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HE +B3，3||H.-E. ！¿? №|- ~~№ ae į Š! ± 2№ |-Sſ += ·<%£ € |ș . . . $3##w }Sſ º ;5%§ 5） }}2| №~--- ™Ni ŽE §}3= 8：:Æ ÈſË ----：í§§}#*- §zażjº/(\ſ|$）:3~^ ººº ！ſºZſ ！Ź，}， %ſ，%ſ,%.]）[\\]]{{[IIIjā № ae aerº， №ºaeaeaeaeaeae × ¶√∞ √≠ ≤ ∞∞∞, ∞; ∞， ∞; ∞, ∞; ∞, ∞; ∞， ∞， ∞; ∞, ∞; ∞, ∞,eae ，Ōe， aelº، ، ، ، ، ،∞ √≠ ≤ ∞， ∞， ∞; ∞， ∞; ∞， ∞， ∞; ∞,©=№ }ÎÏÏĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪĪiſſä ،№→ | �& *** .*---- {*s. • • • • • ►ººººº ！！ ！！ ！！ : « … • • • • • • • • • • • * Chemical Babraſy N. (U/ * E \ e o & v. Olt |-| ºr： |-| -|- |||||||||| |：||||||||||||||| .№ºt | | | | ≤ |ſ| - || ||||||||||||||||||| Oºt.CaeOº ||||||||| | | | | | | | 00：CET)OTT. |ſºſyº，| OLIIOGL05： | 1 || | || || || Cº | || 1 || ~ | | | | | |- C ANALYTICAL CHEMISTRY. BY F. P. TREADWELL, PH.D., Profes&or of Analytical Chemistry in the Polytechnic Institute of Zurich. vs.” A U T H O RIZ E D T R A N S L A TI O N IBY J WILLIAM T. HALL, S.B., Instructor in Chemistry, Massachusetts Institute of Technology. WOLUME I. QUALITATIVE ANALYSIS. SECOND ENGLISH FROM THE FOURTH GERMAN EDITION. THIRD THOUSAND. NEW YORK: JOHN WILEY & SONS. LONDON : CHAPMAN & HALL, LIMITED, 1908. Copyright, 1903, 1906, BY WILLIAM T. HALL. Entered at Stationers' Hall. Öſte §rientific #reas Rubert Bruuuuuun atti, Glomgang $ein ºurk PREFACE TO THE FIRST EDITION. HAVING been repeatedly requested by former pupils to publish the lectures on Analytical Chemistry given by me at this Institute since 1882, and not having time to do it myself, I permitted the “Verein des Polytechniker” to print in manuscript form the notes of one of my students. This output met with such a friendly reception that in 1888 a. second edition became necessary. Subsequently I decided to pub- lish the material in book form, with a radical rearrangement; so that this text-book of Analytical Chemistry represents a somewhat amplified repetition of my lectures. This little book is intended not only for laboratory use, but also for self-study. With each element the mineralogical occurrence, crystalline form, and isomorphous relations are briefly mentioned. Then, after explaining the reactions, I give the methods of separa- tion in the form of tables; because, contrary to the views of many, I have in this way obtained the best results in teaching. These tables are summarized charts with which the student can quickly find his bearings. - Much weight will be placed upon the determination of the sensi- tiveness of the single reactions, as is explained on page 34, because the beginner becomes in this way at once familiar with the solubility of the most important salts, and also with simple stoichiometrical calculations. Thus the approximate solubility of potassium chlor- platinate, for example, is found from the following determination of the sensitiveness of the reaction by which it is formed: If 100 ccm. of the Solution contain 0.156 gm. potassium, one finds that the formation of the chlorplatinate, at ordinary tempera- tures, only takes place on addition of a little alcohol; but on in- iii iv PREFACE TO THE FIRST EDITION. creasing slightly the amount of potassium in the Solution, it takes place immediately. We can, therefore, assume that the solution, which contains 0.156 gm. of potassium per 100 c.c. water, is satu- rated with chlorplatinate; hence the amount of the latter may be calculated: K. : K.PtCl–0.156:2; 78.3: 485.8=0.156: æ; a = 0.97. The result shows that 100 c.c. of water, at ordinary temperatures, dissolve 0.97gm. of K.PtCls, while accurate determinations at 20°C. have given the value 1.12. The difference, about 12 per cent., is explained by the facts that we did not work at exactly 20° C., nor with absolutely pure water; the Solution also contains an excess of chlorplatinic acid, whereby the solubility of the potassium chlor- platinate is diminished; but at the same time the values obtained in this way permit a very good comparison of the Solubilities of the different salts. From the sensitiveness of the reaction between a potassium salt and tartaric acid, the solubility of the potassium acid tartrate may be found to be 0.38; so that the Solubility of the potassium chlorplatinate is to that of the potassium acid tartrate as 0.97 : 0.38; the potassium tartrate is about three times as insoluble as the chlorplatinate, etc. The size of the book does not permit going into the microchemi- cal detection of the different elements. We have, however, in the excellent work of H. Behrens, “Anleitung Zur mikrochemischen Analyse,” a reference book of the highest rank. In publishing this, the first volume of the work, I beg of my colleagues and fellow chemists to kindly inform me of any errors or omissions. F. P. TREADWELL. ZURICH, April 29, 1899. PREFACE TO THE SECOND EDITION. THE material has been carefully revised, and much that is new has been added, but the arrangement of the subject-matter is the S3LT162. The request which I made of my colleagues and fellow chemists to inform me of errors has been in a large measure complied with, especially by Prof. E. Bamberger, Prof. E. Constam, Prof. W. Wis- licenus of Würzburg, and Dr. Schudl in Vienna, for which I wish in this place to express my heartiest thanks. To my assistant, Mr. O. Brunner, I am very grateful for zealous help in reading the proof. F. P. TREADwDLL. ZURICH, September, 1901. TRANSLATOR'S NOTE. THE translator wishes to express his sincere thanks to Professor Treadwell for his kindness and good wishes with regard to the preparation of this translation, to Professor A. A. Noyes, Professor Henry Fay, and Mr. H. S. Walker of the Massachusetts Institute of Technology, for their advice and criticism of the proof, and especially to Dr. W. O. Holway of Newton Centre for his many valuable suggestions. TABLE OF CONTENTS. PART I. INTRODUCTION. & PAGE Qualitative and Quantitative Analysis. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1 Reactions in the Wet Way. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1 Law of Chemical Mass Action. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9 Ion Theory of Arrhenius. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11 Lessening the Solubility of Precipitates. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15 Hydrolysis... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16 Detection of Acids and Bases. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 18 Reactions in the Dry Way.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 19 Concentrations of Reagents... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 30 REACTIONS OF THE METALS (CATIONS). GROUP V (ALKALIES). Potassium. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 36 Sodium. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 40 Ammonium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 44 Magnesium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 49 Separation of the Metals of Group V. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 51 GROUP IV (ALKALINE EARTHs). Calcium... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 54 Strontium. . . . . . . . . . . © e º Q tº $ tº gº e º ſº tº e º º & e º e º & © e º e º 'º G & & e º 'º e g º e º e º ºs e e e 58 Barium. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 59 Separation of the Metals of Group IV. . . . . . . . . . . . . . . . . . . . . . . . © & © tº e º e 62 Spectrum Analysis. . . . . . . . . . . . . . . . . . . . . . . . . tº e º ſº tº e º 'º e º ºs e º e º e º e º e º e 63 GROUP III. Aluminium... . . . . © e º e º º e º sº e ∈ E & © tº C G → G → © º º & © tº º e º e º ºs © & © tº gº e º G º e º te e e ºs 69 Chromium . . . . . . . . e tº e e s e e º e s e º e e s e e º ſº e º e º e s tº e º e º ºs e e e º e s e e s e e s e e s 76 Iron................... tº e e e s tº e º e º e º e º e º e º e º 'º s tº e º ºs e º e º e º ºs e º e º e º e 88 Uranium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 104 Titanium. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 108 Separation of Iron, Aluminium, Chromium and Uranium............. 112 viii TABLE OF CONTENTS. * PAGE Manganese........ tº º tº º º is º ºs é º e º 'º e º & © tº dº e º e º e º e º 'º tº ſº tº & © tº e º º ſº º ſº e g º º 113 Nickel. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . tº & ſº º ºs e º & tº C tº º tº e © & © C & © ſº e º g o º 126 Cobalt . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 132 Zine . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . tº G & º e º dº e º 'º º 138 Separation of Manganese, Nickel, Cobalt, and Zinc. ....... & O & e s tº e s e º 'º 143 Separation of all the Members of Group III........... tº ſº e º e º 'º e º is e 144, 145 GROUP II. Mercury... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . * * * * * * * * * * * * * * * * 146 Lead. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . M56 Bismuth... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 163 Copper . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 168 Cadmium ... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 175 Separation of Hg, Pb, Bi, Cu, and Cd from the preceding Groups and from one another................. ... . . . . . . . . . . . . . . . . . . . . . . . . 179 Arsenic... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 180 Antimony . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 199 Tin... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 210 Separation of the Sulpho-acids from the Sulpho-bases, and from one another. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 226 Gold . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 224 Platinum . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 232 GROUP I Silver. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 237 Mercury (Mercurous Compounds). . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 153 Lead . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 156 §eparation of the Metals of Group I. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 242 REACTIONS OF THE METALLOIDS (ANIONS). Division of the Acids into Groups. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 243, GROUP I. Hydrochloric Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 245 Chlorine . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 253 Hypochlorous Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 255 Hydrobromic Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 258 Bromine... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 261 Hydriodic Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 262 Iodine. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 265 Detection of HCl, HBr, and HI in the presence of one another. . . . . . . . . 268 Hydrocyanic Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .. 269 Dicyanogen. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 275 Hydroferrocyanic Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . tº e º 'º e º & J & tº e º e º & e e 276 Hydroferricyanic Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 279 Sulphocyanic Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 281 * TABLE OF CONTENTS. ix GROUP II. IPAGE Nitrous Acid. . . . . . . . . . . . . . . . . . . . . . . . . . 6 º' tº tº e º tº . . . . . . . . . . . . . . . . . . . . 284 Hydrosulphuric Acid (Hydrogen Sulphide). . . . . . . . . . . . . . . . . . . . . . . . . . 289 $ulphur . . . . . . . . . . . . . . . . . . . . . . . . . . . dº e º 'º º is ſº º tº e º & tº it is º e º º & © & © a . . . . . 293 Acetic Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . & © tº e º e º e ºs 295 Cyanic Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ... . 297 Hypophosphorous Acid. . . . . . . . . . . . . . . . . . . . . . . . . . tº e º e s tº dº ſº e º & G tº e º sº e 299 GROUP III. Sulphurous Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . tº e g a e º e º 'º ſº e . . . . . . . . 300 Carbonic Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 305 Boric Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 309 Oxalic Acid... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . * e º e º is º is tº e º 'º e e º 'º e º º 312 Tartaric Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 314 Citric Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 318 Phosphorous Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 319 Metaphosphoric Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 321 Pyrophosphoric Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 323 Iodic Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 324 GROUP IV. Phosphoric Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 326 Phosphorus . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 331 Arsenious Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 183 Arsenic Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 186 Chromic Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 79 Thiosulphuric Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 333 Detection of Sulphurous and Thiosulphuric Acids in the presence of Hydrogen Sulphide... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 336 GROUP V INitric Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 337 Detection of Nitric Acid in the presence of Nitrous Acid. . . . . . . . . . . . . . 341 Chloric Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 343 Detection of Hydrochloric, Nitric, and Chloric Acids in the presence of one another. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 344 Perchloric Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 345 Persulphuric Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 346 GROUP VI. Sulphuric Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 348 Hydrofluoric Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 349 Hydrofluosilicic Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 354 GROUP VII. Silicic Acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . * * * * * * * c e º is e 355 Treatment of Insoluble Silicates. . . . . . . . . . . . . . . . . . . tº e º s º e º e ge & © tº e º is e 359 Silicon . . . . . . . . • * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * = e s a e s a 362 X n TABLE OF CONTENTS, PART II. COURSE OF ANALYSIS. PAGE) Analysis of Solid and Non-metallic Substances..... * * * * * * * * * * * * * * * * * * 365 Preliminary Examination............................. tº º O & Q - e º ſº tº G & 365 Solution of the Substance. . . . . . . . . . . . ............................. 372 Solubility Table . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 372 Methods of getting Insoluble Substances into Solution. . . . . . . . . . . . . . . . 376 Examination for the Metals (Cathions). . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 380 Examination for the Negative Elements (Anions). . . . . . . . . . . . . . . . . . . . 394 Analysis of Alloys. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 399 Analysis of Liquids. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 400 SUPPLEMENT. REACTIONS OF SOME OF THE RARER METALS. GROUP V (ALKALIES). Caesium. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 407 Rubidium. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 408 Lithium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 409 Detection of Lithium, Rubidium, and Caesium in the presence of Sodium and Potassium... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 411 GROUP III. Beryllium. . . . . . . . . . . . . . . . . * - - - - - - - - - - - - - - - - - - - - - - - © tº ſº º e º e º e e º e º º 412 Zirconium. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 413 Thorium. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ............. & © e º e º O e e e s tº 414 Yttrium... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . tº e o 'º e e º º e º 'º e 416 Cerium... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . is * * * * * * * * * * © & © e o is tº e º e o e 418 Lanthanum . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 421 Didymium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . tº º e º 'º º e º e º e e 428 Analysis of Gadolinite. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . tº º e º O © e o e º e e º e 426 Tantalum. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 429 Niobium. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . & e º 'o tº e e e e º º e º e º e º c e e e e º e 430 GROUP II. Thallium... . . . . . . . . . © o e º ºs e º e º & & © e º e º e º s e e º e º ºr © g c e º O 6 & 2 G s e g º ºs e e º ºs 431 Vanadium... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . © e º e º 'º e º 'º º 434 Molybdenum.... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . tº G & e o O & © e º 'º e o O e º e ... 437 Tungsten... . . . . . e e o e º e e e º e º te e e s e º e º e º e s e º e º e º e º e tº gº o e º O e e e tº ſº e º G tº e 438 Selenium..... e e e s e e e º e º e º O - © e º e º e tº e e e º e º O & © tº 6 e º & Q & Q & Q © tº e o e º e o e o ºs 441 Tellurium. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . tº º e s tº e Q © e º O O & © tº e e e g º e 444 Palladium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . © & © tº º & e º e C & e º º º 447 Rhodium. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 450 3Osmium...... e e e e s e e o e C & e e o O & 0 e º O e º e exe e º e o 'º e º e o e º 0 & 0 & 0 & © tº e e s e e s e 452 / Ruthenium ..... e e e º g tº e º e o e o e º e º e º e º O & © tº e e º e º e º e o e e © e º 'º e º & © e o e º 0. 454 Hridium............. e e s e o e o e e o e o e © Q C & O e º º O © e º tº º O C → Q & Q & © tº o e e º ºs e e º e 456 Separation of the Platinum Metals. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 457 Spectroscopic Chart. . . . . . . . . . . . . . . e e º e º e º e º e º O e º e o ſº tº * * * * * * Frontispiece 1908. TABLE OF INTERNATIONAL ATOMIC. WEIGHTS. Name. Sym. O = 16. H = 1 Name. Sym. O = 16. H= 1. Aluminium . . . . . Al 27.1 26.9 Molybdenum . . . . Mo 96.0 | 95.3 Antimony..... Sb 120.2 |119.3 eodymium .... Nd 143.6 |142.5 Argon . . . . . . . . A | 39.9 || 39.6 Neon . . . . . . . . . . Ne | 20. 19.9 Arsenic . . . . . . . As 75.0 74.4 Nickel . . . . . . . . . Ni 58.7 || 58.3 Barium . . . . . . . Ba |137.4 |136.4 Nitrogen . . . . . . . N | 14.01. 13.91 Bismuth . . . . . . Bi 208.0 |206.3 Osmium . . . . . . . OS 1191. 189.6 Boron . . . . . . . . B | 11. 10.9 Oxygen . . . . . . . . O 16.00. 15.88 Bromine . . . . . . Br 79.96 || 79.36 || Palladium . . . . . . PC 106.5 |105.7 Cadmium . . . . . Col 112.4 |111.6 Phosphorus . . . . P 31.0 30.77 Caesium . . . . . . . Cs 132.9 |131.9 Platinum . . . . . . Pt. 194.8 |193.3 Calcium. . . . . . . Ca. | 40.1 || 39.7 Potassium. . . . . . K 39.15 38.85 Carbon . . . . . . . C | 12.00 | 11.91 || Praseodymium..] Prí140.5 |139.4 Cerium . . . . . . . Ce 140.25 |139.2 Radium. . . . . . . . Ra |225. 223.3 Chlorine . . . . . . Cl 35.45 || 35.18 || Rhodium . . . . . . Rh 103.0 102.2 Chromium . . . . . Cr | 52.1 || 51.7 Rubidium ... . . . Rb 85.5 | 84.9 Cobalt . . . . . . . . Co 59.0 | 58.55 || Ruthenium ... . . . . Ru (101.7 100.9 Columbium. . . . . Cb | 94. 93.3 Samarium ... . . . Sa |150.3 |149.2 Copper . . . . . . . Cu | 63.6 || 63.1 Scandium ... ... Sc 44.1 || 43.8 Dysprosium ...] Dy 162.5 161.2 Selenium ... . . . . Se 79.2 || 78.6 Erbium . . . . . . . E 166. 164.8 Silicon . . . . . . . . . Si 28.4 || 28.2 Europium .....] Eu 152. 150.8 Silver. . . . . . . . . . Ag 107.93|107.11 Fluorine . . . . . . F | 19. 18.9 Sodium . . . . . . . . Na 23.05. 22.88 Gadolinium . . . ] Gd (156. 154.8 Strontium ... . . . Sr 87.6 | 86.94 Gallium . . . . . . . Ga, 70. 69.5 Sulphur . . . . . . . . S 32.06| 31.82 Germanium . . . . Ge | 72.5 | 72. Tantalum . . . . . . Ta. 181. 179.6 Glucinum . . . . . Gl 9. 1 9.03 || Tellurium . . . . . . Te 127.6 |126.6 Gold . . . . . . . . . Au |197.2 |195.7 Terbium . . . . . . . Tb 159.2 |157.7 Helium . . . . . . . He 4. 4. Thallium ... . . . . Tl 204.1 |202.6 Hydrogen |H 1.008] 1.000|| Thorium . . . . . . . Th |232.5 |230.8 Indium . . . . . . . In 115. 114.1 Thulium . . . . . . . Trn 171. 169.7 Iodine . . . . . . . . I 126.97 |126.01 in . . . . . . . . . . . . Sn 119.0 |118. 1 Iridium . . . . . . . Ir |193.0 |191.5 Titanium . . . . . . Ti || 48.1 || 47.7 Iron . . . . . . . . . . Fe 55.9 55.5 Tungsten . . . . . . W 184.0 |182.6 Krypton . . . . . . Kr | 81.8 || 81.2 Uranium . . . . . . . U (238.5 |236.7 Lanthanum . . .] La |138.9 |137.9 Vanadium. . . . . . V 51.2 50.8 Lead . . . . . . . . . Pb |206.9 |205.35 || Xenon . . . . . . . . . Xe |128. 127. Lithium . . . . . . Li 7.03 || 6.98 || Ytterbium . . . . . Yb 173.0 |171.7 Magnesium . . . . . Mg 24.36 24, 18 || Yttrium . . . . . . . Y | 89.0 | 88.3 Manganese . . . . . Mn 55.0 || 54.6 Zinc . . . . . . . . . . . Zn 65.4 64.9 Mercury . . . . . . Hg |200.0 |198.5 Zirconium . . . . . . Zr 90.6 | 89.9 xi QUALITATIVE ANALYSIS. GENERAL PRINCIPLES. By Chemical Analysis is understood all those operations which are performed in order to determine the constituents of a chem- ical compound (or a mixture of chemical compounds). Chemical Analysis is subdivided into Qualitative Analysis and Quantita- tive Analysis. Qualitative Analysis treats of the methods for determining the nature of the constituents of a substance, while Quantitative Analysis treats of the methods for determining in what proportion the constituents are present in any compound or mixture of compounds. In order to recognize a body we change it, usually with the help of a substance of known nature, into a new compound which pos- sesses distinctive properties. This transformation we call a chemi- cal reaction; and the Substance by means of which the reaction is brought about, the reagent. We distinguish between reactions in the wet way and reactions in the dry way. I. REACTIONS IN THE WET WAY. For the purpose of qualitative analysis only such reactions are applicable as are easily perceptible to our senses. A reaction is known to take place— (a) by the formation of a precipitate; (b) by a change of color; (c) by the evolution of a gas. 2 GENERAL PRINCIPLES. A precipitation always takes place when an insoluble substance is formed by means of a chemical decomposition. If dilute sulphuric acid is added to a solution of a barium salt, a white powdery pre- cipitate of barium sulphate is obtained: BaCl, +H,SO. =2HCl- BaSO4. On the other hand, soluble lead salts also give with dilute sulphuric acid a white powdery precipitate: Pb(NO.), + H.SO, F. 2HNO, + PbSO,. Sulphuric acid, therefore, is a reagent for both barium and lead compounds. But in order to determine whether the precipi- tate formed is barium sulphate or lead sulphate it must be sub- jected to a further test, because both substances look alike. If they are heated on charcoal with addition of sodium carbonate, they behave very differently; the lead sulphate is reduced to a metal, and the barium sulphate is only changed to carbonate. It follows from this example that one must never be content with a single reaction for the detection of a body, but its presence must be verified by a confirmatory test. If to a solution of ferrous chloride (obtained by the solution of metallic iron in hydrochloric acid) some sodium hydroxide is added, a greenish-white precipitate of ferrous hydroxide is formed: FeCl,--2NaOH=2NaCl-i-Fe(OH), which, on standing in contact with the air, changes to dark green, almost black, and finally becomes brown, the compound having been oxidized by means of atmospheric oxygen to ferric hydroxide: 2Fe(OH)2+H,O+O=2Fe(OH)s. Green Brown If chlorine water is added to the green solution of ferrous chlo- ride, a change of color takes place, due to the oxidation of the fer- rous salt to a ferric one: FeCl, +Cl=FeCls. Green Yellow If now sodium hydroxide is added to the yellow solution, the brown precipitate of ferric hydroxide is immediately formed: FeCls--3NaOH=3NaCl-i-Fe(OH)2. OXIDATION. 3 Sodium hydroxide is, therefore, a reagent for ferrous as well as for ferric salts; and although the two precipitates cannot be mis- taken for one another, yet, for the sake of certainty, it is advisable that the beginner should make a confirmatory test. We saw that the green ferrous chloride was changed by means of chlorine water into the yellow ferric salt, and recognized that a reaction had taken place by means of the change in color. Color reactions in the wet way occur very often because of an oridation, as in the case just mentioned; but, conversely, they may take place in consequence of a reduction. We shall constantly have to perform oxidations and reductions, so that we will briefly mention the most important methods. Oxidation.—We oxidize a compound when we increase the proportion of its oxygen or acid-forming element (or group). Ferrous oxide on heating in the air becomes ferric oxide: 2Fe0+O=Fe,0s. By dissolving these two oxides in hydrochloric acid the ferrous oxide becomes ferrous chloride and the ferric oxide forms ferric chloride: Fe0+2HCl=FeCl,--H.O, Ferrous oxide Ferrous chloride Fe=O >o +6HCl=2FeCls--3H,0. Fe= O Ferric oxide Ferric chloride Since the ferrous chloride, obtained from ferrous oxide by replacement of oxygen by chlorine, is readily changed into ferric chloride (which we obtained from ferric oxide), the change of ferrous chloride into ferric chloride is also considered an oxidation, although oxygen itself does not take part in the reaction. The most important oxidation agents are– 1. Halogens; 2. Nitric Acid; 3. Hydrogen Peroxide; 4. Potassium Permanganate; 5. Potassium Dichromate. 4 GENERAL PRINCIPLES. 1. The oxidizing action of the halogens depends either upon their direct addition: FeCl, +Cl=FeCls, or the halogen decomposes water, forming, for example, hydro- chloric acid and oxygen: ? H.O.--Cl,-2HC1--O. Thus the oxidation of ferrous sulphate by means of chlorine water may be expressed by the equation & 280, . FeSO Feº *+Cl,--H,SO, = 2HCl-- >SO. \so, 2. The oxidizing action of nitric acid depends upon the formation of the anhydride, which then breaks down into nitric oxide and Oxygen: 2HNO3= H,0+N,0s, N.O. =2NO+3O. Suppose, for example, we wish to oxidize ferrous sulphate to ferric sulphate by means of nitric acid, we must add, then, to every 2FeSO, one more SO. : The SO, necessary for this reaction may be obtained from sulphuric acid (which we add to the solution) by oxidizing its hydrogen atoms: H,SO,--O= H,O+SO4. Consequently the equation in its simplest form would be ;34. H,SO,--O = H,0+Fe2(SO4)s. Since we desire, now, to use nitric acid as the oxidizing agent, two molecules of which furnish three atoms of oxygen, the true reaction is expressed by the equation 6FeSO,--2HNO,--3H,SO. =4H,0+2NO+3Fe,(SO.),. OXIDATION. 5 3. The oxidizing action of hydrogen peroxide depends upon its decomposition into water and oxygen: H.O. = H,0+O; for example, 2FeCl,--2HC1-i-H,O,-2H,0+2FeCls. 4. The oxidizing action of permanganic acid depends upon the splitting off of the anhydride, which is then decomposed into manganous oxide and oxygen: 2HMnO, = H,0+Mn,O, Mn,0,–2MnO--50. On adding to a colorless acid solution of ferrous sulphate some of the intensely red-colored solution of potassium permanganate drop by drop the red color disappears instantly on stirring, and remains permanent only when all of the ferrous salt is changed to ferric salt. The appearance of the red color of the permanganate shows, there- fore, the complete oxidation of the ferrous salt. The reaction. which takes place may be represented by the following equation: 2KMnO,--10FeO= K,0+2MnO--5Fe,Os. There must, however, be sufficient acid present to dissolve the oxides which are formed. The true reaction is, therefore, 2KMnO,--10FeSO,4-8H,SO, =K,SO,4-2MnSO,--5Fe,(SO,),4-8H,O. In alkaline solution the reaction goes somewhat differently: two molecules of the permanganate giving up three atoms of oxygen, being reduced only to MnO,. 2HMnO, = H,0+Mn,Or, Mn,0,-2MnO,--30. 5. The oxidizing action of chromic acid depends, similarly, upon the decomposition of chromic anhydride into oxygen and green chromic oxide: 2CrOs=Cr,Os--3O. Thus ferrous salts, in cold, acid solution, are immediately oxidized to ferric salts: 20rO,--6Fe0=Cr,0,--3Fe,0s. 6 GENERAL PRINCIPLES. For this oxidation potassium dichromate and a mineral acid are employed, the purpose of the latter being, in the first place, to Set free the chromic acid; and, furthermore, to keep the oxides formed in solution. The oxidation of ferrous sulphate to ferric sulphate by means of potassium dichromate is represented by the following equation: K.Cr,0,--6FeSO,--7H,SO, =7H,O+K,SO,--Cr,(SO.),4-3Fe,(SO.),. Reduction.—We reduce a compound by subtracting oxygen or an acid-forming element (or group) from it. Ferric oxide, for example, when heated with charcoal is changed first to ferrous oxide, and finally to metallic iron: Fe,0,--C=CO+2Fe0, 2Fe0+2C=2CO+Fe,. In the same way, the change of a derivative of the higher oxide back to one of the lower oxide is called a reduction. Thus the with- drawal of chlorine from ferric chloride is a reduction: FeCla-Cl–FeCl,. The most important reducing agents are: 1. Nascent Hydrogen; 2. Sulphurous Acid; 3. Hydrogen Sulphide; 4. Stannous Chloride; 5. Hydriodic Acid. 1. Reduction by means of hydrogen may take place in acid, alkaline, or neutral solution. (a) In acid solution, by the employment of zinc, etc.: Zn-H H,SO,- ZnSO,--H,. Thus silver chloride may be readily reduced to metal: 2AgCl--Zn-i-H,SO,- ZnSO,--2HCl·HAga, and arsenious acid may be changed to arsine (arseniuretted hydro- gen): As, O,--6Zn-i-6H.S.O. =6ZnSO,--3H,O+2ASH.a. (b) In alkaline solution, by means of zinc, aluminium, sodium amalgam, or, best of all, Devarda's Alloy (Cu-50, Zn=5, Al=45); REDUCTION. this Alloy is so brittle that it can be easily powdered in a mortar, making it easy to employ small amounts. This reduction also de- pends upon the formation of nascent hydrogen: , Zn-i-2NaOH =Zn(ONa), +H, Al-H3NaOH = Al(ONa), +3H. In the case of Devarda's Alloy, an electrical action comes into play, whereby the reaction is completed much more quickly than by the use of either zinc or aluminium alone. Nitrates and chlorates may be reduced in a few minutes by means of Devarda's Alloy and a few drops of sodium hydroxide; the reaction also takes place in neutral solution, but it takes considerably longer: RNO,--4Zn-H7KOH=4Zn(OK), --2H,0+NH, RCIO,--3Zn-H6KOH=3Zn(OK), H-3H,0+KCl. 2. Reduction by means of sulphurous acid takes place in mod- erately acid solution, and depends upon the fact that SO, readily takes on oxygen and becomes SOs (SO,--O=SOs), which forms with water sulphuric acid. Ferric salts are readily reduced by this reagent: Fe,(SO.)a-H2H,O+SO,-2H,SO,--2FeSO4. In the same way, arsenic anhydride and many other substances are reduced very readily and completely by means of SO2: As O.--2H,0+2SO,-2H,SO,--As,C), An excess of aqueous sulphurous acid is added to the solution which is to be reduced; it is then heated to boiling; and the boiling is continued while a stream of carbonic acid gas is passed through the solution until the excess of sulphurous acid is driven off. 3. Reduction with hydrogen sulphide depends upon its easy decomposition into water and Sulphur: 47 H.S= H,--S. This method of reduction is seldom employed in analytical chemistry because of the difficulty in filtering off the separated sulphur. Many metals are precipitated as sulphides from weakly acid solutions by hydrogen sulphide. If the solution, however, contains oxidizing agents (such as nitric acid, chloric acid, chromic 8 GENERAL PRINCIPLES. acid, etc.), these will be reduced by the hydrogen sulphide with separation of sulphur. The Sulphide obtained will be largely contaminated with sulphur, which renders the Subsequent exami- nation more difficult. If the solution contains no metal which is precipitated by hydrogen Sulphide, but contains oxidizing agents, it will still cause separation of Sulphur. One is often in doubt whether there is not some sulphide mixed with the sulphur, and is therefore obliged to examine the precipitate further, which is often unnecessary if the oxidizing agent is previously destroyed. Hydro- gen sulphide reduces Halogens: H.S.--Cl, =2HC1+S; Nitric Acid: 2HNO,--3H,S =4H,O--2NO+3S; Chloric Acid: HClO,--3H,S =3H,O+HCl·H2S; Ferric Salts: 2FeCls--H.S =2HCl·H2Fe(Cl2 +S; Chromic Acid: 20rO,--3H, S =3H,0+Cr,Os--3S; Permanganic Acid: 2HMnO,--5H,S=6H,0+2MnO-H-5S, and many other substances. 4. Reduction with stannous chloride takes place usually in acid solutions. The reduction depends upon the fact that stannous chloride is readily changed to stannic chloride: SnCl2+Cl, =SnCl4. Ferric salts, chromates, permanganates, mercuric Salts, and many others are reduced in this way: 2FeCla-i-SnCl2=SnCl, +2FeCl,; Yellow Colorless—green 2CrO,--12HCl·H2SnOl,-3SnCl,--6H.O-H2CrCls; Yellow Green 2HgCl,--SnCl2=SnCla-i-Hg,Cl, Colorless White precipitate Hg,Cl,--SnCl2=SnCla-H2Hg. Gray precipitate 5. Hydriodic acid also is a strong reducing agent, being easily decomposed into hydrogen and iodine: HI= H+I. LAW OF MASS ACTION. 9 Ferric Salts, under certain conditions (see page 11), may be quanti- tatively reduced in this way to ferrous salts: 2FeCl,+2HI=2HCl·H2Fe(Cl; +I,. As, however, iodine is an oxidizing agent, under other conditions the ferrous salt can be changed back into ferric salt, so that the reduction then will not be quantitative. But it is very important for the analyst to perform all reactions so that they take place as nearly as possible quantitatively. How this may be brought about by changing the conditions of the experiment is shown by The Law of Chemical Mass Action. If, for example, we allow chlorine to act upon phosphorus tri- chloride in the cold, solid phosphorus pentachloride is formed: PCls--Cl,-PCls. If, however, the phosphorus pentachloride is heated, it is decom- posed into chlorine and phosphorus trichloride: PCls= PCls--Cl,. The reaction accordingly is reversible, which may be represented by writing the equation as follows: PCls z=2 PCls+Cla. We use, instead of the sign of equality, two arrows pointed in oppo- site directions, signifying that the reaction may take place in the sense either of from left to right, or from right to left. Similar to phosphorus trichloride, there are a large number of other sub- stances which on heating decompose into their components, and unite again on cooling. St. Claire Deville (1857) designated this phenomenon Dissocia- tion. Ammonium chloride, on volatilization, is dissociated into ammonia and hydrochloric acid, and the dissociation increases with rise of temperature until, at a sufficiently high heat, it is complete. For every temperature, the ratio of the wrºdissociated to the dissociated part is a constant. Let us return to our first example, the phosphorus pentachloride. If we designate by [PCls] the number of undecomposed gram-mole- IO GENERAL PRINCIPLES. cules of phosphorus pentachloride per liter, by [PCl] the num- ber of gram-molecules of trichloride, and by [Cl2] the number of gram-molecules of chlorine in the same volume, then the quotient [PCl3]X[Cl2] [PCl5] is a constant at any definite temperature. If, therefore, we desire to volatilize phosphorus pentachloride so that the least possible dissociation will take place, the above equation shows us how this may be brought about. If either [PCl3] or [Cl2] be increased, then in order that the value of the fraction [PCl5] remains constant, it is evident that the concentration [PCl5] must become greater; or, in other words, the dissociation of the penta- chloride becomes less and there will be practically no dissociation if the pentachloride is volatilized in an atmosphere of phosphorus trichloride or of chlorine. In this way Wurz obtained for the den- sity of phosphorus pentachloride 6.80–7.42, instead of the calcu- lated value 7.2. Dissociation may be brought about by solution as well as by volatilization; and in this case the dissociation increases with the dilution. For example, if the mineral carnallite (MgCl2. KCl +6H2O), which occurs at Stassfurt, be dissolved in water and the solution allowed to evaporate until a salt begins to crystallize out, we shall find that instead of obtaining crystals of carnallite we obtain those of potassium chloride. Carnallite is dissociated in aqueous solution into easily soluble magnesium chloride and the far less soluble potassium chloride which separates out: MgCl2. KClz=2 MgCl2 +KCl. For every concentration the degree of dissociation is a constant. If we designate by [MgCl2. KCl] the number of gram molecules per liter of undecomposed carnallite, by [MgCl2] the number of gram molecules per liter of magnesium chloride, and those of potassium. chloride by [KCl], the equation holds: [MgCl2]>{[KCl] Rijkaiº-econstant. THEORY OF ELECTROLYTIC DISSOCIATION. I Itſ If we wish, therefore, to recrystallize carnallite we must prevent its dissociation by increasing either [MgCl2] or [KCl].” As a matter of fact, the mineral is recrystallized at Stassfurt from a 23 per cent magnesium chloride solution. The reduction, mentioned on page 9, of ferric salts by means of hydriodic acid represents, similarly, a reversible reaction; 2FeCl3+2HIz=22FeCl2 +2HCl-HI2, and for every degree of concentration the following equation holds: [FeCl2]2×[HCIPX[I2] N [FeClaº XIHIP =a constant. If we desire to make the reaction take place from left to right quantitatively, the concentration of the hydriodic acid should be: increased whereby that of the ferric chloride becomes diminished. In fact, by employing a large excess of HIf the ferric salt may be: quantitatively reduced to ferrous salt, and this reaction is made. use of in the quantitative determination of iron. This law of Chemical Mass Action was first enunciated by Guld- berg and Waage (1867), and holds for all reversible reactions. whether they take place by volatilization or by Solution. Theory of Electrolytic Dissociation. If we insert between the poles of an electric battery a piece of rock salt or some pure distilled water, there will be no electric cur- rent in the circuit; a piece of fine platinum wire placed in the circuit will not be made to glow. The solid rock salt, as well as the dis- tilled water, are non-conductors of electricity; they are non-electro- lytes. If, however, we dissolve rock salt in distilled water, and * In reality an increase of [KCl] will not accomplish the desired result, because potassium chloride is more insoluble than either carnallite or mag- nesium chloride. f Ferric chloride is to be reduced to ferrous chloride; i.e., [FeCl] must be made as small as possible. If we wish to oxidize all of the hydriodic acid, then [FeCls] must be made as large as possible, whereby [HI] becomes prac- tically zero. I 2 GENERAL PRINCIPLES. then insert the solution between the poles of the electric battery, the platinum wire will be brought to a bright glow, showing that the salt solution is a good conductor of electricity—an electro- lyte. It is thereby proved that by dissolving the non-conducting rock salt in non-conducting water an essential change of the former has taken place. We can make the same observation with all acids, bases, and salts. In an anhydrous state they are non-electrolytes, while in aqueous solution,” on the other hand, they are electrolytes. These phenomena are readily explained by the theory of elec- trolytic dissociation proposed by Arrheniusf in 1887. According to this theory, all electrolytes are partially decomposed in aqueous solution into electrically-charged atoms or atom-groups called ions; and the extent of this dissociation increases with dilution, until with very great dilution it is practically complete. For every degree of dilution there exists a certain state of equilibrium between the ions and undissociated molecules. When the non-electrical rock salt is dissolved in water, it breaks up, according to the equation NaCl z=z Na' +Cl’, into positively-charged sodium ions and negatively-charged chlo- rine ions. All salts, acids, and bases behave like rock salt. Thus sodium . sulphate decomposes according to the equation Na2SO4 z=z Na' +Na’ +SO4”, and sodium hydroxide into NaOH => Na' +OH'. By this theory of electrolytic dissociation the phenomena of electrolysis may be explained very simply: If we insert the two poles of a source of electricity into an electrolyte, one of the poles, the anode, is charged with positive electricity, and the other, the cathode, with negative electricity. The electro-positive anode repels the electro-positive ions (cations) and attracts the electro- negative ions (anions); and the latter, as soon as they come in * They are also electrolytes in the fused state. + Z. physik. Chem. I, p. 681. THEORY OF ELECTROLYTIC DISSOCIATION. I3 contact with the anode, give up their negative electricity, become non-electrical and separate out. The same thing happens at the cathode, where the electro-positive ions (cations) are discharged. The amounts of electricity which are neutralized at the electrodes are always renewed by the source of the current, so that the process is continuous. & The ions alone are concerned in electrolysis; the undissociated molecules take no part in the process. Just as the transport of electricity through liquids is brought about only by means of the ions, it is likewise true in the majority of cases that it is the ions which cause reactions to take place in aqueous solutions; in fact, the reactions take place more rapidly in proportion to the concentration of the ions. The more ions there are present in a unit of volume the better the solution conducts electricity, because, according to Faraday's law, equivalent ions transport equal amounts of electricity. - If, for example, the conductivity of two acids at the same con- centration is as 2:1, it follows that the former acid is twice as strongly dissociated as the latter. The first acid, therefore, will react chemically much more energetically, and for this reason we call it the stronger acid. The degree of ionization of an electrolyte can be computed from its electrical conductivity. The strong acids and bases are very largely dissociated, while the weaker acids, such as acetic, carbonic, hydrocyanic, and boric, are dissociated but slightly. This is evident from the table on page 14. Electrolytes having polyvalent ions undergo a progressive dis- sociation on being dissolved in water; thus sulphuric acid in con- centrated solution is dissociated according to the equation H2SO4·2H +HSO4' in H' and HSO4' ions and on diluting the solution the latter are gradually dissociated into H' and SO4” ions, so that in dilute solu- tion the following may be present: A small amount of undisso- ciated H2SO4 molecules, H ions, SO4” ions, and possibly some HSO4' ions. In the case of such weak dibasic acids as carbonic acid or hydro- gen sulphide, the second stage of the dissociation is scarcely reached I 4 GENERAL PRINCIPLES. Degree of Dissociation of Certain Electrolytes. A. ACIDS. e D f e * ; Acid. $. Diºn. ; Acid. Bºº. HCl Y. HF about 13% # Hºo. . 1.3% , 0.15 HNO, about 90% #'s." €6 }% HCIO, HCN 0.05% #3; HaBOs &G 0.013% HaPO, “ 44% #H.so, 6 2 } * 62% HIO, “ 29% s H.SeO, H.PO, “ 20% 32 H.C.O. “ 65% 4 & H.C.H.O. “ 20% B. BASEs. C. SALTs. ; Base. Dissociation. Type of Salt. Dissociation. ROH, NaOH about 94% R*R/ about 85% Ba(OH), “ 75% R“R,’ “ 70% NH,OH & & 1.3% R,'R'' “ 70% R." e R” tº £6 45% even in quite dilute solutions. Carbonic acid is dissociated chiefly according to the equation H2CO3 +2 H +HCO3'. HCO3’z-> H +CO3” takes place only to an inappreciable extent. Exactly the same is true of hydrogen sulphide; it is dissociated to some extent according to the equation H2Sz-> H +SH', whereas the second stage of dissociation hardly takes place at all. It is altogether different in the case of neutral salts of the weak acids. These are largely dissociated, as follows: Na2CO3 z=z Na' +Na’ +CO3”, Na2S z=z Na' +Na' --S”. This explains the fact that such weak acids in the free state in many respects show quite different reactions than their salts. The dissociation DEGREE OF DISSOCIATION OF CERTAIN ELECTROLYTES. I 5 Influence of Changes of Concentration upon the Dissociation of Electrolytes. If we assume 1 gram molecule of a weak electrolyte, such as ammonia, to be dissolved in v liters of solution, the ammonia will be partly dissociated according to the equation NH4OHz-> NH4 +OH' into ammonium and hydroxyl ions. If a gram molecule of the base is dissociated in the sense of the above equation, then the undissociated part will amount to 1-0. The côncentrations per liter are º, A & undissociated dissociated tº & NH4OHa-> NH4 +OH' 1–0. OW. OY Q) ^) Q} and according to the mass-action law & 1.1-q_2: | Cº- (- - ^) º { o:2 *TGia), The constant k is known as the dissociation- or affinity-constant; it is independent of the dilution and is characteristic for every electrolyte. The expression shows, however, that by increasing, v, or diluting, the dissociation (a) will be made larger. If to the solution of the base we add additional ammonium ions by adding n gram molecules of Solid ammonium chloride, then if n is considerably larger than a the degree of dissociation of the base will be greatly diminished, namely, from a to a 1, a value which we can readily compute as follows: In the solution there is present per liter undissociated dissociated NH4OHz-> NHa' --OH' 1–0.1 0.1 + n, 0.1 ^) Q) v' therefore k= (0.1 +n)0:1 (1–0.1) v I6 GENERAL PRINCIPLES. If k and n are known aſ I can be computed: — (n +vk) +M (n +vk)2+40k Q1 = 2 o In the case of N/10 ammonia solution the ammonium hydroxide is dissociated only to an extent of 1.32%, the dissociation constant being 0.000018. If we add to 10 liters of this ammonia solution 2 gram molecules of ammonium chloride (107.08 gms.), then since the ammonium chloride is 93% dissociated at this dilution, we are adding 2X0.93=1.86 NH4 ions. If this value be inserted in the above equation, all becomes 0.00009; in other words, the dissociation of the ammonium hydroxide is diminished by the ad- dition of the ammonium chloride from 1.32% to 0.009%. The solu- tion now contains so few hydroxyl ions that it will not precipitate solutions of magnesium salts (cf. page 50). Similarly the dissociation of weak acids is lessened by the addi- tion of their salts. In the case of the stronger acids and bases, the effect of adding a neutral salt to the solution is not so remarkable, because the stronger acids and bases are dissociated to about the same extent as their salts. Reactions of the Ions. When a solution of hydrochloric acid, or of any soluble chloride is treated with silver nitrate solution, a white curdy precipitate of silver chloride is formed: * R’ +Cl’ +Ag' +NO3" = AgCl--R +NO3'. *—-----' \–—-y——” |-----—” In the solution at the start there are both ions and undissociated molecules present. The electropositive silver ions unite with the electronegative chlorine ions and form electrically neutral silver chloride, which being insoluble is thrown out of solution. This disturbs the equilibrium in the solution, and in order to restore it the reaction between silver and chlorine continues to take place until finally practically all of the chlorine is removed from the solution. In the above equation, if we remove the R ar.d NO3, which take no part in the reaction, the equation becomes in its simplest form Cl’ +Ag' = AgCl. REACTIONS OF THE IONS. 17 The silver ion is a reagent for detecting the presence of chlorine ions. All compounds containing chlorine do not react with silver ions to form insoluble silver chloride, but only those which on being dissolved in water yield chlorine ions. If a solution of potassium chlorate, for example, be treated with silver nitrate solution, there will be no precipitate of silver chloride formed, because the chlorate is not dissociated into chlorine ions, but as follows: KClO3 z-z K' +ClO3. A study of the products obtained on electrolyzing a solution is often a convenient method for determining the species of ions present, but it must be remembered that when certain atomic complexes, such as SO4, NH4, etc., are set free, they undergo second- ary reactions, so that the products obtained on electrolyzing the solution do not correspond to the ions originally present. Thus in the case of the electrolysis of aqueous sulphate solutions the SO4 ions, when discharged by the current, decompose water with the evolution of oxygen, Sulphuric acid remaining in solution: 2SO4+2H2O=2H2SO4+O2. In the same way when ammonium (NH4) is set free by elec- trolysis, ammonia and hydrogen are at once formed. COMPLEX IONS. We learned on page 2 a characteristic reaction for ferrous salts, namely, the precipitation by means of sodium hydroxide. This reaction is one between ions and takes place according to the equa- tion Fe" +Cl2” +2 Na' +2OH' =2|Na' +2C1'--Fe(OH)2, S–-y- —” \--——y--- – -—” \--- —’ or, if we eliminate the ions which take no part, Fe" +2OH' = Fe(OH)2. The hydroxyl ion is a reagent for ferrous ions. If the insoluble ferrous hydroxide is boiled with potassium cyanide solution it dissolves, forming potassium ferrocyanide, K4Fe(CN)6. This salt is not a ferrous salt and does not give the characteristić reactions for Fe ions. It is the potassium salt of hydroferrocyanic acid, H4Fe(CN)6. "I 8 GENERAL PRINCIPLES. On being dissolved in water, the salt undergoes electrolytic dis- sociation, as follows: KAIFe(CN)6]=: 4K --[Fe(CN)6]”. The Fe forms an integral part of the complex ferrocyanide ion; on electrolysis the latter always migrates to the anode and shows characteristic reactions with different reagents (cf. hydroferro- cyanic acid). * Quite a number of such complex ions are met with in the study of analytical chemistry, and in the formulae given in this book such ions, as a rule, will be enclosed in brackets. It is the task of Qualitative Analysis to find out tests for all the ions which may be present and to treat of their separation. It will be 'our aim to transform these ions into insoluble compounds, into gases, or into characteristically colored compounds. Solubility Product. When we desire to transform an ion into an insoluble precipi- tate we must consider how best to carry out the operation. Shall we, in forming insoluble barium sulphate from a solution containing SO4 ions, add simply an equivalent amount of barium chloride, or shall we add an excess of the reagent? No precipitate is absolutely insoluble in water, so that a little barium sulphate will remain in solution. The solution, however, is then saturated with barium sulphate and contains it chiefly in the form of Ba and SO4 ions: BašO4 z-z Ba” +SO4’’. If we designate the concentration of the BaSO4 by [BaSO4], that of the Baby [Baj, and of the SO4 by [SO4], then for any definite temperature the equation holds: [Baj XISO4]=[BaSO4]Xk. If we increase the concentration of the barium ions, [Baj, then the equation shows that [BaSO4] will necessarily become larger, but as the solution is already saturated with barium sulphate this will result in a further precipitation of the salt. The same end would be attained by the addition of an excess of SO4 ions. The product of the concentration of the ions in a saturated solution is known as the solubility product. When this value has not been reached, the solution will dissolve more of the solid substance. In the case of our precipitates the important rule holds: to diminish the solu- REACTIONS OF THE IONS. I9 bility of a salt it is only necessary to increase the concentration of one of the ions into which the salt dissociates. The excess of reagent to be added depends upon the solubility of the precipitate, and upon its tendency to form other atomic complexes. Thus, in precipitating lead or barium sulphates, both of these are soluble in concentrated sulphuric acid. They are least soluble in the presence of a slight excess of sulphuric acid. In order to examine the precipitate further, it must be separated from the Solution by means of filtration, and the general rule holds: The size of the filter should always be governed by the amount of the precipitate and not by the volume of the liquid. In cases where it is a matter of detecting minute traces, it is often necessary to work with large amounts of the original substance, which requires a cor- respondingly large amount of liquid to effect solution. If a few milligrams of an unknown body are precipitated from such a solu- tion, and this is filtered off by means of a large filter, it is clear that the tiny precipitate will be spread over the whole of the large filter, which would make it difficult to work with. Before examining such a precipitate it must be completely freed from all traces of the filtrate. This is accomplished by wash- ing. Washing must be continued until no test can be obtained with the wash-water for a certain substance known to be present in the filtrate. For example, suppose we desire to filter off some suspended barium carbonate from a solution which contains sodium Sulphate; the precipitate must be continuously washed until a sample of the wash-water, acidified with hydrochloric acid, no longer gives a precipitate on the addition of barium chloride. As , a rule, it is not advisable to run the wash-water into the filtrate, because this occasions an unnecessary dilution of the latter. The filter must always be Smaller than the funnel, and the precipitate should not extend higher than to within 5 mm. of the top of the filter. Finally, large precipitates should be avoided as far as possible, for they render exact work more difficult—filtration and washing consuming too much time. Hydrolysis. Hydrolysis is the name given to the decomposing action of water upon many salts. Corresponding to the fact that water is a 194 GENERAL PRINCIPLES. poor conductor of electricity, it follows that water is dissociated only to a slight extent. - t H2O ax H +OH'. According to Kohlrausch and Heidweiler,” the degree of disso- ciation at 25° C. is 1.05×10−7; in other words, 10,000 000 liters of water would furnish 1 gram molecule of dissociated water. Small as this is, it suffices to explain the hydrolysis of: I. The salts of weak acids with strong bases. II. The salts of strong acids with weak bases. III. The salts of weak acids with weak bases. The salts of strong acids with strong bases are not hydrolyzed appreciably. Hydrolysis is shown to take place by the fact that solutions of neutral salts corresponding to I react alkaline, those of II react acid, while those of III are sometimes acid and sometimes alkaline. The cause of hydrolysis is the action of the ions of water upon the ions of the dissolved salt. All monobasic salts in aqueous Solution are largely dissociated into ions (cf. page 14): MSa-M.--S', and the phenomenon of hydrolysis may be represented by the general equation M'4-S-HIH2O]+, MOH+SH. *—y—’ Salt Base Acid According to the mass-action law K= [MOH]X[SH) [M]X[S]X[H2Ol' Since the mass of the water in a dilute solution is practically the same both before and after the reaction has taken place, in the • * Zeitschr. f. phys. Ch., 14, p. 317. REACTIONS OF THE IONS I9b above equation [H2O) may be left out, and then instead of the constant K a new value will be obtained: k, [MQH1XISH) *-TMEs The base MOH and the ac’d SH are subject to electrolytic dis- sociation, and the solution will react acid or basic according as to whether the acid or the base is most readily dissociated, i.e., is the stronger. I. Hydrolysis of Salts of Weak Acids and Strong Bases. The base formed by the hydrolysis is largely dissociated, while the acid is only slightly ionized. The above equation now takes the form _IMIXIOH]x[SH]. Kºś Since, however, the extent of dissociation in the case of strong bases is practically equal to that of the salts (cf. page 14) the value [M] can be cancelled in the numerator and denominator, and we obtain the expression _[OH]X[SH]_[Base|X[Acid] [S] [Salt (1) Kh Owing to the presence of the OH ions in appreciable amount, all salts of this category react alkaline. The alkali salts of hydro- cyanic acid, hydrochlorous acid, carbonic acid, boric acid, and hydrogen sulphide are of this type. II. Hydrolysis of Salts of Strong Acids and Weak Bases. Here the conditions are reversed, and it is the acid which is almost completely dissociated and the base but slightly. In this case the formula becomes [MOH]x[S]x[H]_IMOH]X [H] _[Base|X[Acid (2) Kº–ºff; [M] [Salt] 196 GENERAL PRINCIPLES. Salts of this class, such as those of copper, aluminium, iron, etc., react acid when in aqueous solution. In the case of polyvalent bases both electrolytic and hydrolytic dissociation take place in stages: . . [BCl3 = BC13+Cl’ BCl” ax B"' --Cl’. Hydrolysis BCl” +2H2O air BCl(OH)2+2H' B” +3H2O z-> B(OH)3 +3H'. Ferric, aluminium, chromic, and cupric salts, for example, react acid in aqueous Solution. If such solutions be evaporated to dry- ness considerable hydrochloric acid is volatilized, and the residue obtained is an insoluble basic salt which can only be dissolved by means of acid. III. Hydrolysis of Salts of Weak Acids and Weak Bases. The acids and bases formed by hydrolysis are only slightly dis- sociated, but to different extents. Our formula becomes [Base]X[Acid]_[Base]X[Acid] [M]x[STTTTSaltº (3) Khy, = If the electrolytic dissociation of the acid is greater than that of the base, the solution reacts acid; and conversely, when the base is stronger than the acid the solution of the salt shows an alkaline reaction. Neutral ferric acetate in a boiling, aqueous solution is completely hydrolyzed: Fe(C2H3O2)3 +2H2O z-z Fe(OH)2C2H3O2+2HC2H3O2. Basic ferric acetate is precipitated and can be removed by fil- tering the hot solution. If the solution is allowed to cool, the reaction tends to take place in the reverse direction and some of the basic salt goes into Solution. Heat and dilution always favor hydrolysis. DETECTION OF ACIDS AND BASES. J 94 The hydrolytic action of water, as well as the mass-action law, may be illustrated by the following experiment: A small amount of water is added to a solution of antimony chloride in concentrated hydrochloric acid, causing a heavy precipitation of antimony Oxy- chloride: # -Q!, H =O Sb–Cl–H #Xo = Sb 14-2HCl, —Cl A which again dissolves on adding a little concentrated hydrochloric acid. Further addition of water again precipitates the basic salt, which will dissolve in more of the concentrated acid, etc. It is obvious that by increasing the mass action of the water the re- action goes from left to right, while by increasing the concentration of the hydrochloric acid it goes from right to left. Detection of Acids and Bases.” For the detection of acids and bases it is customary to employ certain substances which change color on coming in contact with one or the other of the above classes of compounds. The acid re- action depends upon the presence of free hydrogen ions, while the alkaline reaction is caused by hydrozyl ions: HClaxcl’ +H" (reacts acid); NaOHz-Na' +OH' (reacts alkaline). For detecting the presence of acids or alkalies in aqueous solu- tions we make use of indicators or substances of organic nature which impart a different color to an alkaline solution than to an acid one. These indicators are, in fact, pseudo acids or bases, being very unstable in the form of free acid or base, and readily undergo a change of constitution, forming closely allied isomeric substances which are not dissociated. The molecule suffers a rearrangement of its atoms; and it has been found that in the case of all colored * The translator has changed the text to conform with the article on The Theories of Indicators, by J. Stieglitz, J. Am. Chem. Soc., XXV, 1112. I96 GENERAL PRINCIPLES. organic compounds the color is due to a particular arrangement of atoms called a chromophor. Methyl orange is an amphoter and is capable of forming salts with both acids and bases, but its indicator characteristics are due to its very weak basic properties. The neutral Solution of its sodium salt is used as an indicator. In this sensitive neutral solution we have a condition of equilibrium between the two isomeric forms of methyl orange as expressed by the equation SO,C.H.N.:NC.H.N(CH3), HOHz-SO.C.H.N.H. N.C.H. :N(CH.),O. H. The formula on the left is yellow in color and its color is due to the azo group N:N, while the other compound is red, having for its chromophor the quinoid group : C6H4:. The sodium salt of methyl orange is yellow and has the formula NaSO3 * C6H4N : N t C6H4NOCH3)2, and when decomposed by acids the free sulphonate sº C6H4NH. N. : C6H4: N(CH3), is formed, which is red. The red quinoid form is ionized as a weak base and forms red salts with acids. It does not form salts readily with weak acids, such as carbonic or acetic acid, because, as we have seen, salts of weak bases and weak acids are hydrolyzed. This is why methyl orange is not a sensitive indicator for weak acids. As a very weak base it will be driven readily out of its red salts by other weak bases; even weak ones and the free base will revert again to its yellow form, the result being that methyl orange is an excellent indicator for weak bases. Phenolphthalein, another valuable indicator, is a very weak acid The free acid, however, is unstable, and when set free from one of its colored salts reverts instantly into a colorless lactoid form: HOOcch.cGººh) :C6H4: O z=2 Qocciºſch,oh) 2. Colorless The red color is in this case also due to the quinoid grouping :C6H4:. In the case of the free acid, the condition of equilibrium HYDROLYSIS. 19ſ favors the lactoid form, and only minimal quantities of the quinoid acid are present. This trace of quinoid acid is ionized and in equilibrium with its ions: HOOCC6H4C(C6H4OH):C6H4:Oz-2 H' +OOCC6H4C(C6H4OH):C6H4:O'. The addition of an alkali causes the hydrogen ions to disappear, more of the quinoid molecules must be ionized to preserve equilib- rium, and the quinoid molecules in turn be reproduced from the lactoid as fast as the former are converted into the salt. Phenol- phthalein is a very sensitive indicator towards acids. but on account of being such a weak acid it does not form stable salts with weak bases. Besides these indicators, others are often employed, among which may be mentioned Litmus and Lackmoid, which are red with acids and blue with alkalies; and Turmeric, which is brown with alkalies and yellow with acids. II. REACTIONS IN THE DRY WAY. These reactions are employed chiefly in the so-called “prelimi- nary examination,” in testing the purity of precipitates, and in the examination of minerals. The most important reactions of this nature consist in the testing of a substance with regard to its— 1. Fusibility; 2. Ability to color the non-luminous Bunsen flame; 3. Volatility; 4. Behavior towards oxidation and reduction. In order to carry out these reactions it is customary to use the non-luminous gas flame; and to understand the operations to be described it is necessary for us to know something about the com- position of illuminating-gas and the nature of the flame. * Illuminating-gas has everywhere a similar composition, dis- regarding certain minor variations, caused by the different kinds of coal that are used, and the temperature at which the gas is prepared. 2 O GENERAL PRINCIPLES. The gas of Zurich averages the following composition: CO2 = 2.0 per cent. C.H.,n = 4.5 “ “ O 0.2 & 4 & & CO = 8.0 “ “ H = 48.0 “ “ CH,-33.0 “ “ N = 4.3 “ “ * 1000 “ “ $ All of these components, except CO, O, N (which are present only in small amounts), are combustible; they are reducing sub- stances. Illuminating-gas ordinarily burns with a luminous flame, and this luminosity is due to the presence of unsaturated hydro- carbons (C.H.,n), principally ethylene, propylene, acetylene, ben- zene, etc. If ethylene is heated to a certain temperature, it is decomposed into methane and carbon: C.H. = CH4+C, and it is this glowing carbon which causes the luminosity of the flame. The other unsaturated hydrocarbons behave like the ethylene. The remaining combustible constituents of illuminating-gas burn with a non-luminous flame. If we bring air into the gas, the flame becomes non-luminous. With the Bunsen burner air is in- troduced by opening the holes at the base of the burner. In such a gas-flame there are, according to Bunsen, the following parts (Fig. 1): I. The inner cone of the flame, aab, in which no combustion takes place, because the temperature here is too low. This part of the flame contains unburned gas mixed with about 62 per cent. of air. II. The flame mantle, indicated by acaba, which is composed of burning gas and air. * III. The luminous tip, at b, which does not appear unless the air-holes are somewhat closed. In these three principal parts of the flame Bunsen distinguishes six reaction-zones: REACTIONS IN THE DRY WAY. 2 I 1. The base of the flame at a'. The temperature here is relatively low, because the burning gas is cooled by the constant current of fresh air, and also because the C burner itself conducts away con- siderable heat. This part of the flame serves to test volatile sub- stances to see whether they impart color to the flame. In case several substances are present which color the flame, it is often possible to observe the colors one after the other, in that the most volatile substance colors the flame first, and later the colors caused by the less volatile ones are seen. This would not be possible at a hotter part of the flame, as all of the substances would then be immedi- ately volatilized, producing a mix- ture of colors. 2. The fusing 20me at 8. This lies at a distance of somewhat more than one third of the height of the flame, and equidistant from the outside and the inside of the mantle, which is broadest at this part. As this is the hottest part of the flame (about 2300° C.), it serves for testing substances as to their fusibility and volatility. 3. The lower oacidizing flame lies in the outer border of the fusing zone at r, and is especially suited for the oxidation of substances dissolved in vitreous fluxes. 4. The upper oasidizing zone, at e, consists of the non-luminous tip of the flame, and acts strongest when the air-holes of the lamp are fully open. It is used for various oxidizing tests, the roast- ing away of volatile products of oxidation, and generally for all processes of oxidation where the very highest temperature is not required. * 5. The lower reducing zone lies at 6, in the inner border of the fusing zone next to the dark cone. As the reducing gases are here 22 GENERAL PRINCIPLES, mixed with oxygen from the air, many substances which are re- duced by the upper reducing flame are here unaffected. This part of the flame is consequently very well adapted for a test which cannot be made with the blowpipe, namely, reduction on the char- coal stick, and in vitreous fluxes. 6. The upper reducing flame is at b, in the luminous tip of the dark inner cone, which may be produced by gradually diminishing the supply of air. If the luminous tip has been made too large, a test-tube or porcelain dish filled with water and placed over it will be blackened, which should never be the case. This luminous tip contains no free oxygen, is rich in Separated incandescent carbon, and has, therefore, a much stronger reducing action than the lower reducing zone. It is used more particularly for the reduction of oxides collected in the form of incrustations. METHODS FOR THE ExAMINATION OF A SUBSTANCE IN THE DRY WAY. 1. Test of the Fusibility. This test is principally made in the examination of minerals, which are introduced into the flame in the loop of a platinum wire (about as thick as a horse-hair). The sample is examined, after heating, by means of a magnifying-glass to see whether the corners are rounded, due to melting. The potentially hottest tem- perature of the fusing zone amounts to about 2300° C.* It will never be possible to reach this temperature with the test, because the substance itself loses heat by radiation. As the amount of heat k)st by radiation is proportional to the surface exposed, it is evident that We will obtain the maximum heat by using a very small sample and holder. For this reason a coarse wire should not be used for this test. We distinguish between the following degrees of heat: 1. Faint red glow... . . . . 525° C. 2 Dark red glow . . . . . . 700° C. Melting-point of aluminium 658° C.# 3. Bright red glow . . . . . 950° C. “ silver. . . . . 955–961° C. O & “ gold... . . . 1064° C. 4. Yellow glow © a e º ºs e tº gº 1100° C. } & 4 “ copper ... 1064–1084°C. 5. Faint white glow..... 1300° C. 6. Full white glow . . . . . 1500° C. & & “ platinum . 1780° C. * This temperature will be considerably lower with too large a supply of air. According to Naumann, the temperature of illuminating-gas with 14 times its volume of air reaches about 1818° C, but the temperature ob- tained is usually lower owing to loss by radiation. The finest platinum wire can be melted by means of the flame, but not when it is as thick as a horse- hair. f Le Chatelier: “High Temperature Measurements.” REACTIONS IN THE DRY WAY. 23 Below 525° C. the following substances melt: tin at 232° C., lead at 327°C., zinc at 420° C. 2. Color Imparted to the Flame. The substance (best in the form of the chloride) is placed in the loop of a fine platinum wire, introduced into the base of the flame, and then finally brought into the fusing zone. 3. Oacidation and Reduction in Vitreous Flua.es or Beads. To make a bead, borax (Na,B,C), 10H2O) or salt of phos- phorus (NaNH4HPO, +4H2O) is used. A piece of very fine platinum wire, about 3 cm. long, is sealed into the end of a glass tube. The wire is heated to redness, and then quickly dipped into the borax or salt of phosphorus, which is held near the flame; whereby a small'amount of the Salt is fused to the end of the wire. By repeated heating and dipping into the salt a bead of sufficient size is obtained. This should be about 1.5 mm. in diameter at the most. It is not advisable to make a loop at the end of the wire, because in this way the exposed surface is unnecessarily increased. There is no danger of the bead falling off, provided the wire is held horizontally in the flame and the bead is not too large. In order to bring the substance in question into the bead, it is only necessary to moisten the latter with the tongue, and then dip it into the finely- powdered substance, which will cause a small amount to adhere to the bead. It is preferable to introduce too little substance into the bead rather than too much, because, in the latter case, the bead will become dark and opaque. The oxidation of the substance in the bead is brought about by heating it in the lower oxidizing flame; while reduction is usually effected by heating in the lower reducing zone, and cooling in the dark inner cone, in order to pre- vent oxidation, which might take place if the substance were cooled in the air. In order to clean the wire, a borax bead is produced on the wire, which is then heated, as is shown in Fig. 2, a, On one side of the bead only, so that the latter runs along the wire in the opposite direction, dissolving off all impurities. By heating the bead from the other side, Fig. 2, b, it is driven toward the end of the wire, from which it can be shaken off by a quick jerk. By repeating this process three times the wire is cleaned with the exception of a Small amount of 24 GENERAL PRINCIPLES. adhering borax-glass, which can be removed by heating the wire in the fusing zone until the sodium flame entirely disappears. 4. Reduction on the Charcoal Stick. These exceedingly beautiful reactions are among the most sensi- tive of those used in analytical chemistry, and should be faithfully practised by every beginner. The cause of their sensitiveness is due to their taking place on the extreme end of a tiny piece of char- coal, that is in a point, so that the sample has no opportunity to FIG, 2. spread itself over a large surface, which is the case with the ordi- nary reactions on charcoal before the blowpipe. In order to carry out these reactions, we use an ordinary round match (not a four-cornered one), which should consist of good straight wood-fibers. It is impregnated with sodium car- bonate in the following manner: A crystal of sodium carbonate (Na,COs. 10H2O) is warmed in the flame, whereby it melts in a part of its water of crystallization. Three fourths of the length of the match is now smeared with this liquid soda, and the match is then slowly rotated on its axis in the flame, until the soda melts and penetrates the charcoal. On withdrawal from the flame there should be no place which continues to glow; should the latter be the case, the stick should be quickly immersed in the soda again. In this way one obtains a solid little piece of charcoal, which can be heated for a long time without burning through. In order to carry out a reduction, a small amount of the body to be examined is mixed on the palm of the hand with an equal REACTIONS IN THE DRY WAY. 25 amount of calcined soda, a small drop of melted soda is added, and the mixture is made into a paste by means of the blade of a pen- knife. The warmed piece of charcoal is then rubbed into the mix- ture, which adheres to it. The sample is first heated in the lower oxidizing flame until it has melted, and then moved into the lower reducing flame. The reduction will be made evident by a violent swelling up of the melt, caused by the evolution of carbon dioxide. As soon as the mass melts quietly the reduction is complete. The substance is allowed to cool in the dark cone, after which it is re- moved from the flame. The metal is now found on the extreme end of the carbonized match, concentrated in a point. This point is broken off, and triturated with a small amount of water in an agate mortar. The excess of sodium carbonate goes into Solution, , part of the charcoal floats on the surface of the water, while the heavier metal sinks to the bottom. In case the reduced metal is iron, nickel, or cobalt, it will not be noticeable to the eye, but it may be taken up with a magnetized knife-blade, to which it will adhere, usually mixed with charcoal. This should be dried by cautious warming, the tuft of metal taken off and rubbed between the thumb and forefinger, and then brought into contact with the knife again, to which only the metal will now adhere. The metal is then transferred to a piece of washed filter-paper about 3–4 mm. wide and 50 mm. long, so that it comes as near as possible to the end of the strip. By means of a capillary tube, a drop of hydrochloric acid and one of nitric acid are added, and the paper is warmed over the flame until the black speck (the metal) has disappeared, when the final test can be made. In order to test for iron, a drop of potassium ferrocyanide is added, whereby the presence of iron is shown by the appearance of a distinct formation of Prussian blue. To test for nickel and cobalt, the metal is dissolved in nitric acid, the excess of acid is evaporated off, and a drop of concentrated hydrochloric acid added, whereby the paper is colored blue if cobalt is present; the nickel shows at the most only a very weak greenish color—usually, however, no color. A little caustic soda solution is now added, and the paper held in the vapors of bromine; in case either nickel or cobalt is present a brownish-black spot appears, due to the formation of either Ni(OH)2 or Co(OH)s. If, however, the metal reduced was malleable, it is usually ob- 26 GENERAL PRINCIPLES. tained in the form of a metallic globule on the end of the match, where it can be examined with the aid of a lens. Copper is not separated as a globule, but as a reddish, sintered mass. By pressing down on malleable metals in the agate mortar they are obtained as a glistening fragment, which can be readily separated from the specifically lighter charcoal by washing. To accomplish this the agate mortar is inclined and a stream of water is directed sideways upon the mass, whereby the charcoal is washed out with the water, and the metal is left clean. It is transferred to a watch-glass and tested as follows: 1. The Metal is White (Pb, Sn, Ag, Pt). The metal is treated with a few drops of nitric acid and carefully warmed. Lead and silver dissolve readily, particularly upon addition of a little water. Silver will be detected by the addition of a drop of hydrochloric acid, whereby white silver chloride, soluble in ammonia, is precipi- tated. The test for lead is dilute sulphuric acid, which precipitates white lead sulphate. If the metal, on treatment with nitric acid, remains unchanged, it is probably platinum. It should be dissolved in aqua regia, evaporated to dryness, dissolved in a little water, and potassium chloride solution added. A yellow, crystalline precipitate confirms the presence of platinum. If the metal, when treated with nitric acid, becomes changed into a white insoluble oxide, it is tin. In this case, another fragment of metal is dissolved in concentrated hydrochloric acid and tested for tin by means of mercuric chloride solution, or bismuth oxide and caustic Soda. 2. The Metal is Yellow to Red (Cu, Au). Copper is readily dissolved in nitric acid, and the solution gives with potassium ferro- cyanide a reddish-brown precipitate. Gold is insoluble in nitric acid, but soluble in aqua regia. The evaporated solution gives a violet-brown color with stannous chloride, due to finely-divided gold. Reduction in a Glass Tube. Besides the borax bead and the charcoal stick, reduction is often effected by means of metallic sodium, potassium, or magnesium. Thus small amounts of phosphorus in anhydrous salts may be detected in the following manner: The Substance to be tested is placed in a glass tube, 3 mm. wide and 50 mm. long, which is closed REACTIONS IN THE DRY WAY. 27 at one end. A small cylinder of potassium or sodium (freed from petroleum by rubbing between filter-paper), or even a piece of magnesium wire, is added to the tube, and the contents then heated until the glass itself begins to soften. The reaction is so violent that the substance seems to take fire. After cooling, the tube is broken in a porcelain mortar, when, by breathing over the mass, the smell of phosphoretted hydrogen may be detected. The halogens, Sulphur, and nitrogen are tested for in a similar way, as will be shown later. Reduction in the Upper Reducing Flame for the Purpose of forming Metallic and Oxide Incrustations. The volatile elements which are reducible by means of hydrogen or charcoal may be detected in this part of the flame with the great- est ease, as, for example, arsenic, antimony, cadmium, bismuth, selenium, and tellurium. The metallic incrustations are obtained by holding in one hand a small portion of the substance on a thin asbestos thread (platinum will be attacked) in the upper reducing zone of a small gas-flame, where the oxide is reduced to volatile metal, and burned in the upper oxidizing flame to oxide. In the other hand, closely over the substance to be tested, is held a glazed porcelain evaporating-dish, filled with water, as is indicated in Fig. 3 at B. The metallic vapors are condensed by the cold dish, and deposited on it in the form of a metallic mirror or film. If, however, the dish is held above the upper oxi- A dizing flame (at A), there is formed a /\ thin, often invisible, oxide incrustation On the bottom. Should it be necessary to treat the metallic incrustation with a large amount of solvent (as is necessary in the detec- tion of Selenium and tellurium), the porcelain dish is replaced by a test-tube half filled with cold water. A some- what larger test-tube is used to hold the solvent, and the Smaller test-tube, on FIG. 3. which the incrustation was deposited, is placed within the larger tube and the liquid warmed if necessary. 28 GENERAL PRINCIPLES. Blowpipe Reduction on Charcoal. These tests are made in the so-called “preliminary examination.” For this purpose a small cavity is made with a penknife in a piece of good charcoal (preferably of linden wood), in which a knife- bladeful of the substance to be tested is placed, previously mixed with twice as much anhydrous sodium carbonate. As charcoal is a porous body, it will readily absorb easily fusible substances, like the alkalies. Other substances are changed, by means of the sodium carbonate used, into carbonates, which are, for the most part, de- composed, on heating, into oxides and carbon dioxide. The oxides of the noble metals are decomposed, without the aid of the charcoal, into oxygen and metal; while those of the remaining metals are either reduced to metal or remain unchanged. .Thus CuO, PbO, Bi,O, Sb,0s, SnO, Fe,0s, NiO, and CoO are reduced either to a fused metallic globule (Pb, Bi, Sb, Sn, Ag, and Au), or to a sintered mass of metal (Cu), or to a glistening metallic fragment (Fe, Ni, Co, Pt). The oxides of zinc, cadmium, and arsenic do not give metallic globules, but are, however, easily reduced to metal. These metals are so volatile that they are changed into vapors, and are carried from the reducing zone of the flame into the oxidizing zone, where they are changed into difficultly-volatile oxides. These oxides, which have characteristic colors, are then deposited on the charcoal outside the cavity. Zinc gives an incrustation which is yellow while hot, and white when cold; that of cadmium is brown; while the oxide of arsenic gives a white and readily volatile incrustation. Furthermore, the volatilization of arsenic gives rise to a characteristic garlic odor. The metals lead, bismuth, and tin give, besides the metallic globule, an oxide incrustation which is typical. At the same time, nitrates, nitrites, chlorates, etc., may be recognized by the fact that they cause a very rapid combustion of the glowing charcoal (deflagration). This deflagration is not to be confused with a decrepitation which takes place on heating sub- stances containing enclosed moisture or gases, such as rock salt, fluor-spar, etc. Crystals of such substances are burst by the quick expansion of the enclosed liquid, and scattered about. Many difficultly-fusible substances do not melt into the charcoal. Thus many silicates form a bead with the soda, which only after DIVISION OF THE METALS INTO GROUPS. 29 continuous heating will give up the alkali and allow it to be absorbed by the charcoal, leaving behind the white infusible silica. Phos- phates and borates act similarly, only these do not leave behind an oxide, but a fused glass. Infusible white oxides, as those of calcium, strontium, magnesium, aluminium, and many of the rare earths (Welsbach mantle, for example), glow very brightly, and in fact more brightly as they are more strongly heated. Division of the Metals into Groups. We may divide the metals, for purposes of analytical chemistry, into five groups: The First Group contains those metals whose chlorides are in- soluble, or difficultly soluble, and whose sulphides are insoluble, in dilute acids. They may, therefore, be precipitated from their solu- tions by means of either hydrochloric acid or hydrogen sulphide. The Second Group contains those metals whose chlorides are soluble, but whose sulphides are insoluble in dilute acids. They may be precipitated from their solutions by means of hydrogen sulphide, but not with hydrochloric acid. The Third Group contains those metals whose sulphides are solu- ble in dilute acids, but are insoluble in water and alkalies; and also those metals whose Sulphides are hydrolytically decomposed into hydrogen sulphide and metallic hydroxide. The members of this group are precipitated completely by hydrogen sulphide only from alkaline solutions. The Fourth Group contains those metals whose sulphides are soluble in water, but whose carbonates are insoluble in the presence of ammonium chloride. They are precipitated by ammonium carbonate in the presence of ammonium chloride, but not by any of the above reagents. The Fifth Group contains magnesium and the alkalies; they are not precipitated by any of the above reagents. In order to carry out an analysis with certainty it is necessary to understand not only the reactions of the different elements, but we must know as well the sensitiveness of each reaction. The analyst should be able to draw a conclusion by the size of the pre- cipitate formed as to the approximate amount which is present in 3ö' * T GENERAL PRINCIPLES. the original substance. This, however, is possible only when the experiments are made with known amounts. Consequently we take reagents of a known strength, and allow them to act on known, amounts of the different substances. According to the suggestion of R. Blochmann (B. B. 1890, p. 31) it is well to make the solutions of the different reagents either double-normal, normal, half-normal, or tenth-normal. For the last eight years the author has used in his laboratory solutions of reagents and Salts according to this principle, and has found that the beginner in this way gets a far better under- standing of the stoichiometrical relations than when solutions of almost any concentration are used, as was formerly the custom. By a normal solution we understand one which contains in one liter one gram-equivalent of the substance in question, referred to a gram atom of hydrogen as a unit. A tenth-normal solution will contain one tenth of a gram-equivalent in a liter, etc. Thus one liter of a normal solution will contain HCl = 36.46 grams, equivalent to one gram-atom of hydrogen. H,SO, 98.08 ſº t –3– = -g- - 49.04 grams, equivalent to one gram-atom of hydrogen. H.PO, =} 98.03 = 32.68 ( & { { ( & & & {{ r & & & ſ 3 3 © NaOH :- 40 06 & 4 {{ {{ {{ f{ {{ & & KMnO, : 158. 11 = 31.62 “ ( & ( & & & €6. €6 &&. 5 5 g R2Cr, O, s== 294.4 = 49.08 “ & & 4.6 ft & & 66 Čg 6 6 e The great advantage of this system is that one always knows how much of one solution should be used in order to react with another quantitatively. Thus 1 cc. of a normal caustic Soda Solution will neutralize 1 cc. of any normal acid, or 2 cc. of any half-normal acid. In the same way 1 cc. of a normal Solution of Sulphuric acid, or of any sulphate, will precipitate quantitatively the barium from 1 cc, of a normal barium chloride Solution. THE MOST COMMON LABORATORY REAGENTS. I, CoNCENTRATED ACIDS. Sp, Gr. Per Cent Acid by Weight. z. Hydrochloric acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1, 189 37.9 2. Nitric acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1. 386 62.64 3 Sulphuric acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1,840 96.0 LABORATORY REAGENTS. 3:La II. DILUTE ACIDs, # NORMAL. 1000 cc. contain: 1. Hydrochloric acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 72.92 grams 2. Nitric acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 126. 10 “ 3. Sulphuric acid... . . . . . . . . . . . . . . . ". . . . . . . . . . . . . . . . . . . . . . 98.08 “ 4. Acetic acid... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 120.08 “ 5. Tartaric acid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 150.06 “ . III ALKALIES, # NORMAL. 1000 cc. contain: 1. Sodium hydroxide. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 80.12 grams 2. Potassium hydroxide. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 112.32 “ 3. Ammonium hydroxide. . . . . . . . . . . . . . . . . . . . . . . . . ... • * * * * * 70. 18 “ 4. Conc. Ammonia (sp. gr. 0.905) contains 27% NH3. IV. SALTs. (a) # Normal. 1000 cc. contain: 1. Ammonium carbonate. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 96.16 grams 2. Ammonium chloride . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 107.06 “ 3. Sodium carbonate. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 106. 10 “ 4. Ammonium sulphide. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 68.22 * ** (b) Normal 1000 cc. contain: 1. Sodium acetate (NaC, H, O, + 3H2O)... . . . . . . . . . . . . . . . . . 136. 14 grams 2. Sodium phosphate (Na2HPO, + 12H,O) sº e tº e º e º ºs e e 119.45 “ © * 74.5 3. Sodium hypochlorite (NaOCl) T2 . . . . . . . . . . . . . . . . . . . . 37.25 “ 4. Sodium nitrite (NaNO) *........................ 34.54 ** 5. Potassium nitrite (KNO) *....................... 42. 59 “ sº º 294.5 6. Potassium bichromate (K,Cr2O.) 6 . . . . . . . . . . . . . . . . 49.08 “ 7. Calcium chloride (CaCl, 4 GH,0) *.*............... 109.51 “ 8. Magnesium sulphºtoſso, + 7.H.O.) * e - G - tº dº º tº e º & 123.28 “ 9. Barium chloride (BaCl, + 2H2O) ** º º ºs e º 'º e º 'º e º is g º 122. 17 “ 10. Ferrie chloride (FeCl) “......................... 54. 11 “ * This reagent decomposes slowly, losing NH, and forming (NH4)2S,. 32 GENERAL PRINCIPLES, 11. 12. 13. 14. 15. 10. 11. 12. 13. 14. 15. 16. . Ammonium oxalate ((NHA)2C2O, -- H2O) . Mercuric chloride (HgCl2) . Sodium thiosulphate (Na2S,O3 + 5H2O) . Sodium bromide (NaBr + 2.H2O) Potassium ferrocyanide (KFe(CN), +3H,0) *..... Lead acetate (Pb(C.H.O.), +3H,0) *............ Stannous chloride (SnCl2 + 2H2O) ** * * g e º e g tº ſº e º e º e Mercurous nitrate (Hg2(NO3)2) * * e º e º e º e º 'º º ſº tº e º 'º º º 291.2 2 . . . . . . . . . . . . . . . Cobalt nitrate (Co(NO3)2 + 6H2O) (c) # Normal. 1000 cc. contain: 142.18 4 . . . . . . . . 271.2 4 * @ e º is tº e º 'º e º ſº tº e º tº tº ſº tº ſº tº º 248.32 2 e e º ſº e º ſº tº º & 139.05 2 . . . . . . . . . . . . . . . . or (NaBr) lºgº tº º e º e º e º e º 'º e º 'º e º e º ºs e º 'º - º º ſº º ſº e º 'º $ tº tº 65.19 . Potassium cyanide (KCN) -2- . . . . . . . . . . . . . . . . . . . . . . . 166 . Potassium iodide (KI) -3-... . . . . . . . . . . . . . . . . . . . . . . . . . 97.25 . Potassium sulphocyanate (KCNS) -2-.. . . . . . . . . . . . . . . 180.17 . Potassium arseniate (KH2AsO.) -a-. . . . . . . . . . . . . . . . . 287.6 . Zinc sulphate (ZnSO, -- 7B,O) TT . . . . . . . . . . . . . . . . . . . Manganese sulphate (MnSO, + 4H2O) *. tº is dº e º e º 'º e º Nickel sulphate (NiSO, F THoº'................. Copper sulphate (CuSO, + 5H2O) º tº e º e º 'º e º e º 'º ſº e º 'º º Chrome alum (CrR(SO4)2 + 12H2O) * tº e º 'º gº tº gº tº e º ſº e º º Alum (AlK(SO,), + 12H2O) tººl º, º ºs º ºs e º 'º º e º 'º, º ſº e º ſº ſº tº e e 484.72 Bismuth nitrate (Bi(NO)a-H 5H,O) 6- . . . . . . . . . . . . . . . 66 6& &&. 4 & (& 66 & 66 €6. &&. (6. €6. 66 DETERMINATION OF THE SENSITIVENESS OF REACTIONS. 33 (d) ºn Normal. 1000 cc, contain: 425.60 1. Uranyl acetate (UO,(C.H.O.)2 + 2H2O) -20-..... . . . . . . . 21.28 grams 2. Silver nitrate (AgNO3) * as e e s e s a e s s e e s e s a tº e o e º e º 'º º & 8 17.0 “ W. SATURATED SOLUTIONS. 100 grams solution contain at 15° C.: 1. Hydrogen sulphide... . . . . . . . . . . . . . . . . . 0.48 grams H.S 2. Barium hydroxide. . . . . . . . . . . . . . . . . . . . 5.95 “ Ba(OH)2 + 8BI,O 3. Lime water. . . . . . . . . . . . . . . . . . . . . . . . . . 0.13 “ CaO 4. Calcium sulphate. . . . . . . . . . . . . . . . . . . . . 0.26 “ CaSO, + 2H2O 5. Chlorine water ... . . . . . . . . . . . . . . . . . . . . 0.73 “ Cl 6, Bromine water. . . . . . . . . . . . . . . . . . . . . . . 3.66 “ Br In the author's laboratory these solutions are prepared in five- liter amounts. In order to do this as quickly as possible, and with sufficient accuracy for qualitative purposes, flasks are used which contain five liters when filled up to a mark on the neck. Similarly, in weighing out the substance, tares are used containing shot, which correspond to the weight of substance which should be dissolved in five liters. In the preparation of the dilute acids and ammonia, Small cylin- ders are used, which are so etched that the correct amount of com- mercial acid can be measured out therein, and this is then diluted to five liters. Thus 278 cc. conc. sulphuric acid will make 5 liters 2 N acid. 726 cc. “ nitric “ “ “ 5 tº 2 N. “ 809 cc. “ hydrochloric “ “ “ 5 “ 2 N. “ 1922 cc. 30% acetic “ “ “ 5 tº 2 N. “ 698.5 cc. conc. ammonia “ “ 5 “ 2 N. “ Determination of the Sensitiveness of Reactions. The more sensitive a reaction is, the Smaller will be the amount of the substance which can be detected in a given volume, in a definite time, with the reagent in question. Let us suppose that the amount of substance taken is dissolved in 100 cc. of liquid the time allowed to be two or three minutes, and the limit of sensitive- ness to be the Smallest amount of substance which can be detected under these conditions. 34 * GENERAL PRINCIPLES. A few examples will make the method clear: Magnesium salts are precipitated by means of Sodium phosphate, in the presence of ammonium chloride and ammonia, in the form of magnesium ammonium phosphate. What is the Sensitiveness of this reaction? We take 1 cc. of our normal magnesium sulphate Solution, add three drops of ammonium chloride Solution, and two or three drops of ammonia and sodium phosphate solution; the characteristic white precipitate is formed immediately. We dilute, now, the normal solution of magnesium sulphate ten times, and repeat the experiment with 1 cc. of the diluted solution. The result will be— 1 cc. of N. Mg solution, 100 cc. = 1.2 grm. Mg, reacts immediately. N 1 cc. “ IO Mg {{ 100 cc. = 0. 12 “ Mg, “ {{ N ** {{ := & 4 & 4 & & 1 cc. 100 Mg 100 cc. = 0.012 Mg, 1 cc. “ N Mg “ 100 cc. = 0.0012 “ Mg, “ after a few sec- 1000 Onds. 1 CC “ 10000 Mg “ 100 cc. = 0.00012 “ Mg, “ after one to two minutes. If, therefore, 100 cc. of a solution contain 0.00012 gram Mg, the magnesium can be detected within from one to two minutes. Should the detection of smaller amounts be desired, the Solution must be concentrated by evaporation. This reaction can be called very sensitive. The following potas- sium reactions are much less delicate: (a) Reaction with Hydrochlorplatinic Acid (page 37). 1 cc. of s KCl solution, 100 cc. = 0.78 gm. K, reacts with a drop of H.PtCld immediately. 1 cc. of ; KC solution, 100 cc. =0.078 gm. K, does not cause precipitation within three minutes. 1 cc. of . RC1 solution, 100 cc. =0.156 gm. K, does not cause precipitation within three minutes; but does, however, on addition of two drops of alcohol. 1 cc. off. KCI solution 100 ce.-0284 gm. K. reacts in- mediately on stirring. DETERMINATION OF THE SENSITIVENESS OF REACTIONS, 35 The Sensitiveness of the reaction lies, therefore, between 0.156 and 0.234 gm. K. per cc. In order to detect smaller amounts of potassium than 0.156–0.234 gm. per 100 cc., the solution must be strongly concentrated by evaporation. (b) Reaction with Tartaric Acid (page 38). 1 cc. of . KCI solution, 100 ce-07s gm. K, reacts imme- diately with two drops of sodium acetate and two drops of a con- centrated solution of tartaric acid. 1 cc. of ; KCl solution, 100 cc. = 0.078 gm. K, reacts after one to two minutes with vigorous shaking. This can be taken as the limit of sensitiveness. If the beginner will test the delicacy of all reactions in this way, he will quickly get a clear insight into the solubility relations of the different salts. REACTIONS OF THE METALS (CATIONS). WE begin with the metals of the Fifth Group, because a knowl- edge of their reactions is necessary in order to understand the re- actions of the other groups. GROUP V. ALKALIES. POTASSIUM, K; SODIUM, Na; AMMONIUM, NH, ; (CAESIUM, Cs; RUBIDIUM, Rb; LITHIUM, Li). The metals of this group decompose water at the ordinary tem- perature with evolution of hydrogen and formation of strongly alkaline-reacting hydroxides, which cannot be freed from their hydrate water even on fusion. The pure oxides (R20) are difficult to prepare, cautious heating of the metal in the air forming per- oxides at the same time. The salts of these metals are colorless for the most part, and readily soluble in water. Of these salts the car- bonates, the tertiary and Secondary phosphates, the cyanides, and the borates react alkaline in aqueous Solution (hydrolysis). The salts of the alkalies are more or less volatile and impart to the non-luminous flame characteristic colors. POTASSIUM, K. At. Wt. 39.15. Sp. Gr. 0.87. M. Pt. (Melting-point) 62.5°C. Occurrence.