used for bleaching purposes, care must be taken that the chlorine, which adheres mechanically, is afterwards completely removed. Linen, calico, or paper when bleached with chlorine, retains chlorine in the pores so persistently, that it cannot be got rid of by repeated washing with water. Alone this chlorine would have no bad effect, but under the action of light it produces with water hydro- chloric acid, which destroys the organic tissues, rendering the fabric so brittle, that it ultimately falls to pieces. The chlorine which is left in the pores of the fabrics, after they have been bleached and thoroughly washed, must be removed by chemical means. For this purpose substances are chosen which are easily oxidized and which themselves exert no injurious action, as calcium sulphite or sodium thiosulphate. The former substance is trans- formed by chlorine in the presence of water into calcium sulphate (gypsum), the latter into sulphur and sodium sulphate. A substance used for this purpose is called an antichlor. Linen bleached in this way is whiter and lasts longer than that bleached by the old method of exposing it to the light and air. The property which chlorine possesses of decomposing and chemically altering organic and particularly organized bodies gives it another practical application. Chlorine acts as a most powerful disinfectant ; it destroys the organic substances which collect in the air of closed spaces, particularly in the chambers of patients suffering from infectious diseases. Experience has shown that the air of sick rooms is purer and healthier when small quantities of chlorine are allowed to evaporate from time to time. This is best done by placing a little chloride of lime in one or more saucers, and moistening it with a little common vinegar. Small quantities of chlorine are then gradually set free. Chlori?ie Hydrate : Cl a + ioH 2 0. Water, which at the ordinary temperature only absorbs a little more than twice its volume of chlorine, takes up large quantities at o°, and forms a chemical compound, having the composition shown above. This compound is produced, as small crystalline plates of a yellow colour, by leading chlorine into ice-cold water, until finally the whole liquid becomes semi-solid. The crystals maybe filtered off at a temperature below 0°, and freed from adhering water by pressing between blotting paper. At a few degrees above o° the hydrate is again decomposed into its constituents. This 104 Text-Book of Inorganic Chemistry. property affords an easy method for the preparation cf liquid chlorine. The chlorine hydrate is pressed into the end of a thick bent glass tube, which stands in a freezing- mixture, and after a few pieces of calcium chloride have been placed above it, the open end c (fig. 38), is fused together by the blow- pipe. If now the limb c is placed in a freezing mixture and the other limb a gently warmed, the chlorine which is set free condenses in c by its own pressure to a yellow liquid, while the water is retained by the calcium chloride. HYDROCHLORIC ACID {Muriatic Add). Composition : HC1. This substance, the only compound of hydrogen and chlorine, occurs in small quantities in the gases emitted from volcanoes. It is a colourless gas which fumes strongly in the air, possesses a suffocating odour and taste, and a strong acid reaction. Its specific gravity compared with air as unity is 1 '26, and its mole- cular weight is therefore 1-26 x 28-88 = 36-39 (p. 51), corresponding to the formula given it above. Under a pressure of 25 atmo- spheres and at a temperature of — 4° it is condensed to a colourless liquid. Water dissolves large quantities of the gas forming ordi- nary (aqueous) hydrochloric acid. Hydrochloric acid is produced by the direct union of chlorine and hydrogen, with a large evolution of heat, when about equal volumes of the gases are mixed and exposed to direct sunlight, or when raised to the temperature of combustion by an electric spark or a flame. If a very thin glass flask (or better, a thin glass bulb) is filled in the dark with a mixture of equal volumes of hydrogen and chlorine, corked firmly, and then exposed to bright sunshine, the moment when the first ray of sunlight strikes the flask or bulb, the union of the two gases at once follows, and the flask is burst by the force of the loud explosion. The combination takes place more slowly in diffused daylight, and is only produced suddenly with an explosion when the two gases are mixed in exactly equal proportions in the dark. This mixture of exactly equal volumes may be obtained by the electrolysis of strong aqueous hydrochloric Hydrochloric Acid. 105 acid by gas-carbon electrodes, and collecting the mixed gases together after all the air has been expelled from the liquid and the porous carbon, and after the liquid over the positive pole has become perfectly saturated with chlorine. When one volume of chlorine unites with one volume of hydro- gen no condensation occurs, as when oxygen and hydrogen com- bine together, but two volumes of hydrochloric acid are produced. One volume of the compound gas consists, therefore, of half a volume of chlorine and half a volume of hydrogen. From this it follows that the molecule of hydrochloric acid consists of £ molecule = I atom of chlorine, and £ molecule = 1 atom of hydrogen ; from this also its density may be calculated, and will be found to agree with experiment, thus : — i vol. chlorine . weighs 2-450 1 „ hydrogen .... „ 0-069 2 „ hydrochloric acid . . „ 2-519 1 >; » „ . . „ 1-26 Again, since 2-519 parts by weight of hydrochloric acid contain 2-45 parts of chlorine and 0-069 P art °f hydrogen, the percentage composition of the gas, according to the usual simple process, is : — Hydrogen . . = 274 parts by weight. Chlorine ... . = 97-26 „ Hydrochloric acid . . . = 100-00 „ In other words, this small quantity of hydrogen suffices to com- pletely change the properties of more than 35 times as much chlorine on chemical union with it. That hydrochloric acid consists of hydrogen and chlorine in equal volumes may be proved by its electrolytic decomposition. A piece of apparatus by which this experiment may be performed is shown in fig. 39. Two glass tubes of equal calibre, which are closed at the top by stop-cocks, are placed in a vertical position, and communicate by a side-tube with a third vertical tube, some- what longer, and enlarged at the upper end into a bulb. In the lower ends of the two tubes are two cylinders of gas-carbon, which are fitted in water-tight by caoutchouc stoppers, and which are placed in connexion with the two poles of a galvanic batter)'. The tubes are filled with strong aqueous hydrochloric acid (or a mixture 106 Text-Book of Inorganic Chemistry. of hydrochloric acid and a strong solution of common salt) through the bulb-tube, the stop-cocks being open. On commencing the experiment, and for some time afterwards, the stop-cocks are left open, so that the liquid in the tube where the chlorine is liberated may become saturated ; the chlorine which is evolved being carried away by a downward draught as shown in the diagram. If finally Fig. 39- the two stop-cocks are closed at the same time, hydrogen collects in one tube and chlorine in the other in equal volumes, and the liquid remains during the experiment at equal heights in the two tubes. The liquid which is displaced by the gases rises in the third tube and collects in the bulb. We see, therefore, from this Hydrochloric Acid. 107 experiment that hydrochloric acid contains chlorine and hydrogen in equal volumes. For the ordinary preparation of hydrochloric acid gas, common salt is almost exclusively used. This salt (sodium chloride), when mixed with concentrated sulphuric acid, gives a large quantity of the gas at the ordinary temperature without heating, and is con- verted into acid sodium sulphate : — Nad + S0 2 |OH _ Ha + sO a {gg a Sodium Sulphuric Acid sodium chloride acid sulphate But if the ordinary crystallized salt, which offers a large surface to the acid, is employed, the mixture always froths over. It is therefore better to use salt which has previously been melted and afterwards broken into moderately large pieces. This evolves the gas more slowly, as its surface is much smaller. The action com- mences at the ordinary temperature, but must afterwards be sup- ported by a gentle heat. If rock-salt is used for the preparation of hydrochloric acid, the sulphuric acid must be mixed beforehand with an equal volume of water, and, to free the gas from aqueous vapour, it must be passed through wash bottles containing concentrated sulphuric acid. Aqueous Hydrochloric Acid. In the ordinary language of the laboratory, hydrochloric acid means the solution of the gas in water. This is produced by pass- ing a stream of the gas through vessels containing water, which must be kept cold as long as absorption is going on. One volume of water at o° can absorb about 500 volumes of hydrochloric acid gas, at 20 a less quantity, but still over 400 volumes. During this absorption a considerable quantity of heat is set free, partly in consequence of the chemical attraction between the gas and the water, and partly because of the change from the gaseous to the liquid state. It is therefore necessary to keep the water cool, especially if a cold saturated solution of the gas is required. This solution of the gas in water (ordinary hydrochloric acid) fumes in the air and possesses the odour of the gas. It reacts and tastes acid even when very largely diluted with water. A completely saturated solution of hydrochloric acid at o c has a specific gravity of 1-21. On heating such a solution considerable 108 Text-Book of Inorganic Cliemislry. quantities of the gas are evolved, and the specific gravity dimin- ishes. When the specific gravity has become 1-145 the liquid boils at 100° Even at this temperature more hydrochloric acid distils over than water ; but on continued boiling the relative quantity of the water passing over with the acid gradually increases, until the boiling point reaches 1 io°, when it remains constant. The liquid which distils over at this temperature has a constant composition and a constant specific gravity of it. This acid, which distils over at a constant temperature and which has a constant composition, does not fume in the air. It contains about 20 per cent, by weight of the gas. It is not, how- ever, a true chemical compound, as its composition is altered if the pressure at which it is made to boil is changed. The satu- rated solution at o° contains about 42 per cent, by weight of the gas. Impure aqueous hydrochloric acid is obtained in large quan- tities as a bye-product in alkali works, in that part of the process in which sodium chloride is converted into sodium sulphate by the action of sulphuric acid. The hydrochloric acid gas, which is given off in immense quantities, is led into large earthenware jars half filled with water, which communicate with one another, so that the gas unabsorbed in the first is absorbed in the second, and so on. The commercial hydrochloric acid so obtained fumes in the air and has generally a specific gravity of about IT 6, which corre- sponds to about 33 per cent, of hydrochloric acid by weight. It is generally coloured yellow owing to the presence of iron (ferric chloride) or organic substances, and contains besides small quanti- ties of sodium chloride and sulphate, and always arsenious chloride if the sulphuric acid employed for its production contained arsenic. The presence of this substance may be recognized by the yellow precipitate of arsenious sulphide, produced when sulphuretted hydrogen is added to the acid. The commercial acid, free from arsenic, may be employed for the preparation of chlorine, and in all cases when the quite pure acid is not required. It may also be employed for the preparation of the pure non-fuming acid of specific gravity it, by distilling in a glass retort until the boiling-point has become constant at 1 10° and then collecting the distillate. At least one-fourth of the whole must be left behind in the retort to retain all the impurities. Hydrochloric acid is one of the strong inorganic acids, and 2HC1 = CuCl 2 + H 2 0. Hydrochloric acid Cupric chloride Water Oxides and Oxy -acids of Chlorine. 109 expels many other acids from their salts. With the metallic oxides it forms metallic chlorides and water — e.g. CuO Cupric oxide In the same manner it decomposes many metallic sulphides — e.g. ferrous sulphide, antimonious sulphide — liberating sulphuretted hydrogen. It also serves as a solvent for metals. Tin, which is oxidized but not dissolved in nitric acid, easily dissolves in hydro- chloric acid, forming stannous chloride and hydrogen : — Sn + 2HCI = SnCl, + H. Hydrochloric Stannous acid chloride With most of the peroxides it produces chlorine — e.g. manganese peroxide, lead peroxide. The presence of hydrochloric acid or of a soluble chloride in a liquid may be recognized by the white precipitate of silver chloride produced when silver nitrate is added to the solution. The pre- cipitate is easily dissolved by ammonia and turns black in the light. All chlorides yield chlorine when heated with manganese peroxide and sulphuric acid, and the free chlorine may be easily recognized by its odour and by its bleaching properties. OXIDES AND OXY-ACIDS OF CHLORINE. The chemical affinity of chlorine for oxygen is so slight that the two elements cannot be made to unite directly with one another. But compounds of chlorine and oxygen are produced when the one element is brought into contact with the other in the nascent state, or by other indirect means. We have seen that chlorine and hydrogen unite together only in one definite proportion, but chlorine and oxygen form com- pounds in five different proportions, of which some are only known as oxygen compounds, others as oxy-acids, or in combination with bases as salts. With the exception of perchloric acid, these com- pounds are very unstable, and easily decompose into their con- stituent elements, often with an explosion. no Text-Book of Inorganic Chemistry. We distinguish the following compounds of chlorine and oxy- gen, in which the chlorine atom may have a valency from one to seven : — Oxides. Oxy-acids. Hypochlorous anhydride . CljO Hypochlorous acid . CIO • H Chlorous anhydride . . C1 2 3 Chlorous acid . . CIO 'OH Chlorine peroxide . . C10 2 Chloric acid . . C10 2 -OH Perchloric acid .CIO, -OH CHLORIC ACID. Composition: C10 2 -OH. We may well commence our description of the compounds of chlorine and oxygen with the most important of them — chloric acid. The anhydride corresponding to this acid is unknown, and the acid itself is only known in an aqueous solution. In order to prepare it a solution of potassium chlorate is mixed with fluosilicic acid. The insoluble potassium fluosilicate is filtered off and the acid liquid neutralized with baryta water ; this solution, when filtered and evaporated, deposits crystals of barium chlorate. An aqueous solution of this salt is made and dilute sulphuric acid carefully added, towards the last drop by drop, as long as a pre- cipitate of barium sulphate is produced. The acid liquid, contain- ing the chloric acid, is then filtered off and evaporated in a vacuum over sulphuric acid ; it must not be heated, as it begins to decom- pose at temperatures over 40 . So prepared, it forms a thick acid liquid which has not yet been obtained in a crystalline form. The salts of the acid — the chlorates — are, however, more stable, and of these the potassium salt is the best known and most impor- tant. They are obtained when chlorine and nascent oxygen are brought into contact with an alkali, under certain conditions. Potassium chlorate — C10 2 .OK — is produced in considerable quantities when chlorine is led into a hot concentrated solution of caustic potash. 1 The substances produced are potassium chlorate, potassium chloride, and water, of which the potassium chlorate 1 The delivery tube must be wide or it will become choked by the potassium chlorate. Chloric Acid. n i crystallizes out first on evaporation owing to its being less soluble in water than potassium chloride. The following equation shows the reaction : — 6KOH + 3 C1 2 = C1CVOK + 5KCI + 3 H 2 0. Potassium Potassium Potassium hydrate chlorate chloride It is probable that potassium hypochlorite is first produced with potassium chloride, and that the former then breaks up into potas- sium chlorate and potassium chloride : — H,0. 2KOH + Cl 2 = ClOK + KCI + Potassium Potassium Potassium hydrate hypochlorite chloride 3C10K = = C10 2 -OK + 2KC1. Potassium Potassium Potassium hypochlorite chlorate chloride Potassium chlorate so prepared always contains more or less potassium chloride mixed with it, from which it may be easily and completely separated by repeated crystallization. The presence of potassium chloride mixed with the chlorate may be easily recog- nized by the white precipitate of silver chloride produced when a few drops of silver nitrate are added to a solution of the salt, silver chlorate being soluble in water. Potassium chlorate crystallizes in thin iridescent plates, without water of crystallization. The salt may also be prepared by the direct oxidation of potassium chloride. If a concentrated solution of potassium chloride is electrolyzed by a powerful current, using platinum electrodes, the nascent oxygen which is set free at the positive pole oxidizes the potassium chloride to potassium chlorate. We have already stated that potassium chlorate is largely employed for the preparation of oxygen (p. 12). The chlorine and oxygen are so feebly held together that the latter gas is com- pletely given off on heating, and potassium chloride remains behind : — 2CICVOK = 2KCI + 30 2 . Potassium Potassium chlorate chloride In the first stage potassium perchlorate is produced, and in such quantities that the method is used for the preparation of this salt :— 2CICVOK = C10 3 -OK + KCI + 0. 2 . Potassium Potassium Potassium chlorate perchlorate chloride H2 Text-Book of Inorganic Chemistry. On heating more strongly this salt is also decomposed into potassium chloride and oxygen. In consequence of the slight affinity of chlorine for oxygen the chloric acid of potassium chlorate and other chlorates acts as a powerful oxidizer, something in the same manner as the nitric acid of the nitrates. A red-hot piece of charcoal, on which a little powdered potassium chlorate is sprinkled, burns brilliantly by combining with the oxygen of the salt. A mixture of dried flowers of sulphur and powdered potassium chlorate not only explodes when heated, but also by percussion or even by the heat produced when it is rubbed in a mortar. Caution must, therefore, be exercised in experimenting with this mixture, especially when charcoal powder has been added. A gunpowder containing potassium chlorate instead of nitre — that is, a mixture of charcoal, potassium chlorate, and sulphur — is so explosive that fire-arms charged with it are burst, and potassium chlorate is, therefore, not employed for the manufacture of such explosive mixtures. Chloric and hydrochloric acids mutually decompose one an- other, even in dilute aqueous solutions. A portion of the oxygen of the chloric acid oxidizes the hydrogen of the hydrochloric acid, and chlorine together with chloric peroxide are given off. 1 The process goes on when aqueous hydrochloric acid is added to a solution of potassium chlorate : — 2C10 2 -OK + 4HCI = 2CIO., + Cl a + 2KCI + 2 H 2 0. Potassium Hydrochloric Chloric Potassium chlorate acid peroxide chloride Concentrated sulphuric acid acts very energetically on potas- sium chlorate. The chloric acid which is first produced at once breaks up into oxygen, and the reddish-brown gas chlorine peroxide, which on heating decomposes further into chlorine and oxygen with an explosion. Chloric acid is a monobasic acid, and all its salts are soluble in water. They may be easily recognized by their oxidizing properties and by the production of chloric peroxide when heated with concentrated sulphuric acid. 1 This mixture of gases was called euchlorine by Davy. — Ed. Perchloric Acid. 113 PERCHLORIC ACID. Composition : C 1 3 • H . This compound is a colourless oily liquid of specific gravity 178, which does not become solid at -34°, fumes in the air, and becomes heated when mixed with water. It is prepared by dis- tilling potassium perchlorate with a large excess of sulphuric acid. Potassium perchlorate, as we have just mentioned (p. in), is obtained on heating potassium chlorate. The heating is stopped after the molten salt has ceased frothing and has become viscid or nearly semi-solid. At this point most of the potassium chlorate has been converted into potassium chloride and perchlorate, and as the latter salt is even more difficultly soluble in water than potassium chlorate, it may be easily purified by dissolving in hot water and allowing the solution to crystallize. The salt requires seventy times its weight of water at the ordinary tempe- rature to dissolve it, and separates out from its solution in rhombic prisms. On heating pure potassium perchlorate in a retort with about four times its weight of concentrated sulphuric acid, perchloric acid distils over as an oily liquid, usually with a yellow colour. When mixed with a small quantity of water it becomes heated and' gives a hydrate of the composition C10 3 • OH + H 2 0, crystallizing in fine needles. If this hydrate is heated to 110° it decomposes into perchloric acid and a hydrate with the composition C10 3 • OH + 2H..O, which remains behind in the retort as an oily liquid, boiling only at 203° From these facts it will be seen that, although perchloric acid contains more oxygen than chloric acid, it is a much more stable compound. This is also shown in its behaviour when treated with hydrochloric acid. A mixture of potassium perchlorate (or per- chloric acid) and hydrochloric acid remains unchanged, while one of potassium chlorate and the same acid evolves chlorine and chloric peroxide. In this way it may be readily seen whether crystalline potassium perchlorate is contaminated with potassium chlorate. Even when only very small quantities of the latter salt are present, the hydrochloric acid becomes yellow and gives off the odour of chlorine. I 114 Text-Book of Inorganic Chemistry. Potassium perchlorate when heated breaks up into potassium chloride and oxygen, but requires a higher temperature than that necessary to decompose potassium chlorate into potassium per- chlorate and oxygen. It deflagrates when thrown on red-hot charcoal in the same manner and for the same reason as potassium chlorate. The free acid attacks many organic bodies most energetically ; it oxidizes them, often with an explosion. Perchloric acid is a monobasic acid, and all its salts are soluble in water. CHLORINE PEROXIDE. Composition : CIO,. This body is only known as an oxide and does not combine with water or with bases. Chlorine peroxide is a most dangerous substance to experiment with, as it easily decomposes with a powerful explosion into its constituents. It is produced, together with perchloric acid, when small pieces of fused potassium chlorate a_e treated with pure con- centrated sulphuric acid and the mixture heated not higher than 40 on a water-bath. It is then evolved as a dark yellow or reddish- brown gas, with a powerful odour, and which often explodes even at this temperature. In a freezing mixture the gas condenses to a red liquid, which boils at about 9 . CHLOROUS ANHYDRIDE. Composition : C1.,0 3 . CHLOROUS ACID. Composition : CIO -OH. Chlorous anhydride is a yellow gas, with an odour similar to that of chlorine. It has a specific gravity of 4-07. Like chlorine, it attacks the organs of respiration, and bleaches even Chlorous Acid. 115 more powerfully than the element itself. Water dissolves about ten times its own volume of the gas, becoming of a yellow colour. This solution contains chlorous acid — CIO -OH, which has, how- ever, never been prepared pure. It bleaches like the gas and pro- duces yellow spots on the skin. Chlorous anhydride may be easily prepared by abstracting oxygen from a dilute aqueous solution of chloric acid. Nitric acid may be best used to set the chloric acid free from potassium chlo- rate and arsenious anhydride for the reduction. The following equations show the reactions : — (i.) CICvOK + N0 2 -OH = CIO,- OH + N0 2 -OK. Potassium Nitric acid Chloric acid Potassium chlorate nitrate (ii.) 2C10 2 ■ OH + As a 3 = Cl,O s + H 2 + As,0 5 . Chloric acid Arsenious Chlorous Arsenic anhydride anhydride anhydride For the preparation of the gas 3 parts of arsenious anhydride and 12 parts of potassium chlorate, both in the state of a fine powder, are mixed, and then 18 parts of pure nitric acid of specific gravity 1-33 and 24 parts of water are added. On gently warming the liquid it becomes coloured yellow, and chlorous anhydride is evolved in considerable quantities. The gas may be dried by pass- ing it over calcium chloride, and may then be collected by dis- placement. At — 1 8° it is condensed to a reddish-brown mobile liquid, which boils a few degrees above 0°. Chlorous anhydride is a very unstable compound. At a little above 50° it decomposes with explosion into its constituents. Great care must therefore be taken in its preparation, and only small quantities should be prepared at once. Its aqueous solution combines slowly with bases to form salts — the chlorites — which are mostly soluble in water, and which are easily decomposed on evaporation into a mixture of chlorate and chloride. The lead salt, obtained by mixing a solution of potassium chlorite with lead acetate, is deposited as yellow scales. The general formula of the chlorites is CIO • OM', and the acid is there- fore monobasic. n6 Text-Book of Inorganic Chemistry. HYPOCHLOROUS ANHYDRIDE. Composition : C1 2 0. This compound is a yellow gas with an odour similar to that of chlorine. It has a specific gravity of 2-97, and condenses at — 2o°to a bright red liquid, boiling at about + 20°. When heated it is easily decomposed into its constituents, usually with a violent explosion ; the liquid decomposes even when shaken. In its preparation great care must therefore be exercised. The gas is prepared by treating mercuric oxide with chlorine, when the reaction goes on as shown in the equation : — 2HgO + 2C1 2 = HgO,HgCl 2 + CLO. Mercuric Mercuric Hypochlorous oxide oxychloride anhydride The ordinary red oxide of mercury, which has been obtained by heating the metal in the air, cannot be employed for the reaction, but the yellow oxide, which has been prepared by precipitating mercuric chloride with caustic soda, and has been carefully washed and dried for some time at 300 , must be used. If dry and pure chlorine is slowly led over such mercuric oxide contained in a long well-cooled tube, it is absorbed and the hypochlorous anhydride gas set free. The gas is easily soluble in water, one volume taking up about 200 volumes of the gas. This solution may be considered to contain HYPOCHLOROUS ACID— (ClOH), which has, however, never been prepared in the pure state. A similar aqueous solution of hypochlorous acid may be also easily prepared by shaking up precipitated mercuric oxide with chlorine water. The yellow colour of the solution quickly disappears, and the mercuric oxychloride which is formed at the same time may be separated by filtration. The salts of this monobasic acid — the hypochlorites — are also unknown in the pure state, but mixed with other bodies they form important substances in the arts and manufactures. Bleaching powder, which is manufactured in large quantities and chiefly used for bleaching purposes, yields a mixed solution of calcium chloride and calcium hypochlorite when dissolved in water. This bleaching powder, or chloride of lime (not calcium chloride) Hypochlorous Acid. 117 as it is sometimes called, may be considered as a peculiar com- pound of calcium, oxygen, and chlorine, having the composition ( CI (CaOCl)Cl or Caj„„,. It is prepared bypassing dry chlorine over dry slaked lime (calcium hydrate). When no more chlorine is absorbed, a white powder, which smells faintly like hypochlorous acid, remains behind, having the above composition. This com- pound, when treated with water, is decomposed into calcium hypo- chlorite and calcium chloride, according to the equation : — 2(CaOCl)Cl = (C10),Ca + CaCl,. Bleaching Calcium Calcium powder hypochlorite chloride But from this mixture pure calcium hypochlorite cannot be sepa- rated. The weakest acids, even carbonic acid, separate the hypo- chlorous acid from this and all other hypochlorites, and the free acid then readily breaks up into chlorine and oxygen. On this depends the bleaching and oxidizing action of bleaching powder. The articles to be bleached are first dipped in a dilute solution of the powder in water, and then submitted to the action of a dilute acid. The effect produced is the same as if the articles were dipped in chlorine water. The reaction takes place according to the two equations : — (C10),Ca + 2HCI = CaCl, + 2CIOH. Calcium Calcium Hypochlorous hypochlorite chloride acid 2C10H + 2HC1 = 2H 2 + 2C1 2 . Hypochlorous acid For a similar reason bleaching powder is well adapted to disinfect sick chambers. If a little of the powder placed in a saucer is moistened with vinegar (dilute acetic acid), it slowly gives oft chlorine (p. 103). If bleaching powder is digested with water at the ordinary temperature and filtered, the alkaline liquid contains, as we have seen, calcium hypochlorite together with calcium chloride and un- changed hydrate. On boiling this liquid the calcium hypochlorite is decomposed into chlorate and chloride : — 3(C10)„Ca = C10 2 -o! Ca + 2CaC1 2- Calcium Calcium Calcium hypochlorite chlorate chloride Ifchlorinegas is led into a dilute solution of potassium or sodium 1 1 8 Text-Book of Inorganic Chemistry. hydrate in the cold, a mixture of the corresponding hypochlorite and chloride is obtained : — 2KOH + Cl 2 = KOC1 + KC1 + H 2 0. Instead of the caustic alkalies, their carbonates may also be used. No evolution of carbonic acid is noticed at first, owing to the for- mation of an acid carbonate. As soon, however, as the chlorine begins to act upon this acid carbonate, carbonic acid is given off. These liquids, which are sometimes called Eaux de Javelle, also bleach powerfully on the addition of a dilute acid. Hypochlorous acid is a powerful oxidizing agent, A solution of bleaching powder or of sodium hypochlorite when added to a solution of manganous sulphate produces first a white precipitate of manganous hypochlorite and hydrate, which rapidly becomes oxidized to the brown manganese peroxide. Similarly, with a solu- tion of lead nitrate lead peroxide is formed. BROMINE. Chemical Symbol : Br. — Atomic Weight: 80. This element, which was discovered by Balard in 1826, occurs in nature combined with sodium and other metals and usually asso- ciated with sodium chloride. Sea-water and many mineral springs contain small quantities of these compounds of bromine. In the mineral kingdom it occurs, particularly in South America, as silver bromide, and has lately been found combined with potassium in the enormous salt deposits of Stassfurt and other places in Northern Germany. The bromine which is now brought into trade at a low price is chiefly prepared from this last-named source. Bromine is a dark red-brown liquid, which is nearly opaque even in thin layers, and which possesses a specific gravity of 3- 1 8. It freezes to a crystalline solid, with a metallic lustre, at — 7"'3, and boils at 63°, forming a dark-brown gas. At the ordinary temperature it vaporizes rapidly, so that if a few drops are placed in a large empty flask, the whole soon becomes filled with the brown vapour. Its vapour density is 5-5, according to which it is more than twice as heavy as chlorine and 80 times as heavy as hydrogen. Bromine. 119 Its molecular weight is therefore 5-5 x 28-88 = 159, and its molecule, like that of hydrogen, contains two atoms. Bromine, like chlorine, dissolves in water, though in rather less quantity ; the solution — bromine water — has a brownish colour. In the cold below 4 bromine combines chemically with water to form Iromine hydrate : Br, + ioH.,0, which is decomposed at 15° or is rather more stable than the corresponding chlorine compound. Bromine has an odour resembling that of chlorine, but more intense (whence its name from fipa>iioc = a. stink). It attacks the eyes and the mucous membrane of the respiratory organs even more powerfully than chlorine, and great caution must therefore be exercised in performing experiments with it. A single drop allowed to fall on the skin produces a painful, slowly-healing wound. Bromine may be prepared in the same manner as chlorine by decomposing hydrobromic acid with manganese peroxide. But as hydrobromic acid cannot be so readily obtained as the commercial hydrochloric acid, and is much more difficult to prepare, a mixture of potassium or sodium bromide with manganese peroxide and moderately dilute sulphuric acid answers the purpose better. The hydrobromic acid which is at first produced is at once oxidized to bromine and water by the manganese peroxide. The distillation is best carried out in a tubulated retort fitting air-tight in a receiver, the uncondensed bromine vapour being led away by a good draught. The reaction is precisely the same as in the preparation of chlorine, substituting bromine for chlorine : — 2KBr + Mn0 2 + 3 so >{oh" 2S0 4oK + sc V°* Mn + 2H *° + Br 2- The bromine may be purified by shaking with water, and may be again freed from this substance by calcium chloride or concentrated sulphuric acid. The extraction of bromine on the large scale is carried out in just the same manner. For this purpose the mother-liquors of the mineral- or sea-waters are employed from which most of the less soluble chlorides have crystallized out. But as these mother- liquors always contain some chlorides, and as the bromides are more easily decomposed by sulphuric acid than the chlorides, only enough sulphuric acid and manganese peroxide is added to the mixture to decompose the bromides. In this way only bromine is obtained. The liberation of chlorine is to be avoided, as it com- 1 20' Text-Book of Inorganic Chemistry. bines with the bromine to form a chloride, which is difficult to separate from the bromine. Small quantities of chlorine may be removed from bromine by distilling with potassium bromide, thus : — BrCl 3 + 3KBr = 3KCI + 2Br 2 . Bromine Potassium Potassium chloride bromide chloride In its chemical behaviour, bromine is very closely related to chlorine. The differences in their chemical action may almost all be ascribed to the weaker affinity of bromine in comparison with that of chlorine. This difference may be well illustrated in their behaviour with regard to hydrogen. Although the affinity of bromine for hydrogen is very great, and although a mixture of bromine and hydrogen may be made to combine by heating, the union of the two substances cannot be produced by the rays of the sun, not even the direct rays. The weaker affinity of bromine is also shown by the fact that chlorine can expel bromine from most of its compounds. The non-metals and metals which burn in chlorine combine under similar circumstances with bromine also. Its affinity for oxygen and carbon is as weak as that of chlorine for these elements. Of the bromides, or the compounds so produced, most are soluble in water ; lead bromide is difficultly soluble, silver bromide quite insoluble. HYDROBROMIC ACID. Composition : HBr. This compound, like hydrochloric acid, is a colourless gas fuming in the air, with a strong acid taste and odour, and very soluble in water. Its specific gravity is 271, corresponding to the molecular weight: 81. The aqueous solution agrees also in its leading properties with aqueous hydrochloric acid ; when com- pletely saturated it fumes in the air, and gives off a portion of its hydrobromic acid gas on heating. As previously mentioned, although sunlight cannot cause the chemical union of bromine gas and hydrogen, they may be made to unite by heating. But it is impossible to prepare pure hydrobromic acid by heating a mixture of potassium or sodium bromide with Hydrobromic Acid. 121 concentrated sulphuric acid, because a portion of the hydrobromic acid which is set free reduces some of the sulphuric acid, forming bromine, water, and sulphurous anhydride, thus : — 2HBr - Hydrobromic acid so J° H - [OH Sulphuric acid Br 2 + 2H a O + SO,. Sulphurous anhydride The hydrobromic acid prepared in this way is, therefore, always coloured red with bromine gas, and is contaminated with sulphurous anhydride. The pure, colourless gas may be easily prepared by decompos- ing phosphorous bromide with water, phosphorous acid being at the same time produced : — 3 H 2 = PHO-[°£ + 3 HBr. PBr 3 Phosphorous bromide PHO (OH Phosphorous acid Hydrobromic acid For this purpose, I part by weight of amorphous phosphorus is mixed with 2 parts of water, and then 10 parts of bromine gradu- Fig. 40. ally added, drop by drop, from a stoppered funnel fitted into the cork of the flask containing the phosphorus and water (fig. 40). At first every drop of bromine gives rise to a powerful action, accom- panied with a flash of light, but afterwards the union of the phos- phorus and bromine goes on more quietly. The hydrobromic acid, of which the liberation may be completed by gentle heating, is freed 1 22 Text-Book of Inorganic Chemistry. from small quantities of bromine by passing it through a (J-tube containing fragments of glass and ordinary phosphorus. Hydrobromic acid is easily and completely decomposed by chlorine. If a cylinder of the dry gas is placed mouth to mouth with one containing dry chlorine, hydrochloric acid and bromine are produced, the latter being recognized by its brown colour. If the chlorine is in excess the brown colour disappears again, and bromine chloride is produced as a reddish-yellow mobile liquid. In the same manner as hydrobromic acid, all bromides are also decomposed by chlorine. If chlorine water is added to a solution of a bromide, bromine will be set free, and may be 1 easily recognized by its yellow or brown colour. If then a drop of carbon disulphide is added and the mixture shaken, this substance will dissolve the bromine and its colour will become more marked. Hydrobromic acid and all soluble bromides give a pale yellow precipitate of silver bromide when mixed with silver nitrate. The precipitate only dissolves in ammonia with difficulty. OXY-ACIDS OF BROMINE. In the same manner as chlorine, bromine cannot combine directly with oxygen. The combination may be brought about in an indirect manner, but no compounds of bromine and oxygen alone are known. Bromic Acid : Br0 2 -OH. — The potassium salt of this acid is produced, together with potassium bromide, when bromine is dissolved in concentrated caustic potash : — 6KOH + 6Br = Br0 2 - OK + sKBr + 3 H 2 0. Caustic Potassium Potassium potash bromate bromide And as the potassium bromate, like potassium chlorate, is much less soluble in water than the bromide, it may be easily obtained pure by repeated crystallization. When heated, the salt breaks up into potassium bromide and oxygen. Bromic acid — which may be separated from potassium bromate according to the method given on p. Ho for chloric acid — is a colourless, powerfully acid liquid, which first reddens and then Iodine. 123 bleaches litmus paper. At ioo° it decomposes into bromine and oxygen. Hypobromous Acid: BrOH, may be obtained by shaking bromine water with precipitated mercuric oxide. The aqueous solution is a pale yellow liquid, which acts as a powerful oxidizer, and which bleaches like hypochlorous acid. If bromine is allowed to evaporate at the ordinary temperature under a bell-jar with slaked lime, bromide of lime is obtained, a substance which resembles chloride of lime, and which also possesses bleaching properties. IODINE. Chemical Symbol : I. — Atomic Weight: 127. Up to the commencement of the present century soda was exclusively prepared by burning sea-plants, extracting the burnt mass (called kelp, or varec) with water, and crystallizing the soda from the solution so obtained. The French chemist Courtois, when attempting to obtain chlorine from the residues of the above pro- cess by heating them with manganese peroxide and sulphuric acid, noticed a splendid violet vapour, the investigation of which led him to the discovery of iodine (181 1). The chemical nature of this element was, however, first made clear by the investigations of Gay-Lussac four years later. It owes its name to the colour of its gas (icoS^s- = violet-coloured). Like chlorine and bromine, iodine does not occur free in nature, but always in combination with some of the metals. Sodium iodide is almost always present in common salt, although in extremely small quantities. Even those mineral springs which contain sodium chloride, and which are, comparatively speaking, rich in sodium iodide, contain such small quantities of iodine that its presence can scarcely be detected without concentration. Sea-water contains even smaller quantities of sodium iodide, and yet this is the source from which by far the greatest proportion of the 100 tons of iodine which is consumed annually is obtained. But if we were obliged to procure iodine directly from sea-water, 124 Text-Book of Inorganic Chemistry. the process would certainly be an unprofitable one ; we therefore extract it in an indirect manner with the aid of organic nature. Sea-plants, and particularly sea-weeds, extract the iodine from the sea-water and concentrate it in their structures, and it is from these plants alone that it is profitable to extract iodine by chemical processes. The manufacture is carried on principally on the west coasts of Britain '(especially Scotland) and France. The sea- weeds are gathered during low water, dried and burnt, and the ashes thus obtained (called kelp) are rich in sodium iodide. Not only the Algae, but also other sea plants, take up iodine from the sea, and the same is true of many marine animals — the sponges, &c. In this way iodine is conveyed into the bodies, and especially into the fat, of many other marine animals, and occurs, for example, in ordinary cod-liver oil. In the mineral kingdom iodine is found as silver iodide in Peru, Mexico, &c, and as sodium iodide in those salt deposits which are due to the evaporation of sea-water. So, for example, in rock salt, and particularly in Chili saltpetre, or sodium nitrate occurring in Chili. From the mother- liquors of the crude sodium nitrate very considerable quantities of iodine are now manufactured. I odine is a solid crystalline substance, with a dark grey colour and metallic lustre. It neither conducts electricity nor heat ; it is brittle, and may be easily powdered, and possesses a peculiar, un- pleasant odour, similar to that of chlorine and bromine, but less intense. It melts at 1 13 , forming a dark brown liquid, boils at about 200°, and gives off a dark violet-coloured vapour. On cooling, the walls of the vessel become covered with innumerable small, lustrous crystals of the solid. Its specific gravity is 4'95, and that of its vapour 87 compared with air as unity. Its molecular weight is thus 87x28-88 = 251 ; whence it follows that its mole- cule, like that of hydrogen, consists of two atoms. Although iodine only boils at about 200 , it is so volatile at the ordinary temperature that when left exposed to the air it soon loses considerably in weight ; and when kept in closed vessels it sub- limes, like camphor, on to the cooler portions of the bottle. It is slightly soluble in water, much less than chlorine and bromine, and forms a yellow-brown solution. It possesses a strong unpleasant taste, colours the skin yellow, and is poisonous. Water containing salts dissolved in it, particularly potassium iodide, dissolves larger quantities of iodine than pure water ; it is Iodine. 1 2 5 also dissolved in considerable quantities by aqueous hydriodic acid. Alcohol and ether, as well as chloroform and carbon di- sulphide, are good solvents of iodine. Its alcoholic solution pos- sesses a dark brown colour and is called tincture of iodine. The solution in ether is also of a brown colour, while that in chloroform or carbon disulphide is dark violet, or when largely diluted a bright pink. Many applications are made of iodine in medicine and in the arts — for example, in photography. Tincture of iodine is applied externally to reduce goitre, enlarged joints, and other similar swel- lings, and potassium iodide is taken internally to produce the same effect. Long before this action of the salt was known, even before iodine itself was discovered, the ashes obtained by burning sponges were used as a medicine for these diseases. It was later found that these ashes contain sodium iodide, and since then potassium iodide or sodium iodide has always been used in these cases. The preparation of iodine from the aqueous extraction of the ashes of sea-plants after most of the soda and common salt have crystallized out, or from the mother-liquors of Chili saltpetre, may be carried out by passing chlorine through the liquid, which separates the iodine. Care must be taken that too much of the gas is not passed, otherwise iodine chloride will be produced, which diminishes the yield. Another method is to heat the crude sub- stances in iron or earthenware vessels with only just enough sul- phuric acid and manganese peroxide to decompose the sodium iodide present. This salt is more easily decomposed than the chloride, and consequently the iodine separates first before the chlorine. The process is similar to that used for the preparation of chlorine (p. ioo), substituting iodine for chlorine :— 2 NaI + Mn0 2 + 3S0 2 |q^ = 2SO *{oNa + SC V°3 Mn + 2H =° + l r The iodine vapour is received in a number of communicating vessels to condense it, then freed from water and purified by re- sublimation. Usually the two methods are employed together. The iodine is first separated by chlorine, excess being avoided, again converted into sodium iodide by caustic soda and then dis- tilled with manganese peroxide and sulphuric acid. Commercial iodine not unfrequently contains iodine chloride and cyanide ; it 126 Text- Book of Inorganic Chemistry. may be purified, in the same manner as bromine (p. 120), by distil- ling with potassium iodide : — ICL, + 3KI = 3KCI + 2l_. Iodine Potassium Potassium chloride iodide chloride In its chemical properties iodine exhibits close analogy with chlorine and bromine. It has, however, less affinity for the metals and for hydrogen than these two elements, which there- fore expel iodine from its soluble salts and from its compound with hydrogen — hydriodic acid. On the other hand, iodine has a stronger affinity for oxygen than the other two halogens ; iodic acid, for example, is a solid substance easily crystallized and much more stable than the corresponding chloric and bromic acids. The behaviour of iodine to starch is very remarkable. Thin starch paste' mixed with an aqueous solution of iodine becomes of a dark blue colour, which, however, vanishes on warming. It appears as if a real chemical compound were produced by this reaction, but no attempt to determine its composition has been yet successful. The colour is only produced by free iodine, and not by solutions of its salts, and starch paste can therefore be mixed with potassium iodide in any proportion without producing any change in colour. But if to the mixture a drop of chlorine water is added, which sets iodine free, a dark blue coloration is at once noticed. This property is employed both to detect traces of starch and of free iodine. If a piece of paper moistened with starch is hung up in a closed cylinder containing a few drops of a liquid with free iodine at the bottom, the paper soon becomes blue owing to the vapours of iodine rising from the liquid. More delicate even than this reaction, and therefore applicable to the detection of infinitesimal quantities of iodine, is the pink or purple colour which traces of iodine impart to chloroform or to carbon disulphide. Chlorine water is not well adapted for the separation of minute quantities of iodine from a liquid, as the slight- est excess unites at once with the separated iodine to form iodine chloride. It is better in such a case to employ fuming nitric acid. A few drops of this acid added to the liquid to be tested and well shaken up with it in a glass cylinder suffice to convert the iodide into a nitrate and set hydriodic acid free ; this at once gives up its hydrogen to the loosely combined oxygen of the nitric peroxide contained in the fuming nitric acid and sets iodine free : — Hydriodic Acid. 127 4HI + N 2 4 = 2 I 2 + 2H„0 + 2NO. Hydriodic N itric Nitric acid peroxide oxide If a little chloroform or carbon disulphide is then added and the liquid well shaken, the former after settling to the bottom of the cylinder becomes pink coloured, even if only traces of iodine are present. When, on the other hand, iodine is completely absent, the chloroform or carbon disulphide remains colourless. HYDRIODIC ACID. Composition : HI. Like hydrochloric and hydrobromic acids, this compound is a colourless gas, fuming in the air, with an acid taste and piercing odour, and is very easily absorbed by water. In the same manner as hydrochloric acid, it consists of one volume hydrogen united with one volume iodine vapour to form two volumes of the com- pound without condensation. Its vapour density (H = 1) is, there- 127 + 1 fore, — = 64, or compared with air as unity 4-4. In consequence of the high specific gravity of iodine gas and the low specific gravity of hydrogen, the weight of the hydrogen contained in hydriodic acid is so small, that it was for some time overlooked. From the formula? we see that 127 + 1 = 128 parts by weight of hydriodic acid only contain one part by weight of hydro- gen, or less than one per cent. ; the exact percentage being 1x100 = 078. 128 ' The chemical affinity of iodine for hydrogen is even less than that of bromine for the same substance, and hydriodic acid cannot therefore be obtained by the direct union of its elements. Neither can it be prepared pure by decomposing potassium iodide with con- centrated sulphuric acid. Large quantities of hydriodic acid are, it is true, set free by this reaction, but the gas is largely contaminated with iodine, because the hydriodic acid is decomposed at the moment of its production by the excess of sulphuric acid into iodine, water, and sulphurous acid. In order to prepare pure hydriodic acid we can only use one 128 Text-Book of Inorganic Chemistry. method — viz. the decomposition of phosphorous iodide by water. In order to form this compound ordinary (yellow) phosphorus and iodine are made to combine in the proportion of one equivalent of the former to three equivalents of the latter— i.e. in the proportion of 31 parts of phosphorus to 3 x 127 = 381 parts of iodine, or 1 part to i2'3 parts. But as it is better to have the phosphorus in slight excess, 1 part of phosphorus is mixed with jo parts of iodine. The proper quantity of iodine is weighed off into a dry flask, the flask filled with carbonic acid, and then the requisite quantity ot phosphorus, cut into small pieces and dried, is gradually added. The two bodies unite immediately with one another, producing flashes of light, to form liquid phosphorous iodide. In order to distribute the excess of phosphorus uniformly through the mass, the compound is gently heated and kept in a molten state for a short time. After the compound has cooled and solidified to a crystalline mass, it is moistened with a little water and gently warmed. Phos- phorous acid and hydriodic acid are produced, and the latter passes away through a tube fitting in an india-rubber stopper in the neck of the flask. It may be led into water if an aqueous solution of the gas is to be prepared. The delivery tube must be wide, otherwise it is apt to become plugged up for the following reason. Phosphorous acid is decomposed on heating into phosphoric acid and phosphoretted hydrogen, and this decomposition is greatest, the smaller the quantity of water present. But phosphoretted hydrogen and hydriodic acid, when in contact with one another, unite to form a solid crystalline substance of the composition : PH 3 -HI = PH,I (phosphonium iodide). In the preparation of hydriodic acid this compound may be easily formed, and the delivery tube may become plugged with crystals of it. Aqueous hydriodic acid is a colourless liquid with a strong acid reaction and odour. If completely saturated with the gas in the cold, it fumes when exposed to the air. When kept in vessels containing air, especially when exposed to the light, the solution soon becomes of a yellow or brown colour, from the iodine liberated by the oxidation of the hydrogen. The weak affinity with which the hydrogen and iodine are combined together is apparent from the fact that when a glass rod is heated in the flame, and plunged into a jar of the gas, violet vapours of iodine are at once produced. If the rod was red-hot, the liberated hydrogen often catches fire. Chlorine and bromine decompose hydriodic acid and all the Oxy-acids of Iodine. 1 29 iodides (except silver iodide) very easily. If a jar of chlorine is inverted over one of hydriodic acid gas, and then the position of the jars reversed, so that the heavier hydriodic acid comes upper- most, a considerable quantity of heat is produced, and iodine is set free, with the production of hydrochloric acid. If the chlorine is in excess, the iodine soon disappears, uniting with the chlorine to form iodine chloride. Hydriodic acid and the soluble iodides produce with silver nitrate a yellow precipitate of silver iodide, which is almost inso- luble in ammonia. OXY-ACIDS OF IODINE. Of the oxygen compounds of iodine only two are well known — iodic acid and periodic acid — which correspond in their composition to chloric acid and perchloric acid. Iodic Acid : IO„-OH. This compound is distinguished from the unstable chloric acid by its comparative stability. It is a solid, and crystallizes in hexagonal plates ; is soluble in water and in alcohol, and possesses a bitter acid taste. It may be heated to above ioo c without being essentially changed ; at a higher temperature it is decomposed into water and iodic anhydride — 1 2 5 . Iodic acid may be prepared by the direct oxidation of iodine by nitric acid, if too large quantities are not employed at one time. About 10 grammes of iodine are finely powdered and gently heated in a capacious flask with twice as much concentrated nitric acid. As soon as no further action takes place the acid liquid is poured off, and the residue digested with a fresh quantity of nitric acid until all iodine has disappeared. As iodic acid is only slightly soluble in nitric acid, the greater portion remains behind in the flask as a crystalline powder ; the acid liquid is poured off and evaporated to dryness, when a further quantity of iodic acid is obtained. Both portions are then dissolved in water, evaporated to dryness, to expel the last traces of nitric acid, and, finally, heated to 100° — 130 in a stream of dry air. An aqueous solution of iodic acid does not yield large crystals of the acid when evaporated down, but such crystals may be K 1 30 Text-Book of Inorganic Chemistry. easily obtained by adding a little nitric or sulphuric acid, and then evaporating. Barium iodate is decomposed when boiled with dilute sulphuric acid into insoluble barium sulphate and iodic acid. If the clear liquid is filtered off, crystals of iodic acid separate on evaporation. At the ordinary temperature sulphuric acid has but little action on the barium salt. To prepare iodic acid by this process, two parts of concentrated sulphuric acid diluted with eight parts of water are allowed to act upon nine parts of the finely-powdered barium salt. If iodine is boiled with caustic potash, the solution contains potassium iodide and potassium iodate — just as when chlorine acts upon a hot concentrated solution of caustic potash — and the iodate, being less soluble than the iodide, crystallizes out first on evapora- tion. In a solution of potassium iodate, barium chloride produces a white precipitate of the difficultly soluble barium iodate. This salt may also be directly prepared by boiling iodine with concen- trated baryta water. Iodic acid when heated up to 130°, or slightly higher, loses water and becomes converted into a compound having the com- position I 3 7 -OH, which may be considered as one molecule of iodic acid united with one molecule of iodic anhydride— i.e. I0 2 OH + I 2 5 . It is distinguished from iodic acid by its insolubility in ordinary alcohol. If iodic acid is heated still higher, up to 170 , it loses a further quantity of water, and is completely converted into the anhydride, I 3 5 . The decomposition is also produced by absolute alcohol, or by a mixture of this substance with sulphuric acid. Iodic Anhydride : I 2 6 or ,,-,'- \ O. This compound is easily soluble in water, and then reproduces iodic acid. When strongly heated — up to about 300 — it is de- composed into iodine and oxygen. Although the oxygen in iodic acid is much more firmly com- bined than in chloric or bromic acid, the acid still easily gives up its oxygen to those substances which exercise a reducing action, and is therefore a powerful oxidizing agent. Sulphurous acid and sulphuretted hydrogen, when added to a dilute aqueous solution of iodic acid, cause an immediate separation of iodine ; nitrous acid, phosphorous acid, and other substances produce the same result. Oxy-acids of Iodine. 1 3 1 With dry hydrochloric acid gas, dry iodic anhydride yields iodine chloride, chlorine, and water : — I 2 5 + 10HCI = 2lCl s + 2C1 2 + sH,0. Iodic acid is a monobasic acid, and produces salts which are mostly insoluble in water. It shows a strong tendency to form so- called acid salts besides the neutral compounds. The acid salts may be considered as neutral salts united with one or two molecules of the acid. We know, for example, the following three potassium compounds : — Normal potassium iodate . . I0 2 OK. Monacid potassium iodate . . I0 2 OK + I0 2 OH. Diacid potassium iodate . . I0 2 OK + 2l0 2 OH. Like the chlorates, the iodates of the metals are decomposed on heating into iodides and oxygen — the iodates of the heavy metals usually giving up iodine as well, and becoming converted into oxides. It is not, however, possible to expel all the oxygen from potassium iodate by heating it even to redness. Iodic acid possesses the remarkable property of uniting chemi- cally with strong acids. If dry and powdered iodic acid is gradu- ally added to about five times its weight of hot concentrated sulphuric acid, a substance separates out on cooling which contains both acids chemically united. Periodic Acid : I0 3 -OH + 2H 2 0, or IO(OH) 5 . The monobasic acid is unknown in the free state ; we are only acquainted with its compound, with two molecules of water, which may be considered as a pentabasic acid. This substance crys- tallizes in colourless rhombic prisms, which are soluble in water, alcohol, and ether, and deliquesce when exposed to the air. The aqueous solution may be boiled without causing decomposition, and the crystals are not altered by standing in a desiccator over sulphuric acid. At a temperature of 133 the acid melts, and begins to be decomposed into water, oxygen, and iodic anhydride at 140°. Four kinds of periodic acid are known in its salts, which are mono-, tri-, tetra-, and pentabasic respectively. These acids may be represented as compounds of the unknown periodic anhydride : k 2 132 Text-Book of Inorganic Chemistry. I 2 7 , containing heptad iodine, with varying proportions of water, thus : — I 2 7 + H 2 = 2l0 3 -OH. TO + 2HO - o^ I0 ^ OH) *- I,0 7 + 3H 2 . . . = 2lO,(OH) 3 . I 2 7 + 5 H 2 = 2lO(OH),. Of these acids the last-named is known in the free state, but the others only in their salts. Pentabasic periodic acid may be obtained in the following way. When chlorine is led into a hot solution of sodium iodate, containing free caustic soda, an acid sodium salt of the pentabasic periodic acid is produced according to the equation : — I0 2 -ONa + 3NaOH + Cl 2 = I°Kq^ s + 2NaCl. This sodium salt, which is nearly insoluble in water, dissolves in nitric acid, and gives with silver nitrate a precipitate of the corresponding silver salt, which when evaporated down with nitric acid, is converted into yellow crystals of the silver compound of the monobasic acid : I0 3 -OAg. Finally, if this salt is boiled with water, the former silver compound is again produced, and the pentabasic acid set free : — 2l0 3 -OAg + 4 H 2 = 10{ ( g^ + IO(OH) 5 . The solution is then filtered and evaporated down to crystallize. The normal silver salt of the pentabasic acid, IO(OAg),, has also been prepared, as well as the sodium salt of the tetrabasic acid : O j |q /ON^V b y heatin £ the above acid salt to 220°— 2lU {(OH) 3 " U jl0 2 (ONa) 2 + 3Jrl 2 U - As will be seen from the above formulas, the composition of the periodates is very various and somewhat complex. Similar com- pounds are known of the oxy-acids of sulphur and phosphorus, which have been subjected to more exact investigation than the periodates. Fluorine. 133 Chlorides of Iodine. Chlorine and iodine unite together to form two compounds — iodine monochloride, and iodine trichloride. Iodine monochloride : I CI, is a red-brown liquid, which crystal- lizes on standing, and which possesses the odour of both iodine and chlorine. The crystals melt at 25° ; they are soluble in water, whence they may be again extracted by shaking with ether. Iodine monochloride is prepared by heating together one part of iodine with four parts of potassium chlorate, or by dissolving iodine in aqua regia, diluting with water and extracting with ether. It then remains behind on allowing the ethereal solution to evapo- rate. It may also be prepared by acting upon one molecule of iodine trichloride with one molecule of iodine : — I'C1 3 + I„ = 3ICI. Iodine trichoride : IC1 3 , crystallizes as yellow needles when chlorine is led over gently heated iodine ; if it contains iodine monochloride its colour is more or less of a brownish tint. On heating, it melts and becomes brown, being decomposed into chlorine and iodine monochloride. FLUORINE. Chemical Symbol: F. — Atomic Weight: 19. This element possesses such powerful chemical affinities for other bodies that it has not yet been isolated. It is therefore unknown to us in the free state, but the chemical nature of its compounds justify us in classing it with chlorine, bromine, and iodine. Attempts to isolate fluorine have mostly failed, because no substance of which vessels might be constructed can withstand its action. Glass and porcelain, as well as silver, gold, and platinum, are at once attacked by it. It is very remarkable that up to the present no compound of fluorine with oxygen is known, neither an acid nor a salt. And there is, in fact, no element of which we know so few compounds as fluorine. It combines neither with chlorine, sulphur, nitrcgen, 134 Text-Book of Inorganic Chemistry. nor carbon, and the only non-metallic elements with which com- pounds of it are known, are hydrogen, phosphorus, boron, and silicon. In the mineral kingdom it occurs exclusively in combination with various metals, especially with calcium, aluminium, and sodium. Calcium fluoride is found crystallized in cubes and is called fluor-spar. This is by far the most abundant compound of fluorine. With aluminium and sodium it forms a double salt having the composition Na 3 AlF 3 = 6NaF,Al 2 F 6 . This mineral occurs chiefly in Greenland, and is known as cryolite. HYDROFLUORIC ACID {Fluoric Acid). Composition : HF. This compound is a colourless gas with a piercing acid odour, and may be condensed by cold to a colourless liquid boiling at 19 and not solidifying at — 34° The gas acts very injuriously upon the lungs and air passages when breathed in small quantities, and in larger quantities may produce death. Like hydrochloric acid, the gas fumes powerfully in the air, and is also energetically absorbed by water with considerable evolution of heat. If led into a vessel of water which is kept cool, as long as absorption takes place a fuming aqueous solution is produced, which gives off hydrofluoric acid when heated. At the same time the boiling-point gradually rises and at last remains constant (under the normal atmospheric pressure) at 120° The aqueous hydrofluoric acid, which distils over at this temperature, contains about 36 per cent, of the gas by weight. As all siliceous substances are decomposed by hydrofluoric acid, with the production of silicon fluoride, the acid cannot be prepared in glass or porcelain vessels. It is best to employ vessels of platinum, which is not in the least attacked by the acid, or of lead, upon which it only slightly acts. A very convenient apparatus is made of lead, and has a platinum tube fastened air- tight into it to conduct the gas into water contained in a platinum vessel (fig. 41). The vessel which serves as the retort is a cylinder of lead, on Hydrofluoric Acid. 135 which fits a cap of the same metal, and which is pierced by a hole at the side to receive the platinum tube. Finely powdered fluor- spar is mixed with concentrated sulphuric acid in the vessel to about the consistency of cream, the cap is placed on and made tight with plaster of Paris, and, finally, the platinum tube is fixed in position and also cemented with plaster. As soon as the plaster has set the cylinder is placed on an iron plate and gently warmed. The tube is so arranged that it just touches the surface of the water which is to absorb the gas, and which is contained in a platinum crucible. The crucible must be surrounded with ice or Fig. 41. cold water in order to keep it cool. By means of this simple apparatus, concentrated aqueous hydrofluoric acid may be quickly obtained at any time. To prepare the gas perfectly pure and absolutely dry it is best to employ the compound which is produced by acting on potassium fluoride with hydrofluoric acid. This substance, which has the composition KHF 2 , decomposes again when heated in a platinum vessel into potassium fluoride and hydrofluoric acid. Fluorine in combination with hydrogen and with most of the metals plays the part of a monad element. But in certain double fluorides {e.g. the potassium compound just referred to) it appears 1 36 Text-Book of Inorganic Chemistry. as if two atoms of fluorine coalesced to form a dyad, and this compound would have a composition analogous to that of potas- sium hydrate, thus : — Potassium hydrate .... KO"H. Hydric potassium fluoride . . . KF"H. By the union of hydrofluoric acid with bases the fluorides are produced. Many of the fluorides are soluble in water, and to these belongs the silver salt. In this respect hydrofluoric acid differs from the otherwise similar acids, hydrochloric, hydrobromic, and hydriodic acids, which form insoluble silver salts. On the other hand, the fluorides of some metals {e.g. calcium), which form very soluble chlorides, bromides, and iodides, are quite insoluble in water. One of the most important properties of hydrofluoric acid, in which it is again distinguished from the other halogen acids, is its decomposing action on silica and the silicates. No other acid, not even sulphuric acid or phosphoric acid, attacks silica, while hydrofluoric acid dissolves it, forming gaseous silicon fluoride and water : — Si0 2 + 4HF = SiF 4 + 2H 2 0. If the silica is united with bases, as, for example, in common glass with lime and soda, the corresponding fluorides are produced by the action of the acid. We make use of this property not only to test for hydrofluoric acid, but also to etch glass — for example, to mark a series of divisions on a glass tube. The operation for either of these purposes is as follows : Fluor-spar, or the substance to be exam- ined for fluorine, is finely powdered, placed in an open platinum crucible, mixed with concentrated sulphuric acid, and, if necessary, very gently warmed. A watch-glass is then taken and its convex side covered with a thin layer of wax, by warming it gently and rubbing it over with a piece of wax. As soon as the glass is cold the characters which are to be etched are written on the waxed surface with a sharply-pointed piece of slate-pencil, which removes the wax at those points where it touches the glass. If the watch- glass so prepared is now placed over the crucible with its convex side downwards, the glass becomes corroded at those points which are not protected by the layer of wax, and only at those points. By the action of the hydrofluoric acid on the glass, silicon fluoride, sodium fluoride, and calcium fluoride are produced. The firsts Hydrofluoric Acid. 137 named substance passes away as a gas, the second is removed by subsequent washing with water carrying the third mechanically with it. If the glass is removed after remaining for a short time over the crucible, washed with water and the wax removed by warming, the characters traced on the glass become distinctly visible, in consequence of the contrast between those parts of the glass which have been etched and those which were unacted upon by the acid. For the production of delicate and exactly equal divisions on glass tubes and other similar etchings, it is best to employ the commercial aqueous acid, which only fumes slightly in the air. The divisions having been marked on the waxed tube, this aqueous acid is painted on the parts to be etched with a camel's hair brush. The fact that hydrofluoric acid etches glass was observed two hundred years ago by Schwankhard of Nuremberg — i.e. he found that a mixture of fluor-spar and sulphuric acid corroded glass. But it was only a hundred years later that Scheele, the discoverer of chlorine, showed that this action is 4 ue to a gas given off by this mixture, and the composition of this gas was first discovered by Ampere at the commencement of the present century. Gutta-percha, like wax, is neither attacked nor dissolved by hydrofluoric acid, and the aqueous acid can therefore be preserved and transmitted from one place to another in bottles of gutta-percha provided with stoppers of the same substance. ELEMENTS OF THE SULPHUR GROUP. To this group belong the three elements, sulphur, selenium, and tellurium. They form a natural group similar to those referred to on p. 65. All three are solid and crystalline ; they can be easily melted and sublimed, and are combustible. Selenium and tellu- rium are distinguished from sulphur by their metallic lustre. Their chemical nature is interesting from the fact that they can form two distinct classes of compounds in which they play different parts. On the one hand, they so closely resemble oxygen that they can partially or entirely displace this element from its com- pounds ; on the other hand, they combine with oxygen itself to. 138 Text-Book of Inorganic Chemistry. form compounds usually possessing acid properties, and in which they have the same relation to oxygen as phosphorus has in its acids and arsenic in its acids. In all cases where they take the place of oxygen, their atomicity, like that of this element, is always two, while in their compounds with oxygen they offer four or six points of attraction to this element. When sulphur, which may be here considered as a representa- tive of the entire group, enters into chemical combination as a dyad, the compounds produced — sulphides — exhibit a close che- mical relationship to the corresponding oxides. Thus : — Water H 2 corresponds to Sulphuretted hydrogen .... H 2 S. Arsenious oxide .... As 2 3 corresponds to Arsenious sulphide .... Potassium hydrate . corresponds to Potassium sulphydrate Cupric oxide .... corresponds to Cupric sulphide. KOH KSH. CuO CuS. Arsenious oxide and arsenious sulpide possess acid properties while potassium hydrate and sulphydrate, cupric oxide, and sulphide, are bases. Potassium arsenite is produced by the union of arsenious oxide and caustic potash, and, in the same manner, from arsenious sulphide and potassium sulphydrate potassium sulpharsenite is formed : — As 2 3 + 2KOH = 2As0-0K + H„0, As a S 3 + 2KSH = 2AsS-SK + H S S. So, too, arsenious oxide and cupric oxide unite to form cupric arsenite, while the union of arsenious sulphide and cupric sulphide produces cupric sulpharsenite ; the sulphur compounds contain- ing in every case sulphur instead of oxygen. But when the elements of the sulphur group are united with oxygen to form the radicals of the oxy-acids, their chemical character is entirely different. In the compounds of dyad sulphur Sulphur. 139 this element can always be displaced by oxygen ; but this is not possible in the second class of compounds, in which the sulphur exists as a tetrad or hexad element combined with oxygen, because, as far as we know, oxygen only exists as a dyad element, and never possesses any higher atomicity. We might imagine the existence of a compound corresponding to sulphurous anhydride— S lv 2 — in which the two atoms of dyad oxygen might be displaced by two atoms of dyad sulphur — S iv S" 2 , and the sulphur which is set free by the action of sulphuretted hydrogen on sulphurous anhydride is perhaps this sulphur : — S lv 2 + 2H 2 S» = S l S\ + 2H 2 0. But we cannot imagine the existence of a similar compound in which the tetrad sulphur of sulphurous anhydride is displaced by tetrad oxygen, O lv O", because, as far as we know, oxygen never plays the part of a tetrad. The composition of thiosulphuric acid proves that the displace- ment of oxygen in the oxy-acids, by dyad sulphur, is not only pos- sible, but actually takes place. This acid is sulphuric acid, S vi 2 j qtt, in which one of the atoms of oxygen united with the hydrogen is displaced by an atom of dyad sulphur, thus : s-o i s ' iH These facts are not only true of sulphur compounds, but may be generally extended to those of selenium and tellurium. Selenium and tellurium may be said to occur as rarely in nature as sulphur does abundantly. These two elements (particularly selenium) very often occur associated with sulphur. And their scarcity is probably the reason why their compounds have not been so thoroughly investigated as those of sulphur. SULPHUR. Chemical Symbol : S. — Atomic Weight: 32. Sulphur occurs free in nature in considerable quantities ; Sicily and other parts of Italy as well as Iceland are particularly rich in native sulphur. It is still more widely, distributed in chemical combination with other bodies. Of the sulphides, the commonest 140 Text-Book of Inorganic Chemistry. are ironpyrites (ferric disulphide, FeS 2 ), copper pyrites (copper-iron sulphide, CuFeS 2 ), galena (lead sulphide, PbS), blende (zinc sul- phide, ZnS),. &c. Sulphur is not less common in combination with calcium and oxygen, as calcium sulphate, in the mineral gypsum (S0 2 -0 2 Ca + 2H 2 0), which often occurs in layers of con- siderable thickness. Other sulphates are also found in nature, among which may be mentioned Glauber's salt (sodium sulphate, S0 2 (ONa) 2 , Epsom salt (magnesium sulphate, S0 2 -0 2 Mg) — two compounds which are contained in mineral waters and in sea water. Other naturally occurring sulphates are, heavy-spar (barium sulphate, S0 2 -0 2 Ba), alunite (basic aluminium-potassium sulphate), &c. Besides all these forms in which sulphur is found in nature, it also forms a constituent of many important organic substances, such as albumen. Few substances are so generally known as sulphur, and yet the outside world is but little acquainted with its remarkable physical and chemical properties. Sulphur is a brittle, crystalline solid, with a pale yellow colour, without taste or odour, quite insoluble in water, and only slightly soluble in alcohol and ether. It is, however, dissolved by carbon disulphide in large quantities, especially when the liquid is heated to its boiling point. At 1 1 5 it melts to a clear mobile liquid of a pale yellow colour, which easily resolidifies to ordinary yellow sulphur, but which if more strongly heated (up to 200 ) undergoes a remarkable change. It does not become more mobile, as might be expected, but more and more viscid, and at the same time darker and darker in colour. At 200° it is a dark-brown liquid so viscid that the vessel containing it may be inverted without any of it running out. When heated above 200°, up to 400 , the liquid gradually becomes mobile again, without appreciably changing its colour, and boils at 448°. The sulphur gas which is so produced and fills the vessel has a dark red- brown colour, somewhat resembling gaseous bromine. Sulphur which has been heated up to 200° possesses altogether other properties than the normal liquid sulphur at 115 . If the latter be cooled, for example, by pouring it into cold water, the original brittle yellow sulphur, soluble in carbon disulphide, is obtained. But if the sulphur at a temperature of 200°, or slightly higher, is poured into cold water, a tough, brown, elastic mass is produced, which may be drawn out into long threads and which only changes into ordinary yellow, brittle sulphur after some time. Sulphur. 141 This tough elastic modification of sulphur, unlike ordinary sulphur, is insoluble in carbon disulphide. Only after some time, or after it has been heated to 1 1 1° and allowed to cool again, does it regain its solubility in this liquid. Sulphur is dimorphous — i.e. it crystallizes in two distinct forms. The transparent yellow crystals, as they occur in nature, belong to the rhombic system. The two accompanying figures (fig. 42) re- present two common forms of native sulphur : the one (a) is simple, the other (b) has numerous secondary faces. Sulphur crystallized in the wet way — e.g. from its solution in carbon disulphide — shows exactly the same forms. Large regular crystals of rhombic sulphur may be obtained by heating a piece of sulphur with carbon disul- phide in a sealed glass tube up to 120 or 130 and allowing it to cool slowly. 1 The crystals are afterwards freed from the carbon Fig- 43- disulphide and dried, when they preserve their transparency and crystalline form unchanged. On the other hand, when sulphur is melted and quickly cooled, it crystallizes in the monoclinic system (fig. 43). These crystals are best obtained by melting sulphur in a crucible at the lowest possible temperature, so that it remains mobile and does not pass into the viscid condition, allowing it to cool until a thin solid crust forms on the surface of the sulphur, then breaking a hole in this crust and pouring out the still liquid sulphur inside. If the crucible is broken up when cold the interior will be found filled with pale yellow needle-like prisms, often as much as an inch long. These crystals soon lose their transparency and become opaque and brittle ; they then consist of a number of minute rhombic crystals similar in form to those in which sulphur occurs in nature, and in which it crystallizes from its solution in carbon disulphide. These 1 The glass tube must be thick and well sealed up in the blowpipe, other- wise a very dangerous explosion may be produced. — Ed. 142 Text-Book of Inorganic Chemistry. two modifications of sulphur are also distinguished from one another by their different specific gravity. The specific gravity of natural or rhombic sulphur is 2-07, while that of the monoclinic modification is i - 96. The vapour density of sulphur at different temperatures is remarkable. If the molecule of gaseous sulphur consisted, accord- ing to the general rule, of two atoms, its molecular weight would be 2 x 32 = 64, and its density --\ =2-216 (p. 51). But the direct 2o'oo determination of the density of sulphur vapour at about 500° — a temperature considerably above its boiling point — gives quite a different number — viz. 6-654, or about three times 2-216. Further experiments undertaken to explain this anomaly showed that between the temperatures 440 and 850 sulphur vapour did not expand regularly with the increase of temperature, and that at 900 to 1,000° its density was only one-third of that at 500 , or about 2-216. From this it appears that there are two polymeric modifica- tions of sulphur in the gaseous state, one with the density 2-216, and the other with a density three times as great — viz. 6-654. The molecule in the one case consists, therefore, of two atoms, and in the other case of six atoms of the element, and the latter molecule breaks up, when heated, into three new molecules, each containing two atoms of sulphur. Most of the sulphur which comes into trade is brought from Italy, and especially from Sicily, where the crude sulphur is partly melted out, and partly distilled in cast-iron vessels, from the earthy impurities. This sulphur is, however, still impure, and must be purified by distillation. For this purpose the arrangement shown in fig. 44 is employed. The sulphur is heated to boiling in the cast-iron retort G, which is built into the masonry, and of which the neck D opens into a large empty bricked chamber A. In the same manner as water vapour when cooled at once below its freezing-point condenses in the form of snow, so the sulphur vapour when issuing into the cold chamber at a temperature below its melting point condenses to sulphur snow, which collects on the walls and floor of the chamber. This light pale yellow sulphur- snow is the flowers of sulphur of commerce. But when the hot sulphur vapour has raised the temperature of the chamber to the melting-point of sulphur (115°) the sulphur melts and collects on the floor as a thin liquid. It may then be allowed Sulphur. 143 to flow out by an opening which can be regulated from the outside. This sulphur is nearly pure. Stick sulphur is easily obtained from it by running the liquid sulphur into wooden moulds (fig. 45). The sulphur gas which streams into the chamber is at a higher temperature than that necessary for the combustion of sulphur ; it therefore unites with the oxygen contained in the air of the chamber and burns to form sulphurous anhydride, evolving at the same time a considerable quantity of heat. In consequence Fig. 44. Fig- 45- of this increase in temperature the whole of the air of the chamber suddenly expands, and the walls would not be able to withstand the increased pressure if the chamber were not furnished with a valve. This valve opens during the combustion of the sulphur and closes afterwards by its own weight. The sulphurous anhydride which is now contained in the chamber is absorbed by the flowers of sulphur, and if this sulphur is afterwards brought into contact with moist air the sulphurous 144 Text-Book of Inorganic Chemistry. anhydride is gradually oxidized to sulphuric acid. The sulphuric acid adheres so persistently to the flowers of sulphur that, even after repeatedly washing with water, sulphuric acid may be always recognized in the wash-water. For this reason flowers of sulphur moistened with water always redden blue litmus paper, which is not the case with stick sulphur. Sulphur is also prepared by heating iron pyrites in a small supply of air, and remelting the sulphur, but the product so ob- tained is nearly always contaminated with arsenic. All elements which unite with oxygen, either more or less energetically, also combine with sulphur and in nearly the same degree. And sulphur can also unite with oxygen itself. At the ordinary temperature the two elements do not unite with one another ; but if sulphur is heated in the air, or in oxygen, it catches fire and burns with a pale blue flame, to form sulphurous anhydride. Besides this body there are several other compounds of sulphur and oxygen, of which sulphuric acid is the best known. Sulphur unites directly with hydrogen, only with difficulty, and the compound of the two elements — sulphuretted hydrogen — is easily decomposed again. Chlorine and sulphur combine with one another when gently heated. Phosphorus unites with sulphur as readily as it does with oxygen, and two of the compounds produced correspond in composition to phosphorus and phosphoric anhy-. drides. Nitrogen and sulphur have only a weak affinity for one another, and only combine under peculiar circumstances. Finally, charcoal (carbon) burns in sulphur gas as it does in oxygen, but requires a higher temperature. The compound produced, carbon disulphide, has a similar composition to the oxygen com- pound, carbonic anhydride. All those metals which combine with oxygen either at the ordinary temperature or when heated burn almost as easily in sulphur gas, forming sulphides, with a similar composition to the corresponding oxides. Thin copper foil brought into the vapour of boiling sulphur burns brilliantly, with a considerable evolution of light and heat, to form molten copper sulphide. The sulphides of those metals of which the oxides are decom- posed by heating — e.g. of silver, gold, platinum — are also reduced when heated, especially if exposed to the air. Mercuric sulphide (cinnabar), which simply volatilizes when heated, forms an excep- tion to this rule, since the corresponding oxide is decomposed into mercury and oxygen by heat. Sulphuretted Hydrogen. 145 Accordingly as the oxides of the metals are soluble or insoluble in water,, so are the corresponding sulphides. Soluble potassium sulphydrate corresponds to soluble potassium hydrate, calcium sulphide to calcium oxide, and insoluble lead sulphide to insoluble lead oxide. The metallic sulphides which are soluble in water possess the property of uniting chemically with more atoms of sulphur than are contained in the normal compounds. Potassium sulphide (K 2 S), for example, can unite with four more atoms of sulphur, pro- ducing potassium pentasulphide, which is also soluble in water : — ICS + 4S = K 2 S 5 . If a solution of this or a similar polysulphide, soluble in water, is mixed with hydrochloric acid, the corresponding chloride is formed, sulphuretted hydrogen is liberated, and the excess of sulphur sepa- rates out as a fine, white, amorphous powder. Milk of sulphur is this white precipitate which has been washed and dried. The equation illustrating its production is as follows : — K„S 5 + zHCl = 2KCI + H.,S + 4S. SULPHURETTED HYDROGEN. Composition : H 2 S. This compound is analogous to water in its composition, and also resembles water in many of its chemical properties. It is a colourless gas with a disgusting odour resembling rotten eggs, has a specific gravity of 1-19, and may be condensed, under pressure, to a mobile liquid, which solidifies at about — 85° The gas is some- what soluble in water, to which it imparts its odour ; the aqueous solution reacting slightly acid with litmus paper. A saturated solution in water at the ordinary temperature contains about three times its volume of the gas. Alcohol dissolves more, up to five times its volume. Sulphuretted hydrogen, when inhaled in considerable quantities, is poisonous. The chemist, to whom the gas is indispensable, becomes more sensitive to the poison the more frequently it is inhaled, and the same is also true of other poisons, such as hydro- cyanic (prussic) acid. In a high state of dilution, sulphuretted L 146 Text- Book of Inorganic Chemistry. hydrogen is present in many mineral waters — the so-called sul- phureous springs {e.g. those of Aix-la-Chapelle, Harrogate). These waters are often used for medicinal purposes. Sulphur and hydrogen unite directly with One another when hydrogen is led over heated sulphur ; but as the affinity between the two substances is small, only a small quantity of the compound is produced. It is better to bring hydrogen in the nascent state into contact with sulphur, as by decomposing a suitable metallic sulphide with a strong acid. The gas is easily obtained in quantity when pieces of ferrous sulphide contained in a Woulff's bottle are acted upon by dilute sulphuric or hydrochloric acid. In order to free the gas from mechanical impurities it is led through a wash-bottle containing water, and is then passed through a calcium chloride tube to dry it. It cannot well be collected over mercury, as it is partially decom- posed by the metal, forming mercuric sulphide. It is best to em- ploy warm boiled water, which dissolves much less of the gas than water at the ordinary temperature. The ferrous sulphide used for the preparation of sulphuretted hydrogen usually contains some free iron, and the gas is therefore generally mixed with a small quantity of free hydrogen. In most cases this is immaterial, but if the gas is wanted quite pure it may be obtained by acting on antimonous sulphide (black antimony) with dry hydrochloric acid gas. The reactions in each of these cases are expressed in the following equations : — FeS + 2HCI = H„S + FeCL,. Ferrous sulphide Ferrous chloride Sb 2 S 3 + 6hci = = 3H 2 S + 2SbCl 3 . Antimonous sulphide Antimonous chloride Sulphuretted hydrogen is combustible, and is easily ignited. If a burning body is brought near a jet of the gas it catches fire and burns with a pale, bluish flame, producing water and sulphurous anhydride. If the supply of oxygen is insufficient, the hydrogen only burns, and yellow sulphur is separated. The combustion is always accompanied with the separation of sulphur, and it is there- fore probable that the gas is first decomposed and its constituents afterwards burnt. In fact, the two elements in sulphuretted hydrogen are only very loosely combined. The compound is not only decomposed Sulphuretted Hydrogen. 147 by a high temperature, but also by all oxidizing agents, even the weakest, sulphur, being usually separated. Even the oxygen of the atmosphere decomposes it in presence of moisture. In a bottle half filled with sulphuretted hydrogen water, especially when often opened, the odour of the gas gradually disappears and the liquid becomes milky from separated sulphur. Sulphuretted hydrogen led through water containing iodine in suspension is rapidly decom- posed, sulphur being set free and hydriodic acid formed ; and this reaction may be utilized to prepare a dilute aqueous solution of hydriodic acid. Bromine and chlorine act in the same manner, but more energetically. Finally, ferric chloride is reduced by sulphuretted hydrogen to ferrous chloride with the formation of hydrochloric acid and the separation of sulphur : — Fe,Cl + H 2 S = 2HCI + 2FeCl 2 + S. Ferric Ferrous chloride chloride The insolubility of the majority of the metallic sulphides in water and the insolubility of some of them in dilute acids makes sulphuretted hydrogen a valuable reagent for the precipitation of many of the metals from the solutions of their salts. If the gas is led through a solution of silver nitrate, or if sulphuretted hydrogen water /is added to this solution, a black precipitate of silver sulphide is formed and the solution then contains dilute nitric acid : — aNCVOAg + H,S = Ag 2 S + 2N0 2 -OH. In the same manner the gas produces a yellow precipitate of arsenious sulphide in an acid solution of arsenious acid, an orange coloured precipitate of antimonous sulphide in a solution of anti- monous chloride, &c. The affinity of sulphur for most of the metals is so strong and for hydrogen so weak that some metals which do not combine directly with oxygen can decompose sulphuretted hydrogen. Silver, for example, becomes brown in an atmosphere containing the compound, owing to the production of silver sulphide. "We say the silver tarnishes. This action is even produced by the minute quantities of the gas exhaled by the human body, contained in our coal-gas, and, therefore, always present in inhabited places. The darkening of oil paintings by age is also produced by the action of sulphuretted hydrogen. The lighter tints in these pic- tures nearly always contain white-lead (lead carbonate), and this becomes gradually converted into black lead sulphide. 148 Text-Book of Inorganic Chemistry. HYDROGEN PERSULPHIDE. The composition of this body, which is richer in sulphur than sulphuretted hydrogen, has not yet been determined with certainty owing to its instability. Its probable composition is H 2 S 2 , corre- sponding to hydrogen peroxide, H 2 2 . If hydrochloric acid is added to a solution of potassium penta- sulphide, decomposition into potassium chloride, sulphuretted hydrogen, and sulphur occurs, as shown on page 145. But if the solution of potassium pentasulphide is poured into concentrated hydrochloric acid, there is produced, besides the substances mentioned above, a heavy oily liquid, with a piercing odour, which sinks to the bottom of the acid liquid. If this liquid is removed by a separating funnel and left to itself, it gradually decomposes into sulphuretted hydrogen and sulphur. If the freshly-prepared oil is removed to a thick glass tube and then sealed up it undergoes the same decomposition. And the greater the quantity of sulphuretted hydrogen set free the greater becomes the pressure until finally the gas is liquefied. The sul- phur which is set free at the same time is usually deposited as distinct crystals in the tube. Such tubes containing liquid sulphu- retted hydrogen are dangerous to handle. It sometimes happens that they suddenly explode after having withstood the pressure of the gas for years. OXYGEN COMPOUNDS OF SULPHUR. Sulphur and oxygen unite together in two proportions and produce : — Sulphurous anhydride . . . SO.,. Sulphuric anhydride . . S0 3 . In the former of these sulphur is a tetrad, in the latter it is a hexad. Both substances unite with bases and form stable salts, and the latter when combined with water yields the most important of the acids of sulphur : — Sulphuric acid . SO J ^ Oxygen Compounds of Sulphur. 149 Besides these compounds, which are by far the most important, a number of other oxy-acids of sulphur are known, partly in the free state and partly combined in their salts. They are all dibasic except hydrosulphurous acid. Among the next important are : — I OT-T Thiosulphuric acid . SO. ' " Dithionic acid . SH I SO, -OH tSOJ-OH I SO,- OH Disulphuric acid . O ; „,, Thiosulphuric acid is a derivative of sulphuric acid ; in it one of the atoms of oxygen united with the hydrogen in sulphuric acid is displaced by an atom of dyad sulphur. The two sulphur atoms in this compound have, therefore, different functions ; the one atom is hexad, and the other, displacing the oxygen, is dyad. Dithionic acid is a compound of two atoms of the monad radical S0 2 • OH ; and in disulphuric acid two atoms of the same radical are united together by an atom of oxygen. Besides these compounds the following are also known, although less thoroughly investigated : — Hydrosulphurous acid . . . SO ^OR H Trithionicacid . . sjfg«;gg Tetrathionic acid . Mso'-OR Pentathionic acid . . . S 3 j c . ( - ) ' 2 /-j R In these formula? R stands for an atom of a monad metal in the salts of these acids, many of which have not been prepared in the free state. Hydrosulphurous acid probably contains tetrad sulphur united with one atom of oxygen and one of hydrogen to form the monad radical (SOH). Tri-, tetra-, and pentathionic acids are similarly constituted to dithionic and disulphuric acids, and contain two atoms of the radical S0 2 -OH united together by one, two, or three atoms of dyad sulphur respectively. 150 Text-Book of Inorganic Chemistry. SULPHUROUS ANHYDRIDE Composition : S0 2 . This compound, which is produced when sulphur burns in the air or oxygen, is a colourless gas with a powerful piercing odour, an acid reaction and taste. Its specific gravity is 2-216, and it contains its own volume of oxygen ; for if sulphur is burnt in a closed volume of oxygen, the volume of the gas remains unchanged. From these data the composition of the gas may be easily calcu- lated, thus : — 2 vols. (1 molecule) sulphurous anhydride weigh 2-216 x 28 ■ 88 = 64 2 „ „ oxygen „ 2x16 =32 1 atom of sulphur 32 — i.e. a molecule of sulphurous anhydride contains one atom of sulphur and two atoms of oxygen. Sulphurous anhydride may be easily condensed by cold or pressure to a colourless mobile liquid, which boils at — io° and solidifies under — 70 . The gas is absorbed by water in considerable quantities. One volume of water at the ordinary temperature absorbs about fifty times its volume of the gas. The aqueous solution possesses the odour of the gas and again gives it up when heated. Sulphurous anhydride is contained in the gases of volcanoes. It is obtained in large quantities for technical purposes by burning sulphur or by roasting sulphurous ores {e.g. iron pyrites). The gas so obtained is of course very impure ; it contains all the nitro- gen present in the air employed. The pure compound is best obtained by reducing sulphuric acid with some suitable metal — usually copper. Pieces of copper are placed in a large flask, covered with concentrated sulphuric acid, and the mixture gently heated until the evolution of gas begins (fig. 46). The two substances do not act upon one another in the cold. The gas is purified by passing it through a wash- bottle containing concentrated sulphuric acid, and is then collected in suitable vessels. As it is absorbed in such large quantities by water it cannot well be collected over this liquid ; mercury may, however, be employed. It can also be easily collected by displace- ment by leading the gas to the bottom of the vessel to be filled Sulphurous Anhydride. 151 and loosely closing the mouth. The heavy gas gradually fills the jar and expels the lighter air before it. Liquid sulphurous anhydride may be easily obtained in quantity by leading the pure gas into a tube surrounded with a freezing mixture of ice and salt ; the tube (as shown in the figure) being contracted at one point so that it may be easily melted and hermeti- cally sealed after a sufficient quantity of the gas has been con- densed. Fig. 4<5- The production of sulphurous anhydride from sulphuric acid and copper may be expressed by the following equations, which indicate that an atom of copper first abstracts an atom of oxygen from the sulphuric acid, forming sulphurous acid and copper oxide, the former breaking up immediately into sulphurous anhydride and water, and the latter uniting with the excess of sulphuric acid to form copper sulphate and water : — S0 4oH + Cu = S0 * + H 2° + Cu0 - SO »loH + Cu0 '' SC V°2 Cu + H 3°- The final products are therefore sulphurous anhydride, copper sulphate, and water, and the whole reaction may be thus ex- pressed : — 152 Text-Book of Inorganic Chemistry. 2SoJq^ + Cu = SO, + SO„-0,Cu + 2H 2 0. The student will notice that this process is essentially different from that which takes place when dilute sulphuric acid acts upon zinc to form zinc sulphate and hydrogen. In the latter case the two atoms of hydrogen in sulphuric acid are simply displaced by the equivalent quantity of zinc (one atom) : — Sa MOH + Zn = S0 2 -0 2 Zn + H 2 ; while in the former case the sulphuric acid loses an atom of oxygen and becomes reduced to sulphurous acid. Only a few metals besides copper can reduce sulphuric acid in this manner ; among them are mercury and silver. Sulphurous anhydride mixed with carbonic oxide and carbonic acid may also be cheaply obtained in large quantities by heating concentrated sulphuric acid with charcoal : : — SO *joH + c = sa i + CO + H 2 0. This method is valuable for those purposes where the admixture with carbonic oxide and acid has no injurious effect— for example, in the manufacture of the sulphites. Sulphurous anhydride is a compound radical called sulfthuryl, and behaves in its compounds like a dyad element. Mixed with oxygen and led over heated platinum in a finely divided state the two substances unite to form sulphuric anhydride : — SO, + O = S0 3 . Its aqueous solution absorbs oxygen from the air and becomes converted into sulphuric acid. It further unites directly with chlorine, when a mixture of the gases is exposed to bright sunlight, producing sulphuryl chloride — S0 2 C1 2 . This compound is a colourless liquid with a powerful odour, which boils at 70 and is decomposed by water into hydrochloric and sulphuric acids. From a saturated solution of the gas in water at o°a crystalline compound separates out, having the composition SO„ + I5H„0. It melts at 4°, and is decomposed into its constituents. Most bodies which easily give off oxygen oxidize sulphurous anhydride, in the presence of water, to sulphuric acid. It is absorbed by manganese peroxide forming manganous sulphate and by lead peroxide producing lead sulphate : — SO.;, + PbO., Sulphurous Anhydride. 153 The production of this lead sulphate is accompanied with a large evolution of light and heat. If a little dried lead peroxide enclosed in a piece of thin muslin is brought into a jar of the dry gas, the black oxide becomes red-hot and is rapidly converted into white lead sulphate. Under ordinary circumstances hydrogen has no action on sul- phurous anhydride, but in the nascent state, especially in the presence of acids, the latter is reduced to sulphuretted hydrogen and sulphur. If a little of the aqueous solution is poured into a flask containing zinc and sulphuric acid, the hydrogen which is evolved soon smells strongly of sulphuretted hydrogen and pro- duces a black precipitate of lead sulphide if led into a solution of lead acetate, at the same time the liquid in the flask becomes milky from separated sulphur. The sulphur which is here set free is produced by the action of sulphuretted hydrogen upon sulphurous anhydride. Whenever these two substances come into contact with one another, whether in the gaseous state or in solution in water, mutual decomposition ensues into sulphur and water : — 2H„S + SO, = 2H 2 + 3S. When the reaction takes place slowly the sulphur is often deposited in the crystalline form, and it is possible that considerable quan- tities of the sulphur occurring free in nature have been produced by this reaction. Sulphurous anhydride bleaches organic colouring matters like chlorine, but the bleaching effect is produced in a different way. We have seen that chlorine bleaches partly by the formation of colourless substitution compounds, and partly by oxidation in the presence of water, or by both processes together, and the bleached colours cannot be therefore restored. But it appears as if the bleaching of sulphurous anhydride were produced by its direct union with the colouring substances. And as sulphurous anhydride can be expelled from its compounds by strong acids, the colour of many bodies bleached with this substance may be restored by treatment with strong acids. A rose which has been bleached by sulphurous anhydride regains its colour when washed with water and placed in dilute sulphuric acid. Similarly, a substance which unites with sulphurous anhydride more powerfully than the colour- ing matter {e.g. a strong base), will also restore the colour. Flannel which has been repeatedly washed regains the original yellow colour of the wool owing to the action of the alkali contained in the 154 Text- Book of Inorganic Chemistry. soap used in cleansing it. Substances upon which chlorine acts too energetically— silk, wool, straw, &c. — are usually bleached by sulphurous anydride. Sulphurous anhydride also possesses powerful antiseptic proper- ties. It destroys the smell of decomposing organic substances and stops the action of those organisms which produce fermentation and putrefaction. Burning sulphur was used as early as Homer's time to disinfect closed spaces. 1 Sulphurous acid — probably SO j q H ~ is known only in the sulphites. It is a weak dibasic acid, and forms, therefore, two series of salts, normal and acid— e.g. : — Normal sodium sulphite .... SO j Qvi a f OH Acid sodium sulphite . . SO j Q^ a The sulphites are formed by the union of sulphurous anhydride with strong bases. They are all decomposed by dilute mineral acids, with liberation of sulphurous anhydride. The presence of sulphurous acid in any of its compounds may be detected by the odour of the gas when liberated by a dilute acid. If the quantity is too small to be detected in this way, the gas may be allowed to act upon iodic acid, when sulphuric acid and free iodine are produced. If a piece of paper moistened with a solution of iodic acid and a little starch is hung up in a cylinder containing the liquid to be examined, the paper soon becomes blue, owing to the union of the liberated iodine with the starch. The sulphites may, however, have a different composition to that given above. They may be formed on the type of the hypothetical f OT-f acid S0 2 j tt , which would be called hydrosulphuric acid. This acid would then be a monobasic acid, but in which the second atom of hydrogen — that united directly to the sulphur — could also be displaced by a metal. What we now call acid sodium sulphate 1 ' Bring sulphur straight and fire ' (the monarch cries) : She hears, and at his word obedient flies. With fire and sulphur, cure of noxious fumes, He purg'd the walls and blood-polluted rooms. Odyssey, xxii. 527-530. (Pope's Translation, ,) Sulphuric Acid. 155 would then be the normal hydrosulphate : SO a ] H, and our normal sulphite would be sodium sodiosulphate : SO„ j N a . These con- siderations are supported by many facts in organic chemistry, and the decision which view is correct cannot be long delayed. Possibly isomeric salts of the two acids exist together. Corresponding to this hypothetical hydrosulphuric acid with one atom less oxygen is : — Hydrosidphurous acid : SO j „ , sometimes called after its dis- coverer, Schiitzenberger's Acid. The acid, which has not yet been obtained in the pure state, is prepared by digesting a concentrated aqueous solution of sulphurous anhydride with metallic zinc, in which the latter dissolves with the evolution of gas. The yellow liquid so obtained, which contains hydrosulphurous acid, possesses powerful reducing and bleaching properties. It rapidly decomposes first into thiosulphuric acid and then into sulphurous anhydride, water, and sulphur. The sodium salt may be prepared by digest- ing together zinc and a concentrated solution of acid sodium sulphite. SULPHURIC ACID. Composition : H 8 S0 4 , or S0 2 \ qt. If a further atom of oxygen is made to unite with sulphurous anhydride, the compound sulphuric anhydride (S0 3 ) is formed. But far more important than this body is the substance produced when it unites with water — viz. sulphuric acid, or oil of vitriol, as it is sometimes called from the old method of preparing it. Sulphuric acid is a compound of hexad sulphur, and contains the radical S0 2 united to two atoms of hydrogen through the interven- tion of two oxygen atoms. Sulphuric acid is a colourless, odourless, viscid liquid, which does not fume in the air, and has a specific gravity of 1-84. It boils 1 56 Text-Book of Inorganic Chemistry. at 338 and can be distilled, but only with partial decomposition, into sulphuric anhydride and water. It mixes with water in all propor- tions, and during mixing evolves a large quantity of heat. Even when very largely diluted with water it reacts and tastes strongly acid. Below 0° it freezes to a crystalline solid. In nature sulphuric acid occurs chiefly in combination with lime as gypsum or anhydrite, widely distributed and in very large quanti- ties, and besides this as heavy spar or barite (barium sulphate), as celestine (strontium sulphate), as Glauber's salt (sodium sulphate), &c. Calcium sulphate is such a common mineral that it might be thought capable of yielding an inexhaustible source of cheap sulphuric acid, and this would be so were it possible to separate the sulphuric acid from its combination with the lime as easily as we can separate carbonic acid from chalk or nitric acid from nitre. But sulphuric acid is so firmly united with the lime in calcium sulphate that we can neither expel it by heating nor by the action of a stronger acid, and gypsum is not therefore adapted for the manu- facture of sulphuric acid, nor has it ever been used for this purpose. The acid may, however, be obtained in many other ways. Sulphur, which burns in oxygen to sulphurous anhydride, is con- verted when heated with concentrated nitric acid or other powerful oxidizing agents into the higher oxide — sulphuric acid. And here the question may present itself, Why is it that while carbon and phos- phorus when burning in the air always produce their higher oxide, sulphur only forms its lower oxide ? The reason is that sulphuric anhydride cannot exist at a high temperature, but is decomposed into oxygen and its lower oxide, and that the temperature of com- bustion of sulphur is higher than the point at which sulphuric anhy- dride is decomposed. If we could by any means lower the tempe- rature of combustion of sulphur, it would undoubtedly produce sulphuric anhydride when burning in air, but up to the present this problem has not been solved, and we must therefore look for some other method of oxidation. The best substance for this purpose, and that exclusively used for the manufacture of sulphuric acid, is nitric acid. In order to understand this process, it must be remembered that nitric acid, which has the composition NO., -OH, yields up a portion of its oxygen and becomes converted into nitric oxide (NO) in the presence of sulphurous anhydride and water. Three molecules of sulphurous anhydride therefore require two Sulphuric Acid. 157 molecules of nitric acid and two of water to be converted into sul- phuric acid : — 3S0 2 - 2\0„-OH + ?H,0 = 3 S0 2 jg£[ + 2NO. But if as in this equation two molecules of nitric acid were con- sumed in the production of every three molecules of sulphuric acid, the price of the latter acid would be scarcely lower than that of the former. As a matter of fact, the nitric acid can oxidize much larger quantities of sulphurous acid than that expressed in the above equation — more than ten times as much. This is rendered possible by the remarkable property which the colourless nitric oxide possesses of uniting with the oxygen of the air, even at the ordinary temperature, and forming red vapours consisting of nitrous anhydride and nitric peroxide, and these gases again give up oxygen to sulphurous anhydride, oxidizing it to sulphuric acid in the pre- sence of water, and being themselves again reduced to nitric oxide. It thus appears as if a small quantity of nitric oxide would be able to convert an unlimited quantity of sulphurous anhydride into sul- phuric acid, if allowed to act upon a mixture of the gas and pure oxygen together with water vapour in the correct proportions. The oxygen would at once convert the nitric oxide (NO) into nitric per- oxide (NO,,), which would be again reduced to nitric oxide in the presence of sulphurous anhydride and water vapour : — NO., + SO a + H 2 = s( V(OH + N0 > and this could then produce further quantities of sulphuric acid, and so on. But such a process cannot be carried on in practice, simply becauses it presupposes an inexhaustible supply of pure oxygen. The manufacturer must make use of the oxygen contained in the air. The air, by the gradual abstraction of its oxygen, becomes largely diluted with nitrogen, and ultimately the small proportion of oxygen present will not unite with the nitric oxide sufficiently quickly. When this point is reached, the valuable nitrogen com- pounds would be wasted if a method had not been discovered of condensing them, and reintroducing them into the chambers. The manufacture of sulphuric acid on a large scale, based upon the above-mentioned chemical principles, is conducted as follows. The sulphurous anhydride is obtained either by burning sulphur 158 Text-Book of Inorganic Chemistry. or by roasting iron pyrites. It is then allowed to stream, mixed with atmospheric air, into a large leaden chamber, where it comes into contact with nitric acid vapours. A small portion of the sul- phurous anhydride is at once oxidized by the nitric acid to sulphuric acid, but by far the greater portion passes on, mixed with nitric peroxide, nitrous anhydride, and nitric oxide, into the next chamber, into which steam is led from a boiler. The reactions mentioned above go on in this chamber, and the dilute sulphuric acid collects on the floor. 1 The motion of the gases through the chambers is produced by a tall shaft placed at the further end, which returns the useless nitrogen to the atmosphere. But before it reaches the shaft it is made to pass through what is called a Gay Lussac's tower to retain the valuable oxides of nitrogen so that they may be again used to produce a further quantity of sulphuric acid. It has been found that strong sulphuric acid can absorb nitric peroxide, and especially nitrous anhydride, in considerable quantities. This property is utilized in the Gay- Lussac's tower by causing strong sulphuric acid to trickle down in a finely divided state over pieces of coke, in order to expose as large a surface of the acid as possible to the ascending gases. The strong acid containing the oxides of nitro- gen in solution (so-called nitrated acid), which collects at the foot of the tower, is pumped back to the other end of the chambers. The gases which it contains are then again liberated by mixing it with dilute acid, and are again introduced into the chambers by allowing it to flow down a second tower (Glovers tower), through which the sulphurous anhydride and air are made to pass. The chambers in which the sulphuric acid is produced are made of lead, because lead is the only substance available which resists to some extent the action of sulphuric acid. The leaden plates constituting the chambers are melted together by means of the oxy-hydrogen blowpipe, and are supported externally by a stout wooden framewotk. The acid which collects on the floor of the chambers — the chamber acid— is somewhat dilute, and only contains little more 1 The reactions given here only explain in a general manner what goes on in the chambers. What really happens is doubtful, and is probably dependent upon the proportions in which the various substances are introduced into the chambers. If the supply of steam is deficient, white crystals having the com- f OH position S0 3 -] j^q are produced, which under the action of water at once break up into sulphuric acid, nitric acid, and nitric oxide.— -Ed. Sulphuric Acid. 159 than 64 per cent, of the pure acid (sp. gr. = 1*5 5). By heating in leaden vessels it is concentrated until it contains about. 78 percent, of the pure acid. This is the brown acid of commerce, and has a specific gravity of about 171. The concentration cannot be carried further in leaden vessels, as at this point the lead begins to be attacked. Further concentration is then carried on in vessels of glass or better of platinum. The sulphuric acid so prepared is usually coloured brown from traces of organic substances, and always contains lead sulphate in solution. If diluted with water or alcohol, the liquid becomes turbid, and ultimately deposits a white precipitate of this lead sulphate, which is more soluble in the con- centrated than in the dilute acid. If iron pyrites is used in the manufacture of sulphuric acid, the acid always contains arsenic, sometimes in not inconsiderable quantities. A sample of the acid containing arsenic when diluted with water, and saturated with sulphuretted hydrogen, first turns yellow, and then deposits a yellow precipitate of arsenious sulphide. By far the greater quantity of sulphuric acid which is brought into trade contains arsenic ; only a comparatively small quantity of the more expensive acid free from arsenic is manufactured from Sici- lian sulphur. The colourless, concentrated, chemically pure sulphuric acid is prepared by distilling the acid free from arsenic in platinum vessels. During distillation it always undergoes a partial decom- position into sulphuric anhydride, which passes over with the distillate and water which remains behind ; this goes on until the acid contains 98 per cent, of the pure compound, which then distils over unchanged. Sulphuric acid is one of the strongest acids, and expels there- fore nearly every other acid from its compounds. It is, like sulphurous acid, a dibasic acid, and its acid salts possess a strong acid reaction and taste, with few exceptions, its salts are soluble in water. Sulphuric acid is distinguished by its powerful attraction for water, with which it unites in several definite proportions, always evolving a large quantity of heat. The compound with one (Ol-F 0H + H,0 — has the specific gravity 178, and solidifies in a crystalline form at + 8°. Unlike water, it does not expand on solidification, and vessels filled with this acid do not therefore burst on freezing. If to this compound a further 1 60 Text-Book of Inorganic Chemistry. quantity of water is added, a further quantity of heat, though less than before, is evolved. And the quantities of heat which are set free when one, two, or three molecules of water unite with one molecule of sulphuric acid are in a certain definite proportion to one another. The strong attraction of sulphuric acid for water is employed to dry those gases upon which the acid has no action (e.g. oxygen, hydrogen, carbonic acid), the gases being simply led through a wash-bottle containing the strong acid. It also serves to dry solid and liquid bodies which are placed in a closed space (a desiccator) with the concentrated acid. The acid then rapidly absorbs all the water vapour which evaporates from the substances to be dried. Ordinary commercial sulphuric acid is always of a brownish colour, due to the decomposition of particles of dust of organic origin. If a splinter of wood is dipped into the concentrated acid it soon becomes brown and then black from the charcoal which the acid has set free from the wood by abstracting oxygen and hydrogen in the form of water. These elements are not present as water in the wood, but the acid compels them to combine to form water, with which it then unites. Sulphuric acid is by no means so permanent as its powerful affinities and high boiling point might lead us to suppose. Even below a red heat, if allowed to drop upon hot bricks or upon pieces of pumice stone contained in a hot platinum retort, it is decom- posed into water, oxygen, and sulphurous anhydride. If the mixture of gases so obtained is led into a solution of caustic soda, the sulphurous anhydride is absorbed, and considerable quantities of pure oxygen may be obtained. The instability of sulphuric acid at high temperatures is the reason why burning sulphur produces sulphurous anhydride and not sulphuric anhydride. Nascent hydrogen reduces concentrated sulphuric acid (not the dilute acid) to sulphuretted hydrogen : — S0 3 {°55 + 4 H 2 = H 2 S + 4 H 2 0. If a few drops of concentrated sulphuric acid are allowed to flow by means of a funnel tube into a flask evolving hydrogen, the hydrogen soon acquires the characteristic odour of sulphuretted hydrogen, and produces a black precipitate of lead sulphide in a solution of lead acetate. Zinc, iron, nickel, manganese, and other metals which possess a strong attraction for oxygen and are Sulphuric Acid. 161 dissolved by dilute sulphuric acid with evolution of hydrogen, are not attacked even when heated with the concentrated acid, probably because their sulphates are so insoluble in the concentrated acid that they incrust the metals and prevent any further action. Other metals, as copper and mercury, which in consequence of a feebler attraction for oxygen are not attacked by dilute sulphuric acid, reduce the concentrated acid when heated with it, forming sulphurous anhydride and a sulphate soluble in concentrated sulphuric acid. In this way sulphuric acid, which we prepare from impure sulphurous anhydride, affords us the best material for the preparation of this substance in the pure state. Sulphuric acid contains two atoms of displaceable hydrogen, and is therefore a dibasic acid. With monad metals it forms two series of salts — the normal and acid — e.g. : — Normal sodium sulphate .... S0 2 Jqjj iOT-f ONa Most of the sulphates are soluble in water, the most important exception being barium sulphate, which is quite insoluble in water and dilute acids, and which is therefore used to detect sulphuric acid or a sulphate in an aqueous solution. On the addition of a few drops of barium chloride to such a solution, the presence of sulphuric acid is at once recognized by a white turbidity or precipitate of barium sulphate which is insoluble in hydrochloric acid. The great progress which has been made during the past fifty years in technical chemistry and the arts generally, is due to a large extent to improvements in the manufacture of sulphuric acid. Hundreds of thousands of tons of sulphuric acid are manufactured annually in England, Germany, and France alone, and by far the greatest part of this immense quantity is used for the production of soda, from which again two of the most indispensable articles of daily life— soap and glass — are obtained. Besides this the acid is also used for an immense number of other chemical pro- cesses, and, in fact, there is scarcely any chemical manufacture in which sulphuric acid does not take a direct or indirect part. It will, therefore, be at once apparent that a reduction in the price of this important substance to one-tenth the amount paid for it a little more than a century ago must have exercised a most beneficial in- fluence on the development of technical chemistry. If as is not only possible but even probable, further improve- M 1 62 Text-Book of Inorganic Chemistry. ments are made in the manufacture of sulphuric acid, and its price becomes still lower, the price of soda, glass, soap, and stearine, of superphosphate manures and hundreds of other things, would at once fall to a considerable extent. SULPHURIC ANHYDRIDE. Composition : S0 3 . This substance, which is sometimes called sulphuric acid, is very unlike sulphuric acid in its properties. It is a solid body, crystal- lizing in white, silky needles. It melts at I5°and boils at 46 , producing vapours which form thick white fumes of sulphuric acid in moist air. The density of its vapour, compared with the air as unity, is 277, corresponding to a molecular weight of 80. Sulphuric anhydride may be easily obtained from sulphurous anhydride and oxygen. A mixture of the two gases remains un- changed even if heated or exposed to sunlight, but if led over a layer of heated platinum, contained in a tube of hard glass, the two substances at once unite with one another, and the presence of sul- phuric anhydride is made manifest by the dense white fumes pro- duced where the gases come into contact with the air. The platinum itself remains quite unchanged ; it acts in some unknown manner upon the gaseous mixture, and its action is an example of what is called, for want of a better name, contact action. Perhaps the phenomenon depends upon the property of platinum to con- dense considerable quantities of oxygen on its surface and so liquefy the gas ; or it may be that the oxygen is converted into its active modification — ozone. Instead of pure platinum it is better to em- ploy platinized asbestos — i.e. asbestos of which the surface has been covered with a thin layer of platinum. 1 Sulphuric anhydride is commonly prepared from Nordhausen or fuming sulphuric acid, which may be considered as a solution of the anhydride in sulphuric acid. When this acid is gently heated in a retort, the anhydride volatilizes and condenses in a cool dry receiver as a colourless crystalline mass. Sulphuric anhydride is especially characterized by its strong attraction for water. If a drop of the molten substance is allowed 1 Prepared by dipping asbestos into platinic chloride solution, drying, and heating to redness. — Ed. Nordhausen or Finning Sulphuric Acid. 163 to fall into a vessel of water, it sets free a large quantity of heat and hisses like a red-hot iron on immersion in water. If a drop of' water happens to fall into a glass vessel containing the anhydride, chemical union at once takes place with an explosion,, and the vessel is always shattered. In all cases the compound produced by its union with water is sulphuric acid. Sulphuric anhydride acts even more powerfully on organic substances than sulphuric acid ; it abstracts water and chars them. A cork or india-rubber stopper cannot, therefore, be used to close a bottle containing the volatile anhydride. Formerly sulphuric anhydride was only employed for certain reactions in the laboratory on a small scale, but it is now used more and more in the arts ; for example, in the manufacture of alizarine from anthracene. It is now prepared in considerable quantities in chemical works, and is an article of commerce. NORDHAUSEN OR FUMING SULPHURIC ACID. This acid is not a definite chemical compound, but consists (as we have previously stated) of sulphuric acid containing more or less sulphuric anhydride in solution. In former times it was principally manufactured at Nordhausen, in the Harz Mountains, though at present none is made there. Fuming sulphuric acid is a thick oily liquid, usually coloured brown from the presence of minute particles of carbon, and is dis- tinguished from ordinary sulphuric acid by the fact that it fumes in the air. These fumes consist of the volatile sulphuric anhydride which it gives off, and which unites with the moisture of the air. The acid is usually prepared from green vitriol (ferrous sul- phate). This salt crystallizes with seven molecules of water, and has the composition: S0 2 -0 2 Fe + 7H a O. When heated in the air it easily loses six molecules of water, but the seventh only at a higher temperature, at which the salt itself begins to be decomposed. The salt, after being dried as thoroughly as possible, is heated in clay retorts provided with receivers of the same material." It is then decomposed into ferric oxide (called caput mortuum or col- cothar), which remains behind in the retorts, and sulphurous anhy- dride with sulphuric anhydride, which distil over, the latter con- densing in the cool receivers. The small quantity of water which the dried salt still contains suffices to convert the sulphuric anhy- m 2 1 64 Text-Book of Inorganic Chemistry. dride in the retort into liquid fuming sulphuric acid. The following equation represents the process : — 2S0 2 -0 2 Fe = Fe 2 3 + S0 2 + S0 3 . A larger quantity of acid is obtained, and without the production of sulphurous anhydride, when the mother-liquor from ferrous sul- phate, which contains the ferric salt, is evaporated and calcined. The mass so obtained consists principally of basic ferric sulphate of the composition : S 2 B Fe 2 = 2S0 3 , Fe 2 O s , which on further heat- ing in retorts breaks up into ferric oxide and sulphuric anhydride. As was stated on p. 162, fuming sulphuric acid when heated in a retort boils and gives off the sulphuric anhydride which it con- tains even below 100° As the quantity of anhydride becomes less and less the boiling point rises, until at last ordinary sulphuric acid remains in the retort. When cooled, fuming sulphuric acid deposits colourless crystals, which melt at 35 , and consist of disulphuric acid, the properties of which are described below (p. 166). Sulphurous anhydride can unite not only with an atom of oxygen but also with an equivalent quantity — i.e. two atoms — of chlorine. The compound so produced is Sidphicryl chloride : S0 2 C1 2 . Besides this body another chlorine compound is known, intermediate between sulphuryl chloride and sulphuric acid, and which may be considered as sulphuric acid with one atom of hydroxyl (OH) displaced by one atom of chlorine. This is Chlor- sulphonic acid : S0 2 \ ^j,. The latter is a monobasic acid, because it contains one atom of hydrogen displaceable by a metal, while the former is an indifferent substance, and cannot enter into com- bination with bases. SULPHURYL CHLORIDE. Composition : S0 2 C1 2 . This substance is a mobile colourless liquid, with a piercing odour. It fumes slightly in the air, has a specific gravity of 17, and boils at 70°. It may be obtained by direct union of chlorine and sulphurous anhydride when a mixture of the two gases in equal volumes is exposed to direct sunlight, or more readily, and in larger quantities, when chlorsulphonic acid is heated in closed Sulphury I Chloride. 165 tubes to 200 or 210 for about twelve hours. Two molecules of chlorsulphonic acid are then decomposed into one molecule of sulphuryl chloride and one of sulphuric acid : — 2S ° 2 {oH = S °2 C1 * + SO =|oH The liquid product, which is usually of a greenish yellow colour, is heated in a retort, and the portion passing over up to 110° again rectified on the water bath. The portion which then distils over at about 70° is nearly pure sulphuryl chloride. In order to separate minute quantities of chlorsulphonic acid and sulphuric anhydride, which it still contains, the distillate is poured into a separating funnel containing pieces of ice. The ice at once acts upon these impurities, while the sulphuryl chloride, which is scarcely attacked, sinks as a heavy oil to the bottom of the funnel. It is then freed from water with phosphoric anhydride, redistilled, and if the first portions of the distillate, which contain sulphurous acid and chlorine, are rejected, is so obtained perfectly pure. Sulphuryl chloride is easily decomposed by water into sulphuric and hydrochloric acids : — SOjCL, + 2H 2 = SOJ'q^ + 2HCI. CHLORSULPHONIC ACID. f CI Composition : S0 2 j qtt This substance is also a colourless liquid, with a specific gravity of 177, and boiling at 153°. It is at once decomposed by water, with a considerable evolution of heat, into sulphuric and hydro- chloric acids : — SO »{oH + H *° = SO =|oH + HCL Chlorsulphonic acid may be produced by the direct union of sulphuric anhydride and hydrochloric acid gas — S0 2 + HC1 = SO a j nn, but is best obtained by heating together molecular weights of sulphuric acid and phosphorus pentachloride, when hydrochloric acid and phosphoric oxychloride are also produced : — S0 2 J8h + PC1 * = SO »{oH + P0Q 3 + HC1 " 1 66 Text-Book of Inorganic Chemistry. It may then be separated from the more volatile phosphoric oxy- chloride (B.P. = 110°) by fractional distillation. Metallic salts of chlorsulphonic acid have not yet been prepared. NITROSULPHONIC ACID. Composition : S0 2 j o^r 2 This compound, which may be considered as sulphuric acid in which one atom of hydroxyl is displaced by the group of atoms NOj, separates out as colourless crystals when nitric peroxide is led into concentrated sulphuric acid or sulphurous anhydride into cold fuming nitric acid. The latter reaction is expressed by the equation : — S0 2 + NO..OH = S0 2 |q° 2 Nitrosulphonic acid is also produced by the action of sulphurous anhydride on a mixture of nitric oxide and oxygen in the presence of a small quantity of water, and occurs therefore in the sulphuric acid chambers when the quantity of steam is insufficient. The so-called white crystals to which we have previously referred (p. 158 note), consist of this compound, and their presence in the chambers is usually considered to indicate irregularities in the manufacture. Nitrosulphonic acid is dissolved by sulphuric acid, but decom- poses when gently heated or when brought into contact with water. In the latter case sulphuric acid and nitrous acid are produced, and the latter then decomposes into nitric oxide and nitric acid. When dissolved in sulphuric acid it requires a greater quantity of water to decompose it than when in the free state. DISULPHURIC ACID. Composition : H 2 S 2 7 = O j en 2 OH The latter formula indicates that two atoms of the sulphuric acid radical (sulphuryl = S0 2 ) are contained in a molecule of di- sulphuric acid, that they are united together by an atom of oxygen, Disulphuric A rid. 1 6j and that the other two unsatisfied bonds of the SO a are connected with two atoms of hydroxy 1 (HO), the hydrogen of which can be displaced by metals ; the acid is therefore dibasic. Disulphuric acid is deposited from fuming sulphuric acid at a low temperature as colourless and often large-sized crystals, which melt at 35 and fume in the air. It is decomposed when brought into contact with water into ordinary sulphuric acid : — JSO s .OH HQ _ 2SO JOH u (S(VOH + HaU ~ 25U MOH and when heated into sulphuric anhydride and sulphuric acid : — ]S0 2 -OH _ SQ SQ JOH u (S0 2 -OH " 3U > + ^(OH Its salts aie more stable than the free acid. The sodium salt is easily obtained by heating acid sodium sulphate as long as water is expelled. The following equation shows the reaction : — or> fONa „ (SCVONa „ ~ 2S0 2 | QH = 0- lS0 ^. 0Na + H 2 0. Acid sodium Sodium sulphate disulphate This salt dissolves in water unchanged. When strongly heated it breaks up into normal sodium sulphate and sulphuric anhydride : — „jS0 -ONa Qri jONa Qr . ;l SO;-ONa " bu MONa + SU THIOSULPHURIC ACID. Thiosulphuric acid is not known in the free state, but only in its salts. It is a dibasic acid, and is to be considered as sulphuric acid, in which one of the atoms of oxygen united with the hydrogen is displaced by dyad sulphur. The composition of the hypothetical p. tt, or H 3 S 3 3 . Before the relations of thio- sulphuric acid to sulphuric acid were known, the acid was compared with dithionic acid, which has the empirical formula, H„S 2 6 . Whence arose the former name for this acid : hyposulphurous acid, 1 based upon the old name for dithionic acid — hyposulphuric acid. 1 This name survives in the commercial designation for the sodium salt : ftyposutyhite o/soda.~ED, 1 68 Text-Book of Inorganic Chemistry. Of the compounds of thiosulphuric acid we are only acquainted with its metallic salts, some of which are stable substances, and of which the sodium salt is most easily obtained in the crystalline state. If we endeavour to set free the thiosulphuric acid from its salts by means of a stronger acid, we find that the acid can only exist for a very short time even in dilute solutions. The- liquid soon becomes turbid from deposited sulphur, and begins to smell of sulphurous anhydride. The acid decomposes in fact into sulphur, sulphurous anhydride, and water : — SoJq^ = S0 2 + S + H 2 0. This reaction may be used for the detection of the thiosulphates, especially as the separated sulphur always possesses a yellow colour. The best known salt is sodium thiosulphate, which with the potassium salt is distinguished by its stability. This impor- tant salt may be obtained in various ways. It is produced if an aqueous solution of neutral sodium sulphite is boiled with sulphur : — soi' ONa + s - so i 0Na bu |ONa + b - bU ^SNa or if sulphurous anhydride is led into a solution of sodium mono- sulphide. Normal sodium sulphite and sulphuretted hydrogen are then first produced : — SO, + Na 2 S + H 2 = SOJ°£J a + H,S, and the latter gas reacts at once on the excess of sulphurous anhy- dride, forming water and free sulphur : — S0„ + 2H 2 S = 2 H 2 + 3S. The greater part of this sulphur sinks to the bottom of the vessel, but a portion unites with the sodium sulphite to form sodium thio- sulphate. Sulphur therefore acts on the sulphites in the same way. as oxygen. Just as a solution of sodium sulphite when exposed to the air absorbs oxygen and becomes converted into sodium sulphate, so the same solution when warmed with powdered sulphur unites with it to form the corresponding thiosulphate. The composition of thiosulphuric acid is also interesting because its molecule con- tains two atoms of sulphur of different valency. The atom in the radical sulphuryl (SO a ) is hexad, while that uniting the atom of hydrogen or monad metal to this radical is dyad. Sodium thiosul- Thiosulphuric Acid. i6g phate, together with polysulphides, is produced when a solution of caustic soda is boiled with an excess of sulphur : — 6NaOH + I2S = SO.jOga + 2NaA + 3HjA Sodium pentasulphide Finally, some sulphides (e.g. that of calcium) are oxidized when exposed to the air and converted into thiosulphates. Calcium sulphide is largely produced as a bye-product in the manufacture of soda, and this alkali-waste when exposed to the air becomes partly converted into calcium thiosulphate, which, when digested with a solution of sodium carbonate or sulphate, yields insoluble calcium carbonate or sulphate, and a solution of sodium thiosulphate. The compounds of thiosulphuric acid with the heavy metals are mostly insoluble and easily decomposed. Sodium thiosulphate produces in a solution of silver nitrate a yellowish white precipitate of silver thiosulphate. This salt, however, rapidly darkens in colour, and at last becomes quite black, owing to the formation of silver sulphide and sulphuric acid : — S0 4sa| + H *° = A & S + SO »{oH In the same way the white lead salt blackens on boiling, and for the same reason. The extraordinary property which sodium thiosulphate, in com- mon with other soluble thiosulphates, possesses of dissolving chlo- ride, bromide, and iodide of silver, with which it forms soluble double salts, is of considerable importance. Use is made of this property chiefly in the production of photographs. After the sen- sitive plate, which is coated with a thin layer of these compounds, has been exposed in the camera, and after the iodide, &c, of silver has been decomposed in those parts on which the light has fallen, it becomes necessary in order to make the picture permanent to remove the iodide, &c, which has remained unchanged. This is effected by dipping the plate into a solution of sodium thiosulphate, which at once removes the unchanged compounds. Many tons of sodium thiosulphate are annually manufactured for this purpose. 170 Text-Book of Inorganic Chemistry. DITHIONIC ACID. This acid, formerly called hyposulphuric acid, is not known to us in the free state, but only in its salts. From the composition of these bodies we consider the acid to be : — H cn (SO„OH n s a ' u « ~JS0 2 0H the latter formula expressing that the molecule of the acid consists of two atoms of the radical : SO.-.OH, each of which contains an atom of displaceable hydrogen, thus making the acid dibasic. Manganous dithionate may be obtained by leading sulphurous anhydride into water containing finely divided manganese peroxide in suspension. The dry peroxide unites at once with sulphurous anhydride to form manganous sulphate : — S0 2 + Mn0 2 = S0 2 -0 2 Mn, but in the presence of water twice as much of the gas takes part in the reaction, and manganous dithionate is formed : — 2S0 2 + Mn0 2 =j^;°-Mn. At the same time a small quantity of the sulphate is always produced, and both salts remain in solution. The filtered liquid is then mixed with a slight excess of baryta water — i.e. until it be- comes faintly alkaline — which decomposes the manganous sul- phate, and precipitates manganous hydrate and barium sulphate. The liquid, which now contains barium dithionate and the excess of barium hydrate, is again filtered, and the slight excess of the latter converted into insoluble barium carbonate by passing a stream of carbonic acid, and at the same time gently warming it. From the clear solution, barium dithionate is deposited in large, colourless crystals on evaporating down, and allowing to cool. If lime water is employed instead of baryta water, calcium dithionate is obtained instead of the barium salt. By decomposing these compounds with the soluble sulphates of other metals, the dithionates of these metals may be easily obtained. All the salts of dithionic acid are soluble in water. The acid itself may be set free from barium dithionate by exactly precipitating the barium with sulphuric acid, or from the lead salt by precipitating the lead with sulphuretted hydrogen, and then filtering off the clear aqueous solution. The water may be Ditliionic Acid. 171 removed to some extent by evaporation in a vacuum over sul- phuric acid ; but if the concentration is driven too far, the acid decomposes into sulphuric acid and sulphurous anhydride : — fSUi ' The same decomposition, which may be recognized by the odour of sulphurous anhydride and by the precipitate produced on the addition of a drop of barium chloride, takes place when the dilute aqueous solution of the acid is boiled. But although easily decomposed, dithionic acid is decidedly more stable than thio- sulphuric acid. Dithionic acid is usually thought to be a dibasic acid, although this has not yet been definitely proved. At present only neutral salts of the acid are known, and until we can prepare the acid salts which dithionic acid in common with all polybasic acids ought to yield, the question of its basicity must remain unsettled. Perhaps these compounds are as easily decomposed as the free acid, and break up on evaporation into the acid sulphate and sulphurous anhydride. It is therefore always possible that the molecule of dithionic acid is only half as large as that generally accepted, that it is a monobasic acid and a compound of pentad sulphur (S0 2 OH). This possibility is not contradicted by the production and chemical behaviour of the acid, nor is it impossible to imagine that sulphur might exist in some compounds as a pentad element. . POLYTHIONIC ACIDS. Under this name may be included the three acids of sulphur which contain more than two atoms of sulphur. These are Trithio?iic Acid : H 2 S 3 O e , Tetrathionic Acid : H 2 S 4 O , and Penta- thionic Acid : H 2 S 6 6 . In these acids it is thought that the two groups of atoms which constitute dithionic acid are united together by one, two, and three atoms of sulphur respectively, as is expressed in the following formulas : — Trithionic acid = H 2 S 3 6 = S j cq 2 qh Tetrathionic acid = H 2 S 4 0,. = S 2 1 qq 2 qjj \J2 Text-Book of Inorganic Chemistry. Pentathionic acid = H 2 S 5 6 = S 3 ||q 2 q^ formulae which involve the supposition that the double as well as the triple atom of sulphur can play the part of a dyad radical just like a simple atom of the element. „ ■ u c n c fSO.,OH Trithionic Acid. hf„b 3 U e = ^Icjo'DH The potassium salt of this acid may be obtained by leading sulphurous anhydride into a solution of potassium thiosulphate, sulphur being at the same time separated : — 2S ° 2 iSK + 3S ° 2 = 2K 3 S 3°» + S - The acid, which may be set free from the potassium salt by the action of fluosilicic acid, very easily decomposes into sulphur, sul- phurous anhydride, and sulphuric acid. Tetrattaionic Acid: H 2 Sj0 6 = S 2 - L« ! nu The sodium salt is obtained by adding iodine to an aqueous solution of sodium thiosulphate : — 2S0 4sNa + ^ = Na A°<> + 2NaI. The barium salt may also be prepared in the same way. If this compound is decomposed with sulphuric acid, free tetrathionic acid is produced, which decomposes at a certain stage of concen- tration into sulphur, sulphurous anhydride, and sulphuric acid. Its salts, which are mostly soluble in water, are decomposed on evaporation into sulphur and the corresponding trithionates. Pentathionic Acid 1 : H 2 S 6 6 = S 3 1 go'^OH The free acid is produced together with large quantities of sulphur when sulphuretted hydrogen is led into an aqueous solution of sulphurous acid : — 5H,S + 5SO., = H 2 S 5 6 + 5S + 4HLO. The barium salt may also be prepared by adding sulphur di- 1 Recent experiments have shown that the existence of pentathionic acid is doubtful. Compounds of Sulphur and Chlorine. 173 chloride to water containing finely divided barium thiosulphate in suspension : — 2S0 2 j°Ba + SC1 2 = BaS 5 O d + BaCl,. The free acid, which may be concentrated in a vacuum to a given strength, soon decomposes into sulphur, sulphurous anhydride, and sulphuric acid in the same manner as the other two acids. COMPOUNDS OF SULPHUR AND CHLORINE. A chloride of sulphur corresponding to sulphuric anhydride {i.e. containing one atom of hexad sulphur united with six atoms of chlorine) has not yet been prepared. That consisting of one atom of sulphur and four of chlorine — sulphur tetrachloride : SC1 4 — is a very unstable compound. More stable but still easily decomposed is sulphur dichoride : SC1 2 ; while the most stable compound of all is Disulphur Dichloride : S ,C1 .. This compound is a transparent dark yellow liquid, fuming slightly in the air and possessing a powerful piercing odour. Its specific gravity is 17, and it boils at 138 . Disulphur dichloride is easily obtained by heating sulphur in a retort and passing dry chlorine gas over it. The sulphur then melts and the two elements unite with one another, the volatile chloride condensing in the cool receiver. The raw product is contaminated with sulphur di- chloride, from which it may be freed by fractional distillation, reserving that portion only which passes over at 138°. Disulphur dichloride is a good solvent for sulphur, of which it dissolves more than half its weight. It is decomposed by water into sulphur, sulphurous anhydride, and hydrochloric acid. Sulphur Dichloride : SC1„, is a heavy brown oil with similar properties to the preceding compound. It is obtained by satura- ting disulphur dichloride with chlorine at the ordinary temperature and then removing the excess of chlorine by passing a stream of dry carbonic anhydride. The compound begins to boil at 64°, due to evolution of free chlorine. It is then decomposed, and partially 174 Text-Book of Inorganic Chemistry. even at lower temperatures, into disulphur dichloride and chlorine. Water slowly decomposes it into sulphur, sulphurous anhydride, and hydrochloric acid. Sulphur dichloride forms double compounds with some other chlorides — e.g. antimony trichloride, arsenic tri- chloride, &c. Sulphur Tetrachloride : SCI ,. This is a mobile, pale brown liquid produced by saturating either of the other chlorides with chlorine gas at — 22° The com- pound is very unstable and breaks up even below o° into sulphur dichloride and free chlorine. Sulphur also unites chemically with bromine and iodine. The most interesting of these compounds, Sulphur hexiodide : SI,., is deposited by allowing a mixed solution of sulphur and iodine in carbon disulphide to gradually evaporate. It is a solid crystal- line substance, of dark grey colour, with a metallic lustre, is so easily decomposed that it gradually loses all its iodine when allowed to lie exposed to the air. SELENIUM. Chemical Symbol : Se. — Atomic Weight : 79. This element was discovered by Berzelius in 1817, and on account of its similarity to the previously discovered tellurium (from tellus, the earth), was named by him selenium (from a-eX^vrj, the moon). Berzelius found the new element in the deposit which had collected on the floor of a sulphuric acid chamber, fed with sulphurous anhydride derived from iron pyrites. Afterwards selenium was also discovered in the flue-dust deposited in the passages through which the acid vapours had to pass to reach the leaden chambers. It was later shown that the element exists combined with various metals — e.g. lead, silver, and mercury — in the mineral kingdom, and just as many samples of iron pyrites contain traces of iron selenide as well as iron sulphide, so too some varieties of native sulphur contain minute quantities of free selenium. Selenium is, however, a rare element, and always difficult to obtain in large quantities. Selenium. 175 Selenium which has been melted and allowed to cool is a dark brown amorphous substance with a lustre nearly metallic. It has a specific gravity of 4-2, melts at 217°, and boils at about 700°. Its vapour is of a dark yellow colour, and condenses either as bright red flowers of selenium or as a dark metallic-like mass. The specific gravity of its vapour at about 860° is 7-67, but diminishes nearly to that corresponding to the molecule Se 2 (5-58) at about 1400°. If selenium is reduced from a solution of selenious acid (e.g. by sulphurous anhydride) it is deposited as a. red amorphous powder. This modification is soluble in carbon disulphide, from which it is again separated on evaporation as dark-red, transparent crystals, isomorphous with monoclinic sulphur. A second crystalline modi- fication is also known which is insoluble in carbon disulphide, has a specific gravity of 4'5, and conducts an electric current feebly; it also possesses the remarkable property of conducting better under the influence of light than in the dark. This form of selenium is prepared by melting the element and keeping the temperature constant at 210% until it has become crystalline in structure. It then conducts twice as well in diffused daylight, and nearly ten times as well in direct sunlight as in the dark. If the intensity of the light is diminished, the conductivity rapidly decreases. The extraction of selenium from flue-dust containing it, or from its ores, is a tedious process. The various methods chiefly consist in oxidizing the selenium to selenious acid and then reducing it again by sulphurous anhydride. Selenium burns when heated in the air with a bright blue flame to form selenious anhydride. It requires, however, a higher temperature for its combustion, and burns more difficultly than sulphur. During combustion it gives off an intense and disgusting odour, which, however, is not peculiar to selenious anhydride, but is probably due to the simultaneous production of small quantities of a lower oxide. COMPOUNDS OF SELENIUM. The compounds of selenium are closely related in their che- mical properties to those of sulphur. Seleniuretted hydrogen corresponds to sulphuretted hydrogen, selenious anhydride to 176 Text-Book of Inorganic Chemistry. sulphurous anhydride, selenic acid to sulphuric acid, and the two chlorides of selenium to disulphur dichloride and sulphur tetra- chloride respectively. But, at the same time, many differences may be noted when similar compounds are compared together. Sulphurous anhydride, for example, is a gas, and does not unite with water to form the corresponding acid, while selenious anhy- dride is a solid crystalline body and easily forms the corresponding selenious acid. Sulphur and sulphurous anhydride are oxidized by nitric acid to sulphuric acid, while nitric acid has no action on selenious acid and only oxidizes selenium to the lower acid. Sul- phuric acid can be distilled unchanged, or at least with very slight decomposition ; selenic acid breaks up at about 280 into selenious acid and oxygen. Sulphur tetrachloride is a very unstable liquid, decomposing even below o° ; while selenium tetrachloride is a crystalline solid which can be heated without decomposition. It will thus be seen that considerable differences exist between sulphur and selenium with regard to their affinities for other bodies. It is clear from the facts mentioned above that selenium has a stronger affinity for chlorine and a weaker affinity for oxygen than sulphur, and this is confirmed by the fact that sulphur catches fire and burns when heated in oxygen or air much more easily than selenium. Still more remarkable is the affinity of selenium for hydrogen, and in this respect again it is at once distinguished from sulphur. Sulphur and hydrogen unite directly with one another only imper- fectly, and the resulting sulphuretted hydrogen is very easily decom- - posed by heat. Selenium, on the other hand, unites with hydrogen at about 500 , forming seleniuretted hydrogen, which is only de- composed into selenium and hydrogen at a high temperature. Selenium, with the similar elements tellurium and sulphur and the metals of the alkalies, are the only elements which unite directly both with hydrogen and oxygen. Those elements which possess a strong affinity for oxygen — e.g. phosphorus, arsenic, silicon, carbon, and most of the metals— cannot unite directly with hydrogen ; while the few other elements distinguished by their affinity for hydrogen — e.g. chlorineand palladium — cannot combine directly with oxygen. Compounds of Selenium. 177 Seleniuretted Hydrogen : H 2 Se. This compound is a colourless gas with an odour resembling sulphuretted hydrogen, and is exceedingly poisonous. Minute quantities in the air produce painful inflammation of the eyes and nasal passages, loss of the sense of smell, dryness in the throat, and a species of catarrh, often lasting for weeks together. Great care must, therefore, be taken in working with the gas, and in all operations during which it may be set free. Seleniuretted hydrogen is produced when dry hydrogen gas is led over selenium heated to the temperature at which it commences to volatilize. It maybe obtained in larger quantities by. acting, upon potassium selenide or ferrous selenide with hydrochloric acid. Like sulphuretted hydrogen, the gas is soluble in water, but in larger quantities. This solution reddens litmus paper, colours the skin red from reduced selenium, and becomes oxidized when ex- posed to the air, selenium being set free. Seleniuretted hydrogen precipitates many metals from solutions of their salts as insoluble selenides, just as sulphuretted hydrogen precipitates insoluble sulphides. Selenious Anhydride : SeO„. This substance is obtained when selenium is heated in a stream of dry oxygen. The selenium burns with a bright blue flame pro- ducing the anhydride, which is deposited on the cooler portions of the tube in long, quadrangular, white needles. It sublimes without first melting. Its vapour resembles that of chlorine in colour and possesses a piercing acid odour, the disgusting odour produced when selenium burns in oxygen or air being probably due to some lower oxide. When boiled with water it unites with this substance and forms the corresponding — Selenious Acid ; SeO -; qjt This dibasic acid is deposited on cooling the hot saturated aqueous solution in transparent crystalline prisms resembling those of potassium nitrate, and with a strong acid taste. When heated these crystals break up into the anhydride and water. The aqueous solution when made acid with a few drops of hydrochloric acid deposits selenium if warmed with sulphurous anhydride or other reducing agents. Particles of dust of organic origin produce the N 178 Text-Book of Inorganic Chemistry. same change, so that the lips of glass bottles containing aqueous selenious acid become coloured red from reduced selenium. With sulphuretted hydrogen selenious acid is decomposed into com- pounds of selenium and sulphur, water being also formed. Selenious acid can also be easily obtained by oxidizing seleniurh with nitric acid ; after the excess of nitric acid has been expelled it remains behind as a white crystalline mass. Selenic acid is never produced by this reaction. OH This acid, which corresponds to sulphuric acid in its composi- tion, and which it resembles in many respects, cannot be obtained by oxidizing selenium with nitric acid. It is, however, easily produced by leading chlorine into a strong aqueous solution of selenious acid, or by the oxidation of the latter substance by potassium dichromate. Potassium selenate may also be easily prepared by fusing together either selenium or a compound containing it with potassium nitrate. An aqueous solution of selenic acid may be concentrated until it attains a boiling-point of 265 and specific gravity of 2-6, but further concentration decom- poses it into oxygen and selenious acid. The concentrated acid is a colourless, oily liquid, which evolves heat when mixed with water, and is not affected by sulphurous anhydride or sul- phuretted hydrogen. It is, however, reduced to selenious acid when boiled with hydrochloric acid, chlorine being evolved. The selenates closely resemble the sulphates in their properties ; barium selenate, like barium sulphate, is insoluble in water and dilute acids. Selenium Bisulphide : SeS.,. This compound separates out as a yellow precipitate when sulphuretted hydrogen is led into an aqueous solution of selenious acid. The yellow precipitate gradually coagulates to a red mass, a colour which it retains on drying. If, on the other hand, seleniuretted hydrogen diluted with hydrogen is led into a saturated aqueous solution of sulphurous anhydride, the yellow precipitate consists chiefly of sulphur diselenide : SSe„. Selenium and sul- phur may be fused together in any proportion. Tellurium. 1 79 Chlorides of Selenium. Diselenium dichloride : Se 2 Cl 2 , is prepared in precisely the same manner as disulphur dichloride, which it closely resembles. Selenium tetrachloride : SeCl„ which is obtained by saturating the preceding compound with chlorine or by distilling selenious anhydride with phosphorus pentachloride, is a solid body, crystal- lizing in colourless cubes. When heated it sublimes without melting. It dissolves in water and then gradually decomposes into selenious and hydrochloric acids. TELLURIUM. Chemical Symbol : Te. — Atomic Weight : 128. This extremely rare element was discovered by Miiller v. Reichenstein in 1782 in gold ores obtained from Transylvania, but was only exactly studied by Klaproth in 1798, and later by Berzelius in 1832. It occurs native in the mineral kingdom, but more commonly combined with gold, silver, lead, and bismuth. The mineral tetradymite (bismuth telluride) contains about 50 per cent, of tellurium. The ores are found in Hungary and Transylvania, and recently in large quantities at several places in North America. Tellurium possesses many of the properties of a metal. It has a perfect metallic lustre, is nearly silver white in colour, crystalline in structure (crystallizing in rhombohedra), is brittle, and therefore easily reduced to powder. It conducts heat well, and electricity to some extent. Tellurium melts at about 500 , and can be sublimed at a higher temperature, producing a bright yellow vapour ; its specific gravity is 671. Carbon disulphide does not dissolve it, but it is soluble in warm concentrated sulphuric acid, forming a magnificent red solution, from which it is again precipitated unchanged as a grey powder when the solution is diluted with water. When strongly heated in the air it burns with a brown flame, bordered with green, forming tellurous anhydride. Hot nitric acid converts it into tellurous acid. Like selenium, tellurium unites directly with hydrogen when heated in a stream of the dry gas, and then produces telluretted hydrogen : H 2 Te. The same substance is also easily obtained by 180 Text-Book of Inorganic Cliemistry. acting on zinc telluride with hydrochloric acid. It is a. colourless gas with a disgusting odour, similar in its properties and reactions to sulphuretted hydrogen and seleniuretted hydrogen, though not so poisonous as the latter. The other compounds of tellurium — tellurous anhydride and acid: Te0 3 and TeC)]^, telluric acid: TeO a | OH , tellurium di- and tetrachloride : TeCl,, and TeCl,, tellurium disulphide : TeS. 2 — are closely allied to the corresponding compounds of selenium. The tellurates are, however, very different in their properties from the sulphates and selenates ; those of the heavy metals being mostly insoluble in water, while barium tellurate is only difficultly soluble in water and dissolves easily in dilute hydrochloric acid. Double salts of telluric acid corresponding to the alums have not been prepared. ELEMENTS OF THE NITROGEN GROUP. This group consists of the elements nitrogen, phosphorus, arsenic, and antimony. These elements are so different from one another both in their physical properties and their chemical beha- viour, as far as the strength of their affinity for other elements is concerned, that it is not at first sight evident why they are placed together in a group. Nitrogen does not catch fire when heated in the air, and it combines so feebly with other substances that it was thought formerly not to unite directly with any other body. Phos- phorus, on the other hand, is an easily fusible solid, distinguished by its strong attraction for oxygen, chlorine, and other elements. Arsenic and antimony possess the general physical properties of the metals ; they are heavy solids, with a metallic lustre. We shall, however, include these two last-named elements under the non- metals, because their chemical compounds with oxygen, hydrogen, and other elements are closely related to those of nitrogen and phosphorus. And this is also the reason why these four so different elements are included in one group. In most of their compounds the elements of the nitrogen group possess a triad or pentad atomicity. Nitrogen, 1 8 1 NITROGEN. Chemical Symbol: N. — Atomic Weight: 14. Nitrogen occurs free in atmospheric air, of which it is the chief constituent; 100 volumes of common air contain about 79 volumes of nitrogen and 21 volumes of oxygen. It also occurs in nature combined with hydrogen as ammonia, and with hydrogen and oxygen as nitric and nitrous acids. From these compounds nitrogen enters into the structure of plants and then of animals. It is a' constituent of albumen, casein, blood, and numerous other organic products. Urea, a substance present in large quantities in urine, contains a considerable proportion of nitrogen. Nitrogen is a colourless gas, without taste or smell, and rather lighter than common air. Its specific gravity is exactly 0-972, which corresponds to a molecular weight of 0-972 x 28-88 = 28-07, the molecule containing therefore two atoms (N 2 ). The gas is only slightly soluble in water. Nitrogen, like hydrogen and oxygen, was formerly called a permanent gas, as up to quite recently it had not been reduced to the liquid form. It has, however, been proved that nitrogen can be condensed to a liquid if the cold and pressure to which it is exposed are sufficiently intense. Nitrogen gas may easily be obtained from common air, by abstracting the oxygen which it contains. This may be done by burning a piece of phosphorus in a bell-jar full of air and standing over water. All the oxygen then unites with the phosphorus, pro- ducing white vapours of phosphoric anhydride, which afterwards dissolve in the water. As the air cools the water rises in the bell- jar to supply the place of the consumed oxygen, and when the air has become clear it consists almost entirely of nitrogen. A better method, especially when a stream of nitrogen is required, is to pass atmospheric air, which has been previously dried and freed from carbonic acid, over red-hot metallic copper. Finely divided copper, such as copper turnings, is heated to redness in a tube of hard glass while a slow current of the purified air is passed over it. The gas which then escapes at the other end of the tube is pure nitrogen, all the oxygen having combined with the copper to form black- copper oxide. Another method for the preparation of nitrogen consists in 1 82 Text-Book of Inorganic Chemistry. heating ammonium nitrite. This salt then breaks up into nitrogen gas and water : — NO-0(NH 4 ) = N 3 + 2 H 2 0. Finally, nitrogen is also set free when chlorine gas is led into strong ammonia. In this reaction, hydrochloric acid is produced, which unites with the excess of ammonia, forming ammonium chloride : — 8NH 3 + sCl 2 = N„ + 6NH 4 C1. If, however, the chlorine is passed until no more free ammonia is present, the gas then begins to decompose the ammonium chloride, and oily drops of nitrogen chloride, a highly explosive substance, are produced. This process cannot therefore be recommended for the preparation of nitrogen. Free nitrogen possesses so weak affinities for other elements that it can only be made to combine directly with a very few of them. With oxygen it only unites at the highest temperatures and in the presence of water vapour, then producing nitric acid with other lower oxides of nitrogen. Flashes of lightning passing through moist air cause the oxygen and nitrogen to unite and form nitric acid, which can be easily detected in rain water after a thunderstorm. The direct combination of nitrogen with hydrogen, chlorine, or sulphur, has not yet been effected, and only a very few of the metals combine directly with nitrogen when heated in the gas. Nitrogen neither supports ordinary combustion nor respiration. A candle immersed in the gas is at once extinguished. At the same time it is not an actual poison, as at once follows from the fact that we respire a mixture of nitrogen and oxygen in common air. An animal when brought into an atmosphere of nitrogen is simply suffocated from lack of oxygen. COMPOUNDS OF NITROGEN AND HYDROGEN. We have already stated that all attempts to unite nitrogen and hydrogen directly with one another have failed. They unite, how- ever, when the two elements are brought into contact with one another in the nascent state. The only compound of nitrogen and hydrogen which can exist Ammonia. 183 in the free state is ammonia. This substance consists of one atom of nitrogen united with three atoms of hydrogen, and in which, therefore, the nitrogen plays the part of a triad element. No compound analogous to hydrochloric acid — i.e. containing one atom of nitrogen united with one atom of hydrogen, nor one analogous to water, containing two atoms of hydrogen and one of nitrogen— appears to exist in the free state. Both are, however, known in many compounds as hypothetical radicals— i.e. unsatu- rated compounds of triad nitrogen. The former (NH) is called imidogen, and the latter (NH a ) amidogen. Finally, the compound of pentad nitrogen with four atoms of hydrogen (NH 4 ), called ammonium, is unknown in the free state, but exists in a large number of compounds very similar to those of the metal potassium. AMMONIA. Composition : NH,. This compound of nitrogen and hydrogen occurs free in nature and is produced during the putrefaction of nitrogenous organic matter in the absence of strong bases. It is formed, for example, in large quantities in putrefying urine, one constituent of which (urea) then combining with water and breaking up into carbonic anhydride and ammonia. Ammonia is a colourless gas with a strong, piercing odour, bringing tears to the eyes. It is extremely soluble in water, 1 volume at the ordinary temperature dissolving more than 800 volumes of the gas. This solution, the liquor ammonia or spirits of hartshorn of the shops, possesses the odour of the gas, and reacts powerfully alkaline — i.e. turns red litmus paper blue. The gas is considerably lighter than air, its exact specific gravity being 0-589. Its molecular weight is thus 0-589x28-88 = 17-01, and its molecule is represented by the formula NH 3 . Under a pressure of seven atmospheres at the ordinary tempera- ture, or at — 40 under the ordinary pressure, dry ammonia gas becomes converted into a colourless liquid which freezes to a crystalline solid at about — 80°. Ammonia is prepared from the ammonium salts (compounds formed by the union of ammonia with acids), and generally from 1 84 Text-Book of Inorganic Chemistry. ammonium chloride or sal-ammoniac. This salt, which can be obtained from ammonia and hydrochloric acid, is decomposed again into these constituents when intimately mixed with strong bases, such as caustic potash, caustic soda or quick-lime, and gently heated. Even when dry ammonium chloride and dry quick- lime are mixed together, ammonia begins to be given off without the mass becoming liquid, whence it follows that the old chemical rule, only liquids act upon one another, is not to be taken literally. The reaction goes on according to the equation : — 2NH,C1 + CaO = CaCl 2 + H 2 + 2NH3. Calcium chloride and water are thus simultaneously produced, and a, portion of the latter passes over with the gas on heating. In practice, ammonium chloride and quick- orslaked-lime, both finely powdered, are intimately mixed together, and the mixture brought into a glass flask or an iron vessel if large quantities are employed. Heat is then applied and the ammonia dried by passing it through a tube containing pieces of quick-lime. It is then either collected over mercury or by upward displacement. The sub- stances generally employed to dry gases — sulphuric acid or calcium chloride — cannot be used for ammonia, as both these bodies absorb large quantities of the gas. If an aqueous solution of the gas is required, it is of course unnecessary to dry it. It is then sufficient to separate the solid particles which may be carried over with the gas. This is best done by passing the gas through a small quantity of concentrated aqueous ammonia contained in a wash-bottle. Large quantities of heat are evolved when ammonia is absorbed by water, principally due to the latent heat set free when the substance is changed from the gaseous to the liquid form. Care must therefore be taken to cool the vessel containing the water during the absorption of the gas. The production of ammonia from nitric acid by the action of nascent hydrogen is interesting theoretically. This reduction may be produced in several ways ; for example, by the action of dilute nitric acidj! on zinc. No evolution of hydrogen is noticed, as when zinc is dissolved in sulphuric acid, but the whole of the hydrogen is used to reduce the nitric acid. The ammonia which is formed unites with the excess of acid, producing ammonium nitrate ; the zinc is also converted into its nitrate, and water is set free : — Ammonia. 185 2NCVOH + 4 H 2 =. N0 2 -0(NH 4 ) + 3H 2 0. On adding an excess of caustic soda to the solution, the ammonia is again liberated, and may be easily recognized by its odour. In order' to produce small quantities of liquid ammonia, advan- tage is taken of the fact that many metallic salts— e.g. calcium chloride, silver chloride — absorb ammonia in large quantities and liberate the gas again on warming. Dry silver chloride saturated with dry ammonia is placed in the closed end, a, of a thick walled bent glass tube, and the open end, c, is then melted together (fig. 47). If now the limb c is placed in a freezing mixture, and a gently warmed, gaseous ammonia is liberated in large quantities from the silver chloride and is con- densed to a colourless liquid in the cold part F ' s ' 47 " of the tube. When the tube cools again, the silver chloride reabsorbs the whole of the ammonia, and the experiment may therefore be repeated several times with the same tube. Although ammonia contains a large quantity of hydrogen, it does not catch fire in the air, but in oxygen gas it burns with a pale yellowish flame, forming water and free nitrogen. At a high temperature it is decomposed into its constituents, hydrogen and nitrogen. If the gas is led through a porcelain tube filled with broken fragments of the same and heated to bright redness, the volume of the gas becomes doubled : — 2NH 3 = N„ + 3H„. 4 VOls. 2 VOls. 6 VOlS. The same decomposition occurs when electric sparks are allowed to pass through the gas for a long time. Ammonia combines readily with acids and produces real salts, which closely resemble the corresponding compounds of potassium. These are the ammonium salts and contain the monad radical ammonium (NHJ, which plays the part of a metal. When, for example, ammonia unites with hydrochloric acid : — NH 3 + HCl" = (NH 4 )CI, we consider that the hydrogen of the acid first unites with the nitrogen, and that the ammonium so produced then forms with the chlorine the haloid salt. In a similar manner ammonia also unites with the oxy-acids, but in these compounds the ammonium is united to the acid radical by an atom of oxygen, thus : — i86 Text-Book of Inorganic Chemistry. N0 3 -OH + NH a = N0 2 -0(NH 4 ) Nitric acid Ammonium nitrate SO l 0H + 2NH -SO l°( NH <) Sulphuric Ammonium acid sulphate SO,{°g + NH 3 =S0 2 )g ( H NHj) _ Sulphuric Acid ammonium acid sulphate Ammonia also combines with the acid anhydrides, in which the hydrogen necessary to convert it into ammonium is absent. Two molecules of ammonia are then decomposed into one molecule of ammonium and one of amidogen : — 2NH3 = NH 4 + NH„ which unite with the anhydride {e.g. C0 2 or SO.) to form ammonium compounds of monobasic amido-acids : — C0 2 + 2NH 3 - COJ^H^ Ammonium carbamate In all cases when gaseous ammonia comes into contact with a gaseous acid, heavy white fumes are produced of the ammonium salt. This affords a ready test for small quantities of free am- monia. If a glass rod dipped in concentrated hydrochloric acid is brought near the surface of a liquid containing free ammonia, white fumes of ammonium chloride are at once produced. If the am- monia is combined with an acid, as an ammonium salt, it must first be set free by warming with caustic soda or potash. Chlorine at once decomposes ammonia, forming hydrochloric acid, which then combines with the excess of ammonia to form ammonium chloride (p. 182). HYDROXYLAMINE— (OXYAMMONIA). Composition : NH..-OH. In hydrogen peroxide we have already become acquainted with a monad radical — hydroxyl — having the composition HO (p. 90). We meet with this hydroxyl in a large number of compounds, par- ticularly in organic chemistry, in which it displaces other monad Hydroxy lamine — (Oxyammonia). 1 87 elements or radicals, especially hydrogen. And we may look upon hydrogen peroxide as water in which one atom of hydrogen is dis- placed by hydroxyl, thus : O ] ^ H In precisely the same manner, hydroxyl can also displace one of the atoms of hydrogen in ammonia, and so form the substance hydroxylamine, or oxyammonia. Oxyammonia has not yet been prepared in the pure state ; it is only known in aqueous solution and in combination with acids. It possesses similar properties to those of ammonia, unites with acids producing the oxyammonium salts, but is less volatile. Its aqueous solution is colourless and reacts strongly alkaline. Oxyammonia is easily produced by the reduction of nitric oxide or higher oxides of nitrogen with nascent hydrogen. For this purpose nitric oxide is led into a mixture of granulated tin and hydrochloric acid, to which a few drops of platinic chloride have been added to accelerate the evolution of hydrogen. The acid liquid then contains stannous chloride together with oxyammonium chloride ; it is freed from tin by passing a stream of sulphuretted hydrogen through it, and the filtered liquid evaporated to dryness. The dry mass is then extracted with absolute alcohol, which dis- solves the whole of the oxyammonium chloride, together with a small quantity of ammonium chloride (easily removed by a few drops of platinic chloride). The clear alcoholic liquid on evaporation then yields colourless crystals of oxyammonium chloride : NH 3 (OH)Cl. From the chloride the sulphate : S0 2 (ONH 3 OHj s , is easily prepared by adding to it the requisite quantity of dilute sulphuric acid and evaporating off the hydrochloric acid thus set free. Free oxyammonia may then be obtained by adding baryta water to the aqueous solution of the sulphate as long as a turbidity of barium sulphate is produced. The filtered alkaline solution then contains free oxyammonia. If this solution is distilled a portion of the oxyammonia passes over with the water vapour, but the greater part is decomposed into ammonia, nitrogen, and water : — 3NH 2 -OH = NH, + N 2 + 3H 2 0. The aqueous solution of oxyammonia powerfully reduces many compounds. Silver and mercury are precipitated from the solutions of their salts in the metallic state, and with copper sulphate it gives an orange-coloured precipitate of cuprous oxide. 1 88 Text-Book of Inorganic Chemistry. Oxyammonia is also produced when some metals, especially zinc, are acted upon by sulphuric acid in the presence of nitric acid. The nascent hydrogen which is then set free reduces a part of the nitric acid to ammonia and oxyammonia. If zinc is used the solution contains the nitrates of zinc, ammonium, and oxyammonium, and on the addition of excess of caustic potash and a few drops of copper sulphate gives an orange- coloured precipitate of cuprous oxide. The oxyammonium salts, like those of ammonium, are all solu- ble in water, and can easily be obtained in the crystalline form. AMIDOGEN : NH„; IMIDOGEN : NH. 2 J These are two unsaturated compounds of triad nitrogen, the first being a monad, and the second a dyad radical. Both occur principally in organic compounds, but amidogen is also known in some inorganic compounds ; for example, in the amido-acids. AMMONIUM: NH, This unsaturated compound of pentad nitrogen has not yet been prepared in the free state. It is the radical of the ammonium salts, but all attempts to separate it, or rather its molecule : (NH,).,, have resulted in a mixture of ammonia and hydrogen. It is very probable that, should this body ever be obtained in the free state, it would be found to possess the properties of a metal. Not only have its compounds the greatest similarity with those of potassium and sodium, but it can also unite in the nascent state with mercury to form an amalgam — a property peculiar to the metals. Ammonium amalgam, with metallic lustre and other re- markable properties, is easily obtained by adding sodium amalgam to a warm, moderatelyconcentrated solution of ammonium chloride. The sodium amalgam immediately swells up to a large extent, at the same time retaining its metallic lustre, and the mass becomes so light that it floats on the surface of the liquid, notwithstanding the mercury which it contains. The reaction simply consists in the substitution of sodium for ammonium in the ammonium chloride, forming sodium chloride, and setting the ammonium free. The amalgam is very unstable, breaking up immediately into ammonia, hydrogen, and mercury, Compounds of Nitrogen and Oxygen. 1 89 the first-named substance being at once recognized by its charac- teristic odour. The low specific gravity of the ammonium amalgam makes it probable that if the ammonium itself were separated, it would be a light liquid or a gas. COMPOUNDS OF NITROGEN AND OXYGEN. Nitrogen and oxygen unite in five different proportions, and produce as many different compounds. These compounds have the following names and symbols. Two of them (nitrous anhy- dride and nitric anhydride), unite with water and form the corre- sponding acids : — . N 2 . NO Nitrous oxide . Nitric oxide Nitrous anhydride Nitric peroxide Nitric anhydride N a 3 Nitrous acid NO-OH n 2 o; N 2 5 Nitric acid NO„-OH NITRIC ACID. Composition: HN0 3 = NO, -OH. Nitric acid occurs free in the atmosphere, especially after thunderstorms, and is produced directly from its elements by the high temperature of the flashes of lightning. Its salts are also produced in nature, when nitrogenous organic substances putrefy in contact with strong bases {e.g. potash, soda, lime). Nitric acid is a colourless, strongly acid liquid, fuming in the air. It has a specific gravity of 1-530, boils at 86°, and becomes solid at - 50° ; it mixes with water, with evolution of heat, in all pro- portions. The acid is not very stable, and is even partially decom- posed by light into oxygen and lower oxides of nitrogen, which impart a yellow tint to it. It destroys the skin and other animal substances and colours them yellow, and, in consequence of the ease with which it breaks up into oxygen and the lower oxides of nitrogen, is an excellent oxidizing agent, especially for the metals. Nitric acid, as we have previously remarked (p. 182), is produced directly from its elements when electric sparks pass through moist air, or when a mixture of hydrogen and excess of oxygen, with 190 Text-Book of Inorganic Chemistry. traces of nitrogen, is exploded, or when hydrogen burns in oxygen mixed with a small quantity of nitrogen. The acid is always obtained from one of its salts — potassium or sodium nitrate — by decomposing it with sulphuric acid. On the small scale pure nitre (potassium nitrate) is powdered and intro- duced into a retort furnished with a cooled receiver. It is then mixed with concentrated sulphuric acid and gently heated, when the whole melts to a homogeneous liquid. As the temperature rises, the nearly colourless acid distils over, and is condensed in the receiver. It is best to employ one molecule of nitre (101), to one molecule of sulphuric acid (98) — i.e. about equal parts by weight. The reaction is then represented by the following equa- tion : — NCyOK ■»• S0 2 |q^ = NO a -OH + S0 2 |q^ Potassium Acid potassium nitrate sulphate If we employ twice as much nitre so as to produce the normal sulphate, the first stage of the reaction goes on as before, but the acid potassium sulphate which remains behind is then mixed with one-half of the nitre employed. At a higher temperature these two substances act upon one another, forming another molecule of nitric acid and normal potassium sulphate : — S0 4oK + NC V 0K = S0 2lOK + NC V 0H > Acid potassium Normal potassium sulphate sulphate The temperature at which this takes place is so high that the nitric acid undergoes a partial decomposition into oxygen, nitric peroxide, and water. The oxygen passes off, but the nitric peroxide, a dark red gas, is mostly absorbed by the nitric acid. The acid then becomes of a reddish colour, and gives off brown vapours. This acid is known as fuming nitric acid. In order to obtain a pure acid, it is best to take a small excess of sulphuric acid, and then redistil the acid with an equal volume of concentrated sulphuric acid. The strong attraction of sulphuric acid for water then frees the nitric acid of this impurity. Commercial nitric acid is now often prepared from Chili salt- petre (NaN0 3 ), this salt being much cheaper than ordinary nitre, and yielding weight for weight a greater quantity of nitric acid. It is usually coloured yellow, and is often contaminated with chlorine, Nitric Acid. 191 iodine, sodium sulphate, and iron. These impurities may be readily separated by repeated distillation, neglecting the first and last portions which come over. It always contains water, the quantity of which may be determined by means of an hydrometer. The greater the quantity of water the lower the specific gravity. The specific gravity of this commerical acid is often as low as I "33, and then contains only about 50 per cent, of the pure acid. The specific gravity of the pure acid is, as we have stated above, l'530. An acid which has a specific gravity of i"424, and which boils at about 120°, contains nearly 70 per cent, of pure acid. The dilute acid, of specific gravity 1-2, which is employed in the laboratory for general purposes, contains only 32 per cent, of the pure acid. If concentrated nitric acid containing more than 70 per cent, of the pure acid is boiled under the ordinary atmospheric pressure, the portion which distils over contains less water than that in the retort until the boiling-point rises to about 120°, when the compo- sition of the distillate and of the residue in the retort remains nearly constant. This acid of constant boiling-point has a specific gravity of 1-42, and contains about 70 per cent, of pure acid. Further, if dilute nitric acid containing less acid than 70 per cent, is distilled, the portion which passes over is at first more dilute than that in the retort, and rises in strength until it contains about 70 per cent, of pure acid, when the composition again remains nearly constant. It must not, however, be supposed that this particular mixture is a chemical compound of water and nitric acid. That it is not so is proved by the fact that if the pressure is changed the composition of the distillate also changes. Nitric acid is a solvent for many substances which are insoluble in water — particularly for the metals. Among the few metals which withstand its action are gold and platinum. Silver is at once attacked by it and converted into the soluble nitrate, a portion of the nitric acid being reduced to nitric oxide, producing red fumes of nitric peroxide in the air. The reaction goes on according to the following equation, which also represents the oxidation of many other metals by nitric acid : — 4 NO,.OH + 3Ag = 3N0 2 -OAg + 2 H 2 + NO. In consequence of this property of nitric acid to dissolve silver but not gold, nitric acid has been long employed to separate the two metals from one another. 192 Text-Book of Inorganic Chemistry. Nitric acid is a monobasic acid, as it only contains one atom of displaceable hydrogen ; it, therefore, only produces normal salts. The nitrates are all soluble in water, and we cannot therefore recognize them, as we can the sulphates, by the production of one which is insoluble in water. All nitrates are decomposed by strong sulphuric acid yielding nitric acid ; and if a solution of a nitrate is mixed with sulphuric acid and copper turnings and then warmed, the nitric acid which is set free acts on the copper and produces nitric oxide, which at once gives rise to brown fumes of nitric peroxide and nitrous anhydride on coming into contact with the oxygen of the air. Another delicate test for a nitrate is mentioned under nitric oxide. Nitric Anhydride : N 2 5 . This substance, as its name implies, is nitric acid minus the ele- ments of water, and may be considered as two molecules of nitric peroxide united by an atom of oxygen, thus mo 5 [0. It is obtained in colourless lustrous crystals belonging to the rhombic system when finely-powdered dry silver nitrate is dropped into a flask containing dry chlorine. The two substances then react upon one another, producing nitric anhydride, silver chloride, and oxygen. On opening the flask afterwards the oxygen and excess of chlorine are got rid of by blowing in dry air, and the anhydride, which melts at 30 , is then easily separated from the silver chloride by gentle heating. If heated slightly above its boiling point (45°), it decomposes with explosion into nitric peroxide and oxygen, and the decomposition also goes on slowly in sealed tubes. It readily dissolves in water, evolving heat, and producing nitric acid : — ^q4o + H 2 = 2N0,-0H. NITRIC OXIDE. Com-bosiiion : NO. This oxide of nitrogen is a colourless gas, slightly soluble in water, and with a specific gravity of ro39. This density corresponds to a molecular weight of ro39 x 28'88 == 30, whence its composition. Nitric Oxide. 193 Nitric oxide possesses none of the acid properties of nitric acid. It is distinguished from all other gaseous bodies by its remarkable property of combining directly with oxygen at the ordinary temperature to form a dark red gas. From this it follows that nitric oxide may at once be recognized by the red colour produced when it comes into contact with common air. This red gas is either nitrous anhydride or nitric peroxide, according to the quantity of oxygen present. Nitric oxide is obtained from nitric acid by abstracting a portion of its oxygen by copper or some other similar metal. Small pieces of copper are placed in a flask and then dilute nitric acid, of specific gravity f23, poured over them. The gas soon begins to be given off, and the mixture evolves large quantities of heat, and if the flask is not cooled from time to time by dipping it into cold water, the reaction may become so violent that the liquid froths over. The nitric oxide which is first produced unites at once with the oxygen of the air in the flask and forms red fumes of nitrous anhydride, but as this compound is absorbed by water, the gas when collected in a jar over water soon becomes colour- less. After a time, when the whole of the air has been expelled from the flask, pure colourless nitric oxide is given off. The reaction may be considered as taking place in two stages. In the first the copper is oxidized to copper oxide, and the nitric acid reduced to nitric oxide and water ; while, in the second, the copper oxide unites with the excess of nitric acid to form copper nitrate. The entire reaction is expressed in the following equation : — 3 Cu + 8NCVOH = 3no^o} Cu + 2N0 + 4H =°- If the mixture of copper and nitric acid becomes too hot the nitric oxide may easily become contaminated with nitrous oxide. Instead of copper we may also use mercury or silver for the preparation of nitric oxide ; but all metals cannot be employed, some give off nitrous anhydride or nitric peroxide. In nitric oxide the constituents are firmly united together, much more firmly than in nitrous oxide, which contains twice as much oxygen. A glowing chip of wood, faintly glowing charcoal, or burning sulphur is at once extinguished in the gas, but strongly glowing charcoal continues to burn brilliantly. Phosphorus can be melted in the gas without catching fire, but ignited phos- phorus goes on burning. In all cases when substances burn in O 194 Text-Book of Inorganic Chemistry. nitric oxide the temperature of combustion must be high enough to decompose the gas into its constituents. Among the more easily combustible substances, the volatile liquid carbon disulphide is distinguished by the readiness with which it burns when mixed with nitric oxide. A small quantity of the disulphide is poured into a jar of the gas closed with a glass plate, the jar slightly agitated so as to volatilize the disulphide, and the mixture then ignited. The carbon disulphide then burns with so brilliant a flame that it can scarcely be borne by the naked eye. Strong as well as dilute nitric acid dissolves considerable quantities of nitric oxide and becomes coloured brown, yellow, green, or blue according to the concentration of the acid. A solution in the concentrated acid (specific gravity 1 35) is brown, in an acid of specific gravity 1-25 is blue. On heating the acid the nitric oxide is again given off. Nitric oxide is also absorbed by concentrated solutions of ferrous sulphate, yielding a solution of a dark brown colour. This property of ferrous sulphate is used to detect small quantities of nitric acid (free or combined) contained in a solution. To the solution which is to be tested for nitric acid or a nitrate a little concentrated sulphuric acid is added, and then a crystal of ferrous sulphate. The sulphuric acid serves to set the weaker nitric acid free, and this is reduced to nitric oxide by the ferrous sulphate ; the gas then unites with the excess of ferrous sulphate to form a dark- coloured mantle around the crystal. NITROUS ANHYDRIDE AND NITRIC PEROXIDE. These two oxides of nitrogen, which are intermediate between nitric acid and nitric oxide in composition, are so similar in the gaseous state, that one can scarcely be distinguished from the other, especially as they are generally produced together. Both are brown gases, with a strong unpleasant odour, and both are produced by the union of nitric oxide with oxygen — when the nitric oxide is in excess, nitrous anhydride is formed, and with excess of oxygen nitric peroxide. Both gases can be easily condensed to the liquid state by lowering the temperature ; liquid nitrous anhydride is blue, while liquid nitric peroxide is brown. A mixture of nitrous anhydride and nitric peroxide is easily Nitrous Anhydride and Nitric Peroxide. 195 obtained by heating arsenious anhydride, white arsenic (in lumps) with nitric acid of specific gravity 1-33. When the reaction has once been started it goes on without further heating until finished, oxygen being abstracted from the nitric acid while the arsenious anhydride is oxidized to arsenic acid. The red gases are led through a tube surrounded by a freezing mixture and are there condensed to a deep blue liquid. This liquid begins to boil at + 2°, and the boiling point gradually rises to 22 . It is a mixture of the volatile nitrous anhydride, and the less volatile nitric peroxide. From this mixture the two substances may be easily obtained in the pure state. Nitrous Anhydride : N,0 3 . This substance is obtained from the blue liquid mentioned above by passing its vapour mixed with nitric oxide through a red- hot glass tube. The nitric peroxide which it contains then unites directly with the nitric oxide to form nitrous anhydride : NO, + NO = N,0 3 , which can be condensed in a freezing mixture to an indigo coloured liquid. This liquid does not solidify at - 30 and begins to boil below o°, while at higher temperatures it easily decomposes into nitric peroxide and nitric oxide. Nitrous anhy- dride dissolves in a small quantity of ice-cold water and forms Nitrous acid : NO -OH, according to the equation : N 2 3 + H o = 2NO-OH. It easily decomposes when slightly warmed into nitric acid, nitric oxide, and water : — 3NO-OH = NO,OH + 2NO + H.,0. More stable than the acid itself are its compounds — the nitrites. Potassium nitrite (NO -OK), which contains one atom of oxygen less than potassium nitrate (N0 3 -OK), is easily produced by fusing the latter substance until the oxygen which is given off begins to be mixed with nitrogen. The potassium nitrite can then be extracted from the powdered mass by means of alcohol, in which it is soluble. If potassium nitrite, or any other salt of nitrous acid, is treated with dilute sulphuric acid, it is at once decomposed ; the nitrous acid which is thus set free imparts a transient blue colour to the liquid, and soon breaks up into water and red vapours of the lower oxides of nitrogen. An aqueous solution of potassium nitrite when mixed with silver nitrate produces a white precipitate of silver nitrite, thus :— o 2 196 Text-Book of Inorganic Chemistry. NO-OK + N0 2 -OAg = NO-OAg + N0 2 -OK. Ammonium nitrite (NO-ONHJ, which, it will be remembered, breaks up into water and nitrogen when heated (p. 182), is also produced in minute quantities when water rapidly evaporates, a portion of the water then uniting with the nitrogen of the atmo- sphere : — ■ 2H 2 + N 2 = NO-ONH 4 . The same compound is also produced when ammonia is oxi- dized by a mixture of ozone and oxygen : — 2NH3 + 3 = NO-ONH, + H 2 0, and the salt is then further oxidized to ammonium nitrate. Wltric Peroxide : (NO,), = N 2 4 . This compound of nitrogen and oxygen possesses all the properties of a compound radical, and resembles chlorine in many respects. A more suitable name for it than nitric peroxide would be nitryl, corresponding to hydroxyl, sulphuryl, &c. Nitric peroxide is best obtained from the crude blue liquid re- ferred to on the preceding page. When this liquid is kept in a freezing mixture and a current of oxygen led through it the nitrous anhydride which it contains becomes oxidized, and pure nitric peroxide is obtained. Nitric peroxide may also be prepared by leading a mixture of nitric oxide and excess of oxygen into a cooled tube, or, finally, by carefully heating dried lead nitrate in a retort, when the salt suffers the following decomposition : — NCVo} Pb = < N0 ^ + O + PbO. Nitric peroxide is a pale brown liquid, of specific gravity 1-45 ; it boils at 22 and then forms a brown gas, which becomes of a darker colour when heated. It freezes at — 20°to a colourless crystalline mass, which again melts at about — 12° The vapour density of nitric peroxide at low temperatures has been found to be 2-65, at high temperatures J '58. The. former number approaches the density required for the formula N,0 4 , which would be (3 * T4 l + oi 4 X l6) = 92 „ = 3-183, while that of the latter is about 28-88 28-88 one-half of this, or corresponds to the molecule N0 2 . From this we conclude that at low temperatures the molecule of nitric peroxide is composed of twice N0 2 or N0 2 -NO x , which, when Nitrous Anhydride and Nitric Peroxide. 197 heated to 100° — 150 , breaks up into separate molecules of the com- position NO,. The same change probably takes place when many elements — e.g. oxygen or chlorine — are raised to a very high tempe- rature. The atoms, which at ordinary temperatures are united to form molecules, then divide into separate atoms and so acquire a far stronger power of affinity than in the ordinary molecular state. In the same way nitric peroxide unites much more readily with other bodies at loo°than at the ordinary temperature ; for example, it can then combine directly with chlorine, bromine, cyanogen, &c. Concentrated, colourless nitric acid dissolves large quantities of nitric peroxide, producing the red or fuming nitric acid. This acid is a much more powerful oxidizing agent than ordinary nitric acid, because the nitric peroxide which it contains parts with a portion of its oxygen far more readily than nitric acid itself. In contact with cold water nitric peroxide is decomposed into nitric and nitrous acids : — (NO,), + H a O = NCyOH + NO-OH; the nitrous acid gradually decomposing into nitric acid, nitric oxide, and water (p. 195). Chlorine, which at low temperatures is without action on nitric peroxide, combines with it when a mixture of the two gases is led through a heated tube, producing nitryl chloride ; NOX1, a yellow liquid boiling at + 5 , which is decomposed by water into nitric and hydrochloric acids. The same compound is also produced by leading a stream of chlorine over gently heated silver nitrate : — NCvOAg + Cl 3 = NOX1 + AgCl + O. If liquid sulphurous acid and nitric peroxide are heated together in sealed tubes, or if nitric oxide is allowed to act upon sulphuric anhydride, a crystalline compound is produced, of the composition : S.,O s -2N0 2 , or O 1 cn a NO 2 ' anc ^ wn ^ cn may be considered either as the dinitryl of disulphuric acid, or as the anhydride of the next following substance, nitrosulphonic acid. This compound melts at 217 , and is decomposed by water into sulphuric acid, nitric oxide, and nitric peroxide. Concentrated sulphuric acid' absorbs nitric peroxide and pro- duces white crystalline nitrosulphonic acid: SOj-I^tt 2 (p. 166), in 198 Text-Book of Inorganic Chemistry. which one atom of hydroxyl is displaced by one of nitric peroxide, nitric acid being probably formed at the same time. This com- pound is also decomposed by water into sulphuric acid, nitric oxide, and nitric peroxide. Aqua Regia. — This liquid, which owes its name to its property of dissolving gold, the king of metals, as well as platinum, is a mixture of three parts of hydrochloric acid and one part of nitric acid. The two acids act upon one another at the ordinary tem- perature, the mixture becoming yellow and giving off minute bubbles of gas smelling like chlorine. This gas, which is evolved in quantity on gently warming, is a mixture of chlorine and .nilrosyl chloride: NOC1. In the nascent state the latter sub- stance gives up its chlorine to the gold or platinum, produc- ing soluble chlorides of these metals. The reaction by which it is produced is shown by the following equation : — NO, -OH + 3HCI = NOC1 + 2H,0 + CI,. NITROUS OXIDE. Composition : N 2 0. This substance is a colourless, odourless gas, with a somewhat sweet taste, and a density of i'52, which corresponds to a molecular weight of 43'9- Under strong pressure (thirty atmospheres at o°) it is condensed to a colourless liquid, which i6 lighter than water, and with which water does not mix. The liquid boils at about — 92 , and when allowed to evaporate, rapidly solidifies to a white crystalline solid. Nitrous oxide is only slightly soluble in water, and can, therefore, be collected over this liquid. The gas is easily prepared by heating crystallized ammonium nitrate in a retort. The salt then melts and completely decom- poses, with much frothing and apparent boiling, into nitrous oxide and water. Thus : — N0 2 -ONH, = N,,0 + 2H 2 0. It is not advisable to decompose the whole of the salt employed, as overheating may cause the nitrous oxide to decompose into its constituents with an explosion. This method quickly yields large quantities of pure nitrous oxide, if care is taken that the ammonium Nitrous Oxide. 199 nitrate contains no ammonium chloride, otherwise the gas becomes contaminated with chlorine compounds. Nitrous oxide is also produced when zinc or one of a few other metals is dissolved in very dilute nitric acid. The gas obtained in this way is, however, never pure, but always contains more or less nitric oxide. The most remarkable property of nitrous oxide is the readiness with which it allows combustible bodies to burn in it. Combustion takes place in the gas almost as energetically as in pure oxygen, much more so than in nitric oxide, although this gas contains twice as much oxygen. From this we conclude that the nitrogen and oxygen are much more loosely combined in nitrous oxide than in nitric oxide. In fact, nitrous oxide is decomposed into its con- stituents even when passed through a red-hot tube. A glowing chip of wood catches fire in nitrous oxide almost as easily as in oxygen. Phosphorus and charcoal burn brilliantly, and sulphur, if strongly ignited, continues to burn with a bright flame. Feebly burning sulphur is extinguished in the gas, because the temperature is not high enough to decompose it into its constituents. That nitrous oxide, although a chemical compound of nitrogen and oxygen, sup- ports combustion much better than the mechanical mixture of the same two substances which make up common air is partly because the compound contains thirty-three per cent, by volume of oxygen, while common air has scarcely more than twenty per cent, by volume, and is partly due to the fact that when nitrous oxide is de- composed a considerable quantity of heat is set free. Nitrous oxide is chemically an indifferent body, and does not enter into combination with any substance. Its peculiar pro- perty of producing intoxication when inhaled is of physiological interest. When breathed in small quantities it produces a pleasur- able nervous excitation, whence its common name of laughing-gas. But if mixed with atmospheric air and inhaled for some time it produces, like chloroform, temporary insensibility, and without the unpleasant after-effects of the latter substance. It is largely used for minor surgical operations, particularly in dentistry, to produce insensibility to pain and to allow the surgeon to operate more easily. 1 1 The composition of nitrous oxide and of nitric oxide may be easily determined by burning a small piece of potassium in a given volume of the gas, standing over mercury. The gas is thus decomposed, its oxygen uniting with the potassium and the nitrogen being set free. Nitrous oxide is then found to contain its own volume of oxygen and nitric oxide one-half its volume, whence the formulae N 2 and NO may be easily deduced.— Ep. 200 Text-Book of Inorganic Chemistry. ATMOSPHERIC AIR. Although the atmosphere is not a chemical compound, it is scarcely of less interest and importance to the chemist than water. The importance of the atmosphere for chemistry and other allied sciences is at once evident when we remember that the air has always played a large part in building up and in disintegrating the solid crust of the earth, that all processes of combustion or oxida- tion on the earth's surface are supported by it, that its constituents nourish both plants and animals, and that the state of health of individuals as well as of entire communities depends upon the purity of the air as well as of the water supplied to them. We have already learnt that atmospheric air deprived of its other admixtures consists of about 21 volumes of oxygen mixed with about 79 volumes of nitrogen. Besides these two gases atmospheric air always contains very varying quantities of water vapour, small quantities of carbonic acid (3 to 4 volumes per 10,000), and traces of ammonia (a few parts per million). The composition of atmospheric air by volume, taking the mean quantity of water vapour, is then : — Nitrogen O tygen Water (mean) . Carbonic acid . Ammonia These numbers lead to several interesting questions, the first of which is whether the nitrogen and oxygen contained in atmo- spheric air are chemically united or only mechanically mixed. If we abstract the water vapour and carbonic acid from the air, the proportion of nitrogen and oxygen then becomes : — Nitrogen .... = 79-04 volumes Oxygen = 20-96 „ ioo-oo volumes, and it has been found that ever since the composition of the air has been accurately known, this proportion has not appreciably varied, and that whether the air is taken from over the sea, from the surface of the earth inland, or from high mountain tops, even 78-35 volumes 20-77 >: 0-84 3) 0-04 si traces ioo-oo volumes. Atmospheric Air. 20 J as high as i4,ooofeet above the sea- level, the proportion of nitrogen and oxygen is always almost exactly the same. This speaks strongly for the supposition that the two gases are chemically united in atmospheric air. But in spite of this we know for certain that the gases are only mixed together — that they do not form a chemical compound in the atmosphere. This conclusion is supported by the followingfacts : — i. If we mix together 21 volumes of oxygen with 79 volumes of nitrogen, we obtain a gas which possesses all the properties of common air deprived of its water and carbonic acid. But we notice no elevation of temperature, which would certainly take place had the gases combined together chemically. ii. If cold boiled water is shaken up with atmospheric air, the water takes up more oxygen than nitrogen, and the air which re- mains behind is then poorer in oxygen than common air. Or, if common air is passed through a porous tube more of the lighter nitrogen diffuses through in a given time than the heavier oxygen. Neither of these partial separations of the nitrogen and oxygen in common air could take place if the two substances were chemically united. iii. Finally, every chemical compound contains its constituents not only in a fixed proportion, but also in weights which are some simple multiple of the atomic weight of its constituents. Atmo- spheric air contains its nitrogen and oxygen in a constant propor- tion by volume and weight, but not in simple atomic proportions. The composition of air by weight is : — Nitrogen 77 parts Oxygen 23 „ 100 parts, and if we divide these numbers by the respective atomic weights of the elements, we obtain the ratio : — - 7 : 2 l = S'55 : i - 44 = 3'85 : 1, 14 10 a ratio which, although it approaches 4 : 1, is far from being a simple one, as it would be in a chemical compound. The difference in the composition of the air, and of a compound containing four atoms of nitrogen and one of oxygen, becomes more apparent if we calculate the percentage composition of this unknown body [N 4 0]. We find it to be : — 202 Text-Book of Inorganic Chemistry. 4N = 56 = 77'8 per cent. = l6 = 22-2 „ 72 IOO'O According to which the air contains 0'8 per cent, nitrogen less, and o-8 per cent, oxygen more than this unknown compound, a difference which is much too great to be ascribed to experi- mental errors, especially as the methods employed for the analysis of such simple gaseous mixtures leave nothing to be desired in exact- ness, and indeed exceed in this respect almost all other analytical methods. A further and much disputed point is whether the atmospheric air has always possessed the same composition which it does at present, and whether its composition will not change in the course of ages. The occurrence of nitrogen in the mineral kingdom as nitre, with that which is contained in the substance of plants and animals, is so extremely minute compared with the immense quantities of nitrogen in the air that little change has probably taken place in the absolute quantity of nitrogen present in the air. But the same cannot be said of the oxygen. Although we have no means of even approximately estimating the temperature of the earth when it was in the liquid state, still it is more than probable that many of the oxygen compounds which now form the crust of the earth were built up from their constituents as the earth gradually cooled to the still extremely high tempera- tures at which the compounds of oxygen with silicon, calcium, hydrogen, &c, can exist. Experiment has shown that water is not only decomposed by the electric current, but also when raised to a very high temperature. If at any time the tempera- ture of the earth was higher than that requisite to decompose water, its atmosphere must have contained, besides nitrogen and hydrogen, much larger quantities of oxygen than after it had cooled down. Only when the temperature fell low enough could the hydrogen and oxygen unite_ together with enormous explosions to form water vapour. Possibly the protuberances of the sun's atmosphere, which are considered to be chiefly glowing hydrogen, are produced by the union of immense quantities of hydrogen and oxygen at heights in the sun's atmosphere where the tempera- ture is low enough to allow these two substances to combine with one another. Atmospheric Air. 203 Whether the amount of carbonic acid present in the atmosphere was ever greater than is now the case, cannot be satisfactorily answered in the present state of our knowledge. We know that during geological time immense - quantities of carbonic acid, as calcium carbonate (limestone), have become a portion of the solid crust of the earth. A further quantity of carbonic acid has also been fixed in the form of our deposits of coal, the remains of extinct plants. At the same time it must not be forgotten that the air is continually receiving carbonic acid from the interior of the earth by volcanoes, active and extinct, and how far these two processes balance one another we are at present unable to say. The luxu- riant vegetation which produced our coal-fields seems to indicate that the amount of carbonic acid then present in the atmosphere was greater than now. The quantity of oxygen contained in the atmosphere is propor- tionally small, and must be continually diminished by the respira- tion of men and animals and by the enormous quantities of coal which are continually being burnt. At the same time the quantity of carbonic acid in the air is increased by these processes. The question thus suggests itself, Can this enormous consumption of oxygen ever deprive the air of such quantities of this gas as to render respiration impossible ? We know the quantity of oxygen which is daily consumed by a man, from which we can easily obtain the quantity of oxygen annually consumed by the entire human race (reckoned at 1,500,000,000) ; and we can further calculate the total quantity of oxygen contained in the air. Now, if we consider that the oxygen consumed by animals, processes of putrefaction and combustion, is even nine times more than that consumed by man, then in 1,800 years the total quantity of oxygen in the air would only be diminished by about one-fifth per cent. But the idea that the air can ever become too poor in oxygen for respiration may be discarded when we remember the action of green plants on atmo- spheric carbonic acid. Green plants require carbonic acid for their existence and growth as much as animals require oxygen ; under the influence of sunlight their green parts decompose the carbonic acid, assimilating the carbon and setting the oxygen free. Thus, by the mutual action of plants and animals, the proportions of oxygen and carbonic acid in the air remain practically unchanged for a long time. Water and air, once considered to be chemical elements, are still important elements in another sense, and especially for the 204 Text- Book of Inorganic Chemistry. science of health. The remains of Roman aqueducts illustrate the importance which was attached to a good supply of pure water 2,000 years ago by a nation without our present scientific know- ledge. But in this respect we are still far behind the Romans. For centuries we have contented ourselves with impure well-water, and only in recent years have cholera, typhus, and other epidemics, together with a general high death-rate, been attributed to an impure and insufficient water supply. The question of ventilation, or the supply of pure fresh air, is even more neglected. It is well known how detrimental it is for a number of persons to remain long in a closed space where the air becomes charged with carbonic acid, water vapour, and other noxious emanations from the body, and it is also well known how easily this great evil may be avoided by leading away the bad air and supplying good air in its place. But, notwithstanding this knowledge, public and private buildings are continually being erected with insufficient ventilation, or, in some cases, with even none at all. COMPOUNDS OF NITROGEN WITH THE HALOGENS. Nitrogen combines with the halogens even less easily than with hydrogen, and never unites directly with them. The halogen com- pounds of nitrogen can only be obtained, like ammonia, by indirect methods ; unlike ammonia, which they resemble in composition, they do not unite with acids, and they are further distinguished by their great instability. The elements in these compounds are so loosely combined, that mere contact with some indifferent liquid or solid often suffices to decompose them with a. powerful explosion. They are therefore dangerous substances to deal with, and their properties ought to be known before commencing experiments with them. The best known of these compounds are nitrogen chloride and nitrogen iodide. Nitrogen Chloride : NC1 3 (?). This compound, which may be considered as ammonia in which the hydrogen is partly or entirely displaced by chlorine, but of which the exact composition is un- Compounds of Nitrogen with the Halogens. 205 known, is obtained by the action of free chlorine on ammonia, with simultaneous production of hydrochloric acid. Nitrogen chloride is a yellow oily liquid, with a specific gravity of r6, and sinks in water. It possesses a powerful, piercing odour, by which minute quantities of it may be detected. The compound is easily obtained by passing chlorine into a saturated solution of ammonium chloride at about 30 , when the following reaction probably takes place : — NH..C1 + 3C1 2 = 4HCI t NCI a . Aqueous ammonia is even more easily decomposed than the chloride, but cannot be employed, as the nitrogen chloride is again decomposed by free ammonia. Owing to the great danger in experimenting with nitrogen chlo- ride, it is not advisable to prepare large quantities of it. By the following method, small quantities, sufficient to observe its odour and explosive properties, may be safely obtained. In a strong leaden dish about eight inches wide and four inches deep is placed a second smaller dish, also of lead, and about three inches wide, and the large dish then filled to about two-thirds with a luke-warm saturated solution of ammonium chloride. A cylin- drical glass jar is also filled with the same liquid, inverted in the dish and suspended by a wire just over the smaller dish. Chlorine gas is now led into the cylinder until the liquid is completely dis- placed, when the apparatus is allowed to stand quietly. As the liquid now gradually rises in the jar, small oily drops of nitrogen chloride appear on the surface, which gradually fall inco the small leaden dish. When nearly all the chlorine is absorbed, the small dish is removed, a portion of the solution poured off, and the com- pound exploded. The explosion is best brought about by tying an oiled feather to the end of a long stick, and then bringing it into contact with the nitrogen chloride ; the whole then instantly explodes, while the liquid is shot up several feet high and the base of the leaden 1 dish pressed quite flat. The force of the explosion is so great that a drop of the chlo- ride of the size of a pea, exploded under water in a porcelain dish with an oiled feather, not only breaks the dish into small fragments, but forces the piece immediately under the drop into the wooden support. The discoverer of nitrogen chloride, the French chemist Dulong, who being unacquainted with its explosive character, attempted to 206 Text-Book of Inorganic Chemistry. distil some in a retort, lost several fingers of one hand by the ex- plosion. , If in any way larger quantities of nitrogen chloride have been obtained than it is considered advisable to explode, the substance may be left to itself under the liquid. After some little time, it will gradually disappear owing to spontaneous decomposition. Nitrogen chloride is also obtained by the electrolytic decom- position of ammonium chloride. If two platinum plates are im- mersed in a warm saturated solution of the salt, and a strong current passed for a short time, the chlorine, which is evolved at the positive plate, becomes converted into nitrogen chloride. A powerful explosion then ensues if the two plates are brought into contact. Nitrogen Iodide : NI 3 (?). This compound is a dark-brown solid, closely resembling precipitated iodine. It is obtained when a saturated alcoholic solution of iodine is mixed with excess of strong ammonia. The precipitate is then brought on a filter and well washed with water. As long as the substance is moist, it bears pressure and contact with all those substances which so easily explode nitrogen chloride ; but if dried, mere contact with a particle of dust or the percussion produced by closing a door is often sufficient to decompose it into its constituents with an explosion. PHOSPHORUS. Chemical Symbol : P. — Atomic Weight: 31. Phosphorus is not found in the free state in nature, but occurs in the mineral kingdom principally as calcium phosphate in the minerals ■phosphorite, apatite, &c, and in coprolites. All soils contain phosphorus, but usually in only minute quantities. From the soil the phosphorus is taken up by plants, which produce very complex organic compounds containing it, and is collected by them in various parts of their structures, particularly in their seeds. These phosphorus compounds then pass into the bodies of animals with the nutriment contained in the plants ; they form an impor- tant constituent of the blood, nerve substance, &c, and undergo an oxidation in the body, being again converted into phosphoric acid Phosphorus. 207 In combination with lime, this phosphoric acid forms the greater part of the mineral substance of the bones, while a portion is expelled from the body in the fasces, and a larger portion in the urine, and is thus returned again to the soil. Phosphorus was first obtained by its discoverer, Brand, of Hamburg (1669), from concentrated urine. A hundred years later Scheele recognized the fact that the urine as well as the bones contain calcium phosphate, and that it is this substance which yields the phosphorus obtained from urine. The name phosphorus is derived from its property of shining in the dark (05s, light, and c[>epeLv, to bear). We are acquainted with phosphorus in at least two different modifications, which are so different in appearance and in their physical as well as in some of their chemical properties that they might have been thought to be different elements if it were not easy to convert one into the other, and if the products of their union with other elements were not the same. The variety longest known, and commonly called ordinary or yellow phosphorus, is, when freshly prepared, a colourless trans- parent solid, which is coloured yellow under the action of light, and has a specific gravity of 1 '83. It may be melted under water at a temperature of 44° to a colourless liquid, easily solidifying again and sometimes assuming a crystalline form. At the ordinary temperature phosphorus is as soft as wax, but becomes brittle at low temperatures. Phosphorus boils at 290 , and when out of contact with air maybe distilled unchanged. Its gas is colourless, and has a specific gravity (at 1,000") of 4-5. From the density of its vapour its molecular weight is 4-5 x 28-88 = 129-9, or > m other words, the molecule of phosphorus in the gaseous state has the composition P 4 and contains four atoms. Phosphorus heated in the air slightly above its melting point catches fire and burns, producing white fumes of phosphoric anhydride. At lower temperatures, when exposed to the air it oxidizes without the production of flame, forming phosphorous anhy- dride. /The white fumes produced when a stick of phosphorus is exposed to the air, and which are luminous in the dark, are due to phosphorous anhydride, and the garlic-like odour noticed at the same time is produced by the same substance. Ordinary phos- phorus is so inflammable that it catches fire when rubbed on a rough surface ; it should never be held, even for a short time, be- tween the fingers, as burning phosphorus produces very painful, 2o8 Text-Book of Inorganic Chemistry. slowly healing wounds. This modification of phosphorus is always preserved under water, in which it is insoluble. And although it possesses a very strong affinity for oxygen, it does not decompose water like potassium and sodium. It is slightly soluble in alcohol and ether, but easily soluble in carbon disulphide, from which solu- tion it again separates in crystals belonging to the regular system. Phosphorus in this modification is extremely poisonous ; even small quantities have been known to produce fatal effects. A second modification of phosphorus is the red or amorphous variety. This is a red-brown amorphous powder, of specific gravity 2'2 (or heavier than common phosphorus). It does not melt, even when heated up to 250°, neither does it produce white fumes when exposed to the air, nor is it luminous in the dark. It catches fire in the air only when heated up to 260°, is insoluble in carbon disulphide, is odourless, and not poisonous. This red modification is produced when ordinary phosphorus is heated in an atmosphere of carbonic acid or hydrogen to a temperature of 240° for some time ; the change goes on more quickly if the phosphorus is heated in sealed tubes to 300 . But in all cases the weight of the phosphorus neither increases nor diminishes. If iodine is mixed with an excess of ordinary phos- phorus and the mixture heated, the portion of phosphorus which does not combine with the iodine is converted into the red variety. A small quantity of iodine suffices to produce the change in a large quantity of phosphorus. The same change is also slowly produced by the action of light. Phosphorus exposed to the light becomes first yellow and then red. In order to purify the red modification from common phosphorus mechanically mixed with it, it may be re- peatedly shaken with carbon disulphide or boiled with caustic soda. It is remarkable that just as ordinary phosphorus is converted into the red variety on heating, so also the latter may be changed into the former by the action of heat. Amorphous phosphorus heated in a retort filled with carbonic acid to 260 — i.e. to a temperature only slightly higher than is required for its production — is completely changed into the common form. And in this case too, no change in weight is observed. Since phosphorus has been employed for important technical purposes, particularly for the manufacture of lucifer matches, large quantities are annually manufactured. The material employed for its extraction is not urine, from which it was first prepared, but is the calcium phosphate of burnt bones, or as one of the several Phosphorus. 209 forms in which the compound occurs in the mineral kingdom. Although the attraction of phosphorus for oxygen is very great, it is exceeded by that of glowing charcoal. Phosphoric acid, in the proper form, when mixed with powdered charcoal and heated to bright redness, is completely reduced with formation of carbonic oxide. But if the compound of phosphoric acid and lime, as it occurs in nature or in burnt bones (calcium phosphate, Ca 3 (POj) 2 ), were employed, no reduction would take place. It is only the free acid, or compounds of calcium containing more phosphoric acid than the normal salt, which are decomposed on heating with charcoal. In order to obtain such a compound the following method is adopted. Finely powdered bone-ash, containing normal calcium phosphate, is warmed with enough dilute sulphuric acid to withdraw two-thirds of the calcium, which unites with the sulphuric acid forming insoluble sulphate, while the remaining third of the calcium remains combined with the whole of the phosphoric acid as a soluble acid phosphate : H 4 Ca(P0 4 ) 2 , or p„ \ ^. _\ commonly called superphosphate of lime : — £glo Ca 3 + 2S0 2 (OH) 2 = 2 S0 2 -0 2 Ca + pg } C Q 9 Ca The clear solution is next drawn oft' and concentrated in leaden pans, and the concentrated solution so obtained is then mixed with powdered charcoal, dried, and heated to low redness. During this process the acid phosphate loses water and becomes converted PO ) into calcium metaphosphate : Ca(P0 3 ) 2 , or p„ 2 \ 2 Ca, a com- pound which contains more phosphoric acid than the normal phosphate, and from which two-thirds of the phosphorus is ob- tained on glowing with charcoal : — PO) 3(PO.,) 2 3 Ca + 10C = P 4 + pQ 6 Ca 3 + 10CO, the calcium being again converted into the normal phosphate. The reduction is carried on in earthenware retorts connected with earthenware vessels containing water. The gaseous phosphorus distils over and condenses under the water, while the carbonic oxide escapes. The raw product is then purified by redistillation from iron retorts and by squeezing the liquid through leather under warm water. Phosphorus as it occurs in trade is usually p 210 Text-Book of Inorganic Chemistry. in the form of sticks ; this form is given it by the highly dangerous practice of sucking up the molten substance under water into conical glass tubes, from which it is easily removed on solidifi- cation. If a stick of phosphorus, one end of which is wrapped in paper, is used to write on a board, the characters are luminous in the dark. Small pieces of phosphorus are by this means rubbed oft on the board and gradually oxidize to phosphorous acid. The luminosity of phosphorus or its oxidation to phosphorous acid is always accompanied with the production of ozone. Various bodies in the gaseous state when mixed with atmospheric air prevent its slow oxidation and at the same time its luminosity ; among these are small quantities of oil of turpentine vapour, ether, ammonia, and various hydrocarbons. Phosphorus is, however, much less luminous in pure oxygen than in common air, or in oxygen diluted with an indifferent gas. Phosphorus unites as easily with the halogens as with oxygen, liberating a large amount of light and heat. With chlorine it forms either phosphorus trichloride or pentachloride according to the relative quantities of phosphorus and chlorine which act upon one another. The affinity of phosphorus for sulphur is so great that by drop- ping flowers of sulphur on to phosphorus melted under water the union takes place with a very considerable elevation of temperature, and dangerous explosions may often be produced. Phosphorus does not appear to unite directly with hydrogen, nitrogen, or carbon. The principal use of phosphorus is in the manufacture of lucifer matches, for which both modifications are employed. Common matches are dipped in a mixture containing common phosphorus, and ignite when rubbed on any rough surface, the ignition being imparted to the wood by previously dipping the matches in melted paraffin. In the other kind of matches, commonly known as safety matches, the matches themselves contain no phosphorus but only some substance which readily gives off oxygen (usually potas- sium chlorate), while one side of the box is covered with finely divided amorphous phosphorus put on with gum water. The matches are then ignited by rubbing them on the prepared surface. Their ignition is due to the change of the part of the amorphous phosphorus rubbed by the match into the common form by the heat of friction ; this then adheres to the match and causes it to Compounds of Phosphorus and Hydrogen. 2 1 1 catch fire. If phosphorus is rubbed up with fat a mixture is pro- duced which is luminous in the dark, and which is used as a poison for rats and mice. COMPOUNDS OF PHOSPHORUS AND HYDROGEN, Two compounds of these elements are known — viz. a gaseous compound having the composition PH 3 and corresponding to- ammonia (NH 3 ), and a liquid, P.jH,, corresponding to- amidogen, (NH 2 ), Caseous Phosphoretted Hydrogen or Phosphine : PH .. This compound is a colourless gas, with an exceedingly un- pleasant garlic odour, and is nearly insoluble in water. Its specific gravity is ri8, which corresponds to a molecular weight of ri8 + 28'88 = 34'o8, and its molecule is therefore represented by the formula PH 3 : The gas is not spontaneously inflammable, but at once ignites when brought into contact with a flame, burning to form water and phosphoric acid. Phosphine is easily obtained by heating phosphorus, with a solution of a strong base, such as potash, soda, lime, or baryta Although phosphorus possesses so strong an attraction for oxygen it cannot decompose water even at the boiling temperature, but the decomposition goes on readily in the presence of a strong soluble base, a portion of the phosphorus uniting with the base to form a hypophosphite. The gas may also be prepared from solid calcium phosphide, which, in contact with warm water, is rapidly decom- posed into phosphine, calcium hypophosphite, and calcium hydrate. But the gas obtained by either of these processes is always spon- taneously inflammable. This is due to the fact that it contains small quantities of the volatile liquid phosphoretted hydrogen, which is spontaneously inflammable, and causes the gas to ignite in the air just as it does when mixed with other inflammable gases, such as carbonic oxide or hydrogen. The following equations re- present the simultaneous production of the two compounds when phosphorus is heated with caustic potash : — 212 Text-Book of Inorganic Chemistry. (i.) 4P + 3KOH + 3H 2 = 3H 2 PO-OK + PH 3 . Potassium hypo- phosphite (ii.) 6P + 4KOH + 4H 2 = 4H.PO-OK + P 2 H>. The preparation of phosphine containing the spontaneously inflammable liquid compound must be performed with care. If phosphorus were heated in a flask half filled with caustic potash and provided with a tightly fitting cork and delivery-tube, the phosphoretted hydrogen would unite with the oxygen in the flask and catch fire, and the heat so produced would be sufficient to cause an explosion. It is better, therefore, to fill the flask nearly up to the neck and to place the cork in loosely until the oxygen contained in the small quantity of air present has been consumed. The delivery-tube must be wide so that it may not become closed with small particles of phosphorus carried over mechanically ; it should dip under warm water. In order to prepare the spontaneously inflammable gas from calcium phosphide, this compound is broken into small pieces of the size of a pea and dropped through a wide tube into a Woulffe's bottle half filled with warm water, the other tubulus of the bottle being provided with a delivery tube. The wide tube must just dip under the liquid and must be closed with a small cork. Before any calcium phosphide is introduced, the air contained in the Woulffe's bottle must be completely displaced by carbonic acid to prevent the gas from catching fire inside the bottle. The spontaneously inflammable gas when preserved over water for some time, and particularly when exposed to the light, is found to lose its property of catching fire when brought into contact with the air, and at the same time the walls of the glass vessel become covered with a yellow film. 1 This phenomenon depends upon the instability of the liquid compound, which easily breaks up, espe- cially in sunlight, into phosphine and phosphorus : — 3 P 2 H 4 = 4PH3 + 2P. The spontaneously inflammable liquid may be easily separated from the gaseous compound by passing the mixture through a U-tube placed in a freezing mixture. By this means the volatile liquid is condensed and the gas is no longer spontaneously inflam- mable. 1 This yellow film is considered by some chemists to contain hydrogen and to consist of a compound of 4 atoms of phosphorus united with 2 atoms of hydrogen : P 4 H 2 . — Ed. Compounds of Phosphorus and Oxygen. 213 Pure phosphine may be prepared by dropping calcium phos- phide into hydrochloric acid instead of into water, or by acting upon phosphorus with alcoholic instead of aqueous caustic potash. Hydrochloric acid and alcohol appear to prevent the formation of the liquid phosphoretted hydrogen, or to decompose it immedi- ately. The pure gas when led into nitric acid mixed with a small quantity of the fuming acid becomes spontaneously inflammable. This is probably due to the abstraction of a portion of its hydrogen) which is oxidized by the nitric acid, with the production of liquid phosphoretted hydrogen. Phosphine is distinguished from its analogue ammonia not only by its insolubility in water, but also by its incapability of uniting either with hydrochloric acid or with oxy-acids to form chemical compounds, but it still possesses weak basic properties. If dry phosphine and dry hydriodic acid are brought into contact with one another they unite to form a compound called fihosfihonium iodide, PH 4 I (p. 128), analogous in composition to ammonium iodide, which crystallizes in cubes, and may be sublimed unchanged. The compound is, however, unstable ; it is decomposed by water into hydriodic acid and phosphine. Liquid Phosphoretted Hydrogen : )'.,H r This compound, which is prepared from the spontaneously in- flammable gas by passing it through a freezing mixture, is a colour- less, powerfully refractive liquid, does not mix with water, and at once catches fire when exposed to the air. The liquid undergoes a gradual decomposition, even in sealed tubes, into phosphine and amorphous phosphorus. The gas finally exercises so great a pressure on the tube as to produce an explosion. It is, therefore, not practicable to preserve the liquid in sealed tubes. COMPOUNDS OF PHOSPHORUS AND OXYGEN. Although phosphorus is so closely allied to nitrogen, we are only acquainted with two oxides of this element — viz. phosphorous anhydride (P„0 3 ) and phosphoric anhydride (P 2 5 ), corresponding to nitrous anhydride and nitric anhydride respectively. Compounds 214 Text-Book of Inorganic Chemistry. of phosphorus having a similar composition to nitric peroxide, nitric oxide, and nitrous oxide, have not yet been prepared. From these two oxygen compounds are derived the two acids, common tribasic phosphoric acid and dibasic phosphorous acid. Besides these, a third (lower) acid is also known, which is called hypophosphorous acid and is monobasic. The following are the formula; of these three acids of phosphorus, all containing pentad phosphorus : — (OH Common phosphoric acid . . . POJOH OH {OPT oh Hypophosphorous acid . . H 2 PO-OH A fourth acid of phosphorus has recently been discovered of the composition : j pofOHl ' anc ' ca " e ^ hypophosphoric acid. PHOSPHORIC ANHYDRIDE : P.,0,. The white fumes produced when phosphorus burns with a flame in dry air consist of this compound. If the phosphorus is burnt in aclosed space — e.g. under a bell-jar — they then condense to a light snow-white powder, which very readily attracts moisture from the air and forms liquid phosphoric acid. When thrown into water the anhydride hisses like a red-hot iron, and a large quantity of heat is at the same time set free. In fact, there are few substances which equal this body in its attraction for water. Unlike nitric anhydride, it is not volatile at ordinary temperatures, and does not, therefore, fume in moist air ; but it may be sublimed when strongly heated in a glass tube. The acid first produced when the anhydride unites with water contains only one atom of hydrogen displaceable by a metal, and is therefore monobasic. This acid is ?netaphosphoric acid: PO,-OH, and its production is shown by the following equation : — P0*}° + H,0 = 2PO,-OH. Phosphoric Metaphosphoric anhydride acid Phosphoric A nhydride. 2 1 5 Besides this phosphoric acid, which corresponds to nitric acid in its composition, we are also acquainted with a tri- and a tetrabasic acid. Tribasic phosphoric acid, which has the composition PO(HO) 3 , and whose normal salts with monobasic metals contain three atoms of the metal replacing the three atoms of hydrogen, is common or orthophosfihoric acid. Only compounds of this tribasic acid occur in nature. Metaphosphoric acid, produced as we have seen from phos- phoric anhydride and water, takes up a molecule of water when boiled and becomes converted into common phosphoric acid : — P0. 2 -OH + H 2 = PO(OH) 3 . The change which has taken place may be readily proved by a few simple reactions. Metaphosphoric acid coagulates white of egg, common phosphoric acid does not ; on adding a little ammonia, the. former gives with silver nitrate a white precipitate of silver metaphosphate : P0 2 -OAg, the latter with the same reagent gives a yellow precipitate of silver orthophosphate : PO(OAg) 3 . The conversion of metaphosphoric acid into common phosphoric acid is therefore complete when the liquid no longer coagulates white of egg and gives a bright yellow precipitate with silver nitrate. Preceding the formation of common phosphoric acid from meta- phosphoric acid, when the solution is warmed, is the production of tetrabasic pyrophosphoric acid, which may be recognized by the fact that this acid does not coagulate white of egg but gives a white precipitate of silver pyrophosphate when mixed with silver nitrate. Its production from the monobasic acid is shown, by the following equation : — 2 P0 2 -OH + H 2 = 0{Pg[gH), Metaphosphoric Pyrophosphoric acid acid All three acids may therefore be represented as compounds of phosphoric anhydride with varying quantities of water, thus ; — PO,) U + H 2 = 2P0 2 OH Metaphosphoric acid P0 4o + 2HO - oI p °(° H) » POJ° + 2H * V _ u lPO(OH) 2 Pyrophosphoric acid P °*}0 + 3 H 2 = 2PO(OH) 3 Orthophosphoric acid 216 Text-Book of Inorganic Chemistry. ORTHO- OR COMMON PHOSPHORIC ACID : PO(OH) 3 . On evaporating the aqueous solution of common phosphoric acid, when the temperature is not allowed to rise much above 1 50 , the acid remains behind as a colourless, thick, viscid liquid, of specific gravity I '88, from which large, colourless crystals occasionally separate out. It readily attracts moisture from the air and deliquesces. When heated more strongly it loses water and becomes first converted into pyro- and then into metaphosphoric acid. Common phosphoric acid is easily prepared by oxidizing phosphorus with nitric acid. Sticks of phosphorus are heated in a retort with pure nitric acid of specific gravity 1 -2, when an energetic reaction takes place and red vapours of the lower oxides of nitrogen are given off. The retort is best provided with a cooled receiver, so that the nitric acid distilling over may be poured back into the retort. As soon as all the phosphorus has disappeared the excess of nitric acid is driven off, which at the same time oxidizes the phosphorous acid, also present, to the higher phosphoric acid. Finally, the viscid liquid is heated in a platinum dish until the temperature reaches 150 . Another and very easy method of preparing phosphoric acid is from phosphorus pentachloride. If this substance is dropped into water in small quantities at a time, it decomposes, with consider- able rise in temperature, into phosphoric and hydrochloric acids, phosphoric oxychloride being produced as an intermediate pro- duct : — PC1 5 + 4H 2 = PO(OH) 3 + 5HCI. Pure phosphoric acid may then be obtained by evaporating the strongly acid liquid. Since commercial phosphorus always contains arsenic, because the sulphuric acid used in the preparation of the acid calcium phosphate is always contaminated with arsenic, the phosphoric acid prepared either from phosphorus or from phosphorus penta- chloride always contains traces of arsenic in the form of arsenic acid. In order to separate this impurity the phosphoric acid, after being freed from the nitric or hydrochloric acid, is diluted with a large quantity of water, gently warmed, and a stream of sulphu- retted hydrogen passed through it as long as insoluble yellow arsenious sulphide separates out. The whole of the arsenic is Ortlw- or Common Phosphoric Acid. 217 precipitated, when a small portion of the clear liquid again saturated with sulphuretted hydrogen and allowed to stand for some hours shows no further turbidity. The clear filtered liquid then yields pure phosphoric acid on evaporation. Phosphoric acid may also be prepared from the acid calcium phosphate (superphosphate of lime), obtained by acting on bone-ash or some other form of calcium phosphate with sulphuric acid. The solution of the superphosphate, which still contains calcium sulphate, is mixed with a solution of ammonium carbonate to pre- cipitate the lime. The filtered liquid, which contains ammonium sulphate and phosphate, is then evaporated to dryness and glowed in order to drive off the ammonium sulphate and to expel the ammonia from the ammonium phosphate. Phosphoric acid pre- pared in this way always contains, however, small quantities of ammonium salts. Phosphoric acid is one of the strongest inorganic acids. At a high temperature it even expels sulphuric acid from its compounds, being less volatile than this acid. Since it is a tribasic acid, it forms three series of salts : — (i.) Normal salts, in which all three of its hydrogen atoms are displaced by a metal. (ii.) Monacid salts, in which two atoms of hydrogen are dis- placed by a metal, and (iii.) Diacid salts, in which only one atom of hydrogen is displaced. In the normal salts the hydrogen may be either displaced by three atoms of a monad metal, by one of a monad and one of a dyad, or by one of triad metal, thus ; — Normal sodium phosphate . . . PO(ONa) 3 Normal sodium-ammonium phosphate . PO QNH ( O Me" Normal ammonium-magnesium phosphate PO J qj^t? Normal ferric phosphate . . . . PO-0 3 Fe'" In the first class of acid salts the two atoms of hydrogen may either be displaced by two atoms of a monad metal or by one atom of a dyad, thus : — Monacid sodium phosphate . . . ^^ | OH Monacid calcium phosphate . . . PO qj_j 2 1 8 Text-Book of Inorganic Chemistry. And, finally, for the second class of acid salts we may take as an example : — Diacid sodium phosphate . . . PO ,qj^ The phosphates of the metals, particularly the normal salts, are mostly insoluble in water, but easily soluble in nitric acid ; only the phosphates of the alkalies can be obtained in a well-defined crystalline state from their solutions. Normal silver phosphate is a bright yellow amorphous precipi- tate, easily soluble in nitric acid. It is obtained on the addition of silver nitrate to a solution of common sodium phosphate. This latter compound, the phosphate of soda of the shops, is not the normal compound, but is monacid sodium phosphate — i.e. it con- tains one atom of hydrogen undisplaced. The reaction by which normal silver phosphate is produced when this salt is mixed with silver nitrate is as follows : — PO loH' a ^ + 3 NC V 0A g= PO(OAg) 3 + 2N0 2 -ONa + N0 2 OH. Whence it is seen that nitric acid is at the same time set free, and if the clear liquid is filtered off, it will be found to possess an acid reaction, and to contain some silver phosphate dissolved in it. We produce, therefore, from the two neutral solutions of mono- hydric sodium phosphate and silver nitrate an acid liquid. The phosphoric acid contained in this and other similar acid compounds can, therefore, only be completely precipitated by silver nitrate by neutralizing with soda or some other suitable base. The insolubility of ammonium-magnesium phosphate in water, especially in presence of ammonia, and the constant composition of this compound, makes it well adapted both to detect phosphoric acid and to determine its quantity. It is precipitated as a white crystalline powder when a solution of sodium or any other soluble phosphate is mixed with magnesium sulphate, ammonium chlo- ride, and ammonia. If monacid sodium phosphate is mixed with magnesium sul- phate the acid magnesium compound is produced. Thus : — P^OH 1 ^ 8 + sc V°s M S = ^joH^" 1 " sc V(ONa) a , which is soluble in water, even in the presence of ammonium chloride!' If, however, ammonia is added in excess, the insoluble ammonium- magnesium phosphate is precipitated : — ■ Pyrophosphoric A cid. 2 1 9 1U |OH + AH » - FU |ONH,, The addition of ammonium chloride is necessary to prevent the precipitation of magnesium hydrate from the excess of magnesium sulphate and ammonia. Magnesium hydrate forms an easily soluble compound with ammonium chloride. An excellent method of detecting and separating phosphoric acid in nitric acid solutions containing iron, lime, &c, is furnished by molybdic acid. If such a solution containing free nitric acid is mixed with a solution of ammonium molybdate in nitric acid, the liquid first becomes yellow and then deposits a yellow crystalline precipitate consisting of a compound of phosphoric and molybdic acids. By this means a soil, for example, may be examined for phosphoric acid, and its quantity, if present, determined. If solid monacid sodium phosphate is gently heated, it first of all loses its water of crystallization. On further heating it gives up the hydrogen which it contains, together with sufficient oxygen to form water, and normal sodium pyrophosphate remains behind : — „ po f(ONa), = jPO(ONa) 2 HQ " ru {OH u {PO(ONa) a + n - iU - When diacid sodium phosphate is heated, water is also given off, but the compound which remains behind is then normal sodium metaphosphate : — PO !(OH) = p °2-° Na + H 'A The same compound is more easily obtained by heating acid ammonium-sodium phosphate {microcosmic salt), which then gives off ammonia and water : — (ONa PO ONH 4 = P0 2 -ONa + NH 3 + H 2 0. (OH Pyrophosphoric Acid: 0|pn)oH^ = H^P^O,. This compound is obtained as a colourless crystalline mass, easily soluble in water, when phosphoric acid is heated slightly above 200 , until a portion when neutralized with ammonia gives a pure white precipitate with silver nitrate. If sodium pyrophos- phate, obtained as mentioned above, is dissolved in water mixed 220 Text-Book of Inorganic Chemistry. with silver nitrate, the white precipitate filtered off, washed well, and then again decomposed by passing a current of sulphuretted hydro- gen through water containing it in suspension, black silver sulphide is formed, and the acid solution contains pyrophosphoric acid. Such an aqueous solution of the acid cannot, however, be kept for any time ; it combines with water and forms common phosphoric acid, slowly when cold, more rap'idly when heated : — Pyrophosphoric acid is a tetrabasic acid, its four atoms of hydrogen may be displaced by the same or by different metals, and besides the normal salts three classes of acid salts are also known accordingly as one, two, or three atoms of hydrogen are displaced. Metaphospboric Acid : P0 2 OH. This acid is obtained in aqueous solution by allowing phos- phoric anhydride to deliquesce in the air, or as a vitreous, amorphous mass by heating common phosphoric acid to redness. In the latter state it forms the glacial phosphoric acid of the pharmacists. In combination with sodium it is produced, as we have already seen, by heating microcosmic salt to redness. The acid and its soluble salts give with silver nitrate an insoluble white precipitate of silver metaphosphate : PO.,-OAg, and are distinguished from the two other phosphoric acids and their salts by their property of coagulating albumen (white of egg). Metaphosphoric acid when in solution gradually absorbs water and is converted into common phosphoric acid. The change goes on slowly in the cold, but more rapidly when the liquid is heated. Solutions of the metaphosphates also take up water in the same way — e.g. : — ONa is a P0 2 -ONa + H 2 = PO|, OH }. The recently discovered hypophosphoric acid : \ pn/nT-n > tetrabasic acid intermediate between phosphoric and phosphorous acids, and is produced with phosphorous and phosphoric acids, when phosphorus oxidizes slowly in moist air. Its solution may be evaporated at a low temperature, but is decomposed on boiling into phosphorous and phosphoric acids. Phosphorous Anhydride. 221 Phosphorous Anhydride : P 2 3 . The white fumes produced when a stick of phosphorus is exposed to the air consist principally of phosphorous acid. The pure anhydride may be obtained by gently heating a stick of dry phosphorus in a glass tube and leading a current of dry air over it. The phosphorus then burns with a pale greenish flame, and the anhydride is deposited in the cool end of the tube as a snow-white powder resembling phosphoric anhydride. Phosphorous anhydride has still a strong attraction for oxygen ; when exposed to the air it catches fire and burns to phosphoric anhydride. When phosphorous anhydride is brought into contact with water, or when phosphorus is allowed to oxidize slowly in moist' air, the corresponding acid — phosphorous acid — is produced. (OT-T OH Although this compound contains three atoms of hydrogen, only two of them can be displaced by metal, and it is therefore a dibasic acid. It is best prepared by dropping phosphorus trichlo- ride into water. The chloride then breaks into hydrochloric and phosphorous acids, thus : — PC1 3 + 3H,0 = 3HCI + HPO-'q^ The strongly acid liquid is freed from the hydrochloric acid by rapid evaporation, when phosphorous acid remains behind as a clear, colourless syrup, which if preserved for some time over sul- phuric acid in a vacuum changes into a crystalline mass. Pure phosphorous acid is obtained as a bye-product in the pre- paration of acetyl chloride, when glacial acetic acid is gradually added to phosphorus trichloride, contained in a retort, and the volatile acetyl chloride together with the excess of acetic acid dis- tilled off. Phosphorous acid then remains behind as a thick colourless liquid, which solidifies on cooling. If we represent the radical of acetic acid by the symbol A, acetic acid then has the formula AOH, and the process is represented by the following equation : — PC1 3 + 3AOH = 3ACI + HPO(OH) 2 . Phosphorous acid melts at 70 ; if the acid is strongly heated it 222 Text-Book of Inorganic Chemistry. breaks up into phosphoric acid and phosphine, the latter catching fire in the air : — 4HPO(OH) 2 = 3PO-(OH) 3 + PH 3 . The phosphites when strongly heated undergo a similar change. The strong attraction of phosphorous acid for oxygen and its readiness to pass into phosphoric makes it a powerful reducing agent. When an aqueous solution is gently warmed with solutions of mercury or silver the compounds of these metals are reduced to the metallic state. Similarly, when the acid is exposed to the air it gradually takes up oxygen and becomes converted into phosphoric acid. Hypophosphorous Acid : (H 2 PO)OH. This acid of phosphorus, which contains less oxygen than phos- phorous acid, is monobasic. Its molecule contains three atoms of hydrogen, but only one of these atoms can be displaced by a metal. The acid is soluble in water, and on evaporating down its aqueous solution remains behind as a syrupy liquid, which crys- tallizes below o c . Most of its salts crystallize easily, and are all soluble in water. The hypophosphites are produced, as we have already seen (p. 21 1), when an aqueous solution of an alkali is heated with phos- phorus. Phosphoretted hydrogen is then evolved, and the hypo- phosphite remains behind in solution. For the preparation of hypophosphorous acid and its salts it is best to employ barium hydrate (baryta-water), since the insolubility of barium carbonate affords a ready means of removing the excess of alkali. After a hot saturated solution of baryta-water has been heated with excess of phosphorus for some time until phosphoretted hydrogen is no longer evolved, the whole is allowed to cool, and carbonic acid passed through the clear liquid until it no longer reacts alkaline ; heat being applied towards the end. The barium carbonate is then filtered off, and the liquid evaporated down, when barium hypophosphite crystallizes out on cooling. The production of this salt is represented by the following equation, which is similar to that by which potassium hypophosphite is formed (p. 212) : — 6P + 2Ba(0H)., + 4H.0 = 2^ 2 pQ}o a Ba" + P„H 4 . Compounds of Phosphorus with the Halogens. 223 Milk of lime (calcium hydrate) may be employed instead of the baryta-water, when calcium hypophosphite is easily obtained in the crystalline form. Calcium hypophosphite is also obtained when calcium phosphide is decomposed by warm water (p. 211). The free acid may be prepared from the barium salt by pre- cipitating the barium with sulphuric acid and filtering off from the insoluble barium sulphate. The aqueous solution can be evapo- rated to a certain degree of concentration without decomposition, but when more strongly heated breaks up (like phosphorous acid) into phosphoric acid and phosphine : — 2H 2 PO-OH = PO(OH) 3 + PH 3 . It possesses a strong attraction for oxygen, and, like phosphorous acid, reduces various metals from their oxides or from solutions of their salts. COMPOUNDS OF PHOSPHORUS WITH THE HALOGENS. We are acquainted with two compounds of phosphorus and chlorine — viz. phosphorus trichloride, PC1 3 , in which the phos- phorus is triad, and phosphorus pentachloride, PC1 5 , in which the phosphorus is pentad. Similar compounds with bromine are also known. A corresponding phosphorus triodide, PI 3 , also exists, but the pentaiodide has not been prepared ; on the other hand, a com- pound containing less iodine, of the composition (PI 2 ) 2 or P 2 I 4 — phosphorus di-iodide — is known. Phosphorus forms only one com- pound with fluorine — the gaseous pentafluoride, PF 5 . The attraction between phosphorus and the halogens is so great that all these compounds, except the fluoride, can be easily pre- pared by the direct union of their constituents : while the strong attraction of phosphorus for oxygen and of the halogens for hydro- gen is the reason why these compounds are easily decomposed by water with a considerable evolution of heat. 224 Text-Book of Inorganic Chemistry. PHOSPHORUS TRICHLORIDE (Phosphorous Chloride). Composition : PC1 3 . This compound is a colourless liquid, fuming in moist air, and which may be prepared in quantity in the following manner. The bottom of a large tubulated retort is covered about half an inch deep with dry sand, and when the air has been expelled by car- bonic acid carefully dried sticks of phosphorus are introduced. The neck of the retort is then connected with a cooled receiver, and a rapid stream of dry chlorine led into it. The phosphorus catches fire in the atmosphere of chlorine, melts, and burns with a pale flame, forming the trichloride which distils over. The process must be so regulated that the chlorine is not in excess, partly by reducing the current of the gas and partly by heating the retort more strongly. An excess of chlorine would produce the solid pentachloride. But if too much phosphorus vapour is produced, a portion distils over into the condenser and receiver. In the pre- paration of large quantities it is best to employ several vessels generating chlorine, when the trichloride can then be distilled in an even stream. The substance so obtained is never pure ; either it contains dissolved pentachloride or free suspended phosphorus. In the latter case redistillation is sufficient to remove most of the phos- phorus : in the former case the liquid must be digested with free phosphorus, which converts the pentachloride into the trichloride, and be again distilled. The phosphorus trichloride then only con- tains small quantities of free dissolved phosphorus, for which it is a good solvent. Phosphorus trichloride is a clear, colourless liquid of 1-58 specific gravity. It boils at 76°, and its vapour has a density of 47, which corresponds to a molecular weight of 137, and its com- position PCI3. It fumes in moist air and is at once decomposed by water into phosphorous and hydrochloric acids with a consider- able rise in temperature. With chlorine it unites readily and forms the solid pentachloride. Phosphorus Pentachloride {Phosphoric Chloride). 