—Sylvite (KCl), isometric, and carnallite (MgCl, RCl·6H,O) at Stassfurt in the presence of halite and anhydrite. Saltpetre (KNO.), orthorhombic prisms. Further, in very many silicates, e.g., monoclinic feldspar (KAlSigOs), and muscovite (KH,Al,SigO2); also in plants in the form of Organic Salts, which yield on combustion potassium carbonate (potash), REACTIONS IN THE WET WAY. Potassium forms very few salts that are difficultly soluble in water. The chlorplatinate, acid tartrate, and perchlorate are the 36 POTASSIUM. 37 least soluble, and are consequently used in the detection of potas- sium. I. Hydrochlorplatinic Acid + (H.PtCl) gives in concentrated solutions a yellow precipitate, H.PtCl4-2KCl=2HC1+K.PtCls, which consists of small regular octahedra (visible with a magnifying- glass). In case the potassium solution is not very concentrated, no precipitation may appear at first; but on rubbing the sides of the beaker or test-tube with a glass rod the formation of the pre- cipitate will be hastened. This is always the case when a crystalline precipitate is formed. The solution before the precipitate separates out is supersaturated, and the formation of crystals is hastened by the mechanical shock. The behavior of the potassium chlorplatinate on ignition is very characteristic; it is decomposed into chlorine, inum, and potassium chloride: R.PtCl,-2KC1+Pt-H2Cl2. º If the products of ignition are treated with water, and the platinum filtered off, the filtrate will again give with hydrochlor- platinic acid the yellow crystalline precipitate of K.PtCls. (Note difference from ammonium chlorplatinate.) Solubility of the Potassium Chlorplatinate in Water. 100 parts of water dissolve: At 0°C. . . . . . . . . . . . . . . . . . . . . . 0.70 parts K.PtCl, “ 10°C. . . . . . . . . . . . . . . . . . . . . . 0.90 “ & & “ 20°C. . . . . . . . . . . . . . . . . . . . . . 1.12 “ { { “100° C. . . . . . . . . . . . . . . . . . . . . . 5. 18 “ ( & In a saturated KCl solution, or in 75 per cent alcohol, the pre- cipitate is practically insoluble. For this reaction it is best to use the chloride. The addition of hydrochlorplatinic acid to potassium iodide solution causes a deep- * Platinic chloride, PtCl, gives no precipitate with potassium salts, or at least only after long standing. The above reagent, hydrochlorplatinic acid, is a dibasic acid, and is obtained by dissolving platinum in aqua regia. The solution is prepared of such a strength that there are ten grams of platinum in every 100 cc. 38 REACTIONS OF THE METALS. reddish-brown color due to the formation of the soluble Salt R2Pt.I. H2PtCl6+8KI=6|KCl +K2Pt.I6+2HI. Similarly potassium cyanide is not precipitated by hydrochlor- platinic acid owing to the formation of complex platinum-cyanogen compounds. In case we desire to test an iodide or cyanide for potassium, the salt should first be changed to chloride by evaporation with concentrated hydrochloric acid. 2. Tartaric Acid (H2C4H4O6) produces, in not too dilute neutral Solutions of potassium salts, a white crystalline precipitate of potas- sium acid tartrate (orthorhombic, hemihedral): soon COOK | CHOH CHOH | +KCl=HCl·H gºon gºon COOH COOH The potassium acid tartrate is readily soluble in mineral acids, but difficultly soluble in acetic acid and water; 100 parts of water at 10° C. dissolve 0.425 gm. of this salt. If sodium acetate is added to the solution, the mineral acid set free by the reaction will be re- placed by acetic acid: CH, CH, | +HCl= | + NaCl, COONa COOH Sodium acetate Acetic acid whereby the reaction is made much more sensitive. Rubbing the sides of the dish will also hasten the formation of the precipitate. It is not advisable to neutralize the mineral acid with caustic soda; for if the latter is added till the solution is neutral, Sodium-potassium tartrate (Rochelle Salt) will be formed, KHC, H,Os-i-NaOH = KNaC.H.O.--H.O, which, like all neutral tartrates, is readily soluble in water; so that no precipitation will take place. On igniting the acid potassium tartrate, empyreumatic vapors (Smelling like burnt Sugar) are given off, and carbon and potassium carbonate are left behind. If the POTASSIUM. 39 mass is now moistened with hydrochloric acid, it will froth strongly. This is a property not only of potassium tartrate, but is common to all salts of organic acids. On ignition they are changed into carbo- nates, and when the acid is non-volatile, carbonization takes place; but with volatile acids there is at the most only a slight carbon- ization. With many metals the carbonate is not left unchanged; frequently it is broken up into carbon dioxide and oxide of the metal; while in the case of reducible metals, the metal itself is left with the carbon. Thus sodium acetate will yield sodium carbonate and acetone, with only a slight carbonization: CH, 4– CH, COONaj | ; j= Na2CO3 + CO. COONa | | - CH, & CH, *śl, On gentle ignition, calcium oxalate yields calcium carbonate and carbon monoxide; the latter burns with a blue flame. GOON | XCa =CaCO,--CO. COO On strong ignition, the calcium carbonate is decomposed into lime and carbon dioxide: CaCO3=CaO + CO,. Tartrates of lead, iron, and many other metals leave behind on ignition carbon and metal. . Bismuth-Sodium. Thiosulphate (Carnot's * reaction).—If to a drop of dilute bismuth nitrate solution a few drops of sodium thiosulphate solution and 10–15 cc. of absolute alcohol are added (turbidity being cleared by the cautious addition of water), this solution will cause a precipitation of yellow potassium-bismuth thiosulphate when added to a solution of a potassium salt: (a) Bi(NO3)a-i-3Na2S,Os-8NaNO3+[Bi(S.O.), Nas; Soluble in alcohol (b) [Bi(S.O.), Nas--3KCl=3NaCl-HBi(S.O.), K. Insoluble in alcohol The presence of ammonium chloride prevents the reaction.f * Zeitschr. f. anal. Ch., 1897, p. 512. t Private communication from W. Wislicenus. 40 REACTIONS OF THE METALS. 4. Hydrofluosilicic Acid (H.SiF6), added in considerable excess to a solution of a potassium Salt, precipitates gelatinous potassium silicofluoride, H.SiF.--2KCl=2HC1-i-K,SiF, which is difficultly soluble in water and dilute acids and insoluble in alcohol. On heating, it is decomposed into volatile silicon fluoride, and potassium fluoride remains behind: R,SiF =2KF--SiF. 5. Perchloric Acid (HClO.) precipitates white crystalline potas- sium perchlorate, HClO,--KCl–HCl·H KClO,; 100 parts of water dissolve at 0° 0.07 parts KClO4, and at 100° 19.8 parts KClO4. REACTIONS IN THE DRY WAY. Potassium compounds color the non-luminous flame violet. The presence of very small amounts of Sodium conceals the violet potassium flame. If, however, the flame is viewed through cobalt glass or indigo solution, the reddish-violet potassium rays pass through, while the yellow Sodium rays are completely absorbed. Spectrum.—Potassium gives a characteristic spectrum. In the red part of the spectrum there is a very bright line, and in the ultra violet there are two faint lines (see chart). SODIUM, Na. At. Wt. 23.05. Sp. Gr. 0.97. M. Pt. 95.6° C. Occurrence.—Sodium occurs very extensively in nature. Its most important mineral is halite, rock salt (NaCl), isometric system. It is found in very large deposits often quite pure, but usually con- taminated with clay, anhydrite, and gypsum, and is present in large amounts in the ocean, and in many Salt Springs. Sodium also occurs in nature in the form of carbonate, as thermonatrite (Na,CO, H.O), orthorhombic; natrite or soda (Na,CO, 10H2O), monoclinic; trona (Na,COs. NaHCOs' 2H2O), monoclinic; as nitrate in Chili saltpetre, or Soda nitre (NaNO3), hexagonal, rhombohedral; as cryolite (Nas/Alfe), triclinic; in many silicates as albite (NaAlSiOs), triclinic; and as tinkal, borax (Na,B,C), 10H2O), monoclinic. sodium. ' 4 I REACTIONS IN THE WET WAY. 1. Potassium Pyroantimonate * (K.H.Sb2O.) produces inneutral or weakly alkaline solutions of sodium salts a heavy, white, crystal- line precipitate, which is formed more quickly by rubbing the sides of the vessel with a glass rod: K.H.Sb,O,--2NaCl-Na,H,Sb,O,--2KCl. The test must not be made in an acid solution, for in that case an amorphous precipitate of pyroantimonic acid will be formed: K.H.Sb,O,--2HCl= H,Sb2O,--2KCl. Furthermore, no other metals than the alkalies should be present, because they also cause precipitates—amorphous ones for the most part. 2. Tartaric Acid and Hydrochlorplatinic Acid do not pre- cipitate sodium salts, the sodium salts of these acids being readily soluble in water. Sodium chlorplatinate is of an orange color, and is readily soluble in 75 per cent alcohol (note difference from potassium). Sodium Peroxide, Na,0,. This substance, which is now used commercially on account of its energetic Oxidizing property, is obtained by burning dry Sodium in the air as a heavy, yellow powder, and shows the fol- lowing characteristic reactions: Behavior to Water.—If a little water is added to some of this substance in a test-tube, a considerable amount of heat will be set free, and oxygen gas will be given off (sufficient to ignite a glowing splintert). Water decomposes the sodium peroxide, according to the equation Na,O,--2H,O=2NaOH-i-H.O. But on account of the heat of the reaction a part of the hydrogen peroxide is decomposed into water and oxygen. If warming up of the Solution is avoided (which can be done * For the preparation of this reagent see page 208. f This will sometimes cause an explosion, because the commercial prod- uct often contains metallic sodium, which with water forms hydrogen; so that both hydrogen and oxygen are set free at the same time. The glowing splinter may then cause an explosion. (Private communication from E. Constam.) 42 REACTIONS OF THE METALS. by throwing the sodium peroxide in Small portions into ice-water), it will dissolve, with scarcely any evolution of oxygen, to a clear, strongly alkaline liquid, which gives, as before, all the reactions of hydrogen peroxide. t If some sodium peroxide is placed on a watch-glass under a bell- jar and near an evaporating-dish containing water, the sodium peroxide in twelve hours will completely change over to a pure white hydrate (Na2O,--2H2O), which will dissolve in water without decomposition at the ordinary temperature. Reactions of Hydrogen Peroxide. (a) In Acid Solution. If the solution obtained by the action of water on sodium per- oxide is used for these tests, it must be acidified with dilute sul- phuric acid, care being taken to keep the solution cool. 1. Titanium Sulphate gives a distinct yellow color, caused by the formation of pertitanic acid. This is the most delicate test for hydrogen peroxide. The titanium sulphate solution for this reac- tion may be prepared by fusing one part of commercial titanium dioxide with 15–20 parts of potassium pyrosulphate and dissolving the fusion, after cooling, in cold dilute sulphuric acid. 2. Chromic Acid.—If the acid solution of hydrogen peroxide is shaken with a little ether (free from alcohol) and a trace of potas- sium dichromate is added, after which the mixture is again shaken, the upper layer of ethereal solution will be colored a beautiful blue, owing to the formation of chromium peroxide. 3. Permanganic Acid in acid solution will be decolorized, with evolution of oxygen: 2KMnO,--4H,SO,--5H,O, =2|KHSO,--2MnSO,--8H,O+5O,. Similar to the permanganate, many other oxides are reduced by hydrogen peroxide, with evolution of Oxygen; e.g., Ag,O, Pb2O, PbO2, MnO2, etc.: Ag,O+H.O. = H,0+O2+ Aga, PbO,--H,O, = H,0+O,--PbO. 4. Potassium Ferricyanide and Ferric Chloride.—If a trace of potassium ferricyanide is added to a very dilute and almost neutral SODIUM. 43', solution of ferric chloride, so that the solution appears a distinct yellow, and an almost neutral solution of hydrogen peroxide is then, added, the mixture will soon assume a green tint, and finally, on standing, Prussian blue will separate out. The potassium ferri- cyanide is reduced by the hydrogen peroxide to potassium ferro- cyanide, which forms with the ferric chloride Prussian blue: 2K, Fe(CN)2 + H2O,-2K.HFe(CN), --O, —and 3.K.HFe(CN)6-1-4FeCla-FeſRe(CN)6], H-9KC1--3HC1. According to Schönbein (J. f. pr. Ch. 79, p. 67, 1860) the smallest traces of hydrogen peroxide may be detected by this reaction (+$g mg. H2O, per liter). As, however, many other Substances (SnCl2, SO, etc.) will reduce potassium ferricyanide to potassium ferrocyanide, this reaction alone cannot always be relied on. 5 Starch Paste and Potassium Iodide.—If to an acid solution containing starch paste and potassium iodide Some hydrogen per- oxide is added, a blue color will at Once appear: 2KI+H,O. =2|KOH-i-I,. By means of this reaction, rāg mg. per liter of hydrogen peroxide may be detected. (b) In Alkaline Solution. 1. Gold Chloride will be reduced, by means of hydrogen per- oxide at ordinary temperatures, to metal, with evolution of Oxygen. The gold usually separates in a very finely-divided state, and ap- pears brown by reflected light and greenish blue by transmitted light: 2AuCls--3H,0,--6NaOH=6NaCl-H6H,O+3O,--2Au. If very dilute gold solutions are used, the metal sometimes separates out in the form of a yellowish film adhering to the sides of the test-tube. 2. Salts of Manganese, Nickel, and Cobalt give dark-colored precipitates: (a) MnOl,+2KOH+H.O. =2KC1-i-H,O+ Mººi. ; (b) 2NiCl, H-4KOH+H,O, =4KC1+ 2N. QH), ; (c) 2000l., +4KOH+ H.O. =4KC1-i- 2C9, Qº), ack 44 REACTIONS OF THE METALs. Hypochlorites give the same reactions with manganese, nickel, and cobalt salts, but they do not give the reaction with gold chlo- ride. The principal reactions of sodium are the REACTIONS IN THE DRY WAY. Sodium salts color the non-luminous gas-flame a monochro- matic yellow, which can be readily distinguished from the yellow flame of the gas alone in the following way: If we illuminate an Orange-colored body (such as a stick of Sealing-wax or a crystal of potassium dichromate) with white light (all glowing solid bodies emit white light), the red and orange rays will be reflected: the body appears orange. If these bodies are illuminated with the monochromatic sodium light, they can now only reflect yellow light: the bodies appear yellow (a delicate test). Spectrum.—A yellow line, coinciding with the D-line of the sun’s spectrum. This is an extremely delicate reaction; the rºmºrrºr part of a milligram of Sodium can be recognized in the spectrum. AMMONIUM, NH. At. Wt. 18.08. Occurrence.—In small amounts as carbonate and nitrite in the air. As ammonium chloride it is found in the fissures of active volcanoes. Ammonium derivatives are formed by the decay of many organic substances containing nitrogen: albumin, urea, etc., /NH, º +H,O=CO,--2NHs. 2 Urea and in a similar way by the dry distillation of many nitrogenous substances, such as coal, horn, hair, etc. Although ammonium itself is known only in the form of its amalgam, yet we are justified in considering it as a metal, in the first place because the electrolysis of ammonium salts causes the setting free of the cation NH,(NH3+H) at the same time that the corresponding anion is set free; and, further, because the ammo- nium salts are isomorphous with potassium salts. AMMONIUM. 45 REACTIONS IN THE WET WAY. Ammonium Salts show in many reactions a striking similarity to potassium salts. 1. Hydrochlorplatinic Acid gives a yellow crystalline precipi- tate: H.PtCla-i-2NH,Cl=(NH,), PtCl4-2HCl. This Salt may be distinguished from the potassium salt.— (a) by its behavior on ignition; platinum alone is left behind: (NH4)2PtCle=2NH,Cl-H2Cl2 +Pt; (b) by its behavior on treating with strong bases, whereby the smell of ammonia may be detected: (NH,),PtCls--2NaOH = Na, PtCle-H2H,0+2NH. 2. Tartaric Acid produces, as with potassium, a white, crys- talline precipitate of ammonium acid tartrate. The addition of sodium acetate, and rubbing the sides of the glass vessel with a stirring-rod, will hasten the formation of the precipitate: soon COONH, | HCOH HCOH | +NH,Cl= | +HC1. HCOH HCOH | | COOH COOH The ammonium acid tartrate, like the corresponding potassium salt, is soluble in alkalies and mineral acids. It may be distin- guished from the potassium salt by its behavior on ignition, carbon alone being left behind, and the residue not effervescing with hydrochloric acid; furthermore, it will give off ammonia on heating with caustic soda Solution. It is much safer to test for ammonium by heating with strong bases than to attempt the formation of a precipitate with the above reagents; sodium hydrate solution is the base most often employed. With this test all ammonium salts will give off ammonia, recognizable by its odor, or by its ability to turn mercurous nitrate paper black, Hg—NO, /Hg 2 +4NH, PH,0-8NHNO, toº XNH-No,4 2Hg; Hg—NOs Hg *— Black 46 REACTIONS OF THE METALS. recognizable also by its turning red litmus paper blue, or by causing dense fumes of ammonium chloride if a glass rod moistened with concentrated hydrochloric acid is held in contact with the ammonia. Vapors: NH,--HCl=NH,Cl. The above-mentioned reactions are not suitable for the detec- tion of very small amounts of ammonia, to the extent, for example, that it occurs in drinking-water. In such a case Nessler's reagent is used (an alkaline solution of mercuric potassium iodide). Large amounts of ammonia produce a brown precipitate, Hg 2Hgki, H&KOH+NHoH=0& XNH-I-3H,0+7KI, Hg of enormous coloring power, so that the minutest trace of ammonia can be recognized by the formation of a distinct yellow color in the solution. In order to detect the presence of ammonia in drinking- water with the help of this reaction, we proceed as follows: s SS t § # Žs : tº:* 2–L. FIG. 4. First of all, the apparatus (Fig. 4) must be freed completely from all traces of ammonia. To accomplish this, about 500 cc. of spring water is placed in the retort (whose neck is bent as in the illustration), 1 cc. of a boiled, saturated soda solution is added, and the solution is distilled. The end of the retort’s neck is pushed into AMMONIUM. 47 the condenser tube; it is best not to use a rubber connector here, as the condensed water serves to make a sufficiently tight connection. The distillation is continued until 50 cc. of the distillate, placed in a small 50 cc. graduate of white glass, will show no sign of a yellow color after adding 1 cc. of the Nessler solution, stirring, and letting it stand for a quarter of an hour. The apparatus is now ready for the experiment proper. The retort is emptied of its contents, 500 cc. of the water to be tested is introduced, 1 cc. of the saturated soda solution is added, and 50 cc. distilled off. The dis- tillate is treated with 1 cc. of Nessler's reagent and stirred. With large amounts of ammonia a yellow color will immediately be noticeable, which on standing becomes orange, while with very large amounts of ammonia a brown precipitate appears. If only traces are present, the solution will assume a yellow tint only after standing some time. The Nessler's reagent * which is used for this experiment is pre- pared as follows: 6 gms. of mercuric chloride are added to 50 cc. of water (which must be absolutely free from ammonia f), and solu- tion is effected by warming to 80° C. in a porcelain dish; 7.4 gms. of potassium iodide are dissolved in 50 cc. of water, added to the above, and the mixture allowed to cool. The clear liquid is then poured off and the residue washed three times with 20 cc. of cold water, in order to dissolve out as much of the chloride as possible. Then 5 gms. more of potassium iodide are added, which causes the mercuric iodide to go into solution. The solution is now poured into a 100-cc. flask, 20 gm. KOH dissolved in a little water are added, and the contents of the flask diluted to 100 cc. When * Cf. L. W. Winkler, Ch. Cent. Bl., 1899, II, p. 320. i Water absolutely free from ammonia, commonly called “best water,” is prepared from ordinary distilled water, or good spring water, by adding some soda and distilling The first portion of the distillate, which always con- tains ammonia is rejected until Nessler’s reagent, applied as above to 50 cc. of the distillate, gives a negative test, which will be after about one fourth of the original amount of water has been distilled off. The portion now coming off will be completely free from ammonia, and is used for the preparation of the reagent. The distillation, however, must be stopped when about five sixths of the water has been distilled, as the remainder may contain ammonia owing to the decomposition of Organic substances in the water by the strongly concentrated soda solution. 48 REACTIONS OF THE METALS. the solution has completely cleared, it is carefully decanted into a clean flask and kept in the dark. ** REACTIONS IN THE DRY WAY. º Thébelºor of ammºnwºsals Qºheating in a £ºosed tube is very characteristic. • ‘s All ammonium salts, except those with non-volatile acids, are com- pletely volatilized, partly without and partly with decomposition. The non-volatile acids whose salts form an exception to this rule are boric, phosphoric, chromic, molybdic, tungstic, and vanadic acids. Only the Salts with the halogen acids are volatile without de- composition; they give a white sublimate. All ammonium salts which suffer decomposition on ignition split off water. Thus the nitrate gives water and nitrous oxide, NH, NO,-2H,0+N.O; the nitrite gives water and nitrogen, NH, NO,-2H,0+N,; the sulphate gives water, nitrogen, ammonia, and sulphur dioxide, 3(NH,),SO, =N,+4NH,--6H,0+3SO,; the oxalate gives water, considerable ammonia, carbon dioxide, car- bon monoxide, and, towards the end of the reaction, dicyanogen, which is best recognized by its odor. With regard to ammonium salts with non-volatile acids, the behavior of the phosphate and the dichromate may be mentioned: 2(NH),PO,-6NH,--2H,0+2HPO, Metaphosphoric acid (NHL),Cr,O,-4H,0+2N+Cr,0s. The chromic oxide remains as a voluminous mass, looking very much like green tea. Ammonium salts do not impart a very char- acteristic color to the flame; the border of the flame is tinged slightly greenish. MAGNESIUM. 49 MAGNESIUM, Mg. At. Wt. 24.3. Occurrence.—Magnesium compounds are found very abun- dantly in nature. The most important minerals are magne- site, MgCOs, rhombohedral, isomorphous with calcite; dolomite, (Ca,Mg)CO, ; brucite, Mg(OH)2, rhombohedral ; carnallite, KMgCl2, Orthorhombic; kieserite, MgSO,--H,O, monoclinic; ep- somite, MgSO4+7H,O, Orthorhombic; spinel, MgAl,O, isometric, isomorphous with magnetite, Fe2O, and with chromite, FeCr2O4. Magnesium also occurs in a great many silicates. Thus almost all the minerals of the olivine group contain more or less mag- nesium. To this group belong forsterite, Mg,SiO, monticellite, CaMgSiO, and olivine, (FeMg)SiO, all orthorhombic. An impor- tant decomposition product of the olivine minerals is serpentine, Mg,FLSi,0,. Almost all the minerals of the pyroxene-amphi- bole group contain magnesium: augite, MgAl,SiOs, hornblende, CaMg,Al,SigO2, and tremolite, CaMg,Si,OI2, all three forming monoclinic crystals. Asbestos is a variety of tremolite with very fine fibers. Meerschaum is a magnesium silicate of the composi- tion H.Mg,SigOlo, and is quite similar to talc, H.,MgaSiO2, Some- times called steatite. Properties of Magnesium.—Magnesium is a silver-white metal of 1.75 sp. gr., and melts at about 630°C. It decomposes water very slowly, forming an oxide (MgO) which is only slightly soluble in water, with a weakly alkaline reaction. Magnesium reacts directly with nitrogen at 300° C., forming magnesium nitride (Mg,N.), which is readily decomposed by water, forming magne- sium hydroxide and ammonia: Mg,N,+6HOH=3Mg(OH),4-2NH4. The salts of magnesium are almost all colorless and soluble in water. The hydroxide, carbonate, phosphate, arseniate, and ar- senite are insoluble. The Sulphide, which can be prepared only in the dry way, is completely decomposed by water into hydroxide and hydrogen sulphide (hydrolysis). If an aqueous solution of magnesium chloride be evaporated to dryness on the water-bath, there is no hydrolytic decomposition, the residual salt, MgCl2+ 6H2O, dissolves in water, forming a perfectly clear Solution. On heating the chloride containing the water of crystallization to 106° and higher, however, a considerable amount of hydrochloric acid fumes escape and a basic magnesium chloride insoluble in water is left behind: 5o REACTIONS OF THE METALS. Mg 3 Cl Mg-Cl §4-#50 — >o +2HCl. Mg ‘oi Mg—Cl When a Saturated solution of magnesium chloride is mixed with magnesium oxide, the mixture soon solidifies, forming a mass hard as stone known as magnesia cement, consisting of magnesium oxychloride. *. 1. Ammonia.-In a neutral solution containing no ammonium salts, ammonia produces a white, gelatinous precipitate of magne- sium hydroxide. The precipitation, however, is never quantitative; in dilute solutions when only a slight excess of ammonia is used only a slight precipitate is formed, but on increasing the excess of ammonia, more and more magnesium is thrown out of solution until finally the greater part of the magnesium, although never all of it, is changed to the insoluble hydroxide. The behavior of magnesium compounds towards ammonia is quite different in case the Solution already contains a considerable amount of ammonium Salts; ammonia then causes no precipitation in the cold even when it is added in great excess; in the latter case, part of the magne- sium will precipitate on boiling the solution. The reaction takes place according to the following equation: In the presence of considerable ammonium salts the reaction takes place in the sense of the equation from right to left and in fact quantitatively; on the other hand, the presence of an excess of ammonia serves to increase the tendency of the reaction to take place in the direction from left to right, while the ammonium chloride formed will invariably cause the reaction to cease before all of the magnesium is precipitated. As we shall find later on, solutions containing bivalent metals of the ammonium sulphide group behave in precisely the same way towards ammonia as those of magnesium, whereas the trivalent metals of this group act differently. The latter are precipitated completely as hydroxides * by ammonia from solutions containing * With the exception of uranium, which is precipitated as ammonium uranate. MAGNESIUM. 5oa their salts even when only a slight excess of the reagent is used, and in spite of the presence of ammonium salts. Why does magnesium behave so differently from aluminium, ferric iron, etc.? Lovén has answered this question and shown it to be due to the different solubilities of the hydroxides in question.* Magnesium hydroxide is appreciably soluble in water, enough dissolving to make the water react alkaline. When magnesium hydroxide is suspended in water, the solution is saturated with it and the dissolved hydroxide is almost completely dissociated: Mg(OH)2 a Mg” + OH' + OH' [Mg] × [OH]? = k. / [Mg(OH)2] If we increase the concentration of [OH], the above equilibrium will be disturbed, and to offset this some undissociated magnesium hydroxide will be formed. Since the solution is already saturated with the latter compound, this means that magnesium hydroxido will be precipitated from the solution. If, however, we diminish in any way the concentration of the hydroxyl ions, that of the magnesium hydroxide will likewise decrease in accordance with the above equation, because the only way to restore equilibrium is for magnesium hydroxide to dissolve; if sufficient hydroxyl is removed finally all of the precipitate will disappear. The solubility product of magnesium hydroxide being rela- tively large, an appreciable amount of hydroxyl ions is required to precipitate it from solutions. The hydroxyl ions may come from ammonia according to the equation NH3 + H2O a NH3 + OH'. Ammonia, however, is dissociated only slightly in aqueous solution, and furthermore the concentration of the hydroxyl ions is greatly diminished if we increase the concentration of the ammo- nium ions by adding an ammonium salt. There will then be extremely few hydroxyl ions in solution.} * Z. anorg. Chem. XI (1896), p. 404. + Cf. page 15. 5ob REACTIONS OF THE METALS. On treating a solution of magnesium chloride with ammonia, the hydroxyl ions from the latter will cause the precipitation of magnesium hydroxide: MgCl2 + 2NH4OH z= Mg(OH)2 + 2NH4Cl. At the same time, however, ammonium chloride is formed so that the concentration of the ammonium ions is increased, and this tends to prevent the further dissociation of ammonia and the precipitation of magnesium. In fact, by adding ammonium chloride to a solution in which magnesium hydroxide is suspended, the concentration of the hydroxyl ions is diminished so that the solubility product of mag- nesium hydroxide is no longer reached and the ammonium chloride solution exerts a solvent action upon the hydroxide. The bivalent metals of the ammonium sulphide group, as mentioned above, act similarly to magnesium in this respect. Trivalent iron and aluminium, on the other hand, are precipitated Quantitatively by a slight excess of ammonia even in the presence of ammonium salts because the hydroxides of these metals are almost entirely insoluble, or in other words their solubility prod- ucts are so low that they are exceeded even by the small amount of hydroxyl ions into which ammonia dissociates in the presence of considerable ammonium chloride. Formerly, the non-precipitation of magnesium by ammonia in the presence of ammonium salts was explained by the assumption that complex Salts such as [MgCl3]NH4 and [MgCl4](NH4)2 were formed, but Lovén 4 has shown that this is not the true explana- tion; there is no evidence of the formation of any such salt under the given conditions. Inasmuch as magnesium, ferrous iron, Mn, Ni, Co, and Zn are not precipitated by ammonia in the presence of ammonium salts, we have a means for separating these from the remaining metals of the ammonium sulphide group (ferric iron, Al, Cr, Ti, U). If the separation takes place in a hot solution more ammonium chloride is required to prevent the magnesium from precipitating than in the cold. * Z. anorg. Chem. XI (1896), p. 404. See also Treadwell, Z. anorg. Chem. 37 (1903), p. 326; Herzebenda, 38 (19.3), p. 138. MAGNESIUM. 51 2. Barium Hydroxide precipitates magnesium almost quanti- tatively as hydroxide from its solutions, but only in the absence of ammonium salts: MgCl,--Ba(OH)2+BaCl, +Mg(OH)2. 3. Ammonium Carbonate precipitates, in the absence of other ammonium Salts, a basic salt (usually only on boiling or after long standing). The composition of the precipitated Salt varies with the temperature and the concentration of the solution, the follow- ing Salt being often obtained: 4MgSO,--4(NHA),CO,--H,O= Mg,(CO)2(OH)2+CO,--4(NH,),SO, 4. Sodium Phosphate is the characteristic reagent for magne- sium. It produces in solutions containing ammonium chloride, and in the presence of ammonia, a white crystalline precipitate (orthorhombic, hemimorphous) of magnesium-ammonium phos- phate, Rubbing the sides of the beaker with a glass rod hastens the forma- tion of the precipitate. From very dilute solutions the precipitate separates out only after standing some time. REACTIONS IN THE DRY WAY. All magnesium salts are more or less changed on heating in the air, leaving behind the oxide or an insoluble basic salt. On charcoal with sodium carbonate before the blowpipe, magnesium compounds are changed to white magnesium oxide, which is strongly luminous when hot. Calcium, strontium, and aluminium compounds behave the same way. The magnesium salts are non-volatile, do not color the flame, and give no flame spectrum, but do give a characteristic spark spectrum. DETECTION AND SEPARATION OF MAGNESIUM, AND THE ALKALIES IN THE PRESENCE OF ONE ANOTHER. These metals are assumed to be present in the form of chlorides. Ammonium is tested for, first of all, by heating a small part of the substance with caustic soda solution. The rest of the substance 52 REACTIONS OF THE METALS, is used in testing for magnesium, potassium, and sodium. It is divided into two parts, one being used in the test for magnesium and the other for sodium and potassium. * In order to test for magnesium, the substance is dissolved in as little water as possible; or, in case we have a solution, it is evapo- rated to dryness and then dissolved in a little water. To the solu- tion, ammonium chloride is added (if not already known to be present), and then ammonia. In case ammonia produces a pre- cipitate of Mg(OH), enough more ammonium chloride must be added to dissolve it. Sodium phosphate solution is now added, and the sides of the beaker rubbed with a glass rod. If there is magnesium enough present to amount to a few tenths of a milli- gram per 100 cc. of the Solution, a precipitate of magnesium-ammo- nium phosphate will Surely appear within two or three minutes. If no precipitate is formed, the beaker is set aside in order to see if Small crystals will be deposited on the bottom and rubbed sides after twelve hours. These crystals are more apparent on pouring off the solution. If the crystals appear on standing, then traces of magnesium are present. Before testing for sodium and potassium the solution must be freed from magnesium; and, as ammonium salts are usually present, it is necessary to evaporate to dryness and drive them off by gentle ignition. Heating to redness must be avoided, in order to prevent loss of sodium and potassium. The residue from igni- tion is dissolved in a little water (it is not necessary to obtain a perfectly clear Solution, as a little basic Salt of magnesium usually remains undissolved), and barium hydrate solution is added until the Solution is strongly alkaline, when it is boiled and the magne- sium hydroxide filtered off. This operation must be carried on in a platinum or good porcelain dish, never in glass, lost some alkali from the glass get into the solution. The filtrate, which is now free from magnesium, is cautiously acidified with hydrochloric acid, froed from the excess of barium hydrate by precipitation with ammonia and ammonium carbonate at the boiling temperature and ſiltering off the crystalline barium carbonate. The filtrate is now evaporated to dryness, the ammonium salts are driven off by igni- tion, and the residue is dissolved in a little water. In order to be sure that the barium is completely removed from the solution, it is treated again with ammonia and ammonium carbonate, filtered, ALKALINE EARTHS. 53 evaporated to dryness, and the ammonium salts driven off as before. The residue is again dissolved in a little water, and the Small amount of black carbonaceous matter (due to carbonization of a small amount of pyridin bases in the ammonia) is filtered off. A Small part of the solution is now tested for potassium on a watch-glass by means of hydrochlorplatinic acid, a yellow crystalline precipi- tate showing the presence of potassium. The larger part of the solution is used in testing for sodium by means of the flame reac- tion, and the potassium pyroantimonate. A white crystalline precipitate with the latter reagent shows the presence of Sodium. GROUP IV. ALKALINE EARTHS. CALCIUM, At. Wt. 40. STRONTIUM, At.Wt. 87.6. BARIUM, At. Wt. 137.4. GENERAL CHARACTERISTIC REACTIONS. The metals of the alkaline-earth group are bivalent and heavier than water, which they decompose slowly at ordinary tempera- tures, with evolution of hydrogen and formation of difficultly solu- ble hydroxides of strongly alkaline reaction. The salts are mostly colorless and insoluble in water. The halogen compounds, nitrates, nitrites, and acetates are soluble in water. The carbonates are insoluble in water and are decomposed on, ignition into carbon dioxide, and white infusible, strongly luminous metallic oxide: CaCO3=CO,--CaO. Barium carbonate is an exception to the general rule. It loses its carbon dioxide only when heated to a white heat, and the oxide is not very luminous. The sulphates and oxalates are very difficultly soluble. The sul- phate of barium is the most insoluble and calcium sulphate the most soluble; of the oxalates, the calcium salt is the most insoluble. The solubility of the strontium Salt is always midway between that of the corresponding calcium and barium salt, for the atomic weight of strontium, of which the solubility is a function, lics midway between the atomic weights of barium and calcium. The halogen salts are volatile and impart a characteristic color to the flame. The metals of this group form oxides of the general type RO, and peroxides corresponding to the formula RO,. The latter, on treat- 54 REACTIONS OF THE METALS. ment with acids, give hydrogen peroxide and salts corresponding to the oxide RO: RO,--2HCl=H.O., +RCl,. CALCIUM, Ca. At. Wt. 40.1. Sp. Gr. 1.58. Occurrence.—Calcium is widely distributed in nature. It is found in enormous deposits in all stratified formations as carbonate (limestone, marble, chalk), often rich in petrification. The car- bonate, CaCO3, is dimorphous, crystallizing in rhombohedrons as calcite and in the Orthorhombic system as aragonite. Calcium also occurs in large masses as Sulphate, partly as monoclinic crys- tallizing gypsum, CaSO4·2H2O, and partly as the anhydrous anhydrite, CaSO4, which crystallizes in the Orthorhombic system. Calcium also occurs as fluoride, fluorite, CaF2, which crystallizes in the isometric system, with perfect octahedral cleavage; as apatite, 3CasP,Os Caº, which belongs to the hexagonal system; and, finally, in innumerable silicates, such as the monoclinic wollastonite, CaSiO3, and the triclinic anorthite, CaAl,Si,Os. The calcium minerals are the chief representatives of several important isomor- phous groups: Calcite Group, Rhombohedral. Aragonite Group, Orthorhombic. Calcite, CaCO, Aragonite, CaCO, Magnesite, MgCOs Strontianite, SrCO, * Ca Witherite, BaCOs Dolomite, Mg } COs Cerussite, PbCO, Siderite, FeCO, Smithsonite, ZnCOs Rhodochrosite,MnCOs Anhydrite Group, Orthorhombic. Apatite Group, Hezagonal. Anhydrite, CaSO, Apatite, 3Ca,P,Os--Ca(CIF) Celestite, SrSO, Pyromorphite, 3Pb,P,Os--PbCl, Barite, BaSO, Mimetite, 3PbsAs,Os--PbCl, Anglesite, PbSO, Vanadinite, 3Pb.V.Os--PbCl, REACTIONS IN THE WET WAY. 1. Ammonia, in case it is free from carbonate, produces no pre- cipitate with calcium salts; on standing in the air, however, it absorbs carbonic acid and causes a turbidity of calcium carbonate. CALCIUM. 55 2. Ammonium Carbonate.—The commercial salt is a mixture of ammonium bicarbonate and ammonium carbamate: * co( co- OH ONH, Ammonium bicarbonate Ammonium carbamate We therefore add ammonia to the reagent, whereby the ammonium bicarbonate is changed to normal carbonate: ONH, 2ONH, coº +NH,OH = H,0+CO tº OH NONH, The commercial Salt and ammonia produce in solutions containing calcium salts, at first a voluminous, flocky precipitate, which on standing gradually becomes crystalline, the change taking place more quickly on heating: CaCl2 +(NH4)2CO3=2NH,Cl--CaCOs. This reaction is noticeably reversible. Consequently, an excess of the precipitant should be added, and the solution boiled only long enough to cause the precipitate to become crystalline. In the presence of considerable ammonium chloride, and with only a small amount of ammonium carbonate, the precipitate often fails to form. 3. Ammonium Oxalate produces in neutral or alkaline solu- tions a precipitate of calcium oxalate, which when formed from cold solutions is composed of extremely fine crystals, hard to filter. while from hot solutions larger crystals are formed: COONH, COO | +CaCl, -2NH,Cl-H XCa. COONH, COO Tº: The ammonium carbamate is changed completely to ammonium car- bonate on warming the aqueous solution to 60° C.: NH /ONH <$” Ho - 66” NôNH, NONH, Ammonium carbamate produces no precipitate with calcium salts in the cold, because calcium carbamate is soluble. If, however, the solution be heated, the calcium is quantitatively precipitated as carbonate: cº Nö HN, Ca + XO = 2NHs + CO2 + CaCOs. cºš H NNH, 56 REACTIONS OF THE METALS. Calcium oxalate is practically insoluble in water and acetic acid, but dissolves readily in mineral acids: COO COOH | Ca+2HCl=CaCl,-- | e COOH COO Ammonia precipitates from such a solution calcium oxalate again, the excess of mineral acid, as well as the oxalic acid which was set free, being changed to ammonium salt, causing the re-formation of the calcium oxalate. Calcium oxalate on being boiled with sodium carbonate solution is easily changed to carbonate: COO | Ca+Na,CO3=Na,C,04+CaCOs. COO t 4. Sulphuric Acid produces a precipitate only in concentrated solutions: CaCl, -i-H,SO, =2HCl--CaSO4. 100 parts water dissolve at 40°C. 0.214 gm. CaSO4+2H2O. If alcohol is added to an aqueous solution of calcium sulphate, all of the calcium salt will be precipitated. Calcium sulphate, dissolves in hot dilute hydrochloric acid, and also in concentrated ammonium sulphate solution, forming [Ca(SO4)2](NH4)2, which is decomposed by water. 5. Calcium Sulphate solution produces no precipitation with calcium salts. (Note distinction from strontium and barium.) 6. Sodium Phosphate (Na,HPO,) produces in neutral solutions a white flocky precipitate of calcium hydrogen phosphate: NaO O cac,+Nºëpo-2Nacºpo. Ho / Hö/ If ammonia is added to the solution at the same time, tertiary calcium phosphate will be precipitated: 2Na,HPO,--2NHa-H3Ca(Cl2=4NaCl--2NH,Cl-i-Casp,0s. Both of these phosphates of calcium are soluble in mineral acids, or even acetic acid. From such solutions ammonia always precipi- tates the tertiary Salt. CALCIUM. 57 7. Chromates of the Alkalies do not precipitate calcium salts. (Note distinction from barium.) 8. Absolute Alcohol, or a mixture of equal parts of absolute alcohol and ether, dissolves both the nitrate and chloride of calcium. All deliquescent salts, with the eacception of potassium carbonate, dissolve in absolute alcohol. All other salts are, as a rule, insoluble, or very difficultly soluble, in absolute alcohol. An exception to this rule is found in the case of mercuric chloride, which is not hygro- scopic, and is much more readily soluble in alcohol than it is in water. 9. Water decomposes the carbide, phosphide, and nitride of calcium at the ordinary temperature, as follows: * (a) The carbide: CaC,--2HOH=Ca(OH)2+C.H. Acetylene Acetylene is evolved by the reaction, a gas with a peculiar smell.” If this gas is conducted into an ammoniacal copper solution, it rapidly produces a red precipitate of copper acetylide. The latter compound is harmless so long as it is moist; but in the dry state it can be readily exploded by a blow, by rubbing, or by simply warm- ing. (b) The phosphidé: Ca,P,4-6HOH=3Ca(OH), H-2PHs. The garlic-smelling phosphine gas which is evolved is sponta- neously combustible, because it always contains a small amount of the spontaneously combustible, liquid phosphuretted hydrogen (P2H4). (c) The nitride: Ca,N,+6HOH=3Ca(OH)2+2NH3. REACTIONS IN THE DRY WAY. Calcium compounds, on being heated with sodium carbonate before the blowpipe, are changed to the white infusible oxide, which glows brightly when hot. The volatile calcium compounds color the non-luminous gas flame brick-red. * Pure acetylene is odorless. Almost all calcium carbide contains a little. calcium phosphide which evolves phosphine on treatment with water. 58 REACTIONS OF THE METALS. Spectrum.—Several lines in the orange and yellow, one in the green, and one in the ultra-violet. The orange and green lines are particularly bright. As a rule, at least three lines are seen—the ever- present Sodium line, and, at about equal distances on the right and left, the bright green and Orange lines (see chart). STRONTIUM, Sr. At. Wt. S7.6. Sp. Gr. 2.5. Melts at red heat. Occurrence.—Strontium occurs quite commonly with calcium, but usually in much smaller amounts. There are only a few real strontium minerals. The most important of these are: Strontianite, SrCOs, Orthorhombic, isomorphous with aragonite; and celestite, SrSO4, orthorhombic, isomorphous with barite. REACTIONS IN THE WET WAY. 1. Ammonia: same as with calcium. 2. Ammonium Carbonate: same as with calcium. 3. Ammonium Oxalate: same as with calcium; but the stron- tium oxalate is somewhat soluble in acetic acid. 4. Dilute Sulphuric Acid produces a white precipitate of stron- tium sulphate: SrCl,--H,SO, =2HCl·HSrSO4. Strontium sulphate is much less soluble in water than calcium sulphate (6900 parts of water at ordinary temperatures dissolve 1 part SrSO.), but much more soluble than barium sulphate. It is soluble in boiling hydrochloric acid, and insoluble in ammonium sulphate. By boiling with ammonium carbonate solution, the strontium sulphate is changed to carbonate: N SrSO,--(NHA),CO3=SrCO,--(NH4)2SO4. 5. Calcium Sulphate solution produces in neutral or weakly acid solutions, after some time, a precipitate of strontium sulphate: SrCl,--CaSO4–SrSO,--CaCl2. 6. Chromates of the Alkalies produce in dilute solutions no precipitate (thus differing from barium); but from quite concen- trated solutions strontium chromate is precipitated. BARIUM. 59 $ 7. Absolute Alcohol dissolves the deliquescent strontium chloride. The nitrate is not deliquescent, and does not dissolve in absolute alcohol. REACTIONS IN THE DRY WAY. Heated on charcoal before the blowpipe the strontium com- pounds behave similarly to the calcium compounds. The volatile strontium salts color the non-luminous gas-flame carmine red. Spectrum.—A number of red lines and a blue line. BARIUM, Ba. At. Wt. 137.40. Sp. Gr. about 4.0. The melting-point is higher than that of platinum. Occurrence.—Like strontium, barium is almost always found associated with calcium, but only in Small amounts. The most important barium minerals are: Witherite, BaCOs, orthorhombic, isomorphous with aragonite; barite, or heavy spar, BaSO4, ortho- rhombic, isomorphous with anhydrite; and the hydrous barium aluminium silicate, harmotome, BaAl,H,SigOls--4H2O. Harmo- tome crystallizes in the monoclinic system, and belongs to the class of zeolites. REACTIONS IN THE WET WAY. 1. Ammonia and Ammonium Carbonate: same as with cal- cium and strontium. 2. Ammonium Oxalate: same as with calcium and strontium, except that the barium oxalate formed is much more soluble in water (1 part dissolves in 2590 parts of cold water), and is com- pletely soluble in acetic acid on boiling. 3. Phosphates of the Alkalies: same as with calcium. 4. Chromates of the Alkalies produce in neutral solutions of barium salts a yellow precipitate of barium chromate (thus differing from calcium and strontium), BaCl,--K.CrO. =2KCl-i-BaCrO4, 6o REACTIONS OF THE METALS. § which is insoluble in water and acetic acid, but readily soluble in mineral acids; consequently barium solutions are not completely precipitated by a bichromate solution, 2BaCl, --K,Cr,O,--H,O=2|BaCrO,--2HCl·H2KCl; for the mineral acid which is set free dissolves half of the barium chromate: BaCrO,--2HCl=BaCl, +H,CrO,. By addition of sodium acetate the solvent action of the mineral acid is diminished, so that the precipitation becomes quantitative: CH,CO.Na+HCl= NaCl-i-CH,CO.H. 5. Dilute Sulphuric Acid produces, in even the most dilute solutions, a precipitate of barium sulphate: BaCl, 4-H,SO,-2HCl·H BaSO,. The precipitate is practically insoluble in water (1 gm. dissolves in 344,000 * liters of water), but it is soluble in hot concentrated sulphuric acid, forming an acid sulphate, /* 4 Baso,4-HsO,-Baº, HXO. which decomposes, on dilution with water, into barium sulphate and sulphuric acid. Barium sulphate is also readily soluble in sodium thiosulphate, and slightly soluble in boiling hydrochloric acid. If barium sulphate is boiled with a solution of sodium car- bonate, it undergoes only a partial decomposition, because the reaction is reversible: BaSO4+Na,COs a 2 BaCO3 +Na,SO4. In order to make this decomposition quantitative, the barium sulphate must be boiled with the sodium carbonate solution, filtered, treated with a new portion of sodium carbonate solution, and the process repeated until the filtrate no longer gives a test for sulphate. The more concentrated the sodium carbonate solution is, the more - * Fr. Kohlrausch and Fr. Rose. BARIUM. 6 I complete will be the decomposition. The highest degree of con- centration will be reached by fusion of the barium sulphate with anhydrous sodium carbonate. Consequently, in order to bring barium sulphate into solution, we proceed as follows: The substance is mixed with four times as much calcined sodium carbonate, and the mixture fused in a plati- num crucible. The product of the fusion is treated with a little water, boiled until the mass is disintegrated, filtered, washed with a solution of sodium carbonate (until the filtrate no longer gives a test for sulphate), and then washed with water. The barium car- bonate, which remains undissolved, can be readily brought into solution by means of dilute hydrochloric acid. If the product of the fusion is treated with considerable water, the sodium sulphate formed reacts with the barium carbonate, forming barium sulphate again. In this case the residue insoluble in water will not dissolve in hydrochloric acid to a clear solution. 6. Hydrofluosilicic Acid produces a white, crystalline precipi- tate of barium silicofluoride: H.SiFa-i-BaCl, -2HC1-HDaSiR. In order to effect complete precipitation, the solution must stand some time. Barium silicofluoride is difficultly soluble in water and dilute acids, and insoluble in alcohol. 7. Absolute Alcohol dissolves neither the nitrate nor the chlo- ride; neither of these salts is deliquescent. 8. Concentrated Hydrochloric Acid and Nitric Acid will pre- cipitate from barium solutions the chloride and nitrate respectively. REACTIONS IN THE DRY WAY. Heated with sodium carbonate on charcoal before the blowpipe, the barium compounds, unlike calcium and strontium, do not give a brightly luminous mass, because the barium carbonate formed is not decomposed at this temperature into the infusible oxide and carbon dioxide, but sinters together without perceptible change. Volatile barium salts color the non-luminous gas-flame yellowish- green. The sulphate is only slightly volatile in the hottest flame, and in the ordinary gas-flame it shows scarcely any coloration. In 62 REACTIONS OF THE METALS. Order to bring about the coloration, it is necessary to change the sul- phate into chloride, by reducing a small particle on a platinum wire in the upper reducing flame to Sulphide and adding a little hydro- chloric acid by means of a capillary tube. The wire is then brought into the flame, and should give the characteristic flame coloration. Spectrum.—A number of lines are to be seen, some in the orange, but more in the green. Separation of Calcium, Strontium, and Barium. In the course of a systematic analysis these three metals are always obtained in the form of their insoluble carbonates, either by precipitation with ammonium carbonate or by fusion of the sulphate, with sodium carbonate. The carbonates are treated in a small porcelain dish with dilute nitric acid (barium nitrate is insoluble in concentrated nitric acid), until all effervescence has ceased, and the contents of the dish cautiously evaporated to dryness (keeping the dish in constant motion); they are then heated until all moisture and nitric acid have been expelled. The residue of nitrates must not be heated hot enough to decompose the nitrates into oxides. A small part of the residue is dissolved in as little water as possible, and treated with calcium sulphate solution, If no precipitate is formed, then only calcium can be present. If a precipitate forms only after standing some time, then barium is absent. If a precipitate is formed immediately, then barium is present and calcium and strontium may be present. In both of the last two cases the nitrates must be examined for calcium, and in the last case for strontium also. To effect this, the absolutely dry nitrates are treated with a small quantity of absolute alcohol, stirred with a glass rod, and the alcoholic solution, which may con- tain calcium nitrate, is poured through a filter (previously moist- ened with alcohol) into a small porcelain crucible. The contents of the crucible are now evaporated to dryness, and any calcium salt adhering is removed by rubbing the inside of the crucible with a piece of filter-paper, which must itself be free from calcium. The filter-paper is fastened to a fine platinum wire, burned to an ash, moistened with hydrochloric acid (by means of a capillary SPECTRUM ANALYSIS. 63 tube), and introduced into the non-luminous flame of the burner. A brick-red coloration shows calcium. If a considerable residue remains after the evaporation of the alcohol, it can, of course, be tested for calcium without using the filter-paper. A portion of the residue should in this case be dis- solved in water, and tested for calcium by means of ammonium oxalate. A white precipitate, insoluble in acetic acid, shows the presence of calcium. In order to detect Strontium, the original residue of nitrates is washed with absolute alcohol (until the filtrate no longer gives a test for calcium with sulphuric acid), dried, mixed with an excess of ammonium chloride, and heated until vapors of ammonium salts are no longer given off. By this operation the nitrates of strontium and barium are changed to chlorides: * 2Sr(NO.), H-6NH,Cl=2SrC1,--10N+2C1+12H,69.f The chlorides thus obtained are treated, as above, with a little abso- lute alcohol, and the alcoholic solution is evaporated and tested for strontium by means of the flame. A carmine-red color shows strontium. t The residue insoluble in alcohol, which should be barium chlo- ride, is washed with absolute alcohol in order to remove all the strontium chloride, and the residue is, in this case also, tested with the flame. A yellowish-green color shows the presence of barium. Traces of alkalies and alkaline earths afe recognized best by means of the spectroscope. \, e will, therefore, at this place give a brief description of spectrum analysis. Spectrum Analysis (Bunsen and Kirchhoff, 1865). If a ray of white light is passed through a glass prism, not only is the direction of the ray changed, but the white light is decom- posed into colors; it suffers dispersion. It will be found that the * It is also easy to convert the nitrates into chlorides by repeated evap oration with hydrochloric acid. f The following reaction also takes place to some extent: Sr(NOS), + 2NH,Cl = SrCl3 + 4H2O + 2NO + N2. 64 REACTIONS OF THE METALS. red rays are deflected least, while the violet rays are most deflected. The picture obtained—the spectrum—if projected on a screen (Fig. 5), does not show the colors sharply separated, but merging into one another. Such a spectrum is called a contin- uous, or uninterrupted, spectrum. 2. Every glowing solid or liquid body emits white light; the spectra ob- tained in all such cases will be a continuous one. Glowing vapors and gases behave quite differ- ently. They do not emit white FIG. 5. light, but light composed of certain definite rays, which are characteristic for each gas and for each vapor. The light emitted from glowing vapors or gases, when decomposed by the prism, yields on the Screen an interrupted spec- trum. If the light is passed through a fine slit before reaching the prism, the spectrum will be found to consist of a greater or less number of colored lines which always appear in the same place with any given substance, provided the prism or its position is not changed. In order to determine the exact position of these lines, we make use of the spectroscope of Bunsen and Kirschhoff (Fig. 6). Fig. 7 shows a cross-section of the apparatus. The substance to be examined is placed in the loop of a platinum SPECTRUM ANALYSIS. 65 wire and introduced at A into the non-luminous gas-flame, by means of which it is volatilized. The rays of light pass through the slit into the tube Sp, reach the prism, by which the rays are refracted into the telescope C (the collimator) and are observed at D. Upon a glass plate at the end of the tube Sk is a transparent scale, which is illuminated by a small flame at B. This tube is so inclined toward B O Sk A , jºb FIG. 7. the face aa of the prism that the rays of light from this tube are totally reflected into the tube C, and reach the eye of the observer, so that the rays from the substance appear at a certain position on the scale. As, however, the position of the lines depends upon the dispersive power of the prism, and upon its angle of refraction, it is clear that the position of the lines will be somewhat different in different spectroscopes. As every ray has a definite wave length, it is better to give the wave lengths of the rays which appear, rather than their position on the scale. Wave lengths are expressed in millionths of a millimeter and designated as pupa. According to the most accurate measurements the Red potassium line Ra = 7700 pupt Violet potassium line Kg = 4044 “ Red lithium line Lia = 6075 “ Yellow sodium line Na = 5895 “ Green thallium line TI = 5350 “ Blue strontium line Srs = 4608 “ Blue indium line In a = 4510 “ Violet indium line In g = 4101 “ 66 REACTIONS OF THE METALS. Let us assume that the foregoing lines were observed at the following positions on the Scale: Ka at 17, Kg at 159, Lia at 32, Na at 50, Tl at 68, Sr., at 101, Ina at 111, and Ing at 149. If now, upon a rectangular system of coordinates, we plot the wave lengths as ordinates and the corresponding scale divisions as abscissae, and connect the points of intersection, we will get a curve which expresses the relation between the wave lengths and the corresponding divisions on the Scale. We can, without serious error, regard the portions of the curve between two points of inter- section as proportional to the corresponding Scale divisions. The use of the curve may be illustrated by a single example: Suppose we observe at the division 60 of the scale a line a. What is the corresponding wave length? The line lies between the sodium and thallium lines, which are of known wave lengths. The sodium line lies at division 50 and corresponds to 5895 pupi & & The thallium line “ “ “ 68 “ { % “ 5350 Consequently 18 scale divisions= 545 “ and 1 “ division = 30.28 “ The line a is at 60=50+10 divisions, and these correspond to Division 50 = 5895 pupi 10 divisions = 302.8 “ Division 60 = 5592.2 “ The scale divisions increase in their value as the wave lengths diminish, so that we subtract the 302.8 from 5895. It is not possible to volatilize all substances in the gas-flame. These non-volatile bodies will not give any flame spectra, but by means of the electric spark they will be volatilized enough to give spark spectra. As the determination of their spark spectra fur- nishes the only method by which the purity of many substances may be tested, we will briefly outline the methods for the produc- duction of spark Spectra: The apparatus invented by Bunsen consists of two platinum wires, which are each attached at one end to conical pieces of char- coal. The latter are soaked with a solution of the body to be tested. The two carbon points are now placed opposite to one another, quite near together; and the other ends of the platinum wires are connected with an induction coil. which causes sparks * SPECTRUM ANALYSIS. 67 to pass between the two carbon points, volatilizing some of the salt. If now the sparks are viewed through a spectroscope, a large number of lines will be seen, of which only a part are produced by the substance itself. Some of the lines are caused by the carbon points and some by the air. In order to determine which lines are caused by the substance which is being tested, the experiment is first performed without using any of the substance, and the spec- trum thus obtained is either drawn or photographed. The experi- ment is then performed with some of the Salt impregnated in the carbon points and the new lines in the spectrum will be caused by the substance itself. Spark spectra can be more conveniently produced by means of the so-called “fulgurator” of Dela- chanel and Mermet, as shown in Fig. 8. The salt solution is in a test-tube, so that the slit of the spectroscope can- not be contaminated with the spat- tered particles of the salt. The small apparatus of H. Dennis, Fig. 9, is also very convenient for pro- ducing spark spectra. The platinum wire a is fused in a glass tube and ends in a point of Ceylon graphite, which extends up out of the arm e of the ap- paratus. The upper pole is not shown in the illustration. In order to fill the tube m with the salt solution of the substance to be tested, s is taken away, and the appara- tus is inclined to the left. The solution is poured in at the upper end of m, when s is again introduced and shoved FIG. 8. FIG. 9. down until it reaches nearly to the bottom. The apparatus is now placed in a vertical position, whereby the liquid in e rises until it is level with the lower end of s. On raising s the liquid will rise up to the upper edge of e. By means of electric sparks from the carbon point, enough of the liquid is evaporated to afford the de- sired spectrum. Care should be taken to prevent any spattering of the substance into the slit of the spectroscope. S 68 REACTIONS OF THE METALS. In order to examine the spectra of permanent gases, small Geissler tubes are used, which are filled with the gas to be exam- ined. Besides flame spectra and spark spectra, there remain ab- Sorption spectra to be mentioned. If white light is passed through a colored solution or gas, certain rays are absorbed by the liquid, So that if the light is now examined in the spectroscope, these rays will be found to be lacking. We See a bright spectrum broken by black bands (absorption bands) which are characteristic for differ- ent substances. Thus Solutions of permanganic acid, neodymium, praseodymium, erbium, and many other SubSntaces give char- acteristic absorption spectra. These absorption spectra are often obtained from colorless solutions. S SSSSSSSSSSSSSSSSSSSSSSSSSSSSSS s SS Şsiss º § º § N. | | S SS Us!!Nss §ss=s º FIG. 10. In order to obtain any of the above spectra sharply defined, it is necessary to have the spectroscope properly adjusted, i.e., the rays must come parallel from the collimator tube into the telescope. The telescope is removed and adjusted for parallel rays by focusing it upon some distant object. It is then replaced, the prism re- moved, and the collimator tube brought exactly opposite the tele- scope, so that the slit of the former can be observed by the latter. The collimator tube must be lengthened or shortened until the picture of the slit is sharply defined. The prism is now replaced and the scale tube adjusted until the scale also can be seen clearly defined. The instrument is now ready for use. The direct-vision pocket spectroscope, with an arrangement of prisms as shown in Fig. 10, is very convenient for ordinary use. ALUMINIUM. 69 GROUP III. AMMONIUM SULPHIDE GROUP. ALUMINIUM, TITANIUM, CHROMIUM, IRON, URANIUM, ZINC, MANGANESE, NICKEL, COBAL.T. (BERYLLIUM, ZIRCONIUM, THORIUM, YTTRIUM, ERBIUM, CERIUM, LANTHANUM, NEO- DYMIUM, PRASEODYMIUM, NIOBIUM, TANTALUM.) ALUMINIUM, Al. At. Wt. 27.1. Sp. Gr. 2.56–2.67. M. Pt. about 658° C. Occurrence.—Aluminium occurs very extensively in nature, principally in the form of silicates, of which the feldspars and micas with their decomposition products are important examples: Si,Os-EKs SiO, a KH, OH &=Al Aksö-Atº Alóð,-H, Si,Osa Al NsióEAl SiO, as Al Orthoclase Muscovite ECaolin (feldspar) (mica) (decomposition product) Impure kaolin is called clay. Among the most important minerals which contain aluminium may be mentioned cryolite, AlFalNas; spinel, Al,O,Mg, or mag- nesium aluminate, which crystallizes in the regular system and is isomorphous with magnetite, Fe,0, Fe, and chromite, Gr,O.Fe; alunite, Al, K(SO4)2(OH)6; aluminium hydroxide as /QH g hydrargillite, Al-OH, monoclinic; Nółł OH A&# 29 bauxite, X. , and diaspore, AlòH, orthorhombic; Alz OH - NOH and corundum, Al,O, (ruby, Sapphire, emery), which is hexagonal, rhombohedral, with a hardness of 9, and isomorphous with Fe,0, and Cr,0s. * Aluminium is trivalent of silver-white color, is only slightly affected by moist air, dissolves readily in hydrochloric acid, diffi- cultly in Sulphuric acid, and very difficultly in nitric acid, but readily in Sodium or potassium hydroxide Solution with evolution of hydrogen and formation of aluminates: 3Al-H6KOH=2Al(OK), +3H, 7o REACTIONS OF THE METALS. Aluminium forms only one oxide, Al,O, and consequently only one series of Salts. The Salts are as a rule colorless, and those which are soluble in water show an acid reaction in aqueous solution, on account of their being hydrolyzed to a considerable extent. This explains the fact that on evaporating a solution of aluminium chlo- ride in water we do not obtain aluminium chloride, but the in- soluble oxide, or hydroxide: AlCl4-3HOH=Al(OH),4-3HCl. The property which aluminium possesses of forming alums is very characteristic. The alums are double salts of aluminium Sulphate with the Sulphates of potassium, caesium, rubidium, or ammonium, of the general formula Al=SO, N + 12H2O. 4 The alums crystallize in the regular system, usually in octa- hedrons, often combined with oo O OO and oo O. The sulphide of aluminium can be prepared only in the dry way. It is a bright-yellow substance, which is hydrolytically decomposed even by cold water into hydrogen sulphide and aluminium hy- droxide: Al,S,--6HOH=3H.S.--2Al(OH),. Towards strong acids aluminium hydroxide plays the part of an alkali, while towards strong bases, on the other hand, it acts as an acid. REACTIONS IN THE WET WAY. 1. Ammonia produces a gelatinous precipitate of aluminium hydroxide, which is somewhat soluble in water, but insoluble in the presence of ammonium Salts: AlCl4-3NH,OH=Al(OH),4-3NH,Gl. The property which the aluminium hydroxide shows of partly dissolving in water, is common to all colloidal Substances. They can exist in an insoluble form as hydrogel, and in a soluble form as hydrosol. ALUMINIUM. t 71 On boiling a solution of the hydrosol, it can usually be changed into hydrogel, or the same change will take place on adding a salt to the solution. The hydrosol form of aluminium hydrate cannot be converted into hydrogel by boiling, but it can be by the addition of Salts,” preferably ammonium salts. If, therefore, we desire to precipitate aluminium from a solution by means of ammonia, we take care that a large eaccess of ammonium chloride is present. The freshly-precipitated aluminium hydroxide is readily soluble in dilute acids; but after standing some time under water, or after long boiling, it becomes more difficultly soluble, so that it is neces- sary to digest it with acid for a long time in order to bring it com- pletely into solution. The hydrate, Al(OH), is probably changed into Al,O(OH), or AlO(OH), which are much more insoluble. All three hydrates occur in nature, Al(OH), as hydrargillite, Al,0(OH), as bauxite, and AlO(OH) as diaspore. The first dissolves in hot hydrochloric acid, but the two latter do not. 2. Potassium or Sodium Hydroxide produces the same precip- itate as ammonia which is, however, in this case soluble in excess of the reagent, forming an alkali aluminate: AlCl4-3KOH=3KCl-i-Al(OH), Al(OH),4-3KOH=3H,0+ Al(OK). In this last reaction the aluminium hydroxide behaves as an acid. If we add dilute acid to a solution of an aluminate, there is formed at first a precipitate of aluminium hydroxide, which dis- solves on further addition of acid: Al(ONa),4-3HCl=3NaCl-i-Al(OH), Al(OH), +3HCl=3H,0+ AlCl, The aluminates are also decomposed by ammonium salts, be- cause ammonium aluminate suffers complete hydrolysis: Al(ONa),4-3NH,Cl= 3NaCl-H Al(ONH.), and Al(ONH,),4-3H,O =3NH,OH-i-Al(OH),. * This principle is illustrated by the technical process of “salting out” colloidal dyes from their solutions. § 72 REACTIONS OF THE METALS. Aluminium hydroxide is solublé in neutral tartrates of the alka- lies, so that in the presence of tartaric acid there will be no precipi- tation on the addition of ammonia. Consequently the aluminium cannot be present in the solution in the form of the aluminium ion, but as a complex negative ion. The aluminum is present then in the solution, not in the form of positive aluminium ions, but in a complex negative ion, perhaps of the salt: tºok FOOOC COOK tion HOHC Or tºo COO OOC COOK N/ A. O H Many other organic oxy-acids and oxy-compounds, such as malic and citric acids, sugars and starches, have the same effect of preventing the precipitation of aluminium hydroxide by ammonia. 3. Ammonium Sulphide causes precipitation of the hydroxide, because the sulphide is completely hydrolyzed by water: 2A1C1,4-3(NH,),S=6NH,Cl-i-Al,S, Al,S,--6HOH=2Al(OH),4-3H,S. 4. Alkali Carbonates precipitate aluminium hydroxide (hy- drolysis): 2AlCl,--3Na,CO, =6NaCl-i-Al2(CO)3. Al,(CO),4-6HOH =3H,CO,--2Al(OH), 3(H.O.--CO.) 5. Barium Carbonate, suspended in water and added to the solu- tion of an aluminium salt, also precipitates the hydroxide: 2AlCl4-3BacO,--6HOH=3|BaCl,--3H,CO,--2Al(OH),. 3(H,0+CO) ALUMINIUM. 73 6. Alkali Acetates produce no precipitation in cold neutral solutions, but, on boiling the solution, a very voluminous precipi- tate of basic aluminium acetate separates Out: AlCl4-3NaC.H.O,-3NaCl-i-Al(C.H.O.),(in the cold), Soluble /Q# * * * Al(C.H.O.),4-2HOH z=2 Alsº, O +2HC.H.O., (on boiling). 2++3\’2 If the solution is allowed to cool, the basic aluminium acetate redissolves. The reaction is, therefore, a reversible one, and goes more completely from left to right according as we increase the amount of water and raise the temperature. 7. Alkali Phosphates, Na,FIPO, give a gelatinous precipitate of aluminium phosphate: 2Na,HPO,--AlCl,-AlPO,--NaH,PO,--3NaCl, or, on addition of ammonia, Na, HPO,--NH,--AlCl,-AlPO,4-2NaCl- NH,Cl. Aluminium phosphate is soluble in mineral acids, insoluble in acetic acid (differing from Ca, Sr, Ba, Mg), but readily soluble in Sodium or potassium hydroxide solutions: AlBO,--6NaOH=Al(ONa),4-Na,PO,4-3H,0. On boiling this alkaline solution (obtained in the last reaction) with ammonium chloride, a precipitate will be formed, consisting of a mixture of aluminium phosphate and aluminium hydroxide; while barium chloride, on the contrary, will precipitate barium phosphate from such a solution and leave the aluminate dissolved. 8. Sodium. Thiosulphate, Na,S.O., completely precipitates the aluminium as hydroxide on boiling: 2AlCl4-3Na,S,O,--3HOH=6NaCl-H3S--3SO,--2Al(OH), 74 REACTIONS OF THE METALS, Detection of Aluminium in the Presence of Organic Substances. The presence of tartaric acid or other non-volatile organic oxy- compounds prevents the precipitation with the above reagents. In order to detect the presence of aluminium in such cases, the organic substance must be first destroyed. This is best accomplished as follows: Some sodium carbonate and a little potassium nitrate are added to the solution, which is then evaporated to dryness in a platinum dish, and the residue ignited, whereby the aluminium becomes aluminate, and the organic substance is destroyed with separation of carbon. If now the residue is treated with nitric acid + and filtered, the aluminium goes into Solution as nitrate and can be precipitated with any of the above reagents. When strongly ignited, aluminium hydroxide loses water and forms the anhydride Al,0s, which is scarcely soluble at all in hydro- chloric and nitric acids. In warm concentrated sulphuric acid, with a little water, it will dissolve after long digestion. The ignited aluminium oxide, as well as the natural corundum, is most readily brought into solution by fusion with potassium pyrosulphate. The fusion is accomplished in the following way: About twelve times as much commercial potassium acid Sulphate as there is oxide to get into Solution is first heated by itself in a large platinum crucible over a small flame. The acid potassium sulphate melts readily, at about 300° C., gives off water (causing frothing), and becomes changed into potassium pyrosulphate: so,3; SO,-OK -Hoi Xo so.3% SO,-OK As soon as the frothing has ceased, the transformation into potas- sium pyrosulphate is complete. The dry oxide is now added to the crucible, and the heating is continued until the melt begins to Solidify * If sufficient nitrate was present, the carbon will be completely burnt to CO, and the residue then often contains undecomposed nitrate or nitrite. It would, therefore, be unwise to treat the residue in the platinum dish with hydrochloric acid, as aqua regia would be formed and the platinum would be dissolved. Consequently the residue is treated with nitric acid (or with hydrochloric acid in a porcelain dish). a ALUMINIUM. 75 (showing that a considerable amount of potassium sulphate, which melts much more difficultly than the pyrosulphate, has been formed), then the contents of the crucible are heated to a higher temperature and the heating is continued until the oxide has dissolved clear in the melt. By heating the pyrosulphate, SO, is given off, which in the nascent state and at the high temperature is very active: K.S.O. = K,SO,--SO,. After the reaction is finished, the melt contains the aluminium as aluminium sulphate in the presence of potassium sulphate, 3K.S.O.--Al,O, = Al,(SO.),4-3K,SO, and both of these substances can be brought into solution by treat- ing them with water. The ignited aluminium oxide can also be brought into solution by fusion with caustic alkalies: * Al,O,--6KOH=2Al(OK),4-3H,0. This last operation is usually carried on in a silver crucible, never in platinum, as the latter would be strongly attacked. Native Al,O, (corundum, ruby, Sapphire, emery) can be com- pletely brought into solution by fusion with potassium pyrosulphate. REACTIONS IN THE DRY WAY. Aluminium compounds, on being heated with sodium carbonate on charcoal before the blowpipe give a white, infusible, brightly glowing oxide, which, when moistened with cobalt nitrate solution and again heated, becomes a blue infusible mass (Thénard's blue). This test is best accomplished by taking a small piece of filter-paper with the aluminium oxide on it, fastening it in a platinum spiral, moistening the paper with a little dilute nitric acid and warming, so that the substance is dissolved as much as possible, impregnated in the fibres of the filter-paper, and spread over a considerable sur- face. A few drops of dilute cobalt nitrate solution are now added, and the mass is strongly heated. Traces of aluminium can be recognized by the mass turning to a blue color. The presence of colored metallic oxides hinders the reaction. Aluminium salts are not volatile, and do not color the flame. By ignition in the air, all aluminium salts, with the exception of the phosphate and silicate, are decomposed, leaving behind the oxide: 2AlCls--Os=Al,Os--3Cl2, 2Al(NO2), as Al,Os--3N,Os, Al,(SO4), - Al,Qs-H3SOs. 76 REACTIONS OF THE METALS. CHROMIUM, Cr. At. Wt. 52.1. Occurrence.—Chromium occurs in nature as chromite, Cr,0s' Fe0, isomorphous with spinel (See aluminium); as monoclinic crocoite, PbCrO4; and as laxmannite, a double compound of lead-copper phosphate and basic lead chromate, (Pb,Cu),(PO4)2. Pb2O(CrO2)2. Furthermore, it is found in Small quantities in many silicates, such as muscovites, biotites, augites, etc., and consequently in their weathering products, as in many kaolins, bauxite, etc. Metallic chromium is a whitish crystalline powder, having a specific gravity of 6.81. With oxygen it forms the following oxides: Chromous oxide Chromic oxide [Cr=O] Cr- O O Cr= O Chromium trioxide Chromium peroxide cº Cr,O, NO Chromium, therefore, exists in its compounds with a valence of either two, three, six, or seven. The oxides CrO and Cr,0s are basic anhydrides, and, on being dissolved in acids, yield the corresponding salts, the chromous and the chromic compounds. Chromium trioxide is the anhydride of the hypothetical chromic acid, H,CrO4, and forms chromates with bases. Chromium heptoxide has never been obtained pure; it is regarded as the anhydride of perchromic acid, HCrO4. I. CHROMOUS COMPOUNDs. Chromous oxide is known only in the form of its hydroxide, Cr(OH), which, on being dried, loses hydrogen and water, leaving behind chromic oxide: *—OHT H,--H,O+Cr2O3. Like the chromous hydroxide, all chromous compounds are CHROMIUM. 77 extremely unstable, being changed readily by contact with the air into chromic compounds. Only the halogen compounds, the phosphate, carbonate, and acetate, are known in the dry state; the Sulphate only in Solution. The acetate, Cr(C.H.O.), -i-H.O, is a reddish-brown, crystalline body, insoluble in water, but readily soluble in hydrochloric acid. This solution, as well as that of all chromous salts, absorbs oxygen with avidity, and is consequently used in gas analysis for the determi- nation of Oxygen in gas mixtures. Solutions of chromous com- pounds may be readily obtained by the reduction of chromic com- pounds with nascent hydrogen (zinc and acid), out of contact with the air. On account of the instability of these compounds the analytical chemist will rarely meet with them, so that we will pass over to the reactions of the chromic compounds. II. CHROMIC COMPOUNDs. All chromic compounds contain chromium as a trivalent element; they are colored either green or violet, and are soluble in water as a rule. The oa;ide, hydroacide, phosphate, anhydrous chloride, and the sulphate (which has been strongly heated in a stream of carbon dioxide gas), are insoluble in water. Violet chromium chloride obtained in the dry way is insoluble in acids; it dissolves readily in water containing a trace of chromous chloride, or in the presence of stannous chloride (tin and hydrochloric acid). By dissolving the grayish-green chromic hydroxide in acids, green solutions are always Obtained, which on long standing become greenish violet or violet but on boiling become green again. Chromic sulphate forms, with Sulphates of potassium, ammonium, cassium, or rubidium, the so- called chrome-alums, which crystallize in the regular system. These alums, like all other chromic salts, react acid in aqueous solution (hydrolysis). Chromic sulphide, Cr,S, can only be obtained in the dry way. On being treated with water it is quantitatively decomposed into hydroxide and hydrogen Sulphide. 78 REACTIONS OF THE METALS. REACTIONS IN THE WET WAY. 1. Ammonia and Ammonium Sulphide produce a grayish- green, gelatinous precipitate of chromic hydroxide: CrCl,--3NH,OH=3NH,Cl-HCr(OH), and 20rCl,--3(NH,),S+6HOH=6NH,Cl-H3H,S+2Cr(OH)2. Chromic hydroxide is somewhat soluble in excess of ammonia, forming a violet-colored Solution, particularly soluble when the ammonia is added to a violet Solution of a chromic salt, in the presence of ammonium Salts. This is caused by the formation of chromic-ammonium compounds, which, however, may be decom- posed by boiling the solution until the excess of ammonia has been driven off, when the chromium is quantitatively precipitated as hydroxide. In order, then, to precipitate the chromium from a solution as hydroxide, it is necessary to precipitate at a boiling temperature, and to use as little ammonia as possible. By ignition of chromic hydroxide, green chromic oxide is ob- tained, which after strong ignition is insoluble in acids. In order to bring it into Solution, it must be fused with potassium pyrosul- phate, as was described in the case of aluminium oxide; or with sodium carbonate and nitre in a platinum crucible, whereby it goes over into readily Sokuble sodium chromate: Gr,Os--2Na,COs--Os=2|Na,CrO4+2CO2. If the product of this last fusion is dissolved in water, acidified with hydrochloric acid, and boiled with alcohol, a green solution of chromic chloride will be obtained (page 84), from which the chromium can be precipitated as hydroxide with ammonia. On fusing with sodium carbonate and potassium nitrate in a platinum crucible, the latter will always be slightly attacked, so that a small amount of platinum will go into Solution with the fused mass; it can be removed, after the treatment with hydrochloric acid, by passing hydrogen sulphide into the boiling Solution, and filtering off the precipitated platinum sulphide. 2. Potassium and Sodium Hydroxides cause the same precipi- CHROMIUM. 79 tation as ammonia; only the precipitate is readily soluble in excess of the reagent, forming a green chromite: GrGls--3KOH=3KC1-i-Cr(OH), and * Cr(OH)2+3KOH z-> 3H,0+Cr(OK)s. Chromic hydroxide behaves here as a weak acid. The reaction is reversible, the presence of considerable water causing the reac- tion to go from right to left, particularly at the boiling tempera- ture. By boiling the dilute solution, complete hydrolysis takes place; the chromium is almost quantitatively precipitated as hy- droxide (differing from aluminium). 3. Alkali Carbonates, Barium Carbonate, Ammonium Sul- phide, and Alkali Thiosulphates precipitate chromic hydroxide, as with aluminium. 4. Alkali Phosphates give a greenish, amorphous precipitation of chromic phosphate: 2Na,HPO,--CrCls—3NaCl-H NaH,PO,--CrPO,. Chromic phosphate is readily soluble in mineral acids and in cold acetic acid. On boiling the acetic acid solution, chromic phosphate separates out again. 5. Alkali Acetates produce in solutions of chromic salts no pre- cipitation, neither in the cold nor on boiling. If, however, consid- erable amounts of aluminium and ferric salts are present at the same time, the chromium will be almost quantitatively precipitated with the iron and aluminium as basic acetate. In case, however, chromium predominates, only a part of the metals will be precipi- tated as basic salts; the filtrate will contain iron and aluminium with chromium. In the presence of chromium, the basic acetate separation is always uncertain. 6. Hydrogen Sulphide produces no precipitate in acid solutions of chromic salts. III. CHROMATES. Chromium trioxide, CrO3, forms red orthorhombic needles, which dissolve readily in water to an orange-red solution. If we neutralize this solution with potassium hydroxide, it becomes yellow, and on evaporation we obtain the beautiful yellow-colored 8o REACTIONS OF THE METALS. K.CrO4, the potassium salt of chromic acid, H,CrO4. If the yellow Solution of potassium chromate be acidified, and then allowed to crystallize, we obtain the beautiful orange-red prisms of triclinic potassium dichromate crystals, K2Cr,07. The aqueous solution of potassium chromate, K.CrO3, contains colorless potassium ions and yellow CrO, ions, ---|- – -- R.CrO4 z=2 K, --CrO4, while the aqueous solution of potassium dichromate, K.Cr,O, contains, in the presence of the colorless potassium ions, the Orange- red colored Cr2O, ions: ++ — — K.Cr,0,zi+2 K, --Cr,O. We are able, therefore, to determine from the color of a chro- mate solution the nature of the chromate ion which is present. Since, now, chromium trioxide dissolves in water with an orange- red color, we must assume that we have in the aqueous solution dichromic acid, or rather the dichromate ion, although the free acid itself has not yet been isolated. The normal chromic acid, H,CrO4, appears not to be capable of existence even in aqueous solution. Remark.-Although we may judge as to the color of the ions from the color of the solution, and often predict what the color of the solid salt will be, yet, on the other hand, we cannot tell what the color of the solution will be from that of the salt itself. Yellow lead iodide dissolves in water to a colorless Solution, and the yellow and red iodides of mercury, although only slightly Soluble, also do not yield colored Solutions. If the solution of a salt is colored, the salt itself will be colored; but the reverse is not always true. All chromates are insoluble in water, except those of the alkalies, calcium, strontium, and magnesium. All chromates dis- solve in nitric acid, except fused lead chromate, which dissolves with difficulty. Formation of Chromates. All chromium compounds may be readily Oxidized to chromates. According to whether the compound is soluble in water or not, different methods are used to effect the Oxidation. * CHROMIUM. 81 The oxidation of soluble chromates in alkaline solutions takes place— (a) By the halogens. If sodium or potassium hydroxide is added in excess to a solution of a chromic salt, and chlorine or bromine is conducted into the solution, the oxidation will be complete in a few minutes; the green chromite solution becomes a bright yellow: 2Cr (OK) 3 + 4KOH + 3Cl2 =6|KCl + 2H2O + 2K2CrO4 and 2Or (OK) 3 + 4KOH + 3Br2 =6|KBr-i- 2H2O + 2K2CrO4. Chromic compounds may be also oxidized by halogens in the presence of Sodium acetate, the reaction going extremely slowly in the cold, but very quickly on warming: CrCl3+8NaC.H.O., +4.H.O.--3Cl=6NaCl-F8HC, H2O,--Na,CrO4. (b) By hypochlorites (sodium hypochlorite, chloride of lime, etc.): 20r(OK), H-H,0+3NaOCl=3NaCl-i-2KOH+2K,CrO,. (c) By lead peroxide. The alkaline solution is boiled with lead peroxide: 2Cr(OK), H-3PbO,-2PbO--Pb(OK), H-2R,CrO. (d) By hydrogen peroxide, 2Cr(OK), H-3H,0,–2KOH+2H2O+2K,CrO, the reaction taking place on warming. (e) By freshly-precipitated manganese dioxide. The oxidation takes place on boiling the neutral or slightly acid solution: 2CrCls--3MnO,--2H,0=3MnOl,--2H,CrO,. Oaxidation in acid solution may be effected only by means of potassium chlorate in the presence of strong nitric acid. In the case of an insoluble chromium compound, such as strongly- ignited chromic oxide, or the mineral chromite, the oxidation is effected by means of fusion with sodium carbonate and either potassium nitrate or chlorate (cf. page 88). The alkali chromates 82 REACTIONS OF THE METALS. W * thus obtained are of a deep-yellow color, and are readily soluble in water. a. If acid is added to the solution of a normal chromate, the color is changed to orange in consequence of the formation of a dichro- mate: /OK +#XO-K so-Ho- O c.6.9% CrO,-OK 2 NOK Conversely, the dichromates of the alkalies are changed back into normal chromates by the addition of caustic alkali or alkali carbonates to the solution: K.Cr,O,--2KOH = H,0+2K,CrO. and K,Cr,O,--K,COs=CO,--2K,CrO4, the solution becoming yellow again. While alkaline solutions of chromic salts readily take on oxygen and become chromates, acid Solutions of chromates, on the other hand, give off oxygen just as readily and become reduced to chro- mic salts again. Chromic acid and chromates are strong oridizing agents in acid solutions. The equation of oxidation, reduced to its simplest form, is as follows: 2CrO3=Cr2O,--Os. This decomposition, which takes place readily on igniting chro- mium trioxide, takes place just as readily in aqueous solution in the presence of readily oxidizable substances; thus ferrous salts are, even in the cold, immediately oxidized to ferric salts: (a) By chromium trioxide: 20r.0s-i-6FeO=3Fe,Os-HCr,0s. (b) By potassium dichromate, R.Cr,O, =K.O. 20rCs, R.O.2CrOs--6FeO=3|Fe,Os--Cr,Os--K,O, CHROMIUM. 83 but it is necessary that sufficient acid be present in order to dis- solve the oxides which are formed: 20rO2+6FeSO4+6H,SO,-6H.O--3Fe2(SO4)2+Cr,(SOA)a, FC,Cr,0,--6FeSO44-7H,SO, =7H2O +K,SO,--3Fe2(SO4)2+Cr,(SO.)a. Similarly sulphurous acid, hydrogen sulphide, and hydriodic acid are oxidized at the ordinary temperature almost immediately; ozalic acid and alcohol on long standing, but quickly on heating; hydrochloric and hydrobromic acids only on warming. Sulphurous acid is oxidized to sulphuric acid: K.Cr,O,--3H,SO,--H,SO, = K,SO,--Cr,(SO.), H-4H.O. Hydrogen Sulphide is oxidized to sulphur, which separates and makes the solution turbid: |- " k K.Cr,O,--3H,S+4H,SO, = K,SO,--Cr,(SO.), H 7H,0+3S. Oxalic acid, H.C.O., is oxidized to carbonic acid: R.Cr,O,--3H,C,0,--4H,SO,-K,SO,--Cr,(SO4),4-7H,O+6CO,. Hydriodic acid is oxidized to iodine: K.Cr,0,--6HI+4H,SO,-K,SO,--Cr,(SO.),4-7FLO-i-61. With all these reactions a change of color from yellow to green takes place, the yellow chromate being reduced to green chromic salt. In the case of the oxidation of hydriodic acid in the cold the solution will not become green, but brown, on account of the sepa- ration of iodine. If, however, the solution is heated to boiling, the iodine is volatilized and the green color appears. Hydrochloric acid also is oxidized to chlorine by means of chromic acid: K.Cr,O,--14HCl=2KCl-i-2CrCls--7H,O--3Cl2. As this reaction takes place only on warming, it furnishes us with a convenient method for preparing small quantitics of chlorine for analytical purposes, because the evolution of chlorine ceases as soon as the lamp is taken away. It is necessary, however, to em- $4 REACTIONS OF THE METALS. ploy an excess of hydrochloric acid, as otherwise no chlorine will be evolved owing to the formation of potassium chlorochromate: /OK * . He yok >o + Hä-H,0+2CrO, CrO, NCI NOK which is decomposed on adding more hydrochloric acid: /OK * +12HCl=2KCl-H2CrCl,--6H,04-6Cl. Cl If alcohol and hydrochloric acid are allowed to act simulta- neously upon a chromate (the reaction takes place on gentle warm- ing without the evolution of chlorine), the alcohol is oxidized to aldehyde: K.Cr,0,--3C.H.OH+8HCl=7H,0+2KCl+2CrCl,--3CH,CHO. Aldehyde This last reaction is taken advantage of when it is desired to reduce a chromate, because the aldehyde (recognizable by its peculiar empyreumatic odor) and the excess of alcohol are easily removed by warming the solution, and the latter then contains simply the chromium and the metal of the chromate as chlorides. By boiling chromates with concentrated Sulphuric acid, reduc- tion takes place with evolution of oxygen: 2K2Cr2O7 --8H2SO4–2K2SO4+2Cr2(SO4)3 +8H2O +3O2. The behavior of free chromic acid towards hydrogen peroacide is characteristic. Neutral hydrogen peroxide added to a solution con- taining free chromic acid gives a blue color when the two liquids meet. On stirring the solution a green color is obtained, due to the mixture of blue perchromic acid and Orange-yellow chromic acid. This green color disappears gradually. The chromic acid is first of all oxidized to blue perchromic acid, but this gradually loses oxygen, forming chromic acid again. On adding an eaccess of hydrogen peroxide; first a blue color is obtained which disappears with evolution of oxygen, and the solution becomes a dirty green in color owing to a partial reduction of the perchromic CHROMIUM. 85 acid to chromic hydroxide, and the latter usually is precipitated of warming. If sufficient mineral acid is present all of the chromic acid is reduced to chromic salt. Since the formation of the intense-blue-colored perchromic acid takes place so readily, we have here a sensitive test for free chromic acid, which is made as follows: One or two cubic centimeters of hydrogen peroxide are treated with a few drops of dilute sulphuric acid and shaken with 2 c.c. of ether, after which a little of the chromate solution is added and the mixture again shaken. In the presence of 1/10 milligram of chromic acid, the upper liquid ether layer is colored intensely blue, and the reaction is noticeable with only 7/1000 milligram of chromic acid (Lehner). Only free chromic acid gives this test. We have, therefore, an excellent means for detecting free chromic acid in the presence of a dichromate. Commercial hydrogen peroxide always reacts acid, and must be neutralized before being used for this test; otherwise dichromates will give the same reaction. If, on the other hand, it is desired to test for chromic acid without regard to whether it is present in the free state or as a chromate, it is evident that the neutralization of the hydrogen peroxide is unnecessary; for in this case it is well to acidify the solution to be tested with sulphuric acid before the addi- tion of the hydrogen peroxide. Most chromates are insoluble in water, and exhibit characteristic colors; therefore the easiest test for chromium is to change it to a chromate. Reactions for the Precipitation of Chromic Acid. 1. Sulphuric Acid.—Dilute sulphuric acid causes, at the most, a change of color from yellow to orange, without any evolution of gas. Concentrated Sulphuric acid causes a change to orange color, and often the separation of red needles of CrO, ; the solution on being heated becomes green, the chromic acid being reduced to chromic salt with evolution of oxygen: 20r.0,--3H,SO,-3H,0+30+ Cr,(SO.). 2. Silver Nitrate produces in neutral chromate solutions a brownish-red precipitate of silver chromate: K.CrO,--2AgNO,-2KNO,--Ag,CrO, 86 REACTIONS OF THE METALS. soluble in ammonia and mineral acids (hydrochloric acid changes it into insoluble silver chloride and chromic acid), insoluble in acetic acid. If to a moderately concentrated solution of potassium dichromate, silver nitrate be added, a reddish-brown precipitate of silver dichromate is formed: K.C.O,4-2AgNO,-Ag:Cr;O,4-2KNO,” which, on being boiled with water, is changed into chromic acid and normal silver chromate: 2Ag,Cr,Oz-H H.O= H,Cr,O,--2Ag,CrO4. 3. Lead Acetate produces in solutions of normal chromates and dichromates a yellow precipitate of lead chromate, which is soluble in nitric acid but insoluble in acetic acid: K.CrO4+Pb(C.H.O.), -2KC, H2O,--PbCrO, and R.Cr,O,--2Pb(C.H.O.), +H,O=2KC.H.O., +2HC.H.O., +2PbCrO4. If lead nitrate is used instead of lead acetate, the precipitation is not complete unless sodium acetate is added. 4. Barium Chloride produces in solutions of normal chromates a yellow precipitate of barium chromate: R.CrO,--BaCl, -2KC1-i-BaCrO4, soluble in mineral acids, insoluble in acetic acid. From solutions of dichromates the precipitation is complete only on addition of an alkali acetate (cf. page 60). 5. Mercurous Nitrate produces in the cold a brown precipitate of basic mercurous chromate: 3K,CrO,--4Hg2(NO2), + H2O=6|KNO,--HgsCr,0s--2HNOs, which on being boiled goes over into fiery-red, neutral mercurous chromate: HgsCr,0s--2HNO2=Hg2(NO2), -i-H2O+3Hg,CrO,. * In the presence of sodium acetate neutral silver chromate is precipi- tated: K2Cr,01-1-4AgNO3 + H2O+2NaC2H5O2 = 2KNO3 +2NaNO3+2HC, H2O,--2Ag:CrO4. CHROMIUM. 87 Behavior of Chromium Trioxide and Chromates on Ignition. As already mentioned, chromium trioſcide is decomposed on igni- tion into chromic oxide and oxygen, 20rOa=Cr2O3+Oa. The chro- mates of ammonium and mercury behave quite similarly. Thus normal ammonium chromate on ignition is changed to chromic oxide, ammonia, nitrogen, and water. We can represent this reaction somewhat as follows: first, the chromate splits up into the basic and acid anhydrides: 2(NH,),CrO,-2(NH,),0+20r.0, The ammonium oxide is broken up into ammonia and water, while the chromium trioxide is decomposed into chromic oxide and OXygen, 2(NH3)2O=4NHs--2H,0,. and 20rCa-Cr,Os--Os. The oxygen, however, is used up immediately in oxidizing a part of the ammonia to water and nitrogen, 2NH,--Os=3H2O+N2, so that the whole reaction takes place according to the equation 2(NH,),CrO,-2NH,--2N--5H,0+Cr,0s. Ammonium Dichromate evolves only water and nitrogen: (NH4)2Cr,0,-4H2O+N,+Cr,0s. This decomposition takes place violently with scintillation. The chromic oxide which remains behind is very voluminous and re- minds one of tea-leaves; consequently it is sometimes called “tea- leaved oxide.” Mercurous Chromate is decomposed on ignition into chromic ozide, metallic mercury, and ozygen: 4Hg2CrO4 --> 2Cr2O3 + 8Hg + 502. 