225 PHOSPHORUS PENTACHLORIDE (Pbosplioric Chloride). Composition : PC1 5 . The solid volatile phosphorus pentachloride is easily prepared from the liquid trichloride by causing the latter to unite directly with two further atoms of chlorine. A wide-necked flask is about half filled with phosphorus trichloride on to which chlorine is led by a wide tube passing through a glass plate closing the neck of the flask. The tube must not touch the liquid or it will become closed by the solid pentachloride. The chlorine is rapidly ab- sorbed, a large quantity of heat is evolved, and the flask must be kept cool throughout the operation. It is necessary to stir the mass from time to time and finally to slightly warm it. The con- version is complete when chlorine is no longer absorbed on allow- ing the solid mass to stand in contact with the gas in closed vessels. Another method of preparation consists in dissolving phos- phorus in carbon disulphide, and then passing chlorine through this solution. Phosphorus pentachloride, which is insoluble in carbon disulphide, then separates out in indistinct crystals. Phosphorus pentachloride is a white crystalline mass fuming in the air. It sublimes without melting at ioo°, and melts under pressure at about 148. Its vapour when inhaled is exceedingly irritating, a fact which should not be forgotten in experimenting with it. All operations with it should be carried on in a good draught of air. Its vapour density is 3-6, corresponding to a molecular weight of 104, which is only one-half that required for the formula PC1 5 (208-5). The reason of this is that one molecule of the pentachloride breaks up when heated (dissociates) into one molecule of the trichloride and one of free chlorine, which occupy twice the volume that one molecule of the pentachloride would take up. On cooling again the two substances recombine and produce phosphorus pentachloride. When the pentachloride is thrown into water it is at once decomposed with a large evolution of heat, while a colourless oil collects at the bottom of the vessel and after a time disappears. This oil is phosphoric oxychloride, Q 226 Text-Book of Inorganic Chemistry POCI3, or the chloride of common phosphoric acid, PO(OH) 3 . It is decomposed by water into phosphoric and hydrochloric acids: — PC1 5 + H 2 = POCI3 + 2HCI. POCl 3 + 3H a O = PO-(OH) 3 + 3HCI. The tendency of phosphorus pentachloride to exchange two atoms of chlorine for one of oxygen makes it a valuable agent for preparing compounds of chlorine which are difficult to obtain in any other way. We have seen (p. 165) that when a mixture of phosphorus pentachloride and sulphuric acid is heated, chlorsul- phonic acid, phosphoric oxychloride, and hydrochloric acid are formed. PHOSPHORIC OXYCHLORIDE (Fbospboryl Chloride). Composition : POCl 3 . This substance, which closely resembles phosphorus trichloride in its physical properties, is gradually produced from phosphorus pentachloride when exposed to moist air. It is, however, better prepared by heating the pentachloride with certain acids or with compounds containing water chemically combined. The best substance to employ for this purpose is ordinary crystallized boric acid, which is then converted into boric anhydride : — 3PC1 5 + 2B(OH) 3 = 3POCI3 + B. 2 3 + 6HC1. Boric acid Boric anhydride Phosphorus pentachloride and boric acid are mixed in the proportions shown in the equation — i.e. about five parts of the former with one part of the latter — the mixture placed in a retort, and then heated. Torrents of hydrochloric acid are given off", and phosphoric oxychloride distils over into the well-cooled receiver. The boric anhydride, which finally remains behind in the retort, may be dissolved up in warm water, allowed to crystallize, and again used for the same purpose. Phosphoric oxychloride is a colourless liquid sinking in water, by which it is soon decomposed. It has a specific gravity of 17, boils at 110 , fumes when exposed to the air, and possesses a piercing odour. It is easily decomposed by water into hydrochloric and phosphoric acids. Corresponding to phosphoric oxychloride is a sulphur compound, phosphoric sidphochloride : PSC1 3 , which is also a colourless liquid, Bromides of Phosphorus. 227 of specific gravity 1 -6 and boiling at 125" It fumes in the air and is slowly decomposed by water into phosphoric acid, hydrochloric acid, and sulphuretted hydrogen. This compound is obtained by heating phosphorus pentachloride with certain metallic sulphides— e.g. antimonous sulphide : — 3PC1 5 + Sb,,S 3 = 3PSCI3 + 2SbCl 3 . Antimonous Antimonous sulphide chloride BROMIDES OF PHOSPHORUS. Phosphorus Tribromide (Phosphorous Bromide) : PBr 3 . This is a heavy colourless liquid, of specific gravity 2-9, boiling at 173°, and closely similar in its chemical properties to the corresponding trichloride. It is best obtained by dissolving a known weight of phosphorus in about eight times as much carbon disulphide, and then dropping in the requisite quantity of bromine — i.e. three atoms of bromine for every atom of phosphorus. As soon as the reaction has ceased and all bromine has disappeared the carbon disulphide is evaporated off on a water bath, and then the phosphorus tribromide distilled over. Phosphorus Pentabromide {Phosphoric Bromide) : PBr 5 . The pentabromide is prepared in precisely the same way as the tribromide, except that five atoms of bromine for every atom of phosphorus are employed. It is a yellow, crystalline solid, less stable than phosphorus pentachloride, decomposing even at 100° into phosphorus tribromide and bromine. When acted upon by water it behaves in a similar manner to the corresponding pentachloride. Phosphoric Oxybromide : POBr 3 , is a crystalline solid, melting at 45 and boiling at 195°. It is obtained, together with acetyl bromide and hydrobromic acid, by acting on phosphorus penta- bromide with acetic acid : — PBr, + AOH = POBr 3 + ABr + HBr. Q 2 228 Text-Book of Inorganic Chemistry. IODIDES OF PHOSPHORUS. We have already remarked that phosphorus penta-iodide is unknown ; only the tri- and di-iodides have been as yet prepared. Both are obtained in the same way — viz. adding the requisite quantity of iodine to a solution of phosphorus in carbon disulphide. Phosphorus Tri-iodide : PI 3 , is soluble in carbon disulphide, but separates out as red crystals when the liquid, after distilling off the disulphide, is placed in a freezing mixture. It melts at 55°, gives off iodine at its boiling temperature, and is easily decomposed by water into phosphorous and hydriodic acids. Phosphorus Di-iodide : P 2 Ij, can be obtained not only by the method indicated above, but also by the direct union of solid phos- phorus and iodine. For this purpose a flask is filled with carbonic acid gas and the phosphorus and iodine introduced in the correct proportions. On contact they combine energetically with one another, forming a dark-coloured liquid, which, when the reaction has been completed by slight rise in temperature, solidifies to a crystalline mass on cooling. When this mass is dissolved in hot carbon disulphide, phosphorus di-iodide is deposited on cooling as bright orange coloured tablets or prisms. The di-iodide melts at uo°to a bright red liquid, and is decomposed by water into hydriodic acid, phosphorous acid, and amorphous phosphorus : — 3P 2 I d + I2H 2 = 4HPO(OH) 2 + 12HI + 2P. COMPOUNDS OF PHOSPHORUS AND SULPHUR. Phosphorus and sulphur unite in several proportions with one another. They have a strong attraction for one another, and their union is accompanied with so considerable an evolution of heat that if large quantities are employed disastrous explosions may be produced. But if amorphous phosphorus is employed, large quantities of these compounds may be easily and safely prepared. Powdered Arsenic. 229 sulphur and amorphous phosphorus are mixed together, and the mixture heated in a closed Hessian crucible. The crucible is first gently heated, and then more strongly after the reaction has taken place. The two substances unite with one another at the moment when the amorphous phosphorus passes into the common form. A large quantity of heat is given off, but no explosion takes place. If sulphur is in excess phosphorus pentasulphide : P 2 S 6 , is produced, while with excess of phosphorus tetraphosphoric trisulphide : P 4 S 3 , is formed. Phosphorus Pentasulphide (Phosphoric Sulphide) : P 2 S 5 , may also be prepared by heating yellow phosphorus with sulphur and carbon disulphide in sealed tubes at 210° On cooling, the compound separates out in long yellow needles. It melts at 210°, boils without decomposition at about 530°, and is decomposed by water into phosphoric acid and sulphuretted hydrogen. Tetraphosphoric Trisulphide : P 4 S 3 , is easily soluble in cold carbon disulphide, from which it crystallizes in long, yellow, rhombic prisms, which melt at 166 and are only slowly decomposed by water into phosphorous acid and sulphuretted hydrogen. Among other sulphides of phosphorus may be mentioned : Phosphorus disulphide : PS 2 , or more probably P 3 S 6 , which is also prepared by heating a mixture of phosphorus, sulphur, and carbon disulphide. ARSENIC. Chemical Symbol : As— Atomic Weight : 75. Arsenic occurs in nature both native and in chemical combina- tion with other elements. Native arsenic is sometimes found in the crystalline'state, but it generally occurs in rough lumps, which easily break up into uneven laminae. More common in nature are its compounds, of which the following are the most important : Arsenical iron, FeAs 2 ; arsenical iroii pyrites or mispickel, FeSAs ; Kupfer-nickel, NiAs ; smaltine or tin white cobalt (Co,Ni,Fe) As„ ; realgar, As 2 S 2 ; orpiment, As 2 S s . Besides these, arsenic is also found in combination with oxygen as white arsenic : As 2 3 , and in 230 Text-Book of Inorganic Chemistry. the form of arsenic acid in various minerals, such as phenacolite, cobalt-bloom, mimetesite, &c. Finally, it is a very remarkable fact that arsenic, probably as arsenious acid, has been discovered in many mineral springs, in those of Ems, Kissingen, Pyrmont, Schwalbach for example. The quantity is, however, so extremely small that the physiological action of the mineral waters can scarcely be ascribed to its influence. Arsenic (commonly called ' metallic ' arsenic to distinguish it from 'white' arsenic, the oxide), when freshly broken, possesses a steel-grey colour, and strong metallic lustre. It is very brittle, and may therefore easily be reduced to powder. Its specific gravity is 57. When heated arsenic sublimes and condenses partly as an amorphous mass, and partly as rhombohedra ; it may be melted under increased pressure. The specific gravity of its vapour at about 860° is io'2, which corresponds to a molecular weight of about 300, or, in other words, its molecule in the gaseous state contains four atoms. In this abnormal property arsenic re- sembles, therefore, phosphorus and sulphur. When exposed to moist air arsenic loses its metallic lustre, and becomes dull in consequence of surface oxidation. It is insoluble in water, but if the water contains air a small quantity of the arsenic is oxidized by the oxygen of the air to arsenious acid, which then dissolves in the water. When heated in the air arsenic volatilizes and burns, forming white fumes of arsenious anhydride. At the same time an unplea- sant garlic-like odour is noticed, which is perhaps peculiar to arsenic vapour, or is perhaps due to a lower oxide. Heated in oxygen, arsenic burns with a large evolution of light and heat, also producing arsenious anhydride. Finely powdered arsenic, when shaken into a jar of chlorine, catches fire and burns to arsenious chloride. Nitric acid oxidizes arsenic to arsenious or arsenic acid accordingly as the arsenic or nitric acid is in excess. In the same way concentrated sulphuric acid also oxidizes arsenic, and is itself reduced to a sulphurous anhydride. Hydrochloric acid scarcely acts upon it. Native arsenic is never pure. It may be obtained pure by sublimation from arsenical iron pyrites, which breaks up when heated into ferrous sulphide and metallic arsenic. Arsenic is chiefly used in the arts for hardening lead in the manufacture of shot. Arseniuretted Hydrogen (Arsine). 231 ARSENIURETTED HYDROGEN (Arsine). Composition : AsH 3 . This substance, corresponding to ammonia in composition, is the only known compound of arsenic and hydrogen. In conse- quence of the weak affinities of the two elements for one another it cannot be prepared directly from its constituents. Arseniuretted hydrogen is a colourless gas, with a repulsive odour resembling garlic, and only slightly soluble in water. It is condensed at - 40 to a colourless liquid, and burns when ignited in the air with a livid flame forming water and arsenious anhydride. The specific gravity of the gas is 27, corresponding to the mole- cular formula : AsH 3 . Arseniuretted hydrogen is one of the most deadly poisons with which we are acquainted ; it is more poisonous than any other compound of arsenic. The chemist Gehlen lost his life by inhaling this gas, and fatal effects have also been produced in some other cases. Experiments with arseniuretted hydrogen must therefore be undertaken with great care, and the conditions under which the gas may be produced ought to be well known. Arseniuretted hydrogen is produced in all cases when nascent hydrogen and arsenic are brought together, especially in acid liquids. In considerable quantities it may be obtained by acting on an alloy of zinc and arsenic (prepared by adding two parts of arsenic to three of melted zinc, and pouring the alloy into cold water) with dilute hydrochloric acid. The reaction is illustrated by the following equation : — As 2 Zn 3 + 3SO.Jq^ = 3S a portion of the boron may be present as acid, and a portion as base ; but until our knowledge is further advanced no constitutional formula can be given with certainty. Boric acid, like all weak acids, can unite with strong acids, and forms compounds with sulphuric and phosphoric anhydrides having the respective composition : B„0 3 ,S0 3 and B,O 3 ,P.,O , in which the boric acid plays the part of a base. The chief use of boric acid is for the preparation of borax. It is an excellent antiseptic, and has been considerably used for this purpose in recent years. Whether, however, large quantities of boric acid can be taken into the system without injurious effects remains as yet doubtful. OTHER COMPOUNDS OF BORON. Boron Trichloride : BC1 3 . — This compound is produced when boron is heated in gaseous chlorine, or by strongly glowing an intimate mixture of boric anhydride and charcoal in a porcelain tube through which a stream of dry chlorine is passed. The reaction which then goes on is expressed by the following equa- tion : — B,0 3 + 3 C + 3CL, = 2BC1 3 + 3CO. The gaseous products are led through a (J -tube surrounded by a freezing mixture, in which the boron trichloride condenses to form a colourless liquid, boiling at 17 , and having a specific gravity Other Compounds of Boron. 265 of 1-35. This liquid fumes strongly in the air, and is rapidly de- composed by water into boric and hydrochloric acids. Boron trichloride unites with ammonia to form a solid white crystalline compound of the composition : 2BC1 3 ,3NH 3 . This compound, when heated, sublimes, and is decomposed at once by water into boric acid, hydrochloric acid, and ammonium chloride. Boron Tribromide : BBr 3 , is a colourless liquid, boiling at 90°, and with similar properties to the chloride. Boron Trifluoride : BF 3 . — This body, unlike the two other corresponding compounds, is a colourless gas, and can only be condensed to a liquid with difficulty. It possesses a piercing suffocating odour, fumes strongly in the air, and on account of its strong attraction for water, chars many organic substances like concentrated sulphuric acid. It is extremely soluble in water, one volume of this liquid dissolving 700 to 1,000 volumes of the gas. The gas may be obtained by heating an intimate mixture of 1 part of boric anhydride (or 2 parts of anhydrous borax) with 2 parts of fluor-spar and 12 parts of concentrated sulphuric acid in a glass flask. The change which then takes place is as follows : — ■ B 2 3 + 3CaF 2 + 3SO,(OH) 2 = 3S0 3 -0 2 Ca + 2BF 3 + 3H.,0. The large excess of sulphuric acid is required to absorb the water produced in the reaction. Boron trifluoride may also be prepared from boric anhydride and fluor-spar without sulphuric acid if a mixture of these two substances is heated to bright red- ness in an iron tube, the calcium then remaining behind as calcium borate : — 2B,0 3 + 3 CaF 2 = 2BF3 + B,(0,Ca) 3 . On evaporating an aqueous solution of the fluoride, boric acid separates out, and a monobasic acid remains, of the composition : HBF 4 or HF,BF 3 — -Jluoboric acid — which may be considered as a double compound of hydrofluoric acid and boric fluoride. The same substance is also more easily obtained by dissolving boric acid in hydrofluoric acid. Fluoboric acid cannot, however, be obtained in the pure state. If the aqueous solution is concentrated by evaporation, hydro- fluoric acid is evolved. Its salts, having the general composition : M'BF 4 (where M' is a monad metal), are mostly soluble in water, the most insoluble being the potassium compound : KBF 4 . On 266 Text-Book of Inorganic Chemistry. mixing aqueous fluoboric acid with a solution of a potassium compound, a gelatinous precipitate of potassium fluoborate is formed, which becomes a white powder when dried. From a hot saturated solution in water, the salt separates out on cooling in brilliant crystals. Boron Sulphide : B 2 S 3 , is a white, vitreous solid, of piercing odour, which can be fused in a stream of hydrogen. It is pro- duced by heating amorphous boron in sulphur vapour or in sul- phuretted hydrogen, and then condenses in a well-cooled receiver. Water easily decomposes it into boric acid and sulphuretted hydrogen. Boron Nitride : BN. — Boron is one of the few elements that unite directly with nitrogen. The nitride is formed when amor- phous boron is heated in a stream of nitrogen or ammonia, also when a mixture of boric anhydride and charcoal is strongly heated in a stream of nitrogen, or by glowing a mixture of dried borax and ammonium chloride. By digesting the residue obtained in the last method with dilute hydrochloric arid, the nitride remains as a white amorphous powder. Boron nitride is a stable compound, and remains unchanged when boiled with water or when glowed in the air or in hydrogen. It is, however, decomposed into boric acid and ammonia when heated to low redness in a current of steam ; the same change is produced by fusion with caustic potash. SILICON. Chemical Symbol ; Si. — Atomic Weight: 28. Silicon occurs most abundantly and is widely distributed in nature, but never in the free state. It is always found, like phos- phorus and boron, united with oxygen. Its only oxide, silica, is known both free and in combination with bases as innumerable minerals. Considering the wide distribution and great variety of the silicates, it may be asserted that silicon plays the same part in inorganic as carbon does in organic nature. Both are tetrad elements in almost all their compounds. Silicon. 267 Silicon may be separated from silica by potassium, but the decomposition is always incomplete, and the product therefore impure. It is better to employ a compound of silicon fluoride with potassium or sodium fluoride — potassium or sodium fluosilicate : K 2 SiF 6 — which can be readily obtained in the pure state. A mixture of one of these salts with sodium chloride and metallic sodium cut into small pieces is thrown into a red-hot iron crucible, which is then closed and kept for a short time at a low red heat. After cooling, the contents of the crucible are boiled with dilute hydrochloric acid. By this process fluorine is abstracted by the sodium from the double fluoride and silicon set free. Amorphous silicon prepared in this way is a dark brown powder, insoluble in water and in nitric or sulphuric acid. Aqueous hydro- fluoric acid dissolves it with evolution of hydrogen, and it also de- composes strong caustic potash, forming potassium silicate. Heated in the air it burns easily with a bright light, to form silica ; the com- bustion, however, is only incomplete, as the silica is fused by the high temperature, and so protects a portion of the silicon from the air. When heated in a stream of hydrochloric acid gas this form of silicon takes up three atoms of chlorine and one of hydrogen, and becomes converted into the substance called silicochloroform : SiHCl 3 (see sequel). If amorphous silicon is strongly heated in a crucible out of access of air, it contracts considerably, becomes of a chocolate brown colour, and now no longer catches fire when heated in the air. It has also lost its property of dissolving in hydrofluoric acid or in caustic potash after this treatment. Silicon may be prepared in the crystalline form by fusing aluminium with thirty times as much sodium fluosilicate. The black residue which remains is then treated first with concentrated hydrochloric acid to remove the aluminium, and then with hydro- fluoric acid. Another and simpler method of preparing crystalline silicon consists in throwing a mixture of three parts of potassium fluosilicate, one part of sodium in small pieces, and one part of granulated zinc, into a red-hot Hessian crucible, and then heating for some time to such a temperature that the zinc remains fused, but is not volatilized. The zinc regulus, which contains the silicon, is afterwards removed from the crucible, thoroughly boiled with water to remove the slag, and then the zinc dissolved out with hydrochloric acid. Crystalline silicon prepared in either of these ways consists either of opaque, lustrous tablets of a dark grey colour, resembling 268 Text-Book of Inorganic Chemistry. graphite, or else of brilliant prismatic crystals of an iron grey colour and considerable hardness. The specific gravity of this modification of silicon is 2-5 ; it is a conductor of electricity, and remains unaltered when heated in the air, but melts at high tem- peratures. It is neither attacked by nitric nor by hydrofluoric acid, but dissolves when heated with strong caustic potash, with evolu- tion of hydrogen, and burns in chlorine to silicon tetrachloride. COMPOUNDS OF SILICON. Silicon unites directly with oxygen and with the halogens to form compounds, in which it exists as a tetrad element. An oxide of silicon corresponding to carbonic oxide, with the composition : SiO, has not yet been prepared. Otherwise its compounds are similar to those of carbon, although the similarity is more in their composition than in their properties. The following pairs of com- pounds correspond to one another : — Silica . SiO a Carbonic acid . . C0. 2 Silicon hydride SiH 4 Methane (marsh-gas) CH 4 Silicon tetrachloride . SiCL, Carbon tetrachloride . CCL. Silicon hexachloride • Si 2 Cl 6 Carbon hexachloride . C.,Cl i; Silicon sulphide . SiS, Carbon sulphide . CS, Silicochloroform SiHCl 3 Chloroform CHCL, Silicoiodoform . SiHI 3 Iodoform . . CHI, Silicoxalic acid JSiO-OH ISiO-OH Oxalic acid (CO-OH I CO -OH SILICA (SILICIC ANHYDRIDE): SiO,, and SILICIC ACID. Silicon forms only one oxide — silicic anhydride — or, as it is briefly called, silica. This united in varying proportions with water forms different varieties of silicic acid. Free silica occurs in nature both crystalline and amorphous. In the crystalline form chiefly as quartz or rock crystal, in six- Silicic Acid. 269 sided prisms bounded at the ends by similar pyramids, and belong- ing therefore to the hexagonal system. Quartz in its purest form is transparent and colourless, but is frequently coloured brownish or violet — in the former case it is known as smoky-quartz, in the latter as amethyst. Ordinary sand consists of particles of silica, which when united together by some cementing substance constitute the different varieties of sandstone. Quartz is also contained in the free state in many important rocks —granite, for example, is made up of separate crystals of quartz, felspar, and mica. The specific gravity of this form of silica is 2-65. Crystalline silica is also found in some rocks in minute crystals belonging to the rhombic system, and has a specific gravity of 2 '3 1 , or lower than that of quartz. This variety, which is called tridy- mite, appears to be the most stable form of silica at a red-heat, as both quartz and amorphous silica pass into it when strongly heated. Besides these two crystalline forms, silica is also found in nature, usually in combination with water, in the amorphous state. The minerals agate, chalcedony, opal, jasper, flint, consist of amorphous silica mixed to some extent with the crystalline modification. Some of these — e.g. opal — are of attractive colours and are used as jewels. Agate and chalcedony are celebrated for their extreme hardness, and are highly valued for the manufacture of mortars. The siliceous sinter, deposited during the evaporation of the water from siliceous springs, and the so-called infusorial earth, or the in- organic remains of certain infusoria, consist essentially of amor- phous silica. This infusorial earth — Kieselguhr of the Germans — occurs largely in certain localities in North Germany as a light, extremely fine powder, mostly of a yellow colour. It is used for many technical purposes, especially for mixing with nitro-glycerine to form dynamite. Silica is likewise found in the animal and vegetable kingdoms. The glassy coating of the stems of certain grasses and other allied plants, cereals, rushes, and especially of the Equisetaceae (horse- tails) consist essentially of silica. The ashes of the feathers of many birds contain as much as 40 per cent, of silica. Much more widely distributed in the mineral kingdom than silica are its compounds with various bases, called silicates, and of which the greater part of the solid crust of the earth consists. Among these the different felspars, double silicates of aluminium and potassium or sodium, together with the clays, schists, &c, produced from them, take the first place. 270 Text-Book of Inorganic Chemistry. The properties of the various modifications of silica differ in many respects. All are insoluble in water and in nitric, hydro- chloric, or sulphuric acid, as well as in aqua regia, but dissolve in hydrofluoric acid, and possess in general none of the properties of an acid. They remain unchanged at the highest temperatures of our furnaces, but melt in the oxy-hydrogen flame and form a clear transparent glass on cooling. The quartz modification is very hard — it easily scratches glass — but the amorphous form is much softer. The specific gravity of the former is 2^65, that of the latter 2-20 ; the specific gravity of both changing to 2-31 when strongly glowed, owing to the conversion into tridymite. Crystalline silica is scarcely attacked even by boiling aqueous caustic potash or soda, but when fused with an excess of either of these substances in the solid form, or with their carbonates, silicates soluble in water are produced (water-glass). Amorphous silica is dissolved more or less readily by aqueous alkalies ; infusorial earth is very readily dissolved. If a concentrated aqueous solution of an alkaline silicate pre- pared by either of the above-mentioned methods is mixed with hydrochloric acid, the silicic acid separates out as a gelatinous mass, and the whole liquid becomes so viscid that the vessel may be inverted without anything running out. This silicic acid is soluble in water, especially in the presence of hydrochloric acid, for if the solution of the silicate is sufficiently dilute, hydrochloric acid produces no precipitate. It is this silicic acid which is con- tained in the water of siliceous springs, and which is deposited as siliceous sinter on evaporation. Pure amorphous silica, as a light, finely divided powder, may be easily obtained by the decomposition of silicon fluoride with water (p. 275). If this gelatinous soluble silicic acid is evaporated to complete dryness on a water-bath, it loses its water, and there remains amorphous silica, which is no longer soluble in water or hydro- chloric acid, but which easily dissolves on warming in aqueous caustic soda or in sodium carbonate. When heated to redness it is no longer soluble in aqueous alkalies, and can now only be con- verted into soluble water-glass by fusion with caustic alkalies or alkaline carbonates. Corresponding to these differences in solubility of the various forms of silica are differences in the readiness with which the various silicates are decomposed. Those which are poor in silica, Silicic Acid. 271 especially those containing- water (the zeolites), when finely divided, are decomposed by hydrochloric acid. Those, on the other hand, which contain more silica—^, the felspars — withstand the action of strong hydrochloric acid, and can be only brought into solution by treatment with hydrofluoric acid or by fusion with sodium (potassium) carbonate in a platinum crucible. In the latter case an alkaline silicate soluble in water is produced, and the bases which were previously united with the silica — e.g. lime, alumina, iron — are converted either into carbonates or into oxides soluble in acids. The fused cold mass is now digested with water and hydrochloric acid added to slight excess, when a clear solu- tion is obtained as soon as the carbonic acid has been expelled by warming. If the quantity of water used was small, the solution will contain gelatinous flocks of silicic acid. The liquid is next evaporated to complete dryness, and best heated for a short time to about 120 , to render the silica completely insoluble, afterwards moistened with strong hydrochloric acid to bring into solution any oxychlorides which might have been formed during evaporation (e.g. of iron or aluminium), and finally extracted with hot water. In this way everything goes into solution except the silica, which can be easily separated by filtration, washed and weighed, while the bases in the filtrate can be estimated by any known method. It is difficult to obtain a. silicic acid of definite composition. If the gelatinous hydrate is dried over concentrated sulphuric acid, a transparent vitreous mass is obtained, of which the composition is approximately represented by the formula : SiO(OH) 2 — i.e. a dibasic acid corresponding to the hypothetical sulphurous acid. Tetra- basic silicic acid of the composition : Si(OH). t has not yet been pre- pared. We know, however, from the composition of numerous silicates that salts of both these acids exist, as well as compounds of a dibasic disilicic acid of the possible formula : O s j ' QH To the salts of the dibasic acid may possibly belong the mineral tuollastonite : SiO • 2 Ca, to those of the tetrabasic acid, olivine : Si(0„Mg) 2 . And a combination of the dibasic and tetrabasic acids may represent serpentine : g j \ (0 2 Mg) 3 . Similarly, those silicates containing three atoms of silicon — the trisilicates — may be considered as compounds of two of these acids 272 Text-Book of Inorganic Cliemisti y. united together. Orthoclase (potash felspar : KAlSi 3 8 ) would thus be a compound of dibasic silicic acid and disilicic acid : — n (SiO- O) [SiO- O Al SiO -! n T- and prehnite : Al.,Ca.,Si 3 O n , a compound of tetrabasic silicic acid and of a hexabasic disilicic acid of the hypothetical constitution : O 1 „.L„ ! — i.e. its formula would become : — lSi(OH) 3 jSi O 1 Si Si (0 3 A1) 2 (0 2 Ca) 2 It must, however, be distinctly understood that these formula; are not the symbolic expression of established facts, but only re- present an attempt to give a simple explanation of the very various and often complex composition of the silicates. The chemical constitution of these compounds cannot be deduced simply from the results of analysis, but require extended researches similar to those which are necessary to establish the constitution of any chemical compound. SILICON HYDRIDE. Composition : SiH 4 . This compound is a colourless combustible gas, insoluble (or nearly so) in water. It may be obtained by acting upon magne- sium silicide with dilute hydrochloric acid : — SiMg 2 + 4HCI = 2MgCl 2 + SiH 4 . This magnesium silicide (SiMg 2 ) is prepared by throwing a mixture of 8 parts of anhydrous magnesium chloride, 7 parts of sodium fluosilicate, 2 parts of sodium chloride, and 4 parts of sodium in small pieces into a red-hot Hessian crucible, which is at once closed. The crude compound thus obtained always contains free magnesium, whence it follows that the silicon hydride prepared from it is always mixed with free hydrogen. In this state silicon hydride catches fire spontaneously in the air, Silicon Tetrachloride. 273 burning with a luminous flame to form silica and water. If a clean porcelain dish is depressed on the flame, brown amorphous silicon is deposited on it. Excluded from the air and heated to faint redness, the gas decomposes, like arsine and stibine, into amorphous silicon and free hydrogen. The. pure compound may be prepared by acting on an organic silicon compound of the composition : SiH(OC 2 H 5 ) 3 , with sodium. Four molecules of this compound are decomposed under the influence of the sodium, which remains itself unchanged, into one molecule of silicon hydride and three of the compound : Si(OC 2 H 5 ) 4 :- 4 SiH(OC,H 5 ), = SiH 4 + 3 Si(OC,H 5 ) 4 . Silicon hydride burns easily in chlorine, producing hydrochloric acid and silicon tetrachloride. With caustic potash, it gives a silicate of potassium and free hydrogen : — SiH 4 + 2KOH + H 2 = SiO(OK), + 4H.. SILICON TETRACHLORIDE. Composition : SiCI 4 . This, the best known chloride of silicon, is a colourless mobile liquid of 1-52 specific gravity and boiling at 58° It fumes in the air, and is energetically decomposed by water into hydrochloric and silicic acids. Silicon tetrachloride is obtained when amorphous silicon is heated in a stream of chlorine gas, or more readily by strongly heat- in^ an intimate mixture of silica and carbon in the same gas. A stiff dough is made of finely divided silica, powdered wood-char- coal, and starch paste, which is then formed into balls. These are allowed to dry in the air, and then heated to bright redness imbedded in charcoal powder. The porous mixture of silica and carbon so obtained is introduced while still hot into a dry porcelain tube, which can be heated to bright redness in a tube-furnace, while a current of dry chlorine passes through it. The reaction is expressed by the equation :— Si0 2 + 2C + 2C1 2 = SiCl 4 + 2CO. The other end of the tube communicates air-tight with a T 274 Text-Book of Inorganic Chemistry. U-tube placed in a freezing mixture in which the silicon tetra- chloride condenses, while the carbonic oxide and excess of chlorine are carried away to the open air. At the bottom of the U-tube is a small vertical tube (provided with a glass stopcock), through which the silicon tetrachloride can be drawn off into bottles or sealed tubes. The corresponding silicon tetrabromide, SiBr 4 , and the letrio- dide, Sil 4 , are prepared in exactly the same way as the chloride. The bromide is a colourless liquid, boiling at 153° and solidifying to a crystalline mass at -12° The iodide is a crystalline solid, boiling at about 290° Silicon Hexachloride : Si 2 Cl 6 . — Besides the above-mentioned tetrachloride of silicon, a second compound — silicon hexachloride : Si 2 Cl 6 — is also known. This is also a colourless liquid, but has a specific gravity of 1-58, boils at 145°, and solidifies at about -4°. It is obtained by heating silicon in the vapour of silicon tetra- chloride or by heating a mixture of silicon tri-iodide and mercuric chloride. At 350 it decomposes into the tetrachloride and free silicon, and this decomposition increases up to 8oo°. But at about i,03o°the hexachloride is again stable — i.e. silicon and the tetra- chloride again unite to form the hexachloride. Water decomposes it into hydrochloric acid and disilicic or silico-oxalic acid : j ~.„ „„ The corresponding silicon hexiodide is a crystalline solid, ob- tained by heating the tetriodide with finely divided silver. Silver iodide is then formed with the hexiodide. SILICON FLUORIDE. Composition : SiF 4 . This substance is a colourless gas, fuming strongly in the air, and with a suffocating odour. Its specific gravity is y6, and it is at once decomposed by water with separation of gelatinous silicic acid. It is produced when hydrofluoric acid acts upon silica in the presence of de-hydrating substances. For its preparation a mixture of powdered fluor-spar and fine Fluo silicic Acid. 275 dry sand, the latter in considerable excess (about equal parts), is heated in a thick glass flask with a large excess of concentrated sulphuric acid, so that a thin liquid of the consistency of cream is produced. Silicon fluoride is then evolved, according to the equa- tion : — Si0 2 + 2CaF 2 + 2SO,(OH) 2 = SiF 4 + 2 S0 2 (0 2 Ca) + 2 H.,0, while the water produced at the same time is absorbed by the excess of sulphuric acid. If the quantity of sand employed in the experiment is insufficient, the glass flask is strongly attacked by the hydrofluoric acid, and may be completely eaten through. FLUOSILICIC ACID. Composition : H 2 SiF 6 . This dibasic acid, which may be considered as a double fluoride of silicon and hydrogen, is produced by the union of silicon fluoride with nascent hydrofluoric acid. These conditions are fulfilled when silicon fluoride comes into contact with water. Decomposition then occurs according to the equation : — SiF 4 + 3H 2 = SiO(OH) 2 + 4HF. or, SiF 4 + 4H 2 = Si(OH) 4 + 4HF. The silicic acid separates as a gelatinous mass, and the hydro- fluoric acid unites at once with another portion of silicon fluoride, forming fluosilicic acid, which remains in solution and can after- wards be separated by filtration. The whole reaction may, there- fore, be thus represented : — 3SiF 4 + 3H..0 = SiO(OH) 2 + 2H 2 SiF 6 , or, 3SiF 4 + 4H 2 = Si(OH) 4 + 2 H 2 SiF a . The glass tube leading the silicon fluoride into the water, even when wide, becomes easily stopped up by the deposited silicic acid, and so an explosion might be produced. In order to avoid such an occurrence a little mercury is placed at the bottom of the vessel in which the fluosilicic acid is to be produced. The tube from the flask evolving the silicon fluoride is made to dip under this mercury, t 2 276 Text-Book of Inorganic Chemistry. and the vessel then filled up with water. The decomposition of the silicon fluoride then only commences above the level of the mercury. The danger of an explosion may also be avoided by connecting the end of the delivery tube with the neck of a glass funnel, and dipping the large open end in the water, so that the gas comes at once into contact with a large surface of water. The aqueous fluosilicic acid, when separated from the silicic acid by filtration, is a strongly acid colourless liquid. It may be concentrated up to a certain point by evaporation at the ordinary temperature of the air. But if this limit is exceeded, or if the liquid is heated, it decomposes into silicon fluoride and hydro- fluoric acid, both of which volatilize. On heating, therefore, a solution of fluosilicic acid in a platinum basin, ultimately nothing remains. By neutralizing fluosilicic acid with bases or carbonates, its salts — the fluosilicates — are obtained, having the general composi- tion : R',SiF 6 , where R' is one atom of a monad metal. These salts are mostly soluble in water, except the barium compound : BaSiF 6 , and those of potassium and sodium. As the strontium salt is easily soluble in water, fluosilicic acid has been used as a means of separating barium and strontium. In concluding this account of the compounds of silicon, brief reference may be made to certain substances which are of interest on account of analogy with similar compounds of carbon usually included under the head of organic chemistry (compare p. 268). In lack of suitable names for these compounds they are usually called after the corresponding carbon compounds ; thus the body corre- sponding to chloroform (CHC1 3 ) is called silico-chloroform. siiico-chioroform : SiHCl 3 , is produced when amorphous silicon is heated in a stream of dry hydrochloric acid gas. At the same time a small quantity of the less volatile tetrachloride is also produced, which can be separated by fractional distillation. This compound is a colourless inflammable liquid, fuming in the air, with a specific gravity of I '65 and boiling at 36 . Chlorine converts it into silicon tetrachloride and hydrochloric acid, and water, even at o°, into hydrochloric acid and a. compound of the composition : Si 2 H 2 3 , corresponding to the hypothetical formic anhydride : SpX O. Carbon. 277 Silico-formic Anhydride : h c|q O, produced as just men- tioned, is a white powder, which when heated in the air burns easily to silica and water. Caustic soda easily dissolves it, with evolution of hydrogen, forming sodium silicate. silicoxallc Acid : ] cjq.oh ' s obtained as a white powder by decomposing silicon hexachloride with water, and is dissolved by caustic soda to form sodium silicate. In composition it corresponds (CO • OFT CO OH w ' t ' 1 w hich, however, it has no other similarity ; it does not even possess the properties of an acid. CARBON. Chemical Symbol: C. — Atomic Weight : 12. Carbon occupies a peculiar position among the other elements. Although it only combines directly with few other elements — with oxygen, sulphur, hydrogen, and iron, but neither with nitrogen nor the halogens — the compounds of carbon far exceed those of any other element both in number and variety. All organic com- pounds, including the constituents and products of all animal and vegetable organisms, as well as the numerous substances which can be artificially prepared from them, are compounds of carbon. Organic chemistry has been defined as the chemistry of the carbon compounds, according to which a description of these sub- stances should find no place in a text-book of inorganic chemistry. But sharp lines of demarcation, which so completely separate sub- stances from one another as to make a strict classification possible, are unknown • in nature. Connecting links are everywhere to be found showing a gradual transition from one class of substances to the other. And although all organic compounds contain carbon, this by no means compels us to refer all carbon compounds to organic chemistry ; many of the carbon compounds, especially those of simpler composition, cannot be easily separated from the inorganic compounds of other elements. Carbon occurs free in nature in three modifications, as different 278 Text-Book of Inorganic Chemistry. from one another in their physical properties as could be imagined. We know it, firstly, as the diamond, crystalline, transparent, cha- racterized by its great hardness and high index of refraction for light, and a non-conductor of electricity ; secondly, as graphite, with a dark grey colour, crystalline, lustrous, a conductor of elec- tricity, and extremely soft ; and, lastly, as charcoal, the chief con- stituent of common coal, amorphous, black, not much harder than graphite, and a non-conductor of electricity. The diamond is among the rarest of substances and is the most precious of gems ; graphite is found in large quantities, but usually in isolated spots ; amorphous carbon, as the various forms of coal, occurs widely distributed and in immense quantities. The last- named variety, although by far the cheapest, is still by far the most valuable and important. It furnishes the most important source of our artificial heat, light, and power, and influences not only the industrial, but also the entire social and political life of the civilized countries in which it occurs. Like their other physical properties, the specific gravity of the three modifications of carbon is very different. That of the diamond is 3'5, of graphite 2 '2, and of amorphous carbon r6 to 2-o. Greater similarity, however, exists between the chemical pro- perties of the three forms. Although they burn in oxygen or air with different degrees of readiness, they all produce the same pro- ducts of combustion. Equal weights of either modification always produce equal weights of carbonic acid. And by the reduction of carbonic acid, carbon is again formed, but only amorphous carbon or graphite, never diamond. All attempts to prepare the diamond by the reduction of carbonic acid or by any other method have, as yet, been unsuccessful. Carbon in all three forms is infusible at all temperatures, but volatilizes at the high temperature of the electric arc ; it is insoluble in any of our ordinary solvents. The only solvent at present known is fused iron, which, however, only dissolves I or 2 per cent, of it. On cooling, most of the carbon separates out in the form of graphite. Should a liquid ever be discovered capable of dissolving carbon at not too high a temperature, the carbon might be again separated in the diamond form. As was stated above, carbon does not unite directly with nitro- gen nor with the halogens, and carbon compounds of these sub- stances can only be obtained indirectly. Even with oxygen, Carbon. 279 sulphur, and hydrogen, carbon has not the slightest tendency to combine, either at the ordinary temperature or at 100". But when heated to redness its affinity for oxygen and sulphur gradually in- creases, so that at a white heat the affinity of carbon for oxygen and sulphur even surpasses that of phosphorus and potassium. Carbon when heated to redness is, therefore, one of the most powerfully reducing substances known, and a very extended use is made of this property in the extraction of metals from their ores. Hydrogen only unites with carbon at the extremely high tempe- rature of the electric arc. When this is produced between carbon poles in an atmosphere of hydrogen, small quantities of a com- pound of the two bodies, called acetylene : C 2 H 2 , are produced. Under certain circumstances carbon can also unite directly with nitrogen when a substance — e.g. potassium — is present with which the compound of the two elements (cyanogen) can combine. From carbon, nitrogen, and potassium carbonate we can thus obtain potassium cyanide. The diamond is always found in alluvial deposits, which have been produced by the disintegration of older rock masses, and from which it is obtained by a process of washing. Diamond-fields, or districts which yield diamonds, occur especially in Brazil, India, and at the Cape, as well as in California, Borneo, and in the Ural Mountains. Diamonds have also been found in Brazil in a matrix of a variety of talc schist rich in quartz — termed itacolinnite, and distinguished by the ease with which thin strips of it may be bent. But whether this is the original matrix of the diamond, or whether the itacolumite was built up from older rock with the diamonds, remains as yet undecided. The diamond crystallizes in the regular system, and its faces are often curved. The crystals may be readily fractured parallel to the faces of the octahedron ; it is distinguished by its great hardness and bright lustre. Diamonds, are usually transparent and colourless, or faintly yellow.- They are sometimes found of a bright yellow or brown" colour, and sometimes even blue, green, and black. Diamonds used as jewels always require cutting and polishing, which, on account of their great hardness, can only be done with diamond dust. Small diamonds useless for other purposes — 280 Text-Book of Inorganic Chemistry. diamond-boart — are crushed in a steel mortar. Notwithstanding the great hardness of the diamond, it is tolerably brittle, and can therefore be easily reduced to powder. The form which the diamond receives when ground for a gem depends partly upon its original form, and partly upon the use to which it is to be put. The most valuable form is that known as the brilliant (fig. 52), which may be described as two cones placed base to base, of which the upper one is cut off at about half its height. Both pyramids are made up of a number of triangular or oblong Fig. 52- Fig- S3- faces (facets). The brilliant form is always held by its centre parts in the setting, so that both upper and lower pyramids are free. Differing from this is the table or rosette form (fig. 53), consisting of a flat base, on which rises a circle of oblong facets, supporting a hexagonal pyramid. The value of a diamond depends partly on its size, partly on its transparency (water), and partly on its play of colours. The weight of a diamond is expressed in carats (0-205 gramme, or 3-165 grains), but the value of a diamond, other things being equal, is not pro- portional to its weight, but to the square of its weight. Thus, if one diamond weighs ten times as much as another, its value will be approximately a hundred times as much. Large diamonds of 100 to 300 carats (20-5 to 61-5 grammes) are very rare, and are mostly crown jewels. To these belong the so- called Star of the South, found in Brazil, and weighing 127 carats, and the Koh-i-Noor, of 106 carats, found in India, and one of the English crown jewels. Of the two largest diamonds known, one weighs 194 carats, is of a yellowish colour, and is fixed at the end of the Russian sceptre ; the other, weighing 277 carats, is in the possession of the Nizam of Hyderabad. Of the formation of the diamond nothing for certain is known, except that it is not produced at a high temperature. Its optical Carbon. 281 properties and microscopical observations seem to show that the diamond is of organic origin. All attempts to prepare diamonds artificially have as yet given no certain result, although it is always more probable that a method might be discovered for converting amorphous carbon or graphite into diamond, or of preparing the diamond in some other way, than that the problem of the alche- mists — the transmutation of common metals into gold — should ever be solved. The present high price of diamonds is due to their scarcity. Should larger quantities of diamonds ever be discovered, or should it ever be possible to prepare them artificially, their value would be considerably diminished. On account of its hardness, the diamond is largely used for cutting hard stones, and for scratching and cutting glass. For simply writing on glass, a splinter of diamond is sufficient, but if the glass is to be cut, a deeper scratch is required, and a natural convex edge of the diamond is necessary. The diamond requires a higher temperature for its combus- tion than amorphous carbon, but burns easily in the oxy-hydrogen blowpipe with excess of oxygen, or when raised to redness by a spiral of platinum wire connected with a galvanic battery in an atmosphere of oxygen. In the finely divided state the diamond burns readily in the air when heated on a piece of platinum foil before the blowpipe. In absence of oxygen, for example, in an atmosphere of hydrogen, the diamond may be heated to a high temperature without any appreciable change. Graphite, plumbago, or blacklead, occurs in nature in nodular or columnar masses, and occasionally in small tablets belonging to the hexagonal system. It is of a dark grey colour, opaque, a good conductor of electricity, feels greasy to the touch, and is so soft that it easily marks paper. Its specific gravity varies from 2-5 to 1 -8, according as it is more or less pure. Graphite usually occurs in isolated spots, but often in considerable quantities. The mines of Cumberland and Westmoreland are the oldest, but are now nearly exhausted. It is also found in Passau and other parts of Germany and Austria, in Bohemia, Greenland, Sicily, and Ceylon, and, in large quantities of great purity, in California. Graphite may be artificially prepared by solution of carbon in molten iron, when it separates out on cooling. Finely powdered and pure graphite can be easily compressed 282 Text-Book of Inorganic Chemistry. at a high temperature to a compact mass of the same specific gravity as that occurring in nature. It is only necessary to remove air as much as possible by an air-pump before compressing. Thin plates or threads for black-lead pencils can be easily sawn from the blocks obtained by this process. Compact graphite burns almost as difficultly and sometimes more difficultly than the diamond. If finely powdered graphite is mixed with three times its weight of powdered potassium chlorate, and with enough concentrated nitric acid to make a thin paste, then heated for several days to 6o°, or placed in direct sunlight until yellow vapours are no longer given off, it becomes changed to a bright yellow, transparent, crystalline substance of the composition : C n H 4 5 . By repeating the process several times the whole may be changed into this sub- stance. It is slightly soluble in water and possesses a faint acid reaction, dissolves readily in aqueous alkalies, and in general has all the properties of an acid. This peculiar acid is called gra- phitic acid. The graphite which occurs in nature is never pure, but contains smaller or larger quantities of inorganic constituents, which remain * behind as ash on burning. Graphite is of considerable value in the arts. It is used for the manufacture of black-lead pencils, and of crucibles which are to withstand a high temperature, and which are made of a mixture of powdered graphite and clay. It is further employed for polishing the grains of gunpowder, for coating moulds of gutta-percha or plaster to make them conductors of electricity in the process of electrotyping, and finally for blacking iron grates and other articles of the same metal, which by receiving a coating of the material are protected to some extent from rusting. Amorphous carbon is produced from organic substances when they are heated to redness, or when they decay out of contact with the air. Wood charcoal may be chosen as representative of amorphous carbon which has been produced by the first-named process. As its name implies, this substance is obtained by char- ring wood out of contact with the air, and is prepared either by piling logs of wood together, covering with sods, and then slowly burning them, or else by heating wood or sawdust in iron retorts. By the first process only charcoal is obtained ; but by the second and more modern, various volatile products of considerable value, Carbon. 283 such as acetic acid (pyroligneous acid), wood naphtha, creosote, &c., are obtained. According to the kind of wood used the charcoal may be more or less porous, and therefore of apparently different specific gravity. The different varieties of wood charcoal are employed for different purposes ; large quantities are used for the manufacture of gunpowder. Charcoal is also used (largely in some countries) as a fuel. When heated to redness in the air, it burns without smoke or flame, and produces a considerably higher temperature than the same weight of wood. Wood charcoal is extremely porous, and possesses the property of absorbing considerable quantities of atmospheric air. Gases which are more easily lique- fied than common air — for example, ammonia, sulphuretted hydro- gen, sulphurous anhydride, and carbonic acid — are absorbed by charcoal in even greater quantities. It thus possesses the power of absorbing noxious gases from the air, and purifies or disinfects the air. Organic substances which may be suspended or dissolved in drinking water and impart a disagreeable taste or smell are removed when such water is shaken up with charcoal, or when the water is passed through a filter of powdered charcoal. The water which runs through such a filter is usually clear, odourless, and pleasant to the taste, and charcoal filters are therefore often employed to purify drinking water. Colouring matter of various kinds is easily removed from a liquid by amorphous charcoal — claret, for example, is easily de- colorized. This property is possessed in a much higher degree by the charcoal obtained from animal substances — called animal char- coal — than by wood charcoal. Afiimal charcoal — known also as bone-black, animal-black, ivory-black, &c. — is obtained by glowing blood, bones, and other animal refuse, excluded from the air. It contains all the inorganic substances — principally calcium phosphate and carbonate — present in the substances from which it was prepared. These two com- pounds can be readily extracted by warming the charcoal with hydrochloric acid, when, after washing with water, a form of amor- phous carbon remains, which far exceeds wood charcoal in porosity and absorption power. Animal charcoal purified in this way is largely used for many technical purposes, especially to decolorize the syrup in sugar refineries. Wood charcoal' contains less. inorganic impurity than animal 284 Text-Book of Inorganic Chemistry. charcoal, but is still far from being pure carbon. When burnt in the air it leaves behind a considerable ash, rich in potas- sium carbonate. Pure amorphous carbon may be obtained by glowing certain organic substances free from nitrogen and inorganic impurities. Pure crystallized sugar or tartaric acid, for example, when glowed in a platinum basin and afterwards heated strongly in a covered platinum crucible, yields a very pure form of amor- phous carbon. Impure amorphous carbon is found in nature as the various forms of coal which have been produced by the decay of vegetable matter out of contact with the air, and sometimes also by elevation in temperature and strong pressure. All varieties contain, besides inorganic constituents — their ash which they leave behind on burn- ing — varying quantities of hydrogen, oxygen, nitrogen, and sulphur, which are mostly driven off when the coal is heated in closed vessels. The heavy form of carbon produced by charring coal in closed vessels is called coke. Coke is an excellent fuel, and pro- duces, when freely supplied with air, a higher temperature than ordinary coal. It is very largely used for smelting iron and other ores. Coke is of a dark grey colour, lustrous, and a conductor of electricity. The different varieties of coal may be classified according as the charring process which they have undergone has been more or less complete — i.e. according to their relative proportions of carbon, hydrogen, and of oxygen and nitrogen. Coal is usually black in colour, mostly of slaty fracture and fatty lustre, and when heated in the air generally evolves combustible gases and therefore burns with a flame. Anthracite is the most highly carbonized form of coal. It is distinguished by its hardness and conchoidal fracture ; it pro- duces very little gas when heated, and therefore burns almost without flame. Anthracite is found in considerable quantities in South Wales and in Pennsylvania, and is highly valued as a smoke- less steam-coal and for metallurgical purposes. It does not soil the fingers like common coal. Bituminous coal in its varied forms comes next to anthracite in percentage of carbon. All the commoner kinds of coal belong to this class. It contains varying proportions of carbon and hydro- gen, but always more hydrogen and less carbon than anthracite. Bituminous coals rich in hydrogen evolve considerable quantities of gas when heated, and are therefore used for the manufacture of Carbon. 285 coal-gas. A variety rich in volatile matter, and especially suited for gas making, is cannel coal. This is much more compact in structure than ordinary bituminous coal, and, like anthracite, has a conchoidal fracture, but its surface is dull, not bright and lustrous like common coal and anthracite. Brown coal or lignite is always of more recent formation than ordinary coal. It is mostly soft and of a brown colour, and often exhibits distinct traces of woody structure not to be found in the older coals. It contains considerably more hydrogen, oxygen, and nitrogen than ordinary coal and less carbon. Jet is a compact, black form of lignite, hard enough to take a high polish, and largely used for ornamental articles. Peat or turf is produced in marshy localities by the gradual decay of grasses, mosses, and other marsh plants. It contains large quantities of inorganic constituents, and leaves a large amount of ash when burnt. It is of much less value as a fuel than any of the varieties of coal mentioned above. The following table l gives the approximate percentage compo- sition of the organic parts of wood, peat, and the different kinds of coal, omitting the ash in all cases : — Oxygen, Carbon Hydrogen Nitrogen, and Sulphur Oak . 5° -I 7 6-o8 4375 Peat .... S9-83 578 34 '39 Lignite (Bovey Tracey) 67 '90 573 26-37 Staffordshire coal . 79 '37 537 15-26 Wigan cannel coal . 82-47 5'67 n-86 Newcastle caking coal 87'9S 5'34 671 South Wales anthracite 9 r8 S 3 '35 4-80 Amorphous carbon in a finely divided state is obtained as soot or lampblack, by the incomplete combustion of volatile organic compounds, rich in carbon, such as turpentine, benzol, or wood containing turpentine and resin. If a cold body is placed in the flame of a candle or oil-lamp, it soon becomes covered with soot or finely divided carbon, just as arsenic is deposited from the flame of burning arsine. Lampblack is used in a more or less pure state for various purposes — the finer varieties for preparing Indian ink, the more coarse as a black paint and for printer's ink. Unless purified by special processes, it always contains varying quantities of hydrocarbons mixed with it. : From Watts's Dictionary of Chemistry. — Ed. Text-Book of Inorganic Chemistry. COMPOUNDS OF CARBON. No element exists in so large a number of compounds as carbon. By far the largest proportion of these belongs to organic chemistry, and we should be justified in referring all compounds of carbon to this division, if it were possible to separate inorganic and organic compounds distinctly from one another, and if it were not abso- lutely indispensable to describe some of the more simple of the carbon compounds under inorganic chemistry. To these belong carbonic acid and oxide, oxalic acid, carbon disulphide and oxysulphide, carbon tetrachloride, marsh gas and cyanogen, together with hydrocyanic acid, the cyanides, and a few other compounds. CARBONIC ANHYDRIDE, or CARBONIC ACID. 1 Composition : CO. r Carbonic acid, which is always produced when carbon in any form or any compound of carbon is burnt with a free supply of air, is a colourless gas, of peculiar prickly smell and faintly acid taste. It occurs very widely distributed in nature, both free and com- bined with various bases. It is a normal constituent of atmo- spheric air (p. 200), and is contained in most natural waters in the free state. In immense quantities combined with lime it forms the different natural varieties of calcium carbonate — common limestone, calcspar, arragonite, marble, chalk, &c. ; further, as magnesium car- bonate in magnesite, as double carbonate of calcium and magnesium in dolomite, as ferrous carbonate in spathic iron ore, as barium and strontium carbonates in witherite and strontianiie respectively, and, finally, as sodium carbonate in trona or soda — to mention only the more important compounds. In many places, carbonic acid is evolved in considerable quan- tities from clefts in the earth, usually in the neighbourhood of active or extinct volcanoes. The gas is, of course, invisible, but can be easily recognized by chemical means. Among the most 1 True carbonic acid, which would result from the union of the anhydride with water, is unknown, and the anhydride (C0 2 ) is therefore v.sually called carbonic acid, for the sake of brevity. — Ed. Carbonic Acid. 287 celebrated localities of these carbonic acid springs are the Grotto del Cane, near Naples, those in Pyrmont, and in various parts of Germany, and the well-known poison valley of Java. Generally, the carbonic acid carries water with it in its upward course to the surface of the earth, and often with considerable force, owing to the pressure of the gas. At Nauheim a spring rises 80 feet high, probably owing to the force with which the compressed carbonic acid escapes. Finally, carbonic acid is produced in such large quantities during the fermentation of sugar in the manufacture of alcoholic liquors, that attempts have been made to utilize the gas — for ex- ample, in the manufacture of white-lead. It is thus clear that there is no lack of material from which carbonic acid may be prepared ; but if the gas is wanted pure, the choice of substances becomes more limited. All limestones which contain traces of organic matter when treated with strong acids yield carbonic acid with an unpleasant odour and taste, because small quantities of these organic substances are carried away with the gas. And carbonic acid prepared from such material should therefore never be used for the manufacture of mineral waters. On the other hand, the gas evolved from calc-spar, marble, magnesite, an d many kinds of dolomite and limestone, as well as from acid sodium carbonate, is always pure and free from any foreign smell or taste. For the preparation of small quantities of carbonic acid Kipfs apparatus is best adapted, especially as with this arrangement a stream of pure carbonic acid can be obtained at any time by simply opening a stop-cock. Carbonic acid may be evolved from mag- nesite or acid sodium carbonate by means of dilute sulphuric acid, and need then only be dried (when necessary) by passing through concentrated sulphuric acid. To prepare carbonic acid from dolo- mite or calcium carbonate, it is better to use hydrochloric acid, as sulphuric acid produces difficultly soluble calcium sulphate, which surrounds the pieces of marble or dolomite and tends to prevent their contact with the acid. If sulphuric acid is preferred, the limestone or dolomite must be finely powdered, and the mixture continuously agitated. The process is represented by the following equation ; — C0 2 Ca + S0 2 (OH) 2 = SGy0 2 Ca + H 2 + C0 2 . When carbonic acid is liberated from limestone or marble with dilute hydrochloric acid, calcium chloride then remains behind in solution, and the equation becomes : — C0 2 Ca + 2HCI = CaCl 2 + H 2 + C0 2 . 288 Text-Book of Inorganic Chemistry. The gas prepared in this way carries some of the volatile acid with it, and should be freed from this by passing through several wash-bottles filled with water. Carbonic acid may be collected over mercury or over water, in which it is only slightly soluble. At the ordinary temperature water dissolves about its own volume of the gas. Since, however, carbonic acid is considerably heavier than air, vessels are best filled with the gas, especially when required dry, by leading the gas with a long glass tube to the bottom of the vessel to be filled (fig. 54). The carbonic acid then gradually displaces the lighter air until in a short time the whole vessel is filled with the gas. Fig. 54. Carbonic acid has a specific gravity of 1-52, corresponding to a molecular weight of 1-52 x 28-88 = 44, and its molecule is therefore C0 2 . It is thus considerably heavier than air, and diffuses from an open jar so slowly that it can be poured like water from one vessel to another. If two jars are taken, one filled with carbonic acid and the other with air, the carbonic acid can, with care, be poured almost completely from one into the other, as may be easily shown by introducing a burning taper. Carbonic acid neither burns nor supports combustion. A burning candle is at once extinguished in a jar of the gas ; even if Carbonic Acid. 289 the air only contains a small quantity of carbonic acid, a candle ceases to bum. Respiration in such an atmosphere is also impos- sible, and carbonic acid, without being actually poisonous, is fatal to the life of animals, because it takes the place of the oxygen necessary to aerate the blood. Death therefore ensues simply from suffocation. That carbonic acid is not a true poison is at once evident from the following considerations : we continuously inhale carbonic acid, largely diluted it is true, into our lungs ; carbonic acid is continually produced in considerable quantities by the slow combustion going on in our bodies, and forms a large proportion of the air exhaled from our lungs ; finally, continued use of effer- vescing drinks, rich in carbonic acid, does not apparently produce any injurious effect. Nitrogen, like carbonic acid, neither burns nor supports com- bustion, nor can it be respired ; the latter gas may, however, be distinguished from the former by the fact that when brought into contact with clear lime-water it at once produces a white tur- • bidity or precipitate of calcium carbonate, insoluble in water, and which again dissolves with effervescence in dilute hydrochloric acid. Baryta-water is even better adapted than lime-water for the detection of carbonic acid. Carbonic acid possesses the properties of a weak acid anhy- dride, and unites therefore with bases. The combination, however, takes place slowly, and but little heat is evolved. If a glass tube closed at one end is filled with carbonic acid and the open end placed in a vessel of caustic potash, the liquid rises in the tube slowly, more quickly if shaken, until ultimately the whole of the carbonic acid disappears. We employ this property of the gas not only to separate it when mixed with other indifferent gases, but also to determine its amount quantitatively. The quantity of carbon present in an organic substance is found by heating it with some substance that readily gives up oxygen (e.g. copper oxide) in a tube of hard glass. By this process the hydrogen present is converted into water, and the carbon into carbonic acid. These compounds are allowed to pass first through a U-tube containing calcium chloride to absorb the water, and then through a series of bulbs (called Liebig's potash-bulbs, fig. 55), which contain strong caustic potash, and absorb the carbonic acid. If the calcium chloride tube and the potash bulbs are separately weighed before and after the experiment, the increase in weight of the former gives the weight of water formed, and of the latter the U 290 Text-Book- of Inorganic Chemistry. weight of carbonic acid. From these weights the actual weights of hydrogen and carbon present in the substance can be easily calculated. The affinity of carbon for oxygen increases as the temperature rises, and at a white heat carbon can even exlract oxygen from potassium oxide or carbonate and set free metallic potassium. But at a red-heat, and when the carbonic acid is in excess, potassium can abstract oxygen from carbonic acid. If a small piece of potassium is heated in a bulb-tube of hard glass to low red- ness while a stream of dry carbonic acid is passed over it, the fused metal soon begins to burn with a dark red light, and becomes covered with a coating of potassium carbonate and amorphous The following equation explains the process ; — 4K + 3CO,, = 2C0(0K) 2 + C. This experiment shows that colourless carbonic acid contains black charcoal, and that it can be reduced to this substance. Carbonic acid is the anhydride of an unknown dibasic acid, [OH Fig. 55- carbon. the composition of which would be : CO J qtt but which has not yet been prepared. The majority of its salts are insoluble in water ; the only soluble carbonates being those of the alkali metals : potassium, sodium, lithium (ammonium, rubidium, and caesium), and of thallium. Normal potassium carbonate (potashes) and normal sodium carbonate (soda) possess an alkaline reaction, because the strong basic properties of these oxides are not completely neutralised by the weak carbonic acid. Its acid salts of these ( OT-T two metals — e.g. acid sodium carbonate : CO q N — react neutral. When carbonic acid is led into aqueous ammonia the chief pro- duct is ammonium carbonate, of the composition : CO J J^xtS 4 : — CO, + 2NH3 + H s - COJg^JJ* This salt cannot, however, be produced from dry ammonia and dry carbonic acid, because the molecule of water necessary Carbonic Acid. 291 for its production is now absent. The two gases unite, however, with considerable evolution of heat, and produce a salt, which is ammonium carbamate, having the composition: CO] ON ?j (p. 186). Carbonic acid can be condensed by a high pressure to a liquid which, when allowed to evaporate in the air, produces a sufficient diminution in temperature to freeze it. For the preparation of liquid carbonic acid, the apparatus \ devised by Natterer is now almost universally em- ployed. This consists of a compression pump, which also acts as a suction pump, and a strong wrought-iron vessel. The latter vessel (fig. 56) is screwed at its base to the top of the cylinder of the pump, and is there provided with a small valve opening upwards at the end of every stroke of the piston, and otherwise kept closed by a weak spiral spring. This valve, therefore, only allows gas to enter the strong vessel, and prevents any from leaving it. At the upper end of this vessel, a narrow tube, n, is screwed on at right angles, which can be set in communication with the interior of the vessel by raising a conical valve attached to the screw, t. This tube is to allow the liquid carbonic acid to escape after about 250 grammes have been pumped in. For this purpose, the receiver is unscrewed, turned upside down, and the screw, t, carefully turned. If a jet of liquid carbonic acid is allowed to flow out of the side tube into a metallic vessel, which is so arranged that it receives a circular motion, and therefore rapidly evaporates, a large portion of the carbonic acid freezes at the low tem- perature produced, and fills the vessel as a light, snow-like mass. Condensed carbonic acid is a colourless, transparent, mobile liquid of 0-95 specific gravity at 0°, and only 078 at + 25 . It therefore expands under the influence of heat not only greater than any other liquid, but also greater than the gases — a very remarkable property. The tension of liquid carbonic acid at different temperatures is about as follows : — Fig. 56. 292 Text-Book of Inorganic Chemistry. At -79° . . i.i atmospheres „ -3o° • - 18 3? „ -io° . . 27 3? . °° . . • • 36 33 „. + IO° . . 1 46 3; ,. +30° . . 74 33 ., +35° ■ • 80 33 During free evaporation in the air its temperature falls to about -79°. The solid, snow-like carbonic acid evaporates but slowly even on a warm hand, and notwithstanding its low temperature produces scarcely a feeling of cold. It does not, in fact, touch the hand at all, but is separated by a layer of the gas which it is continually giving off. The phenomenon is similar to that produced when a drop of cold water is allowed to fall into a red-hot platinum dish ; the drop then remains suspended by its own vapour and does not touch the hot metal at all. If, however, a piece of solid carbonic acid is pressed firmly on the skin, a burning pain is felt and a blister is raised on the spot touched. The action is, in fact, the same as when the skin is burnt with a hot body. If solid carbonic acid is mixed with some volatile liquid — e.g. ether — both evaporate with great rapidity, the mixture begins to boil, and a sufficiently low temperature (about — ioo c ) is produced to freeze mercury in a few seconds. Carbonic acid is largely used for the manufacture of artificial mineral waters. Pure water in which various salts are dissolved, according to the mineral water required, is impregnated with carbonic acid in thick metallic vessels. The gas is pumped in by a special apparatus, and the liquid is kept constantly agitated. In this way not only can many mineral waters be imitated, but other compositions of any required kind can be prepared. Water in which small quantities of salts have been dissolved retains its carbonic acid when exposed to the air longer than pure water. Carbonic acid is also used in the arts for the preparation of white-lead, and of salicylic acid from phenol. Liquid carbonic acid is now employed on a large scale by Krupp at Essen, in Ger- many, for compressing cast-steel during solidification. Carbamic Acid, 293 CARBAMIC ACID. Composition of the hypothetical acid: CO j W^ 1 As previously mentioned (p. 290), dry carbonic acid and dry ammonia unite to form a salt which cannot be ammonium car- bonate, as the molecule of water necessary for the formation is wanting. The compound produced is the ammonium salt of a monobasic acid, which is as little known in the free state as the true carbonic acid : CO X „,,. It has the same relation to this dibasic acid as 183- ■• f CI chlorsulphonic acid (p. 165): SO^Jqit has to sulphuric acid, except that carbamic acid contains an atom of amidogen, instead of one of chlorine, in place of an atom of hydroxyl. The production of ammonium carbamate may be easily under- stood if we imagine one of the two molecules of ammonia which unite with one molecule of carbonic acid to give up an atom of hydrogen to the other. By this, means amidogen is produced on the one hand, and ammonium on the other, which then unite with the carbonic acid as shown by the following equation : — C0 2 + 2NH 3 = Cojggf^ Ammonium carbamate is easily obtained by leading carbonc acid and ammonia, both perfectly dry, into absolute alcohol, when the salt separates as a white, voluminous, crystalline mass. When heated with absolute alcohol in sealed tubes to ioo° and allowed to cool slowly, the salt separates out in large thin crystalline plates. This salt is always contained together with ammonium car- bonate in the commercial carbonate of ammonia. In contact with water it gradually takes up a molecule of this substance, and after some time is converted into ammonium carbonate. It is distin- guished from this latter compound by the fact that its aqueous solution when mixed with a solution of calcium chloride does not at once deposit a white precipitate of calcium carbonate, because the calcium carbamate which is first produced is soluble in water. Calcium carbamate : (CONH : ,).,0 2 Ca + H 2 0, is produced when 294 Text-Book of Inorganic Chemistry. carbonic acid is led into a mixture of thick milk of lime and four times its volume of strong ammonia. The filtered liquid is then cooled to o° and precipitated with absolute alcohol : the amorphous precipitate thus produced soon changes into small prismatic crystals of the above composition. Calcium carbamate is tolerably soluble in water, but the solution soon becomes turbid from the calcium carbonate produced by its decomposition with water. In aqueous ammonia the salt remains longer undecomposed. When gently heated it loses its water of crystallization, and then bears a tempe- rature of 180 without decomposition. Carbamic acid forms more stable organic compounds with the alcoholic radicals, producing then a series of bodies called urethanes. Ammonium carbamate, which contains the constituents of one molecule of water less than ammonium carbonate, contains itself one molecule of water more than carbamide, thus : — Ammonium carbonate .... COJnNT-T 4 Ammonium carbamate . . . . CO,' ATTT 4 I IN rl 2 Carbamide or urea . . . CO MIJ ! Carbamide not yet prepared in the free state is isomeric with urea, a normal constituent of urine ; it may be considered either as the hypothetical carbonic acid in which both atoms of hydroxy] have been displaced by amidogen, or as carbonic anhydride in which one atom of oxygen is displaced by two of amidogen. Am- monium carbamate is converted with loss of water into carbamide, which at once changes into urea, when heated in sealed tubes up to 140°. CARBONIC OXIDE. Composition : CO. The second lower oxide of carbon, which contains the element as a dyad, is, like carbonic acid, a colourless gas, but is distinguished from this compound principally by its combustibility. When heated in the air it catches fire and burns with a blue flame, forming car- bonic acid. Carbonic oxide has a specific gravity of 0-968, corresponding to Carbonic Oxide. 295 a molecular weight of 28, is scarcely soluble in water, and is only condensed to a liquid under the greatest pressure combined with a very low temperature. The gas, although tasteless and odour- less, is very poisonous ; air containing only a few per cent, of it when breathed produces giddiness and intense headache, followed by insensibility, and after a little time by death. Carbonic oxide does not occur free in nature like carbonic acid, but is always produced by the incomplete combustion of carbon or carbonaceous substances — i.e. when these are burnt with an insuffi- cient supply of air. It may also be prepared from carbonic acid if this gas is led over glowing charcoal. The blue flames seen when air flows freely over large masses of glowing carbon — for example, when the door of the firebox of a steam-engine is opened — are due to burning carbonic oxide, produced by the contact of carbonic acid with the red-hot carbon. For the preparation of carbonic oxide it is best to employ either oxalic or formic acid, or else a salt of the last-named acid.. Oxalic acid, which occurs in nature in many plants, especially the sorrels, and which can also easily be prepared artificially, is a solid crystalline substance of strong acid properties,, and having I CO ■ OH the composition p„ ^.y.. This body, when mixed with concen- trated sulphuric acid and warmed, breaks up into water, carbonic acid, and carbonic oxide : — (cO-OH = H *° + C0 * + C0 - The sulphuric acid first removes the water of crystallization from the oxalic acid, and then, on account of its strong, attraction for water, abstracts a further two atoms of hydrogen and one of oxygen, which unite to form water. As, however, the oxalic anhydride : P~ 1 O, which is thus produced, does not appear to exist, this hypo- thetical substance at once breaks up into' carbonic acid and car- bonic oxide : — £°J0 = C0 3 + CO. In order to purify the carbonic oxide from the large quantities of carbonic acid mixed with it, the mixed gases must be passed through a system of wash-bottles containing strong caustic potash, which absorbs the carbonic acid, but allows the carbonic oxide to 296 Text-Book of Inorganic Chemistry. pass on unchanged. If the gases are evolved rapidly, it is difficult to separate the whole of the carbonic acid from the mixture. It is, therefore, better to prepare the carbonic oxide from formic acid, which breaks up when heated with sulphuric acid into carbonic oxide and water only. CO DH and its sodium salt is : j r( ~. .~., T . On heating this compound with concentrated sulphuric acid, formic acid is first set free, and this is then decomposed into carbonic oxide and water. The fol- lowing equation expresses the entire reaction : — {cVoNa + S0 ^(Sh " SO «{o5a. + C0 + H ^ The water produced is, as before, retained by the excess of sulphuric acid. This process easily yields considerable quantities of pure car- bonic oxide, and is the best method to employ, now that formic acid is artificially prepared in large quantities. Carbonic oxide is an indifferent substance in so far that it neither unites with acids nor with bases to form salts. And not- withstanding the tendency, of the unsaturated dyad carbon which it contains to pass into saturated tetrad carbon by union with some other element, it neither unites with oxygen, nor with sulphur, or chlorine at the ordinary temperature. Combination with oxygen to form carbonic acid or with sulphur to form carbon oxysulphide : COS, requires a high temperature ; and to produce its compound with chlorine — carbon oxychloride : COO, — either a high tempe- rature or direct sunlight is necessary. Although carbonic oxide does not unite with aqueous caustic potash, the solid substance combines easily and completely with the gas when the two bodies are heated up to ico° or over. The sole product of this union is potassic formate : — KOH + CO = j" . 0K We can, therefore, obtain carbonic oxide from formic acid by the abstraction of a molecule of water, and reconvert the gas into formic acid by the action of caustic potash. If carbonic oxide is led into a concentrated hydrochloric acid solution of cuprous chloride : Cu„Cl 2 , it is absorbed in considerable Carbon Oxychloride. 297 quantities, and a crystalline compound is produced which contains the elements of cuprous chloride, carbonic oxide, and water. CARBON OXYCHLORIDE (Carbonyl Chloride). Composition : COCl„. This compound, also known as phosgene gas, because formed by the action of light, may be considered as carbonic acid in which one atom of oxygen has been displaced by two atoms of chlorine, and has the same relation to> this substance as sulphuryl chloride (p. 164) to sulphuric anhydride. It is produced, without explosion, when equal volumes of carbonic oxide and chlorine are exposed to direct sunlight, the volume of the mixed gases diminishing to one half. Carbon oxychloride is a colourless gas of peculiar, suffocating odour, easily condensed in an ordinary freezing mixture to a colourless liquid, with a specific gravity of I -43, and boiling at + 8' J . The specific gravity of the gas is 3:46, and its molecular weight 3-46 x 28-88 = 99-9, corresponding to the formula : C0C1 2 . Its com- position is also proved from the fact that two volumes (one molecule) of carbonic oxide and two volumes of chlorine produce two volumes of the gas : — CO + Cl a = COCl 3 2 VOls. 2 Vols. 2 VOls. In contact with water it is at once decomposed into carbonic acid and hydrochloric acid. OXALIC ACID. Composition :\ pn . r> i_r Oxalic acid occurs in nature as the acid potassium salt in the various kinds of sorrel and other plants, and also in combination with lime as calcium oxalate. It crystallizes in colourless rhombic prisms with two molecules of water of crystallization, which can be easily expelled on heating. Oxalic acid possesses a strong acid reaction, and is tolerably 298 Text-Book of Inorganic Chemistry, soluble both in water and alcohol. Nine parts of water dissolve one of the acid at the ordinary temperature, and a much larger quantity when hot. When the anhydrous acid is heated above ioo°, a portion sublimes unchanged, but the greater part is decom- posed into carbonic acid and formic acid : — (CO-OH _ CQ (H jCO-OH " yKJ * + (CO-OH which latter then further decomposes into water and carbonic oxide. Oxalic acid is not now obtained, as formerly, from the acid juice of the sorrel, nor by the expensive process of oxidizing cane- sugar with nitric acid, but is exclusively prepared from some form of cellulose, generally sawdust. The sawdust is mixed with caustic potash, to which a certain, not too large a quantity of caustic soda may be added, and the mass heated until it fuses. The cellulose is thus destroyed — oxidized — with evolution of hydrogen, and a large part of its carbon converted into oxalic acid, which unites with the potassium and sodium. The fused mass is then extracted with water, neutralized with hydrochloric acid, and mixed with a solution of calcium chloride. The insoluble calcium oxalate which is thus produced is filtered off, suspended in water, and digested with the requisite quantity of sulphuric acid. By this means slightly soluble calcium sulphate (gypsum) is formed, and the oxalic acid remaining in solution is deposited on evaporation ; the crude acid being afterwards purified by repeated crystallization. Of theoretical interest is the fact that oxalic acid can be directly obtained from carbonic acid by reduction. If dry carbonic acid is led over potassium amalgam, a brisk reaction takes place, and potassium oxalate is formed. The mercury of the potassium amalgam undergoes no change, but serves simply as a diluent : — ■?cn 4. ,k - (CO-OK 2C0 2 + 2K = | C0 . 0K Oxalic acid is a dibasic acid, and is a chemical compound of two atoms of the monad radical oxatyl : CO-OH, just as a molecule of hydrogen consists of two separate atoms. The oxalates are mostly insoluble in water. One of the most insoluble is calcium oxalate, which is not even dissolved by acetic acid, and which serves, therefore, for the detection and estimation of oxalic acid (and calcium). Oxalic Acid. 299 Oxalic acid in its relations to nitric acid is very stable ; even when the two are boiled together it is only slowly oxidized to carbonic acid. Other oxidizing agents, as potassium permanganate or manganese peroxide, convert it easily and completely into car- bonic acid and water. On heating oxalic acid to the point when it begins to sublime, a portion of it is converted into carbonic acid and formic acid, as previously explained ; and if the oxalic acid is dissolved in aqueous glycerine, and then heated to ioo°, the change into these two sub- stances is so complete and easy that this is employed as the best method for preparing formic acid. The conversion of oxalic acid into carbonic acid, carbonic oxide, and water, when heated with concentrated sulphuric acid, has been already referred to (p. 295). Qxamic Acid and Oxamide. Just as the hypothetical carbonic acid becomes carbamic acid when one of its atoms of hydroxyl is displaced by amidogen, and carbamide on displacement of both atoms by the same radical, so the same change in oxalic acid produces oxamic acid and oxamide respectively : — r, r -a (CO-OH Oxalic acid 1 CO-OH Oxamic acid . . . . {g>;NH a „ ., ICO-NBL 0xamlde JC0-NH 2 Oxamic Acid: - ,-.,-. qtt' 2 is a white, crystalline powder, slightly soluble in water. It is produced by heating acid ammonium oxalate: — (CO-ONH, jCONH, Hn . (CO-OH " (CO-OH + n ^' or by boiling oxamide with aqueous ammonia : — (CO-NH, Hn jCO-NH, jCO-NH* + H * U _ (CO-ONH, and is separated from the solution of its ammonium salt thus obtained by hydrochloric acid. Oxamic acid is a monobasic acid, and is easily dissolved by 300 Text-Book of Inorganic Chemistry. alkalies, but when boiled with these it takes up the elements of water and is converted into oxalic acid and ammonia. Boiled with water, it gives acid ammonium oxalate. Oxamide : Ipn'vn-T 2 ' s a wn i te crystalline powder, tasteless and odourless, insoluble in cold water or alcohol, and only slightly dissolved by these liquids when hot. It is produced when normal ammonium oxalate is heated : — (CO-ONH 4 ICO-NH, „ H n |C0-0NH 4 ~ {CO-NH 2 + 2 " 2U - It is, however, better prepared from an organic compound of oxalic acid — oxalic ether, a colourless liquid with pleasant ethe- real odour and insoluble in water ' — by shaking it with aqueous ammonia. The oily liquid gradually disappears, and a crystalline powder of oxamide takes its place, with the simultaneous formation of alcohol : — JCO-OC.H, , NH _ JCO-NH 2 2 ch-OH lCO-OC,H 5 + - iNH s - (CO-NH 2 + 2l " 2 ■• un Oxalic ether Oxamide Alcohol When carefully heated, oxamide may be sublimed, but quick heating decomposes it into various products, among which cyanogen may be recognized by its odour. Boiling with aqueous alkalies changes it into oxalic acid and ammonia. CARBON DISULPHIDE. Composition : CS„. This compound, the anhydride of a sulpho-acid corresponding to carbonic acid, was discovered by Lampadius towards the end of the preceding century. It is produced in a precisely similar method to carbonic acid — viz. by burning carbon in sulphur gas. A con- siderably higher temperature is, however, necessary to cause carbon to burn in sulphur than in oxygen. Carbon disulphide is a colourless, mobile, very volatile liquid, insoluble in water, in which it sinks, refracting light strongly, and 1 Oxalic ether is prepared by distilling a mixture of dry oxalic acid and absolute alcohol. Carbon Disulphide. 301 possessing an unpleasant ethereal odour. It has a specific gravity of 1-29, boils at 46°, and is miscible with alcohol and ether in all proportions. It is very inflammable, and burns with a bluish flame to carbonic acid and sulphurous acid, with the separation of sul- phur if the air supply is insufficient. Carbon disulphide is prepared by heating freshly glowed wood- charcoal in large tubulated earthenware retorts to bright redness in a suitable furnace, and dropping in pieces of sulphur from time to time. The neck of each retort is connected air-tight with a large receiver in which the uncombined sulphur condenses, and from this the carbon disulphide vapours are led through well-cooled condensers. The crude product contains sulphur, sulphuretted hydrogen, and other impurities, imparting to it a very unpleasant odour, and from which the carbon disulphide cannot be completely freed even by repeated redistillation. It may, however, be effec- tually purified by shaking with mercury and with mercuric chloride, and then distilling. Carbon disulphide is employed for numerous technical purposes, and is now manufactured on a large scale and sold at a low price. It is an excellent solvent for many substances — e.g. for iodine, sul- phur, phosphorus, fatty and ethereal oils, resins, &c. — and in con- sequence of this property is largely used to impregnate india-rubber with sulphur in the manufacture of vulcanised rubber, to free the finer kinds of wool from fat, and to extract fatty oils from seeds containing them. Air containing small quantities of the vapour of carbon disul- phide, when respired, acts in a similar manner to chloroform and produces insensibility ; in larger quantities it is poisonous. At the same time carbon disulphide possesses strong antiseptic properties ; it prevents the putrefaction of meat, and stops the processes of fermentation. Pure carbon disulphide exposed to sunlight soon becomes of a yellow colour and acquires the unpleasant odour of the crude sub- stance. After some time it deposits a brown amorphous substance, which apparently has the composition of carbon monosulphide : CS, and which is insoluble in carbon disulphide and most other liquids. It is probable that the light decomposes the carbon disulphide into sulphur and the still unknown monosulphide, which then polymerises to form this brown substance. Dry chlorine gas decomposes carbon disulphide with formation of disulphur dichloride and carbon tetrachloride. 302 Text-Book of Inorganic Chemistry. Carbon disulphide is the anhydride of a dibasic sulpho-acid — ctt — and as such can unite with aqueous solutions of sodium or potassium sulphides, or other sulphides soluble in water, to form sulphosalts — e.g. with potassium sulphide to form the compound : CS \ „„ — potassium sulphocarbonate. The sulphocarbonates can only be obtained in the solid form with difficulty, as they very easily decompose. The potassium compound separates out from its concentrated solution in yellow deliquescent crystals, containing water. Aqueous alkalies also dissolve carbon disulphide and produce a mixture of sulphocarbonate and carbonate : — 3CS 2 + 6K0H = 2CS |sk + CO loK + 3H 2°- If aqueous ammonia is used, ammonium sulphocarbonate : ShH 4 ' s a ' so produced, together with ammonium sulphocar- bamate : CSJ^^A and ammonium sulphocyanate : CSN-NH 4 . Alcoholic potash when warmed with carbon disulphide dissolves considerable quantities, and after a short time the liquid almost solidifies to a mass of small yellow crystals. This compound is not potassium sulphocarbonate, but is the potassium salt of an organic acid, termed xanthogenic acid. If a sulphocarbonate is mixed with alcohol and then hydro- chloric acid added, free sulphocarbonic acid : CS I ott separates as a dark yellow oil, with a disagreeable odour. The acid readily decomposes into carbonic disulphide and sulphuretted hydrogen. CARBON OXYSULPHIDE. Composition : COS. This compound, which may be considered as carbonic acid in which one atom of oxygen is displaced by sulphur, and is therefore the intermediate compound between carbonic acid and carbon disulphide, is a colourless gas, with an unpleasant odour, combus- Carbon Oxy sulphide. 303 tible and easily inflamed. Water dissolves its own volume of the gas, but decomposes it on standing into carbonic acid and sulphu- retted hydrogen. Carbon oxysulphide is found in some mineral waters, and may be artificially prepared by leading carbonic oxide with excess of sulphur vapour through a glass tube heated to low redness, or by heating a mixture of sulphuric anhydride and carbon disulphide. The best method for its preparation consists in heated potassium sulphocyanate with moderately dilute sulphuric acid, There are then produced, besides carbon oxysulphide, acid ammonium sul- phate and acid potassium sulphate as secondary products, together with some hydrocyanic acid and sulphuretted hydrogen : — CSNK + 2SO s (gg + H 2 - COS + SO.jg^ + SO.jgg Alkalies, like water, decompose carbonic oxysulphide with formation of carbonic acid and sulphuretted hydrogen ; thus, with potash the change is as follows : — COS + 4KOH = COJo£ + K,S + 2H 2 0. A second compound isomeric with carbonic oxysulphide, of the composition : CSO, may exist, but has not yet been prepared. COMPOUNDS OF CARBON AND HYDROGEN. Carbon and hydrogen, although they have little attraction for one another in the free state, unite together to form a large series of compounds — the so-called hydrocarbons. Nearly all these belong to organic chemistry, and we can here only refer to two of the simplest — viz. methane and ethylene. 304 Text-Book of Inorganic Chemistry. METHANE, or MARSH GAS. Composition : CH 4 . This, the most important compound of carbon and hydrogen, is a colourless, odourless gas, insoluble in water, and burning in the air with a pale, non-luminous flame. Its specific gravity is 0-55, and its molecular weight, therefore, 0-55 x 28-88 = 15-9, correspond- ing to the above composition. Its low specific gravity gave rise to its old name light carburetted hydrogen, as distinguished from heavy carburetted hydrogen, or ethylene. Methane is found in nature as a product of the decomposition of organic substances out of contact with air — for example, by the putrefaction of vegetable matter at the bottom of stagnant pools and marshes. The marsh gas thus produced rises to the surface in large bubbles when the mud is stirred up with a stick. This gas is, of course, not pure methane, but contains, besides, some carbonic acid and nitrogen — the former being produced by decom- position, the latter from the air dissolved in the water. Methane is set free in immense quantities in some coal-mines, owing to partial decomposition of the coal. A mixture of this methane with atmospheric air, which is often produced in such mines, explodes when ignited, even more violently than a mixture of hydrogen and air. Two volumes of hydrogen, requiring for its combustion one volume of oxygen, produce two volumes of water- gas ; but two volumes of methane, requiring four volumes of oxygen, produce, on ignition, two volumes of carbonic acid, and four volumes of water-gas : — CH 4 + 20 2 = CO, + 2H 2 0. 2 vols. 4 vols. 2 vols. 4 vols. With a falling barometer — i.e. with reduced air pressure, large quantities of this methane pass from cavities in the coal into the workings of the mines, and, mixing with the air, produce an ex- plosive mixture (fire-damp), which is readily ignited by the lamp of the miner, and then produces those fearful explosions by which so many lives are annually lost. The safety lamp constructed by Davy to avoid these explosions consists of a small oil-lamp completely surrounded by wire gauze, and its action depends upon the fact that an ignited and explosive gaseous mixture when passing through wire gauze has its tempera- • Methane, or Marsh Gas. 305 ture lowered below its point of ignition by the conductivity of the metal for heat, and, therefore, does not burn on the other side of the gauze. Various unforeseen difficulties have, however, prevented its general application, and it is often rendered useless by the miner opening the gauze casing for purposes of his own. Methane is produced, not only by slow putrefaction, but also by the destructive distillation of coal, and is therefoie one of the chief constituents of coal-gas. The well-known lightness of coal- gas and its use for filling balloons are due to this methane, and to the hydrogen which it contains. Other carbon compounds also yield methane when strongly heated. Thus, the vapour of ordinary alcohol : C 2 H,,0, when driven through a red-hot iron tube, yields considerable quantities of methane, mixed however with carbonic acid :— 2C,H,p = 3CH 4 + CO,,. Methane may be obtained fairly pure by strongly heating an intimate mixture of dry sodium acetate and caustic soda, or better soda-lime : in a tube of hard glass. The products are then methane and sodium carbonate : — f CH (CO-ONa + NaOH = CH, + CO(ONa). 2 . Sodium acetate Chemically pure methane can only be obtained from an organic compound called zinc methyl : Zn(CH 3 ) 2 , and which is distin- guished by its great tendency to unite with oxygen. This body is a colourless, volatile liquid, which catches fire at once when exposed to the air and decomposes water with almost explosive violence. If zinc methyl is boiled in a glass flask, out of contact with air, and its vapour then passed into water, it is at once decomposed into methane and zinc hydrate, the latter separating out as a white, gelatinous solid, or remaining dissolved as zinc chloride, if hydro- chloric acid has been previously added to the water : — Zn(CH 3 ) 2 + 2H 2 = 2CH 4 + Zn,OH) 2 . Zinc methyl Methane is a perfectly indifferent compound, and unites with no other substance without decomposition. Sulphuric, nitric, or phosphoric acid are without action on it, and even the powerful oxidizing mixture of potassium dichromate and sulphuric acid, as 1 Quick-lime which has been slaked with caustic soda solution instead of with water. — Ed. X 306 Text-Book of Inorganic Chemistry. well as caustic potash, leave it entirely unchanged. Chlorine is without action upon it in the dark, but if a mixture of the two gases is led into vessels exposed to direct sunlight, one or more atoms of hydrogen, according to the volume of the chlorine, are abstracted from the methane to form hydrochloric acid, while the place of the hydrogen atoms in the methane molecule is taken by an equal number of atoms of chlorine. The end product of the action ef chlorine on methane contains no hydrogen, but as many atoms of chlorine as were originally present of hydrogen. The composition of these different chlorine substitution-products of methane is as follows : — Methane (Methyl hydride) Monochlormethane (Methyl chloride) Dichlormethane (Methylene chloride) Trichlormethane (Chloroform) . Tetrachlormethane (Carbon tetrachloride) These compounds belong to organic chemistry. CH 4 CH 3 C1 CHjCL CHC1 3 ' CCI.j Ethylene, also known as defiant gas, because it forms an oily liquid with chlorine, contains in its molecule the same quantity of hydrogen as methane but twice as much carbon, and has, there- fore, the composition : C 2 H 4 . Ethylene is a colourless gas, of peculiar odour, slightly soluble in water, and condensed to a liquid under considerable pressure and at a low temperature. Its specific gravity is 0-968, and its molecular weight 0-968 x 28-88 = 28-00, corresponding, therefore, to the formula : C a H 4 . It is easily in- flammable and burns with a bright luminous flame. Ethylene is easily obtained from common alcohol. This com- pound, which, at a low red heat, breaks up into carbonic acid and methane, is converted by the abstraction of the elements of water into ethylene. Concentrated sulphuric acid is best employed for this purpose. Four parts by weight of the acid are gradually added to one part of strong alcohol in a large flask provided with a delivery tube. The mixture is then gently heated, and as soon as its temperature has risen to about 1 50 , a copious evolution of ethylene begins. To free the gas from alcohol carried over with it, it is passed through a wash-bottle containing water, and to remove the sulphurous anhydride, produced towards the end of the operation, a second wash -bottle containing caustic potash is used. Ethylene or Okfiant Gas. 307 The process is represented by the following equation, the sulphuric acid remaining unchanged : — C 2 H 6 = C 2 H 4 + H 3 0. If ethylene and chlorine are mixed in a large vessel and ex- posed to diffuse daylight (in direct sunlight the mixture would explode), the two gases unite to produce a colourless, heavy oil, of pleasant ethereal odour, and having the composition : C 2 H 4 C1. 2 — ethylene dichlo?ide. Ethylene also unites directly with bromine, producing ethylene dibromide : CjHjBr.,, a liquid with similar properties to the chloride. The corresponding iodide : C 3 H 4 I 2 , is a crystalline solid, but less stable. Ethylene behaves, therefore, like a dyad element, and. unites without change with two atoms of chlorine, &c, just like copper, zinc, and other metals. If a glass cylinder is half filled with ethylene over water, then an equal volume of chlorine added and a burning taper quickly brought near the open end of the jar, the mixture catches fire, and burns with an exceedingly smoky flame. The products are hydrochloric acid and carbon. Even when excess of chlorine is present no chloride of carbon is formed, a proof that the attraction of carbon for chlorine is not sufficient to cause the elements to unite directly with one another. Since the molecule of ethylene contains two atoms of carbon, instead of one, like methane, it requires a larger quantity of oxygen to burn it than this gas. Two volumes of ethylene require six volumes of oxygen for complete combustion, and produce four volumes of carbonic acid, and four of water-gas : — C 2 H 4 + 30 2 = 2C0 2 + 2H„0. 2 vols. 6 vols. 4 vols. 4 vols. Ethylene and oxygen mixed in this proportion — i.e. 1 : 3, explode violently when ignited, and easily burst a thin glass flask which has been carefully enveloped in towels. At first sight it seems strange that ethylene should burn with a flame so much more luminous than methane, although the latter gas contains half as much carbon and the same quantity of hydro- gen ; and connected with this point is the question, Upon what does the luminosity of a flame depend ? Experiment teaches us that under normal conditions only those flames are luminous which contain a glowing solid body. The x 2 3 o8 Text-Book of Inorganic Chemistry. scarcely luminous hydrogen and oxy-hydrogen flames give out an intense light if a solid substance — e.g. platinum, lime, &c. — is strongly heated in them. In the same way the flame of burning ethylene contains a solid substance — i.e. carbon. And the presence of carbon in the flame may be readily proved by holding a cold, white porcelain dish in it. Those parts of the dish which come into contact with the inner parts of the flame receive a black deposit of soot, just as burning arsine deposits black arsenic on a piece of cold porcelain (p. 232). % The constitution of a luminous flame is not, however, so simple as might be thought from the above statements. Three parts may be distinguished in every luminous flame : an innermost If and cold portion, an intermediate and luminous, and a HI less luminous and hot external mantle, as is illustrated in fig. 57. The gas which streams out of the burner at A does not at once enter into combustion, because no oxygen is present with which it can combine. The inner cone, a'a, consists therefore of unburnt gas, and is cold. This may be readily proved by stretching a fine platinum wire across the flame, which will become red-hot at the two sides of the flame but remain dark, because cold, in the centre. The oxygen which rapidly diffuses from the exterior to the interior portions of the flame not only causes the gases to burn in the exterior mantle, but, by the high temperature produced during this combustion, decom- poses the unburnt gases in the interior of the flame. The decom- position which is thus brought about is the same as that which ethylene suffers when passed through a red-hot tube out of con- tact with the air. It is then changed into methane and carbon : C„H., = CH 4 + C, and it is this carbon which is so liberated, and which is heated in the interior mantle of the flame/ e, g (fig. 57) by the external burning gases, that gives the flame its luminosity. Ethylene is not, therefore, directly burnt in the flame, but rather the products into which it has been changed — viz. methane and carbon. The same process goes on when a candle burns. A burning body brought near to the wick produces the same changes in the fat or wax as destructive distillation, and forms combustible gases which catch fire. The heat of the flame then continually melts Fig. 57 Carbon Tetrachloride. 309 the solid fat at its base, which is sucked up by the porous wick to the flame to be decomposed and burnt. It is not, however, the fat itself which burns, but the gaseous products into which it has been decomposed by the heat of the flame. The further course of the combustion then resembles that of ethylene. A luminous flame may easily be made non-luminous if the combustible gases are mixed with enough oxygen to completely oxidize the whole of the carbon. This can either be done by blowing air or oxygen into the luminous flame (as in the blowpipe), or by previously mixing the gas with air. In the Bunsen burner a mixture of gas and air is obtained in the following manner. The gas is allowed to issue from a small opening over which is screwed a brass tube, with two openings at the base, to admit the air. The gas and air then mix so thoroughly in ascending the tube, that the mixture burns at the top of the tube without smoke or luminosity. The temperature of a flame is sometimes judged from the in- tensity of the light which it emits, and although this may to a certain extent be true for solid bodies (iron is hotter when white- hot than when red-hot, and when brightly glowing than when feebly glowing), in the case of burning gases luminosity and temperature are often inversely proportional to one another. The non-luminous flame of coal gas burning in a Bunsen burner is much hotter than the luminous flame of the same gas burnt in the usual manner. In the latter case the interior of the flame contains unburnt carbon, but in the former all the constituents of the gas are at once burnt. For the same reason the external non- luminous mantle d, b, c, of the gas flame (fig. 57) is much hotter than the interior luminous portion. CARBON TETRACHLORIDE. Of the numerous compounds of carbon and chlorine which are more closely allied to organic than inorganic compounds, we shall here only briefly allude to the most important and at the same time the simplest. Composition: CO,. Carbon tetrachloride is a. colourless liquid sinking in water, but not mixing with it. Its specific gravity is i-6, and its boiling 310 Text-Book of Inorganic Chemistry. point J7°. The liquid possesses a pleasant, ethereal odour, and when breathed acts as an anaesthetic like chloroform. Carbon tetrachloride cannot be obtained by the direct union of carbon and chlorine, but is readily produced from carbon disulphide when this liquid is mixed with an excess of dry chlorine and exposed to the light. A quicker method of preparing it consists in passing chlorine which has bubbled through warm carbon di- sulphide through a red-hot porcelain tube, or by treating this last- named liquid with antimony pentachloride in the presence of chlorine. In all cases the carbon disulphide is converted into carbon tetrachloride and disulphur dichloride. These mixed liquids are then digested with caustic soda, which decomposes the latter, and the carbon tetrachloride then distils over with water vapour on heating. It has already been mentioned (p. 306) that carbon tetrachloride can also be obtained from methane. Carbon tetrachloride is a very stable compound ; it is distin- guished from most other inorganic chlorides by the fact that when boiled with aqueous potash it remains unchanged. Its alcoholic solution is not decomposed when a similar solution of silver nitrate is added to it — silver chloride is not precipitated. If gaseous carbon tetrachloride is led through a red-hot tube it is decomposed into chlorine, and a liquid chloride of carbon con- taining only one half as much chlorine. The molecule of this com- pound does not, however, consist of one atom of carbon and two of chlorine, but contains two atoms of carbon united with four atoms of chlorine, and has, therefore, a similar composition to the tetra- chloride as ethylene has to methane : — Methane . . . CH 4 Tetrachlormethane . . CC1 4 Ethylene . . . C,H, Tetrachlorethylene . C„C1 4 This tetrachlorethylene, like ethylene, unites with two atoms of chlorine, and yields a compound corresponding to ethylene dichloride : — Ethylene dichloride : C 2 H 4 C1„. Tetrachlorethylene dichloride (carbon hexachloride) : C 2 C1 6 . Carbonic hexachloride is a volatile crystalline solid with an odour resembling that of camphor, and which, like camphor, easily sublimes on the walls of the vessels in which it is kept. Carbo7i Tetrabromide : CBr 4 , closely resembles the tetrachloride in its preparation and properties. Cyanogen. 3 1 1 CYANOGEN. Cyanogen is the name given to a compound consisting of one: atom of carbon and one atom of nitrogen, which is so similar in its properties to an element, and especially to the halogens, that it would certainly have been considered as such if it were not pos- sible to produce it from its constituents, and to again decompose it into them. Cyanogen was discovered by Gay-Lussac in 1815 ; its name is derived from Kvdveos, blue, because it is a constituent of Prussian blue, which can be prepared from it. Composition : CN. The molecule of cyanogen consists of two atoms, 1 just as the molecule of hydrogen also contains two atoms, and the composition, of its molecule is therefore expressed by the formula : CN-CN, or. (CN)„. Cyanogen may also be still more briefly expressed by the symbol : Cy. Cyanogen is a colourless gas of a peculiar odour, resembling hydrocyanic acid, and with a specific gravity of r8. This gives a molecular weight of i"8 x 28-88 = 52, corresponding to the formula (CN) 2 . At a low temperature, and under a strong pressure, it may be condensed to a colourless liquid, boiling at — 20° or even to a crystalline solid. Water dissolves about 4J times its volume ; alcohol considerably more. It is easily inflammable, and burns with a beautiful flame of a peach-blossom colour, which, as neither glowing carbon nor nitrogen possess this colour, must be peculiar to glowing cyanogen. Cyanogen is prepared by heating mercuric cyanide. This white crystalline compound, which is easily soluble in water, is obtained by dissolving mercuric oxide in excess of- aqueous hydrocyanic acid, and then crystallizes out on evaporating the clear solution. 1 Strictly, it is, of course, incorrect to speak of an atom of cyanogen, since it is not, like the atoms of the elements, chemically indivisible. It is, however, convenient to evtend the word atom to the compound radicals of organic .and inorganic chemistry, which pliy the part of elementary substances instead of coining a new word such as semi-molecule. When we speak of an atom of cyanogen, ammonium, amidogen, or ethyl, we do not mean an indivi- sible quantity of the substance, but such a quantity which exactly corresponds to the elementary atoms and can displace them in chemical compounds. 3 1 2 Text-Book of Inorgania Chemistry. After being freed from traces of water which it contains, it breaks up when heated to low redness into mercury and cyanogen, just as mercuric oxide decomposes under similar circumstances into ■ mercury and oxygen. The gas can be collected over. mercury. A portion of the cyanogen remains behind in the tube as an amor- phous brown substance called paracyanoge?i, which has the same composition as cyanogen but a higher molecular weight, and of which if is -therefore a polymeric compound. No compound of cyanogen with oxygen or sulphur is known in the free state, and it can only be brought.into combination with these substances indirectly. With the halogens it .combines at once as soon'as they liberate it from its compounds with hydrogen (hydro- cyanic acid) or with the- metals^ (potassium or mercuric cyanide.) . Potassium and sodium burn in cyanogen like they bum in chlorine, and unite directly with the gas, . producing potassium or sodium cyanide. These compounds are. distinguished by their stability; they can" be- heated to redness, out of contact with the air, without decomposition. , • Although carbon and' nitrogen never unite directly with one another when heated together alohe,'6hey may be made to combine in the presence of a third substance which readily unites with cyanogen^f jg\ potassium. Potassium carbonate and carbon, when •heated'- to 'bright redness, form carbonic pxHe and potassium, and if nitrogen, is led over the glowing -mixture, a compound of potas- sium and cyanogen is produced, call'ed potassium cyanide : KCy, -a. white solid,' soluble -in' water like potassium chloride, -and, like this,salt, yielding, 'when decomposed by sulphuric acid, an acid — 'hydrocyanic acid-— corresponding to hydrochloric acid. The same conditions necessary for fhe production of potassium, cyanide are also present when' nitrogenous organic substances, such as blood, hair, horn, are mixed with potassium carbonate and heated. Under' these circumstances, the nitrogen and carbon of ' the organic substances unite with the pqtassium to form potassium cyanide. This process is that "by means of which cyanogen com- pounds are manufactured on a large scale. >.The first compound always prepared is^the* so : called yellow- pritssiatc, or potassium ferrocyanide, which may be considered as a double cyanide of potassium and iron :_ 4.K.Cy,Fe"Cy , and which will be.more'fully discussed under the compounds of iron. 'It is also produced when an ' aqueous solution ' of potassium cyanide acts upon- powdered fe,rrbus sulphide ; a portion' 'of the potassium cyanide is then decom- Hydrocyanic Acid. 313 posed, and forms potassium sulphide and ferrous cyanide. From its solution in water it crystallizes in large yellow tablets. Potassium ferrocyanide forms the starting-point for the prepara- tion of all the other cyanogen compounds — e.g. hydrocyanic acid, potassium cyanide, potassium cyanate, &c. HYDROCYANIC ACID. (Prussic Acid.) Composition ; HCN or HCy. This compound, well known on account of its highly poisonous character, is, when completely anhydrous, a colourless, mobile, volatile liquid, miscible with water in all proportions. It possesses a strong stupefying odour, and has a specific gravity of 07. Its boiling-point is 27 , and at - 15 it solidifies to a crystalline solid. It is easily combustible. Anhydrous hydrocyanic acid may be prepared in the same way as hydrochloric acid by pouring concentrated sulphuric acid, on to potassium cyanide, but the reaction is much too violent, and, on account of the poisonous nature of the gas, far too dangerous to be employed. It is much better to first prepare the aqueous acid and then to abstract water from it. For the same reason it is better to employ for the preparation of the dilute acid r,ot potassium cyanide, but its compound with ferrous cyanide — i.e. potassium ferro-cyanide. To obtain the dilute aqueous acid, 10 parts of powdered potas- sium ferrocyanide are covered with a mixture of 6 parts of sulphuric acid and 30 parts of water in a tubulated retort with its neck in- clined upwards. The end of the neck of the retort is connected by a cork and a bent tube with a condenser, the other end of which dips in a receiver carefully cooled in ice, and in which it is best to place a little water at the commencement. The mixture in the retort is then gradually brought to boiling, when a portion of the water-vapour condenses in the neck of the retort, while the volatile hydrocyanic acid passes over to the well-cooled receiver. The whole apparatus should be placed in a good upward draught, and every precaution taken, not only in the preparation of this highly poisonous volatile compound, but also in all experiments made with i't afterwards. 314 Text-Book of . Inorganic Chemistry. In order to prepare the anhydrous compound from its aqueous solution, a flask which is about one-third filled with the latter is placed in ice and small pieces of fused calcium chloride gradually dropped in, which unites with the water and partially liquefies. On no account must the aqueous acid be poured on to the calcium chloride, as the heat produced by the union of the latter with the water would drive off large quantities of gaseous hydrocyanic acid. After a few hours, the flask is attached to a long condenser and a receiver immersed in a freezing mixture, and then gently heated on a water bath. Anhydrous hydrocyanic acid then distils over as a clear, transparent liquid. Aqueous, and still more anhydrous hydrocyanic acid, undergoes a change when kept for some time in closed vessels, with the separation of a brown amorphous substance, known as ftarahydro- cyanic acid, and probably a polymeric compound of hydrocyanic acid. The decomposition is usually accompanied with an evolution of gas, which sometimes exerts sufficient pressure to destroy the vessels in which it is preserved. It has been found that the addi- tion of a few drops of hydrochloric acid materially diminishes this spontaneous decomposition. Hydrocyanic acid is one of the most powerful poisons known. A few drops of the strong aqueous acid produce almost instanta- neous death. Its vapour is also highly poisonous when respired ; 1 pigeon dies in about half a minute when breathing air mixed with the vapour from anhydrous hydrocyanic acid, even when the tempe- rature of the liquid is as low as 0°. Very largely diluted with water, hydrocyanic acid is a valuable medicine. The strength prescribed by the ' British Pharmacopoeia ' is a 2 per cent, solution ; a 3 per cent, solution is also sometimes used under the name Scheele's prussic acid. Bitter almonds, and the kernels of cherries, plums, &c, as well as the leaves of the common or cherry laurel, contain a crystalline organic compound of complex composition, called amygdaline. When these sub- stances are crushed and moistened with water, the amygdaline comes into contact with a ferment (emulsine) which they also con- tain, and is then decomposed as sugar is decomposed by yeast. One of the products of this decomposition is hydrocyanic acid, which, when the water is afterwards distilled, passes over with the water-vapour into the receiver. The preparations known as bitter- almond-water and laurel-water are obtained in this way. Hydrocyanic acid is so weak an acid that it does- not even red- Hydrocyanic Acid. 315 den blue litmus paper, and is expelled by carbonic acid from its compounds with the alkalies. For this reason, potassium cyanide when exposed to the air smells strongly of hydrocyanic acid. In many of its reactions, it shows a similarity with the hydrogen compounds of the halogens — for example, when mixed with a solution of silver nitrate it gives a white, curdy precipitate of silver cyanide, which, like silver chloride, is insoluble in nitric acid but dissolves in ammonia. Silver chloride may, however, be at once distin- guished from silver cyanide if heated ; the former compound simply melts, but the latter is decomposed into silver and cyanogen gas. As potassium chloride unites with platinic chloride to form a stable crystalline double compound, so also potassium cyanide unites with nearly all the metallic cyanides, and produces from these insoluble substances soluble crystalline double cyanides, of which the stable potassium ferrocyanide may be taken as an example. Chlorine easily decomposes hydrocyanic acid, uniting with its two constituents to form cyanogen chloride and hydrochloric acid. In order to detect hydrocyanic acid in a liquid, advantage is taken of the readiness with which Prussian blue may be formed from it. Prussian blue may be considered as a double compound of ferric and ferrous cyanides, Fe" / Cy 3 and Fe"Cy„, and although insoluble in water and dilute acids, is not produced when a mixture of a ferrous and ferric salt is added to the aqueous acid. If, how- ever, this mixture is made alkaline with caustic potash, potassium ferrocyanide, and a mixture of ferrous and ferric hydrates are pro- duced, and this, when acidulated with hydrochloric acid, forms Prussian blue clue to the union of the potassium ferrocyanide with the ferric chloride. Cyanogen is a monad radical, and is to be considered as a com- pound of triad nitrogen with dyad carbon. According to this, hydro- cyanic acid would be a compound resembling ammonia in which two atoms of hydrogen are displaced by one atom of dyad carbon, and would then possess the formula : N j -^ And the production of silver cyanide from hydrocyanic acid would be represented by the equation : — NJ£ + NCvOAg = N{£ g + NCyOH. 