88 REACTIONS OF THE METALS. The Dichromates of the Alkalies are changed on ignition into normal chromate, chromic oxide, and oxygen: 4K2Cr2O7 = 4K2CrO4+2Cr2O3 + 3O2. REACTIONS OF CHROMIUM IN THE DRY WAY. All chromium compounds color the borax, or salt of phosphorus, bead an emerald green both in the Oxidizing and reducing flames. Heated with sodium carbonate on charcoal before the blowpipe, all chromium compounds yield a green slag, which after long heat- ing is changed to green infusible chromic oxide. By fusing with Soda and nitre in the loop of a platinum wire, all chromium com- pounds yield a yellow melt of alkali chromate: 2Cr2O3 +4Na2CO3 + 3O2=4Na2CrO4+4CO2. If the fused mass is dissolved in water and acidified with acetic acid, the solution will give with silver nitrate a reddish-brown precipitate of silver chromate. This reaction is very delicate and serves for the detection of minute traces of chromium. Cloth which has been dyed with a chromium mordant can be tested in this way; the ash from a thread five cra. long is sufficient to give this test. IRON, Fe. At. Wt. 55.9. Occurrence.—Native iron is rarely found. It occurs in basaltic rocks; also in meteorites, associated with nickel, cobalt, carbon, Sulphur, and phosphorus. The most important iron ores are the oxides and the Sulphides. Of these may be mentioned:— Hematite, Fe,0s, isomorphous with corundum; magnetite, Fe2O4, isomorphous with spinel; göthite, FeFIO, isomorphous with diaspore and manganite; limonite (bog ore), Fe(OH)3, which is used in the purification of illuminating-gas; pyrite, FeS2, which crystallizes in the isometric system; marcasite, FeS2, orthorhombic. Iron disulphide is, therefore, dimorphous. Another important iron ore is siderite, FeCOs, which is rhombohedral. IRON. 89. The metallic iron of commerce is never pure, but contains iron carbide, iron sulphide, iron phosphide, iron silicide, corresponding nanganese compounds and graphite, etc. On dissolving commercial iron in acids (H,SO, HCl), hydrogen, contaminated with small amounts of hydrocarbons, hydrogen sulphide, mercaptans, phosphuretted hydrogen, and silicon hydride are given off, and the latter compounds give to the gas its unpleasant odor. There remains almost always an undissolved residue con- sisting chiefly of carbon. Iron is bivalent or trivalent, and forms the following oxides: * Iron protoxide Iron sesquioxide sº ë #. oxide Ferric . Ferrous-ferric oxide FeO Fe=O Fe=O X) NO Fe=O Dissolving these oxides in acids, we obtain the corresponding iron salts; thus ferrous oxide gives with hydrochloric acid ferrous chloride, Fe0+2HCl= H,O--FeCl,; ferric oxide gives ferric chloride, Fe,O,--6HCl=3H,0+2FeCls, while ferrous-ferric oxide yields- a mixture of ferrous and ferric, chlorides: Fe2O,--8HCl=4H2O+2FeCla--FeCl2. Iron, therefore, forms two series of salts: first, the ferrous, de-, rived from ferrous oxide, containing bivalent iron; second, the ferric, derived from ferric oxide, containing trivalent iron. These two series of salts show quite different behavior towards reagents; therefore, we will treat them separately. * Iron trioxide, Fe0s, containing hexavalent iron has never been isolated. It plays the part of an acid anhydride in ferrates of the general formula. RAFeO4, which are decomposable by water. 90 REACTIONS OF THE METALS. A. FERROUS CoMPOUNDs. Ferrous compounds, which may be prepared by dissolving in acids metallic iron, ferrous oxide, ferrous hydroxide, ferrous carbo- nate, ferrous Sulphide, etc., are usually greenish in the crystallized state, while in the anhydrous condition they are white, yellow, or bluish; in concentrated solution they are green; in dilute solutions almost colorless. Ferrous compounds exhibit a strong tendency to change over into ferric Salts; they are strong reducing agents. FEACTIONS IN THE WET WAY. 1. Ammonia produces in neutral solutions an incomplete pre- cipitation of white ferrous hydroxide: FeCl2 +2NH; +2H2O =2 Fe (OH)2 +2NH4C1. Ferrous salts in this respect are similar to those of magnesium (cf. page 50). In the presence of ammonium chloride the reaction takes place in the direction from left to right; ammonia, therefore, causes no precipitation with ferrous salts out of contact with the air, provided sufficient ammonium chloride is present. On expo- sure to the air, however, a turbidity is soon formed, green at first, then almost black, and finally becoming brown. The small amount of ferrous hydroxide contained in the solution is oxidized by the air, forming at first black ferrous-ferric hydroxide and finally brown ferric hydroxide. t 2. Potassium and Sodium Hydroxides produce, if air is ex- cluded, complete precipitation of white ferrous hydroxide, FeSO,--2KOH=Fe(OH)2+ K.S.O., which is quickly oxidized by the air into ferric hydroxide. 3. Hydrogen Sulphide produces no precipitation in acid solu- tions of ferrous salts; in dilute neutral solutions a small amount of black ferrous sulphide is precipitated; but if the Solution contains considerable alkali acetate, hydrogen sulphide precipitates more of the iron as ferrous sulphide (but not all of it), in spite of the fact that ferrous sulphide is readily soluble in acetic acid. This inter- esting fact is an instructive illustration of the law of chemical mass action. } IRON. 9I The solution of ferrous sulphide in acetic acid may be explained as follows: No substance is absolutely insoluble in water; but the amount dissolved cannot always be detected by chemical means, though it can be by physical methods. This small amount, in the case of such an insoluble body as ferrous sulphide, is practically completely dissociated into its ions. If we allow acetic acid to act upon ferrous Sulphide, the positive charges of the hydrogen ions from the acetic acid will be neutralized by the negative charges of the Sulphur ions from the ferrous sulphide, with the formation of undissociated hydrogen sulphide which escapes as a gas from the Solution: + + — — —- Eº + * + + — - Fe+S + H+C.H.O.·H+C.H.O.-Fe+C.H.O.E.C.H.O.4-H.S. f —-v--—–’ y Ferrous sulphide Acetic acid Ferrous acetate The equilibrium is disturbed by the escape of the gas; to restore it, more ferrous Sulphide must go into solution, so that the above reac- tion is repeated until finally all of the ferrous sulphide has become dissolved. The solution is effected by means of free hydrogen ions. If, now, the concentration of the acetic ions is increased, the extent of the dissociation of the acetic acid will be lessened, in consequence of the law of chemical mass action. This means that there will be fewer hydrogen ions in the solution, so that the ferrous sulphide will be less soluble. The concentration of the acetate ions is in- creased by the addition of largely-dissociated alkali acetate. 4. Ammonium Sulphide precipitates iron completely as black ferrous sulphide: *. FeCl,+(NH,),S=2NH,Cl–H FeS, which is readily soluble in acids with evolution of hydrogen sul- phide. In moist air it turns slightly brown, a part of the sulphur separates out, and a basic ferrous sulphide is formed. 5. Alkali Carbonates precipitate the white carbonate: FeCl,--Na,CO,-2NaCl-i-FeCO, which in contact with the air becomes green, then brown: 4Fe(O +6H2O+O2=4CO2+4Fe(OH)3, 92 REACTIONS OF THE METALS. being converted into ferric hydroxide with loss of carbonic anhy- dride. Ferrous carbonate is soluble in carbonic acid, forming ferrous bicarbonate: FeCO,-- H,CO, == FeH,(CO.), j a compound which is found in many natural waters, but which, like the normal carbonate, is decomposed by atmospheric oxygen with separation of ferric hydroxide: 4FeH2(CO3)2 +2H O -O2=8CO2+4Fe(OH)3. Consequently a mineral water which contains ferrous bicarbonate, if allowed to stand in contact with the air, will become turbid owing to the deposition of ferric hydroxide. To prevent this, the bottle ' must be filled with water and tightly corked so that no trace of air can get in. The ferric hydroxide is insoluble in carbonic acid, At this point we will say a few words with regard to the rusting, of iron. Bright iron rusts if it comes in contact with moist air, as is well known. The process of rusting is a cyclical one, and three factors play an important part: 1. An acid; 2. Water; 3. Oxygen. The process of rusting is always started by an acid (even the weak carbonic acid suffices); the acid changes the metal into a ferrous salt with evolution of hydrogen: 2Fe--2H,CO,-2FeCO,--2H,. \ Water and oxygen now act upon the ferrous salt, causing the iron in this salt to separate out as ferric hydroxide, setting free the same amount of acid which was used in forming the ferrous salt: 2FeCO,--5H,0+O=2Fe(OH),4-2H,CO, The acid which is set free again acts upon the metal, forming more ferrous salt, which is again decomposed, forming more rust. A very small amount of acid, therefore, suffices to rust a large amount IRON. 93 of iron. If the acid is lacking, the iron will not rust. If we desire to prevent this rusting, we must neutralize the acid, e.g., add milk of lime. Iron remains bright under an alkali. That carbonic acid acts upon bright iron can be demonstrated by exposing the metal to water containing carbonic acid but free from air. After several hours a measurable amount of hydrogen gas will be given off, and the solution will contain an equivalent amount of ferrous bicarbonate. 6. Potassium Cyanide precipitates yellowish-brown ferrous cy- anide: FeCl,+2KCN=2KCl+Fe(CN), which is soluble in excess of the reagent, forming potassium ferro- cyanide: Fe(CN),4-4KCN=E,Fe(CN)]. The potassium ferrocyanide (formed in the last reaction) is not a ferrous salt, but a potassium salt. It gives none of the above re- actions for ferrous Salts; SO that the aqueous solution cannot con- tain ferrous ions, but K and Fe(CN), ions. This property of the metallic cyanide forming with cyanides of the alkalies a complex salt, is a very common one. The cyanides of silver, nickel, iron (ferrous and ferric), and cobalt all dissolve in potassium cyanide, forming the following complex Salts: [Ag(CN), K; [Ni(CN), K,; [Fe"(CN)]K,; [Fe’”(CN).IK,; [Co”(CN), K. These compounds must be regarded as Salts of the correspond- ing acids: Ag(CN), H; [Ni(GN), H.; [Fe"(CN).]H,; [Fe"(€N),TH,; [Co”(CN), H, And it is possible, as ā'matter of fact, to isolate the last three acids, though the two former have never been prepared, as they imme- diately break down into metallic cyanide and hydrocyanic acid, just as carbonic acid is decomposed into water and carbon dioxide. With iron, therefore, there are two series of complex cyanogen compounds, the ferrocyanides and the ferricyanides. The ferro- 94 REACTIONS OF THE METALS. cyanic derivatives contain the quadrivalent ferrocyanide group [Fe’’(CN)6]T, and the ferricyanides contain the trivalent ferri- cyanide group [Fe’”(CN), E. By Saturating the free valencies with hydrogen, we obtain the corresponding acids: H [Fe"(CN).Iti and [Fe"(CN), H. NH NH Hydroferrocyanic acid Hydroferricyanic acid, By substituting metals for the hydrogen atoms we obtain the corresponding Salts: FC Zi. /K [Fe"(CN).I.: [Fe’”(CN), K. Nº. NK Potassium ferrocyanide Potassium ferricyanide (yellow prussiate of potash) (red prussiate of potash) As has been mentioned, potassium ferrocyanide yields in aqueous solution the ferrocyanide ion [Fe’’(CN)6] and potassium ions, while potassium ferricyanide yields ferricyanide [Fe’’(CN)6] and potassium ions. The solubility of their alkali and alkaline-earth salts, and the insolubility and color of the Salts of the heavy metals (especially with both ferric and ferrous iron), are very characteristic of ferro- and ferricyanides. 7. Potassium Ferrocyanide, K, Fe(CN)e, produces in solutions of ferrous salts, with complete exclusion of air, a white precipitate of jerrous potassium ferrocyanide or ferrous ferrocyanide, according as to whether one or two molecules of ferrous Salt react with one mole- cule of potassium ferrocyanide: z: /Fe [Fe(CN).I. + SO, Fe=R,SO, +[Fe(CN), HR Nº. NK OI’ K SO, Fe Fe(CN),ſº + 4 =2K.SO, + [Fe(CN).K : NK SO,Fe o IRON. 95 Although both of the above salts are white, a light-blue color is almost always obtained, because it is immediately oxidized some- what by the air, forming the ferric salt of hydroferrocyanic acid (Prussian blue): 3|Fe(CN)6]Fez-H.3H2O+3O=[Fe(CN).], Fe,” +2Fe(OH)2 Prussian blue OI’ TeGN).E.13H,0+30– =[Fe(CN)6]. Fe,” +3|Fe(CN)6]K,--2Fe(OH)3. 8. Potassium Ferricyanide added to solutions of ferrous salts yields blue compounds of different composition according to the relative amounts of the reacting substances and the temperature of the solution. When the reagent is in excess, soluble Turnbull’s blue is formed from neutral solution which is identical with soluble Prussian blue * and has the formula [Fe’’(CN)6KH2]–Fe’”(OH)2. When an excess of ferrous salt is treated with potassium ferri- cyanide, insoluble Turnbull's blue is formed. The latter, in spite of contradiction in the literature, appears to be identical with insoluble Prussian blue and is the ferric salt of ferrocyanic acid and has the formula FeaſTe(CN)6]3+10H2O. The latter compound is always obtained at 15° C. when an excess of acid is not present. From hot, acid solutions complicated mix- tures of different blue compounds are obtained. In forming the blue compounds the ferrous salt is usually oxidized by the potas- sium ferricyanide. * Hofmann, Heine, and Höchtlen, Annalen, 337 (1904), p. 1. 96 REACTIONS OF THE METALS. 9. Potassium Sulphocyanate gives no reaction with ferrous salts (note difference from ferric Salts). As has been stated, ferrous Salts are oxidized by the air to ferric salts; thus ferrous Sulphate is gradually changed into brown, basic ferric sulphate, Feso, o "S.P. fesot 9 p. 29, , Fe—SO, which is insoluble in water. Consequently it often happens that ferrous sulphate will not dissolve in water to a clear solution, but gives a brown turbid solution, becoming clear on the addition of acid; the basic ferric Salt being changed to a soluble neutral salt: Fe=SO, O + H2SO4= H.O.--Fe2(SO4)a. Fe—SO, Such a solution, which then contains ferric sulphate, reacts with potassium sulphocyanate (cf. page 99). In order to free the solution from ferric salt, it may be boiled with metallic iron, with exclusion of air, whereby the ferric Salt is changed into ferrous salt: Fe,(SOI), H-Fe=3FeSO4. By means of strong oxidizing agents, ferrous salts can be quickly and completely changed into ferric Salts, as was shown in the intro- duction (cf. page 4). Detection of Ferrous Oxide in the Presence of Metallic Iron. The mixture is treated with a large excess of a neutral solution of mercuric chloride and warmed on the water-bath, whereby the metallic iron goes into solution as ferrous chloride: 2HgCl,--Fe=FeCl,+Hg,Cl,. The residue is filtered off and the filtrate tested with potassium ferricyanide, a precipitate of Turnbull’s blue showing that metal- lic iron was originally present. The residue is now washed with cold water, until all of the §errous chloride has been dissolved out, and is then treated with IRON. 97 dilute hydrochloric acid.” If the solution now gives a precipitate of Turnbull’s blue with potassium ferricyanide, ferrous oxide was present. B. FERRIC COMPOUNDs. Ferric oxide, Fe,0s, is reddish brown, becomes grayish black On strong ignition, but on being pulverized appears red again. The ferric Salts are usually yellow or brown, but the ferric am- monium alum is pale violet. Ferric salts are yellowish brown in aqueous Solution, and the solution reacts acid (hydrolysis). Dilu- tion and warming favor the hydrolysis, so that all strongly-diluted ferric Salts deposit basic salts on being boiled: Fe, SOAls--H,O z=2 Fe2(SO4)2O+H.SO4. With ferric salts of the weaker acids, often all of the iron is precipitated as a basic salt; thus the acetate, on being boiled in a dilute solution, reacts as follows: Fe(C.H.O.)a-i-2H,0=Fe(OH)2(C.H.O.)+2HC.H.O. By the addition of acid all basic salts may be changed back into neutral salts. REACTIONS OF FERRIC SALTS IN THE WET WAY. 1. Ammonia precipitates brown, gelatinous ferric hydroxide: FeCl,--3NH,OH=3NH,Cl-LFe(OH),. Ferric hydroxide is readily soluble in acids. On ignition it loses water and is changed to oxide, which is very difficultly soluble in dilute acids. It is best brought into solution by long-continued heating with concentrated hydrochloric acid, at a moderately warm temperature. 2. Potassium and Sodium Hydroxide also precipitate ferric hydroxide. —--- * If hydrogen is given off, some metallic iron is still present, so that the treatment with HgCl, must be repeated, 98 REACTIONS OF THE METALS. 3. Sodium Carbonate produces a brown precipitate of basic carbonate, which at the boiling temperature is completely decom- posed hydrolytically into hydroxide and carbon dioxide: 2FeCls--3Na,CO,--3H,0=2Fe(OH)2+ 6NaCl-H3CO,. 4. Zinc Oxide and Mercuric Oxide also precipitate the iron as hydroxide: 2FeCl,--3ZnO--3H,0=3ZnCl2 +2Fe(OH)3. This reaction is frequently used in quantitative analysis. 5. Sodium Phosphate precipitates yellowish-white ferric phos- phate: FeCl,--2Na,HPO,-3NaCl-i-NaH,PO,--FeFO,. Ferric phosphate is insoluble in acetic acid, but readily soluble in mineral acids. The precipitation of iron with sodium hydrogen phosphate is consequently only complete when a large excess of the precipitant is employed, or when sodium acetate is added: FeCla-i-Na,HPO,--NaC.H.O. =3NaCl-H HC, H,0,--FeFO,. In this last case all the iron and all the phosphoric acid are precipitated. The reaction is often made use of in order to quanti- tatively precipitate phosphoric acid. An excess of the disodium phosphate will also cause complete precipitation of iron as phos- phate, if the phosphate solution is previously exactly neutralized with ammonia: Na,HPO,--NH.OH=H.O.--Na,NH,PO, and Na,NH, PO,--FeCls=2NaCl-i-NH,Cl-i-FePO,. If, however, an excess of ammonia is added to the iron solution with the sodium phosphate, the precipitation of iron is incomplete, because the ferric phosphate dissolves in the excess of sodium phos- phate, in the presence of ammonia (or ammonium carbonate), with a brown color and formation of a complex Salt. Ferric phosphate is transformed by ammonia into a brown basic phosphate, and by potassium hydrate almost completely into IRON. 99 ferric hydroxide and potassium phosphate; while by fusion with caustic alkali or alkali carbonate it is completely decomposed. 6. Alkali Acetates produce in cold, neutral solutions a dark- brown coloration, and on boiling the dilute solution all of the iron separates as basic acetate: FeCla-i-3NaC.H.O. =3NaCl--Fe(C.H.O.), (in the cold), Fe(C.H.O.)a--2H2O=2HC,H,O2+Fe(OH)2C.H.O., (on boiling). The presence of organic oxy-acids (tartaric, malic, citric, etc.) and of polyatomic alcohols (glycerine, erythrite, mannite, Sugars, etc.) prevent all of the above-mentioned reactions, because com- plex salts are formed in which the iron is present in the form of a complex anion (cf. aluminium, page 72). 7. Potassium Sulphocyanate, KCNS, produces in solutions of ferric salts a blood-red coloration: FeCl,--3KCNS=Fe(CNS), H-3KCl. This action is reversible; the red color being most intense when an excess of ferric salt, or of potassium sulphocyanate is present. If the solution is shaken with ether, the Fe(CNS), goes into the ether. Ferric sulphocyanate combines readily with potassium sulphocyanate, forming complex potassium iron Sulphocyanate: Fe(CNS),4-3KCNS=[Fe(CNS).]Ks,” analogous to potassium ferricyanide: , [Fe(CN)6]Ks. This complex salt is insoluble in ether, the Fe(CNS), only being soluble therein; so that the red color is due to the formation of the ferric sulphocyanate and not to the complex salt. This reaction is extremely sensitive, but cannot always be relied on. If the solution contains considerable alkali acetate, the colora- tion cannot be recognized. The presence of organic oxy-compounds (tartaric acid, etc.) prevents the reaction in neutral solutions, but * Fe(CNS), Ks + 4H.O. Cf. Rosenheim, Zeit. für anorg. Ch. 1901, Bd. XXVII, p. 298. IOO REACTIONS OF THE METALS. not in acid solutions. In the presence of mercuric chloride the red color disappears entirely; the mercuric chloride reacts with the ferric Sulphocyanate, forming a colorless, soluble mercuric double Salt : 2Fe(CNS),46HgCl,-2Fec,+3(HgCNS), HgCl). 8. Potassium Ferrocyanide, K.Fe(CN)e, produces in neutral or acid solutions of ferric Salts an intense blue precipitation of Prussian blue: z: ºs ºrg # （ºr, º: "…, &rd º A/ Fe [Fe ©E; 㺠-12K64 [Fe (CN).K e l— P X gºre [Fe"(CN)4sre // Cl [Fe"(GN ).f: ër. Nk Ci/ Prussian blue, the ferric salt of hydroferrocyanic acid, is in- soluble in water, but soluble in oxalic acid with a blue color (blue ink). It is also soluble in concentrated hydrochloric acid, but is precipitated again on dilution. As the ferric salt of hydroferro- cyanic acid it behaves like other ferric salts to the hydroxides of the alkalies; ferric hydroxide and the alkali salt of hydroferro- cyanic acid being formed: [Fe(CN)]. Fe,”--12KOH=4Fe(OH),4-3|Fe(CN), K., . similar to (SO.), Fe,4-6KOH=2Fe(OH)2+3(SOJK,. Besides the insoluble Prussian blue, there exists a soluble Prussian blue, which is obtained by adding the ferric salt to an excess of potassium ferrocyanide. The same compound is formed by adding a solution of ferrous salt to more than an equivalent amount of potassium ferricyanide. IRON. IOI According to Hofmann, Heine, and Höchteln, this soluble blue has the formula * [Fe(CN)6]KH2. Fe(OH)2. This substance, although soluble in water with a blue color, is insoluble in salt solutions. If, therefore, we desire to separate it from the solution, we “salt it out,” i.e., we add considerable salt to the solution (preferably potassium chloride), whereby it is changed to the insoluble form, and may be filtered from the solution. 9. Potassium Ferricyanide, KaRe(CN), produces no precipita- tion in solutions of ferric salts, only a brown coloration (differing from ferrous salts): [Fe’”(CN).]Ka--FeClge|Fe’”(CN).]Feſ”--3KCl. 10. Ammonium Sulphide added to an acid solution gives a pre- cipitate of ferrous sulphide, FeS, but if the solution becomes am- moniacal ferric sulphide, Fe2S3, is precipitated. In the ordinary course of analysis the precipitate produced by ammonia and am- monium sulphide consists of ferric sulphide, 2FeCl3 + 3(NH4) 2S == Fe2S3 +6NH4C1, which is soluble in cold, dilute hydrochloric acid, forming ferrous chloride and sulphur, Fe2S3 +4HCl=2FeCl2 +2H2S +S. 11. Hydrogen Sulphide reduces ferric salts to ferrous salts, with separation of Sulphur: 2FeCl,--H.S=2FeCl,+2HC1+S. Besides hydrogen sulphide, many other substances (nascent hydrogen, stannous chloride, sulphurous acid, hydriodic acid, etc.) will reduce ferric salts, as was shown on page 7. 12. Sodium. Thiosulphate, Na,S,Os, colors neutral ferric solu- tions a violet red, but the color disappears quickly and the solution then contains ferrous salt and sodium tetrathiomate : 2Na,S,Os--2FeCls=2|NaCl-H2Fe(Cl2+Na,S,Os. * Annalen, 337 (1904), p. 1. f Louis Gedel, Ueber Schwefeleisen, Karlsruhe, 1905. I O2 REACTIONS OF THE METALS. The composition of the violet-red body which is first formed is unknown; perhaps it is ferric thiosulphate. As we have seen, there exist a number of iron compounds which contain the metal as a complex ion, so that it cannot be de- tected by the ordinary reagents. The complex oxy-organic com- pounds, as well as the ferro- and ferricyanide compounds, belong to this class of compounds. If it is a question of proving the presence of iron in such a com- pound, a different method should be used in the case of an organic oxy-compound from that in the case of a ferro- or ferricyanide. If organic substances are present, the iron is precipitated as sulphide by means of ammonium sulphide; or the organic matter is first removed by ignition, whereby metallic iron and carbon are obtained. In case we have a ferro- or ferricyanide, the iron cannot even be precipitated by means of ammonium sulphide; the compound must be completely destroyed before it will be possible to detect the presence of iron by any of the Ordinary methods. This may be accomplished (a) by ignition, (b) by fusion with potash or soda, or (c) by heating strongly with concentrated sulphuric acid. (a) Decomposition by Ignition.—The ferrocyanides are decom- posed (with evolution of nitrogen) into potassium cyanide and carbide of iron: [Fe(CN), JK, -4KCN+FeC,--N,. The ferricyanides also leave behind iron carbide and potassium cyanide, but evolve cyanogen as well as nitrogen: 2[Fe(CN)6]K3 =6|KCN +2FeC2+(CN)2+2N2. The residue from the ignition is treated first with water, whereby the potassium cyanide goes into solution, leaving behind the iron carbide; this is filtered off and treated with hydrochloric acid. The iron goes into solution as ferrous chloride, hydrocarbons are given off, and there remains some carbon. The above decomposition can be imagined to take place as follows: By heating potassium ferrocyanide, it is decomposed first into IRON. Io.3 potassium cyanide and ferrous cyanide, while the latter on further heating is changed to iron carbide and nitrogen: (a) [Fe(CN)]K, =4KCN+Fe(CN)2; (3) Fe(CN),= FeC, 4-N2. Potassium ferricyanide is decomposed into potassium cyanide and the very unstable ferric cyanide, which splits off cyanogen and becomes ferrous cyanide; the latter is decomposed, as before, into iron carbide and nitrogen: (a) [Fe(CN), K, H3KCN+Fe(CN), ; (6) Fe(CN), - Fe(CN),4-CN; (y) Fe(CN),= FeC,--N,. (b) Decomposition by Means of Fusion with Potash.--The sub- stance is mixed with an equal amount of potash and heated in a porcelain crucible until the fusion is quiet. By this means a mix- ture of potassium cyanide and potassium cyanate (both soluble in water) is obtained in the presence of metallic iron: [Fe(CN), K, H-K,CO,-5KCN+ KCNO+CO,--Fe. The melt is therefore first treated with water, and the iron remaining behind is dissolved in hydrochloric acid. (c) Decomposition by Heating with Concentrated Sulphuric Acid. —By heating with concentrated Sulphuric acid all complex cyano- gen compounds may be decomposed. By this means the metal present is changed into Sulphate, the nitrogen of the cyanogen into ammonium sulphate, while the carbon of the cyanogen escapes as carbon monoxide: [Fe(CN)6]K.--6H,SO,--6H.O = =2K,SO,--FesO4+3(NH3),SO,--6CO, 2[Fe(CN).]K.--12H,SO,-- 12H,0= =3K,SO,--Fe2(SO4)2+6(NHA),SO,--12CO. The treatment with concentrated Sulphuric acid is best accom- plished in a porcelain crucible placed in an inclined position over the flame, and the flame directed against the upper part of the cru- IO4 REACTIONS OF THE METALS. cible. The heating is continued until fumes of Sulphuric acid cease to come off. The residue, which consists of an alkali Sulphate, and ferrous or ferric sulphate (both being in an anhydrous condition), is treated with a little concentrated Sulphuric acid, warmed some- what, and water added little by little. In this way the Sulphate is readily ,brought into Solution; while with water alone Solution takes place only with difficulty. REACTIONS IN THE DRY WAY. The borax (or sodium metaphosphate) bead, containing a small amount of an iron salt, is yellow while hot and colorless when cold in the oxidizing flame, and faint green in the reducing flame. When strongly saturated, however, the bead in the oxidizing flame is brown while hot, yellow when cold; and in the reducing flame it becomes bottle-green. Heated on charcoal with soda before the blowpipe, all iron compounds leave a gray particle of metallic iron, which is usually difficult to See, but can be separated from the charcoal by means of the magnet. The reduction on the charcoal stick, as described on page 24, is a much more delicate test. URANIUM, Ur. At. Wt. 239.6. Occurrence.—Uranium occurs in nature chiefly in the mineral pitch-blende, U.Os; but it is also found in a few rare minerals, uranite, (UO,),Cup,0s + 8H,0; Samarskite (a niobite of iron, yttrium, cerium, and erbium with varying amounts of uranium); and liebigite, U(CO.), :20aCO,--10H.O. Klaproth showed, in 1789, that the mineral pitch-blende con- tained a new metal, which he called uranium. By heating the oxide with reducing agents he obtained a brown, almost copper- red, substance, which he took to be the metal, and it indeed does behave like a metal, dissolving in acids in contact with the air, forming yellowish-green salts. It was not until 1842 that it was shown by Peligot that this reddish-brown body was not the metal uranium, but its dioxide. The hexavalent metal itself was obtained by Peligot, by reducing the tetrachloride with sodium, as a gray powder of specific gravity 18.33 and melting at 1500° C. URANIUM. IOS Out of contact with the air, uranium dioxide (uranyl) dissolves in strong acids, forming uranous salts. UO,--4HCl=2H,0+UCl, UO,--2H,SO,-2H,0+U(SO.), The uranous salts are extremely unstable, and on being ex- posed to the air change rapidly, forming uranyl salts: /Cl UC1,--O--H,O=UO, +2HCl, NCl U(SO.), H-O--H,O=UO,SO,--H,SO,. We will consider the reactions of the uranyl salts only. The uranous compounds are important for the quantitative determi- nation of uranium itself, and will therefore be considered later. Besides uranyl (or uranium dioxide) uranium forms a trioxide, UO, which can be regarded as uranyl oxide, UO,O. It dissolves in acids, forming uranyl Salts: /Cl UO,OH-2HCl =UO, +H.O, NCI UO,O-i-H,SO,-UO,SO,-i-H.O. By igniting the oxides of uranium in contact with air, dark- green urano-uranic oxide, UAOs or (2UO,-UO,), is obtained, which out of contact with the air dissolves in strong acids, forming uranous and uranic salts: (2UO,-UO)+4H,SO,-2UO,SO,--U(SO),4-4H.O. By dissolving in aqua regia, uranyl chloride is obtained: ; 3U,Os-1-18HCl·H2HNO,-9UO,Cl,--2NO+10H,0. All uranyl compounds are colored yellow or yellowish green. Most of them are soluble in water, but the oxides, the sulphide, phosphate, and uranates are insoluble. In mineral acids all uranium compounds are soluble, with the exception of the ferro- cyanide. Ioč REACTIONS OF THE METALS. REACTIONS OF URANYL COMPOUNDS IN THE WET WAY. 1. Potassium Hydroxide precipitates yellow amorphous po- tassium uranate. The reaction takes place in three stages: (a) At first uranyl hydroxide is formed, /Cl /OH Uó, Hä-2Kel-Uó. NCl NOH which immediately splits off water and becomes uranic acid, /OH -H,04 × , /OH |UO,-OH UO, 2 NOH which combines with more potassium hydroxide and forms potas- sium uranate: (c) UO,-OH HOK UO,-OK Y O + -2H,0+ X O . UO,-OH HOK UO,-OK 2. Ammonia precipitates yellow, amorphous ammonium ura- Inate: /NO, UO,-ONH, 2UO, + 6NH,OH=4NH, NO,--3H,0+ - O . NO, UO,-ONH, The alkali uranates are soluble in alkali carbonates, particu- larly in ammonium carbonate, with the formation of complex salts: º UO,-ONH, > O +6(NH,),CO,--3H,0= UO,-ONH, =2|UO,(CO.), (NH,),4-6NH,OH. Consequently, in the presence of Sufficient alkali carbonate, am- monia fails to precipitate uranium. Tartaric and citric acids (and (JRANIUM. Io? other organic substances) also prevent the precipitation with 'am- monia and caustic alkalies, as with iron, chromium, and aluminium. 3. Sodium Carbonate produces in concentrated Solutions an orange-yellow precipitate of uranyl sodium carbonate: UO,(NO.),4-3Na,CO,-2NaNO,--UO,(CO.),Na,. Uranyl sodium carbonate is soluble in considerable water, so that no precipitate is formed from dilute solutions. It is still more soluble in alkali carbonate solution, particularly in a bicarbonate solution. From such solutions sodium hydroxide precipitates the uranate, but ammonia does not. 4. Barium Carbonate precipitates in the cold all of the uranium, probably as uranyl barium carbonate: UO2(NO2)2+3BaCO3=Ba(NO3)2+[UO,(CO2)2]Baz. 5. Ammonium Sulphide precipitates brown uranyl sulphide, UO2(NO2)2+(NH3)2S-2NH4NO2+UO,S, soluble in dilute acids and in ammonium carbonate: UO.S.--3(NH,),CO,-(NHI),S+[UO,(CO2), (NHI). Ammonium sulphide, therefore, produces no precipitate in solu- tions of wranyl salts in the presence of ammonium carbonate. 6. Sodium Phosphate precipitates yellowish-white uranyl phosphate, Na NO, PoéNº.1” SUO,-2NaNo.4 PO.69% “NH NO, / 2 3. *NH while, in the presence of ammonium acetate, uranyl ammonium phosphate is precipitated: NH COONH, 3. /UO =2NaNO,-- --PO ? . *" doorſ" ‘‘‘NNH, Na. CH, Poéºtor | == CH Io& REACTIONS OF THE METALS. Both precipitates are insoluble in acetic acid, but soluble in min- eral acids. 7. Potassium Ferrocyanide produces a brown precipitate, or, In very dilute solutions, a brownish-red coloration, K Mn / § Mn,Ti, +4HNO,-2Mn(NO),4-Mn ºff, Nº. XMn O H zº /OH Nóir NOH * Mn2Os, however, also plays the part of a basic anhydride, forming a sul- phate, Mn2(SO)s, which is decomposed by water into sulphuric acid and Mn,0s. Furthermore, potassium-manganese and ammonium-manganese alums are known. .A.<) 3. II6 REACTIONS OF THE METALS. MnO, stands in the same relation to H.MnO, as CO, to H,CO, as SiO, to H,SiO, and as SnO, to H,SnO, MnO, therefore, behaves like an acid anhydride and is isomorphous with tinstone, SnO, crystallizing as polianite in the tetragonal system. Like SnO, (which See), manganese dioxide behaves partly as an acid anhydride and partly as the anhydride of a base, because it probably forms the chloride MnGl. For if MnO, is treated with cold concentrated hydrochloric acid it dissolves with a brownish- green color, forming manganese tetrachloride, Soluble in ether with a green color. If, therefore, the solution is shaken with ether, the upper layer is colored green. Not only manganous manganites are known, but quite a num- ber of other manganites. Some of these play a very important part in analytical chemistry; as, for example, zinc and calcium bimanganites: OH OH H Miº Mé cº NO Zn, NQSc which are analogous NON /Q /Q/ * to the bicarbonates /Q2Ca Mn-O Mn+=O C=O NOH NOH NOH Zinc bimanganite is formed in the volumetric determination of manganese (see Vol. 2). Calcium bimanganite is of importance technically. Thus the recovery of manganese, by the Weldon process, in the manufacture of chlorine, depends upon the forma- ton of calcium bimanganite. Manganous oxide, MnO, is the only oxide of manganese which in all cases acts as the anhydride of a base. By dissolving this oxide in acids, manganous salts are always obtained, in which the manganese is bivalent. The oxide MnO, has never been isolated, but there are salts (R,MnO, see page 122) known which are derived from it. Mn,0, is a distinct acid anhydride, from which the per- manganates (RMnO4) are derived. In the study of the reactions of manganese we will consider first the manganous compounds, then the manganates and per- manganates. MANGANESE. I 17 A. Manganous Compounds. The manganous compounds are colored pink both in the crystal- line state and in aqueous solution; but in the anhydrous state they are colorless with the exception of the sulphide. REACTIONS IN THE WET WAY. 1. Potassium or Sodium Hydroxides precipitate white man- ganous hydroxide, MnCl,+2KOH=2KC1+Mn(OH), which rapidly becomes brown in the air, owing to the formation of manganous manganites: First of all, a part of the manganous hydroxide is oxidized by the air to manganous acid: OH Mn §4. O= Maº 2 NOH which, on coming in contact with the basic manganous hydroxide, immediately forms salts with it—the manganites, OH O Mº". #3XMn-2H,0+ MéðM. NOH NO Or /OH Mn+O 295, Ho NO ownegºtiºn-ºnor Xin. NoH This oxidation takes place in the air only gradually, but imme- diately in the presence of chlorine, bromine, hypochlorites, hydro- gen peroxide, etc.: OH OH / Mn3+2NaOH+C,-2NaCl-H,0+ Mºr OH Mn3+NaOCl-NaCl4 Mº". NOH OH Mn49: / of 4-H.O.-H.O+ Mºr I 18 REACTIONS OF THE METALS. The formation of manganites is of technical importance, as mentioned on page 116. The residue obtained in the preparation of chlorine from pyrolusite and hydrochloric acid consists chiefly of manganese chloride, and is precipitated by the addition of lime, manganous hydroxide being formed. This mixture of manganous hydroxide and lime is exposed to the action of the atmosphere, whereby manganous acid is formed, which unites with the calcium as the stronger base, forming calcium bimanganite, so that finally all of the manganese is oxidized to manganous acid: OH Mn-O /95 . Ho NON Mn-O NOH On treating the residue, when in the right condition, with hydro- chloric acid again, the same amount of chlorine is obtained as from the original pyrolusite: 2MnO,--8HCl=4H2O+2MnCl,--2Cl, Mn,00aFI,--10HCl=6H,0+2MnCl,+CaCl,--2Cl2. It is, however, necessary to add a little more hydrochloric acid in the latter case, because a part of the acid is used up in Setting the manganous acid free from the manganite. 2. Ammonia precipitates (as with magnesium and ferrous salts) from neutral solutions free from ammonium salts a part of the manganese as the white hydroxide MnOlz-H2NH3+2H2O z=2 Mn(OH)2+2NH4Cl. If sufficient ammonium chloride is present, ammonia causes no precipitation. The greater part of the manganese then remains in solution as manganous chloride, but a small amount exists as the hydroxide. On standing in the air, this dissolved hydroxide is changed into the more difficultly soluble manganous acid which is deposited in brown flocks. The condition of equilibrium in the solution is thereby disturbed, and in Order to restore it more hydrox- ide is formed, and the reaction continues in this way until finally all of the manganese is precipitated. This fact must be considered in the separation of manganese from ferric iron, aluminium, etc. If a solution of ferric and manganous chlorides contains sufficient MANGANESE. II 9 ammonium chloride, none of the manganese and all of the iron will be precipitated on the addition of ammonia, but if the solution stands in contact with the air little by little the manganese will be precipitated. In effecting the separation, therefore, an excess of ammonium chloride should be present, the solution boiled to remove the air as much as possible from the solution, then a slight excess of ammonia should be added and the solution filtered immediately. The separation even then is not quantitative, but is satisfactory for qualitative analysis. 3. Alkali Carbonates precipitate white manganous carbonate, MnCl,+Na,CO,-2NaCl-FMnCO, which after long boiling is changed by the oxygen of the air into hydrated manganese dioxide: /QH MnCO,-i-H,0+0=CO,--Mn3–O . NOH 4. Ammonium Carbonate precipitates even in the presence of ammonium salts the white carbonate (difference from magnesium). 5. Barium Carbonate produces a precipitate only in hot solu- tions. 6. Sodium Phosphate precipitates white tertiary manganous phosphate, 4Na,HPO,--3MnCl,-6NaCl-H2NaH,PO,--Mn,(PO.), soluble in mineral acids and in acetic acid. If to the boiling Solution of this precipitate in acid an excess of ammonia is added, manganous ammonium phosphate will be precipitated, as with magnesium (see page 51): *g, Na,FIPO,--NH2+MnCl,+7H,O=2|NaCl--Mn(NH,)PO,--7H,O. —, The precipitate consists of pink scales, and is very insoluble in Water. * 7. Lead Peroxide and Concentrated Nitric Acid (Volhard’s reaction).*—If a solution containing only traces of manganese is boiled with lead peroxide and concentrated nitric acid, diluted with water, and the residue allowed to settle, the supernatant * Sodium bismuthate (formed by fusing bismuth oxide with sodium per oxide) and dilute nitric acid effect the same oxidation. I 2 O REACTIONS OF THE METALS. liquid acquires a distinct violet-red color, owing to the formation of permanganic acid: § 2MnSO,--5PbO,--6HNO,- =2PbSO,--3Pb(NO,),4-2H,0+2HMnO,. This extremely delicate reaction does not take place in the presence of much hydrochloric acid or chlorides, because the per- manganic acid is thereby destrºyed: 2HMnO4+14HCl=8H2O+MnCl2 +5Cl2: 8. Ammonium Sulphide precipitates from manganese solutions flesh-colored hydrated manganese sulphide: MnCl,+(NH3),S=2NH,Cl-i-MnS+aq. S-Y—’ On boiling with a large excess of ammonium sulphide it is changed into less hydrated green manganese sulphide of the for- mula, 3MnS+H,O. - 9. Potassium Cyanide.—On adding potassium cyanide to a solution of a manganous salt a brownish precipitate appears, which redissolves in an excess of potassium cyanide, forming a brownish- colored solution. From this solution there separates, on standing or warming, a voluminous green precipitate of [Mn(CN)2]K, which dissolves in still more potassium cyanide forming a yellow Solution of [Mn(CN)].K. This salt is very unstable; it can exist only in the presence of considerable potassium cyanide. Consequently, on diluting the yellow solution with water, the green Salt again Separates Out: [Mn(CN).]K, a [Mn(CN), K+3KCN. On heating the dilute solution to boiling, the manganous potas- sium cyanide is completely decomposed into potassium cyanide, prussic acid, and insoluble white manganous hydroxide, which Separates out: [Mn(CN).]K,4-2HOHz 4KCN+2HCN+Mn(OH)2. The manganous potassium cyanide is not precipitated by am- monium sulphide in the presence of considerable potassium cyanide; MANGANESE I 2 I but the diluted solution, on the contrary, readily deposits the man- ganese as Sulphide on boiling: [Mn(CN).IK,4-(NH,),S=4KCN+2NH,CN+MnS. This reaction offers a convenient means for separating manga- nese from nickel, because nickel potassium cyanide is not precipi- tated by ammonium sulphide on boiling the dilute solution. REACTIONS IN THE DRY WAY. The bead of borax, or salt of phosphorus, is amethyst red in the oxidizing flame with Small amounts of manganese, almost brown with larger amounts, and can then be mistaken for the nickel bead. Heated in the reducing flame, the manganese bead becomes color- less, while the nickel bead appears gray. On fusing any manganous compound with caustic alkali or alkali carbonate (on platinum foil) in the air, or, better still, in the presence of an oxidizing agent (such as potassium nitrate, potas- sium chlorate, etc.), a green melt is obtained, owing to the formation of the alkali Salt of manganic acid, as is shown by the following equations: MnO-i-Na,CO,--O,-CO,--Na,MnO, MnO,--Na,CO,--O=CO,-i-Na,MnO, Mn,O,--2Na,CO,--3O=200,--2Na,MnO, Mn,O,--3Na,CO,--5O=3CO,--3Na,MnO, MnSO,--2Na,CO,--O,-200,--Na,SO,-i-Na,MnO,. The oxygen comes either from the air or from the nitrate or chlorate: KNO,-KNO,--O, RCIOs = KCl-l-3O. This reaction is exceedingly delicate, for a fraction of a milli- gram can be recognized by the formation of this green color. By ignition in the air the oxides of manganese are changed to Mn2O4: 3MnO--O=Mn,0, 3MnO,-Mn,0,--O, 3Mn,Os-2Mn,O,--O. II 2.2 REACTIONS OF THE METALS. * B. Manganic and Permanganic Acids. The free manganic acid has never been isolated. If we attempt to form it from the green melt of the alkali manganate by the addi- tion of acid, permanganic acid and hydrated manganese dioxide will be obtained; a part of the unstable manganic acid oxidizes another part of the same to permanganic acid, while the oxidizing part is itself reduced to hydrated manganese dioxide: Q= Mn-Qāī O=Mr.--OH Z | Mn =2HMnO,--H.O--Mn-O . O=*—OH This transformation takes place so readily that the green solution of the manganate is changed to a reddish-violet solution of a per- manganate by simply standing in the air with the help of the car- bonic acid which the air always contains; /OK MnO, NOK /OK OH Máo.4%+H,0–2K GO,4 Mº"+2KMo. NOK “” NOH /OK MnO, NOK The reaction takes place much more rapidly, however, if a few drops of a strong acid are added.* * The above-mentioned reaction, in which an oxygen compound is de- composed into a compound richer in oxygen, and at the same time into one poorer in oxygen, or, in other words, a part of the compound is oxidized at the expense of the oxygen of another part, is so common that we will give a few other typical examples at this place. The hypochlorites are changed by warming the aqueous solution into chlorate and chloride: ! • * * * * * * * * * * * * * * * * * * * * * * * * : * * * * NajQ;C1 = NaClOs + 2NaCl. NaOCl * MANGANESE. “ I 23 Permanganic Acid, HMnO, although much more stable than manganic acid, is only known in aqueous solution; but the anhy- dride Mn,0, on the other hand, has been isolated. On cautiously adding concentrated sulphuric acid to the cooled solution of a per- manganate, oily drops of reddish-brown Mn2O, separate out, which, however, on being warmed (the heat of reaction is sufficient) ex- plode with scintillation: Mn,0,--2H,SO, =2MnSO,4-2H,0+50. The salts of permanganic acid (the permanganates) are all solu- ble in water, with a reddish-violet color, and are very energetic oxidizing agents. According to whether the oxidation takes place in acid or in alkaline solution the permanganic acid is reduced either to MnO or MnO2. * The oxidation in acid solution takes place according to the following scheme: 2KMnO, = K,0+2MnO4-5O. By igniting the chlorate a perchlorate and a chloride are formed: * * * * ~ * * * * * * * * * * * * = ** = - f = * * * * * * * * * * * * * * * * ſº is ºn tº s s as s is sº w ś HöNio – HNO, + 2NO + H.O. Hypophosphorous acid and phosphorous acid are changed on warming into phosphoric acid and phosphine: O., H. PO, H º = H, PO, + PH, and P o, H. = 3H,PO, + PHs. * PO, H, Similarly the alkali thiosulphates and alkali sulphites are changed on igni- tion in the absence of air into Sulphate and Sulphide: Na2SSOs Na2SOs Na,SOs tº- 3Na2SO, + Na2S, and Na,S Os Na,SSOs Na2SOs = 3Na2SO, -i- Na2S I24 REACTIONS OF THE METALS. It is only necessary to provide sufficient acid to dissolve the oxides formed. The oxidation in alkaline solution reduces the permanganic acid only to manganese dioxide: 2KMnO, = K,0+2MnO,--3O. Eacamples of Oaxidation in Acid Solution.—The oxidation of ferrous salts by potassium permanganate has been explained (see page 5); but we will mention here a few other important cases. Almost all hydrogen compounds of the negative elements are oxidized by a Solution of potassium permanganate, hydrochloric and hydrobromic acids on warming, hydriodic acid even in the cold: 2KMnO,--3H,SO,--10HCl=K,SO,--2MnSO,--8H,O--5Cl, 2KMnO,--3H,SO,-- 10HI = K,SO,--2MnSO,--8H,O-H-5T, Hydrogen sulphide is oxidized in the cold with separation of sulphur: 2KMnO,--3H,SO,--5H,S=K,SO,--2MnSO,--8H,0+5S. The hydrogen compounds of phosphorus, arsenic, and antimony are oxidized to the corresponding acids: 8KMnO,--12H,SO,--5PH,-4K,SO,--8MnSO,--12H,0+5H,PO,. Sulphurous acid decolorizes the solution of potassium perman- ganate, Sulphur and dithionic acid being always formed. It is never possible by means of this reaction to oxidize sulphurous acid completely to Sulphuric acid. The proportion of sulphuric acid formed to that of dithionic acid varies with the concentration and with the temperature; consequently Sulphurous acid cannot be de- termined quantitatively by means of potassium permanganate. If sulphurous acid is allowed to act upon manganese dioxide suspended in water, manganese dithionate and manganese sul- phate are formed: MnO,--2SO,-MnS,Os, MnO,-- SO,-MnSO,. MANGANESE. 125 In the cold more of the dithionate is formed, while on warming more sulphate. As a point must be reached in the oxidation of sulphurous acid by potassium permanganate where MnO, begins to act upon the former, there must, therefore, always be some dithionic acid formed. At a certain temperature and with a definite concentration, the reaction can take place as follows: K.O2MnO,O,4+6SO,--2H,0=2KHSO,--2MnSO,--H.S.O. Oxalic Acid is completely oxidized on warming to carbonic anhydride: COOH 2KMnO,--5 | +3H,SO, = K,SO,--2MnSO,--8H,O-H-10CO,. COOH Tartaric acid is also oxidized by permanganates. In the case of all the above reactions, a considerable excess of acid must be present, otherwise the solution will become turbid, Owing to the formation of brown manganous manganite: OH Mº 4KMnO,--11MnSO,-- 14.H.O=4KHSO,--7H,SO,--5 Yº > Mn. Mn-O NOH According to the concentration and temperature at which the reaction takes place other manganites may be formed. Hydrogen Perozide, or the peroxides of the alkalies and alkaline earths, as well as percarbonic acid, all decolorize permanganic acid; the permanganic acid as well as the other bodies are reduced, with evolution of oxygen: 1. 2KMnO,--5H,0.44H.SO,-2KHSO,4-2MnSO,4-8H,0+50. 2. 2KMnO,--5K.C.O.-- 14H,SO, = = 12.KHSO,--2MnSO,--8H,O-H-10CO,--502. Persulphuric acid, which is analogous to percarbonic acid, does not reduce a solution of a permanganate. * * 2KMnO, = K,O.2MnO, O,. I 26 REACTIONS OF THE METALS. Oxidation in Alkaline Solution.—Many organic substances are oxidized by permanganates in alkaline solution with deposition of manganese dioxide. Thus formic acid is oxidized to carbonic acid, ethyl alcohol to aldehyde and acetic acid, cellulose (paper) to oxalic acid principally, so that a Solution of a permanganate cannot be filtered through paper. By boiling a concentrated solution of potassium permanganate with concentrated potassium hydroxide, potassium manganate is formed with evolution of oxygen, so that the color of the solution becomes green: 4KMnO,44KOH-4K2MnO,42H20+O. By heating solid potassium permanganate to 240° C., potassium manganate is formed, also with evolution of oxygen: 2KMnO, = K,MnO,--MnO,--O,. NICKEL, Ni. At. Wt. 58.7. Occurrence.—In the native state nickel occurs only in meteor- ites. It is most frequently found in combination with sulphur, arsenic, and antimony in regular and hexagonal crystallizing minerals, of which the following are the most important: A. Isometric System. Chloanthite Gersdorffite Ullmannite NiAs,; NiAsS; NiSbS. B. Hexagonal System. N iccolite Breithauptite Millerite Ni. As Ni, Sb Ni.S 2 2 7 2 2 y - - -2-2. Nickel also occurs as regular crystals of bunsenite, NiO, iso- morphous with periclasite, MgO, and manganosite. MnO; as gar- nierite, or noumeite SiO,(NiMg)H,--aq a mineral occurring in New Caledonia, from which pure nickel can be prepared; and finally as annabergite, Ni,(AsO4). .8H,O, isomorphous with erythrite. Metallic nickel possesses a silver-white color, with a specific gravity of 8.9 and a melting-point of about 1500° C. It is diffi. cultly soluble in hydrochloric and Sulphuric acids, but readily soluble in nitric acid. It forms two oxides: t Nickeious oxide, NiO, Nickelic oxide, Ni,0,. Green Brownish black NICKEL. 127 By dissolving either of these oxides in acids, salts of bivalent nickel are always obtained: NiO+2HCl=H,0+NiCl, Ni,O,--6HCl=3H,0+2NiCl, HCL, 2Ni2O3 +4H2SO4=4|H2O+4NiSO4+O2. Nickelous oxide behaves as a basic anhydride, while nickelic oxide acts as a peroxide. The crystallized salts of nickel and their aqueous solutions are green, but in the anhydrous condition they are usually yellow. Most of the salts are soluble in water; the sulphide, carbonate, and phosphate are insoluble. REACTIONS IN THE WET WAY. 1. Potassium Hydroxide precipitates apple-green nickelous hydroxide. OH NiC, H2KOH-2KCI+Niği, insoluble in excess of the precipitant, readily soluble in acids. 2. Ammonia precipitates (in neutral Solutions free from ammo- nium salts) a green basic salt, Ni–OH 2NiSO,--2NH,OH=(NH),SO,-- Xo, Ni–OH soluble (with a blue color) in excess of ammonia, forming a complex nickel ammonium salt: Ni,SO,(OH),4-(NH,),SO,--10NH,-2|Ni(NH,).]SO,”--2H,0. In the presence of sufficient ammonium salt ammonia produces no precipitate as with magnesium, ferrous and manganous salts; potassium and sodium hydroxides, however, precipitate the green hydroxide (difference from cobalt, see p. 133). * Just as cyanogen possesses the ability to form with metals complex ions, so do ammonia, water, pyridine, etc. Cf. A. Werner, Zeit. für anorg. Ch., III ff.; A. Werner and A. Miolatti, Zeit. für phys. Ch., XII, XIV, and XXI. I 28 REACTIONS OF THE METALS. The anhydrous chloride and sulphate readily absorb ammonia, forming anhydrous nickel ammonium salts: NiCl2 +6NH3=[Ni(NH3)6]Cl2, NiSO4+6NH3=[Ni(NH3)6]SO4. 3. Potassium and Sodium Carbonates precipitate apple-green nickel carbonate: NiCl, -Na,CO,-2NaCl- NiCO,. 4. Ammonium Carbonate behaves similarly, only the precipi- tate which is formed is soluble in an excess of the precipitant, form- ing nickel ammonia carbonate. 5. Sodium Hypochlorite precipitates in the presence of alkalies all of the nickel as brownish-black nickelic hydroxide, Ni(OH)2. Nickelous hydroxide is first formed by the alkali present, but it is then oxidized by the hypochlorite to nickelic hydroxide: 2Ni(OH),4-NaOCl4-HOH=NaCl-H2Ni(OH),. On adding chlorine or bromine to the nickel Solution to which alkali has been added, nickelic hydroxide is likewise formed:* Ni(OH),4-NaOH--Cl– NaCl- Ni(OH), 6. Barium Carbonate produces in the cold no precipitation; but by continued boiling, all of the nickel is thrown down as basic car- bonate. 7. Hydrogen Sulphide precipitates no nickel from solutions which contain mineral acid or much acetic acid; while from Solu- tions slightly acid with acetic acid and containing an alkali acetate, all the nickel is precipitated as the black Sulphide: NiCl, H2NaC.H.O.--H.S=2NaCl-H2HC.H.O.--NiS. * By carefully evaporating the brown solution of nickel sulphide in am- monium sulphide, Antony and Magri claim to have isolated a sulphide of the formula NiS,. (Chem. Zeit., 1902, p. 37.) *S*, NICKEL. I 2 9 8. Ammonium Sulphide precipitates from neutral solutions the nickel as sulphide: NiCl, H. (NH,),S=2NH,Cl-i-NiS. Nickel sulphide is somewhat soluble (with a brown color) in an excess of ammonium sulphide; particularly so in the presence of free ammonia. The nickel may be precipitated from this brown solu- tion by acidifying with acetic acid and boiling. In the presence of considerable ammonium salts the nickel is not dissolved by colorless ammonium sulphide. If, therefore, we desire to precipitate the nickel as Sulphide from an ammoniacal solution, it is best to add considerable ammonium chloride and then Saturate the solution with hydrogen sulphide. The nickel sulphide thus obtained can be easily filtered and the filtrate is completely free from nickel. The reason why nickel Sulphide is Soluble in ammonium sulphide in the presence of ammonia is, probably, that a complex nickel ammo- nium sulphide is formed, perhaps of the formula Ni(NH.).S. Nickel sulphide is very difficultly soluble in dilute mineral acids, readily soluble, however, in strong nitric acid or aqua regia, with separation of sulphur: 3NiS+6HCl·H2HNO,-3NiCl,--2NO+4H,0+3S. The Sulphur usually separates out as a black film. This is caused by the Sulphur first melting in consequence of the heat of reaction, enclosing small particles of the black sulphide and pro- tecting them from the action of the acid. By continued action of the acid all the Sulphide is dissolved, and the sulphur remains as yellow drops, which are oxidized little by little to sulphuric acid, going into Solution: S+2HNO,-H.SO,4-2NO. 9. Potassium Cyanide produces a bright-green precipitate of nickelous cyanide, NiCl,--2KCN=2KCl-i-Ni(CN), readily soluble in an excess of the precipitant, forming nickel potas- sium cyanide: Ni(CN),4-2KCN=[Ni(CN).]K,. I3o REACTIONS OF THE METALS. This last salt is readily decomposed by dilute mineral acids, with evolution of prussic acid and deposition of nickelous cyanide, which dissolves finally by the addition of more acid: [Ni(CN),JK,--2HCl=[Ni(CN), H,--2KCl. The nickel hydrogen cyanide, which is first formed, is, like car- bonic acid, very unstable, and decomposes according to the follow- ing equation: [Ni(CN), H,-Ni(CN),4-2HCN. Nickel potassium cyanide is not decomposed by ammonium sulphide (difference from manganese and zinc), but is readily de- composed by chlorine, bromine, and hypochlorites: [Ni(CN)]K,48C-2KCI+NiCl, 4CNC. On adding sodium hydroxide to a solution of nickel potassium cyanide and conducting chlorine into the solution, nickelous chloride is first formed, which is decomposed by the caustic soda to nickel hydroxide; and the latter is oxidized by further action of chlorine or bromine, forming black voluminous nickelic hydroxide (cf. p. 128). This reaction is exceedingly delicate and serves for the detection of nickel in the presence of cobalt, as the corresponding complex cobalt cyanide compound is not decomposed under these conditions. A large excess of potassium cyanide must be avoided, because in that case the reaction takes place much more slowly. The re- action depends entirely upon the presence in the Solution of nickel chloride, and this is formed only when the excess of the potassium cyanide has been destroyed. By the further action of chlorine the nickel potassium cyanide is decomposed; i.e., nickel chloride is formed, upon which sodium hydroxide and chlorine act, forming black nickelic hydroxide. The best procedure is as follows: A few drops of the solution to be tested for nickel are taken So that two to three drops of potas- sium cyanide will be sufficient to produce a clear solution, two to three c.c. of double normal sodium hydroxide Solution are added, and chlorine is conducted into the Solution in the cold. Under these conditions a precipitate of Ni(OH), will be surely formed within one to two minutes. NICKEL. I 3 I Bromine acts similarly to chlorine, but it is decidedly preferable to conduct chlorine gas into the solution. 10. Sodium Phosphate precipitates apple-green nickel phos- phate, 3NiCl, H4Na,HPO,-6NaCl-H2NaH,PO,--Ni,(PO.), readily soluble in acids, even acetic. 11. Potassium Nitrite produces in dilute nickel solutions no precipitate (difference from cobalt). In very concentrated solu- tions a brownish-red precipitate of Ni(NO), .4KNO, is thrown down; in the presence of alkaline earth Salts a yellow crystalline precipitate is formed; e.g., Ni(NO.), Ba(NO.),.2KNO, which is very difficultly soluble in cold water, but readily soluble in boiling water, with a green color. REACTIONS IN THE DRY WAY. The borax, or sodium metaphosphate, bead is brown in the oxi- dizing flame, almost the same shade as the strongly-Saturated man- ganese bead; in the reducing flame the bead becomes gray, in con- sequence of the formation of some metallic nickel. On looking at the bead through the microscope the finely-divided metal can be seen suspended in the colorless glass. On heating nickel Salts with sodium carbonate on charcoal a gray scale of metallic nickel is obtained. This reaction is best per- formed with the charcoal stick, as described on page 24. The mag- netic metal obtained in this way is placed on a piece of filter-paper, dissolved in nitric acid, a drop of concentrated hydrochloric acid is added, and the paper carefully dried by moving it back and forth over the flame. If nickel is present the paper appears greenish (colorless with very small amounts of nickel), or bluish if cobalt is also present. The paper is now moistened (where the nickel is) with caustic soda or potash, and is then held in bromine vapors, which are obtained by shaking Some bromine water in a wide- mouthed flask. If nickel or cobalt is present, a black spot will be formed by the above treatment, consisting of the hydroxide of the trivalent metal I32 REACTIONS OF THE METALS. (p. 128). The blackening often does not appear at first; in this case the paper is moistened once more with potassium hydroxide and again treated with bromine. The spot will now appear if nickel is present. COBALT, Co. At. Wt. 59.0. Occurrence.—Like nickel, native cobalt is found only in meteor- ites. It occurs in the earth's crust chiefly as sulphide, arsenide, and as salts of Sulphoarsenious and Sulphoantimonous acids; it is almost always accompanied by nickel and iron. The most important ores are smaltite, (Co,Ni,Fe)As, iso- metric; cobaltite, (Co,Fe)(As-S), isometric; skutterudite, CoAs, isometric; erythrite, Cos(ASO), .8H,O, monoclinic, isomor- phous with vivianite, Fea(PO.),.8H2O, and with annabergite, Ni,(AsO4)2.8H,O. Metallic cobalt is steel gray, possesses the specific gravity of 8.5, dissolves much more readily in dilute acids than nickel, and is, like the latter, magnetic. Cobalt forms, like iron, three oxides: Cobaltous oxide Cobaltous cobaltic oxide Cobaltic oxide CoO, Co,04, Co,0s. By dissolving these three oxides in acids, salts derived from cobaltous oxide are always obtained, containing therefore bivalent eobalt: CoO+2HCl= H,0+ Co.Cl, Co,O,--6HCl=3H,0+2COCl, HCl, Co,0,--8HCl=4H,0+3COCl, -Cl, Simple cobaltic salts are unknown, but many complex com- pounds exist with trivalent cobalt, as, for example, potassium cobaltic nitrite, potassium cobaltic cyanide, and numerous cobaltic ammonium derivatives. Cobaltous compounds in a crystallized state (as well as in aqueous solution) are pink, in the anhydrous condition yellow or green, and blue in aqueous solutions in the presence of hydrochloric acid. The solubility reactions of cobaltous salts are similar to those of manga- nese and nickel. g COBAL.T. 133 REACTIONS IN THE WET WAY. 1. Potassium or Sodium Hydroxide precipitate in the cold a blue basic salt: —Cl co-º-KOH-KCH-Co(3* Cl which on warming is further decomposed by potassium hydroxide, forming pink cobaltous hydroxide: —OH Co of OH + KOH = KCI-H co&#. In the case of a moderately concentrated Solution of the alkali the precipitate of pink cobaltous hydroxide is often produced in the cold, sometimes only after standing some time. The rapidity of the reaction depends entirely upon the concentration of the alkali. Cobaltous hydroxide changes gradually to brown in contact with the air, going over into cobaltic hydroxide: 2Co-º;4-HOH+0– 2c28; —OH NOH In this respect cobalt behaves similar to iron and manganese, and differs from nickel, for the hydroxide of the latter is not oxidized by atmospheric oxygen. On adding chlorine, bromine, hypochlorites, hydrogen peroxide, etc., to an alkaline solution containing cobaltous hydroxide, cobaltic hydroxide is immediately formed, as with nickel and manganese: OH —OH / + NaOH-i-Cl– NaCl- Co-OH, —OH NOH OH /Qā 2003:# HOH-H NaOCl = NaCl- 2004 OH. NOH From ammoniacal cobalt Solutions the above oxidizing agents cause no precipitation, but merely a red coloration; the addition of potassium hydroxide then causes no precipitation (difference from nickel). Remark.—Cobaltous hydroxide, Co(OH)2, behaves under some conditions as a weak acid, for on adding to a cobaltous solution a very concentrated solution of KOH or NaOH the precipitate at Co I34 REACTIONS OF THE METAL first produced dissolves with a blue color * similar to copper. Ty the addition of Rochelle salts to this blue cobalt solution the color either disappears almost entirely or becomes a pale pink, while the similarly treated copper Solution becomes more intensely blue. By the addition of potassium cyanide to the blue cobalt solution it be- comes yellow, and in contact with air turns intensely brown. A copper solution would be decolorized by the addition of potassium cyanide. By pouring a little cobalt Solution (or adding a little solid cobalt carbonate) into a concentrated Solution of caustic soda or potash, to which a little glycerine has been added, a blue solution is formed (the color being intensified by warming), which after standing some time in the air becomes a beautiful green on the addition of hydro- gen peroxide. * 2. Ammonia precipitates, in the absence of ammonium salts, a blue basic salt, soluble, however, in excess of ammonium chloride. Ammonia, therefore, produces no precipitate in solutions which contain sufficient ammonium chloride. The dirty yellow ammonia- cal Solution is little by little turned reddish on exposure to the air, Owing to the formation of very stable cobaltic ammonium derivatives 3. Alkali Carbonates produce a reddish precipitate of basic salt of varying composition. 4. Ammonium Carbonate also precipitates a reddish basic salt, soluble, however, in excess. 5. Barium Carbonate precipitates in the cold, and in the absence of air, no cobalt; but on exposure to the air cobaltic hydroxide is gradually thrown down. The precipitation takes place much more quickly on the addition of hypochlorites or hydrogen peroxide: 2000l,+2BaCO,4-3HOH+NaOCl= =NaCl-H2BaCl, -2CO,--2Co(OH),. If the solution is heated to boiling, all of the cobalt is precipi- tated as a basic salt, even out of contact with the air. - 6. Hydrogen Sulphide produces no precipitate in solutions con taining mineral acids. In neutral Solutions containing an alkali acetate all of the cobalt is precipitated as black sulphide. 7. Ammonium Sulphide precipitates black cobalt sulphide. CoCl,--(NH,),S=2NH,Cl-FCoS, * Ed. Donath, Zeitschr. f. anal. Ch., 40, p. 137 (1901). COBAL.T. I35 insoluble in ammonium sulphide, acetic acid, and very dilute hydro- chloric acid; soluble in concentrated nitric acid and aqua regia, with separation of sulphur: 3CoS--8HNO,-4H,0+2NO+3S4-3Co(NO),. By continued action of strong nitric acid all the sulphur goes into solution. 8. Potassium Cyanide produces in neutral solutions a reddish- brown precipitate, soluble in excess of potassium cyanide in the cold, with a brown color, forming potassium cobaltocyanide: 1. CoC,4-2KCN=Co(CN), 2. Co(CN),4-4KCN=[Co(CN), K. On warming the brown neutral solution for some time it be- comes a bright yellow and reacts alkaline; it now contains potas- sium cobalticyanide, of analogous composition to potassium ferri- cyanide. The formation of the cobaltic salt takes place with the help of atmospheric oxygen: z: Co(CN), i. =* = sma ºn s º ºs ºs ºº sº gº ºs ººººººººº. The reaction takes place more quickly by means of chlorine, bromine, hypochlorites, etc.: /º ºn wº Co(CN), Li-i-Cl-KCl4-[Co(CN)]–K. Nk NK An excess of chlorine, bromine, etc., does not decompose the cobaltic salt (difference from nickel). Potassium cobalticyanide is very much more stable than the cobaltous compound. By adding hydrochloric acid to the brown I 36 REACTIONS OF THE METALS. solution of potassium cobaltocyanide, prussic acid will be given off and yellow cobaltous cyanide will be formed, [Co(CN), K,4-4HCl=4KCl-H4HCN+Co(CN), while potassium cobalticyanide is not decomposed by hydrochloric acid. Potassium cobalticyanide forms, with most of the heavy metals, difficultly soluble or insoluble salts possessing characteristic colors. Thus, it produces with cobaltous salts pinkish-red cobaltous cobalti- cyanide: [Co(CN).]K, [Co(CN)]^* 3COCl. = Co-H6KGl, [Co(CN ).]Ks [Co (CN )ek Co and with nickel salts greenish nickel cobalticyanide. If, therefore, a cobalt Solution contains nickel it gives, on pre- cipitating and redissolving with potassium cyanide, boiling, and adding hydrochloric acid, a greenish precipitate of nickelous cobalti- cyanide: 2[Co(CN)]K,4-3|Ni(CN), K,4-12HCl=12KCl-H +12HCN+[Co(CN),Ni,. 9. Potassium Nitrite produces in concentrated solutions of cobalt salts, with the addition of acetic acid, an immediate precipi- tation of yellow crystalline potassium cobaltic nitrite. "If the solu- tion is dilute, the precipitate appears only after standing for some time, but more quickly on rubbing the sides of the beaker. The reaction takes place in the following stages: 1. CoCl,--2KNO, a Co(NO.),4-2KCl. 2. 2KNO,--2HC.H.O. =2KC.H.O,--2HNO,. The free nitrous acid oxidizes the cobaltous nitrite to cobaltie nitrite, †† . NO —NQ, L H NO, ^Nº. 2 which now combines with more potassium nitrite: 4. Co(NO.),4-3KNO,-Co(NO.).K. \? COBAL.T. I37 This reaction offers an excellent means for detecting the presence of cobalt in nickel salts. 10. Ammonium Sulphocyanate (Vogel's reaction).* If a con- centrated solution of ammonium sulphocyanate is added to a cobal- tous solution, the latter becomes a beautiful blue, owing to the formation of ammonium cobaltous Sulphocyanate: (a) CoCl4-2NH,CNS=2NH,Cl-HCo(CNS), (b) Co(CNS),4-2NH,CNS= [Co(ºN S).](NH,),. On adding water the blue color disappears and the pink color of the cobaltous salt takes its place. If, now, amyl alcohol f is added (or a mixture of equal parts amyl alcohol and ether), and the solu- tion shaken, the upper alcoholic layer is colored blue. This reaction is so sensitive that the blue color is recognizable when the solution contains only +}r of a milligram of cobalt. The blue Solution also shows a characteristic absorption spectrum.j Nickel salts produce no coloration of the amyl alcohol. If, however, iron is present, the red Fe(CNS, is formed, which likewise colors the amyl alcohol, mak- ing the blue color (due to the cobalt) very indistinct, so that, under some conditions, it can no longer be detected. If, however, some sodium carbonate solution is added drop by drop, the iron will be precipitated as ferric hydroxide, while the blue color produced by the cobalt is unaffected. Detection of Traces of Cobalt in Nickel Salts. In order to test a nickel salt for cobalt, a concentrated solution of ammonium sulphocyanate is added to the solution of a consider- able amount of the salt, a few cubic centimeters of a mixture of amyl alcohol and ether are added and the Solution shaken. After the latter has been allowed to settle, if the upper alcohol-ether layer is colorless, then the nickel Salt contains neither iron nor cobalt; if the layer is reddish, iron is present. In this case a few drops of * Ber. deutsch. Chem. Ges. 12, 2314. † T. T. Morrell first showed that cobalt salts give a blue color with am- monium sulphocyanate, disappearing on the addition of water, but reap- pearing when alcohol is added. Zeit. anal. Chem. 16, 251. † Wolff, Zeitschr. anal. Chem. 18, 38. 138 REACTIONS OF THE METALS. Soda solution are added, and the solution again shaken; if cobalt is present the alcohol-ether layer is now distinctly blue. REACTIONS IN THE DRY WAY. The bead produced by borax or sodium metaphosphate is blue in both the oxidizing and reducing flames. By holding the bead in the upper reducing flame for a long time it is possible to reduce the cobalt to metal, when it appears, like nickel, gray. On the charcoal stick cobalt compounds yield gray metallic cobalt, which can be removed by means of a magnetized knife-blade, as described on page 25, placed on filter-paper, dissolved in hydro- chloric acid and dried. The paper is then colored blue by cobalt (difference from nickel). If, now, sodium hydroxide is added and the paper exposed to the action of bromine vapors, black cobaltic hydroxide, Co(OH), is formed. ZINC, Zn. At. Wt., 65.4. Occurrence.—Smithsonite, ZnCO, isomorphous with calcite, CaCO, etc.; sphalerite, ZnS, isometric; calamine, Zn,SiO,--H.O, orthorhombic, hemimorphic; zincite, ZnO, hexagonal; and frank- linite, [FeC.,(Fe,Mn,7n), isometric. Metallic zinc is bluish white, with sp. gr. 6.9; it melts at 420° C. and boils at about 950° C. At low temperatures and at about 200° C. it is so brittle that it can be pulverized, but at 110°– 150° C. it is ductile and can be drawn out into wire and rolled into foil. Zinc is readily soluble in all acids; in hydrochloric, sulphuric, and acetic acids with evolution of hydrogen: Zn-i-2HCl=ZnCl,+H,. Nitric acid dissolves it forming the nitrate, but hydrogen is not evolved on account of its being used up in reducing the excess of nitric acid. The reduction products are different according to the concentration of the acid used; with concentrated acid, nitric oxide, NO, is formed, while dilute acid is reduced to ammonia: 2HNO,--6H=4H,0+2NO (concentrated), HNO,--8H=3H,0+NH, (dilute). ZINC. I39 The solution of zinc in concentrated nitric acid may be supposed to take place as follows: First of all, the nitrate is formed, with evolution of hydrogen, 3Zn-i-6HNO,-3Zn(NO),4-3H, and the latter is used up as quickly as it is formed by reducing the nitric acid to nitric oxide: 2HNO,4-6H =4H,0+2NO. The whole reaction, then, may be represented by the equation 3Zn-L8HNO,-3Zn(NO.),4-4H.O--2NO. A similar reaction takes place in the solution of zinc in dilute nitric acid: 4Zn-H 10HNO,-4Zn(NO),4-NH.NO,--3H,O. Like aluminium, zinc dissolves in caustic soda or potash with evolution of hydrogen and the formation of a zincate: Zn-L2KOH=Zn(OK),4-H. Zinc forms only one oxide, ZnO. It is a white infusible powder, which becomes yellow when heated, but turns white again on cooling. Zinc oxide dissolves readily in acids, forming zinc salts: ZnO-i-H...SO,-H,0+ZnSO,. There exists only one series of zinc salts, and the zinc is always bivalent. Most of the Salts are white. The chloride, nitrate, sul- phaſeº"and acetate are soluble in water; the remainder dissolve readily in mineral acids. REACTIONS IN THE WET WAY. 1. Potassium and Sodium Hydrates precipitate white gelati- nous zinc hydroxide. ZnCl,--2KOH=2KCl-i-Zn(OH), easily soluble in excess of the precipitant, forming a zincate: * Zn(OH)2+2KOH-2H,0+ZnT3:. e-ºrding to Hantzsch the zinc is not present as zincate, but probably as hydroxide in the hydrosol form. Zeit. f. anorg. Ch., XXX (1902), p. 289. I4o REACTIONS OF THE METALS. Zinc hydroxide, therefore, behaves sometimes as a base and Sometimes as an acid, like aluminium hydroxide. On boiling a diluted solution of a zincate, hydrolysis takes place and zinc hydroxide is precipitated: —OH OK Il —OH" Znoj. +2HOHz−22KOH-i-Z If, however, the Solution contains considerable sodium or potas- sium hydroxide, there will be no precipitation. 2. Ammonia precipitates from neutral solutions, free from am- monium salts, zinc hydroxide, readily Soluble in ammonium salts, as in the case of magnesium, nickel, manganese, or iron: ZnCl2 +2NH3+2H2O z=2 Zn(OH)2 +2NH4Cl. Zinc hydroxide is also soluble in an excess of ammonia, due to the formation of complex zinc ammonium hydroxide: Zn(OH),4-6NH,-[Zn(NH,), (OH),. In the presence of ammonium salts the corresponding salt is formed: Zn(OH),4-2NH,Cl-H4NH,-[Zn(NH,).]Cl, H2H,O. 3. Alkali Carbonates precipitate a white, basic carbonate, of variable composition, as is the case with magnesium (p. 51). 4. Ammonium Carbonate does the same, except that the pre- cipitate is soluble in an excess of the reagent. The presence of ammonium salts prevents the precipitation. 5. Barium Carbonate precipitates no zinc in the cold, but on boiling all the zinc is precipitated as basic carbonate. 6. Sodium Phosphate precipitates gelatinous, tertiary zinc phosphate, which soon becomes crystalline, and is soluble in am- monia and in acids: 4Na, HPO,--3ZnCl,-6NaCl-i-2NaH,PO,--Zn,(PO.), 7. Hydrogen Sulphide precipitates, from neutral solutions of a zinc salt of a mineral acid, the zinc incompletely as Sulphide: ZnCl,--H,Sze 2HCl- ZnS. ZINC. I41 Zinc sulphide is soluble in mineral acids; consequently the above reaction is reversible. The more dilute the solution, the more com- plete will be the precipitation; but it will never be quantitative. Zinc sulphide is insoluble in acetic acid. If, therefore, we add Some sodium acetate to the solution, and then pass into it hydrogen sulphide, all the zinc will be thrown down as Sulphide: ZnCl,--2NaC.H.O.--H.S=2NaCl-i-2HC.H.O,-- ZnS. 8. Ammonium Sulphide precipitates from neutral or alkaline solutions all the zinc as amorphous Sulphide: ZnCl,--(NH,),S=2NH,Cl-HZnS. Zinc sulphide is a precipitate hard to filter; it runs through the filter-paper, particularly on washing. This is a peculiarity of almost all metallic Sulphides and of many other amorphous bodies, such as aluminium hydroxide, titanic acid, tungstic acid, and many others. Such bodies exist in a soluble form as hydrosol, and in an insoluble form as hydrogel. The hydrosol can be changed into hydrogel by precipitation with concentrated Salt solutions, or by boiling, or by the addition of acids. In order, then, to obtain zinc sulphide in a form which can be filtered, it is precipitated from a boiling Solution, and in the presence of a considerable amount of Salts,” preferably ammonium salts. The precipitate is washed with a solution of ammonium chloride to which a little ammonium sulphide has been added. 9. Potassium Cyanide produces a white precipitate of zinc cyanide, soluble in an excess of the precipitant, ZnCl,--2KCN=2KCl-i-Zn(CN), and Zn(CN),4-2KCN=Zn(CN),K. Zinc potassium cyanide is readily decomposed by acids and alkali sulphides: Zn(CN),K,4-2HCl=2KCl4-2HCN-I-Zn(CN), Zn(CN),K,4-(NH,),S=2KCN+2NH,CN+ZnS. * Of course acids cannot be added in this case, as they would dissolve the zinc sulphide. I42 REACTIONS OF THE METALS. 10. Potassium Ferrocyanide precipitates white zinc ferrocy.n- ide, which is changed by an excess of the potassium ferrocyanide into insoluble zinc potassium ferrocyanide: ſº (a) K.Fe(CN)--2ZnCl,-4KCl-- [Fe(CN), #, FC [Fe(CN) #% d'e 6.RTZI). NZ. (b) 3Zn,Fe(CN),4-K Fe(CN),=2 Zn. |Fe(CN)]^2n NK The formation of the last-named body plays an important part in the volumetric determination of zinc, according to the method of Galetti. IREACTIONS IN THE DRY WAY. Heated with sodium carbonate on charcoal before the blowpipe, it is not possible to obtain metallic zinc on account of its volatility; but an incrustation of oxide is obtained which is yellow while hot and white when cold. Zinc oxide (or such compounds of zinc as are changed over to oxide on ignition), when moistened with cobalt nitrate yields a green infusible mass: Rinnmann's green. This reaction is per- formed exactly as with aluminium (p. 75). Separation of the Metals of Group III from the Alkalies and Alkaline Earths. The separation of the members of the ammonium sulphide group from the alkalies and alkaline earths is effected by means of ammo- nium sulphide and ammonium chloride. If, however, the solution contains phosphoric acid, oxalic acid, or considerable boric acid, ammonia and ammonium sulphide will each precipitate calcium, strontium, barium, and magnesium as phosphate, oxalate, or borate together with the members of the ammonium sulphide group. This special case will be treated in Part II. t The separation of the metals of Group III is effected according to Tables II, III and IV, pages 143, 144, and 145. § TABLE II. SEPARATION OF MANGANESE, NICKEL, COBALT, AND ZINC. The solution containing ammonium chloride is treated with ammonia until alkaline, and hydrogen sulphide is passed in until it becomes saturated. The precipitate is filtered off and washed three times with water to which a little ammonium sulphide has been added. The precipitate is then removed to a porcelain dish and treated with cold double-normal hydro- chloric acid, being stirred with a glass rod until no more hydrogen sulphide is given off, when it is filtered. g SoLUTION (MnCl2,ZnCl2, and traces of NiCl2). RESIDUE (CoS,NiS). The residue is tested for cobalt by heating a small por- tion with the borax bead; a blue bead shows cobalt. In order to test for nickel, a portion is dissolved in a few drops of aqua regia, evaporated just to dryness, and dissolved in a little water; potassium cyanide is added drop by drop until the precipitate produced is entirely dissolved, two to three c.c. of sodium hydroxide are added and chlorine conducted into the cold solution. If nickel i8 present, black Ni(OH)3 is formed within a few minutes. N.B.-If the borax bead was brown in the cobalt test, a further test for nickel is unnecessary. Very small amounts of cobalt may be present. In order to detect the cobalt, a considerable amount of the precipitate is dissolved in aqua regia, evaporated just to dryness, dissolved in as little water as possible, a concentrated solution of potassium nitrite is added, the solution is acidified with acetic acid and allowed to stand twelve hours. A yellow crystalline precipitate of [Co(NO2)6]K3 shows cobalt to be present (cf. p. 136). Vogel's reaction (p. 137) is still better for detecting small amounts of cobalt. For this purpose the aqueous so- lution, obtained as before, is treated with a concentrated so- lution of NH4SCN and shaken with a few c.c. of a mixture of equal parts amyl alcohol and ether. If the upper layer of alcohol-ether solution is colored blue, cobalt is present. The solution is boiled until hydrogen sulphide ceases to be given off, then precipitated with strong caustic potash and filtered. PRECIPITATE [Mn(OH)2 and a little Ni(OH)2]. SoLUTION [Zn(OK)2]. (a) A portion of the precipitate is dissolved in a little hydrochloric acid, evaporated care- fully to dryness, and dissolved in a little water; a few drops of potassium cyanide solution are added; the solution is diluted largely with water, treated with (NHA)2S and boiled. A flesh-colored precipitate of MnS shows manganese. (b) A small portion of the precipitate is dis- solved in concentrated nitric acid and a knife- bladeful of lead peroxide is added; the mixture is then boiled and diluted with water. The supernatant solution is colored red if man- gamese is present. (c) A small portion of the precipitate is fused on platinum foil with sodium carbonate and potassium carbonate. The melt is colored green if manganese is present. On dissolving in a little water, and acidify- ing with a few drops of acetic acid, a red solu- tion is obtained. The solution is acidi- fied with acetic acid, and hydrogen sulp hide is passed into the solution. 4 white precipitate of ZnS shows 2inc. S The precipitate is filtered, a small portion is transferred to a piece of filter-paper and placed in a platinum spiral. It is then treated with nitric acid, evaporated to dryness, moistened with dilute cobalt nitrate so- lution and ignited. A green infusible mass (Rinnmann's green) shows 2inc. A RESIDUE [CoS,NiS). It is then treated with cold double-normal hydrochloric acid in a porcelain dish, being stirred until no more hydrogen sulphide is given off, when it is filtered. B. SOLUTION [FeCl, UO,Cl, AlCls, CrCls, MnO, ZnCl2 (traces of NiCl,)]. To be examined by Table II. The solution is evaporated to a small volume, oxidized with two to three c.c. of concentrated nitric acid, treated with concentrated sodium hydroxide until strongly alkaline, boiled and filtered. C PRECIPITATE [Fe(OH)3, Cr(OH)4, U.O.Na, Mn(OH), Ni(OH)2]. D. Solution [Al(ONa),Zn(ONa),). The precipitate is dissolved in as little concentrated hydrochloric acid as possible, diluted with a little hot water, and boiled several minutes. Ammonium chloride is added, the boiling solution treated with a slight ex- cess of ammonia and filtered as quickly as possible. E. PRECIPITATE [Fe(OH)4, F. SOLUTION MnOl, Cr(OH)3, (NHA), U.O.] and traces of NiCl2. The solution is acidified with hydro- chloric acid, an excess of ammonia is added, and the solution filtered. The precipitate is dissolved in as little || is *: TABLE III. FIRST METHOD OF SEPARATING ALL THE MEMBERS OF THE AMMONIUM SULPHIDE GROUP. (To be used in the presence of considerable cobalt.) The neutral solution, which should contain ammonium chloride, is treated at the boiling temperature with ammonium sulphide added drop by dro urated with hydrogen sulphide. filtered off and washed. until no further precipitation takes place; or it is treated with ammonia in excess and sat- The precipitate, which consists of FeS, Cr(QH),Al(OH), UO.S, MnS, NiS, CoS, and ZnS, is concentrated HCl as possible, treated by Table II with a large excess of ammonium car- y º bonate, warmed slightly (should not be boiled long!), and filtered. I. PRECIPITATE J. SOLUTION [Fe(OH)3, Cr(OH)3]. [UO2(CO3)3(NHA),. according to Table I. || ing to Table I. Iron and chro- U ra n i u m is mium are tested for || tested for accord- G. PRECIPITATE H. SoluTION [Al(OH)8]. [Zn(NH3)3Cl2]. A white gelat- The solution is inous precipitate | a c i di fi e d with shows the pres- || acetic acid and ence of alumin- || hydrogen sulphide ium. It is con- || passed into it. A firmed by the Thé- || white recipitate nard’s blue re- || of ZnS sº zinc. action (p. 75). It is confirmed by the R in n- mann’s green re- action (Table II). § § TABLE IV. SECOND METHOD OF SEPARATING ALL THE MEMBERS OF THE AMMONIUM SULPHIDE GROUP. (To be used when no cobalt is present, or only a small amount.) The metals are assumed to be present as chlorides. The iron is oxidized by means of nitric acid,” and the solution boiled for some time to expel all free chlorine; t 20 c.c. NH,Cl solution are added; the solution is precipitated at the boiling tem- perature by as little ammonia as possible and filtered quickly. A. PRECIPITATE [Fe(OH)3, Al(OH)3, Cr(OH)3, U.O.(NHA)2].j: B Solution [MnOl., ZnCl2 NiCl2, CoCl2]. The precipitate is dissolved in as little concentrated HCl as The solution is precipitated with colorless ammo- C dº.t;§ & )s, D. Solution [Al(ONa).]. [MnS, ZnS, NiS, CoS] Examined according Contains members of Examined accord- || Acidified with HCl and precipitated | to Table II. the following groups. ossible, diluted with a little water, precipitated with an excess of || nium sulphide and filtered. aOH and ſiltered. E. PRECIPITATE F. SOLUTION . P IX [Te(OH MnS, ZnS, NiS. CoSl. º gº ing to Table III, E. with ammonia; a white gelatinous pre- cipitate shows aluminium. Confirm by Thénard’s blue reaction (see page 75). * The solution should be tested, to see if all of the ferrous iron has been oxidized, by adding a drop of the solution to a drop of potassium ferricyanide on the cover of a porcelain crucible. There should be no trace of Turnbull’s blue apparent. f Free chlorine is formed by the action of nitric acid upon hydrochloric acid or upon the chlorides present. If it is not O expelled, manganese, nickel, and cobalt will be precipitated as Mé , Ni(OH)3 and Co(OH)s. OH f If considerable manganese is present, some will always be carried down with this precipitate. It is, therefore, advisable to dissolve the precipitate produced by ammonia in hydrochloric acid and repeat the precipitation with ammonia. The fil- trate from the second precipitation is added to B. Cobalt acts still worse than manganese. If considerable cobalt is present, even four repeated precipitations with am- monia are insufficient to completely separate the cobalt from the iron, because the ammoniacal solution absorbs oxygen rap- idly, precipitating the cobalt as Co(OH)3. In this case it is always better to effect the separation according to Table III. I46 REACTIONS OF THE METALs. METALS OF GROUP II. HYDROGEN SULPHIDE GROUP. MERCURY, LEAD, COPPER, BISMUTH, CADMIUM, ARSENIC, ANTIMONY, TIN (GOLD, PLATINUM, SELENIUM, TELLURIUM, VANADIUM, TUNGSTEN, MOLYBDENUM, THALLIUM). MERCURY, Hg. At. Wt. 200.0. Occurrence.—Mercury occurs in nature chiefly in the form of rhombohedral cinnabar, HgS; and deposits of cinnabar form the chief source of all our mercury. According to G. F. Becker,” cinnabar is deposited from Solutions of its sulpho salt. The richest deposits are those of New Almaden in California, where it occurs with Serpentine; of Almaden in Spain, Idria in Carniola, and Moschellandsberg in the Palatinate of the Rhine. With the cinna- bar small amounts of native mercury are often found. Mercury is also an important constituent of many varieties of tetrahedrite. Metallic mercury is the only one of the metals which is liquid at ordinary temperatures. Its specific gravity is 13.595; it solidifies at –39.4°C., and boils at 357° C, slightly above the boiling-point of sulphuric acid, which is 338°C. Mercury is insoluble in hydro- chloric and dilute sulphuric acids, but is soluble in hot concentrated sulphuric acid with evolution of sulphur dioxide, forming mercurous or mercuric Sulphate according to whether the metal or the acid is present in excess. The formation of mercuric sulphate takes place according to the equations 1. Hg-i-H,SO,-HgC)+H,0+SO,. 2. HgC)+H,SO,-HgSO,--H,O. Hydrobromic acid hardly attacks the metal at all, while in hydriodic acid the metal dissolves readily with evolution of hydrogen: Hg-i-4HI=HgI,H,--H,. The proper souvent for mercury is nitric acid. * Geology of the Quicksilver Deposits of the Pacific Slope. Washington, 1888. MERCURY. I47 If the metal is treated with hot concentrated nitric acid, mer- curic nitrate is formed: 3Hg+8HNO,-3Hg(NO),44H,0+2No. If, however, cold nitric acid is allowed to act upon an excess of mercury in the cold, mercurous nitrate is obtained: Hg(NO.), H-Hg=Hg,(NO3)2. Mercury is attacked by chlorine, forming calomel (mercurous chloride): 2Hg-HCl,-Hg,Cl,. Two oxides of mercury are known: 1. Yellow or red mercuric oxide, HgC). 2. Black mercurous oxide, Hg,O. These oxides are basic anhydrides, from which two series of salts are derived: (a) The mercuric salts, which contain the group Hg 3, and (b) the mercurous salts, which contain the group Hg— Hg— We will consider first the more stable mercuric salts. Mercuric Salts. Mercuric salts are mostly colorless. The iodide is red or yellow. By heating the red tetragonal crystals of mercuric iodide a yellow sublimate of orthorhombic needles is obtained, which gradually changes back to the red tetragonal modification; very quickly, almost instantly, if the yellow crystals are rubbed. This is a general property of dimorphous bodies; the more symmetrical form is almost always the more stable. The sulphide is black or red. Mercuric chloride is soluble in water, 100 parts of water dis- Solving at 10° 20° 50° 80° 100° 6.57 7.39 11.34 24.3 53.96 grams HgCl, In water containing hydrochloric acid, mercuric chloride is much more soluble than in pure water; and in fact the solubility in- I48 REACTIONS OF THE METALS. creases with the concentration of the hydrochloric acid, probably due to the formation of the complex acid (HgCl)H, The aqueous solution of mercuric chloride is a poor conductor of electricity; it is only dissociated to a slight extent and acts quite differently in many cases from a Solution of the nitrate, which is a good conductor of electricity and therefore contains a good many mercuric ions. The cyanide differs from the nitrate even more, as we shall See. Mercuric bromide is difficultly soluble in water (94 parts of water dissolve at 9°C. only 1 part of the bromide), but is readily soluble in alcohol, and still more soluble in ether. The iodide is more difficultly soluble. The halogen compounds of mercury readily form complex com- pounds with the halogen compounds of the alkalies, which are very stable. Mercury compounds are furthermore characterized by the readi- ness with which they undergo hydrolysis, forming insoluble basic salts. Thus the sulphate is decomposed when diluted largely with water (particularly on warming) into a yellow insoluble basic salt: Hg—O HgSO, Xo HgSO,--2H,0=2H,SO,--Hg SO,. HgSO, Xo Hg—O The basic sulphate dissolves readily in hydrochloric acid. The nitrate also is readily hydrolyzed into more or less insoluble basic salts according to the “mass action ” of the water. The neutral Salt forms with water a basic Salt according to the equation NO OH Hg6.3.4. HoH = HNo.4 Hgº, OT Hg—OH 2Hg(NO),4-2HOHz-> o( +3HNO3. Hg—NO, MERCURY. I49 REACTIONS IN THE WET WAY. A solution of mercuric chloride and one of mercuric nitrate are used for these reactions. 1. Potassium Hydroxide precipitates yellow mercuric oxide: C1 , KOH [The hydroxides of the ‘noble metals are exceedingly unstable; they lose water, as a rule, even in aqueous solution, forming the anhydrous oxide] Hg334 2KOH=2KNO,4-H,0+Hg0. 3 On adding a lesser amount of caustic potash to a solution of mercuric chloride, a reddish-brown precipitate of basic chloride is obtained: Cl Cl Hgó Hº NCI KOH Y + =2KGl+H.O-- Cl ROH / Hs. Hg Sl NCI OT 3HgCl, i-4KOH=4KCl-i-2H,0+Hg,O,Gl, 8XHg The mercuric oxide and the basic salts are readily soluble in acids. 2. Ammonia produces in a solution of mercuric chloride a white precipitate of mercuric amidochloride: ICT HINH —N Hg Gi-F § =NH,Cl4-Hg gº. This compound, the so-called “infusible precipitate,” volatilizes I5o REACTIONS OF THE METALS. before it melts. It is soluble in acids, and in hot ammonium chlo- ride, forming the “fusible precipitate’” —NH * Hg. Ci *NHo-Hº-Nă ă. If ammonia is allowed to act upon mercuric nitrate a white Oxyamido compound is always formed: Tº LANH.4-HO-3NHNO o( XH NO _Nö-4NH,4-H,0=3NH,NO,4- 2TiN V3. Hg Nö. Hg 3 3. Potassium Iodide produces a red precipitate of mercuric iodide, \ HgCl, E2KI–2KCI+HgI, soluble in excess of potassium iodide, forming a colorless complex salt: HgI, +2KI=(HgI)K,. The solution of the latter salt contains no mercuric ions, for it gives no precipitate with caustic soda or potash. The alkaline solution is the so-called “Nessler's reagent,” and serves for the detection of very slight traces of ammonia. There is formed in this reaction the brown-colored body Hg O& "YNH,-1, NHg/ which is soluble in an excess of the “Nessler's reagent,” with an intense yellow color (cf. p. 47). 4. Alkali Carbonates precipitate from both the chloride and the nitrate a reddish-brown basic carbonate in the cold, 29–H5–0–Hgs 4HgCl, 1-4Na,CO,-8NaCl-i-3CO,--Cº-Q X), NO—Hg—O—Hg which on boiling loses carbon dioxide and is changed into yellow mercuric oxide. MERCURY. I51 5. Alkali Bicarbonates produce no precipitation in a solution of mercuric chloride, but do cause precipitation from mercuric nitrate: 4Hg(NO.), H 8NaHCO,-8NaNO,--4H.O--7CO,--Hg,O.COs. 6. Hydrogen Sulphide produces in solutions of mercuric salts a precipitate which is at first white, then yellow, brown, and finally black. The white precipitate is formed according to the following equation: —C. Hg or #X º & Grijºs-Hothº, —Cl H/ Hg & *QC (white) By the further action of hydrogen sulphide, black mercuric sulphide is finally obtained: Hg,Cl,S.--H.S=2HCl·H 3HgS. Mercuric sulphide is insoluble in dilute boiling acids. Hot con- centrated nitric acid transforms it gradually into white Hg,S,(NO), 9HgS+8HNO,-2NO+3S--4H,0+3Hg,S,(NO), which by long boiling is changed into the soluble nitrate. It dissolves readily in aqua regia, forming the chloride with separation of sulphur: 3HgS+6HCl·H2HNO,-3HgCl, H3S4-2NO+4H,O. ..Mercuric sulphide is insoluble in caustic soda and potash solu- tions, and in ammonium sulphide, but it dissolves readily in sodium or potassium sulphide: HgS+ K.S=Hg(SK). By dilution with water this compound is completely hydrolyzed into mercuric sulphide and potassium hydroxide: Hg(SK),4-H,0= KOH+KSH-i-HgS. Therefore it is always necessary to dissolve the mercuric sulphide with considerable potassium sulphide, or with little potassium sul- I 52 REACTIONS OF THE METALS. phide and considerable caustic potash, in order to prevent this hydrolysis. * The fact that Hg(SK), is so readily hydrolyzed explains the formation of cinnabar in nature. In the interior of the earth the sulpho-compound is formed, which is brought by springs to the surface and there undergoes the above decomposition. * 7. Potassium Cyanide produces in a solution of mercuric chlo- ride no precipitation, because the cyanide, as well as the chloride, forms readily soluble complex compounds with alkali chlorides, The following are known: HgCl, + KCl =(HgCl)K, HgCl, +2KCl = (HgCl)K, * Hg(CN ),4- KCl = (HgCN),Cl)K, Hg(CN),4-2KCl =(HgCN),Cl)K, Hg(CN),4-2KCN=(HgCN),)K,. In a concentrated solution of mercuric nitrate, potassium cya- nide produces a precipitate of mercuric cyanide, soluble in consider- able water and in potassium cyanide: Hg(NO), H2KCN=2KNO,--Hg(CN), Mercuric cyanide is the only cyanide of the heavy metals that is soluble in water. It dissolves mercuric oxide perceptibly, forming the complex compound (HgCN),O. Mercuric cyanide is not pre- cipitated by alkali carbonates or by caustic alkalies, because the mercuric oxide is soluble in mercuric cyanide. It is not decom- posed by dilute sulphuric acid, although it is by the halogen acids— most difficultly by hydrochloric acid, and most readily by hydriodic acid; hydrogen Sulphide decomposes it with precipitation of mer- curic Sulphide: Hg(CN),4-H.S=2HCN+HgS. 8. Neutral Alkali Chromates precipitate yellow mercuric chromate from both the chloride and nitrate solutions. On long standing or by boiling, the precipitate becomes red; a basic salt being probably formed. 9. Alkali Dichromates throw down a yellowish-brown pre- cipitate from the nitrate solution, but not from that of the chloride. MERCURY. I53 10. Ferrous Sulphate reduces mercuric nitrate on boiling to metallic mercury: 3Hg(NO,),4-6FeSO,-2Fe(NO),4-2Fe,(SO.),4-3Hg. Mercuric chloride and cyanide are not reduced by ferrous Sulphate. 