3 16 Text-Boak of Inorganic Chemistry. At present we only know one kind of hydrocyanic acid, but it is possible, and indeed probable, that a second isomeric compound exists, containing tetrad carbon as the grouping element. In this compound, three of the four bonds of the carbon atom would be satisfied with the atom of nitrogen, and one with the atom of hydrogen, thus : C \ u- This hydrocyanic acid would probably be only slightly, if at all, poisonous. Cyanides are known which from their peculiar chemical behaviour appear to contain two modifications of cyanogen, from which the existence of a second modification pf hydrocyanic acid may be deduced. Possible, but not probable, is the existence of a third isomeric hydrocyanic acid containing pentad nitrogen as grouping element saturated with one atom of tetrad carbon and one of hydrogen, and of which the composition would be expressed by the formula : iN IH CYANIC ACID. Composition : CONH. Although cyanogen cannot unite directly with oxygen as it does with chlorine, potassium cyanide can, under favourable circum- stances, easily take up an atom of oxygen and become oxidized to potassium cyanate. But the compound called cyanic acid does not possess a similar composition to chloric acid ; such a compound is, in fact, as much unknown as compounds of cyanogen corresponding to perchloric or chlorous acid, or any of the oxides of chlorine. Cyanogen appears in fact to unite with oxygen and hydrogen, or a metal in only one proportion — viz. as cyanic acid : CONH. Potassium cyanate, although similar in composition to potassium hypochlorite, ClOK, has by no means the same constitution. In the hypochlorite the atom of oxygen unites together the chlorine and potassium, but in the cyanate the oxygen atom is not united to the cyanogen as a whole, but to the atom of dyad carbon, form- ing carbonic oxide, which then plays the part of the carbon in potassium cyanide. The relation of potassium cyanide to potas- sium cyanate is best expressed by the following rational for- mulas : — Cyanic Acid. Potassium cyanide n( c " • 1n |k Potassium cyanate nJCO 317 whence it follows that 'the constitution of cyanic acid is N ■ K, The change of potassium cyanide into the cyanate, which takes place so readily, consists therefore of the oxidation of the atom of carbon and in its change from the dyad to the tetrad state. Potassium cyanate, a crystalline salt, soluble in alcohol and water, is readily obtained by dropping lead oxide into fused potas- sium cyanide when the oxide is reduced to metallic lead. The salt will be more fully described under the potassium compounds. Cyanic acid cannot, however, be set free from this salt by a stronger acid, because whenever the acid comes into contact with water it is at once decomposed into carbonic acid and ammonia : — ■ N 1H° + H *° = C ° 3 + NH a- Cyanic acid may be obtained by heating a polymeric compound — cyamiric acid, C 3 N 3 3 H 3 — which is a white solid, obtained from urea. Cyanuric acid is not itself volatile, but when heated breaks up into cyanic acid, which distils over as a colourless liquid, with an odour resembling acetic acid. If kept below o° it may be pre- served for some time unchanged, but at a few degrees above this temperature the liquid becomes heated, and changes into a white solid, resembling porcelain. This compound — cyamelide—\s iso- meric with cyanic acid, and is again converted into it on heating. The compounds of cyanogen, especially those belonging to organic chemistry, are remarkable for the readiness with which they change into isomeric and polymeric compounds. SULPHOCYANIC ACID. Composition : CSNH. Just as fused potassium cyanide unites readily with oxygen, even that of the air, so also it possesses a strong attraction for sul- phur. Sulphur dropped into fused potassium cyanide readily com- bines with it, producing potassium sulphocyanate : CSNK, which 3 1 8 Text-Book of Ino7'ganic Chemistry. from its analogy with potassium cyanate may be considered to ( CS" possess the same constitution : N \ y- The corresponding hydro- gen compound, sulphocyanic acid : CSNH, is more stable than cyanic acid ; an aqueous solution of it may be prepared by dis- tilling potassium sulphocyanate with dilute sulphuric acid. The anhydrous acid is best obtained by passing dry sulphuretted hydro- gen over powdered mercuric sulphocyanate. It then condenses in the receiver as a colourless oily liquid, with a piercing odour, also resembling acetic acid, but when kept for a short time soon de- composes into hydrocyanic acid and a yellow crystalline solid nearly insoluble in water — persulphocyanic acid: C 2 H 2 N 2 S 3 . Sulphocyanic acid or any soluble sulphocyanate can be at once recognized by the bright red colour which it gives with ferric chloride. This reaction is quite as delicate and distinct as the blue precipitate produced from potassium ferrocyanide and ferric chloride. Chlorides of Cyanogen. We are acquainted with two polymeric compounds of cyanogen and chlorine, one a volatile liquid and the other a crystalline solid. These compounds appear to stand in the same relation to one another as cyanic acid to cyanuric acid. Liquid Cyanogen Chloride : CNC1, is a colourless liquid, boiling at 12°, solidifying at —y°, and with an exceedingly sharp, tear-producing odour. It is very poisonous. This compound is obtained by the action of chlorine on dilute hydrocyanic acid or on metallic cyanides in the presence of water. The simplest method of preparation is to lead a slow stream of chlorine into a solution of potassium cyanide with twice its weight of water kept cooled in ice. The solution is placed in a retort, which it only fills to about one-fourth, and the neck of the retort is connected with a tube filled partly with bright copper turnings (to absorb the excess of chlorine), and partly with calcium chloride (to retain moisture). This tube communicates with a (J-tube placed in a freezing mixture in which the cyanogen chloride condenses. A portion of the volatile compound distils over as soon as it is formed ; the residue is after- wards obtained by heating the retort to 50° Pure liquid cyanogen chloride may be preserved unchanged in sealed tubes, but when impure it soon changes into a solid poly- Coal- Gas. 319 meric modification. Water and alcohol dissolve considerable quantities without decomposition, and it can be again expelled from these solutions on heating. Alkalies decompose it. Liquid cyanogen chloride unites chemically with the chlorides of other elements — e.g. of antimony, boron, titanium. Solid Cyanogen Chloride, probably CgN^Clj, is produced by the action of chlorine on anhydrous hydrocyanic acid in bright sunlight, when it separates out in colourless lustrous crystals. Like the liquid chloride, it has a piercing odour. It melts at 145°, boils at 190°, and is decomposed on boiling with alkalies, or even water, into cyanuric and hydrochloric acids. Cyanogen Bromide : CNBr, and Cyanogen Iodide : CNI, are volatile crystalline solids, and are produced by treating mercuric cyanide or potassium cyanide with bromine or iodine respectively. Both possess a similar penetrating odour to that of the two chlorides. Cyanogen iodide is often contained in commercial iodine. COAL-GAS. 1 Coal-gas is a complex mixture of various gases produced by the destructive or dry distillation of coal. The coal when distilled also yields a number of liquid and solid products, which are con- tained in coal-tar— & thick, oily, black liquid ; and in the ammo- niacal-liquor — an aqueous liquid, smelling strongly of ammonia and ammonium sulphide. The apparatus used in the manufacture of coal-gas consists of the retorts in which the coal is heated, the condensers for removing the coal-tar and ammoniacal-liquor, the purifiers for separating gaseous impurities, and gasometers for storing the gas. The retorts are made of fire-clay, and so that their cross section has a D-shape ; they are closed at one end and are arranged in the furnace in a horizontal position, with their open ends projecting slightly outwards. After the charge of coal has been introduced, the open ends are closed by an iron plate, which can be bolted on, and from which a vertical iron tube passes upwards. This iron 1 By the Editor. 320 Text-Book of Inorganic Chemistry. tube, which conveys away all the products of distillation from the retorts, dips under the surface of "water contained in a large iron tube running in a horizontal position above the retorts and called the hydraulic main. Most of the tar and much of the ammonia is here condensed, and when the liquid reaches a certain height it runs off through an opening into the tar-well — a large reservoir receiving all the liquid products of distillation. The gas next passes through a series of upright iron tubes, called atmospheric condensers, which are freely exposed to the air, and so cool the gas and condense a further quantity of liquid products. But in order to completely remove these substances, as well as some of the gaseous impurities, the gas is. passed through the scrubbers, or tall vessels filled with fragments of coke, and kept continually moist by a descending spray of cold water. • The gas enters at the base of the scrubber and escapes near the top. It now only remains to remove traces of gaseous impurities, especially sulphuretted hydrogen, carbon disulphide, and other sulphur compounds, which impart an unpleasant odour to the gas, and which, when burnt, produce sulphurous and sulphuric acids. This final purification is effected in the purifiers, which consist of large chambers containing thin layers of dry slaked-lime or of ferric hydrate on shelves. The lime then absorbs the sulphuretted hydrogen and other sulphur compounds, together with any car- bonic acid which may be present, and is converted into spenl- lim£ — a dirty-green amorphous mass of various compounds of calcium, with a very offensive odour. If ferric hydrate is employed it is converted by sulphuretted hydrogen into ferrous sulphide and free sulphur, and the former compound is partially oxidized on afterwards exposing the mass to the air. But when the quantity of free sulphur becomes large, the ferric hydrate can no longer effect the purification, and is then sent to the sulphuric acid works to burn out the sulphur and again obtain pure ferric oxide. Ferric hydrate does not remove carbonic acid, and it is, therefore, better to mix it with a certain proportion of slaked-lime in order to get rid of this impurity. The presence of sulphuretted hydrogen in coal- gas may be detected by allowing a jet of the gas to play on a piece of paper impregnated with a solution of lead acetate. If the paper does not become blackened after a few seconds, the gas may be said to be free from this impurity. This test does not, however, detect the presence of carbon disulphide or other sulphur com- pounds. Coal-Gas. 321 The gases contained in coal-gas may be divided into three classes : — (i.) Diluents, which have practically no illuminating power, but which are of special value when the gas is used for heating purposes. These consist of hydrogen, marsh gas, and car- bonic oxide, and constitute from 85 to 90 percent, by volume of the entire gas. (ii.) Illuminants, to which the illuminating power of the gas is due, and which consist of hydrocarbons containing more carbon in their molecule than marsh gas. They are some- times called heavy hydrocarbons. These illuminants are principally ethylene (C„H 4 ), propylene (C 3 H 6 ), butylene (C 4 H ? ), acetylene (C 2 H„), and benzene (C„H,,). (iii.) Impurities, which consist of sulphuretted hydrogen, carbon disulphide and other sulphur compounds, together with traces of carbonic acid, oxygen, and nitrogen, the two latter being derived from the air. The composition of coal-gas depends upon the temperature at which the coal is distilled and upon the kind of coal used. The gas also varies in composition at different stages of the distillation. If the temperature is high, the heavy hydrocarbons are decomposed and a large volume of gas of low illuminating power is obtained. And if cannel coal is employed the gas contains a much larger proportion of these heavy hydrocarbons than that obtained from ordinary bituminous coal. The following table gives the compo- sition of two samples of coal-gas. No. 1 is ordinary London gas, made from a mixture of bituminous coal and a small quantity of cannel ; No. 2 is gas made entirely from cannel coal and supplied to the Houses of Parliament : — No. , No. 2 Marsh Gas , . . . . 35 ' s 9 41-88 Hvdrogen 50-68 41-72 Carbonic Oxide .... 3-98 4-98 Heavy Hydrocarbons 4-08 8-72 Nitrogen .... 4 '93 271 Oxvgen 0*26 none Carbonic Acid traces none 99*82 IOO "OI The greater part of the nitrogen which the coal contains is converted into ammonia during the distillation. An aqueous Y 322 Text-Book of Inorganic Chemistry. solution of this ammonia constitutes the ammoniacal-liquor which collects in the tar-well. The ammonia is contained in this liquid partly in the free state, and partly in combination as ammonium carbonate, sulphide, cyanide, &c. This ammoniacal-liquor is now the chief source of ammonia and its compounds. The coal-tar which is also found in the tar-well is a highly complex mixture of a large number of organic compounds. Among these may be mentioned the liquids — benzene and its homologues, aniline and phenol (carbolic acid); and among the solids — paraffin, napthalene, and anthracene. When the tar is distilled the more volatile liquids pass over and produce a combustible mixture of various compounds, called coal-tar naphtha. Coal-tar is used for a variety of purposes, some of which depend upon the antiseptic properties of the phenol which it contains, but by far the largest quantities are consumed in the manufacture of the beautiful coal- tar or aniline colours. After the distillation is finished a considerable portion of the carbon present in the coal remains behind in the retorts as a grey porous mass, containing also the whole of the ash of the coal. This constitutes coke. At the same time, owing to the decomposition of some of the heavy hydrocarbons during distillation, a layer of ex- tremely hard and compact carbon collects on the upper surface of the retorts. This form of carbon is called gas-carbon \ it is used in the construction of some galvanic batteries and for highly refrac- tory crucibles. TITANIUM. Chemical Symbol: Ti. — Atomic Weight: 48. This element, which in many respects is closely allied to silicon, is also always found in nature, combined with oxygen. Titanium dioxide or titanic anhydride : TiO a , either free or combined with bases, is the form in which the element always occurs. Free titanic anhydride is found in three varieties : rutile, anatase, and brookite, which, although chemically the same, have different crystalline and other physical characteristics. Rutile crystallizes according to the tetragonal system, anatase and brookite according to the rhombic, and titanic anhydride is therefore trimor- phous. Of the other naturally occurring compounds of titanium, Titanium. 323 titanic iron, probably ferrous titanate and tit unite or sphene, calcium titanate and silicate, are the most important. Free titanium is obtained in the following manner. Two iron boats, one containing dry potassium fluotitanate : K s TiF, ; , and the other metallic sodium, are brought into a tube of hard glass, which is then filled with pure hydrogen. The boat containing the fluo- titanate is first heated to redness, then that containing the sodium, so that the vapour of the metal is carried by the hydrogen over the hot titanium salt. By this process sodium and potassium fluorides are produced and free titanium. If the two former com- pounds are afterwards extracted by hot water the titanium remains behind as a dark -grey amorphous powder, resembling iron which has been reduced in hydrogen. The element, in this pulverulent form, burns brilliantly when heated in oxygen or the air, forming titanic anhydride : TiO,. It dissolves easily in hydrochloric acid, with evolution of hydrogen, probably producing a solution of titanium dichloride : TiCl,. It burns when heated in chlorine, giving titanium tetrachloride : TiCl 4 . The material used for the preparation of titanium compounds is nearly always rutile. This substance, like quartz, is insoluble in hydrochloric acid, but may be brought into solution either by fusing with caustic potash or potassium carbonate, or with acid potassium sulphate, or, finally, by heating with concentrated sul- phuric acid until the greater part of the liquid has been volatilized. Titanic Anhydride : TiO a . — The titanium sulphate produced by either of the last two methods dissolves in cold water, and from this solution titanic acid is precipitated in while flocculent masses on the addition of alkalies or alkaline carbonates, and is not re- dissolved by an excess of these substances. The dilute aqueous solution of titanium sulphate is decomposed on boiling, and deposits titanic acid as a white powder, which is only slightly soluble in concentrated sulphuric acid. Potassium titanate, obtained by fusing powdered rutile with three times its weight of potassium carbonate, does not dissolve when the fused mass is afterwards treated with water, but forms acid potassium titanate, which is insoluble in water, and which can therefore be purified by washing. This acid salt dissolves, how- ever, in cold hydrochloric acid, and when the solution is diluted with water and heated to boiling, titanic acid is precipitated as a white powder. 324 Text-Book of Inorganic Chemistry. This form of titanic acid is scarcely soluble in water containing hydrochloric acid or ammonium chloride, but when washed with pure water passes in a milky state through the filter. The titanic acid precipitated by boiling aqueous solutions con- ' taining hydrochloric or sulphuric acid always contains traces of the acid mixed with it, from which, however, it may be completely freed by digesting with ammonia, which has no action on the titanic acid. The white precipitate produced by ammonia in a cold hydro- chloric acid solution of titanic acid is said to have the composition : Ti(OH) 4 . This compound is easily soluble in dilute acids, but is converted on warming into an acid containing less water — meta- titanic acid, of the probable composition : TiO(OH) 2 , which is only slightly soluble in acids. The titanic acid precipitated by boiling a sulphuric acid solution is insoluble in acids except concentrated sulphuric acid. If metallic zinc is added to a hydrochloric acid solution of titanic acid, the nascent hydrogen which is produced reduces the titanic acid to the hydrated sesquioxide. In consequence of this reduction the liquid becomes of a violet-blue colour, and after a time, if the solution was not too dilute, a dark blue or violet blue precipitate separates out, which again becomes white when exposed to the air from absorption of oxygen. Titanium Sesquioxide : Ti^Oj, is obtained in the anhydrous state by glowing titanic anhydride in a stream of hydrogen. It is a black powder insoluble in nitric or hydrochloric acid, but dissolving in sulphuric acid to form a violet solution. When strongly glowed in the air it becomes white, being reconverted into titanic anhydride. Titanium Tetrachloride : TiCl 4 . — This compound, which closely resembles silicon tetrachloride, except that it is less volatile, is obtained, like silicon chloride, by heating an intimate mixture of rutile and charcoal in chlorine. It then distils over into a cooled receiver as a heavy liquid, usually with a yellowish tinge due to traces of ferric chloride. When purified by redistillation its specific gravity is 176, and boiling point 136 . It fumes strongly in moist air and has a vapour density of 6'8, corresponding to the molecular formula : TiCl 4 . Water at once decomposes it, with considerable evolution of heat, into hydrochloric and titanic acids. Titanium chloride absorbs considerable quantities of ammonia Molybden um. 325 yielding a reddish-brown powder of the composition : TiCl 4 ,4NH 3 , which is easily decomposed by moist air. Titanium Telrafluoride : TiF 4 , is a colourless, fuming liquid, obtained by heating a mixture of titanic anhydride and fluor-spar with concentrated sulphuric acid. Fluotitanic Acid: H 2 TiF fi , which is analogous in composition to fluosilicic acid, is obtained by dissolving titanic acid in hydro- fluoric acid. If the aqueous solution is neutralized with potassium carbonate, or if potassium titanate is dissolved in hydrofluoric acid, or, finally, if rutile is fused with acid potassium fluoride, potassium fluotitanate : K 2 TiF 6 , is easily obtained in the crystalline form. Titanium belongs to those elements which unite with nitrogen to form compounds stable at a red heat. If titanium chloride is saturated with dry ammonia, and the resulting solid : TiCl.,, 4NH3 heated to redness in gaseous ammonia, titanium nitride : Ti 3 N 4 , remains behind as a copper-coloured mass, with metallic lustre, and which evolves ammonia when fused with solid caustic potash. In the slag from blast furnaces, where iron ores containing titanium have been used, bright-red, lustrous crystals in cubes are often found, and sometimes in considerable quantities. They consist of a chemical compound of titanium nitride and cyanide. MOLYBDENUM. Chemical Symbol : Mo. — Atomic Weight : 96. Molybdenum occurs chiefly in nature as molybdenite, the di- sulphide : MoS,, and more rarely as wulfenite, lead molybdate : Mo0 2 '0 2 Pb. Molybdenite, which closely resembles graphite in its appearance, is a soft mineral, feeling greasy to the touch, and easily marking paper. It is, however, heavier than graphite, and is further distinguished by giving an olive-green streak on porcelain, instead of a grey one like graphite. Molybdenum is readily obtained by heating one of its chlorides or oxides in hydrogen, when the walls of the glass tube used for the reduction often become covered with a brilliant, metallic-like mirror of the free substance. In its physical properties, molyb- denum far more closely resembles the metals than the non-metals. 326 Text-Book of Inorganic Chemistry. It is of a silver-white colour, with a strong metallic lustre, very hard, melts with extreme difficulty, and is possibly infusible when quite pure. Its specific gravity is 8'6. In the air it remains un- changed, and only combines with oxygen when heated to low red- ness. Hydrochloric acid or dilute sulphuric acid has no action on it. Concentrated sulphuric acid converts it into a brown mass. Aqua regia easily dissolves it, and nitric acid converts it into one of the oxides. For the preparation of molybdenum compounds from molyb- denite, the mineral is powdered, and placed in an open, inclined crucible, which is then surrounded with pieces of glowing char- coal and heated to low redness, with frequent stirring, as long as sulphurous anhydride is evolved. Crude and impure molybdic anhydride then remains behind as a yellow powder. This com- pound is insoluble in water, but dissolves in aqueous ammonia, leaving the impurities and the unoxidized molybdenite behind. A further quantity of molybdenum may generally be obtained from the residue by roasting again as before. On evaporating down the ammoniacal solution ammonium molybdate crystallizes out, and from a solution of this salt, cautious addition of dilute hydrochloric or nitric acid precipitates molybdic acid as a white powder, scarcely soluble in water, but easily soluble in an excess of acid. Molybdenum forms several compounds with oxygen. We dis- tinguish, for example, molybdenum sesquioxide : Mo,0 3 , obtained by reduction from molybdic anhydride as a black powder insoluble in acids, also molybdenum dioxide : Mo0 2 , which remains behind as a red crystalline mass when ammonium molybdate is heated in a closed crucible, and Molybdic Anhydride : Mo0 3 , the best known oxide. Molyb- dic anhydride or molybdenum trioxide remains as delicate white crystals, yellowish when hot, on heating ammonium molybdate in an open crucible. It melts at a red heat, and sublimes in small, thin, glistening scales. Acids do not dissolve it, but with alkalies it forms soluble crystalline compounds — the molybdates. Molybdic Acid : Mo0 2 (OH) 2 , is a dibasic acid, and separates from a solution of any of its salts on careful addition of dilute hydrochloric acid as a white crystalline powder, easily soluble in excess of acid. If zinc is added to a hydrochloric acid solution of molybdic acid, a blue colour is first produced which soon becomes Tungsten. 327 green, and finally dark-brown. These changes in colour are due to the reduction of the molybdic acid by the nascent hydrogen. Among the salts of molybdic acid that of ammonium is of considerable importance in analytical chemistry ; it is used for the detection and estimation of phosphoric acid in acid liquids containing other compounds. If such a liquid containing phos- phoric acid — for example, the aqueous or acid extract of a soil or manure — is mixed with an excess of a solution of ammonium molybdate strongly acidified with nitric acid and then warmed, a yellow crystalline precipitate of the composition : 1 1 MoO„, PO(ONH 4 ) 3 ,6H 2 0, and containing the whole of the phosphoric acid, is thrown down. This precipitate is insoluble in water or dilute acids, but easily soluble in ammonia. If this yellow compound is boiled with aqua regia, the ammonia is destroyed, and on allow- ing the solution to evaporate at the ordinary temperature, phospho- molybdic acid: 1 iMo0 3 ,PO(OH) 3 ,I2H.,0, crystallizes out in yellow prisms. The acid is easily soluble in water, and produces, with acid solutions of ammonium salts or nitrogenous organic bases, yellow compounds which are insoluble in dilute acids. Ammonium molybdate, when mixed with an acid solution of arsenic acid or an arsenate, produces a similar yellow insoluble salt. Several chlorides of molybdenum are known as : MoCl^, M0CI3, MoCl 4 , MoCl 5 , and MoCl e ; besides these two oxychlorides of hexad molybdenum also exist, viz. : MoOCl 4 and Mo0 2 Cl 2 . TUNGSTEN. Chemical Symbol : \V. — Atomic Weight: 184. This element resembles molybdenum in many respects, but is distinguished from it by the high specific gravity of the free element and of its compounds. Tungsten never occurs free in nature, and its compounds are only found in small quantities. The best known of the tungsten minerals are wolfram (ferrous and manganous tungstate) and schee- lite (calcium tungstate), the former of a dark grey colour with a specific gravity of 7-5, the latter consisting of white crystals also remarkable for its high specific gravity. 328 Text-Book of Inorganic Chemistry. Tungsten, prepared by the reduction of tungstic anhydride in a stream of hydrogen at a bright red heat, forms small grey crystal- line particles, which under the burnisher become lustrous like iron. Its specific gravity is about 19, and thus nearly approaches that of gold. It is only difficultly fusible, especially in large quantities, is brittle, and so hard that it scratches glass. At the ordinary temperature it remains unchanged in the air ; but when it is heated in the pulverulent form it burns easily, and pro- duces tungstic anhydride. Hot nitric acid oxidizes it to the same product. Of the oxides of tungsten, the dioxide, WO,, is known as a brown powder insoluble in hydrochloric or nitric acid. Moie im- portant than this is the trioxide — Tungstic Anhydride : W0 3 , which remains behind as a yellow powder, insoluble in water and acids, when powdered wolfram is heated with nitric acid, while the iron and manganese go into solution. It may also be obtained by acting on scheelite with nitric or hydrochloric acid. This yellow tungstic anhydride dis- solves readily in alkalies (including ammonia), forming soluble crystalline salts — the tungstates. From a cold solution of these salts dilute acids precipitate white tungstic acid of the composition : WO(OH)„ which, when dried over sulphuric acid, passes into the dibasic acid : W0 2 (OH) 2 , and when heated to the boiling point of water is converted into yellow tungstic anhydride. If tungstic acid is precipitated from a solution of a tungstate by hydrochloric acid, and then a piece of zinc added to the acid liquid, a beautiful dark blue colour is produced, due to a soluble oxide of uncertain composition, and which, when exposed to the air, again oxidizes to tungstic acid. The normal tungstates are salts of the dibasic acid : \VO,(OH).,. Others, however, possess a more complex composition ; for exam- ple, the metatungstates of which the sodium salt has the formula : Na.,W 4 0i3 + ioH 2 0. The common sodium tungstate of commerce is still more complex, and has the composition : Na 10 W lo O 41 . Corresponding to phosphomolybdic acid is a phosphotungstic acid of similar composition. And a remarkable double compound of tungstic and silicic acids is known — silicotungstic acid, the soluble and crystalline compounds of which are obtained by boil- ing acid alkaline tungstates with gelatinous silicic acid. Like molybdenum, tungsten unites in several proportions with Vanadium. 329 chlorine. The hexachloride : WC1, ; , corresponding to tungstic anhydride, forms a dark violet mass, melting at 275° to a red liquid, and boiling at 346°. It is decomposed by water into tungstic and hydrochloric acids. This compound may be obtained by moderately heating free tungsten in a stream of pure chlorine. Chlorides of composition : WC1 0) WC1,„ WC1 2 , are also known. Corresponding to tungsten hexachloride are two oxychlorides : WOCl 4 and W0 2 C1 2 , both solid crystalline substances. Tungsten, when alloyed, in small quantities, with iron, increases its hardness ; steel containing about 5 per cent, of tungsten known as tungsten-steel — is distinguished by its great hardness. VANADIUM. Chemical Symbol : V. — Atomic Weight : 5 r -3. Vanadium belongs to the rarer elements, although it is widely distributed in minute quantities ; it occurs especially in certain ores of iron. Vanadic acid is usually prepared from these iron ores, the mineral vanadinite — lead vanadate — being too rare to serve for the preparation of vanadium compounds. Considerable quantities of vanadium compounds are now obtained in England from a mineral called mottramite, a double vanadate of lead and copper, occurring in Lancashire. Vanadium, obtained by glowing the dichloride in pure hydrogen, is a grey powder of specific gravity 5'5, which remains unchanged in the air, and can be raised to redness out of contact with air with- out melting, but when heated in the air burns to form vanadic anhydride. It also unites readily with nitrogen at a red heat. It is insoluble in hydrochloric acid, but dissolves easily in nitric acid to a blue solution. Fused with caustic soda, hydrogen is evolved, and sodium vanadate produced. Of the oxides of vanadium the best known is : — Vanadic Anhydride : V 2 5 , analogous in composition to phos- phoric anhydride, and, like this, forming an acid, but little known in the free state, which is tribasic. Vanadic anhydride, obtained by gently roasting ammonium vanadate, is a reddish powder, which melts when heated to redness, and crystallizes on cooling. It is slightly soluble in water, giving a solution of a yellow colour, but 330 Text-Book of Inorganic Chemistry. dissolves easily in acids or alkalies, forming with the latter soluble crystalline vanadates. Metallic zinc brought into contact with a hydrochloric acid solution of vanadic acid produces first a blue colour, which after- wards changes to green. The extraction of vanadic anhydride from iron ores containing vanadium may be carried out in the following manner. The ores are finely powdered, mixed with nitre and heated. The mass, which contains potassium vanadate, is then extracted with water, nearly neutralized with nitric acid and then precipitated with barium chloride. Insoluble barium vanadate is thus produced, which is washed with water, and then decomposed by boiling with sulphuric acid. The acid liquid is then separated from the barium sulphate, neutralized with ammonia, concentrated, and solid pieces of ammonia chloride added, when ammonium vanadate, which is soluble in water, but insoluble in ammonium chloride, gradually separates out. This salt, when washed with ammonium chloride solution, dried and gently heated in an open crucible, leaves vanadic anhydride behind. NIOBIUM AND TANTALUM. These rare elements are but little known either in the free state or in combination with other substances. Their compounds are most nearly related to those of vanadium. They both form anhy- drides, corresponding to vanadic anhydride, and having therefore the composition : Nb 2 5 niobic anhydride and Ta 2 6 tantalic an- hydride, in which the elements are pentads. Niobium and tantalum nearly always occur associated with one another in nature ; they are found in the minerals columbite and tantalite, which are essentially ferrous niobate and tantalate, in samarskite, a niobate and tantalate of the metals of the yttrium and cerium groups, and in some other still rarer minerals, from which the oxides are best extracted by fusion with acid potassium sulphate. ., Niobic and tantalic anhydrides are white powders, yellow when hot, and insoluble in water, acids, or alkalies. They are principally distinguished from one another by their very different specific gravity. 33i THE METALS. The classification of the elements into the two great divisions, metals and non-metals, dates from a time when a number of physical properties were ascribed to the metals, which were thought to be peculiar to them and which were considered suit- able for their distinction from the non-metals. Such characteristic and peculiar properties of the metals were, for example, their lustre (whence the term 'metallic lustre'), their conductivity for heat and electricity, a high specific gravity (exceeding 6), opacity to light, and others. We now know that not one of these physical characteristics belongs exclusively to the metals. Iodine, tellurium, and graphite possess metallic lustre. Tellurium conducts heat, and graphite, as well as selenium (in a less degree) conduct electricity. On the other hand, metals are known — e.g. potassium and sodium — of which the specific gravity is not only less than 6, but which are even lighter than water. The exact meaning, therefore, of the words metal and non-metal cannot be sharply defined — it is impossible to say what properties are exclusively peculiar to the former and what to the latter, just as it is impossible to exactly classify any series of natural substances. This, however, by no means prevents us from grouping to- gether, according to artificial rules, certain classes of natural phenomena and substances which possess some similarity with one another ; and by so doing we not only obtain a clearer con- ception of the entire subject, but also make it easier for the student. It cannot be denied that a more comprehensive view of the more than sixty elements known to modern chemistry is obtained if we classify those resembling one another into certain groups. This has already been done with the elements of the sulphur and nitrogen groups, and with the halogens. For the same reason, it is therefore advisable to retain the division of the whole of the ele- ments into non-metals and metals ; and since the application of 332 Text-Book of Inorganic Chemistry. certain physical properties for a sharp classification has proved useless, we must seek for other peculiarities which shall serve- to separate oxygen, hydrogen, carbon, the halogens, the members of the sulphur and nitrogen groups, &c, from potassium, aluminium, zinc, lead, silver, &c., when writing or speaking of chemical facts. An artificial classification of the elements cannot be obtained from their physical properties, and we must therefore base our divisions on their chemical behaviour. For this purpose we can use the chemical character of their compounds with oxygen. The larger number of the elements when united with oxygen produce oxides of a basic character, the other elements produce almost exclusively oxides with acid properties. To the latter class belong those elements which we have as yet discussed, a?td which are called non-metals. Those elements of which the oxides usually possess basic properties are the metals. That this division of the elements into non-metals and metals is an artificial and not a natural classification, and that it cannot therefore be consistently carried out, is proved by the following considerations : — (i.) Oxygen belongs to neither class ; but is reckoned with the non-metals, (ii.) Hydrogen often plays the part of a metal, and its oxide, water, behaves sometimes as a base and sometimes as an acid, (iii.) Antimony, although included under the non-metals, forms a basic oxide with oxygen — antimonous oxide, (iv.) The metals manganese and chromium form oxides of acid properties — manganic and chromic anhydrides — as well as others of basic properties, (v.) The basic oxides of many metals — e.g. lead oxide, zinc oxide — play the part of acids in some com- pounds. Lead oxide, which combines with nitric acid to form lead nitrate, also unites with potash and forms a salt — potassium plumbite. In the former compound it plays the part of a base, in the latter that of an acid. It is clear, therefore, that the division of the elements accord- ingly as they produce acid or basic oxides cannot, like every other artificial classification, be carried out in every particular ; it is only suited to separate certain elements which differ from one another in certain important particulars. The Chemical Constitution of Salts. With the definition of a metal is connected that of a salt, since all salts of inorganic chemistry (if we consider ammonium as a Constitution of Salts. 333 metal) are metallic compounds ; and a short, accurate definition of a salt is scarcely easier than that of a metal. On p. 78 it has been stated that salts are produced by the union of acids with bases, which tells us how they are produced, but not of what they actually consist. Salts are binary compounds consisting of an electro-positive constituent — e.g. a metal, and an electro-negative l constituent — e.g. a halogen or acid radical. They are usually decomposed by electrolysis into these two constituents — the former separating at the negative, the latter at the positive pole. Common salt (sodium chloride) when fused or dissolved in water is decomposed by an electric current into chlorine and sodium, of which the former goes to the positive pole and the latter to the negative. In the presence of water the sodium is decomposed, so that instead of sodium we obtain hydrogen and sodium hydrate at the negative pole. Salts of this kind contain a metal in direct union with a nega- tive element, the halogen. They are called haloid-salts because only the halogens and the compound substances resembling them, cyanogen, sulphocyanogen, &c, but not the other non-metals, can thus unite directly with the metals to form salts. Another class of salts, called ampkid-salts, consist also of an electro-positive and electro-negative constituent, which, however, are not directly united with one another, but are combined through the intervention of one or more atoms of oxygen or sul- phur. If the uniting element is oxygen, the compound is called an oxysalt, if sulphur a sulphosalt. A haloid-salt may be converted into an oxysalt by the intro- duction of an atom of oxygen — for example, in the conversion of sodium chloride into sodium hypochlorite : NaCl + O = NaOCl, in which the sodium and chlorine are not now directly united with one another, but through the atom of oxygen, common to both. The negative portion, and sometimes the positive portion also, of the oxysalts usually consists of a compound radical. Thus, potassium nitrate : N0 2 -OK, contains the monad radical nitric peroxide (nitryl) : N0 2 , sodium sulphate: S0 2 Q Na contains the dyad radical sulphurous anhydride (sulphuryl) : S0 2 , and sodium 1 Using the obsolete but convenient nomenclature of the electro-chemical theory.— Ed. 334 Text-Book of Inorganic Chemistry. fONa phosphate : PCM ONa the triad radical : PO, which has not yet lONa been isolated. And these three acid radicals are all united to their respective metals — the positive portion of the salt, by one, two, or three atoms of oxygen respectively. 1 Two metals may also be united with one another by oxygen, as in potassium plumbite : Pbj _.„ in which the lead is the negative element, and the potas- sium the positive. As an example of the sulphosalts may be mentioned the com- pound produced by the combination of carbon disulphide with 1SK q,- corre- sponding to ordinary potassium carbonate : COj JxJ; in which compounds the respective oxygen and sulphur atoms may displace one another. For although a compound of the composition: o-^- is at present unknown, the previously described (p. 168) sodium thiosulphate : S0 2 j o»j is an example of these little known inorganic compounds. In organic chemistry they are much more numerously .represented. Besides these two elements, oxygen and sulphur, it also appears that in certain cases the' element fluorine can assume a dyad character, and so serve to unite the negative and positive portions of a salt (p. 136). We also distinguish salts called normal acid, and basic salts. Monobasic acids— i.e. those containing only one atom of hydrogen displaceable by one atom of a monad metal, e.g. nitric acid : N0 2 -OH — can only form normal salts with monacid bases — i.e. those containing only one atom of hydrogen displaceable by an acid radical, e.g. caustic potash : KOH. Acid salts can only be produced from polybasic acids— i.e. those which contain more than one atom of displaceable hydrogen, as sulphuric or phosphoric acid, when a portion of their hydrogen is displaced by a metal (comp. p. 56). In the same manner, basic ( O T-T salts are formed from polyacid bases — e.g. lead hydrate \ Pb- n „ 1 The electrolysis of such an oxysalt has already been explained (p. 79). — Ed. Constitution of Salts. 335 (OH and ferric hydrate : Fe OH when a portion only of their hydro- lOH gen is displaced by an acid radical or by a haloid. The following are some examples of basic salts : — Basic lead nitrate .... FblS^ * Basic lead chloride . . . . Pb CI OH fON0 2 Basic ferric nitrate .... FejOH (oh Basic ferric sulphate .... Fe|2|?° 2 Basic ferric chloride . . . . Fe j £7, Many compounds which are called basic compounds, but which do not contain the acid radical or haloid and the metal in equiva- lent proportions, are best considered as mixtures of a normal salt and a metallic hydrate. Thus the basic antimonous chloride, referred to on p. 257, and sometimes represented by the formula : Sb - p, contains varying quantities of oxygen and chlorine united with the same quantity of antimony, and is, therefore, better con- sidered as a mixture of antimonous chloride : SbCI 3 , and anti- monous oxide : Sb 2 3 . (OH Atriacid base — e.g. aluminium hydrate : Al- OH requires three (OH molecules of monobasic acid in order to produce a normal salt, while at the same time three molecules of water are also produced : Al(OH) 3 + 3N0 2 OH=A10 3 (N0 2 ) 3 + 3H 2 0; and in the same way a molecule of the same base will require only one molecule of a tribasic acid: Al(OH) 3 + PO(OH) 3 = A10 3 -PO + 3H 2 0. If, how- ever, we have a dibasic acid united to the same triacid base, we shall require three molecules of the former to two of the latter. Thus, normal aluminium sulphate is produced in this way : — 2Al(OH), + 3S0 2 (0H), = A1 2 6 (S0 2 ) 3 + 6H 2 0. Normal salts may also be produced by the union of two mole- cules of a dibasic acid with one of a triacid, and one of a monacid 336 Text-Book of Inorganic Chemistry. base. Common alum: ,,„ 3 [ 2SO,, for example, is a normal sulphate of the two metals aluminium and potassium, and is thus formed : — Al(OH) 3 + KOH * 2SO,(OH) 2 = ^Uo, + 4H,0. From these examples it will be seen that a salt may be defined either as an acid in which the hydrogen is displaced by a metal, or as a base in which its hydrogen is displaced by an acid radical, or its hydroxyl by a haloid element or some similar group of elements. The normal salts often have no action on vegetable colours — i.e. are said to be neutral, and hence this class of salts is sometimes called neutral salts. But some normal salts have a strong alkaline or acid reaction, and even some acid salts react occasionally alka- line, while, on the other hand, some acid salts possess a neutral reaction. Red litmus paper is strongly blued by normal potassium ( OK carbonate: CO J OK and by normal potassium phosphate : PO-! OK 1 0K I OK (OK and even faintly by monacid potassium phosphate : PO \ OK On (OH the other hand, acid sodium carbonate : CO i (-,„ reacts neutral. The cause of this is that the strongly basic properties of potassium hydrate are not neutralized by the weak carbonic acid, and that phosphoric acid, although one of the strongest acids, does not completely neutralize the potash in the monacid phosphate. Double Salts. — Compounds in which the chemicai affinities of the combining atoms are completely satisfied are called saturated compounds. Such, for example, are potassium chloride : KC1, and platinum tetrachloride : PtCl,. And it might be thought that these two compounds could not be capable of any further combination. But in spite of this they both unite to form a douhle, and very stable salt of the composition : 2KCl,PtCI 4 . And a large number of other similar double salts are also known. In most of these cases where perfectly saturated compounds can unite with one another chemically and form new compounds, we are unacquainted with the forces which cause this union. Such compounds have been called molecular compounds — i.e. compounds Constitution of Salts. 337 produced by the chemical union of molecules. This expression, of course, explains nothing, but simply states that the molecules of saturated compounds are capable of uniting with one another to produce new compounds. The constitution of the haloid double salts may be expressed in the following manner. From a consideration of the composition of numerous double fluorides, it appears that two atoms of monad fluorine can coalesce to form what plays the part of a dyad atom. Thus in the compound KF, HF (comp. p. 136), we can consider that the two monad atoms of fluorine unite to form one dyad atom of double the atomic weight, which maybe indicated by the symbol *. This double compound then becomes KFH, analogous to KOH. Just in the same manner the double chloride of potassium and platinum, .which may be written : 2KCl,(PtCl 2 )Cl 2 , may be considered as two atoms of potassium united with the dyad radical PtCl 2 by the other four atoms of monad chlorine, which have coalesced to form two dyad atoms, thus : — k 2 ci 2 , (Ptcyci, = £fj)ptci 2 . In the same manner the double iodide of potassium and mercury: 2KI, Hgl 2 may be written : -^t [ Hg, and the double cyanide of po- tassium and silver, K€-yAg. The constitution of the potassium ferrocyanides, as well as other similar compounds, may be inter- preted in the same way. The oxysalts also unite together to form double salts. Zinc sulphate, for example, unites with potassium sulphate to form the compound : S0 2 2 Zn + S0 2 (OK) 2 + 6H 2 0, which is produced from zinc sulphate by the substitution of a molecule of water of consti- tution, by potassium sulphate. SPECTRUM ANALYSIS. It has been previously mentioned (p. 308), that a non-luminous flame becomes luminous when a solid body is raised to incande- scence in it, and then emits white light. But if metallic compounds are introduced into such a flame and there volatilize, for which the metallic chlorides are best adapted, the heated vapour of the metal imparts a colour to the flame, peculiar to each. Thus z 338 Text-Book of Inorganic Chemistry. compounds of sodium colour the flame an intense yellow, those of potassium violet, and those of lithium crimson. But owing to the similarity of many of these colours when viewed by the naked eye, they are not always suitable for directly determining which metal may or may not be present. It is often difficult to say by mere observation whether a given flame is coloured red by strontium or calcium, or whether a flame coloured yellow by sodium, also con- tains potassium, since the bright yellow sodium flame can com- pletely mask the violet colour due to potassium compounds, even when the latter substance is present in considerable quantities. But something quite different is observed if, instead of simply looking at a flame with the naked eye, we allow the light to fall through a narrow slit on to a triangular prism of glass and then magnify the bent rays of light by a telescope. Light of different colours in passing through the prism is differently bent, and ordinary white light gives a band of colours extending from red to violet — a so-called continuous spectrum. The light from a glow- ing gas is broken up into its constituent parts in the same way : the spectrum in this case will not, however, be a continuous one but will consist of bright lines separated by more or less dark spaces. And these bright lines always appear in those parts of the spectrum to which they belong ; thus the red lines will always be in the red, the blue in the blue portion, and so on — each gaseous element having its own peculiar spectrum. The coloured table of spectra forming the frontispiece shows firstly the continuous spectrum of sunlight and then the spectra of ten metals, which differ much from one another, both with regard to the number and to the position of their lines. Sodium, lithium, thallium, and indium possess at the most two lines each, while caesium, rubidium, strontium, calcium, and barium have much more complex spectra. Caesium, strontium, and indium have all blue lines, but their position, when compared with the same scale, is different. As is shown in the table, the blue strontium line is most to the left, then follow the two blue lines of caesium more to the right, and still further to the right is the blue indium line. A substance which is supposed to contain thallium, and which like thallium gives a green line but in a different position — either more to the right or to the left— certainly does not contain this metal. Such an observation would rather point to the existence of a new (hitherto unknown) element ; for, as previously remarked, the position of the bright lines in the spectrum of any one element Spectrum Analysis. 339 is always the same under all conditions. In this way the new metals, rubidium, caesium, thallium, indium, and gallium, have been dis- covered by the aid of spectrum analysis, and it is extremely probable that others will be added to this list. And although, as previously intimated, potassium compounds, even when present in consider- able quantity, cannot be recognized by the naked eye in a flame in which sodium is volatilized at the same time, if we allow the light to fall through a slit and decompose it by the prism, the lines of both metals are easily seen in the spectrum. The apparatus for spectrum analysis — called a spectroscope — as devised by Bunsen and Kirchhoff, to whom this delicate method Fig. 58. of chemical analysis is due, is shown in one of its simplest forms in fig. 58. On an iron stand are fastened three brass tubes, and between these is placed the glass prism. At the outside end of the tube A is a narrow slit, which can be widened or contracted at will, and through which the light from the heated substance passes on to the glass prism ; the spectrum so produced is then observed through the telescope B. The third tube, C, contains a small reduced photographic scale, illuminated by a small gas flame 340 Text-Book of Inorganic Chemistry. placed in front of it, and an image of which is thus reflected from one face of the prism into the telescope B. This serves to fix the position of any line in a given spectrum. All extraneous light must, of course, be excluded from the prism by covering it with a black- ened box or cloth. A little practice with the spectroscope soon enables one to determine the presence or absence of any metal in a given sub- stance. Those metals which are only volatilized with difficulty (e.g. iron), and which therefore do hot become gaseous in the gas flame, are investigated by allowing powerful electric sparks to pass between poles of the metal, or between small carbon poles saturated with a solution of its chloride. Gases yield spectra in the same way. Hydrogen, when heated by the passage of electric sparks, appears bright red, and its spectrum consists essentially of a bright red, a green, and a blue line. The results of spectrum analysis have also made it possible to determine the composition of the sun and other fixed stars. And in no heavenly body has an element been discovered which is un- known on the earth. The great delicacy of spectrum analysis makes it especially valuable for chemical purposes. Most minute quantities of various metals, which it was previously quite impossible to detect, can by recognized with the greatest certainty by spectrum analysis. It has been calculated from experiments that one hundred-thousandth part of a milligramme of lithium chloride, and even one three- millionth part of a milligramme of sodium chloride may be easily detected by their spectra. CLASSIFICATION OF THE METALS. It has been repeatedly attempted to discover some physical or chemical properties of the metals which might serve to classify them into a series of divisions and sub-divisions. But, however desirable such a classification may be in simplifying the study of this large group of elements, all such attempts have only strength- ened the conviction that the metals cannot be exactly divided upon any natural system into a number of distinct classes, and that, if such a classification is made, it must be according to some Classification of the Metals. 341 artificial rules. The properties of the oxides and sulphides of the metals may serve as the basis of such an artificial classification. One class of the metals forms oxides, or rather hydrates, which are easily soluble in water, the solutions of which have a powerful alkaline reaction ; these hydrates are called the alkalies, and the corresponding metals the metals of the alkalies. The following five metals belong to this class — viz. potassium, sodium, lithium, rubidium, and ccesium, and the hypothetical compound metal ammonium. A second class also yields soluble oxides, though far less soluble than the alkalies. These oxides also react faintly alkaline, but are at the same time allied to the oxides of the next — the so- called earths — and their corresponding metals are therefore called the metals of the alkaline earths. In this class the metals calcium, strontium, barimn, and magnesium are included. The last-named metal, of which the oxide is only very slightly soluble in water, forms a connecting link with the next group. The third group, of which the oxides are completely insoluble in water, and are called earths, are the metals of the earths. The chief representative of this group is aluminium, with which are associated the rarer metals, beryllium, indium, gallium, yttrium, erbium, cerium, lanthanum, didymium, and zirconium, as well as some still rarer elements. The sulphides of this group (except indium) are not produced on the addition of sulphuretted hydrogen or an alkaline sulphide to a solution of one of their salts. The precipitate produced on adding sodium sulphide or ammonium sulphide to their solutions usually consists of the insoluble hydrate, while sulphuretted hydrogen is. given off: — A1 2 C1 6 + 3Na 2 S + 6H 2 = 2A1(0H) 3 + 6NaCl + 3H 2 S. A fourth class includes those metals of which the sulphides are insoluble in water, but mostly soluble in dilute acids. These com- pounds are precipitated on adding an alkaline sulphide, such as ammonium sulphide, to a solution of one of their salts. They are not, however, produced when sulphuretted hydrogen is led through an acid solution of their salts. The metals of this group are : iron, manganese, chromium, uranium, cobalt, nickel, and zinc. The fifth group includes those metals of which the sulphides are insoluble both in water and in dilute acids, and are therefore precipitated when sulphuretted hydrogen is led through an acid 342 Text-Book of Inorganic Chemistry. solution. They are : lead, thallium, bismuth, cadmium, tin, and copper. The metals belonging to the fourth and fifth groups are some- times called the heavy metals. The heaviest or so-called noble metals are : mercury , silver , gold, platinum, iridium, palladium, rhodium, ruthenium, and osmium. The sulphides of this group are also insoluble in dilute acids. We now proceed to the discussion of the metals arranged in these six classes. The Metals of the Alkalies. At the commencement of this century it was universally believed that every metal must possess a comparatively high specific gravity — i.e. be considerably heavier than water (see p. 331). But when the metals of the alkalies were discovered by Davy this opinion was shown to be incorrect. Potassium and sodium are lighter than water, and lithium, which was discovered later, is only about one- half as heavy. The metals of the alkalies are characterized by their strong chemical affinities for negative elements, and are therefore well adapted for withdrawing oxygen, sulphur, chlorine, &c, from com- bination with other elements. They belong to our most powerful reducing agents. Water is decomposed by them even at the ordi- nary temperature of the air. Their oxides and-sulphides are easily soluble in water and are powerful bases. With the halogens they produce, like most other metals, neutral compounds, all of which are soluble in water. POTASSIUM. Chemical Symbol: K. — Atomic Weight: 39. The strong attraction of the metals of the alkalies, and especially of potassium, for oxygen, makes it impossible for them to occur in the free state in nature. Potassium is chiefly found in the mineral kingdom in combination with oxygen and silica as a silicate, and in this form is a constituent of potash felspar ; it also occurs, united with chlorine, as potassium chloride. The upper layer of the salt deposits at Stassfurth in Germany consists principally of Potassium. 343 potassium chloride, which is largely used for the manufacture of other potassium compounds. Potash felspar (orthoclase), which may be considered as a double compound of aluminium silicate and potassium silicate, and which is one of the constituents of granite, gradually becomes disintegrated under the influence of frost, and is then decomposed by water containing carbonic acid into insoluble aluminium silicate (clay) and soluble potassium silicate, or carbonate, which fertilizes the soil. Nearly all land plants require potassium compounds for their growth and even for their existence, just as we require salt in our food. These potassium compounds are absorbed by their roots and converted in their structure into potassium salts of organic acids. And when the plants are burnt nearly the whole of the potassium which they contain remains behind as potassium car- bonate or potash. This was formerly the material from which potassium compounds were almost exclusively prepared. Potassium is a silver-white metal with a bright lustre. Its specific. gravity is 0-87, it melts at 62 , and can be distilled at a red heat in a stream of hydrogen, giving a vapour of a bright green colour. At the ordinary temperature the metal is soft and can be easily moulded or cut with a knife, but at o° it is hard and brittle. In consequence of its strong attraction for oxygen, the freshly-cut surface is only lustrous for a moment, and immediately afterwards becomes covered with a thin layer of oxide of a bluish colour. The white lustrous surface can only be preserved by melting the metal in a glass tube filled with hydrogen, then exhausting the gas, sealing up hermetically, and allowing the liquid metal to spread over the tube. When slowly cooled, the metal can some- times be obtained in the crystalline form. Potassium is now never prepared by the process by which Davy first obtained it — the decomposition of fused caustic potash by an electric current — as this only yields a very small quantity at one time. It is far better to reduce potassium carbonate with charcoal at a red-heat. The chemist is no longer compelled to prepare his own potassium, as it is now manufactured in chemical works at a far cheaper rate than it can be prepared in small quantities. Nevertheless, potassium is still an expensive sub- stance, principally because it is only used in small quantities in the arts. Its price is about 6s. per ounce. To obtain a good yield of potassium from potassium carbonate 344 Text-Booh of Inorganic Chemistry. and carbon, it is very necessary that the two substances should be mixed together as intimately as possible. And as this can scarcely be done by mechanical means, it is better to employ some organic salt of potassium rich in carbon, of which purified tartar, acid potassium tartrate, is the most suitable. The salt is heated in a covered crucible, first gently and then to redness, and the carbona- ceous residue, consisting of an intimate mixture of potassium car- bonate and charcoal, is then introduced into an iron retort. The Fig. 59- retort is afterwards heated to whiteness in a furnace, and the potas- sium vapour which distils over condensed in a suitable receiver. One of the wrought-iron bottles in which mercury is imported answers well for the retort. It lies in the furnace in a horizontal position, supported on iron bars, as shown in fig. 59. A short wide iron tube, open at both ends, is screwed air-tight into the re- tort and leads to the upper part of the receiver, which contains a little naphtha, and is cooled externally by ice. The potassium then condenses in the receiver and collects under the naphtha. The operation is not without danger, for the iron tube may easily Potassium. 34S become plugged with substances less volatile than the potassium, and so produce a violent explosion. Great care must therefore be taken to keep this tube free from obstruction. This can be done by frequently inserting an iron wire bent into a corkscrew shape through a hole in the condenser into the tube and moving it back- wards and forwards. 1 The potassium which condenses in the receiver is afterwards rectified by distilling it in an earthenware retort, or in one of hard glass, covered externally with a layer of clay. It is thus obtained in small pellets of about the size of a hazel-nut, which, if kept under naphtha in well-closed bottles, can be preserved for a long time. They are generally coated with a thin brownish crust. If potassium is exposed to the air it takes up oxygen, water, and carbonic acid, with evolution of heat, and becomes converted into a mixture of potassium hydrate (caustic potash) and potassium car- bonate, which gradually becomes a syrupy solution of potassium carbonate : — 2K + 3H a O = 2KOH + H 2 and 2KOH + C0 3 = CO(OK) 2 + H 2 0. When heated in the air potassium burns with a violet flame and produces its peroxide. The metal decomposes water at the ordinary temperature with a large evolution of heat, and appears to burn with a violet flame when thrown upon water. What really burns is the hydrogen which the potassium liberates from the water in its conversion into potassium hydrate, and which is ignited by the large amount of heat evolved during the reaction. The violet colour of the burning hydrogen is due to small particles of potassium with which it is mixed. Potassium also burns in chlorine, and forms potassium chloride. At the ordinary temperature it does not unite with hydrogen. But at 200 , and especially between 300 and 400 , it begins to absorb the gas and becomes changed into a lustrous, brittle, crystalline sub- stance, which catches fire on exposure to the air, and is again de- composed at a higher temperature. This compound may have the composition : KH. 1 Violent explosions are also sometimes produced in the manufacture of potassium by the formation of a peculiar black compound of potassium and carbon which very easily decomposes. These explosions may be prevented if the potassium vapour is rapidly cooled by attaching a flattened receiver to the iron deliver}' tube. This exposes a large surface to the air and so rapidly cools the hot vapour. — ED. 34-6' Text-Book of Inorganic Chemistry. COMPOUNDS OF POTASSIUM. Potassium unites with oxygen in two proportions, to form potassium oxide : K 2 0, and potassium peroxide : KO,,. In the former of these, and all its other compounds except the peroxide, the metal plays the part of a monad ; in the latter it appears to have a higher atomicity. Both compounds are only of theoretical interest. If potassium is heated in dry oxygen or air, it forms not the oxide but the peroxide as a yellow powder, which melts at a higher temperature and solidifies on cooling to a crystalline mass. This peroxide is decomposed by water, or even by exposure to moist air, forming caustic potash, and giving off oxygen. Potassium oxide : K 2 0, which cannot be obtained by heating potassium in oxygen, is produced when caustic potash is fused with as much potassium as it already contains : KOH + K = K 2 + H. It is a white powder, which melts at a red-heat, energetically attracts moisture from the air, and in contact with water becomes converted into potassium hydrate, with a large evolution of heat : K 2 + H 2 = 2KOH. Potassium Hydrate, Caustic Potash: KOH. — This important compound is a white crystalline solid, melting to a clear liquid at a red-heat. It is strongly alkaline and caustic, deliquesces in the air, and is easily soluble both in water and alcohol. An aqueous solution of pure potassium hydrate may be obtained by dropping small pieces of potassium into water. And from this solution the solid may be prepared by rapidly evaporating the liquid in a silver basin, silver being much less attacked by caustic potash than platinum. The ordinary method for the preparation of caustic potash on the large scale consists in dissolving potassium carbonate in about eight or ten times as much water, heating the solution in an iron vessel to boiling, and then gradually adding slaked-lime which has been previously mixed with water to a thin paste. The lime then combines with the carbonic acid, forming calcium carbo- nate, which is insoluble in water, and caustic potash remains in solution. A more concentrated solution of potassium carbonate cannot be employed, because it would only be incompletely de- Compounds of Potassium. 347 composed by the lime. Strong caustic potash decomposes calcium carbonate when boiled with it, forming potassium carbonate and lime. The quantity of quick-lime ' (calcium oxide) which is re- quired to convert a given weight of potassium carbonate into potassium hydrate is calculated according to the following equa- tion : — CO(OK) 2 + CaO + H 2 = CO-0 3 Ca + 2KOH, and weighed out. This lime is next slaked, mixed with water, and the milk of lime dropped into the boiling potassium carbonate until a small quantity of the clear liquid no longer effervesces with hydrochloric acid — a proof that it no longer contains potassium carbonate. The fire is then removed, the vessel covered with an iron lid and allowed to cool. As soon as the calcium carbonate has settled, and the liquid become clear, it is run off with a syphon, and rapidly evaporated in a clean iron or silver dish, when solid crystalline potassium hydrate remains behind. Potassium hydrate appears to form several crystalline com- pounds with water, and the solid commercial hydrate usually contains more or less water, up to 20 per cent. This water may be removed by heating the compound in a silver dish until it fuses quietly. A saturated aqueous solution of potassium hydrate de- posits crystals on cooling, consisting of one molecule of the hydrate united with two of water : KOH + 2H 2 0. Caustic potash prepared in this way is never pure. It contains firstly the impurities present in the commercial potassium car- bonate, such as potassium chloride, sulphate, &c. ; secondly, the impurities of the crude lime ; and, thirdly, potassium carbonate, due to the absorption of carbonic acid from the air during the evaporation of the solution. To prepare the pure substance from the crude hydrate, it is broken into pieces and shaken up in a closed vessel with cold alcohol. This liquid dissolves the potas- sium hydrate, but leaves all the impurities, including the potassium carbonate, behind undissolved. The clear solution is then run off, diluted with water, and heated, first to drive off the alcohol and then more strongly to expel the water. If the alcoholic solution were heated alone without the addition of water, the liquid would be decomposed under the influence of the potash and the oxygen of the air, and would acquire a brown colour and unpleasant odour. Small quantities of chemically pure potassium hydrate may be- 348 Text-Book of Inorganic Chemistry. prepared by boiling a solution of pure potassium carbonate with pure lime obtained from Iceland spar or white marble. But as caustic potash, as well as its solution, readily dissolves both silica and alumina, it can neither be prepared nor evaporated in vessels of glass or porcelain ; it energetically attacks both these substances. The fused caustic potash of commerce is the hydrate which has been melted and strongly heated as long as it loses water. It is sometimes brought into trade in the form of rough flat pieces, showing a crystalline fracture, or else as sticks of about the thickness of a black-lead pencil, which are obtained by casting the fused substance in iron moulds made in two pieces (see fig. 60). Fig. 60. Solid caustic potash rapidly deliquesces in the air, absorbing water and carbonic acid, and producing a colourless liquid with a strong alkaline reaction. This liquid gives off its carbonic acid with effervescence when treated with an acid. Aqueous caustic potash gives an oily, slippery feeling to the fingers, and possesses a sharp alkaline taste. It powerfully attacks the epidermis, and is therefore largely used as a strong caustic. The solution dissolves grease and fats, with which the alkali unites and forms soluble compounds called soaps ; a dilute potash lye (or solution of the alkali) is therefore employed for cleansing purposes. Caustic potash, both in the solid form and in solution, is l.argely used in the laboratory. In its attraction for acids it Caustic Potash. 349 excels all other bases, 1 and is, therefore, the most suitable sub- stance to separate them from their salts. If a solution of a salt of copper, for example, is mixed with caustic potash, the correspond- ing potassium salt is produced, and copper hydrate separates as a precipitate : — SOj-OjCu + 2-KOH = S0 2 (OK) 2 + Cu(OH) 2 . Caustic potash is also largely used in the arts, but is gradually becoming displaced by the cheaper caustic soda, except when the latter is too feeble to effect the required change — for example, in the conversion of cellulose into oxalic acid (see p. 298). 2 The salts of potassium can be easily prepared by the union of the base with the corresponding acid. In this small text -book we can only refer to the most important of this numerous class of compounds. (OTC q^ is easily obtained by neutral- izing potassium carbonate with sulphuric acid, and crystallizes when a concentrated solution is allowed to cool in hard four-sided pyramids or prisms, without water of crystallization. It possesses a saline, bitter taste, and melts when strongly heated without change. It is not so easily soluble in water as the other potassium salts : 100 parts of water at 12° only dissolve twelve parts of the salt, and at ioo c only twenty-six parts. Alcohol does not dissolve it in the least. f OH Acid Potassium. Sulphate: S0 2 1q K is prepared by heating the normal sulphate with the requisite quantity of sulphuric acid, and is a bye -product when nitric acid is manufactured from potassium nitrate (p. 190). It is easily soluble in water, and crystallizes from the hot concentrated solution in rhombohedra ; it reacts and tastes acid. When heated to about 200°, it melts quietly; at a higher temperature it gives off water, and is converted into potassium disulphate like the sodium salt (p. 167) : — so fOH _ fS0 2 OK + HQ 2bU 2|OK _ U (S0 2 OK + ti " u - 1 Except the oxides of the rare metals rubidium and caesium. — Ed. 2 Owing to the large quantities of caustic putash now sent into trade from the German manufactories, its present retail price scarcely exceeds that of caustic soda. — Ed. 350 Text-Book of Inorganic Chemistry. At a still higher temperature this salt splits up into sulphuric anhydride and the normal sulphate. Hence the use of acid potassium sulphate to decompose certain minerals which are not attacked by acids. Potassium Nitrate {Nitre, Saltpetre)] N0 2 -OK, crystallizes from a hot solution in long, furrowed, rhombic prisms, with a saline cool taste, and may be easily prepared by neutralizing nitric acid with potassium carbonate and evaporating down. Nitre is also produced naturally in countries where the soil is rich in potassium salts and where rain seldom falls, or only at periodic intervals (Egypt, Persia, India, &c). In these countries the surface of the ground after the rains gradually becomes covered with a white efflorescence, which consists chiefly of potassium nitrate. The salt is then obtained by simple solution in water and re-crystallization. This nitre owes its production to the potassium salts and nitrogenous organic matter present in the soil. The latter gradually undergoes decay, and its nitrogen is converted into ammonia, which, in the presence of the strong base, becomes oxidized to nitric acid, forming potassium nitrate. This process has been imitated in some European countries, and large quantities of nitre have been artificially produced in the following manner. Heaps of earth mixed with some potassium carbonate and lime were made under open sheds and then re- peatedly watered with liquid manure or other nitrogenous organic refuse. The mass was frequently turned so as to expose it as much as possible to the oxidizing influence of the air. After a considerable period (often one or more years), when all the nitro- gen had been converted into nitrates, the heaps were lixiviated with water to dissolve these soluble salts, and the solution mixed with potassium carbonate to convert the calcium and magnesium nitrates into the soluble potassium salt. The clear solution was then run off from the insoluble calcium and magnesium carbonates and evaporated down to crystallize. 1 Considerable quantities of potassium nitrate are now prepared from the sodium nitrate found so largely in Chili. If a solution 1 This production of nitre from nitrogenous organic substances is exceed- ingly costly and wasteful, since the value of the nitrogen in the form of am- monia for increasing the fertility of the soil is far greater than that of the nitre obtained from it, and the process has now been almost entirely abandoned. — ED. Potassium Nitrate. 351 of this salt is mixed with one of potassium chloride, and the mixed solution evaporated down under pressure, double decom- position occurs into sodium chloride (common salt) and potassium nitrate. And since sodium chloride is only slightly more soluble in hot than in cold water, this salt first separates out on evaporating, while potassium nitrate, which is much more soluble in hot than in cold water, crystallizes out on cooling the hot decanted liquid. Potassium nitrate prepared in any of these ways is always mixed with impurities, especially common salt, from which it can be freed by repeated crystallization. This purification depends upon the greater solubility of saltpetre in hot than in cold water. 100 parts of water dissolve only 13 parts of the salt at 0°, and 29 parts at 18°, but at 97° the same quantity of water dissolves 236 parts. And a boiling-hot saturated solution of saltpetre, of which the boiling-point is 11 6°, is said to contain 335 parts of the salt dissolved in 100 parts of water. Nitre is used for various purposes in medicine, and sometimes for the manufacture of nitric acid, but far larger quantities are ab- sorbed in the manufacture of gunpowder, in which the nitre is the oxidizing substance. The well-known explosive action of gun- powder is due to the sudden production of large volumes of hot gases from a small volume of a solid. If powdered nitre is sprinkled on a piece of glowing charcoal, the latter burns brilliantly to carbonic acid and carbonic oxide, while the nitre is converted into potassium carbonate and free nitrogen. The same process goes on when an intimate mixture of powdered nitre and charcoal, in the correct proportions, is ignited, and may be represented by the following equation : — 2N0 2 -OK + 3C = C0 2 + CO + N 2 + C0Jq]| Gunpowder contains, however, sulphur as well as nitre and charcoal, and it was formerly thought that its decomposition when burnt was expressed by the equation : — 2 NOyOK + 3C + S = 3CO, + N 2 + K 2 S, which represents the various constituents in the proportions in which they are actually contained in gunpowder. But it has been found that other substances, besides those shown in the above equation, are also produced, and that the sulphide formed is potassium disulphide and not the monosulphide. The substances into which gunpowder is converted when burnt in a closed space are essentially : — 352 Text- Book of Inorganic Chemistry. (i.) Solids, includingpotassium carbonate, potassium sulphate, and potassium disulphide, and which are about 57 per cent, of the whole. (ii.) Gases which consist of carbonic acid, nitrogen, and carbonic oxide, and constitute the remaining 43 per cent, by weight. The volume of these gases, measured at the normal temperature and pressure, is nearly 300 times that of the powder, and at the high temperature of the explosion nearly 2,000 times. Hence the force of the explosion. Other substances are also formed when gun- powder is exploded, but only in very small quantities. And the actual chemical changes which go on when the powder is burnt can only be represented by a somewhat complex equation. The gunpowder of different countries, and even of the same country, varies considerably. The mean composition of English powder is about as follows : — Nitre . . . . 75-3 Charcoal . . . 13-4 Sulphur . . . 11-3 ioo-o In the manufacture of a suitable gunpowder it is not only necessary to mix the constituents intimately together, but the mixture must be afterwards granulated, an operation which is performed by suitable machinery. The size of the grain depends upon the purpose for which the powder is to be used : it varies from about the size of a millet seed for ordinary purposes and small arms, up to blocks as large as a cubic inch (pebble powder) for large cannon. By diminishing the quantity of nitre, and increasing that of charcoal or sulphur, the rate at which the powder burns is dimin- ished. And as the gases produced by the burning of the powder cannot support combustion, but, on the contrary, extinguish bodies in the act of burning, this property has been successfully utilized to extinguish fires in closed places. A mixture for this purpose may contain, for example, 60 to 66 per cent, nitre, 30 to 36 per cent, sulphur, and at most 4 per cent, charcoal. On burning, it produces sulphurous anhydride with carbonic acid and nitrogen, all of which extinguish burning bodies. Potassium Nitrite : NO -OK. — If potassium nitrate is melted at a low red heat, it is decomposed into oxygen which escapes, and potassium nitrite which remains behind, usually mixed either with Potassium Carbonate. 353. undecomposed nitrate, or, if too strongly heated, with caustic potash. The salt is easily soluble in water, deliquesces in the air, and evolves copious red vapours of nitrous acid when treated with moderately dilute sulphuric acid. (OK n i, deliquesces in the air and is easily soluble in water. The aqueous solution reacts strongly alkaline, and when evaporated down leaves the salt as a white solid which melts unchanged at high temperatures. An extremely con- centrated, warm solution deposits crystals containing about 16 per { OK cent, of water of the composition: 2CO j qtP + 3H s O. Potassium carbonate is a constituent of the ashes which remain when wood and land plants generally are burnt, which, however, do not contain more than 20 per cent, of it. The ashes consist chiefly of the sulphates, phosphates, and silicates of potash, lime, and magnesia, together with some chlorides. Crude potash is obtained by extracting the ashes with water, and evaporating the filtered solution to dryness. It contains at most 60 per cent, of potassium carbonate mixed with the other soluble salts contained in the ashes. By digesting this crude potash with an equal weight of water for several days, with frequent stirring, the easily soluble carbonate is dissolved, leaving most of the other salts behind. The clear solution evaporated to dryness gives the fiearlash of commerce. This is still very far from pure potassium carbonate, and in fact we know of no method of completely purifying it. Chemically pure potassium carbonate may, however, be easily obtained by heating a pure potassium salt of an organic acid in a platinum vessel. Vessels of porcelain or glass cannot be employed, as the hot potassium carbonate abstracts silica from them, and so becomes impure. For the preparation of the pure salt, acid potassium tartrate is commonly used. This salt can be readily obtained, and, as it is only difficultly soluble in water, can be easily purified. The heated mass, when extracted with water, and the clear solution evaporated in a platinum dish, leaves pure potassium carbonate as a snow-white solid. Since the discovery of the potassium chloride beds at Stassfurth, potassium carbonate is manufactured from the chloride, just as soda is obtained from common salt by Leblanc's process (seep. 371). *AA 354 Text-Book of Inorganic Chemistry. Potassium carbonate is one of the most important of the potassium salts. It is used not only for the preparation of caustic potash and many other potassium compounds, but is also largely employed in the arts, especially in. the manufacture of glass and of soft soap. (OFT OK is obtained by leading carbonic acid into a concentrated aqueous solution of the normal salt, when it separates out as a crystalline powder : — CO {ol + C0 » + H *° ' 2CO |oH It requires about four times its weight of water for solution, reacts neutral, and is decomposed at about 8o° into the normal salt and carbonic acid. Potassium Chlorate: C10 s -OK. — This salt, of which the pre- paration has been already described (p. 1 10), crystallizes when its hot saturated solution is cooled in lustrous scales or larger plates. It melts at 334 , and begins to be decomposed, with evolution of oxygen, at 352 . It requires 16 parts of water for solution at the ordinary temperature, but only 2 parts at 100°. In consequence of the readiness with which it parts with oxygen, it is an excellent oxidizing agent, and yields considerable quantities of chlorine when warmed with hydrochloric acid (p. 112). Potassium Perchloraie : C10 3 -OK, has been already described. It belongs to the potassium salts which are difficultly soluble in water, and can, therefore, be readily purified by crystallization. The salt is insoluble in alcohol. Potassium Hypochlorite : ClOK. — This salt, which is only known in its aqueous solution and which is characterized by its powerful bleaching properties, is obtained by passing chlorine into a cold, dilute solution of potassium .carbonate until large quantities of carbonic acid are evolved. Potassium chloride and hypochlorite with acid potassium carbonate are first produced : — 2CO |oK + 2C1 + H *° = 2CO |oH + KC1 + C10K ' and carbonic acid is, therefore, not evolved at once, but only when the chlorine begins to attack the acid potassium carbonate. Potassium Chlorate. 355 This solution of potassium hypochlorite is known as eau de Javelle. Strong acids liberate considerable quantities of chlorine from it. f CO OK Potassium Oxalate : j ^X'oK + H2 °' Tlie normal salt is 'ob- tained by neutralizing oxalic acid, or the acid'salt, with potassium carbonate, and easily crystallizes from its aqueous solution. The f CO OT-T acid salt: -,,.-.' qtj- + H 2 (salts of sorrel), occurs in some plants, especially in the leaves of the sorrels, from which it may be obtained in large crystals by pressing the leaves and evaporating the acid juice. It has a strong acid reaction, and unites with ferric oxide to form a soluble double salt. Hence its use to remove ink or iron stains from paper or linen. After the stains have been removed, the fabric must be carefully washed, first with water and then with dilute soda, to prevent corrosion by the oxalic acid contained in the acid salt. Besides this acid salt, a so-called diacid salt (potassium quadroxalate) is also known. The salt crystallizes well, and is to be considered as a double compound of acid potassium oxalate and oxalic acid : — jCO-OK (CO-OH H0 (CO-OH + (CO-OH + - n * u " Potassium Silicate (potash water-glass). — Silica and caustic potash or potassium carbonate may be fused together in varying proportions, and produce, when the quantity of silica is not too great, a glass which is distinguished from ordinary glass by its solubility in water. Hence the name tuater-glass. It may be obtained by melting together equal weights of pearlash and sand until the mass fuses quietly, or by dissolving infusorial earth in caustic potash. By the latter operation a solution is at once obtained. In England considerable quantities of water glass are manufactured by heating powdered flints with caustic potash (or soda) under pressure, when the flints readily dissolve. On evaporating a solution of water-glass, an amorphous transparent mass remains behind, which, when exposed to the air, gradually becomes covered with an opaque coating of silicic acid and potassium carbonate. The solid fused water-glass undergoes the same decomposition. An aqueous solution of water-glass has many technical appli- cations. It is used, for example, to render theatrical scenes and a a 2 356 Text-Bool: of Inorganic Chemistry. other stage appliances uninflammable. Such articles when coated with a thin layer of water-glass are no longer combustible ; the thin layer of potassium silicate protects their surface like a varnish from the oxygen of the air. If brought into contact with a burn- ing body they become charred but do not catch fire. Water-glass is also largely used for protecting buildings from damp, for preventing the decay of woodwork in damp situations, and very largely in the manufacture of the cheaper kinds of soap. In a very dilute aqueous solution potassium silicate forms a valuable manure for many plants, especially for those which require potash salts and silica for their growth. Pasture lands are improved by occasionally watering with a dilute solution of potash water-glass. It has already been mentioned that potassium silicate is a constituent of many minerals — for example, of felspar. Potassium Chloride : KC1, crystallizes like sodium chloride in the regular system, and usually in cubes without water of crystallization. The crystals contain, however, water mixed mechanically, which causes them to decrepitate when heated. It has a saline taste, and is tolerably soluble in water : 100 parts of water dissolve about 33 parts of the salt at the ordinary tempera- ture, and hot water only a little more. It melts without change when heated to redness, and then volatilizes in not inconsiderable quantities. Potassium chloride is a constituent of sea-water, is contained in the ashes of land plants, and is also found in large quantities as a mineral. It is the chief constituent of the salt beds in the neighbourhood of Stassfurth in Northern Germany. The mineral sylvin consists almost exclusively of it. Potassium Bromide : KBr, closely resembling potassium chloride, may be prepared by neutralizing potassium carbonate with hydrobromic acid, or by saturating caustic potash with bro- mine, evaporating to dryness, and then glowing to destroy the potassium bromate produced at the same time. It is now largely used in medicine for neuralgic affections. Potassium iodide : KI. — This salt resembles potassium chlo- ride in its external properties. Like the chloride, it crystallizes in cubes without water. It is, however, much more easily soluble in Potassium Iodide. 357 water than potassium chloride, and thus remains in the mother- liquor when a mixed solution of the two salts is evaporated down. At the ordinary temperature loo parts of water dissolve more than 140 parts of potassium iodide ; the saturated solution contains, at its boiling-point (120 ), more than 200 parts of potassium iodide for every 100 parts of water. Potassium iodide is of such importance in medicine and photo- graphy that many attempts have been made to discover profitable and ready methods for preparing it. The simplest method — viz. neutralization of caustic potash or potassium carbonate with hydri- odic acid, just as potassium chloride maybe obtained — is impracti- cable, because hydriodic acid is very much more difficult to obtain in quantity than hydrochloric acid. Two methods are chiefly used for the preparation of potassium iodide. One consists in gradually, adding iodine to warm caustic potash until the liquid becomes yellow from free iodine, evaporating the liquid, which now contains potassium iodide and iodate (p. 130), to dryness, and glowing, in order to convert the potassium iodate into potassium iodide and oxygen. It has been found, however, that a much higher temperature is required to decompose potassium iodate than potassium chlorate, and at this high temperature some of the potassium iodide is volatilized and lost. It is therefore customary to add a little powdered wood charcoal to the liquid containing the mixed salts. When the saline mass is afterwards glowed, this charcoal abstracts all the oxygen from the iodate and is converted into carbonic acid. The residue is then extracted with water, and the clear solution evaporated down to crystallize. A second method largely used for the preparation of potassium iodide is as follows. One part of iron filings is mixed with water and digested with three parts of iodine, when the two substances readily unite with one another, with considerable rise in temperature, and produce the easily soluble salt — ferrous iodide : Fel 2 . As soon as all the iodine has disappeared, the green solution is filtered, and one more part of iodine added, which produces the compound ferroso-ferric iodide, Fel 2 , FeJ 6 . The liquid is then exactly precipitated with potassium carbonate, when ferroso-ferric hydrate easily separates and can be readily washed, and the clear solution, which now contains potassium iodide, is evaporated down to crystallize. From 4 parts of iodine, 5 parts of potassium iodide may be thus obtained. The ferrous iodide might be at once decomposed by potassium carbonate, but the precipitate then 358 Text-Book of Inorganic Chemistry. consists of ferrous carbonate, which does not readily settle, and can only be washed with difficulty, and therefore requires large quantities of water and a long time to extract all the potassium iodide. An aqueous solution of potassium iodide dissolves large quan- tities of free iodine; the saturated solution contains about twice as much iodine as was present in the original salt, producing the com- pound : KI 3 , which may be obtained as dark, lustrous needles, on careful evaporation. This compound is very deliquescent, and easily decomposes into potassium iodide and iodine. Potassium iodide, together with the sodium salt, is present in many mineral waters, and especially in the ashes of sea plants (pp. 123, et seq.). It is largely used in medicine as a remedy for goitre and swellings, in cases of syphilis, &c.,and is, in fact, one of the most highly prized medicines. Large quantities of the salt are also used in -photography. Potassium Fluoride : KF, is easily obtained by neutralizing hydrofluoric acid with potassium carbonate. It crystallizes at the ordinary temperature with two molecules of water, but above 35 without water, in cubes ; the salt easily deliquesces in moist air, and strongly attacks glass. Potassium fluoride unites with other fluorides and with hydro- fluoric acid, and produces double salts, in which the fluorine appears to play the' part of a dyad. Acid Potassium 'Fluoride : HKF„. — This salt is easily prepared by neutralizing a given volume of aqueous hydrofluoric acid with potassium carbonate, and then adding an equal volume of the acid. It crystallizes out when the solution is evaporated down in a platinum basin. The crystals are anhydrous ; they easily dissolve in water, but difficultly in that containing hydrofluoric acid. It requires heating to redness before it is decomposed into potassium fluoride and hydrofluoric acid. Potassium Fluoborate: KF,BF 3 = KBF 4 , is precipitated aS a gelatinous, iridescent mass, which dries to a white impalpable powder when potassium fluoride, or any other soluble potassium salt, is mixed with aqueous fluoboric acid. It is difficultly soluble in cold water, of which it requires 70 times its weight for solu- tion. Potassium Cyanide. 359 Potassium Fluosilicate : 2KF,SiF 4 = K 2 SiF 6 , closely resembles the fluorborate, but is still less soluble in water. Potassium fluo- silicate is therefore sometimes used as a test for potassium. Potassium Cyanide : KCN = KCy. — This extremely poisonous salt is decomposed by all acids, including carbonic acid, 'giving free hydrocyanic acid. It may be obtained pure by passing hydro- cyanic acid gas into well-cooled alcoholic potash, until the mass becomes pasty from separation of the potassium cyanide, which is almost insoluble in alcohol. The crystalline mass is then brought on a filter, quickly washed with strong alcohol, as much of the liquid removed as possible by a filter-pump, and finally dried in a vacuum over sulphuric acid. The less pure compact potassium cyanide of commerce, which occurs as flat pieces with a crystalline fracture, is prepared by a much less expensive process. Potassium ferrocyanide : K 4 FeCy + 3H a O, which may be considered as a double cyanide of ferrous iron and potassium : 4-KCy, FeCy 2 + 3H 2 0, is first thoroughly freed of its water of crystallization by powdering it and heating it in a shallow iron dish. The dried salt is then placed in an iron crucible, which it must not more than two-thirds fill, and heated to fusion in a furnace. The ferrous cyanide is then decomposed into nitrogen gas which is given off, and black iron carbide which remains behind. The potassium cyanide melts unchanged, pro- vided it is not exposed to the air, from which (when in the fused state) it readily absorbs oxygen, and becomes partly converted into potassium cyanate. As soon as the salt fuses quietly, the crucible is removed from the furnace, and repeatedly knocked on some hard body that the iron carbide may settle to the bottom and the potassium cyanide become clear. When the fused salt has sufficiently clarified, which may be easily known by removing a drop on the end of a glass rod, the liquid is carefully poured on to an iron slab, broken into pieces when cold, and preserved in well- closed vessels. In this method of preparation two of the six atoms of cyanogen contained in a molecule of potassium ferrocyanide are lost : the two atoms combined with the iron are decomposed. In order to obtain this cyanogen as well, it is best to use Liebig's • method, which consists in mixing every 8 parts of the dry potassium ferro- cyanide with 3 parts of potassium carbonate. When this mixture is melted, the ferrous cyanide of the potassium ferrocyanide and 360 Text-Book of Inorganic Chemistry. the potassium carbonate first form ferrous carbonate and potassium cyanide : — FeCy 2 + Co|q|5 = 2KCy + CO-0 2 Fe. But the ferrous carbonate is at once broken up at the high tem- perature into carbonic acid and ferrous oxide : — CO-0 2 Fe = C0 2 + FeO, the oxide being finally reduced to metallic iron by a portion of the potassium cyanide, which is converted into potassium cyanate. Thus, of the two atoms of cyanogen which were united with the iron, one is obtained as potassium cyanide and one as potassium cyanate. As this small quantity of potassium cyanate is without im- portance for most of the technical purposes for which potassium cyanide is used, the commercial salt is now always manufactured by Liebig's process. ' The quantity obtained is not only larger, but the product is also whiter. Potassium cyanide is very easily soluble in water, and deli- quesces in moist air. It cannot be crystallized from its aqueous solution, much less obtained pure, as it gradually decomposes with water into potassium formate and ammonia : — KCN + 2H,0 = {co-OK + NH 3" It is scarcely soluble in alcohol. The aqueous solution reacts alkaline, and plays the part of a base with many other insoluble cyanides. Just in the same way as potassium hydrate dissolves lead oxide, alumina, and silica, producing salts in which these substances take the part of an acid, so, too, potassium cyanide dissolves lead cyanide, silver cyanide, and numerous other similar insoluble compounds, forming salts which are soluble in water, and can usually be obtained in the crystalline state. Neither solid potassium cyanide nor its aqueous solution unites directly with oxygen. But in the fused state its attraction for oxygen is so great that most of the metallic oxides are decomposed by it. The substance produced by this oxidation is potassium cyanate. The salt also easily combines directly with sulphur, forming potassium sulphocyanate. Union of the two substances even takes place when flowers of sulphur are stirred up with a ,warm aqueous solution of potassium cyanide. Potassium Cyanate. 361 Potassium Cyanate : CONK. — This white crystalline salt is easily soluble both in water and alcohol, and may be purified by crystallizing from the latter liquid. It is obtained by adding an oxide of lead to fused potassium cyanide, when the oxide is reduced to metallic lead and the cyanide oxidized to cyanate. It is, however, best prepared directly from potassium ferrocyanide by oxidization with manganese peroxide. Well dried and finely powdered potassium ferrocyanide is thoroughly mixed with an equal weight of dry commercial manga- nese peroxide containing about 75 per cent, of the pure compound, and the mixture gently heated with continual stirring in a flat iron dish. After a little time the mass begins to glow ; the heating is continued as long as this lasts and until a fragment when dissolved in water gives no precipitate of Prussian blue on mixing the clear solution with ferric chloride. The mass is next repeatedly digested with strong alcohol (82 per cent.), and potassium cyanate then separates from the hot saturated solution on cooling as thin tablets. After filtering off the crystals, they must be rapidly dried, as potassium cyanate readily decomposes with water, even at the ordinary temperature, into potassium and ammonium carbonates : — 2CONK + 4H s O = CO(OK) 2 + CO(ONH 4 ) 2 . Its aqueous solution, when mixed with ammonium sulphate, forms potassium sulphate and ammonium cyanate, which latter changes immediately into the isomeric compound urea. Addition of a strong acid to a solution of potassium cyanate does not liberate free cyanic acid, but decomposes it into carbonic acid and ammonia (p. 317). Potassium Sulphocyanate: CSNK, may be obtained by adding sulphur to fused potassium cyanide, or, better, by heating a mix- ture of 46 parts of potassium ferrocyanide, 17 parts of potassium carbonate, and 32 parts of sulphur in a HeSsian crucible until the mass fuses quietly, and until a drop dissolved in water no longer gives a blue precipitate with ferric chloride, but only a red colour. The fused mass is broken in pieces when cold and digested with hot alcohol, which dissolves the potassium sulpho- cyanate and leaves ferrous sulphide behind. From the hot solution it crystallizes on cooling in long colourless crystals, resembling nitre, and possessing, like this salt, a cooling taste. 362 Text-Book of Inorganic Chemistry. Water dissolves large quantities of it, so that it deliquesces in moist air. A considerable fall in, temperature is produced when the salt is dissolved in water. Potassium sulphocyanate, or any soluble sulphocyanate, gives no precipitate with ferric salts, but a deep red colour. Sulphur Compounds of Potassium. The alkali metals unite with sulphur in several proportions, forming definite chemical compounds, which may be compared with the compound of potassium iodide with iodine. Just as potas- sium iodide unites with two atoms of iodine (p. 359), so also potas- sium sulphide unites with one, two, three, and four atoms of sulphur, producing the compounds : — Potassium disulphide K,S. 2 Potassium trisulphide K 2 S 3 Potassium tetrasulphide ..... K 2 S 4 Potassium pentasulphide . . . K^S^. The two atoms of potassium do not appear to combine with more than five atoms of sulphur. Potassium Monosulphide : K,S, is a strong sulpho-base, easily soluble in water, and deliquescing in the air. It has a strong ' alkaline reaction and taste. This compound may be obtained by reducing potassium sulphate in a stream of hydrogen at a red- heat : — S0 3 (OK) 2 + 4 H 2 = K 2 S + 4 H„0, or, better, by glowing a mixture of one molecule of the salt with four molecules of carbon : — SO„(OK). 2 + 4C = K. 2 S + 4CO. A mixture of 3 parts of powdered potassium sulphate and 1 part of wood charcoal is placed in a covered Hessian crucible, which it must not more than two-thirds fill, and heated until the mixture has ceased frothing and fuses quietly. When the crucible is cold it is broken, and the potassium sulphide obtained as a red crystalline mass. It soon loses its colour in the air, from which it abstracts not only water, but also oxygen and carbonic acid. An aqueous solution of potassium monosulphide is readily ob tained in the following way. A given volume of moderately con- centrated caustic potash is divided into two equal parts. One of Sulphur Compounds of Potassium. 363 these is then completely saturated with sulphuretted hydrogen gas, which forms the sulphydrate : KSH, and the other part is then added. One molecule of potassium sulphydrate is decomposed by one molecule of caustic potash into one molecule of potassium monosulphide and one of water : — KSH + KOH = K 2 S + H 2 0. This solution is decomposed when treated with dilute acids, and then evolves sulphuretted hydrogen without depositing sulphur. It absorbs oxygen from the air, and is converted into potassium thiosulphate : — ■ K 2 S + 3O = SO*{sK Potassium monosulphide easily dissolves the sulphides of arsenic, antimony, and carbon and other similar sulpho-acids, with which it unites chemically and forms soluble sulpho-salts. Potassium Sulphydrate : KSH or „ [ S. This compound, which corresponds to caustic potash, is obtained by leading sulphuretted hydrogen into caustic potash as long as the gas is absorbed. If the solution of caustic potash is sufficiently concentrated, the compound separates out on cooling in long colourless crystals, containing water of crystallization. In its chemical behaviour it closely re- sembles the monosulphide. The polysulphides of potassium are obtained by adding the requisite quantity of sulphur to the fused monosulphide, or to its concentrated aqueous solution. If more sulphur is added to fused potassium monosulphide than corresponds to 4 molecules of the former to 1 molecule of the latter, there is produced Potassium Pentasulphide : K 2 S 5 , and the excess of sulphur volatilizes. This compound has a yellow-brown colour ; it easily dissolves in water, and when acted on by hydrochloric acid yields sulphuretted hydrogen and large quantities of milk of sulphur (p. 145). When heated more strongly it loses sulphur and be- comes Potassium Trisulphide : K 2 S 3 , a compound closely resembling the pentasulphide. liver of Sulphur (Potassa sulphurata^) — This pharmaceu- tical preparation, which owes its name to its colour, consists essen- 364 Text-Book of Inorganic Chemistry. tially of a mixture of potassium pentasulphide and potassium sulphate. It is prepared by gradually heating a mixture of about equal parts of potassium carbonate and sulphur in a large earthen- ware crucible. At first considerable quantities of carbonic acid are evolved, which causes the mixture to froth up ; as soon as this has ceased, the crucible is closed and the temperature gradually raised until all fuses quietly. The liquid is then poured on to a sheet of iron, broken into pieces when cold, and quickly brought into well-closed bottles, as it readily attracts moisture from the air. The following equation represents the reaction : — 4CO(OK) 3 + 16S = 3k' 2 S 5 + SO„(OK)„ + 4CO,. The preparation, when well made, should dissolve in water to a nearly clear solution, and when mixed with hydrochloric acid should give sulphuretted hydrogen and large quantities of free sulphur. A solution containing potassium pentasulphide and thiosul- phate may be obtained by gradually adding flowers of sulphur to boiling caustic potash as long as the sulphur is dissolved : — 6KOH + 12S = 2K,S 5 + SO,jg^ + 3H..O. Detection of Potassium Compounds. Potassium compounds when heated in the Bunsen burner impart a characteristic violet colour to the flame. In the presence of sodium, even in small quantities, the intense yellow light emitted by this substance obscures the feeble potassium flame, and it is therefore not seen by the naked eye. But if the flame is viewed through a piece of blue cobalt glass, the yellow sodium light is absorbed by the blue of the glass, while the violet potassium light passes through the glass unchanged. The most certain way of detecting traces of potassium is by its spectrum, which consists essentially of two lines in the red and one in the violet. {See Table of Spectra). If a solution of potassium chloride or of any other soluble potassium compound (except the iodide) is mixed with platinic chloride, a yellow crystalline precipitate of a double chloride of potassium and platinum : 2KCI, PtCl 4 , or potassium clitorplatinate : K 2 PtCl c , is produced, which is slightly soluble in water, but in- soluble in alcohol. The heavy metals, as well as the metals of .the earths and alkaline earths, if present, must have been pre- Detection of Potassium Compounds. 365 viously removed from the solution, and ammonium compounds destroyed by glowing. Sodium salts do not give this precipitate, as the corresponding sodium salt is soluble both in water and alcohol. This double chloride of potassium and platinum is decomposed on glowing into metallic platinum, gaseous chlorine, and potassium chloride. The last-named compound can be ex- tracted with water, and easily obtained in cubic crystals. If a concentrated solution of a potassium salt is mixed with a concentrated solution of tartaric acid in excess, and the mixture well shaken or stirred with a glass rod, a white crystalline pre- cipitate of acid potassium tartrate is produced, which is first deposited on those parts of the glass vessel which have been touched by the rod. Sodium salts do not produce this precipitate, as the corresponding acid sodium tartrate is easily soluble in water. SODIUM. Chemical Symbol : Na. — Atomic Weight : 23. Sodium, which is closely allied to potassium, is also found in nature in the form of a double silicate of aluminium as soda felspar or albite. It further occurs as the mineral cryolite : 6NaF, A1 2 F„, which is worked both for soda and for alumina. But by far its most important natural compound is the chloride — common salt or rock-salt. The preparation of sodium is carried on in just the same way as that of potassium, except that some powdered chalk is usually added to the mixed sodium carbonate and charcoal, to prevent fusion at the high temperature necessary for reduction. Since the affinity of sodium for oxygen is less than that of potassium, its reduction, according to the method described on p. 344, is easier and the yield is greater. There is also less danger of the formation of explosive compounds, as in the preparation of potassium. This is one reason why sodium is so much cheaper than potassium. It is chiefly owing to the researches of St. Claire Deville, sup- ported by Napoleon III., who, in attempting to prepare cheap aluminium by means of sodium, so perfected the method for the extraction of this metal that its price has fallen from <^s. per oz. to 8j. per lb. during the last thirty-five years. The present low price 356 Text-Book of Inorganic Chemistry. of sodium is also partly due to the large quantities which are now used for the manufacture of metallic magnesium. Commercial sodium comes into trade in the form of bars. When freshly cut the surface is of a silver-white colour, but it very rapidly tarnishes in the air, becoming covered with a layer of its oxide. At the ordinary temperature it is soft like wax, and can be easily moulded or pressed into any shape. The metal is lighter than water, but somewhat heavier than potassium — its specific gravity being 0-978. Its melting-point, which is also higher than that of potassium, is 96°. It catches fire when heated in the air, and burns with an intensely yellow flame. Water is easily decom- posed by it with evolution of hydrogen, and the sodium hydrate which is produced at the same time dissolves in the water and gives it an alkaline reaction. The liberated hydrogen does not, however, ignite like that set free by the action of potassium on water, because the quantity of heat set free when sodium decomposes water is less than when the same decomposition is effected by potassium — the affinity of the former metal for oxygen being less than that of the latter. The liberated heat would, however, raise the temperature of the hydrogen to its ignition point, if it were not continually cooled by the movement of the fragment of sodium on the surface of the cold water. If hot water, at about 80°, instead of cold water, is used for the experiment, the hydrogen catches fire at once and burns with a yellow flame. Or, if the surface of the water is covered with a piece of filter-paper the movement of the pellet of sodium is prevented, and it remains in contact with only a small quantity of water ; this then becomes so strongly heated that the hydrogen catches fire. Potassium and sodium unite together and form an alloy which is liquid at the temperature of the air, and closely resembles mercury. This is a striking example of the law that the melting- point of an alloy is lower than that of its constituents. Sodium, like potassium, unites chemically with hydrogen when heated up to about 300 in the gas. This compound is again de- composed at about 420°. If dry ammonia is led over gently heated sodium the gas is decomposed with the formation of sodium amide : Na-NH 2 , while hydrogen is set free. When fused, this compound has a greenish-blue colour, but solidifies to a pink crystalline mass. Compounds of Sodium. 367 COMPOUNDS OF SODIUM. With few exceptions, the compounds of sodium closely resemble those of potassium, and we shall, therefore, only describe those in detail which differ essentially from the corresponding potassium compounds. Sodium, when heated in dry oxygen or air, produces sodium peroxide : Na,O s , similar in its properties to the potassium com- pound. Sodium oxide : Na 2 0, is obtained, like potassium oxide, by heating sodium hydrate with sodium. Sodium Hydrate {Caustic Soda) : NaOH. — This compound is so similar to potassium hydrate that the two cannot be externally distinguished from one another. The general behaviour and the preparation of the two substances is also the same, except that caustic soda is a less powerful base than caustic potash. In pre- paring caustic soda from crystalline sodium carbonate, according to the method described on p. 346, it must be remembered that the sodium salt contains about 60 per cent, of water, and not more than five parts of water should, therefore, be used to dissolve one part of the salt. Sodium Sulphate : S0 2 j q vr„ + loH 2 0. — This salt, com- monly known as Glauber's salt, is occasionally found in the solid state in nature, but usually in solution in various mineral springs. Many of these springs, especially those of Marienbad, Carlsbad, and others, are highly valued on account of the mild aperient action of the sodium sulphate which they contain. Sodium sul- phate may be easily obtained by neutralizing sodium carbonate with sulphuric acid, or by heating this acid with common salt. This latter process is largely carried on in alkali (soda) works, where common salt is first converted into sodium sulphate prior to its change into sodium carbonate. Sodium sulph ate-crystallizes with ten molecules (i.e. 56 per cent, of water) in large colourless prisms, which, when heated, melt in their water of crystallization, and finally lose all their water. It is easily soluble in water but insoluble in alcohol. At 0°, 100 parts of water dissolve 12 parts of the salt, at 18° 48 parts, at 25° 100 368 Text-Book of Inorganic Chemistry. parts, and at 33° as much as 322 parts. But above 33° the solu- bility diminishes, because the compound with ten molecules of water cannot exist above this temperature. It is then converted into the anhydrous salt, which is less soluble in water, and separates from the solution above 33 in rhombohedra. If a glass flask is filled with a warm saturated solution of sodium sulphate and the neck closed with a loosely-fitting plug of cotton-wool, the solution remains clear and liquid after it has been allowed to cool quietly. And no change is produced if the plug is carefully removed and the end of a glass rod which has been heated in a gas flame and allowed to cool is dipped into the liquid. But the slightest touch with the other end of the rod, by which particles of dust are brought into contact with the liquid, at once causes the crystalliza- tion of the supersaturated solution, with a corresponding rise in temperature due to the liberation of latent heat. (QTT Oft is prepared in the same way as the potassium salt, and is deposited from its warm solution in large transparent crystals, without water. At the ordinary tempe- rature it crystallizes with 1 molecule of water. It decomposes when heated into water and sodium disulphate (p. 167), and the latter compound is converted into normal sodium sulphate and sulphuric anhydride at a higher temperature. Sodium Sulphite: SO j nja a is obtained from the acid salt by adding an equal quantity of sodium carbonate as was used in its preparation ; it crystallizes without water from its warm solution, and with 7 molecules in the cold. (OH ONa P re P are d by saturating a solution of sodium carbonate with sulphurous anhydride. Both salts are easily soluble in water, and both are largely used by brewers for antiseptic purposes — partly to cleanse stale casks, and partly as an addition to the beer to prevent acid fermentation. Sodium Thiosulphate (Sodium Hyposulphite) : SO, ,- N + 5H 2 0, is obtained by evaporating its aqueous solution in large transparent crystals, which melt in their water of crystallization at Sodium Nitrate. 369 56°, and lose all their water at 100° At higher temperature's it is decomposed. Its preparation and its property of dissolving the halogen salts of silver have been already referred to (pp. 168, 169). Large quantitie are used for photographic purposes. Owing to its decomposition by chlorine, it is also employed to remove the last traces of chlorine from articles which have been bleached with this substance, and is hence often called an antichlor. A cheap mode of preparing it consists in decomposing a solution of oxidized alkali-waste (p. 169), which contains calcium thiosulphate, with sodium carbonate ; insoluble calcium carbonate and a solu- tion of sodium thiosulphate are thus obtained. Sodium Nitrate {Chili Saltpetre): NO a -ONa, crystallizes in rhombohedra which so closely resemble cubes that it has been called cubic nitre. It is more easily soluble in water than potassium nitrate, and only requires about its own weight of water for solution at the ordinary temperature. The salt is found in large deposits in nature, especially in South America on the borders of Chili and Peru, and is an important article of commerce. This crude nitrate of soda is, however, far from pure ; it contains sodium chloride, iodide and sulphate, together with earthy impurities, and may be purified by crystallization. The mother-liquors are used for the preparation of iodine (p. T24). Pure sodium nitrate maybe pre- pared by neutralizing caustic soda or sodium carbonate with nitric acid. It has already been mentioned that sodium nitrate is slightly hygroscopic, so that gunpowder prepared from it becomes moist and useless. Nevertheless, sodium nitrate can be readily converted into potassium nitrate, and so used indirectly for the preparation of gunpowder. It is also used for the manufacture of nitric acid. Jon a Sodium Carbonate, or Soda: COj,-,,^., + loH 2 0. — Just as land-plants require potassium compounds for their growth, so are sodium compounds necessary for sea-plants. The former derive their potash from the soil, the latter obtain their necessary soda from the sea. And when sea-plants are burnt their ashes contain sodium carbonate, in the same way as the ashes of land-plants contain the corresponding potassium carbonate. Until some eighty or ninety years ago, most of the soda used in the arts was obtained in this way from sea-plants, and, at that time, h U 370 Text-Book of Inorganic Chemistry. large quantities of potash were imported into France and Western Europe from Russia, Germany, and America. But at the end of the preceding century, during the French revolution, this importa- tion of potash into France almost ceased, and many important industries depending upon a supply of alkali seriously diminished. A commission was then appointed by the French Government to examine into any processes for converting common salt into soda, and they decided in favour of the method discovered by Leblanc — a process which has remained practically the same from then to the present day. Since then this branch of industry has attained immense dimensions, and the price of soda has fallen so low that in many cases it has displaced the more expensive potash, except when this alkali alone can be used — e.g. in the manufacture of potash glass; &C. 1 The quantity of soda which occurs in the mineral kingdom, partly in the crystalline state and partly in solution, in the soda- lakes of Egypt and North America, is only of local importance, and is inappreciably small in comparison with the immense quan- tities of this salt which, are used for so many purposes. The conversion of common salt into soda, when casually con- sidered, does not appear a difficult problem. It might be thought that on heating calcium carbonate with sodium chloride the two salts would yield sodium carbonate and calcium chloride. But even supposing this reaction to take place when the dry sub- stances are heated, it would not be possible to separate the sodium carbonate from the calcium chloride, because, on the addition of water, both would be dissolved, and would again produce sodium chloride and insoluble calcium carbonate. Such a process could only be successful if the calcium chloride were converted into a compound insoluble in water. And, as a matter of fact, the process introduced by Leblanc depends upon the conversion of a sodium salt and calcium carbonate into sodium carbonate and an insoluble calcium compound, so that when the mixture is digested with water, sodium carbonate alone dissolves. Leblanc used sodium sulphate for this purpose, and this salt must first be obtained from sodium chloride. The process of manufacturing soda, or the alkali manufacture, as it is usually called, is not a simple one ; various conditions are necessary for carrying it on successfully, and the process consists of several distinct chemical operations. The manufacture requires, 1 See foot-note on p. 349. Sodium Carbonate. 371 in the first place, a large supply of common salt and limestone, and, secondly, of sulphuric acid. And as the immense quantities of sulphuric acid which are necessary for the production of the sodium sulphate cannot well be obtained from other manufactories, the acid is made on the spot, and sulphuric acid chambers with other requisites for manufacturing the acid are one necessity of every alkali works. This sulphuric acid is then used in the first or salt-cake process, to convert the sodium chloride into sodium sulphate or salt-cake ; large quantities of hydrochloric acid being then obtained as a bye-product (p. 108). After all the preliminary conditions have been fulfilled which are necessary to obtain the sodium sulphate or salt-cake on a sufficiently large scale, the actual operation of soda manufacture, or the conversion of sodium sulphate into sodium carbonate, can be commenced. But it is not the object of this small text-book to describe all the details of the manufacture of alkali, nor indeed can technical chemistry, as little as practical chemistry, be learnt from books. A number of practical details, often apparently inex- plicable, must be attended to if the manufacture is to be a success, and these details cannot be learnt from books, but a knowledge of them must be acquired in the manufactory itself. And this is true not only for the manufacture of soda but also for every other branch of chemical industry. There is always a gap between theoretical and practical chemistry, which remains unbridged for the man who works by rule-of-thumb, but which is readily accounted for by the scientific chemist who has learnt to observe accurately and to solve chemical problems by experiment. It is therefore sufficient if the student understands the chemical reactions upon which the change of sodium sulphate into sodium carbonate depends. In the second process of alkali manufacture — the black-ash process — the sodium sulphate (saltvcake) is mixed with calcium carbonate (limestone) and small coal, the mixture heated in rotating furnaces, and the black mass [black-ash) which remains behind lixiviated with water to dissolve the sodium carbonate. The salt then crystallizes out on evaporating this solution. The chemical reactions which produce this change are as follows. Sodium sulphate, when heated with the carbon of the coal, is converted into sodium sulphide and carbonic oxide : — S0 2 (ONa) 2 + 4C = Na 2 S + 4CO, B B 2 372 Text-Book of Inorganic Chemistry. and this sodium sulphide is at once decomposed by the calcium carbonate, producing sodium carbonate and calcium sulphide : — Na 3 S + CO-0 2 Ca = CO(ONa) 2 + CaS. On lixiviating this mass with water, the soluble sodium carbonate dissolves, and the insoluble calcium sulphide remains behind. This insoluble residue of impure calcium sulphide constitutes alkali-waste, and when exposed to the air, gradually forms poly- sulphides of calcium and other salts which dissolve in the drainage water, and which liberate sulphuretted hydrogen when coming into contact with acid liquors. Hence, not only is the loss of material in the waste very considerable (it contains the whole of the sulphur present in the sulphuric acid employed), but it is also an intolerable nuisance to the neighbourhood of alkali works. Many proposals have been made to recover the sulphur from this waste, some of which have proved partially successful. In practice, a much larger quantity of limestone is employed than that represented by the above equation. This excess of limestone becomes converted into quick-lime during the heating, and then produces a corresponding quantity of caustic soda, which remains in the mother-liquor after the sodium carbonate has crystallized out. In some alkali works considerable quantities of caustic soda are manufactured in this way. Soda is also manufactured to some extent from cryolite, a double fluoride of sodium and aluminium : 3NaF, A1F 3 (p. 134). This compound, when mixed with powdered limestone and strongly heated, decomposes into soluble sodium aluminate and insoluble calcium fluoride : — Na 3 AlF 6 + sCO-0 2 Ca = Na 3 A10 3 + 3CaF 3 + 3CO3. And when a stream of carbonic acid is led through the solution of sodium aluminate produced on digesting the mass with water, sodium carbonate is formed and aluminium hydrate precipitated. The direct conversion of common salt into soda has been very successfully worked in recent years, and the quantity of soda annually produced by this ammonia-soda process is rapidly in- creasing. It has been found that a mixed solution of common salt and acid ammonium carbonate is converted under a pressure of about two atmospheres into ammonium chloride and acid sodium carbonate : — NaCl + Co(°^ = C0J°g a + NH 4 C1. Sodium Carbonate. 373 The process which is based upon this reaction is carried out in the following manner. A saturated solution of common salt is mixed with ammonia, and then saturated with carbonic acid under a pressure of two atmospheres. The acid sodium carbonate, which is only difficultly soluble in water, then separates out, is filtered off under pressure, and converted into the normal carbo- nate by heating. The carbonic acid which is then given off is again used to reconvert the ammonia which is recovered from the ammonium chloride into ammonium carbonate. The ammonia-soda process is not only simpler than the older method, but it only yields one bye-product — viz. calcium or magnesium chloride, arising from the decomposition of the ammo- nium chloride by lime or magnesia. Leblanc's process, on the other hand, gives a number of objectionable residues, from which only a portion of the sulphur can be again obtained. Still the ammonia process can never entirely displace the black-ash process, because we cannot dispense with hydrochloric acid which is pro- duced in such large quantities as a bye-product. Although the compounds of sodium and potassium as a general rule are very similar to one another, the two carbonates are strikingly different, especially in their , behaviour with water. Potassium carbonate is very soluble in water, and can only be obtained in the crystalline form with difficulty ; it even attracts moisture from the air and deliquesces. Sodium carbonate, on the other hand, is much less soluble in water, and can be easily ob- tained in large crystals, which when exposed to the air rapidly lose water and effloresce. A dry fresh crystal placed on the scale-pan of a balance loses weight rapidly and continuously. The ordinary crystals of sodium carbonate are large trans- parent monoclinic tablets containing 10 molecules or 63 per cent, of water. When these crystals are heated, they melt at 50° in their water of crystallization, to form a clear liquid. At a higher temperature water is expelled, and finally the anhydrous salt remains behind. Sodium carbonate when crystallized from solu- tions above 50° contains only 7 molecules of water. The degree of solubility of sodium carbonate in water at different temperatures varies considerably. The solubility increases with the temperature up to about 35°, but then begins to diminish, with the separation of a salt containing less water. The crystals when exposed to the air become covered with a 374 Text-Book of Inorganic Chemistry. thin, loosely attached layer of a compound containing less water, into which the salt is finally completely converted. If a crystal of soda which has been long exposed to the air is broken, a small nucleus of the transparent undecomposed salt may often be found in the centre. When strongly heated in a platinum crucible, anhydrous sodium carbonate melts to a clear liquid, without undergoing decomposition ; carbonic acid is not expelled. An aqueous solu- tion of sodium carbonate, like one of potassium carbonate, reacts strongly alkaline. f ONa Acid Sodium Carbonate : CO \ .-.tt . — This salt, which is com- monly known ' as bicarbonate of soda, is obtained by leading a stream of carbonic acid through a strong solution of the normal carbonate. The acid carbonate, owing to its smaller solubility, then separates out as a crystalline powder. It is usually prepared by saturating a mixture of I part of the crystalline normal carbonate and 3 parts of the anhydrous salt with carbonic acid. The salt crystallizes without water in small distorted monoclinic plates, and requires about 1 1 parts of water for solution at the ordinary tempe- rature. Its solution reacts neutral, but loses carbonic acid when boiled, and is first converted into the so-called sesquicarbonate of soda, which is a hydrated double compound of the normal and acid carbonates, of the composition : — C0 J0Na + 2C0 J0H + 2H *°- This double compound is found in nature as small hard crystals, which are permanent in the air. It separates on the banks of the soda-lakes of Egypt and America, and is brought into trade under the name of Trona. Acid sodium carbonate is used to some extent in medicine, especially in the preparation of Seidlitz powders. It is found in many mineral springs, and particularly in those which are alkaline. Phosphates of Sodium. — Of the various compounds of phos- phoric acid and sodium, the best known and most important is Monacid Sodium Phosphate : PO j qi^ 2 + I2H 2 0, which is obtainedby nearly neutralizing caustic soda or sodium carbonate with ordinary phosphoric acid. On evaporating this solution down, the Phosphates of Sodium. 375 salt separates out in large transparent crystals which easily effloresce in the air. Its solution possesses a faint alkaline reaction although it is an acid salt. The hydrated salt readily loses its 60 per cent, of water of crystallization when heated to 100° ; and at a red heat it is decomposed into water and sodium pyrophosphate (p. 219). Normal Sodium Phosphate: PO(ONa) 3 + i2H.,0, crystallizes in thin six-sided prisms when a solution of the preceding salt is mixed with sufficient caustic soda and evaporated down. The solution has a strong alkaline reaction ; it absorbs carbonic acid from the air, and becomes converted into sodium carbonate and the monacid phosphate. Diacid Sodium Phosphate : ^^ > \/A\3\ + H 2 0, is obtained when a solution of the monacid phosphate is mixed with as much phosphoric acid as it already contains. The solution has an acid reaction ; when evaporated down it yields the salt in rhombic prisms, which are easily soluble in water. Very different from these three salts is the sodium compound of monobasic metaphosphoric acid : Sodium Metaphosphate : P0 2 -ONa, which is obtained by heat- ing diacid sodium phosphate :— PO |(OHJ = P0 2 -ONa + H 2 (p. 219), as an amorphous vitreous mass, and which cannot be ob- tained in the crystalline form. Sodium Pyrophosphate : Q\t,c\(CY!Z '■? + ioH 2 0, is prepared by heating the monacid sodium phosphate to redness (p. 219), and sepa- rates from an aqueous solution of the residue in large crystals, which are permanent in the air. Its solution may be boiled without change, but if nitric acid is added and the solution then boiled, the salt is decomposed into sodium nitrate and diacid sodium phosphate. Sodium Arsenate: AsOjkri "+I2H 2 0, which is isomor- phous with the corresponding common sodium phosphate, is de- posited from its aqueous solution in large well-developed crystals. It is largely used in calico-printing and dyeing. The normal and diacid arsenates closely resemble the corresponding phos- phates. 376 Text- Bo ok of Inorganic Chemistry. Sodium Borate : B 4 5 (ONa) 2 + ioH 2 (p. 264). — This salt, which is commonly known as borax, and is found in nature as the mineral tinkal, is prepared by neutralizing a hot solution of sodium carbonate with boric acid. When the solution is allowed to evaporate slowly the salt separates out as large transparent mono- clinic prisms of the above composition, which only slightly change in the air. At the ordinary temperature 100 parts of water only dissolve about 7 parts of the salt, but the same quantity of water dissolves more than 200 parts of the salt at ioo°. The solution possesses a faint alkaline reaction. From a concentrated aqueous solution of borax at 70° or 8o°, the salt crystallizes in octahedra with only 5 molecules of water. This octahedral borax is distin- guished from ordinary prismatic borax by its greater hardness. Borax loses its water when heated and swells up into a porous mass ; at higher temperatures the salt melts to a clear liquid which solidifies on cooling to a hard transparent glass. Fused borax possesses the property of dissolving metallic oxides, often in con- ' siderable quantities, and the variously coloured glasses which are thus produced afford a test for various metals. Thus cobalt oxide imparts a blue colour to fused borax ; manganic oxide gives a violet, but manganous oxide a colourless glass ; while cupric oxide and chromic oxide yield green-coloured glasses. Borax is employed in the arts as well as in the laboratory. Its property of dissolving metallic oxides makes it useful in soldering two pieces of metal together when the surfaces to be united must be perfectly free from oxide. Considerable quantities of borax are used for glazing porcelain and in the manufacture of enamels. It is also a valuable medicine. The compounds of sodium with the halogens and with sulphur closely resemble the corresponding compounds of potassium, and of these the chloride is the only one which need be referred to in detail. Sodium chloride {Common Salt) : NaCl. — This salt not only occurs widely distributed in nature, but is also found in immense quantities, partly as solid rock-salt, and partly in aqueous solution in brine springs and in the sea. Common salt crystallizes in the regular system, and generally in cubes. It contains no water of crystallization, but usually includes small quantities of water mechanically, which cause the Sodium Chloride. 377 crystals to decrepitate when heated. It possesses a well-known saline taste, and is only slightly more soluble in hot than in cold water. 100 parts of water dissolve 36 parts of salt at the ordinary temperature, and only 39 parts at ioo°. The concentrated aqueous solution, therefore, scarcely contains 27 per cent, of the salt. It is insoluble in strong alcohol. When an aqueous solution of common salt is evaporated down, the crystallization commences first where the concentration is greatest — i.e. on the surface — and the small cubes which are thus produced gradually increase in weight, and sink to the bottom of the vessel. The extraction of sal: from the brine springs of Cheshire and Worcestershire is an extremely simple process, and consists in eva- porating the brine in shallow iron pans. In countries where fuel is dear and the brine less concentrated — e.g. in Germany, the solution is allowed to trickle several times over tall walls of faggots of thorn, which exposes a large surface to the air, and causes a rapid evapo- ration of water. After a concentrated solution has been obtained by this process (called graduation), the strong brine is boiled down in iron pans. In the same way, salt has also been obtained from sea-water, although this liquid only contains about 3-5 per cent, of it. In warm southern climates, for example, in the south of France and on the .coast of Africa, shallow basins are filled by the sea at high water, then closed, and the water evaporated by the heat of the sun and by the warm winds. Salt obtained from sea- water in its crude state, and containing various impurities, is commonly known as bay-salt. Brine springs are produced by water coming into contact with deposits of rock-salt in the earth, and artificial springs are very often formed by drilling down to the layer of salt and pumping in water. In some cases the rock-salt is mined in the ordinary way, and brought to the surface in the solid state. The chief deposits of salt occur in Galicia, in various parts of Germany, and at North- wich and Droitwich in England. Very important salt deposits have been recently discovered at Stassfurt, near Magdeburg, where the common salt is associated with potassium chloride and other com- pounds. Rock-salt can be used in its crude state for most technical pur- poses, but not as a condiment, as it possesses an impure taste due to various impurities, which may, however, be separated from it by recrystallizing it from its solution in water. 378 Text-Book of Inorganic Chemistry. Commercial common salt often differs both in its external appearance and in its taste. In some cases the cubes in which it crystallizes are arranged together in large pyramids, in other cases they are separate and form a granular crystalline powder. One sample may become moist in the air, and possess an intense saline taste, while another remains dry, and tastes much less salt. The form in which the salt crystallizes depends partly on the temperature of evaporation, and partly upon the presence or absence of foreign substances, such as sodium sulphate, magnesium chloride, or calcium chloride. The two latter salts are those which cause the salt to deliquesce in the air, and which impart the sharp saline taste that some samples possess. Pure sodium chloride does not become moist when exposed to the air, and has a pure saline taste. Common salt is almost as essential for the life and health of men and animals as the air which they breathe, and if a State imposes a tax upon this most necessary food (as in Germany), while luxuries such as tobacco remain almost untaxed, it proves that its rulers lack scientific training, and that they do not under- stand the simplest and most important questions bearing upon the material well-being of the people. Detection of Sodium Compounds. The intensely yellow colour which sodium compounds impart to a Bunsen flame, together with the single yellow line of the sodium spectrum (see table of spectra) are so characteristic that it is scarcely possible to overlook the presence of even small quantities of the metal. Sodium salts may be distinguished from those of potassium by the solubility of sodium chlorplatinate : Na 2 PtCl e , both in water and in alcohol. The two carbonates are also very different. Potassium carbonate deliquesces in the air, but the sodium salt effloresces. LITHIUM. Chemical Symbol : Li. — Atomic Weight: 7. Lithium is a rare element, but on account of the extensive use of some of its compounds in medicine is of considerable value. It Lithium. 379 is also widely distributed in nature, and is found, though always in small quantities, in the minerals : lepidolite, petallite, triphylline, &c, as well as in numerous mineral springs, especially in one near Redruth in Cornwall. The solubility of lithium carbonate in water affords a basis for the separation of this salt from the carbonates of the alkaline earths, while, on the other hand, its nearly insoluble phosphate permits its separation from the easily soluble phosphates of the alkalies. The preparation and purification of lithium compounds depend on the properties of these two salts. Metallic lithium, which may be prepared by electrolysis of its fused chloride, is a silver-white tough metal, harder than potassium or sodium, but softer, than lead. It is distinguished by its low specific gravity (0-59), and is, in fact, the lightest solid known. Its melting-point is 180°, but, unlike potassium and sodium, it is not volatile at a red-heat, and can only be distilled at very high tem- peratures. In the air it oxidizes less easily than potassium and sodium, but its freshly-cut surface soon becomes covered with a coating of oxide. When heated in the air above its melting-point, it catches fire and burns brightly. Water is easily decomposed by it at the ordinary temperature, but the hydrogen is not inflamed. It burns readily in chlorine, producing its chloride. Of the compounds of lithium : lithium oxide and lithium per- oxide have been but little investigated. Lithium Hydrate (Lithia): LiOH, which may be prepared by boiling the carbonate with slaked-lime, remains behind on evapor- ating the alkaline solution in a silver dish as a white, transparent, easily fusible mass. It resembles caustic soda in external appear- ance, but is less soluble in water and does not deliquesce in the air. Lithium Sulphate: S0 2 | q L | + H 2 0, crystallizes in rhombic prisms and is easily soluble in water. — Lithium Nitrate : N0 2 -OLi, crystallizing without water in rhombohedra, is very easily soluble in water, and deliquesces in moist air. Lithium Phosphate: 2PO(OLi) 3 + H 2 0, is very difficultly solu- ble in water, and is precipitated on mixing a solution of mon- -acid sodium phosphate with one of a soluble lithium salt as a white, heavy crystalline powder, the quantity of which is increased 380 Text-Book of Inorganic Chemistry^ on neutralizing the liquid with a, little caustic soda, or ammonia. The salt requires about 2,500 parts of water for its solution at the ordinary temperature, and is still less soluble in water containing ammonia. Lithium Carbonate : CO ^r \. — This compound, which is also difficultly soluble in water in comparison with the carbonates of potassium and sodium, is deposited as a white powder on adding sodium carbonate • to a concentrated solution of lithium chloride. It dissolves in about 100 parts of water at the ordinary temperature, yielding a solution which is faintly alkaline. Hot water dissolves more, but it is insoluble in alcohol. Lithium carbonate possesses the remarkable property of forming a soluble compound with the insoluble uric acid, and is therefore largely taken as a medicine to remove this substance from the body, when it separates out during the progress of certain diseases (gouty affections, stone, &c). The lithium carbonate is usually given in water containing free carbonic acid (lithia-water), in which the salt dissolves more readily than in pure water. Lithium Chloride : LiCl, crystallizes from very concentrated solutions in cubes or octahedra. It is very soluble in water and deliquesces in the air. It is also soluble in alcohol, and even in a mixture of alcohol and ether. When fused in an open vessel it loses some chlorine and becomes partially converted into a basic chloride. Detection of lithium Compounds. The smallest quantity of a lithium compound when introduced into a non-luminous gas flame gives a splendid crimson colour, the spectrum of which is one single line in the red (see table). The precipitates produced by sodium carbonate or phosphate in strong solutions of lithium compounds also serve to distinguish them from those of potassium and sodium. RUBIDIUM AND CESIUM. Chemical SymMs: {^.-Atomic Weights : (^ These two rare elements are as yet only of chemical interest. Both were discovered by Bunsen and Kirchhoff during their Rubidium and Ccesium. 381 researches in spectrum analysis. The two metals are found in minute quantities in certain minerals, in the ashes of some plants, and in some mineral springs. One of the richest sources of the two chlorides is the mineral water of Diirkheim in the Bavarian Palatinate, in which caesium was first discovered by Bunsen, but even this only contains one part of caesium chloride in 6,000,000 parts of the water. Ricbidium, which is obtained, like potassium and sodium, by glowing an intimate mixture of its carbonate and charcoal, is a white metal with a faint yellowish tint. It is as soft as wax even at 10°, melts at 38 , and has a specific gravity of 1-5. It at once oxidizes when exposed to the air, and then becomes so strongly heated that it soon catches fire. When brought into contact with water, decomposition at once takes place and the liberated hydro- gen is ignited. The compounds of rubidium so closelv resemble those of potas- sium that it is very difficult to separate the two from one another. The double chlorides of the metals with platinum are best adapted for this separation, the rubidium compound (2RbCl, PtCLJ being much less soluble in water than the corresponding potassium com- pound. Rubidium is more electro-positive than potassium. Com- pounds of rubidium when held on a platinum wire in a non- luminous gas flame impart a violet colour to it resembling that of potassium compounds. The spectrum of this flame is, however, quite different, and is characterized by two lines in the red and two in the violet. Ccesium, which is even rarer than rubidium, has been recently obtained in the metallic state by the electrolysis of a mixture of caesium and barium cyanides. It is a white, very soft metal, melting at 27 or a summer temperature, and of r88 specific gravity. It decomposes water with explosive violence, and is even more electro-positive than rubidium. Caesium is, therefore, the most electro-positive of all metals, and has a stronger affinity for oxygen, chlorine, &c, than any other metal. Its compounds closely resemble those of rubidium and potassium. In order to separate it from these two metals the solubility of its carbonate in absolute alcohol, in which the carbonates of rubidium and potassium do not dissolve, may be made use of. The spectrum of caesium compounds is especially characterized by two lines in the bright blue. 382 Text-Book of Inorganic Chemistry. AMMONIUM. It has already been remarked (p. 188) that we are not acquainted with the substance called ammonium in the free state ; but the fact that it forms an amalgam with mercury and the close relation of its compounds to those of potassium, justify us in deciding in favour of its metallic nature, and in placing it together with the metals of the alkalies. Its composition is expressed by the formula : NH 4 , or perhaps more correctly as N 2 H 8 = (NH 4 ) 3 . Ammonium oxide and hydrate are just, as little known as ammonium itself. It has sometimes been stated that the aqueous solution of ammonia in water contains ammonium hydrate : NHj + H„0 = NH,OH, but this is as little true as the assertion that a solution of carbonic anhydride in water contains carbonic acid : CO j Y.Vt. The mere fact that it is only necessary to blow a current of air or some other indifferent gas through these solutions to expel the whole of the gas, sufficiently proves that they are not true chemical compounds. The ammonium salts are easily obtained by neutralizing ammonia or ammonium carbonate with the corresponding acid. Like those of potassium, they are nearly all soluble in water, and mostly crystallize well. Ammonium Sulphate : SO„j qvttt 4 - — This salt, which is iso- morphous with potassium sulphate, can easily be obtained in the crystalline state. It melts at 140 , and decomposes at higher tem- peratures into ammonia, water, nitrogen, and ammonium sulphite, which sublimes with some unchanged ammonium sulphate. It is manufactured on a large scale by leading the ammonia evolved on heating the ammoniacal-liquor of the gas workswith milk of lime into dilute sulphuric acid until saturation, and then evaporating (ATT „„„ is also known. Besides these compounds, normal salts, in which one of the atoms of ammonium in ammonium sulphate is displaced by potassium or Ammonium Carbonate. 383 sodium, also exist; they have the composition: SoJSKtt and - (U±\H 4 so (ONa „ 1 bu "(NH0 4 + 2ti 2°- Ammonium Nitrate: NO, -ONH 4 , crystallizes in six-sided prisms, without water, and is isomorphous with potassium nitrate. It is easily soluble in water, absorbing heat, and deliquesces in the air. When heated it first melts and is then converted into nitrous oxide and water (p. 198). Ammonium Phosphates. — The compounds of phosphoric acid with ammonia closely resemble those of sodium and potassium. A salt corresponding to the monacid phosphate: Pol)5v/ i ' i I OH in which one atom of ammonium is displaced by one of sodium, is the compound called microcosmic salt, or monacid sodium-ammo- fOH nium phosphate : PO ON H 4 + 4H„0. It is easily obtained by dis- lONa solving 6 or 7 parts of common sodium phosphate and 1 part of ammonium chloride in 2 parts of boiling water, and separates out on cooling in well-defined crystals. It may be freed from the sodium chloride which it contains by recrystallizing with the addi- tion of a little ammonia. This salt loses its water when heated, and then gives off am- monia, leaving sodium metaphosphate : PO, -ONa, as an easily fusible mass. For this reason the salt is used in blowpipe analysis. A fragment heated on a platinum wire in a Bunsen burner gives a glass bead of sodium metaphosphate, which, like borax, can dissolve various metallic oxides and become coloured by them. Silica is not dissolved by the fused salt, and may thus be readily detected. Ammonium Carbonate. — On the union of ammonia and carbonic acid, different compounds are produced according to the conditions under which the combination takes place. It has already been stated (p. 290) that if the two gases act upon one another in the absence of water, ammonium carbamate : CO }t^h 1 These may also be regarded as double salts of the two sulphates, thus : c „ f OK , c/ ~, fONH 4 , c „ fONa, Qrv fONH,, „ _ , 4 S0 2 | 0K +S0 3 | ONH * and SO a | ONa +S0 2 -j NH 4 +4H 2 ' but are bes ' represented by the above formulae. — Ed. 384 Text-Book of Inorganic Chemistry. and not ammonium carbonate is formed. The same salt is also produced when aqueous ammonia is saturated with carbonic acid, although the chief product of this reaction is ammonium carbo- nate. And it is highly probable that the normal as well as the acid salt which crystallize from concentrated solutions (the former when ammonia is in excess, the latter with excess of carbonic acid) contain ammonium carbamate. The commercial ammonium carbonate {carbonate of ammonia) is manufactured by heating a mixture of ammonium chloride or sulphate with powdered chalk ; it then sublimes as a translucent, crystalline solid, with a strong ammoniacal odour. When exposed to the air, it crumbles to a white crystalline powder, consisting chiefly of acid ammonium carbonate. This substance was pre- viously thought to be a mixture of the normal and acid ammonium carbonates, of the composition : CO -j q,, „ 4 + CO j q N tt hence its old name of sesquicarbonate, but is now known to contain considerable quantities of ammonium carbamate, and is therefore a mixture of several salts. The presence of the ammonium carbamate may be readily shown by shaking the freshly prepared solution with calcium chloride, and filtering off the clear solution from the precipitated calcium carbonate. This solution contains calcium carbamate, which gradually decomposes on standing, more quickly if heated, with the formation of insoluble calcium carbonate. A solution of ammonium carbonate free from the carbamate does not give the same result ; nor does a solution of the commercial salt which has been prepared for some time, since ammonium carbamate when dissolved in water gradually, or quickly if warmed, takes up the elements of water, and becomes converted into ammonium carbonate. Ammonium carbonate is produced during the destructive dis- tillation of nitrogenous organic substance — e.g. coal. And although the quantity of nitrogen contained in coal is very small, and only a part of this nitrogen is converted into ammonia, still the immense quantities of coal which are used for the manufacture of coal-gas are so great that very nearly all the ammonium compounds now brought into trade are derived from the ammoniacal liquors of the gas-works. Ammonium carbonate is also produced by the putrefaction of organic compounds containing nitrogen. Large quantities are always present in putrid urine and sewage, and are produced from Ammonium Chloride. 385 the decomposition of the urea which they contain. This compound takes up two molecules of water during putrefaction, and becomes converted into ammonium carbonate : — CH 4 N 2 + 2H 2 = Cojg^H* Urea Ammonium Chloride {Sal-ammoniac) : NH 4 C1. — This salt is brought into trade in two forms. First, when crystallized from water as a white powder of small octahedra, or cubes, grouped together ; and secondly, when obtained by sublimation as a fibrous, crystalline, and compact mass. It possesses a saline taste, and is soluble in about 2-5 parts of water at the ordinary temperature, or in rather more than its own weight at 100° It is nearly insoluble in alcohol. When heated it volatilizes, without melting and with- out decomposition. At a higher temperature its vapour dissociates into hydrochloric acid and ammonia, which reunite on cooling. Informer times large quantities of sal-ammoniac were imported from Egypt, where, for lack of other fuel, dried camel's dung was burnt. The nitrogen and common salt contained in the dung gave rise on burning to ammonium chloride, which was deposited with the soot in the chimneys, and was afterwards purified by some crude method. Beautiful crystals of ammonium chloride are sometimes found in the cavities of solidified lava, and it was formerly thought that the salt had been sublimed with the lava from the interior of the earth. But it has been proved by Bunsen, from observations of volcanic phenomena in Iceland, that crystals of ammonium chloride are only found where the glowing lava has flowed over ground covered with vegetation. The ammonia produced by the dry dis- tillation of the plants would then unite with the free hydrochloric acid contained in the lava, and so produce ammonium chloride. Ammonium chloride, like all other ammonium compounds, is, now almost exclusively obtained from the ammoniacal liquor of the gas-works (p. 322). The ammonia which is evolved on heating the gas-liquor with milk of lime is passed into dilute hydrochloric acid until saturated, and the solution so obtained then evaporated down. The crude salt is afterwards purified, either by crystalliza- tion or by sublimation. Sal-ammoniac is largely used in the arts for the preparation of ammonia (p. 184), as a valuable medicine, in dyeing, and for many other purposes. c c 386 Text-Book of Inorganic Chemistry. Ammonium Bromide: NH 4 Br, closely resembles the chloride, but is still more soluble in water. It is largely used in photography and in medicine. Ammonium Sulphide : (NH 4 )„S, and Ammonium Sulphy drate : (NHJHS. — Both compounds may- be obtained in the crystalline form by allowing dry ammonia and dry sulphuretted hydrogen to act upon one another at a low temperature (- 15°). When the ammonia is in excess the former compound is produced, with excess of the sulphuretted hydrogen, the latter. Ammonium sulphydrate may also be prepared by leading sulphuretted hydrogen into a saturated solution of ammonia in absolute alcohol, when the compound crystallizes out. The two compounds are usually pre- pared in aqueous solution. If aqueous ammonia is completely saturated with sulphuretted hydrogen, ammonium sulphydrate is formed, and if to this liquid the same quantity of ammonia is added as was used for its preparation, it is converted into ammonium sulphide. Both are colourless liquids, soon becoming yellow when exposed to the air, and smelling both of ammonia and sulphuretted hydro- gen. The yellow colour which the solutions acquire when exposed to the air is due to partial oxidation of the sulphuretted hydrogen which they contain. During this process water and free sulphur are produced, and the latter substance then dissolves in the unde- composed sulphide. Ammonium sulphide, like potassium sulphide, can dissolve large quantities of sulphur, and so produce polysul- phides corresponding to those of potassium. Ammonium sulphide is much used in analytical chemistry, partly to precipitate the sulphides of those metals which are insoluble in water, but which are not formed by the action of sulphuretted hydrogen on an acid solution, and partly to dis- solve those insoluble sulphides which have the character of sulpho-acids (e.g. antimony sulphide). With such sulphides it forms soluble sulphosalts, and they may be thus separated from other insoluble sulphides (e.g. lead sulphide) which have not this property. Oxy ammonium Salts. — Just as ammonia unites with acids and forms the ammonium salts, in the same manner the derivative of ammonia — oxyammonia or hydroxylamine : NH 2 -OH (p. 186) — Detection of Ammonium Compounds. 387 produces with acids the salts of the radical oxyammonium. The radical itself: NH 3 OH or (NH 3 OH) 2 , is as little known in the free state as is ammonium. Oxyammonium Sulphate : SO^ \ njxrtr Y> yX crystallizes from its aqueous solution in large colourless prisms, and is insoluble in alcohol. Oxyammonium Chloride : NH 3 OH ■ CI, is soluble both in water and in absolute alcohol, and by this latter property can be separated from ammonium chloride. It crystallizes from water in tablets, from alcohol in monoclinic prisms, melts at 151°, and is decom- posed at a slightly higher temperature into nitrogen, hydrochloric acid, ammonium chloride, and water. Detection of Ammonium Compounds, The similarity between the ammonium and potassium com- pounds is so close that on mixing solutions of ammonium chloride and platinum chloride a yellow crystalline precipitate of ammonium chlorplatinate : (NHj) 2 PtCl 6 , is produced, just as difficultly soluble in water, and just as insoluble in alcohol as the potassium com- pound. If, however, the ammonium compound is heated, ammo- nium chloride and chlorine are evolved, and pure spongy platinum remains behind, from which water extracts nothing. Compounds of ammonium are best recognized by their beha- viour with caustic alkalies (potash, soda, or lime), which liberate ammonia even in the cold, but more readily when gently warmed. This free ammonia can be detected by its odour, and by the white fumes of ammonium chloride produced when a glass rod, moistened with hydrochloric acid, is brought near the liquid. Minute traces of free ammonia or ammonium compounds may be detected by Nesslef 's solution — an alkaline solution of potassium- mercuric iodide. When a few drops of this solution are added to a dilute solution containing ammonia or a salt of ammonium, a reddish-brown precipitate or corresponding colour is at once pro- duced. In this way it is possible to detect ~ of a milligramme of ammonium chloride dissolved in 50 c.c. of water. 388 Text-Book of Inorganic Chemistry. METALS OF THE ALKALINE EARTHS. The four metals belonging to this group : calcium, strontium, barium, and magnesium, are distinguished from the metals of the alkalies by their weaker attraction for oxygen, and further by the fact that their oxides, or rather hydrates, are much less soluble than those of the alkali-metals, and do not deliquesce in the air. Their sulphates (except that of magnesium) are also much less soluble in water than the alkaline sulphates, and their carbonates are quite insoluble. By far the most important of this group of metals is calcium, the one we have placed first. CALCIUM. Chemical Symbol : Ca. — Atomic Weight : 40. Calcium, with oxygen, silicon, and aluminium, belongs to those elements which are most widely distributed in nature. It is never found in the metallic state, but usually in combination with oxygen as some salt, never as the free oxide. In combination with fluorine it forms the important mineral fluor-spar, and occurs as its chloride in many mineral waters and in the sea. The most important oxygen compounds of calcium found in nature are the very various forms of its carbonate, its sulphate (as gypsum), its phosphate, and its silicate, which is a constituent of many rocks. Metallic calcium, which is difficult to prepare pure, is only of theoretical interest. It may be obtained by heating calcium iodide with sodium in a well-closed iron crucible, or by heating a mixture of 3 parts of calcium chloride, 4 parts of zinc, and 1 part of sodium in a Hessian crucible. An alloy of zinc and calcium is thus obtained, from which the zinc may be volatilized by strong heating in a crucible of gas-carbon, placed inside a larger one of clay, then leaving a regulus of metallic calcium behind. The best method of all of obtaining pure calcium is by the electrolysis of its fused chloride. Calcium is a brilliant pale yellow metal, of about the same hard- Calcium. 389 ness as lead. Its specific gravity is i-6 ; it is malleable, but becomes brittle when hammered. When exposed to moist air, calcium soon becomes covered with a grey coating of oxide and is gradually converted into calcium hydrate. It decomposes water at the ordinary temperature, but the hydrogen does not catch fire. The metal burns brilliantly when heated to redness in the air, or in chlorine. Dilute nitric acid acts violently on it, but the concen- trated acid does not attack it. Calcium, like the other metals of the group, is nearly always a dyad element ; it appears to possess a valency of four in the per- oxide, and possibly also in bleaching-powder. Calcium Oxide, Quick-lime : CaO. — This oxide can be far more easily obtained than the oxide of potassium or sodium. It suffices to heat the carbonate strongly in the air in order to expel the whole of the carbonic acid, and leave the oxide behind. The product so obtained is then more or less pure according to the purity of the material used for its preparation. Crystallized calcite, especially transparent Iceland-spar, yields oxide chemically pure, and white marble nearly so. It is remarkable that calcium carbonate is only completely de- composed by heating when a second indifferent gas is present. If heated in an atmosphere of carbonic acid, decomposition goes on until a certain pressure (depending on the temperature) is reached, and then ceases. Marble or Iceland-spar when heated in a closed crucible is only partially decomposed, because it is always sur- rounded with an atmosphere of carbonic acid. Complete decom- position can, however, be easily brought about by placing a piece of charcoal on the bottom of the crucible. The glowing charcoal then converts the carbonic acid into carbonic oxide, into which the former gas can diffuse until all the calcium carbonate is decom- posed. Calcium oxide or quick-lime is a white amorphous mass, usually retaining the form of the calcium carbonate from which it has been prepared. It is infusible, even in the oxy-hydrogen blowpipe, but becomes then so strongly heated as to emit a most intense white light (lime-light). When exposed to the air, it absorbs water and carbonic acid, and crumbles to a fine white powder, which effer- vesces with hydrochloric acid. If water is poured on to quick-lime, the porous mass first absorbs considerable quantities of the liquid like a sponge. After a short 390 Text-Book of Inorganic Chemistry. time chemical action sets in, and the calcium oxide unites with water to form calcium hydrate : — CaO + H,,0 = Ca(OH) 2 . The reaction is accompanied with so large an evolution of heat that a portion of the water is converted into steam. At the same time the piece of quick-lime swells up and is finally converted into a dry, white, soft powder of the following compound, calcium hydrate. (C\xs qtt — This hydrate is slightly soluble in water, I part requiring 600 parts of cold water to dissolve it. In hot water it is less soluble than in cold, whence it follows that a cold saturated solution when boiled becomes turbid, and calcium hydrate separates in the crystalline form, but redis- solves as the solution again cools. The solution has a distinct alkaline reaction. Slaked-lime mixed with water forms a turbid solution of the hydrate with undissolved suspended particles. Such a mixture is called milk of lime. If this is filtered, the suspended particles remain behind on the filter, and clear alkaline lime-water runs through. Lime-water must be preserved in well-stoppered bottles, as it absorbs carbonic acid from the air and becomes turbid from the insoluble calcium carbonate so produced. This property of lime-water makes it useful for detecting carbonic acid, even when mixed with large quantities of other gases — e.g. in coal-gas. Calcium hydrate easily loses its water when heated to redness — more easily than calcium carbonate parts with its carbonic acid — and is reconverted into calcium oxide. Slaked-lime is largely employed for numerous and highly im- portant technical purposes. Immense quantities are used for the preparation of mortar. For this purpose it has been employed for thousands of years, and it may probably be considered as one of the first chemical compounds artificially prepared by man. It would be interesting to know how and when the knowledge was acquired that limestone, when heated to redness, becomes changed into a new substance, and that this quick-lime when made into a paste with water gradually hardens in the air, or if placed between stones, cements them together. But questions of this nature must, for the present, remain unanswered, since the preparation of mortar from lime was known even in pre-historic times. Calcium Hydrate. 391 Ordinary mo?-ta?- is prepared by mixing slaked-lime and sand together with water to form a paste, which is spread between the bricks or stones to be cemented together. The hardening of the mortar only takes place slowly in the course of years — many years if the walls are thick, and is the more complete the older the walls. And not only does the mortar itself gradually become hard, but it also attaches itself firmly to the stones between which it is placed, and it often happens when old buildings are pulled down that the stones themselves break more easily than the mortar which holds them together. The same effect may be noticed in some of the calcareous conglomerates or pudding-stones, which often consist of quartz pebbles cemented together with limestone, and which if broken by the blow of a hammer usually fracture across the pebbles and not at the joints. The hardening of mortar is due to the action of the carbonic acid of the air, which gradually unites with the calcium hydrate to form calcium carbonate, while water is expelled :— Ca |oH + C0 - = c °-°3 Ca + H *°- The sand which the mortar also contains serves to make the mass more porous, so that the carbonic acid can penetrate into the interior of the walls ; it also causes the mortar to adhere more firmly to the bricks or stones. That old buildings — e.g. those of the Romans or of the Middle Ages— possess greater solidity than our modern structures is not only because a better kind of mortar was used, but is a natural consequence of the longer action of the carbonic acid. A mortar which has only been exposed to the air for a few years, and which still contains unchanged calcium hydrate, is of course less hard than one on which the carbonic acid has acted for centuries, and in which the whole of the calcium hydrate has been converted into calcium carbonate. It is considered unhealthy to live, and especially to sleep in the rooms of newly built houses, because the moisture which is contained in the walls continually saturates the air of the rooms. The water mechanically mixed with the mortar soon evaporates from the porous walls, but the walls still remain moist from the water which is gradually set free chemically during the union of the calcium hydrate with carbonic acid. The hardening of the mortar, and consequent drying of the walls, can be accelerated by burning coke or charcoal in open grates in the closed rooms. The 392 Text-Book of Inorganic Chemistry. large quantities of carbonic acid so produced then quickly pene- trate into the walls, and the water which is set free is evaporated by the heat ; and to allow this water vapour to escape, it is necessary to open the doors and windows occasionally. In this way the walls of new buildings may be dried in as many hours as they would otherwise require weeks or even months. The manufacture of the large quantities of quick-lime which are used for the preparation of mortar is carried on in furnaces specially constructed for the purpose — called lime-kilns. Two forms of these kilns are shown in figs. 61 and 62. Both are built of firebricks in the form shown. In the one kiln (fig. 61), which is but little used, large pieces of limestone are built below in the Fig. 61. Fig. 62. form of an arch, and the whole kiln filled up with the same material, leaving free spaces for the hot gases to pass through the entire mass. A fire is then made in the arch below, and continued until the whole of the limestone has been converted into quick-lime. After cooling, the lime is extracted, and the process repeated. Fig. 62 represents a much more economical form, as it is con- tinuous in its action. The kiln is filled with pieces of limestone, and a fire kept up continuously at the side, a. The lime is raked out from the bottom of the kiln as it is produced, and fresh lime- stone introduced at the top to supply its place. A form of kiln largely used in England is similar in shape to fig. 62, but without a side opening for the fire. The kiln is filled with alternate layers of limestone and small coal, and the lowest layer of coal ignited. Calcium Hydrate. 393 As the coal gradually burns away, the lime is abstracted from the base, and fresh alternate layers of limestone and coal are intro- duced from the top. Some varieties of limestone contain other compounds than calcium carbonate, such as magnesium carbonate, alumina, and silica. If these are burnt, a lime is obtained which only becomes slightly heated when slaked, and which is unsuitable for the pre- paration of ordinary mortar. In the case of an argillaceous limestone — i.e. one mixed with clay (aluminium silicate), the lime decomposes some of the clay when moistened with water, producing compounds of lime and alumina and of lime and silica, which are insoluble in water, and which, in contact with water, gradually become hard and solid. If such an argillaceous limestone or a mixture of limestone and clay is burnt, then powdered and mixed with water, it gradually hardens, owing to the formation of these insoluble compounds, and becomes harder and harder the longer it remains in contact with water. Mortars of this description, known as hydraulic mortars, are therefore especially suited for buildings which are always covered with water. Hydraulic mortars are prepared from various materials and in varying proportions — some hardening slowly and some quickly. Therapidity with which the mortar hardens dependsuponthepropor- tion of clay present : thus Roman cement, containing about 30 per cent, of clay, hardens in a few hours ; while Portland cement, largely made in England by mixing chalk with clay and then carefully burning, contains less clay and hardens more slowly. A hydraulic mortar may also be prepared by mixing quick-lime with soluble silicic acid or with an aluminium silicate which has been acted on by hydrochloric acid ; such a silicate is decomposed by the lime in the presence of water. Calcium hydrate is also used for a large number of other tech- nical purposes — for example, for the preparation of ammonia from ammonium chloride, for the manufacture of the caustic alkalies and of bleaching powder, for the saponification of fats, for the purifica- tion of coal-gas, and the manufacture of hard glass, in tanning to remove hair from the skins, &c. Besides calcium oxide, a second, far less important oxygen compound of calcium is known, viz. : — Calcium peroxide : Ca0 2 , which cannot, like barium peroxide, be prepared by heating calcium oxide in oxygen. It is, however, 394 Text-Book of Inorganic Chemistry. easily obtained by adding hydrogen peroxide to lime-water, when it is precipitated in crystalline plates. When gently heated it readily loses one-half of its oxygen. Of the other compounds of calcium, the following are of most general and chemical importance. Calcium Sulphate: SO.j-OXa. — This slightly soluble salt is largely distributed in nature in different forms, and usually accom- panies common salt. The anhydrous compound, of the above composition, is known as anhydrite, and occurs either in crystals belonging to the rhombic system or more generally in granular or fibrous crystalline masses. Far commoner than this is the compound of calcium sulphate with two molecules of water: SGy0 2 Ca + 2H 2 0, which in its most general form is called gypsum. This compound is found in nature in very various modifications, but usually as a compact crystalline mass, generally of a greyish colour : gypsum proper or common gypsicm. Colourless, transparent crystals are also often found : they crystallize in the monoclinic system, and twins are common. A granular, crystalline pure variety of gypsum, highly valued as an ornamental stone, is known as alabaster. Finally gypsum is often found as fibrous masses : fibrous gypsum ; or in transparent plates, which very easily split into thin leaves, as selenice. Selenite is largely used for optical purposes. Gypsum may be readily pre- pared artificially by mixing a solution of calcium chloride with sulphuric acid or a sulphate, when it falls as a crystalline preci- pitate. Crystalline gypsum is so soft that it can be scratched with the nail. As mentioned above, it is only slightly soluble in water, either hot or cold — one part requires nearly 400 parts of water for solution. The solubility is increased by the addition of common salt, so that one part of gypsum dissolves in about 120 parts of a saturated solution of common salt. The two molecules of water which gypsum contains are only loosely united to it. It parts with the greater portion when heated up to 1 20 , but the residue is only expelled above 200 . Gypsum which has been heated to a temperature not much above 120 (Plaster of Paris) acquires the property of again uniting with water when mixed with it, and of setting to a hard mass. Plaster of Paris, prepared as above, and ground to a fine powder, is largely used for various technical purposes. When mixed with water it Calcium Sulphate. 395 soon sets hard, and can therefore be employed as a cement, which will not, however, withstand the action of water. As the hardening (with an evolution of heat) sets in soon after the plaster is mixed with water, only a small quantity must be prepared at a time ; but if a strong solution of common alum is used instead of water the setting takes place more slowly, and so allows more time for mani- pulation. If a stronger cement is required, the plaster of Paris is first mixed with iron filings and then with water. Such a mixture is sometimes used to cement the iron parts of a building into the walls. It is a well-known, but remarkable fact, that if gypsum is heated considerably above I20°it loses the property of uniting rapidly with water. And in the same manner, the naturally occur- ring anhydrous calcium sulphate (anhydrite) cannot unite with water to form a hard mass. Both forms combine, however, slowly with water when left in contact with it for a long time, and finally produce ordinary gypsum. As the mixture of plaster of Paris and water readily adapts itself to any surface with which it is brought into contact, it is used for producing copies of coins and medals, as well as of busts and statuettes (plaster casts). It was also previously used in the process of stereotyping, in which a cast is taken ofthetype, and from this a second cast with fusible metal, which can then be used for printing just as the original type. Plaster casts for this purpose have now been to a large extent superseded by paper pulp, which has the special advantage that it is not so brittle as the dry plaster. If plaster of Paris is mixed with a solution of glue or gelatine in- stead of with water, it solidifies more slowly, and is hard enough to take a good polish. The mixture may be coloured with various metallic oxides, and is used to prepare artificial marbles, &c. Large quantities of plaster of Paris are used to coat the interior walls and ceilings of our rooms, and it is also employed in medicine and agriculture. The surgeon uses it for the preparation of plaster bandages, which preserve a joint surrounded with them from move- ment ; and the agriculturist employes it to impart ammonia from the air to plants sprinkled with it. Plaster of Paris has the power of absorbing ammonia, and attracts this substance when exposed to the air, yielding it again slowly to the plants with which it is in contact. Naturally occurring gypsum and anhydrite do not possess this property. 396 Text-Book of Inorganic Chemistry. Calcium Nitrate: N( y j 0,Ca + 4H 2 0, may be easily obtained by neutralizing nitric acid with calcium carbonate, and crystallizes from its aqueous solution when evaporated down to a syrup. It is easily soluble both in water and alcohol. The salt, when freed from its water of crystallization by heating, forms a porous white deliquescent mass. Calcium nitrate is produced when lime— e.g. that contained in the mortar of walls — is exposed to the vapours of ammonia, which, in the presence of the strong base, is then oxidized to nitric acid by the oxygen of the air. In buildings where considerable quantities of ammonia are pro- duced — e.g. in stables and cow-stalls — the quantity of calcium nitrate accumulating in the walls may often be sufficient to crystallize out in dry weather, and its deliquescent nature is one reason why the walls of such buildings are usually moist. Even the human body emits small quantities of ammonia, which is also absorbed in the walls and converted into calcium nitrate. For this reason the rubbish of old walls when extracted with water often yields con- siderable quantities of calcium nitrate. POl Calcium Phosphate: Ca 3 P = 8 , or p„h0 6 Ca3. — This salt is widely distributed in nature, and is sometimes found in consider- able quantities ; and since its immense importance in agriculture has been recognized, new localities of its occurrence have been and probably will be discovered. To the minerals which consist essentially of calcium phosphate belong phosphorite, found in compact masses in some parts of Spain, Germany, and other places, and sombrerite, imported in large quantities from some of the islands of the Antilles, particu- larly Sombrero. In combination with calcium fluoride and chloride, it forms the mineral apatite: 3(PO).,0 6 Ca 3 + Ca(F,Cl)„, which is found either crystalline or in a massive form. Finally calcium phosphate is the chief mineral constituent of bones, which when burnt leave a residue of calcium phosphate and carbonate, bone- ash. The excrements of many animals, especially of those which are carnivorous — coprolites (petrified droppings), urine, guano, &c. — are all more or less rich in calcium phosphate, and are therefore valuable manures. The calcium phosphate which we require for the growth of our. Calcium Phosphate. 397 bodies, principally for the formation of bone, we derive directly and indirectly from plants, and these again, to whom it is just as neces- sary, absorb it by their roots from the soil. Calcium phosphate is as necessary as carbonic acid for the growth of most plants, par- ticularly cereals. A grain of wheat, barley, or rye, planted in a soil absolutely free from phosphates, will, it is true, germinate and grow to a certain stage, but will produce no fruit. And if we sow cereals year after year on the same plot of ground, and so gradually make it poorer and poorer in calcium phosphate without returning any to the soil, the crops become less and less, and finally the ground is unfruitful. For centuries it has been known to the farmer that he can pre- serve the fertility of his land by the addition of animal excrements, without suspecting what was the action of these substances. But Liebig, the founder of scientific agriculture, showed that this manure again gives to the land the mineral constituents (amongst other things) which have been abstracted from it by the harvest, and that the productiveness of the soil can be still further increased by the addition of those mineral substances which the plants re- quire for vigorous growth. The scientific agriculturist is, therefore, no longer satisfied with farmyard manure (which is indeed often insufficient in quantity), but adds to his land artificial manures, prepared in chemical factories from suitable materials. To these substances belongs calcium phosphate, and this compound, in the form of bone-meal, or of minerals containing it, has been successfully added to soils originally poor in calcium phosphate or from which this substance has been gradually abstracted by continuous har- vests. Calcium phosphate is, however, insoluble in water, and is only slowly dissolved by water containing carbonic acid, hence it can be only slowly .absorbed by the roots of the plants. A soluble calcium phosphate acts more quickly ; and the farmer requires the phosphate in this form if he is to materially increase the fertility of his land. The chemical process by which this change is effected is a very simple one. Phosphorite, bone-ash, or some other min- eral containing calcium phosphate is powdered and treated with about a half or two-thirds of its weight of sulphuric acid. A por- tion of the calcium is then converted into calcium sulphate, setting free some phosphoric acid which forms a soluble acid salt with the undecomposed calcium phosphate : — 398 Text-Book of Inorganic Chemistry. PO]-O Ca 3 + SO,(OH), = 2SO,-O s Ca + pgjgjg 4 The mixture of calcium sulphate and acid calcium phosphate is not separated, but is brought in this form as a manure into trade, and is called superphosphate. The acid calcium phosphate is soluble in water, and when mixed with the soil is at once absorbed by the roots of the plants. Chemically pure normal calcium phosphate can be obtained by mixing a solution of calcium chloride with one of normal sodium phosphate, or with the common acid phosphate if a little ammonia is afterwards added. The gelatinous precipitate dries to a white earthy powder. It is insoluble in water, but soluble in hydrochloric acid, from which solution it is again precipitated unchanged by ammonia. Monacid Calcium Phosphate : HCaP0 4 , or PO ] o'tV 3 ' is preci- pitated as a gelatinous mass, often becoming crystalline on stand- ing, when a solution of calcium chloride is mixed with one of monacid sodium phosphate and acetic acid : — P0(g£ a >* + CaCl 2 = P0{g£a Like the normal salt, it is insoluble in water. 2NaCl. Diacid Calcium Phosphate: H 4 CaP 2 8 , or pf!fr> r ma y ^ e obtained from the normal or monacid salt by the addition of phos- phoric or hydrochloric acid, and evaporating down. It then crystallizes out in colourless tablets, which, with a small quantity of water, give a crystalline precipitate of the monacid phosphate, but which dissolve in a larger quantity of the solvent. It is the chief constituent of the manure called superphosphate of lime, or simply superphosphate. Calcium Carbonate: CO-O s Ca. — This compound is very widely distributed on the surface of the earth, and in very different forms. It is prepared artificially by precipitating a solution of a calcium salt with sodium carbonate,' and then forms a white amorphous powder, which soon becomes crystalline. It is in- soluble in pure water, even when hot, but dissolves in water con- taining carbonic acid with the formation of an acid carbonate. Calcium Carbonate. 399 Calcium carbonate easily dissolves in dilute hydrochloric, nitric, or acetic acid, with evolution of carbonic acid ; and when heated to redness loses its carbonic acid, and is transformed into calcium oxide. Calcium carbonate is dimorphous : as calcite it crystallizes in the hexagonal system, usually in rhombohedra, while as arragonite it belongs to the rhombic system. Both forms may be produced artificially. If calcium carbonate is precipitated from a solution at the ordinary temperature, the precipitate is at first amorphous, but soon becomes crystalline, and then consists of minute rhom- bohedra like calcite ; the same crystals are also formed when the clear solution of calcium carbonate in carbonic acid is allowed to evaporate slowly. But if the calcium carbonate is precipitated from a boiling-hot solution instead of from a cold one, the crys- talline precipitate then consists of small rhombic prisms, resembling those of arragonite. The calcium carbonate deposited from hot springs (e.g. Carlsbad) usually has this form. The numerous forms in which calcium carbonate occurs in nature are designated by different mineralogical names. Calcite or calc-spar is nearly pure calcium carbonate crystallized in rhombohedra or in some other hemihedral form belonging to the hexagonal system. A variety of calcite, which is distinguished by its purity and transparency, and which is largely used in optical instruments to produce double refraction, is called Iceland-spar. Arragonite is the second form in which calcium carbonate crystallizes. It is usually found' in rhombic prisms, and is consider- ably rarer than calcite. Marble is a granular, crystalline, compact form, produced from limestone by the action of underground heat and pressure. White marble often looks like loaf-sugar, and consists of small crystals of calcite. Marble is found in various colours from white to black. Compact limestone, of various colours according to the impurities which it contains, forms large mountain masses. It has nearly all been produced by the agency of marine animals, whose remains are often found petrified in it. Oolitic limestone, freestone, consists of numerous small globular masses of calcium carbonate cemented together, and is so called because it resembles the roe of a fish. Lithographic stone is a compact, fine-grained form of limestone, found near Solenhofen, in Bavaria, and largely used in the processes of lithography. 400 Text-Book of Inorganic Chemistry. Chalk is a soft, earthy variety of calcium carbonate, and from its white colour is apparently very pure. It consists chiefly of the minute cases of various species of foraminifera, and has been formed in deep seas. That it still contains organic traces of these animals is readily proved by dissolving it in dilute hydrochloric acid. The carbonic acid evolved possesses an unpleasant odour, and a brown organic residue remains behind. Calcareous tufa or travertine is a porous form of calcium car- bonate, produced by the gradual evaporation of a carbonic acid solution of the carbonate. Of a similar origin are the stalagtites found hanging from the roofs of limestone caves, and other similar deposits. Finally, calcium carbonate is a constituent of many products of animal life — e.g. of egg-shells, and of the shells of oysters, and other mollusca. On some parts of the coast, where lime is scarce, oyster-shells are burnt to obtain quick-lime. The technical applications of calcium carbonate are exceedingly numerous and important. Marble is a valuable material for the builder and sculptor. Chalk is used for writing, and, when finely powdered, as a pigment. Compact limestones are employed for the manufacture of glass, and as a flux to produce a fusible slag in the extraction of many metals from their ores. And, finally, im- mense quantities of limestone are used for the manufacture of quick -lime and then of mortar. Bleaching Powder. {Chloride of Lime.) — The chief compound contained in this important substance has (p. 117) the composition : Ca(OCl)Cl or Caj r , but the pure compound corresponding to this composition has not yet been prepared. Bleaching powder is a soft, white, amorphous powder, with a faint odour of hypo- chlorous acid, soluble in water with decomposition into calcium chloride and calcium hypochlorite and easily decomposed by all acids. Bleaching powder is manufactured on a large scale by saturat- ing dry slaked-lime with chlorine. The gas is generated in leaden retorts and led into chambers containing shelves on which the slaked-lime is placed in thin layers. As soon as the lime ceases to absorb the gas, the process is finished ; the chambers are then opened and the bleaching powder preserved in well-closed vessels. Bleaching Powder. 401 A solution of bleaching powder in water does not itself bleach, or only very slowly as it is decomposed by the carbonic acid of the air. If, however, a dilute acid is added, the bleaching-powder is decomposed, yielding hypochlorous acid, and then free chlorine. The articles to be bleached are first thoroughly cleansed, dipped intoadilute solution of bleaching powder, and then into a very dilute solution of hydrochloric acid, which, with the bleaching powder they contain, liberates free chlorine (p. 117). Thorough washing is necessary afterwards, as any small traces of chlorine which remain behind gradually rot the stuff. But as it is impossible to completely remove the last traces of chlorine, it is customary to dip the articles in a solution of some substance which acts chemically upon it — called an antichlor. Such a substance is sodium thiosulphate (P- 37o). _ The formula given above for bleaching powder and the assump- tion that it is decomposed by water are supported by many facts. The action of this substance on bodies which easily take up chlorine and oxygen is quite different in the absence and in the presence of water. Absolute alcohol and dry bleaching powder react ener- getically upon one another and produce first ethyl hypochlorite, and then acetic ether, but if aqueous alcohol is used, the chief product is chloroform. We have already seen (p. 117) that a solution of bleaching powder yields calcium chloride and chlorate when boiled. Calcium Silicate: SiO-0 2 Ca. — The normal salt of this com- position is produced as a gelatinous precipitate, drying to a white amorphous powder, when calcium chloride in solution is added to a solution of sodium silicate. This compound is found in nature as the mineral wollastonite, and united with other silicates as double salts. To these belong many of the zeolites, lime-garnet, datolite, &c. Similar double silicates, but not crystalline, are the chief varieties of GLASS. This important substance is essentially a double silicate of potash or soda and lime, in which the lime is sometimes displaced by lead oxide. It is distinguished by its transparency, hardness, fusibility and resistance to many solvents which attack other substances, even the metals, and is therefore indispensable to the chemist. D D 402 Text-Book of Inorganic Chemistry. Glass is manufactured by fusing together a mixture of calcium carbonate, potassium (or sodium) carbonate and silica at a high temperature, in specially constructed furnaces and crucibles. At a bright-red heat the silica expels the carbonic acid from the carbonates, producing a double silicate of calcium and potassium or sodium, which on continued heating forms a homogeneous liquid, and on cooling solidifies to a transparent glass without a trace of crystalline structure. The purer the materials employed, the purer and less coloured is the glass obtained. In order to give the required shape to the glass it is partly cast and partly blown. Large slabs of thick glass (plate-glass) are obtained by pouring the liquid substance on to horizontal slabs of iron, and the com- moner kinds of glass vessels (bottles, tumblers, &c), especially the thicker ones, are cast in moulds, just as liquid iron is cast. The majority of glass vessels are, however, made of what is called blown glass. The workman takes a mass of semi-fluid glass on the end of an iron tube, and by blowing, combined with various movements which cannot be well described, gives to it almost any desired shape. Most of the various shaped vessels used in daily life, and in the laboratory — wine-glasses, bottles, funnels, flasks, &c. — are made in this way. Glass is a brittle substance, and if of considerable thickness, or rapidly cooled, easily breaks when subjected to rapid changes in temperature. In order to allow the molecules of the glass to arrange themselves regularly, that the tension on cooling may be the same throughout the entire thickness, it must be annealed, or allowed to cool slowly. For this purpose, the glass vessels, after they have been made, are brought into the annealing furnaces, where the glass is heated nearly to the temperature at which it becomes soft, and then allowed to cool very slowly during several days or even weeks. The influence which rapid cooling has on the properties of the glass is well seen in what are called Rupert's drops, which are made by dropping liquid glass into cold water. Owing to the rapid cooling, they are hard, and stand pressure or even a blow. But if slightly scratched, or if a small piece is broken off the thin end, the whole drop breaks up into small pieces of glass. A few years ago a variety of quickly annealed glass was intro- duced into trade, and highly recommended for various purposes. This toughened glass, as it is called, is harder than common glass, and bears rapid changes of temperature better. It is prepared by dipping the freshly blown hot glass into a bath of hot oil or paraffin, Glass. 403 a process which resembles that by which Rupert's drops are formed, except that the cooling is less sudden. Vessels prepared in this way may be dropped from a considerable height on to the floor without breaking, and a glass dish filled with cold water may be placed on glowing charcoal, but, like Rupert's drops, they will not stand the slightest scratch. A sharp grain of sand is often sufficient to cause their violent fracture into a number of small pieces, and it sometimes happens that vessels of toughened glass, which have been used for years, and have withstood considerable changes of temperature without damage, suddenly break up into a thousand pieces. Toughened glass vessels, from which so much was expected at the time of their introduction, should therefore never be used, especially for chemical apparatus. Various kinds of glass are distinguished from one another according to their constituents, and according to the purity of the- substances employed in their manufacture. Potash, or hard-glass, is a double silicate of calcium and potassium, and is largely manu- factured in Bohemia. It is less acted on by chemical reagents than any other kind of glass, and when the proportion of silica is high, is difficultly fusible. The glass tubes used for combustions in organic analysis are made from this difficultly fusible glass. Crown, or soda-glass, contains soda instead of potash, and is the common window-glass. It is more easily fusible than potash-glass, and more readily attacked by acids. Bottle-glass is a coarse form of crown-glass, made with impure materials, and more or less coloured by the presence of iron. It also contains not inconsider- able quantities of aluminium silicate. Flint-glass is a double silicate of lead and potassium — i.e. it is potash-glass in which the calcium has been displaced by lead. It is prepared by fusing together silica, potassium carbonate, and lead oxide (litharge), and is distinguished by its lustre and high index of refraction. Flint- glass is the most easily fusible of all glasses, and the most easily attacked by acids ; large quantities of it are manufactured in England and France. Flint-glass is especially valuable for optical purposes, owing to its high index of refraction. The refractive power of the glass can be increased by using a larger proportion of lead oxide ; such glass, called strass, is used for the manufacture of artificial gems. Finally, the various slags obtained in metallurgical operations are also a species of glass. In smelting some of the heavy metals, especially iron, calcium carbonate is often added as a flux to form 404 Text-Book of Inorganic Cliemistry. a fusible slag, which removes the earthy impurities, swims on the surface of the heavier metal, and so protects it from the oxygen of the air. The slags consist of the silicates of the alkalies, of lime, alumina, magnesia, iron, &c, and when cold yield a strongly coloured, often opaque, vitreous mass. Glasses of various colours are produced by the addition of different metallic oxides to the liquid glass. The silica of the glass unites, like borax, with these oxides, forming coloured silicates. Blue glass is produced by the addition of small quantities of cobalt ; green by uranium or chromium ; olive-green by ferrous oxide ; yellow or brown by ferric oxide ; amethyst by manganous oxide ; red by cuprous oxide, or a trace of gold, &c. Coloured glass often consists ofcolourless glass with the coloured glass on one side only. Such glass is said to be flashed. White, opaque, or milk-glass, is made by adding some white infusible powder, such as bone-ash, or cryolite, to the molten glass. COl Calcium Oxalate : ^.^ r 2 Ca + H 2 0, is contained in many plants, especially in the sorrels, and is obtained by precipitating a solution of calcium chloride with ammonium oxalate. The preci- pitate is crystalline, and loses its water of crystallization at 100°. The salt is insoluble both in water and in dilute acetic acid, and serves, therefore, for the estimation of calcium and of oxalic acid. calcium Chloride : CaCl 2 . — This salt, which is distinguished by the energy with which it attracts water, and which is, therefore, largely used for drying purposes in chemistry, is obtained as a bye- product in many operations — e.g. in the preparation of ammonia from ammonium chloride and quick-lime (p. 184), and of carbonic acid from calcium carbonate and hydrochloric acid. The acid liquid resulting from the latter process when evaporated down to a small bulk deposits large rhombic crystals of the composition : Ca0 2 + 6H 2 0. These crystals deliquesce in the air and dissolve readily in water with a considerable fall in temperature; and when mixed with snow or powdered ice the temperature may fall as low as — 45° If this crystalline compound is heated, it melts in its water of crystallization and loses 4 molecules of water, leaving a white porous mass, which readily absorbs water again. At a red heat the salt loses all its water and fuses to a clear liquid. This fused mass is poured on to a clean iron plate, broken into Calcium Chloride. 405 pieces when cold, and so brought into trade. Fused calcium chloride dissolves in water with evolution of heat ; it is also solu- ble in alcohol. It energetically absorbs dry ammonia gas, pro- ducing a compound : CaCl 2 + 8NH 3 , which is decomposed into its components either when heated or when brought into contact with water. Calcium chloride cannot, therefore, be used to dry gaseous ammonia. Calcium Fluoride : CaF \, is the mineral fluor-spar. It crystal- lizes in the regular system, but is also found in compact masses. When finely powdered it is used for the preparation of hydrofluoric acid, and as a flux to produce a fusible slag — hence its name (from fluo, I flow). Calcium fluoride is quite insoluble in water, and is therefore precipitated when a solution of a calcium salt is mixed with one of a soluble fluoride. Calcium Sulphide : CaS, is obtained as a yellowish-white earthy mass by glowing quick-lime in a stream of sulphuretted hydrogen, or, less pure, by igniting a mixture of powdered calcium sulphate and charcoal, when calcium sulphide and carbonic oxide are rmed. The compound so produced is insoluble in water, but gradually decomposes in contact with this liquid into calcium hydrate and sulphydrate : — 2CaS + 2H a O = Ca(OH) 3 + Ca(SH) 2 . Calcium sulphide, like the other sulphides of the alkaline earths, shines in the dark after it has been exposed to a bright light — in other words, it phosphoresces. It is the chief constituent of Bal- main's luminous paint, in which form it is used to render match- boxes, buoys, and other objects luminous in the dark. Calcium Sulphydrate : Ca(SH) 2 , is easily obtained in solution by saturating milk of lime with sulphuretted hydrogen. It possesses the property of softening hair and changing it into a gelatinous mass. Mixed with other substances, it is used by some Oriental nations as a depilatory. Calcium Pentasulphide : CaS 5 . — A mixture of this salt and calcium thiosulphate may be prepared by boiling milk of lime with an excess of flowers of sulphur, and filtering off the reddish-coloured liquid. Like potassium pentasulphide, hydrochloric acid decom- poses it with a copious precipitate of milk of sulphur (p, 145). 4-06 Text-Book of Inorganic Chemistry. Detection of Calcium Compounds. Calcium compounds when introduced into a non-luminous gas- flame give a reddish colour, somewhat yellower than that produced by the compounds of strontium. The spectra of the two flames are, however, so different, that they may be very easily distinguished from one another. If the compound is an insoluble one, it must be first moistened with hydrochloric acid before it is introduced into the flame. Calcium salts closely resemble those of strontium and barium. The sulphates are all difficultly soluble in water, that of calcium being the most soluble. The carbonates are all equally insoluble in water, whence it follows that they are all precipitated together by sodium or ammonium carbonate. The oxalates are also all insoluble in water, but that of calcium is distinguished from the other two by its insolubility in acetic acid. Of the anhydrous nitrates only that of calcium is soluble in absolute alcohol, a fact which is utilized to separate calcium from strontium and barium. STRONTIUM. Chemical Symbol : Sr. — Atomic Weight: 87-5. The compounds of this element closely resemble those of calcium and barium, especially the latter, and strontium with its compounds forms a natural transition from calcium to barium. The most important minerals containing strontium are celestine or strontium sulphate, and strontianite, strontium carbonate. The latter mineral is called from Strontian in Scotland, where it was first found in quantity, hence the name given to the element. Recently, large deposits of strontianite have been found in Westphalia, and this mineral is now almost exclusively used for he preparation of strontium compounds. The metal may be obtained by the electrolysis of the fused chloride. Its specific gravity is 2'5, it is harder than lead, and is malleable. Strontium, like calcium, is a yellow metal and melts at a red heat. Strontium Oxide : SrO, cannot be prepared, like calcium oxide, by glowing the carbonate, as this compound is only slowly Strontium. 407 and incompletely decomposed by heat. It is obtained, like barium oxide, by heating its nitrate to redness, and forms a grey, porous mass. This oxide unites with water and produces the hydrate : Sr(OH) 2 , which crystallizes from water as the compound Sr(OH) 2 + 8H,0. Strontium hydrate is more soluble in water than the calcium compound, but less than that of barium ; 1 part dissolves in about 60 parts of cold water. The solution is strongly alkaline. Strontium Peroxide : SrO, is obtained, like the calcium com- pound, by precipitating a solution of strontium hydrate with hydrogen peroxide. Strontium oxide does not unite directly with oxygen. Strontium Sulphate : S0 2 -0 2 Sr, is found in nature, as the mineral celestine, in transparent, rhombic prisms, often of consider- able beauty. The salt may be obtained as a white powder by precipitating a solution of a strontium salt with sulphuric acid or a soluble sulphate. It is slightly soluble in water, but less so than calcium sulphate ; 1 part of the salt requires about 6,000 parts of water for solution. Strontium Nitrate : ,-q 2 \ 2 Sr, is distinguished from calcium nitrate by its insolubility in alcohol, and in other respects resembles the barium salt. It is easily obtained by dissolving strontianite in dilute nitric acid and evaporating down. As strontium compounds impart a red colour to flames into which they are introduced, its nitrate is used instead of potassium nitrate to produce an inflam- mable powder which burns with a red light. Strontium Carbonate : CO • 2 Sr, is obtained artificially like the calcium compound, which it closely resembles. It occurs in nature as the mineral strontianite, either in rhombic prisms or in indistinct crystalline masses. Strontium Chloride : SrCl 2 + 6H 2 0, crystallizes in needles, is easily soluble in water and slightly deliquesces in the air. Like calcium chloride, it is soluble in strong alcohol. Strontium Sulphide : SrS, is obtained from celestine by heating a mixture of the powdered mineral with charcoal, and forms the starting-point in the preparation of strontium compounds from this source. Like calcium sulphide, it is luminous in the dark. 408 Text-Book of Inorganic Chemistry. Detection of Strontium Compounds. Strontium compounds impart an intense crimson colour to the non-luminous flame, which can at once be distinguished from the lithium or calcium flame by the spectroscope. The spectrum of strontium is especially characterized by a fine line in the bright blue. Strontium salts are distinguished from those of calcium by the fact that calcium sulphate produces a precipitate of strontium sulphate on standing. They give no precipitate with potassium chromate, especially if the solution contains free acetic acid, although the chromate is less soluble than the calcium salt. Strontium nitrate is insoluble in strong alcohol and thus differs from calcium nitrate. This property affords a ready means of quantitatively separating the two salts. BARIUM. Chemical Symbol : Ba. — Atomic Weight : 137. This element, which derives its name from the high specific gravity of its compounds (ffapus = heavy), is found in nature like strontium chiefly as the sulphate, barite or heavy spar, and as the carbonate, witherite. The former compound is the source from which nearly all barium compounds are prepared. Metallic barium is exceedingly difficult to prepare. It is said to have been obtained by electrolysis of the fused chloride, and is described as a yellow metal, melting at a bright red heat and de- composing water at the ordinary temperature. Barium, like calcium and strontium, is a dyad element in nearly all its compounds ; in the peroxide : Ba0 2 , it is a tetrad. Barium Oxide {Baryta) : BaO, cannot be prepared by heating the carbonate, as the carbonic acid is more firmly united even than in strontium carbonate, and is only very slightly expelled even at the highest temperatures. Barium oxide is best prepared by heat- ing barium nitrate. A Hessian crucible is about half filled with the salt and heated first gently to expel most of the nitric peroxide, and then, when the frothing has subsided, to bright redness. The oxide which remains behind is a grey porous mass, which has Barium. 409 apparently been fused, but is really infusible at this temperature. Occasionally it has a greenish colour, due to manganese derived from the crucible which has been converted into barium manga- nate. When mixed with water, it evolves a considerable amount of heat and is converted into barium hydrate. When gently heated in the air or oxygen, it absorbs considerable quantities of oxygen and is changed into barium peroxide. Barium Hydrate: Ba(OH) 2 — The preparation of barium hydrate from barium oxide is tedious and unsuitable if large quantities are required. It is better to employ the sulphide, which is easily obtained from the sulphate by heating with powdered charcoal. For this purpose a solution of the sulphide in water is boiled with copper oxide until a portion of the liquid no longer blackens a solution of lead acetate. The clear solution is then filtered off from the insoluble copper sulphide. The reaction is expressed in the following equation : — BaS + CuO + H 2 = Ba(OH) 2 + CuS. From the hot solution so obtained, crystals of the compound : Ba(OH) 2 + 8H IJ separate out on cooling in tablets or prisms. Barium hydrate may also be prepared from the sulphide by first converting this salt into the carbonate by moist carbonic acid- sulphuretted hydrogen being set free — and then acting on the car- bonate with superheated steam — i.e. steam at high pressure. The carbonic acid then passes away with the steam, and barium hydrate remains behind. The crystalline compound parts with its water of crystallization when heated, but, unlike calcium hydrate, cannot be converted into the oxide by heat. Barium hydrate dissolves in about 20 parts of cold water, but requires less than 1 part of hot water for solution. It is thus much more easily soluble than calcium or strontium hydrate. The solu- tion {baryta-water) is strongly alkaline and precipitates the weaker bases from their solutions like caustic potash. When exposed to the carbonic acid of the air it becomes turbid more rapidly than lime-water. Baryta is a weaker base than potash or soda, and is therefore precipitated by these alkalies, not by ammonia, from concentrated solutions of its salts. Barium oxide possesses the remarkable property of forming an insoluble compound with cane-sugar ; and on this account large quantities are used by sugar refiners. Strontium oxide also pos- sesses this property. 41 o Text-Book of Inorganic Chemistry. Barium Peroxide : Ba0 2 .— Barium oxide, when heated to faint redness in a stream of air or oxygen, increases in weight owing to absorption of oxygen, and becomes gradually converted into barium peroxide, which cannot be distinguished by its external properties from the oxide itself. At a slightly higher temperature than that required for its production, but still far below bright red- ness, it again parts with an atom of oxygen, and is reconverted into the oxide, which can again unite with oxygen, and so on. Experiments made on a large scale to ascertain whether in this way large quantities of pure oxygen might not be obtained from the air have not been altogether successful. It appears that the oftener the oxide is converted into peroxide; the less becomes the energy with which it unites with oxygen, until it finally loses this property altogether. The cause of this is possibly because the oxide gradually becomes more compact, and so exposes less surface to the action of the gas. Recently it has been discovered that barium peroxide parts with its oxygen at the same temperature at which it is produced if the pressure is reduced. Or, in other words, to obtain a supply of pure oxygen, the barium oxide is kept at a constant temperature, and the pressure only varied. Possibly some use may be made of this interesting fact. Barium peroxide prepared as above described is never pure, but always contains more or less barium oxide. It may be purified in the following way. It has already been stated (p. 90) that barium peroxide is distinguished from manganese and other peroxides by the fact that it gives hydrogen peroxide, not chlorine, when treated with hydrochloric acid, and its purification depends upon this reaction. The impure barium peroxide, obtained by heating barium oxide in a stream of air, is finely powdered and rubbed up with water in a mortar, with which it unites chemically. The mixture is next gradually added to dilute hydrochloric acid, until the liquid only reacts faintly acid, and the clear solution, con- taining barium chloride and hydrogen peroxide besides the excess of acid, filtered off. On adding an excess of baryta-water to this liquid, the pure hydrated peroxide separates out in lustrous tablets of the composition : Ba0 3 + 8H 2 0. These crystals when washed, and dried between filter paper and over sulphuric acid, gradually lose all their water and leave pure barium peroxide behind. The pure compound remains unchanged in the air, gives up one-half of its oxygen when heated, and unites with water to form the hydrated compound. Compounds of Barium 4 1 1 Barium Sulphate : S0 2 -0.;Ba. — This salt, the most impor- tant naturally occurring compound of barium, is usually found in compact masses, and sometimes in tabular crystals belonging to the rhombic system. It has a specific gravity of 4'4, and is known as barite or heavy-spar. It may be readily prepared arti- ficially, as a heavy, white precipitate, by adding sulphuric acid or a sulphate to a soluble barium salt. Barium sulphate is quite insoluble both in water and dilute acids, and is therefore used to precipitate barium or sulphuric acid from a liquid. Concentrated sulphuric acid dissolves it to some extent, but on dilution with water it is again precipitated, like lead sulphate. Precipitated barium sulphate is used as a white pigment under the name of permanent white ; it is also largely used to adulterate white-lead, for which purpose its high specific gravity, combined with its cheapness, makes it admirably adapted. Barium Nitrate : N( -) 2 r O a Ba, is prepared by dissolving with- erite or barium sulphide in nitric acid, and is deposited from the hot solution in brilliant, heavy octahedra, without water of crystal- lization. The salt is tolerably easily soluble in water, especially when hot, but less soluble in the presence of free acid ; hence, on adding nitric acid to an aqueous solution, a precipitate is produced. It is insoluble in alcohol. Barium nitrate is poisonous, like nearly all the barium salts. Barium Carbonate : CO • 2 Ba, occurs in nature as the mineral witherite, and is thrown down as a white precipitate on adding sodium carbonate to a solution of a barium salt. It is insoluble in water, but slightly soluble in the presence of carbonic acid ; and, notwithstanding its insolubility, is poisonous. Barium Chloride : BaCl 2 + 2H 2 0. — This compound is one of the least soluble of the barium salts. It is prepared in a similar manner to the nitrate, and crystallizes in rhombic plates. At ico° it loses its water of crystallization and melts at a red heat. Alcohol does not dissolve it — a reaction which is employed to separate barium from calcium and strontium. Its aqueous solution is largely used in analytical chemistry as a test for various acids. 412 Text-Book of Inorganic Chemistry. Barium Sulphide : BaS ; and Barium Sulphydrate : Ba(SH) 2 . Barium sulphide is obtained as a white amorphous mass when barium oxide is heated in a stream of sulphuretted hydrogen, or, less pure, by reducing the sulphate with charcoal. An intimate mixture of finely powdered heavy-spar and charcoal is mixed with starch paste and formed into balls. These are then dried in the air, placed between alternate layers of charcoal in a furnace, the bottom layer of which is ignited. When cold, they are extracted and digested with water. The solution, especially if prepared hot, does not contain barium sulphate, but a mixture of barium sulphydrate and hydrate, which crystallize together : — • 2BaS + 2H 2 = Ba(SH) 2 + Ba(OH) 2 . If sulphuretted hydrogen is led into this solution, the hydrate is converted into sulphydrate, and this salt alone crystallizes out on evaporation. As heavy-spar is a much commoner mineral than witherite, barium sulphide is the compound from which nearly all other barium salts are manufactured. Detection of Barium Compounds. Compounds of barium when introduced into a Bunsen flame produce a yellowish-green colour, the spectrum of which is more complicated than that of either calcium or strontium, and is dis- tinguished by several lines in the green. In their chemical characters the barium compounds closely resemble those of calcium and strontium, but it is still easy to dis- tinguish them from one another. Barium sulphate is so insoluble in water, that on adding a solution of calcium sulphate to one of a barium salt an immediate turbidity is produced. In the same manner calcium and strontium chromates are much more soluble than barium chromate, and potassium chromate added to a dilute solution of barium chloride produces a pale yellow precipitate of barium chromate, but no precipitate is produced with calcium (or strontium) chloride. Barium can be separated from calcium or strontium by digesting the perfectly dry chlorides with absolute alcohol. Calcium and strontium chlorides then dissolve while barium chloride remains undissolved. Magnesium. 413 MAGNESIUM. Chemical Symbol : Mg. — Atomic Weight : 24. The compounds of magnesium are not less widely distributed than those of calcium, but are not found in such large quantities. Among the more important minerals containing magnesium are : magnesite (magnesium carbonate: CO -0 2 Mg), 'dolomite (mag- nesium-calcium carbonate : CO • 0„Mg + CO • 0.,Ca); talc, steatite or soap-stone, serpentine, and meerschaum (all essentially magnesium silicate). Other silicates which are rich in magnesium are : augite, hornblende, asbestos, olivine, and biotite or magnesium mica. Mag- nesium also occurs in combination with sulphuric acid as Epsom salt in many mineral waters and in the sea. Magnesium is a malleable silver-white metal of specific gravity = 175. It melts at a low red-heat, and at a higher temperature can be distilled. It remains almost unchanged when exposed to the air, and only becomes covered with a thin, grey coating of the oxide which preserves the metal from further oxidation. When heated in the air above its melting-point it catches fire and burns with a brilliant white light, producing its only oxide — magnesia. A wire or thin ribbon of magnesium very easily takes fire and burns when one end is held in the flame of a Bunsen burner. The light of burning magnesium is so intense and so rich in chemical rays that a mixture of chlorine and hydrogen explodes when ex- posed to it, just as it does in bright sunlight. Magnesium does not decompose water at the ordinary tempe- rature, or at all events only very slowly ; at about 30° a feeble evolution of hydrogen begins, which becomes more rapid as the temperature rises. Dilute acids, as well as ammonium chloride or carbonate, readily dissolve it. Magnesium may be prepared either by the electrolysis of the fused chloride, or else by decomposing the chloride with sodium. The electrolysis of the chloride may be performed in a porcelain crucible partially divided into two divisions by a piece of unglazed porcelain (fig. 63). When the crucible has been filled with the fused chloride, the two electrodes; which are made of gas-carbon and fitted in the lid (fig. 64), are introduced and connected with the poles of a powerful battery (about 8 to 10 Bunsen cells). The metal which is then produced is lighter than the fused chloride, and would rise to the surface and burn in the air if the negative pole were not 414 Text-Book of Inorganic Chemistry. provided with a series of notches which catch the globules of metal in their ascent. For the reduction of magnesium from the chloride by sodium, a mixture is prepared of six parts of anhydrous magnesium chloride, Fig. 63. Kg. 64. one part of fluor-spar, one part of fused potassium-sodium chloride, and one part of sodium cut into small pieces, and the whole thrown into a red-hot Hessian crucible. As soon as the first violent reac- tion is over, the closed crucible is removed from the furnace, and when it is no longer red hot, its contents are stirred with a piece of tobacco-pipe to unite the smaller globules of metal into larger ones. When the crucible is cold it is broken open, and the globules of metal picked out and washed with water. Magnesium is a dyad in all its compounds ; it unites in only one proportion with oxygen. Magnesium Oxide. Magnesia : MgO. This,, the only oxide of magnesium, is a white, light powder, which is infusible at the highest temperatures of our furnaces. When strongly heated it becomes much denser, and its specific gravity increases from yi to yb. It is especially distinguished from the oxides of calcium, strontium, and barium by its almost complete insolubility in water. But although one part of magnesia requires more than 50,000 parts of water for solution, it still pro- duces a distinctly blue spot when moistened with water and placed on red litmus paper. It unites with water with a very slight evolu- tion of heat, forming the hydrate, from which the water is again easily expelled on heating. It attracts carbonic acid from the air (when in the pulverulent state), and, being a strong base, forms compounds with all acids. Magnesium Sulphate. 415 Magnesium oxide is obtained by gently glowing the basic car- bonate, known to pharmacists as magnesia alba, and then gives a fine white powder called magnesia usta or calcined magnesia. The oxide is produced as a solid, compact mass on glowing the compact natural form of magnesium carbonate, magnesite. Mag- nesium oxide gradually hardens in contact with water, and forms a compact mass of the hydrate. Magnesium Hydrate : Mg(OH),, is precipitated on adding caustic soda or ammonia to a solution of magnesium sulphate ; when washed and dried, it yields a white powder. It may be heated to too" without losing its water of hydration, and slowly attracts carbonic acid from the air. It is found in nature as the mineral brucite. Magnesium Sulphate. Epsom Salt: S0 2 -O s Mg + 7H..O. — This salt differs from the sulphates of calcium, strontium, and barium, by its easy solubility in water, and by the readiness with which it crystallizes. Considerable quantities of magnesium sul- phate united with one molecule of water are found in the Stassfurt beds, and are known as Kieserite : S0 2 - 2 Mg + H„0. Magnesium sulphate is further contained in sea-water and in many highly valued mineral waters, among which may be men- tioned those of Epsom near London, Seidlitz, Friedrichshall in Germany, and Hunyadi Janos in Hungary. The salt may be obtained artificially by decomposing magnesite or dolomite with dilute sulphuric acid, and is a bye-product in manufactories of aerated waters when either of these substances is used to prepare carbonic acid. If dolomite is used, calcium sulphate remains un- dissolved. Magnesium sulphate crystallizes in colourless rhombic prisms, and possesses an unpleasant saline and bitter taste. It is easily soluble in water, insoluble in alcohol, and when heated melts in its water of crystallization, of which it loses six molecules under 1 50°. The last molecule is only expelled at a temperature above 200°. This molecule of water — so-called water of constitution — can be displaced by salts, producing double compounds, containing six molecules of water. Thus, when solutions of magnesium and potassium sulphates are mixed, a double salt of the composition : S0 2 -0 2 Mg+S0 2 (OK) 2 + 6H 2 0, crystallizes out. In this respect magnesium sulphate is closely allied to other sulphates of dyad 41 6 Text-Book of Inorganic Chemistry. metals which also crystallize with seven molecules of water (the vitriols), and thus connects the metals of the alkaline earths with the heavy metals. Owing to its well-known action as a mild and safe purgative, magnesium sulphate is a valuable and much used medicine, and mineral waters containing it are among those of which the largest quantities are consumed. Magnesium Phosphates. The normal salt : (PO) a O Mg s , is obtained as a white precipi- tate on mixing a solution of normal sodium phosphate with one of magnesium sulphate. If common (monacid) sodium phosphate is added to magnesium sulphate the monacid phosphate : PO ] ntr + 7H 2 0, is produced, which separates from"concentrated solutions as a white, amorphous powder, and is slowly deposited from dilute solutions in the crystalline form. The salt is difficultly soluble in water, but easily dissolves in acids and in salts of ammonium. If ammonia is added to a solution of magnesium phosphate in am- monium chloride, or to a mixture of magnesium sulphate, am- monium chloride, and sodium phosphate, a crystalline precipitate at once separates which is insoluble both in ammonium salts and in water containing free ammonia. This compound is a double phosphate of magnesium and ammonium and has the composition : P< ~MoNH + 6H 2 0. If heated it loses its water of crystalliza- tion and ammonia, while magnesium pyrophosphate remains be- hind, according to the equation : — 2 ^ u |oNH 4 - u }PO-0„Mg + H » u + 2JNH 3- The insolubility of this ammonium-magnesium phosphate renders it a suitable compound for quantitatively estimating both magnesium and phosphoric acid. The salt is often found in decomposing urine, and is sometimes deposited in the bladder as a constituent of urinary calculi. It is also contained in guano and occasionally in old dung-heaps, often forming large, well-developed crystals (struvite). Ammonium-Magnesium Arsenate: AsOj.~.'i,Jf + 6HjO, is Magnesium Carbonate. 417 prepared like the corresponding phosphate by mixing solutions of magnesium sulphate, ammonia, and sodium arsenate. Its pro- perties closely resemble those of the phosphate with which it is isomorphous. When heated it also leaves a residue of magnesium pyrarsenate. Magnesium Carbonate : CO-0 2 Mg. — This insoluble salt is found in nature as magnesite, partly crystallized in rhombohedra isomorphous with calcite, and partly as compact crystalline masses. Dolomite is a double carbonate of calcium and magnesium car- bonate. When a magnesium salt is precipitated with sodium car- bonate, a basic, not the normal salt is produced. This basic carbonate, or compound of normal magnesium carbonate and magnesium hydrate, has a varying composition, but may be gene- rally represented either as : 3CO-0 2 Mg+ Mg(OH) 2 or2CO-0 2 Mg + Mg(OH) 8 , with varying quantities of water. Acid sodium car- bonate is produced at the same time in which a small quantity of the magnesium carbonate remains dissolved. This white pre- cipitate is brought into trade as magtiesia alba. Magnesium carbonate forms double salts with the carbonates of the alkalies, which may be obtained by digesting magnesia alba with the acid alkaline carbonates. Magnesium Borate is deposited on boiling a solution of magne- sium sulphate with one of borax, but dissolves again on cooling. Magnesium borate in combination with magnesium chloride forms the mineral boraciie, crystallizing in the regular system. Magnesium Silicate : Si(0 2 Mg) 2 . — The normal salt of this com- position occurs in nature as the mineral olivine. The minerals talc, steatite, serpentine, and meerschaum are compounds of mag- nesium silicate with varying quantities of water. Augite, hornblende, and asbestos are double silicates of magnesium and calcium. Magnesium Chloride : MgCl 2 . — This salt is contained in sea- water and many mineral springs. It is obtained pure by dissolving magnesia alba in hydrochloric acid, adding chlorine water to oxidize the iron, and then digesting the solution with a slight excess of magnesia alba to precipitate the iron and alumina. The clear solution when evaporated down to a small bulk deposits very de- liquescent crystals of the composition : MgCl 2 + 6H 2 0. The water E E 4 1 8 Text-Book of Inorganic Chemistry. cannot be expelled from this hydrated compound by heating, since' the chloride then decomposes into magnesia and hydrochloric acid : — MgCl 2 + H.O = MgO + 2HC1, and a basic chloride is formed. This decomposition can, however, ^be prevented by the addition of ammonium chloride, as a mixture of magnesia and ammonium chloride when heated produces mag- nesium chloride, ammonia and water : — MgO + 2NH 4 C1 = MgCl 2 + 2NH3 + H„0. The concentrated aqueous solution is therefore mixed with enough ammonium chloride to form the double compound : MgCl 3 + NH 4 C1, and then thoroughly dried. The dried mass is afterwards gently heated in a Hessian crucible until vapours of ammonium chloride are no longer given off and the whole fuses quietly. The fused salt is poured on a clean slab of iron, rapidly broken into pieces and preserved in well-closed vessels. Some basic salt usually remains behind in the crucible. Anhydrous magnesium chloride is a white crystalline mass, deliquescing in the air. It can be distilled in a stream of hydrogen at a bright-red heat, and gives a colourless, soft distillate, which solidifies to a crystalline mass of brilliant plates. The salt is easily soluble both in water and alcohol. Magnesium chloride forms double salts not only with am- monium chloride, but also with potassium and other chlorides. Potassium-magnesium chloride: KC1, MgCl 2 , +6H 2 is found in considerable quantities at Stassfurt, and is known as carnallite. Detection of Magnesium Compounds. Magnesium compounds do not colour a non-luminous gas flame when introduced into it, neither do they produce any characteristic spectrum. In this respect, magnesium differs from the other metals of the alkaline earths, and resembles those of the earths. It is further distinguished from the former by the solubility of its sulphate, and by the fact that solutions of its salts are not pre- cipitated by ammonium carbonate, although magnesium carbonate is nearly as insoluble in water as the carbonates of calcium, strontium, and barium. The reason of this is that all insoluble magnesium compounds, except the double phosphate and arsenate of ammonium and magnesium, form soluble double compounds with ammonium salts. Hence a solution of magnesium sulphate Metals of the Earths. 419 to which ammonium chloride has been added gives no precipitate of magnesium hydrate with ammonia. From these reactions, magnesium salts may be just as readily detected and separated as those of the other alkaline earths. From calcium, which commonly accompanies it, it may be separated by adding ammonium chloride and ammonia to the mixed solution, and then ammonium oxalate or carbonate. All the calcium is then precipitated as calcium oxalate or carbonate, and all the magnesium is contained in the filtered solution, from which it may be thrown down with sodium phosphate as ammonium-magnesium phosphate. It is remarkable that ammonia (in the absence of ammonium salts) partly precipitates magnesia from neutral solutions, while in solutions of calcium, strontium, and barium, it produces no pre- cipitate. METALS OF THE EARTHS. The chief representative of these metals is aluminium, 1 but the group also includes a number of metals of less importance, and much less known — viz. beryllium, gallium, and indium ; yttrium, terbium, erbium and ytterbium ; cerium, lanthanum, and didy- mium ; thorium and zirconium. All these metals have a strong attraction for oxygen, but considerably less than the metals of the alkaline-earths. Their basic oxides' are sesquioxides — except those of thorium and zirconium. Notwithstanding their strong attrac- tion for oxygen, they appear to possess only a weak affinity for sul- phur, and their sulphides (except that of indium) cannot be obtained in the wet way. If ammonium or sodium sulphide is added to a solution of one of their salts, a precipitate of the hydrate, not the sulphide, is produced, and sulphuretted hydrogen is set free : — A1 3 C1 6 + 3(NH 4 ) 2 S + 6H 2 = 2 A1(0H) 3 + 6NH«C1 + 3 H 2 S. ALUMINIUM. Chemical Symbol : Al. — Atomic Weight : 27. This element is never found free in nature, but combined with oxygen and silica it is one of the chief constituents of the super- 1 The author includes chromium with these metals, but the editor prefers to place it after manganese. The editor has also made some slight alterations — with additions— in treating of the rarer of these metals. 420 Text-Book of Inorganic Chemistry. ficial layer of the crust of the earth. Felspar, granite, mica, and many other rock masses contain aluminium silicate. In combina- tion with fluorine and sodium fluoride it forms the mineral cryolite. Aluminium is a tin-white metal of specific gravity 2 -6, tough, but malleable ; it may be rolled into thin sheets and drawn into fine wire. It melts at about 700° — i.e. at a higher temperature than zinc, but lower than silver ; it is not magnetic, but conducts an electric current better than iron. The metal remains unchanged in dry as well as in moist air, becoming covered with an extremely thin layer of oxide — even at its melting point it only oxidizes slowly. If, however, a large surface is heated in the' air — e.g. a piece of aluminium foil in the lamp, it burns brilliantly. Concentrated nitric acid does not,act upon aluminium, but dilute sulphuric or hydrochloric acid dissolves it readily with copious evolution of hydrogen. It easily dissolves in a solution of caustic potash or soda, evolving hydrogen and producing soluble com- pounds of alumina and potash or soda. The molten alkalies do not act upon it, but when fused with potassium carbonate, carbon is reduced. The strong affinity of aluminium for oxygen and the halogens makes its preparation somewhat difficult. Its oxide is not reduced when glowed with charcoal, nor when heated in a stream of hydro- gen. It may, however, be easily prepared, either by the electrolysis of fused sodium-aluminium chloride : 2NaCl, AljCl,., or by decom- posing the chloride with sodium. The latter process can be carried out on a small scale by leading the vapour of aluminium chloride over fused sodium with a stream of hydrogen. After the violent reaction is over, the reduced aluminium is found under the fused sodium-aluminium chloride produced at the same time, if the sodium was not in excess. On the large scale, aluminium is manufactured in the following manner. The double chloride of sodium and aluminium is first ■prepared by heating a mixture of alumina, common salt, and char- coal in a stream of chlorine. 100 parts of the sublimed double chloride are next mixed with 35 parts of sodium and 40 parts of cryolite (to serve as a flux), and heated on the hearth of a reverbe- ratory furnace. The aluminium is then reduced, and collects on the hearth under the fused slag. The high tenacity of aluminium, its property of not rusting, like iron, in the air, and its low specific gravity, would make it an ex- tremely useful metal could it be prepared in larger quantities and Aluminium. 421 at a lower price than is at present the case. It is now little used, except for scientific purposes and for the manufacture of aluminium bronze. We .are only acquainted with one oxide of aluminium, and only with one compound with chlorine and the other halogens. Aluminium Oxide, Alumina : Al a 3 . — This compound occurs in nature in crystals belonging to the hexagonal system — partly transparent, of a blue or red colour (probably due to traces of cobalt and chromium) as the two precious stones, the sapphire and ruby — and partly less pure, as grey, usually opaque, crystals, called corundum. Crystalline alumina is distinguished by its hardness — next to the diamond it is the hardest substance known, and is, therefore, highly valued for grinding and polishing. For this pur- pose massive crystalline alumina, called emery, is used. Emery is found in some parts of Asia Minor, especially on the island of Naros, as well as in Massachusetts. The oxide may be easily prepared as a white amorphous powder by glowing either its hydrate or its double sulphate with ammonium — called ammonia alum. Like all sesquioxides, alumina is a weak base, and is not soluble in acids either in the natural or artificial form. It can only be brought into solution by fusion with acid potassium sulphate, or with alkaline hydrates or carbonates. Alumina forms salts not only with strong acids but also with strong bases, in which it plays itself the part of an acid. Some of these salts — e.g. those of potas- sium and sodium — are soluble in water, hence it follows that aluminium hydrate when separated from aluminium salts by caustic soda (or potash) dissolves in an excess of the reagent. Alumina is infusible at the highest temperatures of our furnaces, but melts in the oxy-hydrogen flame to a colourless glass, which crystallizes on cooling. By the previous addition of minute quan- tities of potassium dichromate, artificial rubies may be thus obtained. Aluminium Hydrate : Al 2 (OH) 6 ,or Al(OH) 3 . — This compound, occurring in nature as hydrargillite or gibbsite, is obtained as a white gelatinous precipitate when a soluble aluminium salt is pre- cipitated with ammonia. The precipitate, which is difficult to wash, dries to a translucent mass, and afterwards forms a light, white amorphous powder, insoluble in water. Whengently heated 422 Text-Book of Inorganic Chemistry. it loses two molecules of water and is changed into a hydrate of the composition : ,,„ .-.,_, which is also found in the mineral king- dom as diaspore. Aluminium hydrate is easily dissolved both by acids and alkalies (except ammonia). In both cases salts are produced, the alumina playing the part of base or of an acid. It possesses the property of uniting with organic colours and precipitating them from their solutions, and is, therefore, a valuable mordant. Calico and other stuffs, which do not themselves unite with the colours with which they are to be dyed, retain them firmly if impregnated with alu- minium acetate — a salt which readily deposits aluminium hydrate. Organic colours precipitated with aluminium hydrate are termed lakes. Alumina and its compounds are thus largely used in dyeing. Aluminium Salts.— Aluminium, as we have just mentioned, forms two classes of salts : those in which it plays the part of a base, and those in which it is an acid. The former are the common salts of aluminium ; the latter are called aluminates. They are found in the crystalline form in nature, and can also be prepared artificially. Two classes of aluminates are known, one in which the aluminic acid is tribasic, and one in which it is dibasic. These two acids are simply the two hydrates : Al(OH) 3 and | . ,,-. „„ The naturally occurring aluminates (often called spinelles) belong chiefly to the latter class. The following are the more important of these salts. ( AlO ■ OK Potassium Aluminate : ,,„ n] . and sodium aluminate : Al(ONa) 3 , are obtained by dissolving aluminium hydrate in caustic potash or soda, and separate out on evaporation or on the addition of alcohol. Magnesium aluminate — the mineral spinelle :,,,, 2 Mg, zinc aluminate — gahnite: ,, ,.-. j- O^Zn, and ferrous alumhiate — iron-spinelle : .,.-. [ 2 Fe, are aluminates occurring in nature. Aluminium Sulphate: (S0 2 ) 3 C A1, + l8H,0. — The normal salt of the above composition is occasionally found in nature, and is prepared by dissolving aluminium hydrate in sulphuric acid, or Aluminium Sulphate. 423 commercially by acting on clay (aluminium silicate), with concen- trated sulphuric acid. The mixture of aluminium sulphate and silicic acid obtained by the latter process is known in trade as alum-cake. The salt crystallizes from concentrated solutions in fine needles or plates, which deliquesce in the air. Besides the normal salt, numerous basic compounds are known. One of these is found in nature as the mineral aluminite, and has the composi- tion : (S0 2 ) 3 O e Al 2 + 4Al(OH) 3 + 2iH 2 0. Aluminium sulphate forms crystalline double salts with the sulphates of the alkalies, of thallium, and of silver. These double sulphates all crystallize in the regular system, and are all isomor- phous with one another : they are called alums. The longest known of these compounds is potash-alum, potassium-aluminium sulphate: |q4q|^ 1 + I2H 2 0, or KA1(S0 4 ) 2 + I2H 2 0, to which reference has been previously made (p. 336). It may, like all the other alums, be considered as a double sulphate of potas- sium and aluminium, and its formula would then be : S0 2 (OK) 2 + (SO 2 ) 3 O Al 2 + 2 4 H 2 O. Potash-alum, of the above composition, crystallizes from its aqueous solution in large, colourless, and transparent regular octahedra, often combined with faces of the cube. If its warm solution is mixed with potassium carbonate (or hydrate) so that the precipitate at first produced is redissolved, the solution, which now contains a basic salt, deposits crystals of ordinary alum in the form of cubes, not octahedra. This is called cubic alum. Alum possesses a sweet, astringent taste, has an acid reaction, and is easily soluble in water, especially warm. It is insoluble in alcohol. When the crystals are exposed to the air they become covered with a white opaque crust, which is said to be due to the formation of a basic salt from the ammonia of the air, and not to loss of water. If alum is heated it melts in its water of crystalliza- tion, and as the liquid loses water it becomes thicker and thicker and froths considerably. Finally, the anhydrous salt remains behind as a spongy mass, known in medicine as burnt alum. Alum may be prepared directly from its constituents by mixing solutions of aluminium and potassium sulphates, and evaporating down. The aluminium sulphate used for the manufacture of alum is principally obtained in two different ways — either by treating kaolin or some other clay (aluminium silicate) with concentrated 424 Text-Book of Inorganic Chemistry. sulphuric acid, or by roasting alum-shale and extracting with water. The alum-shale does not contain aluminium sulphate ready formed, but consists of a shale (essentially solidified clay) mixed with small crystals of iron pyrites, and often with some combustible sub- stances. When roasted in the air oxidation takes place, with the production of ferric oxide (or ferrous sulphate), sulphurous anhy- dride and sulphuric acid, the latter acid decomposing the alumi- nium silicate, forming aluminium sulphate. The mass is extracted with water, the requisite quantity of potassium sulphate added, and evaporated down to crystallize. Some varieties of alum-shale are so easily decomposed, and contain the iron pyrites so finely divided, that it suffices to moisten them with water and leave them exposed to the air for some time to effect the required decomposi- tion. The value of an alum for dyeing purposes is greater the less the amount of combined iron it contains, because this substance spoils the colour. Alum particularly free from iron is prepared from a naturally occurring basic aluminium sulphate, called alunite. This substance is much less widely distributed in nature than alum-shale, and is chiefly found near Rome, in Tuscany, and Hungary. The alum prepared from it is known in trade as Roman alum. Alunite, which is insoluble in water, may be con- sidered as a compound of one molecule of anhydrous alum united with two of aluminium hydrate, and is therefore represented by the formula: so^foK^ + zA1 (OH) s . When gently heated, it yields water, insoluble alumina, and soluble potash alum. The last named compound is then extracted with water, and usually crystal- lizes in cubes. If a solution of aluminium sulphate is mixed with one of sodium, ammonium, rubidium, caesium or thallium sulphate, alums are produced of the same crystalline form as potash-alum and of analogous composition. An alum of silver has also been obtained. These different alums are distinguished by the names soda-, ammonia-, &c, alum. Formerly, large quantities of ammonia- alum were manufactured in England from the ammoniacal liquors of the gas-works ; the manufacture has now almost entirely ceased, owing partly to the immense quantities of cheap potassium compounds obtained from the Stassfurt beds. Ammonia-alum cannot be distinguished from potash-alum in its external cha- racters. Ahem. 425 Besides the substitution of the potash by other monoxides, the alumina in alum may also be displaced by other sesquioxides with- out altering the shape or general character of the compounds ; and alums can be prepared in which the aluminium is displaced by iron, manganese, or chromium respectively. These alums are known by the names of the triad metals which they contain, and have the following composition : — t ) SO., jO.Fe u n Iron-alum . . . . ^j Q fc + i 2 H 2 0. Manganese-alum . . . fo'loK 1 " + I2H 2°- Chrome-alum .... sSMoK* + I2H 2°- These alums are coloured : iron-alum usually violet, chrome- alum, dark-green or purple. Like alum proper, the potassium can be displaced by other monad metals, and we are thus ac- quainted with more than twenty isomorphous compounds belong- ing to this class. Aluminium Phosphate: P0-0 3 A1 + 4H„0. — The normal salt is deposited as a white gelatinous precipitate, when a solution of sodium phosphate is added to one of aluminium sulphate or of alum. It is easily soluble both in acids and alkalies (except ammonia). Various basic phosphates are known, among which is the mineral wavellite. Aluminium Silicate : (SiO),0 6 Al, 2 . — A silicate having this composition is thrown down as bulky precipitate on adding excess of a hot dilute solution of aluminium sulphate to a neutral solu- tion of sodium silicate. A salt of the composition : Al 2 3 ,Si0 2 is found crystallized in the mineral kingdom as andalusite and disthene. In combination with other silicates, aluminium silicate is one of the most important constituents of the crust of the earth. Among these may be mentioned the felspars of which the com- position has been already given (p. 272). The commonest of these is potash-felspar or orthoclase : KAlSi 3 O g , less common are soda- felspar or albite, of similar composition, and lime-felspar or anor- thite. Most of the naturally occurring double silicates are anhy- drous but some — the zeolites — contain water. To this class belong the minerals natrolite and analcime, which are hydrated 426 Text-Book of Inorganic Chemistry. double silicates of aluminium and sodium. They belong to the few silicates which are decomposed by hydrochloric acid. The felspars and most of the other silicates are quite unacted upon by hydrochloric acid, and can only be decomposed by treatment with hydrofluoric acid, or by fusion with sodium (or potassium) carbonate. One of the most important hydrated silicates is clay, which in its purest form is kaolin or china-clay. This is hydrated alu- minium silicate of the approximate composition : Al 2 O s , 2Si0 2 + 2H„0. It is produced by the gradual weathering of felspar, which is often found mixed with it. The process consists in a gradual change of the potassium silicate by the carbonic acid of the air into soluble potassium carbonate (see p. 344), leaving insoluble alu- minium silicate, which then unites chemically with water. Kaolin is largely used for the manufacture of porcelain and earthenware. Clay when finely powdered and mixed with water forms a plastic mass which readily allows itself to be moulded into various shapes, and if the clay is afterwards heated, it bakes together to a hard but porous mass. It is because of these two properties that clay is so largely used for the manufacture of articles of china, stoneware, and earthenware. But vessels prepared from clay alone are porous, and cannot be used to contain liquids : to make them available for this purpose they are either coated with a thin layer of glass (the glaze), or else the clay is previously mixed with some fusible material (the frit), then burnt and glazed. In the manufacture of true porcelain the frit employed is usually either felspar or a mixture of felspar and quartz, which is carefully mixed with the kaolin and then burnt at a high temperature. In true porcelain, the frit penetrates through the entile mass— giving it its well-known translucent appearance. Porcelain is usually glazed by dipping it in finely powdered felspar suspended in water, and then again heating. Stoneware of very various qualities is a coarser kind of porce- lain made with a more easily fusible frit which does not penetrate through the entire mass of the clay. Stoneware is glazed with a more fusible glaze than felspar. Earthenware or Fayence, which is made from variously coloured clays, is porous, and of a yellow or reddish colour. It is glazed either with lead oxide, forming an easily fusible lead silicate, or with common salt. The salt is usually thrown into the kiln Ultramarine. 427 when the burning is nearly over. It then volatilizes, and is de- composed with the water-vapour present and the silica of the clay, forming hydrochloric acid and a sodium silicate, which acts as a glaze. A clay rich in calcium carbonate, and which is sometimes mixed with clay in making earthenware, is known as marl. Intimate mixtures of clay with ferric oxide of various colours from yellow to brown are used as pigments under the name oi ochres. Fuller's earth and bole are fine clays used for cleansing clothes and other purposes. Ultramarine. — The highly valued ornamental stone called lapis lazuli, and occurring especially in China and Thibet, was previously powdered and sold under the name of ultramarine for its weight in gold. This substance is essentially a compound of sodium and aluminium silicates with sodium pentasulphide. Fifty years ago Gmelin discovered a method for the artificial manufac- ture of ultramarine, and since then his processes have been so perfected, that the artificial substance can now be made of a finer colour than that found naturally. This discovery has, therefore, been well described as one of the most brilliant victories of chemical science during the present century. The artificial preparation of ultramarine appears at first ex- tremely simple, but in order to prepare a pigment of fine colour a number of practical details must be attended to. These can only be acquired by long experience, and are often preserved as trade secrets. Ultramarine is obtained by heating clay in a closed crucible with some mixture which produces sodium sulphide, such as sulphur and sodium carbonate, or sodium sulphate and charcoal. The product first produced is green — ultramarine-green — but if mixed with sulphur and again gently heated, changes to the well- known blue colour of ultramarine proper. The change in colour is probably due to the formation of a polysulphide of sodium. If heated too strongly, the colour is destroyed. The chemical constitution of ultramarine is still uncertain. It is thought that the blue compound contains sodium pentasulphide, and the green the disulphide. The sodium sulphide cannot be dis- placed by potassium sulphide. The products are finely ground, brought into trade in various tints, and largely used for many purposes. 428 Text-Book .of Inorganic Chemistry Ultramarine is not so permanent as smalt (cobalt blue), nor as indigo, but is not decomposed by soaps and alkalies like Prussian blue. Acids, however, at once decompose ultramarine, and it may be recognized by the reaction that when mixed with hydrochloric acid, sulphuretted hydrogen is evolved, much white sulphur sepa- rates — from the sodium pentasulphide — and the blue colour en- tirely disappears. Aluminium chloride : A1 2 C1,.. — This compound may be ob- tained by burning the metal in a stream of chlorine. It is, however, usually prepared by intimately mixing alumina and powdered charcoal (best sugar charcoal ') with starch paste, and then proceeding as described under the preparation of silicon tetrachloride (see p. 273). On glowing in a stream of dry chlorine, aluminium chloride sublimes in yellowish crystalline plates. The yellow colour is due to traces of ferric chloride, which may be removed by sublimation over hot aluminium, when the volatile ferric chloride is converted into non-volatile ferrous chloride. Pure crystalline aluminium chloride is colourless, sublimes when gently heated, and rapidly attracts water from the air. The aqueous solution of the chloride, which may be obtained by dissolving aluminium hydrate in hydrochloric acid, when evaporated down deposits colourless deliquescent crystals of the composition: AI 2 Cl e + I2H,0. The water cannot be expelled from these crystals by heating so as to leave the anhydrous chloride, because the aluminium chloride' is then decomposed with the formation of hydrochloric acid, and a basic chloride or even alumina. Aluminium chloride unites with sodium and other chlorides, producing fairly stable double salts. Aluminium Fluoride : A1 2 F„, is obtained by dissolving alumi- nium hydrate in hydrofluoric acid, or by glowing a mixture of alumina and fluor-spar in hydrochloric acid gas. It is insoluble in water, sublimes at a high temperature in colourless cubes, and in combination with sodium fluoride forms the mineral cryolite : Aluminium Sulphide : A1.,S 3 , is produced when alumina is glowed in the vapour of carbon disulphide. It is a bright yellow vitreous mass, which with water, or even moist air, rapidly decom- poses into sulphuretted hydrogen and aluminium hydrate. 1 Or a very pure form of lamp-black, known as gas-black. — Ed. Beryllium. 429 Detection of Aluminium Compounds. Aluminium salts are distinguished from those of the alkalies and alkaline earths by the production of a white gelatinous pre- cipitate of aluminium hydrate with ammonia — a precipitate which is insoluble in ammonium salts and is therefore produced in their presence (comp. magnesium). They are further distinguished by yielding a white precipitate of aluminium hydrate, not sulphide, with ammonium sulphide. Like all sesquioxides, alumina is a weak base, and does not form compounds with weak acids — e.g. carbonic acid. If a solu- tion of aluminium chloride is mixed with one of sodium carbonate, aluminium hydrate and sodium chloride are formed, and carbonic acid set free. If the sodium carbonate is in excess, acid sodium carbonate is produced. The weak basic properties of alumina are further shown by its behaviour with caustic soda (or potash). Caustic soda precipitates aluminium hydrate from solutions of aluminium salts, but the precipitate dissolves in an excess of the alkali, forming sodium aluminate — a salt in which alumina plays the part of an acid (p. 423). Aluminium compounds when heated on charcoal before the blowpipe leave a white infusible mass, which, when moistened with a drop of cobalt nitrate and again heated becomes of a fine blue colour. BERYLLIUM. Chemical Symbol : Be. — Atomic Weight: 13-6. This somewhat rare element is chiefly found in nature as a double silicate of aluminium and beryllium : Al^C^, Be 2 3 , 6Si0 2 , or (SiO) 3 \r\-a wri ' cn ) m ' ts rou gh opaque form, is called beryl; when purer and transparent it is either of a pale sea-green colour : aquamarine, or else of a bright green : emerald, both of which varieties have considerable value as precious stones. It further occurs as silicate simply in the mineral phenakite : 2Be 2 O s , 3SiO„, or Si 3 (0 3 Be) 4 , and as double oxide of aluminium in chrysoberyl : Be 2 3 , A1 2 3 . The metal may be obtained by leading the vapour of the 43Q Text-Book of Inorganic Chemistry. chloride over fused sodium contained in iron boats, and afterwards washing with water. It is a white metal like aluminium, with a specific gravity of 17, and melts at a red-heat under partial oxida- tion. It is without action on water at the ordinary temperature, but decomposes it slowly on boiling. Dilute hydrochloric or sulphuric acid easily dissolves it with rapid evolution of hydrogen. Concentrated nitric acid scarcely acts on it. It dissolves also in the caustic alkalies (not in ammonia), forming soluble compounds resembling the aluminates. Beryllium Oxide or Beryllia : Be 2 3 , is a white, amorphous powder of 3'o specific gravity, obtained by heating Beryllium Hydrate: Be 2 (OH) 6 . — This compound is thrown down as a white gelatinous precipitate when ammonia is added to a solution of a beryllium salt. The precipitate forms a white powder when washed and dried ; it is soluble in caustic potash or soda and in ammonium carbonate, from which solution it is again precipitated on boiling as a carbonate. The salts of beryllium possess a pure sweet taste, whence the name gli/cinum, which is sometimes given to the metal ; they are mostly soluble in water. Beryllium Sulphate : (S0 2 ) 3 6 Be. 2 + I2H. 2 0, easily crystallizes when its aqueous solution is evaporated down. It forms a double compound with potassium sulphate, which is difficult to obtain in the crystalline form, and has a different composition to common alum, with which it is therefore not isomorphous. Beryllium Chloride : Be,Cl 6 , is prepared like aluminium chloride, which it resembles, but is less volatile. It forms, when pure, white, crystalline, highly deliquescent needles. Beryllium Carbonate. — On adding sodium carbonate to a solu- tion of a beryllium salt a precipitate of a basic carbonate is pro- duced, which dissolves with difficulty in an excess of the reagent. This carbonate is much more easily soluble in ammonium carbo- nate, a fact which enables us to detect beryllium and to separate it from aluminium. Water containing carbonic acid also slightly dissolves the basic carbonate. From the above properties of beryllium and its compounds it will be seen that the metal occupies a peculiar position between magnesium and aluminium, both of which it resembles in some Gallium. 43 1 points. And it has long been a matter of discussion whether the element should be classed with the alkaline-earths or with the earths. In the former case it would be a dyad with an atomic weight of 9-1, and its oxide would be written : BeO ; in the latter the element is a triad, its atomic weight is half as much again, or 13-6, and its oxide becomes : Be 2 3 . Recently the specific heat of the metal has been determined and found to be 0-445, which when multiplied by the atomic weight (i3'6) gives 6-05, proving, according to Dulong and Petit's law (p. 70), that this is the true atomic weight and not 9/1. GALLIUM. Chemical Symbol : Ga. — Atomic Weight : 69-8. This element is contained in some samples of blende found in the Pyrenees, but only in extremely minute quantities (about 0-002 per cent.) Its spectrum contains two characteristic lines in the violet which led to its discovery by the French chemist, Lecoq de Boisbaudran in 1875. Gallium is a white lustrous metal of low melting-point (30°) ; when once melted it remains liquid, like mercury, even at o°. Its specific gravity is 5-9. The metal remains unaltered in the air, and only becomes covered with a thin layer of oxide when heated nearly to redness. It is not attacked by water at the ordinary temperature, but dilute hydrochloric acid as well as the alkalies dissolve it with evolution of hydrogen. The compounds of gallium resemble those of aluminium. Gal- lium oxide : Ga 2 O s and the hydrate are white amorphous powders, insoluble in water. Gallium chloride : Ga 2 Cl 6 , which may be ob- tained by heating the metal in a stream of dry chlorine, is soluble in water and easily sublimes. A dichloride : GaCl 2 is also known. Ammonium-gallium alum: cr\\ ONH + 12 ^fi> crystallizes in regular octahedra. 