11. Stannous Chloride reduces mercuric salts, at first to in- soluble mercurous chloride (calomel), 2HgCl, -SnCl,-SnCl4-Hg,Cl, and by further action to metal, Hg,Cl,--Smel,-SnCl,--2Hg. Metallic mercury separates out in the form of a gray powder. By decanting the solution, and boiling the residue with dilute. hydrochloric acid, the mercury appears in tiny globules. 12. Copper, Zinc, and Iron precipitate mercury from Solutions. of its salts: HgCl, H Fe =FeCl,--Hg,” HgCl, H-Cu,-Cu,Cl,+Hg. On placing a drop of mercury Solution (whether of a mercurous or a mercuric Salt) upon a piece of bright copper-foil. a gray spot is formed, which, after being dried, becomes bright as silver on. rubbing. Mercurous Salts. The mercurous salts all contain the bivalent mercurous group. Hg— ſº and are changed more or less readily into mercuric salts, Hg— splitting off one atom of mercury from the molecule. Mercurous salts containing, oxygen, like mercuric salts, are readily hydrolyzed in dilute aqueous solutions; thus the nitrate is decomposed accord- ing to the equation t Hg—NO, Hg—NO, # * +HOH=HNO,-- g—NO, tº Hg—OH. * This reaction is employed for detecting metallic iron in the presence of: Fe0 (cf. p. 96). ºv I54 REACTIONS OF THE METALS. Mercurous chloride (calomel) is insoluble in water and hydro- chloric acid, but soluble in nitric acid and aqua regia. REACTIONS IN THE WET WAY. 1. Caustic Potash precipitates black mercurous oxide: Hg—NO, Hg | +2KOH=2KNO,-- >0+H.O. Hg—NO, Hg 2. Ammonia produces a black precipitate of mercuric amido salt with metallic mercury: Hg—NO, /Hgs 2# +4NHa-i-H.O =3NHANOa-i- oq XNH,-No,4 2Hg. g—NO, Hg J Y It can easily be shown that this precipitate contains metallic mercury by rubbing a piece of pure gold over it, when silver-lustrous gold amalgam will be formed. Mercurous chloride gives with ammonia a mercuric amine with separation of metallic mercury: ! Hg—Cl § NH, * t” | +2NH, -NH.Ql-H Hg +Hg. Hg—Cl Cl \—Y—’ 2. § boiling the black precipitate with dilute hydrochloric acid OT *. concentrated ammonium chloride solution, the mercuric amine goes into solution, leaving behind drops of mercury. 3. Alkali Carbonates give, first, a yellow precipitation of the carbonate, which quickly become gray, owing to the formation of mercuric oxide, metallic mercury, and carbon dioxide: Hg,(NO),4-Na,CO,-2NaNO,--Hg,COs, and Hg,CO,-HgO--Hg+CO,. 4. Ammonium Carbonate yields the same precipitate as am- monia. § MERCURY. I55 5. Hydrogen Sulphide immediately throws down a black pre- cipitate of mercuric sulphide and mercury (difference from mer- curic salts): Hg,(NOs),-- H.S =2HNO2+HgS+Hg. The black precipitate does not dissolve completely in potassium sulphide, the mercury remaining insoluble, but in alkali polysul- phides it dissolves. 6. Hydrochloric Acid and Soluble Chlorides precipitate white mercurous chloride (calomel), Hg—NO, HCl Hg–Cl + := 2HNO2+ | 5 Hg–NO, HCl Hg–Cl insoluble in water and dilute acids, soluble in strong nitric acid and aqua regia. On boiling for a long time with water, calomel be- comes gray, owing to a partial decomposition into mercuric chlo- ride and mercury. On boiling with concentrated sulphuric acid, mercuric sulphate is formed with evolution of sulphur dioxide and hydrochloric acid: (a) Hg,Cl,--H..SO,-2HCl·H Hg,SO,. (b) Hg,SO,--2H,SO, =2H,0+SO,--2HgSO,. 7. Neutral Potassium Chromate precipitates red mercurous chromate on boiling (cf. p. 86): Hg,(NOs),--K,CrO, =2KNOs--Hg,CrO,. 8. Potassium Iodide precipitates green mercurous iodide, .* Hg,(NOs),--2KI=2KNOs--Hg.I., soluble in an excess of the precipitant, with the formation of mer- curic potassium iodide and separation of mercury: Hg—I | +2KI=(HgI)K. --Hg. Hg—I 9. Potassium Cyanide precipitates metallic mercury, mercuric cyanide being formed at the same time: Hg,(NO),4- 2KCN=2KNO,-- Hg(CN),4- Hg. I 56 REACTIONS OF THE METALS. 10. Stannous Chloride precipitates gray metallic mercury; Hg—NO, | +SnCl,--2HCl=2HNO,--SmC1,4-2Hg. Hg—NO, REACTIONS OF MERCURY IN THE DRY WAY. Almost all mercury compounds Sublime on being heated in the closed tube. Mercuric chloride melts first, then vaporizes, forming a crystal- line deposit on the cold sides of the tube. Mercurous chloride sublimes; the sublimate is almost white, but there is a slight grayish tint owing to the decomposition of a small part of the substance into mercuric chloride and mercury. Mercuric iodide yields a yellow sublimate, which becomes red on being rubbed with a glass rod. Mercury compounds containing oxygen (all more or less un- stable) yield mercury. The sulphide gives a black sublimate. All compounds of mercury, when mixed with sodium carbonate and heated in a closed tube, yield a gray mirror, consisting of Small globules of mercury. In order to make the drops more apparent a piece of filter-paper is placed over a glass rod, with which the mirror is rubbed. The small drops then run together into large ones, Stick to the paper, and can be removed from the glass. LEAD, Pb. At. Wt. 206.9. Occurrence.—Galena, PbS, isometric; cerussite, PbCO, ortho- rhombic and isomorphous with aragonite, CaCO,; anglesite, PbSO, orthorhombic, isomorphous with anhydride, CaSO, celestite, SrSO, and barite, BaSO4; pyromorphite, Pbs(PO4)2Cl, hexagonal; mime- tite, Pbs(ASO),Cl; vanadinite, Pbs(VO.),Cl. The last three min- erals are isomorphous and belong to the apatite group. Other minerals which may be mentioned are wulfenite, PbMoC), tet- ragonal, isomorphous with stolzite, Pb WO4, and the monoclinic crocoite, PbCrO4. Lead is a bluish-gray metal of specific gravity 11.36–11.39, melting at 334° C. and boiling at 1600° C. It is attacked by all acids. As, however, most lead Salts are difficultly soluble in water, it usually becomes coated with a layer of salt, which protects it DEAD. I57 from further action of the acid. Thus lead is immediately at- tacked by dilute Sulphuric acid according to the equation Pb-H H.SO,-PbSO,4-H, But, as the lead sulphate formed is insoluble in dilute sulphuric acid, the reaction quickly ceases. Upon this principle rests the use of “lead chambers ” in the manufacture of sulphuric acid, and the use of “ lead pans” for the concentration of the dilute “chamber acid.” It has, however, been found from experience that the sulphuric acid should not be concentrated too much in lead pans—stopping when a 78–82 per cent. acid is obtained. The protecting layer of lead sulphate is soluble in hot concentrated sulphuric acid, forming soluble lead bisulphate, * O—H sok PbSO,--H.SO,- ">r, Pb-i-3H, SO,-SO,--2H,0+ XPb. O so..., H Lead behaves quite similarly on treatment with hydrochlorie acid. On the surface a protecting coating of lead chloride is obtained, which is soluble in hot concentrated hydrochloric acid, forming PbCl,H. Lead is therefore soluble in concentrated hydrochloric acid: Pb-H3HGl=PbCl, H+H,. Hydrofluoric acid attacks lead similarly, forming a protecting layer of lead fluoride, which is insoluble in hydrofluoric acid. Con- sequently lead retorts can be used for the distillation of hydro- fluoric acid and in the preparation of hydrofluoric acid by means of fluorite and sulphuric acid. I58 REACTIONS OF THE METALS. Nitric acid is the proper solvent for lead. Lead nitrate is in- soluble in strong nitric acid, so that lead does not dissolve in con- centrated nitric acid; the solution must be sufficiently dilute to prevent the separation of the lead nitrate formed. Lead forms the following oxides: Lead suboxide, Lead oxide (litharge), Lead sesquioxide, Pb,0; Pb=O; b,0s; Minium (red lead), Lead peroxide, Pb,04; PbO,. Of these oxides, PbO alone is the anhydride of a base; * from it the salts of lead are derived, in which the lead is bivalent. This lead monoxide (litharge) is a yellow powder, which melts at about 980° C., and solidifies on slow cooling, forming tetragonal crystals (needles). It is slightly soluble in water with an alkaline reaction, and is readily soluble in dilute nitric acid. Lead suboxide, Pb,O, is formed as a black velvety powder on heating the oxalate to about 300° C.: 2PbCO,-3CO, HCO+Pb,0. On heating the suboxide in the air, it becomes readily oxidized to lead monoxide. Lead dioxide, PbO, must be considered as the anhydride of the acids: H 28; OH PbToi Pb-O , NOH NOH Orthoplumbic acid, Metaplumbic acid just as SiO, SnO, CO, MnO, are anhydrides of silicic, stannic, 2OH carbonic, and manganous acids. The acid Pºé. is formed by the oxidation of lead hydroxide, Pb(OH), in alkaline solution by means of hypochlorites, chloriné, bromine, hydrogen peroxide, or potassium persulphate: —OH /OH Pb ºf H2NaOH+ Cl,-H,0+2NaCl-- Pººr * Although Pb(C,EI,O.), is known. DEAD. I 59 The brown metaplumbic acid which separates out goes over at 100° C. into the anhydride; and the latter on ignition loses oxygen, changing into yellow lead monoxide. The other two oxides of lead, Pb,O, and Pb,O, must be regarded as salts of the plumbic acids; O Pb,O, as the salt of metaplumbic acid: Pºp!, ; and Pb,O, as Nö/ 23)Pb the salt of the hypothetical orthoplumbic acid: PbTö wº Nºrb Pb,O, is obtained as a yellow precipitate on gently oxidizing an alkaline solution of lead monoxide by means of hypochlorites, halo- gens, hydrogen peroxide, or persulphates: —OH +2NaOH+Cl,-2NaCl-i-3H,0+Pb,O, and the red minium, Pb,O, by igniting lead oxide or lead car- bonate for some time in the air at about 430° C.: 3PbO4-O =Pb,0,. Both bodies behave chemically as salts; for, on treating with nitric acid, brown plumbic acid and lead nitrate are formed; which corresponds to the action of nitric acid on, say, lead carbonate: /QN /OH Pb-ºrbi-2HNO,-Pb(NO), Pio NO NOH O OH coºper 2HNO,-Pb(NO),4- coºHot CO) /3>Pb /OH Pº, HHNO-Pb(NO),4-H,0+Pho NöXPb NOH These salt-like oxides * are completely analogous to those of manganese; on treatment with hydrochloric acid they yield chlo- * Besides the lead salts of plumbic acid, alkali and alkaline earth salts are known. I6o REACTIONS OF THE METALS. rine, the plumbic acid, at first set free, behaving like a per- oxide: PbO, --4HCl=2H,0+ PbCl, HCl, Pb,O,4-6HCl=3H,0+2PbCl, HCl, Pb,O,--8HCl=4H,0+3PbCl, HCl, REACTIONS IN THIE WET WAY. Lead salts are difficultly soluble or insoluble in water; but all dissolve in dilute nitric acid, excepting, perhaps, fused lead chro- mate, which is very difficultly soluble. 1. Potassium and Sodium Hydroxides precipitate white lead hydroxide: O OH Pº 842KOH-2KNo.4 Pb(#. soluble in an excess of the precipitant, forming a plumbite: * OH OK Pb(6; OK” +2KOH=2H,0+Pb Pb(OH), is also slightly soluble in water, except when it contains carbonic acid. The aqueous solution of lead hydroxide is slightly alkaline. 2. Ammonia precipitates the white hydroxide, insoluble in excess of the reagent. 3. Alkali Carbonates precipitate white basic lead carbonate. Alkali bicarbonates precipitate the normal carbonate. 4. Sodium Phosphate precipitates white lead phosphate: 3Pb(NO),4-4Na,HPO,-2NaH,PO,--6NaNO,--Pb,(PO,), insoluble in acetic acid, readily soluble in nitric acid, caustic soda or potash. 5. Potassium Cyanide precipitates white lead cyanide, insoluble in an exocSS. ,--, 6. Hydrochloric Acid or Soluble Chlorides precipitate from { }oderºy concentrated Solutions flocculent, white lead chloride: Pb(NO),4-2HCl=2HNO,--PbCl, * According to Hantzsch, Zeitschr. f. anorg. Ch., XXX (1902), p. 289, the /H O . alkaline solution contains the lead as Pb DEAD. 161 difficultly soluble in cold water (135 parts of water dissolve 1 part of PbCl2, but much more soluble in hot water; on cooling the solution, lead chloride separates in the form of glistening needles or plates. Lead chloride is much more soluble in concentrated hydro- chloric acid and in a concentrated solution of a chloride of an alkali than it is in water, as it forms complex compounds with these bodies, which are, however, decomposed on dilution with water, with separation of lead chloride. 7. Potassium Iodide precipitates yellow lead iodide: Pb(NO.),4-2KI=2KNO,--Pbſ, The iodide is much more insoluble in water than the chloride (194 parts of boiling water dissolve only 1 part of lead iodide), forming a colorless solution from which lead iodide separates on cooling, in the form of gold-yellow plates. The iodide dissolves to a considerable extent in hydriodic acid, and in a solution of an alkali iodide, forming lead hydriodic acid; Pbſ,H or one of its salts (such as Pb.I.K), all of which are decom- posed on dilution, with deposition of lead iodide. , 8. Alkali Chromates produce a yellow precipitate of lead chro- mate: d Pb(NO,),4-K,CrO,-2KNO,--PbCrO, 3.I) ... " 2Pb(C.H.O.), i-K,Cr,O,--H,0=2KC.H.O,--2HC.H.O.--2PbCrO,. Lead chromate is insoluble in acetic acid, but soluble in nitric acid and in caustic alkali. 9. Hydrogen Sulphide produces in the most dilute lead solutions (from slightly acid solutions, as well as from neutral or alkaline ones) a black precipitate of lead sulphide: Pb(NO,),4-H, S=2HNO,--PbS. From hydrochloric acid solutions an Orange-red precipitate of lead sulphochloride is at first obtained: Pb–Cl 2PbCl, L.H.S=2HCl-H X, Pb–Cl which is decomposed immediately by more hydrogen sulphide, forming the black lead sulphide. In this behavior it is quite similar to mercuric salts (see page 151). 162 REACTIONS OF THE METALS. Lead sulphide is soluble in dilute, boiling, double-normal nitric acid, forming lead nitrate, with separation of sulphur: 3PbS+8HNO,-3Pb(NO.),4-4H.O-H2NO+3S. The reaction usually goes a little further; some of the sulphur is oxidized to sulphuric acid, forming insoluble lead sulphate. The amount of sulphuric acid formed (and therefore of the lead sulphate also) increases with the concentration of the acid. Lead sulphide is also soluble in strong hydrochloric acid: PbS4-2HCl=PbCl, L H.S. 10. Sulphuric Acid and Soluble Sulphates cause in solutions of lead salts the separation of white difficultly-soluble lead sulphate: Pb(NO.),4-H.SO,-2HNO,--PbSO,. One part of the salt dissolves at the ordinary temperature in 22,800 parts of water; in water containing a little sulphuric acid it is still less soluble, while in alcohol it is insoluble. Lead sulphate dis- solves perceptibly in nitric acid; in hot concentrated hydrochloric acid it is completely soluble. On cooling the hydrochloric acid solution, lead chloride separates out in needles. Lead sulphate is also soluble in concentrated Sulphuric acid, readily on boiling and considerably in the cold, forming acid lead sulphate which is decomposed by water into Sulphuric acid and lead sulphate. Almost all the sulphuric acid of commerce con- tains some dissolved lead sulphate. In order to detect this, 200–300 c.c. of the concentrated acid should be diluted with an equal volume of water and allowed to stand twelve hours, whereby the lead sulphate separates as a white powder. Besides being soluble in acids, lead sulphate is easily soluble in caustic alkalies, and in Solutions containing the ammonium salts of many organic acids. This last property is of great importance for the analytical chemist, as it offers a means for Separating lead Sulphate from barium sulphate, silica, etc., which remain undis- Solved. Ammonium acetate or ammonium tartrate is usually used as the solvent. The Solution of lead sulphate by means of a concentrated solu- BJSMUTH. I63. tion of ammonium acetate takes place perhaps according to the following equation: /O—NH, SO, 2PbSO,--2NH,C,E,O,- X3>Pb +Pb(§ 2 2 O So-NH, By diluting the solution with water, the reaction takes place in the opposite direction, and lead sulphate is precipitated; which is also effected by the addition of sulphuric acid. According to Kahlenberg and Schreiner (Zeitschr. f. phys. Ch. 17) lead sulphate dissolves in ammonium tartrate in the presence of ammonia with the formation of COONH, NH,OOC | | CHOH HOHC | | CHOH HOHC | - | COO—Pb—O—Pb—OOC REACTIONS IN THE DRY WAY. Heated with sodium carbonate on charcoal, all lead compounds yield a malleable button, surrounded with an incrustation of the yellow oxide. On the charcoal stick also, the malleable button is readily obtained. Lead glasses turn black on heating in the reducing flame, owing to the separation of lead. BISMUTH, Bi. At. Wt. 208.5. Occurrence.—Bismuth usually occurs native with nickel and cobalt ores. The following ores are of no great importance: Bis- mite, Bi,0,; bismuthenite, Bi,S,; emplectite, Bi,S,Cu,; bismutite, 3|CO, [BiOH]-5Bi(OH), Bismuth is a brittle, reddish-white metal of sp. gr. 9.8; which crystallizes in the hexagonal system, melts at 268°C. and vaporizos at about 1600° C. The proper solvent for bismuth (as is the casc with most other metals) is nitric acid. Hydrochloric acid does not ** attack bismuth, and sulphuric acid dissolves it only on warming. ** IO4 REACTIONS OF THE METALs. Bismuth forms two oxides: Bismuth trioxide and Bismuth pentoxide. Biº-O O=Biº- º sº Bizºo yellow. O=Biºč, brown. wº-ºº-ºº: Bismuth is either trivalent or pentavalent like nitrogen, phos- phorus, arsenic, and antimony. Bismuth trioxide is a basic anhydride,” from which the salts are derived. Bismuth pentoxide, a brownish body, acts as an acid anhydride, forming an acid HBiO, corresponding to metaphos- phoric acid. Salts of this acid have never been prepared in a pure state. On igniting, Bi,0s loses oxygen, forming yellow Bi,0, It dissolves in hydrochloric acid with evolution of chlorine, forming a salt of trivalent bismuth: Bi,Os--10HCl=5H,0+2BiCl,--2Cl, Bismuth Salts are mostly colorless, and are all insoluble in con- siderable water, on account of being hydrolyzed into an insoluble basic salt; thus the chloride is quantitatively decomposed into bismuth oxychloride: BiC, H.H.Oa2HCH-Biº, insoluble in tartaric acid (difference from antimony). Bismuth oxychloride is readily soluble in hydrochloric acid, the above equation taking place from right to left. The reaction there- fore is reversible and the relative amounts of water and hydrochloric acid present determine in which direction the reaction will go. On adding water to a slightly acid solution of BiCl, a white precipitate of the oxychloride appears immediately. On carefully adding hydrochloric acid, the precipitate again dissolves, but may be re- precipitated by the addition of more water. All the other com- pounds of bismuth act as the chloride. The nitrate yields, at first, an amorphous precipitate of BiONO,. /NO º #Xo =2HNO.4 Biºlo., 3 3 *Bismuth trioxide acts as a weak acid under some circumstances (cf. foot-note p. 165). BISMUTH. 1.65 which becomes more basic on further addition of water, becoming crystalline at the same time: O OH BićNo, Biº, I.F.HoH = HNo.1">dº Bºo. Bi-O This last compound is the bismuth submitrate, which is so much used in medicine. REACTIONS IN THE WET WAY. 1. Potassium Hydroxide precipitates, in the cold, white bis- muth hydroxide: BiCl, H3KOH=3KCl+Bi(OH), which, on boiling, becomes faint yellow, because it loses water and forms a less hydrated hydroxide: Bi(OH)=H,0+ Bión. Both of these hydroxides are insoluble in an excess of the pre- cipitant,” but are readily soluble in acids. On adding to the alkaline Solution, in which the hydroxide is suspended, chlorine, bromine, hypochlorites, or hydrogen peroxide, the white or yellowish precipitate becomes brown, owing to the formation of bismuthic acid: O O Bić +2Naoh-oo-ºnacº Ho-biº OH NOH 2. Ammonia precipitates a white basic salt (not the hydroxide), the composition of which varies with the concentration and with the temperature. * 3. Alkali Carbonates precipitate, according to the temperature and concentration, a number of basic carbonates; one of which is formed according to the following equation: ſº ./CO 2BiCl,--3Na,CO,4-HOH=6NaCl-KCO,-- 2B& *In very concentrated KOH, Bi(OH)3 dissolves on warming. On cool- ing, part of the Bi(OH), is precipitated, and on dilution all of it. In this case the hydroxide acts as a weak acid, like antimony trioxide. I66 REACTIONS OF THE METALS. 4. Sodium Phosphate precipitates the white, granular phos- phate, insoluble in dilute nitric acid, difficultly soluble in hydro- chloric: 2Na,HPO,--BiCl, H3NaCl- NaH,PO,--BiPO,. 5. Potassium Cyanide precipitates the white hydroxide (not the cyanide). The cyanide is at first formed, but is hydrolyzed: (a) BiCl, H3KCN=3KCl-i-Bi(CN), (b) Bi(CN), H 3HOH=3HCN-I-Bi(OH), 6. Potassium Dichromate precipitates yellow bismuthyl di- chromate: *** CrO,-OK BiCl, H.O CrO,-O(Bi=O)f Xo + + =4HCl*-i-2KCl-H YO CrO,-OK BiCl, H.O CrO,-O(Bi=O) Soluble in mineral acids, insoluble in caustic alkalies (difference from lead). 7. Hydrogen Sulphide precipitates brown bismuth sulphide: 2BiCl, F3H,S=6HCl- Bi,S, insoluble in cold dilute mineral acids and alkaline sulphides, soluble in hot dilute nitric acid, and in boiling, concentrated hydrochloric acid, g 8. Alkali Stannites (an alkaline solution of stannous chloride) cause a black precipitation of metallic bismuth. This very sen- sitive reaction is performed as follows: To a few drops of stannous chloride, caustic alkali is added until the white precipitate at first produced dissolves clear, and the bismuth solution is then added in the cold; on shaking, a black precipitate immediately appears. The following reactions take place in this test: (a) SnCl,+2KOH=2KCl+Sn(OH), (white precipitate). (b) Sn(OH), H2KOH=2H,0+Sn(OK), (potassium stannite). (c) 2BiCl, 1-6KOH=6KCl-H2Bi(OH), (d) 2Bi(OH),4-3Sn(OK), -3H,0+3SnO(OK),4-Bi, Potassium stannate. * An excess of K,Cr,O, should be employed to prevent the solvent action of the hydrochloric acid; the hydrochloric acid is then used up as follows: R,Cr,O, 4-2HCl = H,Cr,0, + 2KCl. # The univalent group (BiO) is called bismuthyl. * Wanino & Treubert, B B., 1898, p. 1113. BISMUTH. 167 So that the whole reaction may be represented as follows: 3SnCl4-2BiCl, H 18KOH=12KCI+9H,0+3SnO(OK),4-Bi, In making this test, a too concentrated caustic alkali Solution should be avoided; and the solution must be kept cold, otherwise the stannite itself may give a black precipitate. 1. If too much caustic potash is used, metallic tin will separate out: 29 2Sn(OK),4-HOH= shºt 2KOH-H Sn. OK - 2. If too little caustic potash is used, black stannous oxide will be thrown down in the cold, after long standing; quickly on boiling: —OK Sn 3:4-HOH-2KOH+Sno. 9. Potassium Iodide precipitates black bismuth iodide: BiCl, -3KI=3KCl--BiL, soluble in excess of the reagent, with a yellow to orange color: Biſs--KI = BiLiK. By diluting this last solution with not too much water, the black iodide is reprecipitated, which, on the addition of more water, is changed into orange-colored basic iodide: Biſ.--H,0=2HI+ BiOI. 10. Metallic Zinc precipitates metallic bismuth: 2BiCl, 4-3Zn=3ZnCl,--Bi,. REACTIONS IN THE DRY WAY. Bismuth salts color the non-luminous flame a pale greenish white. Heated with soda on charcoal before the blowpipe, a brittle button of the metal is obtained, Surrounded by a yellow incrusta- tion of bismuth oxide. On heating a compound of bismuth in the upper reducing flame (page 27) of the Bunsen burner, the bismuth is reduced to metal, I68 REACTIONS OF THE METALS. which is volatilized and burnt to oxide in the upper oxidizing flame. On holding a porcelain evaporating-dish (glazed on the outside and filled with water) just above the oxidizing flame, a barely visible deposit is obtained, which, on being treated with hydriodic acid,” is changed to Scarlet bismuth hydriodic acid: Bi,0,--8HI=3H,0+2Bil,H. By breathing on this deposit, the color disappears, but reap- pears as soon as the moisture has evaporated. On exposure to fumes of ammonia (by blowing the vapors away from the stopper of an ammonia bottle) the deposit is colored a beautiful orange, owing to the formation of the ammonium salt of the bismuth hydriodic acid: * BiLH+NH,-BiL(NH.), which also becomes invisible on being breathed upon. By moistening this coating with an alkaline solution of stannous chloride, black metallic bismuth is deposited. COPPER, Cu. At. Wt. 63.60. Occurrence. Native copper, Cuprite, Chalcocite, Chalcopyrite, Cu Cu,0 Cu.S CuFeS, Isometric Isometric Orthoriombie Tetragonal Malachite, Azurite, Atacamite. Cu–OH Cu—OH Cu—Cl Xco, Xco, Xo +H,O. tº e Cu–OH Cu Gu-OH g Monoclini;fCºlomorpheus YCO, Orthorhombic • Cu–OH Monoclinic Copper is a ductile metal possessing a peculiar red color, and a Sp. Gr. of 8.94. It melts at about 1064°C., forming a bluish-green liquid. *s The proper solvent for copper is nitric acid: 3Cu+8HNO,-4H,0+2NO+3Cu(NO.), * The hydriodic acid is easiest obtained by moistening a piece of asbestos, held in the loop of a platinum wire, in a solution of alcoholic iodine solution and then setting fire to the moist asbestos. By holding the burning asbestos under the dish, enough hydriodic acid is formed to change the bismuth oxide into the red compound. COPPER. 169 . Bright copper is not dissolved by hydrochloric acid; but hot hydrobromic acid dissolves it with evolution of hydrogen, forming cuprous hydrobromic acid: (a) Cu, +2HBr=Cu,Br,4-H. (b) Cu, Br,4-4HBr a 2 Cu, Br. H. At the beginning of this reaction the solution usually turns dark violet on account of the formation of the cupric salt of cuprous hydrobromic acid, owing to the copper being somewhat oxidized on the surface. In this case, however, the solution soon becomes colorless; the cupric salt is reduced by the nascent hydrogen. On adding water to the clear solution cuprous bromide is precipitated. Copper is not attacked by dilute sulphuric acid, but it dissolves in hot concentrated Sulphuric acid, forming cupric Sulphate with evolution of sulphur dioxide. & The copper is at first oxidized (at the cost of the sulphuric acid) to cupric oxide, and this latter dissolves in more of the acid: (a) Cu+H,SO,-H,0+SO,--CuO. (b) CuO +H,SO,-H,0+CuSO,. Copper forms two oxides: red cuprous oxide, Cu,0, and black cupric oxide, CuO. Both oxides are basic anhydrides, forming cuprous and cupric salts. The cuprous series contains the bivalent cuprous group: Cu— | , while the cupric series contains the simple, bivalent copper Cu— atom Cuk. A. Cuprous Compounds. The cuprous compounds are extremely unstable, being oxidized quickly to cupric compounds. The only known cuprous salts are those with the halogens; * these are colorless, insoluble in water, but readily soluble in concentrated halogen acids forming colorless solutions. Such solutions contain the unstable cuprous halogen acids, probably of the formula [Cu, N.H., in which “X” is either chlorine, bromine, or iodine. Salts are known which are derived from these acids, e.g., [Cu,CljR. * Cu,SO, is said to exist in solution. 17o REACTIONS OF THE METALS. The cuprous halogen acids are changed dark on contact with air -—the chloride, brownish black; the bromide, dark violet; probably due to the formation of cupric salts of the cuprous halogen acids. The behavior of the cuprous halogen acids toward carbon mon- oxide is very important; the latter is readily absorbed, forming an unstable compound: Cu2Cl2 +2CO +2H2O z=2. Cu2Cl2-2CO-2H2O. By boiling the solution the compound is decomposed into cuprous chloride and carbon monoxide; cuprous chloride is used in gas analysis for the absorption of this gas. _REACTIONS IN THE WET WAY. A solution of cuprous chloride in hydrochloric acid should be used, which may be prepared as follows: Two gms. of cupric oxide are dissolved in 25 cc. of hydrochloric acid (Sp. Gr. 1.12);" he solution is poured into a flask and 0.58 gram of copper fili 3 added; a copper spiral is placed in the flask, one end of w...ch reaches into its neck; the flask is then stoppered and allowed to stand a few days. The originally dark solution will gradually be- come colorless, when it is ready to be used for the following reac- tions: 1. Potassium Hydroxide produces in the cold a yellow precipi- tate of cuprous hydroxide: Cu—Cl Cu–OH | +2KOH=2KCI-E 2 Cu——Cl Cu–OH which loses water at the boiling temperature, changing to red cu- prous oxide: Cu,(OH), H.H.O.--Cu,O. 2. Hydrogen Sulphide precipitates black cuprous Sulphide: Cu,Cl, 4-H, S=2.HCl·HCu,S, soluble in warm dilute nitric acid, forming blue cupric nitrate, with separation of Sulphur: --- 3Cu,S+16HNO,-8H,0+4NO+3S--6Cu(NO), * COPPER. 171 3. Potassium Cyanide prêpitates white cuprous cyanide: Cu,C,4-2KCN=2KC1+Cu,(CN), soluble in excess, forming colorless cuprous potassium cyanide: Cu,(CN),4-6KCN a [Cu, (CN), Ks. This solution contains no cuprous ions, but ions of cuprous cyanide [Cu,(CN), , and potassium, and gives no precipitation with potas- sium hydroxide or hydrogen sulphide. This property is made use of in the separation of copper from cadmium. `-- g Potassium cuprous cyanide, on being diluted with considerable water, forms successively the following compounds: [Cu,(CN).]K, [Cu,(CN), H.O]K, and finally Cu,(CN), All of these compounds, even in the solid state, are decomposed by hydrogen sulphide with precipitation of black cuprous sulphide. Coosanuently, in order to prevent the precipitation of copper by hydro- gén, phide, considerable potassium cyanide must be added, more than enough to form the salt [Cu,(CN)s]Ks. B. Cupric Compounds. Cupric salts are either blue or green in aqueous solution; in the anhydrous state they are white or yellow. The chloride, nitrate, sulphate, and acetate are soluble in water; most of the remaining salts are insoluble in water, but readily solu- ble in acids. FEACTIONS IN THE WET WAY. A solution of copper sulphate should be used. 1. Potassium Hydroxide produces in the cold a blue precipitate of cupric hydroxide:* * CuSO,--2KOH=K,SO,--Cu(OH), which on boiling becomes changed into brownish-black cupric oxide. * CuO.H. dissolves in very concentrated KOH or NaOH, particularly on warming, with a blue color. Cf. page 133, Remark. I 72 REACTIONS OF THE METALS. 2. Ammonia.—On adding ammonia cautiously to the solution of a cupric salt, a green powdery precipitate of a basic salt is ob- tained, which is extremely soluble in an excess of the reagent, form- ing an azure-blue Solution: Cu–OH CuSO, N (a) +2NH.OH=(NH),SO,-- SO,. CuSO, Cu—OH (b) Cu,(OH),SO,--(NH),SO,--6NH,-2(Cu(NH,), (SO). H.O. On adding alcohol to the concentrated blue solution, the above body is precipitated as a blue-violet crystalline substance, which gradually loses ammonia on being warmed, leaving behind the cupric salt. On conducting ammonia gas over an anhydrous cop- per salt, the ammonia is eagerly absorbed, whh the formation of a complex cupric ammonium salt: t CuCl,--óNH,-[Cu(NH);]Cl, These compounds (which contain as a maximum 6NH, to one atom of copper) are completely analogous to the corresponding compounds of nickel, cobalt, and zinc. By the precipitation of the ammoniacal solution with alcohol, the compound with 4NH, to one atom of copper is always obtained. 3. Potassium Cyanide produces, at first, yellow cupric cyanide, which immediately loses dicyanogen, forming white cuprous cyan- ide. The latter, as we have already seen, forms soluble potassium cuprous cyanide with more potassium cyanide: (a) 20uSO,--4KCN=2Cu(CN), - KSO. /CN CuSON Cu—CN (*) cºst") tº on C CN u—CN (c) Cu,(CN),4-6KCN=[Cu,(CN)]K. On adding sufficient potassium cyanide to the blue ammoniacal cupric solution, the complex compound will be decolorized, forming cuprous potassium cyanide: 2Cu(NH);so, H,049KCN- FC =[Cu,(CN)s] } NH, -NH.CNO+2KSO,4-6NH,4-Ho COPPER. I 73 This apparently complicated reaction is in reality quite simple, when we look at the separate steps. First, ammonia is split off, forming the cupric cyanide: & 2Cu(NH);SO, H,0+4KCN-2Cu(CN),4-8NH,42KSO,4-2H,0. The cupric cyanide then decomposes, as just stated, into cuprous cyanide and dicyanogen: 2Cu(CN), -Cu,(CN),4-(CN),. The cuprous cyanide then dissolves in an excess of the potas- sium cyanide, forming cuprous potassium cyanide (equation (c) above); and the dicyanogen acts upon the ammonia, forming ammonium cyanide and ammonium cyanate (just as chlorine acts upon caustic potash, forming potassium chloride and potassium hypochlorite): af 2KOH+2Cl–KCl+KOC]+H,0; 2NH,OH+2CN=NH,CN+NH,CNO+H,O. 4. Potassium Sulphocyanate, KCNS, produces black cupric sulphocyanate: CuSO,--2KCNS=K_SO,--Cu(CNS), which is gradually changed into white cuprous sulphocyanate, or immediately on adding Sulphurous acid: /CNS + O+H,SO,-H,SO,--2HCNS+ cºs H/T.” “T” u—CNS *NCNS Cuprous sulphocyanate is insoluble in water, dilute hydro- chloric acid, and sulphuric acid. 5. Alkali Xanthogenates produce in solutions of cupric salts, at first, a brownish-black precipitate of cupric Xanthogenate, which splits off ethyl-xanthogen-disulphide, forming finally yellow cu- prous xanthogenate: /gºh. /QC.H. . . NS @ 20-s, roso-Nºsot }>0. NOCH, Sodium xanthogenate Cupric Xanthogenate 174 REACTIONS OF THE METALS. OC, H, ozººi. Cºgs. cºs C.H.QSC s cºs śc ºs-Cu SN – * C=S. Nöch, ch.67 ºn, Choº The reagent, Sodium Xanthogenate, is readily obtained by mix- ing carbon disulphide with alcoholic sodium hydroxide: OC.H cº +choNa-cés” SN a The alkali xanthogenates are not used as reagents in testing for copper, but cupric salts are used in testing for Xanthogenates. The reaction is made use of in the detection of carbon disulphide in gas mixtures; the gases are allowed to act upon alcoholic sodium hy- droxide, whereby sodium Xanthogenate is formed if carbon disul- phide is present, and the Solution after neutralizing with acetic acid is tested for xanthogenate by means of a solution of cupric salt. 6. Hydrogen Sulphide precipitates from neutral or very slightly acid solutions colloidal, black cupric Sulphide; which, as part of it is usually in the form of hydroSol, has a tendency to run through the filter. In order to prevent this, the hydrosol is changed into hydro- gel by the addition of considerable acid. In order to precipitate copper quantitatively from a solution, a large amount of dilute hydrochloric acid (10–20 cc. of double normal acid to 100 cc. of the copper solution) is added and hydrogen sulphide then passed into the solution: CuSO,--H.S=H.SO,--CuS. Copper sulphide is soluble in hot dilute nitric acid, but insoluble in boiling dilute sulphuric acid (difference from cadmium); soluble in potassium cyanide, forming potassium cuprous cyanide. From a solution of the latter salt the copper cannot be precipitated by hydrogen sulphide. Copper sulphide is appreciably soluble in ammonium sulphide, but is insoluble in potassium or Sodium sulphide (difference from mercury). CADMIUM. I 75 7. Potassium Ferrocyanide produces in neutral and acid solu- tions an amorphous precipitate of reddish-brown cupric ferro- cyanide: Fe(CN).K. +2CuSO4– (Fe(CN).Kºi. 2K,SO, insoluble in dilute acids, but soluble in ammonia with a blue color. It is also decomposed by potassium hydroxide; in the cold, light- blue cupric oxide and potassium ferrocyanide are formed, while, on warming, black cupric oxide is obtained (difference from uranium, which yields the yellow uranate both with ammonia and sodium or potassium hydroxide). REACTIONS IN THE DRY WAY. The borax, or salt of phosphorus, bead is green in the oxidation flame when strongly saturated with the copper salt; blue if con- taining only a small amount. The reducing flame decolorizes the bead unless too much copper is present; in such cases it is reddish brown and opaque. Owing to the separation of copper. Traces of copper may be determined with certainty as follows: To the slightly bluish bead produced by the oxidizing flame, a trace of tin or of a tin compound is added; it is heated in the oxidizing flame until the tin has completely dissolved, when it is slowly brought into the reducing flame and then quickly removed. The bead now appears colorless when hot, but ruby-red and transparent when cold. If, however, the bead is kept too long in the reducing flame, it remains colorless; but the ruby-red color may be produced by cautious oxidation. This reaction is very sensitive, and can also be used for the detection of tin. Heated with charcoal before the blowpipe (or better still with the charcoal stick), spongy metal is obtained. Copper salts color the flame blue or green. CADMIUM, Cd. At. Wt. 112.4. Occurrence.—Cadmium is usually associated with zinc in its ores. It is also found as greenockite, CdS, hexagonal; and as a regular crystallizing oxide, CdC).* * With calamine in the zinc deposits of Monte Poni. Sardinia. Ch. Ztg, 1901, p. 561. 176 REACTIONS OF THE METALS. Cadmium is a silver-white, ductile metal of sp. gr. 8.6. It melts at 320° C. and boils at 770° C. Heated in the air, it burns to brown cadmium oxide. The proper solvent for cadmium is nitric acid. Dilute hydro- chloric and sulphuric acids dissolve it but slowly, with evolution of hydrogen. Cadmium forms two oxides: Cadmium suboxide, Cd,0, and Cadmium oxide, CdC). Black Brown to black Cadmium suboxide (whose existence is doubted) is formed with cadmium oxide in small amounts when the metal is burned in the air. It is also said to be formed. Similar to lead suboxide. by gently heating cadmium oxalate away from air. There are no cadmium salts derived from this oxide, Cadmium forms only one series of salts, in which cadmium is bivalent. Cadmium salts are mostly colorless, though the sulphide is yel- low to orange. Most of the salts are insoluble in water, but readily soluble in mineral acids. The chloride, nitrate, and Sulphate are soluble in water. REACTIONS IN THE WET WAY. 1, Potassium Hydroxide precipitates white, amorphous cad- mium hydroxide, insoluble in an excess of the reagent (difference from zinc and lead): CdCl4-2KOH=2KCI+Cd(OH),. On gently igniting the hydroxide the brown oxide is obtained, which becomes darker on stronger ignition. The ignition of cad- mium nitrate yields the black crystalline oxide. 2. Ammonia also precipitates the white hydroxide, soluble in excess (difference from lead), forming complex cadmium ammo- nium compounds, as with zinc, nickel, etc. On diluting with water and boiling, cadmium hydroxide is reprecipitated from the solution of the cadmium ammonium compound. 3. Alkali and Ammonium Carbonates precipitate the white basic carbonate insoluble in excess. CADMIUM. 177 4. Potassium Cyanide precipitates white, amorphous cadmium cyanide, % Cd(NO.),4-2KCN=2KNO,--Cd(CN), readily soluble in excess: Gd(CN),4-2KCN=Cd(CN),K. From the solution of cadmium potassium cyanide the above- mentioned reagents produce no precipitation; but hydrogen sul- phide decomposes it readily, forming a precipitate of yellow cad- mium sulphide (difference from copper): Cd(CN),K,4-H,S=2KCN+2HCN+CdS. 5. Hydrogen Sulphide produces precipitates varying in color from a canary-yellow orange to almost brown, according to the conditions. In neutral solution, whether hot or cold, light-yellow cadmium sulphide is obtained in a condition hard to filter. From acid solutions (containing in 100 c.c. from 2 to 10 c.c. of conc. H2SO4, or from 2 to 5 c.c. of conc. HCl) yellow precipitates which turn orange in color are at once thrown down and are easy to filter. The latter precipitates are not pure CdS, but always contain more or less CdZOlºs or Cd5(SO4)S. For this reason cadmium should not be determined quantitatively as the sulphide. Cadmium sulphide is insoluble in alkaline sulphides (difference from arsenic), but is soluble in considerable hydrochloric acid, warm dilute nitric acid, and hot dilute sulphuric acid (difference from copper). 6. Ammonium Sulphide produces in ammoniacal Solutions yellow colloidal cadmium sulphide, which has a tendency to pass through the filter. The presence of a concentrated Salt solution prevents its doing this (cf. p. 141). 178 REACTIONS OF THE METALS. IREACTIONS IN THE DRY WAY. Cadmium compounds, heated on charcoal with soda, give a brown incrustation of cadmium oxide. If a compound of cadmium oxide is reduced in the upper re- ducing flame of the Bunsen burner, the cadmium oxide is changed to metal, which volatilizes, and, in the upper oxidizing flame, goes back to oxide, which will be deposited as a brown coating if a glazed porcelain dish filled with water is held just above the flame. This oxide always contains some suboxide mixed with it, and has the property of reducing silver oxide to metal; so that if the coating of oxide is moistened with silver nitrate solution a black deposit of metallic silver will be obtained: Cd,0+2AgNO,-Cd(NO.), H–CdC)+Ag. This reaction is very sensitive. If it is desired to test the precipitate produced by hydrogen sul- phide for cadmium in this way, the sample is first roasted in the oxidizing flame and then treated as just described. SEPARATION OF MERCURY, LEAD, COPPER, BISMUTH, AND CADMIUM FROM THE PREVIOUS GROUPS AND FROM ONE ANOTHER. In order to separate these metals from the previous groups, the solution is acidified with nitric acid + (for 100 c.c. of solution, 10 to 15 c.c. of double normal acid should be used) and saturated with hydrogen sulphide in the cold; then diluted with an equal volume of water, f again saturated with hydrogen sulphide, filtered and washed with water containing hydrogen sulphide. * Nitric acid is used only when lead is present, which can always be de- termined by the preliminary examination; in all other cases it is far better to acidify with hydrochloric or sulphuric acid (cf. p. 7). The reason so much acid is used is to prevent the precipitation of zinc. In the presence of much copper and little zinc, hydrogen sulphide precipitates all of the zinc as sulphide if the solution is only slightly acid with mineral acids; if, however, enough acid is present, none of the zinc will be precipitated. After all the copper has been precipitated, the zinc will not be precipitated by dilution with water and treatment with more hydrogen sulphide. f The above mentioned dilution with water is in all cases necessary, as cadmium will otherwise (often) not be precipitated, and will come down on adding ammonium sulphide with the members of Group III. If the solu- tion is diluted as above, all the cadmium will be precipitated. § TABLE W. SEPARATION OF MERCURY, LEAD, COPPER, BISMUTH, AND CADMIUM. The sulphides obtained by hydrogen sulphide (HgS, PbS, CuS, Bi,Ss, and CdS) are treated in a porcelain dish with nitric acid (sp. gr. 1.2) and warmed; all of the sulphides dissolve except HgS. The residue is filtered off and washed with water. Solution [Pb(NO3)2, Cu(NO3)2, Bi(NO3)s, Cd(NO3)2]. Evaporate to a small volume, add a few cubic centimeters of dilute sulphuric acid,” evaporate until fumes of sulphuric acid come off, cool, dilute with a little water,i and filter. SoLUTION [CuSO4, Bi,(SOA)s, CdSO,. The solution is saturated with ammonia and filtered. RESIDUE (BiSO,OH) SoLUTION [Cu(NH3),SO, Cd(NH3), SOJ. The , white precipitate shows bismuth if the anal- ysis has been done cor- rectly. It is confirmed (always) by washing the Fº dissolving in a ittle HCl, and adding an alkaline solution of potas- sium stannite. A black precipitated metallic bis- muth shows bismuth. —s A blue solution shows copper. RCN is added until the solution is decolor- ized, and hydrogen sulphide is passed into the solution; a yellow precipitate shows cadmium. Cadmium is confirmed by the flame reaction. A little of the filtered and washed precipitate is placed on an asbestos thread, roasted carefully in the upper oxidizing flame, and the oxide in- º produced on a porcelain dish (see p. 178). On moistening the brown coating with silver nitrate it becomes bluish-black if cadmium is present. * If considerable lead is present, the addition of dilute sulphuric acid produces a precipitate immediately. If only little lead is present, the RESIDUE (HgS). Dissolve in aqua regia, evaporate al- most to dryness, add water and filter off any sulphur. Add st a n n o us chloride to the fil- trate; a white precip- itate of Hg2Cl2, which, on further addition of SnCl2, becomes gray, shows mercury. RESIDUE (PbSO). The formation of a white residue of PbSO, show s lead. It is con- firmed by the charcoal – stic k test; a malleable button, soluble in nitric acid, shows lead. Too much water should not be added, otherwise basic b precipiº appears only after evaporation and dilution (see p. 162) # A white and brown precipitate of Fe(OH)8] and Al(OH)8] is often obtained completely washed. ismuth sulphate may be precipitated. . ere, if the precipitate produced by hydrogen sulphide was in- 18o REACTIONS OF THE METALS. The precipitate thus obtained contains mercury, lead, copper, bismuth, and cadmium as sulphides, and may be analyzed accord- ing to Table V, page 179. The filtrate from the hydrogen sulphide precipitate contains the members of the previous groups. The following metals are also precipitated from acid solutions by hydrogen Sulphide, but their sulphides are sulphoanhydrides, and are, therefore, soluble in alkali sulphides, forming sulpho salts. To this group belong arsenic, antimony, tin (gold, platinum, tungsten, molybdenum, vanadium, Selenium, and tellurium). ARSENIC, As. At. Wt. 75.0. Occurrence.—Arsenic is widely distributed in nature, being found in Small amounts in almost all sulphides, as, for example, sphalerite and pyrites; therefore almost all the zinc and sulphuric acid of commerce contain arsenic. Arsenic occurs native in kidney-shaped masses; also in the form of its oxide, As, O, as isometric arsenolite and orthorhombic claudetite, it being dimorphous. Mimetite, (AsO4)2Pb:Cl, hexagonal, isomorphous with apatite, pyromorphite, and vanadinite, is a well-known mineral containing arsenic oxide. The most important sources of arsenic are the sulphides, ar- Senides, and Sulpho salts: realgar, AS,S, monoclinic; orpiment, AS.S, monoclinic; arsenopyrite, Fe(ASS), orthorhombic; niccolite, NiAs, hexagonal; 16llingite, FeAs, Orthorhombic; smaltite, (Co,Ni,Fe)As, isometric; and proustite, As(SAg), rhombohedral. Metallic arsenic is a steel-gray, brittle substance, with a sp. gr. of 5.73. On being heated it sublimes, giving off a characteristic garlic odor. The merest trace of arsenic may be recognized by this odor. The molecule of arsenic contains, like phosphorus, four atoms, (ASI). Arsenic is insoluble in hydrochloric acid, but readily soluble in nitric acid and in aqua regia. Dilute nitric acid dissolves arsenic, forming arsenious acid: As, -2HNO,--2H,0=2AS(OH),4-2NO. Concentrated nitric acid and aqua regia dissolve it, forming arsenic acid: 3As, 10HNO,--4H,0=6AsO(OH), -10NO. Arsenic belongs to the same natural group of elements as nitro- gen and phosphorus, and forms, as they do, two oxides. ARSENIC. ſº 181 Arsenic trioxide and Arsenic pentoxide –O A <3 No o/”YO sº º which almost always act as acid anhydrides. A. Arsenious Compounds. Arsenic trioxide is formed by the combustion of arsenic in the air as white, glistening crystals of regular octahedrons. If the vapors of the trioxide are allowed to cool slowly, they solidify to an amorphous glass (arsenic glass), which gradually becomes crystal- line (white and opaque, like porcelain). Arsenic trioxide is known in three different modifications: 1. Isometric arsenic trioxide, white arsenic; 2. Monoclinic arsenic trioxide, and 3. Amorphous, glassy arsenic trioxide. The monoclinic modification is difficultly soluble in water (80 parts of cold water dissolve one part of AS,0s); while the amor- phous, glassy modification is much more soluble (25 parts of cold water dissolve one part of arsenic trioxide). By treatment of the ordinary modification (white arsenic) with water, it is not wet by the latter, appearing like meal, which is very characteristic. The trioxide dissolves quite readily in hydrochloric acid, par- ticularly on warming, from which Solution it often separates out, on cooling, in a beautiful, crystalline, anhydrous condition. Acting as an acid anhydride it dissolves readily in alkalies, forming easily soluble arsenites (salts of arsenious acid): As O,--6NH,OH=3H,0+2AsONH,), ; * 29; *The tri-metal arsenites derived from the ortho acid As3 OH are usually * /QAg OH unstable. Only the silver salt *ś is well known. The alkali arsenites g are really derived from OH OH *::::H As ÓH As of: eta- arsenious acid Xo Xo –OH As—OH and others As of: O Pyroarsenious acid As—OH Xo As—OH NOH 182 REACTIONS OF THE METALS. As, O,--6KOH=3H,0+2AS(OK), ; As, O,--3Na,CO,-3CO,--2AsONa), Free arsenious acid, As(OH), has never been isolated; as a very weak acid it breaks down, like carbonic acid, into water and the anhydride. Arsenic combines with chlorine directly, like phosphorus, form- ing the chloride, AsCls, which behaves exactly like the chloride of arsenious acid, similar to PCls. It is a colorless liquid, boiling at 134°C., and is decomposed quantitatively, like all acid chlorides, with water: /g HQ; /Q; AS3Cl–H HOH z-2 3HCl–H AS3–OH C| HOH NOH /9}, HQ; /9; Pé-Cl-- HOH z-> 3HCI+ P& OH Cl HOH OH The aqueous solution of arsenic trichloride, and the solution of the trioxide in dilute hydrochloric acid, contain the arsenic as arsenious acid. As the concentration of the hydrochloric acid increases, the amount of arsenic trichloride increases, until in very concentrated hydrochloric acid the arsenic is present almost entirely as trichlo- ride. By boiling a solution of arsenic trichloride in hydrochloric acid, arsenic trichloride is given off as a gas. If hydrochloric acid is conducted into the Solution at the same time (so that the con- centration of the hydrochloric acid is kept as large as possible), all the arsenic can be volatilized from the solution as arsenious chloride. On evaporating a hydrochloric acid solution of arsenious acid, arsenious chloride constantly escapes, so that all the arsenic may be volatilized. If, however, the arsenic is present in the form of With sodium and potassium, only the two following types are known: Asºk and AsO.K.H. As tº-3 (ONH.), With ammonium, only the ammonium pyroarsenite, Xo , is known. As == (ON H.) 2 The salts, Asºk and As E8N a, in alkaline solution behave as salts of the ortho acid: As(OK); and As(ONa)s. ARSENIC, 183 arsenic acid, no arsenic is lost during the evaporation of the solu- tion. * REACTIONS OF ARSENIOUS ACID IN THE WET WAY. The arsenites of the alkalies are soluble in water; the remaining arsenites are insoluble in water, but soluble in acids. 1. Hydrogen Sulphide precipitates from acid solutions yellow, flocculent arsenic trisulphide: 2AsOH),4-3H,S=6H,0+As,S,, 2ASCI, +3H,S=6HCl·HAs,S,. Arsenious sulphide is insoluble in acids; even quite concentrated boiling hydrochloric acid * (1:1) does not dissolve it. Concen- trated nitric acid oxidizes it to arsenic acid and sulphuric acid: 3As;S,4-28HNO,4-4H,0=9H,SO,--28NO+6H,AsO. The sulphide is more soluble in ammoniacal hydrogen per- oxide: f As,S,--14H,0,--12NH,OH=20H,0+3(NH,),SO,4-2(NH,),AsO. It is also dissolved by alkalies, ammonium carbonate, and alkali sulphides: As S,--6KOH=3H,0+Asé–OK+ As SK NOK' NSK As.13NH2.co-sco, Aº§: Azº: S = Sº- S- ***-**śrass}} /SN H, As,S,--3(NH,),S=2AS$ SNH, SNH, Just as the anhydride, As, O, can be referred to the acid, As(OH), so the sulphoanhydride, As.S., can be referred to the sulphoarse- nious acid, As(SH), which is not capable of existence in the free state, but is known in the form of its salts. If one of the latter salts * By long heating with concentrated hydrochloric acid, it is slowly changed to volatile AsCls and H.S. I84 REACTIONS OF THE METALS. is acidified, then sulphoarsenious acid is set free; but it immediately loses H.S, forming the insoluble sulphoanhydride: (a) 2AsOSNH),4-6HCl=6NH,Cl4-2AsOSH), ; (b) 2AS(SH)a=3H,S-i-As,Sa. On treating a mixture of Sulpho- and oxyarsenites with acid, arsenic trisulphide is also precipitated: As(ONa). tº- As (śNaj. +6HCl=6NaCl-i-3H,0+As,S,.” This property of forming Sulpho-salts accounts for the fact that hydrogen Sulphide produces no precipitation from normal arsenites, and only a partial precipitation, as AS,Ss, from mono- and dimet- allic salts: As(OK)a-i-3H,SześH,0+ As(SK)s; 6AsO.·HK,-- 15H.S=18H,0+As,S,--4As(SK)s; 6AsO.H.K+12H.S=18H,0+2As,S,--2AsOSK). • * Consequently, in order to precipitate arsenic completely as tri- sulphide, it is always necessary that the solution should contain enough free acid to prevent the formation of soluble sulpho-salts. 2. Silver Nitrate produces in neutral solutions of arsenites a yellow precipitate of silver Orthoarsenite (difference from arsenic acid): As(OK)s-H 3AgNO3=3KNOs--As(OAg)a, soluble in nitric acid and ammonia: As(OAg)s--3HNO. =3AgNO3+ As(OH)s; As(OAg)a-H 3NHa = As(ONHaAg)s. The normal (ortho) arsenites are unknown in the solid state; only the meta- and pyroarsenites have been isolated (cf. page 181, foot-note), which behave in aqueous solutions as acid salts of the ortho-acid. Consequently, silver nitrate produces only an incom- plete preciptiation in aqueous Solutions of the mono- and dimetallic nSalts: OH 3A ºr 3AgNO3=3KNO,--2AsOOH), + As(OAg)s. OK Tº The precipitation is quantitative only when the solution is dilute; from a concentrated solution H.S escapes, so that more H2S must be conducted into the solution in order to precipitate all the arsenic. ARSENIC. 185 In order to make the precipitation quantitative, an alkali (pref- erably ammonia) must be added. As, however, the solution al- ready reacts alkaline, it is difficult to reach the exact neutral point. Usually too much alkali is added. It is, therefore, more practicable (if it is desired to make the precipitation quantitative) to add the ammonia to the silver nitrate until the precipitate of silver oxide just dissolves; but as a rule this is unnecessary, as it is only needful to observe the color of the precipitate in order to determine whether an arsenite or arsenate is present. If, however, the arsenite solu- tion is made weakly acid with nitric acid,” and the silver nitrate solution (which has been treated with ammonia) is then added, the arsenite will be, practically, completely precipitated, even if an excess of the reagent is employed: /Q#, AgNHS2, As –OH-i-AgNHANO,-3NH, NO,-i-As(OAg),. NOH AgNHNO. In case the solution to be tested contains also a chloride, it should be acidified with nitric acid and the chloride precipitated as silver chloride by an excess of silver nitrate, and filtered off. To the filtrate, dilute ammonia should be cautiously added. At the neutral zone formed by the ammonia above the acid solution, a yellow precipitate of silver arsenite will appear. This reaction is very sensitive. 3. Magnesium Ammonium Chloride produces no precipitation in dilute arsenite solutions in the presence of ammonia (difference. from arsenic acid). 4. Iodine Solution is decolorized by arsenious acid, the latter being oxidized to arsenic acid; the reaction does not take place. quantitatively in acid solution: OH A&# NaHCO,--I,-NaI+KI+CO,--AsO(OH),. OK * As otherwise the solution will be alkaline, dissolving a part of the silver- arsenite: OH Aé; +3AgNHSNOs--H,0=2NH, NO, +KNOs--4NH.OH+As(OAg). A. I 86 REACTIONS OF THE METALS. The Solution must be neutralized with sodium bicarbonate and not with normal Sodium hydroxide, because the latter also decolor- izes iodine: - 6NaOH+31, =5NaI+NaIO,--3H,0. 5. Stannous Chloride (Bettendorff's Test).-On adding to con- centrated hydrochloric acid a few drops of an arsenite solution, and then } c.c. of a Saturated Solution of stannous chloride in hydro- chloric acid, the solution quickly becomes brown and then black, owing to the deposition of metallic arsenic. The reaction takes place more readily on warming, but a dilute aqueous solution will not give the reaction. In concentrated hydrochloric acid, however, the arsenic is all present as trichloride, and this is reduced by the stannous chloride, while arsenious acid is not: 2ASC1,4-3SnCl,-3SnCl,--As, . B. Compounds of Arsenic Pentoxide. Arsenic pentoxide, which may be obtained by heating arsenic acid, is a white, fusible substance, and is changed by strong ignition into arsenic trioxide: AS,0s=AS,0s-HO,. Arsenic pentoxide is quite soluble in water, forming arsenic acid; AsO.·3H,0–2Asíš’”. Arsenic acid itself may be obtained in the solid state in the form of orthorhombic prisms corresponding to the formula 2AsO(OH),4-H,0. At 100° C. water escapes, orthoarsenic acid, AsO(OH)a, being left behind as a crystalline powder. By gentle ignition more water is given off, forming pyroarsenic acid, As,C), H, which on further ignition is changed to metarsenic acid, AsO.H. In this respect arsenic acid acts exactly like phos- phoric acid. Both the pyro- and the meta-acids readily take on water, and are changed back to the Ortho acid. , The salts of arsenic acid are called arseniates. ARSENIC. 187 As with orthophosphoric acid, mono-, di-, and trimetallic salts are known: O O O ^Na ^Na nd ^Na Nº Nºa' a. Nº. OH OH ONa The arseniates of the alkalies are soluble in water; the others are insoluble in water but easily soluble in acids. REACTIONS IN THE WET WAY. 1. Hydrogen Sulphide.—If hydrogen sulphide is conducted into a cold, fairly acid solution of arsenic acid, the liquid remains clear for a long time; it gradually becomes turbid, the arsenic acid being reduced to arsenious acid with deposition of sulphur, and the arse- nious acid is then quickly precipitated as trisulphide: 1. AsO.H.--H.S=H,0+S--H,AsO,. 2. 2H,AsO,4-3H,S=6H,O+As,S,. If hydrogen sulphide is passed into the same solution after it has been heated, the reduction takes place more rapidly and the arse- nious sulphide, therefore, is more readily formed. In the presence of considerable hydrochloric acid, and by passing hydrogen sulphide rapidly into the cold solution, all the arsenic is precipitated as arsenic pentasulphide: 2H,AsO,--5H,S=8H,O--As, Ss. If a fairly concentrated hydrochloric acid solution is treated while hot with hydrogen Sulphide, a mixture of tri- and penta- sulphides is obtained. This peculiar behavior may be explained by assuming that the solution in concentrated hydrochloric acid contains the arsenic as pentachloride,” which is not reduced readily by hydrogen sulphide, so that arsenic pentasulphide' is precipitated; while on addition of water the pentachloride is changed to arsenic acid, which is re- * Arsenic pentachloride has never been isolated; but it is possible that it can exist only in hydrochloric acid solution. The reaction would then be I88 REACTIONS OF THE METALS. duced by hydrogen sulphide to arsenious acid before it is precipi- tated. & In order to precipitate the arsenic quickly by hydrogen sulphide from a solution of arsenic acid, or an arseniate, without employing considerable hydrochloric acid, it is only necessary to reduce the arsenic acid, in the first place, by boiling with sulphurous acid; boil off the excess of the latter, and then conduct hydrogen sulphide into the Solution, whereby an immediate precipitation of arsenious Sul- phide takes place. Arsenic pentasulphide, like the trisulphide, is insoluble in boiling concentrated hydrochloric acid, but readily soluble in alkalies, ammonium carbonate, and alkali Sulphides: /S /S As SS-H 6NaOH=3H,0+ASE(SNa),4-ASE(ONa),; /S /S As S.--3(NHA),CO,-3CO,--ASE(SNH),4-ASE(ONH,),; /S As S.--3(NH,),S=2AS=(SNH). By acidifying these solutions, arsenic pentasulphide is repre- cipitated: * 2ASS(SNH),4-6HCl=6NH,Cl4-3H,S+As,S.; AsS(SNH),4-AsS(ONH,),4-6HCl=6NH,Cl-i-3H,0+As,S. Arsenic pentasulphide is oxidized by fuming nitric acid to sulphuric and arsenic acids; also by solution in ammoniacal hydro- gen peroxide: As Sg-i-20H.O.--16NH,OH=28H,0+5(NH,),SO,--2(NH3),AsO. 2. Silver Nitrate precipitates from neutral solutions chocolate- brown silver arseniate (difference from arsenious and phosphoric acids); AsO(ONa),—H3AgNO,-3NaNO,-i-AsO(OAg), soluble in acids and in ammonia. | f ! \ ARSENIC. 189 3. Magnesium Chloride precipitates, in the presence of ammonia and ammonium chloride, a white, crystalline precipitate of magne- sium ammonium arseniate: O Zö AsO,Na,FI+MgCl, -NH,-2NaCl-As. 5-Mg. NöNH, This precipitate is insoluble in ammonia water, and is, therefore, used for the quantitative determination of arsenic. By ignition it is changed into magnesium pyroarseniate: 2AsO.MgNH,-H.O--2NH,--Mg,As,O,. 4. Ammonium Molybdate in considerable excess precipitates, in a boiling nitric acid solution, yellow crystalline ammonium arsenomolybdate: H,AsO,--12(NH,),MoC),4-21HNO,- =12H,0+21NH, NO,--(NH,),AsO,-12MoC),. Y— Arsenic acid, like phosphoric acid, forms with molybdic acid a tribasic complex salt, the arsenomolybdic acid, of the structure /(MoQ.-QH * O=As3(MoC),), OH. N(Mooj -óH This tribasic acid is soluble in nitric acid, but the ammonium and potassium salts are insoluble. The solution of ammonium molybdate for use as a reagent may be prepared as follows: 150 grams of commercial ammonium molybdate, [(NH,),Mo,O,4-i-4H,O], are dissolved in a liter of dis- tilled water and poured into a liter of nitric acid (sp. gr. 1.2). At first white molybdic acid will be precipitated. but this dissolves finally in the nitric acid, forming a clear solution. If arsenic acid is added to this solution, the soluble arseno- molybdic acid is first formed, which immediately reacts with the ammonium nitrate present (formed from the ammonium molyb- date and nitric acid), producing the insoluble ammonium salt of the arseno-molybdic acid. 190 / REACTIONS OF THE METALs. But this ammonium salt is. Soluble in a solution containing an alkali arseniate, forming complex acids containing more arsenic in the molecule, whose ammonium salts are soluble in nitric acid. Consequently a large excess of ammonium molybdate should be used if it is desired to precipitate arsenic acid.* The ammonium salt of arseno-molybdic acid is also readily Soluble in alkalies and in ammonia: (NH),AsO.· 12MoC),4-24NH,OH= =12H,0+12(NHA),MoQ.--(NHA),AsO, from which solution arsenic acid may be precipitated as white crystalline magnesium ammonium arseniate. As we shall see later, phosphoric acid behaves in exactly the same way towards magnesium Salts and ammonium molybdate. If, therefore, phosphoric acid and arsenic acid are both present, it is necessary to precipitate first the arsenic with hydrogen sulphide, filter, and oxidize the precipitated arsenic Sulphide to arsenic acid with fuming nitric acid. In such a solution a precipitate produced by means of ammonium molybdate or magnesium chloride must be caused by arsenic acid. In the same way a precipitate pro- duced in the filtrate from the hydrogen sulphide precipitate must be caused by phosphoric acid. C. Reactions which may be obtained with either Arsenious or Arsenic Compounds. 1. The Berzelius-Marsh Test for Arsenic.—All compounds containing arsenic may be reduced, in acid solution, by means of nascent hydrogen to arsine, AsFis: As O, +6H, -3H,0+2AsH,; . As,C), +8H, -5H,0+2AsH, As,S, +6H, -3H,S +2AsH. The Sulphides are reduced very slowly, but the oxides are reduced quickly even at ordinary temperatures. To produce nascent hydrogen, zinc and sulphuric acid are used. * The addition of a concentrated ammonium nitrate solution increases the sensitiveness of this reaction. ARSENIC. I91 This very poisonous arsine possesses a property which enables us to detect with certainty the merest trace of arsenic—as little as 0.0007 milligram As. By conducting the gas through a heated glass tube filled with hydrogen, it is decomposed into hydrogen and metallic arsenic; and the latter is deposited as a brownish-black mirror on the sides of the glass tube, just beyond the place where it was heated. This test is extremely sensitive, and must be made with great precaution, as almost all reagents, especially commercial zinc and sulphuric acid, are likely to contain traces of arsenic. In case these are used without previous testing, arsenic is likely to be found even although it may not have been present in the substance itself. The Marsh test is particularly useful for detecting the presence of very small amounts of arsenic which could not be found by any of the previously-mentioned reactions. In cases of poisoning, and for detecting the presence of arsenic in wall-papers, this test, or a modification of it, is always used; we will, therefore, discuss it in detail. Formation and Properties of Arsine. (a) Formation.—Arseniuretted hydrogen, or arsine, is produced, as above mentioned, by the reduction of compounds containing arsenic by means of nascent hydrogen. For developing the latter pure zinc and pure sulphuric acid should be used. If other metals and other acids are used (e.g., zinc and hydrochloric acid, iron and sulphuric acid), the arsenic compound will be reduced; but if iron is used, a part of the arsenic is changed to solid As, H, which remains in the flask and consequently escapes detection. If zinc and hydro- chloric acid are used, a high temperature is necessary in order to accomplish the reduction; while with zinc and sulphuric acid the reaction takes place readily at the ordinary temperature. Chemi- cally-pure zinc dissolves with difficulty in chemically-pure sul- phuric acid, so that it is customary to accelerate the reaction by adding a little piece of platinum (or a drop of hydrochlorplatinic acid), which makes a galvanic element, and the solution of the zinc takes place much more rapidly. This, however, is not permissible, for arsine produced in this way always contains some solid arseniu- retted hydrogen (Bernstein, 1870). I92 REACTIONS OF THE METALS. Arsenic, arsenious oxide, arsenic pentoxide, and arsenic trisul- phide are readily reduced in alkaline solution by Sodium amalgam, aluminium, or Devarda's alloy and caustic potash, forming arsine. The reduction takes place quickly, and the arsine may be detected by the Gutzeit reaction. The presence of organic matter in Solu- tion prevents the reaction, and must therefore be previously oxi- dized with nitric and sulphuric acids. Arsine is also obtained by dissolving many arsenides in hydro- chloric or Sulphuric acid: * Zn,As, i-6HCl=32ncl,--2AsH. The arsenides of iron are attacked by acids only with difficulty, except when an excess of iron is present, when, with the help of the nascent hydrogen, they are decomposed, forming Solid and gaseous arseniuretted hydrogen. Consequently iron sulphide containing arsenic, on treatment with acids, always yields hydrogen sulphide contaminated with arsine.* (b) Properties.—Arsine is a colorless, unpleasant-smelling, ex- tremely-poisonous gas, which, on being heated away from the air, is decomposed into arsenic and hydrogen: 2ASH,-As, H-3H, By heating in the air, it is oxidized to water and arsenic trioxide. Solid iodine changes it to arsenious iodide and hydriodic acid: AsH,--3I, + Asſa-i-3HI. This reaction takes place on conducting arsine over solid iodine. This property serves to free hydrogen sulphide from arsine, as hydro- gen Sulphide does not act upon Solid iodine, but only upon aqueous iodine solutions. Arsine is not attacked by hydrogen sulphide at ordinary temperatures, but at 230° C. Sulphide of arsenic and hydrogen are formed. Arsine is a strong reducing agent: silver salts are reduced to metal (see page 197). w * Certain microbes, namely, Penicillium brevicaule, when provided with nutriment containing only traces of arsenic, have the power of forming vola- tile arsenic compounds of a garlic odor, and this may be used as an extremely Sensitive test for arsenic. (Chem. Centralbl., 1902, I. p. 1245.) ARSENIC, I93 Directions for Performing the Berzelius-Marsh Test. The apparatus devised by G. Lockemann,” shown in Fig. 11, may be used to advantage. In the flask K, of 100 to 150 c.c. capacity, three or four grams of zinc f alloyed with copper are placed with about 20 c.c. of sul- phuric acid free from arsenic (1 vol. conc. H2SO4 diluted with 8 vols. of water). A steady stream of hydrogen is at once evolved, and in twenty minutes the air will be entirely driven out of the apparatus. When the gas escaping at b is found to be pure (by collecting a little in a small tube and lighting it) the hydrogen is lighted at b. The flame will be about two or three millimeters high and should remain so during the whole of the experiment; if it becomes higher the solution in K is cooled by placing the flask in cold water, and, conversely, when the flame is too low a little sulphuric acid is added, or the flask is placed in warm water. First of all the reagents, zinc and sulphuric acid, are tested to see that they are free from arsenic. The hard-glass tube is heated at B just before the restriction in the tube, which is 5 mm. long and 1.5–2 mm. wide. If at the end of twenty minutes there is no arsenic mirror formed in this capillary, the reagents are free from arsenic. * Melt 20 grams of zinc (free from arsenic) with a trace of pure copper and pour the fusion into cold water. f Zeit. f. angew. Ch., 1905, pp. 427 and 491. I94 REACTIONS OF THE METALS. The sulphuric acid solution to be tested for arsenic, and which must be free from Organic substances, sulphides, chlorides, nitrates, or other oxidizing agent, is then placed in the graduated funnel T, and is added little by little to the flask K without in any way inter- rupting the current of hydrogen. Just before adding the solution to the flask, the two burners at A are lighted and the glass tube thereby heated to dull redness. The gas as it escapes from the flask K passes through the drying-tube C containing crystallized calcium chloride, and then passes into the tube A, where any arsine is quantitatively decomposed into arsenic and hydrogen, The arsenic is deposited on the cold walls of the capillary. The end of the capillary is cooled in order to form a sharply defined mirror by winding around it a piece of rubber tubing, as shown in Fig. 11, and allowing water to drop upon it from the dish W during the experiment. All of the arsenic, unless the amount present was unusually large, will be deposited at the end of 25 minutes, and by comparing the mirror with a series of standards the amount can be estimated accurately (see page 197). Remark.-If the tube A is not heated at all, but the gas ignited at b as above described, the arsenic may be deposited upon a cold porcelain dish by holding the dish in the flame. The deposit is readily soluble in sodium hypochlorite solu- tion (difference from anti- mony). In this form the test was used by James Marsh in 1836. v. Confirmatory Test.—In the small glass tube open at both ends (see Fig. 12) is found the arsenic mirror. The tube is FIG. 12. held in an inclined position and heated by means of a small flame, whereby the arsenic is changed to arsenic trioxide, giving off the characteristic garlic odor, which can be detected at the upper end of the tube if only rºw of a milligram of arsenic trioxide is formed. After the tube is cooled, the arsenic trioxide is to be found at a in the form of small glistening octahedrons, which can be seen with the magnifying-glass or often with the naked eye. .* ARSENIC. I95. These three facts—formation of the mirror, the garlic odor, and the octahedrons—suffice to make us sure of the presence of ar- senic; but the more proofs we have, the more certain we are of the accuracy of the result. If the octahedrons have been recognized, the capillary end of the tube is sealed with a flame, and 1–2 drops of pure concentrated hydrochloric acid are introduced into the tube. with the help of a dropper, and the tube is moved so that the arsenic trioxide is moistened by the acid; 6–10 drops of distilled. water are added and hydrogen sulphide is conducted into the tube, whereby yellow arsenious sulphide is formed. The hydrogen sulphide in this case is generated from a solution of sodium sulphide by adding a little Sulphuric acid, as illustrated. in Fig. 13. The upper part of the test-tube contains a wad of cotton wool, which prevents any of the solution in the tube from being mechanically carried over into the tube containing the arsenic. After this confirmatory test has been performed, the tube is washed, first with water, then with ammonia, and finally with water again; it is carefully dried by drawing warm air through the tube; and weighed after it has cooled. U As an example of the practical application of this delicate test, we will describe the method to be employed in the detection of ar- Senic in wall-papers, etc. Many wall-papers contain small amounts of arsenic, and it is desired to know how much is present in a definite area of the paper, say a Square meter. The amount of arsenic contained in wall-papers is usually so small that weighing the mirror produced would not be accurate. It is best, therefore, to prepare a number of mirrors from FIG. 13 I96 REACTIONS OF THE METALS. known amounts of arsenic, to establish a scale for determining how much is contained in the given wall-paper or fabric.” We must first extract the arsenic without loss from the paper. To do this, a square decimeter of the paper is cut into small pieces, placed in a glazed porcelain dish, and 1–5 c.c. of concentrated sul- phuric acid, to which about ºff of its volume of concentrated nitric acid has been added, is poured upon it. The paper is thoroughly soaked with acid (being stirred into it by a thick glass rod), so that the paper or cloth finally absorbs the whole amount. The dish is then heated over a small flame until its contents become charred and crumbly. Sulphuric acid fumes are given off thickly by this time. After cooling, about 6 c.c. of water are added, the charred mass is thoroughly stirred, the liquid is boiled to expel any sulphurous acid which may have been formed, and filtered hot by the aid of suction. A weighed tube of 25–30 c.c. capacity, having a side arm attach- ment, can be used to advantage for receiving the filtrate. The residue is washed on the filter (by successive applications of small amounts of hot water) until the glass is nearly full, when it is again weighed. By deducting the weight of the empty vessel, the weight of the liquid is obtained. During this operation, the Marsh appara- tus should be made ready for the test. A few drops of the well-mixed solution are now added through T to the reduction flask B. If no mirror appears within three or four minutes, #–% of the filtrate is added little by little; and if no mirror appears after five minutes, the whole filtrate. The whole fil- trate is not added at once, because if too strong a mirror is obtained, it is much more difficult to estimate the amount of arsenic present. After twenty-five minutes all the arsenic will be deposited if not more than 0.05 milligrams of arsenic are in solution. If a mirror of sufficient amount was obtained in fifteen minutes from only a frac- tion of the whole solution, no more should be added, but the opera- tion should be continued for ten minutes more, the flame extin- guished, and the tube allowed to cool while hydrogen continues to pass through it. The mirror is then compared with the scale, and the remaining part of the filtrate is weighed in order to determine how much was used for the test. If sufficient material is at hand, a duplicate experiment should * C R. Sanger, Amer. Acad. of Arts and Sciences, XXVI, p. 24. ARSENIC. 197 be made with a new tube and a new sample. The results of a few such determinations are given in the following table: Cm.” of weight of wº. of Extract weight of Tºº! Weight | Mg. As, O, Paper used. Extract. taken. Mirror. *::::::: II]. per m.” xtract. 100 31.63 31.63 0 0 0. 100 30. 11 10.23 0.015 0.044 4.4 9.87 0.013 0.0399 3.99 100 28.72 8.32 0.045 0.155 15.5 7.53 0.042 0.163 16.3 50 30.22 2.64 0.015 0.172 34.4 3.22 0.020 0.187 37.4 The comparison of the mirrors is best made in transmitted light. The normal mirrors are prepared as follows: 1 gram of pure sublimed arsenic trioxide is dissolved in a little sodium carbonate solution, acidified with dilute sulphuric acid, and diluted to a liter. Ten c.c. of this solution, of which 1 c.c. contains 1 milligram of As, O, are again diluted to a liter, so that a solution is now obtained of which 1 c.c. contains 0.01 mg. of AS,0,. One c.c., 2 c.c., 3 c.c., 4 c.c., and 5 c.c. are taken from the solution and introduced separately into the Marsh apparatus, and the corresponding mirrors are obtained in different tubes. Two tubes are made from each amount of arse- nic, as the mirrors are not always the same. These mirrors may be kept in the dark for some time; but on exposure to the light they fade perceptibly. Mirrors which are sealed up with hydrogen do not keep so well. 2. The Gutzeit Test for Arsenic depends upon the behavior of arseniuretted hydrogen towards a concentrated Solution of silver nitrate (1 : 1) (according to Eidenbenz, a crystal of solid silver nitrate should be used). The silver nitrate is at first colored yellow and then black, the following reaction taking place: 1. 6AgNO,-i-AsFis-AsAga-3AgNO,--3HNO,. Yellow. —” 2. AsAgs-3AgNO,--3HOH=As(OH)2+3HNO,--6Ag. 198 REACTIONS OF THE METALS. The test is carried out as follows: A small amount of the sub- AgNO3 stance is put in a small test-tube, Fig. 14, a ſ\Paper few grains of zinc and a little dilute sulphuric acid are added, and a wad of cotton is placed ğ|Cotton near the top of the tube as a filter. Over the mouth of the tube a piece of filter-paper is placed, with a crystal of silver nitrate on top. If arsenic is present, the silver nitrate is at first turned yellow, but it becomes black very quickly. This reaction is often used for quickly testing commercial acid for arsenic, but it is not so reliable as the Bettendorff test (page 186), because phosphine * and stibine give a similar reaction with silver nitrate, while they are not reduced by stannous chloride. If arsine is allowed to act upon a dilute solution of silver nitrate, the yellow compound ASAga-3AgNO, is not formed, for it is imme- diately decomposed hydrolytically, according to equation (2), page 197, the reaction going quantitatively if the nitric acid formed is neutralized by ammonia. -, If the precipitated silver is filtered off, and ammonia then poured on top of the filtrate, the neutral Zone will appear yellow owing to the formation of silver arsenite. 3. The Reinsch Test is very easy to make, but not so sensitive as those just mentioned. *. If a strip of polished copper foil is added to a solution of arse- nious acid, the copper is colored gray owing to the deposition of the arsenic on the copper, forming copper arsenide of the formula AS,Cus. From concentrated solutions the arsenic separates out in the cold, but from dilute solutions only on warming. If considerable arsenic is present, the gray copper arsenide drops off from the copper. Antimony is also precipitated on copper from its solutions, so that the deposit must be tested for arsenic in the dry way. Arsenic acid is also reduced by copper, but only on warming. * Commercial zinc often contains a small quantity of phosphorus. ANTIMONY. I99 The Reinsch test is often used in testing wall-papers for arsenic. The pieces of paper are treated with a little hydrochloric acid (1:2), a piece of copper foil added, and warmed. A gray deposit on the copper indicates the presence of arsenic. REACTIONS IN THE DRY WAY. Metallic arsenic burns, giving off a garlic odor. Mixed with sodium carbonate and heated on charcoal, all arsenic compounds give this odor. Oxygen compounds of arsenic are easily reduced to metal in the upper reducing flame. On holding a porcelain dish (glazed on the outside), filled with water, directly over the sample, the metallic vapors are condensed on the dish, forming a brownish-black coating which is soluble in sodium hypochlorite solution, disappearing in- stantly, the arsenic being oxidized to arsenic acid: As, -5NaOCl--3H,0=5NaCl--2H,ASO,. If the porcelain dish is not held closely above the reducing flame, but above the upper oxidizing flame, the arsenic vapors are burned with a bluish flame to white arsenious oxide which deposits on the dish. If this deposit is moistened with silver nitrate, and ammonia vapors blown upon it, a yellow coloration due to Aga ASO, is formed, which disappears if more ammonia is allowed to act upon it (differ- ence from antimony): As O,--6AgNO,--3H,0=2Ag,AsO,--6HNO,. The ammonia serves to neutralize the nitric acid formed by the reaction, but the precipitate dissolves in excess of ammonia as well as in nitric acid. ANTIMONY, Sb. At. Wt. 120.2. Occurrence.—Antimony seldom occurs free in nature, although large amounts of the metal have been found recently in Australia. The most important compounds containing antimony are (as with arsenic) the sulphur compounds. Stibnite, Sb,S, orthorhombic, is 2OO REACTIONS OF THE METALS. found in Japan in beautiful crystals. The occurrence of kermesite, SbFS * º is very interesting, as this compound is often met with in Sb2_S analysis. t Of the oxygen compounds the dimorphous antimony trioxide is known as isometric senarmontite and Orthorhombic valentinite. Antimony also occurs in many sulpho salts, of which the tribasic silver Sulphoantimonate, or pyrargyrite, Sb(SAg)s, may be men- tioned. Antimony is a silver-white, brittle metal, of sp. gr. 6.7–6.8. It melts at 630°C., and distils in a stream of hydrogen at about 1350° C. It burns readily in the air to antimony trioxide. The solvent for antimony is aqua regia, by which it is converted into chloride. Nitric acid attacks antimony, changing it into Sb,0s and Sb,0s, which dissolve slightly in concentrated acid, but are insoluble in dilute acid. Antimony forms three oxides: SbFO O=Shº-O * Xº, O and Sb,0, Sb2_O O=Sb2–O Antimony trioxide Antimony pentoxide Antimony tetroxide of which the latter may be regarded as antimonyl antimonate, —O Sb-0 —O(SbO) * timony trioxide as a rule shows basic properties, while antimony pentoxide has more the character of an acid anhydride. , and is a very indifferent substance chemically. An- A. Compounds of Antimony Trioxide. By burning the metal in the air, the trioxide is obtained, which on stronger ignition in the presence of air is changed to the inert Sb,O,. The trioxide is dissolved by concentrated hydrochloric acid, forming antimony trichloride, a compound which (like bismuth chloride) is readily changed into a basic salt by the action of water, the decomposition of which depends upon the “masses " of the ANTIMONY. 2O I reacting substances. Thus an oxychloride SB-3 is known, which is formed according to the following equation: I / g H =O sieg: #X) =2HCl+Sb 1. In the presence of a large amount of water some oxide is formed with the oxychloride: 2SbCl,--3H,0ze 6HCl·HSb,0,. A mixture of oxychloride and oxide is known as “algaroth” powder, Sb,0,.2SbOCl. By boiling with considerable water the oxide alone is obtained. Antimony trioxide forms three hydroxides, which behave as very weak acids: —OH stºº O , and the hypothetical Sb of OH Sb–OH Metantimonous acid: —OH Orthoantimonous acid Pyroantimonous acid - Salts of the metantimonous acid are known, although the free acid itself has never been isolated. On boiling the oxide Sb,O, with concentrated caustic soda or potash, it goes into solution, but on dilution with considerable hot water Sb,0, separates out again. On filtering this off, tetragonal crystals of sºn, are deposited on cooling; which are, however, very unstable, and are decom- posed by standing in the air into Sodium carbonate and antimony trioxide. By dissolving antimony trioxide in strong alkali, the orthoantimonate must be formed first: /QN 8, Sb,O,--6NaOH=2Sb6–ONa+3H,O, NóNa which is hydrolytically decomposed on dilution into metantimo– nite and alkali hydroxide: =O Sb(ONa),4-H,0–2NaOH+Sb.I.N. 2O2 REACTIONS OF THE METALS. The latter is decomposed by more water into trioxide and alkali hydroxide; so that on adding to a solution of the trichloride either Sodium hydroxide or carbonate, an almost quantitative precipita- tion of Sb,O, will be obtained: 2SbCl, E6NaOH =6NaCl-i-3H,0+Sb,0,; 2SbCl, H3Na,CO,-6NaCl-3CO,--Sb,0,. Antimony oxychloride, Sbº, and Sodium metantimonite, Sbº a contain, as do many other compounds, the univalent grOup, Sb2O, which is known as the antimonyl group. Antimony Oxychloride, therefore, can be regarded as antimonyl chloride, and Sodium metantimonate as antimonyl oxide of sodium. Antimonyl nitrate, stºo , is also known, and antimonyl sulphate, 3 (SbO),SO,. All these compounds are easily hydrolyzed into acid and oxide, so that they are rarely met with in the course of analysis, with the exception of antimonyl chloride. The antimonyl compounds of certain organic acids (such as tartaric acid) are very much more stable. On boiling antimony trioxide with a solution of potassium acid tartrate, it goes readily into solution, forming the so-called “tartar emetic,” COOK took *. CHOH Sb-O CHOH 2 + O =H.O--2 2 gºon b=O gºon COOH ČOO(SbOjº Tartar emetic Antimonyi potassium tartrate which is comparatively soluble in water. 100 parts of water dissolve at 8.7°C. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.26 parts of salt 21°C.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 7.94 “ “ “ 31°C.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12.20 “ “ “ 50°C................................ 18.18 “ “ “ 75°C.. . . . . . . . . . . . . . . . . . . . . . . . . . . . ... 31.21 “ “ ANTIMONY. 2O3 Not only antimonyl oxide, Sb,O, but all antimonyl compounds are soluble in alkali tartrate solution, forming tartar emetic; thus, for example, antimonyl chloride readily dissolves in Rochelle salt, or tartaric acid: COOK COOK | | CHOH CHOH | +Cl(SbO)=NaCl-H CHOH CHOH | | COONa COO(SbO) C.H.O.K.H. SbOCl= HCl-H C.H.O.30 Tartar emetic * is the most important antimony compound of commerce. Consequently it will be worth while to say a few words with regard to its behavior toward acids. If an aqueous solution of potassium antimonyl tartrate is treated with hydrochloric acid, a white precipitate of antimonyl " chloride is formed, f * C.H.0.3 of 2HCl=KCl+H,C,EI,O,--SbOGl, which readily dissolves in more hydrochloric acid, SbOCl-H2HCl e H.O.--SbCl, but by the addition of more water is reprecipitated, etc. Sulphuric and nitric acids precipitate from a solution of potas- sium antimonyl tartrate, orthoantimonous acid, for the antimonyl compound, which is at first formed, is immediately decomposed by water, H , ~ =O (a) cho.3 of 2HNO,- KNO,-- C.H.O.3+ Sb No.; (b) shºot 2HOH = HNO,--Sb(OH), REACTIONS OF ANTIMONOUS COMPOUNDS IN THE WET WAY. 1. Water precipitates at first a basic salt which is changed into oxide by more water. 2. Sodium Hydroxide, Ammonia, and Alkali Carbonates pre- cipitate amorphous hydrated oxide. *Tartar emetic crystallizes with half a molecule of water: K(SbO)C, H, O,-- * H.O. 2O4. REACTIONS OF THE METALS. 3. Hydrogen Sulphide precipitates, from solutions which are not too acid, flocculent, Orange-red antimony trisulphide: 2SbCl, H3H,Sze 6HCl·HSb,S,. As is indicated in the equation, the antimony trisulphide is soluble in hydrochloric acid; in concentrated acid (1 : 1) it is readily soluble (difference from arsenic); so that in order to precipitate antimony by means of hydrogen Sulphide, the solution must be quite dilute. In case we have a solution of antimony chloride in hydrochloric acid, we do not attempt to dilute very much at first, for the oxychloride will precipitate out if we do; but we conduct hydrogen Sulphide into the Solution for some time, then dilute with water, and Saturate once more with hydrogen sulphide. If we do not proceed in this way, and attempt to precipitate from a solution which is too acid, the clear filtrate will become turbid as soon as it comes in contact with more water (from a moist beaker, perhaps). In such a case the solution must be largely diluted with water, then filtered again. e Antimony trisulphide is soluble in ammonium sulphide, forming a sulpho salt: Sb,S,--3(NH,),S=2Sb(SNH), * - If yellow ammonium sulphide is employed, the ammonium salt of sulphoantimonic acid is obtained: Sb,S,--2(NH,),S,-SbS(SNH),4-SbS,NH,. If the solution of ammonium sulphoantimonate is boiled for a long time in the air, the red-colored oxysulphide is often precipitated: 2Sb(SNH),4-8O=2(NH,),S,O,--2NH,--H,0+Sb,S,O. By boiling antimony chloride with sodium thiosulphate, the oxysulphide is also obtained, 2SbCl, -3Na,S,O,-6NaCl-H4SO,--Sb,S,O, which, on being warmed with ammonium sulphide, redissolves, forming the sulpho salt. * The triammonium salt has never been isolated, the mono salt, Sb-g NH 47 alone being known in the solid state. In solution, however, particularly in the presence of considerable ammonium sulphide, the ion SbS, must be pres- ent. ANTIMONY. 2O5 Antimony trisulphide is also soluble in caustic alkali, forming sulpho and oxysulpho salts:* —O - SbS,4-2KOH=H,0+Sb ski- Sb ŠK. These sulpho salts are decomposed by acids, precipitating anti- mony trisulphide, with evolution of hydrogen sulphide, 2Sb(SNH),4-6HCl=6NH,Cl-H2Sb(SH), and the unstable sulpho acid is further decomposed: 2Sb(SH), -3H, S-I-Sb,S,. The solution of the sulphide in caustic potash is also precipi- tated by the addition of acid: Sb ski-Sb-sk-2HCl=2KCl+H,0+SbS, 4. Zinc precipitates from solutions of antimony compounds metallic antimony. If a piece of platinum foil and a little zinc are placed in an antimony solution containing hydrochloric acid, so that the two metals touch one another, the antimony is deposited on the platinum in the form of a black stain which does not disappear on removal of the zinc (difference from tin). - 5. Potassium Iodide does not set free iodine when treated with an antimonous solution (difference from antimonic compounds). B. Antimonic Compounds. Antimony pentoxide, Sb,0s, is formed as a yellow powder by the oxidation of antimony by means of concentrated nitric acid and gently igniting the reaction product (antimonic acid). On strong ignition it loses oxygen and goes over into the very stable antimonyl antimonate Sb,0,. * It is frequently stated that a mixture of sulpho and oxysalts is formed by the solution of Sb,S, in caustic potash. This can hardly be right, because the antimonites of the alkalies are readily decomposed by water into caustic potash and hydrated oxide, the latter separating out. The above solution can be diluted largely without becoming turbid; so it can contain no anti- monite. g 2 off REACTIONS OF THE METALS. The pentoxide dissolves in concentrated hydrochloric acid, forming the pentachloride Sb,Os--10HCl=5H,0+2SbCls. If this solution is treated with water, a white precipitate of antimonic oxychloride, SbO,Cl, is formed, which, by the addition of more water, is changed on warming into antimonic acid: SbCls+2H,O =24HCl·HSbO,Cl, SbO,Cl+2HOH = HCl·HSbO,H, Tartaric acid prevents the precipitation of the oxychloride, as with SbOCl (p. 203). Antimony pentoxide is an acid anhydride, and, like the corresponding P.O.s, can be referred to three acids, o-s, š, 0–sºn, =sb_3# NOH Metantimonic acid Orthoantimonic acid O _ct –OH O=SB 5; Pyroantimonic acid which have all been isolated. The salts of the metantimonic and pyroantimonic acids are the most common. The trimetallic salts of the ortho acid are unknown, but the monometallic salts are known to exist. All antimonates, being salts of a weak acid, are very unstable, being easily hydrolyzed by water. If antimony pentoxide is fused with an excess of caustic potash, the product of the fusion must contain the trimetallic salt of ortho- antimonic acid. If, however, the melt is dissolved in a little water and allowed to crystallize, deliquescent crystals of potassium pyro- antimonate, K.Sb,O, are formed. The ortho salt, which is at first formed, is decomposed by water as follows: OK —OK o-stº O=Sh_5}: §4-HOH = 2KOH+ Xo º Nök O=Sb ðī: ANTIMONY. 2 o'7 By the action of considerable cold water (or more quickly by rapid boiling with less water) this deliquescent Salt is gradually changed into the acid salt, losing KOH, _cº-OK gº —OK =Sb Ök HOH O=Sb ðī; Xo + z-2 2KOH-i- Xo y —OK HOH —OH which is a granular powder, difficultly soluble in cold water.* It dissolves to a considerable extent in water at about 40°–50° C.; this salt is used in testing for Sodium, as the Sodium Salt is very much more insoluble in water. On boiling the granular potassium salt for a long time with con- siderable water, it gradually takes on water, forming the easily soluble monometallic Salt of Orthoantimonic acid, _ct –OK Xo Hoe (-6.) —OH NöH which is obtained, on evaporating the Solution, as a gummy mass of the composition 2KH.SbO4+ H.O, but on boiling the aqueous solution for a long time, more KOH is lost, with the formation of amorphous orthoantimonic acid: OK OH o=sbőfi-HOH = KOH4. ( ==<) NöH NöH All antimonates are decomposed by acids, amorphous antimonic acid separating out. The gummy, monometallic Salts give an amorphous precipita- tion with sodium salts, gradually becoming crystalline, while the potassium pyroantimonate gives a crystalline precipitate imme- diately. * The granular powder has the composition K.H.Sb,O′ + 6H,O. 208 REACTIONS OF THE METALS. This sensitive reagent for Sodium is prepared as follows: Commercial potassium antimonate (which usually contains con- siderable antimonite as impurity) is treated with concentrated nitric acid and boiled until no more red fumes are given off; the acid is decanted, and the undissolved, heavy, powdery antimonic acid is washed by decantation with water and then boiled several minutes with a quite concentrated solution of caustic potash. In this way a pure, granular acid potassium salt is obtained. The liquid is cooled, the alkali poured off, 5 or 6 c.c. of water added, quickly heated to boiling, cooled, and filtered. The filtrate thus obtained reacts immediately with a normal Solution of a sodium Salt. REACTIONS OF ANTIMONIC COMPOUNDS IN THE WET WAY. A solution of the purified potassium antimonate in hydrochloric acid may be used for these reactions. 1. Hydrogen Sulphide precipitates from fairly acid solutions the orange-red pentasulphide e 2SbCl,--5H.S=10HCl·HSb,Ss. Antimony pentasulphide is soluble in strong hydrochloric acid, forming antimony trichloride, with deposition of Sulphur and evo- lution of hydrogen Sulphide: Sb.S.--6HC1–3H.S+S,4-2SbCl, It also dissolves (like the trisulphide) in alkali Sulphides, and in alkalies, but not in ammonium carbonate. By treatment with an alkali sulphide, a sulpho salt is obtained, Sb,S,--3(NHA),S=2SbS(SNH), which is decomposed by the addition of acids, forming the insoluble pentasulphide with evolution of hydrogen Sulphide: 2SbS(NHI),4-6HCl=6NH,Cl-H3H,S+Sb,Ss. Alkalies dissolve the pentasulphide, forming Sulpho and oxy- sulpho Salts: g Sb,S,4-6KOH=SbS(SK),4-SbS(OK),4-3H,0. ANTIMONY. . 2 og 2. Hydriodic Acid reduces antimonic compounds in acid solu- tions, with separation of iodine (difference from antimonous com- pounds): SbCl, H2HI=2HCI+SbCl, L.I. 3. Nascent Hydrogen.—By treating any solution which con- tains antimony with nascent hydrogen, stibine is formed: Sb,O,--6H,-3H,0+2SbH, Sb,O,--8H,-5H,0+2SbH, If the stibine is generated in a Marsh apparatus (cf. p. 193), and the gas is conducted through a red-hot glass tube, a mirror of metallic antimony will be deposited, as with arsenic. But as stibine is much more unstable than arsine, and the antimony itself is much less volatile, the mirror is found nearer the heated place than is the case with arsenic—sometimes before the hottest part of the tube is reached—as the decomposition of the stibine takes place at a much lower temperature than with arsine. If the stibine is allowed to escape from the tube with the hydro- gen, it burns with a pale greenish-white flame to water and anti- mony trioxide. If a piece of glazed porcelain is held directly over the flame, a deposition of metallic antimony is obtained which is un- affected by a solution of sodium hypochlorite (difference from arsenic). If stibine is allowed to act upon a solution of silver nitrate,” a black precipitation of silver antimonide is thrown down: SbH,--3AgNO,-SbAga-i-3HNO,. REACTIONS OF ANTIMONY IN THE DRY WAY. Antimony compounds impart to the flame a pale greenish-white color. Heated with sodium carbonate on charcoal, a brittle metal- lic button is obtained, surrounded by a white incrustation. Compounds containing oxygen are reduced in the upper reducing flame to metal, which is volatile and burns in the upper oxidizing flame to trioxide; the latter can be deposited on a glazed porcelain surface. If the deposit is moistened with silver nitrate solution, * Solid silver nitrate is turned yellow at first, then black; exactly the same as by arsine (cf. p. 198). 