432 Text-Book of Inorganic Chemist ty INDIUM. Chemical Symbol : In. — Atomic Weight: 113-4. Indium, like gallium, belongs to the rarest metals ; it is also found in very small quantities in some zinc ores, and when these are worked for zinc alloys itself with this metal. The zinc from Freiberg in Saxony contains from o'05 to cri per cent, of indium. If such zinc is treated with insufficient hydrochloric acid to com- pletely dissolve it, the indium remains behind with other metals,, and is purified by different processes. Indium is a white lustrous metal, of specific gravity 7-4, and melting at 176 . It dissolves slowly in hydrochloric or dilute sul- phuric acid, but easily in nitric acid. The metal remains unchanged in the air, but burns, when heated to redness, with a violet light, and producing brown vapours. Indium is a triad in all its com- pounds. Indium Oxide : ln.,0 3 , is a bright yellow powder, which becomes dark brown when heated. Indium Hydrate: In(OH) 3 , is thrown down from indium salts on addition of ammonia, as a white gelatinous precipitate, resem- bling aluminium hydrate ; it also dissolves in caustic soda or potash. Indium Sulphate : (SO a ) 3 8 In„, remains as a gummy mass when its solution is evaporated down. It unites with ammonium sulphate, but not with potassium sulphate, and forms ammonium- indium alum: en iONH + I2 H 2 0. Indium Chloride : InCl 3 , is prepared by heating either the metal or a mixture of the oxide and charcoal in a stream of dry chlorine, when it sublimes in white crystalline plates. It is very deliques- cent. Indium Sulphide : In 3 S 3 , is produced when sodium and sulphur are gently heated together, and then forms a brown infusible mass. Sulphuretted hydrogen precipitates the same compound from neu- tral solutions of indium salts, or when sodium acetate has been added to the solution. In the presence of free hydrochloric acid no precipitate is obtained. Indium, which is otherwise closely allied to aluminium and gallium, in this respect resembles zinc. Traces of an indium compound, if introduced in the non- luminous gas flame, may be at once detected by its spectrum, Yttrium, Terbium, Erbium, and Ytterbium, 433 which is characterized by one line in the dark blue. It was this line which led to the discovery of the metal, and from which it derives its name. YTTRIUM, TERBIUM, ERBIUM, and YTTERBIUM. Chemical Symbols : Y, Tr, Er, Yb. Atomic Weights : 90, 148, 166, 173. These four rare elements, which have not yet been prepared in> the free state, are only found in a few extremely scarce minerals — e.g. in gadolinite as silicates, in ytlrotantalite and samarskite as tantalates and niobates. Of their oxides, yttrium and ytterbium oxides : Y„0 3 and Yb 2 3 , are white, terbium oxide : Tr 2 3 , is yellow, and erbium oxide : Er 2 3 , pink. The salts of all four so closely resemble one another that they can only be partially separated with considerable difficulty. The sulphates and nitrates are soluble in water, the former compounds being more soluble in cold than hot water. The chlorides are non- volatile, and can be prepared anhydrous like magnesium chloride. Finally, the oxalates and carbonates are insoluble, so that they can be separated from aluminium and iron (in the ferric state) by pre- cipitating with oxalic acid. The salts of erbium are of a pink colour, and yield a very charac- teristic absorption spectrum, by means of which small quantities of the element may be detected. CERIUM, LANTHANUM, and DIDYMIUM. Chemical Symbols : Ce, La, Di. Atomic Weights : 141, 139, 146. These three elements, which may be prepared by the electro- lysis of their fused chlorides, always occur associated with one another in nature— usually as a hydrated silicate : cerite, which is decomposed by hydrochloric or sulphuric acid. Cerium forms two oxides — viz. a sesquioxide : Ce. 2 3 , and a per- oxide : Ce0 2 . Both are white compounds, but the latter becomes F F 434 Text-Book of Inorganic Chemistry. yellow when heated in the air. The two other oxides, lanthanum oxide : La,0 3 , and didymium oxide : Di 2 3 , correspond to the ses- quioxide of cerium. Didymium also forms a brown peroxide. The salts of the three metals resemble one another so closely that it is difficult to separate them. The cerium may be got rid of as the peroxide when the alkaline solution containing cerium hydrate in suspension is acted on by chlorine, but the separation of lantha- num and didymium is much more difficult. The sulphates, nitrates, and chlorides closely resemble those of the preceding group. The oxalates are also insoluble in water and dilute acids. The sulphates are distinguished by forming a double sulphate with potassium (not an alum), which is quite insoluble in a saturated solution of potassium sulphate, and so enables us to separate this group of metals from the preceding group. Didymium compounds are pink, like those of erbium, but with a bluer tinge. They also yield a characteristic absorption spec- trum, of which the darkest band is one in the yellow. THORIUM. Chemical Symbol : Th. — Atomic Weight : 232-5. This extremely rare element is only found in two scarce minerals, both of which are silicates — the one anhydrous, called iliorite, and the other hydrated, orangite. The metal can be extracted by acting on its volatile chloride : ThCl 4 , with sodium. The only oxide has the composition : Th0 2 . Its sulphate, chloride, and nitrate are all soluble in water. ZIRCONIUM. Chemical Symbol: Zr. — Atomic Weight: 90. Zirconium is also one of the rarer elements, and chiefly found in its silicate, the minerals zircon and hyacinth : Si(0 4 Zr). The metal is obtained in brilliant brittle scales, by strongly heating potassium fluozirconate.with aluminium in a graphite crucible, and Zirconium. 43 5 afterwards extracting the excess of aluminium with hydrochloric acid. It has a specific gravity of 4-1, is very difficultly fusible, and is not attacked by oxygen at a red heat, but burns in the oxy-hydro- gen flame. Hydrochloric, nitric, or sulphuric acid scarcely attacks it, but it dissolves in hydrofluoric acid with evolution of hydrogen. The mineral zircon is not attacked by any acid, but is decom- posed by fusing it either with sodium carbonate, or, better, with acid potassium fluoride ; in the latter case potassium fluosilicate and fluozirconate are formed, of which the latter dissolves when the fused mass is boiled with water, while the former remains undis- solved. From this solution ammonia precipitates Zirconium Hydrate : Zr(OH) 4 , as a white voluminous mass, which loses water when heated, and leaves zirconium oxide or zirconia : ZrO„ as a white power, insoluble in dilute acids. Caustic soda also precipitates the hydrate like ammonia, which is soluble in an excess of this reagent like aluminium and beryllium hydrates. Zirconium Sulpliate : en". 01 ^ r > obtained by dissolving the hydrate in sulphuric acid, crystallizes with difficulty from its aqueous solution. Zirconium Chloride : ZrCl 4 , prepared like aluminium chloride, by glowing a mixture of zirconia and charcoal in dry chlorine, is a white crystalline substance, volatile without decomposition, and soluble in water. The same solution is obtained by dissolving zirconium hydrate in hydrochloric acid, from which, however, an oxy chloride : ZrOCl 2 , with varying quantities of water, crystallizes in colourless silky prisms on concentration. Zirconium fluoride : ZrF 4 , sublimes as a colourless crystalline mass when an intimate mixture of powdered zircon and fluor-spar is heated in a stream of hydrochloric acid. It is insoluble in water and hydrochloric acid, but dissolves in hydrofluoric acid. On mixing this solution with potassium fluoride and evaporating down, crystals of potassium fluozirconate: ZrF t + 2KF, or K 2 ZrF 6 are obtained, isomorphous with the corresponding compounds of silicon, titanium, and tin. Zirconium (like thorium) is a tetrad in nearly all its com- pounds ; it is connected with silicon and titanium on the one hand, and with tin on the other, as is readily seen from the above description of its compounds. 436 Text-Book of Inorganic Chemistry. HEAVY METALS. IRON. Chemical Symbol : Fe. — Atomic Weight : 56. Iron belongs to the most widely distributed, and to the most important elements. It is found in the free state in meteorites — larger or smaller masses of iron which occasionally reach the earth from extra-terrestrial space. These meteorites always contain nickel, often as much as 10 per cent, and may be distinguished from metallic iron of other origin by peculiar markings (first noticed by Widmanstadt) which are produced when the surface is polished and treated with nitric acid. The most important compounds of iron, and those used ex- clusively for the extraction of the metal, are its oxygen compounds. They are not only very widely distributed, but are often found in immense quantities. Among these are : specular-iron, ferric oxide : Fe 2 3 , beautifully crystallized in the hexagonal system, and espe- cially abundant on the Island of Elba. Ferric oxide, as granular, columnar, or earthy masses forms the important ore, red haematite, which when combined with water is known as brown haematite or limonite. Red haematite is principally found in Lancashire and Cumberland in England, while brown haematite is found either associated with carboniferous rocks in the Forest of Dean and Glamorganshire, or else with oolitic rocks, as the earthy haematite of Northamptonshire and Lincolnshire. The bog iron-ores of Ireland and North Germany also consist of limonite, but of a much more recent date. Magnetic iron ore : Fe 3 0„ when pure, is the richest, and one of the most valuable of iron ores. It is found crystalline, but usually in rough masses, or as sand, and is the chief ore from which the Swedish iron is extracted. Spathic iron ore or siderite is impure ferrous carbonate : CO • 0,Fe, and is found in small quantities as such in England, but chiefly occurs in some German localities, and in Styria. Spathic iron ore is gene- rally found in England, intimately associated with clay as clay iron- stone, or argillaceous iron ore — one of the most important, if not Iron. 437 the most important English ores of iron. It occurs in large quan- tities in various parts of the kingdom. Other compounds of iron which are found in large quantities in nature but which cannot be used as iron ores are : Iron pyrites : FeS 2 , distinguished by its bright yellow colour and the readiness Fig. 6 5 . with which it crystallizes ; magnetic pyrites : Fe 3 S 4 or Fe 7 S 8 , richer in iron ; copper pyrites : CuFeS., ; arsenical pyrites : FeAsS, &c. Iron is extracted from its ores by reduction with carbon at a high temperature. The furnaces used for iron smelting, called blast-furnaces, are usually built in the form represented in fig. 65. They are of strong masonry, lined with fire-bricks, and reach a height of sixty or eighty feet.- The iron ore, which, if clay iron ore, 438 Text-Book of Inorganic Chemistry. must first be roasted to expel carbonic acid and water, is introduced at the top with alternate layers of coke, or in some countries (Sweden) charcoal. And if the ore does not contain the requisite constituents to form a fusible slag, either limestone or silica in some form must be added for this purpose. At the lower part of the furnace a blast of air is blown in from several iron openings, called tuyeres, which are kept at a low temperature by a stream of cold water circulating through them. Above the blast about D in the figure, the iron is reduced to a spongy mass, partly by the hot carbonic oxide, and partly by the free carbon. This reduced iron takes up carbon, &c, in its descent, becomes more fusible, and finally forms a liquid mass on the hearth at G. At the same time the silica, lime, alumina, &c, unite to form a fusible slag (p. 404), which is lighter than the iron, and by floating on its surface pro- tects it from the oxidizing action of the blast. When the furnace is at work, a continuous stream of molten slag flows out by a side opening — over a raised portion — the dam-plate (not shown in the figure). The inflammable gases which escape from the top of the furnace consist chiefly of carbonic oxide, produced by the reduction of the carbonic acid, and of nitrogen. In modern furnaces these gases are not allowed to burn in this wasteful manner at the mouth of the furnace, but pass out at a side opening near the top, and are then used to heat the blast of air entering at the base of the furnace. A large economy of fuel is effected by using a hot instead of a cold blast. In such furnaces, the top is closed with a valve arrangement {cup and cone) through which the materials are introduced when necessary. At stated intervals, when sufficient molten iron has collected on the hearth G, a hole at the base, which is closed with sand or clay, is opened with a long iron rod, and the iron flows out into suitable moulds, pro- ducing rough bars — called pigs. Such a furnace is kept in con- tinuous action for several years. Cast-iron obtained in this way is still very impure ; it may be purified from the slag which it contains mechanically by re-melting with a suitable flux. Even then it contains small quantities of sulphur, phosphorus, and of silicon reduced from the slag, as well as up to four per cent, of carbon, partly as graphite and partly in chemical combination with the iron. Two. varieties of cast-iron are distinguished — viz. white cast-iron and grey cast-iron. The former is generally produced when the temperature of the furnace is low, and when the metal is rapidly cooled. It contains nearly Iron. 439 the whole of its carbon combined with the iron, is silver white in colour, whiter than pure iron, and with a lamellar, crystalline fracture. The latter variety — grey cast-iron — is produced with a high temperature and slow cooling, and contains most of its carbon in the state of graphite. Grey cast-iron is darker in colour than white iron, and has a granular crystalline fracture. Mottled cast-iron is an intermediate variety between these two forms. When cast-iron is dissolved in dilute hydrochloric acid, the carbon present as graphite remains behind, while that in chemical com- bination with the iron combines with the nascent hydrogen and forms volatile hydrocarbons. Spiegeleisen and ferro-manganese are varieties of cast-iron containing manganese, and are largely used for the manufacture of steel. Cast-iron melts at a lower temperature than any other form of the metal, but is brittle, and cannot be worked up into any shape. It is, therefore, only used for those articles which can be cast in moulds. But as the percentage of carbon diminishes the malle- ability increases, and the iron becomes more difficultly fusible. Iron containing the smallest quantity of carbon (not more than o'j per cent.) is called wrought-iron, while intermediate between wrought- and cast-iron are the various kinds of steel. In order, therefore, to convert cast-iron into wrought-iron it is necessary to abstract the greater part of the carbon which it contains. This is done by two processes, both depending on the same principle, called puddling and fining. In the former process, which is most generally used in England, the crude pig-iron is melted with slags rich in ferric oxide on thje hearth of a reverberatory furnace with side openings. A considerable quantity of its silicon and carbon is re- moved by contact with the hot air blowing over it, and by the oxy- gen of the slag. Long iron bars are then inserted by the workmen at the side doors, and the molten metal stirred and worked up so as to expose it still further to the oxidizing action of the air. As the process goes on, the iron gradually becomes more difficultly fusible, and at last forms a tough mass, in which state it is removed from the furnace and hammered to remove the fusible slag which it still contains. It is then further hammered or rolled into bars or plates. In the process oi fining, the cast-iron is first melted on a hearth, and then a blast of air blown through it from several tuyeres. A portion of the iron is oxidized to the magnetic oxide, and nearly the whole of the carbon and silicon is removed. After the mass 440 Text-Book of Inorganic Chemistry. has become pasty it is worked, like puddled iron, by hammering and rolling. Wrought-iron prepared in either of these ways is never quite free from carbon, but only contains from o'i to o - 5 per cent. Traces of silicon are also usually present, and, what is of more importance, traces of sulphur and phosphorus. Iron containing even a small quantity of sulphur is brittle when hot, or is said to be red-short, while that containing phosphorus to the extent of more than a half per cent, is brittle when cold, or, as it is said, is cold-short. Wrought-iron has a pale grey colour, is very tenacious, ductile, and malleable. Its specific gravity is about 7'8. At a red-heat it be- comes pasty, and can then be readily welded together and fashioned into various shapes. The smith, when he wishes to weld together two pieces of iron, makes them both red-hot in the forge, sprinkles them with sand, and then unites them into one solid piece by hammering. Without the addition of sand or some such substance (borax), the two surfaces of iron would be unable to unite with one another, because of the layer of oxide with which they become coated when heated. The silica of the sand combines with the iron oxide, forming an easily fusible slag, which is removed by hammering, leaving the two surfaces of bare metal, which unite together under the pressure of the blows. Wrought-iron, during its process of manufacture, acquires a fibrous structure, which greatly increases its resistance to breaking or tearing. And it is because of this property that wrought-iron is used for so many purposes for which brittle cast-iron, with its crystalline, granular structure, is unsuitable. An axle of a wheel, for example, if made of cast-iron, would break after a few turns on a rough road, but one of wrought-iron stands for years ; and if it occasionally happens that an axle of a railway carriage breaks, it is always found that the fibrous wrought-iron has gradually become granular and crystalline by the long-continued vibrations. Steel is a modification of iron intermediate between cast- and wrought-iron both in its properties and composition, as far as carbon is concerned. Steel contains about i 'o per cent, of combined carbon; it unites the hardness and fusibility of cast-iron with the ductility and malleability of wrought-iron. Steel can be produced either from wrought- or cast-iron. The old, and up to twenty-five years ago the only known, method of obtaining steel was from wrought- iron, by imbedding bars of the metal in charcoal-powder contained in closed iron boxes, and then heating to bright redness for about Iron. 441 a week. By this process carbon is gradually taken up by the iron in some unknown manner, and the latter becomes completely con- verted into steel. The process is called cementation, and the product cement-steel. Cement-steel can be rendered more uniform in struc- ture by working under the hammer at a red-heat, or by fusing and casting into blocks. It then forms what is called cast-steel. Steel can now be much more economically obtained directly from cast-iron. Two processes are chiefly employed to effect this change — known as the Bessemer and Siemens-Martin processes respectively. In the Bessemer process five to ten tons of cast-iron are fused and poured into the converter — a large egg-shaped vessel lined with fire-clay, with a narrow mouth and with an opening at the base through which air can be blown. A current of air is then blown through the molten metal, which oxidizes nearly the whole of the carbon and silicon. As soon as this point is reached the blast is stopped, a small quantity of spiegeleisen is added to impart car- bon and manganese, and the metal at once run into moulds. The cast-iron employed to manufacture Bessemer steel must be free from sulphur and phosphorus if a converter with the usual lining of 'Siliceous fire-clay is used, as these impurities are not removed during the process. It has, however, been recently discovered that if the lining of the converter contains free lime, this substance unites with the sulphur and phosphorus and a steel free from these impurities is obtained, even when an impure cast-iron is used. Steel is also made from cast-iron by the Siemens-Martin process, which consists in fusing cast-iron in a specially constructed furnace, and adding wrought-iron to it until the requisite composition has been obtained. Some spiegeleisen is then added as in the Bessemer process, and the steel cast into moulds. Steel is of a pale grey colour and takes a high polish. It is more easily fusible than wrought-iron, and its structure is finely granular, not fibrous, in consequence of which it is more brittle than wrought- iron. Steel is especially distinguished from wrought-iron by its hardness, which it especially acquires when heated to redness and rapidly cooled in water; it is then also exceedingly brittle and highly elastic. Glass-hard steel— i.e. steel hardened in this way, has but few uses ; in order to adapt the steel to the wants of daily life it must be toughened or tempered. This is done by again heating and again rapidly cooling, and the higher the temperature of the second heating up to a certain limit, the tougher and less hard the steel. If the steel is polished it gradually becomes 442 Text-Book of Inorganic Chemistry. covered with a thin coating of oxide, the colour of which varies from a pale yellow to a dark blue, according to the temperature (240° to 300 ) to which it has been heated. Cutting tools of various kinds are tempered in this way according to the purpose for which they are to be used. The specific gravity of steel is less after hardening (7-6) than it is before (yS). The preparation of chemically pure iron in the compact state is exceedingly difficult, if not impossible. If pure iron is fused in a crucible of gas-carbon it combines with some of the carbon ; if in one of porcelain, it reduces small quantities of silicon and aluminium with which it unites chemically ; and approximately pure iron can only be obtained by fusing in a lime crucible with the oxy-hydrogen blowpipe. Pure iron in the form of a powder may be obtained by reducing pure ferric oxide or ferrous chloride in a stream of hydrogen. It then remains as a dark grey powder {reduced iron), which is so finely divided that it takes fire in the air when it has not been strongly heated in the stream of hydrogen. Iron in all its forms is strongly attracted by the magnet : we say it is magnetic. Its property in this respect far exceeds that of any substance, and is only feebly approached by the magnetic power of the metals nickel and cobalt. Soft iron becomes mag- netic when brought near a magnet, but loses its magnetism as soon as the magnet is removed. Steel, on the other hand, does not so readily become magnetised, but retains its magnetism for some time. Artificial magnets are made of straight or curved steel bars. A steel magnet gradually loses its magnetism when warmed, but partially regains it on cooling. If, however, the magnet is heated to redness it permanently loses its magnetism. Notwithstanding the strong attraction of iron for oxygen, the metal remains unchanged in dry air and oxygen, but moist air, and especially water containing dissolved air, gradually oxidizes it — we say, the iron rusts. A drop of water placed on a polished strip of iron and allowed to evaporate leaves behind a dull reddish stain of oxide — iron-rust. This rust is not the oxide but the hydrate, and is produced by the simultaneous action of the oxygen and the water. Ferric hydrate produced in this way can unite chemically with ammonia ; hence it follows that iron-rust which has been long exposed to the air gives off ammonia when heated. Many metals which, like zinc, rust in the air, are protected from further oxidation by the compact layer of oxide which is first pro- Iron. 443 duced. But in the case of iron the layer of rust is less compact, and the process goes on, comparatively quickly, until the whole of ths metal is oxidized. For this reason sheet iron cannot be used to protect the roofs of houses, but zinc or lead must be employed instead. Iron may be protected from rusting by coating it with a layer of some metal which oxidizes slowly in the air, such as tin {tin-plate), zinc {galvanized iron), or lead {terne-plate). The process in all cases consists in dipping the clean iron plates into the molten metal, and has nothing to do with electricity. Recently a process has been introduced of rendering iron articles proof against rust by coating them with a layer of the magnetic oxide. For this pur- pose, the articles are heated to redness in a current of steam, a por- tion of which oxidizes the surface of the iron while hydrogen is set free. It has also been found that iron does not rust in water which has been made feebly alkaline by sodium carbonate. Iron is the most valuable of all metals ; it is more valuable than gold, because its properties in its varied forms as cast-iron, wrought- iron, and steel, make it applicable to so many useful purposes for which no other metal can be used. And it has been well said that without iron there would be no civilization. The metal would be still more valuable if some simple means were known by which it could be converted into a passive state, so that, like platinum, it should remain unchanged in the air and in water containing air. Such a discovery would completely revolutionize many branches of industry. We can, in fact, actually produce this change in the properties of the metal. A rod of iron dipped into red fuming nitric acid becomes thus passive, and now no longer dissolves in ordinary nitric acid, nor rusts in the air. But if touched or scratched with a piece of ordinary iron it loses these peculiar properties at once. The problem of how to make this passive state permanent remains as yet unsolved. At a red-heat iron is more rapidly oxidized than by moist air at the ordinary temperature. A red-hot rod of iron becomes covered with a coating of oxide, which falls off when the iron is hammered on the anvil by the smith. This scale, as it is called, consists essentially of the magnetic oxide : Fe 3 4 , and forms black, .dull scales when cold. The particles of iron which fly off and burn as sparks when a flint and steel are struck together, also form the same oxide, and the fused nodules produced when steel burns in oxygen (p. 14) have the same composition. 444 Text-Book of Inorganic Chemistry. Red-hot iron absorbs considerable quantities of carbonic oxide, which it parts with to the air when in contact with it. This happens, for example, when iron stoves are heated to redness. And the injurious action of such stoves when too strongly heated is largely due to the carbonic oxide which then diffuses into the air of the room. The iron filings produced by the mechanic in working up iron articles are largely used for a number of purposes, owing to the larger surface which they expose to the action of various liquids than larger pieces of the metal. Many compounds of iron are prepared from them, and considerable quantities are employed for the commercial reduction of nitrobenzol to aniline in the presence of acetic acid. Iron in the state of a fine powder {reduced iron) is produced, as previously stated, by reducing ferric oxide in a stream of hydrogen. Iron is attacked and dissolved by nearly all acids ; dilute sulphuric, hydrochloric, or acetic acid dissolve it, with evolution of hydrogen, producing ferrous sulphate, chloride, or acetate respec- tively. Ordinary concentrated nitric acid dissolves iron, with the formation of ferric nitrate and lower oxides of nitrogen. Very dilute nitric acid dissolves it, like sulphuric acid ; the hydrogen which is produced by the decomposition of the water is not, how- ever, set free, but in its nascent state gradually reduces some of the nitric acid to ammonia. That the solution contains ammonium nitrate as well as ferrous nitrate, may be readily proved by the ammonia evolved when it is mixed with caustic soda. Iron occurs in nearly all its compounds as a dyad and triad (or tetrad) element. It is a dyad in the ferrous compounds, ferrous oxide : FeO, ferrous chloride : FeCI,, &c, and a triad (or tetrad) n the ferric compounds — ferric oxide : Fe„0 3 , ferric chloride : Fe 2 Cl 6 , &c. Ferrous oxide and ferric oxide both possess basic properties ; the latter, like all sesquioxides, is a weak base — weaker than ferrous oxide. It is, however, more stable than ferrous oxide, as it is not changed by the oxygen of the air. A third oxide of iron is also known, containing more oxygen than ferric oxide, and having the properties of an acid. This ferric acid, probably Fe0 2 (OH) 2 , appears only to exist in union with strong bases, and probably contains hexad iron, like sulphuric acid contains hexad sulphur. Ferrous Oxide. 445 Terrous Oxide : FeO. — On account of the strong attraction of this compound for oxygen it is difficult to prepare, and has never been obtained pure. It is produced, with some metallic iron, when ferrous oxalate is gently heated without access of air, or as a black, glistening powder when the mixture of carbonic oxide and carbonic acid obtained by heating oxalic acid with sulphuric acid (p. 295) is led over hot ferric oxide. The substance so obtained easily dissolves in hydrochloric or nitric acid, without evolving hydrogen, and when heated in the air is first oxidized to the mag- netic oxide, and then to ferric oxide. Ferrous Hydrate : Fe(OH) 2 , is obtained as a white precipitate on mixing solutions of ferrous sulphate and caustic soda both free from air. In consequence of the readiness with which it attracts oxygen from the air, the white precipitate rapidly becomes green, and finally yields a brown mass of ferric hydrate. Ferrous hydrate dissolves in salts of ammonium, and is, therefore, not precipitated by ammonia from a solution to which ammonium chloride has been added. Ferrous Sulphate, Green Vitriol: S0 2 '0 2 Fe + 7K.fi. — This, the best known salt of iron, and that from which most of the other salts are prepared, is obtained by dissolving iron in dilute sulphuric acid : — Fe + S0. 2 (OH) 2 = SO,-0 2 Fe + H 2 , or, as a bye-product -in making sulphuretted hydrogen, when ferrous sulphide is dissolved in the same acid : — FeS + S0 2 (OH) 3 = SO s -O a Fe + H,,S. It is also produced by gently roasting iron pyrites in the air, when a portion of the sulphur burns to sulphurous acid, and another portion is oxidized to sulphuric acid, which unites with the iron to form ferrous sulphate. Some varieties of pyrites — e.g. marcasite are oxidized at the ordinary temperature when moistened with water and exposed to the air, and the decomposed mass yields a green solution of ferrous sulphate when treated with water. The salt separates from its solution in bluish green crystals with 7 molecules of water ; it is easily soluble in water, especially hot. It readily loses 6 molecules of water when heated, but the seventh molecule, the water of constitution (p. 86), is only expelled at about 300°, with a partial decomposition of the salt. The anhydrous salt 446 Text-Book of Inorganic Chemistry. is white, but it regains its green colour when treated with water. If heated to redness; the anhydrous salt decomposes into ferric oxide and sulphurous and sulphuric anhydrides (p. 164). Both solid ferrous sulphate and its solution are partly oxidized when exposed to the air, with the formation of basic ferric sulphate of a reddish colour ; and when crystals of the substance which have been exposed to the air are dissolved in water, this insoluble basic salt remains undissolved : — 6S0 2 -0 2 Fe" + 3O + 3H.,0 = 2(SCy 3 O d Fe 2 " + 2Fe' // (OH) 3 . Ferrous sulphate unites with the sulphates of the alkalies (potassium and ammonium), and forms double sulphates containing 6 molecules of water — e.g. SO„-0.,Fe + S0 2 (OK) 2 + 6aq., corre- sponding to the magnesium compound (p. 415). These double sulphates oxidize less readily than ferrous sulphate itself. The salt is an important article of commerce, and is employed for the preparation of ferric acetate (used in dyeing), to produce a black colour with galls (writing ink), a blue colour with potassium ferro- cyanide (Prussian blue), and in the preparation of indigo solutions. It is also used as a disinfectant, especially to preserve wood from decay. Ferrous Carbonate : CO-0 2 Fe. — This compound is found in nature as spathic iron ore or siderite, sometimes crystallized in rhombohedra which are isomorphous with calcite, but usually in compact amorphous masses. Mixed with clay it forms clay iron- stone — one of the most important of the English ores of iron. The salt is obtained as a voluminous white precipitate on mixing a solution of ferrous sulphate with one of sodium carbonate. This white precipitate gradually absorbs oxygen from the air and changes colour, becoming first of a dirty green and then of a brown colour, and is finally completely converted into ferric hydrate, while car- bonic acid is set free — no carbonate containing iron in the ferric state is known. If the moist precipitate is rubbed up with sugar, it is somewhat protected from the action of the air, and a brownish mass is obtained which is used in medicine under the name oiferri carbonas saccharata. Like calcium carbonate and other carbonates, ferrous carbonate, though insoluble in pure water, dissolves in water containing car- bonic acid. Mineral springs containing ferrous carbonate dissolved in carbonic acid are called chalybeate springs. The water has a bitter, inky taste, and when exposed to the air becomes turbid from Ferric Oxide. 447 ferric hydrate. The dissolved ferrous carbonate is decomposed by the oxygen of the air into ferric hydrate and carbonic acid : — 2CO-0„Fe" + O + 3H 2 = 2Fe'"(OH) 3 + 2CO r POl Ferrous Phosphate : p( -. 6 Fe 3 + 8H 4 0, is found in nature as the mineral vivianite, sometimes crystalline, but usually as an earthy powder of a dull blue colour. The compound is precipitated when a solution of sodium phosphate is added to one of ferrous sulphate, and is used in medicine. It is at first white, but becomes blue or green by partial oxidation. By heating iron with phosphoric acid, a solution is obtained which after some time deposits colour- ( O Fe less needles of the composition : PO 1 qIt + H 2 0. Terric Oxide : Fe.^Oj. — This compound occurs in immense quantities in the mineral kingdom, either in lustrous crystals belonging to the hexagonal system — specular iron, or more com- monly in amorphous masses, with or without water, as the important iron ores included under the name of hamatite. It may be ob- tained artificially in the crystalline state by strongly glowing a mixture of dry ferrous sulphate and common salt, or in the amor- phous state by heating dry ferrous sulphate or ferric hydrate. The red amorphous powder obtained by either of the last-named methods is used for polishing and as a pigment, and is known as colcothar, caput mortuum, or rouge. A fine variety of rouge is obtained by glowing ferrous oxalate in the aiv. Ferric oxide is completely insoluble in water, and after heating to redness only dissolves in acids with difficulty. Ferric Hydrate : Fe(OH) 3 , is thrown down as a red-brown flocculent precipitate when ammonia is added to a solution of a ferric salt. When washed with water and dried, it forms a brown amorphous mass, insoluble in water. Some important ores of iron, and the rust which forms when iron is exposed to the air, con- sist essentially of ferric hydrate. Freshly precipitated ferric hydrate dissolves not only in acids, but also in a solution of ferric chloride, producing a dark red solu- tion. If this solution of basic chloride is placed in a vessel of which the bottom is closed with a piece of parchment paper and then floated on water, hydrochloric acid diffuses through, and there remains in the dialyzer a solution of colloid ferric hydrate as a 448 Text-Book of Inorganic Chemistry. blood-red liquid {dialyzed iron). The addition of a small quantity of sulphuric acid, or of an alkali or salt, at once causes the precipi- tation of the ferric hydrate from this solution as a gelatinous mass. The use of ferric hydrate as an antidote in cases of poisoning by arsenic has been already referred to (p. 236). Like alumina, ferric oxide is a weak base, and its soluble salts can only be crystallized with difficulty. The normal salts are white or yellowish when anhydrous, red when hydrated ; the basic salts are yellow or red. Ferric Sulphate : (S0 2 ) 3 6 Fe,. — This salt is produced by oxidizing a solution of ferrous sulphate with nitric acid, with the addition of the one molecule of sulphuric acid which one molecule of ferric sulphate contains more than two molecules of ferrous sulphate : — 2SO,-0 2 Fe" + S0 2 (OH)„ + O = (S0 2 ) 3 6 Fe./" + H 2 0. When the nitric acid is dropped into the hot solution of ferrous sulphate and sulphuric acid, the liquid first becomes dark brown and nearly black from the compound of ferrous sulphate with the nitric oxide produced by the reduction of the nitric acid (p. 194), until, when sufficient nitric acid has been added, the whole of the nitric oxide is evolved with effervescence, and the dark-red liquid now contains ferric sulphate. If no sulphuric acid is added to the ferrous sulphate, a larger quantity of nitric acid is required, and the solution then contains a mixture of ferric sulphate and ferric nitrate. On evaporating the aqueous solution of the salt, ferric sulphate remains as a white amorphous powder, which attracts water from the air, and deliquesces to a dark red liquid, with a bitter astringent taste. An excess of ammonia when added to its aqueous solution precipitates ferric hydrate, but if less ammonia is added than is required to completely decompose it, a. brown precipitate of a basic salt is produced. The oxidation of ferrous sulphate in the air produces a similar compound (p. 446). A solution of ferric sulphate, when mixed with one of the potas- sium sulphates and the mixed solution allowed to evaporate slowly, deposits yellowish crystals of potash-iron alum. The correspond- ing ammonium compound — ammonia-iron alum : o|y ! [ /Vvju + Triferric Tetroxide. 449 12H,0, crystallizes more readily. The crystals are usually of a violet colour, possibly due to organic compounds contained in the ammonia. Ferric Nihate : (N0 2 ) 3 3 Fe, is prepared by dissolving iron in moderately concentrated nitric acid. From the red solution so obtained, the salt can only be crystallized with difficulty. The solid compound deliquesces in the air. Ferric Phosphate {Orthophosphale) : (PO)0 3 Fe, is obtained as a yellow amorphous precipitate when ferric chloride is mixed with sodium phosphate. It is insoluble in water and acetic acid, but dissolves in strong inorganic acids. Ferric Pyrophosphate closely resembles the preceding salt, and is insoluble both in water and acetic acid, but forms a soluble com- pound with sodium phosphate or pyrophosphate. The soluble double pyrophosphate is prepared by adding sodium pyrophosphate to the ferric salt until it is just dissolved ; from this solution alcohol precipitates the double salt. Its aqueous solution is colourless, and does not possess the strong astringent taste of the other soluble ferric salts. It is used in medicine. Ferric carbonate does not appear to exist. Ferric oxide is too weak a- base to unite with carbonic acid. On adding a solution of sodium carbonate to one of ferric chloride, a brown precipitate of ferric hydrate is produced, and the carbonic acid which is set free partly unites with the excess of sodium carbonate to form the acid carbonate. Triferric Tetroxide, Magnetic Oxide of Iron : Fe 3 4 = Fe"0, Fe 2 w 3 . — This oxide is found in nature as black, well- developed, regular octahedra, especially in chlorite rocks. It further occurs as crystalline masses, and more largely in rough granular deposits. Magnetic oxide of iron, so called because pieces of it sometimes form natural magnets, is one of the most valuable ores of iron. Small lustrous octahedra of the oxide may be obtained artificially by heating bright iron wire to redness in a porcelain tube, and leading a current of steam over it. The water is then decomposed according to the equation : — 3 Fe + 4H a O' = Fe 3 4 + 4H 2 . G G 450 Text-Book of Inorganic Chemistry. Magnetic oxide of iron is soluble in hydrochloric acid, but the solution does not then contain a chloride corresponding to the oxide, but a mixture of ferric and ferrous chlorides. If, ammonia is gradually added to this solution, a precipitate of ferric hydrate is first formed, and then one of ferrous hydrate. But if the solu- tion is poured, with stirring, into an excess of ammonia, a dark- green precipitate is produced, which becomes black and granular when boiled. This is the hydrate of the magnetic oxide, and re- mains unchanged in the air. The same compound may also be obtained by converting two-thirds of a solution of ferrous chloride into ferric chloride by passing chlorine, then mixing this, when the excess of chlorine has been driven off, with the other third of the solution of ferrous chloride, and finally pouring the mixed solu- tion into an excess of ammonia and boiling. FeiTtc Acid. — The free acid, which would have the composi- tion : Fe vi 2 (OH) 2 , has not yet been prepared ; its potassium salt, potassium ferrate : Fe0 2 (OK) 2 , is obtained by heating a mixture of reduced iron with twice its weight of potassium nitrate. The oxidation of the iron takes place with a considerable evolution of heat, and on extracting the cold mass with water the potassium ferrate dissolves with a red colour. On evaporating this solution in a vacuum, dark-red crystals are deposited. The salt may also be obtained by passing chlorine through concentrated cooled caustic potash containing freshly precipitated ferric hydrate in suspension : — Fe(OH) 3 + 5KOH + 3CI = Fe0 2 (OK) 2 + 3KCI + 4H 2 0. Iron forms two compounds with each of the halogens — chlorine, bromine, and iodine — corresponding to ferrous and ferric oxide respectively. Ferrous chloride : FeCl 2 . — The anhydrous compound is ob- tained by heating iron wire in a stream of hydrochloric acid gas, when it sublimes at a high temperature in white, glistening iri- descent crystals which feel soft and talc-like to the touch. It melts at a red-heat, and then solidifies to a crystalline mass. It is easily soluble in water, and on evaporating the aqueous solution without access of air, green crystals of the compound : FeCl 2 + 4H 2 separate out. The same compound may be obtained by dissolving iron in hydrochloric acid and evaporating down. A solution of ferrous chloride, when exposed to the air, is gradually oxidized to basic ferric chloride. Ferric Chloride. 451 Ferric Chloride : Fe 2 Cl B . — Although ferrous chloride readily unites with a further quantity of chlorine, an atom of iron can never decompose more than two molecules of hydrochloric acid. And when iron is dissolved in the acid, ferrous chloride is always produced, never ferric chloride. Anhydrous ferric chloride may be obtained by heating iron in dry chlorine, and then sublimes in nearly black, lustrous, crystalline plates. It is far more volatile, and therefore sublimes much more easily than ferrous chloride. When exposed to the air it becomes yellow, and deliquesces, forming a yellow or reddish solution ; it is also soluble in alcohol and ether, and is even abstracted from its aqueous solution when shaken with ether. The aqueous solution, which is also obtained by dissolving iron in aqua regia, or by passing chlorine through a solution of ferrous chloride, deposits reddish, deliquescent crystals of the compound: Fe 2 Cl 6 + I2H 2 0, when evaporated to a syrup. This compound, when heated, decomposes into water and hydro- chloric acid, which pass off with some anhydrous ferric chloride, while ferric oxide remains behind. An aqueous solution of ferric chloride can dissolve freshly precipitated ferric hydrate in con- siderable quantities, producing a dark-red liquid, from which ferric hydrate is again precipitated on diluting with water and heating. Ferric chloride unites with the chlorides of potassium and ammonium to form double salts, which are deposited in bright-red crystals of the composition: Fe 2 Cl 6 + 4KC1 + 2H 2 and Fe 2 Cl 6 + 4NH 4 C1 + 2H.,0, when the mixed solutions are evaporated down. The bromides of iron closely resemble the chlorides. Ferrous Iodide : Fel 2 . — This compound is obtained by heating iron filings with iodine in a porcelain crucible. A small quantity of iodine is first added, then an excess, and the heating continued as long as vapours of iodine are evolved. The salt then remains as a dark green lamellar mass, easily solub'e in water. An aqueous solution of ferrous iodide may also be obtained by digest- ing iron filings with water and iodine — the iron being in excess. On evaporating this solution, crystals of the composition : Fel 2 + 4H 2 0, are deposited. This solution dissolves more iodine when digested with it, and the dark-red solution so produced probably contains ferric iodide : Fe 2 I 6 , which has not yet been obtained in the solid form. The compounds of iron with sulphur do not exactly correspond CC2 45 2 Text-Book of Inorganic Chemistry. with its oxides. A sesquisulphide of the composition : Fe.^S^and corresponding to ferric oxide, is not yet known for certain. On the other hand, the sulphide most widely distributed in nature is the disulphide : FeS 2 , to which no corresponding oxide is known. Ferrous Sulphide : FeS, is prepared as a black crystalline mass, by heating three parts of iron filings with nearly two parts of flowers of sulphur. If an excess of sulphur is employed, no compound richer in sulphur is obtained, but the excess is volatilized. Iron and sulphur unite together even at ordinary temperatures if the finely powdered mixture is moistened with water. When larger quantities are used the heat of combination is often suffi- cient to violently expel a portion of the mixture from the vessel in which the reaction takes place. Ferrous sulphide is insoluble in water, but is decomposed by acids. With hydrochloric or dilute sulphuric acid it yields con- siderable quantities of sulphuretted hydrogen, and is therefore largely used to prepare this gas. Ferric Bisulphide : FeS,. — This compound is widely distri- buted in the mineral kingdom, either in bright yellow crystalline masses, or in well-developed lustrous cubes, and other forms belonging to the regular system. In these forms it is known as iron-pyrites. It may also be prepared in the form of a yellow powder by gently heating powdered ferrous sulphide with about its own weight of sulphur, until the excess of this latter substance is expelled. Iron-pyrites is so hard that it strikes sparks with steel ; unlike ferrous sulphide, it is not attacked by dilute acids, and is not therefore adapted for the preparation of sulphuretted hydrogen. When heated strongly it parts with some of its sulphur, leaving compounds of ferrous sulphide with more or less sulphur, accord- ing to the degree and duration of the heating. If iron-pyrites is heated in the air, a large portion of the sulphur is oxidized to sulphurous anhydride, and for this reason immense quantities of pyrites are employed in the manufacture of sulphuric acid. Many kinds of iron-pyrites contain small quantities of copper-pyrites, and traces of the noble metals, and the residues of these pyrites after roasting are often worked for copper and the other metals. Besides iron-pyrites, which crystallizes in the regular system, a second modification of ferric disulphide is known, the crystals of which belong to the rhombic system, and which is distinguished Compounds of Iron and Cyanogen. 453 from iron-pyrites by the readiness and rapidity with which it oxidizes when exposed to moist air — evolving at the same time a consider- able amount of heat. This variety of pyrites is called marcasite ; and the chief product of its oxidation is ferrous sulphate, which is obtained in considerable quantities from it. A third sulphide of iron is also found in nature, which contains less sulphur than iron-pyrites, and not unfrequently has the com- position : Fe 7 S a , or, if considered as a compound of ferrous sul- phide and iron-pyrites : 6FeS,FeS 2 . This bronze-coloured mineral is known as magnetic pyrites, because it is attracted by the magnet. It crystallizes in the hexagonal system. COMPOUNDS OF IRON AND CYANOGEN. Although the compounds of iron with cyanogen corresponding to ferrous and ferric chlorides have not yet been prepared pure, the double salts which ferrous and ferric cyanide yield with the cyanides of other metals are among the best known and most stable of chemical compounds. Various views are held as to the chemical constitution of these substances, but in the present state of our knowledge it is perhaps best to write their formulas em- pirically, without entering into the question of their constitution. The best known of these, and that from which all the others are prepared, is Potassium Ferrocyanide : K 4 Fe"Cy 6 + 3H 2 = 4KCy, Fe"Cy 2 + 3H 2 0. — This compound, known in trade as yellow prussiate of potash, or simply as yellow prussiate, crystallizes from its aqueous solution in large, pale-yellow, quadratic plates, which are soft, and therefore difficult to powder. Heated to 100° or a little above, it loses its water of crystallization, and then becomes white and brittle, so that it can be easily pulverized. It is soluble in water, but insoluble in. alcohol ; 100 parts of water dissolve 25 parts at the ordinary temperature and 50 parts at the boiling-point. Potassium ferrocyanide is obtained from potassium cyanide — a salt which is extremely poisonous, and which is readily decom- posed by dilute acids. From this it might be expected that the ferrocyanide would also be poisonous, but experiment has proved that considerable quantities may be taken internally without any in- jurious or even unpleasant effects. When taken into the system it 454 Text-Book of Inorganic Chemistry. soon gets into the blood, and in a short time can be detected in the urine, in which it is expelled unchanged from the body. The salt is readily obtained from potassium cyanide by digest- ing an aqueous solution with iron filings or powdered ferrous sulphide. In the former case, caustic potash and hydrogen are also produced ; in the latter potassium sulphide : — 6KCy + Fe + 2H 2 = K 4 FeCy 6 + 2KOH + H 2 and 6KCy + FeS = K 4 FeCy 6 + K 2 S. The same process goes on in the manufacture of potassium ferrocyanide on a large scale, except that instead of using potassium cyanide itself materials are employed which, when heated with potassium carbonate, produce this substance. For this purpose, organic compounds rich in nitrogen, such as blood, horn, hair, leather, are mixed with potash and iron filings, and fused at a red heat. During the process the carbon and nitrogen of the organic substances unite with the potassium and form potassium cyanide ; and when the fused mass is afterwards extracted with water this substance acts upon the iron, which has been partly converted into ferrous sulphide by the sulphur of the crude materials, in the manner explained above. The salt which separates out on evapo- ration is purified by recrystallization ; it often contains potassium sulphate. Potassium ferrocyanide is not acted upon by carbonic acid nor even by dilute acids, and the potassium cyanide in combination with the ferrous cyanide has, therefore, acquired other properties, and has become more stable. The iron, too, has lost many of the properties which are possessed by the ferrous salts ; ammonia does not precipitate ferrous hydrate, nor does ammonium sulphide throw down ferrous sulphide. And that potassium ferrocyanide is not an ordinary double salt of 4 molecules of potassium cyanide and I molecule of ferrous cyanide is proved by the fact that it is not decomposed by dilute acids like other true double cyanides. The double cyanide of nickel and potassium : 2KCy,NiCy 2 , a compound soluble in water, at once gives a precipitate of nickel cyanide with dilute hydrochloric acid, while the potassium cyanide is decomposed into potassium chloride and hydrocyanic acid. But no ferrous cyanide is separated by hydrochloric acid from the apparently analogous compound potassium ferrocyanide. Potassium Ferrocyanide. 455 Many attempts have been made to explain this peculiar be- haviour of potassium ferrocyanides and other similar compounds, but as yet without success. That potassium ferrocyanide when heated with sulphuric acid leaves a compound containing 1 mole- cule of ferrous cyanide and 1 molecule of potassium cyanide, seems to prove that one of the four molecules of potassium cyanide in the ferrocyanide is more firmly combined than the other three. An aqueous solution of potassium ferrocyanide is decomposed by strong acids (hydrochloric or sulphuric acid), with the forma- tion of hydroferrocyanic acid : H 4 FeCy e , and the potassium salt of the acid employed. If finely powdered potassium ferrocyanide is heated with dilute sulphuric acid, hydrocyanic acidis evolved (p. 313), and there remains potassium sulphate with the compound we have just mentioned : KCy,FeCy 2 . This is a white substance, turning blue in the air. Potassium ferrocyanide, when powdered and heated with concentrated sulphuric acid, undergoes a further decomposi- tion. There are produced ferrous, potassium, and ammonium sulphates with the evolution of a gas which consists chiefly of car- bonic oxide. But since traces of other gases are mixed with the carbonic oxide, this method is not suitable for preparing the pure substance. When heated in an iron crucible, potassium ferrocyanide is decomposed into iron carbide and potassium cyanide (p. 359). Fused with manganese peroxide or red-lead, potassium cyanate and ferric oxide are produced (p. 361). And when mixed with potassium carbonate and sulphur and fused, potassium sulpho- cyanate results (p^ 361). If chlorine is led through a solution or potassium ferrocyanide, the iron is oxidized to the ferric state and some of the potassium converted into chloride, producing a com- pound called potassium ferricyanide : K 3 Fe" / Cy 6 = 3KCy,Fe /// Cy 3 (P- 457)- Potassium ferrocyanide, like all other soluble ferrocyanides, is distinguished by its reactions with ferric chloride and soluble copper salts. When highly diluted it gives with the former a dark blue precipitate of Prussian blue, and with the latter a chocolate brown precipitate of copper ferrocyanide. With a soluble lead salt a solution- of potassium ferrocyanide gives a white precipitate of the corresponding lead salt. The following equations represent these reactions : — 456 Text-Book of Inorganic Chemistry. 2K 4 Fe"Cy 6 + Fe 2 '"Cl 6 = 6KC1 + 2^ e '"| Cy 4 ,Fe"Cy 2 . Prussian blue K 4 FeCy 6 + 2CuCl 2 = 4KCI + Cu 2 FeCy 6 . Copper ferrocyanide K 4 FeCy e + 2(N0 2 ).,0 2 Pb = 4N0 2 -OK + Pb 2 FeCy 6 . Lead ferrocyanide The majority of the ferrocyanides are insoluble in water ; those of the alkalies and alkaline earths are soluble, but that of barium only with difficulty. Sodium Ferrocyanide: Na 4 FeCy 6 + I2H 2 0. — This salt cor- responds closely to the potassium compound, except that it contains considerably more water of crystallization. It may be obtained from a Prussian blue which contains an atom of hydrogen in place of potassium in the usual compound, and which is pre- cipitated from an acid solution. If this Prussian blue is boiled with caustic soda, ferric hydrate and soluble sodium ferrocyanide are produced : — ^ e |cy 4 ,FeCy 2 +- 4NaOH = Fe(OH) 3 + Na 4 FeCy 6 , + H 2 0. Other soluble ferrocyanides may be obtained in the same way. Ammonium Ferrocyanide : NH 4 FeCy 6 + 3H 2 0, crystallizes in bright yellow quadratic prisms, and is isomorphous with the potas- sium salt. Fe'"l Prussian Blue : ^ \ Cy 4 ,FeCy 2 . Various dark blue compounds are known under the name of Prussian blue, but the commonest insoluble compound has the above composition ; it is obtained on adding potassium ferro- cyanide to a solution of ferric chloride or some other ferric salt. The compound may be considered as potassium ferrocyanide in which three atoms of potassium have been displaced by one of triad iron. The atom of potassium which it contains may be dis- placed by hydrogen without any change in the physical properties of the substance. Such a compound may be obtained by precipi- tating hydroferrocyanic acid with ferric chloride, or by adding Prussian Blue. 457 potassium ferrocyanide to a solution of ferric chloride containing free hydrochloric acid. Prussian blue can only be washed on the filter with difficulty ; when dry.it forms a deep blue, hygroscopic powder, which under pressure assumes a lustre like copper. It is slightly soluble in dilute acids, and when warmed with concentrated acids is decom- posed. Oxalic acid dissolves the compound when free from potassium, and gives a deep blue liquid, which is used as a blue ink. Alkalies and alkaline carbonates and even the alkalies of soap decompose it, separating ferric hydrate and reproducing potassium ferrocyanide : — £ e jCy 4 ,FeCy 2 + 3KOH = K 4 FeCy 6 + Fe(OH) 3 . Prussian blue cannot, therefore, be used for dyeing or printing articles which are to be washed with soap. A variety of Prussian blue, soluble in water to a deep blue solution, is obtained when ferric chloride is added to a large excess of potassium ferrocyanide. This soluble Prussian blue is probably a compound of ordinary Prussian blue and potassium ferrocyanide. Hydroferrocyanic Acid : H 4 FeCy 6 . — This acid compound is obtained as small white crystalline plates when a concentrated solution of potassium ferrocyanide free from air is mixed with strong hydrochloric acid in the cold : — K 4 FeCy 6 + 4HCI = 4KCI + H 4 FeCy 6 . The crystals separate more readily if a little ether is added. They are afterwards brought upon a filter and washed with water con- taining ether. The compound is soluble in alcohol and in water, but is insoluble in hydrochloric acid and ether. It behaves like a tetrabasic acid, and when mixed with potassium carbonate again produces potassium ferrocyanide. It is rapidly oxidized in the air and becomes of a blue colour. Potassium Ferricyanide : K 3 Fe'"Cy 6 = 3KCy,Fe'"Cy s .— This compound, which is known in trade as red prussiate, differs from the ferrocyanide by containing its iron in the ferric state. It is produced by abstracting potassium from the ferrocyanide, when the cyanogen thus set free unites with the iron and changes it from the ferrous to the ferric state. This is best done by leading 458 Text-Book of Inorganic Chemistry. chlorine into a solution of potassium ferrocyanide until a portion of the liquid no longer gives a blue precipitate with ferric chlo- ride :— K 4 FeCy 6 + CI = KC1 + K 3 FeCy 6 . The salt separates out on evaporating its aqueous solution in dark-red, lustrous, rhombic prisms, easily soluble in water and in- soluble in alcohol. Its aqueous solution is of a dirty green colour and decomposes on standing, especially when exposed to sunlight, depositing a blue precipitate. The readiness with which it again forms potassium ferrocyanide, especially in the presence of alka- lies, make it a powerful oxidizing agent. Potassium ferricyanide gives no blue precipitate with ferric salts, but only a brown colour, but with ferrous salts it gives a fine blue precipitate, closely resembling Prussian blue, and called — Fe' Turnbull's Blue : ,- [ Cy 3 ,Fe"'Cy 3 . — The precipitate is formed according to the equation : — ■ 3KCy,FeCy 3 + FeCl 2 = 2KCI + „- e \ Cy 3 ,FeCy 3 . The compound, like Prussian blue, is decomposed by caustic alkalies, reproducing potassium ferricyanide, while ferrous hydrate is precipitated. At the same time, and especially with excess of alkali, a portion of the ferrous hydrate is oxidized to ferric hydrate by the potassium ferricyanide, which is reduced to the ferro- cyanide. I-tydroferricyanic Acid : H 3 FeCy 6 . — This compound separates, like the corresponding hydroferrocyanic acid, on mixing a solution of potassium ferricyanide with strong hydrochloric acid in the cold. It forms brown crystals, which are easily decomposed. sodium Nitroprusside : Na 2 Fe'"Cy 5 (NO) + 2rl 2 0. — This salt, which contains one atom of sodium less than the ferricyanide, and one atom of cyanogen displaced by the monad radical : NO (nitryl), crystallizes more easily than that of potassium. It may be prepared in the following manner. Two parts of potassium ferrocyanide are gently warmed with 3 parts of strong nitric acid which has been diluted with an equal volume of water, until a portion of the liquid no longer gives a blue precipitate either with ferric chloride or with ferrous sulphate. The liquid is then allowed Detection of Iron Compounds. 459 to cool, so that potassium nitrate which has been produced shall crystallize out ; the acid liquid is then poured off, diluted with water, and neutralized with sodium carbonate. On evaporating this solution sodium nitroprusside crystallizes out on cooling in large, dark-red, transparent crystals, easily soluble in water and in dilute alcohol. The aqueous solution gradually decomposes, especially when exposed to light, and deposits a blue precipitate. Sodium nitroprusside does not precipitate ferric salts, but is especially characterized by the dark purple colour which it produces with alkaline sulphides, and even with sulphu- retted hydrogen. The colour soon changes to a blue, and after- wards becomes a dirty brown. This reaction is so delicate that by its means the presence of the most minute traces of an alkaline sulphide or of a nitroprusside may be detected. Detection of Iron Compounds. The presence of small quantities of iron in any form may be re- cognized by the colour of the borax bead produced in the blow- pipe flame. In the outer or oxidizing flame, the ferric oxide colours the bead yellow, the colour becoming fainter on cooling. In the inner or reducing flame, the bead becomes of an olive-green colour, due to ferrous oxide. These colours cannot be observed in the presence of some other metals — e.g. cobalt — which obscure the colours produced by the oxides of iron. An aqueous solution of a ferrous salt gives no precipitate with sulphuretted hydrogen, but a black precipitate of ferrous sulphide with ammonium sulphide. Caustic soda or ammonia precipitates ferrous hydrate from the same solution ; the precipitate is first white, but soon oxidizes and becomes green and then brown. Potassium ferrocyanide gives a white precipitate rapidly oxidizing to Prussian blue, while potassium ferricyanide gives a blue pre- cipitate (Turnbull's blue) at once. When a stream of sulphuretted hydrogen is led through a solution of a ferric salt, a white precipitate of sulphur falls down, and the ferric salt is reduced to the ferrous state : — Fe 2 Cl 6 + H 2 S = 2FeCl 2 + 2HCI + S. And the black precipitate which ammonium sulphide produces in ferric salts is a mixture of ferrous sulphide and sulphur. Caustic soda or ammonia, "when added to a solution of a ferric salt, pre- 460 Text-Book of Inorganic Chemistry. cipitates brown ferric hydrate ; the presence of potassium tartrate prevents the precipitation. If a ferric salt containing phosphoric acid is precipitated with ammonia, the precipitate is not ferric hydrate, but ferric phosphate. Potassium ferrocyanide when added to a ferric salt gives a dark- blue precipitate of Prussian blue. Potassium ferricyanide produces no precipitate, but only a brown colour. Iron, in the form of a ferric salt, may be separated from the metals of the alkaline earths by precipitating with ammonium chloride and ammonia. Ammonium sulphide precipitates iron in any form, but does not throw down the alkaline earths. Iron is separated from aluminium by digesting with an excess of caustic soda, which dissolves aluminium hydrate, but not ferric hydrate. MANGANESE. Chemical Symbol: Mn. — Atomic Weight: 55. Manganese is closely related to iron in its properties, and has about the same atomic weight. Its compounds are also nearly always found associated with iron in nature. The most important minerals containing manganese are its oxygen compounds, and of these pyrolusite (manganese peroxide : Mn0 2 ) is found in by far the largest quantities. Other manganese minerals are brau- nite (manganic oxide : Mn 2 O a ), manganite (manganic hydrate : I MnO-OH^' h ausmann it e (trimanganic tetroxide : Mn 3 4 ), man- ganese-spar (manganous carbonate: CO-0 2 Mn), psilomelane and •wad (impure manganic perhydrate). Manganese also occurs in nature as its silicate, and in combination with sulphur as various sulphides. The metal itself is of but little interest. It may be obtained by heating a mixture of precipitated manganous carbonate and oil to redness, and then again heating this mixture with charcoal powder and anhydrous borax to the highest temperature of a wind furnace. The regulus of metal so obtained, which contains carbon and other impurities, is of a grey colour with a reddish tinge, very hard and brittle, and with a high melting-point. It has a specific gravity of about 7 - 2, and oxidizes when exposed to moist air, The metal Very weak base Indifferent Dibasic acid Monobasic acid. Manganese. 46 r dissolves easily in dilute sulphuric, hydrochloric, or even nitric acid, producing the corresponding manganous salts. Manganese unites with oxygen in no less than five proportions, forming oxides of very different chemical properties, and in which the manganese plays the part of a dyad, triad, tetrad, hexad, and heptad element. These compounds are the following : — Manganous oxide : MnO . . . Strong base Manganic oxide : Mn s O, Manganese peroxide : Mn0 2 Manganic acid : H 2 MnO, Permanganic acid : HMn0 4 Of these oxides, manganous oxide corresponds to ferrous oxide, manganic oxide to ferric oxide, and an intermediate compound of these two : Mn 3 4 , to the magnetic oxide of iron. An oxide of iron corresponding to manganese peroxide is unknown, but we are ac- quainted with the corresponding sulphide in iron-pyrites. Finally, manganic acid corresponds to ferric acid, but no compound of iron is known similar to permanganic acid. manganous Oxide : MnO. — This oxide remains as a greenish powder when manganous carbonate is glowed out of contact with air, or in a stream of hydrogen. It may also be obtained by heat- ing a mixture of manganous chloride and sodium carbonate. Man- ganous oxide is a strong base, and dissolves easily in hydrochloric, sulphuric, or nitric acid. When heated in the air it absorbs oxygen and is converted into trimanganic tetroxide : Mn 3 4 . Unlike the iron compound, it is not reduced to the metallic state by hydrogen, even at a high temperature. Manganous Hydrate : Mn(OH) 2 , is deposited as a white pre- cipitate when caustic soda is added to a solution of a manganous salt. It rapidly darkens, especially when exposed to the air, and finally becomes completely converted into dark brown manganic hydrate. Ammonia produces the same precipitate when the solution does not contain ammonium salts ; manganous hydrate forms soluble compounds with these substances. The most important of the manganous compounds — i.e. those in which the metal is a dyad — are the following:— manganous Sulphate : SO„-0 2 Mn + 7H 2 0.— This salt is pro^ 462 Text-Book of Inorganic Cliemistry, duced, with evolution of oxygen, when manganese peroxide is heated with strong sulphuric acid : — Mn0 2 + S0 2 (OH) 2 = S0 2 -0 2 Mn + H^O + O. The filtered solution, when sufficiently concentrated and freed from iron, deposits the salt in clear, rose-coloured, tabular crystals. At the ordinary temperature the salt crystallizes with only five molecules of water, and at higher temperatures other hydrated salts may be obtained. Like magnesium and ferrous sulphates, the salt unites with the alkaline sulphates, forming double sulphates with six molecules of water. Manganous Nitrate : (N0 2 ) 2 -0 2 Mn. — If manganous carbonate is dissolved in nitric acid and the solution evaporated down, man- ganous nitrate is obtained, and the salt is not oxidized as in the case of iron. Manganous nitrate is highly deliquescent, and can therefore only be crystallized with difficulty ; when dried and heated it darkens in colour, and manganese peroxide separates out. Manganous Carbonate: CO-0 2 Mn, is thrown down as a white precipitate when sodium carbonate is added to a solution of a manganous salt. It remains unchanged in the air when moist and only slightly darkens when dried. The salt is therefore much more stable than ferrous carbonate. Manganous carbonate occurs in nature as the mineral manganese-spar, crystallizing in rhombo- hedra, isomorphous with spathic iron ore and calcite. Manganous Chloride : MnCl 2 + 4H 2 0. — Of the two chlorides of iron, ferric chloride is the more stable, and ferrous chloride readily combines with chlorine. But manganous chloride is by far the most stable chloride of the metal. The compound is pro- duced when manganous carbonate is dissolved in hydrochloric acid, and is also a bye-product in the preparation of chlorine from man- ganese peroxide. As already stated (p. 98), manganese tetra- chloride : MnCl 4 , is first produced, but is so unstable that it at once breaks up into manganous chloride and free chlorine. But as the manganese peroxide used for the preparation of chlorine always contains iron, the manganous chloride produced when it is dis- solved in hydrochloric acid always contains ferric chloride. Various methods are used to separate the iron, one of which consists in heating the dried residue in a Hessian crucible to low redness : the ferric chloride is then partly volatilized and partly Manganous Compounds. 463 con verted into insoluble basic chloride, but the manganous chloride remains unchanged. The mass is afterwards powdered and ex- tracted with hot water ; and on evaporating the violet-coloured solution, manganous chloride with four molecules of water separates out in pink-coloured crystals. It is very easily soluble in water, and deliquesces in moist air ; when heated, it parts with all its water, accompanied by a trace of hydrochloric acid. Manganous chloride unites with ammonium chloride and forms a crystalline double salt of the composition : MnCl 2 ,2NH 4 Cl + H 2 0. Manganous Sulphide : MnS, is found in the mineral kingdom as alabandite, in black cubes, or in crystalline masses. It is obtained artificially as a flesh-coloured precipitate by mixing a solution of a manganous salt with one of sodium sulphide. It easily dissolves in dilute acids with evolution of sulphuretted hydrogen. — Manganese Bisulphide : MnS 2 , has not been prepared artificially, but occurs in nature as the mineral hauerite in reddish-brown crystals belonging to the regular system. Manganic Oxide : Mn 2 3 , occurs in nature as the mineral braunite in black lustrous crystals belonging to the quadratic system. Manganic Hydrate : j M r\ r\\i is found also in nature as manganite in dark-brown rhombic prisms or crystalline masses. It is prepared as a dark-brown amorphous precipitate when a solution of manganous chloride is mixed with ammonium chloride and ammonia and exposed to the air. Unlike manganous hydrate, it is insoluble in ammonium chloride. Manganic oxide is a weak base, and difficultly combines with acids. Its salts decompose easily, especially when warmed, and are converted into manganous compounds. Manganic Sulphate : (SO 2 ) 3 O Mn 2 '", may be obtained as a dark-green amorphous powder by digesting pure manganese per- hydrate with concentrated sulphuric acid. It deliquesces when exposed to the air, forming a red liquid, and is a difficult compound to prepare. Somewhat more stable is its double compound with potassium sulphate, manganese alum: eg | ok + I2 H 2 0, which is obtained by mixing the solution of the two sulphates. 464 Text-Book of Inorganic Chemistry. Manganic Chloride is a very unstable compound, and unknown in the solid state. When manganic hydrate is treated with hydro- chloric acid in the cold a dark brown liquid is obtained, which probably contains manganic chloride, but which continually evolves chlorine. Trimanganic Tetroxide : Mn 3 4 = MnO,Mn 2 0„ is found in nature as the mineral haus7nannite in brown quadratic crystals. It is the most stable oxide of manganese, and may be heated to redness in the air without change. If manganese peroxide is strongly heated as long as oxygen is evolved, this compound remains as a reddish powder, and when manganous oxide or car- bonate is heated in the air it absorbs oxygen and produces the same substance. Hydrochloric acid dissolves trimanganic tetroxide, forming manganous chloride, while chlorine is evolved. Manganese Peroxide : MnO„, is the commonest and most important of the manganese minerals, and as such is known by the name fiyrolusite. It crystallizes in grey rhombic prisms, with a metallic lustre, but is usually found in rough masses often fibrous in structure. The chief European localities where it is found are Thuringia, and on the Lahn, in Germany, and in Bohemia, France, and Spain. Manganese peroxide gives a grey streak on paper or porcelain, and may thus be distinguished from other manganese minerals, which give a brown streak. The compound may be artificially prepared by carefully heating manganous carbonate with potassium chlorate and extraction with water, or by heating manganous nitrate or the hydrated peroxide. Manganese peroxide is an indifferent substance, insoluble in water and in nitric acid, even when concentrated. Strong sulphuric acid dissolves it with formation of manganous sulphate and free oxygen ; hydrochloric acid converts it into manganous chloride and chlorine. Substances which are easily oxidized, such as sugar, oxalic acid, &c, reduce manganese peroxide in the presence of sulphuric acid. Oxalic acid is thus completely converted into carbonic acid and water, while the peroxide becomes manganous sulphate. A process based upon this reaction is used to determine the. quantity of the peroxide contained in commercial pyrolusite. Every molecule of manganese peroxide oxidizes a molecule of oxalic acid to carbonic acid, and it is, therefore, only necessary to know the weight of the latter substance, which may be either Manganese Peroxide, 465 collected and weighed, or estimated from the loss in weight, in order to determine the value of any particular sample of the ore : I CO -OH + Mn0 * + s °2(OH) 2 = 2C0 2 + S0 2 -0 2 Mn + 2H 2 0. Manganese peroxide is largely used for many technical purposes, chiefly for the manufacture of chlorine and bleaching powder. When heated, it liberates considerable quantities of oxygen, and becomes converted into trimanganic tetroxide. Manganese Perhydrate, of uncertain composition, is obtained as a dark-brown amorphous powder when a solution of manganous chloride is precipitated by sodium hypochlorite, or when chlorine is passed through a solution containing suspended manganous hydrate. The hydrate also separates when carbonic acid is passed through a solution of potassium manganate, which is then changed into potassium permanganate. The same change takes place if a very dilute solution of potassium manganate is boiled. (OH Qu — The acid of the above com- position has not yet been prepared ; in all attempts to separate it, it breaks up into permanganic acid, water, and manganese perhy- drate. From the composition of its salts we infer that it is a dibasic acid, having the above formula. (OK (-.XT- — This salt is obtained by the oxidation of any oxide of manganese in the presence of free alkali — for example, when any compound of manganese is fused with a mixture of potassium (or sodium) carbonate, and a little nitre. It may be obtained purer by the following process. Five parts of powdered caustic potash are mixed with a little water and 4 parts of powdered manganese peroxide stirred up with it ; this mixture is next heated nearly to boiling, and 3^ parts of powdered potassium chlorate added. After evaporating the mass to dry- ness, it is heated in a crucible to low redness for one and a half hours, and the cold mass extracted with a little cold water. The intensely green solution of potassium manganate so obtained deposits dark green, almost black, rhombic crystals of the salt when evaporated in a vacuum over sulphuric acid. The reaction is represented by the following equation : — 3Mn0 2 + 6KOH + C10 2 -OK = 3Mn0 2 (OK) 2 + KG + 3H 2 0. Potassium manganate is an unstable compound ; its aqueous H H 466 Text-Book of Inorganic Chemistry. solution changes on standing, especially when warmed, and be- comes red from formation of potassium permanganate — hence the name cha?neleon ?nineral sometimes given to the salt. Permanganic Acid: Mn0 3 -OH. — This monobasic acid is only known in solution, and is so unstable that it rapidly decomposes. On the other hand, the anhydride — permanganic anhydride : Mn 2 0, = M n o 3 f O; ma y De obtained by adding potassium perman- ganate in small quantities at a time to well-cooled strong sulphuric acid. It then separates from the olive-green liquid as drops of a heavy green oil, which soon decompose with evolution of oxygen. Potassium Permanganate: Mn0 3 -OK. — This compound, the best known of the permanganates, is obtained by boiling the green solution of potassium manganate, or better, by leading a stream of carbonic acid through it. In the former case there are produced potassium permanganate, manganese perhydrate, and caustic potash ; in the latter case, potassium carbonate is formed instead of the hydrate : — 3Mn0 2 (OK) 2 + 2H 2 = 2Mn0 3 -OK + Mn0 2 + 4KOH. 3Mn0 2 (OK) 2 + 2CO~ = 2Mn0 3 -OK + Mn0 2 + 2CO-(OK) 2 . For the preparation of the salt a solution of potassium man- ganate is warmed and a stream of carbonic acid led through it until the solution has become of a bright red colour. The clear solution is then run off from the precipitate and evaporated down, when potassium permanganate separates on cooling in long, dark- red, almost black, needles, with a metallic lustre. The salt is iso-morphous with potassium perchlorate. It is tolerably easily soluble in water ; at the ordinary temperature it requires 16 parts, but much less quantities at the boiling-point. Potassium permanganate is one of the most powerful oxidizing agents with which we are acquainted, and attacks a number of bodies which are not affected by other oxidizing substances. It liberates chlorine from hydrochloric acid, oxidizes oxalic acid to carbonic acid, ferrous salts to ferric salts, sulphurous acid to sulphuric acid, &c. If the oxidizing solution of potassium permanganate contains free sulphuric acid, manganous sulphate is formed, and five atoms of oxygen are set freefrom each two molecules of the permanganate for the purposes of oxidation ; but if the solution is neutral, hydrated Detection of Manganese Compounds. 4.67 manganese peroxide is produced, and only three molecules of oxygen are liberated from two molecules of the salt : — 2Mn0 3 -OK + 3SO s (OH), = 2S0 2 -0,Mn + SO,(OK) 2 + 3H,0 + 5O. 2Mn0 3 -OK + H,,0 = 2Mn0 2 +2KOH + 36. If finely powdered potassium permanganate is dropped into concentrated sulphuric acid, a green solution is obtained, and oxy- gen accompanied by considerable quantities of ozone is evolved. Owing to its powerful oxidizing properties; potassium permanga- nate cannot be filtered through paper ; the paper is at once turned brown, and a reduced ■ solution of a green or brown colour runs through. And the facility with which potassium permanganate oxidizes organic substances generally makes it well adapted for purposes of disinfection. A solution of the crude salt or of the sodium compound is largely used in England under the name of Condfs fluid. The other permanganates resemble that of potassium ; they are all of a red colour and soluble in water. Detection of Manganese Compounds. The amethyst colour imparted by a small quantity of a manga- nese compound to a bead of borax in the oxidizing flame of the blowpipe, and the disappearance of this colour when the bead is brought into the reducing flame, serve to detect the metal even in the presence of iron. Still more delicate is the green colour due to a manganate produced when a minute quantity of any man- ganese compound is fused on a piece of platinum foil with sodium carbonate and a little potassium nitrate. The test may be con- firmed by dissolving the green mass in water, acidulating and warming, when the solution becomes pink (permanganate). The flesh-coloured precipitate of manganous sulphide produced when ammonium sulphide is added to a solution of a manganous salt, and the darkening of the white precipitate (manganous hydrate) thrown down by ammonia on exposure to the air, are characteristic reactions of manganese compounds. But the colour of the sulphide is obscured by a small quantity of a salt of iron, which gives black ferrous sulphide with the ammonium sulphide. Ferric salts are at once distinguished from those of manganese by the white precipitate of manganous ferrocyanide, produced when a solution of potassium ferrocyanide is added to one of a manganous salt. Iron and manganese, which so often occur h h 2 468 Text-Book of Inorganic Chemistry. together, may be approximately separated by the addition of ammonium chloride and ammonia, which only precipitates ferric hydrate. The solution must be quickly filtered, otherwise man- ganic hydrate is produced by oxidation, and this is not dissolved by ammonium chloride. A more exact method consists in adding an excess of sodium acetate to the dilute solution of the two salts and then boiling, when the whole of the iron is precipitated as basic ferric acetate, while manganous acetate remains in solution. CHROMIUM. Chemical Symbol : Cr. — Atomic Weight: 52'2. Chromium, so-called on account of the various and beautiful colours of its compounds (xP^^ a - colour;, is not very widely distributed in nature, and never occurs in the metallic state. By far the most important mineral containing chromium, and that from which its compounds are always prepared, is chrome-iron ore : FeO,Cr,O s . Metallic chromium is only of scientific interest, and has not yet oeen applied to any useful purpose. It may be obtained, according to Wohler, by mixing one part of violet chromic chloride with two parts of a fused mixture of potassium and sodium chlorides, and heating this with two parts of granulated zinc in a Hessian crucible under a layer of potassium-sodium chloride. The crucible is heated up to the boiling-point of zinc, recognized by the bubbling noise, kept at this temperature for ten minutes, then removed from the furnace and knocked several times on the floor to cause the metal to unite at the bottom of the crucible. On breaking open the crucible when cold, a regulus of an alloy of zinc and chromium is obtained, which is treated with dilute nitric acid until all the zinc is dissolved. The chromium then remains as a pale grey, crystalline powder. According to Bunsen, chromium may be obtained by electrolyzing a concentrated solution of chromic chloride. It then forms brittle lustrous scales, of the colour of iron. The metal is not oxidized by moist air so rapidly as iron ; its specific gravity is about 6'5. When heated in the air it becomes covered with a thin film of the oxide, imparting to it various colours, like steel ; it is only slowly oxidized in dry air, even at a red-heat. Hydrochloric acid dissolves the metal with evolution of hydrogen ; Chromium. 469 sulphuric acid acts similarly on warming. Nitric acid, even when hot and concentrated, does not act on it. Chromium, like manganese, forms several compounds with oxygen. We are acquainted with — chromous oxide : CrO, chromic oxide : Cr 2 3 , chromic anhydride : Cr0 3 , and fierchromic acid : HCr0 4 (?). The first and last of these compounds are unstable, and scarcely known. Chromic oxide, like alumina and other sesquioxides, is a weak base, while chromic anhydride possesses strong acid properties. CHROMIC OXIDE AND CHROMIC SALTS. Chromic Oxide : Cr„0 3 . — This compound, which is isomor- phous with alumina and ferric oxide, is amorphous or crystalline according to its mode of preparation, but always possesses a green colour. Chromic oxide may be obtained by a variety of methods, of which the following are the most important : — (i.) By heating chromic hydrate. (ii.) By heating mercurous dichromate to redness : — °{c!&OhI = Cr »°3 + 2U S + 20 2 , when it remains as a dark green powder. (iii.) By igniting ammonium dichromate : — °{crOl.oSr?: - Cr »°» + 4H.0 + N„ when the chromic oxide remains behind in olive green scales. (iv.) By very strongly heating a mixture of common salt and potassium dichromate covered with a layer of common salt in a Hessian crucible. On extracting the fused mass with water, the chromic oxide remains as glistening iridescent crystals : — °|cS;Ck - c '.°. + K *° + 30, the common salt only serving to assist the crystallization of the chromic oxide. (v.) Finally, by passing the vapour of chromyl chloride through a porcelain tube heated to redness : — 2Cr0 2 Cl 2 == Cr 2 3 + 2CL, + O, 470 Text-Book of Inorganic Chemistry. when chlorine and oxygen are given off, and the chromic oxide is deposited on the walls of the tube in lustrous dark-green crystals, which are so dark in colour as to appear almost black. These crystals are as hard as corundum, and have a specific gravity of 5 2. Strongly ignited chromic oxide is insoluble in acids, even in concentrated sulphuric acid. It may be heated either in the air or in hydrogen to the highest temperatures of our furnaces without change ; but is reduced when strongly heated with charcoal. When heated with potassium (or sodium) carbonate and exposed to the air, or with the addition of a little nitre, chromic oxide gives soluble potassium (or sodium) chromate. Insoluble compounds containing chromium — e.g. chrome-iron ore — are thus brought into solution. Chromic oxide easily dissolves in molten glass as well as in borax, and the chromic silicate or borate so produced possesses a beautiful gretn colour. On account of this property and because it remains unchanged at a high temperature, chromic oxide is largely used for colouring glass green and for painting on porce- lain. Chromic oxide, like alumina and other sesquioxides, is a weak base. It unites with acids and with bases, forming, in the latter case, compounds which correspond to the aluminates. These compounds, sometimes called chromites, include chrome-iron ore : FeO,Ci\,0 3 = ( . ,-, 0., Fe", referred to above, and which is distin- guished by the difficulty with which it is attacked by most acids. Among other chromites are the soluble green compounds pro- duced by digesting freshly precipitated chromic hydrate with caustic potash or soda. These compounds are, however, less stable than the corresponding aluminates, and are decomposed on boiling their solutions. It is remarkable that chromic hydrate when mixed with ferric hydrate is not dissolved by caustic al- kalies. chromic Hydrate : Cr(OH) 3 . — This substance separates when a solution of a chromic salt is mixed with ammonia as a bluish-green bulky precipitate, which when dried forms a dirty-green powder. If dried at ioo° it possesses the above composition. When gradually heated more strongly it begins to glow throughout its entire mass at one particular temperature, and is then converted into dark-green chromic oxide, which is now insoluble in acids. Freshly precipitated Chromic Sulphate. 471 chromic hydrate easily dissolves in acids, and produces two kinds of differently coloured chromic salts ; the solutions of the one kind are green, and those of the other violet. The violet salts change into green by continued boiling, but on allowing the latter to stand for some time a change takes place from the green to the violet. The violet salts crystallize -easily ; but the green solutions, when evaporated down, leave a green amorphous mass. 1 Chromic hydrate, precipitated by ammonia from a solution of a violet salt, is somewhat soluble in an excess of the reagent, but is rendered completely insoluble on boiling. A hydrate of chromium of a fine green colour is known in trade as Guignet's green, and has the composition : Cr 2 O s , 2H 2 = ^"Mr^OJ-n *' ' s 0Dtame d by fusing a mixture of potassium dichromate and crystallized boric acid, and then extracting the fused mass with water. During the reaction, potassium and chromic borates are produced, and the latter compound is decom- posed when treated with water. Chromic Sulphate: (S0 2 ) 3 6 Cr,, + I5H 2 0. — This salt crys- tallizes in violet octahedra, when chromic hydrate is dissolved in sulphuric acid and allowed to stand for some time. If its solution is mixed with potassium sulphate, the two sulphates unite to form chrome-alum : „„' [ „?, r + I2H 2 0, which crystallizes out in fine dark -purple octahedra when the solution is allowed to evaporate slowly. Chrome-alum is, however, best prepared by the reduction of a mixture of potassium dichromate and sulphuric acid by some easily oxidized substance, such as sulphurous acid or alcohol. The chromic acid which is thus set free by the sulphuric acid parts with oxygen, and is reduced to chromic oxide. One of the best reducing substances to employ is ordinary alcohol, which then becomes oxidized to the volatile compounds — aldehyde and acetic acid. To obtain chrome-alum in this way, 2 parts of a saturated solution of potassium dichromate are mixed with 3 parts of con- centrated sulphuric acid in a large flask, and alcohol added drop by drop from a separating funnel, with continued agitation, until ' The chemical difference between the violet and green modifications is at present unknown. It has sometimes been thought that the green varieties contain a mixture of an acid and basic salt. — Ed. 47 2 Text-Book of Inorganic Chemistry. the red liquid has changed to a dark green. On allowing the green liquid to stand for some time in an open dish it gradually becomes of a violet colour, and, in consequence of evaporation, small octahedra of chrome-alum begin to be deposited. If these small crystals are allowed to lie in the solution, which is kept in a cellar where the temperature is nearly uniform, and are turned daily, they gradually grow to large regular octahedra. The above reaction takes place according to the following equation : — .-.(CrO^-OK c „ ,,-.,tt\ SO„) 0,Cr lTJ „ -.' °(Cr(VOK + 4&O a (OH) a = ^so^ok + 4 ^° * 30, in which the oxidation of the alcohol by the three atoms of oxygen is not shown. As in ordinary alum, the potassium in chrome-alum may be displaced by any of the alkali-metals, or by thallium, without changing its crystalline form. Chromic Nitrate : (N0 2 ) 3 3 Cr + 9H 2 0, crystallizes with diffi- culty from the violet solution obtained by dissolving chromic hydrate in nitric acid. Chromic Phosphate : PO ■ 3 Cr + 6H 2 0, falls as a dirty-green precipitate when a solution of sodium phosphate is added to one of chrome-alum. It becomes crystalline on standing under the liquid. Chromic Chloride : CrXl e . — This chloride sublimes in beau- tiful peach-blossom coloured scales, when a mixture of chromic oxide and charcoal is heated in a stream of dry chlorine ; it is distinguished from the corresponding chlorides of aluminium and iron by its insolubility in water. When heated in the air it loses chlorine, and takes up oxygen, being first converted into an oxy- chloride and then into chromic oxide. A simple method of preparing considerable quantities is as follows. The dried mixture of chromic oxide, charcoal, and starch paste is placed in a Hessian crucible, on which is cemented a second crucible of similar size placed upside down, and through the bottom of which a hole is bored to receive the glass tube bringing the chlorine. The two crucibles are placed in a char- coal furnace so that the lower one is surrounded by the glow- ing charcoal, and is heated to redness, while the upper one is only moderately heated. Chlorine is then led into the upper crucible through a glass tube which fits the opening loosely, and Chromic Chloride. 473 the lower crucible is raised to a bright red-heat. The carbonic oxide and the excess of chlorine escape through the opening at the top, and the chromic chloride condenses on the sides of the upper and cooler crucible : a portion of the chloride remains in the lower crucible loosely attached to the carbonaceous residue, from which it may be easily removed when cold. Chromic chloride when pure is quite insoluble in water, but dissolves at once to form a green liquid when the water contains a trace of crystalline chromous chloride. Mere contact with the chromous chloride suffices to convert the insoluble violet chromic chloride into the soluble green variety. By dissolving chromic hydrate in hydrochloric acid and eva- porating down, a green hydrated chromic chloride is obtained, which loses water and hydrochloric acid when heated, and leaves a basic chromic chloride behind. Chromous Chloride : CrCl 2 . — If chromic chloride is heated to low redness in a stream of hydrogen it loses chlorine, and is con- verted into a white crystalline mass of chromous chloride. This chloride dissolves in water to a blue liquid, which then absorbs oxygen when exposed to the air, and produces an oxychloride. Chromous chloride, like ferrous chloride, can dissolve considerable quantities of nitric oxide, forming a dark-brown solution. The merest trace of this chloride can convert considerable quantities of the violet chromic chloride into the soluble green modification (see above). When caustic potash is added to a solution of chromous chloride, a brown precipitate is produced, which is probably of chromous hydrate : Cr(OH) 2 , but which possesses so strong an attraction for oxygen, that it decomposes water and forms tri- chromic tetroxide : Cr 3 4 , with evolution of hydrogen. CHROMIC ACID AND THE CHROMATES. Chromic Anhydride : Cr0 3 . True chromic acid is unknown in the free state ; when separated from its salts it breaks up into water and chromic anhydride. 474 Text-Book of Inorganic Chemistry. Chromic anhydride, commonly known as chromic acid, is easily obtained by adding i£ volume of concentrated sulphuric acid to i volume of a saturated solution of potassium dichromate at 5o c . The mixture becomes strongly heated, and deposits on cooling long bright red prisms of the chromic acid liberated by the sulphuric acid. When quite cold the mother-liquor, containing acid potassium sulphate, is poured off, and the crystals dried first on an asbestos filter, and then by placing them on a porous plate of unglazed porcelain, which absorbs the remaining traces of water. By recrystallizing from a small quantity of water and repeating the drying process, the crystals . may be obtained pure and free from •sulphuric acid. Chromic acid is easily soluble in water and deliquesces in moist air. The solution is of a reddish-brown or yellow colour according to its strength, and reacts strongly acid. Chromic acid contains one half of its oxygen very loosely combined, and is decomposed into chromic oxide and oxygen not only when heated alone : 2Cr0 3 = Cr„0 3 + 3O, but also when heated with sulphuric acid, chromic sulphate being formed in this case. In the same way it evolves chlorine with strong hydrochloric acid : — 2CrO a + 12HCI = Cr.,Cl a + 6H 2 + 3CL/ It oxidizes organic substances so powerfully that few can with- stand its action. If the solution is filtered through paper, or if the crystals are dried with paper, the chromic acid at once becomes of a green colour from chromic oxide. Alcohol when dropped on the crystalline substance is so rapidly oxidized that it catches fire. If a solution of chromic acid is mixed with hydrogen peroxide a compound of a beautiful blue colour is formed, which easily dissolves in ether. It has been suggested that this compound is perchromic acid, corresponding to permanganic acid, but its com- position is doubtful (p. 91). The Chromates. (OTT (-.t, is, like sulphuric acid, a dibasic acid, but is distinguished from this and all other 1 An excellent method of obtaining a long-continued, regular stream of chlorine consists in gently heating potassium dichromate with an excess of hydrochloric acid. — Ed. Chromic Acid and the Chromates. 475 dibasic acids by the fact that it never produces acid salts. That acid salts of chromic acid might be produced cannot be doubted, but we are as yet unacquainted with the conditions under which they can exist. They decompose as soon as produced into the (CYO -OH r- ^ 2 Vvu and water. If, for Cr0 2 -OH ' f OK example, normal potassium chromate : Cr0 2 i „„ is mixed with a solution of chromic acid, or with a stronger acid, then it is probable that acid potassium chromate is first produced : — r n (OK r n (OH „ r n (OH CrC M0K + Cr0 *|0H = 2CrC MoK or r n i 0K en !°H r n jOH cri (OH Cr0 2 | OK + SO MQH = Cr0 2 | QK + S0 2 | QK and that two molecules of the acid chromate immediately afterwards break up into the normal dichromate and water : — 2Cr0 f° H - 0- ,Cr0 *- 0K + HO 2UU ^OK ~ U lCr0 2 -OK + n * U - These soluble dichromates are again converted into the normal chromates on the addition of a base : — o (cS:Ck + 2K0H = 2Cr °4CI + H *°- Similar relations to that between chromic acid and dichromic acid also exist between sulphuric and disulphuric acids (p. 167), and between phosphoric and pyrophosphoric acids (p. 219) ; the only difference being that the normal salts of disulphuric and pyrophos- phoric acids are produced from the acid sulphates and acid phos- phates respectively at a high temperature, while the change from the acid chromates to the normal dichromates takes place at the ordinary temperature. It is possible that acid chromates may exist at a low temperature. Potassium Dichromate : K a Cr 2 7 = O -j (- q 2 . q lj — This salt is manufactured on a large scale from chrarae-iron ore : FeO, Cr^Oj, and is the starting-point for the preparation of all other chromium compounds. It is obtained when the finely powdered ore is mixed with potassium carbonate and nitre and fused for some time. The iron is then oxidized to ferric oxide and the chromium 476 Text-Book of Inorganic Chemistry. to chromic acid, which unites with the potassium and forms yellow potassium chromate. The latter compound goes into solution on digesting the fused mass with water. On the large scale the mixture of the pulverized ore and potassium carbonate, without nitre, is heated in a reverberatory furnace and the oxidation effected by excess of air. A certain quantity of lime is usually added, which keeps the mass porous and prevents complete fusion. The yellow solution of potassium chromate obtained by afterwards extracting the mass with water is then acidulated with sulphuric acid, which both destroys any excess of potassium carbonate and converts the normal chromate into red potassium dichromate, and the solution is allowed to crystallize. The chromate is converted into the dichro- mate because the latter salt crystallizes much more easily from its solutions, and can, therefore, be more readily purified. Potassium dichromate, purified by recrystallization, forms orange-red tabular crystals of the above composition, and without water of crystallization. It dissolves in ten parts of water at the ordinary temperature, but is more soluble in hot water ; the solu- tion reacts faintly acid. When heated to low redness it melts unchanged to a dark red liquid, and at very high temperatures is decomposed. Concentrated sulphuric acid separates crystalline chromic acid from its aqueous solution. Sulphuretted hydrogen reduces a solution of potassium dichromate and produces a preci- pitate of chromic hydrate and sulphur, while the potassium is converted into sulphide or sulphydrate, which remains in solu- tion : — O jcrO^.'oK + 4H * S = 2 Cr (° H )3 + 3S + K,S + H 2 0. No precipitate of chromic hydrate is produced if hydrochloric or sulphuric acid is previously mixed with the solution of potassium dichromate ; the red colour of the solution then changes to a green from the chromic chloride or sulphate produced. If a solution of potassium dichromate is mixed with caustic potash or potassium carbonate until the red colour has become a pure yellow, and the solution evaporated down, Potassium Chromate : CrO a (OK) 2) separates out in yellow crystals, isomorphous with potassium sulphate. The salt is very easily soluble in water, and cannot therefore be so easily crys- tallized as the dichromate. Its yellow solution reacts slightly alkaline, and becomes red on the addition of an acid owing to the Chromic Acid and the Chromates. 477 formation of potassium dichromate, which in its other properties it closely resembles. Sodiwii Chromate. — Both sodium chromate and dichromate closely resemble the corresponding potassium salts, except that they are much more soluble in water, and even deliquesce in the air. Ammonium Dichromate : Oi r .,-.- ' -.», „ l crystallizes in yellow scales, and otherwise resembles the potassium salt. It is easily soluble in water, and gives green chromic oxide when heated (p. 469). Ammonium Chromate : Cr0 2 (ONH 4 )„ crystallizes in yellow needles on evaporating a mixture of aqueous chromic acid and excess of ammonia. The crystals remain unchanged in the air. Calcium and Strontium Chromates crystallize in yellow prisms, and are much less soluble in water than the potassium salt. Barium Chromate : Cr0 3 • 2 Ba, is nearly insoluble in water, and separates as a pale-yellow precipitate when a solution of potassium chromate or dichromate is added to one of barium chloride. It is used as a yellow pigment. The chromates of the heavy metals are coloured substances, insoluble in water. Chlorcnromic Acid and Chromyl Chloride. Chromic acid, which resembles sulphuric acid in so many (CI points, also forms a compound, chromyl chloride : CrO J r , corre- 'Cl spondingto sulphuryl chloride : SO.,-, c , in which the two atoms of hydroxyl of the hypothetical acid are displaced by chlorine. A (CI second compound, corresponding to chlorsulphomc acid : SO, j qu and in which only one atom of hydroxyl is displaced, is known in the salts of chlorchromic acid — e.g. potassium chlorchromate : CrO '■ k T .r The acid itself has not yet been prepared. To obtain this potassium salt, 3 parts of finely powdered potassium dichro- mate are mixed with water to a paste, then gently warmed with 4 parts of concentrated hydrochloric acid until a faint odour of 478 Text-Book of Inorganic Chemistry. chlorine is perceptible. The salt afterwards separates in large red, tabular crystals as the clear, red solution cools : — °|crC::Sl + 2HC1 - 2Cr °.(0K + H *°- The salt is decomposed by water, but dissolves in hydrochloric acid unchanged. When gently heated it evolves chlorine copiously. fCl Chromyl chloride : CrO, \ ~, is a blood-red, heavy liquid of i '9 specific gravity, and boiling at 116°. It possesses a powerful piercing odour, and fumes in the air. In order to prepare this compound, 10 parts of common salt are fused together with 12 parts of potassium dichromate, the fused mass poured on to a clean iron plate and broken into pieces when cold. It is then in- troduced into a large retort connected with a condenser, and 40 parts of concentrated sulphuric acid, containing some Nordhausen acid, added. A brisk reaction ensues, which must afterwards be supported by a gentle heat, and the compound distils over into the well-cooled receiver. It is purified by repeated distillation, finally in a stream of carbonic acid gas. The reaction is represented by the following equation : — °{ck>;:8k + 4NaC1 + 6S °.i§h - 2Cr0 ^ + 2S0 2 (°K + 4 S0 2 {g£a + 3 h A According to which 1 molecule of potassium dichromate and 4 molecules of sodium chloride give, when heated with an excess of sulphuric acid (about 9 molecules), 2 molecules of chromyl chloride, 2 of acid potassium sulphate, 4 of acid sodium sulphate, and 3 of water, for the retention of which at least 3 molecules of sulphuric acid are required. Chromyl chloride is decomposed by water into chromic and hydrochloric acids. It acts energetically on phosphorus, sulphur, and many organic substances with the production of light and heat. When its vapour is led through a red-hot tube it is decomposed into chlorine, oxygen, and chromic oxide. If the requisite quantity of chromyl chloride is dropped into a solution of normal potassium chromate, containing a little acetic acid, the liquid becomes filled with crystals of potassium chlorchromate : — Cr0 2 Cl 2 4 Cr0 2 (OK) 2 = 2Cr0 2 {g K Detection of Chromium Compounds. 479 Detection of Chromium Compounds. Compounds of chromium, in any form, may be readily and cer- tainly detected by simple blowpipe tests. When heated with sodium carbonate and a little nitre on a piece of platinum foil, they all produce yellow sodium chromate, and the fused mass, when dis- solved in water and acidulated with acetic acid, gives a yellow pre- cipitate of lead chromate on the addition of a few drops of a solution of lead acetate. A trace of a chromium compound imparts a green colour to a borax bead (due to chromium borate), both in the oxidizing and reducing flame, but which becomes converted into yellow sodium chromate when heated for a long time in the oxidizing flame. An aqueous solution of a chromate is recognized by the yellow precipitate of lead chromate produced on the addition of a solution of lead acetate ; and by the change of the yellow or red solution to a green salt of chromic oxide when acted upon by sulphuretted hydrogen, sulphurous acid, alcohol, or other reducing agents in the presence of hydrochloric or sulphuric acid. The soluble salts of chromic oxide all yield chromic hydrate with ammonium sulphide, ammonia, or caustic soda, and are not precipitated by sulphuretted hydrogen. In this respect the salts of chromic oxide resemble those of alumina. Freshly precipitated chromium hydrate, like aluminium hydrate, is easily dissolved by caustic soda, but the hydrate is reprecipitated on boiling, or when allowed to stand. Compounds of chromic oxide may therefore be separated from those of alumina by adding excess of caustic soda and boiling ; they may be separated from salts of iron by fusing with sodium carbonate and nitre, when the iron is converted into insoluble ferric oxide, while the chromium takes the form of a soluble chromate, which may be extracted by digesting the fused mass with water. URANIUM. Chemical Symbol : U. — Atomic Weight: 240. This somewhat rare element is never found free in nature ; it occurs almost exclusively as the oxide : U 3 O s = UO.,,2U0 3 , in the mineral pitchblende. The metal may be obtained by fusing 480 Text-Book of Inorganic Chemistry. uranous chloride with sodium, best in an iron crucible, and then forms either a black powder or a hard metallic button of a grey colour. It melts at a red heat, and has a specific gravity of 187. The metal does not decompose water even at 100° but oxidizes when heated in the air, and is easily dissolved by dilute acids. Uranium unites in two proportions with oxygen, forming itra- nous oxide : U0 2 , and uremic oxide : U0 3 , and a third oxide also exists of the composition : U 3 8 , which may be considered as a compound of the first two: U0 2 + 2U0 3 . Pitchblende consists chiefly of this oxide, but contains also compounds of lead, copper, arsenic, iron, manganese, cobalt, and other metals. To obtain pure uranium compounds from this mixture the following process may be adopted. The powdered mineral is heated with mode- rately concentrated sulphuric acid and small quantities of nitric acid added from time to time. When only a white insoluble residue (of lead sulphate, &c.) remains, the liquid is evaporated down to expel most of the acid, then diluted with water and the clear solution saturated with sulphuretted hydrogen to precipitate lead, copper, arsenic, &c. The solution is next filtered off from these sulphides, the iron oxidized again by nitric acid, and ammonia added in excess. The dark-yellow precipitate so obtained, which chiefly consists of ferric hydrate and ammonium uranate, is then well washed and treated with a warm concentrated solution of ammonium carbonate containing free ammonia, until all the uranium is dissolved and the precipitate consists entirely of dark-brown ferric hydrate. The yellow solution, which must be quickly filtered, deposits small, hard, yellow crystals of ammonium- uranic carbonate, and a further deposit of the same salt is obtained as a yellow powder when the solution is boiled. This compound when heated to redness leaves the pure dark-green oxide : U 3 8 . Uremic Nitrate: N( -. 2 \ 2 (U0 2 ) + 6H 2 0, is obtained in large yellow deliquescent prisms when the preceding or any other oxide of uranium is dissolved in nitric acid and the solution evaporated down. If the salt is heated to 250° until nitrous fumes are no longer evolved, it becomes changed into Uranic Oxide: U0 3 = (U0 2 )0, which remains behind as a yellow powder, and which when strongly heated loses oxygen and reproduces uranoso-uranic oxide : U 3 8 . Uranic Sulphate: SO„'0 2 (U0 2 ) + 3H 2 0, may be obtained by Uranium. 48 1 decomposing the nitrate with sulphuric acid, and crystallizes in pale yellow needles. Uranic Carbonate, — Ammonium carbonate, when added to a solution of a uranium compound, produces a yellow precipitate, which, however, is not uranium carbonate, but a double carbonate of ammonium. This salt is soluble in an excess of ammonium carbonate, and if this solution is allowed to evaporate in the air, or if gently warmed, it deposits small yellow glistening crystals of the compound: CO-0,(UOo) + 2CO-(ONH 4 ) 2 , which is only slightly soluble in water. This double salt is used to separate uranium from other metals. Uranic Phosphate : PO j q'4 -' + 4H,0, is thrown down as a yellow, crystalline precipitate on adding a solution of phosphoric acid or a soluble phosphate to one of uranic acetate. It is in- soluble in acetic acid. If the solution of uranium contains ammo- nium chloride, the precipitate which is produced on the addition of a phosphate has the composition : PO ~?^ T ~J « This com- (OJN H 4 pound is also insoluble in acetic acid, and is used to estimate phos- phoric acid, especially in manures. The mineral autunite is a double phosphate of uranium and calcium. From the composition of the uranium compounds given above, it is clear that uranic oxide is not a triacid base like alumina, ferric oxide, and chromic oxide, but that it always combines only with two molecules of a monobasic acid, like lime, magnesia, and the oxides of other dyad metals. Two atoms of oxygen are firmly united with an atom of uranium to form a dyad radical : UO.,, called uranyl, and the salts of the oxide — e.g. uranic nitrate, are properly called uranyl salts — e.g. uranyl nitrate. This rule, however, is not followed, because no compound of the composition : (N0 2 ) 6 6 U, is known. If, then, uranic oxide : U0 3 , is to be considered as uranyl oxide : (UO.JO, we should expect that on dissolving this oxide in hydrochloric acid a chloride of the composition : (UO„)Cl 2 , would be obtained, according to the equation : — (U0 2 )0 + 2HCI = (U0 2 )C1 2 + H 2 0, and as a matter of fact the compound : Uranyl chloride, which 1 1 482 Text-Book of Inorganic Chemistry. crystallizes with difficulty from the yellow solution obtained by dissolving uranic oxide in hydrochloric acid, has this composition. Uranic oxide unites with strong bases to form a. series of com- pounds called uranates, in which it plays the part of an acid. All these salts, even those of the alkalies, are insoluble in water. Potassium nra?iate : O-jjj.-. 2 ^^ corresponding to potassium dichromate, is obtained as a yellow precipitate on adding caustic potash to a solution of uranium nitrate. It bears a high tempera- ture without decomposition. Sodium ura?iate and Ammonium uranate closely resemble the potassium salt, and are obtained by adding caustic soda or ammonia to a solution of uranium nitrate. Ammonium uranate dissolves in ammonium carbonate in the cold, and when ignited in the air leaves uranoso-uranic oxide as a green powder. Uranyl Sulphide : (U0 2 )S, is precipitated as a black amor- phous powder, nearly insoluble in water, when ammonium sulphide is added to a solution of an uranic salt. It is soluble in ammonium carbonate, and is therefore not thrown down if the solution to which the ammonium sulphide is added contains this salt. Uranous Oxide : UO„, which is produced by glowing uranic oxide or uranoso-uranic oxide in a stream of hydrogen, is a black or dark brown crystalline powder, insoluble in dilute hydrochloric acid, but dissolving in nitric acid and in concentrated sulphuric acid. If a solution of uranous oxide in sulphuric acid, which is of a green colour and contains uranous sulphate : (SO„). 2 4 U, is mixed with caustic soda, a brown precipitate of uranous hydrate : U(OH) 4 , is produced, which turns black on boiling. Vranous Chloride : UC1,, may be prepared by igniting a mixture of any oxide of uranium and charcoal in a stream of chlorine, and then sublimes in lustrous dark green octahedra. It is readily soluble in water and deliquesces in moist air. Uranoso-uranic Oxide : U 3 8 = U0 2 ,2U0 3 , is obtained as a dark green powder, by glowing either uranous or uranic oxide in the air. In the former case oxygen is absorbed, in the latter it is given off. It is easily soluble in nitric acid, but only dissolves in Cobalt and Nickel. 483 sulplvuric or hydrochloric acid with difficulty, and withstands a high temperature without change. Glass (or a borax bead) is coloured a bright green by this oxide, while uranic oxide imparts a greenish yellow tint, which fluoresces strongly in bright light. COBALT AND NICKEL. These twin metals, as they may be called, always occur in association with one another in nature, the one being in excess in some minerals, and the other in others, Both are constant constituents of metallic meteorites, though the nickel is always in excess. The two metals have nearly the same atomic weight, and their compounds are so much alike, chemically, that for a long time no method was known of separating them. Their salts are chiefly distinguished from one another by their colour : those of cobalt are red or blue, and those of nickel green. The similarity between the two metals extends even to their names. The external appearance of the cobalt ores led to the belief that a valuable metal might be extracted from them by the ordinary metallurgical processes. But when all attempts in this direction proved futile, the German miners imagined the ores to be possessed of an evil spirit or Kobold ( = sprite, goblin), and this word then gave the name cobalt to the metal afterwards discovered in the ores. In the same way, many attempts were made to ex- tract copper from the commonest ore of nickel — Kupfer-nickel — which is itself of a copper colour ; but when all these endeavours were in vain the ore was thrown on one side and called Nickel (an abusive word for an obstinate person in some South German dialects), a name which has now been retained for the metal itself. Of the two metals, cobalt was first used in the arts, and in the form of its oxide, after it had been discovered that this compound dissolves in molten glass with a dark-blue colour, and that such glass when finely powdered yields a fine blue pigment {smalt). In the manufacture of smalt the nickel compounds are separated under the name of cobalt-speiss, and this was formerly thrown away as useless. At one time this speiss was used in Hesse for mending the roads, but when its value became known the road scrapings were carefully collected and worked for nickel. 484 Text-Book of Inorganic Chemistry. Cobalt is only used technically in the form of its oxide, not as the metal ; while nickel, on the other hand, is almost exclusively used in the metallic state for the preparation of alloys and for nickel-plating. COBALT. Chemical Symbol : Co. — Atomic Weight : 59. Cobalt is only found native in metallic meteorites, and then in very small quantities. The most important minerals which contain the metal are smaltite or sfieiss-cobalt : CoAs„, and glance-cobalt : CoS 2 ,CoAs. 2 = CoAsS. From these ores, which usually contain copper, iron, and nickel, cobalt or a pure cobalt compound may be obtained by a variety of methods. To prepare cobalt from either of the ^.bove ores on the small scale, the following pro- cess may be adopted. The ore is first freed from some of its sulphur and arsenic by roasting in the air, then dissolved in hydrochloric acid, with the addition of small quantities of nitric acid, until completely decomposed. Sulphuretted hydrogen is next passed through the acid liquid to precipitate arsenic, copper, lead, &c, and after the iron has been again oxidized by nitric acid, a solution of sodium carbonate added until just neutral, by which all the iron is precipitated as ferric hydrate. The filtered solution, on further addition of sodium carbonate, gives a precipitate of the mixed carbonates of cobalt and nickel, which, when washed and digested with a strong solution of oxalic acid, are converted into the insoluble oxalates, both of which dissolve in strong ammonia. This solution of the oxalates in ammonia is then allowed to stand in the air for some time, so that the ammonia may slowly evaporate, when the whole of the nickel is precipitated as oxalate, while the cobalt remains in solution. The clear purple solution, which now only contains cobalt, is evaporated to dryness, the residue glowed, and then dissolved in hydrochloric acid, and after the excess of hydrochloric acid has been removed by heating, precipitated with a solution of oxalic acid. The cobalt ( CO oxalate : \ r( -, OXo, so obtained is decomposed when heated into carbonic acid and cobalt, the latter fusing at a sufficiently high temperature. To reduce the metal from the oxalate, a narrow, Compounds of Cobalt. 485 deep, unglazed, porcelain crucible is filled with the dry salt, then covered, placed in a larger Hessian crucible lined with charcoal, and exposed to a high temperature in a wind furnace. The metal then melts to a steel-grey button, with a tinge of red. It has a specific gravity of 8-6, is malleable, and can be highly polished. It is only fused with difficulty, is more compact than iron, and is slightly magnetic. Cobalt is less readily oxidized than iron, but when heated in the air becomes covered with a film of the-oxide ; it dissolves in hydrochloric acid and dilute sulphuric acid, but less readily than iron, and then produces cobaltous chloride and sul- phate respectively ; nitric acid also dissolves it, forming, cobaltous nitrate. Cobalt unites in two proportions with oxygen, forming cobaltous oxide : CoO, and cobaltic oxide : Co 2 3 . The former is a strong base, while the latter possesses more the character of a peroxide. Cobaltous Oxide : CoO ; Cobaltous Hydrate : Co(OH) 2 . When caustic soda is added to a solution of cobaltous chloride, a avender coloured precipitate of a basic chloride is formed ; the precipitate changes to pink cobaltous hydrate when boiled, but soon absorbs oxygen from the air, and becomes of a brownish colour. When this precipitate is heated out of contact with the air, it loses water and is converted into cobaltous oxide, forming a dull green powder. The same compound is produced when cobaltous carbonate is heated out of contact with air. When heated in the air, cobaltous oxide is converted into the three-quarter oxide — tri- cobaltic tetroxide : Co 3 4 , corresponding in composition to the magnetic oxide of iron. Cobaltous oxide dissolves in hydrochloric, sulphuric, or nitric acid, and produces bright red solutions. Cobaltic Oxide : Co.,0 3 ; Cobaltic Hydrate: Co(OH) 3 . Cobaltic oxide is a black powder, insoluble in water, and is obtained by carefully heating cobaltous nitrate or cobaltic hydrate. The hydrate is produced when chlorine is led into an alkaline solu- tion containing cobaltous hydrate in suspension. Such a solution is easily obtained by precipitating a cobaltous salt with excess of caustic soda. The reaction is as follows : — Co(OH)., + NaOH + CI = Co(OH) 3 + NaCl. Cobaltic hydrate may also be obtained by precipitating a solu- tion of a cobaltous salt with sodium hypochlorite, or a solution of bleaching powder. Like the oxide, it is also of a black colour, and 486 Text-Book of Inorganic Chemistry. both compounds dissolve easily in hydrochloric acid. This solution appears at first to contain cobaltic chloride : CoCl 3 , as but little chlorine is set free in the cold ; if, however, the liquid is gently warmed, chlorine is liberated in abundance, and the solution now contains red cobaltous chloride. Cobaltic oxide, when strongly heated in the air, loses oxygen and is converted into the oxide, Co 3 4 . Cobaltous Sulphate : S0 3 -0 2 Co + 7H„0, is easily soluble in water, and separates from its aqueous solution in dark-red crystals, isomorphous with ferrous sulphate. Cobaltous Nitrate : (N0 2 ) 2 2 Co + 6H,0, crystallizes with diffi- culty in dark-red prisms, which deliquesce in the air. When heated it easily parts with oxygen and all its nitrogen as nitric peroxide, and leaves a residue of cobaltic oxide, which becomes cobaltoso- cobaltic oxide : Co 3 Oj, when more strongly heated. Cobaltous Nitrite is an unstable compound, and has not been prepared pure. If a neutral aqueous solution of a cobaltous salt is mixed with one of potassium nitrite in excess, a yellow, crystal- line precipitate soon separates, which is a double nitrite of potas- sium and dyad cobalt, and has the composition : — (NO)„O..Co" + 2NO-OK + H 2 0. This compound is slightly soluble in cold water, and dissolves more readily when warm, forming a red solution. A second double nitrite, but containing triad cobalt, is obtained when potassium nitrite is added in excess to a solution of a cobal- tous salt strongly acidified with acetic acid. A double salt of the composition : — (NO) 3 3 Co'" + 3NO-OK, with varying quantities of water, then separates also in the form of a yellow crystalline precipitate. This compound is only slightly soluble in water, but quite insoluble in water containing a small quantity of potassium nitrite or some other potassium salt. The corresponding nickel compound is easily soluble in water, and potassium nitrite is therefore used to separate cobalt and nickel. Cobaltous Phosphate is thrown down as a bright-red precipitate on adding a solution of sodium phosphate to a cobaltous salt. Cobaltous Carbonate. — When sodium carbonate is mixed with Compounds of Cobalt. 487 a cobaltous salt a pink precipitate is produced. This is not the normal carbonate, but a basic salt of the composition : 2CO • O .Co + 3Co(OH), + H a O. ( CO Cobaltous Oxalate : j p^O^Co, is slightly soluble in water, and separates when oxalic acid is added to a cobaltous salt as a pink crystalline powder. CobaltDus Chloride : CoCl., + 6H..O, crystallizes on evaporating its aqueous solution in fine red prisms. When heated the salt loses water, and becomes of a blue colour. The same change is produced by concentrated hydrochloric acid. It forms a crystalline double salt with ammonium chloride of the composition : CoCl 2 + NH 4 C1 + 6H.,0. Cobaltic chloride, which in the free state is so unstable, unites with ammonia and forms a number of stable and most remarkable compounds, whose constitution still remains an unsolved problem. If an aqueous solution of cobaltous chloride is mixed with an excess of strong ammonia, the clear liquid exposed to the air until its brown colour has changed to a red, and then mixed with con- centrated cold hydrochloric acid, a brick red crystalline powder of the composition : CoCl„5NH 3 + H,0, and called Roseocobaltic chloride, separates out. The cobaltous chloride by standing in the air has become oxidized to a compound of triad cobalt. If the solution is precipitated with sulphuric acid, instead of with hydrochloric acid, a bright red crystalline pre- cipitate of roseocobaltic sulphate is produced. The dark red solution of roseocobaltic chloride becomes of a purple colour when boiled, and then deposits small, lustrous, purple crystals of Purpureocobaltic chloride : CoCl 3 ,5NH 3 , on cooling. This compound differs only from roseocobaltic chloride by contain- ing one molecule of water less. With the formation of the purpureocobaltic chloride, a second compound called Liiteocoballic chloride : CoCl 3 ,6NH 3 , is also pro- duced, which is difficultly soluble in cold water, and is deposited when the hot solution cools as dark yellow crystals. If the aqueous solutions of these chlorides are digested with silver oxide, silver chloride is produced, together with a strongly alkaline liquid, from which, however, the bases cannot be separated 488 Text-Book of Inorganic Chemistry. as they decompose, with evolution of ammonia, when the solutions are concentrated. In what form the five or six molecules of ammonia exist in these compounds it is still uncertain. But the constitution of these compounds is certainly not that expressed by the above empirical formulae, which represent them as double compounds of cobaltic chloride and ammonia. Cobaltons Cyanide : CoCy 2 . — If potassium cyanide is added to a cobaltous salt, a dirty-red precipitate of cobaltous cyanide is thrown down. This salt is insoluble in water, but readily dissolves in an excess of potassium cyanide, forming a double salt : 2KCy, CoCy 2 , which is decomposed by hydrochloric acid. If, however, the clear solution of this double salt is boiled with an excess of potassium cyanide, hydrogen is evolved, the cobalt oxidized, and Potassium Cobalticyanide : K 3 Co" / Cy 6 = 3KCy,CoCy 3 , pro- duced according to the equation : — 2KCy,Co'Cy, + 2KCy + H,0 = 3KCy,CoCy 3 + KOH + H. The salt is deposited on afterwards evaporating the solution in yellow rhombic crystals of the above composition and without water. It is easily soluble in water, acids do not precipitate cobaltic cyanide, and it is generally similar to potassium ferri- cyanide, with which it is isomorphous. In the same way, the cobalt which it contains is neither precipitated with caustic soda nor with ammonium sulphide, but hydrochloric or sulphuric acid when added to its concentrated aqueous solution separates hydro- cobalticycmic acid: H 3 CoCy G , which separates out as white needles on evaporation and addition of alcohol. In solutions of the salts of the heavy metals, potassium cobalticyanide produces insoluble precipitates of its various salts. The cupric compound, for ex- ample, is a pale-blue precipitate. Cobalt unites with sulphur in several proportions, and produces the following compounds. Cobaltous Sulphide : CoS, is obtained as a black amorphous precipitate when ammonium sulphide is added to a cobaltous salt, and is insoluble both in water and in dilute acids. The com- pound may also be obtained as a grey compact mass by heating a mixture of cobalt and sulphur. Tricobaltic Telrasulfihide : Co 3 S 4 = CoS,Co.,S 3 , occurs in nature Compounds of Cobalt. 