2 IO REACTIONS OF THE METALS. and ammonia blown upon it, it becomes black, owing to the separa- tion of metallic silver: * Sb,O,--4AgNO,--4NH,--2H,0=4NH, NO,--Sb,Os--4Ag. TIN, Sn. At. Wt. 119.0. Occurrence.—Tin does not occur free in nature, but mostly in the form of the dioxide SnO, as tetragonal tinstone, or cassiterite, isomorphous with rutile (TiO), zircon (SiO,ZrO.), and polianite (MnO,). * Tin is a silver-white metal, of sp. gr. 7.29; it melts at 233° and boils at about 1500° C. At ordinary temperatures tin is malleable and ductile, but at very low temperatures and near the melting-point it is so brittle that it can be powdered. In order to pulverize tin, it should be . heated in a porcelain dish till it melts; the flame is then removed and the substance is quickly crushed with a pestle. It soon cools, , yielding a fine powder. Tin is soluble in hot concentrated hydrochloric acid with evolu- tion of hydrogen: Sn+2HCl–SnCl4-H, In the presence of platinum the Solution takes place more quickly and at a lower temperature. Dilute hydrochloric acid dissolves tin, but very slowly. Nitric acid, of sp. gr. 1.2 to 1.3, does not dissolve tin, but oxi- dizes it to metastannic acid: 3Sn+4HNO,4-H,0=3SnO(OH),4-4NO. Cold dilute nitric acid dissolves the metal very slowly, without any evolution of gas, forming ammonium and stannous nitrates. The reaction takes place in two phases, hydrogen and 'stannous nitrate being first formed: Sn+2HNO,-Sn(NO), H H. The nascent hydrogen then reduces the nitric acid to ammonia, HNO,--8H=3H,0+NH, TIN. 2 I It which combines with the excess of the acid, forming ammonium nitrate. The whole reaction may, therefore, be expressed as follows: 4Sn+10HNO,-4Sn(NO.), H-NH,NO,4-3H,0. Aqua regia dissolves tin, forming stannic chloride: 3Sn+4HNO,--12HCl=4NO+8H,0+3SnCl. Tin dissolves in dilute sulphuric acid very slowly, but readily in hot concentrated acid, forming stannic sulphate, with evolution of sulphur dioxide: Sn+4H,SO,-2SO,--4H.O.--Sn(SO.), Tin forms two oxides: —O SnO and Sn−5. Stannous oxide Stannic oxide Salts are known corresponding to both these oxides—stannous and stannic salts. The former contain bivalent tin and the latter quadrivalent tin. Stannous Compounds. Stannous oxide (according to the way it is prepared) is either an olive-green or a black powder, which, on being warmed in the air, like all stannous compounds readily changes to stannic oxide. By dissolving stannous oxide (or, better still, the metal itself) in hydro- chloric acid, stannous chloride is obtained, which is the most im- portant of all the stannous salts. This salt, with two molecules of water of crystallization, SnCl,--2H,0, is the so-called “tin salt” of COIOIOleI’Ce. Fresh crystals of “tin salt” will dissolve clear in a little water; if more water is added the solution becomes turbid, owing to the formation of a basic salt, /Cl —OH, Snºit HOHz-> HCl–H Sn of * which is readily soluble in hydrochloric acid. 2 I 2 REACTIONS OF THE METALS. The clear concentrated solution also becomes turbid on standing in the air, in consequence of the formation of the same basic Salt with loss of chlorine: —Cl Sn Tº Lo-HOH-2Sn-º-C, *–Cl The chlorine, however, is not set free, but unites with some of the unchanged Stannous chloride, forming stannic chloride: SnCl,--Cl,-SnCl. If tin tetrachloride is treated with metallic tin, the latter goes into solution and the former is reduced to stannous chloride: SnCl,--Sn=2SnCl, Consequently, in order to maintain a solution of stannous chlo- ride, hydrochloric acid should be added to prevent the formation of the basic salt, and metallic tin to keep the solution in the stan- nous condition. g” Such a solution constantly grows more concentrated, owing to the gradual solution of the tin. In order to keep a solution of stannous chloride at a definite concentration (only necessary for purposes of quantitative analysis) the hydrochloric acid solution is kept away from air in an atmosphere of carbon dioxide without the addition of metallic tin. Nearly all stannous compounds are colorless; the oxide (as already mentioned) is black and the sulphide dark brown. REACTIONS IN THE WET WAY. 1. Potassium and Sodium Hydroxides produce a white precipi- tate of gelatinous stannous hydroxide, SnCl,+2KOH=2KC1+Sn(OH), which is readily soluble in an excess of the precipitant, forming po- tassium stannite: * * According to Hantzsch theºlution contains: | /O Sn—OK. 2eitschr. f. anorg Ch. (1902)., p. 289. TIN. 2 I 3 The hydroxide is also readily soluble in hydrochloric acid. Stannous hydroxide, therefore, behaves partly as an acid and partly as a base (like the hydroxides of zinc, aluminium, and chromium). The alkaline solution of an alkali stannate is often changed brownish black or black (particularly on warming, or when very concentrated caustic alkali is used), owing to the separation of either metallic tin or stannous oxide: From dilute potassium hydroxide solutions there gradually sepa- rates on standing, or more rapidly on heating, the black monoxide, —OH te Sn Tâk–KOH +Sno; and from quite concentrated alkali the precipitate is almost wholly black metallic tin: 2Sn 3:42H,0-K Sn(OH)8]+Sn. 2. Ammonia and Alkali Carbonates precipitate the white hydroxide, which is not absolutely insoluble in an excess of the precipitant: SnCl2 +2NH4OH=2NH4Cl-HSn(OH)2, SnCl2 +Na2CO3 + H2O=2|NaCl-HCO2+Sn(OH)2. 3. Hydrogen Sulphide produces (in solutions which are not too acid) a brown precipitate of stannous sulphide, H2S +SnCl2 a.22HCl·HSnS, readily soluble in strong hydrochloric acid; therefore no stannous sulphide is precipitated if the solution is very acid. On diluting a strongly acid solution with water, however, stannous sulphide is completely precipitated on Saturating the solution with hydrogen sulphide gas. Stanfous Sulphide is insoluble in ammonia and ammonium car- bonate (difference from arsenic); also in colorless ammonium sul- phide (difference from arsenic and antimony); but is readily soluble in yellow ammonium sulphide, forming ammonium sulphostannate: S SnS+(NH),S,- sºn H. NSN.H. If the solution of ammonium sulphostannate is acidified with any acid, yellow stannic sulphide is precipitated: /$ tº- Sné–SNH,--2HCl=2NH,Cl-i-H.S +SnS,. NSN.H. 2I4 REACTIONS OF THE METALS. 4. Mercuric Chloride produces in solutions of stannous salts a white precipitate of mercurous chloride: /Cl g—Cl Hg H \ X: SnCl,+SnCl,-- tº Hg. Cl Hg—Cl But if the stannous chloride is present in excess, the mercurous chloride will be further reduced to gray mercury: Hg,Cl,+SnCl,+SnCl,--2Hg. 5. The Gold Test is much more sensitive. If to a solution of chloride of gold a Solution containing a trace of stannous chloride. is added, finely divided metallic gold will be precipitated, 2AuCl4-3SnCl,-3SnCl,--Au, which appears brown by transmitted light, and bluish-green by reflected light. 6. Metallic Zinc precipitates tin from both stannous and stannic solutions as a spongy mass, which adheres to the zinc: SnCl,--2Zn=2ZnCl,--Sm. The finely-divided, spongy metal is easily soluble in strong hydrochloric acid; the experiment must not, therefore, be made in strongly acid solutions, as the tin is loosened from the zinc by the violent evolution of hydrogen and is redissolved by the acid. The test is best made by adding a drop of the (not too acid) solution to a piece of platinum foil, and then placing a piece of bright zinc so that it comes in contact with both the solution and the platinum. The tin is precipitated partly on the zinc and partly on the plati- num,” in the form of a gray stain, which disappears from the latter as Soon as the zinc is removed, provided the solution is still acid (difference from antimony). If the zinc is kept in contact with the acid until the evolution of hydrogen ceases, the tin stain will not disappear from the platinum, because all the acid has been used up. On adding a few drops of concentrated hydrochloric acid to the * In weakly acid solutions tin is precipitated chiefly on the zinc.; in strongly acid solutions, chiefly on the platinum. TIN. 215 platinum, the stain quickly disappears with an evolution of hydro- gen. The reason why the tin is deposited on the platinum notwith- standing the presence of acid is that a galvanic current is formed by the contact of the zinc with the platinum, which flows from the zinc to the platinum; the platinum thus serves as a cathode, and the zinc is deposited upon it. On removing the zinc the current stops and the stain disappears. Stannic Compounds. The stannic compounds (which are all colorless, with the excep- tion of the yellow sulphide SnS,) cannot be obtained by the solution of the oxide SnO, from which they are derived, because the oxide is with difficulty attacked by acids. They are obtained indirectly from metallic tin or from stannous compounds. The simple stannic compounds are all, more or less readily, com- pletely hydrolyzed by water, so that the analyst almost never meets with them. The nitrate Sn(NO.), and the sulphate Sn(SO.), are quickly decomposed, in the cold, into acid and Stannic hydroxide. The halogen compounds are much more stable, and are decomposed only by boiling with considerable water. For the following reac- tions, therefore, we will assume that we have a solution of stannic chloride to work with. Stannic chloride is a colorless liquid, which fumes in the air and boils at 120° C. On adding a little water it solidifies, forming crystals of monoclinic hydrates, .* SnCl4-3H,0 * , SnCl,--5H,O, SnCl,--8H,O of which the salt with 5H,O is used commercially as a mordant in dyeing. *. On adding more water to these hydrates they dissolve, form- ing a clear solution, which on boiling (the freshly-prepared dilute solution) gradually becomes turbid, owing to the precipitation of voluminous stannic hydroxide: SnCl4+4HOH z=24HC1-i-Sn(OH)4. If the solution is very dilute it becomes turbid in the cold. The stannic acid thus formed is not precipitated quantitatively, either 2 I 6 REACTIONS OF THE METALS. in the cold or on boiling, because a considerable amount remains in the hydrosol form. By “salting out” the hot solution (best with ammonium nitrate), the stannic acid may be completely precipi- tated. g A solution of stannic chloride can be most readily obtained for analytical purposes by chlorinating or brominating a solution of stannous chloride. On adding chlorine to a solution of stannous chloride, stannic chloride is formed in the cold: SnCl,--Cl,-SnCl, As, however, chlorine is colorless in a dilute solution, it is diffi- cult to tell when the oxidation is complete; it is more easily ascer- tained if bromine is used. On adding strong bromine water (with constant stirring) to a solution of stannous chloride, the brown color will disappear as long as any stannous chloride remains unchanged, and the solution be- comes colored by the bromine only when the oxidation is complete. The solution then contains a mixture of Stannic chloride and stannic bromide: 2SnCl,--2Br, -SnCl,--SnBr. The excess of bromine can be removed by passing a current of air through the solution. It is not advisable to accomplish this by boiling, for a basic Salt may be formed. Just as platinum tetrachloride combines with hydrochloric acid to form hydrochlorplatinic acid, so tin tetrachloride unites with hydrochloric acid, forming hydrochlorstannic acid, PtCl, H., SnCl4II, Hydrochiorºſatſaic acid Hydrochlorºmſ. acid and yields, like the former, beautifully crystalline, easily soluble salts with the alkalies, of which the ammonium salt (NH),SnCl, is an article of commerce being known as “pink salt.” The above- mentioned stannic chloride, SnCl4, is sometimes designated as a-stannic chloride, to distinguish it from a compound (which we shall soon study) known as b-stannic chloride (stannyl chloride). TIN. 2 I 7 Reactions of the a-Stannic Compounds. 1. Hydrochloric and Sulphuric Acids produce in moderately concentrated Solutions of stannic chloride no precipitation, even on long standing (difference from 3-stannic compounds). In very dilute sulphuric acid solutions a precipitate of basic sulphate is sometimes obtained. In very dilute hydrochloric acid solutions, also, a slight turbidity is often formed, which becomes considerable on boiling the Solution: SnCl4-4HOH z-z 4HCl·HSn(OH,). 2. Potassium and Sodium Sulphates produce no precipitation in the cold (difference from stannyl chloride), but on boiling all the tin is precipitated as hydroxide. 3. Potassium or Sodium Hydroxide.—On adding caustic alkali to a solution of a stannic salt, a voluminous, gelatinous, white pre- cipitate is obtained: SnCl4 +4KOH=4KCl +Sn(OH)4. The precipitate has the above formula when dried in the air, and the formula SnO(OH)2 if dried over sulphuric acid. The precipitate dissolves readily in an excess of alkali hydroxide, forming salts which are not derived from either of the above com- pounds, but from H2.[Sn(OH)6], which has itself never been isolated: Sn(OH)4--2KOH =K2|Sn(OH)6]. The hydroxide dissolves in ammonia also, but only in the ab- sence of ammonium salts. By dissolving in alkali, stannic hydroxide behaves as an acid, and according to Bellucci and Parravano,” the hexa-oxystannic acid stands in the same relation to hydrochloric stannic acid as hexa-oxyplatinic acid to hydrochlorplatinic acid: H2PtCl6] H2(SnCl6] The salts of hexa-oxystannic acid are designated briefly as stan- nates, or cº-stannates, to distinguish them from the 3-stannates or metastannates, which are derived from the polymer (H2SnO3)5 (see below). * Zeitschr. f. anorg. Chem., 45 (1905), p. 156. 2 18 REACTIONS OF THE METALS. The ready solubility of a-stannic acid in cold dilute mineral acids is very characteristic. It dissolves promptly in hydrochloric, nitric, and sulphuric acids, behaving, in this respect, as a base. By boiling the dilute acid solution (particularly the sulphuric acid solution) stannic acid is reprecipitated, which is soluble in cold dilute acids provided the boiling has not been continued too long. In the latter case the 3-stannic acid is formed, which is insoluble in dilute acids. Potassium Carbonate precipitates stannic acid from stannic chloride solutions, which is completely soluble in an excess of the reagent (difference from 3-stannic acid): SnCl4+2K2CO3 + H2O=4KCl +2CO2+Sn(OH)4. 5. Sodium Carbonate behaves similarly, but the precipitate is not so easily Soluble in an excess. 6. Ammonia precipitates stannic acid from a solution of stannic chloride; tartaric acid prevents the precipitation (difference from 8-stannic acid). Reactions of the 9-Stannic Compounds. (Metastammates.) By the oxidation of metallic tin with hot nitric acid of sp. gr. 1.3 stannic nitrate is first formed, which, by boiling with water, is completely hydrolyzed, forming nitric acid and metastannic acid. Metastannic acid is a white powder insoluble in nitric acid, which when dried over sulphuric acid has the formula H2SnO3. This is of the same empirical composition as the hydroxide pre- cipitated by treating a stannic chloride solution with alkalies, though differing essentially from it in many reactions. While the a-stannic acid (as already mentioned) is easily soluble in dilute mineral acids, the b-stannic acid is almost insoluble therein. 1. If the 3-stannic acid is treated for a short time with concen- trated hydrochloric acid, a chloride is formed which is insoluble in hydrochloric acid, but readily soluble in water. The solution con- tains the so-called 3-stannic chloride (though the designation stan- nyl chloride would be more suitable) of the composition, Sn,0,0,00H),.” * R. Engel, Chem. Zeitg., 1897, pp. 309 and 859. TIN. 219 2. On treating the aqueous solution of stannyl chloride with hydrochloric acid, almost all the tin is reprecipitated in the form of a highly chlorinated compound of the composition: Sn;O.Cl,(OH)2+4H,0.* 3. If an aqueous solution of stannyl chloride is heated to boiling, almost all the tin is precipitated as b-stannic acid, which is insoluble in dilute acids. This differing behavior of the two acids, as well as of the two chlorides, can be explained as follows: Silicic acid, which is closely related to stannic acid, exists in innumerable silicates in different polymeric forms; as, for example, in the cases of the minerals of the pyroxene and amphibole groups, Wollastonite, CaSiO, and Tremolite, CaMg,Si,Ola, is a derivative of ordinary is a derivative of the poly- metasilicic acid, meric silicic acid, #3 Si-O-ST3:#f /O | | -—OH O O y Si ºf Ho–4 a 1–OH H Si-O-Si ºff and it is highly probable that the stannic acid can exist in analogous polymers. One of these polymers seems to possess the composi- tion [SnO,H,]s, to which we can ascribe the following structural formula: HO—Sn—OH /N O o OH HO—s. A gºº. ñó—sº *—ofi O O Ho— –OH #ö Sn—O—Sn 5; If such a compound is treated with hydrochloric acid, the hy- droxyl groups will, first, be replaced by chlorine, and a compound * Weber, J., 1869, 244 and P. A. 122,358. + Groth. Tabellarische Uebersicht d. Min. 1898, p. 148. 22 O REACTIONS OF THE METALS. will be obtained containing tin, oxyggn, and chlorine, e.g., Sn;O.Clio. This hypothetical compound of fe b-stannic acidis hydrolyzed, forming different chlorides of varying solubilities. Thus R. Engel found that the chloride Sn;OCl,(OH)s is soluble in water; and Weber showed that from an aqueous solution of the latter, hydro- chloric acid precipitates the compound Sn;O.Cl,(OH)--4H,O. The reaction which takes place on dissolving the 3-stannic acid in hydrochloric acid and water may be satisfactorily expressed by the following equations: Sn;O,(OH)10+10HCl=10H,0+Sn;O.Clio (insoluble in HCl); B-stannic acid OT Metastannic acid Sn;O.Clo-F8H,Oz-> 8HC1-i-Sn;O.Cl,(OH), (soluble in water); Sn;O.C.,(OH)8–1-2HCl zº. 2B,0+Sn;O.Cl,(OH), (insoluble in HCl). On boiling the aqueous Solution, complete hydrolysis takes place: Sn;OCl,(OH)s+2HOH=2HCl·HSn;O,(OH)10. If the 3-stannic acid is treated for a long time with concentrated hydrochloric acid, the ring Sng|Os is finally broken down, so that the tin goes into solution in the form of ordinary o-stannic chloride: Sn;O,(OH)10+20HCl=15H,0+5SnCl. Further reactions of stannyl chloride (3-stannic chloride) are: 4. Sulphuric Acid precipitates from Solutions of stannyl chlo- ride white stannyl sulphate, which on being washed with water is completely changed to 3-stannic acid (difference from a-stannic chloride). 5. Potassium and Sodium Sulphates behave in the same way as sulphuric acid. 6. Potassium Hydroxide throws down in solutions of stannyl chloride a voluminous precipitate of 3-stannic acid, which does not dissolve in an excess of the concentrated precipitant, but forms a b-stannate easily soluble in water and dilute caustic potash: ` (a) Sn,0,0,00H),4-2KOH=2KCl-HSn,0,00H),0; 6-stannic acid (b) Sn,0s(OH)10+2KOH=2H,O--Sn;O,(OK),(OH)s. potassium b-stannate TIN. 22 I By long treatment of the potassium /3-stannate with concentrated caustic potash, it gradually goes into solution, forming a-potassium stannate. This change takes place more readily by fusing g-stannic acid with solid potassium hydroxide in a silver crucible. If a dilute solution of a mineral acid is added to the ſ-potassium stannate, a voluminous precipitate is formed, consisting partly of ſº-stannic acid (insoluble in an excess of mineral acids) and partly of a-stannic acid (readily soluble in an excess of the acid). The latter compound is formed when a very concentrated solution of caustic potash was used in forming the potassium salt. 7. Ammonia also precipitates fl-stannic acid, even in the pres- ence of tartaric acid (difference from a-stannic chloride). As we have seen, the a-compounds may be readily changed into /3-compounds and conversely. The dilute aqueous solutions of the a-compounds are gradually changed, at the Ordinary temperature, into fl-compounds, but more quickly on boiling ; thus stannic chloride changes to stannyl chloride: 5SnCl2 + 13H2O = 18HCl + Sn;0;Cl2(OH)s. The fl-compounds are changed into a-compounds by boiling with concentrated hydrochloric acid or concentrated caustic potash. 8. Hydrogen Sulphide precipitates (from not too acid solutions) yellow stannic sulphide from both the a- and the A-compounds:* SnCl4 + 2 H2Sze 4HCl + SnS2; Sn;05Cl2(OH)3 + 10H2Sz-22HCl + 13H2O + 5SnS2. Stannic sulphide is soluble in hydrochloric acid; hydrogen sul- phide will cause no precipitation, therefore, if the solution is very acid. If such a solution, Saturated with hydrogen sulphide, is largely diluted, the sulphide will precipitate out. The yellow sulphide is the sulpho-anhydride of the sulpho acid; it dissolves, therefore, in alkali sulphide, forming salts soluble in Water: * From 6-stannic solutions hydrogen sulphide produces a precipitate, but very slowly, the SnS, remaining largely in the hydrosol form. By the addition of Salts it is coagulated, and separates out in a flocculent form, usu- ally mixed with Á-stannic acid (cf. Zeitschr. f. anorg, Ch. XXVIII, p. 140). If the stannyl chloride solution is heated on the water-bath in a pressure flask, the tin is quickly precipitated as greenish-yellow sulphide. 222 REACTIONS OF THE METALS. S sus, (NH9s-sºnh, * NSN.H. Acids precipitate from such a solution the yellow stannic sulphide S - snênh,42HG-2NHCH-Hsi-sºº. NSN.H. 2° *—S The Sulphide is insoluble in ammonia and ammonium carbo- nate (difference from arsenic). By means of concentrated nitric acid it is easily oxidized to 3-stannic acid; or by roasting in the air it can be completely changed to tin dioxide. The sulphide obtained in the dryway, known as “mosaic gold,” is not attacked by nitric acid, and is also insoluble in alkalisulphides. It dissolves on treatment with aqua regia, forming stannic chlo- ride with separation of sulphur. It is most readily brought into solution by fusing with sodium carbonate and Sulphur (see below). 9. Mercuric Chloride produces no precipitation in solutions of stannic Salts. Tin dioxide as it occurs in nature, and the artificially produced oxide, after strong ignition, are both insoluble in all acids. They can be brought into Solution by the following methods: 1. Fusion with sodium carbonate and sulphur; 2. “ “ caustic potash or soda; 3. “ “ potassium cyanide; 4. Reduction of hydrogen at a high heat. (1) Fusion with Sodium Carbonate and Sulphur.—The dry sub- stance is placed in a small porcelain crucible, mixed with six times as much calcined Sodium carbonate and Sulphur (equal parts mixed together), covered, and heated over a small flame until the excess of sulphur has distilled off and burned. This operation requires about twenty minutes. The crucible is then allowed to cool, its contents are treated with warm water, and the solution filtered if necessary: S 2SnO,--2Na,CO,--9S=3SO,-- as ºr 2CO,. a, TIN. 223 If iron, lead, copper, or any other metal forming sulphides in- soluble in water and ammonium polysulphide are present, these remain undissolved as Sulphides, and are separated from the tin by filtration. (2) Fusion with Sodium Hydroacide.—The sodium hydroxide is melted in a silver crucible until all water has been driven off (the fusion becoming quiet) and allowed to cool somewhat; the finely- powdered, dry substance is added, and the mixture is again heated until the fusion is clear. After cooling, the product is dissolved in Water: SnO,--2NaOH=SnO(ONa),4-H,0. Stannic oxide is not completely attacked by fusion with sodium or potassium carbonate. (3) Fusion with Potassium Cyanide.—Some potassium cyanide is melted in a porcelain crucible, the powdered stannic oxide is added, and the mixture is fused until the separated tin has melted together: SnO,--2KCN=2KCNO+Sn. After cooling, the mass is treated with water, the tin filtered off, flattened into foil, and then dissolved in concentrated hydrochloric acid. (4) Reduction in a Stream of Hydrogen.—The substance is put in a platinum boat, placed in a glass tube open at both ends, and made of difficultly-fusible glass; hydrogen is conducted into the tube in the cold until the air has been driven out, when it is heated (at a faint-red heat) until no more water is given off: SnO,--4H=2H,O--Sn. The metal is then dissolved in hydrochloric acid. REACTIONS IN THE DRY WAY. Heated with soda (or, better still, potassium cyanide) on char- coal, usually only a small malleable button is obtained, which, on taking away the flame, is immediately covered with a white coating of oxide. This can be observed when the flame is allowed to play upon the fusion. If the product is crushed in an agate mortar, a small flake of metallic tin is obtained, which can be distinguished 224 REACTIONS OF THE METALS. from silver and lead by its insolubility in nitric acid, and by its solubility in hydrochloric acid. This reaction is particularly suited for the charcoal-stick test. The borax bead which has been colored pale blue by copper, becomes a transparent ruby red in the reduc- ſº ing flame if a trace of tin is added. This is a very sensitive reac- tion. For the separation of the sulpho-acids from the Sulpho-bases, See Table VI, page 226. Gold, Au. At. Wt. 1972. Occurrence.—Gold usually occurs native in quartz and in river sands; also as telluride of gold in Sylvanite, (AuſAg), Tes, and in nagyagite, (PbAu),(TeSSb), and is found in Small amounts in many pyrite and other sulphide ores. Metallic gold is of a yellow color, has a sp. gr. of 19.33, and melts at 1064°C. without being oxidized. It is the most ductile of all metals, and may be hammered into exceedingly thin leaves, which are transparent, with a bluish-green color. The proper solvent for gold is aqua regia, but it is also soluble in bromine and chlorine water, forming a trihalogen compound: 2Au-F2HNO,--6HCl=4H,0+2NO+2AuCl, Au--3Br-AuDr. Gold is not attacked by mineral acids. It forms two oxides: Aurous oxide, and Auric oxide, Au,O, Au,Os. Both of these are exceedingly unstable; on gentle ignition they lose oxygen and are changed to metal (a property common to all “noble’’ metals). All gold salts are unstable; even the most stable salt of all, the chloride, AuCls, is changed by gentle ignition into yellowish-white aurous chloride, AuCl: AuCl,-AuCl4-Cl, On stronger ignition the last atom of chlorine is lost, and the yellow metal itself is left behind. Aurous chloride, AuCl, is insoluble in water, but on being boiled GOLD. 225 with water for some time, or very gradually in the cold, it is changed to auric chloride, with deposition of metal: 3AuCl–AuCla-i-Aug. The solution obtained by dissolving gold in aqua regia always contains auric chloride, so that only the reactions of auric com- pounds are of interest to the analytical chemist. Auric chloride unites with hydrochloric acid forming hydrochlorauric acid, AuCl3+HCl=AuCl4EI, which yields beautifully crystalline salts. Gold chloride is soluble in ether and can be extracted from its aqueous solutions by means of this solvent. Auric salts are mostly yellow and readily soluble in water. The sulphide is black and soluble only in aqua regia. REACTIONS IN THE WET WAY. A solution of auric chloride in hydrochloric acid should be used for these reactions. 1. Potassium or Sodium Hydroxide.—If caustic alkali is cau- tiously added to a concentrated gold Solution, a reddish-brown, voluminous precipitate of auric hydroxide is obtained, which looks exactly like ferric hydroxide. If more caustic alkali is added, how- ever, the gold hydroxide redissolves, forming alkali aurate: (a) AuCl,--3KOH=3KCI-H Au(OH), ; (b) AuſoH), KOH-2H,0+Auºk. If the bright-yellow solution of potassium aurate is carefully acidified with nitric acid, a precipitate of reddish-brown auric acid is thrown down, which is soluble in nitric acid, but is reprecipitated, for the most part, by dilution and boiling. As a rule, potassium hydroxide yields no precipitate in solutions of gold chloride, because the gold solution is usually so dilute that the amount of alkali added is sufficient to form the aurate at once. 2. Ammonia throws down a precipitate of dirty-yellow, fulmi- nating gold, [AuCl,]H+ 6NH,--3H,0=4NH,Cl-H [Auº OH,” 2 which explodes in a dry condition on warming or by concussion. * The formula of fulminating gold is usually expressed as AuNH,NH, -i-3H,O. All the metals of the hydrogen sulphide group are assumed to be present in the form of freshly-precipitated sulphides. The washed precipitate is placed in a porcelain dish, treated with yellow ammonium sulphide,” warmed gently for a short time (with stirring), and then filtered. SoLUTION [AsS,(NH4)3, SbS,(NH4)3, SnS3(NH3)2]. The solution is diluted with considerable water, and hydrochloric acid added until It is then boiled, the precipitated sulphides allowed to settle, the supernatent liquid decanted off, and the residue filtered and washed. The arsenic, antimony, and tin sulphides are then analyzed by one of the two follow- 1. The three sulphides (contaminated with considerable sulphur) are boiled with con- centrated hydrochloric acid (1:1) [until the vapors cease to blacken lead acetate paper] . TABLE VI. SEPARATION OF THE SULPHO ACIDS FROM THE SULPHO BASES, AND FROM ONE ANOTHER. RESIDUE may contain HgS, PbS, Bi,S, CuS, CdS, and should be examined accord- ing to Table V, page 179. the solution reacts acid with litmus. ing methods: and then filtered. RESIDUE [As2S3 + S]. This is treated with fuming nitric acid in a covered beaker, boiled until completely dissolved and brown fumes cease to be given off, then evaporated to a small volume, an excess of ammonia and some magnesium chloride or mag- nesium sulphate solution added and the mixture vigorously stirred. If consider- able arsenic is present, a precipitate of white crystalline MgNH4AsO4 + Aq will be formed. If only a small amount of arsenic is present, the precipitate will form only on standing. If no precipitate is formed after twelve hours, no arsenic is present. SoLUTION [SbOls + SnCl4]. The solution is concentrated to a small volume, a few drops are placed on some platinum foil, and a piece of bright zinc is put in the liquid, so that it comes in contact with both the solution and the platinum. After a few seconds the zinc is remoyed, and the platinum is examined to see whether a black stain has formed upon it which is insoluble in hy- drochloric acid. Such a stain shows antimony. The zinc is again added to the solution and al- lowed to remain until the evolution of hydrogen al- most entirely ceases, when the solution is carefully washed off with distilled water, keeping the zinc and platinum in contact. The zinc is then removed as completely as possible (any adhering tin being scraped off by the foil), dissolved in a few drops of concentrated hydrochloric acid, placed in a small test-tube, and a drop of mercuric chloride added. A white precipitate, which finally turns gray, shows tin to be present. § § 2. The mixture of arsenic, antimony, and tin sulphides (contaminated with sulphur) is heated with a concentrated solution of ammonium carbonate and filtered: RESIDUE [Sb2S3, SnS, S]. The sulphides are dissolved in * hydrochloric acid and analyzed as above. SOLUTION [AsS,(NHA)s + AsSO3(NH3)3]. The solution is acidified with hydrochloric acid, which causes yellow arsenic sulphide to be precipi- tated, showing the presence of arsenic. In order to confirm this, the sulphide is dissolved in fuming nitric acid (as above) and the precipitate of mag- nesium ammonium arsenate is produced. The transformation of the arsenic sulphide into arsenic acid can be better accomplished by an am- moniacal solution of hydrogen peroxide: AsS,(NHA); + AsSO3(NH); + 20H.O., -i- 10NH,OH == 2AsO.(NHA)s + 5(NH4)2SO, + 25H,O. From this solution magnesium chloride precipi- tates Mg|NH,AsO,. The magnesium precipitate may be tested in the dry way for arsenic by permitting the white incrus- tation to form on a porcelain dish (see p. 199) and testing this coating with silver nitrate and am- II].OD18. A yellow deposit shows arsenic to be present. * Yellow ammonium sulphide is used, because SnS is insoluble in colorless ammonium sulphide. 228 REACTIONS OF THE METALS. The most important reactions for the detection of gold are those which depend upon the extreme readiness with which the auric compounds are reduced. Auric compounds are strong oa'idizing agents. 3. Ferrous Salts precipitate at ordinary temperatures from neutral or acid solutions all the gold as a brown powder (difference from platinum): AuCl4-3FeSO,-Fe,(SO.), H-FeCl,-- Au. 4. Oxalic Acid precipitates all of the gold in the cold, but more quickly on warming (difference from platinum): COOH 2AuCl4-3 || =6HCl·H6CO,-- Au,. COOH The presence of considerable hydrochloric acid prevents this re- action. 5. Arsine and Stibine precipitate gold completely: 2AuCl,--Asha-i-3H,0=As(OH),4-6HCl·H Au,; 2AuCl,--SbH, -SbCl, H3HCl-i-Au,. 6. Sulphurous Acid reduces gold solutions: 2AuCl4-3H,SO,--3H,0=3H,SO,--6HCl- Au,. 7. Stannous chloride-causes the following reaction to take place: 2AuCl3+3SnCl2=3SnCl4 +Aug.” If the solution tested is strongly acid with hydrochloric acid, the precipitate is pure gold and has the characteristic dark-brown color of the finely-divided metal. In very dilute weakly acid solu- tions the so-called Purple of Cassius is thrown down which con- sists of colloidal gold and tin hydroxide. Purple of Cassius is soluble in ammonia and in dilute caustic potash solution, forming reddish liquids. These solutions when cold remain clear for a long time and can be boiled without decom- position. As the Solution is concentrated a flocculent precipitate is formed which will dissolve on the addition of more ammonia. The brown coloration can be distinctly seen if 3 mg. gold are dis- solved in 100 c.c. of the solution. * Theodor Doring, Chem. Centralbl., 1900, I, p. 735. GOLD. 229 8. Hydrogen Peroxide * in alkaline solution immediately pre- cipitates the gold as finely-divided metal: 2AuCl,--3H,O,--6KOH=6|KCl-i-6H,0+ 3O,--2Au. The precipitated metal appears brownish-black by reflected light, but bluish-green by transmitted light; +30 mg. gold in 10 c. c. of liquid suffice to give a reddish coloration with a bluish shimmer. 9. Zinc.—The following gold test is very sensitive.† A few drops of a dilute gold Solution are taken; a drop of arsenic acid, two to three drops of ferric chloride, and two to three drops of hydro- chloric acid are added; the mixture is then diluted to 100 c.c., and a piece of zinc is dropped in. Around the zinc the solution assumes a purple color, which, by moving the zinc in the solution, is dis- seminated through it, making it appear pink or purple. If the solution contains fºr mg. of gold, within fifteen minutes beautiful reddish color will be noticed. Besides the above reagents many others, such as formaldehyde in the presence of alkali, hydrazine sulphate, etc., are capable of reducing gold from its solutions. 10. Hydrogen Sulphide precipitates, in the cold, black gold di- sulphide from gold solutions: 8AuCl,--9H.S.--4H,O=24HCl- H,SO,--4Au,S,. Gold disulphide is insoluble in acids; but is readily soluble in aqua regia, forming auric chloride. The disulphide is difficultly soluble in ammonium sulphide, but more readily soluble in potassium sulphide, forming a sulpho salt: Aus. HR.S.-2Au-š. From this solution hydrochloric acid precipitates a yellowish- brown sulphide: 2Auiº+2HC-2KCI+H.S.-Aus. (). sºms *sº From a hot solution hydrogen sulphide precipitates brown, metallic gold: 8AuCl4-3H, SH-12H,0=24HCl·H 3H,SO,--8Au. * Vanino and Seemann, B. B. 1899, p. 1968. # Pharm. Chem. Centralbl., 27, p. 321. 23O REACTIONS OF THE METALS. The metallic gold is soluble in hot potassium or sodium poly- Sulphide, forming a Sulpho salt: 2Au:+K.S.-2Au ş". On account of its softness, gold is always alloyed with silver and copper when used for coins or for jewelry. If such an alloy is treated with nitric acid, the copper and silver are dissolved out, the gold remaining, usually as a brownish powder. The gold is filtered off through a small filter; the filter is dried, rolled together and enveloped in platinum wire; the paper is lighted and allowed to burn quietly. [The ash must not be too strongly heated, for the gold would then melt and alloy with the platinum wire.] The ash is melted with a little Soda on a charcoal stick, a gold button form- ing with the characteristic yellow color. The gold button can be pressed into a leaf in the agate mortar, transferred to a watch-glass, and dissolved in a little aqua regia, forming auric chloride. The solution is carefully evaporated to dryness, the residue dissolved in a little water, and a dilute Solution of stannous chloride added, when the presence of gold is shown by the formation of the purple of Cassius. The hydrogen peroxide or zinc tests are still more delicate (see p. 229). If it is a question of detecting very small amounts of gold (as in the case of many copper coins), the above method is unsuitable. In such a case the gold and silver are extracted by melting them together with lead, and the lead is then separated by oxidation and removal of the lead oxide. The process is as follows: five to ten gm. of the auriferous copper (or more in some cases) are treated with 120 gm. of pure lead in a flat dish of infusible stone (a scorifying- dish), and melted in a muffle, with access of air. The copper and a part of the lead are oxidized, and the oxide unites with the silica of the dish to form a readily fusible slag, which eventually covers the unaffected lead and the dissolved silver and gold. This operation is known as scorification. When this point is reached, the molten mass is poured into an iron scorification pan, previously well chalked. As soon as the mass becomes cool, it is removed from the pan, and the slag is hammered off from the enclosed lead button, which is weighed. The button is then placed in a cupel (a sort of crucible made of bone ash), of about the same weight as the lead button or GOLD. 23 I a little heavier. The cupel is then placed in the muffle and again heated, with ready access of air. The lead melts and is oxidized; the resulting lead oxide melts at 980° C. and is absorbed by the porous cupel,” while a kernel of silver and gold remains. The metallic kernel is flattened to foil and treated with nitric acid, which dissolves the silver, leaving the gold, usually in the form of powder. The gold is filtered off, dried, and melted, as above de- scribed, upon the charcoal stick. If the alloy of gold and silver (ob- tained after cupellation) contains three parts of silver to one part of gold, the gold remains after separation with nitric acid as a thin- as-paper, coherent, brownish mass (which becomes hard on ignit- ing), with the characteristic gold color. If the proportion of silver to gold is greater than 3:1, the separation by means of nitric acid will be complete and the gold will be left as a powder. If the ratio of silver to gold is less than 3:1, the separation by means of nitric acid is incomplete, and the gold residue usually appears yellow, and still contains some silver. In this case, more silver is added, and one gm. of lead; the mixture is once more subjected to cupellation, when the subsequent separation by means of nitric acid will be complete. In order to detect very small amounts of gold in ores, a similar procedure is used. If one does not possess a muffle furnace, the tiresome wet process must be used. For example, if it is desired to detect the presence of gold in pyrites, a large amount of the ore is roasted in the air until all the sulphur has been burned off, then treated with bromine water, and allowed to stand twelve hours. The solution (which must now contain all of the gold as auric bromide) is filtered and the excess of bromine boiled off. Ferrous sulphate and a little sulphuric acid are now added, and the solution is again boiled and filtered through a small filter. The residue on the filter is washed and dried and then melted on the charcoal stick. According to the methods described, above a fractional part of a milligram of gold can be detected with certainty. REACTIONS IN THE DRY WAY. All compounds of gold, when heated with soda on the charcoal stick, yield a malleable, metallic button, Soluble only in aqua regia. * Considerable litharge is also volatilized. 232 REACTIONS OF THE METALS. The solution in the latter reagent should be evaporated, the residue dissolved in water and tested with stannous chloride, hydrogen peroxide, or zinc. PLATINUM, Pt. At. Wt. 194.9. Occurrence.—Platinum is found free in nature, usually accom- panied by the other so-called platinum metals. Metallic platinum is grayish white, of sp. gr. 21.48; in a finely- divided state it is grayish-black. It melts at about 1780. Platinum is not attacked by mineral acids; it dissolves only in aqua regia, forming hydrochlorplatinic acid, H.PtCls (not platinum chloride, PtCl). If, however, the platinum is alloyed with silver, it dissolves in nitric acid to a yellow liquid, provided sufficient silver is present. Like tin, platinum forms two oxides: Platinum monoxide, and Platinum dioxide, Pt O, PtC,. Both oxides may be obtained by the careful ignition of the cor- responding hydroxides. They are exceedingly unstable, being de- composed by gentle ignition into metal and oxygen; all the re- maining platinum compounds behave similarly. The most important of the platinum compounds are the chlo- rides. By dissolving platinum in aqua regia, hydrochlorplatinic acid is always obtained, from which the di- and tetrachlorides may be derived, which form with hydrochloric acid the complex acids PtCl,--2HCl=H.PtCls (Hydrochlorplatinic acid—orange red crys- tals), and PtCl,--2HCl=H.PtCl, (Hydrochlorplatinous acid—only known in Solution). The aqueous solution of hydrochlorplatinic acid is yellowish- orange ; on the other hand, a solution of hydrochlorplatinous acid, containing the same amount of platinum as in the other case, is dark brown. •. { º The potassium and ammonium salts of hydrochlorplatinous acid are soluble in water; the corresponding salts of hydrochlor- PLATINUM. 233. platinic acid are difficultly soluble in water and insoluble in 75 per cent. alcohol. REACTIONS IN THE WET WAY. A solution of hydrochlorplatinic acid should be used for these reactions. 1. Ammonium and Potassium Chlorides produce in concen- trated solutions a yellow precipitate: H.PtCl4-2KCl=2HCl·H K.PtCl, (cf. p. 37); H.PtCl4-2NH,Cl=2HCl·H (NHA), PtCl, (cf. p. 45.). Both salts are difficultly soluble in water, but practically in- soluble in 75 per cent. alcohol and in concentrated Solutions of potassium and ammonium chlorides. This last property is taken advantage of in order to separate platinum from gold and other metals. 2. Alkali Iodides give a brownish-red coloration when added to a solution of hydrochlorp’atinic acid: H2PtCl6 --8FI =6|KCl +K2Pt.I6+2HI. 3. Hydrogen Sulphide precipitates in the cold very slowly, but quickly on warming, dark-brown platinum disulphide: H.PtCl4-2H,S=6HCl·H Pts,. Platinum sulphide is insoluble in mineral acids, but readily soluble in aqua regia. It is difficultly soluble in alkali sulphides, but more readily soluble in alkali polysulphides, forming a sulpho salt, which is decomposable by acids, with precipitation of platinum sulphide. 4. Ferrous Salts do not reduce hydrochlorplatinic acid in the presence of acids (difference from gold), but cause precipitation of all the platinum with ferric hydroxide (on warming) in a solution which has been neutralized with sodium carbonate: H.PtCl4-3Na,CO,--4Fe(OH),4-H,O= =3CO,--6NaCl-i-4Fe(OH),4-Pt. * 5. Oxalic Acid does not precipitate platinum (difference from gold). 234 REACTIONS OF THE METALs. 6. Formic Acid precipitates from neutral, boiling solutions all the platinum in the form of a black powder: H | H.PtCl4-2C=O =6HCl·H2CO,--Pt. NOH Formic acid An acid solution must be neutralized with soda before making this test. 7. Stannous Chloride reduces hydrochlorplatinic acid to hydrochlorplatinous acid only, not to metal: H.PtCl4-SnCl,-SnCl4-H, PtCl. 8. Glycerine and Sodium Hydroxide reduce hydrochlorplatinic acid on warming, with the Separation of the black, powdery metal; gºod CHOH +3Na,PtCl4-16NaOH= tion Glycerine COONa =18NaCl-i- + Na,CO,-- 12H,0+3Pt. COONa. Sodium oxalate 9. Zinc reduces hydrochlorplatinic acid to metal: H.PtCl4-3Zn=3ZnCl,+H,--Pt. Preparation of Hydrochlorplatinic Acid for Use as a Reagent. Since hydrochlorplatinic acid is used not only for the qualitative detection of ammonium and potassium, but also for their quantita- tive separation, we will describe the methods for preparing a solu- tion of this reagent. There are two cases to be considered, depending upon whether we start with metallic platinum or with platinum residues (pre- cipitates of K.PtCls, etc.). 1. Preparation of Hydrochlorplatinic Acid from Metallic Plati- ‘mum.—Almost all commercial platinum contains iridium; and although pure iridium is practically insoluble in aqua regia, it dis- PLATINUM. 235 solves considerably in this reagent if it is alloyed with platinum. Furthermore, platinum forms with aqua regia not only hydrochlor- platinic acid, but also hydrochlorplatinous acid (the most harmful of all impurities for this reagent) and nitroso-platinic chloride, (PtCl2)(NO), These facts must be borne in mind in preparing this reagent. First of all, the platinum filings are purified by boiling with con- centrated hydrochloric acid and washing with water. The platinum is then placed in a capacious flask, concentrated hydrochloric acid is poured over it, and nitric acid is added little by little, with con- tinuous, gentle heating on the water-bath. All the platinum and Some iridium are thus brought into solution, while Small amounts of the latter metal usually remain undissolved as a black powder. The solution is poured off (without stopping to filter) into a porcelain evaporating-dish and evaporated to syrup consistency. It is then dissolved in water, Sodium formate and Sodium carbonate are added until the solution is slightly alkaline; it is then heated to boiling, whereby the platinum and iridium are precipitated in a few minutes as a black powder. This operation should be performed in a large evaporating-dish, on account of the violent effervescence due to the escape of carbon dioxide. The supernatent liquid is now poured off and the residue washed several times with hydrochloric acid to remove the sodium salt, and finally with water in order to remove the acid. The powder, which contains iridium and plati- num (in the presence of one another, but not alloyed together), is dried, strongly ignited in a porcelain crucible over the blast-lamp (whereby the iridium is made insoluble in the aqua regia), and weighed. The ignited gray metal is dissolved (at as low a tem- perature as possible) in hydrochloric acid, with a gradual addition of nitric acid. Considerable quantities of nitroso-platinic choride are formed by this operation. On evaporating with water, this compound is decomposed into hydrochlorplatinic acid, with evolu- tion of the oxides of nitrogen: PtCl(NO),4-H,0=NO,--NO+[PtCl,]H,. As, however, a part of the NO, (or N.G.) remains in solution, some more nitric and nitrous acids are formed by the action of water, N.O,--H.O=HNO,--HNO, 236 REACTIONS OF THE METALS, which yields nitrosyl chloride with the hydrochloric acid present, and causes the formation of more nitroSo-platinic chloride. It is necessary, therefore, to evaporate the solution alternately with hydrochloric acid and with water until no more nitrous fumes are given off. The Solution thus obtained always contains some hydrochlorplatinous acid, and is intensely brown. In order to change the latter compound into hydrochlorplatinic acid, the fairly- warm solution is Saturated with chlorine gas (whereby its color be- comes much lighter) and then evaporated (at as low temperature as possible) till it becomes of syrup consistency. After cool- ing, the syrup crystallizes to a yellowish-brown mass, which may be dissolved in a little cold water, and the insoluble iridium filtered off. w If there is a considerable amount of the latter metal, it is ignited in a porcelain crucible and weighed. The weight of the iridium should be deducted from the previous weight of the mixture, in order to find out how much platinum remains in solution. The filtered solution is now diluted with water until 100 c.c. of the solution contains 10 gm. of platinum. r 2. Preparation of Hydrochlorplatinic Acid from Platinum Resi- dues.—These residues consist of potassium platinic chloride and the alcoholic wash-waters. By evaporating an alcoholic Solution of hydrochlorplatinic acid, hydrochlorplatinous acid and ethylene are formed, which yield ethylene platinous chloride; and the latter gives no precipitation with potassium or ammonium Salts: O / H.PtCl4-2CH,CH,OH=CH-C-H+4HCl·H (C.H.)PtCl2+H.O. Alcohol Aldehyde On evaporating an alcoholic solution of this soluble organic platinum compound, it is changed into an insoluble powder, explo- sive when dry, insoluble in acids, and completely decomposed by strong ignition only. To separate platinum, therefore, from platinum residues, the alcoholic solution is, first of all, evaporated to dryness, taken up in water, and the solution poured into caustic soda (sp. gr. 1.2), to which 8 per cent. of glycerine * has been added, the liquid heated tº _ – 2 + Zeitschr, für anal. Chemie, XVIII, p. 509. º º * Q * • , & SILVER. 237 sº to boiling, causing the platinum to be precipitated as a heavy black powder, \ 2C.H.O,4-6H,PtCl4-6H,0=36HCl·H2CO,4-2H,C,0,--6Pt, Glycerine Oxalic acid which is washed, first, with water, then with hydrochloric acid, and finally with water again. The powder is dried, ignited (to destroy any of the organic compound), weighed, and transformed, as before, into hydrochlorplatinic acid. REACTIONS IN THE DRY WAY. All platinum compounds, when heated with soda on charcoal, are reduced to the gray, Spongy metal, which assumes a metallic lustre on being rubbed with a pestle in an agate mortar. It can be distinguished from gold by its color, and from lead, tin, and silver by its infusibility and insolubility. Separation of Gold from Platinum. The platinum is precipitated with a solution of ammonium chloride; after being filtered the Solution is treated with ferrous sulphate to precipitate the gold. GROUP I. (HYDROCHLORIC ACID GROUP.) To this group belong silver, mercury (in mercurous com- pounds), lead, thallium, and (under Some circumstances) tungsten. SILVER, Ag. At. Wt. 107.93. Occurrence.—Silver occurs both native and combined (chiefly with sulphur, arsenic, and antimony). Of the silver-bearing minerals the following may be mentioned; Horn silver, Argentite, Pyrargyrite, and Proustite, AgCl, Ag,S, Sb(SAg), As(SAg), Silver is also found with tetrahedrite, and with galena. Metallic silver is of a pure white color, has the sp. gr. 10.5. melts at 954°C., and absorbs oxygen in the molten state, which it gives up (with tiny explosions) on cooling. 238 REACTIONS OF THE METALS. The proper solvent for silver is nitric acid. It is insoluble in dilute hydrochloric and Sulphuric acids, but dissolves readily in boil- ing sulphuric acid, with evolution of Sulphur dioxide: Ag,--2H,SO, =2H,0+SO,--Ag,SO,. The solubility of silver in concentrated sulphuric acid is made use of in separating silver from gold and platinum in alloys. Silver forms three oxides: Silver suboxide, Silver oxide, Silver peroxide, Ag,O, Ag,O, Ag,O, Of these oxides, Ag,O alone is a basic anhydride; only one series of salts is known. A Silver oxide is a brownish-black powder, which on being heated to 300° is completely decomposed into metal and oxygen. Most of the silver salts are colorless; the following, however, are colored: the bromide (yellow), the iodide (yellow), the sulphide (black), the phosphate (yellow), the arsenite (yellow), the arsenate (brown), and the chromate (reddish-brown). Most of the salts are insoluble in water, and are blackened on exposure to the light. The nitrate, chlorate, perchlorate, fluoride, nitrite, sulphate, and acetate are soluble in water, the three latter with difficulty. REACTIONS IN THE WET WAY. 1. Potassium and Sodium Hydroxides precipitate brown silver oxide, 2AgNO,--2KOH=2KNO,--H,0+ Ag,O, insoluble in an excess of the precipitant, but readily soluble iºnitric acid and in ammonia. If the solution in ammonia is allowed to stand, black detonating silver is deposited, [AgNH,0. 2. Ammonia.-If aneytral solution of a silver salt is cautiously treated with ammonia, the first drops produce a white precipitate, which changes quickly to the brown oxide, Ag,O. The greater part of the silver, however, remains in Solution as silver ammonium SILVER. 239 nitrate, [Ag(NH3)2]NO3,” and the oxide itself is dissolved by an excess of ammonia: Ag,0+2NH,--H,O=2|AgNH.jQH. 3. Sodium Carbonate precipitates white silver carbonate, which becomes yellow on being boiled, being slightly decomposed into oxide, with loss of carbon dioxide: 2AgNO,--Na,CO. =2|NaNO,--Ag,COs; Ag,CO, = Ag,O+CO,. 4. Ammonium-Carbonate produces the same precipitate, but it is soluble in an excess of this reagent. 5. Sodium Phosphate throws down in neutral silver solutions a yellow precipitate of silver phosphate: 3AgNO,--2Na,HPO,-3NaNO,--NaH,PO,--Ag, PO,. Silver phosphate is easily soluble in nitric acid and in ammonia. The solution of the phosphate in ammonia is due to the formation of a complex silver ammonium phosphate, perhaps Ag, PO,--3NH, -[AgNH, PO,. By-mºtº * the ammoniacal solution with nitric acid, or the Leesº * Z e & * sº © "nitric acid solution with ammonia, the silver phosphate is repre- cipitated. 6. Hydrochloric Acid and Soluble Chlorides precipitate from neutral and acid solutions white, curdy silver chloride: AgNO,--HCl=AgCl-i-HNO,. Silver chloride is perceptibly soluble in pure water, particularly on boiling, but it is quite insoluble if an excess of silver nitrate or of hydrochloric acid is present (mass action, cf. p. 15, under Barium Sulphate). Silver chloride dissolves to a considerable extent in a large ex- * Cf. A. Reychler, Chem. Centralbl., 1904, I, p. 252. \ 24O REACTIONS OF THE METALS, cess of hydrochloric acid or of alkali chloride, while it is almost insoluble in dilute nitric acid. It is very soluble in ammonia, AgCl-H2NH3=[Ag(NH3)2]Cl, but is reprecipitated on addition of nitric acid to this solution: [Ag(NH3)2]Cl, H 2HNO3=2NH4NOa--AgCl, Silver chloride is also readily Soluble in potassium cyanide solu- tion, AgCl-H2KCN=Ag(CN), K+KCl, and if such a solution is acidified, silver cyanide will be precipi- tated, with evolution of prussic acid: Ag(CN),K+HNO,-KNO,--HCN+AgCN. In the absence of acid, silver chloride is also readily soluble IIl 8, solution of Sodium thiosulphate: 2AgCl-H3Na,S,Os-[Ag,(S.O.), Nai-H2NaCl. On boiling this solution, silver Sulphide is not precipitated. Silver chloride is slowly attacked by boilingºentrated sul- phuric acid, with evolution of hydrochloric acid, and the forma- tion of crystalline silver Sulphate, insoluble in Sulphuric acid. By boiling with caustic Soda or caustic potash solution, silver chloride is only partially decomposed; in the cold it is unaffected. Sodium carbonate solution does not affect it; but by fusing with sodium carbonate it is completely decomposed: 2AgCl--Na,CO,-2NaCl-HCO,--O--Ag. By fusing silver chloride itself a yellow liquid is obtained, which on cooling solidifies to a horny mass. 7. Potassium Iodide precipitates yellow, curdy, silver iodide, almost insoluble in ammonia, but easily Soluble in potassium cya- nide and sodium thisosulphate Solutions. SILVER. 24I 8. Ferrous Sulphate precipitates gray metallic silver from boil- ing solutions: 3AgNO,--3FeSO,-Fe,(SO),4-Fe(NO.),4-3Ag. Frequently a basic ferric salt is precipitated at the same time, particularly from very dilute solutions. 9. Zinc, having a greater solution pressure than silver, precipi- tates the latter from neutral solutions. Similarly, if insoluble silver chloride is covered with dilute sulphuric acid and a piece of Zinc placed in contact with the chloride, the latter will be reduced to metal. 2AgCl-HZn=ZnCl2 +Ag2. 10. Hydrogen Sulphide precipitates from neutral, ammoniacal, and acid solutions black silver sulphide, 2AgNO,--H.S=2HNO,--Ag,S, insoluble in ammonia, alkali Sulphides, and dilute potassium cyanide , solution. Silver sulphide is perceptibly soluble in a concentrated solution of potassium cyanide,” and easily soluble in hot dilute nitric acid, with separation of sulphur and formation of silver nitrate. 11. Potassium Chromate precipitates brownish-red silver chro- mate, soluble in ammonia and in nitric acid. 2AgNO,--K,CrO, = 2KNO,--Ag,CrO,. 12. Potassium Dichromate precipitates reddish-brown silver dichromate soluble in ammonia and in nitric acid: 2AgNOs-i-K,Cr,O, =2KNO,--Ag,Cr,0,. * REACTIONS IN THE DRY WAY. Fused with soda on charcoal, all silver compounds yield a white, malleable, metallic button without incrustation (difference from lead and tin), readily soluble in nitric acid (difference from tin). * If a considerable amount of concentrated potassium cyanide solution is mixed with a silver solution, and dilute hydrogen sulphide added, no precip- itate is formed at first; but gradually silver sulphide is formed. and with more hydrogen sulphide the silver is all precipitated. 242 SEPARATION OF METALS OF GROUP I. The Solution is not precipitated by very dilute sulphuric acid, but is immediately precipitated by hydrochloric acid (difference from lead). The reactions of lead and mercurous compounds have already been described (pp. 156, 153). Separation of Silver, Lead, and Mercurous Compounds. The chlorides of these three metals are precipitated by hydro- chloric acid. [Lead is partly found in the filtrate, and is, therefore, always tested for in the hydrogen sulphide group.] The chlorides are filtered, washed with a little cold water, boiled with a consider- able quantity of water, and filtered hot. RESIDUE: [AgCl, Hg,0l.,]. SoluTION: PbCl,. It is treated with ammonia and On cooling the solution, crystals filtered: * of lead chloride will separate out if Residue: Solution: considerable lead is present. The NH, AgNH3Cl. solution is treated with dilute sul- Hg. Gº-Hg, - - - - * * & It is acidified phuric acid, whereby difficultly solu- black, showing with HNO3, when ble white lead sulphate is formed; presence of mer- a white precip- or it is treated with potassium di- cury. itate of AgCl chromate solution, when a yellow shows the pres- precipitate shows the presence of ence of silver. lead. * According to A. Thiel (Allgem. chem. Zeit., 1904) a small amount of silver chloride in the presence of considerable mercurous chloride cannot be dissolved by means of ammonia. Thiel, therefore, first treats the chlorides of silver and mercury with bromine water, thereby oxidizing the mercurous Salt, Hg,Cl, +Bra-HgCl2 +HgBr, which goes into solution, leaving the silver chloride behind. The latter is filtered off, dissolved in ammonia, and tested as above. REACTIONS OF THE METALLOIDS (ANIONS). In order to detect a metalloid, it is almost always necessary to get it in the form of an acid. The acids, as well as their salts, are always more or less dissociated into electropositive cathions, and electro-negative anions. Up to this point we have concerned our- selves chiefly with the study of the cations. In the following section of the book we will study the behavior of the anions. The presence of hydrogen ions in aqueous Solutions is the char— acteristic of all acids; upon it alone the acid reaction depends. In the Introduction (p. 19e) we have mentioned methyl orange as a typical reagent which can be used in testing for hydrogen ions. There are other organic dyestuffs which undergo a corresponding change of color on coming in contact with hydrogen ions, and which may be used as acid indicators. Among these substances are litmus and lacmoid (colored red by acids), congo red (changing to blue). etc. DIVISION OF THE ACIDS INTO GROUPS. The classification of the acids which we prefer was first pub- lished by R. Bunsen in 1878 (in manuscript form) for the use of his students; it has been adopted (with his consent) by W. Meyer and the author in their “Tabellen zur qualitativen Analyse.” This classification is based upon the different solubilities of the barium and silver salts of the different acids. GROUP I includes those acids whose silver salts are insoluble in water and in nitric acid, but whose barium salts are soluble in water. 243 244 REACTIONS OF THE METALLOIDS. To this group belong hydrochloric, hydrobromic, hydriodic, hydroferrocyanic, hydroferricyanic, sulphocyanic, and hypochlo- rous acids. GROUP II contains those acids whose silver salts are soluble in nitric acid, but are insoluble, or difficultly soluble, in water, and whose barium salts are soluble in water. To this group belong hydrosulphuric, hydroselenic, hydro- telluric, nitrous, acetic, cyanic, and hypophosphorous acids. GROUP III. comprises those acids whose Silver Salts are white and soluble in nitric acid, but whose barium Salts are difficultly soluble or insoluble in water, but soluble in nitric acid. To this group belong Sulphurous, Selenous, tellurous, phos- phorous, carbonic, oxalic, iodic, boric, molybdic (selenic and tel- luric), tartaric, citric, meta- and pyrophosphoric acids. GROUP IV contains those acids whose Silver Salts are colored and soluble in nitric acid, but whose barium Salts are insoluble in water and soluble in nitric acid. To this group belong phosphoric, arsenic, arsenious, vanadic, thiosulphuric, chromic, and periodic acids. GROUP V comprises those acids whose silver and barium salts are soluble in ^Dater. To this group belong nitric, chloric, perchloric, and the man- ganic acids. GROUP VI comprises those acids whose silver salts are soluble in water, but whose barium salts are insoluble in nitric acid. To this group belong sulphuric, hydrofluoric, and hydrofluo- silicic acids. w GROUP VII contains non-volatile acids, which form Soluble Salts with the alka- lies alone. To this group belong silicic, tungstic, titanic, niobic, tantalic, and zirconic acids. HYDROCHLORIC ACID. 245 GROUP I. Silver Nitrate produces a precipitate insoluble in nitric acid. Barium Chloride causes no precipitation. HYDROCHLORIC ACID, HCl. Occurrence.—Hydrochloric acid is found free in nature, but in small quantities (for example, in the exhalations of active volca- noes); its salts, however, are exceedingly common, especially those with the alkalies. (See these.) Preparation.—Hydrochloric acid is set free by the action of con- centrated sulphuric acid upon any chloride. Ordinary rock salt is usually used, it being the cheapest chloride. If sulphuric acid is allowed to act upon ordinary salt, a consider- able evolution of hydrochloric acid takes place even in the cold, with the formation of monosodium sulphate: NaCl-H H.SO,-NaHSO,--HCl, And, on warming, the monosodium sulphate reacts with more sodium chloride: NaHSO,--NaCl-Na,SO,--HCl. Hydrochloric acid may also be prepared by the action of water on many acid chlorides: PCl,--3H,0=PO,H,4-3HCl. Properties.—Hydrochloric acid is a colorless gas, with a suffo- cating odor, and forms dense clouds in moist air. It is readily solu- ble in water (one volume of water dissolves, at 18°C., 451 volumes of hydrochloric acid). The sp. gr. of the Saturated aqueous solu- tion is 1.21, and 100 gm. of this Solution contains 42.3 gm. of hydrochloric acid gas. The ordinary commercial hydrochloric acid is of sp. gr. 1.18 to 1.19, and contains 36 to 38 per cent of the gas. The aqueous solution of hydrochloric acid is one of our strongest acids. In dilute solution it is almost entirely dissociated, + *Eº HCl z- H+Cl, so that such a solution is a good conductor of electricity. 246 REACTIONS OF THE METALLOIDS. The behavior of hydrochloric acid on oacidation is extremely charac- teristic; the hydrogen is completely burned to water, and chlorine is set free: 2CIH+O= H,0+Cl, This oxidation will not take place on exposure to atmospheric, or even pure, Oxygen, but only by nascent oxygen under certain definite conditions. Thus the following classes of substances will oxidize concen- trated hydrochloric acid (sometimes in the cold): The perozides of the heavy metals,” all nitrates, nitrites, chlorates, hypochlorites, chro- nates, selenates, and tellurates. I. Oxidation of Hydrochloric Acid by Means of Peroxides. (a) By manganese dioacide: 2HCl·HMnO,-H,0+MnO--Cl, One atom of oxygen from the manganese dioxide molecule is availa- ble for oxidation. The manganese dioxide is reduced to manga- nese monoxide; and the latter, being the anhydride of a base, unites with more of the hydrochloric acid to form manganous chlo- ride: MnO--2HCl=MnCl,--H,0, so that the whole reaction may be represented by the following equation: MnO,--4HCl=MnCl,--2H,0+Cl, Some other acid, preferably sulphuric, can be used for the neu- tralization of the manganese monoxide, so that the yield of chlorine from a given quantity of hydrochloric acid will be twice as large: MnO,--2HCl-i-H,SO,-MnSO,--2H,0+Cl,. In the former case only half of the chlorine from the hydro- chloric acid was set free; in the latter the total chlorine becomes available. *The peroxides of the light metals do not yield chlorine, but hydrogen peroxide: BaO2 + 2HCl E BaCl, + H.O. HYDROCHLORIC ACID, 247 (b) By lead perozide: (a) 201H+PbO,-H,0+PbO-HCl, (b) PbO4-2HCl=H,0+ PbCl, Total reaction: PbO,--4HCl=2H,0+PbCl, HCl, (c) By chromium trioacide: (a) §4. 6HCl=3H,0+Cr,0,--3Cl.; (b) Cr,0,--6HCl=3H,O--2CrCl, Total reaction: 20r.0,--12HCl=6H,0+2CrCl,--3Cl. 2. Oxidation of Hydrochloric Acid by Nitric Acid, etc. As we saw in the Introduction (p. 4), the oxidizing action of nitric acid depends upon the formation of the anhydride, which is then decomposed into nitric oxide and oxygen: 2HNO,-H,0+ N.O.; N.O. =2NO+30. If concentrated nitric acid is allowed to act upon concentrated hydrochloric acid, the latter is oxidized with evolution of chlorine: H.O.N.O,O,--6HCl=4H,0+2NO+3Cl. \——” 2HNO, A part of the chlorine, however, combines with the nitric oxide, forming nitrosyl chloride, NO+Cl=NOCl, so that the reaction, in reality, is as follows: HNO,4-3HCl=2H,0+ NOCl-HCl, A mixture of one molecule of nitric acid with three of hydro- chloric acid is known as º Aqua Regia.-The acids are usually mixed, not according to their weights, but according to their volumes. 248 REACTIONS OF THE METALLOIDS. Aqua regia is, therefore, chlorine water, with the distinction that the chlorine exists in the nascent state; which explains why aqua regia is a much more energetic reagent than ordinary chlorine water. Nitrous acid, chloric acid, hypochlorous acid, selenic and telluric acids all act similarly with hydrochloric acid. Hydrochloric acid is monovalent; its salts are called chlorides. Solubility of Chlorides. Nearly all chlorides are soluble in water, but the following are insoluble: Mercurous chloride, Hg,Cl, Aurous chloride, AuCl. Silver chloride, AgCl. Platinous chloride, PtCl, Cuprous chloride, Cu,Cl, Bismuth oxychloride, BiOCl. Lead chloride, PbCl, Antimony oxychloride, SbOCl. Thallium chloride, TICl. Mercuric oxychloride, Hg,Cl,0. Of these chlorides which are insoluble in water all, with the exception of silver chloride, aurous chloride, and platinous chloride, are soluble in strong hydrochloric acid. They are all soluble in aqua regia, with the exception of silver chloride. By boiling these insoluble chlorides in a concentrated solution of sodium carbonate, all, with the exception of silver chloride again, are readily decomposed; e.g.: Hg,Cl2 +Na,CO3=2|NaCl-i-CO,--Hg,O. By filtration a chloride solution is obtained which is free from heavy metal. By fusing with sodium carbonate, even silver chloride is decom- posed, 2AgCl-i-Na,CO. =2|NaCl--CO,--O--Aga, and silver chloride may also be decomposed by nascent hydrogen (cf. p. 241). The deliquescent chlorides (lithium, calcium, and stronium) are all soluble in absolute alcohol and in amyl alcohol. The chlorides of potassium, sodium, and barium are almost en- tirely insoluble in concentrated hydrochloric acid; they can, there- fore, be easily separated from the remaining chlorides which are HYDROCHLORIC ACID. 249. soluble in water by saturating the solution with hydrochloric acid gaS. Most all chlorides are insoluble in ether, with the exception of mercuric, stannous, and stannic chlorides. REACTIONS OF CHLORIDES IN THE WET WAY. A neutral solution of an alkali chloride should be used for these reactions. Like hydrochloric acid, most chlorides are practically com- pletely dissociated in dilute aqueous solution: + *E* RClz-2 R+Cl. The chlorine of hydrochloric acid (or the soluble chloride) is present in the form of the chlorine ion. It is, therefore, a matter of indifference which chloride we take for the following reactions, provided there is no independent reaction taking place due to the presence of the cation. If an alkali chloride is used such reactions will not take place, consequently a neutral alkali Salt is usually used for the anion test. 1. Dilute Sulphuric Acid (1 : 10) produces no reaction, even on warming. 2. Concentrated Sulphuric Acid * decomposes the chloride almost completely in the cold, completely on warming. Sulphate and colorless hydrochloric acid gas result from this reaction; and the latter is easily recognized by its odor, by the clouds which it forms in moist air or with ammonia vapors (obtained by holding a glass rod wet with ammonia near the test-tube), and by its turning moist blue litmus-paper red. Water is not made turbid by hydro- chloric acid (difference from hydrofluosilicic acid). Silver chloride and mercurous chloride are decomposed with difficulty by sulphuric acid, the latter with evolution of sulphur dioxide; the mercurous Sulphate (which is at first formed) is oxidized (at the expense of the Oxygen of the sulphuric acid) to nercuric sulphate: (a) Hg,Cl,--H,SO,-Hg,SO,--2HCl; (b) Hg,SO,--2H,SO,-2HgSO,--2H,0+SO,. *This reaction is best performed with a solid chloride in a small test-tube. 25o REACTIONS OF THE METALLOIDS. 3. Silver Nitrate produces a white, curdy precipitate of silver chloride, RCl+AgNO,-RNO,--AgCl, insoluble in nitric acid, soluble in ammonia, potassium cyanide, and sodium thiosulphate solutions. (See Silver.) From a solution of silver chloride in ammonia, acids reprecipi- tate silver chloride. From a solution in potassium cyanide, acids precipitate silver cyanide. If it is desired to test a solution of fer- rous sulphate for the presence of a small amount of chloride, it must be strongly acidified with nitric acid, as otherwise a precipitate of metallic silver will be obtained, which may cause confusion (cf. p. 241). The best way to test the solution of ferrous sulphate for hydrochloric acid is to add sodium carbonate solution until alka- line, boil, and filter. In the filtrate the acids originally present are now in the form of their sodium salts, in the presence of an excess of sodium carbonate; and the latter should be neutralized with nitric acid before the silver nitrate is added. The detection of chloride when present in the form of chloride of a heavy metal is accomplished in a similar manner; and, with the exception of silver chloride, any insoluble chloride may be decom- posed in the same way, by boiling with sodium carbonate solution. In order to confirm the presence of chlorine in silver chloride, it is treated with dilute sulphuric acid, a piece of zinc added, and after a little while the solution is poured off from the deposited silver and tested with silver nitrate: 2AgCl-HZn=ZnCl2+Aga. 4. Potassium Dichromate and Sulphuric Acid.—If a dry chlo- ride is mixed with potassium dichromate, concentrated sulphuric acid added, and the mixture heated in a small retort, brownish vapors are given off which are condensed, in the receiver, to a brown liquid (chromyl chloride, CrO,Cl): K.Cr,0,--4NaCl-H.3H,SO,-3H,0+2Na,SO,--K,SO,--2CrO.C.,. Chromyl chloride is an acid chloride, and is, therefore, decomposed by water into chromic and hydrochloric acids: CrO.Cl,+2H,O=H,CrO,--2HCl. HYDROCHLORIC ACID. 25 I On adding caustic soda or potash, an alkali chloride and a yellow alkali chromate are obtained. If the solution is then acidified, Some ether and a little hydrogen peroxide added, and the liquid shaken, the upper ether layer will be colored blue, showing the presence of chromium; and the presence of chromium indicates that a chloride was originally present (difference from bromine and iodine). Behavior of Chlorides on Ignition. The chlorides of the alkalies and alkaline earths melt (without perceptible decomposition) on being heated in the air. The chlorides of the sesquioxides are decomposed, more or less completely, on being ignited in the air. Thus, ferric chloride is almost quantitatively decomposed into ferric oxide, with loss of chlorine: 2FeCl,--3O=Fe,O,--3Cl. The chlorides of gold and of the platinum metals are readily decomposed into chlorine and metal: 2AuCls=2Au-H3Cl2; PtCl,-Pt-i-2Cl, The remaining chlorides are mostly volatile, without perceptible decomposition. Détection of Chlorine in Non-electrolytes. Besides uniting with metals and with hydrogen, chlorine also forms compounds with the metalloids; e.g., PC, PCls, AsCl, AsCls, SbCl, SbCls, CC1, SiCl, etc. All these compounds, which may be regarded as acid chlorides, are decomposed by water with the formation of hydrochloric acid (which can be tested for in the usual way). Thus, phosphorous trichloride yields with water phosphorous acid and hydrochloric acid, PC,4-3HOH=H.PO,--3HCl, and phosphorous pentachloride yields phosphoric acid and hydro- chloric acid: PCls--4HOH=H.PO,--5HCl. 252 REACTIONS OF THE METALLOIDS. The remaining acid chlorides are decomposed in a similar way by water at the ordinary temperature, with the exception of carbon tetrachloride, which is only decomposed by heating with water in a closed tube as follows: CC1,4-2H,0=CO,--4HCl. Chlorine acts upon a great many hydrocarbons, forming substi- tution products which are non-electrolytes, and consequently will not give the chloride tests; for example, if chloroform, CHCl, is shaken with a solution of aqueous silver nitrate, it will not yield a precipitate of silver chloride. In order to test such compounds for chlorine (as is frequently necessary in the study of organic com- pounds), the chlorine must be changed to hydrochloric acid by one of the following methods: 1. Carius’ Method.—By heating the compound in a sealed glass tube with concentrated nitric acid, in the presence of silver nitrate, the compound is completely decomposed, and all the chlorine is changed to silver chloride, which can be filtered off; and, after treatment with zinc and dilute Sulphuric acid, can be tested as above. 2. By Heating with Lime.—A layer of granular lime (free from chloride), then a mixture of the substance to be tested with lime, and finally another layer of lime are placed in a small glass tube, which should be about 25 cm. long and about 1 cm. wide. By gently tapping the tube, a canal is opened between the upper wall of the tube and the substance, through which the gases evolved may escape. The tube is then placed horizontally in a small com- bustion furnace and heated (first the front layer of lime, then the back layer, and finally the entire contents of the tube) to a dull red heat. By this means the organic substance will be completely decom- posed, and the chlorine will be found combined with the lime in the form of calcium chloride. After cooling, the contents of the tube should be dissolved in dilute nitric acid, the carbon filtered off, and the filtrate tested with silver nitrate for chlorine ions. 3. Treatment with Sodium.—A small amount of the dry sub- stance to be tested is placed in a small test-tube, a small piece of sodium (freed from petroleum) is added, and the metal covered FREE CHLORINE. 253 with another layer of the substance. The tube is then heated in the gas-flame, the decomposition taking place Suddenly with in- candescence. The still hot tube is transferred to a small beaker containing water [which breaks the tube] and sodium chloride dis- solves with other sodium compounds. The solution is filtered, acidified with nitric acid, and then tested with silver nitrate for halogens. FREE CHLORINE. Chlorine (produced by the oxidation of hydrochloric acid or by igniting various chlorides) is a greenish-yellow gas, with a suffo- cating odor. It is absorbed by water (one volume of water absorbs at 10°C. about 2.7 volumes of chlorine gas), forming chlorine water, a yellowish-green liquid, and a poor conductor of electricity (al- though better than pure water, showing that some ions are present). Chlorine probably decomposes water to a slight extent, forming hydrochloric acid and hypochlorous acid: HOH+Cl, a HCl·H HOCl. If the acids formed are removed by the addition of alkali, the reaction proceeds quantitatively in the direction from left to right: Cl,--2NaOH = NaCl-H NaOCl-H H.O. Chlorine water is a strong oxidizing agent (cf. p. 4), which can be explained by assuming that water is decomposed by the chlorine, forming hydrochloric acid and setting free oxygen: H,0+Cl,-2HCl·HO. A better explanation is the following: First of all, hydrochloric and hypochlorous acids are formed, and the latter is then decom- posed into hydrochloric acid and oxygen: H,04-Cl,—HCl·H HOCl; HOCl= HCl–H O. This decomposition of chlorine water takes place slowly in the dark, but more rapidly in the light, and in the presence of oxidiz- able substances. The bleaching action of chlorine water depends 254 REACTIONS OF THE METALLOIDS. upon this oxidizing action. Thus, indigo solution is decolorized, the indigo being oxidized to isatin: Cishion,0,-HO, =2CsPIs NO,. Indigo Isatin If a solution of potassium iodide is treated with chlorine water, iodine is set free, and the Solution turns yellow to brown: 2KI+Cl,+2KCl+I, If the yellow solution is shaken with carbon disulphide, or chlo- roform, the iodine is absorbed by these reagents and they assume a reddish-violet color. By the addition of more chlorine water the solution becomes colorless, owing to the oxidation of the iodine to colorless iodic acid: I, +6H.O.--5C1,-10HCl·H2HIOa. The free iodine can also be detected by the addition of some starch paste (instead of carbon disulphide, etc.), which is turned blue by iodine. Silver Nitrate gives a white precipitate of silver chloride in chlorine water; but this precipitation is not quantitative, for one sixth of the chlorine is changed into soluble silver chlorate: 6Cl·H 6AgNO,--3H,0=5AgCl- AgClO,--6HNO,. On adding a slight excess of sulphurous acid to chlorine water the chlorine is readily and completely changed into hydrochloric acid: Cl, H H.O.--H.SO,-H,SO,--2HCl. From this solution the chlorine can be precipitated quantita- tively by a silver solution. Chlorine may also be easily changed into the form of a chloride by the addition of ammonia. At first, ammonium chloride and ammonium hypochlorite are formed: Cl2+2NH,OH=NH,Cl-i-NH,ClO4-H,0. The ammonium hypochlorite immediately oxidizes more of the am- monia, with evolution of nitrogen: 3NH.OC14-2NH,-3H,0+N,+3NH,Cl. HYPOCHLOROUS ACID. 255 The total reaction is, therefore, 3Cl2 + 6NH,OH+2NH,-6H,0+ 6NH,Cl+N. Chlorine can be changed to a chloride by the action of hydrogen peroxide in the presence of sodium or potassium hydroxide: Cl2+2NaOH=NaCl-H NaOCl-H H.O, and NaClO--H,O,-H,0+ NaCl--O,. Metallic Mercury is attacked by chlorine at the Ordinary tem- perature, forming insoluble mercurous chloride. If therefore chlorine water is shaken with metallic mercury until it no longer Smells of chlorine, a neutral Solution is obtained which contains no chlorine. If hydrochloric acid is present, the solution now reacts acid, and gives a precipitate with silver nitrate, for me— tallic mercury is not attacked by hydrochloric acid. This reaction is made use of in testing for hydrochloric acid in the presence of chlorine. HYPOCHLoRous ACID, HOCl. Preparation.—A solution of free hypochlorous acid is obtained by shaking chlorine water with yellow mercuric oxide until the solution no longer Smells of chlorine: Hg—Cl 2Hgo426,4-H,0- X)+2HOCl. Hg—Cl Brown, insoluble mercuric oxychloride is formed by the reaction, and the solution contains hypochlorous acid. If the solution is poured off from the insoluble basic mercuric salt, and distilled, a pure solution of hypochlorous acid will be obtained; which, how- ever, cannot be kept long in the light, for it decomposes into hydro- chloric acid and oxygen: 2HOCl=2HCl·HO,. Hypochlorous acid is a vigorous bleaching agent; litmus and indigo are quickly decolorized. The alkali salts of hypochlorous acid (hypochlorites) are ob- 256 REACTIONS OF THE METALLOIDS, tained by neutralizing the acid with dilute sodium or potassium hydroxide; or, more conveniently, by the action of chlorine on a dilute caustic alkali solution: Cl, H2KOH=KCI+KOCl4-H,0. The ammonium salt cannot be prepared, because it is imme- diately decomposed, acting upon the excess of ammonia present (cf. p. 254). All hypochlorites are readily changed, on warming, into chlo- rate * and chloride (cf. p. 122, foot note): |KO Cl flººr RClOs; FCO Cl consequently hypochlorites must always be prepared in cold dilute Solution. The most important commercial hypochlorite is the so-called “chloride of lime,” which is obtained by passing chlorine gas over lime at the ordinary temperature. REACTIONS IN THE WET WAY. All hypochlorites are soluble in water, and are decomposed by acids (carbonic acid even). º Hydrochloric Acid is oxidized by hypochlorites with evolution of chlorine. NaOCl-H2HCl= NaCl-HH,0+Cl, Sulphuric Acid decomposes hypochlorites, setting free hypo- chlorous acid: NaOCl-HH,SO,-NaHSO,--HOCl; and carbonic acid acts similarly: 2NaOCl4-H,CO,-Na,CO,--2HOCl. It is due to the fact that the hypochlorites are so readily decom- posed with the formation of chlorine that they act as strong bleach- ing agents; indigo Solution (a solution of indigo in Sulphuric acid) being immediately decolorized. Hypochlorites act as oxidizing agents not only in acid solutions * In the presence of 40% or more of ca 1stic potash the potassium hypo- chlorite on being heated decomposes into chloride with evolution of oxygen and the formation of no perchlorate (F. Winteler, Zeitschr. f. angew. Chem., 33 (1902), p. 778). HYPOCHLOROUS ACID, 257 but also in alkaline solutions at ordinary temperatures (difference from chlorates); many metallic hydroxides being oxidized by them to higher hydroxides. Thus, ferrous hydroxide is readily oxidized to reddish-brown ferric hydroxide; manganous, nickelous, and cobaltous hydroxides are oxidized to brownish-black hydroxides: 2Fe(OH),4-NaOCl-H H.O = NaCl-H2Fe(OH), ; Mn(OH),4-NaOCl= NaCl-i-MnO, H.; 2Ni(OH),4-NaOCl.--H,0=NaCl-H2Ni(OH), 200(OH), -NaOCl-i-H,O=NaCl--2Co(OH),. Iodo-Starch Paper is turned blue by hypochlorites in weakly alkaline solutions, owing to the separation of iodine: 2KI+NaOCl.--H,0=2|KOH+NaCl-HI, This reaction is not quantitative, because the iodine reacts, to some extent, upon the potassium hydroxide, forming potassium hypoiodite and potassium iodide: 2I-L2KOH a KI+ KIO--H,O. But as this reaction is reversible, the extent to which it takes place must depend upon the concentration of the potassium hy- droxide; and in dilute solutions enough iodine will be present to produce the blue coloration with the starch. Metallic Mercury.—If a solution containing free hypochlorous acid is shaken with metallic mercury, brown basic mercuric chlo- ride is formed, insoluble in water, but soluble in hydrochloric acid: 2Hg+2HOCl=Hg,Cl,0+H.O. Under these same conditions free chlorine forms, on being shaken with mercury, white mercurous chloride, which is insoluble in hydrochloric acid. This property is utilized in detecting hypochlorous acid in the presence of free chlorine. The mixture is shaken with mercury until a little of the solution no longer turns iodo-starch paper blue; the liquid is then carefully poured off, hydrochloric acid is added to the residue, and the mixture is shaken, when the basic chloride produced by hypochlorous acid goes into solution: Hg,Cl,0+2HCl=H,O-H2HgCl, while mercurous chloride remains undissolved. 258 REACTIONS OF THE METALLOIDS. If hydrogen sulphide is passed into the filtered solution, the formation of mercuric sulphide shows that hypochlorous acid was originally present. The salts of hypochlorous acid behave differently towards mer- cury than is the case with the free acid; they form insoluble mer- curic oxide and a soluble chloride: Hg-i-NaOCl=HgC)+NaCl. Silver Nitrate causes in solutions of hypochlorites an incom- plete precipitation of silver chloride. One third of the chlorine remains in solution in the form of chlorate: 3NaOCl-i-3AgNO,-3NaNO,--AgClO,--2AgCl. First of all, the unstable silver hypochlorite is formed: 3NaOCl-i-3AgNO,-3NaNO,--3AgOCl, which is immediately decomposed into chlorate and chloride: 3AgOCl=AgClO,--2AgCl. Hypochlorous acid is distinguished from chlorine by its be- havior towards mercury; from hydrochloric acid by its oxidizing action; and from chloric acid by its being partly precipitated by silver nitrate, and by its oxidizing action in alkaline Solutions. HYDROBROMIC ACID, HBr. Occurrence.—Bromine compounds are constantly found in nature in company with those of chlorine; so that salts of hydrobromic acid are found in the ocean and in many mineral waters. Preparation.—Hydrobromic acid is formed by the action of sulphuric acid upon a bromide: 2NaBr-i-H,SO,-Na,SO,--2HBr. But the hydrobromic acid obtained by means of this reaction is never pure, being contaminated with bromine; a part of the hydro- bromic acid is oxidized by the Sulphuric acid: 2BrFI+H.SO,O= H,0+H.SO,--Brº. HYDROBROMIC ACID. 259. The more concentrated the Sulphuric acid used, the larger will be the yield of bromine. J If dilute sulphuric acid (3H,SO, :1H,0) is used, the hydro- bromic acid obtained is almost free from bromine. Pure hydrobromic acid may be obtained by the action of an acid bromide upon water: PBr,4-3HOH=P(OH),4-3HBr. Properties.—Hydrobromic acid (like hydrochloric acid) is a col- orless gas with a suffocating odor, having the property of fuming in moist air and forming clouds of ammonium bromide with vapors of ammonia. It is very Soluble in water. The concentrated solu- tion has the sp. gr. 1.78 and contains 82 per cent of hydrobromic acid. The affinity of bromine for hydrogen is greater than that of iodine, but less than that of chlorine. While hydrochloric acid can be kept in aqueous solution for an indefinitely long time, a Solution of hydrobromic acid soon turns brown, owing to the separation of bromine. Hydrobromic acid is. oxidized by atmospheric oxygen: 2HBr-i-O=H,0+ Br. As hydrobromic acid is so easily decomposed into its elements, it attacks metals readily, even metals which are unattacked by: hydrochloric acid. Thus, bright copper is promptly dissolved by hydrobromic acid, with evolution of hydrogen and formation of cuprous bromide: Cu, +2HBr=Cu,Br,4-H. As hydrobromic acid is oxidized by atmospheric oxygen, it, follows that it is much more readily oxidized by nascent oxygen. Thus, it is oxidized, with separation of bromine, by peroxides, nitrates, nitrites, chromates, etc., provided a concentrated solution of hydrobromic acid is used. Hydrobromic acid is a monobasic acid; its salts are called bromides. & The solubility of a bromide is about the same as that of the: corresponding chlorine compound. 26o REACTIONS OF THE METALLOIDS. REACTIONS, IN THE WET WAY. 1. Dilute Sulphuric Acid (1 : 10) evolves no hydrobromic acid from bromides in the cold, but does so on warming from bromides of the alkalies. 2. Concentrated Sulphuric Acid causes evolution of hydro- bromic acid and bromine from all bromides. The solution is brown, and, on warming, yellowish-brown vapors are given off (difference from hydrochloric acid); which, as they contain hydrobromic acid, fume in moist air, have a Suffocating odor, and do not render water turbid (difference from hydrofluosilicic acid). 3. Silver Nitrate produces a curdy, yellowish precipitate of silver bromide, insoluble in nitric acid, but soluble in ammonia, potassium cyanide, and Sodium thiosulphate. If silver chloride is digested with potassium bromide, the former will be changed into silver bromide: AgCl-i-KBr=AgBr-i-KCl. If, however, silver bromide is heated and treated with chlorine, it is readily changed into silver chloride: AgBr-i-Cl–AgCl-i-Br. 4. Chlorine Water, on being added to solutions of soluble bro- mides, sets free bromine, which is soluble in carbon disulphide or chloroform, forming a brown Solution; but it is changed, by an excess of chlorine water, into yellowish chloride of bromine (BrCl) (difference from iodine). 5. Potassium Dichromate, in the presence of dilute sulphuric acid, effects no separation of bromine from aqueous solutions of bromides; on shaking the solution with carbon disulphide, the latter remains colorless (difference from iodine). 6. Potassium Dichromate and Concentrated Sulphuric Acid— On mixing a solid bromide with solid potassium dichromate, cover- ing the mixture with concentrated Sulphuric acid and distilling, a brown distillate is obtained (as with a chloride), which, however, consists of bromine and contains no chromium: K.Cr,0,--6NaBr-1-7H,SO,- =K,SO,4-3Na,SO,--Cr,(SO.),4-7H,0+3Br, FREE BROMINE. 261 On adding dilute sodium hydroxide to the distillate, a colorless (or sometimes a faint yellow) solution is obtained; which, after being acidified with sulphuric acid, does not give the chromium reaction with dilute sulphuric acid and hydrogen peroxide * (differ- ence from chlorine). 7. Nitrous Acid does not cause the separation of bromine from a dilute bromide solution in the cold (difference from iodine). Detection of Bromine in Non-electrolytes. The method of procedure is exactly the same as was given in the case of chlorine in a non-electrolyte (see page 251). FREE BROMINE. Free bromine (which may be obtained by the oxidation of hydro- bromic acid) is a brown liquid at the ordinary temperature, and dissolves in water, forming a colored solution. The cold saturated solution of bromine contains 2–3 per cent of dissolved bromine. Concentrated hydrochloric acid at the ordinary temperatures dis- solves much more bromine, the Saturated solution containing about 13 per cent of the substance. Bromine, like chlorine, acts as a strong bleaching agent (oxidiz- ing the color) and combines directly with metallic mercury, forming insoluble mercurous bromide. The detection of hydrobromic acid in the presence of bromine is accomplished in precisely the same way as was described for the detection of hydrochloric acid in the presence of free chlorine (cf. p. 255). Bromine acts upon the alkalies exactly as chlorine does: (a) In a cold dilute solution: Br,4-2NaOH=NaBr-H NaOBr-H H.O. Sodium hypobromite (b) On warming: 3Br,4-6NaOH=5NaBr-i-NaBrO,--3H,O. Sodium bromate (c) Upon ammonia: 3Br,4-6NH,OH--2NH,-6H,0+ 6NH,Br-i-N,. * The solution becomes brown owing to the separation of bromine: 262 REACTIONS OF THE METALLOIDS, HYDRIODIC ACID, HI. Occurrence.—Iodine occurs in nature as the iodide and as the iodate; * most frequently as the former, accompanying (in small amounts) chlorine and bromine; e.g., in the ocean, in mineral waters, etc. Preparation.—Hydriodic acid may be obtained pure by the action of an acid iodide upon water: PI--3H,0=PO.H,--3HI. If we attempt to prepare hydriodic acid by the action of sul- phuric acid upon iodides, even from dilute solutions, it is more diffi- cult than in the case of hydrobromic acid to obtain a pure pro- duct, on account of the slight affinity that iodine has for hydrogen. The hydriodic acid thus obtained always contains a large amount of iodine, together with the reduction products of sulphuric acid; the latter varying in composition according to the concentration of the acid employed and of the iodide solution. Thus, with con- siderable concentrated sulphuric acid, sulphur dioxide is obtained: (a) 2NaI+H,SO,-Na,SO,4-2HI. (b) H.SO,O--2HI= H,0+H,SO,--I, \–-y—’ H,0+SO, But if, on the other hand, considerable hydriodic acid is present, the sulphuric acid is reduced down to hydrogen sulphide: H.SO,--8HI=4H,0+H.S.--4I,. On heating considerable potassium iodide with concentrated sulphuric acid in a small test-tube, grayish, solid iodine separates out, which is volatile on warming, forming violet vapors. Further- more, hydrogen sulphide is given off at the same time, as can be readily shown by the lead acetate paper test (see Hydrogen Sul- phide). Properties.—Hydriodic acid is a colorless gas, with a suffocating odor; it fumes in moist air, and is readily soluble in water, forming a strongly-fuming liquid of sp. gr. 1.99–2.0. On account of the slight aſſinity which iodine possesses for hydrogen, an aqueous solu- •º rººm-º: *To the extent of 0.5 per cent. in Chili saltpeter. HYDRIODIC ACID. 263 tion of hydriodic acid is even more difficult to keep than a solution of hydrobromic acid; it soon turns brown, owing to the separation of iodine: º 2HI+O (air) = H,0+I,. If hydrogen sulphide is conducted into the brown solution, it is decolorized, with separation of sulphur: I,--H.S=2HI+S. On account of the instability of hydriodic acid, it attacks metals which are unaffected by both hydrochloric and hydrobromic acids; thus mercury is easily dissolved, with evolution of hydrogen: Hg--4HI=[HgI.]H,--H,. Hydriodic acid (like hydrochloric and hydrobromic acids) is oxi- dized by peroxides, nitrates, nitrites, chromates, etc., with separa- tion of iodine; only in this case the oxidation of the hydriodic acid takes place much more readily, so that a very dilute solution is oxi- dized by nitrous and chromic acids even in the cold. (See next pages.) Hydriodic acid is a monovalent acid; its salts are called iodides. The solubilities of the iodine compounds are almost exactly analogous to the corresponding bromine and chlorine compounds. Mercuric iodide, HgI2, and palladium iodide, PdL, are, however, insoluble in water, while the corresponding chlorine and bromine compounds are soluble. REACTIONS IN THE WET WAY. 1. Dilute Sulphuric Acid (1 : 10) attacks the iodides of the alkalies perceptibly, but only on warming. 2. Concentrated Sulphuric Acid reacts in the cold (cf. p. 262). 3. Silver Nitrate produces a yellow, curdy precipitate of silver iodide, insoluble in nitric acid, and only slightly soluble in ammonia, but readily soluble in potassium cyanide and sodium thiosulphate. By the action of chlorine, silver iodide is readily changed into silver chloride: AgI+Cl=AgCl-HI. 264 REACTIONS OF THE METALLOIDS. On the other hand, if chloride or bromide of silver is treated with potassium iodide it will be changed into silver iodide: AgCl--KI = KCl-i-AgI, AgBr-i-KI=KBr-i-AgI. 4. Lead Salts precipitate yellow lead iodide, soluble in consid- erable hot water and forming a colorless solution which deposits gold-yellow plates of PbT, on cooling. 5. Palladium Chloride (it is best to use sodium palladium chlo- ride) precipitates, from dilute solutions of an iodide, black palla- dium iodide (difference from chlorine and bromine), PdCl,Na,--2KI=2NaCl-H2KCl4-PdL, which is readily soluble in an excess of potassium iodide. 6. Cupric Salts are reduced by iodides, causing the separation of a brownish mixture of cuprous iodide and iodine: 2CuSO,--4KI=2K,SO,--Cu,I,--I. If sulphurous acid is then added to the solution, a nearly white deposit of cuprous iodide is obtained, owing to the deposited iodine being changed to hydriodic acid by the sulphurous acid: H.SO,--H,0+I, +H,SO,--2HI. 7. Nitrous Acid.—If a dilute solution of an iodide is treated with nitrous acid, iodine separates out and the solution becomes yellowish to brown in color (difference from chlorides and iodides): 2HNO,--2HI=2NO+I, +2H,0. This extremely sensitive reaction is best performed as follows: The solution to be tested is treated with a few drops of a solu- tion of nitrous acid in concentrated sulphuric acid,” a little carbon *This solution is prepared by heating arsenic trioxide with nitric acid (sp. gr. 1.30–1.35) and conducting the gases evolved (NO2 and NO) into sulphuric acid (sp. gr. 1.75–1.80): AS2O: + 2HNO, = As2O5 + H.O + NO2 + NO, 2So. 85 + No. 4 No-Ho +2so. 8 so. This solution of nitrosyl sulphuric acid is sometimes called “nitrose.” It can be kept for some time, but water decomposes it into nitrous and sulphuric acids: SO.Lºo + HOH = H.So. 4 HNo. FREE IODINE. 265 disulphide or chloroform is added and the mixture shaken, when the carbon disulphide (or the chloroform) becomes reddish-violet in color. The iodine set free may also be tested for with starch paste (see p. 267). 8. Potassium Dichromate, in the presence of dilute sulphuric acid, causes the separation of iodine, in the cold, from dilute iodide solutions; the iodine can be more easily recognized by shaking the solution with chloroform or carbon disulphide (difference from bromine): K.Cr,O,--6KI+7H,SO,-4K,SO,--Cr,(SO.),4-312-1-7H2O. By heating a mixture of solid iodide and solid potassium di- chromate with concentrated Sulphuric acid, iodine is set free (ac- cording to the above equation), which distils over with the steam, and can be condensed in the receiver. No chromium is carried over by this reaction (difference from chlorine). 9. Mercuric Chloride produces scarlet mercuric iodide, soluble in an excess of potassium iodide: HgCl, -2KI=2KCl-i-HgI,; HgI, +2KI=[HgIJK,. 10. Chlorine Water sets free iodine from iodides, 2KI+Cl2=2|KCl-HI2, which colors carbon disulphide reddish-violet, or starch-paste blue. By adding an excess of chlorine water, the violet color disap- pears, the iodine being oxidized to colorless iodic acid: I2+6H2O+5Cl2= 10HCl·H-2HIO3. Detection of Iodine in Non-electrolytes. The processes to be employed are the same as those described for detecting chlorine in non-electrolytes (see page 251). FREE IoDINE. Free iodine forms scales resembling graphite in appearance, and has a sp. gr. of 4.94 at 17°C. It melts at 114°C. (at the same tem- perature as sulphur), but begins to volatilize at ordinary tempera- tures, and is completely transformed into violet vapors at 200° C. 266 REACTIONS OF THE METALLOIDS. It is only slightly soluble in water (100 parts of water dissolve 0.02 parts of iodine), soluble to a considerable extent in alcohol and ether, forming a brown solution, and much more soluble in carbon disulphide and in chloroform; so that all the iodine from an aqueous Solution can be removed by shaking a few times with either of these Solvents. Iodine is still more soluble in hydriodic acid, or in a Solution of an alkali iodide, forming a triiodide: RI+I, +IIs. Commercial iodine always contains water, chlorine, bromine, and often cyanogen (iodine cyanide) as impurities. An aqueous Solution of iodine is a weak oxidizing agent. If hydrogen Sulphide is passed through an aqueous solution of iodine, it becomes colorless and turbid, owing to the separation of Sulphur: H2S +I2=2HI+.S. Solid iodine is not acted upon at ordinary temperatures by hydrogen sulphide; heat is necessary to produce the endothermic hydriodic acid. In aqueous solution the necessary heat energy is furnished by the solution of the hydriodic acid formed in water. The fact that solid iodine is not acted upon by hydrogen Sulphide, while it decomposes arsine, is utilized in the preparation of hydrogen sulphide free from arsenic from pyrites containing arsenic (cf. p. 192). The mixture of hydrogen sulphide and arseniuretted hydrogen is passed over iodine, and the latter only is removed. Sodium. Thiosulphate decolorizes iodine solutions, forming sodium tetrathionate and Sodium iodide: 2Na2S2O3 +I2 =2|NaI + Na2S4O6. Chlorine and bromine react in exactly the same way upon sodium thiosulphate when they are not present in excess. In the latter case the reaction goes further and the tetrathionate is oxi- dized to sulphate and sulphuric acid with deposition of Sulphur, and the sulphur itself is gradually oxidized to Sulphuric acid by the halogens: * 2Na2S4O6+12H2O+8Cl2=16HCl·H2Na2SO4+4H2SO4+S2. S2 +8H2O+6Ol2=12HCl +2H2SO4. Other weak oxidizing agents, such as ferric and cupric salts, act upon thiosulphate similar to iodine (see Thiosulphuric acid). FREE IODINE. 267 Starch Paste.—Free iodine colors starch paste blue, but only in the presence of hydriodic acid or a soluble iodide. Opinions differ concerning the composition of “iodide of starch.” Some hold that it is a compound, while others regard it as a Solid Solution. Ac- cording to Mylius,” iodide of starch is the hydriodic acid compound of an iodine addition-product of starch, containing about 18 per cent iodine, corresponding to the formula C, H,60,0I, HI. This com- pound acts as an acid. If iodide of starch is produced in a neutral solution, in the presence of iodides, a salt of the above acid is formed, of which Mylius isolated the barium salt. Iodide of starch, then, can be regarded as a double salt, similar to carnallite, MgCl, KCl. In dilute solutions it must be dissociated into its components, e.g., the potassium Salt C.O.O.M.I, KI z-2 C.H.O.J.-H. KI, and if we assume that the compound C, H,60,0I, is colorless, the aqueous solution of starch iodide will be colorless; but on increasing the concentration of the alkali iodide, the double salt will be less dissociated and the blue color of the undissociated compound will appear, which corresponds with the facts. If a dilute aqueous iodine Solution (obtained by shaking iodine with water) is added drop by drop to a dilute aqueous starch solution, a blue color will appear at the spot where the two solutions first come in contact, but this color will disappear on stirring. If some potassium (or other) iodide is added to the colorless Solution of the starch and iodine, a permanent blue coloration will at once appear. The temporary appearance of the blue color, immediately on adding the iodine solution, is probably due to the fact that the iodine first forms a substitution product with the starch, setting free hydriodic acid, which furnishes the conditions for the formation of the iodide of starch. The fact that a starch Solution containing an iodide is much more sensitive than one in pure water has been known for a long time. *Mylius, B. B., 20, 688, and C. Lonnes, Zeitschr, für anal. Chemie, XXXIII, 409. 268 REACTIONS OF THE METALLOIDS. * Detection of Hydrochloric, Hydrobrcmic, and Hydriodic Acids in the Presence of One Another. The solution to be tested should contain the alkali salts of the above acids. Half of this solution is taken for the bromide and iodide test, while the other half is retained for the chloride test. (a) Detection of Bromine and Iodine. 1. The solution is acidified with dilute sulphuric acid, about one c.c. of colorless carbon disulphide, or chloroform, and a drop of chlo- rine water are added, and the mixture is shaken. If iodine is pres- ent (even in the presence of bromine), the carbon disulphide will be colored reddish-violet. In order to detect the bromine, more chlorine water is added, with repeated shaking of the liquid, until the reddish-violet color has disappeared, showing that the iodine has been completely oxi- dized to iodic acid; the brown color of the bromine will then ap- pear dissolved in the carbon disulphide, becoming a pale yellow on further addition of chlorine water. 2. Instead of using chlorine water, it is often better to test for iodine (especially when only Small amounts are present, as in min- eral waters) with nitrous acid. The solution to be tested for iodine and bromine is acidified slightly with dilute sulphuric acid, carbon disulphide and a few drops of a solution of nitrous acid in Sulphuric acid are added, and the mixture shaken. If the carbon disulphide is colored reddish-violet, iodine is present. The aqueous solution is now poured off from the carbon disulphide (through a moistened filter, in order to remove any suspended drops of colored carbon disulphide), chlorine water is added, and the Solution shaken with fresh carbon disulphide. If the latter now assumes a brown color, bromine is present. (b) Detection of Chlorine. The simplest and surest way of separating chlorine from bromine and iodine is by fractional precipitation with silver nitrate. If the solution containing salts of the three halogens is treated with dilute silver nitrate, drop by drop, the iodine will be first precipi- tated as yellow silver iodide, then the bromine as a slightly-yellow HYDROCYANIC ACID, 269 silver salt, and finally the chlorine as pure white silver chloride. To separate the chlorine from the other two halogens, a few drops of the solution to be tested are taken and acidified with nitric acid. A drop of silver nitrate (1 : 100) is added, and the mixture shaken and boiled, which causes the precipitate to collect together. If bromine or iodine is present, the precipitate is yellow. The pre- cipitate is filtered off and again treated with the dilute silver nitrate, etc., until a pure white * precipitate of silver chloride is obtained, in case chlorine is present. HYDROCyANIC ACID (PRUSSIC ACID), HCN. Occurrence.—The compound of hydrogen with the monovalent radical —CHN, cyanogen seldom occurs in nature. It is found in all parts of a tree growing in Java (Pangium Edule), particularly in the seed kernels. Its compounds are found in many plants as a glu- coside (amygdaline), which yields, on hydrolysis, sugar, benzalde- hyde, and prussic acid: C, H, NO,1-1-2H,0=20.H.O.--C.H.CHO-i-HCN. mygdaline Sugar Benzaldehyde This amygdaline is found in bitter almonds, in the fruit kernels of cherries, apricots, peaches, etc., and in the leaves of the common laurel tree (Prunus Laurocerasus), Amygdaline is usually accompanied by a ferment, so that, on macerating the parts of the plant which contain the amygdaline, an aqueous solution of prussic acid is obtained (bitter-almond water). Preparation.—If ammonia is passed over glowing carbon, ammo- nium cyanide is formed; SO that this salt, as well as other cyanogen compounds, is found in the “gas-water” obtained by the dry dis- tillation of coal. Hydrocyanic acid may also be prepared by adding acid to many cyanogen compounds. If yellow prussiate of potash is treated with dilute sulphuric acid and distilled, prussic acid is evolved, which, after standing over solid calcium chloride, may be obtained in the anhydrous condition as a colorless, exceedingly poisonous liquid, smelling of bitter almonds, and boiling at 26.5°C. : Fe(CN).E. *E**** Fe(CN).E; tºº * By filtering off the precipitate it is easy to tell whether it is pure white, for the slightest tinge of yellow will show against the white paper. 27o REACTIONS OF THE METALLOIDS. Properties.—The liquid (as well as the gaseous) hydrocyanic acid burns with a reddish flame, and mixes in all proportions with Water, alcohol, and ether. Aqueous hydrocyanic acid cannot be kept indefinitely; a brown deposit soon appears, with the formation of formic acid and ammo- Illa . CHN--2HOH=NH,--HCOOH. Formic acid If a little mineral acid is added to the aqueous solution, it may be kept much longer; but, even then, ammonia and formic acid will be formed after a long time.* Hydrocyanic acid in aqueous solution is a very poor conductor of electricity; in other words, it is a weak acid, and only dissociated to a slight extent. The salts of hydrocyanic acid, the cyanides, are very similar in their properties to the corresponding halogen compounds, but are distinguished by their ability to form stable complex salts, which contain scarcely any cyanogen ions in aqueous Solution, and conse- quently do not give Some of the reactions for hydrocyanic acid. Solubility of Cyanides.—The cyanides of the alkalies and alka- line earths are readily soluble in water, but are hydrolytically de- composed into metallic hydroxide and hydrocyanic acid, and their ions: RCN-H HOH a KOH-- HCN. But as hydrocyanic acid is only slightly dissociated, the aqueous solution of an alkali cyanide behaves like a solution of alkali hy- droxide containing free prussic acid; the Smell of the latter can be readily detected. The remaining cyanides, with the exception of mercuric cyanide, are insoluble in water. REACTIONS IN THE WET WAY. 1. Dilute Sulphuric Acid decomposes solutions of all soluble cyanides, with the exception of mercuric cyanide, Setting free hydro- cyanic acid in the cold (recognizable by its odor). The insoluble cyanides are decomposed by dilute sulphuric acid only on warming. 2. Concentrated Sulphuric Acid decomposes all cyanides on warming, the complex cyanides as well as the simple ones. The * Cold concentrated hydrochloric acid converts hydrocyanic acid into formamide, HCN +H.O =HCONH2, but on warming this compound is de- composed into carbon monoxide and ammonia. HYDROCYANIC ACID. 27 I metals are then obtained in the form of sulphates, the carbon of the cyanogen is changed to carbon monoxide, and the nitrogen into ammonia (or rather ammonium sulphate): Ni(CN),4-2H.SO,4-2H,0=NiSO,4-(NH,),SO,--2CO. With mercuric cyanide, besides carbon monoxide, sulphur diox- ide and carbon dioxide are obtained; for mercuric cyanide is de- composed at the temperature of boiling sulphuric acid into mercury and cyanogen; and the former dissolves in the hot sulphuric acid, with formation of mercuric Sulphate and evolution of sulphur di- oxide: Hg(CN),4-3H,SO,--H,O=(NH,),SO,--CO+CO,--SO,--HgSO,. 3. Silver Nitrate.—If silver nitrate is added to a solution of an alkali cyanide drop by drop, a precipitate is formed on the addition of each drop, which, however, redissolves on stirring the liquid, for silver cyanide is soluble in an excess of alkali cyanide: (a) KCN+AgNO,-KNO,--AgCN; (b) AgCN+ KCN=[Ag(CN), K. The complex salt, silver potassium cyanide, is decomposed by further addition of silver nitrate, being finally completely trans- formed into insoluble silver cyanide: [Ag(CN), K+ AgNO,-KNO,--2AgCN. Consequently the precipitation is only complete when an excess of silver nitrate is added. Silver cyanide is insoluble in water and dilute nitric acid, per- ceptibly soluble in concentrated nitric acid, and readily soluble in ammonia, sodium thiosulphate, and potassium cyanide. Dilute nitric acid reprecipitates silver cyanide from the solution in ammo- nia or potassium cyanide. Concentrated hydrochloric acid decomposes silver cyanide, on warming, into silver chloride, with evolution of hydrocyanic acid (difference from silver chloride, bromide, or iodide). By igniting silver cyanide, cyanogen gas is obtained, with metal- lic silver, and brown, difficultly-volatile paracyanogen, which, on 272 REACTIONS OF THE METALLOIDS, further heating, is completely volatilized, leaving behind pure silver: 2AgCN = Ag,4-(CN), Mercuric cyanide is a non-electrolyte, and consequently behaves quite differently from the other cyanides. Silver nitrate gives no precipitation with mercuric cyanide, but combines with the latter, forming a soluble double salt, Hg(CN), AgNO,-2H,O. Mercuric cyanide is not decomposed by dilute sulphuric acid, but easily so by hydrochloric, hydrobromic, or hydriodic acids, with evolution of hydrocyanic acid, and formation of chloride, bromide, or iodide of mercury. Hydrogen Sulphide also decomposes mer- curic cyanide completely, forming mercuric sulphide and hydro- cyanic acid. Alkali hydroxides do not precipitate mercuric cyanide, for mer- curic oxide is soluble in alkali cyanides: HgC)+2KCN+H,O=Hg(CN), +2KOH. Mercuric oxide is also soluble in mercuric cyanide: Hgo-HgCN)-03; N. By means of the formation of Prussian blue or ferric sulpho- cyanate, hydrocyanic acid can be detected much better than by the silver nitrate reaction. 4. Prussian Blue Reaction.—Prussian blue is formed by the action of ferric Salts upon potassium ferrocyanide (cf. p. 100): 3K,Fe(CN)4-4FeCl,-12KCl- FeIFe(CN). In order, therefore, to apply this reaction to potassium cyanide, etc., it is necessary first to transform the cyanide into potassium ferro- cyanide. This may be accomplished by the addition of a ferrous salt, whereby ferrous cyanide is first formed, which dissolves in an excess of potassium cyanide, forming potassium ferrocyanide:* (a) FeSO,--2KCN = K,SO,-- Fe(CN),; (b) Fe(CN),4-4KCN=[Fe(CN)]K,. * When an organic substance is treated with metallic sodium as described on page 252, any nitrogen present is converted into Sodium cyanide and Prussian blue may be formed from it as described above. In case the organic substance is very volatile it should be dropped upon the molten sodium. HYDROCYANIC ACID. 273 Potassium ferrocyanide is formed even more readily by the action of potassium cyanide upon ferrous hydroxide: Fe(OH),4-2KCN=Fe(CN),4-2KOH; Fe(CN),4-4KCN=[Fe(CN), K. Evidently for the formation of the potassium ferrocyanide a little iron and considerable potassium cyanide are required. Con- sequently, to bring about the reaction, a little ferrous Sulphate is added to the alkaline Solution of an alkali cyanide, and the mix- ture boiled. A little hydrochloric acid is then added, whereby a clear solution is obtained, which gives, with a little ferric chloride, the blue precipitate. If only traces of hydrocyanic acid were pres- ent, the Solution appears green at first, but, after standing some time, “flocks” of Prussian blue will be precipitated. 5. The Ferric Sulphocyanate Reaction.—Potassium sulpho- cyanate produces a red coloration with a ferric salt, owing to the formation of soluble ferric sulphocyanate (cf. p. 99): 3KCNS+FeCl,-3KCl4-Fe(CNS). The cyanide, therefore, must be changed to sulphocyanate in order to apply this reaction, which can be done by heating with sulphur: KCN-H S = RCNS: or, better, by treatment with an alkaline polysulphide; RCN-I-(NH),S,-(NH,),S+ KCNS. To the concentrated solution of the cyanide (in a porcelain dish) a little yellow ammonium sulphide is added, and the mixture evap- orated on the water-bath to dryness. The residue is acidified with a little hydrochloric acid,” and a drop of ferric chloride is added, whereby characteristic blood-red coloration will be produced if only traces of cyanide are present. 6. Mercurous Nitrate produces a gray precipitate of metallic mercury when added to a solution of an alkali cyanide (difference from a chloride, bromide, or iodide): Hg,(NO),4-2KCN=2KNO,--Hg(CN),4-Hg. *It is necessary to acidify in order to destroy the (NH),S, which would precipitate black Fes with FeCls and the red coloration would not appear. 274 REACTIONS OF THE METALLOIDS. Behavior of Cyanides on Ignition. The cyanides of the alkalies and alkaline earths fuse without decomposition when heated out of contact with the air; heated in contact with air, they absorb oxygen with avidity, forming a cy- anate: RCN-H O = KCNO. Consequently the alkali cyanides are strong reducing agents (cf. p. 223). * The cyanides of the bivalent heavy metals are decomposed on ignition, out of contact with the air, into nitrogen, and metallic car- bide; the latter often being further decomposed into metal and carbon: Fe(CN), - FeC,--N,; Pb(CN),=Pb-i-C,--N,. The cyanides of the trivalent metals are unknown in the free state; those of the noble metals are decomposed, by igniting, into metal and dicyanogen: 2AgCN=Ag,4-(CN), ; Hg(CN),=IHg-i- (CN), It is a characteristic property of the cyanides of the heavy metals that they are readily Soluble in alkali cyanide solutions, forming very stable complex compounds, which are to be regarded as salts of the following acids: RI(CN), H, RH(CN), H, Rºn (CN),H, and Rn(CN), H. The first two of the above acids are so unstable that they are decomposed, as soon as they are set free, into hydrocyanic acid and cyanide: [R(CN), H=HCN4–RCN; [R(CN), H,-2HCN+R(CN), Consequently all cyanides which are derived from these acids evolve hydrocyanic acid when treated with dilute hydrochloric or sulphuric acid in the cold. Such compounds are: [Ag(CN), K, [Au(CN).]K, [Ni(CN), K, [Zn(CN), K, [CdCN), K.], etc. DICYANOGEN. 275 These salts must be regarded as complex compounds, for their aqueous solutions contain almost no heavy metal ions; they are not precipitated by caustic alkali, alkali carbonate, or ammonia. From this fact it follows that the oxides of these metals are soluble in cyanides of the alkalies, forming the following complex Salts: Ag,0+4KCN+H,O=2|KOH+2|Ag(CN), K; NiO +4KCN+H,0=2|KOH+[Ni(CN),JK,; ZnO +4KCN+H,0=2|KOH+[Zi(CN), K, CdO +4KCN+H,0=2|KOH+[CdCN), K. Hydrogen sulphide decomposes the silver and cadmium salts without difficulty, the zinc Salt slightly, and the nickel salt not at all. The acids of the general formula Rºſſ(CN),H, and RH(CN), H, are, in contrast with the above acids, quite stable in the free state, and can be usually obtained, without the evolution of hydrocyanic acid, by acidifying a solution of one of their salts with cold dilute mineral acid; but, on warming the acid Solution, hydrocyanic acid is given off. As typical representatives of these acids we have hydroferri- and hydroferrocyanic acids, and hydrocobalticyanic acid.* We will consider hydroferro- and hydroferricyanic acids sepa- rately; but before doing so we will briefly describe DICYANOGEN, which is obtained by heating the cyanides of the noble metals, as a colorless gas with a disagreeable odor; it burns with a reddish flame, and is soluble in water (25 parts water dissolve 100 parts of dicy- anogen). The aqueous Solution cannot be kept very long, as brownish “flocks” separate out little by little (azulmic acid) and the solution then contains ammonium cyanide, ammonium car- bonate, ammonium oxalate, and urea. Just as chlorine acts upon caustic alkalies, forming chloride and hypochlorite, so dicyanogen reacts with them, forming a cyanide and a cyanate: 2C1+2KOH=KC-H H.O.--KOCl; 2CN+2KOH=KCN+H,0+ KOCN. * Hydrocobaltocyanic acid is extremely unstable, like hydromangano- and hydromanganicyanic acids. Their salts evolve HCN when treated with cold, dilute, mineral acids. Y. * \ t \ * \ 276 REACTIONS OF THE METALLOIDS. HYDROFERRocyANIC ACID [Fe’’(CN).H. Hydroferrocyanic acid is a white, solid substance, which is readily soluble in water and in alcohol, the solution soon becoming blue on exposure to air. The Salts of this acid are much more stable than the acid itself, being all prepared from the potassium salt, the yellow prussiate of potash. This potassium salt, the most impor- tant ferrocyanide of commerce, is obtained by the fusion of organic substances containing nitrogen and Sulphur (blood, etc.) with potash and metallic iron, and by lixiviating the product of the fusion with water. In the melt, iron Sulphide and potassium cyanide are found, which, on treatment with water, are changed to potassium ferro- cyanide and potassium sulphide, FeS+6KCN=[Fe(CN), K,4-K.S, and, on evaporating the solution, the former Salt separates out (with three molecules of water of crystallization) in the form of large, yellow, tetragonal octahedrons. Recently this salt has been obtained as a by-product in the manufacture of illuminating-gas, Prussian blue and ammonium sulphocyanate being formed from the purification of the gas. The following equations will give Some idea of the formation of potassium ferrocyanide in the gas-house: 1. Fe,(CN),s--6Ca(OH),=4Fe(OH),4-3|Fe(CN).]Ca,; 2. [Fe(CN), Ca, H2KCl=[Fe(CN)]CaK,--CaCl, ; Very difficultly soluble 3. [Fe(CN)]CaK,--K,CO,-CaCO,--[Fe(CN)]K. Solubility of Ferrocyanides.—The ferrocyanides of the alkalies and alkaline earths are soluble in water; but the remaining salts dissolve with difficulty (if at all) in water and in cold dilute acids. REACTIONS IN THE WET WAY. 1. Dilute Sulphuric Acid.—The ferrocyanides are not decom- posed by cold dilute sulphuric acid, but break up at the boiling temperature with evolution of hydrocyanic acid: 2.Ée 2[Fe’’(CN), K,4-3H, SO,-[Fe" (CN), ek+ 3K,SO,--6HCN. |K HYDROFERROCYANIC ACID. 277 2. Concentrated Sulphuric Acid decomposes ferrocyanides completely, on warming, with evolution of carbon monoxide, which burns with a blue flame: * [Fe"(CN), K,4-6H,SO,4-6H,0– =FeSO,--2K,SO,--3(NH3),SO,--6CO. 3. Silver Nitrate produces a white precipitate of silver ferro- cyanide, [Fe"(CN)6]K.--4AgNO,-4KNO,--[Fe’’(CN)]Ag., insoluble in dilute nitric acid and ammonia, but soluble in potassium cyanide solution. On treatment with concentrated nitric acid, it is changed to orange silver ferricyanide, and is then soluble in ammo- nia. 4. Barium Chloride gives no precipitation. 5. Ferric Salts produce a precipitate of Prussian Blue in neu- tral or acid solutions (cf. page 100). * 6. Ferrous Salts yield a light blue precipitation, which changes to a darker blue on exposure to the air (cf. page 94). 7. Cupric and Uranyl Salts produce brown precipitates. In order to detect ferrocyanic acid in an insoluble ferrocyanide, the latter is boiled with caustic Soda or potash, when insoluble me- tallic hydroxide and an alkali ferrocyanide are usually formed. Thus, Prussian blue yields ferric hydroxide and potassium ferro- cyanide: Rºose: Fe 3. rºos). +12KOH=4Fe(OH),4-3|Fe’’(CN).]K,. e Fe’’(CN)-Fe The insoluble hydroxide is filtered off, the filtrate is acidified with dilute hydrochloric acid and treated with ferric chloride, whereby Prussian blue is again formed. *SO, is also liberated by this reaction, as a part of the ferrous Sulphate is oxidized by the sulphuric acid to ferric sulphate: 2FeSO, + 2H2SO, - Fe3(SO)3 + 2H2O + SO2. 278 REACTIONS OF THE METALLOIDS. Prussian blue is often used in wall-papers as a pigment. If it is desired to detect the presence of this compound in a wall-paper, about 100 cm”, of the paper are cut into small pieces, boiled with caustic potash solution, filtered, and the filtrate treated according to the process just described. In a few hours a distinct blue preci- pitate of Prussian blue will be visible in the bottom of the test-tube, if originally present. Some insoluble ferrocyanides do not yield the hydroxide of the metal on treatment with caustic alkali. Thus the brown uranyl ferrocyanide yields insoluble yellow potassium uranate and soluble potassium ferrocyanide (cf. page 108). Insoluble zinc ferrocyanide is completely soluble in caustic alkali, forming an alkali zincate and soluble ferrocyanide: [Fe"(CN), Zn, I-8KOH=[Fe"(CN), K,4-2Zn(OK),4-4H,O. In order to separate the zinc from the ferrocyanide, carbonic acid is conducted into the solution, the latter is boiled, and the insoluble zinc carbonate is filtered off. The filtrate then contains potassium ferrocyanide, which can be detected as above. 8. Lead Salts precipitate white lead ferrocyanide insoluble in dilute nitric acid. Detection of Hydrocyanic Acid in the Presence of Hydroferrocyanic Acid. As the soluble ferrocyanides evolve no hydrocyanic acid on treatment with cold dilute hydrochloric acid, while the soluble cyanides do, it is evident that hydrocyanic acid can be detected in the presence of hydroferrocyanic acid. For this purpose, the mix- ture of the solid salts is placed in a small porcelain evaporating-dish, a little dilute hydrochloric acid is added, the dish is covered with another porcelain dish whose inner surface is wet with yellow sul- phide, and the dishes are allowed to remain in this position for a few minutes. The volatile hydrocyanic acid produces ammonium cyanate in the upper dish. The latter is acidified with hydrochloric acid, and a drop of ferric chloride is added. A blood-red coloration of ferric sulphocyanate shows the presence of hydrocyanic acid in the origi- nal compound. HYDROFERRICYANIC ACID. 279 Behavior of Ferrocyanides on Ignition. On being ignited, the ferrocyanides yield iron carbide, cyanide, and nitrogen: [Fe(CN), K, H4KCN+FeC,--N,; [Fe(CN), Ag,-4AgCN+FeC,--N. In the latter case, the silver cyanide is further decomposed into metal and dicyanogen: £ 2AgCN=Ag,--(CN),. HYDROFERRICYANIC ACID, [Fe’”(CN), H, Hydroferricyanic acid forms brown needles, readily soluble in Water. Its salts, the ferricyanides, are very stable, and are obtained by the oacidation of the corresponding ferrocyanides. The most im- portant of all these Salts, potassium ferricyanide (red prussiate of potash), K. Fe(CN)], is obtained by the oxidation of potassium ferrocyanide with chlorine: [Fe"(CN).f34-Cl-KCl HFe"(CN).IK, Bromine, hydrogen peroxide, etc., may be used instead of chlorine. Solubility of Ferricyanides.—The ferricyanides of the alkalies and alkaline earths, and the ferric salt of hydroferricyanic acid, are soluble in water, but the remaining salts are insoluble even in dilute acids. REACTIONS IN THE WET WAY. 1. Dilute Sulphuric Acid evolves no hydrocyanic acid in the cold (difference from cyanides), but does so on warming with the acid. 2. Concentrated Sulphuric Acid decomposes all ferricyanides, on warming, with the formation of sulphates and carbon monoxide: 2[Fe"(CN), K.--12H,SO,--12H,O= =Fe,(SO.),4-3K,SO,--6(NH,),SO,--12CO. 28o REACTIONS OF THE METALLOIDS. 3. Silver Nitrate produces orange silver ferricyanide: [Fe’”(CN)]K.--3AgNO,-3KNO,--[Fe’”(CN), Ags, Soluble in ammonia, but insoluble in nitric acid. 4. Barium Chloride gives no precipitation. 5. Ferrous Salts produce, in neutral and acid solutions, a pre- cipitate of Turnbull’s blue (cf. page 94). 6. Ferric Salts produce no precipitation, but a brown colora- tion. 7. Cupric Salts yield green cupric ferricyanide: 2[Fe’”(CN), K.--3CuSO,-3K,SO,--[Fe’”(CN),Cu, 8. Behavior of the Ferricyanides in Alkaline Solutions.— Hydroferricyanic acid is a strong oxidizing agent in alkaline solu- tions, being readily reduced to hydroferrocyanic acid by hydrogen sulphide, hydriodic acid, Sulphurous acid, ferrous hydroxide, man- ganous hydroxide, lead oxide, starch, cellulose (paper), etc.; e.g.: 2[Fe’”(CN), K.--K.S=2|Fe’’(CN), K.--S; 2[Fe(CN)].K.--2KI=2|Fe"(CN), K,4-I, 2[Fe!”(CN), K.--KSO,--2KOH=2|Fe"(CN), K,4-K,SO,4-H,0; 2[Fe"(CN), K,4-PbO4-2KOH = 2 Fe"(CN), K,4-PbO,4-H,0; [Fe"(CN).IK,4-Fe(OH), H KOH=[Fe"(CN).IK,4-Fe(OH), , , , 3. The ferricyanides are even reduced by ammonia, forming nitro- gen: * - 12|Fe’”(CN), K, H-16NH,- =9|Fe"(CN)]K.--3|Fe"(CN), (NH),4-2N,. On account of this easy reducibility of hydroferricyanic acid, it is often difficult, Sometimes impossible, to detect its presence, par- ticularly in an insoluble compound. If Turnbull’s blue is boiled with caustic potash, the residue will consist of a mixture of ferrous and ferric hydroxides, and the solution will contain potassium fer- rocyanide. SULPHOCYANIC ACID. 281 The behavior of cyanides toward suspended yellow mercuric oxide is very important. Almost all cyanides, simple or complex, with the exception of potassium cobalticyanide, are completely decom- posed by this reagent. Mercuric cyanide and an oxide of the other metal are formed, which, if insoluble, may be separated from the mercuric cyanide by filtration. Thus potassium ferrocyanide is decomposed by mercuric oxide as follows: [Fe"(CN)]K,--3Hg0+3H,O=Fe(OH), F4KOH--3Hg(CN),. Prussian blue as follows: [Fe"(CN), Fe,”--9Hg0+9H,O= =3Fe(OH),4-4Fe(OH),4-9Hg(CN), This decomposition of the cyanides by mercuric oxide is often used in quantitative analysis for the separation of metallic cyanides. Behavior of the Ferricyanides on Ignition. The ferricyanides are decomposed into iron carbide, cyanide, dicyanogen, and nitrogen: 2[Fe’”(CN), K, H2Fe(C,4-6KCN+2N,+(CN),. By heating a ferricyanide in a closed tube, dicyanogen therefore is given off, which burns with a reddish flame. *: SULPHOCYANIC (THIOCYANIC) ACID, HCNS. Thiocyanic acid is found in Small amounts, in the form of its. sodium salt, in Saliva and urine. The free acid is a colorless, unstable liquid, with a penetrating odor. It can be kept better in aqueous solution than in the anhy- drous state, but its salts, the Sulphocyanates, are much more stable than the acid itself. The alkali Salts can be prepared from the corresponding cyanides by heating with sulphur: RCN-H S= KCNS. They may also be prepared at ordinary temperatures by treat- ing hydrocyanic acid or an alkali cyanide with an alkali polysul- phide: RCN--(NH,),S,-(NH,),S+ KCNS. 282 REACTIONS OF THE METALLOIDS. Solubility.—Most of the thiocyanates are soluble in water; ex- ceptions are the silver, mercury, copper, and gold salts. Lead thiocyanate is difficultly soluble in water; and on boiling with water it is decomposed. REACTIONS IN THE WET WAY. 1. Dilute Sulphuric Acid (double normal) causes no reaction. 2. Moderately Concentrated Sulphuric Acid (5H,SO :4H,O) decomposes the thiocyanates, with evolution of carbon oxysulphide, which burns with a blue flame: KCNS+2H.SO,--H,0=KHSO,--(NH,)HSO,--COS. 3. Concentrated Sulphuric Acid violently decomposes thio- cyanates, with evolution of very disagreeably-smelling vapors, COS, HCOOH, CO, SO, and deposition of sulphur. 4. Silver Nitrate precipitates white, curdy, silver sulpho- cyanate: KCNS+ AgNO,-KNO,--AgCNS, insoluble in dilute nitric acid, soluble in ammonia. 5. Ferric Salts produce a blood-red coloration, due to the for- mation of ferric sulphocyanate: 3KCNS+FeCl,-3KCl+Fe(CNS), very soluble in ether (cf. page 99). 6. Mercuric Nitrate precipitates white mercuric thiocyanate: Hg(NO,),4-2KCNS=2KNO,--Hg(CNS), very difficultly soluble in water, but readily soluble in an excess of potassium sulphocyanate: Hg(CNS),4-KCNS=Hg(CNS), K. If the dry powder is heated, the salt expands greatly (Pharaoh’s serpents). 7. Mercuric Chloride gives a precipitate only after long stand- ing. 8. Mercurous Nitrate produces a gray to black precipitate. On adding mercurous nitrate drop by drop to a fairly-concentrated SULPHOCYANIC ACID. 283 solution of potassium sulphocyanate, a gray precipitate of metallic mercury is first obtained, and the solution contains mercuric po- tassium sulphocyanate: Hg—NO, g | +3KCNS=2KNO,--[Hg(CNS), K+Hg. Hg—NO, If the addition of mercurous nitrate is continued until no more mercury is precipitated, and the solution then filtered, the filtrate will contain mercuric potassium sulphocyanate; but, on adding still more mercurous nitrate, pure white, mercurous sulphocyanate is precipitated: 2[Hg(CNS), K+3Hg,(NO,),=2KNO,--2Hg(NO.),4-3Hg,(CNS),. If, on the other hand, a very dilute solution of potassium sulpho- cyanate is added to a very dilute solution of mercurous nitrate, the white precipitate of mercurous sulphocyanate is obtained directly: Hg—NO, Hg—CNS | +2KCNS=2KNO,-- || Hg—NO, Hg—CNS. 9. Cupric Salts.-On adding a few drops of a solution contain- ing a cupric salt to one of an alkali sulphocyanate, the solution is colored emerald-green; and, on further addition of the copper Solu- tion, black cupric sulphocyanate is precipitated. If sulphurous acid is added, white cuprous Sulphocyanate is deposited, 2CuSO,--SO,--2H,0+2KCNS=Cu,(CNS),4-2KHSO,--H,SO, insoluble in dilute hydrochloric and sulphuric acids. 10. Cobalt Salts.-If a solution containing an alkali Sulpho- cyanate is treated with a small amount of a cobalt salt, and the solution shaken with a mixture of equal parts amyl alcohol and ether, the upper layer of alcohol ether separates out azure-blue in color (cf. page 137). This reaction is analogous to that of cyanic acid upon cobalt salts (cf. page 298). Behavior of Sulphocyanates on Ignition. The sulphocyanates of the alkalies melt readily, and are colored successively yellow, brown, green, and finally blue, becoming white again on cooling. 284 REACTIONS OF THE METALLOIDS. The sulphocyanates of the heavy metals are decomposed into sulphide, splitting off carbon disulphide, dicyanogen, and nitrogen. Thus cuprous thiocyanate is decomposed according to the following equation: CuSON 4 =4Cu,S+2CS,--3(CN), -N. CuSON The mercuric sulphocyanates swell tremendously on being heated (Pharaoh’s serpents). GROUP II. Silver Nitrate produces a precipitate soluble in nitric acid. Barium Chloride causes no precipitation. NITROUs ACID (HNO.). Occurrence.—Nitrous acid never occurs free in nature except in the form of its salts, the nitrites. It is found in the air, as ammo- nium nitrite, in many soils and waters, particularly in those which are contaminated with ammonia or decaying substances. Ammonia is oxidized by the action of micro-organisms (monas nitrificans) to nitrous acid, which combines with more ammonia to form ammonium nitrite. Preparation of Nitrous Acid and its Salts.-Nitrous acid is formed by the gentle reduction of nitric acid. If zinc is allowed to act upon dilute nitric acid for a short time, the latter is reduced to nitrous acid, * ... we HNO,--H,-H,0+HNO, but the reduction can easily go a little farther, forming NO, N,0, and N, ; while by long-continued action of the zinc, hydroxyl amine, NH,0H, and ammonia even, are formed. If nitric acid of sp. gr. 1.3 is heated with arsenious acid, starch etc., a mixture of nitric oxide and nitrogen peroxide is obtained, which, on cooling to — 21°C., condenses to a bluish-green liquid, N.O., the anhydride of nitrous acid. If the anhydride is treated with ice-cold water, a bluish-green liquid is obtained, which contains nitrous acid, but always in com- NITROUS ACID, 285 pany with nitric acid; for N.O. unites with water, forming nitrous and nitric acids, and nitric oxide: 2N,0,4-H,0=HNO,--HNO,--2NO. At a higher temperature nitrous acid is gradually changed into nitric acid: 3HNO,-HNO,4-2NO+H,O. Pure nitrous acid, therefore, is not known. If the above mixture of nitric oxide and nitrogen peroxide is con- ducted into concentrated Sulphuric acid, the two gases are readily absorbed, forming nitrosyl Sulphuric acid: /OH SO, NOH NO /OH =H,0+2SO, /OH Noo NO(NO) SO, sºsºsºsºms Nitrosyl sulphuric acid 2 This solution of nitrosyl Sulphuric acid in sulphuric acid is some- times called “nitrose.” If this Solution is added to cold water, sulphuric and nitrous acids are formed: /OH SO, +HOH=HNO,--SO,(OH), NO(NO) This solution of nitrosyl sulphuric acid can be kept indefinitely, so that it is a convenient reagent for the immediate production of nitrous acid at any time. The salts of nitrous acid, the nitrites, are much more stable than the free acid, and may be obtained by the ignition of nitrates: 2NaNO3=2|NaNO2 +O2. Nitrites prepared in this way always contain some oxide and some nitrate as impurity.” In order to obtain a pure nitrite, silver nitrite is treated with a calculated amount of a metallic chloride: * If the nitrate is heated with a metal, e.g., lead, the reduction takes place at a lower temperature and is almost quantitative. 236 REACTIONS OF THE METALLOIDS. The soluble nitrite can be separated from the insoluble silver chloride by filtration. Solubility of Nitrites.—All nitrites are soluble in water; but silver nitrite and potassium cobaltic nitrite are difficultly soluble. REACTIONS IN THE WET WAY. As all nitrites are soluble in water, the reactions which serve for their detection cannot be those of precipitation, but rather those in which a change of color takes place, owing to an oxidation or a re- duction. Nitrous acid sometimes acts as an oxidizing agent, and some- times as a reducing agent. 1. Dilute Sulphuric Acid decomposes all nitrites in the cold, setting free brown vapors: (a) NaNO,--H,SO,-NaHSO,--HNO,; (b) 3HNO,-HNO,--2NO+H,0; (c) NO+O (air)=NO, (brown gas). 2. Concentrated Sulphuric Acid reacts exactly the same, but much more violently. 3. Silver Nitrate precipitates from nitrite solutions crystals of silver nitrite in the form of needles, which are very difficultly soluble in cold water (300 parts of water dissolve one part of silver nitrite at the ordinary temperature). In boiling water, silver nitrite is considerably more Soluble. 4. Cobalt Salts produce (with an excess of potassium nitrite and acetic acid) a yellow crystalline precipitate of potassium cobaltic nitrite (cf. page 136). 5. Indigo Solution is completely decolorized by warming with nitrous acid. 6. Hydriodic Acid is oxidized by nitrous acid with separation of iodine: 2HI+2HNO,-2H,0+2NO+I, If, therefore, a nitrite is added to a solution of potassium iodide NITROUS ACID. 287 and the solution acidified with sulphuric or acetic acid,” the Solu- tion becomes yellow, owing to the separation of iodine. If the so- lution is now shaken with chloroform or carbon disulphide, the latter will be colored reddish-violet. Or if a little starch paste is added, it will be colored blue by the iodine. This exceedingly delicate reaction is also caused by the action of a great many other oxidizing agents; and it can only be used for the detection of nitrous acid when it is known that all such oxidiz- ing substances are absent. 7. Ferrous Salts are oxidized to ferric salts, with evolution of nitric oxide: 2FeSO,--H.SO,--2HNO,-2H,0+2NO+Fe,(SO.), The nitric oxide dissolves, in the cold, in the excess of ferrous salt, forming a brown compound of a varying composition: (FeSO.),(N O),. To perform this reaction a concentrated solution of ferrous sul- phate is taken, slightly acidified, and the Solution to be tested care- fully poured on top. At the Zone of contact between the two solu- tions the dark-brown coloration will be apparent. Nitric acid gives the same reaction, but only on the addition of concentrated sulphuric acid. 8. Potassium Permanganate.—If nitrous acid is added to a warm acid solution (about 40°C.) of potassium permanganate, the latter will become decolorized, owing to the oxidation of the former to nitric acid: 2KMnO,--5HNO,--3H,SO,-K,SO,--2MnSO,--5HNO,--3H,O. In this reaction nitrous acid acts as a reducing agent. 9. Detection of Small Amounts of Nitrous Acid by the Peter Griess Method (B. B. 12, 427).-As was already mentioned, nitrous * In the presence of considerable alkali acetate, there is no separation of iodine on the addition of acetic acid; but there is on the addition of a drop of a stronger acid. This is a beautiful illustration of the “Law of Mass Ac- tion.” The dissociation of the acetic acid is lessened by the presence of the alkali acetate (or ammonium acetate), so that not enough H ions are present to decompose the nitrite, although enough are present to redden litmus dis- tinctly. (Private communication from E. Bamberger.) 288 REACTIONS OF THE METALLOIDS. acid is often found in drinking-water, in consequence of the oxida- tion of ammonia which results from the decay of organic substances (urine, etc.). In order to detect the small amounts of nitrous acid which may be present in drinking-water, of the above reactions only that of potassium iodide and starch is delicate enough. But as hydrogen peroxide and ferric Salts are also likely to be present, both of which cause the separation of iodine from an acid solution of potassium iodide, it is evident that dependence upon this reaction alone would often lead to error. Consequently, in order to detect the presence of traces of nitrous acid we make use of a reaction which was first proposed by Peter Griess, and which is caused by nitrous acid only. It depends upon the formation of an intensely-colored azo dyestuff. Peter Griess used as his reagent phenylene diamine, whereby a yellow dye, Bismarck brown, is formed. Recently, according to the suggestion of Ilosvay v. Ilosva,” an acetic acid solution of sulpha- nilic acid and of a-naphthylamine is used instead. According to Lunge, f it is best to mix the Solutions of the last two reagents. The reagent is prepared as follows: 1. 0.5 gm. of sulphanilic acid is dissolved in 150 c.c. of dilute acetic acid. 2. 0.2 gm. of Solid a-naphthylamine is boiled with 20 c.c. of water, the colorless solution is poured off from the bluish-violet residue, and 150 c.c. of dilute acetic acid are added. The two solutions are then mixed and can be kept for a long time, provided the mixture is not brought in contact with free nitrous acid. If the solution should become reddish, it can be de- colorized by shaking with zinc dust and filtering. Procedure.—About 50 c.c. of the water to be tested are treated with 2 c.c. of the above reagent. The liquid is well stirred and allowed to stand five to ten minutes, when it will be colored a dis- tinct red if the slightest traces of nitrous acid are present. 10. Diphenylamine, dissolved in concentrated Sulphuric acid, is colored intensely blue by nitrous acid. Nitric acid and many other oxidizing substances, such as selenic acid, chloric acid, ferric chloride, etc., will give the same reaction (cf. Nitric Acid). * Bull. chim. [3] 2, p 317. f Zeitschrift für angew. Chemie, 1889, Heft 23. HYDROSULPHURIC ACID. 289 11. Brucine dissolved in concentrated sulphuric acid (accord- ing to G. Lunge and A. Lwoff +) gives no reddish coloration when treated with nitrosyl Sulphuric acid. Dry, several times recrystallized, silver nitrite, containing 70.05 per cent. silver (theory 70.09), did give with brucine, in an atmosphere of carbon dioxide, a weak but nevertheless distinct test for nitric acid, probably due to the presence of traces of nitrate re- maining in the silver nitrite. On dissolving 15 mgm. of this same nitrite in water, adding an equivalent amount of sodium chloride and diluting to one liter, a solution of sodium nitrite was obtained, of which one cubic centimeter added drop by drop with constant stirring to about four cubic centimeters of concentrated sulphuric acid yielded a solution of nitrosyl sulphuric acid, showing no sign of a red coloration with a drop of brucine reagent. The test was immediately obtained, however, on adding a trace of nitric acid to this solution. Brucine, therefore, is a reagent by which nitric acid can be detected in the presence of nitrous acid. For a description of the behavior of nitrous acid toward ammo- nuim salts and urea, see page 341. HYDRosULPHURIC ACID (HYDROGEN SULPHIDE), H.S. Occurrence and Preparation.—Hydrogen sulphide is found in volcanic regions, in many mineral waters (the so-called “sulphur’” waters), and, in general, wherever substances containing sulphur are subject to decay, or when they come in contact with decaying substances. Sulphates are easily changed into Sulphides by the action of micro-organisms which are present in the air; and this is the reason why many mineral waters containing sulphates smell of hydrogen Sulphide after standing some time in a corked flask. If, however, the flask and the cork are sterilized, the water can be kept indefinitely. The formation of hydrogen sulphide from sulphates takes place as follows: By means of carbonaceous matter (dust, etc.) the sulphates are reduced at first to Sulphides, Na,SO,--C,-2CO,--Na,S, * Zeitschrift für angew. Chemie, 1894, p. 345. 290 REACTIONS OF THE METALLOIDS. which are then decomposed by carbonic acid: Na,S+H,CO,-Na,CO,--H.S. Just as hydrogen sulphide is made from Sodium sulphate by the action of organic matter in the flask, so in nature the same process brings about the presence of hydrogen Sulphide in many mineral WaterS. For laboratory purposes, hydrogen sulphide is similarly pre- pared by the action of dilute sulphuric or hydrochloric acid upon a sulphide (usually iron Sulphide, on account of its cheapness and stability). Properties.—Hydrogen sulphide is a colorless gas, with an odor like that of rotten eggs; it is absorbed by water at the ordinary temperatures (one volume water absorbs two to three times its own volume). The aqueous solution reacts slightly acid toward litmus (toward methyl orange it has almost no action), and becomes tur- bid on standing in the air, owing to its oxidation by atmospheric OXygen: 2H2S +O2=2H2O +S2. Hydrogen sulphide burns in the air with a bluish flame to water and sulphur dioxide: 2H2S + OS =2|H2O +2SO2. The salts of hydrosulphuric acid are called sulphides. g Solubility of Sulphides.—The sulphides of the alkalies and the hydro- and polysulphides of the alkaline earths are soluble in water. The monosulphides of the alkaline earths, particularly calcium sul- phide, CaS, are very difficultly soluble in water, but they are gradu- ally changed from contact with water into soluble hydrosulphides: —OH —SH The remaining sulphides are insoluble in water. Of these latter FeS, MnS, and ZnS are decomposed by dilute hydrochloric acid with evolution of hydrogen sulphide; others require concentrated hydrochloric acid, e.g., Sb,S, SnS, PbS, NiS, CoS, CdS; while the remaining are insoluble in concentrated hydrochloric acid, but are all soluble in aqua regia with separation of sulphur. HYDROSULPHURIC ACID. 2.91. REACTIONS IN THE WET WAY. Free hydrogen sulphide, as has been stated, is a very weak acid, being even weaker than carbonic acid. In aqueous solution, there-- fore, it is very slightly dissociated into H and SH ions: + * H.Sa H+SH. The soluble neutral salts R,S are decomposed into metal and, sulphur ions, + + — — R,S a R+R+S, but under the influence of water, some of the bivalent sulphur ions. ~go over into the monovalent HS ions, + + — — —H - - + Rs. HOH = RSHEROH: but some bivalent Sulphur ions remain in solution, and, in fact, more in concentrated Solutions than in dilute ones. As, therefore, an aqueous solution of a sulphide contains both S. ions and SH ions, while the solutions of the free acid only contain. SH ions, it is plain why in many reactions the former react in a somewhat different way from the latter. 1. Dilute Sulphuric Acid decomposes all soluble, and some in- soluble, sulphides, with evolution of hydrogen sulphide. 2. Concentrated Sulphuric Acid decomposes all sulphides, on warming, with evolution of Sulphur dioxide and deposition of sul- phur: Na,S+2H,SO,-Na,SO,--2H,0+SO,--S. But even the sulphur goes over into sulphur dioxide after being heated with the sulphuric acid for Some time: 2H,SO,--S=2H,0+3SO,. 3. Silver Nitrate produces, in solutions of hydrogen sulphide and of soluble sulphides, a black precipitate of silver sulphide, 2AgNO,-i-H.S=Ag,S+2HNO, insoluble in cold nitric acid, in which, however, it dissolves on, warming. 292 REACTIONS OF THE METALLOIDS. 4. Barium Chloride causes no precipitation. 5, Lead Salts (best a solution containing an excess of alkali) produce a black precipitate of lead sulphide. All sulphides which are decomposed by hydrochloric acid evolve hydrogen sulphide, which, on coming in contact with a piece of filter-paper moistened with an alkaline lead solution, colors the latter black. An insoluble Sulphide (pyrite, arsenic sulphide, mercuric sulphide, etc.) evolves hydrogen sulphide with hydrochloric acid and nascent hydrogen. In this case some powdered zinc is placed in a small test-tube, a little of the insoluble sulphide added, then more zinc.; the mixture is next warmed with concentrated hydrochloric acid and the evolved gas tested for hydrogen sulphide with lead paper. Traces of hydro- gen sulphide can be detected with certainty by this reaction. 6. Sodium Nitroprusside, [Fe(CN),NO]Na,--2H,O, is colored reddish-violet by S ions, but not by SH ions. Consequently hy- drogen sulphide itself does not give this reaction, except upon the addition of caustic alkali. This reaction is very sensitive, but not so delicate as the one with an alkaline Solution of a lead salt. 7. Methylene Blue.—This reaction (which is recommended by Emil Fisher *) is the most sensitive of all reactions for detecting the presence of hydrogen sulphide. It is particularly suited for detect- ing the presence of traces of hydrogen sulphide in mineral waters, even when all other tests give negative results. The solution to be tested for hydrogen sulphide is treated with one tenth of its volume of concentrated hydrochloric acid, a little para-amido-dimethyl-aniline sulphate is added from the point of a knife-brade, stirred into the liquid, and as soon as it has dissolved, one or two drops of a dilute ferric chloride solution is added. If only 0.0000182 gm. of hydrogen sulphide in a liter is present, the blue color is distinctly apparent after half an hour's standing, while the above tests would afford negative results. 8. Oxidizing Agents, such as the halogens, nitric acid, chro. mates, permanganates, ferric Salts, etc., decompose hydrogen sul- phide with separation of sulphur. In order to detect the presence of sulphur in insoluble sulphides they may be fused with a little caustic soda (on the cover of a porce- lain crucible), when soluble sodium sulphide is formed: NiS+2NaOH = H,0+NiO--Na,S. * B. B. 16, 2234. SULPHUR. 293 Some sulphate is always formed by this treatment; but the aqueous Solution of the melt will always contain enough alkali sul- phide for any of the previous tests. 9. Metallic Silver is blackened by both free hydrogen sulphide and Soluble sulphides: & Ag,-i-H.S.--O (air) = H,0+ Ag,S; Ag,--Na,S+H,0+O (air) =2NaOH+ Ag, S. If oxygen and water are not present, the above reactions will not take place. A piece of bright silver suspended for fourteen hours in a Sulphur spring showed no sign of darkening until it had been ex- posed to the air for a short time Absolutely dry hydrogen Sulphide, in the presence of absolutely dry oxygen, acts upon silver at ordinary temperatures only very slowly; it acts instantly if a trace of water is present. Behavior of the Sulphides on Ignition. Most sulphides remain unchanged when heated out of contact with the air; arsenic and mercuric sulphides sublime. \ The polysulphides lose sulphur, which sublimes. The sulphides of gold and platinum lose sulphur, leaving the metal behind. All sulphides when heated in the air give off Sulphur dioxide, which can be recognized by its odor. The Detection of Sulphur in Non-electrolytes is usually accom- plished by the oxidation of the sulphur to sulphuric acid (which see); this can be done by fusion with an oxidizing flux (potassium carbo- nate and potassium nitrate). SULPHUR, S. Occurrence.—Sulphur is found native in volcanic regions in the form of orthorhombic pyramids, and in the neighborhood of sul- phur waters, being formed from the oxidation of some of the hydro- gen sulphide which escapes. Preparation and Properties.—Like the halogens, sulphur is formed by the oxidation of its hydrogen compound: 201H+O= H,0+Cl, SH,--O = H,0+S. 2.94 REACTIONS OF THE METALLOIDS. By heating polysulphides or the Sulphides of the noble metals (gold and platinum), Sulphur is also obtained. Sulphur exists in three allotropic modifications: 1. As Orthorhombic Sulphur, with a melting-point of 114° C., obtained by crystallization from Solutions. 2. As Monoclinic Sulphur, with a melting-point of 120° C., obtained by the solidification of molten sulphur. 3. As Amorphous Sulphur, obtained by quickly cooling the molten sulphur after it has become viscous by heating to 250° C., or after it has become a thin liquid after heating to a higher tem- perature. Monoclinic sulphur changes gradually into orthorhombic octa- hedrons; or, in other words, the unsymmetrical form changes into the symmetrical form. This is a general phenomenon: If a substance eacists in two or more crystallographic forms, the more symmetrical form is always the more stable, and the less sym- metrical always has a tendency to go over into the more symmetrical form. Thus the unsymmetrical, orthorhombic, yellow mercuric iodide is changed, by rubbing with a glass rod, into symmetrical, tetragonal, red mercuric iodide (cf. page 147); and, furthermore, the Orthorhombic form of calcium carbonate, aragonite, changes into hexagonal calcite. Both of the crystalline modifications of sulphur are soluble in carbon disulphide; and, by evaporating the solution, the sulphur always recrystallizes in the form of octahedrons. Amorphous sul- phur is insoluble in carbon disulphide. Commercial “flowers of Sulphur’ is a mixture of crystalline and amorphous Sulphur, and therefore is only partly soluble in carbon disulphide. Sulphur burns in the air to Sulphur dioxide, and, in the presence of “contact substances,” such as platinum, oxide of iron, chromic oxide, etc., it is burned to Sulphur trioxide also. Consequently the gas from pyrites burners always contains a mixture of the two gases. Sulphur is insoluble in water, but soluble in hot caustic alkali, forming a thiosulphate and a sulphide: 6NaOH-i-4S=3H,0+Na,S,O,--2Na,S.* * By further action of sulphur the Na,S is changed into Na,S, Na,S, etc. ACETIC ACID. 295 This action is entirely analogous to the formation of hypochlo- rite and chloride by the action of chlorine upon cold dilute caustic alkali: /* 2NaOH+Cl,-H,0+NaOCl.--NaCl. Sulphur dissolves on warming with an alkali sulphide, forming polysulphides: Na,S+S=Na,S, ; Na,S+4S=Na,Ss. Sulphur is also soluble in alkali sulphites, forming a thiosulphate: Na,SO,--S=Na,S,0s. AcETIC ACID, CH,COOH. * -ºr- Acetic acid is found in the Sap of many plants, partly free and partly in the form of its potassium or calcium salt. It is formed by the dry distillation of wood or by the oxidation of alcohol. Anhydrous acetic acid (glacial acetic acid) solidifies below + 16° C., forming colorless, glistening plates. It has a penetrating odor, similar to that of sulphur dioxide, and is miscible with water, alcohol, and ether in every proportion. It boils at 118°C. The aqueous Solution reacts acid. It is a monovalent acid, and its salts, the acetates, are as a rule readily soluble in water; the silver salt is difficultly soluble. The most important commercial salts of this acid are sodium acetate, and lead acetate (sugar of lead), Pb(C.H.O.),4-3H,0. Neutral lead acetate dissolves lead oxide with the formation of basic salts: Pb(C.H.O.), HPbO4-H,0–2PbTāo; 2++3\" 2 —O—Pb–C.H.O Pb(C.H.O.). --2PbO = Pb 2:3×2.5% (C.H.O.), H- —O—Pb–C.H.O, * The soluble basic salts are decomposed by considerable water, with precipitation of PbO; and, similarly, lead acetate yields a precipitate with considerable water, which is soluble in acetic acid. 296 REACTIONS OF THE METALLOIDS. REACTIONS IN THE WET WAY. A solution of sodium acetate should be used for the following reactions: *g 1. Dilute Sulphuric Acid sets acetic acid free from its salts; it is quite volatile and can be recognized by its odor. 2. Concentrated Sulphuric Acid also sets acetic acid free. If alcohol is added at the same time and the mixture warmed, ethyl acetate is formed, CH,COOH-CH.OH=H,0+CH,COOC.H.” Ethylacetate which can be recognized by its pleasant, fruity odor. 3. Silver Nitrate produces, in fairly concentrated solutions, a white crystalline precipitate of silver acetate (100 parts of water dissolve 1.04 parts at 20° C. and 2.52 parts at 80°C.). 4. Ferric Chloride colors neutral solutions dark brown, and by boiling the diluted Solution, all the iron is precipitated as basic acetate (cf. page 97). REACTIONS IN THE DRY WAY. All acetates are decomposed on ignition, leaving behind either the carbonate, oxide, or the metal itself, and with the evolution of combustible vapors and gases. The acetates of the alkalies are decomposed into carbonate and ºne: / Ma "2" §§-Nacot CH-20-GH, The acetates of the alkaline earths always leave the metal in the form of its oxide, while the acetates of the noble metals leave a resi- due of the metal itself. Cacodyl Reaction.—If a dry acetate (best an alkali acetate) is heated with arsenic trioxide, a very repulsive-smelling and ex- tremely poisonous gas is formed, called cacodyl oxide: 4CH,COONa+As,O,-2Na,CO,-- ÉXAs—o–A-3;+ 2CO,. 3 3 * Without the sulphuric acid or other dehydrating agent the reaction fakes place scarcely at all. CYANIC ACID. 297 In spite of the sensitiveness of this reaction, it cannot always be relied on, for many other organic acids, such as butyric and valeri- anic acids, give similar reactions. CYANIC ACID, HCNO. This very unstable acid is obtained by heating its polymer, (HCNO), ; it is a colorless liquid with a very penetrating, dis- agreeable odor, which immediately decomposes in aqueous solution into carbonic acid and ammonia, HCNO+H,0=CO,--NH, and these two products unite, forming monoammonium carbonate: CO,--NH,--H,O=NH, HCO,. The salts of cyanic acid, the cyanates, are much more stable than the free acid, and may be obtained by the oxidation of cya- nides. By simply fusing potassium cyanide in the air, perceptible quantities of potassium cyanate are formed. If, however, potassium cyanide is heated with oxidizing substances, or those which can be readily reduced, it is easy to change the cyanide completely over to cyanate. The cyanates of the alkalies are stable in the dry state, but take on moisture from the air and are gradually changed into monoalkali carbonate and ammonia: RCNO+2H,0=KHCO,--NH,. #. Solubility of Cyanates.—The cyanates of the alkalies and alkaline earths are soluble in water. Silver, mercurous, lead, and copper cyanates are insoluble in water. All cyanates are soluble in nitric acid. REACTIONS IN THE WET WAY. A freshly-prepared, cold Solution of potassium cyanate should be used for these reactions. 1. Dilute Sulphuric Acid immediately sets cyanic acid free, which decomposes (as given above) into carbonic acid and ammonia. Consequently a strong evolution of carbon dioxide takes place on adding the Sulphuric acid. The carbon dioxide always contains 298 REACTIONS OF THE METALLOIDS. small amounts of undecomposed cyanic acid, which is recognizable by its very penetrating odor. The solution will contain ammonium sulphate; so that, if it is warmed with caustic soda, ammonia will be given off. 2. Concentrated Sulphuric Acid reacts similarly. 3. Silver Nitrate precipitates white, curdy silver cyanate, RCNO+AgNO,-KNO,--AgCNO, soluble in ammonia and in nitric acid (difference from silver cya- nide). 4. Barium Chloride produces no precipitation. 5. Cobalt Acetate is colored azure blue by a solution of potas- sium cyanate. The blue potassium cobaltocyanate, CoſCNO],K, discovered by Blomstrand,” is formed by this reaction, and is obtained in the form of tetragonal crystals of a dark azure blue color. This blue compound dissolves in water with a blue color. If, however, the blue compound is subjected to the action of consider- able water, the color disappears, the double salt being dissociated into its components: Co(CNO),K, a Co(CNO),4-2KCNO. If more potassium cyanate is added to the solution which has become colorless, the blue color reappears. The same result is reached by adding alcohol. Almost all commercial potassium cyanide contains some cya- nate. In order to detect the presence of cyanate in the commercial salt, the cyanide must first be expelled, for the cobalt test does not take place in the presence of cyanide. According to E. A. Schneider,t the following process may be used: Three to five gm. of the potassium cyanide to be tested is dis- solved in 30 to 50 c.c. of cold water, and carbon dioxide passed into the solution for one to one and one half hours; the hydrocyanic * Journal für Praktische Chemie [2], 3, 206. f B. B., 1895, p. 1540. HYPOPHOSPHOROUS ACID. 299 acid is expelled, and monopotassium carbonate is formed, while the potassium cyanate is not affected perceptibly: KGN-i-H,CO,-HCN4 KHCO,. One c.c. of the solution is now treated with 25 c.c. of absolute alcohol (in order to precipitate potassium carbonate) and the pre- cipitate is filtered off. The alcoholic filtrate is treated with a few drops of acetic acid and then with a few drops of alcoholic cobalt acetate Solution. If the original cyanide contained 0.5 per cent of potassium cya- nate, the blue color can be distinctly recognized.* ^: HYPoPHosphorous AcID, PTH NOH The monobasic hypophosphorous acid is obtained by the de- composition of its barium salt with Sulphuric, or of its calcium salt with oxalic, acid. The salts of hypophosphorous acid are obtained by boiling phosphorus with dilute alkali, whereby phosphine is given off: 2P,4-3Ba(OH),4-6H.O=3Ba(PH,O.), H-2PH, ; P,4-3KOH+3H,0=3KPH,0,--PH,. Solubility of Hypophosphites.—All hypophosphites are soluble in water. REACTIONS IN THE WET WAY. 1. Dilute Sulphuric Acid.—No reaction. 2. Concentrated Sulphuric Acid reacts with hypophosphites only on warming, and is reduced to sulphur dioxide, which can be recognized by its odor. 3. Silver Nitrate is reduced to metallic silver, sometimes with and sometimes without the evolution of hydrogen, according to the relative amounts of the substances reacting: 2NaPI2FO2 +2AgNO3 +4H2O=2H3PO4-i-2NaNO3 +Ag2 +3H2; NaH2PO2+4AgNO3 +2H2O=H3PO4+NaNO3 +2Ag2+3HNO3. * Experiments made in Zürich by P. Rieder confirm this statement, but the sulphocyanates of the alkalies give the same reaction. 3oo REACTIONS OF THE METALLOIDS. 4. Barium Chloride causes no precipitation. 5. Copper,” Mercury, and Gold Salts are reduced to metal. 6. Concentrated Caustic Potash.-By boiling with concen- trated caustic alkali, the hypophosphites are oxidized, with evolu- tion of hydrogen, to phosphates: PH,O.K+2KOH=K, PO,--2H, 7. Nascent Hydrogen (zinc and dilute sulphuric acid) reduces hypophosphorous acid to phosphine (see Phosphorous Acid). REACTIONS IN THE DRY WAY. By ignition phosphate and phosphine are obtained: 2H,PO, H=H,PO,--PH, 4PH,O,Na = Na, P.O.--H,0+2PH,; 2(PH,O.),Ca-Ca,P,O,--H,0+2PH,. GROUP III. Silver Nitrate produces a white precipitate, soluble in nitric acid. Barium Chloride does the same. SULPHUROUS ACID, H2SO3. Occurrence and Preparation.—Sulphur dioxide, the anhydride of sulphurous acid, is found in the exhalations of active volcanoes, and is formed by the combustion of Sulphur or sulphides in the air, S+O, =SO, 2FeS,--11O=Fe,O,--4SO, or by the reduction of sulphuric acid on heating with sulphur, sul- phides, carbon, organic substances, and metals: 2H,SO,--S =2H,0+3SO, ; 2H.SO,--C =2H,0+CO,--2SO,; 2H,SO,--Cu-2H,0+CuSO,--SO,. * With copper the reduction can go so far that copper hydride is formed. Cf. Würtz, Compt. rend. 18, 102. SULPHUROUS ACID. 3OI Mercury, silver, tin, etc., act the same as copper. Sulphur dioxide is also formed by the decomposition of sulphites and thiosulphates with stronger acids: Na,SO,--H,SO,-Na,SO,--H,0+SO,; Naso,4-H so.--Naso,4-S-H,0+so, Sulphurous acid may be prepared for laboratory purposes by placing a concentrated solution of sodium bisulphite in a flask and allowing concentrated sulphuric acid to drop upon it. A steady stream of sulphur dioxide will be evolved without warming. Sulphur dioxide is a colorless gas, having the penetrating odor peculiar to burning sulphur, and is readily soluble in water and alcohol: one vol. water at 15° C. dissolves 45.36 vol. SO2; one vol. alcohol at 15° C. dissolves 116 vol. SO2. The aqueous solution contains sulphurous acid, H.SO,. If we attempt to isolate the acid, it decomposes immediately into water and sulphur dioxide; consequently the free acid is known only in aqueous solution. By neutralization of this solution with alkali hydroxides or carbonates, the comparatively stable sulphites are obtained. In solution, sulphites are gradually changed into sul- phates. Solubility of Sulphites.—The sulphites of the alkalies are readily soluble in water; the remaining Sulphites are difficultly soluble or insoluble in water, but are all soluble in hydrochloric acid. REACTIONS IN THE WET WAY. 1. Dilute Sulphuric Acid evolves sulphur dioxide, in the cold, from all sulphides, the gas being easily recognized by its odor. 2. Concentrated Sulphuric Acid reacts in the same way, but much more energetically. 3. Silver Nitrate produces, in neutral solutions of sulphites or in an aqueous solution of sulphurous acid, a white crystalline pre- cipitate of silver sulphite, Na,SO,--2AgNO,-2NaNO,--Ag,SO, soluble in an excess of the precipitant, forming silver sodium sul- phite: Ag,SO,--Na,SO,-2NaAgSO,. 3O2 REACTIONS' OF THE METALLOIDS. By boiling this solution the silver is precipitated as a gray metal: sº ºr sº sº sº ºr m sº mº m * * * * * * * * **** =Na,SO4+SO,--Ag,. ... = * * * * * * * s * * * * s a sm ºm e ºs * * * * * If the precipitate of silver sulphite suspended in water is boiled, half the silver is reduced to metal, while the other half goes into solution of sulphate: iAg,SO, AgSO, O = Ag,SO4+ SO,-- Age. Silver sulphite is soluble also in ammonia and in nitric acid. 4. Barium Chloride produces no precipitation in an aqueous solution of sulphurous acid, but in neutral sulphite solutions white barium sulphite is precipitated, Na,SO,--BaCl,-2NaCl- BaSO, readily soluble in cold dilute nitric acid. By boiling the solution, barium sulphate is formed, which separates out. As the sulphites in aqueous solution are gradually changed to Sulphates, commercial sulphites often contain sulphate. In this case the precipitate pro- duced by barium chloride contains barium Sulphate, which is insolu- ble in dilute nitric or hydrochloric acid. If the barium sulphate is filtered off and the filtrate treated with chlorine or bromine water, a white precipitate of barium sulphate is formed, provided a sul- phite was originally present: H,SO,--BaCl, -H,0+Cl,-4HCl·H BaSO,. 5. Strontium and Calcium Salts behave similarly to the barium salt. The sulphites of the alkaline earths vary in their solubilities in sulphurous acid and in water. Calcium sulphite readily dissolves in an excess of sulphurous acid, forming monocalcium sulphite: CaSO,--H,SO, =CaFI,(SO.), SULPHUROUS ACID. 3O3 On boiling this solution, sulphur dioxide escapes, and calcium sulphite is reprecipitated. The strontium salt also dissolves in sulphurous acid, but more difficultly; the barium salt is practically insoluble in Sulphurous acid. Solubility of THE ALKALINE-EARTH SULPHITES IN WATER. One part calcium sulphite dissolves in 800 parts water at 18°C. “ “ strontium. “ & & “ 30,000 “ “ “ 18° C. & ſº barium & & & & & & 46,000 ( & {{ ( & 18° C. é & Advantage is taken of the difficult solubility of strontium sul- phite in detecting sulphurous acid in the presence of thiosulphuric acid (which see). 6. Lead Salts precipitate white lead sulphite, soluble in cold dilute nitric acid; but on boiling the solution lead sulphate is precipitated. º 7. Sodium Nitroprusside and Zinc Sulphate.—If a neutral sulphite solution is treated with a dilute solution of sodium nitro- prusside, a faint pink coloration is produced. If, however, consid- erable zinc sulphate is added, the coloration becomes a distinct red. The reaction is still more sensitive if a little potassium ferrocyanide is added, a red precipitate being formed (difference from thiosul- phuric acid). This reaction, although very delicate, is not so reli- able as the precipitation with strontium chloride. Sulphurous acid is a strong reducing agent. 8. Iodine Solutions are decolorized by Sulphurous acid: H.SO,--H,0+I, +, 2III+H,SO,. 9. Acid Potassium Permanganate Solutions are also decolor- ized, sulphuric and dithionic acids being formed in varying amounts according to the temperature and concentration. Under certain conditions the reaction can take place according to the following equations: 2KMnO,4-2H,0+6SO,-2KHSO,--2MnSO,--H.S,O. Dithionic acid Under other conditions, however, the Sulphurous acid can be completely oxidized to Sulphuric acid. Consequently, sulphurous 3O4 REACTIONS OF THE METALLOIDS. acid cannot be determined quantitatively by means of potassium permanganate. 10. Chromic Acid is reduced to green chromic salt: 2CrO,--3SO,-3SO,--Cr,O,-[Cr,(SO.),). 11. Mercuric Chloride is unaffected by sulphurous acid at ordinary temperatures; but, on boiling, it is reduced to mercurous chloride, I 2HgCl, -H,SO,--H,O=2HCl·HH,SO,--Hg,Cl, and on adding more Sulphurous acid, the mercurous salt is reduced to gray metal. 12. Mercurous Nitrate is immediately acted upon by free sul- phurous acid and by alkali Sulphite solutions, with the formation of a black precipitate. 13. Gold Solutions are also reduced. 14. Nascent Hydrogen reduces sulphurous acid to hydrogen sulphide, which can be recognized by its odor and by its turning lead acetate paper black. The reduction is best effected with zinc and dilute hydrochloric acid. REACTIONS IN THE DRY WAY. The sulphites of the alkalies, when heated out of contact with the air, are changed to Sulphate and Sulphide: Na,SO, Na,SO, Na,SO, Na,SIOOO =3Na,SO,--Na,S. By heating an alkali sulphite in the closed tube this reaction takes place, and there is no sublimate of Sulphur (difference from thiosulphates). If the melt is treated with hydrochloric acid after cooling, hydrogen sulphide is given off freely. The remaining sulphites are changed, on being heated out of contact with the air, into sulphur dioxide and oxide of the metal: CaSO,-CaO +SO,; Ag,SO,-Ag,--SO,--O CARBONIC ACID. 305 If any sulphite is heated with sodium carbonate on charcoal, sodium sulphide is formed. If the melt is placed upon a bright silver coin and moistened with water, the silver is blackened, owing to the formation of black silver sulphide (Hepar reaction): 2Na,SO,--3C=3CO,--2Na,S, and Na,S+Ag,4-H,0+O *=2NaOH+Ag,S. This Hepar reaction takes place with all sulphur compounds, and therefore shows simply the presence of sulphur. /OH CARBONIC ACID, C+O NO Like sulphurous acid, pure carbonic acid does not exist; being known only in aqueous solution, which is decomposed, on boiling, into its anhydride, a gas which escapes from the hot solution. This anhydride, CO, is formed by the combustion of carbon and of carbonaceous matter of all kinds, and is found therefore very widely distributed in nature (in small amounts in the atmosphere, f and in enormous amounts in volcanic regions, streaming out from fissures in the earth). Carbon dioxide occurs also in many mineral waters, and (in the liquid state) is found enclosed in quartz, feld- spar, etc. As carbonate it exists in enormous quantities as lime- stone, marble, aragonite, dolomite, etc. Carbon dioxide is a color- less, odorless, slightly acid-tasting gas, with sp. gr. 1.52. Being, therefore, one and one half times as heavy as air, it can be poured from one vessel into another. It is comparatively soluble in water; one volume of water dissolves at 0° C. almost twice its volume of carbon dioxide; at ordinary temperatures it dissolves its own volume of gas. Carbon dioxide does not support combustion; a burning candle goes out in air containing 8–10 per cent. Of this gas.j. * Atmospheric oxygen. f Pure air contains 0.35–0.40 per cent. of CO2. In dwelling places the amount increases considerably, owing to breathing and other forms of com- bustion. If 3–4 per cent. is present, as is the case in mines sometimes, breath- ing becomes difficult, and the miner's lamp begins to burn faintly; while when 8–10 per cent. is present the lamp goes out. † Carbon dioxide not only fails to support combustion, but it tends to prevent it; hence its use in fire-extinguishers. Being formed by the com- bustion of carbon, the mass-action principle shows why this should be the case.—Translatcr. 306 REACTIONS OF THE METALLOIDS. In aqueous solution carbonic acid reacts acid; but its con- ductivity is extremely small, it being slightly dissociated. It is not decomposed into H and CO, ions, but according to the following equation, + *- H,CO, a H+ HCO, into H and monovalent HCO, ions, and this only to a slight extent; by increasing, therefore, the concentration of the H ions (addition of a stronger acid) the carbonic acid is scarcely dissociated at all. If the solution then becomes supersaturated with H,CO, the ex- cess is decomposed into water and carbon dioxide, and the latter escapes from the Solution, The salts of carbonic acid, the carbonates, are formed: 1. By passing carbon dioxide gas into a Solution of a metallic hydroxide: 2NaOH+CO,-H,0+ Na,CO,; Ba(OH),4-CO,-H,0+ BaCO, . 2. By the action of carbon dioxide upon cyanides, Sulphides, and borates of the alkalies and alkaline earths. 3. By the ignition of Salts of organic acids (cf. page 39). An illustration of the preparation of large amounts of carbonate is the production of potash by burning parts of plants (wood, for example, or the residue from the manufacture of beet Sugar, the latter being particularly rich in potassium salts). Solubility of Carbonates.—Of the normal carbonates, only those of the alkalies are soluble in water; and their aqueous Solutions react alkaline, owing to hydrolytic decomposition: + + — — —H - + + * - Na,CO,--HOH = Na+Na+OH-i-HCO,. The aqueous solution of the carbonates of the alkalies, therefore, behaves as if it were a solution of caustic alkali and monoalkali car- bonate (or bicarbonate):. + + — – H - + — + - Na,CO,--HOH = NaOH+Na(HCO). Many carbonates dissolve in an excess of carbonic acid. forming bicarbonates, particularly the alkaline-earth carbonates: CaCO,--H,CO,-CaFI,(CO.), CARBONIC ACID. 3o 7 By boiling a solution of calcium bicarbonate, the latter is de- composed into water and carbon dioxide, so that calcium carbo- nate is reprecipitated: CaFL(CO),=H,0+CO,--CaCO,. Almost every sample of drinking-water contains calcium or magnesium bicarbonate; they become turbid, therefore, on boiling (boiler scale). Dilute, cold mineral acids decompose all carbonates with effervescence (due to evolution of carbon dioxide gas). The native carbonates of magnesium and iron (magnesite, siderite, and dolomite) do not effervesce if a lump of the mineral is treated with cold dilute mineral acids, but when reduced to a fine powder they are more readily acted upon; on warming, all carbo- nates dissolve readily. REACTIONS IN THE WET WAY. 1. Dilute Sulphuric Acid decomposes all carbonates with effervescence, with the exception of magnesite, siderite, and dolo- mite; the reaction takes place in the cold. The experiment is best performed by placing a little of the substance in a test-tube, covering it with dilute sulphuric acid, and warming. As the carbon dioxide is heavier than air, the gas can be poured into a second test-tube containing a little barium hydroxide or lime water. If now the second test-tube is shaken, a turbidity due to the formation of cal- cium or barium carbonate will be formed, provided a carbonate was originally present. A blank test should always be made by pouring some barium hydroxide into a test-tube and shaking it. A turbidity of the latter solution can always be detected, owing to the presence of carbon dioxide in the air. If the turbidity produced in the first case is stronger than that formed in the blank test, carbon dioxide was given off from the substance tested; but if no distinct difference can be detected, the experiment should be repeated in an atmosphere free from carbon dioxide. To accomplish this, the test-tube containing the original sub- stance is provided with a rubber stopper containing two holes. Through one of the holes the stem of a small separatory funnel is inserted, through which the Sulphuric acid can be added; in the 308 REACTIONS OF THE METALLOIDS, other hole a right-angled glass tube is introduced, so that the end of the tube is level with the lower surface of the rubber stopper. This tube should be loosely filled with cotton wool,” and connected (by means of rubber tubing) with a second test-tube also provided with a two-holed rubber stopper. Through one of the latter holes a right-angled tube is introduced, extending to the bottom of the tube; in the other hole is placed the stem of a funnel, which contains a folded filter-paper. A stream of air free from carbon dioxide is now conducted through the apparatus for half an hour, the air passing through a Solution of caustic potash and then through a tube filled with Soda lime before reaching the apparatus. A little barium hydroxide solution is then filtered into the second test-tube, and a little more of the purified air passed through the apparatus in Order to make sure that the latter is absolutely free from carbon dioxide (the baryta water should remain clear). Dilute sulphuric acid is now added to the substance, and the purified air is conducted through the apparatus very slowly (two to three bubbles a second). If the solution of barium hydroxide becomes turbid, the presence of a carbonate is assured. 2. Concentrated Sulphuric Acid reacts in the same way as dilute sulphuric acid, only more violently. 3. Silver Nitrate precipitates white silver carbonate, which becomes yellow on the addition of an excess of the reagent. On boiling with considerable water, the carbonate is partly decom- posed into brown silver oxide and carbon dioxide; but the carbonic acid is not completely set free except on heating to 200° C. Silver carbonate is very soluble in ammonia and in nitric acid. 4. Barium Chloride precipitates white, voluminous barium carbonate, in the cold, which gradually on standing, but more quickly on warming, becomes crystalline and denser. Behavior of Carbonates on Ignition. The carbonates of the alkalies melt, with but slight decom- position. Barium carbonate is not decomposed by the blowpipe when placed on charcoal and does not melt; only at a white heat * The wad of cotton wool serves as a filter, preventing any Sulphuric acid from being carried over into the second test-tube, which would itself cause a turbidity if it came in contact with the barium hydroxide solution. BORIC ACID. 309 is it decomposed into infusible barium oxide and carbon dioxide. All remaining carbonates are decomposed at the temperature of the blowpipe into oxide and carbon dioxide. The oxides of the noble metals are further decomposed into metal and Oxygen. 29; BORIC (BoFACIC) ACID, Bé-OH.* NOH Occurrence.—Boric acid is found native as Sassolite; in the form of its sodium salt, as borax or tinkal;f as boracite, 2Mg, BAOls--MgCl, and in many silicates, such as axinite, tourmaline, datolite, etc. Crystallized boric acid forms colorless plates, with a mother-of- pearl lustre, which are soluble in water (100 parts water dissolve four parts of boric acid at 15° C., and 33 parts at 100° C.). The aqueous solution reacts acid, and is a poor conductor of electricity. By heating boric acid to 100° C., it loses one molecule of water and is changed to metaboric acid, HBO,. The latter loses more water when heated to 160° C., forming pyroboric acid, H.B.O.; which, on ignition, loses all its water, being changed to the anhy- dride of boric acid, boron trioxide, which remains as a difficultly volatile hydroscopic glass. The salts of boric acid, the borates, are derived from the meta- and pyroboric acids. The salts of the ortho acid, H.BO, are not known in the pure state. Solubility of Borates.—The borates of the alkalies dissolve in water, and the solution reacts alkaline. A concentrated solution of borax behaves as if it contained sodium metaborate, free boric acid, and a small amount of caustic alkali: Na,B,0,--3H,O a 2NaI3O,--2H,BO, ; NaBO,4-2HOH = NaOH+H,BO,. The more dilute the Solution, the greater the extent to which the hydrolysis represented by the second equation will take place; so that a very dilute solution of borax will react as if it contained simply sodium hydroxide and free boric acid. * In a few exceptional cases boron acts as a metal, forming B(HSO), (TO),SO, BPO., etc. The last compound is insoluble in water and dilute acids, but dissolves readily in caustic alkalies. f Na,C,0,.103,0. 3 Io REACTIONS OF THE METALLOIDS. A solution of an alkali borate will behave differently towards reagents, therefore, according to its concentration and tempera- ture. The remaining borates are difficultly soluble in water, but readily Soluble in acids and in ammonium chloride solution. REACTIONS IN THE WET WAY. For these reactions a solution of borax should be used. 1. Dilute Sulphuric Acid, -No reaction. 2. Concentrated Sulphuric Acid.—No visible reaction. Most borates are decomposed by Sulphuric acid, setting free boric acid, and the latter is capable of coloring the non-luminous gas-flame with a characteristic green tinge. If, therefore, a little Solid borate is placed in the loop of a plati- num wire, moistened with concentrated Sulphuric acid, and heated at the edge of the Bunsen flame, the characteristic green coloration will be noticed. A great many natural silicates containing boric acid, when tested in the above manner, will not give this flame coloration. To pro- duce this coloration the mineral is mixed with calcium fluoride and potassium acid Sulphate, placed in the loop of a platinum wire, and heated on the outer edge of the flame, when the latter will be colored distinctly green, owing to the formation of volatile boron fluoride. 3. Concentrated Sulphuric Acid and Alcohol.—If an alkali or alkaline-earth borate is treated in a platinum crucible with alcohol (best methyl alcohol), then with concentrated sulphuric acid, the mixture stirred and the alcohol lighted, a green-bordered flame will appear, due to the formation of boric acid methyl (ethyl) ether, B(OCH.), 4. Silver Nitrate produces, in moderately concentrated cold borax solutions, a white precipitate of silver metaborate: Na,B,C),4-3H,0+2AgNO,-2NaNO,--2H,BO,--2AgBO,. On warming, a brown precipitate of silver oxide is obtained: 2AgBO,--3H,0=2H,BO,--Ag,O. From very dilute solutions, in the cold, silver nitrate produces a brown precipitate of silver oxide. Silver borate is soluble in ammonia and in nitric acid. BORIC ACID. 3 II 5. Barium Chloride produces, in fairly concentrated solutions, a white precipitate of barium metaborate: Na,B,C),4-BaCl, -3H,0=2NaCl--2H, BO,--Ba(BO.), soluble in an excess of barium chloride and in ammonium chloride. 6. Calcium and Lead Salts behave similarly to barium chloride. 7. Turmeric.—If a piece of turmeric paper is placed in a solu- tion of free boric acid, apparently no change will take place, but if the paper is dried, it becomes reddish brown. If the brown paper is again dipped in the Solution of boric acid, the color remains; which is also true if the paper is dipped in a dilute sulphuric or hydrochloric acid solution (difference from the alkali test with turmeric paper). If the reddish-brown paper is moistened with caustic soda or potash solution, the paper becomes bluish black; or, if only a small amount of boric acid is present, grayish blue. Borate solutions, when acidi- fied with dilute hydrochloric acid, give the same reaction. This sensitive and convenient test for boric acid must be used with caution, for acid solutions of zirconic, titanic, tantalic, niobic, and molybdic acids also color turmeric paper brown. The reaction is much more sensitive if, instead of using the turmeric paper itself, an alcoholic extract is made from a few pieces of turmuric root. Two or three drops of the yellow solution thus obtained are placed in a porcelain dish, the Solution to be tested for boric acid is added, acidified with acetic acid, and evapo- rated to dryness on the water-bath. If as much as Tào mg. of B.O. is present, the residue is colored a distinct reddish brown, while Túrr mg. suffices to cause a visible reaction.* 8. Mercuric Chloride produces a red precipitate. Behavior of Borates on Ignition. The hydrated borates of the alkalies melt with effervescence, forming a colorless glass. This glass has the property of dissolving many metallic oxides when heated, whereby the often very characteristically-colored metaborates are formed (borax beads); thus copper oxide is dis- solved, forming a blue glass: Na,B,0,--CuO=2NaBO,--Cu(BO.), * A. Henz. 3I 2 REACTIONS OF THE METALLOIDS. If this bead is heated in the reducing flame (i.e., with carbon) two things can happen: (a) The colored cupric salt is reduced to colorless cuprous salt: 4NaBO,--2Cu(BO.), --C=CO+Na,B,C), +2NaBO,--Cu,(BO.), (b) The cupric salt is reduced to metallic copper, so that the bead appears reddish brown and opaque: 4NaBO,--2Cu(BO.), H-C=CO,--2Na,B,0,--Cu, COOH OXALIC ACID, º COOH Occurrence and Preparation.—Oxalic acid occurs, in the form of . its acid potassium and calcium salts, in the sap of many plants. It is prepared in large amounts by fusing Sawdust with caustic alkali. The resulting potassium Salt is precipitated with milk of lime, forming the insoluble calcium salt; and the latter is decom- posed with sulphuric acid. Oxalic acid is also formed by the oxida- tion of innumerable organic substances (such as Sugar, starch, cellu- lose (paper)), by means of concentrated nitric acid. It crystallizes from aqueous Solutions in the form of colorless COOH monoclinic prisms, & +2H,0. OOH By allowing the hydrated acid to stand over sulphuric acid the water is lost, and the anhydrous acid remains, which, when heated to about 150° C., sublimes, forming needles. If heated still higher, it is completely decomposed into water, carbon dioxide, and carbon monoxide: COOH | | = H,0+CO,--CO. COOH The crystallized, hydrated acid is soluble in water, alcohol, or ether: 100 parts water at 20° C. dissolve 11.1 parts oxalic acid; 100 “ alcohol at 15° C. dissolve 33.2 parts oxalic acid; 100 “ ether at 15° C. dissolve 1.5 parts oxalic acid. * OXALIC ACID. 3I.3 Oxalic acid is a fairly-strong, dibasic acid, and forms neutral, acid, and peracid salts: COOK COOK COOK COOH & | | | +2H,O. OOK COOH COOH COOH Potassium Potassium U- Potassium tetroxalate oxalate bin oxalate Solubility.—The oxalates are mostly insoluble in water, with the exception of the oxalates of the alkalies and of magnesium. In an excess of an alkali oxalate many of the insoluble oxalates dissolve, forming double salts. All oxalates dissolve readily in mineral acids. REACTIONS IN THE WET WAY. A solution of ammonium oxalate may be used for the following reactions: 1. Dilute Sulphuric Acid.—No reaction.* 2. Concentrated Sulphuric Acid, on warming, acts as a dehy- drating agent, causing the evolution of equal volumes of carbon monoxide and carbon dioxide; the latter burns with a blue flame: COOH | —-H.O = CO +@ COOH ſº so. /*, *. Jºe, 4– 3. Silver Nitrate precipitates white, curdy silver oxalate, COONH, | +2AgNO,-2NH, NO,--Ag,C,0, COONH, almost insoluble in water, but readily soluble in ammonia and in nitric acid. 4. Barium Chloride precipitates white barium oxalate. soluble in oxalic and acetic acids. 5. Calcium Chloride precipitates white calcium oxalate, in- * In the presence of manganese dioxide, all oxalates evolve CQ, with di- lute H,SO4: COOH | + MnO, + H,SO, - MnSO, + 2H,O + 2CO,. COOH In the same way CO, is given off by the action of KMnO, and dilute H,SO, at about 60° C. 2KMnO, + 5H.C.O., + 3H,SO. = K,SO, + 2MnSO, + 8H,0 + 10CO,. 3I4 REACTIONS OF THE METALLOIDS. wº Soluble in Oxalic acid, ammonium oxalate, and acetic acid, but readily soluble in hydrochloric and nitric acids. It is the most insoluble of all oxalates. 6. Lead Salts precipitate white lead oxalate, soluble in nitric acid. Behavior of Oxalates on Ignition. All oxalates are decomposed on ignition with slight carboniza- tion. The Oxalates of the alkalies and barium oxalate are decom- posed to carbonate, with the evolution of carbon monoxide. The Oxalates of the noble metals, and of iron, nickel, cobalt, copper, etc., leave the metal itself. COOH | CHOH TARTARIC ACID, , (or H.C.H.O.). CHOH boon Occurrence.—Tartaric acid occurs partly free and partly as its acid potassium salt in many fruit Saps, particularly in that of the grape. The free acid crystallizes in clear, monoclinic prisms, without water of crystallization. Its aqueous solution is optically active, turning the plane of polarized light to the right.* Tartaric acid is very readily soluble in water (100 parts water dissolve at 15° C. 132 parts of tartaric acid) and alcohol, but it is insoluble in ether. The salts are called tartrates. Solubility.—The neutral alkali tartrates are very soluble in water, as also is acid sodium tartrate, while the acid potassium and the acid ammonium tartrates are difficultly soluble in water. The remaining tartrates are difficultly soluble in water, but all dissolve, more or less readily, in neutral alkali tartrate Solution, forming complex salts. The most important commercial salts of this acid are “cream of tartar,” “Rochelle salt,” and ‘‘tartar emetic.” *Three other modifications of this acid exist, possessing the same chem- ical formula, but differing in their physical properties. One of these turns the plane of polarized light to the left, and the others are optically inactive. TARTARIC ACID. 3I 5 REACTIONS IN THE WET WAY. A solution of Rochelle Salt (sodium potassium tartrate) may be used for these reactions. 1. Dilute Sulphuric Acid.—No reaction. * 2. Concentrated Sulphuric Acid causes carbonization on warm- ing, with evolution of Sulphur dioxide. * & 3. Silver Nitrate produces no precipitation in a solution of free tartaric acid, but in the Solution of a neutral tartrate, a white, curdy precipitate is immediately formed, R.C.H.O.--2AgNO,-2KNO,--Ag,C,EI.O., readily soluble in nitric acid and in ammonia. By warming the ammoniacal silver Solution, metallic silver is deposited. This very important reaction for the detection of tartaric acid is performed in the following manner: The pure tartrate solution is treated with nitrate of silver until no further precipitation takes place, when dilute ammonia is added drop by drop until the precipitate just dissolves; and the test- tube containing the Solution is then placed in water which has a temperature of 60–70° C. After standing for fifteen minutes (twenty at the latest), the silver will be deposited in the form of a beautiful mirror on the sides of the test-tube. This very delicate reaction cannot be performed with certainty in the presence of other acids. In this case the tartaric acid should first be precipi- tated as acid potassium tartrate, by treating the solution (which has been concéhtated as much as possible) with solid potassium car- bonate until alkaline, whereby the tartaric acid is changed to soluble potassium tartrate. If the Solution is carefully acidified with concentrated acetic acid, a precipitate of acid potassium tar- trate will form at once if considerable tarfrate is present; * this is filtered off, washed with a little cold water, and dissolved in as little caustic soda solution as possible. In this way a solution is obtained --> * If no precipitate is formed on the addition of the acetic acid, a little alcohol is added, which causes the precipitate to form at once. It is filtered off, washed with diluted alcohol, dried, dissolved in dilute sodium hydroxide, and treated as above. If the alcohol is not removed by drying, a mirror is sometimes formed when tartaric acid is absent. 3I6 REACTIONS OF THE METALLOIDS. which will readily give the silver mirror on the addition of silver nitrate and treatment as above. 4. Calcium and Barium Chloride.—If to a concentrated solu- tion of neutral alkali tartrate, in the absence of ammonium salts, cal- cium chloride is added drop by drop, a white amorphous precipi- tate is formed which redissolves, forming readily soluble calcium alkali tartrate: COOK COOK KOOC | | 2CHOH diſon Honº | + CaCl, - 2KCl + | CHOH CHOH HOHC | | COOK too-o-oo: Only after the addition of enough calcium chloride to completely decompose the alkali tartrate is a permanent precipitate formed, which at first is flocculent, but soon becomes crystalline, consisting of neutral calcium tartrate: R.C.H.O.--CaCl, -2KCl+CaC.H.O. In dilute solutions the first addition of calcium chloride often produces no precipitation; but after standing Some time (or more quickly on rubbing the sides of the test-tube with a glass rod) the crystalline precipitate is deposited, CaC.H.O.