489 in steel-grey octahedra of the regular system as the mineral linnaite. Cobalt Bisulphide : CoS. z , is not known in the free state, but only in combination with cobalt diarsenide as the mineral glance cobalt. Cobalt Diarsenide : CoAs 2 , and Cobalt Triarsenide : CoAs 3 , both occur in the mineral kingdom, the former as speisscobalt and the latter as skutterudite. The chief application of cobalt in the arts is for the preparation of the stable blue pigment known as smalt. Smalt is a potash glass which contains cobalt oxide instead of lime. It is prepared by fusing the partially roasted cobalt ores (zaffre) with quartz and potassium carbonate. The potassium silicate then unites with the cobalt silicate, producing a dark-blue glass. At the same time the compounds of sulphur and arsenic with the other metals (especially nickel) accompanying the cobalt collect at the base of the crucible, producing a speiss which is worked for nickel. The blue cobalt glass is ground to an impalpable powder, and used in this form as a pigment. Various shades of blue are pro- duced, depending chiefly on the amount of cobalt introduced into the glass, which seldom exceeds 7 or 8 per cent. Smalt is one of the most beautiful of blue pigments, and is, at the same time, the most stable. It is not bleached in light, nor is it attacked by acids or alkalies. It is not, however, available for all purposes where a blue pigment is required, and since the intro- duction of so much cheap ultramarine the demand for the more expensive but more stable smalt has considerably diminished. Detection of Cobalt Compounds. The intense blue colour which traces of any cobalt compound impart to a bead of borax, in common with other kinds of glass, and the fact that the blue colour is permanent both in the reduc- ing and in the oxidizing flame, makes it easy to detect these com- pounds. The salts may also be recognized by their red colour when hydrated and blue colour when anhydrous. Sulphuretted hydrogen produces no precipitate in acid solutions of cobalt salts, and in this way cobalt may be easily separated from copper, lead, and all those metals whose sulphides are precipitated from acid liquids by sulphuretted hydrogen. Cobaltous sulphide, like the 49° Text-Book of Inorganic Chemistry. corresponding sulphide of nicke!, after precipitation with am- monium sulphide, does not redissolve in dilute hydrochloric acid, and may thus be separated from the sulphides of iron, manganese and zinc, which are easily soluble in dilute acids. NICKEL. Chemical Symbol : Ni. — Atomic Weight: 58-5. This metal, like cobalt, is only found native in metallic meteor- ites, but always in larger quantities than cobalt. The chief com- pound in which nickel occurs in nature is Kupfernickel or niccolite : NiAs, so called on account of its reddish colour. The metal is prepared from this and from cobalt speiss by removing the cobalt by a series of processes and then reducing with charcoal. Metallic nickel was formerly almost exclusively brought into trade as small cubes, obtained by heating the compressed oxide in charcoal ; it is now largely used in the form of rolled plates for the processes of nickel-plating. This commercial metal is far from pure, and usually contains copper and iron as well as arsenic and cobalt. A purer metal may be obtained by heating the carbonate or oxide in a stream of hydrogen, and then fusing the spongy mass so produced under borax at a high temperature. Large deposits of nickel, in the form of a hydrated double sili- cate of nickel and magnesium, have recently been discovered in New Caledonia, and considerable quantities of the metal are now obtained from this source. Nickel is a white metal with a yellowish tinge and nearly 9-0 specific gravity. It is malleable and ductile and, like cobalt, faintly magnetic. Nickel is very difficultly fusible, and only changes slightly in moist air ; when heated in the air it becomes covered with a thin layer of oxide. Hydrochloric and dilute sulphuric acid dissolve the metal slowly, with evolution of hydrogen ; nitric acid, when not too concentrated, converts it into the nitrate. Just as the hydrated cobalt salts nearly all possess a red colour, the nickel salts are distinguished by their green colour. Nickel forms two oxides : nickelous and nickelic oxide, corresponding to the two oxides of cobalt. Nickelous Oxide : NiO, is a dirty-green powder obtained by Nickel. 491 heating its hydrate or carbonate. It dissolves in acids and forms green solutions. Kickelous Hydrate : Ni(OH).,, is thrown down as an apple-green precipitate when caustic soda is added to a nickel salt. It dries to a dark -green powder. The same precipitate is produced by ammonia, but dissolves in excess of the reagent to a blue liquid. Nickelic Oxide : Ni 2 3 , is a black powder produced on care- fully heating nickel nitrate. The corresponding hydrate NifOH 3 ) is obtained as a black amorphous powder when a salt of nickel is precipitated with sodium hypochlorite. Both nickelic oxide and hydrate dissolve in hydrochloric acid with evolution of chlorine and formation of nickel chloride ; and both break up when heated into nickelous oxide and oxygen. Nickel Sulphate: S0 2 -0 2 Ni + 7H 2 0, crystallizes from its solutions in emerald-green prisms. Like other sulphates of the same composition, it forms double salts with the alkaline sulphates, of which the ammonium compound: S0 2 -0 2 Ni + SO a (ONH 4 ) a + 6H„0, is largely used in the arts for nickel-plating. Nickel Nitrate : (N0 2 ) 2 2 Ni + 6H 2 0, forms long green prisms, easily soluble in water, and deliquescing in the air. Nickel Carbojiate. — A basic carbonate is' thrown down as a green precipitate when sodium carbonate is added to a nickel salt. The precipitate, like other insoluble nickel compounds, dissolves in ammonia or ammonium carbonate to a blue liquid. Nickel Chloride : NiCl 2 + 6H 2 0, is deposited from its aqueous solution when strongly concentrated as small green crystals, which, when heated, lose water and become yellow. Nickel Cyanide : NiCy 2 , is thrown down as a dirty green pre- cipitate when a solution of potassium cyanide is carefully added to a salt of nickel. The salt is insoluble in water, but dissolves easily in an excess of potassium cyanide to form a double compound of the composition : 2KCy,NiCy 2 + H 2 0, which crystallizes in yellow prisms when the solution is evaporated down. Nickel cyanide produces similar double salts with the cyanides of sodium, barium, calcium, &c, from all of which it is again precipitated by hydro- chloric acid. Potassium -nickel cyanide is not changed when boiled 492 Text-Book of Inorganic Chemistry. with an excess of potassium cyanide ; no compound corresponding to potassium cobalticyanide is produced, nor does such a compound appear to exist. Nickel Sulpiride : NiS, is found in the mineral kingdom in fine hair-like crystals as the mineral millei'ite. The same compound is produced as a black precipitate by adding sodium sulphide to a solution of a nickel salt. If precipitated with ammonium sulphide a portion of the precipitate dissolves and forms a dark brown solution. Nickel Alloys. The chief use of nickel in the arts is for the manufacture of various alloys, the most important of which is that known as Germaii-silver,argentan, orm'ckel-szh>t.'r,a.nd\vhich contains copper, zinc, and nickel in varying proportions. If the quantity of nickel is considerable the alloy has a white colour, with a faint yellowish tinge. German-silver is used for a variety of purposes, and forms the basis of the better kind of silver-plated ware. The five- and ten-pfennig pieces of the Germans, corresponding to our half- pence and pence, consist of an alloy of three parts of copper and one part of nickel. Alloys of nickel are also used for coinage pur- poses in Switzerland, Belgium, America, and Jamaica. In recent years considerable quantities of nickel have been used for the process of nickel-plating, which consists in coating metallic articles with a thin layer of nickel by an electric current. The coating of nickel much improves the appearance, and at the same time only rusts slowly in the air. Nickel-plating is especially used for articles made of iron, which, when well deposited, it effec- tually protects from rust. Detection of Nickel Compounds. Nickel oxide dissolves in molten glass and borax, like cobalt oxide, but the colour is far less characteristic. In the oxidizing flame the borax bead is of a yellowish-brown coloui, which becomes grey and opaque in the reducing flame from finely-divided metallic nickel. A solution of a nickel compound is characterized by the black precipitate of nickel sulphide produced on the addition of ammo- nium sulphide, and especially that the precipitate is slightly soluble, forming a dirty-brown solution. In other respects nickel compounds closely resemble those of cobalt, except in their colour. Zin 493 The salts of the two metals may be separated by the addition of acetic acid and potassium nitrite, when the cobalt is completely precipitated but the nickel remains in solution (p. 486). Salts of calcium, barium, and strontium must not be present, otherwise some of the nickel is precipitated with the cobalt. ZINC. Chemical Symbol : Zn. — Atomic Weight: 65. Zinc is never found free in nature, but chiefly occurs in com- bination with carbonic acid as smithsonite or calamine (zinc carbonate: CO-0„Zn), and united with sulphur as zinc blende (zinc sulphide : ZnS). A further important ore of zinc is a silicate called siliceous calamine or willemite. Zinc is volatile at a bright red-heat, and its oxide cannot therefore be reduced, like iron and other metals, with charcoal in an ordinary furnace. All the zinc would then be volatilized and burnt. The zinc oxide, obtained by roasting any of the above ores in the air, is therefore heated with charcoal in specially constructed vessels of fireclay, resembling retorts, and in the cooler parts of which the volatilized zinc condenses. Commercial zinc is never pure ; even when freed from other less volatile metals by repeated distillation, it always contains more or less arsenic. This impurity may, however, be removed by stirring small quantities of nitre into the fused metal ; a con- siderable quantity of the zinc is then oxidized together with the arsenic, which is converted into potassium arsenate. If all the arsenic has not been removed during the first operation, the pro- cess must be repeated. The presence of arsenic in zinc is best detected by the aid of Marsh's test (p. 239). Zinc possesses very remarkable physical properties, which are not found in the same degree in any other metal. Under some circumstances it is tough and hard, under others soft and malleable, or it may be so brittle as to be easily powdered, and finally it is known as a gas. Zinc acquires these very various properties at different temperatures. The bluish-white metal, with a lamellar, crystalline fracture and a specific gravity of 7-1 5, is hard and tough at the ordinary temperature, and can neither be hammered nor rolled into plates ; for this reason zinc was for a long time only 494 Text-Book of Inorganic Chemistry. used for the preparation of brass and other alloys. The metal became, however, much more valuable when it was discovered that at the boiling-point of water — i.e. between 100° and 150 — it becomes soft and malleable, and can be hammered or rolled to thin plates. At a higher temperature, 200 and above, it is as brittle as antimony, and can be rubbed to powder in a hot mortar. Zinc melts at about 412 , and boils at a bright-red heat — i.e. about l,ooo°. Zinc, when heated in a crucible to above its melting-point and then exposed to the air by opening the crucible, burns with a bright bluish flame, forming zinc oxide, which becomes distributed through the air of the room in light white flocks — lana fihilosophica of the alchemists. The non-volatile zinc oxide which remains behind in the crucible has a fine yellow colour, but becomes white when cold. Zinc oxidizes in moist air even more readily than iron, and becomes covered with a thin film of zinc oxide or basic zinc car- bonate. But this film forms a compact layer over the metal, and so protects it from further oxidation that it scarcely increases in weight when exposed for a long time to the air. For this reason zinc is used for a variety of purposes for which iron, that rusts so easily and continuously, could not be employed. Large quantities of zinc are consurned in the process of galvanizing iron, which consists in coating iron plates or vessels with a layer of zinc, by immersing them in the molten metal. And one of the most important and useful alloys — brass — contains a large proportion of zinc. Zinc is easily dissolved by acids : hydrochloric and dilute sulphuric acid evolve hydrogen when acting on the metal ; nitric acid produces nitric oxide, or if dilute, ammonia, and oxyammonia when sulphuric acid is also present (p. 188). Zinc, like iron, easily decomposes water at a low red heat. And, since zinc oxide forms compounds with the alkalies which are soluble in water, the metal also dissolves in caustic alkalies, and in ammonia when warmed, and then liberates hydrogen. Zinc only forms a single compound with oxygen, sulphur, or the halogens, respectively ; it is a dyad in all its compounds. Zinc Oxide: ZnO. — This compound, the only oxide of zinc, is a soft white powder,infusible at a red heat. It is obtained by burning zinc in the air or by heating its hydrate, carbonate, or nitrate Compounds of Zinc. 495 " When heated it becomes of a fine yellow colour, but is again white when cold. Zinc oxide is used in medicine for some purposes, but is chiefly employed as a white pigment in the place of white-lead. This zinc- white does not cover so well as white-lead, but possesses the advantage of not blackening in the air. Sulphuretted hydrogen, which is contained in traces in the air of inhabited places, gradually converts both the zinc and lead compounds into sulphides, of which lead sulphide is black, but zinc sulphide is white, and is therefore not noticed. Zinc Hydrate : Zn(OH) 2 , is obtained as a gelatinous, white precipitate when a small quantity of caustic soda is added to a solution of a zinc compound. When washed and dried it remains as a white powder. The hydrate dissolves in excess of caustic soda and produces a soluble salt, sodium zincate, in which zinc oxide plays the part of an acid. It also dissolves in ammonia. Zinc Sulphate, white vitriol: S0 2 -0,Zn + 7H 2 0, is prepared by dissolving zinc or zinc oxide in dilute sulphuric acid, or on the large scale by gently roasting blende (zinc sulphide) and extracting with water. It crystallizes in large rhombic prisms, and when heated melts in its water of crystallization. This fused mass is sometimes cast into moulds, broken into pieces when cold, and so brought into the market for dyeing purposes. Zinc sulphate forms a basic sulphate when its solution is heated with zinc hydrate ; it unites with alkaline sulphates to form crystalline double salts — e.g. SCX(OK), + S0 2 • 0. 2 Zn + 6H 2 0. Zinc Nitrate : (NO,) 2 0.,Zn + 6H 2 0, crystallizes from its concen- trated aqueous solution in four-sided prisms. It is very easily soluble in water and deliquesces in the air ; when heated it melts in its water of crystallization. Zinc Carbonate: CO-O.^Zn, occurs in nature as the mineral calamine and occasionally in rhombohedra isomorphous with calcite, as zinc-spar. The precipitate obtained when sodium carbonate is added to a solution of a zinc salt is always a basic salt, or, in other words, a compound of zinc carbonate and hydrate. Ztnc Chloride : ZnCL,. — This salt may be obtained either by heating zinc in chlorine, or by evaporating an aqueous solution and 496 Text-Book of Inorganic Chemistry. distilling the residue. The distillate obtained by either of these processes is a white translucent mass. It eagerly attracts water from the air and deliquesces and dissolves readily both in water and alcohol. Zinc chloride melts at 260° and does not appreciably volatilize at 400 ; it is therefore often used instead of an oil-bath to heat substances to a high ahd constant temperature. Its boil- ing-point is about 700°. An aqueous solution of zinc chloride is obtained by dissolving zinc in hydrochloric acid. If this solution is evaporated, a part of the salt is decomposed by the water into hydrochloric acid, which passes away, and zinc oxide, which unites with the rest of the zinc chloride to form a basic salt. The concentrated solution of zinc chloride is powerfully caustic and cannot be filtered through paper. It also possesses strong antiseptic properties, and is sometimes used to preserve wood from decay. If a concentrated solution of zinc chloride is made into a paste with zinc oxide, the mixture becomes warm and soon sets to a snow-white hard mass of an oxy-chloride. If finely powdered glass is added to the mixture, a very hard cement is obtained ; such mixtures are employed as a cheap stopping for decayed teeth. Zinc Sulphide : ZnS. — The zinc sulphide occurring in nature as blende is of a yellow or brown colour, but that obtained artificially by precipitating a solution of a zinc salt with ammonium sulphide is a pure white powder. Hydrochloric acid readily dissolves it, with evolution of sulphuretted hydrogen and formation of zinc chloride. Zinc sulphide is not precipitated when sulphuretted hydrogen is led into a solution of a zinc salt acidulated with hydrochloric or sulphuric acid ; but the gas throws down all the zinc from a solution of its acetate. Detection of Zinc Compounds. If any zinc compound is mixed with sodium carbonate and heated on a piece of charcoal in the inner blowpipe flame, the zinc is reduced, volatilized and burnt, and produces an incrustation on the charcoal which is yellow when hot, and white when cold. Zinc compounds are further characterized by the white precipi- tate which they give with ammonium sulphide. Zinc may be easily separated from iron, nickel, and cobalt by the addition of caustic soda in excess. Zinc oxide then alone dissolves, and can Cadmium. 497 afterwards be precipitated as white zinc sulphide by passing sulphu- retted hydrogen through the alkaline solution. Zinc hydrate and carbonate dissolve also in caustic potash, in ammonia, or in am- monium carbonate, and produce soluble double compounds. CADMIUM. Chemical Symbol : Cd. — Atomic Weight: 112. Like indium, cadmium is nearly always found associated with zinc in nature, but zinc ores usually contain far more cadmium than indium — often as much as five per cent. Cadmium is ob- tained as a bye-product in the extraction of zinc, and as it is more volatile than this metal, is found chiefly in the first portions which come over when the zinc is distilled. Cadmium sulphide is occa- sionally found as the rare mineral, greenockite. Cadmium is a tin-white metal, of specific gravity 8-6, melting at 3 1 5 and boiling at about 770°. Unlike zinc, it is soft at the ordinary temperature, and can be rolled into sheets or drawn into wire. The metal remains unaltered in the air, but burns brilliantly at a red-heat, producing brown vapours of cadmium oxide. It dissolves slowly in hydrochloric or dilute sulphuric acid, with evo- lution of hydrogen ; nitric acid dissolves it more readily. All compounds of cadmium contain the metal in the form of a dyad, in which respect it resembles zinc. Cadmium Oxide : CdO, is obtained as a brown amorphous powder, soluble in acids, when the metal is heated in the air or when the nitrate is glowed. Cadmium Hydrate : Cd(OH) 2 , is thrown down as a white precipitate on adding caustic soda to a solution of a cadmium salt. The precipitate is insoluble in excess of caustic soda, and so differs from zinc hydrate ; it is, however, like the zinc compound, easily soluble in ammonia. The cadmium salts, with the exception of the sulphide, are all colourless. Cadmium Sulphate : 3S0 2 2 Cd + 8H 2 0, forms large, colourless, easily soluble crystals. Cadmium Nitrate : (NO.,). 2 2 Cd 4H,0, crys- tallizes in deliquescent prisms. Cadmium Carbonate : CO • 2 Cd, is precipitated as a white powder by adding sodium carbonate to a solution of a cadmium salt. Cadmium Chloride: CdCl 2 + 2H 2 0, K K 498 Text-Book of Inorganic Chemistry. crystallizes in colourless prisms, which effloresce when exposed to the air and are easily soluble in water. Cadmium Sulphide : CdS, is thrown down when sulphuretted hydrogen is led into a solution of a cadmium salt as a bright yellow precipitate, which is insoluble in dilute, but dissolves in concentrated hydrochloric acid. It is distinguished from arsenious sulphide, which has a similar colour, by its insolubility in alkaline sulphides, including ammonium sul- phide. Cadmium sulphide is used as a stable yellow pigment. Detection of Cadmium Compounds. Compounds of cadmium when heated on charcoal in the inner flame of the blowpipe are reduced to the metallic state, but the metal at once volatilizes and burns in the outer flame, pro- ducing a brown incrustation of cadmium oxide on the charcoal. When a current of sulphuretted hydrogen is let through an acid solution of a cadmium salt, a yellow precipitate of cadmium sulphide is produced, insoluble in alkaline sulphides (distinction from arsenic). This reaction serves to separate cadmium from zinc, which is not precipitated from its acid solutions by sulphu- retted hydrogen. If ammonia is added to a solution of a cadmium salt, a white precipitate of cadmium hydrate is formed, which dis- solves in an excess of the reagent. Caustic soda produces the same precipitate, but does not redissolve it when an excess is added. Cadmium may be separated from copper by boiling the mixed sulphides with dilute sulphuric acid ; the cadmium sulphide then dissolves and the cupric sulphide remains behind. LEAD. Chemical Symbol : Pb. — Atomic Weight : 207. The most important form in which lead occurs in nature is as the sulphide : PbS — the mineral galena. Lead is also found in smaller quantities as the sulphate — anglesite, as the carbonate — cerussite, and less commonly as the chromate, phosphate, and molybdate. Two processes are chiefly used for the extraction of lead from galena accordingly as the ore is more or less pure. The impure Lead. 499 ores are either roasted and' reduced with carbon or else simply smelted with iron. In the latter case the iron unites with the sulphur, for which it has a stronger attraction than the lead, and forms a fusible slag of ferrous sulphide with the impurities, which covers the heavier reduced lead. In the second method — the so- called air reduction process — the sulphur is indirectly oxidized by the oxygen of the air. The ore is first carefully roasted in a specially constructed furnace, so that one portion is converted into lead sulphate, another portion into lead oxide, and the rest remains as unchanged lead sulphide. The whole is afterwards well mixed and heated more strongly with exclusion of air, when the sulphur in the unchanged sulphide is burnt by the oxygen contained in the lead oxide and sulphate, and the reduced lead runs from the hearth of the furnace into moulds :— PbS PbS + 2PbO + SO.,-0.,Pb 3Pb + SO,. 2Pb + 2SO, Lead when thus smelted often contains silver. If the quantity is sufficient to pay for its extraction, 1 the lead is converted by oxidation into lead oxide, the silver remaining unchanged. This process, called cupellation, is performed in an oxidizing reverbera- tory furnace, of which the German form is shown in fig. 66. This furnace consists of a porous clay hearth, on which the argentiferous lead, A, is placed, and which is kept fused by the heated gases 1 I.e. about one-tenth per cent. K K 2 SOO Text-Book of Inorganic Chemistry. from the fire, F, while a continuous blast of air is blown through the openings a, a, over the surface of the metal in order to oxidize the lead to lead oxide. The lead oxide melts as it is produced, and is partly absorbed by the porous clay, but mostly flows away by side openings into vessels provided to receive it. The silver, however, remains unoxidized, and the lead oxide is therefore free from this metal, except towards the end of the operation. Finally, when all the lead has been converted into lead oxide, and the silver is coated with a thin film of the oxide, beautiful iridescent tints are observed on its surface, due to the colours of thin plates ; immediately afterwards the bare surface of the molten silver flashes out, and the operation is over. In the English form of cupel furnace, the hearth is made of bone ash, and the cover is not moveable. The lead oxide (litharge), of which large quantities are produced by cupellation, is afterwards again converted into metallic lead by reduction with charcoal or coal. Two other processes are now largely used for extracting silver from lead, when the quantity is insufficient to pay the expense of cupelling. The first of these, known as Pattinsoris process, depends upon the fact that when fused argentiferous lead is allowed to cool slowly, the crystals of lead which first separate out are almost entirely free from silver ; and if these crystals are removed and the operation repeated several times, a sample of lead is at last obtained which is rich enough in silver to pay for cupelling. A second method, called Parke? process, is based upon the property of zinc to form an alloy with silver, but not with lead. If a small quantity of zinc is added to the molten argentiferous lead, and the mass allowed to cool, the alloy of zinc and silver rises to the sur- face and can be removed during solidification. The zinc can then be abstracted from this alloy by dissolving it in dilute sulphuric acid, or by distilling it off. Lead is a lustrous metal of bluish-grey colour, and soft enough to be cut with a knife, and to mark paper. It is malleable and ductile, but possesses little tenacity. The metal has a specific gravity of 11-4, melts at 32 5 , and volatilizes at a white heat, though it cannot be distilled. When exposed to moist air, it becomes dull and coated with a thin grey layer of lead carbonate. Pure water containing air also attacks it and dissolves a small quantity of the lead. But if the water contains traces of certain salts dis- solved in it — e.g. calcium sulphate — the leaden pipes through which the water passes are not attacked. All lead compounds are Oxides of Lead. 501 poisonous, and especially when dissolved in water, and hence drinking water containing only traces of dissolved lead (probably as its acid carbonate) is injurious to health. The presence of lead in a sample of drinking water may be readily proved by the addition of sulphuretted hydrogen water, or by leading a stream of the gas through it, when if lead is present the water becomes of a brown colour from the lead sulphide formed. Hydrochloric acid and dilute sulphuric acid scarcely act upon lead, either in the cold or when warmed. Hot concentrated sul- phuric acid dissolves lead slightly, but only so slightly that sul- phuric acid may be evaporated to a certain strength in leaden dishes (see p. 159). The best solvent for lead is nitric acid, which dissolves it readily, producing lead nitrate and nitric peroxide. From this salt, or from a solution of lead acetate, zinc precipitates metallic lead. The lead is then deposited in lustrous crystalline plates, and, as the metal in this, form possesses some similarity with leaves and boughs, it is sometimes called a 'lead-tree.' Lead unites with oxygen (not with sulphur or the halogens) in three proportions, and produces the compounds — lead suboxide : Pb.,0, lead oxide : PbO, and lead peroxide : PbO,. A fourth oxide is red-lead: Pb 3 4 or Pb,0 5 , which, however, must be considered as a compound of lead oxide and lead peroxide. lead Suboxide : Pb,0. — This compound, which is only of scientific interest, is a dark-grey indifferent powder, and is pro- duced when lead oxalate is heated to 300 in an oil-bath as long as gas is given off. The decomposition which the oxalate then under- goes is represented in the equation : — 2-!^Q0 2 Pb = Pb 2 + 3C0 2 + CO. Lead suboxide, when heated in the air, is converted into lead oxide ; it does not unite with acids, but is decomposed by them into lead oxide and lead. lead Oxide, Litharge, Massicot : PbO, occurs as a yellow or reddish powder, or as crystalline scales, according to the method used for its preparation. It easily melts, and solidifies on cooling to a crystalline mass. Lead oxide may be obtained by heating either lead hydrate, carbonate, or nitrate, or by the oxidation of molten lead in the air. Most of the litharge brought into trade is derived from the cupellation of argentiferous lead. 502 Text-Book of Inorganic Chemistry. Lead oxide is insoluble in water, but dissolves easily in nitric or acetic acid. It is a powerful base, and as most of its salts are insoluble in water, a solution of the soluble acetate is used in the laboratory for the detection of acids. Litharge is largely used in the arts : considerable quantities are consumed in the manufacture of flint-glass, and as a cheap glaze for stone and earthenware. Nearly all other lead compounds are prepared from it — e.g. lead acetate and nitrate, red-lead, lead- plaster, &c. lead Hydrate : Pb(OH) 2 , is produced as a white, voluminous precipitate, when a solution of lead acetate is mixed with a large excess of caustic soda, until the solution is strongly alkaline, and until a portion of the hydrate has been redissolved by the soda. If a smaller quantity of caustic soda is used than is required to com- pletely precipitate the lead, the lead hydrate is mixed with basic lead acetate ; and as sodium acetate itself has an alkaline reaction, the liquid must be strongly alkaline before the soda is in excess. Lead hydrate preserves its white colour when washed and dried. It is quite insoluble in water, but unites with acids more readily than lead oxide, and even attracts carbonic acid from the air. When heated it loses water and yields lead oxide. Red-lead: Pb 3 0, = 2PbO,PbO„.— If litharge (lead oxide) is heated in the air to about 400 , but not much above, it absorbs oxygen, and is converted into a bright brick-red crystalline powder. This substance is called red-lead or minium, and is used both as a pigment and for a number of other technical purposes. This oxide cannot unite with acids and form salts like lead oxide ; and although dilute nitric acid acts upon it even in the cold, it only abstracts lead oxide from it, and leaves brown lead peroxide be- hind. Red-lead evolves chlorine when warmed with hydrochloric acid, owing to the lead peroxide which it contains. When strongly heated it loses oxygen and is converted into yellow lead oxide. Lead Peroxide : Pb0 2 .-^The preparation of this substance from red-lead by the action of dilute nitric acid has just been de- scribed. When repeatedly boiled and washed with water and dried, it remains as a brown amorphous powder, with a purple tinge — whence it is sometimes called the puce-coloured oxide of lead. Lead peroxide may also be obtained by precipitating a solution of Salts of Lead. 503 lead nitrate with sodium hypochlorite. No solvent is known for lead peroxide ; like manganese peroxide, which it resembles in other points, it is not attacked by nitric acid. When warmed with hydrochloric acid it evolves chlorine ; and when heated alone or with sulphuric acid oxygen is given off, and in the latter case, white insoluble lead sulphate remains behind. It unites with dry sul- phurous anhydride, with considerable evolution of light and heat, and produces the same substance (p. 152). If lead peroxide is warmed with nitric acid, and finely powdered sugar gradually added, the sugar is oxidized to carbonic acid and water by a half of the oxygen contained in the peroxide, and a clear solution of lead nitrate is soon obtained. Salts of lead. — All the salts of lead contain the metal as a dyad, and correspond to lead oxide ; none have yet been prepared containing lead in the tetrad form, although it is possible that a tetrachloride corresponding to the peroxide might be produced by the action of hydrochloric acid on the peroxide ; but if produced it is so unstable that its isolation is exceedingly difficult. Most of the lead salts are insoluble in water. Those which dissolve in water include the pitrate and chlorate, as well as the acetate. And although the acetate belongs strictly to organic chemistry, it is the most important soluble lead salt, and may well be included with the inorganic compounds. Many salts of lead are largely used both in the arts and in medicine. All lead compounds, and especially those which are soluble in water, are very poisonous. The soluble compounds are characterized by their sweet taste. lead Sulphate : S0 2 -0 2 Pb, occurs in nature as the mineral angleszte, and may be prepared artificially by precipitating a soluble lead salt with dilute sulphuric acid or with a solution of a sulphate. It is only very slightly soluble in water and dilute acids, but hot concentrated sulphuric acid dissolves not inconsiderable quantities of it. Strong hydrochloric acid converts it partially into lead chloride, with the liberation of the corresponding quantity of sul- phuric acid. Lead sulphate dissolves readily in basic ammonium tartrate, for example, when digested with tartaric acid and then ammonia added in excess. Ammonium sulphate also forms a soluble crystalline double salt with it, which is, however, decom- posed by water, with separation of lead sulphate. 504 Text-Book of Inorganic Chemistry. Lead Nitrate : (N0 2 ) 2 2 Pb, is obtained by dissolving lead or lead oxide in a slight excess of nitric acid, and crystallizes from the solution containing free acid as colourless, transparent, hard octahedra, without water of crystallization. It is soluble in water, but insoluble in strong nitric acid, by which it is precipitated from its concentrated aqueous solution. If the aqueous solution is boiled with powdered lead oxide, some of this substance dissolves and produces a difficultly soluble basic salt of the composition : N0 2 -OPb(OH), which crystallizes out on cooling. Lead nitrate decomposes when heated into lead oxide, oxygen, and nitric oxide (p. 196). Lead Phosphate : (PO),0 6 Pb 3 , is thrown down as a white pre- cipitate when, sodium phosphate is added to a solution of a lead salt. It is insoluble in water and acetic acid, and forms double compounds with other lead salts — e.g. the nitrate and chloride. The mineral pyromorphite is a double compound of lead phosphate and chloride: 3(PO). 2 6 Pb 3 + PbCl 2 , and is isomorphous with apatite (p. 396). Lead Silicates. — Silica and lead oxide may be fused together in almost any proportion, producing a vitreous mass of varying composition, which is an essential constituent of flint or lead-glass, and strass (p. 403). A lead silicate of definite composition has not yet been prepared. lead Carbonate : CO-0 2 Pb. — The normal salt of the above composition is found in nature as the mineral cerussite, isomor- phous with arragonite, and may be prepared artificially as a white powder, insoluble in water, when a solution of lead nitrate is pre- cipitated with ammonium carbonate. Basic lead Carbonate, 'White lead : 2CO ■ 2 Pb + Pb(OH) 2 . — Far more important than the preceding compound is a basic car- bonate of lead, very largely used in the arts as a white pigment. White lead is manufactured by a variety of methods, of which the following two are the most important. The French or Etiglish process depends upon the fact that a stream of carbonic acid, when led into a solution of basic lead acetate, precipitates the whole of the lead oxide which the salt contains in excess of the normal compound as white lead. Finely White Lead. 505 powdered litharge (lead oxide) is boiled with dilute acetic acid until no more is dissolved, when the normal acetate is first obtained, and then the basic salt. On passing a stream of carbonic acid through the clear solution, a precipitate of white lead is produced with a solution of the normal acetate. This solution can be again converted into the basic acetate by boiling with litharge, and from this carbonic acid again throws down a further quantity of white lead, and so on. Thus, with a small quantity of acetic acid a large quantity of white lead can be obtained. A second and older method, called the Dutch process, because first carried on in Holland, consists also in decomposing a basic lead acetate with carbonic acid, but the arrangement by which this reaction is brought about is different. Plates of lead coiled into a spiral or cast leaden gratings are placed in earthenware pots con- taining a layer of strong vinegar at the bottom. The pots are then loosely closed with a leaden lid and imbedded in horse-dung or spent tan. After several weeks the vessels are withdrawn and opened, when the leaden plates are found to be strongly corroded and covered with a loose layer of white lead. This is removed, and the plates returned to the pots until they are completely corroded away. The process by which the white lead is produced is a simple one. The warmth produced by the decomposition of the dung or spent tan gradually evaporates the vinegar which attacks the lead, and, in the presence of the oxygen of the air, produces basic lead acetate. The acetate is then decomposed as fast as it is produced by the carbonic acid evolved during the putrefaction of the dung or spent tan, producing white lead, which remains loosely adhering to the leaden plates. White lead prepared in this way is more highly valued as a pigment than that obtained by any other process. The colour is purer and it covers better. lead Acetate, Sugar of Lead: A 2 -0 2 Pb + 3H,0 = (C,H 3 0). 2 3 Pb + 3H,0.' — Normal lead acetate is obtained by dissolving litharge in a small excess of acetic acid, and crystallizes in colourless, trans- parent, glistening prisms when the solution is evaporated down. Of all lead salts it is the most soluble in water ; it also dissolves in alcohol. 1 The abbreviated symbol A stands for the radical of acetic acid (acetyl) . C5H5O, and has the same relation to acetic acid : AOH, as nitryl : X0 2 , has to nitric acid : N0 2 'OH. 506 Text-Book of Inorganic Cliemistry. If a solution of the normal acetate is boiled with lead oxide it combines with this substance and produces various basic acetates, of which those containing the greatest quantity of the oxide are insoluble in water. The monobasic acetate of the composition : A,-0,Pb + Pb(OH),, = A.OPb(OH), is easily soluble in water and reacts alkaline. It parts with one half of its lead when carbonic acid is led through its solution. Lead Chromate : CrO., -O.Pb, is found in nature as the mineral crocoisite, and may be obtained as a bright yellow powder by precipitating a soluble lead compound with potassium chromate or dichromate. This precipitated lead chromate when washed and dried is largely used as a yellow pigment under the name of chrome-yellow. It melts when heated, and at higher temperatures evolves oxygen. If chrome-yellow is digested with caustic potash or soda, or if boiled with normal potassium chromate, it loses chromic acid and is converted into a basic lead chromate of a bright red colour and the composition: CrO,-0,Pb + PbO. This substance is the chrome-red of commerce, and is also largely used as a pigment. lead Chloride : PbCL. — This salt, which is only slightly soluble in cold water but more in hot water, is deposited as a heavy white precipitate when hydrochloric acid or a soluble chloride is added to a solution of lead acetate. It crystallizes from its hot saturated solution in colourless lustrous scales or needles. When strongly heated, the salt melts, and solidifies on cooling to a horny mass which may be cut with a knife. Lead chloride unites with lead oxide in varying proportions and produces different basic chlorides or oxychlorides. One of these is Cassel yellow ; it is prepared by heating a mixture of litharge and sal-ammoniac, and is used as a pigment. Another oxychloride, of the composition : PbCl,,PbO, is the mineral mat- lockite. If an aqueous solution of lead chloride is boiled with lead hydrate, the whole of the former salt is precipitated as insoluble oxychloride. lead Iodide : Pbl.,, is thrown down as a pale yellow precipi- tate when potassium iodide is added to a solution of lead acetate. It is nearly insoluble in cold water, but boiling water dissolves it slightly. This hot solution is colourless, but on cooling deposits Alloys of Lead. 507 the dissolved lead iodide in lustrous golden scales. Lead iodide forms a double compound with potassium iodide. Lead Cyanide : PbCy„, is a white flocculent precipitate obtained by adding potassium cyanide to lead acetate. It readily dissolves in an excess of potassium cyanide, and the solution then contains the double salt : 2KCy,PbCy,. Xiead Sulphide : PbS, occurs in nature as galena, the most abundant ore of lead. Galena is found either in compact masses of a grey colour, or else crystallized as cubes, usually with sub- ordinate faces of the octahedron. The same compound may be obtained as a black amorphous precipitate by leading a stream of sulphuretted hydrogen into a solution of a lead salt. The precipitate retains its colour when dry, and if heated out of access of air, melts without decomposition ; but if roasted in the air both sulphur and lead undergo oxidation, and there are produced sulphurous anhydride with lead oxide and sulphate. Lead sulphide is scarcely attacked by hydrochloric acid. Dilute nitric acid dissolves it with formation of lead nitrate and free sulphur. Concentrated nitric acid converts it into the white sulphate, as the sulphur is then oxidized. Numerous double sulphides of lead and other metals, especially antimony, are found in nature; the minerals zinkenite, jamesonite, and boulangerite, are all double sulphides of lead and antimony in varying proportions. Besides the salts of lead oxide in which this compound plays the part of a base, others are known in which it is combined with strong bases as an acid. These compounds — called plumbates — are produced by digesting lead hydrate with the aqueous alkalies or with lime- or baryta-water, in which the hydrate easily dis- solves. The salts so obtained — e.g. potassium plumbate, of the probable composition Pb(OK), — have not yet been prepared pure enough to determine their exact composition. Alloys of lead. — Metallic leadjs not only used alone for the manufacture of water-pipes, for covering roofs, and for sulphuric acid chambers, but it is also a constituent of many important alloys. One of these is type-metal, an alloy of lead and antimony. German type-metal contains from 10 to 25 parts of antimony to 100 parts of 508 Text-Book of Inorganic Chemistry. lead, while that used in England contains about 30 parts of antimony and 30 parts of tin to 100 parts of lead, and sometimes a little copper as well. The larger the quantity of antimony, the harder and more brittle is the alloy produced. The addition of a small quantity of arsenic (at most 0'8 per cent.) to lead makes it suitable for the manufacture of shot. The alloys of lead and tin are de- scribed under the latter metal. Detection of Lead Compounds. Scarcely any other metal can be so easily detected and separated as lead. Any compound of lead when mixed with sodium carbonate and heated on charcoal in the reducing flame of the blowpipe yields small soft globules of metallic lead, which mark paper, and which are converted into yellow lead oxide in the oxidizing flame. In the borax bead, as in other forms of molten glass, lead oxide dissolves without imparting any colour. Sulphuretted hydrogen when passed through a neutral or acid solution of a lead salt precipitates black lead sulphide. The same solution gives a white precipitate of lead sulphate with dilute sulphuric acid, a white precipitate of lead chloride with hydro- chloric acid, which dissolves in hot water, and crystallizes out in needles on cooling, a yellow precipitate of lead iodide with potas- sium iodide, and a white precipitate of lead hydrate with caustic soda, soluble in excess of the reagent. Ammonia produces no pre- cipitate when added to a solution of lead acetate or only a slight one, after some time, but at once throws down basic lead nitrate when added to a solution of lead nitrate. The yellow precipitate which is obtained when potassium dichromate is added to a lead salt is also produced with a bismuth salt, but lead chromate may be distinguished from the bismuth salt by its solubility in caustic soda. The separation of lead from the other metals precipitated from their acid solutions by sulphuretted hydrogen is best effected by the addition of sulphuric acid and alcohol. Lead sulphate is slightly soluble in water, less in dilute sulphuric acid, and insoluble in alcohol. Thallium. 509 THALLIUM. Chemical Symbol : Tl. — Atomic Weight : 204. This remarkable element occupies an altogether anomalous position among other metals, for it unites in itself the chemical and physical properties of two groups of metals which are as different as possible from one another. On the one hand it has much similarity with lead and silver, and on the other with the alkali- metals, especially potassium. It resembles lead in possessing nearly the same atomic weight and specific gravity, and in its difficultly soluble chloride and insoluble iodide. Its sulphate is also difficultly soluble in water, though not to such an extent as lead sulphate ; its neutral solutions are precipitated by sulphuretted hydrogen, and the metal is set free from its soluble salts by metallic zinc. It differs from lead, but resembles silver and potassium, in its atomicity, which in the above compounds is always unity, and is similar to potassium in many other points. Although it does not decompose water like potassium, it easily dissolves in water con- taining air, and produces thallium hydrate, which is readily soluble in water, and of which the solution reacts strongly alkaline. Still more remarkable is the fact that a metal so similar to lead as thallium should possess a carbonate soluble in water, and even reacting alkaline. Twenty- five years ago, when thallium was un- known, it was universally held that only the alkali-metals yield soluble carbonates, and it was not thought possible that the car- bonate of a heavy metal could be soluble in water ; still less was it imagined that the sulphate of a heavy metal could displace potassium sulphate in common alum without changing the physical properties of the salt. We now know that thallous carbonate is soluble in water, giving an alkaline solution, and that thallous sulphate can unite with aluminium sulphate and produce a true alum. And further, thallous chloride, just like potassium chloride, can unite with platinum chloride, and produce a double salt inso- luble in water. Thallium is a rare element, and although present in many ores and mineral waters, is only found in minute quantities. The mineral richest in thallium is crookesite (called after Crookes, the discoverer of the element), which contains about 18 per cent, of 510 Text-Book of Inorganic Chemistry. the metal, but which is extremely rare. Thallium is best prepared from some varieties of iron, or copper pyrites. When pyrites con- taining traces of the metal is" used for the manufacture of sulphuric acid, the thallium passes over with the gases, and is then contained in the flue-dust or in the deposit which forms on the floor of the leaden chambers. This flue-dust or deposit when boiled with dilute sulphuric acid yields an impure solution of thallous sulphate, which when evaporated down and mixed with hydrochloric acid gives a precipitate of thallous chloride. This precipitate is filtered off, well washed with cold water, and then decomposed with con- centrated sulphuric acid, which converts it into thallous sul- phate. On dissolving this salt in water and dipping a plate of zinc (free from lead) into it, the thallium is precipitated as a crystalline, spongy mass, which is then pressed together and fused in a crucible. Thallium so obtained is of a tin-white colour when freshly cut ; it is very soft, and, like lead, may be cut with a knife or even scratched with the nail, and easily marks paper. It has a specific gravity of n - 9, melts at 290 , and volatilizes at a red heat. Thal- lium readily oxidizes when exposed to the air and becomes covered with a thin layer of oxide, but, notwithstanding its strong attraction for oxygen, it does not decompose water even at the boiling tem- perature. It preserves its bright surface under water which has been freed from dissolved air by boiling, and is therefore best kept under this liquid. But if the water is exposed to the air the thallium gradually becomes converted into thallous hydrate and thallous carbonate, both of which dissolve in the water. It easily dissolves in dilute nitric or sulphuric acid, but less readily in hydrochloric acid on account of the insolubility of its chloride. With oxygen, sulphur, and the halogens, thallium forms two series of compounds— the thallous and thallic compounds. In the former it is a monad element, and in the latter a triad. The former (thallous) series of compounds are the more stable, and thallous oxide is a much stronger base than thallic oxide. All thallium compounds are poisonous. If thallium is allowed to remain for a long time under water which has access to air or oxygen free from carbonic acid, the aqueous solution becomes strongly alkaline from the thallous hydrate : TIOH, which has been produced. When this solution is evaporated down, the hydrate crystallizes out in yellow rhombic prisms with 1 molecule of water. The same substance may be Compounds of Thallium. 5 1 1 more easily obtained by precipitating a solution of thallous sul- phate with baryta-water, filtering off the barium sulphate produced, and evaporating the clear solution. It is soluble both in water and alcohol. When heated to 100° it decomposes into water and thalloics oxide : T1„0, which forms a black powder, melting at about 300 to a dark yellow liquid. The oxide attracts moisture from the air and is again converted into the hydrate. Thallous Sulphate : SO.,(OTl)„ crystallizes in colourless rhombic prisms isomorphous with potassium sulphate, and unites SO 1 O Al with aluminium sulphate to form thallium-alum: en" f DTI + {ritr nT1 is also known. Thallous Phosphate : PO(OTl) 3 . — If a neutral solution of a thallous salt is mixed with sodium phosphate' and a few drops of ammonia added, thallous phosphate soon begins to crystallize out, and afterwards nearly fills the vessel. The crystals are long, glistening prisms, of the above composition, and almost insoluble in water containing a trace of ammonia. Thallous Carbonate : CO(OTl) 2 , is obtained by saturating a solution of thallous hydrate with carbonic acid, and crystallizes in colourless lustrous prisms when the solution is evaporated down. It is tolerably easily soluble in water, and the solution reacts alka- line. It loses carbonic acid when heated, and is converted into thallous oxide. Thallous Chloride : T1C1, is thrown down as a white curdy precipitate when a solution of a thallous salt is mixed with hydro- chloric acid. It is scarcely soluble in cold water, but dissolves more readily when heated. It easily melts when heated and then volatilizes. Thallous chloride when mixed with platinic chloride forms a yellow insoluble double chloride of the composition : 2TlCl,PtCl 4 , corresponding to the compound of potassium and platinic chlorides. Thallous Iodide : Til, like lead and silver iodides, is nearly insoluble in water, and is produced as a yellow amorphous pre- cipitate when potassium iodide is added to a solution of a thallous salt. 5 1 2 Text-Book of Inorganic Chemistry. Thallons Sulphide : T1_,S, is obtained as a black precipitate when excess of ammonium sulphide is added to a solution of a thallous salt, and is insoluble in excess of the reagent. Sulphu- retted hydrogen produces the same precipitate in a solution of thallous acetate. Thallic Compounds. — Thallic hydrate and oxide, unlike the corresponding thallous compounds, are insoluble in water, but are dissolved by acids to form soluble thallic salts. Thallic hydrate : TIO(OH), is produced as a brown amorphous powder when freshly precipitated thallous chloride is digested with an excess of sodium hypochlorite. At ioo° it breaks up into water and thallic oxide : T1. 2 3 , a black powder insoluble in water, but dissolving in hydro- chloric acid to form thallic chloride. When heated with concen- trated sulphuric acid, oxygen is given off and thallous sulphate remains behind ; when heated alone it breaks up into thallous oxide and oxygen. Thallic Sulphate : (SO a ) 3 0,.Tl/" + 7H.,0, crystallizes in colour- less thin scales when a solution of thallic hydrate in dilute sul- phuric acid is evaporated down. It loses only 6 molecules of water when heated to 220 , and is decomposed by water even in the cold, with separation of thallic hydrate. It forms a difficultly soluble double salt with potassium sulphate, which, however, is not a true alum. Thallic Chloride : T1C1 3 , is produced when thallous chloride is gently heated in a stream of chlorine, or when thallic hydrate is dissolved in hydrochloric acid. On evaporating its solution it crystallizes in easily soluble, deliquescent prisms, which decom- pose, when heated, into thallous chloride and free chlorine. Thallic Sulphide : Tl.^Sj, may be obtained by fusing thallium with an excess of sulphur. It then forms a black soft mass, which easily melts, and is brittle at low temperatures. Detection of Thallium Compounds. The compounds of thallium possess such characteristic proper- ties, that it is an easy matter to detect them, and to separate them from the compounds of other metals. The non-precipitation of thallium compounds by sulphuretted hydrogen from their acid Bismuth. 513 solutions, the solubility of thallous sulphate in hot water, the in- solubility of the chloride and iodide, and the solubility of thallous hydrate and carbonate in water, offer the best means of detection and separation. Minute traces of thallium can be readily recognized with the spectroscope ; its spectrum consists of a single intense green line (see table). It was this property of thallium which led to its dis- covery, and from which it derived its name (ddXKos = a green twig). BISMUTH. Chemical Symbol ': Bi. — Atomic Weight : 210. Bismuth is only sparsely distributed in nature, and nearly always occurs in the free state. From the earthy impurities which accompany it, the metal is separated either by simply melting it in inclined iron tubes, or else by roasting in the air, and then heating with carbon and a slag in suitable crucibles. Commercial bismuth always contains traces of other metals, including arsenic. It may be purified by fusion with nitre, which converts the arsenic into potassium arsenate and oxidizes other impurities. To prepare pure bismuth, a solution of its nitrate is precipitated with much water, which gives an insoluble basic nitrate, and this, when washed, dried, and reduced with a mixture of sodium carbonate and char- coal, yields the pure metal. Bismuth is a greyish-white metal, with a reddish tinge ; it is brittle, and can therefore be easily powdered, and crystallizes very easily. If bismuth is fused in a crucible, and when partly solidi- fied, the crust at the top broken, the still liquid metal poured out, and the crucible broken open when cold, its sides are found covered with beautiful iridescent crystals of bismuth, just as crystals of monoclinic sulphur are obtained in the same way. The colour of the crystals is due to a thin layer of oxide, with which the metal becomes covered when the air enters the crucible. The crystals are rhombohedra, the angle of which so closely approaches that of the cube, that bismuth was formerly thought to crystallize in the regular system. Bismuth has a specific gravity of 9-8, and melts at as low a tem- perature as 264° ; it shares with water the property of expanding considerably on passing from the liquid to the solid state. At high L L 514 Text-Book of Inorganic Chemistry. temperatures it volatilizes, but cannot be distilled, and burns when heated strongly in the air, producing yellow bismuth oxide. Hydro- chloric acid does not attack the metal, and sulphuric acid scarcely dissolves it, but it is at once dissolved by nitric acid, or by aqua regia, and converted into bismuth nitrate or chloride respectively. Bismuth unites with oxygen in three proportions and forms the compounds bismuth suboxide : BiO, bismuth oxide : Bi 2 3 , and bismuth pentoxide or bismuthic anhydride (unknown in the free state) : Bi 2 0.. Of these, bismuth oxide is the only important com- pound. Bismuth Suboxide : BiO, separates as a grey precipitate when caustic soda is added to admixture of bismuth chloride and stannous chloride. By the action of the caustic soda, stannous hydrate and bismuth hydrate are first produced, and the former then abstracts oxygen from the latter, forming stannic hydrate, which remains dissolved in the excess of soda, and bismuth sub- oxide. Like other suboxides, it is converted by acids into a bismuth salt and free bismuth. Bismuth Oxide : Bi,0 3 , remains as a yellow powder, similar to lead oxide, when basic bismuth nitrate is heated to redness. It melts at a higher temperature, and then solidifies to a crystalline mass on cooling. It dissolves easily in hydrochloric, nitric, or sul- phuric acid, producing the corresponding bismuth salts. Bismuth Hydrate : BiO(OH), is thrown down as a white preci- pitate when a solution of a bismuth salt is mixed with caustic soda. It dissolves easily in acids, but not in an excess of caustic soda. When dry it forms a white powder, which is converted into yellow bismuth oxide on heating. Bismuthic Acid: Bi 2 6 + xH„0, has not been much investigated, and is possibly a peroxide. It is obtained as a red heavy powder when chlorine is led through dilute caustic soda containing white bismuth hydrate in suspension, and is afterwards freed from alkali by washing with a dilute acid. It evolves oxygen when heated and sets free chlorine from hydrochloric acid. Bismuth oxide, like other sesquioxides, is only a weak base, and is partly separated from its salts by water, which in this case plays Bismuth. 5 1 5 the part of a base. Only a small number of bismuth salts are known. Bismuth Nitrate : (NO^OgBi + 5 H a O.— Metallic bismuth dis- solves easily in nitric acid, with evolution of nitric peroxide, and on evaporating the clear solution large colourless crystals of bismuth nitrate of the above composition separate out. The salt dissolves in a small quantity of water, especially if a little nitric acid is added, but a large quantity decomposes the salt like antimony trichloride, and produces a white crystalline precipitate consisting of delicate silky scales. This substance is basic bismuth nitrate, of the com- position : — N0 2 .OBi'(OH),, or (NO,) 3 3 Bi + 2Bi(OH) 3 . It is largely used in medicine in cases of cholera and chronic diarrhoea, and is also employed as a cosmetic. Bismuth Chroniate is deposited as a bright yellow precipitate, resembling lead chromate, when a solution of bismuth nitrate is mixed with one of potassium chromate. It is distinguished from the lead salt by its insolubility in a large excess of caustic soda. Bismuth Chloride : BiCl 3 . — If bismuth is heated in a retort and dry chlorine led into it, it burns and forms bismuth chloride, which distils over as a viscid liquid, and solidifies on cooling to a white crystalline mass. It boils at about 430 . The same com- pound may also be obtained by dissolving bismuth in aqua regia, evaporating to dryness, and distilling. It deliquesces in the air, and is decomposed by water, like the nitrate, producing a white crystalline powder of insoluble basic bismuth chloride (bismuth oxychloride) : BiOCl. Bismuth Sulphide : Bi,S 3 , occurs in nature as the mineral bismuthinite, and is thrown down as a dark-brown precipitate when sulphuretted hydrogen is led into a solution of a bismuth salt. It may also be prepared by fusing together bismuth and sulphur. Detection of Bismuth Compounds. — Bismuth, like lead, may be easily reduced by heating any of its compounds with sodium carbonate on charcoal in the reducing flame of the blowpipe, and the globules of metal are easily oxidized in the outer flame to yellow l l 2 5 16 Text-Book of Inorganic Chemistry. bismuth oxide, resembling lead oxide. The globules of reduced bismuth are, however, brittle, while those of lead are soft and malleable. Solutions of bismuth salts give, like lead, a black precipitate with sulphuretted hydrogen, but a solution of a bismuth salt gives a white precipitate of a basic salt when poured into water, and yields no precipitate when mixed with sulphuric or hydrochloric acid. Further, the precipitate of bismuth hydrate obtained with caustic soda is not soluble in excess. Bismuth oxychloride may be distinguished from the corres- ponding compound of antimony by its insolubility in a strong solution of tartaric acid, in which the latter easily dissolves. TIN. Chemical Symbol ■ Sn. — Atomic Weight : 118. Tin belongs to those few metals which have been known from the earliest times, and this is all the more remarkable as the metal never occurs native, and its ores are only sparingly found in nature. For the extraction of tin only one mineral is used — viz. cassiteiite or tinstone, consisting of stannic oxide, with the composition : SnO,. This mineral is found in large quantities in Cornwall, in the Erzgebirge (Saxony and Bohemia), and other European localities, as well as in Banca, the Malay Peninsula, Peru, and Australia. The extraction of the metal from its ore is a very simple process. The tin-stone is first roasted to get rid of arsenic and sulphur, then stamped to powder and washed with water to remove the lighter earthy impurities, and finally smelted with anthracite or powdered charcoal in a reverberatory furnace. The reduced tin, which sinks to the bottom of the hearth, is then withdrawn and refined by re- melting in iron pots and stirring with a green pole. Commercial tin is always more or less impure, and often contains arsenic. The pure metal is easily obtained by heating pure stannic oxide and charcoal in a crucible, with the addition of a little borax or some similar flux. Tin is a white metal, resembling silver, with a brilliant lustre, and a specific gravity of 7-3. It is soft and very malleable, so that it can be hammered out into thin foil (tin-foil), but possesses little tenacity. Up to ioo° its malleability increases, but when heated Tin. 5 1 7 above this it becomes less malleable, and is even brittle at 200 . Advantage is taken of this property to produce what is called grain-tin — i.e. tin which has been heated to about 200° and broken to pieces under the hammer. Tin melts at 228", and when broken shows a crystalline fracture. It is well known that a bar of tin when bent emits a peculiar noise — the so-called cry of tin, which is probably produced by the friction of the crystals of tin over one another. Commercial tin, when grasped in the hand, imparts to the skin a peculiar odour, not possessed by the metal itself. At the ordinary temperature, tin remains unchanged in the air, but becomes covered with a thin crust of white oxide when exposed to the air in the fused state. At a white heat it burns with a bright white light. Hot concentrated hydrochloric acid dissolves tin with evolution of hydrogen and produces stannous chloride ; the metal also dissolves in hot strong sulphuric acid, forming stannous sulphate and giving off sulphurous acid. Aqua regia dissolves the metal and forms stannic chloride. Ordinary concentrated nitric acid oxidizes the metal, with the production of large quantities of brown fumes and formation of white stannic oxide, insoluble both in water and nitric acid. Finally, tin is also dissolved when heated with caustic potash ; hydrogen is then set free and a solution of potassium stannate produced. The compounds of tin are divided into two classes — the stan- nous and the stannic compounds. In the former the metal is a dyad, in the latter a tetrad. Stannous Compounds. Stannous Oxide : SnO. — The precipitate of white stannous hydrate obtained when caustic soda is added to stannous chloride easily dissolves in an excess of soda, and produces sodium stannite. If the alkaline solution is boiled it undergoes decomposition, and a black crystalline precipitate of stannous oxide is thrown down. After washing and drying, the compound remains unchanged in the air. If, however, the solution of stannous chloride is precipi- tated with sodium carbonate instead of caustic soda, the white stannous hydrate : Sn(OH) 2 , which is now produced (together with free carbonic acid) behaves differently and oxidizes in the air to stannic oxide. The oxysalts of stannous oxide are but little known. Stannous sulphate : S0 2 - CXSn, may be prepared by dissolving tin in hot con- 5 1 8 Text-Book of Inorganic Chemistry. centrated sulphuric acid, and is then deposited in crystalline plates when the solution cools. When its aqueous solution is boiled it undergoes decomposition, with formation of a basic salt. Stannous Chloride : SnCl,,. — Tin when heated in dry hydro- chloric acid gas is gradually converted into a white crystalline mass of stannous chloride, which melts at 250 and sublimes at a higher temperature. Stannous chloride is, however, usually prepared by dissolving tin in hydrochloric acid, an operation which may be con- siderably accelerated by the addition of a few scraps of platinum foil or of a few drops of platinum chloride, which is at once reduced to metallic platinum. The contact of the platinum with the tin makes the latter far more electro-positive and causes it to dissolve more quickly. 1 On evaporating the solution of stannous chloride the salt crystallizes out with two molecules of water as the com- pound : SnCl 2 + 2H 2 0. This salt is manufactured on a large scale, and is brought into trade under the name of tm-salt. The readi- ness with which stannous chloride passes into stannic chloride makes it a powerful reducing agent. Mercuric chloride is at once reduced by it to mercurous chloride, and on warming to metallic mercury, while the stannous chloride becomes stannic chloride. An aqueous solution, of stannous chloride becoiries turbid when ex- posed to the air, and gradually deposits a basic stannic chloride. Stannous chloride is largely used in dyeing and for purposes of reduction. Stannous Sulphide : SnS. — On heating a mixture of tin-filings with an excess of sulphur in a closed crucible the two substances unite and form .dark-grey, lustrous scales of stannous sulphide. The same compound is obtained as a dark-brown precipitate when sulphuretted hydrogen is led into a solution of stannous chloride. The precipitate dissolves in hydrochloric acid with evolution of 1 The action of the platinum, which is deposited in a finely divided state on the tin, may be thus explained. Each particle of the platinum forms a galvanic couple with a corresponding particle of tin, and a series of tiny electric currents circulate between the particles of the two metals and the liquid in their neigh- bourhood. The production of these currents necessitates a more rapid solution of the tin than when the platinum is absent, and they are consequently not pro- duced. Other similar couples, as this arrangement has been called, are also occasionally used. Thus, a copper-zinc couple, prepared by treating zinc foil with a very dilute solution of copper sulphate, reduces many substances which zinc alone does not attack, or only very slowly. — Ed. Stannic Compounds. 5 ! 9 sulphuretted hydrogen, but is insoluble in pure, freshly-prepared ammonium sulphide, and is so distinguished from stannic sulphide, which at once gives a soluble compound with ammonium sulphide. Stannous sulphide at once dissolves in ordinary yellow ammonium sulphide, because this compound contains dissolved sulphur, which converts the stannous sulphide into stannic sulphide, and this dis- solves in the ammonium sulphide. Stannic Compounds, Stannic Oxide : SnO, 2 , constitutes the mineral cassiterite or tin- stone, and is found either in compact masses or else in brown quad- ratic crystals. It may be obtained as a white amorphous powder by burning tin in the air or by heating stannic hydrate. In the crystalline state it is formed when the vapours of stannic chloride and water are led through a red-hot porcelain tube. Stannic oxide is insoluble in acids, and is not attacked when fused with acid potassium sulphate, but may be converted into soluble sodium stannate by fusion with caustic soda or sodium carbonate. Stannic oxide has far more the character of an acid than of a base. Stannic Hydrate. — Several hydrates of stannic oxide are known, and are all weak acids and very weak bases. If tin is warmed with moderately strong nitric acid, a violent reaction ensues, tor- rents of nitric peroxide are evolved, and a white insoluble powder of metastannic acid, of the probable composition : Sn(OH),, re- mains behind. This hydrate is insoluble in water and nitric acid, but is converted by strong hydrochloric acid into a gelatinous mass soluble in water. The solution in water contains stannic chloride, and on the addition of sulphuric acid a white precipitate of stannic sulphate is obtained. This is an unstable compound, from which water again abstracts all the sulphuric acid. Metastannic acid, when heated to 100°, loses water and becomes converted into a compound having the composition : SnO(OH)„, but which is different from stannic acid ; at a red heat both this compound and metastannic acid are converted into stannic oxide. Meta- stannic acid when digested with cold caustic soda forms sodium metastannate, a slightly soluble salt of which the composition is uncertain. But if the hydrate is boiled with caustic soda an easily soluble salt is formed, which can be obtained in the crystalline form, and which is known as sodium stannate, of the composition : SnO(ONa).j. The same salt is also produced when 520 Text- Book of Inorganic Chemistry, stannic oxide is fused with caustic soda. It is largely used in calico-printing. The corresponding potassium stannaie : SnO(OK) 2 + 3H.,0, is obtained in the same way as the sodium compound. It is insoluble in caustic potash, but soluble in water, and is deposited in colourless crystals ; its solution reacts strongly alkaline. A second variety of stannic hydrate is known under the name of stannic acid, the composition of which is probably : SnO(OH).,, corresponding to the above stannates. It is obtained as a white, voluminous precipitate when stannic chloride is mixed with sodium carbonate. Stannic acid is insoluble in water, but dissolves easily in strong hydrochloric or nitric acid as well as in caustic soda. The acid solution when boiled gradually deposits stannic acid, provided too much free acid is not present. Stannic acid, when kept for some time under water, becomes insoluble in nitric acid. Stannic chloride : SnCl 4 . — This volatile liquid, which was previously known as Spiritus Libavii fumans, is obtained when tin is heated in a retort and chlorine passed over it. The dis- tillate when purified by redistillation is a colourless heavy liquid of specific gravity 2-28, and boiling at 114 . It fumes strongly in moist air, almost as much as sulphuric anhydride, and hisses when dropped into water, forming the hydrated salt: SnCl 4 + 3H.,0, which crystallizes out. A dilute aqueous solution is decomposed on boiling, with separation of stannic hydrate. A solution of stannic chloride is best obtained by dissolving tin in aqua regia. Stannic chloride unites with the alkaline chlorides and forms crystalline double salts. The double compound with ammonium chloride: SnCl 4 + 2NH 4 C1, is used by calico-printers under the name of pink-salt. Stannic Sulphide : SnS„, is obtained as a yellow amorphous precipitate when sulphuretted hydrogen is led through a not too acid solution of stannic chloride. It is insoluble in water, but dissolves in strong hydrochloric acid. Hot concentrated nitric acid oxidizes it to stannic hydrate and sulphuric acid. Stannic sulphide possesses the character of a sulpho-acid, and as such easily dissolves in ammonium sulphide, producing soluble ammonium sulphostannate, which is easily decomposed by acids with re-sepa- ration of stannic sulphide. Stannic sulphide is prepared in the dry way as golden scales Stannic Compounds. 521 by heating together tin and sulphur with volatile substances — such as mercury and ammonium chloride — which, by their volatilization, appear to prevent the temperature rising high enough to decompose stannic sulphide into stannous sulphide and sulphur. A solid amalgam of tin is first prepared by adding six parts of mercury to twelve parts of molten tin. This is powdered when cold, mixed with seven parts of flowers of sulphur and six parts of ammonium chloride, and the mixture introduced into a glass flask, the mouth of which is loosely closed by a piece of chalk. The flask is next placed in a crucible surrounded with sand and heated in a good draught to low redness. During the process mercuric chloride, ammonium and stannic chlorides sublime, and at the bottom of the flask is found the stannic sulphide as a loose mass of a golden-yellow colour. The substance so prepared is called mosaic gold; it is much more stable than the precipitated sulphide, being scarcely attacked by strong hydrochloric or nitric acid. Mosaic gold is used for a number of technical purposes, but chiefly for cheap gilding. More important than the technical applications of the com- pounds of tin are those of the metal itself, either alone or alloyed with other metals. Considerable quantities of tin are consumed in the manufacture of tin-plate — i.e. iron plates coated with a thin layer of tin ; copper vessels are tinned internally, and tin-foil is largely used for packing small articles. Tin-foil usually contains traces of copper, which gives it greater tenacity. Not unfrequently it also contains lead, which may produce very deleterious effects if articles of food are packed in the foil. Not only is lead more easily attacked by the acids in the food than tin, but the compounds of lead are far more injurious when introduced into the system than those of tin. Alloys of Tin. — Different alloys of lead and tin— known as ■pewter — are used for the manufacture of various vessels, and as a basis for the cheaper silver-plated articles ; common pewter con- tains about four parts of tin to one of lead. Common soft-solder, chiefly used for soldering tin-plate, is an alloy of 2 parts of tin with 1 part of lead, and melts at about 180 , or of equal parts of the two metals, with a melting-point of about 200°. Britannia-metal, a better kind of pewter, is essentially an alloy of about 9 parts of tin with 1 part of antimony. The various alloys of tin with 522 Text-Book of Inorganic Chemistry. copper, bronze, &c, will be described under the latter metal ; an amalgam of tin is used for silvering looking-glasses. By melt- ing together tin with other easily fusible metals — e.g. lead, bis- muth, and cadmium — alloys are obtained with very low melting points, and which are largely used in the process of stereotyping. An alloy of I part of lead, I part of tin, and 2 parts of bismuth {fusible metal), melts at 95 ; and one consisting of 4 parts of bismuth, 2 parts of lead, 1 part of tin, and 1 part of cadmium ( Wood's metal), melts at about 68°. Detection of Tin Compounds. Tin and antimony are the only elements which are converted by nitric acid into insoluble white oxides ; of these the antimony compound may be distinguished by its solubility in a solution of tartaric acid, which leaves the stannic oxide undissolved. Tin itself may be distinguished from antimony by its solubility in hydro- chloric acid, which does not attack the latter substance. Stannous compounds, in consequence of the readiness with which they pass into the stannic form, are powerful reducing agents. Stannous chloride precipitates mercury and gold from solutions of their salts, the latter even in the cold. With sulphuretted hydrogen stannous salts give a dark-brown precipitate of stannous sulphide, and the same reagent throws down yellow stannic sulphide when led through solutions of the stannic salts. Both sulphides dissolve in yellow ammonium sulphide. Stannic oxide has more the character of an acid than a base ; even stannous oxide possesses faint acid properties, and unites both with strong acids and with strong bases to form salts. Metallic tin, as white malleable globules, may be reduced from its compounds when a compound of the metal is mixed with sodium carbonate and heated on charcoal in the inner blowpipe flame, though the reduction is rather more difficult than that of lead. At the same time, an indistinct white incrustation is formed on the charcoal. Compounds of tin do not tinge a bead of borax, but if a minute trace of copper is added, and the bead held in the reducing flame of the blowpipe, it easily becomes of a fine red colour, due to reduced cuprous oxide. Copper. 523 COPPER. Chemical Symbol: Cu. — Atomic Weight: 63-4. Copper, which has been known from the earliest times, is one of the few metals which possess a decided tint ; it is of a bright red colour. The metal is sometimes found in the free state in nature, often well crystallized in octahedra, but more commonly occurs in combination with other elements. The most important ores of copper are : cuprite or red copper ore (cuprous oxide : Cu a O), copper-glance (cuprous sulphide : Cu. 2 S), copper-pyrites (sulphide of copper and iron : CuFeS^), purple copper ore of similar composition, and tetrahedt ite or fahlers (a mixed sulphide of copper, antimony, arsenic, &c). Besides these, some oxy-salts of copper are also found in the mineral kingdom, among which may be mentioned the basic carbonates : malachite and azurite. The extraction of copper from its oxygen compounds (cuprite, malachite, &c.) is very simple, and analogous to the method used for the reduction of iron. The roasted ores are heated in a blast furnace with some form of carbon, and the reduced copper collects or. the hearth under the slag. Only small quantities of copper are extracted from these comparatively rare ores, far more is obtained from copper-pyrites, a double sulphide of copper and iron. The method employed for the separation of copper from these ores is a complex one ; a whole series cf processes is necessary to com- pletely separate both the iron and the sulphur. These processes, which depend upon the attraction of copper for sulphur and iron for oxygen, consist in alternately roasting and melting the ores. The roasting gradually removes the sulphur and oxidizes the iron, which unites with silica to form a fusible slag. From the copper sulphide which then remains, metallic copper is obtained by oxi- dizing the sulphur. Commercial copper is never quite pure, although the impurities are usually insignificant. The pure metal may be obtained as a red powder by heating pure copper oxide in a stream of hydrogen. Copper melts only at a high temperature, above i,ooo°, and burns when heated in the oxy-hydrogen flame. Its specific gravity is about 89. The metal takes a good polish, is malleable and ductile, and very tenacious. Bright copper remains unchanged in dry air, but in the presence of moisture becomes covered with a . green layer of basic copper carbonate. This green rust of copper 524 Text-Book of Inorganic Chemistry. is often called verdigris, which is really a basic acetate of the metal. Copper is more rapidly corroded if the surface is moistened with a dilute acid and exposed to the air ; it is also attacked by ammonia or a solution of common salt in the presence of air. In the absence of air neither hydrochloric nor dilute sulphuric acid has any action on the metal. Hydrochloric acid does not dissolve it even on heating, but hot concentrated sulphuric acid oxidizes it to copper sulphate, and sulphurous anhydride is set free (p. 151). Unlike iron, zinc, and other metals, copper cannot decompose water and form copper oxide and hydrogen ; even when copper is heated to redness in steam no change is produced. If copper is heated to redness in the air it becomes covered with a layer of black cupric oxide which separates as scales when the metal is hammered. This copper scale consists chiefly of cupric oxide, but contains some metallic copper and some cuprous oxide. The best solvent for copper is nitric acid, which dissolves it with evolution of nitric oxide (p. 193), and produces a blue solution of cupric nitrate. Copper is deposited from its solutions by more positive metals — e.g. zinc or iron — and also separates in a compact form at the negative pole when an electric current is passed through a solution of the sulphate. The process of producing electrotypes or copies of various objects in copper depends upon this deposition of the metal by an electric current. Compounds of Copper. Copper forms with oxygen; sulphur, and the halogens, two series of compounds ; the one series contains a single atom of dyad copper, and the other 2 atoms of copper which also play the part of a dyad. The two series of compounds are called cupric and cuprous compounds respectively — e.g. cupric chloride : CuCl 2 , and cuprous chloride : Ca,Cl 2 . Of these two series, the cupric com- pounds are the most important, and shall therefore be first dis- cussed. All compounds of copper are poisonous, especially those which are soluble in water. Cupric Compounds. Cupric Oxide, Black Oxide of Copper : CuO, is a black amor- phous powder, obtained by glowing finely divided copper in the air, or better by heating cupric nitrate. Copper oxide is largely used in organic analysis to oxidize organic substances (p. 289). For this purpose it does not require to be pure, but should be in the form of Cupric Compounds. 525 scales, and is best prepared from copper scale. A layer of the scale, which has been freed from mechanical impurities, is moistened with nitric acid (containing no chlorine) in a Hessian crucible, then a second layer added and also moistened with the acid until the crucible is about three parts full. The whole is then heated to low redness until all the nitric acid is destroyed and expelled. The nitric peroxide set free during the decomposition of the cupric nitrate suffices to oxidize the copper and cuprous oxide, contained in the scale, to cupric oxide, and if the temperature has not been raised high enough to cause the oxide to cake together, it is ob- tained in the form of scales well suited for organic analysis. Cupric oxide may be obtained in the wet way as a black pre- cipitate, by adding hot caustic soda to a hot solution of a cupric salt. Cupric oxide is completely insoluble in water, but readily dissolves in hydrochloric, nitric, or sulphuric acid, without any evolution of gas, then forming the corresponding cupric compounds. When heated in a stream of hydrogen it easily parts with its oxygen and yields metallic copper as a red powder. Cupric Hydrate : Cu(OH). 2 . — If a solution of a cupric salt is mixed with an excess of caustic soda at the ordinary temperature, a blue voluminous precipitate of copper hydrate is obtained, in which the copper oxide and water are so loosely combined that the compound is decomposed at 100°, and black cupric oxide even separates slowly at the ordinary temperature. If insufficient caustic soda is employed a basic salt is obtained, which is also of a blue colour, but which remains unchanged when boiled. Cupric Sulpbate, Blue Vitriol: S0 2 -0 2 Cu + 5H..O, is de- posited, when its aqueous solution is evaporated, as large blue crystals belonging to the triclinic system, and is the salt from which nearly all the other compounds of copper are obtained. Copper sulphate may be prepared by dissolving copper oxide in sulphuric acid, or by heating copper with concentrated sulphuric acid. Considerable quantities of copper sulphate are produced by the oxidation of copper-pyrites when roasted in the air, but always mixed with more or less ferrous sulphate. It cannot be completely freed from this impurity by simple recrystallization as the two salts crystallize together, but if sufficient nitric acid is added to oxidize the iron to the ferric state, the solution evaporated to dryness, and the dry mass extracted with water, most of the iron remains behind as a basic salt. Copper is precipitated from the blue solution of 526 Text-Book of Inorganic Chemistry. cupric sulphate, produced in mines where copper-pyrites is worked, by adding scrap iron. The metal so obtained is known as ceme7it- copper. Cupric sulphate is easily soluble in water ; one part of the salt dissolves in 2-5 parts of cold water or in o # 5 part of boiling water. The blue crystals slightly effloresce when exposed to the air, and lose 4 molecules of their water of crystallization at 100°, but the fifth molecule of water (water of constitution) is only expelled at about 200 , The anhydrous salt is white, but regains its blue colour in contact with water, being reconverted into the hydrated com- pound. The energy with which anhydrous cupric sulphate unites with water is so great that we employ it to abstract the last traces of water from 99 per cent, alcohol, and to remove traces of water from many other liquids in which the salt does not dissolve. Cupric sulphate unites with the sulphates of the alkalies and forms double salts, which may easily be obtained in the crystalline form combined with 4 molecules of water. They are of a lighter colour than cupric sulphate and less soluble in water. Cupric Ammonium Sulphates. — If a concentrated aqueous solu- tion of cupric sulphate is mixed with strong ammonia until the precipitate of cupric hydrate which is first produced is redissolved, a deep blue liquid is produced, which, on standing, or better, when covered with a layer of alcohol, deposits a layer of deep blue crys- tals of the composition: SO„-0 2 Cu + 4NH 3 + H.,0. These crys- tals gradually decompose and become opaque when exposed to the air, and when heated to 1 50 lose their water with 2 molecules of ammonia, leaving a green powder of the composition : SO , ■ 0.,Cu + 2NH3. The formation and constitution of these two compounds may be represented as follows (see p. 563). On the addition of ammonia to cupric sulphate, ammonium sulphate and cupric hydrate are first produced, and the latter compound then appears to unite with ammonia and form cupri-diammonium hydrate : — r fOH NH OH-NH.l,, Cu| QH + 2NH3 = OH . NH ^[Cu, or the hydrate of a hypothetical diammonium in which 2 atoms of hydrogen are displaced by 1 atom of dyad copper. This sub- stance does not appear to exist in the free state, but unites at once with the ammonium sulphate, and produces the blue salt, diammon-cupri-diammonium sulphate, thus : — Cupric Compounds. 527 c n |ONH 4 OH-NH 3 } r „_. fONH,.NH,l^ xi ^ Ammonium Cupri-diammonium Diammon-cupri- sulphate hydrate diammonium sulphate which may be considered as ammonium sulphate in which two atoms of hydrogen have been displaced by the dyad radical cupri- diammonium: jTTj'rCu. On heating the blue compound to 150° it parts with two molecules of ammonia, and leaves the green powder mentioned above, which is cupri-diammonium sulphate : — cri (ONH.-NH,) ,. cr . (ONH,l„ „„ S0 ^ONH3-NH'f Cu = SC MoNH3, Cu + 2NH 3' Diammon cupri- Cupri-d ammonium diammonium sulphate sulphate The constitution of the ammoniacal compounds of cobalt (p. 487) is possibly similar to that of the above salts. Cuprie Nitrate: (NOJ.,0 2 Cu + 3H.,0, is easily obtained by dissolving copper in nitric acid, and separates out when the solu- tion is concentrated as dark blue deliquescent crystals, easily soluble both in water and in alcohol. Cupric Phosphate : (PO),0„Cu, + jH 2 0, is thrown down as a blue-green precipitate when ordinary sodium phosphate is added to a solution of cupric sulphate. Various phosphates of copper occur in the mineral kingdom. Cupric Arsenite, probably : As | .-> "y, is obtained as a green precipitate varying in tint when a solution of sodium arsenite is poured into one of copper sulphate, with constant stirring. This compound is used as a pigment under the name of Scheele's green. A second green pigment — known as Schweinfurt-green — is a double compound of copper arsenate and acetate. Both these compounds are highly poisonous, and their use is therefore very limited. Cupric Carbonates. — All the known carbonates of copper are basic compounds, even when normal copper sulphate is mixed with normal sodium carbonate, the bluish-green gelatinous precipitate, which after a time changes to a bright green powder, consists of a basic carbonate of the composition: CO0 2 Cu + Cu(OH) 2 = 528 Text-Book of Inorganic Chemistry. CO J (~>p U r)H ^^ same compound is found in nature, espiceally in Siberia, as the beautiful mineral, malachite, which in its massive form is used for the manufacture of vases, tops of tables, &c. A second basic carbonate of copper is the dark blue mineral asurite, of the composition : 2CO -0 2 Cu + Cu(OH).,. Cupric Acetates. — These compounds belong strictly to organic chemistry, but it may be remarked that normal cupric acetate is soluble in water and is deposited from its solution in blue crystals. The basic acetates are insoluble or difficultly soluble in water. One of these, or rather a mixture of several, is the poisonous substance called verdigris. Cupric Chloride : CuCl 2 + 2H 2 0,is obtained by dissolving cupric oxide in hydrochloric acid, and on concentrating the solution is. deposited in small, green, rhombic prisms. When heated the salt loses its water and the anhydrous chloride remains behind as a yellowish-brown powder, which rapidly attracts moisture from the air and deliquesces. When more strongly heated, it is decomposed into cuprous chloride and free chlorine. Cupric chloride unites with cupric hydrate and forms basic compounds. One of these is the mineral atacamite, and has the composition: CuCL, + 3Cu(OH) 2 . These insoluble basic chlorides may be obtained by mixing a solution of cupric chloride with a quantity of caustic soda insufficient to convert all the copper into cupric hydrate. Cupric Bromide : CuBr 2 , closely resembles the chloride. Cupric Iodide, strange to say, has not been prepared. If potassium iodide is added to a solution of cupric sulphate, a white precipitate of cuprous iodide is produced and iodine is set free. Cupric Sulphide : CuS, is thrown down as a black precipitate when sulphuretted hydrogen is led through a solution of a cupric salt. In the moist state it readily attracts oxygen from the air and is converted into cupric sulphate, and must therefore be washed with water containing sulphuretted hydrogen and dried in a vacuum. When heated in a stream of hydrogen it parts with one-half of its sulphur and is converted into cuprous sulphide : Cu 2 S. Cupric Cuprous Compounds. 5 2 9 sulphide is found in nature as the mineral covellite in dark-blue hexagonal prisms. Cuprous Compounds. Cuprous Oxide : Cu.O. — This compound is found in nature as the mineral cuprite or red copper ore either in compact crys- talline masses, or else in small regular octahedra. The crystalline compound may be artificially prepared by fusing solid cuprous chloride with dry sodium carbonate. It then remains as a red crystalline powder when the fused mass is afterwards extracted with water. It may also be prepared in the same form but in smaller crystals, by gently warming a solution of cupric sulphate with an excess of caustic soda and grape sugar. The grape sugar is then oxidized by half the oxygen contained in the copper oxide. Cuprous oxide remains unchanged in the air and is quite insoluble in water. When treated with hydrochloric acid, it is converted into white cuprous chloride, which dissolves in excess of the acid, and forms a colourless solution in the absence of air. Dilute sulphuric acid converts it into cupric sulphate with the separation of one-half of the copper as a red powder. Other oxy-acids behave in the same way, except nitric acid, which completely dissolves it. Cuprous Hydrate: Cu 2 (OH),. — When a solution of cuprous chloride in hydrochloric acid is poured into an excess of caustic soda, cuprous hydrate is thrown down as a yellow crystalline precipitate. The compound is very unstable, rapidly absorbing oxygen from the air, and producing blue cupric hydrate. The cuprous oxy-salts are very unstable, and are scarcely known. The corresponding compounds with the halogens are, however, more permanent. Cuprous Chloride : Cu 3 Cl 2 . — This substance is a white crys- talline powder, insoluble in water, but soluble in hydrochloric acid, from which it crystallizes in white tetrahedra. Cuprous chloride may be obtained by boiling cupric chloride with excess of hydro- chloric acid and metallic copper in a stream of carbonic acid to exclude the air, the oxygen of .which would oxidize the cuprous chloride thus produced. On afterwards pouring the clear acid solution into a large quantity of cold boiled water, the cuprous chloride is deposited as a white crystalline powder. Cuprous chlo- ' M M 530 Text-Book of Inorganic Chemistry. ride may also be prepared by adding a solution of stannous chloride containing free hydrochloric acid to a solution of cupric chloride. It then crystallizes out also as white tetrahedra, while the stannic chloride, produced at the same time, remains dissolved in the acid liquid. Cuprous chloride when exposed to the air in the moist state becomes of a green colour, and at ioo c large quantities of oxygen are absorbed, forming the oxy-chloride : CuCl.„CuO. But if this compound is further heated (up to about 400 ), it again parts with its oxygen and leaves cuprous chloride. By this means, consider- able quantities of oxygen may be obtained from the air. Ammonia, as well as hydrochloric acid, dissolves considerable quantities of cuprous chloride, and both solutions absorb carbonic oxide. The compound which is produced is deposited in colour- less scales when a strong solution of cuprous chloride in hydrochloric acid is saturated with the gas. The ammoniacal solution also absorbs and unites with certain gaseous hydrocarbons, especially acetylene. Cuprous Iodide : Cu 2 I 2 , is a white powder insoluble in water and dilute acids, and is deposited when potassium iodide is added to a solution of cupric sulphate containing sulphurous acid or ferrous sulphate, these two substances acting as reducing agents. The insolubility of cuprous iodide makes this salt useful for sepa- rating iodine from solutions of its salts. Unlike cuprous chloride, the iodide remains unaltered when exposed to the air. Cuprous Sulphide : Cu 2 S, is obtained by heating finely-divided copper with sulphur, or by heating cupric sulphide in a stream of hydrogen. It forms a dark-grey mass, which melts at a red-heat, and when slowly cooled may be obtained in a crystalline form. Cuprous sulphide is found in nature as copper-glance, which crys- tallizes in fine rhombic prisms. It also occurs united with other sulphides — e.g. with iron sulphide in copper-pyrites and purple- copper ore, as well as with the sulphides of silver, antimony, arsenic, &c, in other minerals. Alloys of Copper. — Few metals are so largely used as copper, both in the form of its salts, and as alloys with other metals. The most important of these alloys are those with zinc, tin, nickel, and silver respectively. Brass is an alloy of copper and zinc obtained by melting the Alloys of Copper. 531 two metals together, and is of a darker or lighter colour accordingly as it contains more or less copper. Ordinary brass contains about 70 parts of copper to 30 parts of zinc. A larger proportion of zinc makes an alloy of a pale yellow colour, known as Muntz-metal, and largely used for sheathing wooden-ships. By increasing the percentage of copper and adding a little tin, a more malleable alloy is obtained, which can be beaten into thin sheets, and is called tombac or Dutch metal. Gun-metal, bronze and bell-metal are all alloys of copper and tin, with about 80 to 90 per cent, of copper, to which a little lead or zinc is often added. Bronze contains less copper than gun- metal, and bell-metal less than bronze. Speculum-metal is a hard white alloy of about 1 part of tin and 2 parts of copper, to which some arsenic is not unfrequently added. It takes a high polish, and is used for the mirrors in reflecting telescopes. An alloy of copper and nickel is used for small coins in Germany, Belgium, and some other countries ; in Switzerland the small coins are made of an alloy of copper, nickel, zinc, and silver. German-silver or nickel-silver is an alloy of copper, nickel, and zinc (p. 492). All the silver used in the art$ consists of an alloy of copper and silver ; and, finally, an alloy of copper and aluminium, containing about 9 parts of the former metal to 1 part of the latter, forms aluminium-bronze or aluminium-gold, a tena- cious, bright yellow alloy, largely used in the manufacture of small articles. Detection of Copper Compounds. Copper belongs to those metals of which the sulphides are in- soluble in dilute acids, and are separated when a stream of sulphu- retted hydrogen is led through an acid solution of one of their salts. The cupric sulphide then obtained is black and resembles lead, bismuth, and some other sulphides. This precipitation of cupric sulphide from acid solutions serves to distinguish copper from nickel, which it otherwise resembles in some points. A character- istic test for a solution of a copper salt is the bluish precipitate produced on the addition of ammonia and the dark blue solution obtained when the ammonia is added in excess. A solution of a ■ copper salt, even when extremely dilute, gives, on the addition of potassium ferrocyanide, a reddish brown precipitate of cupric ferro- cyanide : Ca,FeCy e ; only uranium compounds form a precipitate of this colour with potassium ferrocyanide. M M 2 532 Text-Book of Inorganic Chemistry. Small quantities of a compound of copper may be easily de- tected by heating with borax on a platinum wire before the blow- pipe. A green bead is then formed when heated in the outer flame, which becomes colourless or of a dirty red colour (from cuprous oxide) in the reducing flame. The red cuprous oxide may be more readily obtained by the addition of a trace of stannic oxide, or any compound of tin, and then heating in the reducing flame. 533 NOBLE METALS. This group includes mercury, silver, gold, platinum, and five other rare metals allied to platinum ; they have been called noble because they do not tarnish in the air at the ordinary temperature. They are further distinguished by the fact that their oxides are all decomposed when heated into the metal and free oxygen, with the exception of those of osmium and ruthenium, which volatilize un- changed. MERCURY. Chemical Symbol : Hg. — Atomic Weight : 200. This metal, with the exception of gallium, is the only one which is liquid at the ordinary temperature, and as it possesses a white colour like silver, is sometimes called quicksilver. It is only sparsely distributed in nature, but large quantities are found in the few localities where it occurs. Mercury is sometimes found free in the mineral kingdom or alloyed with silver or gold, but far more commonly in combination with sulphur as cinnabar : HgS. The mercury mines of Idria in Austria and of Almaden in Spain have been worked for centuries, and considerable quantities of the metal are also obtained from China. In late years large deposits of cin- nabar have been discovered in California. To extract mercury from cinnabar on a small scale the pow- dered ore is mixed with quicklime and heated in a retort, the neck of which dips under water. During the heating calcium sulphide and mercuric oxide are formed, and the latter compound then breaks up into mercury and oxygen. The mercury then condenses under the water, and the oxygen partly escapes, but principally serves to convert the calcium sulphide into calcium sulphate. On the large scale mercury is obtained from cinnabar by roasting the ore with excess of air in specially constructed furnaces, 534 Text-Book of Inorganic Chemistry. and passing the vapours of sulphurous anhydride and mercury which are thus produced through a series of chambers or through rows of earthenware vessels, where the mercury condenses. Commercial mercury, even when it has been distilled, is never pure, but contains mechanical impurities as well as traces of zinc, lead, and other less volatile metals dissolved in it. The presence of these impurities may be easily detected by allowing a drop of the metal to flow over the hand, when it runs off clear if pure, but leaves a dirty, dull tail on the skin when contaminated with other metals. Mechanical impurities may be separated by filtering the mercury through a filter of writing-paper with a pin-hole at the base. To separate the dissolved metals the mercury is placed in a flat dish and covered with a layer of ordinary nitric acid. The acid then oxidizes some of the mercury and produces mercurous nitrate, which then transfers part of its oxygen to the metals mixed with the mercury. After the whole has stood for several days, with frequent stirring, the acid liquid is poured off, the metal washed with water, then run through a separating-funnel and dried. This method separates most of the impurities contained in the metal, but chemically pure mercury can only be obtained by other and more tedious processes. Mercury has a specific gravity of 13-55 > ll solidifies at -40° to a crystallline malleable mass, with a specific gravity of 14-4. Although the metal only boils at 357 , it volatilizes even at the ordinary temperature of the air ; if a piece of gold leaf is suspended in a jar containing a little mercury it soon becomes white, and amalgamated from the mercury with which it has combined. Mercury remains unchanged in the air at the ordinary temper- ature, but when heated to about 300 it combines with oxygen and becomes changed into red mercuric oxide, which coats the metal. It is completely insoluble in water and is not attacked by hydro- chloric or dilute sulphuric acid. Concentrated sulphuric acid is also without action upon it at the ordinary temperature, but when heated dissolves it, producing mercuric sulphate and sulphurous anhydride. The best solvent for mercury is nitric acid, with which it forms either mercuric or mercurous nitrate according to different conditions. The vapours of metallic mercury and mercury compounds, especially those soluble in water, are all poisonous. Even the lower animals cannot exist in an atmosphere containing traces of mercury vapour. Entomological collections may be effectually Mercury. 535 preserved from destruction by minute animals if a few drops of mercury are placed in the cases containing the specimens. Compounds of Mercury. Mercury, like copper, forms two series of compounds — the one containing a single atom of the dyad metal, and the other con- taining two atoms of mercury, which also perform the functions of a dyad. These two series of compounds are the mercuric axiA mer- curous compounds respectively. Many are highly valued and powerful medicines. Mercuric Compounds. Mercuric Oxide : HgO. — If mercury is heated in the air to about 300 , it gradually becomes oxidized and converted into a red powder of mercuric oxide (p. 10). This change goes on too slowly to prepare any quantity of the oxide, and it is preferable to obtain it from its nitrate. For this purpose dry mercuric nitrate is mixed with an equal weight of mercury, and the mixture gently heated as long as acid vapours are given off. The oxide then remains as a heavy,' red, crystalline powder, of specific gravity 1 1 -2, and which becomes of a paler colour when rubbed. Besides this crystalline form of mercuric oxide, a second modi- fication may be obtained as an amorphous yellow powder, by precipitating a solution of mercuric nitrate with caustic soda. Mercury, like the other noble metals, is precipitated from its solu- tions by the alkalies and other strong bases as the oxide and not as the hydrate. Mercuric hydrate is unknown. Mercuric oxide, as previously stated (p. 10), breaks up into mercury and oxygen when heated. Mercuric Sulphate: SO a ;0.,Hg, is obtained by dissolving mercuric oxide or metallic mercury in sulphuric acid, in the latter case with evolution of sulphurous anhydride. It unites with a small quantity of water to form a colourless crystalline compound containing 1 molecule of water. With a large quantity of water it is decomposed into free sulphuric acid and a basic sulphate of the composition : SO, 2 '0 2 Hg + 2HgO, which separates as a heavy lemon-coloured powder. This compound was formerly used in medicine under the name of turpeth mineral. Mercuric sulphate 536 Text-Book of Inorganic Chemistry. is manufactured on a large scale, and used for the preparation of the two chlorides of mercury. Mercuric Nitrate: (NO.J a O,Hg + 8H0O, may be prepared by acting on mercury with an excess of nitric acid, until a drop of the solution no longer gives a precipitate with hydrochloric acid — a proof that it contains no mercurous nitrate. The acid solution, when sufficiently concentrated, and when cooled to — 15 , deposits large, colourless, rhombic crystals of the above composition. This salt is unstable ; it melts in its water of crystallization at 7°, and changes into a basic salt of the composition : (NO.,) 2 2 Hg + HgO + 2H.,0, which is also deposited when a hot concentrated solution of mercuric nitrate is allowed to cool. Both these salts are decomposed by water, yielding a still more basic compound : (NO„).,0,Hg +- 2HgO + H 2 0, which separates as yellow crystals insoluble in water. When any of these compounds are gently heated, brown nitrous fumes are evolved and red mercuric oxide remains behind. Mercuric Chromate is thrown down as a brick -red precipitate when potassium chromate is added to a solution of mercuric nitrate. Mercuric Chloride : HgCl 2 . — This important salt, which is commonly called corrosive sublimate, crystallizes from its aqueous solution in colourless needles of specific gravity 5-4. It melts to a clear liquid at 288 , boils at about 300 , and can be sublimed un- changed. The salt requires 14 parts of cold water for solution, but less than 2 parts of boiling water. It is more easily soluble in alcohol and in ether. Mercuric chloride may be obtained by dissolving the metal in aqua regia, when it crystallizes from the concentrated acid solution. It is, however, more commonly prepared in the dry way by sub- limation. For this purpose a mixture of common salt and mercuric sulphate is heated in a glass retort or flask, on a sand-bath, when the mercuric chloride thus produced condenses in the upper and cooler parts of the vessel as a crystalline cake, which is removed by breaking open the vessel when cold. The operation must be carried on under a good draught on account of the poisonous nature of the vapours of the chloride, some of which always escape into the air. The following reaction represents the process : — SO s -0,Hg.+ 2 NaCl = HgCl 2 + S0 2 (ONa) 2 . Mercuric Compounds. 537 Mercuric chloride readily parts with one-half of its chlorine when heated with substances which easily combine with this ele- ment, and is converted into mercurous chloride. Thus, an intimate mixture of arsenic and mercuric chloride gives arsenious chloride when heated (p. 247). Corrosive sublimate is an important medicine, but extremely poisonous, and can therefore only be given in small doses ( jg to £ grain, or 4 to 8 milligrammes). Its solution coagulates albumen, and possesses powerful antiseptic properties. Wood which has been impregnated with a solution of corrosive sublimate is pro- tected both from decay and from the attacks of insects. It has been proposed to saturate the wood of wooden bedsteads with its solution, tut the process is not to be recommended on account of the poisonous nature of the dust given off by the dry wood. Caustic soda when added to a solution of mercuric chloride gives a yellow precipitate of mercuric oxide, but ammonia when added to the same solution behaves differently, and produces a white precipitate of the composition : xtuVi f Hgo, which is to be considered as a diammonium dichloride, in which one-half of the hydrogen (4 atoms) has been displaced by mercury (2 atoms), and is, therefore, dimercuri-diammonium chloride. This substance is used in medicine under the name of infusible white ■precipitate. If a mixture of mercuric and ammonium chlorides is precipi- tated with sodium carbonate, a similar white compound is obtained of varying composition, and which melts, when heated, to a yellow- ish liquid. This substance is called fusible white precipitate. When a solution of mercuric chloride is gradually added to a boiling mixture of ammonium chloride and ammonia, so that the precipitate which is produced redissolves, the solution deposits on cooling colourless rhombohedra of the composition : nh 3 q| Hg i.e. mercuri-diammonium chloride. This compound is closely related to the infusible white precipitate, from which it may be obtained by boiling with ammonium chloride : — Mercuric Bromide : HgBr.,, resembles the chloride, but is more difficultly soluble in water. 538 Text-Book of Inorganic Chemistry. Mercuric Iodide : Hgl.,. — When solutions of mercuric chloride and potassium iodide are mixed, a yellow precipitate of mercuric iodide is first produced, which soon turns red, and is'soluble either in excess of mercuric chloride or of potassium iodide. On cooling its warm saturated solution in potassium iodide or in alcohol, the salt crystallizes in fine red tetragonal pyramids. It is insoluble in water. When heated the salt becomes yellow, then melts and sub- limes in bright yellow crystals belonging to the rhombic system. This yellow modification of mercuric iodide readily passes into the red form ; mere contact with the point of a needle is sufficient to convert an entire quantity of the yellow crystals into the other modification. A solution of mercuric iodide in potassium iodide, and made alkaline by caustic potash, is used for the detection of minute quantities of ammonia or ammonium salts, and is called Nessler's solution (p. 387). Mercuric Cyanide : HgCy,,. — Red mercuric oxide readily dis- solves when digested with dilute hydrocyanic acid, and mercuric cyanide crystallizes from the clear solution in brilliant colourless quadratic prisms. The crystals contain much enclosed water, but none combined as water of crystallization ; the substance is ex- tremely poisoncus. Mercuric cyanide is decomposed on heating into mercury and cyanogen, a portion of the latter gas taking the form of brown paracyanogen (p. 31 1). If on digesting mercuric oxide with hydrocyanic acid the oxide is in excess, an alkaline solution is obtained, from which crystals of a basic cyanide : HgCy, + HgO, may be separated. Mercuric Sulphide, Cinnabar : HgS. — This compound, distin- guished by its beautiful red colour, is the commonest form in which mercury occurs in nature, and is the substance exclusively used for the extraction of the metal. The mineral cinnabar is sometimes found in small, transparent, red crystals, belonging to the hexagonal system, but generally occurs in compact dark-red masses. It volatilizes unchanged when heated, and in this respect differs from the corresponding red oxide. Cinnabar may be obtained artificially by rubbing together a mixture of 25 parts of mercury with 4 parts of flowers of sulphur. This first produces an amorphous black sulphide, which when heated Mercurous Compounds. 539 in a suitable vessel yields a sublimate of cinnabar as a translucent, red crystalline mass of fibrous structure. Mercuric sulphide is produced in the wet way by leading sulphu- retted hydrogen into a solution of a mercuric salt. A black preci- pitate is then produced, which has the same composition as cinnabar, and into which it may be converted by sublimation. When the sulphuretted hydrogen is led into the mercuric salt a white precipitate is first thrown down, which soon becomes yellow, then red, and finally black. These intermediate substances are double compounds, in varying proportions, of mercuric sulphide and the mercuric compound used for the precipitation. Mercuric sulphide is insoluble both in nitric acid and in hydrochloric acid, and only dissolves in aqua regia. This pro- perty enables us to separate it from the sulphides of other metals ■ — e.g. copper — which dissolve in nitric acid. The sulphide in the form of powdered cinnabar is the valuable pigment vermilion, which is prepared either by grinding sublimed cinnabar or by various wet processes. The colour is improved by continued rub- bing or by allowing the powder to stand for some time in contact with a solution of potassium pentasulphide at a temperature of '50 . Mercurous Compounds. Mercurous Oxide : Hg 2 0. — This compound is much less stable than mercuric oxide ; it may be obtained as a black amorphous precipitate by mixing a solution of mercurous nitrate with a slight excess of caustic soda. When rubbed, or when heated to 100°, it decomposes into mercuric oxide and free mercury. The same change is also produced by the action of light, and it must, there- fore, be filtered, washed, and dried in the dark. Undecomposed mercurous oxide dissolves completely in dilute nitric acid. Mercurous Nitrate: (N0. 2 ).,0 2 Hg 2 + 2H„0. — If mercury is covered with a layer of moderately dilute nitric acid and allowed to stand for some time at the ordinary temperature, colourless transparent crystals gradually separate out in the form of rhombic plates. This salt is mercurous nitrate. The crystals easily dis- solve in the mother-liquor when warmed, and reseparate when the solution again cools. The salt is decomposed by water, with the separation of a yellow basic salt : (NOo) 2 2 Hg 2 + Hg 2 + H 2 0. 54° Text-Book of Inorganic Chemistry. A second basic salt, containing less mercurous oxide, is obtained by warming nitric acid with excess of mercury ; it crystallizes from its hot saturated solution in transparent colourless prisms. On the addition of ammonia to a solution of mercurous nitrate a black precipitate, resembling mercurous oxide, is obtained. This substance contains, however, the elements of ammonia and has a varying composition ; its chief constituent is probably a compound of the composition : iy /->'- ' onh r ^£t — dimercuroso-diammoniiim nitrate — i.e. diammonium nitrate in which four atoms of hydrogen are displaced by four monad atoms of mercury. It was formerly used in medicine under the name of Mercurius soiubilis Hahne- Mercurous Chromate : Cr0 2 • O a Hg„, is thrown down as a red precipitate on adding a solution of potassium chromate or dichro- mate to one of mercurous nitrate ; it is quite insoluble in water, and is used for the separation and quantitative estimation of chromic acid. The dry salt decomposes when heated into mercury and oxygen, which are given off, and green chromic oxide, which remains behind. From the weight of chromic oxide so obtained that of the chromic acid precipitated may be easily calculated. mercurous Chloride, Calomel : Hg„Cl 2 . — This salt, which like corrosive sublimate is a valuable medicine, is distinguished by its complete insolubility in water. It is found in nature as the mineral horn-quicksilver, and may be prepared artificially either in the wet way (by precipitation) or in the dry way (by sublimation). By the former method, it is obtained as a white curdy precipitate, when hydrochloric acid or a soluble chloride is added to a solution of mercurous nitrate. This precipitate when washed and dried forms a heavy white powder. Calomel is also obtained in the dry way when a mixture of mercuric chloride and mercury is heated. In grinding the two substances together it is well to add a little alcohol to prevent the formation of any dust of the poisonous mercuric chloride. Thin glass flasks are then about one-third filled with this mixture, and heated on a sand-bath ; vapours of mercury and mercuric chloride are thus formed, and unite together to produce mercurous chloride, which sublimes and condenses in the upper and cooler parts of the flasks. The operation must be conducted under a hood with a good draught on account of the Amalgams. 54 1 extremely poisonous nature of the vapours of mercuric chloride, some of which always escape from the flasks. This process is not, however, used for the manufacture of calo- mel, but a mixture of mercuric sulphate and common salt (which yields mercuric chloride when heated) is ground up with an equal quantity of mercury to that already contained in the sulphate, and then heated in glass or earthenware vessels. Mercurous chloride then sublimes and sodium sulphate remains behind. Calomel prepared in either of these ways always contains more or less corrosive sublimate, which may be extracted from it by digesting the powdered substance with warm alcohol. This puri- fication is especially necessary if the calomel is to be employed for medicinal purposes. Calomel is much less poisonous than corro- sive sublimate, and can therefore be given in larger doses. Mercurous chloride when treated with ammonia becomes con- verted into a black powder of the composition : N ' Hg 4 — di- mercuroso-diamrnoniuin chloride, and in this way may be distin- guished from silver chloride, which dissolves in ammonia. Mercurous Bromide : Hg^Br^, closely resembles the chloride in its properties. Mercurous Iodide : Hg.i.,, is thrown down as a dirty green pre- cipitate when a solution of potassium iodide is mixed with one of mercurous nitrate : it may also be obtained by rubbing together mercury and iodine in the correct proportions. The compound turns black in the light, and when heated is decomposed into mer- curic iodide and mercury. Mercurous Cyanide has not yet been prepared. Mercurous Sulphide : Hg,S, is possibly produced when sulphu- retted hydrogen is led into a solution of mercurous nitrate ; it is, however, very unstable, and soon decomposes into mercuric sul- phide and mercury. Amalgams. Mercury readily unites with most metals,. and produces alloys — generally known as amalgams. Often mere contact at the ordi- nary temperature suffices to produce an amalgam, as in the case of silver, gold, and the alkali metals. Most amalgams may be 54 2 Text-Book of Inorganic Chemistry. obtained in the solid crystalline form, and are all soluble in ex- cess of mercury. An amalgam of iron can only be obtained with difficulty, and iron vessels are therefore employed to contain the metal. Sodium amalgam may be easily obtained as a solid crystalline mass by adding 4 parts of sodium, in small quantities at a time, to 100 parts of warm mercury contained in a mortar. As each small piece of sodium is pressed under the mercury with the pestle, the two metals unite with a hissing noise, and a sufficient evolution of heat to volatilize some of the mercury. The amalgam remains liquid as long as it is warm, but becomes a crystalline solid when cold. Mercury unites still more energetically with potassium. A piece of gold when dipped into mercury becomes at once covered with a white layer of gold amalgam, and soon completely dissolves. Silver also easily dissolves in an excess of mercury. The readiness with which these two metals unite with mercury is utilized to separate them from their ores. These ores when shaken with mercury yield amalgams, from which the mercury can be after- wards distilled, leaving silver or gold behind (amalgamation pro- cess). One of the most useful amalgams is that of tin, which is used for silvering ordinary looking-glasses (p. 522). Metallic mercury is used for a number of technical and scientific purposes — for example, in the construction of barometers and ther- mometers, and for the mercury air-pump (Sprengel's pump). It is also used in the laboratory to collect gases which are dissolved by water but which do not act upon mercury — e.g. ammonia, hydro- chloric acid. Chlorine cannot be collected over mercury, as the two substances at once unite with one another and produce a chloride. Detection of Mercury Compounds. All solid compounds of mercury when mixed with quick-lime or dry sodium carbonate and heated in a small tube are decomposed, and yield metallic mercury, which condenses in the cooler parts of the tube in minute glistening globules. Soluble mercuric salts may be recognized by the yellow pre- cipitate of mercuric oxide produced by caustic soda and the white precipitate formed on the addition of ammonia. Black mercuric sulphide, which is produced when a stream of sulphuretted hydrogen is passed through a solution of a mercuric salt, is distinguished from other black sulphides prodtrced in the same way (e.g. those Silver. 543 of copper, lead, bismuth) by remaining unchanged when warmed with nitric acid. The soluble mercurous salts, of which only one or two are known, give a white precipitate of mercurous chloride with hydrochloric acid, which blackens on the addition of ammonia. Caustic soda and ammonia, as well as ammonium or sodium carbpnate, all produce black precipitates with a solution of mer- curous nitrate. If a solution of stannous chloride is added to a solution of a mercuric or mercurous salt, a white precipitate of mercurous chloride is produced in both cases, in the former with the simultaneous production of stannic chloride. And if this pre- cipitate of mercurous chloride is mixed with a further quantity of stannous chloride and some hydrochloric acid, it is reduced to a grey powder of metallic mercury, the particles of which unite when boiled and produce small globules of the metal. SILVER. Chemical Symbol : Ag. — Atomic Weight: 108. Silver is tolerably widely distributed in nature, but is seldom found in large quantities. It occurs as native silver crystallized in regular octahedra, and united with chlorine as horn-silver (AgCl), as well as with the other ha'ogens bromine and iodine, or even with all three together : it is also occasionally found mixed with mercury as amalgam. More important than these are the com- pounds of silver with sulphur, silver-glance (Ag 2 S), and those con- taining silver sulphide united with other sulphides (e.g. those of arsenic, antimony, and copper). These double sulphides form the minerals dark-red silver ore or pyrargyrite (3Ag.,S,Sb 2 S 3 ), light- red silver ore or proustite (3Ag 2 S,As 2 S 3 ), and silver copper-glance (Ag. ; S,Cu 5 S). A considerable quantity of silver is extracted from argentiferous lead by the process of cupellation, which has been already de- scribed under lead (p. 500). From ores containing little lead the silver is extracted by the process of amalgamation, which varies in different countries according to the nature of the ore and according as fuel is cheap or dear. The rich ores of Nevada and Colorado "in the United States are roasted with common salt, by which the silver loses its sulphur and is converted into silver chloride. The roasted ores are then mixed with water and scrap iron and rotated 544 Text-Book of Inorganic Chemistry. in barrels for some time, when the iron gradually reduces the silver to the metallic state. As soon as this change has been produced, mercury is introduced into the barrels, and the silver amalgam so obtained run off, dried, and distilled. Mercury then passes over into a condenser and metallic silver remains behind. In Mexico and Chili, where fuel is scarce and dear, the silver is converted into chloride by a mechanical process which consists in mixing the ground ore with common salt and salts of copper and iron by the treading of mules ; mercury in excess is next added, part of which reduces the silver chloride with formation of calomel, the remainder forming an amalgam which is afterwards distilled. Various wet processes are also used for the extraction of silver from very poor ores. Pure silver may be obtained by precipitating a solution of a silver salt with hydrochloric acid, and fusing the pure dry silver chloride so produced with dry sodium carbonate in an unglazed porcelain crucible. When cold, the silver is found at the bottom of the crucible covered with a layer of fused sodium chloride. Silver is a white, lustrous metal, with a specific gravity of I0'6. It is soft, and so malleable that it can be beaten out into extremely thin leaves (silver leaf). Silver melts at a lower temperature than copper, and although it does not oxidize in the air, possesses the property of dissolving oxygen when in the liquid state. This oxy- gen, which is only dissolved in the silver not chemically combined with it, is again given off when the silver solidifies, producing the phenomenon known as the spitting of silver. Pure silver is said to absorb as much as 22 times its volume of oxygen. The property is only possessed by the pure metal : the addition of a small quantity of lead prevents the absorption of the gas and the subse- quent spitting. Silver combines far more readily with sulphur than with oxygen. Sulphuretted hydrogen is decomposed by silver, and the metal be- comes covered with a thin brown or black layer of silver sulphide. Hydrochloric acid does not attack the metal, but it is readily dis- solved by nitric acid with evolution of brown nitrous fumes, and formation of silver nitrate. Concentrated sulphuric acid dissolves silver in the same way as copper and mercury, producing silver sulphate and sulphurous anhydride. Silver is precipitated from solutions of its salts by other more positive metals— £.£-. zinc, iron, and copper — and also by certain reducing substances, such as ammonium sulphite, grape sugar, aldehyde, and tartaric acid. Silver. 