--4H.O. Calcium tar- trate is very difficultly soluble in water; 100 parts water at 15°C. dissolve 0.0159 part of the crystalline salt, and 100 parts boiling water dissolve 0.0285 part of the salt. The precipitate is soluble in acetic acid (difference from calcium oxalate) and also in a solu- tion of concentrated caustic alkali (free from carbonate), probably forming a complex Salt: COO N COOK KOOC | N º CHOH tion Honº *! 6a+2KOHz- H,0+ | | XHOH CHOH HOHC | | COO / ko-º-o-º-oo: On boiling this solution, calcium tartrate is reprecipitated in the form of a voluminous gelatinous precipitate, which again goes into TARTARIC ACID, 31 7 solution on cooling. The presence of ammonium chloride retards the formation of the calcium tartrate, but does not prevent it; after standing some time, the precipitate settles out in the form of a heavy crystallizing powder (difference from citric acid). 5. Potassium Salts produce no precipitation in neutral Solutions of alkali tartrates; but if the solution is acidified with acetic acid, a precipitate of crystalline acid potassium tartrate* is formed imme- diately or after standing some time, according to the concentra- tion of the Solution: COONa COOK | tion CHOH | + KC1-i-CH,COOH=CH,COONa+NaCl-H CHOH CHOH | § | COONa COOH The acid potassium tartrate is difficultly soluble in water (100 parts water dissolve 0.45 part of salt) and in acetic acid, but is readily soluble in mineral acids or in caustic alkali and alkali car- bonate Solutions. If a concentrated solution of free tartaric acid is treated with potassium chloride, a precipitate of the acid potassium tartrate is formed in spite of the presence of the hydrochloric acid which is set free. From dilute Solutions the precipitate appears only after adding sodium acetate (cf. page 38). 6. Lead Acetate produces in neutral solutions a white, floccu- lent precipitate of lead tartrate, easily soluble in nitric acid and in ammonia. REACTIONS IN THE DRY WAY. If tartaric acid itself is heated to 135°C., it fuses, and on stronger ignition it is decomposed, leaving a residue of carbon and giving off empyreumatic odors (smell of burnt sugar). The alkali tartrates are also decomposed by ignition, leaving a residue of carbon and alkali carbonate, which effervesces on treatment with acid. Ammonium tartrate leaves a residue of carbon, which does not effervesce on treatment with acids. The tartrates of the alkaline * Cream of tartar, 3.18 REACTIONS OF THE METALLOIDS. earths leave behind a mixture of carbon and carbonate; on very strong ignition the latter is changed to oxide. The tartrates of those metals whose oxides are reduced by car- bon are left in the form of metal (Ag, Pb, Fe, Ni, Co, etc.). CH,-COOH CITRIC ACID, C(OH)—COOH. | CH,-COOH Citric acid is found in nature in the juices of many fruits. It is a tribasic acid, readily soluble in water and in alcohol, but diffi- cultly soluble in ether. Its salts are called citrates. Solubility.—The citrates of the alkalies are readily soluble in water, and form, with the insoluble citrates of the heavy metals, readily soluble complex Salts, whose solutions are not precipitated by alkali hydroxides, alkali carbonates, ammonia, etc. REACTIONS IN THE WET WAY. A solution of potassium citrate may be used. 1. Dilute Sulphuric Acid.—No reaction. 2. With Concentrated Sulphuric Acid carbonization and evo- lution of sulphur dioxide take place on warming. 3. Silver Nitrate produces in neutral Solutions a flocculent pre- cipitate of silver citrate, Ag,C,EISO, readily soluble in nitric acid and in ammonia. On heating the ammoniacal Solution to 60° C., no silver mirror is formed; but on heating the Solution to boiling, the silver is gradually deposited. 4. Barium and Calcium Chloride give no precipitation in neu- tral solutions (difference from tartaric acid). If, however, caustic soda solution is added to the solution which contains an excess of calcium chloride, a flocculent precipitate of tertiary calcium citrate is at once formed, insoluble in caustic alkali, but readily Soluble in ammonium chloride. On boiling the Solution in ammonium chlo- ride, crystalline calcium citrate is precipitated, which is now in- soluble in ammonium chloride. 5. Lime Water in excess produces no precipitation in solutions of neutral citrates; but on boiling a flocculent precipitate of cal- cium citrate is formed, which almost entirely redissolves on cooling. PHOSPHOROUS ACID. 3 IQ 6. Lead Acetate precipitates from solutions of the free acid, and those containing neutral salts, amorphous (C6H5O7)2Pb3+ H2O. 7. Mercuric Sulphate.—Denigès' reagent: 5gm. HgC) dissolved in 100 c.c. of water and 20 c.c. conc. H2SO4. The solution of the citrate is treated with 1/20 as much reagent and heated to boiling, after which a few drops of N/10 KMnO4 solution are added. A white crystalline precipitate is formed. The test is sensitive and can be made in the presence of tartaric, malic, oxalic, sulphuric, and phosphoric acids.” REACTIONS IN THE DRY WAY. The citrates, on ignition, behave exactly like the tartrates. /O —OH —OH NH Formation.—By the slow combustion of phosphorus in the air phosphorous trioxide is formed, being the anhydride of phosphorous acid, which it forms on treatment with cold water: P2O3 +3H2O=2H3PO3. Phosphorous acid is formed much more readily by the action of water on the trihalogen compounds of phosphorus: PCl3+3HOH=3HCl·H H3PO3. The hydrochloric acid is removed by evaporation, and the last traces of uncombined water by heating to 180° C. If the mass is then allowed to cool, it solidifies to a crystalline, hygroscopic sub- stance which melts at 70° C. By neutralizing the Solution of phosphorous acid with bases, the phosphites are obtained. It is never possible, however, to re- place more than two of the hydrogen atoms with metal; so that phosphorous acid f is considered a dibasic acid. Solubility.—Only the phosphites of the alkalies are soluble in water, but they are all soluble in acid. PHOSPHOROUs ACID, P REACTIONS IN THE WET WAY. A solution of sodium phosphite should be used. 1. Dilute Sulphuric Acid.—No reaction. 2. Concentrated Sulphuric Acid causes no reaction except on * Comptes rend., 130, p. 32. f Certain organic compounds are known, however, which are derived from tribasic phosphorous acid, P(OH)a. 32 o REACTIONS OF THE METALLOIDS. * 3. Silver Nitrate produces at first a white precipitate of silver phosphite, Na,HPO,--2AgNO,-2NaNO,--Ag,HPO, which in the case of a concentrated Solution is changed in the cold to metallic silver; while in dilute solutions this reduction takes place only on warming: Ag,HPO,--H,O=H.PO,--Ag, 4. Barium Chloride precipitates white barium phosphite, soluble in all acids. g 5. Lead Acetate precipitates white lead phosphite, insoluble in acetic acid. 6. Mercuric Chloride is slowly reduced by phosphorous acid in the cold, but more quickly on warming, to mercurous chloride: 2HgCl, -H, PO,--H,O=H.PO,--2HCl-i-Hg,Cl, If the phosphorous acid is present in excess, the reduction in the hot solution (not in the cold) goes further, and gray metallic mercury is deposited: Hg,Cl,+H,PO,--H,O= H,PO,--2HCl--2Hg. 7. Nascent Hydrogen (zinc and sulphuric acid) reduces phos- phorous acid to phosphine: H.PO,4-6H=3H,0+ PH, If the phosphine is allowed to act upon a concentrated solution of silver nitrate (1:1), or better still, upon Solid silver nitrate, the latter is colored yellow, as with arsine: PH,--6AgNO,-PAg, .3AgNO,--3HNO,. S--———- By the addition of water this yellow compound is decomposed with separation of grayish-white silver: PAg, .3AgNO,4-3H,0=3HNO,--H,PO,--6Ag. The phosphorous acid is, however, immediately oxidized by the nitric acid to phosphoric acid: 3H,PO,4-2HNO,-H,0+2NO+3H,PO,. METAPHOSPHORIC ACID, 32 I The mixture of phosphine and hydrogen burns with an emerald- green flame. 8. Sulphurous Acid is reduced by phosphorous acid to hydrogen sulphide: 3H,PO,4-H.SO,-3H,PO,--H.S. 9. Concentrated Potassium Hydroxide Solution changes a phosphite over to phosphate, with evolution of hydrogen, K.HPO,--KOH=K, PO,4-H, but with dilute caustic potash the hydrogen evolution is very slight. REACTIONS IN THE DRY WAY. By ignition, phosphorous acid (like hypochlorous acid) is changed at the cost of its own oxygen to the higher compound, while the oxidizing part of the acid is reduced to its hydrogen compound: 3HOCl=HCIO,4-2HCl; 4H,PO,-3H,PO,--PH,. The phosphites behave similarly: 8Na,HPO,-4Na,PO,--Na,P.O.--H,0+2PH,. /Q METAPHOSPHORIC ACID, P-O. & NOH The monobasic metaphosphoric acid is obtained by treating phosphorus pentoxide with cold water, P.O.--H,0=2HPO, and also by the strong ignition of Orthophosphoric acid; /Q# O=P-OH=H,0+HPO,. NOH Metaphosphoric acid is a colorless, glassy, hygroscopic mass. On boiling its aqueous Solution,” it adds water to the molecule, and is changed to orthophosphoric acid: HPO,4-H,0=H,PO,. * Or slowly in the cold. 322 REACTIONS OF THE METALLOIDS. The metaphosphates are readily obtained by heating the mono- metallic salts of Orthophosphoric acid, NaH,PO,-H,0+NaPO, or by igniting sodium ammonium phosphate: NaNH.HPO,-H,0+NH,--NaPO,. The meta salts are changed into Orthophosphates by boiling the aqueous solution in the presence of mineral acid. Solubility.—The metaphosphates of the alkalies and of magne- sium are soluble in water; but the remaining Salts are difficultly soluble or insoluble in water, and readily Soluble in nitric acid, excess of metaphosphoric acid, and excess of alkali metaphosphate. REACTIONS IN THE WET WAY. 1. Sulphuric Acid causes no visible reaction. 2. Silver Nitrate precipitates white silver metaphosphate, soluble in ammonia and in nitric acid: NaPO,--AgNO,-NaNO,--AgPO,. 3. Barium Chloride precipitates voluminous barium metaphos- phate, which is soluble in an excess of sodium metaphosphate, from which solution ammonia causes no precipitation. Barium sodium . dimetaphosphate (or a similar polymetaphosphate) is probably formed. 4. Magnesium Salts cause no precipitate from moderately dilute solutions, even on boiling (difference from Orthophosphoric acid). 5. Ammonium Molybdate produces no precipitate in the cold; but on boiling the acid Solution metaphosphoric acid is changed to orthophosphoric acid, and the characteristic precipitate of ammo- nium phosphomolybdate is formed. 6. Albumin Solution is coagulated by an aqueous solution of the free acid (difference from pyro- and Orthophosphoric acids), but not by a solution of alkali metaphosphate, except on the addition of acetic acid. 7. Nascent Hydrogen does not reduce metaphosphoric acid (difference from phosphorous acid). PYROPHOSPHORIC ACID. 323, Behavior on Ignition. The alkali metaphosphates form, on fusion, a glassy mass, which has the property of dissolving many metallic oxides, with the forma- tion of orthophosphates of characteristic colors. (See Phosphoric Acid.) By fusion with soda, orthophosphates are formed from metaphosphates. PYROPHOSPHORIC ACID, H.P.O. The tetrabasic pyrophosphoric acid is formed by heating ortho- phosphoric acid to 213°C. It is a soft, glassy mass, readily soluble. in water; and in solution it gradually adds water to the molecule and is changed to phosphoric acid, the change taking place quickly on boiling the Solution. The salts of pyrophosphoric acid, the pyrophosphates, are ob- tained by igniting the dimetallic phosphates: 2Na,HPO,-H,O-i-Na,P.O. Solubility.—The pyrophosphates of the alkalies are soluble in water; the remaining pyrophosphates are difficultly soluble or in- soluble in water, but are all Soluble in acids, and some are soluble- in an excess of sodium pyrophosphate. REACTIONS IN THE WET WAY. 1. Sulphuric Acid.—No reaction. 2. Silver Nitrate gives a white precipitate, soluble in ammonia and in nitric acid. - 3. Barium Chloride causes a white, amorphous precipitate, soluble in acids. 4. Magnesium Chloride produces a white precipitate which is. soluble in an excess of the magnesium salt, as well as in an excess of sodium pyrophosphate. By boiling this solution a precipitate is formed, which does not disappear on cooling. 5. Ammonium Molybdate produces no precipitation in the cold; but, on warming, yellow ammonium phosphomolybdate is precipi- tated. 324 REACTIONS OF THE METALLOIDS. 6. Albumin is not coagulated by free pyrophosphoric acid (dif- ference from metaphosphoric acid). BEHAVIOR IN THE DRY WAY. All pyrophosphates on being fused with sodium carbonate are changed to Orthophosphates: Na,P.O.--Na,CO,-CO,--2Na,PO. IODIC ACID. Occurrence.—In sea-water and in Chili Saltpetre as potassium iodate. Formation.—By oxidizing iodine with fuming nitric acid or by the action of chlorine upon iodine suspended in water: 3L-1-10HNO,-6HIO,--10NO+2H,0; I2 +6H2O+5Cl2= 10HCl·H-2HIO3. The most important iodate, KIO3, is obtained by the action of iodine upon a slightly acid solution of potassium chlorate: 5KClO3 +3I2+3H2O=5KIO3 +HIO3+5.HCl. Iodates are also formed by the action of iodine upon alkali hydroxide solutions: 3I2 +6KOH=5KI+KIO3 +3H2O. In alkaline solutions iodides are oxidized to iodates by hypo- chlorites and potassium permanganate. Solubility.—The iodates of the alkalies are soluble in water, but the remaining iodates are difficultly soluble or insoluble. REACTIONS IN THE WET WAY. 1. Sulphuric Acid.—Neither dilute nor concentrated sulphuric acid decomposes iodic acid; but if reducing substances are present at the same time (such as hydriodic acid, hydrogen sulphide, ferrous salts, etc.), the hydriodic acid is reduced, with separation of iodine: KIO,--5KI+3H;SO,-3K,SO,--3H,0+3I2. 2. Silver Nitrate precipitates white, curdy silver iodate, AgIO, readily soluble in ammonia. but difficultly soluble in nitric acid. IODIC ACID. * 325 3. Barium Chloride precipitates white barium iodate, difficultly soluble in hot water (100 parts of boiling water dissolve 0.6 part of the salt), and only slowly soluble in hydrochloric or nitric acids. 4. Lead Acetate precipitates lead iodate, difficultly soluble in water and only slightly soluble in nitric acid. 5. Reducing Agents. (a) Hydriodic acid reduces iodic acid, with separation of iodine: HIO,--5HI=3H,0+3I2. If the solution is concentrated, the iodine separates out as a brown powder; dilute solutions are colored yellow. The iodine may be absorbed with a reddish-violet color by shaking the solution with chloroform or carbon disulphide. (b) Sulphurous acid also causes separation of iodine; unless a large excess of sulphurous acid is added. The reaction takes place in three stages: I. KIO,--3SO,--3H,0=KI+3H,SO.; II. KI+H,SO,-HI+ KHSO, ; KIO,--H,SO,-HIO,-- KHSO,; III. HIO,--5HI=3H,0+3I2. If too much sulphurous acid is added, there will be no permanent deposition of iodine, for it will be changed to hydriodic acid: SO,--I, +2H,O=H,SO,--2HI. (c) Zinc dust (or, better, Devarda's alloy) reduces neutral iodate solutions to iodide. REACTIONS IN THE DRY WAY. Heated on charcoal the iodates deflagrate, but not so strongly as the chlorates; they are all decomposed on being heated, some with and some without the Separation of iodine. Thus all neutral iodates of the alkalies are easily decomposed into iodide and oxygen, while the biiodates set free iodine at the same time: 2KIO3=2KI+3O2; 4[KIO3. HIO3]=4KI+11O2+2H2O+2I2. 326 REACTIONS OF THE METALLOIDS. GROUP IV. Silver Nitrate produces a colored precipitate in neutral solutions, soluble in nitric acid. Barium Chloride also produces a precipitate which is soluble in nitric acid. /Q —OH. PHosphoric AcID, PT: NOH Orthophosphoric acid is obtained by the oxidation of phos- phorus by means of nitric acid, or by boiling the meta- and pyro- phosphoric acids with water. It is a tribasic acid, and forms salts in which either one, two, or three of its hydrogen atoms are replaced by metals: ONa. ONa ONa) o–Póñ. o–Pºona, =P^N. NöH NOH NóNa Mono-, di-, and trisodium phosphate Solubility.—The phosphates of the alkalies are soluble in water, and so are the primary salts of the alkaline earths. Their secondary phosphates are very difficultly soluble, while their tertiary phos- phates (as well as all other phosphates) are insoluble. All phos- phates dissolve in acids. * REACTIONS IN THE WET WAY. A solution of disodium phosphate should be used for these re- actions. º 1. Sulphuric Acid, dilute or concentrated, produces no visible change. 2. Silver Nitrate produces a yellow precipitate of silver phos- phate (difference from meta- and pyrophosphoric acids), 2Na,FIPO,--3AgNO,-3NaNO,--NaH,PO,--AgPO, readily soluble in nitric acid and in ammonia. The precipitate, therefore, can only be formed in neutral Solution. & PHOSPHORIC ACID. 327 3. Barium Chloride precipitates white, amorphous barium phosphate: Na,HPO,--BaCl, -2NaCl-FBaFIPO,. In the presence of ammonia the tertiary Salt is precipitated: 2Na,FIPO,--3BacL--2NH,-4NaCl-H2NH,Cl-i-Ba,(PO.), The barium phosphates (as well as those of the other alkaline earths) are easily dissolved by acids, even acetic acid (difference from aluminium and ferric phosphates). From these acid Solutions, ammonia reprecipitates the phosphate: Ba,(PO.), -6HCl=2H,PO,--3BaCl, By adding ammonia the phosphoric acid is changed to ammo- nium phosphate, which precipitates the barium chloride, forming insoluble barium phosphate. 4. Magnesia Mixture (a mixture of ammonium chloride, am- monia, and magnesium chloride) precipitates from very dilute solutions white, crystalline magnesium ammonium phosphate, MgNHAPO,--6H.O, Na,HPO,--MgCl, -NH,-2NaCl-i-MgNH,PO, which is soluble in all acids, but practically insoluble in dilute 2% per cent. ammonia. This is a very sensitive reaction (cf. page 34). 5. Ferric Chloride.—If a solution of sodium phosphate is treated with ferric chloride, a yellowish-white precipitate of ferric phos- phate is formed: Na,HPO,--FeCl, z= 2NaCl-i-HC1-i-FePO,. As hydrochloric acid is set free by this reaction, the precipita- tion of the phosphoric acid cannot be quantitative, ferric phosphate being soluble in hydrochloric acid. If some sodium acetate is added to the solution, the amount of free hydrochloric acid is diminished; in its place an equivalent amount of acetic acid is formed, which does not dissolve ferric phosphate; in that case all the phosphoric acid will be precipitated: \ Na,FIPO,--CH,COONa+FeCl,-3NaCl-i-CH,COOH-i-FePO, 328 REACTIONS OF THE METALLOIDS. Ferric phosphate is not inappreciably soluble in ferric chloride and ferric acetate Solutions; consequently the precipitate is pro- duced in a boiling Solution, and as slight an excess as possible of ferric chloride is added. In this way the excess of the ferric chlo- ride will be precipitated as basic ferric acetate with the ferric phos- phate, so that it can exert no solvent action. If the solution is filtered hot, a filtrate free from iron and phosphoric acid is obtained. As ferric phosphate is insoluble in acetic acid, it is evident that phosphoric acid can be completely precipitated from solutions of those phosphates which are soluble in acetic acid by adding ferric chloride, and, thus, separated from the metals with which it was originally combined (e.g., calcium, strontium, barium, and magne- sium phosphates). To accomplish this, the phosphate is dissolved in as little hydro- chloric acid as possible, ammonium carbonate is added until a slight permanent precipitate is obtained, which is dissolved by the addi- tion of one or two more drops of hydrochloric acid; ammonium ace- tate is then added, and ferric chloride drop by drop until the solution above the yellowish-white precipitate of ferric phosphate is colored distinctly brown. The solution is diluted considerably with water, heated to boiling, and filtered hot. In order to detect phosphoric acid in the precipitate, it is dissolved in dilute nitric acid, evaporated to a small volume, and treated with ammonium molybdate, whereby yellow, crystalline ammonium phosphomolybdate is formed (see below); or the precipitate is dissolved in dilute hydrochlorié acid, two gns, of tartaric acid are added, and ammonia to alkaline reac- tion (the iron will not be precipitated, cf. page 99). . From the ammoniacal solution the phosphoric acid may be precipitated with magnesia mixture as magnesium ammonium phosphate. 6. Ammonium Molybdate, in large excess, precipitates from nitric acid solutions in the cold on standing (more quickly on slightly warming) a yellow, crystalline precipitate of ammonium phosphomolybdate: H.PO,-- 12(NH,),MoQ.--21HNO,- =(NH,),PO,-12MoC), -21NH, NO,-- 12H,O. Y. This reaction is completely analogous to the reaction with arsenic acid (cf. page 189), with the difference that the arsenic compound PHOSPHORIC ACID. 329 is only formed quickly at the boiling temperature. The presence of ammonium nitrate greatly facilitates the formation of this precipi- tate. Ammonium phosphomolybdate is readily soluble in alkalies and in ammonia, (NH,),PO, 12MoC),4-24NH,OH = =(NH,),PO,-- 12(NH), MoC),4-12H,O, also in an excess of alkali phosphate solutions, forming compounds which contain less molybdenum. It is, therefore, always necessary to prevent the formation of such compounds by the addition of a large excess of ammonium molybdate. Detection of Phosphorus in Iron and Steel.—Phosphorus is present in iron and steel as iron phosphide, but only to a slight extent (usually less than 0.1 per cent.). The detection of the phosphorus is effected by its oxidation to phosphoric acid and using one of the above reactions. As, however, very Small amounts of phosphorus are present, it is necessary to start with a large amount of the origi- nal substance in order to obtain a perceptible phosphorus test. It is best to proceed as follows: Five to ten grams of the iron or steel are dissolved in about 60 c.c. of nitric acid (sp. gr. 1.2),” the solution evaporated to dryness and then ignited over a free flame (with con- stant stirring) until no more red fumes are given off. All organic matteristhereby destroyed, and the silicic acid is dehydrated. After cooling, the oxides are dissolved in 50–60 c.c. of concentrated hydro- chloric acid (warming gently), the excess of the acid is removed by evaporation; water is added, and the silica filtered off. In the filtrate all the iron and all the phosphoric acid will be found, and the latter may be detected by either the molybdate or the magnesia- mixture reaction. In order to detect the phosphoric acid accord- ing to the former method, the filtrate obtained after the removal of the silica is evaporated to dryness, dissolved in as little nitric acid as possible (sp. gr. 1.2), 50 c.c. of ammonium molybdate solution and 15–20 c.c. of a 75 per cent. ammonium nitrate solution added, the mixture warmed gently and allowed to stand twelve hours. A yellow, crystalline precipitate shows the presence of phosphorus. *If the iron were dissolved in HCl or H.S.O., part or even all of the phos- phorus would escape as phosphine. Nitric acid oxidizes all of the phos- phorus to phosphoric acid. 33O REACTIONS OF THE METALLOIDS. In order to detect the phosphorus according to the magnesia- mixture method, it is necessary first to remove the greater part of the iron. For this purpose the hydrochloric acid filtrate is neu- tralized with ammonia, a saturated solution of Sulphur dioxide is added, and the mixture boiled, whereby the previously dark-colored solution is either decolorized or becomes a light green. Twenty c.c. of concentrated hydrochloric acid are added, and the solution is boiled until the excess of sulphur dioxide is expelled. By this oper- ation all the ferric salt has been reduced to ferrous Salt. A few drops of chlorine water are now added (which forms a little ferric salt), the solution is neutralized with ammonia and diluted to about a liter; three c.c. of a saturated Solution of ammonium acetate are added, also five c.c. of acetic acid, and the solution heated to boiling, whereby all the ferric salt and all the phosphoric acid will be precipi- tated in the form of ferric phosphate and basic ferric acetate, while the greater part of the iron remains in solution as ferrous salt. The slightly-brown precipitate is filtered off through a small plaited filter, washed with hot water, dissolved in dilute hydrochloric acid, evaporated almost to dryness, two gm. of citric (or tartaric) acid added (which should be dissolved in as little water as possible), the solution saturated with ammonia, and the phosphoric acid precipi- tated by the addition of magnesia mixture. A white crystalline precipitate shows the presence of phosphoric acid. 7. Lead Acetate precipitates white lead phosphate, almost in- soluble in acetic acid: 's 2Na,FIPO,--3Pb(C.H.O.),= =4C.H.O.Na+2C.H.O.H--Pb,(PO.), 8. Nascent Hydrogen does not reduce phosphoric acid (differ- ence from phosphorous and hypophosphorous acids). 9. Metastannic Acid.—If metallic tin is added to a nitric acid solution of phosphoric acid, or a phosphate, the tin is changed to metastannic acid, which unites with the phosphoric acid, forming an insoluble compound (probably a complex phosphorous stannic acid). This reaction is often used to separate phosphoric acid from other metals. 10. Mercurous Nitrate precipitates from solutions which are almost neutral white mercurous phosphate, soluble in nitric acid, but insoluble in acetic acid. PHOSPHORUS. 33 I REACTIONS IN THE DRY WAY. The tertiary salts of the alkalies melt without decomposition; the Secondary Salts lose water and are changed to pyrophosphates; while the primary salts form a glassy metaphosphate. The so-called “salt of phosphorus,” or “microcosmic salt,” NaNH.HPO,--4H,0, which is much used as a reagent, loses water and ammonia on being fused, forming a clear glass of sodium metaphosphate: NaNH,HPO,-4H,0=5H,0+NH,--NaPO,. If the salt is heated in the loop of a platinum wire, a clear bead is obtained—the so-called “salt of phosphorus” bead. Just as metaphosphoric acid unites with water, on boiling its solution, forming orthophosphoric acid, HPO,4-H,0=H,PO, so sodium metaphosphate dissolves, on fusing, a great many metal- lic oxides, forming characteristically-colored Orthophosphates, NaPOs-HCuO=NaCuPO, (blue bead), which may be changed in the reducing flame to metaphosphate again: **- s , NaCuPO,--C=CO+Cu+NaPO,. * * ---— Brownish red opaque bead Many anhydrous phosphates are reduced by heating with mag- nesium to phosphides, which, on being breathed upon, give the peculiar odor of phosphine: Ca,(PO.),4-8Mg=8MgO+Ca,P,; Ca,P,--6H.O=3Ca(OH),4-2PH,. PHosphorus, P. At. Wt. 31. Mol. Wt. Pi—124. Occurrence.—Phosphorus is found in nature only in the form of phosphates, of which calcium phosphate is the most important. It occurs as apatite, Cas(PO4)2(ClF), in hexagonal crystals, and in an 332 REACTIONS OF THE METALLOIDS. impure state as phosphorite, which is used extensively as a fertilizer. Calcium phosphate also is an important constituent of bones and the seeds of plants. A very interesting occurrence of phosphorus is pyromorphite (cf. page 156), isomorphous with apatite, vanadinite, and mimete- site. Properties.—Phosphorus exists in three allotropic forms: (a) As ordinary or colorless phosphorus; (b) As red crystalline phosphorus; (c) As black crystalline phosphorus. Ordinary phosphorus is poisonous, is colorless when pure (it becomes yellow on exposure to the light, and is coated with a layer of red phosphorus), melts at 44°C., and ignites at 60°C. in the air, so that it must be kept covered with water, in which it is insoluble. It is readily soluble in carbon disulphide, and slightly soluble in ether. It is easily oxidized by nitric acid to phosphoric acid: 3P,--20HNO,--8H,O=12H,PO,--20NO. Red phosphorus is crystalline (hexagonal, rhombohedral), and is formed by heating ordinary phosphorus to about 250° C. out of contact with the air. It is not poisonous, is insoluble in carbon di- sulphide, and does not ignite until heated to 256°C. It is non- luminous in the dark, does not oxidize in the air, but is readily oxi- dized by nitric acid to phosphoric acid. - Black phosphorus is obtained when red phosphorus and lead are heated together in a sealed tube to a red heat and the mass treated with dilute nitric acid after it is cold; the lead dissolves and leaves the phosphorus as black phosphorus. By heating to 360° C. it is changed to ordinary phosphorus again. Phosphorus is found in a great many organic substances. In order to detect its presence, the compound is heated in a sealed tube with fuming nitric acid, which destroys the organic matter and oxi- dizes the phosphorus to phosphoric acid, (detected by any of the above reactions). Arsenious, arsenic, and chromic acids. which also belong to this group, have already been described on pages 183, 186, and 85. THIOSULPHURIC ACID. 333 THIOSULPHURIC ACID, so.T. This very unstable acid is soon decomposed, even in dilute aqueous solution, into Sulphurous acid and Sulphur: H.S,0,-H,SO,--S. If the aqueous solution of a thiosulphate is treated with dilute hydrochloric or sulphuric acid, the solution remains clear for a short time; but it soon becomes turbid, owing to the deposition of sul- phur, which in this case (unlike most precipitated Sulphur) appears yellow. The salts of thiosulphuric acid, the thiosulphates, are much more stable than the free acid. Formation of Thiosulphates. 1. By boiling sulphur with an alkali or alkaline-earth hydroxide: 4S+ 6NaOH=2Na,S+Na,S,O,--3H,0. This reaction is analogous to the action of the halogens or of phosphorus upon hydroxides, forming chloride and hypochlorite, phosphide (phosphine), and hypophosphite, etc. (cf. pages 253 and 299). 2. By boiling sulphites with sulphur: Na,SO,--S=Na,S,O,. 3. By treating alkali polysulphides with alkali sulphite in the cold: Na,Ss--4Na,SO,-4Na,S,O,--Na,S. 4. By the oxidation of polysulphides: 2Na2S2 + 3O2=2|Na2S2O3. This last reaction takes place on boiling the solution of poly- sulphide in the air, or very slowly on standing. Yellow ammonium polysulphide is changed, on standing in the air, into ammonium thiosulphate with deposition of Sulphur. 334 REACTIONS OF THE METALLOIDS. The Sulphites can be kept well in aqueous solutions provided they are not subjected to the action of carbon dioxide. ' They are gradually decomposed by the latter, with separation of sulphur. The most important commercial thiosulphate is the sodium salt Na,S,O, 5H,0, the well-known “hypo” of photographers. Solubility.—The thiosulphates of the alkalies are readily soluble in water, the remaining ones are difficultly soluble; many of them dissolve in an excess of Sodium thiosulphate, forming complex Salts. REACTIONS IN THE WET WAY. A solution of sodium thiosulphate should be used. 1. Sulphuric Acid.—Both dilute and concentrated sulphuric acid decompose thiosulphates, with deposition of yellow sulphur. 2. Silver Nitrate produces a white precipitate, which rapidly becomes yellow, brown, and finally black, owing to the formation of silver sulphide: Na,S,Os-H2AgNO,-2NaNO,--Ag, S.O.; Ag,S,Os-H H.O=H.SO,--Ag,S. Silver thiosulphate is soluble in an excess of the reagent. Diffi- cultly soluble AgNaS,O, is at first formed, Ag,S,Os-H Na,S,O, =2AgNaS,Os, which combines with more thiosulphate, forming a soluble complex salt: 2NaAgS,Os-H Na,S,Os-[Ag,(S.O.)alNaa. But by boiling the solution, silver sulphide is precipitated: 2NaſAgS2O3]=Na2SO4+SO2 +S-H-Ag2S. Many other metals behave like silver,” especially those of the hydrogen sulphide group. Thus copper, mercurous, and tin salts are precipitated as Sulphides by boiling the acid solutions with Sodium thiosulphate. 3. Barium Chloride in excess produces a white, crystalline pre- cipitate of barium thiosulphate,f difficultly soluble in cold water * Zeitschr. f. anorg, Ch., XXVIII, p 223 (1902). t Rubbing the sides of the test-tube hastens the formation of this pre- cipitate. THIOSULPHURIC ACID. 335 (480 parts of water at 18° C. dissolve one part of BaS,03), but fairly soluble in hot water. 4. Strontium Chloride produces a white, crystalline precipitate, but only in very concentrated solutions (3.7 parts of water at 18°C. dissolve one part of SrS,Os). 5. Lead Acetate precipitates white lead thiosulphate, Soluble in an excess of the alkali thiosulphate. On boiling the solution a voluminous precipitate, consisting of lead sulphate and lead sul- phide, is formed. 6. Iodine Solution is decolorized by solution of thiosulphates: 2Na,S,Os--I, +2NaI+Na,S,0s. Sodium tetrathionate This reaction is explained, according to the dissociation theory, as follows: 2S,0s"+I, +S,O′--I'--I’. The negative anions, S.O., carry half of their electric charge to the undissociated, and consequently uncharged, iodine atoms; which changes the latter to ions, while the two S.O., ions unite, forming S.O., ions. Chlorine and bromine act quite differently upon thiosulphates. If chlorine (or bromine) is conducted into a solution of sodium thio- sulphate, a considerable precipitation of sulphur takes place, which, upon further action of the halogen, disappears: Na,S,Os-H,O+Cl,-2NaCl-i-H,SO,--S; S+4H,0+301,–6HCl·H,SO, Other weak oxidizing agents act in the same Way as iodine. Thus, gy 7. Ferric Chloride produces, in solutions of sodium thiosulphate, at first a dark-violet coloration (perhaps ferric thiosulphate), which disappears after some time, leaving a colorless solution containing ferrous chloride and sodium tetrathionate: 2Na,S,Os-H2Fe(Cls-2NaCl-i-2FeCl,+Na,S,O, 336 REACTIONS OF THE METALLOIDS, Similarly, 8. Cupric Salts are reduced to colorless cuprous compounds, with the formation of sodium tetrathionate: 2Na,S,O,--2CuSO,-Na,SO,--Cu,SO,--Na,S,O. The unstable cuprous Sulphate immediately acts upon more thiosulphate, forming Sodium cuprous thiosulphate: Cu,SO,--2Na,S,O,-Na,SO,--[Cu,(S.O.), Na,. If the colorless Solution of the cuprous salt is treated with caustic potash Solution, yellow cuprous hydroxide is in some cases imme- diately formed, in other cases only on standing or on warming. The precipitate becomes darker colored on being boiled. If the solution is acidified and boiled, black cuprous sulphide is precipitated. The colorless Solution of the cuprous salt gives also a white (usually almost a light red) precipitate with potassium ferrocyanide or cuprous ferrocyanide. 9. Nascent Hydrogen (zinc and hydrochloric acid) causes the evolution of hydrogen sulphide. 10. Zinc Salts produce no precipitate (difference from sul- phides). 11. Zinc Sulphate and Sodium Nitroprusside produce no red coloration (difference from Sulphites). Detection of Sulphurous and Thiosulphuric Acids in the Presence of Hydrogen Sulphide. The three acids are assumed to be present together in solution in the form of their alkali salts. The fairly concentrated solution is treated with zinc sulphate, and the zinc sulphide formed is filtered off. The filtrate is treated with strontium nitrate solution and allowed to stand over night, when the strontium sulphite is filtered off and washed with a little cold water. If the strontium sulphite is treated on the filter with dilute hydrochloric acid, sulphurous acid goes into solution, which can be detected by its property of decolor- izing an iodine solution. In the filtrate from the strontium sulphite, the thiosulphate remains; it can be detected by acidifying with hydrochloric acid and warming, when sulphur will be deposited. NITRIC ACID. 337 Solubility of Sulphites and Thiosulphates of the Alkaline * Earths in Water. Sulphite. Thiosulphate. Calcium. . . . . . . . . . . . . . . 1:800 1: 2 Strontium . . . . . . . . . . . . . 1: 30000 1: 3.7 Barium . . . . . . . . . . . . . . . 1: 46000 1: 480 REACTIONS IN THE DRY WAY. The thiosulphites of the alkalies, on being heated out of con- tact with the air, are changed into Sulphate and polysulphide, and the latter into Sulphide and Sulphur: Na, SSO, Na, S|SO Na. S Sö. =3Na,SO,--Na,Ss Na, SS|OOO and Na,Ss=Na,S+4S. If this reaction is performed in a closed tube, a sublimate of sul- phur is obtained (difference from sulphites); and the residue yields hydrogen sulphide if treated with acid. GROUP V. Silver Nitrate produces no precipitate in acid or neutral solu- tions. {} Barium Chloride, also, causes no precipitation. NITRIC ACID, HNO,. Occurrence.—Nitric acid is found in the form of nitrates in small amounts almost everywhere in nature; thus the ammoni m salt is found in the atmosphere and in Soils; the calcium Salt it, found in old masonry; while the Sodium Salt is found in rainless localities, particularly in Chili (Chili Saltpetre). Nitric acid is the final product of the Oxidation of ammonia; it is found wherever nitrogenous organic Substances have been jubjected to decay, forming ammonia. With the help of micro-organisms * Autenrieth and windaus, Zeit. f. anal. Chem., 1898, p. 295. 338 REACTIONS OF THE METALLOIDS. (Monas nitrificans, according to Winogradsky) the ammonia is changed first to nitrous acid, 2NH3+3O2=2H2O+2HNO2, which, by further oxidation, is changed to nitric acid: 2HNO2+O2=2HNO3. Properties.—Pure nitric acid is a colorless liquid, with a specific gravity of 1.54 at 20° C. At 86° C. it begins to boil, with decom- position, giving off its anhydride, which suffers further decompo- sition into nitrogen peroxide, NO, (brown fumes), and oxygen. By the constant loss of N.O.s, the nitric acid becomes constantly more dilute and the boiling-point constantly rises, until at 120.5° C. it remains constant; when nitric acid of specific gravity 1.414 distils over, forming a 68 per cent. acid. If a more dilute acid is subjected to distillation, water is at first given off, the boiling-point con- stantly rising until 120.5° C. is reached, when a 68 per cent. acid again distils unchanged. Red, fuming nitric acid is obtained by conducting NO, into the colorless, concentrated acid. In its most concentrated condition it possesses a specific gravity of 1.55. If the fuming acid is treated with water, it is colored green, and vapors of nitric oxide are given off, which are colored brown on coming in contact with the air. The dissolved NO, (or, better, N.O.), being a mixed anhydride is changed into nitric and nitrous acids, O N–O Nºll -HNo.4 HNo. NO and the nitrous acid, owing to the heat of reaction, is partly changed into nitric acid, with evolution of nitric oxide, 3HNO,-H,0+HNO,4-2NO. and NO+O=NO, (brown vapors). Nitric acid is a strong oxidizing agent (cf. page 4). It is mono- basic, and, next to the halogen acids, is the strongest acid. It forms NITRIC ACID, 339. stable salts, which are all soluble in water ; but a few of them are changed by water into basic salts (cf. bismuth and mercuric Salts), insoluble in water, but soluble in nitric acid. REACTIONS IN THE WET WAY. As nitric acid does not form insoluble salts, it cannot be detected by means of precipitation; its characteristic reactions depend upon its oxidizing action. Great care must be exercised before deciding whether this acid is present, for other oxidizing substances give similar (in some cases the same) reactions. 1. Dilute Sulphuric Acid gives no reaction (difference from nitrous acid). 2. Concentrated Sulphuric Acid when heated with any nitrate causes evolution of yellow to brown vapors of NO, with a charac- teristic penetrating odor. 3. Silver Nitrate and Barium Chloride cause no precipitation. 4. Ferrous Salts are oxidized by nitric acid, which is itself re- duced to nitric oxide, NO. If the reaction takes place in the cold, the latter combines with the excess of ferrous salt, forming a dark- brown, very unstable compound, FeX. NO. This compound is decomposed, on warming, into ferrous salt and nitric oxide (which escapes) the brown color disappearing. If the amount of nitric acid present is more than sufficient to completely oxidize the ferrous salt to ferric salt, the brown color will only appear as representing a transient intermediate product; for ferric salts do not unite with nitric oxide. The oxidation takes place according to the following reaction, 6FeSO,--3H,SO,--2HNO,-3Fe,(SO,),4-4H,0+2NO, and is best carried out in the following manner: A little of the substance to be tested for nitric acid is placed in a test-tube and dissolved in as little water as possible; a cold satu- rated solution of ferrous Sulphate is added, the two solutions mixed, and the sulphuric acid carefully poured down the sides of the tube. If nitric acid is present, the contact zone will be colored distinctly brown. Nitrous acid gives the same reaction, with the difference that it takes place without the addition of concentrated sulphuric acid. 34o REACTIONS OF THE METALLOIDS, *s-, 5. Indigo Solution is decolorized by warming with nitric acid (as well as by other oxidizing agents). 6. Potassium Iodide is not decomposed by pure, dilute nitric acid (difference from nitrous acid). If to the Solution of a nitrate we add potassium iodide, a few drops of an acid (best acetic acid), and a little zinc, the nitric acid is reduced to nitrous acid, which then reacts with hydriodic acid so that the Solution becomes yellow on account of the separation of iodine. By shaking the solution with carbon disulphide, the latter will be colored reddish violet,” or the iodine may be detected by adding a little starch paste. The reactions which take place may be represented by the following equations: Zn-H2HC.H.O. =Zn(C.H.O.),4-H, 2HNO,--2H,-2H,0+2HNO,; 2HNO,--2HI=2H,0+2NO+I, t 7. Diphenylamine Reaction (the Lunge test f).-The reagent is prepared by dissolving 0.5 gm. of diphenylamine in 100 c.c. of pure concentrated Sulphuric acid and adding 20 c.c. of water. Procedure.—A few cubic centimeters of the diphenylamine solu- tion are placed in a test-tube and carefully covered with the solution to be tested for nitric acid. If the latter is present, there is formed at the Zone of contact between the two liquids a ring of a beautiful blue color. This very sensitive reaction is, unfortunately, also caused by nitrous, chloric, and selenic acids, ferric chloride, and many other Oxidizing agents. In the absence of ferric and selenic salts, it is useful for detecting * If it is desired to detect the iodine by means of carbon disulphide, the latter wmder no circumstances should be added before the zinc has been al- lowed to act upon the acid solution of the nitrate and potassium iodide. In such a case, there will often be no separation of iodine, because the nascent hydrogen is used up reducing the carbon disulphide to thioformaldehyde and hydrogen sulphide, and the latter reacts with any iodine which may be formed, changing it back to hydriodic acid: CS, + 2H, - CH.S + H2S; H.S + 2I = 2HI + S. f Lunge, Zeitschr. für angew. Chemie, 1894, Heft 12. NITRIC ACID. 34 I the presence of Small amounts of nitrogen acids in Sulphuric acid. In this case the concentrated sulphuric acid which is to be tested is first poured into the test-tube, and the specifically lighter diphenyl solution is then poured on top. If one c.c. of an acid which contains only ºr milligram of nitrogen in a liter is used, the reaction will cause a noticeable coloration. 8. Brucine Reaction.—The reagent * is prepared by dissolving 0.2 gm. of brucine in 100 c.c. of pure concentrated sulphuric acid. Procedure.—The solution to be tested for nitric acid is mixed with three times its volume of pure concentrated sulphuric acid, and one c.c. of brucine solution is added. If nitric acid is present, a red coloration quickly appears, which quickly changes to orange, then slowly to lemon or gold yellow, and finally becomes greenish yellow. Nitrous acid does not give this reaction provided it is present as “nitrose,” i.e., dissolved in concentrated sulphuric acid. Aqueous solutions of nitrites always yield a small amount of nitric acid when acidified with sulphuric acid, and consequently give the brucine reaction. 9. Zinc in Alkaline Solution reduces nitric acid to ammonia. If a nitrate solution is boiled with zinc dust and an alkali, a considerable evolution of ammonia takes place: KNO,4-4Zn--7KOH=4Zn(OH),4-2H,0+NH, /... . . .2 º - Devarda's alloy reacts much móré guickly with a drop of caustic soda. This reaction is particularly suited for the detection of nitric acid in the presence of chloric acid (cf. page 344). Detection of Nitric Acid in the Presence of Nitrous Acid. With the exception of the Lunge–Lwoff method, there is no absolutely reliable qualitative test for the detection of traces of nitric acid in the presence of large amounts of nitrous acid in aqueous solution. A number of methods have been proposed which depend upon the destruction of the nitrous acid by diazotizing, but they all yield only approximate results; because, in order to destroy the nitrous acid, it is necessary first to set the acid itself free by the addition of another acid, which always causes a part of the nitrous * Lunge, Zeitschr, für angew. Chemie, 1894, Heft 12. º 342 REACTIONS OF THE METALLOIDS. acid to be changed to nitric acid; so that the latter will be detected even when no nitric acid was originally present. Large amounts of nitric acid in the presence of nitrous acid may be detected by the method proposed by Piccini, in which a concen- trated solution containing salts of both acids is treated with a con- centrated Solution of urea, and then covered (by means of a pipette) with dilute sulphuric acid. A lively evolution of nitrogen ensues, which ceases in a few minutes: CONH),4 2HNO,-CO,--3H,0+2N,. Tea, This reaction, however, does not take place quickly enough to prevent traces of nitric acid being formed according to the following equation 3HNO,- H,0+ HNO,--2NO. The odor of nitrous fumes is always perceptible in the escaping nitrogen,* which can also usually be detected by means of iodo- starch paper. The nitric acid which remains in the solution can be detected by means of the diphenylamine reaction. The nitrous acid may also be destroyed by boiling an alkaline nitrite solution with neutral ammonium chloride; but traces of nitric acid are always in evidence at the same time.f If the diphenylamine reaction gives a very intense coloration after the destruction of the nitrous acid by means of urea, the pres- ence of nitric acid in the original compound is assured; but if the reaction shows that only a trace of nitric acid is present, it is prob- ably due simply to Small amounts of nitric acid formed by the destruction of the nitrous acid. REACTIONS IN THE DRY WAY. By the ignition of nitrates of the alkalies, they are changed into nitrites with loss of oxygen, and the latter are decomposed on stronger ignition into oxide: KNO,-KNO,--O; 2KNO,-K,0+2NO+O. * Even at 0° and in an atmosphere of carbon dioxide. f By evaporation with ammonium carbonate the decomposition scarcely takes place at all. CHLORIC ACiD, 343 All nitrates deflagrate on being heated on charcoal; i.e., the charcoal burns at the expense of the oxygen of the nitric acid, with vivid scintillation. CHLORIC ACID, HClO,. Free chloric acid is extremely unstable, and is decomposed, at 40°C., into perchloric acid with loss of chlorine and oxygen: HOCO, H ClO,O= H,0+2010,--HClO, HCIO, and 2ClO2 = Cl2 +2O2. The salts of the monobasic chloric acid, the chlorates, are quite stable, and are all soluble in water. REACTIONS IN THE WET WAY. 1. Dilute Sulphuric Acid sets free chloric acid from chlorates, which, as above stated, is gradually decomposed, with loss of chlorine and oxygen, into perchloric acid. The Solution, therefore, acts as a bleaching agent, particularly on warming. The neutral salts do not bleach (difference from hypochlorites). 2: Concentrated Sulphuric Acid decomposes all chlorates, setting free greenish-yellow chlorine dioxide gas, which violently explodes on warming: 3KClO,--3H2SO4=3KHSO4+HClO4+2CIO,4-H,0. 3. Silver Nitrate and Barium Chloride do not cause precipita- tion. 4. Nascent Hydrogen reduces chlorates to chlorides in acid, alkaline, and neutral Solutions. & The reduction in acid solution is effected by means of zinc and dilute sulphuric acid, or by means of Sulphurous acid: HClO,--3SO,--3H,0=3H, SO,--HGl. 344 REACTIONS OF THE METALLOIDS. The reduction in alkaline or neutral * solution is brought about by boiling the solution with zinc dust, or, better, by means of Devarda's alloy (cf. page 7): KCIO,--3Zn-i-3H,0=3Zn(OH),4-KCl. The residue of zinc dust (or copper, if the alloy is used) is fil- tered off, the solution acidified with nitric acid, and silver nitrate added, when the characteristic, curdy precipitate of silver chloride is formed. 5. Concentrated Hydrochloric Acid decomposes all chlorates, with evolution of chlorine: KClO,--6HCl=KCl-i-3H,0+ 3Olaf 6. Ferrous Salts.—By boiling chlorates with ferrous salts in the presence of dilute Sulphuric acid, the chlorate is quickly reduced to chloride (difference from perchloric acid): KCIO,--3H, SO,--6FeSO,-3H,0+ KCl-i-3Fe,(SO,)s. 7. Diphenylamine reacts the same as with nitric acid. Detection of Hydrochloric, Nitric, and Chloric Acids in the Presence of One Another. I. First, the presence of hydrochloric acid is confirmed by pre- cipitating a part of the solution with silver nitrate; a white precipi- tate of silver chloride shows the presence of hydrochloric acid. The remainder of the Solution is treated with silver sulphate solution until no further precipitation of silver chloride takes place, and the precipitate is filtered off. The filtrate is boiled with a little caustic potash (in order to expel any ammonia from ammonium Salts which may be present), treated with zinc dust (or Devarda's alloy), and again boiled; if nitric acid is present, ammonia will be given off. The residue is filtered off, the filtrate acidified with nitric acid and treated with silver nitrate. If a precipitate of silver chloride is now obtained, chloric acid was originally present. * The reaction takes place very slowly in neutral Solutions. f This equation is not correct, for some ClO, is always mixed with the Cl. The following equation expresses this: KClO,--2HCl=KCl +Cl+CIO,--H,O. PERCHLORIC ACID, 345 II. Or, a small part of the solution is tested for hydrochloric acid by adding silver nitrate in excess, the precipitate is filtered off, the filtrate treated with sulphurous acid, and again tested with silver nitrate, when a precipitate of silver chloride shows the pres- ence of chloric acid. A second portion of the Solution is tested, as above, for nitric acid. sº REACTIONS IN THE DEY WAY. On ignition, all chlorates are decomposed, forming a chloride, with loss of oxygen. By heating on charcoal, deflagration takes place. PERCHLORIC ACID, HClO,. Free perchloric acid is obtained by the distillation of potassium perchlorate with concentrated sulphuric acid. In this way the solid, crystalline hydrate HClO,--H,O is obtained, which, on heat- ing to 110°C., is broken up, the anhydrous, strongly-fuming-in-the- air, liquid acid distilling off, while the oily hydrate HClO,4-2H,O. remains behind, which itself distils at 203°C. The free acid is very dangerous, and often explodes sponta- neously. In aqueous solution, however, it can be kept without danger. The salts of this monobasic acid are remarkably stable. The potassium salt is obtained from potassium chlorate. On melting the latter compound, at first a lively stream of oxygen is given off which, however, soon lessens. The melt quickly be- comes viscous, and consists of potassium chloride and potassium. perchlorate, 2KClO,-KC1-i-KClO,--O, and the latter may be separated from the much more soluble po- tassium chloride by recrystallization. Solubility.—All perchlorates are soluble in water. REACTIONS IN THE WET WAY. Perchloric acid is not attacked by concentrated sulphuric acid, nor reduced to chloride by zinc dust, Devarda's alloy, sulphurous. acid, or acid solutions of ferrous Salts. 346 REACTIONS OF THE METALLOIDS. Potassium Salts precipitate the relatively insoluble, white, crystalline KCIO, (cf. page 40). Silver Nitrate and Barium Chloride produce no precipitation. REACTIONS IN THE DRY WAY. The perchlorates deflagrate on being heated on charcoal; by fusing they lose oxygen, leaving chloride behind, which when dis- solved in water gives all the reactions for hydrochloric acid. PERSULPHURIC ACID, H.S,O,. Pure persulphuric acid itself has never been isolated, its solution in Sulphuric acid being alone known. It was first prepared by M. Marshall, who electrolyzed tolerably dilute sulphuric acid, keep- ing it very cold. In this process the ions H and HSO, of sulphuric acid unite together at the anode. forming persulphuric acid: H.SO, H HSO, = | + | . H.SO, H HSO, The preparation of ammonium persulphate, from which all other persulphates are made, is entirely analogous. The most important Salts of persulphuric acid are those of am- monium, potassium, and barium. (NHA),S,Os is readily soluble in water, and forms monoclinic crystals; K.S.O.s is difficultly soluble in cold water, but much more soluble in hot water, from which solution it is obtained by rapid cooling in the form of long crystals; BaS,Os--4H.O is made by rubbing ammonium persulphate with barium hydroxide, and is fairly soluble in water. REACTIONS. A solution of ammonium persulphate may be used. 1. All persulphates are decomposed in aqueous solution (slowly in the cold, but more quickly on warming), forming Sulphate, free sulphuric acid, and Oxygen: KSO, HNo._ te (b) Bağ #XO-Baso. HsO4·0. PERSULPHURIC ACID. 347 A large proportion of the oxygen escapes as ozone, which can be detected by its odor, or by its property of turning iodo-starch paper blue. A dilute solution of ammonium persulphate decomposes slowly at 20°C., without evolution of oxygen, part of the nitrogen being oxidized to nitric acid: 8(NH,),S,Os-1-6H.O=14(NH,)HSO,--2H,SO,--2HNO,. 2. Dilute Sulphuric Acid acts the same as water. 3. Concentrated Sulphuric Acid.—If a solid persulphate is dissolved in concentrated Sulphuric acid at 0° C., a liquid is ob- tained which possesses very strong oxidizing properties. The mix- ture is known as Caro's acid.* For further particulars concerning this acid, the student is referred to A. Baeyer's interesting work, B. B. XXXIV (1901), page 853. 4. Silver Nitrate precipitates black silver peroxide: 2AgNO,--K.S. Os-i-2H,0=2KHSO,--2HNO,--Ag,O,. If, however, the concentrated solution of ammonium persul- phate is treated with ammonia and a very little silver nitrate, a lively evolution of nitrogen takes place, and the Solution becomes heated to boiling. Silver peroxide is formed first, and oxidizes the ammonia to water, setting free nitrogen (Catalysis).f 5. Manganese, Cobalt, Nickel, and Lead Salts are oxidized in the presence of alkali to black peroxides : Mn(OH), --K,S,Os--H,O=2KHSO,--H,MnOs. In this last reaction persulphuric acid reacts exactly similar to hydrogen peroxide. It may be distinguished, however, from the latter by the fact that it does not decolorize a solution of potas- sium permanganate, does not produce a yellow coloration with titanium sulphate, and does not react with chromic acid to form chromium peroxide (cf. page 84). Ferrous salts are readily oxi- dized to ferric salts, and cerous salts are changed to yellow ceric salts by persulphates, but the latter are not decolorized by an excess of the persulphate, while they are by hydrogen peroxide, * Zeitschr. f. angew. Ch., 1898, p. 845. f Zeitschr. f. phys. Chem., XXXVII (1901), p. 255. 348 REACTIONS OF THE METALLOIDS, 6. Barium Chloride does not give a precipitate immediately in a freshly-prepared cold Solution of a persulphate; but, on standing some time, or on boiling, insoluble barium sulphate is precipitated. GROUP VI. Silver Nitrate produces no precipitate. Barium Chloride produces a white precipitate, almost insoluble in acids. SULPHURIC ACID, H,SO,. Pure sulphuric acid at ordinary temperatures is a colorless, oily liquid of specific gravity 1.8384; at low temperatures it is a solid. If the acid is subjected to distillation, it is always partially decom- posed; heavy, white vapors of SO, are given off first, and at 338°C. a 98 per cent. acid distils over. Ordinary commercial sulphuric acid has a specific gravity of 1.83–1.84, and contains 93–96 per cent. H.SO,. It often contains lead sulphate, Selenic acid, platinum, pal- ladium, arsenious acid, the nitrogen acids. and Small amounts of organic matter (whereby it is often colored brown) as impurities. Concentrated sulphuric acid is very hygroscopic, and is, there- fore, used for drying gases, etc. The anhydride of sulphuric acid. SO, dissolves in concentrated sulphuric acid, forming pyrosulphuric acid, H.S.O. which is solid at ordinary temperatures, melts at 35° C., and loses SO, at higher temperatures. It fumes strongly, and is called, therefore, fuming sulphuric acid. Sulphuric acid is dibasic and forms both neutral and acid salts. Solubility.—Most sulphates are soluble in water; calcium sul- phate is difficultly soluble, strontium and lead sulphates are very difficultly soluble, while barium sulphate is practically insoluble in water, There are also a number of basic sulphates, Hg, Bi, Gr, which are insoluble in water, but are, as a rule, Soluble in acid. REACTIONS IN THE WET WAY. 1. Sulphuric Acid naturally gives no reaction. 2. Silver Nitrate causes no precipitation in dilute solutions, but in concentrated solutions a white crystalline precipitate is formed HYDROFLUORIC ACID. 349 (100 parts of water dissolve at 18°C. only 0.58 part of silver sul- phate). 3. Barium Chloride precipitates, from even the most dilute solutions, white barium sulphate, insoluble in acids. 4. Lead Acetate precipitates white lead sulphate, soluble in concentrated Sulphuric acid, ammonium acetate, and ammonium tartrate solutions (cf. page 162). In order to detect the presence of the SO, ion in insoluble sul- phates, they are treated with sodium carbonate, whereby insoluble carbonate and soluble Sodium sulphate are formed. Lead sulphate and calcium Sulphate are easily decomposed by boiling with sodium carbonate solution, but barium and strontium sulphates are only incompletely decomposed by this treatment; they are much more readily attacked by fusing with four times as much sodium carbonate (cf. page 61). 5. By Nascent Hydrogen (zinc and acid) the Sulphates are not reduced. REACTIONS IN THE DEY WAY. The neutral salts of the alkalies melt with difficulty without being decomposed, while the acid salts of the alkalies readily give off water and SO, (cf. page 74). The sulphates of the alkaline earths and of lead do not undergo decomposition on ignition, the remaining Sulphates are more or less decomposed. All sulphates are reduced to sodium sulphide when heated with sodium carbonate on charcoal; if the product is placed upon a bright silver coin and moistened, a black stain of silver sulphide results; e.g.: (a) CaSO,--Na,CO. = CaCO,-- Na,SO, ; (b) Na,SO,--2C=2CO,--Na,S; (c) Na,S+ Ag,4-H,0+O=2NaOH-i-Ag,S. This reaction is called the Hepar reaction. HYDROFLUORIC ACID, HF. Occurrence.—Hydrofluoric acid occurs in nature only in the form of fluorides, of which the most important is fluorite, CaF, crystalliz- ing in the isometric system. It is also found as cryolite, (AlF.)Na, 35o REACTIONS OF THE METALLOIDS. in Greenland, and in many silicates, such as turmaline, topaz, lepidolite, apophyllite, apatite, etc. Properties.—Hydrofluoric acid at temperatures above 20° C. is a colorless gas which changes at 19.4° C. to a mobile, fuming liquid. The vapors possess a penetrating odor and are poisonous. When in contact with the skin it produces painful burns. On heating the concentrated aqueous Solution, the gas HF at first distils and then the 36 per cent. acid unchanged. Hydrofluoric acid is distinguished from all other acids by its ability to dissolve silicic acid, a property which is utilized technic- ally for etching glass, and in the analytical laboratory for detecting fluorine and silicic acid, as well as for decomposing silicates. On account of this action upon glass the acid must be kept in platinum, wax, or hard rubber, and prepared in platinum or lead vessels. The action of hydrofluoric acid upon silicic acid takes place according to the equation SiO2+4HF=2H2O+SiF4, and the velocity of the reaction depends upon the fineness and nature of the material. Precipitated and ignited silica reacts violently with concentrated hydrofluoric acid, while quartz powder dissolves but slowly. Most silicates are attacked more readily than quartZ. It is a weak, monobasic acid, reddens litmus paper, and colors Brazil-wood paper yellow (as do other weak acids, e.g., carbonic and acetic). Solutions of the alkali fluorides react strongly alkaline. The property of forming acid salts is characteristic of this acid, as in the case of hydrocyanic acid; thus the following are known: HIAgF2], HKF2], HINaF2], HINH4F2], Nag|Feſ'6], Nag[AlF6], etc. The free acids corresponding to the above salts either do not exist or are extremely unstable. Silver hydrogen fluoride, HIAgF2], is decomposed on gentle heating into AgF and HF, while the corre- sponding alkali compounds are decomposed only on strong ignition; for this reason the latter are suitable for attacking difficultly decom- posable silicates, zircon and titanium minerals, etc., which are only partially attacked by free hydrofluoric acid, HYDROFLUORIC ACID. 35 I , Solubility.—The fluorides of the alkalies, of silver, aluminium, tin, and mercury are soluble in water, while those of the alkaline earths, of lead, copper, and zinc, are insoluble, or at least very difficultly soluble. REACTIONS IN THE WET WAY. For reactions 1, 2, and 3, powdered calcium fluoride may be used, but for reactions 4, 5, 6, and 7, a solution of Sodium fluoride is neces- sary. * 1. Dilute Sulphuric Acid causes only a slight reaction. 2. Concentrated Sulphuric Acid reacts readily on warming, setting free hydrofluoric acid: CaF,--H,SO,-CaSO,--2HF. If this reaction is performed in a test-tube, the hydrofluoric acid will attack the glass, forming volatile silicon fluoride and salts of hydrofluosilicic acid; but the latter are decomposed by the con- centrated sulphuric acid into Sulphate, hydrofluoric acid, and sili- con fluoride, Na,CaSiO44-28HF=14H,0+Na,SiF.--CaSiF.--4SiF, Söda glass . and Na,SiF.--H,SO,-Na,SO,--2HF-H SiR, CaSiF.--H,SO, -CaSO,--2HF--SiF. The silicon fluoride formed by this reaction is a colorless gas with a penetrating odor, and is decomposed by water, forming gelatinous silicic acid and hydrofluoric acid: (1) SiF4+4H2O=Si(OH)4-H4HF. Silicon fluoride, however, readily combines with hydrofluorie acid to form hydrofluosilicic acid: (2) SiF4+2HF = H2(SiF6], and the latter compound is not decomposed by water. The whole 35 I a REACTIONS OF THE METALLOIDS. reaction, therefore, which takes place between silicon tetrafluoride and water is expressed by the sum of the two equations, as follows: 3SiF4+4H2O=Si(OH)4-1-2H2(SiF6]. If, therefore, a fluoride be heated in a glass test-tube with con- centrated sulphuric acid, and the escaping vapors allowed to act upon water, by placing a moist glass tube rod in the tube, the .water adhering to the rod will become turbid. Remark.-Although the above test rarely fails when relatively large amounts of fluoride are used, it will not be obtained in the case of certain minerals containing fluorine, such as topaz, turma- line, etc. The test may fail, furthermore, if the fluoride is mixed with a large excess of that modification of silicic acid which is most readily attacked by hydrofluoric acid. According to Daniel,” this ' is due to the formation of a stable oxyfluoride probably of the formula SiOF2. The silicon tetrafluoride at first formed combines with the excess of amorphous silicic acid present, as follows: SiF4+ SiO2=2SiOF2, but this reaction will take place only very slowly if at all with quartz powder or with the silica of a silicate such as glass. A positive result will be obtained invariably when the tetra- fluoride test is made in a platinum vessel with a relatively large amount of fluoride and comparatively little amorphous silicic acid or silicate (large amounts of quartz do not influence the reaction); the test will be negative, on the other hand, if made in platinum with no silicic acid, or, Strange to say, when only quartz is present with the fluoride. The reason for this different behavior lies in the difficulty with which quartz is attacked by hydrofluoric acid. Daniel recommends the following method for performing the test : The substance to be tested for fluorine is mixed with about * Zeitschr. f. anorg. Ch., 38 (1904), p. 299. HYDROFLUORIC ACID. 35Ib three times as much (by volume) ignited quartz powder, and mixed to a paste in a test-tube with concentrated Sulphuric acid. The test-tube is then closed with a cork in which one hole has been bored and an opening cut in the side. Through the hole in the cork is passed a glass rod blackened with asphalt paint, and on the bottom of the rod hangs a drop of water suspended; this rod is pushed down until it is only a distance equal to about 1% the diameter of the test-tube from the paste in the bottom. The tube and its contents are gently heated over a small flame, and if a fluoride is present a white film of Si(OH)4 will be formed in the drop of water and will be shown plainly in contrast to the black rod. In a tube with 1 cm. diameter fluorine equivalent to 1 milligram of calcium fluoride may be detected, while with a tube of only 0.5 cm. diameter as little as 0.1 milligram of calcium fluoride will give the test. When using a tube of Small diameter, it is best to add the sulphuric acid through a small capillary pipette to avoid wetting the sides of the tube. If the substance contains considerable amorphous silica, or when an oxyfluoride, such as topaz, is present, which is hard to decompose with sulphuric acid, the test will fail, and it is then necessary to make use of the etching test. 3. The Etching Test.—The substance to be tested for fluoride is placed in a platinum crucible, treated with concentrated sulphuric acid, and the crucible covered with a watch-glass whose convex side , is provided with a thin coating of beeswax upon which a few letters have been written. On warming the contents of the crucible, the glass will be etched at the places where the escaping gas comes in contact with it, if a fluoride was originally present. By covering the upper concave side of the watch-glass with a little cold water, the wax coating will not melt during the experiment. If it is desired to detect the presence of a trace of fluorine, the crucible is allowed to stand covered with the watch-glass for twelve hours and then heated for a few minutes. The presence of only 0.0003 gm. of calcium fluoride is sufficient to give this test, provided a crucible of the right size is used. If the fluoride contained silica (as in topaz, turmaline, and other minerals), the etching test will be negative, for even if the fluorine 352 REACTIONS OF THE METALLOIDS. escapes, it will be in the form of silicon fluoride, which does not attack glass. In order to detect Small amounts of fluoride in silicates, it is necessary first to transform the fluorine into calcium fluoride and to subject the latter compound to the test. In order to separate the fluorine as calcium fluoride, the follow- ing process is used: The finely-pulverized silicate is mixed with 6–8 times as much sodium carbonate, fused in a platinum crucible, and the melt treated with water after it is cold. A Solution is then obtained in which all of the fluorine is present as Sodium fluoride, together with sodium silicate. The silica is removed by adding considerable ammonium carbonate to the Solution, warming it slightly, and allowing it to stand twelve hours. After filtering off the silica, the solution is evaporated to a small volume, and a little phenolphthalèin added, which will cause the solution to become reddish. Hydrochloric acid is carefully stirred in until the solution becomes colorless, and it is heated to boiling, when the color will reappear. The solution is again decolorized with hydrochloric acid after it has become cold, and the process repeated until the solution becomes only faintly colored on boiling it. Calcium chloride solution is now added and the solution again boiled. The precipitate formed consists of calcium carbonate and calcium fluoride, which is filtered off, washed, dried, and ignited in a platinum crucible. The ash is then treated with dilute acetic acid, evaporated to dryness, triturated with water and the undissolved calcium fluoride filtered off. After drying the precipitate and burning the filter, it is ready for the etching test. 4. Silver Nitrate causes no precipitation from solutions of soluble fluorides. 5. Barium Acetate precipitates barium fluoride soluble in an excess of mineral acid and in ammonium salts. Traces of fluorine present as preservative in foods, liquors. etc., may be detected by adding a little potassium sulphate (about 0.3 gm.) to the solution, heating it to boiling and slowly introducing 10 c.c. of 10 per cent. barium acetate solution. The precipitate of barium sulphate and fluoride is then subjected to the etching test.* 6. Calcium Chloride gives a white, slimy precipitate, difficultly soluble in hydrochloric and nitric acids, but almost entirely insolu- * Cf. Blarez, Chem. News, 91,39; also Woodman and Talbot, J. Am. Chem. Soc. XXVII, p. 1437. HYDROFLUORIC ACID, 353 - ble in acetic acid. On account of its slimy consistency, the precipi- tated calcium fluoride is extremely hard to filter; but by precipi- tating it in the presence of calcium carbonate a mixture is obtained which can be readily filtered. After igniting and treating with acetic acid, the precipitate is changed to Soluble calcium acetate and insoluble calcium fluoride; it is now much denser and can be filtered readily. 7. Ferric Chloride produces in concentrated solutions of alkali fluorides a white, crystalline precipitate corresponding to the gen- eral formula [FeF.]X. These salts, which are analogous to cryo- lite, [AlF.]Nas, are difficultly soluble in water, and their saturated, aqueous solutions do not give the iron reaction upon the addition of potassium sulphocyanate, except after the addition of acid. These complex fluorides also are incompletely decomposed by ammonia, basic fluorides being formed. Methods for Getting Insoluble Fluorides into Solution. (a) Calcium fluoride alone cannot be completely decomposed by fusing with sodium carbonate. The aqueous Solution of the melt always contains a considerable amount of Sodium fluoride, but never the total amount of the fluorine. If, however, the fluoride is mixed with silica or a silicate, complete decomposition can be effected by fusing with Sodium carbonate, 6CaF,--6SiO,-20aSiF.--4CaSiO, and * GaSiF.--4Na,CO,-CaCO,--Na,SiO,--3CO,--6NaF. On treating the melt with water, sodium fluoride and sodium sili- cate go into solution; the silicic acid can be removed from the solu- tion by the process described on page 352. (b) All fluorides are decomposed by heating with concentrated sulphuric acid, being changed to Sulphates. REACTIONS IN THE DRY WAY. Most fluorides are unchanged by ignition. By heating them with silica in moist air, they are all more or less completely decom- posed: CaF,4-H,0+ SiO,-CaSiO,--2HF. 354 - REACTIONS OF THE METALLOIDS. The acid fluorides give off hydrofluoric acid on ignition, whereby the glass tube in which they are heated becomes etched. HYDROFLUOSILICIC ACID, H.SiF. As we have seen, this acid is formed by the action of silicon fluo- ride upon water: 3SiF,4-4H,0=2H,SiF.--Si(OH), If the silicic acid is filtered off, a strongly-acid solution is ob- tained containing hydrofluosilicic acid. By evaporating the solu- tion, the acid is decomposed into silicon fluoride and hydrofluoric acid, H,SiF,-SiF,4-2HF, so that hydrofluosilicic acid itself is only known in aqueous solution, although its salts are very stable. Solubility.—Most silicofluorides are soluble in water; the potas- sium and barium salts form exceptions, being difficultly soluble in water and insoluble in alcohol. REACTIONS IN THE WET WAY. A solution of sodium silicofluoride should be used. 1. Dilute Sulphuric Acid causes only a very slight decompo- sition. 2. Concentrated Sulphuric Acid decomposes all silicofluorides, evolving silicon fluoride and hydrofluoric acid: Na,SiF--H,SO,-Na,SO,--SiF,--2HF. If the reaction is performed in a platinum crucible, the escaping gaS will etch glass, and will cause a drop of water to become turbid. 3. Silver Nitrate produces no precipitation. 4. Barium Chloride gives a crystalline precipitate (1 part of the salt dissolves at 17°C. in 3731 parts of water). 5. Potassium Chloride produces, from solutions which are not too dilute, a gelatinous precipitate of potassium silicofluoride, which is difficultly soluble in water (833 parts of water dissolve at 17° C. SILICIC ACID. 355 one part of the salt) and much more insoluble in an excess of potas- sium chloride or in alcohol, but soluble in ammonium chloride. 6. Ammonia decomposes all soluble silicofluorides, with separa- tion of silicic acid: Na2SiF6+4NH4OH=2NaP+4NH4F +Si(OH)4. 7. Potassium and Sodium Hydroxides react in the same way as ammonia, but the silicic acid remains in Solution as alkali sili- cate. z REACTIONS IN THE DRY WAY. All silicofluorides are decomposed on being heated into fluoride of the metal and silicon fluoride: K.SiF,-2KF-H SiR,. The escaping gas renders a drop of water turbid, and the residue gives all the reactions of a fluoride. GROUP VII. NON-VOLATILE ACIDS WHICH FORM SOLUBLE SALTS WITH THE ALKALIES. OH zº 29 SILICIC ACID, Si and Si3 OH.” —OH NOH NOH Occurrence.—The above acids, from which very stable salts are derived, are not known in the free state, as is the case with car- bonic and sulphurous acids; although there are indeed amorphous, natural minerals consisting of hydrated silica with varying amounts of water: water opal with about 36 per cent, water, ordinary opal with from 3 to 13 per cent. water, and hyalite with about 3 per cent. water; but none of these substances represents a compound of constant composition. The anhydride SiO, occurs in rhombohedral crystals as quartz, whose prismatic faces are almost always striated horizontally; and as tridymite, also crystallizing in the hexagonal system. The amorphous silicic acid is often found mixed with the crystallized anhydride as flint, agate, chalcedony, jasper, etc. Silicic acid is, however, most frequently found in the form of its salts, the silicates. * To this group also belong titanic, zirconic, tantalic, niobic, and tungstic acids. 356 REACTIONS OF THE METALLOIDS. Preparation and Properties.—Silicic acid can be very readily ob- tained pure by the hydrolysis of its fluoride, 3SiF.--4H,O=2H,SiF.--Si(OH), or by the decomposition of alkali silicates (water-glass) with acids: Na,SiO,--2HCl=2NaCl-H SiO, H, The silicic acid thus obtained forms an amorphous, gelatinous mass, appreciably Soluble in water and acids, and readily soluble in even dilute Solutions of caustic alkalies or alkaline carbonates. Thus freshly-precipitated silicic acid will be readily and completely dissolved by a short digestion with 5 per cent. (or even 1 per cent.) sodium carbonate Solution on the water-bath. On being dried, silicic acid gradually loses water, and at a gentle red heat is changed into the form of its anhydride. According to the extent to which the dehydration has gone, the Solubility of the silicic acid diminishes both in acids and in alkalies. 1. Air-dried silicic acid, with 16.65 per cent. of water, corre- sponding to the formula 3SiO,-2H,0, is perceptibly soluble in acids, and completely dissolved by digestion for 4 to # an hour with 1 per cent. Soda Solution on the water-bath. - 2. Silicic acid dried at 100° with 13.60 per cent. of water, corre- sponding to the formula 2SiO, H.O, is practically insoluble in acids, but can be dissolved by digesting for 4 hour with 1 per cent. Sodium carbonate solution upon the water-bath, or more readily by boiling. 3. Silicic acid dried at 200°, with 5.66 per cent. of water, cor- responding to the formula 5SiO, H.O, and the acid dried at 300°, with 3.40 per cent. of water, corresponding to the formula 9SiO, H.O, dissolve slowly by digestion with 1 per cent. Sodium carbonate Solu- tion on the water-bath. 4. The anhydride obtained by gentle ignition to a faint-red heat is only partly dissolved by 1 per cent. or by 5 per cent. Sodium carbonate after half an hour's digestion on the water-bath; but is dissolved after boiling for two hours with the Sodium carbonate solution. 5. The strongly-ignited anhydride is only soluble in 5 per cent. boiling sodium carbonate solution after repeated boiling for a long SILICIC ACID. 357 time, but is readily dissolved by boiling with concentrated caustic soda or potash. 6. The native anhydride, quartz, after being powdered in an agate mortar, is practically insoluble in 5 per cent. Sodium carbonate Solu- tion, and extremely difficultly soluble in boiling caustic alkali. If it is in the form of an extremely fine powder, it can be dissolved by boiling with 5 per cent. Sodium carbonate solution (Lunge and Mill- berg). It follows from the above that the solubility of silicic acid (and of its anhydride) in alkali carbonates depends largely upon the fine- ness of the material. Silicic acid, as well as its anhydride, is soluble in aqueous hydro- fluoric acid, forming hydrofluosilicic acid: SiO,--6HF=2H,0+H,SiF. By evaporating this solution hydrofluoric acid is evolved; and silicon fluoride, with small amounts of silicic acid, is left behind. In order, then, to completely volatilize silicic acid by means of hydro- fluoric acid, the hydrolytic action of water must be prevented, which is effected by the addition of a little concentrated sulphuric acid. The process is as follows: A little water, not more than ; c.c. of concentrated sulphuric acid, and then the hydrofluoric acid, are added to the silica, and the mixture evaporated on the water-bath until the mass no longer smells of hydrofluoric acid, when the sulphuric acid is driven off by heating the crucible in an inclined position over a small flame. If considerable silicic acid is present, the operation must be repeated two, or perhaps three, times. The Salts of silicic acid, the silicates, are exceedingly numerous, and are usually very stable. Many of them are so stable that they are not attacked by concentrated acids, while others are easily de- composed thereby. The different silicates are classified according to their solubility into A. Water-soluble silicates. B. Water-insoluble silicates, which are again divided into (a) Silicates decomposable by acids; (b) Silicates undecomposable by acids. 358 REACTIONS OF THE METALLOIDS. A. Water-soluble Silicates. The silicates which are soluble in water, or “water-glasses,” are obtained by fusing silica or a silicate with caustic alkali or alkali carbonate: SiO,--Na,CO,-Na,SiO,--CO,. 1. Behavior Towards Acids.-If a solution of water-glass is treated with dilute hydrochloric acid, a gelatinous precipitate is formed (if the Solution is not too dilute) consisting of silicic acid: Na,SiO,--2HCl=2NaCl-H,SiO, This precipitation is by no means quantitative; considerable amounts of silicic acid always remain in Solution, and under some conditions all of it can remain dissolved in the acid. If the water- glass Solution is quickly poured into an excess of not too dilute hydrochloric acid, there will be no precipitation of silicic acid, but after standing some time the whole Solution is changed to a jelly. Possibly the soluble Orthosilicic acid is first formed, but it gradually loses water and is changed into difficultly soluble metasilicic acid. The silicic acid which is precipitated upon the addition of acid is, therefore, considerably soluble in dilute acids. In order to com- pletely separate the silicic acid from a Solution of water-glass, the hydrated acid must be changed into the less hydrated acid, 2SiO, H.O, by heating at 100° C. (cf. page 356). For this purpose the water-glass solution is acidified with hydrochloric acid (or nitric or sulphuric acid) and evaporated on the water-bath to complete dryness (the mass must no longer smell of acid). The dry residue is treated with water to which a little acid has been added, and after filtration very little silicic acid remains dissolved.* 2. Behavior Toward Ammonium Salts.-If a solution of water- glass is treated with an ammonium salt, the silicic acid will, for the most part, be precipitated as hydroxide; the precipitation is not quite quantitative, but more complete than is obtained by the addition of cold dilute acid: Na,SiO,-i-2NH,Cl-H2H,0=2NaCl-H2NH,OH-i-H.SiO, Na,SiO,--(NHA),CO,--2H,0=Na,CO,--2NH4OH-i-H,SiO,. * The small amount of silica remaining in solution can be removed almost entirely by a second evaporation of the filtrate. Cf. Hillebrand, J. Am. Chem. Soc., 1902, p. 362. SILICIC ACID. 359 The precipitation by means of ammonium carbonate is less com- plete than in the case of an ammonium salt of a stronger acid, on account of the formation of sodium carbonate, which exerts a solvent action upon the precipitate. We are, nevertheless, often compelled to use this reagent, as, for example, when it is desired to test a sili- cate for chlorine or fluorine (cf. page 352). Silicic acid is more completely precipitated by zinc ammonium hydroxide than by ammonium carbonate, * Na,SiO,--[Zn(NH,).](OH), –2NaOH+ 6NH,-- ZnSiO, because the zinc silicate formed by the reaction is much more diffi- cultly soluble in dilute alkaline solution than is the free silicic acid. The separation of silicic acid from a solution of water-glass by means of ammonium carbonate may be illustrated by a common case. Many rocks (particularly the zircon-syenite of Norway and Greenland, many granites and basalts) contain Small amounts of so- dalite, NaCl. 3NaAlSiO, a chloride silicate of the leucite group. In order to detect the chlorine in such a rock, the following process may be used: The finely-powdered silicate is fused with six times as much sodium carbonate in a platinum crucible, and the product of the fusion is treated with cold water and filtered. The filtrate contains all of the chlorine as Sodium chloride in the presence of sodium silicate. The Solution is treated with ammonium car- bonate, warmed gently, allowed to stand twelve hours, and the precipitated silicic acid filtered off. In order to separate the rest of the silicic acid, a little zinc ammonium hydroxide * is added and the solution boiled until it no longer Smells of ammonia. The pre- cipitated zinc silicate and zinc oxide are filtered off, the solution acidified with nitric acid and tested with silver nitrate for chlorine. B. Silicates Insoluble in Water. (a) Decomposable by Acids. A large number of native silicates are decomposed by evapora- tion with hydrochloric acid, the silica being deposited sometimes in the form of a jelly and sometimes in the form of a powdery mass. * Prepared by dissolving pure zinc in nitric acid, treating the solution with potassium hydroxide solution until it is neutral, and dissolving the fil- tered and washed zinc hydroxide in ammonia. 36o REACTIONS OF THE METALLOIDS. All zeolites, and a number of artificial silicates (such as Portland and Roman cements) belong to this class of silicates. In order to remove all of the silicic acid from these silicates, the finely-powdered mineral is treated with dilute hydrochloric acid, evaporated on the water-bath to dryness, the mass moistened with concentrated hydrochloric acid,” allowed to stand for 20 minutes, hot water added, and the Solution boiled and filtered. The silicic acid is left on the filter, and the filtrate contains the metals as chlo- rides.f The purity of the residual silicic acid must always be tested. For this purpose the well-washed precipitate is placed, together with the filter-paper, in a clean platinum crucible, held in an inclined position on a triangle, and the filter-paper carefully burnt. The residue is treated with about 2 c.c. of water, a drop of concentrated sulphuric acid, and about 3–5 c.c. of pure hydrofluoric acid, and evaporated as far as possible on the water-bath. The excess of sul- phuric acid is then removed by cautious heating over the free flame. If the silicic acid were pure, nothing should remain after the evapo- ration of the sulphuric acid. Almost always a small residue of aluminium and ferric oxides remains, which in most cases can be neglected. If considerable residue is left, it should always be tested for titanic acid and barium sulphate. (3) Silicates Undecomposable by Acids. Most silicates, the feldspars, micas, artificial glasses, porcelain. etc., belong to this class. In order to remove the silicic acid from such substances, they must be 1. Fused with an alkali carbonate, 2. Fused with lead oxide or boron trioxide, or 3. Heated with Sulphuric and hydrofluoric acids. 1. Fusion with an Alkali Carbonate.—This method is commonly used when it is desired to detect the presence of silicic acid and of the bases with the exception of the alkalies. The finely-powdered substance is mixed with 4–6 times as much * The moistening with concentrated hydrochloric acid serves to convert any oxides or basic salts of iron, aluminium, magnesium, etc., into soluble chlorides. f Cf. page 358.-Foot note. ^ SILICIC ACID. 361 calcined sodium carbonate (or a mixture of equal parts of Sodium and potassium carbonates, which melts lower than Sodium carbo- nate alone), placed in a platinum crucible and heated carefully at first to avoid spattering from too violent evolution of carbon dioxide. The temperature is gradually increased until the full heat of the burner is reached, and the fusion continued until the molten mass is quiet, when it is heated for about a quarter of an hour over the blast lamp. The crucible is then cooled quickly by dipping it while still glowing into cold distilled water, whereby the fused mass can usually be removed from the crucible without difficulty.* The product of the fusion is treated as previously described under (a), being now converted into Sodium silicate or a silicate soluble in acid: C ºliºst 3Na,CO,-2Na,SiO,--CaCO,--2AlO(ONa)+2CO,. northite By evaporation with hydrochloric acid, the following reaction takes place: 2Na,SiO,--CaCO,--2A]O(ONa)+14HCl= =6NaCl-HCaCl, H2A]Cl,--2H,SiO,--CO,--5H,O. 2. Fusion with Lead Oacide or Borom Trioſcide.—This method is very rarely used in qualitative analysis, so that it will not be necessary to describe it here. It plays an important part in quan- titative analysis, and will, therefore, be described in the second volume of this book. 3. Decomposition by Hydrofluoric Acid.—This method is princi- pally used when a silicate is to be examined for alkalies, titanic acid or barium. The finely-powdered silicate is placed in a platinum dish, about 2 c.c. of pure Sulphuric acid (1 vol. concentrated acid and 2 vols. of water) and about 5 c.c. of freshly-distilled hydro- fluoric acid are added, and the mixture evaporated on the water, bath, stirring the mass from time to time with a thick platinum wire until it no longer Smells of hydrofluoric acid. 5 c.c. more of hydrofluoric acid are added and again evaporated when the dish is heated very carefully over the free flame under a good hood. until the greater part of the sulphuric acid is expelled f After * For another excellent method for removing the melt from the crucible see Talbot's Quantitative Chemical Analysis, p 32 f The mass should not fuse, for a part of the sulphate may then be changed to an oxide insoluble in water The sulphates of iron and aluminium, for example, are decomposed on ignition. 362 REACTIONS OF THE METALLOIDS. cooling, the mass is treated with water, when usually everything will by degrees go into Solution. If a residue remains, it is tested for barium sulphate and titanic acid. The solution can be used for the alkali tests, or for the tests for the other metals if it is desired. REACTIONS IN THE DRY WAY. If silicic acid or a silicate is heated in the salt of phosphorus bead, the metallic oxide will dissolve, while the silicic acid itself will be left as a white gelatinous mass, suspended in the bead (skeleton bead). This reaction is, however, not infallible, for certain silicates of the Zeolite group dissolve in the bead without the formation of the skeleton. SILICON, Si. At. Wt. 28.4. Silicon exists in two modifications, one of which is crystalline, while the other is amorphous. Amorphous silicon is a dark-brown powder, which can be oxidized by heating in the air while the crys- talline modification remains unchanged on ignition in pure air or in oxygen, but if the air contains carbon dioxide, it is oxidized to sili- con dioxide with deposition of carbon: CO,--Si-SiO,--C. Crystallized silicon is not attacked by any acid, but is readily dissolved by boiling with concentrated caustic alkali with evolution of hydrogen: Si-H2NaOH+H,0=Na,SiO,--2H, Silicon unites with many metals, forming silicides. The silicides of the light metals, magnesium, calcium, etc., are decomposed by dilute hydrochloric acid with the formation of spontaneously-com- bustible silicon hydride *: SiMg, +4HCl=2MgCl, --SiH, In order to detect the presence of silicon in such a compound it is treated with nitric acid, which oxidizes the greater part of the silicon to silicic acid. * Silicon hydride in the pure state is not spontaneously combustible, ex- cept when mixed with hydrogen, which is almost always the case. SILICON, 363 If it is a question of detecting the presence of silicon in the differ- ent kinds of irons (steel, cast iron, etc.), a large amount of material should be taken, for the amount of iron silicide present is usually very small. Five or ten grams of the material to be tested (best in the form of borings) are placed in a large beaker and treated with 60 c.c. of nitric acid (sp. gr. 1.2). A violent reaction at once takes place with evolution of brown nitrous fumes. As soon as this action lessens. it is heated to boiling, and the heating continued until no more brown fumes are given off, when it is poured into a 200-c.c. evaporating-dish of Berlin porcelain and evaporated as far as possi- ble upon the water-bath. It is then heated, with constant stirring, over a free flame until completely dry. The mass is then ignited until the nitrate is completely changed to oxide, when no more brown fumes will be evolved. After cooling, the mass is dissolved in about 50 c.c. of concentrated hydrochloric acid, heated with con- stant stirring almost to boiling, evaporated almost to dryness, taken up in water, filtered, and the residue tested for silicic acid, by seeing whether it is volatile with sulphuric and hydrofluoric acids. In the case of pig iron, the silicic acid obtained is usually consid- erably contaminated with graphite, which can be removed by long ignition in a platinum crucible before treating with hydrofluoric and sulphuric acids. Other silicides, such as carborundum. SiC. are not decomposed by nitric acid, they can be fused with caustic alkali in a silver crucible, SiC+4KOH+2H,0=K,SiO,--K,CO,--4H,. and on acidifying the melt, the silicic acid separates out.* This last method is often used for getting many metallic silicides into solution. Many copper silicon alloys are scarcely attacked by even aqua regia. If, however, they are fused with caustic alkali in a silver crucible, potassium silicate, metallic copper, and hydro- gen are formed: SiCu,4-2KOH+H,0=K,SiO,--Cu,4-2H, By treating the melt with water, the soluble potassium silicate may be separated from the copper. * Carborundum in the form of a fine powder is also easily decomposed by fusing with potassium carbonate. On removing the cover of the platinum crucible the blue flame of burning carbon monoxide is apparent. v.** E’AERT II. CO UAESAE OF AAVA_Z, YSIS. THE purpose of a qualitative analysis is not simply to find out what elements are contained in a given substance, but it should also give a good idea of the relative amounts that are present. Manganese chloride, for example, is made from pyrolusite, and almost always contains traces of calcium, magnesium, nickel, cobalt, and iron. If the analyst should report that “the analyzed substance consists of chlorides of calcium, magnesium, nickel. cobalt, iron, and manga- nese,” it is evident that one would get but a poor idea of the nature of the substance. The report should read: “The substance exam- ined was manganese chloride, and contained traces of calcium, mag. nesium. etc., as impurities.” º In order to be able to estimate the relative amounts of the differ. ent components of a substance, it is necessary to start with a known amount (usually 3–1 gm.) and compare the size of the precipitates produced. It will be impossible for the beginner to estimate the amount of a precipitate obtained, if he has studied the reactions of the elements with unknown amounts of the different substances. lf he has learned to work with a known amount of material, he will soon be able to judge from the size of a precipitate the amount of element to which it corresponds. Every analysis should be divided into three parts: I. The preliminary eaſamination. II. The eacamination for the metals (cations). III. The earamination for the negative elements (anions). The substance analyzed can be A. Solid and non-metallic. B. A metal or an alloy. C. A solution (liquid). D. A gas. 364 PRELIMINARY EXAMINATION. 365 The whole amount of the substance at hand should never be used for the first analysis, but a portion should always be reserved for unforeseen accidents. The portion taken for analysis should be divided into two parts after the preliminary examination, the first part being used for the tests for the electro-positive and the other part for the tests for the electro-negative elements. Before beginning an analysis, the substance should be carefully examined with the naked eye and with the microscope, and the results noted. Oftentimes the Smell, color, and crystalline form suffice to give important clues as to the nature of the substance. A. THE SUBSTANCE IS SOLID AND NON-METALLIC. I. PRELIMINARY EXAMINATION. This should never be omitted, for it often shows how the subse- quent analysis may be considerably shortened, and in some cases makes the further examination unnecessary. It consists only of making the following few simple tests: 1. Heating in the Closed Tube.—By a closed tube is understood a small test-tube, or a glass tube sealed at One end. A small amount of the substance to be tested is placed in the tube so that none of it remains adhering to the sides. The tube is held in a nearly hori- zontal position and cautiously heated in the flame, while the ob- server notes carefully whether any changes take place. The Substance is Volatile. (a) The substance sublimes completely without any deposition of water; it contains no non-volatile Substance. The sublimate is white. The halogen compounds with ammonia, mercurous chloride and bromide, mercuric amidochloride, arsenic trioxide and arsenic pentoxide * may be present. The sublimate is colored— Gray: all oxygen compounds of mercury, cyanide of mercury,i free iodine, and arsenic. Yellow: arsenic sulphide, Sulphur, mercuric iodide.j: Grayish black: mercuric sulphide. ,” (b) The substance is completely volatile, with separation of * Arsenic pentoxide melts before being changed into the trioxide. f Mercuric cyanide leaves a brown mass, the paracyanide, which only disappears after long-continued heating. f Mercuric iodide becomes red immediately on being rubbed with a glass rod. 366 COURSE OF ANALYSIS. Water and gaseous products: most ammonium compounds (with the exception of those of the halogens) and free oxalic acid.