545 Some of these substances cause a deposit of the silver as a uniform coating on the walls of the vessel used for the experiment, and pro- duce a mirror of metallic silver. All commercial silver, including silver coins, always contains from 5 to 10 per cent of copper. To prepare pure silver from this alloy, the metal may be dissolved in nitric acid, the solution filtered from a trace of gold which remains behind as a black powder, and the filtrate precipitated with hydrochloric acid or a solution of com- mon salt. The silver chloride is then washed, dried, mixed with dry sodium carbonate, and heated in an unglazed crucible. Sodium chloride and silver carbonate are then produced, and the latter com- pound is decomposed at a red-heat into carbonic acid, oxygen, and silver. After cooling, a button of pure silver is found at the bottom of the crucible covered with a layer of fused sodium chloride. Compounds of Silver. Silver is one of the few metals which exists as a monad element in nearly all its compounds. In this respect it resembles the alkali metals and thallium, and like sodium it also forms a per- oxide, containing the metal as a dyad. Its sulphate unites with aluminium sulphate when heated in sealed tubes, and produces silver alum, isomorphous with the alums of potassium and thallium. Silver Oxide : Ag„0, is precipitated as a brown amorphous powder when silver nitrate is mixed with pure caustic soda or some other soluble base. As ordinary caustic soda always contains sodium chloride, the silver oxide obtained by it is mixed with silver chloride, and it is therefore preferable to use baryta-water. Silver oxide is not quite insoluble in water ; it imparts a metallic taste and a faint alkaline reaction to it. When exposed to the light it slowly decomposes at the ordinary temperature into silver and oxygen — a change which takes place rapidly when the compound is heated. Hydrogen reduces it completely even at 100°. Silver oxide unites readily with all acids, but few of its salts are soluble. If freshly precipitated, silver oxide is digested with strong am- monia ; it is converted into a black powder, which may also be obtained by adding caustic potash to an ammoniacal solution of silver nitrate. This substance is one of the most unstable com- pounds known : it explodes violently when slightly pressed, even 546 Text-Book of Inorganic Chemistry. when moist. It is known as Berthollefs fulminating silver, but we are quite ignorant of its composition. Possibly it is silver amide : NH 2 Ag, or silver nitride : NAg 3 . Silver Peroxide : AgO or Ag 2 2 . — This little-studied compound is obtained in the form of black shining octahedra when ozone acts on a moist silver plate, or when an electric current is passed through a strong solution of silver nitrate. It is then deposited on the positive pole, for which it is best to employ a stout platinum wire. The compound is less stable than silver oxide ; it decomposes when heated over ioo° into silver oxide and oxygen, and when inflam- mable substances, such as sulphur or phosphorus, are brought into contact with it, they are oxidized with an explosion. When warmed with nitric acid it produces silver nitrate, and oxygen is evolved. Cold nitric acid dissolves it and produces a brown liquid. Silver Sulphate : S0 2 (OAg) 2 . — When silver is warmed with con- centrated sulphuric acid, the metal gradually dissolves, with evolution of sulphurous acid, and silver sulphate is deposited as small glisten- ing crystals. The salt is difficultly soluble in water, and may there- fore be prepared by adding a solution of sodium sulphate to one of silver nitrate. As the salt is far more soluble in hot than in cold water, it is deposited when the above solutions are made hot, mixed, and the mixture allowed to cool. It crystallizes in small rhombic prisms, which are isomorphous with anhydrous sodium sulphate. At a low red-heat, silver sulphate melts unchanged, but at higher temperatures it is decomposed into silver, sulphurous and sulphuric anhydrides, and oxygen. Silver Nitrate : N0 2 -OAg. — This, the most important salt of silver, is obtained by dissolving pure silver in nitric acid and evapo- rating down the acid liquid. It crystallizes in large colourless rhombic plates, is easily soluble in water, and the solution reacts neutral. Alcohol also dissolves considerable quantities of the salt. Silver nitrate melts at about 220 , and solidifies on cooling to a crystalline mass. At a higher temperature it is converted into silver nitrite, which when more strongly heated is completely de- composed, metallic silver remaining behind. If silver nitrate is carefully melted at the lowest possible temperature and poured into moulds, similar to those .used for caustic potash (p. 348), it is obtained in the form of white solid Silver Nitrate. 547 sticks, which are used as a caustic in surgery under the name of lunar caustic. Silver nitrate destroys those bodies which cause the putrefaction and decomposition of organic substances, and also coagulates albumen ; hence the spots touched with the caustic become covered with a hard crust of coagulated albumen and fibrin, which soon blackens, especially when exposed to the light. The blackening is due to finely divided metallic silver reduced from the nitrate by the organic matter. A solution of silver nitrate also blackens when exposed to the light in contact with organic sub- stances, and such a solution mixed with gum water is used for marking linen and cotton articles. Characters written with this marking-ink gradually blacken when warmed or when exposed to light, and the reduced silver adheres so firmly to the fibres of the stuff, that it is not removed by washing with soap and soda. Such marks may, however, be obliterated by moistening them with a solu- tion of potassium cyanide and then very carefully washing with water, on account of the extremely poisonous character of this compound. Potassium cyanide dissolves finely divided silver in the presence of water, producing soluble potassium-silver cyanide with some caustic potash, and free hydrogen : — 2KCy + Ag + H 2 = KCy,AgCy + KOH + H. Silver nitrate has many applications in the arts. Large quanti- ties are used in photography for the production of the halogen compounds of silver which are readily decomposed by light. Silver Phosphate : PO(OAg) 3 , is produced as a bright yellow precipitate by mixing a solution of common (monacid) sodium phosphate with one of silver nitrate : — PO(ONa)„(OH) + 3 N0 2 -OAg = PO(OAg) 3 + 2NO a ONa + NOo-OH. The filtered solution then reacts acid from free nitric acid, and contains some silver dissolved as acid silver phosphate. Silver Arsenite : As(OAg) 3 , closely resembles the yellow phos- phate. Silver Arsenate : AsO(OAg) 3 , is a. reddish-brown powder, and like the phosphate and arsenite, easily dissolves in nitric acid. Silver Carbonate : CO(OAg),, is thrown down on mixing a solution of silver nitrate with one of sodium carbonate as a yellowish-white precipitate, which quickly darkens in the light. 548 Text-Book of Inorganic Chemistry. The dry salt loses its carbonic acid at a temperature slightly above ioo°. Silver Chromate : Cr0 2 (OAg) 2 , is a dark-red precipitate ob- tained by mixing solutions of potassium chromate and silver nitrate. Silver Chloride, Bromide, and Iodide. These three halogen compounds of silver closely resemble one another, especially in their chemical properties. They all occur in nature, and often associated with one another. They are obtained as curdy precipitates of a white or pale yellow colour, and insoluble in dilute acids when the corresponding halogen acid or haloid salt is added to a solution of silver nitrate. Silver chloride is of a pure white colour, the bromide is faintly yellow, and the iodide has a more decided yellow tint. They all melt when heated without de- composition, and are reduced to the metallic state by zinc when this metal is dipped into one of the fused salts or when zinc is added to the salt suspended in water containing hydrochloric acid. All three dissolve easily in an aqueous solution of sodium thiosul- phate or of potassium cyanide, but are distinguished from one another by their solubility in ammonia. Silver Chloride : AgCl, when freshly precipitated, easily dissolves in a small quantity of ammonia ; Silver Bromide : AgBr dissolves only with difficulty and requires a large quantity of the reagent ; while Silver Iodide : Agl, is almost insoluble in ammonia, but is dissolved by potassium iodide, forming a soluble double salt of the composition: KI,AgI, which crystallizes out when the solution is evaporated down. Silver is scarcely attacked by hydrochloric or hydrobromic acid, but strong hydriodic acid easily converts the metal into silver iodide with evolution of hydrogen. Silver bromide and iodide are less stable compounds than silver chloride: Both the former are decomposed when heated in a stream of chlorine ; silver chloride is then produced, with separation of bro- mine or iodine respectively. But if a mixed solution of potassium iodide and chloride is added to a small quantity of silver nitrate, the compound first produced is not silver chloride but silver iodide. In the same manner, silver chloride is completely decomposed into silver bromide when boiled with a solution of potassium bromide, and both silver chloride and bromide become converted into silver iodide when treated with potassium iodide. From these results it Halogen Compounds of Silver. 549 has sometimes been imagined that the affinity of silver for chlorine in an aqueous solution is less than for bromine ; but if we consider the three equations which represent these changes : — AgCl + KBr = AgBr + KC1 AgBr + KI Agl + KBr AgCl + KI = Agl + KC1 we see that potassium chloride is formed in the first and third equations and potassium bromide in the second, and potassium chloride is a much more stable compound than either the iodide or or the bromide, while the bromide is more stable than the iodide. The cause of the reactions is therefore due not to a stronger affinity of silver for bromine or iodine than for chlorine, but to a stronger affinity of potassium for chlorine than for bromine and iodine, and for bromine than for iodine. The sensibility of the halogen compounds of silver to light, under the influence of which, in the presence of reducing agents, they are reduced to sub-salts and then to metallic silver, makes them very suitable for the production of photographs. The com- pounds, usually silver bromide or iodide, are suspended in some suitable medium (collodion or gelatine), and glass plates coated with the mixture. These plates are then exposed in the camera to the light reflected from the object of which a picture is required, when the silver iodide or bromide is more or- less reduced according to the intensity of the light falling upon it, and according to the length of time which the plate is exposed. On removing the plate from the camera no change can be seen, and no metallic silver has yet been produced : the light has only formed a sub-bromide or iodide of the same colour as the original salt. But if the exposed plate is brought into a solution of some reducing substance, such as pyrogallic acid or ferrous oxalate, the reducing action which the light has begun is carried further and the sub-bromide or -iodide is now converted into black metallic silver. This process is called developing. In order to preserve this picture and to prevent the undecomposed silver salt from blackening when exposed to the light, it must now be fixed — a process which consists in dissolving out the undecomposed bromide or iodide by a solution of sodium thiosulphate or potassium cyanide. The picture so obtained shows the bright parts of the object dark and the dark parts light, and is called a negative ; an ordinary photograph on paper is obtained from this by placing the glass negative over a sheet of paper con- 550 Text-Book of Inorganic Chemistry. taining silver chloride in its pores and then exposing to the light. The print so obtained is then treated with various solutions con- taining gold, to make the colour more pleasing, and finally fixed with sodium thiosulphate. Not only does the intensity of the light influence its reducing action on the halogen compounds, but light of different colour produces very different effects. Thus the blue and violet end of the spectrum has a far more powerful re- ducing action than the green, yellow, and red ; and the invisible rays which lie beyond the visible violet (the ultra-violet or actinic rays) act most powerfully of all. Hence all photographic processes must be conducted in a room which is only illuminated with yellow or red light. Silver Fluoride : AgF, is distinguished from the other halogen compounds of silver by its solubility in water. It may be easily obtained by dissolving silver carbonate in hydrofluoric acid, and crystallizes from its aqueous solution with two molecules of water. Silver Cyanide : AgCy, is a white curdy precipitate produced by carefully adding a solution of potassium cyanide to silver nitrate. It is insoluble in water and dilute acids, but easily dissolves in ammonia. An excess of potassium cyanide also dissolves it, and produces a soluble double cyanide of the composition : KCy,AgCy, which may be easily obtained in the crystalline form. Unlike the chloride, bromide, and iodide, it is decomposed when heated, yielding cyanogen gas. The grey substance which then remains behind is a mixture of metallic silver and paracyanogen. The soluble compound : KCy,AgCy, is largely used in silver plating. The silver which is deposited at the negative pole when an electric current is passed through a solution of this double cyanide is in a more compact form, and adheres firmer to the substance to be plated than that deposited in a similar manner from any other solution of silver. Silver Sulphide : Ag 2 S. — Silver possesses a strong affinity for sulphur, and even decomposes sulphuretted hydrogen in the presence of air, but remains unchanged in the pure dry gas. The compound may be obtained by melting silver and sulphur together, or by pre- cipitating a salt of silver with sulphuretted hydrogen. When prepared by the latter method it is a black powder easily soluble in nitric acid. Silver sulphide occurs in the mineral kingdom as Alloys of Silver. 551 silver-glance, and in combination with the sulphides of arsenic or antimony as light and dark red silver ore respectively. Alloys of Silver. The numerous articles which are made of silver do not consist of the pure metal, but of an alloy of copper and silver. The pure metal is too soft for ordinary use and would wear away rapidly. But the addition of a few parts per cent, of copper considerably increases it in hardness and makes it more suitable for ordinary purposes. The two metals fuse together in all proportions, and those alloys which only contain 20 per cent, of silver have a reddish colour, while those containing 80 to go per cent, are white. English silver coins contain 92^5 per cent, of silver and 7-5 per cent, of copper, while the silver coinage of France, Germany, and Austria, consists of an alloy of 90 parts of silver with 10 parts of copper. The percentage of silver contained in a given alloy is usually expressed in parts per thousand or per mille. Thus the fineness of the silver used in England for coinage would be 925, and that used in France and Germany for the same purpose would be 900. The fineness of the silver alloys which are to be employed for different purposes is regulated by law in some countries. The quantity of silver contained in an alloy with copper was formerly exclusively determined by cupellation. In this process the sample of silver is mixed with lead and heated on a small flat crucible made of bone-ash (the cupel) in a current of air ; the lead and the copper are then oxidized, and their oxides are absorbed by the porous bone-ash of the cupel. Finally, there remains behind a button of pure silver, from the weight of which the fineness of the sample may be easily calculated. A quicker and more accurate process is the wet way introduced by Gay-Lussac. In this method the silver is dissolved in nitric acid and a normal solution of common salt added until the whole of the silver is precipitated as silver chloride. For the details of this method the student must consult a text-book of analytical chemistry. Detection of Silver Compounds. The reduction of any compound of silver when mixed with sodium carbonate and heated on charcoal before the blowpipe, ar.d the production of a white globule of metallic silver which does not oxidize in the air, the precipitation of black silver sulphide when 552 Text- Book of Inorganic Chemistry. sulphuretted hydrogen is led through an acid solution of a silver salt, and, above all, the formation of silver chloride by hydro- chloric acid, make it very easy to detect the presence of silver and to separate it from other metals. Silver chloride may be distin- guished from lead chloride by the solubility of the latter salt in hot water, and from mercurous chloride by the action of ammonia. This dissolves the silver chloride, but converts the mercurous salt into a black compound (p. 543). GOLD. Chemical Symbol : Au. — Atomic Weight : 197. Gold is chiefly found in nature in the free state, sometimes pure and well crystallized, but usually containing silver. It occurs in quartz rocks or in alluvial quartz sand which has been deposited by the disintegration of the older rocks. These auriferous sands, which often cover a considerable extent of country, are washed by suitable arrangements, on a small or large scale, when the lighter particles of sand and mud are carried away and the heavier particles of gold remain behind. The gold is then brought into the market in this form, or if in a finely divided state is extracted with mercury, which is afterwards removed by distillation. Gold is now chiefly obtained from California and Australia, but is found in smaller quantities in other parts — e.g. in the Ural Mountains, in Hungary and Transylvania, and in Africa. Small quantities of gold have also been found in Cornwall and Wales. Chemical compounds of gold rarely occur in nature — sylvanite or graphic tellurium, found in Transylvania, is a compound of gold and tellurium. Many silver ores and some argentiferous lead ores, as well as some specimens of copper-pyrites, contain small quantities of gold. In the silver ores the gold is extracted with the silver, and the two metals are afterwards separated by nitric acid. Gold belongs to the few metals which have a decided tint, and is distinguished from all others by its yellow colour. It crystallizes in the regular system, is very soft, considerably more so than silver, and very malleable. In extremely thin leaves (gold-leaf), it trans- mits bluish-green light. Gold melts more difficultly than silver, but somewhat more easily than copper, and volatilizes at very high temperatures. Its specific gravity is 19-3. Gold. 553 Gold does not directly unite with oxygen at any temperature, but easily combines with chlorine. Hydrochloric, sulphuric, or nitric acid has no action on the metal, but aqua regia dissolves gold and produces a yellow solution of auric chloride. From this solu- tion the metal may be again precipitated by various reducing agents, such as ferrous sulphate, oxalic acid, sulphurous acid, or metallic iron. The best substance to employ for the precipitation of gold from a solution of its chloride is a cold saturated solution of ferrous sulphate, which is then oxidized to a mixture of ferric sulphate and ferric chloride. The gold is thrown down as a brown powder, which acquires the colour and lustre of the metal when rubbed. This method is used to separate pure gold from a solution con- taining copper and other metals. Compounds of Cold. Gold forms two series of compounds in which the element exists, as a triad and as a monad respectively, and which are distinguished as the auric and aurous compounds. The metal only possesses weak affinities for other elements, and its two oxides are such weak bases that their salts are hardly known'. The most important auric compound is Auric Chloride : AuCl 3 , which is obtained by dissolving pure gold in aqua regia and remains as a reddish brown deliquescent mass when the solution is evaporated to dryness. It not only dis- solves in water, but also in alcohol and ether, and to such an extent in the latter liquid that the chloride is abstracted from its aqueous solution when this is shaken with ether. A solution of auric chlo- ride produces a purple stain of reduced gold when a drop is allowed to fall on the skin, and gives a purple precipitate when mixed with stannous chloride — see Purple of Cassius (p. 556). When heated, the salt decomposes into chlorine and aurous chloride. Auric chloride forms soluble crystalline double salts with hy- drochloric acid and soluble chlorides. The hydrochloric acid compound: AuCl s ,HCl + 3H,0, or HAuCl 4 + 3H 2 0, is deposited from a hot acid solution of auric chloride in long yellow needles. On mixing a solution of auric chloride with one of the chlorides of potassium, sodium, or ammonium, and evaporating down, crystals of the compounds : 2(AuCl 3 ,KCl) + H 2 0, or 2KAuCl 4 + H 2 ; AuCL, NaCl + 2H.,0, or NaAuCl 4 + 2H 2 0, and 2AuCl 3 ,NH 4 Cl + SH 2 0, or 2NH 4 AuCl 4 + 5H..0 are obtained. The sodium compound is much 554 Text-Book of Inorganic Chemistry. less deliquescent than pure auric chloride, and is the form in which this salt is usually brought into trade. It is sold as non-deliquescent chloride of gold m 15 grain or 1 gramme tubes. Auric Oxide : Au 2 3 , is obtained by digesting a solution of auric chloride with an excess of magnesia. The brown powder so produced consists of a compound of auric oxide and magnesia (magnesium aurate) with the excess of magnesia. When this powder is treated with strong nitric acid, the magnesia is all dis- solved and auric oxide remains behind as a brown powder. It is decomposed by light or heat into gold and oxygen. If the brown compound of auric oxide and magnesia is digested with dilute in- stead of concentrated nitric acid, auric hydrate : Au(OH) 3 , remains behind as a reddish yellow powder. Auric oxide is so weak a base that neither it nor the hydrate can unite with oxyacids to form salts. On the other hand, it can combine with bases and form salts in which it plays the part of an acid. Hence auric oxide may be considered as the anhydride of an acid — auric acid. Potassium Aurate : AuO ■ OK + 3H 2 0. — Auric oxide or hydrate easily dissolves in caustic potash, and when the solution is eva- porated, finally in a vacuum, potassium aurate separates in bright yellow needles. Its solution reacts alkaline and gives precipitates of insoluble aurates when mixed with the salts of other metals. If a solution of auric chloride is digested with an excess of ammonia a yellowish brown powder is produced, which violently explodes when slightly warmed or rubbed. The composition of this fulminating gold 'is unknown. Auric Cyanide : AuCy 3 + 3H 2 0. — Nearly all the cyanides of the heavy metals are insoluble in water, but mercuric cyanide and auric cyanide are exceptions to this rule. If a solution of auric chloride is mixed with a hot concentrated solution of potassium cyanide, the compound ■potassium-auric cyanide : KCy,AuCy 3 + H 2 0, or KAuCy 4 + H„0, crystallizes out on cooling in colourless plates. This substance, when dissolved in water, gives a white curdy precipitate with a solution of silver nitrate of the corre- sponding silver salt. If this silver-auric cyanide is carefully washed, mixed with water, and then a quantity of hydrochloric acid added insufficient to produce complete decomposition, silver chloride, hydrocyanic acid, and auric cyanide are produced, and Compounds of Gold. 555 when the clear liquid is evaporated in a vacuum over sulphuric acid this salt remains behind as a white crystalline mass. By recrystallizing from its alcoholic solution, the auric cyanide may be obtained as large crystalline tablets or plates. The salt is easily decomposed when heated. Auric Sulphide : Au 2 S 3 (?). — It is still doubtful whether the black precipitate which is produced when sulphuretted hydrogen is led through a cold solution of auric chloride consists of auric sulphide, or is a compound of aurous and auric sulphides. The precipitate easily dissolves in solutions of alkaline sulphides. If gold is fused with potassium polysulphide (liver of sulphur), a double sulphide of potassium and gold is formed, which dissolves when the cold mass is digested with water. From this solution acids precipitate a sulphide of gold. The following more important aurous compounds may be de- scribed. Aurous Chloride : AuCl. — Auric chloride when heated decom- poses into chlorine and aurous chloride or gold, according to the temperature employed. If the temperature does not exceed 200°, aurous chloride remains behind as a dirty-white powder which is insoluble in water. It is decomposed into auric chloride and metallic gold when boiled with water. Aurous Oxide : Au 2 0. — Aurous chloride when digested with caustic potash is converted into a dark-violet coloured powder of aurous oxide. The compound is insoluble in water, and is not changed by sulphuric or nitric acid. Hydrochloric acid decom- poses it into auric chloride and gold. Salts of aurous oxide are not known if we except the crystalline salt which separates from a mixed solution of auric chloride and sodium thiosulphate on the addition of alcohol, and which may be a double salt of aurous and sodium thiosulphates. Aurous Cyanide : AuCy. — Fulminating gold readily dissolves in a warm solution of potassium cyanide, and the solution deposits large prisms of potassinm-aurous cyanide : KCy,AuCy or KAuCy,,, on cooling. Hydrochloric acid separates aurous cyanide from this salt, and the cyanide may be obtained by evaporating to dryness and extracting with water, when it remains as a yellow crystalline powder insoluble in water. Potassium cyanide easily re-dissolves it ; it is decomposed when heated into gold and cyanogen. 556 Text-Book of Inorganic Chemistry. Purple of Cassius is a compound of gold, tin, and oxygen, of varying and, therefore, uncertain composition. It is thrown down as a beautiful purple precipitate by adding a mixture of stannous and stannic chlorides to a dilute solution of auric chloride. If concentrated solutions of these salts are mixed, the precipitate possesses a brown colour. The compound may also be prepared by acting on an alloy of gold, silver, and tin with nitric acid. The silver then dissolves and purple of Cassius remains behind. Purple of Cassius is used in painting on porcelain and glass to produce a beautiful red tint. Alloys of Gold. Pure gold is used for the manufacture of gold-leaf, and in the finely divided state, as obtained by precipitation with ferrous sul- phate, for gilding porcelain. The gold is mixed with a suitable material and painted on the porcelain, and when the articles are afterwards heated in a kiln the metal burns into the glaze. Articles made of the base metals may be covered with a thin layer of gold (or gilt) in various ways. According to one method a gold amalgam is rubbed on the surface to be gilt, the mercury afterwards expelled by heating, and the rough surface polished by burnishing. In the wet way, a thin coating of gold is obtained by simple immersion in an alkaline solution of auric chloride, or better by a separate electric current. In the process of electro-gilding the object to be gilt is dipped in a warm solution of one of the double cyanides of gold and potassium and made the negative pole of an electric current. The coating of gold obtained in this way is far more durable and compact than that produced by any other process. The gold which is used by the goldsmith and for coinage is an alloy of the metal with copper and sometimes silver. This alloy not only melts more easily, but is also harder and does not so readily wear away when used. In England the fineness of a sample of gold is expressed in carats ; pure gold being said to be 24 carats. Thus 18 carat gold consists of 18 parts of gold alloyed with 6 parts of copper. English sovereigns and half-sovereigns are made of 22 carat gold, or contain 916-66 parts of pure gold per thousand (or mille). The gold used for coinage in Germany and some other countries has a fineness of 900, or contains 900 parts of pure gold per thousand, which corresponds to 2r6 carat gold; 15 or 12 carat gold is usually employed for the manufacture of jewellery. 557 PLATINUM METALS. This group of the noble metals includes, besides platinum, the five metals : iridium, palladium, rhodium, ruthenium, and osmium. They are almost always associated with one another in nature and are found native accompanying platinum. The so-called platinum ore contains them as alloys together with varying quantities of gold, copper, and iron. This ore is found in the Ural, in South America, California, Borneo, Australia, and a few other localities, and, like gold, usually in alluvial sands. These six metals closely resemble one another in many points, but in some respects differ strikingly. Such a difference is noticed, for example, when we compare together their specific gravities and their atomic weights. It is then seen that the three : platinum, iridium, and osmium, possess a specific gravity and atomic weight which are about double, those of the other three : palladium, rhodium, and ruthenium. These differences are exhibited in the following table : — Platinum Specific gravity . 21'5 Atomic weight 195 Palladium Specific gravity . 1 1-4 Atomic weight 106 Iridium . . 22 '4 193 Rhodium . 121 I04 Osmium . 22'5 199 Ruthenium . • 12-3 I04 A second difference between these metals is their behaviour with regard to oxygen. Most of them have so weak an affinity for oxygen that they not only do not unite directly with this element, but the majority of their oxides, when prepared in an indirect manner, lose their oxygen when heated. On the other hand, osmium and ruthenium combine energetically with oxygen when simply heated in the air and produce volatile compounds of an acid character. Of the six rare platinum metals, platinum itself is found in far larger quantities than the other five. On this account, and because the metal is used for many purposes in the arts, the compounds of platinum are better known than those of the other metals. But the majority of these compounds are only of chemical interest. 558 Text-Book of Inorganic Chemistry. PLATINUM. Chemical Symbol : Pt. — Atomic Weight : 195. In the extraction of platinum the ore is. first warmed with dilute aqua regia, to dissolve any gold, and then treated with stronger aqua regia, often under pressure, when platinum with traces of iridium and rhodium go into solution as chlorides. The residue consists chiefly of the earthy impurities and of an alloy of osmium and iridium {osmiridiwn) which is not attacked by aqua regia. The solution is evaporated to dryness, then digested with water, and the clear liquid mixed with a solution of ammonium chloride. A yellow crystalline double chloride of platinum and ammonium js thus produced which is difficultly soluble in water and insoluble in alcohol. The salt is filtered off, washed, dried, and then heated to redness, when metallic platinum remains behind as a grey spongy mass, called spongy-platinum. This is the material used for the preparation of vessels made of platinum ; it still contains traces of iridium (and rhodium). Very various and often complicated processes»are used to completely separate platinum from all the other platinum metals and to prepare it perfectly pure, for a de- scription of which the student must consult a larger book. Small quantities of pure platinum may be prepared by the following method, due to Bunsen. The solution of the chlorides is mixed with a large excess of caustic soda until the precipitate which is first produced is redissolved and then boiled with the addition of a few drops of alcohol. During this boiling, the platinum tetrachloride remains unchanged, but the tetrachlorides of the other metals are reduced to a dichloride or a sesqui-chloride, and these do not yield an insoluble compound with ammonium chloride. When, therefore, the alkaline solution is afterwards acidulated with hydrochloric acid and mixed with an excess of ammonium chloride, the platinum alone is precipitated, and the other metals remain in solution. Platinum is a greyish -white metal of specific gravity 21-5 ; it is malleable and ductile, and very tough, and indeed, scarcely less so than iron. Small quantities of iridium render it less malleable, but at the same time make it harder and tougher and more suitable for the manufacture of crucibles, dishes, &c. It cannot be fused at the temperature of our furnaces, but readily melts in the oxy- hydrogen flame, when heated in a lime crucible (p. 24). Thin Platinum. 559 platinum wire melts easily in the oxy-hydrogen blowpipe, and is even partly volatilized. Platinum shares with iron the property of welding at a high temperature. Two pieces of platinum heated to whiteness, but still far below the melting point of the metal, may be welded together under the hammer. In the same way, the spongy platinum which remains behind when ammonium-platinum chloride is heated, may be welded into a compact mass under strong pressure at a high temperature. Such platinum, however, still contains traces of air, and crucibles made of it often become covered with small excrescences due to the expansion of this confined air. Platinum may be obtained in a still more finely divided state than it exists in spongy platinum. We are acquainted with the metal in the form of a loose very fine black powder, called pla- tinum black. This form of platinum may be obtained by the re- duction of an alkaline solution of platinic, or better, platinous chloride. For this purpose platinous chloride is dissolved in an excess of caustic potash (or soda), the mixture heated to boiling, and small quantities of alcohol gradually added. The metal is then reduced and the alcohol oxidized to aldehyde and acetic acid. This black powder, when carefully washed and dried at a low temperature, absorbs and condenses considerable quantities of various gases, especially oxygen, and is therefore a powerful oxi- dizing agent. Spongy platinum also possesses this property of condensinggases, but in a less degree. If a jet of hydrogen is allowed to play on a small piece of spongy platinum in the air, the platinum first becomes red hot and then ignites the hydrogen. Dobereiner's hydrogen lamp depends upon this property of spongy platinum. The oxygen which is condensed in the spongy platinum has stronger affinities than the ordinary gas, and as platinum also condenses hydrogen, the two substances are brought close enough together to cause their union. Compact metallic platinum offers a far less surface to the gases than the spongy metal ; it acts therefore less energetically, and not at all at the ordinary temperature. But if a spiral of platinum wire or a piece of thin platinum foil is heated to low redness and suspended in a vessel containing a little ether, the vapour of the ether, which is contained in the vessel mixed with air, gradually becomes oxidized to aldehyde and acetic acid, with so much evolution of heat that the platinum spiral remains glowing, and 560 Text-Book of Inorganic Chemistry. that the ether sometimes catches fire. It has already been men- tioned (p. 162) that compact platinum, when heated, can convert a mixture of sulphurous anhydride and oxygen into sulphuric .anhy- dride without itself undergoing any change. Platinum black, con- taining the metal in its most finely divided state, acts more power- fully than either the spongy or compact metal. It can absorb as much as 200 times its volume of oxygen, or even more, and we may therefore suppose that the gas is condensed by it to the liquid or even solid state. If platinum black is moistened with alcohol and exposed to the air, the latter substance becomes oxidized to aldehyde and acetic acid, and so much heat is set free that the metal is raised to redness and continues to glow for some time. Platinum is not attacked by either nitric, hydrochloric, or sul- phuric acid, but dissolves in aqua regia with formation of platinic chloride. Molten caustic alkalies, especially lithia, attack platinum at a red heat, and a mixture of fused caustic potash and nitre acts more powerfully. Sodium or potassium carbonate is entirely with- out action on the metal even at a white heat. Sulphur, phosphorus, and arsenic also slowly attack compact platinum, and the spongy metal readily melts when heated with these substances. A mixture of silicates and charcoal also attacks platinum, with formation of a compound of platinum and silicon, and platinum crucibles are therefore soon destroyed if brought into direct contact with glowing charcoal. Many of the easily fusible heavy metals,' such as lead, form easily fusible alloys with platinum. If, for example, a piece of lead is melted in a platinum crucible, an easily fusible alloy of platinum and lead is formed, and a hole is produced in the bottom of the crucible. Care must, therefore, be taken that these substances (sul- phur, phosphorus, arsenic, and the easily fusible heavy metals), or mixtures which might produce them, are never strongly heated in platinum vessels. The properties which are possessed by platinum — its infusibility at any temperature of our furnaces, that it does not unite with oxygen at any temperature, and that it resists the action of all acids including hydrofluoric acid, but excepting aqua regia — make the metal simply indispensable to the chemist. We employ it in the form of wire, foil, crucibles and dishes, tongs, tubes, retorts, &c. Considerable quantities of platinum are now used in the arts for the manufacture of large retorts and coolers in the purification of sulphuric acid. Platinum coins were also formerly used in Russia, Compounds of Platinum. 561 but have now been withdrawn from circulation. Notwithstanding its scarcity, platinum is about one-third cheaper than gold, but is more than seven times as dear as silver. Compounds of Platinum. Platinum unites with other elements and produces two series of compounds — the platinous and platinic compounds — in which the metal is a dyad and a tetrad respectively. The compounds of platinum thus resemble those of carbon and tin. The most interesting of these substances are the chlorides and the peculiar platinum bases. Platinic Chloride : PlCl 4 . — On dissolving platinum in aqua regia and evaporating the solution to dryness with an excess of hydrochloric acid, a brownish-red crystalline and deliquescent mass remains behind of the composition: 2HCl,PtCl 4 or H 2 PtCl 6 + 6H.,0. If, however, excess of the acid is avoided, platinic chloride : PtCl 4 + 5H 2 0, is obtained in fine red crystals. Both dissolve in water and produce a reddish-yellow liquid with an acid reaction. 1 Alcohol and ether also dissolve platinic chloride in the cold, but are oxidized with reduction of the chloride when the solution is warmed. When carefully heated platinic chloride loses one-half of its chlorine and becomes converted into platinous chloride ; if more strongly heated, it is completely converted into platinum and chlorine. Platinic chloride readily combines with the chlorides of the alkali-metals, of thallium and of the metals of the alkaline earths, producing double chlorides in which the chlorine combined with the platinum is twice as much as that united with the other metals. These salts are mostly difficultly soluble in water and insoluble in alcohol, and separate or crystallize out when solutions of the two chlorides are mixed. They usually crystallize in regular octahedra. These double chlorides have all a similar composition to the hydrochloric acid compound referred to above : 2HCl,PtCl 4 or H„PtCl„, and may be considered as salts of this chlorplatinic acid. Potassium chlorplatinate : 2KCl,PtCI 4 or K,PtCI 6 , is a pale yellow crystalline precipitate produced when solutions of potassium and 1 The brownish colour which a. solution of platinic chloride sometimes possesses is due to the presence of traces of iridium and other impurities.— Ed. O O 562 Text-Book of Inorganic Chemistry. platinic chlorides are mixed, and which is only completely precipi- tated on the addition of alcohol. When heated it breaks up into platinum chlorine and potassium chloride, the last of which may be extracted by boiling the residue with water. Ammonium chlorfilati- nate: 2NH 4 Cl,PtCl 4 or (NH 4 ) 2 PtCl d , closely resembles the potassium salt. When gently heated it leaves a residue of pure platinum as a grey spongy mass. Sodium chlorplatinate : 2NaCl,PtCl 4 or Na„PtCl 6 + 6H 2 0, is the only one of the double salts which crys- tallizes with water, and which dissolves both in water and alcohol. These properties enable us to completely separate sodium com- pounds from those of potassium. The sodium and potassium salts are first converted into chlorides, an excess of platinic chloride then added to this solution and evaporated to dryness. The residue is then extracted with strong alcohol when the sodium chlorplatinate and the excess of platinic chloride dissolve, leaving the insoluble potassium salt behind. If an excess of platinic chloride had not been added, some sodium chloride, which is insoluble in alcohol, would remain with the potassium chlorplatinate. Sodium chlorplatinate crystallizes from its concentrated aqueous solution in reddish-yellow prisms. Platinous chloride : PtCL,. — If platinic chloride is carefully heated to 230 in a porcelain basin with continued stirring until no further odour of chlorine is perceived, it is completely converted into platinous chloride and chlorine. This latter compound is an olive-green powder, insoluble in water, but dissolving in hydro- chloric acid to form a dark-brown solution. It is converted by aqua regia into platinic chloride. Platinous, like platinic chloride, unites with other chlorides (of potassium, ammonium, &c.) and produces crystalline double salts, which are soluble in water. Its behaviour with ammonia shall be described under the account of the platinum bases. Platinic Oxide : PtO,, and Platinic Hydrate : Pt(OH) 4 . — If a solution of platinic chloride is mixed with one of sodium car- bonate, the precipitate which is produced is not like platinic carbonate but a compound of soda and platinic oxide. If the whole is then evaporated to dryness and extracted with dilute acetic acid, platinic hydrate remains behind as a compact reddish- brown powder. The compound does not dissolve in this acid, but produces soluble platinic salts with the stronger oxy-acids, which Compounds cf Platinum. 563 have as yet been but little investigated. The hydrate also unites with strong bases and produces compounds (platinates), in which' it plays the part of an acid. "When gently heated the hydrate is converted into a black powder of platinic oxide, which is insoluble in acids. Platlnous Oxide : PtO, and Platinous Hydrate : Pt(OH).,. — Platinous chloride when digested with caustic potash is converted into a dark-brown powder of platinous hydrate, which is insoluble in water, but dissolves both in strong caustic potash and in acids, producing brown solutions in both cases. When the hydrate is gently heated, platinous oxide remains as a black powder insoluble in acids. Platinic Sulphide : PtS , and Platinous Sulphide : PtS, are black precipitates produced when sulphuretted hydrogen is led through a solution of platinic or platinous chloride respectively. They both behave as sulpho-acids and dissolve, for example, in potassium sulphide to form red solutions. Platinous Cyanide: PtCy r — If platinous chloride is digested with sufficient potassium cyanide to dissolve it and the solution evaporated down, crystals of potassium-platinous cyanide : 2KCy, PtCy, or K„PtCy 4 , separate out. The crystals show a beautiful dichro'ism — i.e. they appear of different colours when viewed in different directions. From a solution of this salt, dilute acids separate platinous cyanide as a greenish-yellow powder insoluble in water and acids. It unites with other cyanides and forms other double salts, similar in composition to the potassium compound, and which are partly soluble and partly insoluble in water. Ammoniacal Compounds of Platinum. Under this name is included a number of very interesting com- pounds, produced by the action of ammonia on platinous chloride and by further changes in the primary compound. In these com- pounds a portion of the hydrogen of ammonium is displaced by dyad platinum : Pt", or by the dyad radicals : (PtCl)", (PtO)", or even by ammonium itself. Substitution of the hydrogen in ammonia or in an ammonium compound by a dyad radical may take place in two different ways. 002 564 Text-Book of Inorganic Chemistry. Either two atoms of hydrogen may be displaced in one molecule of ammonia by the dyad radical, or one atom of hydrogen may be symmetrically displaced from each of two molecules of ammonia by the radical. The compounds of the first kind are unstable, and readily pass into ordinary compounds of ammonia ; such, for example, is cyanic acid (see p. 316). The second class of compounds — the diammonium compounds — are more stable and do not easily pass into compounds of ordinary ammonia. The ammoniacal compounds of copper, some of those of mercury, and possibly those of cobalt, belong to this class of diam- monium compounds, but far more important representatives of the same group are those which contain platinum. The simplest of these is: platinoso-diammonium chloride: Pt"J atttV, — i.e. a compound consisting of two molecules of ammonium chloride in which two atoms of hydrogen (one from each molecule) are displaced by one atom of dyad platinum. And the following compounds may serve as types of this very numerous group : — (i.) Platinoso diammonium) Pt// (NH.CI * chloridej " - rl JNH^Cl (ii.) Diammon - platinoso-'l p t „(NH 2 (NH 4 )Cl diammonium chloride J ' ' |NH 2 (NH 4 )C1 (iii.) Diammon ■ platinoso-1 p ,(NH 2 (NH 4 )OH diammonium hydrate ) ■ ' n {NH 2 (NH 4 )OH (iv.) Chlorplatin-diammoniuml prl „fNH 3 Cl chloride f * rtu » JNH 3 C1 (v.) Diammon chlorplatin-i p r , „ (NH 5 (NH 4 )C1 diammonium chloride I ' rl ^ }NH 2 (NH 4 )C1 (vi.) Oxyplatin - diammonium ) p _„ f NH 3 -ON0 2 nitrate}' ■ rtu JNH 3 -ON0 2 (vii.) Diammon oxyplatin- 1 p n „ (NH.,(NH 4 ) -ONO, diammonium nitrate} ' * rtu }NH 2 (NH 4 )-ON0 2 These salts are partly crystalline and soluble in water. The following is a brief account of the methods used for obtaining them. If ammonia is gradually added to a solution of platinous chlo- ride in dilute hydrochloric acid, a green crystalline salt separates out, which is a double compound of diammon-platinoso-diammonium chloride with platinous chloride : — Ammoniacal Compounds of Platinum. 565 p „(NH 2 (NH 4 )Cl p j rt {NH 2 (NH 4 )CI + rtL1 - > - This is commonly known as the green salt of Magnus, because it was discovered by this chemist ; it may also be prepared from its •separate constituents. If this double salt, or if platinous chlo- ride (dissolved in hydrochloric acid) is heated with an excess of ammonia, the compound diammon-platinoso-diammonium chloride: (ii.) : P£ "|nh 2 (NH 4 )C1 crv stallizes out in fine needles. The fol- lowing equation illustrates its production : — pta 2 + 4NH3 = pt"{gg$!g$j If the two atoms of chlorine which this salt contains are displaced by hydroxyl, the corresponding base (iii.) : Pt"j jJjHnh'IoH is produced ; this crystallizes from its solutions and caustic potash expels no ammonia from it. It is a powerful base, and forms numerous salts with acids. When this base is heated, it parts with 2 molecules of ammonia and is converted into platinoso-diammonium >jxj 3 r)ij a substance which is no longer soluble in water, and which gives insoluble explosive compounds with acids. Similarly, if the chloride (ii.) is heated to a temperature not exceed- ing 270°, it also loses 2 molecules of ammonia, and becomes platinoso-diammonium chloride (i.) : ^"InhVi a sa ^' wrnc h s only slightly soluble in water, and which readily again takes up 2 molecules of ammonia when boiled with this substance. If chlorine is led into water which contains platinoso-diammo- nium chloride in suspension, each molecule of the salt combines with 2 atoms of chlorine and the dyad platinum is changed into the dyad radical (PtCl 2 )", containing the metal in the tetrad form. The salt then produced is chlorplatin-diammonium chloride (iv.), according to the equation : — JNH3CI . r, _ / Pti . C irl NH 3Cl Ft (NH 3 C1 + Us " ( -' JNH 3 C1 This salt, which thus contains its 4 atoms of chlorine differently united, is only slightly soluble in water, and when boiled with ammonia takes up 2 molecules and becomes diammon-chlorplatin- 566 Text-Book of Inorganic Chemistry. diammonium chloride (v.) : (Ptdy'lNH^NrnCl a salt which may be also produced by leading chlorine into a solution of diam- mon-platinoso-diammonium chloride (ii.). Oxyplatin-diammonium nitrate (vi.) : (Pt iv O)'M N u 3 _ q n q 2 ma y be obtained from chlorplatin-diammonium chloride (iv.) by long boiling with an excess of silver nitrate, and when heated with ammonia yields diammon-oxyplatin-diammonium nitrate (vii.) : (PtO)'M mtj fNfWoNO 2 wn ' c h crystallizes well from hot water. Similarly constituted diammonium compounds containing a dyad radical are known in organic chemistry. PALLADIUM. Chemical Symbol ; Pd. — Atomic Weight: 106. Platinum ores only contain a very small quantity of palladium — at most one or two per cent., and the metal is usually prepared from a gold ore which contains it, and which occurs in Brazil. This ore is fused with silver, then granulated and digested with nitric acid, which extracts the silver and palladium and leaves the gold behind. The solution is mixed with sodium chloride to preci- pitate the silver and the palladium then thrown down from the filtrate as palladous cyanide by the addition of mercuric cyanide. The cyanide yields the metal when heated to redness, which may then be fused in the oxy-hydrogen flame. Palladium melts more easily than platinum, and is, in fact, the most easily fusible of all the plati- num metals, but can only be melted with considerable difficulty in a wind furnace. Palladium resembles platinum in its external appearance, but is only about half as heavy : its specific gravity is 1 1 -4. It is further distinguished from platinum by its solubility in nitric acid, and it is even attacked by warm concentrated sulphuric acid. It is malleable and ductile, and may be welded at a white heat. When heated in the air the metal tarnishes, but regains its original lustre when the temperature is raised to a bright red heat. Palladium Hydride : Pd 2 H. — The most remarkable property of palladium, a property which is not possessed by any other metal Palladium. 567 to a similar extent, is its power of absorbing and retaining large quantities of hydrogen. The metallic appearance of the compound of palladium and hydrogen reminds one of an alloy, and appears to prove that hydrogen is a gaseous metal. Its composition is Pd 2 H, which corresponds to 600 volumes of hydrogen absorbed by one volume of the metal. The compound is formed at the ordi- nary temperature when palladium is brought into an atmosphere of hydrogen, but the most favourable temperature for its formation is 100 . Above 130 and under the ordinary atmospheric pressure it begins to decompose with evolution of hydrogen, but the gas can only be completely expelled by protracted glowing. Besides the hydrogen which is chemically united with the metal, it can also absorb the gas mechanically. If water is electrolysed and palla- dium, especially if in the spongy form, used for the negative pole, the metal can absorb nearly 1,000 times its volume of the gas. But it readily loses the excess of hydrogen which it then contains beyond that combined chemically, and at ioo° the whole of the hydrogen mechanically absorbed is expelled. If a piece of palla- dium which has thus been saturated with hydrogen is exposed to the air, the gas is rapidly oxidized to water, with a considerable evolution of heat ; and if placed in a solution of potassium ferricy- anide the salt is rapidly reduced to potassium ferrocyanide. Such palladium hydride when saturated with hydrogen is, in fact, one of the most powerful of reducing agents. It may be preserved for a long time unchanged if quickly placed in cold boiled water or alcohol. This property which palladium possesses of combining both chemically and mechanically with hydrogen, explains why hydrogen when led through a palladium tube heated to 130 diffuses through the walls as if they were porous. At a red-heat hydrogen diffuses still more readily through the metal. Palladium plays the part of a dyad and tetrad in its compounds. We are acquainted with, a palladous oxide : PdO, and a corre- sponding palladous chloride : PdCl.,, with a palladic oxide : PdO„ and with palladic chloride : PdCl 4 . On dissolving the metal in nitric acid, a brown solution is pro- duced which contains Palladous nitrate: (N0 2 ) 2 2 Pd, and which yields yellowish brown deliquescent prisms when evaporated to a syrup. The solution serves to detect iodine and to separate this substance from bromine and chlorine, since palladous iodide is quite insoluble in water, while the corresponding bromide and 568 Text-Book of Inorganic Chemistry. chloride dissolve easily. If the dry salt is carefully heated as long as nitrous fumes are evolved, pelliculous oxide : PdO, remains behind as a black powder scarcely soluble in acids. If palladium is dissolved in aqua regia, a dark-brown liquid results, which contains palladic chloride : PdCl 4 ; and if this solution mixed with the alkaline chlorides, similar double salts to those containing platinic chloride are produced. Potassium Chlorpalladinate : 2KCl,PdCl 4 or K,PdCl 6 , crystal- lizes in dark red octahedra. When small quantities of caustic potash are added to this salt a dark yellow powder of palladic hydrate containing potash is produced. And this when washed with warm water leaves a residue of black palladic oxide : PdO.,, which is scarcely attacked by oxy-acids. When the solution of palladium in aqua regia is evaporated down, it loses chlorine and is converted into palladous chloride, which crystallizes out when the solution is slowly evaporated as brownish-red prisms of the composition : PdCl 2 + 2H„0, and which do not deliquesce in the air. This chloride also produces crys- talline double salts with the alkaline chlorides. A solution of palladous chloride in water is a very delicate reagent for carbonic oxide. The smallest quantities of this gas when brought into contact with the solution cause reduction and separation of black palladium. Palladous Iodide : Pdl 2 , is quite insoluble in water and in dilute acids ; it is obtained as a black precipitate when a soluble iodide is added to a solution of palladous nitrate. The correspond- ing chloride and bromide are soluble in water. Palladous chloride unites with ammonia and produces com- pounds similar to the ammoniacal compounds of platinum (p. 563). IRIDIUM. Chemical Symbol : Ir. — Atomic Weight: 193. Iridium resembles platinum in many respects, and only differs from this metal in some of its less important properties. Its specific gravity is 22-4 ; and it is harder and more brittle, and much more Iridium. 569 difficultly fusible than platinum. In the compact state iridium is not acted upon by aqua regia, it is only dissolved when in a finely divided state as iridium-black. The metal may be oxidized by heating it strongly with a mixture of caustic potash and nitre. Iridium is contained in the so-called platinum residues which remain after the ore has been digested with aqua regia. It is usually present united with osmium as an alloy of osmiridium in the form of very hard granules. This alloy may be decomposed by mixing it with common salt and heating the mixture in a stream of moist chlorine. The osmic chloride which is then produced is decomposed by the water contained in the moist gas, and converted into osmic acid, which is a volatile compound and which condenses in a receiver attached to the other end of the tube in which the mixture is heated ; at the same time the iridium becomes sodium- iridium chloride. The residue remaining in the tube and con- taining this salt is extracted with water, mixed with excess of sodium carbonate and evaporated to dryness. The dry mass when heated to redness yields iridium sesquioxide, which remains when the sodium chloride and carbonate are dissolved in water, and which is easily reduced to metallic iridium when gently heated in a stream of hydrogen. Besides the compounds of iridium with oxygen and chlorine which correspond to those of platinum with the same elements — i.e. iridic oxide : IrO.„ iridions oxide : IrO, iridic chloride : IrCl„ and iridous chloride : IrCl 2 , it also forms iridium sesquioxide : lr 2 3 , and iridium sesquichloride : Ir 2 Cl 6 . The former of these com- pounds remains as a black powder when sodium-iridium chloride is fused with sodium carbonate, and the fused mass extracted with water (see above). Iridium sesquichloride is obtained when iridium is strongly heated in chlorine as a yellowish-brown amorphous mass, insoluble in water and acids. Iridic Chloride : IrCl 4 , is obtained by dissolving the sesquioxide in aqua regia. It is much darker in colour than platinic chloride, and its solution is of a dark-brown colour. It produces double salts with the alkaline chlorides similar to those of platinum, but much darker in colour. Potassiiim-iridium chloride: 2KCl,IrCl 4 or K 2 IrCl 6 , forms small dark red, nearly black octahedra, insoluble in cold water and only dissolving slightly when warmed. The ammonium salt is very similar. 570 Text-Book of Inorganic Chemistry. RHODIUM. Chemical Symbol: Rh. — Atomic Weight: 104. This rare element only occurs in very small quantities in plati- num ores, and has therefore been but little studied. In its chemi- cal properties it is most closely related to iridium. Rhodium is a malleable and ductile metal, more easily fusible than iridium, but less so than platinum, and is insoluble in aqua regia, except when alloyed with platinum, in which case it dissolves with this metal. When heated in the air it oxidizes slightly. Its oxides and chlo- rides correspond to those of iridium ; the most stable are rhodium sesquioxide : Rh.,0 3 , and rhodium sesquichloride : Rh 2 Cl d . OSMIUM. Chemical Syinbol : Os. — Atomic Weight: 199. The separation of this metal from iridium has been described on the preceding page. If the vapcurs of osmic acid are led into caustic potash, the resulting liquid evaporated to dryness with much ammonium chloride, and the dry residue heated to redness, metallic osmium remains behind as a black powder when the fused mass is extracted with water. The metal remains unmelted at the highest temperatures ; even in the oxy-hydrogen flame it only becomes a little more dense. It may be obtained as a hard crystalline- powder when fused with tin, and the tin afterwards removed by hydrochloric acid. Osmium has a specific gravity of 22-5, and is apparently the densest of all substances. Aqua regia, or even nitric acid, oxi- dizes osmium to osmic acid. In the finely divided state, the metal burns when heated in the air, and produces the same substance. The compounds of osmium with oxygen and chlorine mostly correspond to those of platinum, thus : osmous oxide : OsO, and osmic oxide : Os0 2 , osmous chloride : OsCl 2 , and osmic chloride : OsCl 4 , and besides these an osmous anhydride : Os0 3 . But by far the most important and interesting of the osmium compounds is Osmic Anhydride or Osmic Acid : Os0 4 , in which the metal behaves as an octad element. This compound is produced when osmium is heated in a stream of oxygen, and is then deposited in Ruthenium. 57 1 the cooler parts of the tube in colourless crystals. It is also ob- tained when the metal or its alloy, osmiridium, is heated in a stream of moist chlorine. The corresponding chloride, which is then pre- sumedly first produced, but which is not known in the free state, is at once decomposed into osmic acid by the moisture of the damp gas. It melts under ioo c to a colourless liquid, boils at a slightly higher temperature, and sublimes in colourless glistening needles. Osmic acid possesses a piercing odour, somewhat resembling chlorine, and powerfully attacks the lungs and eyes. Great care must therefore be exercised in experimenting with this substance. Osmic acid, or more correctly, osmic anhydride, dissolves in water, but a compound with this liquid, which would be the true osmic acid, has not yet been prepared. Most reducing substances sepa- rate metallic osmium from its solution, a change which also goes on when the liquid is exposed to the light. A dilute solution is used by biologists for hardening organic tissues. — Potassium osmate is produced when osmic acid vapours are led into caustic potash, and separates from its solution as a crystalline powder on the addi- tion of alcohol. RUTHENIUM. Chemical Symbol : Ru.— Atomic Weight: 104. This metal, which resembles iridium and osmium, is less diffi- cultly fusible than the latter, but more so than the former metal. It has a specific gravity of 12-3, is scarcely soluble in aqua regia and insoluble in other acids. When fused with caustic potash and nitre it is converted into soluble potassium ruthenate, and when chlorine is led into a solution of this salt the temperature becomes high enough to volatilize the oxide ruthenic acid : Ru0 4 , which is then formed. This compound then condenses in a cooled receiver as a yellow crystalline mass. In this respect, ruthenium resembles osmium, and the two metals are, at present, the only elements known of which one atom can unite with four atoms of oxygen. APPENDIX. 1 Atomic and Molecular Weights. The atomic weights of the elements play so important a part both in theo- retical and in practical chemistry, that a brief account of the methods used to obtain these numbers may not be out of place, even in an elementary text-book of the science. In deciding upon the number which shall be considered as the atomic weight of an element, two distinct questions must be experimentally answered. In the first place, the weight of the element must be found which unites with a given weight of some other element or elements, of which the atomic weight or weights are accurately known ; and, secondly, it must be decided whether the number, so obtained, is really the atomic weight, or whether it is only some simple multiple or fraction of it. It is hardly necessary to remind the student that the atomic weights are simply relative and not absolute numbers, and that, since the atomic weight of hydrogen is taken as unity, the atomic weights of the other elements simply express how much heavier their atoms are than those of hydrogen. The first thing to be done in determining the atomic weight of an element is, therefore, to find the ratio (chemical equivalent) in which it combines with other elements. For this purpose, the element is either made to unite with other elements of which the atomic weights are accurately known, or else one of its compounds with these elements is decomposed and its composition accu- rately determined. Such experiments appear simple enough theoretically, but in practice are surrounded by many and often almost insuperable difficulties if accurate results are to be obtained. The element or its compound must be prepared in a state of absolute purity, and only perfectly pure chemicals must be used in all succeeding operations. The greatest care must be taken that no impurities are introduced either from the air or from the vessels employed, and, what is more difficult to avoid, that a constant error does not occur through the entire series of experiments and so vitiate them all. To eliminate this constant error as far as possible, it is advisable to use not one but several compounds of the element for these determinations, and to employ material which has been obtained from various localities. The chemist must have an accurate balance By the Editor. 574 Text-Book of Inorganic Chemistry. at his disposal, and must know how to use it properly ; he must find before- hand whether his set of weights are exact ratios of one another, for example, whether the io-gramme piece is exactly ten times as heavy as the i -gramme, and one-tenth as heavy as the ioo-gramme piece, and if they differ from this ratio, as the best weights always do, he must find how much they differ and make the necessary corrections in his weighings. Finally, he must add to all his results the weight of air displaced by the substance weighed, which will vary according to the specific gravity of the substance, and according to the height of the barometer, tHe am -runt of moisture contained in the air and the tempera- ture at the time of weighing. The above brief description will give some idea of the delicacy of the opera- tions which are recessary for a correct determination of the chemical equivalent of an element, and it may be safely affirmed that such investigation demand greater skill on the part of the operator than almost any other chemical experi- ments. Supposing, then, that the ratio in which the element combines with other elements for its equivalent) has been accurately determined, it remains to find its true atomic weight. Let us imagine, for example, that a given weight of pure carbon (the diamond) was burnt in a stream of pure oxygen, that the car- bonic acid so produced was carefully collected and weighed by absorbing it in a series of potash-tubes, and that it was then found that 3 '66 grammes of cir- bonic acid were produced for every gramme of carbon burnt. It would then be known that i gramme of carbon, when burning to carbonic acid, combines with 3 "66— i 'oo = 2*66 grammes of oxygen, or that 6 grammes of carbon com- bine with 16 grammes of oxygen. If, then, the atomic weight of oxygen is j 6, that of carbon must be 6, or some simple multiple of this number. It would be 6 if the compound contained an equal number of atoms of each constituent, 12 if it contained twice as many oxygen atoms as those of carbon, and 3 if twice as many carbon as oxygen atoms were present. Or, again, suppose that silicon chloride were decomposed in some suitable manner so that its chlorine was converted into silver chloride, and that, knowing the atomic weights of silver and chlorine and the composition of silver chloride, it was found that 100 grammes of silicon chloride contain 83*53 grammes of chlorine. Such a result would show that 100 — 83*53 = 16*47 parts of silicon combine with 83*53 P arts of chlorine, or that 7 parts of silicon unite with 35*5 parts (one atom) of chlorine. Hence the atomic weight of silicon might be 7 or 14 or 28 or 3*5. In order to decide which of these numbers correctly represents, the atomic weight, further experiments are necessary and other considerations must be brought to bear on the problem. The further con- siderations which assist us in obtaining a solution of this question may be classified under three heads, viz. : — • 1. The relations between the molecular weights of gases and their density (Avogadro's law) ; ^. The relations between the specific heats of the solid elements and their atomic weights (Dulong and Petit's law) ; and 3. The relations between the composition of a solid compound and its crys- talline form (isomorphism). Appendix. 575 Let us new see how we can use these three laws to fix the atomic weights of the elements. I. The determination of the atomic weight of an element from the density of its gaseous compounds. The important law, first enunciated by Avogadro and therefore called after him, states that equal volumes of all gases whether simple or compound con- tain an equal number of molecules, or, in other words, that the density of a gas is proportional to its molecular weight (see p. 50). Avogadro's law is proved not only because it agrees with all known chemical facts, but also because it expliins the physical properties of gases in the only possible manner. In the kinetic theory of gases, as worked out by Maxwell, Clansius, and others, it is supposed that a gas consists of a large number of small particles (molecules) moving at a high but variable velocity, and, therefore, constantly colliding with one another and with the walls of the vessel which contains the gas. This collision of the molecules with the walls of the vessel produces the pressure which the gas exerts, or the pressure which must be applied to the gas to prevent it from expanding. If, now, we call m the mass of each molecule and v the the mean velocity with which the molecules are travelling, the aver- age kinetic energy possessed by each molecule is : $ m z/ 2 . But if two sub- stances are at the same temperature there is no change in the temperature of either when they are brought into contact with one another ; there is, it is true, an interchange of heat between them, but each gains as much heat as the other loses. Hence, if two average molecules of the same or different gases at the same temperature are brought together, no change in their temperature is pro- duced, and no change in their kinetic energy (measured as heat) takes place — i.e. the kinetic energy of each is equal, or : — 4 otxV = i m 2 v a * . . . . (1) Again, the total pressure which is exerted on the vessel containing the gas is represented by the sum of the kinetic energy of all the molecules present in the vessel, and if we call this number n, and the total pressure P we have : — P = ^nmv*, or since we only measure the pressure (p) in one of the three directions in which it is exerted : — p = ^nmv 2 . . . . (2) And by experiment, the pressure exerted by equal volumes of all gases at the same temperature is equal, hence : p = p 2 , and \n± fflj v x 2 = ^n 2 m. i v i 2 . But we have already seen (1) that at the same temperature, 4*2^,2 = tyn % v£. Therefore :— «i = n z> or the number of molecules contained in equal volumes of all gases at the same temperature and pressure is the same. Further, the density of a gas is clearly the product of the number of mole- cules contained in unit volume into the mass of each— i.e. d 1 = n l m 1 and d % ~ n 3 m 2 , and since n L -n 2 , we get : d v : . d 3 = m x m 2 , 576 Text-Book of -Inorganic Chemistry. or the density of a gas is proportional to the mass of its molecules or its molecular ■weight. 1 This is a brief risumi of the physical proof of Avogadro's law with which we started. It must now be decided with what standard the molecular weights are to be compared. We know by experiment that a given volume of oxygen, for example, weighs 16 times more than the same volume of hydrogen at the same temperature and under the same pressure, and, therefore, that the molecule of oxygen must be 16 times as heavy as the molecule of hydrogen, but we do not know how many atoms are contained in the molecule of each gas. That the molecule of hydrogen contains at least two atoms is clear from the fol- lowing. In the union of hydrogen and chlorine equal volumes of the two gases combine to form the same volume of hydrochloric acid. Or, one molecule of hydrogen unites with one molecule of chlorine and produces two molecules of hydrochloric acid. Each molecule of hydrogen (and of chlorine) is therefore broken up into two parts which cannot be less than the atom, but may contain more than one atom. Hence, therefore, the molecule of hy- drogen is considered to consist of 2 atoms ; it is not impossible, but highly improbable, that its molecule may contain some multiple of this number of atoms. The atomic weight of hydrogen being taken as unity, and its molecule con- sisting of 2 atoms, its molecular weight becomes 2. And since air is 14*44 times as heavy as hydrogen, the molecular weight of any gas is found by mul- tiplying its density (compared with air as unity) by 2 x i4'44 = 28'88(p. 15). In this way the molecular weights of a number of the elements and of many com- pounds have been determined. But in making these determinations our ex- periments are confined to those substances which are volatile at comparatively low temperatures ; and, in the case of a compound, to those which do not completely decompose when heated above their boiling-point, or which do not dissociate or decompose when heated, and recombine again on cooling. And, from the experimental difficulties which surround the determination of many vapour densities as well as the deviation from Boyle's and Charles' laws at temperatures only slightly above the boiling-point of a substance, it follows that the molecular weights obtained in this way are never so correct as those derived from the ordinary analysis of the substance. The density of the gas only tells us whether the molecular weight is a multiple or fraction of a certain number, and does not give us the exact number itself. We have seen (p. 574) that, by- chemical analysis, silicon chloride contains 7 parts of silicon united with 35 '5 parts of chlorine, and its molecular weight is therefore either 42-5 or some exact multiple of this. The density of this gaseous compound has been found to be 5"94, which, when multiplied by 28 '88, gives i7i'S, consequently its true molecular weight is 42-5 X 4 = i7o'o, not 171 -5, nor any other multiple of 4 2 '5- The following table gives the vapour densities (d) of those of the elements for which this has been determined ; the third column contains the molecular 1 Maxwell, Theory of Heat, pp. 289 et sea. Appe?idix. 577 weight (28'88flf) to which this points ; and the fourth the true molecular weight (At);- Element d 28-88 xd M Hydrogen . ... o"o693 2'0O Ho = 2 Oxygen ...... 1-1056 3 J '93 Oo = 32 Nitrogen ...... 0-9713 28-05 No = 28 Chlorine 2 '45 7076 CI, = 71 Bromine 5 '54 160-00 Br. 2 = 160 Iodine .... 872 251-8 I, = 254 Selenium (at 1,400°) . S -68 164-0 Se 2 = 158 Sulphur /at 1,000°) . 2 '24 647 S 3 = 64 ,, (at 500°) 678 I9S8 S = 192 Tellurium . . . . 9-08 262-2 Te 2 = 256 Mercury ..... 6-99 201 -8 Hg = 200 Cadmium . 3 '94 113-8 Cd = 112 Phosphorus 4-48 129-4 P 4 = 124 Arsenic 10-42 300-9 Ab 4 300 From this table it will be seen that the majority of the elements contain two atoms in the molecule. To this rule, the hexatomic sulphur, at temperatures just above its boiling point, forms an exception, as well as the four elements — mercury, cadmium, phosphorus, and arsenic. Mercury and cadmium contain only one atom in their molecules, an assumption which is strongly supported in the case of mercury by certain physical properties of its gas ; while the mole- cules of phosphorus and arsenic are made up of four atoms each. We possess, then, in Avogadro's law a valuable and certain means for find- ing the molecular weights of substances which can be converted into the gaseous state without decomposition ; but it remains to be seen how we can utilize the same generalization for determining the atomic weight of an element. In the first place, it is evident from the preceding table that the molecular weight of an element has no definite relation to its atomic weight, and we are not therefore justified in making any deduction from the molecular weight of an element as to its atomic weight. But the true atomic weight of an element may often be found from the molecular weights of its compounds. The atom of an element may be defined as the smallest quantity which can exist in any compound, and among a large number of compounds of one element it is extremely probable that the molecule of one or more of them will contain only one atom. If, therefore, we have determined the vapour densities, and consequently the molecular weights of a number of compounds of one particular element, we may take the atomic weight as the smallest quantity of the element which occurs in the molecule of any of these compounds. The result so obtained is never absolutely certain, because a compound might be discovered containing a less quantity of the element than its received atomic weight, and it has but little value when the number of compounds compared together is small, but approaches certainty when the densities of a large number of compounds are known. The number obtained from this comparison is the maximum value possible for the atomic weight — it may be smaller, but it can- not be larger. P P 578 Text-Book of. Inorganic Chemistry. A few examples may serve to explain more clearly how the atomic weight of an element may thus be found. Among those elements of which numerous gaseous compounds are known is chlorine. If we examine all these com- pounds, we find that none of them contains less than 35'5 parts of chlorine in its molecule, and we therefore consider this number as the atomic weight of the element. The following table gives the percentage composition of four compounds of chlorine, their vapour density [d), the molecular weight indicated by this density 28 "88 x d ) , and the true nolecular composition (M):- Compound Percentage composition d 28-88x0' M Hydroch oric acid Mercuric chloride Phosphorous chloride Silicon chloride |H = 2741 ICI =97-26! iHg =73-801 ICI =26-20! J P =22 56) ICI = 77-44 f jSi =16-471 ICI =83-53 1 1-247 9-80 4-88 5-94 36-0 283-0 140-9 171-6 |H =i-o ICI =35-5 J Hg = 200 ICI =71 = 351x2 jP =3t ICI =106-5 = 354x3 (Si =28 ICI =142 = 35^x4 The table shows that the chlorine contained in the molecules of the four compounds is in the proportion of 1:2:3:4, and as on examining a large number of other compounds of the element none has been found with less chlorine in its molecule than hydrochloric acid, we consider that this compound contains one atom of chlorine, mercuric chloride two atoms, &c. At present, therefore, the atomic weight of chlorine is taken as 35-5, but if, at any future time, a compound of chlorine should be discovered with only 1775 parts of chloiine in its molecule, the present atomic weight would have to be reduced to this number. In the same way, from the comparison of a still larger num- ber of vapour densities, the atomic weight cf carbon has been fixed at 12, and that of oxygen at 16. And the corresponding numbers for many other elements have been determined with more or less certainty, according to the number of compounds examined. The atomic weights obtained from Avogadro's law are increased in value by their agreement with those resulting from other considerations. They agree with those derived from the specific heats of the elements in the solid state, and from the general properties of the element and its compounds. II. The determination of the atomic weight of an element from its specific heat in the solid state. The specific heat of a substance is the quantity of heat required to raise 1 gramme of the substance i° C. as compared with the quantity of heat required to raise 1 gramme of water through the same increment of temperature. The specific heats (or capacities for heat) of the various elements differ very con- siderably from one another. But if, instead of comparing together the quantities of heat required to raise equal weights of the elements 1° we compare those necessary to similarly increase the temperature of the atomic weights of the elements (in the solid state), we arrive at a different result : we then find that all these quantities of heat are nearly equal to one another. This important law, to Appendix. S79 which we have already referred (p. 70), was discovered by Dulong 'and Pelit, and is therefore usually named after them. The law may also be expressed by saying that the capacity for heat of all the elementary atoms in the solid state {atomic heat) is the same, or, that the product of the specific heat of a solid element and its atomic weight is a constant. The following table gives the specific heats of some of the elements in the solid state (s), the temperature (t) at Which these quantities have been determined, the atomic weights (a. w.) of these elements and their atomic heats (a. h.) ; or the product of the specific heat and atomic weight in each case : — Element * t a. TV. t*. k. Lithium 0-941 64 7'0 6-6 Boron ... 0*366 233 II 'O 4'° Carbon (Diamond) d'459 985 I2'0 5 '5 Silicon ... 0*203. 232 28-0 57 Phosphorus . 0*189 19 31-0 3'9 Sulphur 0*171 50 32'0 5 '5 Potassium o'i66 -34 39'° 6'5 Iron 0'TI2 +31 56-0 S'3 Copper . 0-095 58 63 '4 6'o Silver . 0-056 50 108 -o 60 Tin 0-056 5° 118-0 6-6 Platinum 0-0324 55 198-0 6-4 Lead 0315 55 207 "o ,6-5 This table shows us, in the first place, that the atomic heats of the very various elements which it includes are all approximately equal, but we also see that the numbers differ considerably among themselves. These differences are principally due to the change of specific heat with temperature, which nearly always increases as the temperature rises. The most striking examples of this are seen in the elements carbon, boron, and silicon. At the ordinary tempera- tures the specific heats of these three substances are : — Carbon = 0-132 (at 33°) Boron = 0-238 (at 26°) Silicon = C170 (at 22 ) numbers which do not at all agree with the law — e.g. the atomic heat of carbon is then o'i32x 12 = 1-58. But as the temperature rises the specific heat of each rapidly increases until it becomes, to a certain extent, constant, and the atomic heat is then found to be normal. Thus, at high temperatures, the three elements have the following specific heats : — Carbon = 0-459 ( at 9 8 5°) Boron = 0-366 (at 233°) Silicon = 0-203 (at 232°) For carbon and silicon these numbers correspond to the normal atomic heat, and that for boron probably would do so at a slightly higher tempe- rature. TJulong and Petit's law is therefore only approximately true, or the specific p p 2 580 Text-Book of Inorganic Chemistry. heat is a function not only of the atomic weight, but also of other properties of solids with which we are at present but little acquainted. But notwithstanding this, the law enables us to decide between one of two or more numbers, differing considerably from one another, in determining the atomic weight of an element. All that is necessary to find the required atomic weight is to divide the mean atomic heat (6'4) by the specific heat of the element in the solid state. The quotient is then the approximate atomic weight, which has to be corrected by ordinary analytical results. Two examples may be given of how the specific heat of a solid element controls its atomic weight. When the metal indium was discovered by spectrum analysis it was found in association with zinc, a metal which it closely resembles both in the free state and in its compounds. The only oxide which indium forms was therefore thought to correspond to the only oxide of zinc — i.e. to be a monoxide of the formula : InO ; and this gave an atomic weight of 75 '6. But soon afterwards the specific heat of the metal was determined and was found to be C057, which gives an approximate atomic weight of — i- =ii 2| and hence its true atomic weight is, not 75'6, 0-057 but y$'6 x i'5 = ii3"4. From this it also follows that the formula of its oxide becomes ln„0-, and that the composition of all its other compounds must be correspondingly changed. Two atoms of indium now correspond to three of what were formerly considered as the atoms of the element. A precisely analogous case is to be found in the atomic weight of the metal beryllium, which, during late years, has been taken as 9*1, and its oxide considered as a monoxide with the formula : BeO. Recently, however, the specific heat of the metal has been determined and found to be 0^445, which points to an atomic weightof — — = 14. The true atomic weight thus becomes 9*1 x i'5 = i3'65, o'44S and the oxide is a sesquioxide : Be 2 5 . III. The determination of the atomic weight of an element from the crys- talline form of its solid compounds. It was discovered by Mitscherlich that many substances which crystallize in the same form with the same or nearly the same angles have a similar chemical composition. Such substances are said to be isomorphous with one another, and this law of isomorphism may be used to determine the atomic weight of an element in the following way. If two substances crystallize in the same form and the composition of one is known, the composition of the other may be inferred to be similar, and thus the atomic weight of any particular element which it contains may be fixed. The compounds alumina, ferric oxide, and chromic oxide all crystallize in the same form of the hexagonal system, and from the composition of one of them we can therefore deduce that of the other two : if ferric oxide = Fe 2 5 , then alumina = A1 2 5 , and chromic oxide = Cr 2 5 ; and from this ratio the atomic weights of aluminium and chromium may be deduced if that of iron is known. The atomic weights obtained by isomorphism are not, however, so trust- worthy as those based upon Avogadro's or Dulong and Petit's law, and for the following reasons. That the form in which a substance crystallizes does not Appendix. 581 entirely depend upon its chemical composition is at once proved by the fact that the same compound may be dimorphous, or may crystallize in two distinct and entirely different forms — e.g. calcium carbonate, arsenious anhydride, sul- phur, &c. In the next place, the composition of two compounds may be exactly similar, but they need not crystallize in the same form — e.g. the nitrates and the corresponding metaphosphates. And, finally, two substances may crys- tallize in precisely the same form without having a similar composition — e.g. sodium nitrate {NaN0 3 ) and calcite (CaC0 3 ) ; sodium sulphate (NaoS0 4 ), and barium permanganate (BaM 2 8 ) ; cadmium sulphate (3CdS0 4 + 8H 2 0), and didymium sulphate (Di 2 3S0 4 + 8H 2 0). The evidence derived from isomor- phism is, therefore, only of value so far as it confirms the results derived from the other two laws, or when these two laws cannot be applied. Thus the isomorphism of the compounds : — KoSiF 6 + H 2 K 2 TiF 6 + H 2 K 2 ZrF G + H 2 KoSnF. + H 2 afforded valuable evidence concerning atomic weights of silicon, titanium, and zirconium, until they were confirmed by the other two more certain laws. The following is a list of some of the more important isomorphous groups, showing the chief elements which are connected together by this law : — Groups. Elements Connected, (i.) Phosphates P and As Arsenates lii.) Carbonates of the composition :RC0 3 , Ca, Mg, Zn, Fe, Mn where R = Ca, Mg, Fe," Zn, Mn. (iii.) Double sulphates of the general formula: M' 2 S0 4 + M"S0 4 + 6H 2 0, where M' =K, NH 4 Fe",Co, Ni, Zn, Mg, Mn'' M'" = Fe", Co, Ni, Mg, Zn, Mn" (iv.) The Alums of the general formula: M'M'" 3 S0 4 +i2H 2 0, K,(NH 4 ),Rb,Cs,Na,Tl, Ag where M' = K,NH 4 , Rb, Cs, Na, Tl, Ag Al, Fe'", Mn'", Cr" M'" = A1, Fe"', Mn'"Cr (v.) Chromates, Manganates, Sulphates, Cr Ti , Mn*', S, Se, and Selenates (vi.) KMn0 4 and KCIO4 Mn™, CI'" (vii.) Ag,S0 4 , Na 2 S0 4 Ag, Na (viii.) K 2 SiF ,H 2 O ; K 2 TiF 6 ,H 2 ; Si, Ti, Zr, Sn K„ZrF 6 , H 2 ; K 2 SnF 6 ,H 2 0. (ix.) Apatite : 3 Ca 3 P 2 8 , Ca(Cl,F) 2 P, As, V Pyromorphite : 3Pb 3 P 2 8 ,PbCl 2 Mimetesite : 3Pb 3 As 2 8 ,PbCl 2 Pb, Ca Vanadinite : 3Pb 5 V 2 8 ,PbCl 2 582 Text-Book of Inorganic Chemistry. Relations between the Atomic Weights. It has been already remarked (pp. 65 et seq.) that certain of the elements which possess similar chemical properties and which usually occur associated together in nature may be arranged in groups of three ; and that when this is done the atomic weight of one is about the mean of the other two. The fol- lowing have been mentioned as examples of this triple grouping :— (i.) Chlorine = 35-5 1 Bromine = 8o*o[Mean = 81 '25 Iodine — 127'oj (ii. ) Sulphur — 32 'o\ Selenium = 79*0 ^Mean = 80 'o Tellurium = 128 *o ) (iii.) Lithium = yo] Sodium = 23 *o [■ Mean = 23 'o Potassium = 39' o J These relations cannot be accidental, and it certainly does appear that the properties of an element and of its compounds are related in some way or other to its atomic weight. Of the various numerical relations which have been supposed to exist be- tween the atomic weights we shall here only refer to one hypothesis, which was put forward by Prout, and is commonly known as Prout's law. With regard to the connexion between the atomic weight of an element and its properties, the most important theory is that known as the Periodic law. This theory, which has been developed in late years by Newlands, Lothar Meyer, and Mendelejeff, supposes that the properties of the elements change periodically as we pass from one to the other when arranged according to their atomic weights. Prout ' s Law. — The hypothesis which is known under this name supposes that the atomic weights are all multiples of that of hydrogen by whole numbers, and that the elements may therefore be considered as compounds of hydrogen of varying degrees of complexity. Prout's hypothesis was advanced at a time (1815) when the atomic weights were much less accurately known than they are at present, and when they were usually represented by the nearest whole numbers to which the results pointed. It soon became clear, however, that the atomic weights of certain of the elements, notably chlorine, could not be re- presented by whole numbers, and in order to make the hypothesis applicable to these it was proposed to call the atomic weight of hydrogen 2, or even 4. But even with such devices as this, accurate determinations of the atomic weights of many of the elements by which every possible error was avoided or allowed for have shown that in most cases the hypothesis is certainly not true. The atomic weights which are given on p. 60, and used in the text of this book, have been mostly taken as whole numbers to avoid fractions ; but for some elements much more accurate results than this have been obtained by careful investigators and have been confirmed by others. The following table Atpendix. 533 gives a list of these more accurate atomic weights, which shows that only in a few cases is Prout's hypothesis approximately true. In the second column the ratio between the atomic weights of hydrogen and oxygen is 1 : 15*96, and in this column the atomic weights are calculated for these numbers. We may, however, consider the atomic weight of oxygen as 16, when all the other atomic weights will be slightly larger. These numbers are given in the third column. It will be noticed that the atomic weights there more nearly approach whole numbers, and hence Prout's law has not yet been completely forsaken by some chemists. Table of Accurately Determined Atomic Weights. Element O = 15-96 = 16 Aluminium .... 27-02 27-09 Antimony . 119-8 I20'I Barium 136-81 I37't5 Boron 10-92 10-95 Bromine 79-76 79-96 Cadmium 11 1 75 It2'03 Cassium I3 2 ' 6 5 132-98 Calcium 39 '95 40-05 Carbon 1 1 '97 12 -OO Chlorine 3S'37 35 '46 Copper 63-18 6 3'34 Hydrogen . I'OO 1 0025 Indium 1 13 '4 "37 Iodine 126-55 126-87 Iridium I9 2 '5 192-98 Iron . 55 '9 56-14 Lead . 206-43 206-95 Lithium 7-01 7 '03 Mercury 199-7 200 "8 Nitrogen 14-01 14 '°5 Oxygen 15 '96 16-00 Phosphorus 30-96 31-03 * Platinum I94'5 194-99 Potassium 39-02 39-12 Rubidium 85-2 85-4 Selenium 78-8 790 Silver . 107-67 107-94 Sodium 22-99 2 3'°S Strontium 87'3 87-5 Sulphur 31-98 32"o6 Vanadium 5i'i 51 '2 The Periodic Law. If the elements are a-ranged according to their atomic weights, commencing with the smallest, it is found that those elements with similar chemical and physical properties occur at regular intervals. Various tables have been drawn up to exhibit these periodic changes in the properties of the elements, one of which 584 Text-Book of Inorganic Chemistry. is given below, and the series of facts thus shown to be connected with the atomic weight are collectively known as the periodic law. We owe this useful generalization principally to the labours of Lothar Mejer and Mendelejeff: — I II III IV V VI VII VIII H 1 | : Li 7 Be ? ■ B 11 C 12 N 14 16 F 19 -"■ Na 23 Mg 24 Al 27 H 28 P 31 s 32 CI 35 'S K ■39 Ca 40 Ti 48 V 5i C r 52 Mn S5 . Fe 56 Ni 58-5 Co 59 [Cu] 63 Zn 65 Ga 70 j As \ 7S Se 79 Br 80 A Rb 85 Sr 88 Y ? Zr 90 Nb 94 Mo 96 Ru 103 Rh 104 Pd 106 / 108 Cd 112 In "3 Sn 118 Sb 120 Te ? I 127 t Cs 133 Ba 137 La 138 Ce 141 Di 145 1 Er 166 Yb 172 Ta 182 W 184 Os ? Ir 193 Pt IIS '(J [Au] •199 Hg 200 Tl 204 Pb 207 Bi 210 l! Th 231 1 u 240 Referring to the accompanying table, which commences with hydrogen and ends with uranium, we find that, omitting hydrogen, the first seven elements correspond in their properties to the next following seven : thus, Li corre- sponds to Na, C to Si, N to P, &c. Following on in the series we see again that the next seven (less one) also correspond in a regular manner to the pre- ceding seven. In this way the elements become classified into various groups, which are shown in the vertical columns of the table, and which contain closely- related elements. Group I., for example, contains Li, Na, K, Rb, Cs, and the allied metals Ag and Cu, as a sub-group ; Group II. includes the metals of the alkaline earths, together with Mg, Zn, Cd, and Hg. The elements in each of these groups are not only allied chemically but also often connected by similar Appendix. 585 physical properties ; they further possess the same atomicity in some of their compounds. The elements included in the first group are monads, those in the second group are dyads, and so on. This brief description of the periodic law, which space does not allow us to expand more fully, shows clearly that the properties possessed by an element are closely connected with its atomic weight ; and that the method in which the elements are arranged is to some extent correct, is corroborated by other con- siderations. It will be noticed that there are several gaps in the above table' — e.g. between Ca and Ti, Ga and As, Mo and Ru, neglecting those among the elements of higher atomic weights. These gaps are supposed to be occupied by undiscovered elements, a supposition which received a remarkable confir- mation in the recent discovery of the metal gallium. Mendelejeff had pre- viously predicted that such a metal might exist, and had described its properties and approximate atomic weight under the name of ekaluminium. And when a few years afterwards the metal now called gallium was discovered, its pro- perties were found to agree almost exactly with Mendelejeff 's hypothetical ek- aluminium. The periodic law has also done good service in correcting some uncertain atomic weights. Up to a recent date the metals of the cerium and yttrium groups were considered as dyads and their oxides as monoxides ; but Mendelejeff showed that the atomic weights which they then possessed did not agree with their properties according to the periodic law, and proposed to multiply these numbers by I '5, so that the oxides would become sesquioxides. The atomic weights which he had proposed for the cerium metals were after- wards found to be correct as soon as their specific heats were determined. In the same way the atomic weight of uranium was formerly thought to be 120, but was supposed by Mendelejeff to be 240 — a supposition which was also confirmed by a determination of the specific heat of the metal. On the other hand, it is none the less true that the periodic law fails to find a position for some elements of which the atomic weight has been confirmed by Dulong and Petit's law, or by careful analysis. According to the specific heat of beryllium its atomic weight should be 13 '6 and not 9/1 (p. 430), but with the former atomic weight it would fall between caibon and nitrogen, and would be entirely out of place. Again, the element tellurium ought to come before iodine in the tables, but a recent determination of its atomic weight has con- firmed the number 128. And, finally, there ought not to be so many undis- covered or missing elements between didymium (145) and erbium (166). The periodic law is, therefore, a valuable but still imperfect representation of certain chemical facts, and we must be careful not to accept the values for the atomic weights to which it points unless confirmed by evidence derived from other sources. It must for the present be still regarded as a hypothesis, until it has been more fully developed and until the anomalies which it contains have been explained. 586 Text-Book of Inorganic Chemistry. TABLES. I. — Weights and Measures. i lb. Avoirdupois . i oz. Avoirdupois i lb. Troy i oz. Troy i gallon ... i pint i gallon of water at 62 F. (i6°'7 C.) weighs 1 cubic foot of water at 62 n F, weighs .... 16 oz. Av. 7,000 grains 437 '5 grains 12 oz. Troy 5760 grains 480 grains 8 pints 2 77' 2 S cu bic inches 20 fluid oz. 70, 000 grains 62"326 lb. Av. 1 metre (m 1 foot 1 inch . 1 litre (1.) 1 pint 1 fluid oz. 1 kilogramme (kg. ) 1 gramme 1 lb. Avoirdupois 1 oz. Avoirdupois 1 grain , 10 decimetres (dm.) 100 centimetres (cm.) 1000 millimetres (mm. ) 39 "37 inches 30*48 centimetres 25 "40 millimetres 1000 cubic centimetres (1 176 pints 568 cubic centimetres 28-4 1,000 grammes (g.) 2*205 lb, Av. 10 decigrammes (dg. ) 100 centigrammes (eg. ) 1000 milligrammes (rag.] I 5'43 2 grains 453*6 grammes 28'35 ., 64'8 milligrammes Appendix. 587 II. — Weight of one cubic centimetre of Dry Air in milli- grammes at varying temperatures and pressures. (Temperature in ° C. and Pressure in mm. of Mercury.) t i>=745 750 755 760 765 770 775 1-268 1-276 1-285 1-293 1-301 I-3IO 1-319 I 1*264 1-271 1-280 1-288 1-296 i'3°5 1 '3 I 4 2 1-259 1-267 1 '275 1-283 1-292 1-300 1-310 3 !' 2 54 1-262 1 -271 1-279 1-287 1-296 I'3°5 4 1-250 1-258 1-266 !'Z7S 1-283 1 -291 1.300 5 1 '245 1 '253 1 -261 1-270 1-2,-8 1-286 1-295 6 1-240 1-248 1-256 1-265 1-273 1-281 1-290 7 1-236 I.-244 1-252 1 '261 1-269 1-277 1-286 8 1-232 1-240 1*248 1-256 1-264 1-273 1-282 9 1-228 1-236 1-244 1-252 1-260 1-268 1-277 10 1-223 1-231 1-239 1-247 1 "255 1-263 1-272 11 1-219 1-227 1 '23S !' 2 43 1-251 1 '259 1 -268 12 1-215 1-223 1-231 1-239 1-247 1 '255 1-263 13 1 -210 1-218 1-226 1 '234 1*242 1-250 1 '259 14 1-206' I "214 1-222 1-230 1-238 1-246 1 '254 IS I -202 I -210 1-218 1-226 i" 2 34 1-242 1-250 16 IT98 1-206 1-214 I "222 1-230 1-238 1-246 17 1 -194 I "202 I -210 1-218 1-226 1 '234 1*242 18 1-189 i -197 1-205 I -213 I "221 1-229 1 '237 19 1 -185 I '193 I'20I I'209 I "217 1-225 1 '233 20 1-181 I -189 I-I97 I-205 I -213 1-221 1-229 21 1 -177 1-185 f'93 I '20I 1-209 1 -217 1-225 22 1 '173 1-181 1-189 1-197 1-205 1-213 I "221 23 1 -169 1 -176 1-184 1*192 I"200 1-208 I "2l6 24 1-165 1-172 1-180 1-188 I "196 1-204 I '212 25 i'i6i 1-168 1 -176 1-184 I-I92 1 -200 1-208 III. — ]Veighl of one cubic centimetre and volume of one gramme of Water at varying temperatures. (Mean of Several Determinations.) Temp. Weight of 1 c.c. Volume of 1 gr. Temp, Weight of 1 c.c. Volume of 1 gr. in grammes in c.c. in grammes in c.c. "999878 I'OOOI22 13 O-999430 1 "000570 I "999933 I -000067 J 4 0-999297 1 "000703 2 0-999972 I "000028 15 C999154 1-000847 3 o-999993 1*000007 16 0-999004 i '000997 4 I "OOOOOO I 'OOOOOO 17 0-998839 1-001162 5 0-999992 I -000008 18 0*998663 1-001339 6 0-999969 I -000031 19 0-99 8 47S 1-001527 7 0-999933 I -000067 20 0-998272 1*00x731 8 0-999882 I'ooonS 21 0*998065 1-001939 9 0-999819 I -00018 1 22 0-997849 I "002I z>6 10 0-999739 I -000261 23 O '997623 1-002383 11 0-999650 1-000350 24 0*997386 I -002621 12 0-999544 I -000456 25 0-997140 I "002868 Text-Book of Inorganic Chemistry. IV. — Tension of Aqueous Vapours at varying temperatures. (Tension in mm. of Mercury.) Temp. Tension | Temp. Tension Temp. Tension Temp. Tension o -IS i '44 + 7 7 '47 + 17 14-40 +50 91-98 — IO 2'IS 8 7 '99 18 IS'33 60 148-88 - 5 3-16 9 8'SS 19 16-32 70 233 '3 1 o 4'57 10 9-14 20 17-36 80 354 '87 + i 4-91 ir 9 '77 21 18-47 90 525 '47 2 5' 2 7 12 10-43 22 19-63 IOO 760-00 3 5-66 13 n'14 23 20-86 IIO 1075-4 4 6 "07 J 4 n-88 24 22-15 120 1491-3 5 6'St 15 12-67 25 23'52 130 2030-3 6 6 '97 16 13 '5 1 26 24-96 140 2717-6 V. — Boiling-point of Water (t) under varying barometric pressure A /. /• t. >• <■ A t. 740 99.26 749 99 '59 758 99 '93 767 100-26 74i 99 '3° 75° 99 '63 759 99-96 768 100-29 742 99'33 7Si 99-67 760 IOO'OO 769 100-33 743 99 '37 75 2 99-71 761 100-04 770 100-36 744 99-41 753 99'74 762 100-07 771 100-40 74S 99 '44 754 9978 763 roo'ii 772 100 '44 746 99-48 755 99-82 764 100-15 773 100-47 747 99'5 2 756 99'85 76S 100-18 774 100-51 778 99 '56 757 99-89 766 IOO'22 775 i°°'55 VI. — Corrections to reduce the Readings of a Barometer with a Glass Scale at different temperatures to 0°. {Landolt and Bornslein.) (The Corrections are in mm., and are to be deducted from the Readings.) Temp. 740 mm. 750 mm. 760 mm. 770 mm. 780 mm. I 0T3 0-13 C13 0-13 0-13 2 0*26 0-26 0-26 C27 C27 3 0-38 °'39 °'39 0-40 0-40 4 O'SI 052 o-53 o'54 °'54 5 0-64 0-65 0'66 0-67 0-67 6 o'77 078 079 o-8o o-8i 7 0-90 0*91 0-92 °'93 0-94 8 I '02 1 '04 1-05 1 '07 1-08 Appendix. 589 VI. — Continued. Temp. 740 mm. 750 mm. 760 mm. 770 mm. 780 mm. 9 "5 117 I'l8 I'20 I '21 10 1-28 1-30 I'3I 1 '33 1 '35 11 1-41 1 '43 I '45 1 '47 1-48 12 1 '54 i'56 158 i'6o 1*62 13 1-66 1 '69 171 i'73 I- 7S h 179 1-82 T'84 1-87 1-89 IS I*Q2 1-95 1 '97 2'00 2*02 16 2'Ot; 2-08 2"IO 2-13 2 - l6 17 2'l8 2'2I 2 '24 2"26 2 '29 18 2-30 2 '34 2 '37 2*40 2 '43 19 2 '43 2-47 2-50 2'53 2-56 20 2-56 2'6o 2-63 266 270 21 2*69 272 276 2-80 2-8 3 22 2-82 2-85 2-89 2 '93 2 '97 23 2 '94 2-98 3-02 3-06 3-10 2 4 3'°7 3 -ii 3-16 3-20 3' 2 4 25 3-20 3' 2 4 3' 2 9 3 '33 3 '37 VII. — Specific Gravity and Composition of Dilute and Concentrated Acids. Sulphuric Acid Nitric Acid Acetic Acid Hydrochl Dric Acid Per cent. H..S0 4 Specific Per cent. HNO, Specific Per cent. HC,H,0. Specific Per cent. HC1 Specific by weight gravity by weight by weight by weight gravity 5 1-031 5 1-030 5 I "0067 5 I '027 10 I '069 IO 1 '059 10 1 "0142 10 1-050 iS I 1 106 15 1-089 15 I -0214 15 I '°75 20 I '146 20 I'I20 20 I "0284 20 I'lOI 2 5 1-184 2 5 i'iS3 2 5 1 -0350 25 1 -126 3° 1*224 3° 1-185 3° I "0412 3° 1-151 35 I '266 35 1-218 35 I '0470 35 1 '177 40 I "306 40 1-251 40 I '0523 40 1 -200 45 1-051 45 1-284 45 1-0571 42-4 1 '210 5° 1 '399 5° I '319 So I '0615 55 1-448 55 I '346 55 1-0653 60 1-502 60 1 '374 60 1-0685 65 1 '557 65 1 '395 65 1*0712 70 1 -615 70 i'4 2 3 70 I '°733 75 1 '675 75 1-442 75 I "0746 80 1-733 80 1-460 80 I "0748 85 1783 85 1-478 85 1-0739 90 1-819 90 i'49S 90 1-0713 95 1-839 95 i'5 r 3 95 I "0660 100 1-838 100 1 '53° 100 1 '0553 590 Text-Book of Inorganic Chemistry, VIII. — Specific Gravity and Composition of Alkaline Solutions. Causti Potash Caustic Soda Ammonia Per cent. KOH Specific Per cent. NaOH Specific Per cent. NH, Specific by weight gravity by weight gravity by weight 5 1*041 5 1-058 5 0-979 10 1-083 10 1 ■Hi-' 10 °'959 15 1-128 IS 1-170 15 0-941 20 I -177 20 1-225 20 0-925 2S 1-230 25 1-279 25 0911 3^ 1-288 30 I- 332 3° 0-898 35 1 '349 35 I-384 33 0-891 40 1-412 40 1 '437 36 0-884 45 1 '475 45 1-488 ' 5° 1 '539 50 1 '54° 55 1-604 55 1 '591 60 1-667 60 1 '643 65 1729 65 1 -695 70 1-790 70 1748 To convert a density expressed in Twaddell's degrees into true specific gravity : — Divide by 200 and add 1. Thus, 25 Twaddell = 3 i'j ll 5 + 1 = IT25. IX. — Heat of Formation and Solution of some Important Compounds. {After Thomsen.) Note. — The quantity of heat (H) is expressed in the number of kilogrammes of water which would be raised 1° C. in temperature on the formation of the molecular weight of the compound in grammes. Thus : — (H 2 , O) = + 68-4 means that on the formation of 18 grammes of water, or on the combustion of 2 grammes of hydrogen, the quantity of heat set free would raise 68-4 kilos, of water i°. If heat is absorbed, instead of liberated, the number is expressed as a negative quantity. Compound H Remarks (H 2 ,0) . + 68-4 Liquid water (H 2 ,<3 3 , aq.) + 4-T3 Formation of H.,0 2 in dilute solution (H»,S) + 4 '7 Gaseous HqS (H-'S, aq.) + 4-6 Solution of H 2 S in water (H,C1) + 22 "o Gaseous HC1 (HC1, aq.) + 17 '3 Solution of HC1 in water (H.Br) . ¥ 8-4 Gaseous HBr (HBr, aq.) + 19-9 Solution of HBr in water (H,I) — 6-o Gaseous HI HI, aa.) . + IQ'2 Solution of HI in water (H,C1,0", aq.) . + 23 '9 Formation of HCIO5 in dilute solution (H,Br,0 5 , aq.). + 12 "4 HBr0 5 ,, I Appendix. S9i I X. — Continued. Compound (H,I,0 5 , aq.) (S.O 2 ) (SO 2 , aq.). (S.O 5 ) (SO 3 , aq.). (H 2 ,S,0<) . (H 2 SO", aq.) (N,H 5 ) . NHS, aq .) N,H*,Clj (N 2 ,0) . (N.O) (H.N.O 2 , aq.) (N.O 2 ) (H.N.O 3 ) . (HNO 3 , aq.) (P 2 ,0 5 ) (H\P,0 4 ). (P,C1 5 ) . (P,C1 3 ) (As.CF) Sb.CP) (CO) (C.O 2 ) (C,W) . (C 2 ,H*) (C 2 ,H 2 ) (C 2 N 2 ) . (K,C1 aq.) (K,Br, aq.) (K,I, aq.) (K.CI.O 5 ). (K f SO") . (K?S0 4 ,H : SOj K.N.O 5 ) . (KNO„ aq.) (Na.Cl, aq.) (Ca.Cl 2 , aq. ) (Ca.C.O 5 ). (Ba.Cl 2 , aq. ) (Ba.SO 4 ) . (A1-'C1 B ) . (A1 ? C1 6 , aq.) (Fe 2 ,Cl 6 . aq.) (Fe.Cl 2 , aq.) (Zn,S,0\ aq.) (Cu.O) . (Cu,S,0 4 ). Cu.S.O', aq.) (Ag.Cl) . (Ag.Br) (Ag.D , , (AgN.O 3 , aq.) H Remarks + S5'8 Formation of HIO5 in dilute solution + 71*1 Gaseous SO9 + 7-7 Solution of S0 2 in water + 103 "2 Liquid SO, + 39-2 Solution of SO, in much water + 192*9 Liquid H,S0 4 + 17-9 Solution of H 2 S0 4 in much water + z 1 '9 Gaseous NH, + 8-4 Solution of N H, in water + 75-8 Solid NHjCl - iys Gaseous N,0 — 21'6 Gaseous NO + 3° '8 Formation of HN0 2 in dilute solution — 2'0 Gaseous NO^ + 41-6 Liquid HNO, + TS Solution of HNO, in water + 369-9 Solid P 2 5 + 302*6 Crystalline H,P0 4 + 105-0 Solid PCI, + 75 '3 Liquid PCI, + 7' '4 Liquid AsCl 3 + 914 Liquid SbCl, + 29-0 Gaseous CO from amorphous C + 97-0 co 2 + •21 -8 Gaseous CH 4 ' - 2-7 C„H 4 - 48-2 G>H„ - 657 „ (CN f ; + IOI '2 Formation of KC1 in dilute solution + 9C2 KBr + 75 "0 KI + 95 '9 Crystalline KCIO, + 344' 6 K,SOj + 166 Formation of KHS0 4 from K„S0 4 and H„S0 4 + "9'5 Crystalline KNO s - 8-5 Solution of KNO5 in water + 9 6 'S Formation of NaCl in dilute solution + 187-2 CaCL, + 270*4 CaCO, as calcite + 196-8 Formation of BaCl 2 in dilute solution + 338-I BaS0 4 as an amorphous powder + 322-0 Anhydrous A1 2 C1 6 (solid) + '537 Solution of A1 2 C1 6 in water + 2 5S'4 Formation of Fe-Cl 6 in dilute solution + IOO'O FeClo + 248-5 ZnS0 4 + 37-2 Solid CuO + 182-6 Anhydrous CuS0 4 (solid) + 198-4 Formation of CuS0 4 in dilute solution + 29-4 Solid AgCl + 22 '7 „ AgBr + 13-8 „ Agl + 233 Formation of AgNO, in dilute solution INDEX. AC I Acms. 56 Acid hydrates, 56 Affinity, 45, 51 Agate, 269 Air, 200 Air bath, 89 Alabandite, 463 Alabaster, 394 Albite, 365, 425 Albumen, 140 Algaroth powder, 257 Alkali manufacture, 370 — waste, 169, 372 Alloys, 58 Alum, 423 — cake, 423 — shale, 424 Alumina, 421 Aluminates, 422 Aluminic acid, 422 Aluminite, 423 Aluminium, 419 — bronze, 531 — chloride, 428 — detection of, 429 — fluoride, 428 — gold, S3- — hydrate, 421 — oxide, 421 — phosphate, 425 — salts of, 422 — silicate, 425 — sulphate, 422 — • sulphide, 428 Alunite, 140, 424 Amalgamation process, 543 ANT Amalgams, 58, 541 Amethyst, 269 Amidogen, 183, 188 Ammonia, 183 — soda process, 372 Ammonium, 183, 188, 382 — bromide, 386 — carbonate, 383 — chloride, 385 — chlorplatinate, 562 — chromate, 477 — detection of, 387 — dichromate, 477 — ferrocyanide, 456 — nitrate, 383 — phosphates, 383 — sulphate, 382 — sulphide, 386 — sulphydrate, 386 — sulphostannate, 520 Amorphous carbon, 282 Analysis, 4 Anatase, 322 Analcime, 425 Andalusite, 425 Anglesite, 498, 503 Anhydrides, 56 Anhydrite, 394 Animal charcoal, 283 Anorthite, 425 Anthracite, 284 Antichlor, 103, 369 Antimonic acid, 254 — anhydride, 254 — chloride, 257 Antimonite, 250, 258 Q Q 594 Index. Aritimoniuretted hydrogen, 252 Antimonous chloride, 256 — sulphate, 254 — sulphide, 258 Antimony, 250 — compounds of, with oxygen, 25 3 the halogens, 256 sulphur, 258 — glass, 259 — pentachloride, 257 — pentasulphide, 259 — pentoxide, 254 — sulphochloride, 257 — tetroxide, 256 — tribromide, 257 — trichloride, 256 — triiodide, 258 — trioxide, 253 — trisulphide, 258 Apatite, 206, 396 Appendix, 573 Aquamarine, 429 Aqua regia, 198 Argentan, 492 Argillaceous iron ore, 436 Arragonite, 286, 399 Arsenic, 229 — anhydride and acid, 244 Arsenical pyrites, 229, 437 Arsenic bloom, 233 — compounds of, with oxygen, 233 the halogens, 246 sulphur, 248 — detection of, 237 — disulphide, 248 — pentasulphide, 250 — trisulphide, 248 Arsenious acid, 235 — anhydride, 233 — bromide, 247 — chloride, 247 — iodide, 247 — fluoride, 248 — sulphide, 248 Arseniuretted hydrogen, 231 Arsine, 231 Asbestos, 413, 417 — platinized, 162 Atacamite, 528 Atmospheric air, 200 Atomic heat, 579 — weights, table of, 60 — theory, 44 — weights, determination of, 573 relations between, 65, 582 Atomicity, 47 Atoms, 45 Augite, 413, 417 Auric acid, 554 — chloride, 553 — cyanide, 554 — oxide and hydrate, 554 — sulphide, 555 Aurous chloride, 555 — cyanide, 555 — oxide, 555 Avogadro's law, 50, 575 Azurite, 523, 528 Barium, 408 — carbonate, 41 1 — chloride, 41 1 — chromate, 477 — detection of, 412 — hydrate, 409 — nitrate, 41 1 — oxide, 40S — peroxide, 410 — sulphate 411 — sulphide, 412 Baryta, 408 Baryta-water, 409 Basic hydrates, 56 Bay salt, 377 Bell -metal, 531 Beryl, 429 Beryllium, 429 — carbonate, 430 — chloride, 430 — hydrate, 430 — oxide, 430 — sulphate, 430 Bessemer process, 441 Biotite, 413 Bismuth, 513 — chloride, 515 — chromate, 5 T 5 — detection of, 515 — nitrate, 515 Index. ■595 BIS Bismuth oxide and hydrate, 514 — suboxide, 514 — sulphide, 515 Bismuthic acid, 514 Bismuthinite, 515 Black ash, 371 — lead, 281 Blast furnace, 437 Bleaching powder, 102, 116, 400 Blende, 140, 493, 496 Blue vitriol, 525 Bog iron ore, 436 Bole, 427 Bone ash, 396 — black, 283 Boracite, 260, 417 Borax, 376 Boric acid and anhydride, 262 Boron, 260 — nitride, 266 — sulphide, 266 — tribromide, 265 — trichloride, 264 — trifluoride, 265 Boronatro-calcite, 260 Boulangerite, 507 Brass, 530 Braunite, 460, 463 Breithauptite, 250 Britannia metal, 252, 521 Bromic acid, 122 Bromide of lime, 123 Bromides, 122 Bromine, 118 Bronze, 531 Brookite, 322 Brown coal, 285 Brucite, 415 Burnt alum, 423 Bunsen burner, 309 — pump, 33 Cadmium, 497 — - compounds of, 497 — detection of, 498 Csesium, 380 — compounds of, 381 Calamine, 493, 495 ■Calcareous tufa, 400 CHA Calcite, 286, 399 Calcium, 388 — carbonate, 398 — chloride, 404 — chromate, 477 — detection of, 406 — diacid phosphate, 39S — fluoride, 405 — hydrate, 390 — monacid phosphate, 398 — nitrate, 396 — oxalate, 404 — oxide, 389 — pentasulphide, 405 — peroxide, 393 - — phosphate, 396 — silicate, 401 — sulphate, 394 — sulphide, 405 — sulphydrate, 405 Calomel, 540 Calcspar, see Calcite Cannel coal, 285 Caput mortuum, 163, 447 Carbamic acid, 293 Carbon, 277 — disulphide, 300 — compounds of, 286 with hydrogen, 303 — hexachloride, 310 — oxychloride, 297 — oxysulphide, 302 — tetrabromide, 310 — tetrachloride, 309 Carbonic acid or anhydride, 286 — oxide, 294 Carbonyl chloride, 297 Carnallite, 418 Cassel yellow, 506 Cassiterite, 516, 519 Cast-iron, 438 Caustic potash, 346 — soda, 367 Celestine, 406 Cement copper, 526 — steel, 441 Cerite, 433 Cerium, 433 Cerussite, 498, 504 Chalcedony, 269 Q Q2 596 Index. CHA Chalcostibite, 250 Chalk, 286, 400 Chalybeate springs, 446 Chameleon mineral, 466 Charcoal, 282 Chemical affinity, 4, 5 1 — calculations, 61 — combination, laws of, 37 — elements, remarks on, 63 — equivalents, 42 — nomenclature, 54 — processes, 3 — properties, 77 — symbols and formula?, 58 Chili saltpetre, 369 China-clay, 426 Chlorchromic acid, 477 Chloric acid, no Chloride of lime, 102, 116, 400 Chlorides, 109 Chlorine, 97 -- hydrate, 103 — oxides and oxy-acids, 109 — peroxide, 114 Chlorous anhydride and acid, 1 14 Chlorplatinic acid, 561 Chlorsulphonic acid, 165 Chromates, 474 Chrome alum, 425, 471 — iron ore, 468 -— red, 506 — yellow, 506 Chromic anhydride and acid, 473 — chloride, 472 — hydrate, 470 — nitrate, 472 — oxide, 469 — phosphate, 472 — ; sulphate, 47 1 Chromites, 470 Chromium, 468 — detection of, 479 Chromous chloride, 473 — hydrate, 473 Chromyl chloride, 478 Chrysoberyl, 429 Cinnabar, 533, 538 Clay, 426 Clay ironstone, 436, 446 Coal, 284 CUP Coal gas, 319 Cobalt and nickel, 483 Cobalt, 484 — arsenides, 489 — bloom, 230 — detection of, 489 Cobaltic chloride, 486 ammoniacal compounds of, 487 — oxide and hydrate, 485 Cobaltous carbonate, 486 — chloride, 487 — cyanide, 488 — nitrate, 486 — nitrite, 486 — oxalate, 487 — oxide and hydrate, 485 — phosphate, 486 — sulphate, 486 — sulphide, 488 Coke, 284 Colcothar, 163, 447 Cold-short, 440 Columbite, 330 Common salt, 365, 376 Combustion, 29 Compound radical, 61 Condy's fluid, 467 Contact action, 162 Copper, 523 — alloys of, 530 — detection of, 531 Copper-glance, 523, 530 Copper-pyrites, 140, 437, 523, 530 Coprolites, 206, 396 Corrosive sublimate, 536 Corundum, 421 Couple, galvanic, 518 Covellite, 529 Crocoisite, 506 Crocus antimonii, 259 Cryolite, 134, 365, 420, 428 Crystals, how obtained, 75 Crystalline systems, 71 Crystallography, 70 Cubic alum, 423 — nitre, 369 Cupellation, 499, 551 Cupric acetates, 528 — ammonium sulphate, 526 Index. 597 Cupric arsenite, 527 — bromide, 528 — carbonates, 527 — chloride, 528 — - compounds, 524 — hydrate, 525 — nitrate, 527 — oxide, 524 — phosphate, 527 — sulphate, 525 — sulphide, 528 Cuprite, 523, 529 Cuprous chloride, 529 — compounds, 529 — hydrate, 529 — iodide, 530 — oxide, 529 — sulphide, 530 Cyamelide, 317 Cyanic acid, 316 Cyanogen, 311 — bromide, 319 — chlorides, 318 — iodide, 319 Cyanuric acid, 317 Dulong and Petit's law, 70, 578 Dutch metal, 531 Dyad, 47 Dyscrasite, 250 Earthenware, 426 Eaux de Javelle, 118, 355 Electro-chemical theory, 46 Electrodes, 6 Electro-gilding, 556 Electrolysis, 6 Electro-plating 550 Elements, 5 — general remarks on, 63 Emerald, 429 Emery, 421 Epsom salt, 140, 415 Equivalent weights, 42 Erbium, 433 Etching on glass, 136 Ethylene, 306 — dibromide, 307 — dichloride, 307 Eudiometer, 25 Dalton's atomic theory, 45 Dark-red silver ore, 250, 543, 551 Datolite, 260 Decrepitation, 87 Deliquescence, 85 Density of gases, 5 1 Desiccator, 87 Dialysed iron, 448 Diamond, 279 Diaspore, 422 , Dibasic acids, 56 Dichromic acid, 475 Didymium, 433 Dimorphism, 75 Displaceable hydrogen, 56 Diselenium dichloride, 179 Disthene, 425 Disulphur dichloride, 173 Disulphuric acid, 166 Dithionic acid, 1 70 Divalent elements, 47 Dolomite, 286, 413, 417 Drummond's light, 25 Fahlerz, 523 Fayence, 426 Felspar, 420, 425 Ferric acid, 450 — chloride, 451 — disulphide, 452 — hydrate, 447 — nitrate, 449 — oxide, 447 — phosphate, 449 — sulphate, 448 Ferromanganese, 439 Ferrous and ferric compounds, 444 — carbonate, 446 — chloride, 450 — hydrate, 445 — iodide, 45 1 — oxide, 445 — phosphate, 447 — sulphate, 445 — sulphide, 452 Fining, 439 Flame, structure of, 308 598 Inaex. FLI Flint, 269 Flowers, of sulphur, 142 Fluoboric acid, 265 Fluorine, 133 Fluor-spar, 134, 388, 405 Fluosilicic acid, 275 Fluotitanic acid, 325 Fuller's earth, 427 Fulminating gold, 554 — silver, 546 Fuming sulphuric acid, 163 Fur of kettles, 83 Fusible metal, 522 Gadolinite, 433 Galena, 140, 498, 507 Gahnite, 422 Gallium, 431 Galvanic couple, 518 Galvanized iron, 443, 494 Gasometer, 15 Gay-Lussac's law, 43 German silver, 492, 531 Gibbsite, 421 Glance cobalt, 484, 489 Glass, 401 Glauber's salt, 140, 367 Glucinum, see Beryllium Gold, 552 — alloys of, 556 — compounds of, 553 Grain-tin, 517 Granite, 269, 420 Graphic tellurium, 552 Graphite, 281 Graphitic acid, 282 Green vitriol, 445 Greenockite, 497 Guignet's green, 471 Gun-metal, 531 Gunpowder, 351 Gypsum, 140, 388, 394 Haematite, 436, 447 Halogens, 96 Haloid-acids, 57 Haloid salts, 57, 96 Hard water, 83 TRO Hartshorn, spirits of, 183 Hauerite, 463 Hausmannite, 460, 464 Heavy carburetted hydrogen," 304- Heavy metals, 436 Heavy spar, 140, 408 Hexads, 47 Hornblende, 413, 417 Horn-silver, 543 Hyacinth, 434 Hydrargillite, 421 Hydrates, 56, 87 Hydraulic mortar, 393 Hydriodic acid, 127 Hydrobromic acid, 120 Hydrochloric acid, 104 Hydrocobalticyanic acid, 488: Hydrocyanic acid, 313 Hydroferricyanic acid, 458 Hydroferrocyanic acid, 457 Hydrofluoric acid, 134 Hydrogen, 16 Hydrogen peroxide, 90 — persulphide, 148 Hydrosulphurous acid, 155 Hydroxylamine, 186 Hypobromous acid, 123 Hypochlorous anhydride and acid,, 116 Hypophosphoric acid, 220 Hypophosphorous acid, 222 Hyposulphuric acid, 170 Hyposulphurous acid, 167 Ice, 80, Si Iceland-spar, 399 Imidogen, 183, 188 Indium, 432 — compounds of, 432 Infusorial earth, 269 Iodic acid, 129 — anhydride, 130 Iodides, 129 Iodine, 123 — chlorides, 133 — oxy-acids, 129 Iridium, 568 — - chloride, 569 Iron, 436 Index. 599 Iron alum, 425,- 448 — compounds, 444 — compounds of, with cyanogen, ■453 — detection of, 459 — pyrites, 140, 437, 452 — rust, 442 — spinelle, 422 Isomeric compounds, 38 Isomorphism, 75 Jamesonite, 507 Jasper, 269 Jet, 285 Kaolin, 426 Kermes mineral, 259 Kieselguhr, 269 Kieserite, 415 Kupfernickel, 229, 490 Lakes, 422 Lampblack, 285 Lanthanum, 433 Lapis lazuli, 427 Lavoisier, 30, 31 Law, Avogadro's, 50, 575 — Gay-Lussac's, 43 — Dulong and Petit's, 70, 578 — of isomorphism, 75, 580 Lead, 498 — acetate, 505 — alloys of, 507 — carbonates, 5°4 — chloride, 5°6 — chromate, 5°6 — cyanide, 507 — detection of, 508 — hydrate, S 02 — iodide, 5°6 — nitrate, 504 — oxide, 5°l — peroxide, 502 — phosphate, 504 — salts of, 503 — suboxide, 501 — silicates, 504 MAG Lead sulphate, 503 — sulphide, 507 Lepidolite, 379 Light, influence of, on chemical affinity, 53 — carburetted hydrogen, 304 — red silver ore, 543, 551 Lignite, 285 Lime, 389 — chloride of, 102, 1 16, 400 — felspar, 425 — light, 25 — milk of, 390 — water, 390 Limestone, 286, 399 Limonite, 436 Linnteite, 489 Liquor ammonite, 183 Litharge, 501 Lithium, 378 — carbonate, 380 — chloride, 380 — detection of, 380 — oxide and hydrate, 379 — phosphate, 379 — sulphate, 379 Lithographic stone, 399 Liver of sulphur, 363 Lunar caustic, 547 Luteocobaltic chloride, 487 Magnesia, 414 — alba, 417 — calcined, 415 — usta, 415 Magnesite, 286, 413, 417 Magnesium, 413 — arsenate, 416 — borate, 417 — carbonate, 417 — chloride, 417 — detection of, 418 — hydrate, 415 — oxide, 414 — phosphates, 416 — silicate, 417 — sulphate, 415 Magnetic oxide of iron, 436, 449 — pyrites, 437, 453 6oo Index. Malachite, 523, 528 Manganese, 460 — alum, 425, 463 ■ — detection of, 467 — disulphide, 463 — oxides, 461 — perhydrate, 465 — peroxide, 464 — spar, 460, 462 Manganic acid, 465 — chloride, 464 — hydrate, 463 — oxide, 463 — sulphate, 463 Manganite, 460, 463 Manganous carbonate, 462 — chloride, 462 — nitrate, 462 — oxide and hydrate, 46 1 — sulphate, 461 — sulphide, 463 Marble, 286, 399 Marcasite, 453 Marking-ink, 547 Marl, 427 Marsh gas, 304 Marsh's test, 239 Mass, influence of, on chemical affinity, 53 Massicot, 501 Matches, lucifer, 210 Matlockite, 506 Meerschaum, 413, 417 Mercuric chloride, 536 — chromate, 536 — compounds, 535 — cyanide, 538 — iodide, 538 — nitrate, 536 — oxide, 535 — sulphate, 535 — sulphide, 538 Mercurous chloride, 540 — chromate, 540 — ■ compounds, 539 — iodide, 541 — nitrate, 539 — oxide, 539 — sulphide, 541 Mercury, 533 NIC Mercury, detection of, 542 Metaboric acid, 263 Metals, 331 — classification of, 340 — of the alkalies, 342 — of the alkaline earths, 388 — of the earths, 419 — heavy, 436 — noble, 533 — platinum, 557 Metantimonic acid, 255 Metaphosphoric acid, 220 Metastannic acid, 519 Methane, 304 Mica, 420 A'licrocosmic salt, 3S3 Milk of lime, 390 — sulphur, 145 Millerite, 492 Mimetesite, 230 Mineral waters, 82 Minium, 502 Mispickel, 229 Molecular weights, how found, 51, 573 Molecule, 48 Molybdenite, 325 Molybdenum, 325 — - chlorides, 327 — oxides, 326 Molybdic anhydride and acid, 326 Monads, 47 Monobasic acids, 56 Monovalent elements, 47 Mortar, 391 — hydraulic, 393 Mosaic gold, 521 Mottramite, 329 Multiple proportions, law of, 42 Muntz-metal, 531 Nascent state, 50 Native; 63 Natrolite, 425 Nessler's solution, 387, 538 Niccolite, 490 Nickel, 490 — alloys of, 492 — carbonate, 491 Index. 60 1 Nickel chloride, 491 — - cyanide, 491 — detection of, 492 — nitrate, 491 — sulphate, 491 — - sulphide, 492 — silver, 492 Nickelic oxide and hydrate, 49 s Nickelous oxide and hydrate, 490 Niobium, 330 Nitrated acid, 158 Nitrates, 192 Nitre, 350 Nitric acid, 189 — anhydride, 192 — oxide, 192 — peroxide, 196 Nitrites, 195 Nitrogen, 181 — compounds of, with oxygen, i8< — chloride, 204 — iodide, 206 Nitrosulphonic acid, 166, 197 Nitrosyl chloride, 198 Nitrous anhydride and acid, 195 — oxide, 198 Nitryl chloride, 197 Noble metals, 533 Nomenclature, chemical, 54 Nordhausen sulphuric acid, 163 Olefiant gas, 306 Olivine, 271, 413, 417 Opal, 269 Orangite, 434 Orpiment, 229, 24S Orthoclase, 272, 425 Orthophosphoric acid, 216 Osmic acid, 57° Osmiridlum, 558, 569 Osmium, 570 Oxalic acid, 297 Oxamic acid, 299 Oxamide, 300 Oxidation, 36 Oxides, 55 Oxy-acids, 56 Oxyammonia, 186 Oxyammonium chloride, 387 Oxyammonium sulphate, 387 — salts, 386 Oxychlorides, 58 Oxygen, 8 Oxy-hydrogen blowpipe, ?\ Ozone, 91 Palladic chloride, 568 — oxide, 568 Palladium, 566 — hydride, 566 Palladous chloride, 56S — iodide, 568 — nitrate, 567 — oxide, 568 Paracyanogen, 312 Parkes's process, 500 Passive iron, 443 Pattinson's process, 500 Pearlash, 353 Peat, 285 Pentads, 47 Pentavalent elements, 47 Pentathionic acid, 172 Perchloric acid, 113 Perchromic acid, 474 Periodic acid, 131 — law, 583 Permanganic anhydride and acid, 466 Peroxides, 55 Persulphocyanic acid, 318 Petallite, 379 Phenacolite, 230 Phenakite, 429 Phlogiston theory, 31 Phosgene gas, 297 Phosphates, 217 Phosphine, 211 Phosphonium iodide, 128, 213 Phosphoretted hydrogen, 211 Phosphoric acid, 216 — anhydride, 214 — bromide, 227 — chloride, 225 — oxybromide, 227 — oxychloride, 226 — sulphide, 229 Phosphorite, 206, 396 602 Index. PHO Phosphorous anhydride and acid, 221 — bromide, 227 — chloride, 224 Phosphomolybdic acid, 327 Phosphotungstic acid, 328 Phosphorus, 206 — compounds of, with hydrogen, 211 oxygen, 213 the halogens, 223 sulphur, 228 — diiodide, 228 — disulphide, 229 — pentabromide, 227 — pentachloride, 225 — pentasulphide, 229 — tribromide, 227 — trichloride, 224 — triiodide, 228 Phosphoryl chloride, 226 Photography, 549 Physical processes, 3 — properties, 67 Pink salt, 520 Pitchblende, 479 Plaster of Paris, 394 Platinic chloride, 561 — oxide and hydrate, 562 — sulphide, 562 Platinous chloride, 562 — cyanide, 563 — oxide and hydrate, 563 — sulphide, 563 Platinum, 558 — ammoniacal compounds of, 563 — black, 559 — metal--, 557 — spongy, 558 Plumbago, 281 Poles or electrodes, 6 Polythionic acids, 171 Porcelain, 426 Portland cement, 393 Potash, 353 — alum, 423 — felspar, 425 Potassium, 342 — acid carbonate, 354 — acid fluoride, 358 PRO Potassium acid sulphate, 349. — aluminate, 422 — aurate, 554 — bromide, 356 — carbonate, 353 — chlorate, 354 — chlorchromate, 477 — chloride, 356 — chlorpalladinate, 568 - chlorplatinate, 364, 56 1 — chromate, 476 — cobalticyanide, 488 — compounds of, 346 — cyanate, 361 — cyanide, 359 — detection of, 364 — dichromate, 475 — ferrate, 450 — ferricyanide, 457 — ferrocyanide, 453 — fluoborate, 358 — fluoride, 358 — fluosilicate, 359 — hydrate, 346 — hypochlorite, 354 — iodide, 356 — iridium chloride, 559 — manganate, 465 — monosulphide, 362 — nitrate, 350 — nitrite, 352 — osmate, 57 1 — oxalate, 355 — oxide, 346 — perchlorate, 354 — permanganate, 466 — peroxide, 346 — plumbate, 507 — polysulphides, 363 — silicate, 355 — stannate, 520 — sulphate, 349 — sulphocyanate, 361 — sulphydrate, 363 — uranate, 482 Prehnite, 272 Priestley, 8 Proportions, combining, 3S Proustite, 543 Front's hypothesis, 582 Index. 603 Prussian blue, 456 . Prussic acid, 313 Psilomelane, 460 Puddling, 439 Purple copper ore, 523, 530 — of Cassius, 553, 556 Pyrargyrite, 250, 543 Pyrites, copper, 140, 523, 530 — iron, 140, 437, 452 Pyroboric acid, 263 Pyrolusite, 460, 464 Pyromorphite, 504 Pyrophosphoric acid, 219 Quantity (mass), influence of, on chemical affinity, 53 Quartz, 268 Quicklime, 389 Quicksilver, see Mercury Radical, compound, 61 Rain water, 83 Realgar, 229, 248 Red arsenic glass, 248 — copper ore, 523, 529 — lead, 502 — prussiate, 457 — short, 440 Reduced iron, 442 Reduction, 36 Rhodium, 570 Rock salt, 365, 376 Roman alum, 424 — cement, 393 Roseocobaltic chloride, 487 Rouge, 447 Rubidium, 380 — compounds, 381 Ruby, 421 — sulphur, 248 Ruthenium, 571 Rutile, 322 Sal-ammoniac, 385 Salt-cake, 371 Saltpetre, 350 Salts, constitution of, 332 Salts of sorrel, 355 Samarskite, 330, 433 Sandstone, 269 Sassolite, 260 Scheele, 8, 31, 97 Scheele's green, 235, 527 Scheelite, 327 Schlippe's salt, 259 Schiitzenberger's acid, 155 Schweinfurth green, 235, 527 Selenic acid, 178 Selenious anhydride and acid, 177 Selenite, 394 Selenium, 174 — chlorides, 179 — compounds, 175 — disulphide, 178 Senarmontite, 253 Serpentine, 271, 413, 417 Sesqui-oxides, 55 Siderite, 436, 446 Siemens-Martin process, 44 1 Silica, 268 Siliceous calamine, 493 — sinter, 269 Silicic acid, 271 Silico-chloroform, 276 Silico-formic anhydride, 277 Silicon, 266 — compounds, 268 — fluoride, 274 — hexachloride, 274 — hexiodide, 274 — hydride, 272 — tetrabromide, 274 — tetrachloride', 273 — tetriodide, 274 Silicotungstic acid, 328 Silicoxalic acid, 277 Silver, 543 — alloys of, 551 — arsenate, 547 — arsenite, 547 — bromide, 548 — carbonate, 547 — chloride, 548 — chromate, 548 — compounds, 545 — copper-glance, 543 — cyanide, 550 604 Index Silver, detection of, 551 — fluoride, 550 Silver-glance, 543, 551 Silver iodide, 548 — nitrate, 546 — oxide, 545 — peroxide, 546 — phosphate, 547 — sulphate, 546 — sulphide, 55° Skulterudite, 489 Slaked lime, 390 Smalt, 489 Smaltite, 229, 4S4 Smithsonite, 493 Soapstone, 413 Soda, 369 — felspar, 425 ■ — sesquicarbonate of, 374 Sodium, 365 — acid carbonate, 374 — acid sulphate, 36S — acid sulphite, 368 — aluminate, 422 — arsenate, 375 — auric chloride, 553 — borate, 376 — carbonate, 369 — chloride, 376 — chlorplatinate, 562 — chromate, 477 — detection of, 378 — diacid phosphate, 375 — ferrocyanide, 456 — hydrate, 367 — metaphosphate, 375 — metastannate, 519 — monacid phosphate, 374 — nitrate, 369 — nitroprusside, 458 — oxide, 367 — peroxide, 367 — phosphates, 375 — pyrophosphate, 375 — stannate, 519 — sulphate, 367 — sulphite, 368 — thiosulphate, 368 Solder, 521 Solution, 76 SUL Solution, supersaturated, 84 Sombrerite, 396 Spathic iron ore, 286, 436, 446 Spectrum analysis, 337 Specular iron, 436, 447 Speculum metal, 531 Speiss-cobalt, 484, 489 Sphene, 323 Spiegeleisen, 439 Spinelle, 422 Spongy platinum, 558 Stahl's phlogiston theory, 31 Stannic acid, 520 — chloride,, 520 — compounds, 519 — hydrate, 519 — oxide, 519 — sulphide, 520 Stannous chloride, 518 — compounds, 517 — oxide and hydrate, 517 — sulphate, 517 — sulphide, 518 Steatite, 413, 417 Steel, 440 Stibnite, 258 Stoneware, 426 Strontianite, 286, 406 Strontium, 406 — carbonate, 407 — chloride, 407 — chromate, 477 — detection of, 408 — hydrate, 407 — nitrate, 407 — oxide, 406 — peroxide, 407 — sulphate, 407 — sulphide, 407 Struvite, 416 Suboxides, 55 Sugar of lead, 505 Sulphantimonic acid, 259 Sulphites, 154 Sulpho-acids, -bases, and -salts, 57 Sulphocarbonic acid, 302 Sulphocyanic acid, 317 Sulphur, 139 — compounds of, with oxygen, 148 — dichloride, 173 Index. 605. SUL Sulphur diselenide, 178 — flowers of, 142 — group of elements, 137 — hexiodide, 174 — milk of, 145 — tetrachloride, 174 Sulphuric acid, 155 — fuming or Nordhausen, 163 Sulphuretted hydrogen, 145 Sulphurous acid, 154 — anhydride, 150 Sulphuryl, 152 — chloride, 164 Superphosphate of lime, 209, 398 Supersaturated solutions, 84 Sylvanite, 552 Tables, 5S6 Talc, 413, 417 Tantalite, 330 Tantalum, 330 Tartar emetic, 254 Tellurium, 179 Temperature, influence of, on che- mical affinity, 52 Terbium, 433 Terne-plate, 443 Tetrads, 47 Tetradymite, 179 Tetrahedrite, 523 Tetrathionic acid, 172 Tetravalent elements, 47 Thallic compounds, 512 Thallium, 509 — compounds, 511 — detection of, 512 Thallous compounds, 511 Theory, atomic, 44 — electro-chemical, 46 •— phlogiston, 31 Thiosulphuric acid, 167 'thorium, 434 Tin, 516 — alloys of, 521 — compounds, 517 — detection of, 522 Tin-foil, 521 Tin-plate, 443, 521 Tin-salt, 518 VAN Tin-stone, 516, 519 Tinkal, 260, 376 Titanic anhydride and acid, 323 — iron, 3 2 3 Titanite, 323 Titanium, 322 — nitride, 325 — sesquioxide, 324 — tetrachloride, 324 — tetrafluoride, 325 Tombac, 531 Triads, 47 Tribasic acids, 56 Tridymite, 269 Triferric tetroxide, 449 Trimanganic tetroxide, 464 Triphylline, 379 Trithionic acid, 172 Trivalent elements, 47 Trona, 286, 374 Tungsten, 327 — chlorides, 329 — steel, 329 Tungstic anhydride and acid, 32S Turnbull's blue, 458 Turpeth mineral, 535 Type metal, 252, 507 Ultramarine, 427 Uranates, 482 Uranic carbonate, 481 — nitrate, 480 — oxide, 480 — phosphate, 481 — sulphate, 480 Uranium, 479 Uranoso-uranic oxide, 482 Uranous chloride, 482 — oxide, 482 — sulphate, 482 Uranyl, 481 — chloride, 481 — sulphide, 482 Valency, 47 Valentinite, 253 Vanadinite, 329 Vanadium, 329 €o6 Index. Verdigris, 524, 528 Vermilion, 539 Vivianite, 447 Volumes, law of, 43 Yellow prussiate, 453 Ytterbium, 433 Yttrium, 433 Yttrotantalite, 433 Wad, 460 Water, 80 — hard, 83 — of constitution, 86 — of crystallization, 85 Waters, mineral, 82 Wavellite, 425 Weights, atomic, 60 — equivalent, 42 White arsenic, 230, 233 — lead, 504 — precipitate, 537 — vitriol, 495 Willemite, 493 Witherite, 286, 408 Wolfram, 327 Woulff's bottle, 19 Wrought iron, 440 Wulfenite, 325 Zaffre, 489 Zeolites, 425 Zinc, 493 — blende, 493, 496 — carbonate, 495 — chloride, 495 — detection of, 496 — hydrate, 495 — nitrate, 495 — oxide, 494 — spar, 495 — sulphate, 495 — sulphide, 496 — white, 495 Zinkenite, 507 Zircon, 434 Zirconium, 434 — compounds, 435 LONDON : PRINTED CY SPOTTISWOODE AND CO., NEW-STREET SQUARE AND PARLIAMENT STREET TABLE OF SPECTRA.