* The Substance is Only Partly Volatile. In this case gases and vapors may be evolved: Oxygen from peroxides, nitrates, chlorates, iodates, etc. Carbon dioxide from carbonates and organic substances; in the latter case it is usually accompanied with the separation of carbon and evolution of empyreumatic, combustible vapors. Chlorine from chlorides of platinum, gold, copper, iron, etc. Iodine from iodides, in the presence of oxidizing substances. Sulphur from many sulphides and thiosulphates. Arsenic from arsenites t and arseniates, in the presence of carbon or organic Substances. Water from substances containing water of crystallization, from acid Salts, organic substances, or from the phosphate, borate, chromate, vanadate, and tungstate of ammonium. The water given off condenses in the cooler part of the tube and should be tested with litmus-paper. If it reacts alkaline, it comes from ammonium compounds; if acid, it results from easily decom- posable salts of the stronger acids. Many fluorides when heated with water give off hydrofluoric acid, which etches the glass. If a sublimate is formed, the following experiment is made: 2. A Portion is Mixed with Three Times as Much Calcined Sodium Carbonate and Heated in the Closed Tube.—If ammonium salts are present, the Smell of ammonia can be detected. Mercury compounds give a deposit of gray metal (cf. page 156); arsenic and its oxygen compounds also usually yield the gray metal,f (but no globules) accompanied by a garlic odor. 3. The Substance is Tested in the Bead.—A bead of borax, OI salt of phosphorus, is produced in the loop of a very thin platinum wire (as described on page 23) and introduced with a little of the substance into the oxidizing flame, and the color of the bead ob- served both when it is hot and when it is cold, after which it is * By very cautious heating, oxalic acid may be sublimed; it usually de- composes, however, into water, carbon monoxide, and carbon dioxide. f Arsenites are reduced without the aid of charcoal: Jix 5As(OK)a=3KaAsO,--3K,0+ Asa. f The oxygen compounds of arsenic do not give the metal when heated with pure soda. Commercial soda, however, is usually contaminated with paper threads, which cause the reduction. PRELIMINARY EXAMINATION. 367 heated in the reducing flame. Borax is usually used for this experi- ment, except when it is desired to test for silicic or titanic acids, or when the substance is white, in which case salt of phosphorus is used, Only colored oxides are capable of coloring the borax bead.* The following substances impart a characteristic color to the bead: iron, manganese, nickel, cobalt, chromium, uranium, copper (didymium, cerium, Vanadium, titanium, and tungsten). As the coloration varies with the temperature and with the amount of substance used, the results to be expected, with the neces- sary conditions, are summarized in the following table. The follow- ing abbreviations will be used: h-hot; c = cold; h-c =hot and cold; s.S. =slightly Saturated; Sat. =Saturated. WITH BORAx. WITH SALT of PHosphorus. Color of the Bead. || In the Oxidizing | In the Reducing In the Oxidizing In the Reducing 8, Iſle. Flame, Plame. Flame. SiO, (without | SiO, (without SiO2 (usually SiO2 (usually skeleton), alka- skeleton), alka-with skeleton), with skeleton), line earths, Hg, line earths and alkaline earths alkaline earths Colorless || Pb, Bi, Sb, Cd, earths, Mn, Di, and earths (sat. and earths, Mn, Zn, Sn, Ti Ce, Cu (s.s.) =turbid) Di, Ce, Cu (s.s.) W, Mo, Fe * (S.S.—c) W, Ti Ag, Pb, Bi, Sb, Ag, Pd, Bi, Sb, Gray Čd, Zn, Ni "Cd, Zn, Ni Fe (S.S.—h), Ag Fe (s.s.—h), Ag Yellow § & !",*|Ti (h), w (h) º º º Fe (h y *ms y y y º + , Ti (h tº º, Vaiśāoš º Fººtiº (brown) (brown) green | C, @.Vä, Fºgº, lº” Sº º Cu (h) Cr, Vá (h) (cºsat.) Vd (c), Mo (c) - (h— Blue º º Co (h–c) º, Co (h–c), W (c) Mn(h—c), Di f - & s Violet || (h—c), and Ni Mn º Di Ti (c) (with cobalt) Cu (sat.), opaque; * º as in the tº-º very slightly Fe (h—sat. orax bead; Red Fe ſh–sat.), sºld “g d. ) Ti and W in the Ce (h) -: with a trace of presence of Sn, ruby red * iron =blood red transparent. * Some are reduced to metal, so that the bead appears gray in the re- ducing flame (see table). CuSO, is white when anhydrous, but becomes blue immediately on the addition of water. 368 course OF ANALYSIS. 4. A Small Sample is Heated by Itself upon Charcoal before the Blowpipe, and if deflagration takes place a nitrate, nitrite, chlo- rate, iodate, etc., may be present. 5. The Substance is Heated with Soda upon Charcoal before the Blowpipe.—As much of the Substance as can be taken up on the end of a knife-blade is mixed with twice as much sodium carbonate (as described on page 28), placed in the cavity of a piece of charcoal and heated in the reducing flame of the blowpipe. There is obtained f As malleable button: Au, Ag, Sn, Cu, which can be pressed flat in an agate mor- tar. (a) Metal without incrustation. As gray metallic particles: Pt, Fe, Ni, and Co. Pt may be pressed flat in an agate mortar; Fe, Ni, and Co are magnetic and lar attracted by a magnet (cf. page 25). ſ As a brittle metallic button: Sb (white in- crustation), Bi (yellow incrustation). The button may be reduced to a powder by grinding in an agate mortar. As a malleable button: Pb (yellow incrus- Utation). (b) Metal with incrustation....- White, yellow when hot: Zn. (c) Incrustation without metal. } Brown: Col. White: As (garlic odor). (d) White, infusible, strongly luminous mass. . . . . . . . . . . Ca, Sr, Mg, Al, and rare earths. (e) Sulphur compounds are reduced to sulphides. If the melt is placed on a bright silver coin and moistened with water, the silver is blackened (Hepar reaction). 6. The Substance is Tested to See Whether it Imparts Any Color to the Non-luminous Flame.—A Small sample is introduced by means of a platinum wire into the base of the flame (cf. page 21), and then into the fusing zone. It is afterwards moistened with di- lute hydrochloric acid and the experiment repeated. The follow- ing indications may be obtained: PRELIMINARY EXAMINATION. 369 Sodium gives a yellow mono-chromatic flame; a piece of sealing-wax or a crystal of potassium dichromate appears yellow when illuminated by this flame. Potassium (caesium and rubidium) gives a violet flame which is completely obliterated by the sodium flame. If the flame is ob- served through cobalt glass, the Sodium flame disappears and the potassium flame appears pink. Lithium gives a carmine-red flame (or a red line in the spectro- Scope). Strontium also gives a carmine-red flame (which the spectro- scope shows to consist of several lines in the Orange, and a bright line in the blue). Calcium gives a brick-red flame (in the spectroscope an orange and a green line are seen, both about an equal distance away from the sodium line). Barium gives a greenish-yellow flame. In the case of barium Sulphate the green flame is either only in- distinctly visible or not at all. In order to detect barium in this case, a small portion of the substance is heated in the upper reducing flame; after cooling it is moistened with hydrochloric acid (odor of hydrogen sulphide) and again heated, when the barium flame can be easily Secrl. Thallium gives an emerald-green flame. Another portion of the substance should then be tested for boric acid, by treating with concentrated Sulphuric acid and bringing near the flame. A green color indicates the presence of boric acid, but if copper is present this test cannot be relied on.* w Lead, Arsenic, Antimony color the flame light blue, and copper compounds color the flame either green or blue. Preliminary Examination for the Electro-negative Elements (Anions). 1. Dilute Sulphuric Acid (double normal).-About a gram of the substance is treated in a small test-tube with dilute sulphuric acid, and One notes whether a reaction takes place in the cold or not (evolution of a gas). The following gases can be recognized: HCN from cyanides (odor); * By heating the solid substance with potassium ethyl sulphate in a test- tube, boric acid is converted into B(OC.H.), which burns with a green flame. Copper chloride does not interfere with this test. W. Castellina, Chem. Centralbl., 1905, I, p. 1619. 370 COURSE OF ANALYSIS. H.S from soluble sulphides (odor, and blackening of lead ace. tate paper); NO, from nitrites (brown fumes); SO, without separation of sulphur from sulphites (odor of burn. ing Sulphur); SO, accompanied by separation of sulphur from thiosulphates. the deposited sulphur is yellow, particularly after warming; CO, from carbonates or cyanates (barium hydroxide solution is rendered turbid). By boiling with dilute sulphuric acid, soluble ferro- and ferri- cyanides are decomposed and evolve hydrocyanic acid; acetates, set free acetic acid; hypochlorites evolve chlorine (which also takes place in the cold); while the peroxides of the alkalies and alkaline earths are decomposed with evolution of oxygen. 2. Concentrated Sulphuric Acid.—If the substance does not react with dilute sulphuric acid, a new portion is treated with 3 or 4 c.c. of concentrated sulphuric acid and heated. If the substance reacted with dilute sulphuric acid, it will react violently with con- centrated sulphuric acid and the gas will come off so quickly that it will carry small particles of the sulphuric acid with it, which makes the gas appear to have a penetrating odor and may lead to a mis- taken conclusion, especially as it will also cause barium hydroxide solution to become turbid. In such a case, dilute sulphuric acid is added drop by drop to a new portion of the substance until no further action takes place when five c.c. of concentrated Sulphuric acid are added and the mixture warmed. Gases and vapors may be evolved, which are (a) Colorless. HCl from chlorides, fuming in the air, with a penetrating odor. The fumes do not cause a turbidity with water. SiF,” from fluorides, fuming in the air, with a penetrating odor, and causing a turbidity on coming in contact with water. SO, without separation of sulphur. If there was no evolution of sulphur dioxide on treatment of the substance with dilute sul- * SiF, is formed on account of the experiment being performed in glass. In platinum, HF would be evolved in the absence of silica, which does not render water turbid. PRELIMINARY EXAMINATION. 37 i phuric acid, the Sulphur dioxide which now escapes must come from the Sulphuric acid itself; a metal, Sulphur, a Sulphide, carbon, or non-volatile organic matter, such as tartaric acid, citric acid, sugar, starch, etc., must be present. If non-volatile organic matter is present, carbonization will take place on warming. SO, with separation of sulphur indicates the presence of a sulphocyanate, in case there was no action with dilute sulphuric acid. CO from oxalates and other organic substances, and cyanates. It is an odorless gas, which does not fume in the air and burns with a blue flame. (b) Colored. Cl, a yellow gas with a Suffocating odor, turns iodo-starch paper blue, and indicates the presence of both a chloride and an oxidiz- ing substance. ClO, a yellow gas, very similar to chlorine, but which explodes violently on being heated, indicates a chlorate. If the substance deflagrates on being heated on charcoal, only a Small portion of the substance should be used for this test; but if no explo- sion takes place on warming, more of the substance should be added. HBr from bromides has a penetrating odor, fumes in the air, and is always colored yellowish brown by the presence of small amounts of bromine. The sulphuric acid is at first colored brown in the case of a colorless bromide, but becomes colorless on being boiled. CrO.C., brown (similar to bromine), results from the presence of a chloride and chromic acid. I, violet. In the case of a colorless iodide, the Sulphuric acid is at first colored brown by Small amounts of iodide, or gray, Solid iodine is deposited if considerable iodide is present, which vola- tilizes on warming, forming violet vapors. If considerable iodide is used for this test, the sulphuric acid is reduced to SO, or even H.S (cf. page 262). Mn,O, violet, is formed from permanganic acid, and is decom- posed with scintillation, often exploding, on being warmed. NO, brown, with a penetrating odor, comes from nitrates. 372 COURSE OF ANALYSIS. After the preceding tests have been made, the next step is the Solution of the Substance. As solvents the following are used: 1. Water; 2. Hydrochloric acid; 3. Nitric acid; 4. Aqua regia. In the majority of cases the first three solvents suffice, aqua regia being only necessary in a comparatively few cases, as will be seen from the following table: SOLUBILITY TABLE. SUBSTANCES SOLUELE IN WATER. Of Group I (page 243) the following are soluble: 1. Chlorides.—All except AgCl, Cu,Cl, Hg,Cl, PtCl, AuCl, BiOCl, SbOCl, Mg,OCl, PbCl, and TICl are difficultly soluble. 2. Bromides.—The same as the chlorides. 3. Iodides.—All except AgI, Hg, I, HgI, Cu,I, PdL, TII; PbI, is very difficultly soluble. 4. Cyanides.—Only the cyanides of the alkalies, alkaline earths, and mercury. 5. Ferrocyanides.—Only those of the alkalies and alkaline earths. 6. Ferricyanides.—Same as the ferrocyanides. 7. Cobalticyanides.—Only those of the alkalies, alkaline earths, and the ferric, mercuric, and lead Salts. 8. Sulphocyanates.—Those of the alkalies, alkaline earths, iron, cupric copper, and mercuric mercury. 9. Hypochlorites.—All. Of Group II (page 244) the following are soluble: 10. Nitrites.—All. Silver nitrite is difficultly soluble. 11. Acetates.—Silver and mercurous acetates and certain basic acetates are difficultly Soluble. 12. Cyanates.—Those of the alkalies, alkaline earths, and most of the remaining ones. Silver and lead cyanates are insoluble. & * SOLUBILITY TABLE. 373 13. Sulphides.—Only those of the alkalies and alkaline earths. CaS is difficultly soluble. 14. Hypophosphites.—All. Of Group III (page 244) the following are soluble: 15. Sulphites.—Those of the alkalies, and the bisulphites of the alkaline earths. 16. Carbonates.—Those of the alkalies, and the bicarbonates of Ca, Sr, Ba, Mg, Fe, Mn. 17. Oxalates.—Those of the alkalies; the remainder are diffi- cultly soluble or insoluble. Most oxalates, however, with the ex- ception of Ba, Ca, and Sr oxalates, form soluble complex Salts with alkali oxalates. 18. Iodates.—Only those of the alkalies. 19. Borates.—Those of the alkalies. The remaining borates are all difficultly soluble in water, but Soluble in ammonium chloride as a rule. 20. Molybdates.—Only those of the alkalies. 21. Selenites.—Those of the alkalies are readily soluble, the re- maining ones are difficultly soluble. 22. Selenates.—All except the barium and lead salts. 23. Tellurites.—Only those of the alkalies. 24. Tellurates.—Only those of the alkalies. 25. Tartrates.—The normal tartrates of the alkalies, and lith- ium and Sodium bitartrates. The remaining tartrates are insolu- ble in water, but are usually soluble in an excess of alkali tartrate solution, forming complex Salts. 26. Citrates.—Only those of the alkalies are readily soluble in water. The insoluble citrates usually dissolve in an excess of alkali citrate solution. 27. Pyrophosphates.—Only those of the alkalies. 28. Metaphosphates.—Only those of the alkalies. Of Group IV (page 244) the following are soluble: 29. Phosphates.—Only those of the aſkalies. 30. Arsenites.—Only those of the alkalies. 31. Arseniates.—Only those of the alkalies. 32. Thiosulphates.—Almost all are Soluble, though the silver and barium salts are difficultly soluble. 33. Chromates.—Those of the alkalies, Ca, Sr, Mg, Zn, Mn, Fe, 374 COURSE OF ANALYSIS. and Cu are soluble; the others are difficultly soluble or insolu- ble. 34. Vanadates.—The Orthovanadates are unstable; the pyro-, meta-, and polyvanadates are soluble in water, as a rule. The lead and mercurous salts are insoluble. 35. Periodates.—All more or less soluble in water, except silver periodate, which is insoluble. Of Group V (page 244), the following are soluble: 36. Nitrates.—All except a few basic salts. 37. Chlorates.—All. 38. Perchlorates.—All. 39. Manganates and Permanganates.—All. Of Group VII (page 244), the following are soluble: 40. Sulphates.—All except the Ca, Ba, Sr, and Pb salts, and a few basic sulphates. 41. Fluorides.—Those of the alkalies, silver, and mercury; the remaining fluorides are difficultly soluble or insoluble in water. Of Group VII (page 244), the following are soluble: 42. Silicates.—Only those of the alkalies. 43. Tungstates.—Only those of the alkalies. Of the salts insoluble in water, all dissolve in acid (hydrochloric or nitric) except AgCl, AgBr, AgI, AgCN, AuCl, PtCl, BaSO, SrSO, PbSO, HgS, Prussian blue, CaF, SnS, (mosaic gold), SiO, many silicates, fused PbCrO4, and the strongly-ignited oxides: Al2O3, Cr2O3, TiO2, SnO2, Sb2O3.* TiO2, SnO2, and Sb2O3 can be dissolved by long continued boiling with concentrated hydrochloric acid. Of the salts insoluble in acids, the following dissolve in aqua regia: PtCl2, AuCl, HgS, Sb2O3, SnS2, and Prussian blue (after long treat- ment). The following compounds are not dissolved by aqua regia: AgCl, AgBr, AgI, AgCN, BaSO4, SrSO4, PbSO4, CaF2,f fused PbCrO4, Al2O3, Cr2O3, native TiO2 (rutile, anatas, brookite), native SnO2 (cassiterite, tinstone), SiO2, Si, many silicates, C, carborundum, and strongly-ignited iridium (rhodium, ruthenium, and osmium). In order to bring such substances in Solution it is necessary to subject them to a special treatment. The process to be chosen depends largely upon the nature of the insoluble substance, so that * The oxides of antimony are changed to Sb,0, after long ignition in the air. f Calcium fluoride will be dissolved by the long continued action of aqua regia. TREATMENT OF INSOLUBLE SUBSTANCES. 375. a few general tests are necessary before going farther. Very often the preliminary examination will have been sufficient, but it is always well to perform the following simple experiments: 1. A small portion of the residue is heated on the charcoal stick to see whether a metallic button can be produced. (a) No metallic button is produced. The absence of silver, lead, and tin is thereby assured. (b) A metallic button is formed. The button is flattened in an agate mortar, and its solubility in acids is tested. (a) The metal dissolves in nitric acid forming a clear Solution, showing the absence of tin. A little hydrochloric acid is added to the nitric acid solution; a curdy precipitate is formed if the metal is silver, consisting of silver chloride, insoluble in water, but soluble in ammonia. If the nitric acid solution becomes turbid on the addition of sulphuric acid, lead is present. (6) The metal does not dissolve in nitric acid forming a clear solution, but leaves a white, insoluble powder: metastannic acid. A new button is treated with concentrated hydrochloric acid, when it will completely dissolve if silver is absent. Mercuric chloride produces a white precipitate of mercurous chloride in the hydro- chloric acid solution: tin is present. 2. A second portion of the residue is heated in a small test-tube with concentrated sulphuric acid and tested to see whether the escaping gas renders a drop of water turbid. A turbidity shows the presence of an insoluble fluoride (CaF.). 3. Still another portion of the residue is heated (with the help of a platinum wire) in the upper reducing flame of the gas-burner, allowed to cool in the inner mantle, moistened with dilute hydro- chloric acid, and it is noticed whether the odor of hydrogen sulphide is given off. It is then tested to see whether it will now impart a characteristic coloration to the flame. The presence of a sulphate is betrayed by the odor of hydrogen Sulphide, and the flame test shows whether barium alone or a mixture of barium, calcium, and strontium is present. 4. Another portion of the residue is tested in the salt of phos- phorus bead; silicic acid or a silicate usually gives a skeleton bead (cf. page 362). As the skeleton bead is not always obtained even when silica is 376 COURSE OF ANALYSIS. present, a further test for silicic acid is often necessary. For this purpose a small portion of the Substance is mixed with a little cal- cium fluóride (itself free from silicate), treated with concentrated sulphuric acid and warmed gently on platinum. The evolution of a gas which causes a drop of water to become turbid shows the pres- ence of silicic acid. * 5. The salt of phosphorus bead is now heated in the reducing flame to test for the presence of titanium, which causes the bead to become violet. The violet color appears more quickly on the addi- tion of a little piece of tin-foil. If iron is present at the same time, as is always true in the case of rutile, the bead is colored brownish red in the reducing flame. 6. The presence of chromium is often detected by the green color of the residue. In the case of chromite (gray or black residue) a portion is fused with Sodium carbonate and potassium nitrate in the platinum spiral (cf. page 88), when a yellow melt is obtained if chromium is present, which yields a reddish-brown precipitate with silver nitrate (after being dissolved in water and acidified with acetic acid) of silver chromate. 7. If the residue is gray or black, it may also consist of carbon. A small portion is heated upon a piece of platinum-foil, when the mass will glow and leave a lighter-colored ash if carbon is present. In doubtful cases a little potassium chlorate is melted in a test-tube, and a little of the residue is added, when a distinct glowing or a little explosion will take place if carbon is present. It is necessary to avoid the addition of shreds of filter-paper in this test. 8. Silicon and silicides (carborundum, etc.) are seldom met with, and show the greatest stability toward the above mentioned reagents. By fusing with caustic alkali in a silver crucible, however, they are readily decomposed with evolution of hydrogen. After dissolving the melt in water and acidifying, gelatinous silicic separates out, particularly after evaporation. METHODS FOR GETTING SUBSTANCES INTO SOLUTION WHICH ARE INSOLUBLE IN ALL ACIDs. 1. Insoluble Halogen Compounds (the silver compounds alone come into consideration) can be brought into solution by melting the mass, and by adding a little dilute sulphuric acid, and a piece of zinc so that it comes in contact with both the acid and the insolu- TREATMENT OF INSOLUBLE SUBSTANCES. 377 ble substance. After a while the acid is poured off; it contains the halogen acid in the presence of zinc sulphate, and is kept for the subsequent acid test. The residue consists of metallic silver. It is washed with water, dissolved in dilute nitric acid, filtered from any sulphate, silicate, etc., and the solution tested for silver with hydrochloric acid. 2. Insoluble Sulphates of the Alkaline Earths are brought into solution by fusing in a platinum crucible with 4–5 times as much calcined sodium carbonate, or with a mixture of equal parts of Sodium and potassium carbonates. The finely-powdered substance is mixed in the crucible with the Sodium carbonate, the mixture is covered with a thin layer of more carbonate, the crucible is covered and heated at first over a small flame in order to drive off the moisture which the carbonate always contains; after which the temperature is raised until the mass fuses to a thin liquid, this temperature being maintained for about a quarter of an hour. The glowing crucible is then plunged into some cold distilled water, and the fused mass will usually contract so that it can be readily loosened from the sides of the crucible. The mass is warmed with a little water on the water bath until it disintegrates, and no more hard lumps can be felt with a glass rod, when it is filtered. The filtrate will contain the Sulphate as Sodium sulphate, and the residue will now consist of carbonates of the alkaline earths. It is washed a few times with strong Soda solution, then with a 5 per cent. Soda solution until no more Sulphuric acid can be de- tected in the filtrate. It is finally washed with hot water until the wash-water no longer reacts alkaline (cf. page 61), dissolved in nitric acid, and analyzed as described on page 62. 3. Lead Sulphate is boiled with a concentrated soda solution, filtered and washed first with Soda Solution, and then with water. Calcium sulphate is also decomposed completely by boiling with soda solution, as is strontium Sulphate (though less readily), but barium sulphate is only incompletely decomposed.* 4. Silicic Acid and Silicates should be fused with sodium car- bonate, exactly as described on page 360. 5. Metastannic Acid, as obtained by the oxidation of tin with nitric acid, is readily dissolved by boiling with a little concentrated * Lead sulphate dissolves readily in hot ammonium acetate solution (con- veniently prepared by adding acetic acid, sp. gr. 1.04, to ammonia, sp. gr. 0.96, until the solution no longer reacts alkaline). This solution gives lead sulphide on adding ammonia and ammonium sulphide, lead chromate on adding potassium chromate solution, and lead sulphate on adding a little sulphuric acid. It shows the presence of sulphuric acid by giving a precipi- tate with barium chloride. 378 COURSE OF ANALYSIS. hydrochloric acid, and then treating with considerable cold water (cf. page 218). Tin dioxide, as it occurs in nature (tinstone), as well as the strongly ignited metastannic acid, cannot be brought into solution in this way. One of the methods described on page 222 (usually the Sodium carbonate and sulphur method) must be used. 6. Insoluble Fluorides are first heated with concentrated sul- phuric acid, and the Sulphate formed is brought into solution by the method described above under 2. 7. Titanium Dioxide is fused with potassium pyrosulphate in a platinum crucible (cf. pages 74 and 109), or with sodium carbonate, the melt is treated with cold water, and the residue dissolves in hydrochloric acid (cf. page 113). 8. Chromium Trioxide and Chromite are fused with sodium carbonate and a little potassium nitrate (cf. page 88). 9. The Insoluble Complex Cyanides are completely decom- posed by boiling with caustic soda in a porcelain dish. After boiling with the alkali, the product is diluted with water and filtered. The filtrate will contain the acid in the form of its sodium salt; and, in Some cases, may also contain aluminium and zinc. A part of the filtrate, therefore, is acidified with acetic acid, saturated with ammonia, and filtered in case a precipitate of Al(OH), is formed. This filtrate is then treated with ammonium sulphide in order to precipitate any zinc, which may be present, as sulphide. The soluble complex cyanides are decomposed before the analy- sis by heating them with concentrated Sulphuric acid (cf. page 103). SOLUTION OF THE SUBSTANCE. When a substance is dissolved, whether in water or in acids, phenomena are often observed which may be of great importance as concerns the Subsequent analysis. Then the color, reaction of the solution towards indicators, or the evolution of gases will lead to important conclusions. The substance is first tested with regard to its solubility in water, by taking about 3 gram of its fine powder, adding a little cold water, and noticing whether any bubbles of gas are given off. A gas is evolved when there is present: (a) Peroxides of the Alkalies or Alkaline Earths, which are partly decomposed into hydroxide and oxygen: TREATMENT OF INSOLUBLE SUBSTANCES. 379 } Na,O,--H,0=2NaOH--O; BaO,4-H,0=Ba(OH),4-O. The escaping gas is tested for oxygen by means of a glowing splinter. $º. In the alkaline solution (red litmus is changed to blue) Some un- decomposed peroxide will still be found. The solution is diluted considerably with water, cooled, and carefully acidified with sulphuric acid, a few cubic centimeters of ether and a little potassium dichromate solution added, and the mixture shaken. If a peroxide is present, the upper ether layer will now be colored blue. A still better method for detecting the hydro- gen peroxide, formed by the action of the Sulphuric acid upon the peroxide, consists in adding a few drops of titanium Sulphate Solu- tion; a distinct yellow color will be noticed if only traces of hydro- gen peroxide are present (cf. page 111). (b) Carbides of the Alkaline Earths (calcium carbide). These are decomposed into acetylene (which has a peculiar odor, and burns with a luminous flame) and calcium hydroxide: CaC,--2H,0=Ca(OH), H-C.H. (c) Nitrides of the Alkaline Earths (magnesium nitride). Magnesium nitride is decomposed by water into magnesium hydroxide and ammonia: Mg,N,+6HOH=3Mg(OH),4-2NH, If considerable water is added, there is no gas evolution, because the ammonia will be absorbed by the water; but on boiling the solu- tion, ammonia will be given off, which can be readily recognized by its odor. (d) Phosphides of the Alkalies and Alkaline Earths.-These are decomposed by water, Setting free spontaneously-combustible phosphine: Ca,P,--4H.O=P.H,--2Ca(OH), Very small quantities of the phosphide can be recognized by the characteristic garlic Ordor. (e) Many Chlorides, Bromides, and Iodides of the Negative 38o COURSE OF ANALYSIS. Elements, PC, PCls, etc., are decomposed into the halogen hydride and the oxygen acid of the negative element: PCls+4H,O=5HCl·HH,PO,. (f) A few Sulphides which are seldom met with (MgS, Al,S,- etc.)—These are decomposed by water with loss of hydrogen sul- phide, which can be detected by its odor, and by its blackening lead acetate paper: MgS+2HOH=Mg(OH),4-H.S. After any reaction caused by the first addition of water is over, about 10–15 c.c. more are added, then heated to boiling and allowed to cool. If the substance dissolves completely, forming a clear solution, it is evident that it is unnecessary to test for any insoluble bodies in the subsequent analysis. 4. If a residue remains, it is possible that a part of the substance has dissolved in the water. To determine whether this is the case, the liquid is decanted through a filter, and a few drops are carefully evaporated to dryness on platinum-foil (or a watch-glass). If the foil is heated too hot, volatile compounds may escape unnoticed. If a residue remains after evaporation, it is evident that a part of the original substance is soluble in water. The original residue is then treated several times with small amounts of water, and the aqueous extraction thus obtained is analyzed by itself. The part remaining undissolved is now treated with acid; hydrochloric acid being usually used (unless the preliminary examination has shown the presence of either lead or silver, when nitric acid should be used). The solution in acid is best effected by treating the residue with #–1 c.c. of concentrated acid (it being noticed whether there is any evolution of a gas), warming, and then diluting with water, to dis- solve any chlorides insoluble in hydrochloric acid. It must be remembered, however, that bismuth and antimony Salts form in- soluble basic chlorides on dilution with water, so that too much water should not be added. If a residue remains after treatment with acid, it is brought into solution by one of the methods described on pages 376–378. II. ExAMINATION FOR THE METALS (CATIONs). The different solutions are analyzed separately according to the table on pages 382, 383. EXAMINATION AND SEPARATION OF THE GROUPS. 381 NOTES ON THE GENERAL TABLE FOR THE EXAMI- NATION AND SEPARATION OF THE GROUPS. 1. Introduction of Hydrogen Sulphide.—The acid solution (usu- ally about 50–100 c.c. in volume) is placed in an Erlenmeyer flask, of about 300 c.c. capacity, which is provided with a double-bored stopper. Through one of the holes a right-angled piece of glass tubing is introduced so that it just reaches the lower surface of the stopper, while through the other hole another right-angled glass tube is fixed so that it almost reaches the bottom of the flask. The longer tube is raised till it does not touch the liquid, and the solution is heated to boiling, a strong stream of water vapor coming out through both of the tubes. The longer tube is then connected with a Kipp H.S.-generator by means of a rubber tube which is already filled with hydrogen Sulphide gas. The apparatus is ar- ranged so that the gas from the generator passes first through a flask filled with water, and then through a tube filled with cotton (to remove any liquid which may be mechanically carried over from the wash-bottle). A steady stream of hydrogen sulphide is con- ducted through the apparatus, the flame is removed from under the flask, and the shorter tube is closed by means of a piece of rubber tubing which contains a solid glass rod. The longer tube is intro- duced into the liquid, the flask is shaken well for a-few minutes and allowed to cool somewhat. The short tube is then opened, the stop- cock of the generator is closed, and an equal volume of cold distilled water is added.* The Solution is again Saturated with hydrogen sulphide, the short tube is once more closed and the contents of the flask well shaken for two or three minutes, after which time the pre- cipitation is complete and the precipitate can be immediately fil- tered off.i. If the preliminary examination has shown that oxidizing sub- stances are present (cf. page 7) or arsenic acid, it is best to reduce these with sulphurous acid before treating with hydrogen sulphide. For this purpose, the weakly acid solution is concentrated as much as possible in an Erlenmeyer flask, 100 c.c. of a saturated sulphurousf * Cf. page 179: Separation of Hg, Pb, etc., from previous groups. † Cf. Gräbe, B. B. 31, page 2981 (1898). i Ferric salts are not completely reduced by SO, in strongly acid solutions. When they are present, the solution is neutralized with ammonia until a ermanent precipitate is formed and then the SO, water is added. It is still 5. to pass SO, gas into the Solution. TABLE VII. GENERAL TABLE FOR THE EXAMINATION AND SEPARATION OF THE GROUPS. The solution is treated with hydrochloric acid until no more precipitation takes place, warmed gently, and filtered. PRECIPITATE. FILTRATE (can contain all of the remaining metals). R RRECIPITATE. Hydrogen sulphide is conducted into the solution (which should contain 10–15 c.c. of double normal acid to every 100 c.c. of solution) until it is saturated: f it is diluted with an equal volume of water, more hydrogen sulphide led into it, and then filtered. FILTRATE (can contain the metals of Groups III and TV in the presence of the alkalies). HgS, black As,Ss, yellow Y PbS, black As,Ss, yellow Bi,Ss, brown Sb2S3, orange CuS, black Sb2S3, orange CdS, yellow SnS, brown or orange SnS, yellow insoluble in # y ;: IDS In , Pts, blac (NH4)2S, and Moš, brown a2S, SeS, orange TeS, black | } ; 2 º This precipitate is examined ac- cording to Table IX. The filtrate is freed from H2S by boiling, 2 c.c. of concentrated HNOs are added, and the solution evaporated to a small volume. A small portion is now taken, an excess of ammonium molybdate added, and the mixture warmed gently; a yellow precipitate shows the presence of HaPO,.f. If the preliminary examination has shown the presence of non-volatile organic matter, or of oxalic acid, it must be destroyed before the test for HAPO, is made. For this pur- pose the filtrate is evaporated to dryness and gently ignited. After cooling, a little concentrated HCl is added, the solution is boiled, a little water added, and any carbon or silica which may be left is fil- tered off.] A portion of this filtrate is then tested with ammo- nium molybdate for HAPQ,. . According to whether phosphoric acid is present or not, the method A or B is used. ‘. : | A Phosphoric Acid is Absent. The filtrate is diluted to about 100 c.c. with water, NH3 is added until the solution reacts alkaline, and then colorless (NH4)2S until no further precipitation takes place. PRECIPITATE. FILTRATE. Can contain FeS, black NiS, black CoS, black MnS, flesh color ZnS. white Al(OH), white Cr(OH)3, green UO,S, brown Ti(OH), white The precipi- tate is analyzed according to Table X. Can contain Ca, Sr, Ba, Mg, and Alkalies. The filtrate is concen- trated, HCl added, the solution boiled and the deposited S is filtered off. NHs and (NHA),COs are added, the liquid is heated to boiling" and filtered. PRECIPITATE, FILTRATE. Can con- Can con- tain tain BaQ0s ) 8 || Mg and the SrCOs 3 || Alkalies. CaCO, ) > It is ana- It is ex- ||lyzed a c- amined ac- || cording to cording to || Tabl II. Table XI. 8,OHe H Phosphoric Acid is Present. The filtrate is evaporated to dryness, 10 c.c. of concen- trated HNO3 are added, the solution is again evaporated to dryness, and this process is repeated once more. 10 cc. of concentrated HNO3 and little by little about 1 gram of tinfoil are added, the liquid boiled to a small volume, poured into 100 c.c. of water contained in a narrow cylinder, and al- lowed to stand quietly over night; . In the morning the liquid is siphoned off and is now free, from H3PO4. It is saturated with #s in order to , precipitate Pb, Cu, etc. (which are almost always present with the tin as im- purities), and filtered. PRECIPITATE. FILTRATE. Contains || Is exam- the impuri- || ined ac- ties which || cording were intro- || to A. duced with the tin, and should be discarded. * Besides these substances, HCl can also precipitate from an alkaline solution SiOs H2, WO4H2, SbOCl, Sb2Ss, Sb2S3, As2Ss, SnS2, and others. In such case, the precipitate is examined according to Table IX. f Cf. note 1, page 381. f Cf. note 2, page 384. § Cf. note 3, page 384 | Cf. Note 4, page 385. Cf. page 55, § 2. 384 COURSE OF ANALYSIS, acid solution are added, a stopper is introduced into a flask provided with inlet and outlet tubes, and the solution is boiled for a quarter of an hour; 10 c.c. of concentrated hydrochloric acid are then quickly added, and the boiling is continued while carbon dioxide is being passed into the Solution, until the excess of sulphurous acid is completely removed.* The solution is then treated with hydro- gen sulphide as above described. In this way a precipitate con- taining no free sulphur is formed by the hydrogen sulphide. 2. The test for phosphoric acid at this place must never be omitted. As is evident from the table, it is desired to separate the members of Group III from those of Group IV by precipitation with ammonia and ammonium sulphide. If phosphoric acid is pros- ent in the Solution (or oxalic acid) the addition of ammonia can cause the precipitation of calcium, strontium, barium, and magne- sium phosphates (or of calcium, strontium, and barium oxalates) with the members of Group III. It is, therefore, necessary to remove the phosphoric acid before the precipitation with ammonia and ammonium sulphide (cf. page 330) by evaporation with tin and nitric acid. In the absence of chromium, the phosphoric acid may be removed by precipitation as basic ferric phosphate (cf. page 327). 3. Oxalic acid is usually detected in the preliminary examination by heating the Substance with concentrated Sulphuric acid, whereby inflammable carbon monoxide is produced. As, however, formates and other organic substances also give off carbon monoxide on being treated with concentrated Sulphuric acid, this test simply indicates that oxalic acid may be present, and is not conclusive. On the other hand, the carbon monoxide is often not detected in the pre- liminary examination; e.g., if considerable chloride (or other salt) is present causing a gas to be evolved on treatment with concentrated sulphuric acid which does not support combustion. In such a case the carbon monoxide would not burn. If carbon monoxide was not detected in the preliminary examination, and there was no evidence of carbonization, the substance may be tested as follows for oxalic acid: A small portion of the solution is treated with Sodium carbon- * If the SO, were not removed the following reaction would take place: SO,--2H, S =2H,O +3S. | If Pb, Ba, or Sr salts are present, they will be precipitated as sulphates. In this case the solution is filterod before treating with H.S, and the precipi- tate examined by itself for Ba, Sr, and Pb. EXAMINATION AND SEPARATION OF THE GROUPS. 385 ate in a porcelain dish until the solution reacts strongly alkaline, then boiled and filtered. The filtrate is acidified with acetic acid, boiled to drive off carbon dioxide, calcium chloride and ammonia are added, and the Solution again acidified with acetic acid; a white crystalline precipitate insoluble in acetic acid shows the presence of Oxalic acid. 4. If carbonization and evolution of vapors with a burnt odor are caused by heating the original Substance, the presence of non- volatile organic matter is assured; e.g., tartaric and citric acids, sugar, starch, etc. Such substances prevent the precipitation of alu- minium, iron, etc., by means of ammonia and ammonium sulphide (cf. page 72), and consequently must be destroyed by ignition. After ignition, the carbonates or oxides formed are extracted with hy- drochloric acid and filtered off. A black carbonaceous residue almost always remains at this point, which is washed well with water, dried, and the carbon completely burned away, when a residue of Al,O, Cr,O, Fe,0, insoluble in hydrochloric acid, often remains with some silica. This is examined by itself. In order to get it into solution, it should be fused with potassium pyrosulphate (cf. page 74). The metals are then left in the form of soluble sulphates, while the silica remains undissolved, and can be tested according to page 375, § 4. * TABLE VIII. EXAMINATION OF GROUP I. The precipitate produced by HCl can contain AgCl, Hg,Cl, PbCl, (TICI).” It is washed with cold water, then boiled with a little water and filtered hot. RESIDUE. Solution. Can contain AgCl, Hg,Cl2, and a little PbCl, If considerable lead is It is washed with boiling water until the PbCl, | present, needles of PbCl, is all removed, and the residue treated with am- || separate out on cooling. monia on the filter.f * If only a little lead is RESIDUE. Solution. present, there will be no deposition of PbCl. In all c as e s the HgNH,Cl + Hg Contains Ag(NH3)2Cl. solution is treated with (black) shows Hg to || It is acidified º, K2Cr,0, solution; a yel- to be present. HNO; ; a white curdy low precipitate of PbCrO, precipitate of AgCl shows Pb to be present. shows Ag to be pres- ent. * For the detection of thallium consult the supplement. t Cf. foot note to page 245. The freshly-precipitated sulphides are carefully washed with water containing H.S, and the precipitate dried by suction. It is then washed into a porcelain dish, covered with yellow (N H.),S,” warmed a few minutes with constant stirring (long boiling must be avoided) (cf. page 204), and filtered. RESIDUE can contain Hgs, PbS, Bi,Say CuS, CdS. It is washed first with water containing ammonium sulphide, then with H.S water, and the liquid sucked off as much as possible; it is then warmed with nitric acid (sp. gr. 1.2) in a porcelain dish; all of the sulphides dissolve with the exception of HgS. filtered and washed. SoLUTION [AsS,(NHA)a, SbS,(NH,)s, º SnSs The solution is largely diluted with water and HCl added to acid reaction. boiled, the sulphides allowed to settle, the upper liquid is poured off, and the residue 4/2_* It is then For the separation of arsenic, antimony, and tin sulphides, one of the two following TABLE DK. EXAMINATION OF GROUP II. This residue is filtered and washed with water. methods can be used: 1. The residue, which always contains considerable free sulphur, is boiled with concentrated HCl (1:1) until no more H2S is given off, when it is filtered. RESIDUE (HgS). This is dis-l solved in aqua regia, evaporated almost to dry n e s s, water is added, and the depos- ited sul — p hur fil- tered off. SoLUTION [Pb(NO3)2, Cu(NO3)2, Bi(NO2)3, CdGNO3)2]. It is evaporated to a small volume, a few cubic centimeters of dilute H2SO,t are added, and the solution evaporated until fumes of sulphuric acid are given off. After cooling, a little water f is added and the solution filtered. RESIDUE (PbSO,). The for- mation of w h i t e PbSO, in - dicates Pb to be pres– ent. § SoLUTION [CuSO, Bi,(SO)s, CdSO,. The solution is saturated with NH3 and filtered. RESIDUE (BiSO,OH). The white pre- cipitate should not contain any other metal t h a n bis- SoLUTION [Cu(NH3),SO, Ca(NH3)6SO, j. A blue color shows the presence of Cu, KCN is added until the solution is decolorized, and H.S is con- RESIDUE (As2S3 +S). It is treated, together with the filter, with fum- ing HNO3 in a covered beaker, until it has com- |; dissolved and no | more brown, vapors are given off. The solution is | evaporated to a small yol- ume, an excess of NH3 is added, then some “mag- nesia mixture,” and the so- lution is vigorously stirred. If considerable arsenic is present, a white crystal- line precipitate of MgRNH4AsO4 + aq is immediately formed. × SOLUTION (SbCls, SnCl,) The solution is concen— trated to a small volume, a few drops are placed on plat- inum foil and a piece of zinc added, so that it comes in contact with both the plati- num and the solution. After a few seconds the zinc is re- moved and the platinum is examined to see whether a black stain has been formed upon it, showing the pres- ence of antimony. The zinc is again allowed to act upon the solution, un- til the evolution of hydrogen % § Y ceases almost entirely, when the solution is washed away with distilled water, being careful to keep the zinc and platinum in contact with one another. The zine is then removed, scraping off with the foil any tin which may adhere to it, which is dissolved in concentrated HCl, the solution placed in a small test-tube and a few drops of HgCl2 added. - A white precipitate finally be- coming gray shows Sm to be present. 2. The mixture of the three sulphides and sulphur is warmed with a solution of ammonium carbonate and The residue is dis- solved in concentrated HCl and treated as under sub 1 ×. * Yellow ammonium sulphide is used, because SnS is insoluble in colorless ammonium sulphide (cf. # If considerable Pb is present, a white precipitate of PbSO4 will appear on the addition of the nitric acid solution. If only a little Pb is present, the precipitate expelled by evaporation and the solution has been diluted with water (cf. page 162). # Too much water should not be added, as otherwise Bi may be precipitated as basic sulphate (cf. page 165). § In the presence of Au, is then made with a portion of the residue, but the larger part is tested for Au, Pt, Sn, and Pb (formed by the oxidation of the sulphide by HNO3). Pt and considerable Sn their sulphides will be mixed with the HgS. age 213). 2SO4 to the does not appear until the HNO3 has been The mercury test 4- The dried precipitate is fused with soda and potassium cyanide in porcelain, the melt extracted with water, and the reduced metals are first treated with nitric acid. which is tested for with H2SO4. This dissolves the lead, Concentrated HCl is next added to dissolve any Sn, for which the solution is tested. If a gray residue remains, it is dissolved in aqua regia, and tested for Au, and Pt by adding ammonium chloride evaporating to dryness on the water-bath, taking up the platinum precipitate in a very little water, filtering, and testing for Au with ferrous sulphate. g = * | A white precipitate of Al(OH)3 or a brown precipitate of Fe(OH)3 is often formed here if the H2S precipi- The fil- tra t e is t r e a t e d with stan- nous chlo- ride solu- t i o n; a white pre- cipitate of Hg2Cl3, which be- comes gray on further addition of SnCl2, shows Hg to be pres- ent. ed by the char coal- stick reac- tion; t he formation of a malle- able metal- lic button shows Pb to be pres- ent. muth, if the work to this point h a s been c are ful ly done. Ś Its pres– ence is confirmed by washing the pre- cipitate, dissolving it in a little HCl, and treating this solution with an al- kaline solution of potassium stannite. A black precipi- tate (metallic B i) shows the presence of B. ducted into the solution; a yellow precipitate shows Cal to be present. Its presence is confirmed by the flame reaction. A little of the filtered and washed precipitate is care- fully roasted on an asbestos thread in the upper oxidiz- ing flame, and the oxide coating is produced on a porcelain dish (cf. page 178). If the brown deposit be- comes bluish black on be— ing moistened with AgNOs solution, Cd is present. The precipitate is only ormed after standing some time in case only little arsenic is present. If no precipitate is formed after standing 12 hours, arsenic is absent. filtered. RESIDUE (Sb2S3, SnS2, S). tate was incompletely washed. SOLUTION (AsS,(NH.), +AsSO3(NH4)3]. The solution is acidified with HCl, whereby yellow arsenic sulphide is precipi- tated, showing the presence of arsenic. To confirm this, the precipitate is treated with fuming HNOs and the arsenic acid precipitated as MgRNH4AsO4 as described under 1. An ammoniacal solution of H2O2 is better § fuming #Nóa (cf. page 7). The magnesium ammoni- um arsenate can be tested in the dry way by produc- ing the oxide coating on por- celain (cf. §. 199) and treating it with AgNO3 + A. 3. A yellow stain shows 8, The precipitate produced by ammonia and ammonium sulphide in the presence of ammonium chloride is washed four times with water containing (NH)2S, three times with pure water, and then treated in a porcelain dish under constant TABLE X (A). EXAMINATION OF GROUP III. stirring with cold, double-normal hydrochloric acid, until no more hydrogen sulphide is given off, when it is filtered. SOLUTION [FeCl, AlCls, CrCl3, UO,Cl, MnOl., ZnCl2 (traces of NiCl,)]. RESIDUE [CoS, Ni S]. The well-washed residue is tested in the borax bead. A blue color shows Co to be present. To test for nickel, a small por- tion is dissolved in a little aqua regia, the solution evaporated just to dryness, dissolved in a little water, and KCN solution added drop by drop until the precipitate first formed redissolves complete- ly; 2 or 3 c.c. of caustic soda solu- tion are then added, and chlorine conducted into the cold solution (cf. page 130). If nickel is pres- ent, a black precipitale of Ni(OH), is formed within a few minutes. N.B.-If the bead was brown in the cobalt test, the presence of nickel is thereby assured, and fur- ther testing is unnecessary. Small amounts of Co may be present. In order to detect this, a consid- erable amount of the precipitate is dissolved in aqua regia, evapo- rated just to dryness, dissolved in a little water, a concentrated so- It,is evaporated to a small volume, oxidized with 1–2 c.c. of conc. HNO3, NaOH added until strongly alkaline, boiled, diluted with an equal volume of water, and filtered. RESIDUE [Fe(OH)2, Cr(OH), U.O.Na, Mn(OH)2, and traces of Ni(OH2)]. SoLUTION [Al(ONa)a, Zn(ONa),). It is washed with hot water, dissolved in as little con- centrated HCl as possible, diluted a little with hot water, boiled a few minutes,” treated with 10 c.c. NH,Cl solution, precipitated hot with NH, until the latter is present in ex- cess, and filtered as quickly as possible. RESIDUE [Fe(OH)2, Cr(OH), U.O.Na,]. SOLUTION MnCl2(NHA), and traces of NiCl2. This is dissolved in as little concen- trated HCl as possible, considerable (NHA),COs solution is added, it is heated just to the boiling-point, and filtered. RESIDUE SOLUTION Fe(OH)3, Cr(OH)s. [UO,(CO2), (NH,),. Test for Iron: A small por- || The so- tion is dissolved in a few drops || 1 u t i on of HCl, it is diluted with || is acidi- water, and a few drops of |fied with In case Ni is present, a few drops of KCN solution are added, the n (NH4)2S, a n d the solution boiled; a flesh- c o 1 or e d pre- ci p it at e of j MnS, show s It is acidified with HCl, NH3 is added in excess, and it is filtered: RESIDUE SOLUTION [Al (OH)3.][Zn(NH3)3Cl2]. A white gelatin ous precipitate shows the presence of Al It is con- firmed by the Thé – nard – blue rea c ti on (cf. p a ge 75). The solu- tion is acid- ified with acetic acid a n d H.S conducted into it. A white pre- cipitate of ZnS shows t he pres– ence of Zn. It is con- firmed by º: § the Rinn- m a n n 's green re- action (cf. Ta b le II (page 142). If the boiling were omitted, the greater part, and in some lution of potassium nitrite added, the Solution is acidified with ace- tic acid and allowed to stand 12 hours. A yellow crystalline pre- cipitate of Co(NO2) sks shows Co to be present. A still better test for cobalt is the Vogel reaction (cf. page 137). The solution is neutralized with Na2CO3, until a permanent precip- itate is formed, considerable am- monium sulphocyanate is added, and also a little amylalcohol with an equal volume of ether The mixture is well shaken. The al- cohol-ether layer becomes colored blue if Co is present. K, Fe(CN), solution are added; a dark-blue precipitate of Prussian blue shows that Fe f is present. Test for Chromium: A second portion of the residue is fused with sodium carbo- nate and potassium nitrate in a platinum spiral. A f t e r cooling, the melt is flattened on a porcelain plate by means of a pestle, dissolved in a little H2O, acidified with ace- tic acid and a few drops of AgNO3 added; a red precipi- tate of silver chromate shows that Cr is present. HCl and K, Fe(CN), solution is added ; a brown pre- cipitate or coloration shows the presence of the presence of Mn. This is c on - firmed by fus- ing a little of the precipitate with so di u m carbonate and potassium n i- trate on plati- num; a green melt, which yields a green solution, be - coming red on addition of ace- tic acid, shows the presence of Mn. * The boiling serves to convert any manganic chloride into manganous chloride. cases, all, of the manganese would be precipitated by ammonia (cf. behavior of manganous salts toward, KOH and NH3 pages 117 and 118). f In order to determine whether the Iron was originally present in the ferrous or ferric condition the aqueous, or of the original substance is tested with KalPe(CN) in which form the iron originally was present. forming FeCl3 and CrCl3 hº K4Te(CN)6 (cf. pages 95 and 100). - - - Thus if we have a mixture of ferrous sulphate and considerable - water the iron is oxidized and precipitated as basic sulphate, together with a part of the Cr as Cr(OH)s; both substances are soluble in HCl, K2Cr2O7 + 6FeSO4 + 4H2O -: 3Fe2O(SO4)2 + 2Cr(OH)s + 2KOH. - ydrochloric acid solution It is, however, sometimes, impossible to determine K2Cr2O7, on treatment with Exam in ed according to Table X (A). . The H2S is completely removed by boiling, an excess of bromine water is added, the excess is removed by boiling, the solu- tion is poured into a small Erlenmeyer flask, almost neutralized with Na2CO3 and an excess of BaCOs suspended in water is added. The contents of the flask are well shaken, the flask is corked up, and the precipitate allowed to settle. After standing about an hour the solution is decanted through a filter. SolūTION [MnOlz, ZnCl2, BaCl2 and traces of NiCl2]. PRECIPITATE [Fe(OH)3. Cr(OH)3 Al(OH)3. [UO2CO3].3Baz, BaCO3]. - It is washed with water until the filtrate gives no precipitate with (NH4)2S, dissolved in a little, dilute HCl the Ba precipitated hot with dilute H2SO4 and filtered. The filtrate is evaporated to a small volume, treated with an excess of ammonium carbonate, heated just to the boiling-point, a little NHS added, and filtered. RESIDUE [Fe(OH)3 Cr(OH)3, Al(OH)3]. SoLUTION [UO2(COs)3] (NH4)4.” Fe(OH)3 and Mºjöğ). Discarded. | The yellow filtrate is carefully acidified with dilute H2SO4, NH3 is added till it reacts alkaline, and it is then filtered. RESIDUE Al(OH)3. SoLUTION (Na2CrO4). The white pre- cipitate shows Al to be present. It is confirmed by the Thénard’s blue reaction. The yellow solution is divided into two parts. The first part is acidified with acetic acid and treated with AgNO3 solution; a red- dish-brown precipitate of Ag:CrO4 shows Cr to be present. The other part is acidified with dilute HgSO4, an excess of hydrogen peroxide is added, and the solution is shaken with ether. The latter is colored blue if Cr iS present. * TABLE X (B). EXAMINATION OF GROUP III. The precipitate produced by NHs, and (NHA)2S (in the presence of NH,Cl) is washed four times with water containing (NH4)2S and three times with pure water; it is then treated with cold, double-normal HCl in a porcelain dish (stirring constantly) until no more #3 is given off, when it is filtered. RESIDUE [CoS, NiS). Solution [FeCl2, AlCls CrCl3, UO2Cl2, MnOlz, ZnCl2 with traces of NiCl2]. The solution is heated to boil- ing and the barium completely precipitated by the addition of di- lute H2SO4. The BaSO4 is filtered off. The solution is treated with NaOH until it reacts strongly al- kaline; it is warmed gently and filtered. It is dissolved in a little dilute H2 SO4, and a few drops of the solution formed is tested with KAFe(CN)6, for Fe; a precipitate or coloration of Prussian RESIDUE [Mn(OH)2, SoLUTION blue shows Fe to be present. The remainder of the solution is treated with a traces of Ni(OH), Zn(oNa). concentrated caustic soda solution until it reacts distinctly alkaline; a few crystals 3. of KMnO4 are added, and the solution boiled. The chromium is by this, means oxidized to Na2CrO4 with deposition of MnO(OH)2. Alcohol is added drop by The precipitate is dissolved in HCl, boiled, neutralized with NH3, and if Ni is present, a few drops of KCN so- lution are added, SODOle (NH4)2S, and bo i le d ; a pink, flesh-colored pre- cipitate shows the presence of Mn. It is confirmed by f us in g with soda and nitre (cf. Table X (A)). The solution is acidified with acetic acid, and conducted in- to it; a white precipitate of ZnS shows Zºn. to be present. It is con- firmed by the Ri n n mann's green reac- |tion (cf. Table X (A)). drop to the hot violet solution until it becomes colorless, when it is diluted with a little water and filtered. RESIDUE. Solution [Na2CrO4, Al(ONa)3]. * The solution always contains some chromium, but it does not interfere with the test for uranium. § EXAMINATION AND SEPARATION OF THE GROUPS. 391 TABLE XI. The precipitate produced by ammonia and ammonium carbonate is well washed with hot water, dissolved in as little dilute HNO, as possible, and carefully evaporated in a porcelain dish to dryness. (The residue must on no account be ignited.) A small portion of the dry residue is dissolved in a little water, filtered from threads of filter-paper, and treated with calcium sulphate solution. (a) No precipitate is formed; Ba and Sr are absent; Ca is present. (b) A precipitate is formed slowly; Ba is absent; Sr, and perhaps Ca, is present. (c) A ºpiate is formed immediately; Ba is present, and perhaps Ca a DCI YE”. If the calcium sulphate solution produces no precipitation, only Ca is present, and a further test is unnecessary. If, however, it produces a pre- cipitate, the original residue must be analyzed according to one of the two following methods: 4. The completely dry nitrates are treated with a few drops of absolute alcohol, stirred, and the alcoholic solution filtered through a filter wet with absolute alcohol into a porcelain crucible. RESIDUE [Ba(NO3)2, Sr(NO3)2, SOLUTION. [Ca(NO3), with perhaps some Ca(NO3)2]. and traces of Sr(NO3)2]. Two to three c.c. of absolute alcohol are poured on the residue, decanted through the filter, and the process repeated three to four times in order to remove all of the Ca(NO3)2. The É Se, which is now free from calcium, is evaporated to dryness, trèated with 3 c.c. of concentrated HCl, evaporated again to dryness, and the process repeated. The residue of dry chlorides thus obtained is treated with a few drops of absolute alcohol, stirred, and decanted through a filter wet with absolute alcohol. RESIDUE [BaCl2, with perhaps some SrCl2]. SoLUTION [SrCl2].' The alcoholic solu- tion is evaporated to dryness and the resi- due tested in the flame. carmine-red flame The remaining SrCl2 is removed by de- canting several times with alcohol and the residue is tested in the flame; a yel- lowish-green flame shows Ba. The rest of the residue is dissolved in || A a little water and treated with K2Cr2O7 || shows Sr. solution and sodium acetate; a yellow This is confirmed precipitate of BaCrO4 shows Ba to be || with the spectroscope present. (cf. chart). The solution is evaporated to dry- ness, the residue ignited. moistened with HCl and test- ed in the flame; a brick-red flame shows Ca. This is con- firmed with the Spectroscope (cf. chart). The calcium may also be detected by adding one or two drops of sulphuric acid, the sulphate being insoluble in alcohol. B. The dry nitrates are treated with 3–1 c.c. of concentrated HNO, stirred well with a glass rod, and filtered through a small asbestos filter by means of suction. RESIDUE [Ba(NO3)2, Sr(NO3)2]. SOLUTION Ca(NO3)2. All of the Ca(NO3)2 is removed by washing with concentrated HNO3 the residue is dissolved by pouring a little water through the filter. evaporated to dryness, a few cubic centimeters of concentrated HCl added, and evaporated to dryness. This process is repeated twice more, in order to convert the nitrates completely into chlo- rides. The dry chlorides are treated with 3–1 c.c. of concentrated HCl, half as much absolute alcohol is added, and the solution is filtered through asbestos. RESIDUE (BaCl2). SolūTION (SrCl3). This is evaporated to dryness and tested as under A. The SrCl3 is all removed by washing with absolute alcohol and the residue is tested as under 4. It is evaporated to dryness a n d tested as under A. 392 COURSE OF ANALYSIS. * NOTES ON TABLE VIII. If the original Solution shows an alkaline reaction (with phenolph- thalèin or litmus), a precipitate may be formed on the addition of hydrochloric acid even when none of the metals of the first group are present; in which case the precipitate must be examined by a special method. Thus a solution of water-glass gives a white, gelatinous pre- cipitate of silicic acid on adding hydrochloric acid, but also on the addition of nitric or Sulphuric acid. In the same way, a solution of the alkali tungstates gives a white amorphous precipitate of tungstic acid upon the addition of acids; the precipitate becomes yellow on warming (cf. page 440). Solutions containing the sulpho salts of arsenic, antimony, and tin are decomposed by the addition of dilute acid, with precipitation of insoluble arsenic, antimony, and tin Sulphides. The water-glass solution can be freed from its silicic acid by evaporation with acid, as described on page 358, and the filtrate analyzed in the usual way for the metals. In exactly the same way tungstic acid can be removed from a solution of an alkali tungstate. If sulpho salts are present, the solution is decomposed by the addition of a dilute acid, and the precipitate is examined according to Table IX, while the solution is tested for alkalies and alkaline earths according to Tables XI and XII. Only in case non-volatile organic matter is present is it neces- sary to test the solution for iron and aluminium. § EXAMINATION OF GROUP V (Mg and the Alkalies). The filtrate from the (NH4)2CO3 precipitate can contain magnesium and the alkalies, in the presence of considerable amounts The solution is divided into two parts; one part is used The larger part of the hydrochloric acid solution, which is now free from ammonium salts but can contain magnesium,” is carefully treated with BaCl, solution in order to change any sulphate which may be present into chloride, and Ba(OH)2 is then added, Ammonia and ammonium carbonate are added, the solution warmed to render the BaCO3 crystalline, and filtered. The filtrate will then contain small amounts of barium, owing to the solubility of BaCO3 in ammonium salts; it is tºº to dryness, the ammonium salts jić by gentle ignition, the residue dissolved in a little water, a few drops of ammonium carbonate and ammonia added, then it is warmed and filtered. The filtrate, which is now free from barium, is evaporated to dryness, the small amounts of ammonium salts are expelled, and a small portion of the residue tested to see what color it imparts to the flame. t (a) If the flame is violet, K is present and Na is absent. (b) If the flame is colored yellow permanently, Na is present. It is observed through cobalt, glass; if the flame appears pink, K is also present. e he rest of the residue is dissolved in a little water, the deposited carbon is filtered off, the solution divided into two parts. One part is tested for potassium, and the other part for sodium. KaBLaSb2O7 produces a white crystalline precipitate. of ammonium salts. TEST FOR MAGNESIUM. * TABLE XII. It is evaporated to dryness, gently ignited in order to expel the ammonium salts, dissolved in a little dilute HCl, and filtered from the carbonaceous residue (cf. page 53). for the Mg test, and the other part for the K and Na tests. TESTs FOR POTASSIUM AND SODIUM. The smaller portion of the hydrochloric acid solution is treated with a little NH3, and, in case a precipitate is formed, ammonium chloride is added until the solution is clear. Sodium phosphate is then added, the solution is stirred and saturated with ammonia. A white crystalline precipi- tate of magnesium ammo- nium phosphate shows Mg to be present. without filtering, until the solution is strongly alkaline. RESIDUE [Mg(OH)2, BaSOJ. It is then filtered. SoLUTION (NaCl, KCl, BaCl, , Ba(OH)2]. Discarded. TESTs FoR POTASSIUM. 1. HaPtCls produces a łºść crystalline precipitate of KaptCl6. 2. Tartaric acid, in the presence of sodium acetate, produces a white crystalline precipitate of KHC4H4O6. TEST, FoR SoDIUM. * If Mg is absent, the treatment with BaCl2 + Ba(OH)2 should be omitted. The analysis is started at t. 394 COURSE OF ANAL YSIS. EXAMINATION FORTHE NEGATIVE ELEMENTS (ANIONS). The tests for the acids (anions) are always made after the analy- sis for the metals (cations), for the preliminary examination (heat- ing in the closed tube and with dilute and concentrated sulphuric acid) and the solubility, combined with the knowledge of the metals that are present, suffice to tell us what acids may and what acids cannot be present. In order to avoid side-reactions, the acids are first obtained in the form of the neutral alkali Salts before proceeding to test for them. PREPARATION OF THE SOLUTION FOR THE ANALYSIS FOR ACIDS. Two cases can be distinguished: A. The original substance contains no heavy metal (i.e., only alkalies or alkaline earths are present). (a) The substance is soluble in water. The solution is tested with litmus paper to see whether it is acid, alkaline, or neutral. An Alkaline Reaction shows the possible presence of alkali cyan- ides, alkali nitrites,” borates, tertiary phosphates, alkali Sulphides, Sulpho salts of the alkalies, alkali silicates, etc. An Acid Reaction is shown by many acid salts. The solution is divided into two parts. If it is neutral, it is analyzed directly for the acids; if it is alkaline half of it is exactly neutralized f with acetic acid and the other half with nitric acid; if it is acid, it is neutralized with sodium carbonate solution. (b) The substance is insoluble or very difficultly soluble in water, but readily soluble in dilute acids. In this case only the acids of Groups III and IV need be tested for. The dry substance is boiled with a little concentrated sodium carbonate solution and filtered. The filtrate contains all the acids in the form of their sodium salts, with the exception of phosphoric acid, which has already been tested for in the analysis for the metals. The solution is now neutralized with dilute nitric acid. (c) The substance is insoluble in water and in dilute acids. The following substances may be present: BaSO4, SrSO, (CaSO.), * Entirely pure alkali nitrites are neutral. The alkaline reaction of the commercial salts is due to the presence of alkali oxide, or silicate. * † Sulphosalts, silicates, stannites, stannates, aluminates, molybdates, tungstates, etc., will yield precipitates at this point which must be examined according to Table XIII. EXAMINATION FOR THE NEGATIVE ELEMENTS. 395 CaF, and silicates, which often contain H.P.O., HBO, H.SO, HF, and HCl. The substance is fused with sodium carbonate in a platinum crucible, extracted with water, and the aqueous Solution thus ob- tained is used for the analysis for acids, after being neutralized. If the substance is partly soluble in water and in acids, it is first treated with water and then with soda solution, while the residue is fused with sodium carbonate. The three Solutions thus obtained arc analyzed separately. B. The substance contains heavy metals. (a) It is Soluble in Water or in dilute acids, and contains no non-volatile organic matter (no carbonization in the closed-tube test). The solid substance is treated with sufficient concentrated soda solution to make the resulting solution weakly alkaline, and filtered. If ammonium salts are present, it is first boiled with the solution of sodium carbonate until the vapors from the y ition no longer smell of ammonia, when it is filtered. The resulting solution is divided into t to parts, one part being acidified with acetic acid, and the other with nitric acid. (b) The Substance is soluble in Water or Dilute Acids and Contains Non-volatile Organic Matter.—If metals of the ammo- nium Sulphide and hydrogen Sulphide groups are both present, hydrogen Sulphide is conducted into the weakly acid solution until it is saturated, the precipitate is filtered off, ammonia is added to the filtrate until it is slightly alkaline, it is filtered again, and this last filtrate is acidified with acetic acid and evaporated to a small volume. The deposited Sulphur is filtered off, the Solution treated with Solid potassium carbonate, filtered if necessary, carefully acidi- fied with nitric acid, vigorously stirred, and if any acid potassium tartrate is formed, it is filtered off and tested as described under tartaric acid. The filtrate is tested for the remaining acids. (c) The Substance is Insoluble in Strong Acids.-Besides the salts mentioned under A (c), the following may be present: AgCl, AgBr, AgI, AgCN, PbSO, silicates (ferro- and ferricyanides). If silver is present, the halogen acids must be looked for. The insoluble silver Salt is reduced by zinc and sulphuric acid, the residue filtered off, and the filtrate examined according to Table XV for HCl, HI, HBr, and HCN. The original neutral, aqueous solution, or the solution which is acid with acetic or nitric acid, is treated with an excess DISSOLVES. THE PRECIPITATE REMAINs UNDIssolvKD. All acids of Acids of Groups I, III, and IV may be present. The precipitate is warmed with Groups I, III, dilute nitric acid. and TV are ab- || THE PRECIPITATE sent. DISSOLVES. THE PRECIPITATE DOES NOT DISSOLVE. One or m or e Present: Acids of Group I. acids of Groups The precipitate is filtered off and the filtrate treated drop by º and IV is pres– drop with dilute ammonia. “The a c id s of A TURBID Zone: Is FoEMED. No TURBID ZONE IS FORMED Group I are ab- w Sent. Present: Groups * III or IV. Absent: Groups III Yellow zone, Aga/AsOs, Ag3RO, and IV. Red “ Ag,CrO, Brown “ Aga AsO, f White “ silver salts of H,SOs, H.C.H.O., H2C2O, HPOs, H.P.O. † It must not be forgotten that NH3 produces a brownish precipitate of Ag2O in silver Solutions; in case the zone appears brown, the solu- tion is gently shaken, whereby the turbidity disappears; more NHS is added, and with a little experience it can be determined whether the tur- bidity is AgºO or Ags.AsO4. TABLE XIII. THE SILVER NITRATE TEST. * of silver nitrate. NO PRECIPITATE Is ForMED. . A PRECIPITATE Is FoEMED. All acids of Groups I, III, and TV are ab- sent (cf. pages 243–244). If the solution was very concentrated, the precipitate may be the silver salt of HC2H5O2, HNO2, or H.SO,. It is, therefore, treated with water and warmed. THE PRECIPITATE * Boric acid would not be precipitated in case the free acid were neutralized with ammonia. It must be tested for separately. # § § The original, aqueous, neutral solution, or the solution neutralized with acetic acid, is treated with an excess of barium chloride solution. The precipitate is filtered off and the filtrate neutralized with pure KOH or NaOH (free from carbonate, etc.). No PRECIPITATE IS FORMED. A PRECIPITATE Is ForMED. Absent: Acids of Groups Present: Acids of Groups III and IV. II and IV. TABLE XIV. THE BARIUM CHLORIDE TEST. No PRECIPITATE IS FORMED. • A PRECIPITATE Is FoEMED. * Absent: All a c id s of A little dilute HCl is added and the solution warmed.* IV. and VI. Groups III, IV, an THE PRECIPITATE DISSOLVES. THE PRECIPITATE REMAINs. A b sent : Group VI (H,SO, and HF). p Present: Group VI. * Carbonates, sulphites, and thiosulphates will be destroyed; they must afterwards be tested for specially, 398 COURSE OF ANALYSIS. If the insoluble substance contains lead, it is boiled with sodium carbonate solution and filtered; the filtrate is acidified with hydro- chloric acid, and tested with BaCl, for H.SO, If silicates are present, H.PO, HF, HBO, HCl, and H.SO, must also be tested for. In whatever way a Solution is prepared, its behavior toward silver nitrate and barium chloride is determined, in order to ascer- tain to what groups the acids present belong. After this is determined, the different tests for the individual members are made. TABLE XV. ExAMINATION OF GROUP 1. HCN is first tested for, by placing a little of the solution on a watch-glass, adding a few drops of yellow ammonium sulphide, evaporating carefully to dryness, acidifying the dry mass with HCl, and adding a drop of FeCl, solu- tion. If a blood-red coloration is produced, HCN is present, in which case a larger portion of the neutral solution is treated with nickel sulphate * so- lution in excess, and filtered. PRECIPITATE SOLUTION. Ni(CN), . . The solution, which is now free from hydrocyanic acid, Discarded. is treated with a little caustic soda solution (free from halo- gen), boiled, and the precipitate of Ni(OH), filtered off. The filtrate is divided into two parts, one part being used for the HBr and HI tests and the other for the HCl test. Tests for HI and H.Br. Test, for HC1. The solution is acidified with dilute H2SO4, chlorine water is added drop by drop, and the solution is shaken with CS, or CHCl2. If the lattſ r is colored violet, HI is present. By , further addi- tion of chlorine water, the CS, or CHCl, is decolorized completely if HBr is absent, but turned yellowish-brown if HBr is present If too." much chlorine water is used a wine-yellow color is pro- duced. The solution is slightly acidified with HNO3, and precipitated by the addi- tion of dilute AgNO, drop by drop. AgI and AgBr are first precipitated (yel- low). This is filtered off, and more AgNO3 is added. If the precipitate still ap- pears yellow, it is filtered through a new filter, and AgNO3 is again added to the filtrate until a white precipita t e of AgCl is formed if IICl is present. * If hydroferricyanic acid is present a little ferrous sulphate is also added. Hydro- ferrocyanic acid is completely precipitated by nickel sulphate. EXAMINATION FOR THE NEGATIVE ELEMENTS. 399 GROUP II. The members of this group are almost always detected in the preliminary examination. The special tests for these acids are described on page 284 et seq. GROUP III. SO, CO, H.C.O., are recognized in the preliminary examination. HPO, MH, P.O., HBO, and H,C,ELO, are tested for separately by the special reactions described on page 300 et seq. GROUP IV. CrO, H.PO, and H.S.O, are detected in the preliminary exami- nation, and in the analysis for metals. GROUP V. HClO, and HNO, are usually detected in the preliminary exami- nation. Their presence is, however, always confirmed by the reactions described on pages 337 and 343. GROUPS VI AND VII. These acids are usually detected in the preliminary examination. Their presence is confirmed by the tests described under H.SO, HF, and SiO,. IB. THE SUBSTANCE IS A METAL OR AN ALI/OY. The analysis of a metallic alloy is much simpler than that of a mixture of salts, because there are no acids to test for. Of the electro-negative elements, usually only phosphorus, silicon, carbon, and sulphur have to be considered. As all metals, with the exception of gold, platinum, tin, and antimony, are soluble in nitric acid, alloys are usually brought into solution by dissolving therein, and the use of aqua regia is only necessary in a few cases. Many alloys rich in silicon (e.g. copper silicide) are extremely difficultly soluble even in aqua regia, and are best brought into solution by fusing with caustic alkali in a silver crucible, and afterwards dissolving the melt in nitric acid. It is not advisable to dissolve an alloy in hydrochloric acid, for phosphides, carbides, silicides, sulphides, and arsenides, which are often present in Small amounts, are evolved in the form of their 4oo COURSE OF ANALYSIS. hydrogen compound, and thus escape detection. For the analysis of ordinary alloys, the following procedure is used: One to two grams of the alloy (best in the form of borings) are placed in a 200 c.c. porcelain dish under a good hood and treated with about 20 c.c. of nitric acid, sp. gr. 1.25–1.30 (1 vol. conc. HNO,-- 1 vol. H.O). After the first violent reaction is over, the acid is carefully evaporated (with constant stirring) almost to dryness, being careful to avoid overheating; * a little water is added and warmed. (a) The mass dissolves completely. The alloy contains neither tin nor antimony; it is analyzed according to Table XVI. (b) The mass does not dissolve completely, but a white, greenish residue remains; it is analyzed according to Table XVII. C. THE SUBSTANCE IS A LIQUID. The color, odor, and reaction towards litmus enable one to draw important conclusions. (a) The Solution reacts neutral; it contains no free acid, free base, acid Salt, no salt which shows an acid or alkaline reaction on account of hydrolysis, nor any insoluble salt. & First of all, it is determined whether there are any solid sub- stances dissolved in the liquid by evaporating a small portion to dryness at as low a temperature as possible (so as not to lose any volatile substances). If a residue remains, it is examined according to A, page 365. (b) The solution reacts alkaline. An alkaline reaction may be due to the presence of hydroxides of the alkalies or alkaline earths, peroxides, carbonates, borates, cyanides, silicates, sulphides (zin- cates, aluminates, molybdates, tungstates) of the alkalies, or am- monia. If the solution, for example, contains hydroxides or car- bonates of the alkalies, it is evident that substances which are precipitated by them cannot be present at the same time, except, in some cases, in the form of complex ions (cyanides), tartrates, etc. g- The solution is at once tested for peroxides, hydroxides, and carbonates, as well as for the sulphides of the alkalies. To test for peroxides (H,0,) a small portion of the solution is warmed with a few drops of cobalt nitrate solution, a black precipi- * Otherwise insoluble basic salís are 1.kely to be formed. If this be the case, as is often shown by the dark color of the residue, a little conc. HNOs is added, the liquid is warmed gently and then diluted with water. º A. The alloy after being evaporated with nitric acid dissolves in water, forming a clear solution. The alloy contains neither tin nor antimony. e * º g g e Two c.c. of H2SO, are added and the solution evaporated until white fumes of sulphuric acid come off thickly. RESIDUE. SoLUTION. tº Can contain Bi, Cu, Cd, Fe, Mn, Ni, Co., Zn, Al, P.Os. Pºit e It is saturated with H.S and filtered. 4". PRECIPITATE. SoLUTION. It is w a s h ed º with water Bi,Ss Col.S. CuS. Feso, MnSO4, N iso, Coso, 4. Al2(SO4)3H3PO4. d and the It is di i. d i y dilut The excess of H2S is expelled by boiling, a few drops of concentrated HNOs are adde 1S CIISSO1VeG II). W8.III] * || to oxidize the iron, 10 c.c. of NH4Cl solution, and NH3 solution are added drop by drop ºn't of HNO3, the deposited sul hur fill tº the boiling liquid until it smells of NH3 permanently, when it is quickly filtered. €3. CQn- || tered off, an excess of NHs added, firmed by || and filtered SoluTION. © PRECIPITATE. & * tº: PRECIPITATE. SOLUTION. [Fe (OH)3. Al(OH)3. (AlFe)PO3]. ſºciolºgº º [Ni(NH3)6]Cl2, COal-S Ü I C It is dissolved in a little HNO3, This is saturated with §§ and filtered, the fil- t e St. and a little of the solution test: |trate being rejected.* The precipitate, consisting malle a b le gºngºso, ºrº. *; ; ||ººd?ºisiented with tº sº 6. * but to n, Bi(OH), white. | NHs) } Fºussian šue shºws presence of cold dilute and filtered. soluble in It is dissolved 1I] A biºd €. RESIDUE. Solution. HNO a little HCl a n d || Solution shows Another portion of the solu- 3 treated with an ||94. The sºlu- tion is tested with ammonium ſe MnCl2, ZnCl2. shows Pb & e tion is decolor- molybdate solution. A yellow NiS, CoS, The H2S is boiled off and the solu- to be pres- alkaline solution 1zed by the ad- crystalline precipitate shows the examined | tion treated with an excess of KöH ent of SnCl2. A black | dition of KCN, presence of P. º: and filtered. sº recipitate of |º (º); The principal part of the so- % * RESIDUE. Solution. i.,O Bish added. A yel- || lution is treated with an excess (A). -- 1293 Or Pl SDOWS | low precipitate | of KOH and filtered. "The fill Mn(OH)2. Zn(OK)2. . . Bi to be present. || of CdS shows | trate, which may contain any Al A small portion || Acetic acid is added and H2S conducted into the solution. A white precipitate of ZnS shows Zn to be present. TABLE XVI. the presence of Ca. as Al(OK)3. is acidified with HCl treated with NH3. white, gelatinous precipitate shows the presence of Al. is fused with soda and nitre on plat- inum foil. A green melt shows the presence of Mn. * If neither iron nor aluminium were present in the alloy, this filtrate is tested for phosphoric acid. tered from deposited sulphur if necessary, and tested with ammonium molybdate. * It is evaporated to a small volume, fil- TABL B. The alloy, after being evaporated with nitric acid, does not all dissolve in water. E XVII. The alloy contains tin or anti- SnO, Sb,0s, (Sb.O.), P.O., Bi,0s, and traces of Cu, Pb, Fe, etc. It is washed with water, KOH added to alkaline reaction, and then 5–10 c.c. of K.S solution. This is warmed some time on the water-bath and filtered. RESIDUE. Solution. - Bi,S, PbS, CuS, etc. R2SnSs, K.SbSs, K.PO,. It is dissolved in HNO3, It is diluted with water, acidified with dilute HCl, and filtered. RH,PO, KCI. The solution is evaporated to a small volume, saturated with ammonia and treated with magnesia mixture. A white crystalline precipitate shows HaPO, or P to be pres- ent.f Examined ac- cording to Table XVI. mony, or both.* It is filtered. RESIDUE. SOLUTION." filtered from sulphur, and the filtrate added to the PRECIPITATE. SoLUTION. e & 1 previous solution. SnS, Sb,Say or Sb2S3. tested according to Table and Sb. It is dissolved in concentrated HCl, boiled, diluted with a little water, filtered from deposited sulphur, and X for Sm, * In the case of many alloys containing silicon (iron and steel), silicic acid separates out on the addition of HNOs. in very small amounts, and is detected as described under silicic acid. t"I'his is confirmed by dissolving the precipitate in HNOs and testing with ammonium molybdate solution. This is usually present f ‘A EXAMINATION FOR THE NEGATIVE ELEMENTS. 4O3 tate shows the presence of H.O.” Or the solution can be tested by adding Some titanium sulphate solution and acidifying carefully with cold dilute-Sulphuric acid; a yellow coloration shows the pres- ence of H.O.f A still more sensitive reagent, according to Schöne (B.B. 7.1695), is a very dilute solution of FeCl,--K, Fe(CN)4. If the slightest trace of H.O., is present in the solution, the red solution becomes greenish, and after a time Prussian blue separates out. In order to detect the presence of hydroxides and carbonates in the presence of H.O., a portion of the solution is boiled for a long time in a porcelain dish in order to destroy the peroxide, when barium chloride is added until no more precipitate is formed. If the Solution then shows an alkaline reaction, the presence of hydrox- ides i is assured. If the precipitate produced by barium chloride dissolves in acid with effervescence, and the gas evolved renders barium hydroxide solution turbid, carbonates are present. If the solution smells of ammonia, a small portion is evaporated to dryness in order to see whether other compounds are present, and the residue is examined according to A, page 365. (c) The solution reacts acid; it can then contain substances which are soluble in water and in acids, as well as free acids. A small portion is evaporated to dryness in order to see whether any non-volatile matter is present. . If no residue is obtained, the Solu- tion is neutralized with soda and tested for acids. If a residue is obtained, it is examined according to A, page 365. D. The substance to be analyzed is a gas. This case will be considered in Volume II under Gas Analysis. * If the alkaline solution contains hydrochlorites or sulphides, these will also give a black precipitate with cobalt nitrate; the above reaction serves only to detect H.O., in the absence of hypochlorites or sulphides If the latter bodies are present, H2O, cannot be present, because hypochlorites are reduced to chlorides and sulphides oxidized to sulphates by H2O2. The presence of hypochlorites is usually detected by the odor; by acidi- fying with dilute H.SO, chlorine can be smelled. Sulphides give off H.S on being acidified. Hypochlorites and sulphides cannot exist together in the same solution. f H.O., can also be detected by the chromic acid reaction, but this test is less certain than that with titanium sulphate. f Either as such in the original Solution or by hydrolysis of peroxides. This test for OH ions in the presence of carbonates has been used for years by the author. It is reliable, though care must be taken to add an excess of BaCl, as otherwise an alkaline reaction may be due to BaCOs, which is more soluble in water than in BaCl, Solution. SUPELEMENT. REACTIONS OF SOME OF THE RARER METALS. REACTIONS OF SOME OF THE RARER METALS. IN the treatment of the rare metals, the same method and order will be used as with common metals. THE ALKALI. GROUP. CAESIUM, RUBIDIUM, LITHIUM. CAESIUM, Cs. At. Wt. 132.9. Occurrence.—Caesium and rubidium are not rare elements, strictly speaking, for they are found almost everywhere, but always only in very small amounts. Thus cæsium replaces potassium in many feldspars and micas, and is found in many rocks which carry these minerals, as well as in mineral waters which ooze from them. Caesium and rubidium were discovered in the mother liquor. of Durkheimer brine in the year 1860 by Bunsen and Kirchoff by means of the spectroscope. Pollucite, a mineral closely related to leucite, found at Elba, and crystallizing in the regular system, is a typical capsium mineral. Its composition is [SiO2),Al,Cs, H. Caesium and rubidium in all their reactions behave almost exactly- like potassium. The principal difference between these three metals lies in the different solubilities of their corresponding salts, as will be seen by the table given on page 410. REACTIONS IN THE WET WAY. A solution of cassium chloride should be used. 1. H.PtCl, produces a yellow crystalline precipitate which is of a brighter color than the corresponding potassium salt and is much. 4O7 408 REACTIONS OF SOME OF THE RARER METALS. more insoluble; 100 parts of water dissolve at 0° C. only 0.024 part and at 100° C. 0.377 part of the salt. 2. Tartaric Acid produces, as with potassium and rubidium, a white, crystalline precipitate, CSHCH, Os; 100 parts of water dis- Solve at 25°C, 9.7 parts, and at 100° 97.1 parts of the salt. 3. H. SnClo (a solution of SnCl, in concentrated HCl) produces in concentrated solutions a white, crystalline precipitate of Cs,SnCl, (octahedrons). Ammonium salts give the same reaction, but potas- sium and rubidium Salts do not. REACTIONS IN THE DRY WAY. Caesium compounds color the flame reddish violet, very similar to the potassium flame. Spectrum.—A number of lines in the orange, yellow, and green, and two particularly bright lines in the blue, one of which almost coincides with the blue strontium line. See chart. RUBIDIUM, Rb. At. Wt. 85.5. Occurrence.—Rubidium almost always accompanies caesium and is found in many mineral waters; in carnallite from Stass- furt; in lepidolite, Si,O,Al,(Li,K,INa),(F,0H,) *; in triphylite, PO,(FeMn)(LiCsRb); and in spodumene, (SiO,),Al(LiNa), a mineral of the pyroxene group. As far as the author knows, there is no typical rubidium mineral. REACTIONS IN THE wit way. 1. H.PtC1, produces, as with cassium and potassium salts, a white, crystalline precipitate of Rb, PtCl4, which is more difficultly soluble than the corresponding potassium Salt, but more soluble than the caesium salt; 100 parts of water dissolve at 0°C. 0.134 part, and at 100° 0.634 part of the salt. 2. H. SnCl, produces a white precipitate only in very concen- trated solutions. The salt is much more soluble than the corre- sponding caesium salt, but this reaction is not suitable for separating the two metals. . 3. Tartaric Acid produces a precipitate of RbHC, H, O, only in concentrated solutions: 100 parts of water dissolve at 25° C. 1.18 * Lepidolite from Rozena contains about 0.54 per cent. Rb and 0.0014 per cent. Cs. Af IITHIUM. - 4O9 parts, and at 100° C. 94.1 parts of the salt. The corresponding capsium salt is more soluble, while that of potassium is less soluble. REACTIONS IN THE DRY WAY. Flame Coloration.—Similar to capsium. Spectrum.—Besides several lines in the red part of the spectrum, the rubidium flame is characterized by two bright lines in the ultra- violet. LITHIUM, Li. At. Wt. 7.03. Occurrence.—Lithium is found to a greater extent in nature than caesium and rubidium; in triphylite, PO,(FeMn)(Li,Cs, Rb); in peta- lite, Si,Oloal(Li,Na,FI), a mineral of the feldspar group (also called castorite); in amblygonite, Li(Alf)PO4, monoclinic; in lepidolite, Si,O,Al,(Li,K,INa),(F,0H), ; in many varieties of turmaline, and muscovite, in epidote and Orthoclase, and consequently in many mineral-spring waters. In some cases as much as 36 milligrams are contained in a liter of spring water. Lithium is the lightest of all metals, and floats on petroleum. It oxidizes quickly in the air, and decomposes water at ordia: peratures, forming LiOH, which dissolves slowly in the wºe solution reacts alkaline and absorbs carbon dioxide from iſ avidity, forming difficultly soluble Li,CO. Lithium chloride is soluble, even in the anhydrous state, in a mixture of alcohol and ether as well as in amyl alcohol (difference from the remaining metals of this group). REACTIONS IN THE WET WAY. A solution of lithium chloride should be used. 1. H.PtC1, produces no precipitation. 2. Tartaric Acid produces no precipitation. 3. Na, HPO, produces from moderately concentrated solutions a white precipitate of tertiary lithium phosphate on boiling. The precipitation is only quantitative when the Solution is made alka- line with caustic soda, evaporated to dryness, and taken up in water containing ammonia: Na,HPO,--3LiCl-NaOH=3NaCl-i-Li,PO,4-H,O. 4IO REACTIONS OF SOME OF THE RARER METALS. Lithium phosphate is fusible (difference from magnesium and the alkaline earths). 4. (NHA),COs. If ammonium carbonate and ammonia are added to a concentrated lithium solution, lithium carbonate is pre- cipitated in the form of a white powder. The salt, contrary to the other alkali carbonates, is very difficultly soluble in water: 100 parts of water dissolve at 13°C. 1.31 parts of Li,CO. In the presence of considerable alkali chloride or of ammonium chloride, no precipi- tation takes place. REACTIONS IN THE DRY WAY. Flame Coloration.—Pure lithium salts impart a magnificent carmine-red coloration to the gas-flame. If considerable amounts of sodium salts are present at the same time, the lithium flame is completely masked; but if the flame is observed through cobalt glass the red color becomes distinctly visible. Spectrum.—An intensive red line, which is nearer to the sodium line than is the red potassium line. SUMMARY OF THE ALFCALI METALS. See chart of spectra. LITHIUM. Sod1UM. PoTASSIUM. RUBIDIUM. CAESIUM. 7,03 23.05 39, 15 85.4 133. 180° C. 96° C. 62.5° C. 38.5° C. 26–27° C. . . insoluble insoluble insoluble X2PtC 6 com- p ou n d in Water. 100 parts of water at 20°C. Y. dissolve . . . . considerable | considerable 1. 12 0.141 0.079 100 parts of w a t e r a t 100° C. dis- Solve . . . . . . . J considerable | considerable 5. 13 0.634 0.377 Solubility of acid tart rates 100 parts water dissolve— at 10° C. | | considerable | considerable 0.425 -*. tº- at 25° C. sºmºsº -º-º: - 1.18 9.7 Solubility of al- TIII). S 100 parts water jºlve at 13.5 27 17° C. . . . . . * º 3. 2. 0.619 Sº...”. readily insolubl lubl lubl Chior) Cies 1 In 11\SOIU1016 insoluble insoluble insoluble alcohol-ether. ) soluble e § Solubility of the carbonates in insoluble insoluble insoluble insoluble soluble absolute alco Ol. . . . . . . . . . DETECTION OF LITHIUM, RUBIDIUM, AND CAESIUM. 4 II DETECTION OF LITHIUM, RUBIDIUM, AND CAESIUM IN THE PRESENCE OF CONSIDERABLE SODIUM AND POTASSIUM. The solution containing the metals as chlorides is evaporated almost to dryness, treated with 90 per cent. alcohol (which is well rubbed into the residue) and filtered. The alcoholic solution con- tains all of the Li, Rb, and Cs, with small amounts of K and Na. It is again evaporated almost to dryness and once more extracted with alcohol.” This alcoholic solution is now evaporated to dryness, the residue treated with concentrated hydrochloric acid, f again evaporated, gently ignited over the free flame, and, after cooling, the residue is rubbed with ether-alcohol mixture by means of a glass rod, after which the solution is filtered through a filter that is moistened with the ether-alcohol mixture. The filtrate which contains the lithium is evaporated to dryness and then tested for lithium by means of the flame reaction, and in the spectroscope. The residue insoluble in alcohol-ether is dissolved in a little water, precipitated with hydrochlorplatinic acid and filtered. The precipitate is repeatedly treated with boiling water in a porcelain dish, decanting off the liquid each time through the filter. First of all, the potassium salt dissolves in the hot water, form- ing a yellow solution. The treatment with hot water is continued until the residue is of a bright-yellow color. It is then dried, placed in a porcelain boat, and heated in a glass tube made of difficultly fusible glass in a stream of dry hydrogen, whereby the alkali chlor- platinates are reduced to chloride and platinum: X, PtCl4-4H=4HCl·H2XCl+Pt. After cooling, the residue is treated with a little water, the plati- num filtered off, and the solution is evaporated to dryness and tested in the spectroscope for Cs and Rb. tº: traces of the rare alkalies are present, the residue must be ex- tracted several times with alcohol f This is done in order to change LiOH into LiCl Some of the former is formed by evaporating the solution, and it is insoluble in alcohol-ether. 4 I 2 REACTIONS OF SOME OF THE RARER METALS. In order to detect lithium, casium, and rubidium in a silicate undecomposable by acids (lepidolite, for example) the finely-pow- dered mineral is decomposed with hydrofluoric and sulphuric acids as described on page 361, the sulphates changed to chlorides by the addition of barium chloride, and the solution freed from other metals as described on page 52 in the case of magnesium, after which the above separation is made. METALS OF THE (NHos GROUP. S- ~! BERYLLIUM, ZIRCONIUM, THORIUM, YTTRIUM, ERBIUM CERIUM, LANTHANUM, DIDYMIUM, TANIALUM, NIOBIUM. —A BERYLLIUM, Be. At. Wt. 9.1. Occurrence.—Chrysoberyl, (AlO.),Be; phenacite, SiO,Be,; beryl, Si6O1sAl,Bes; euclase, SiO,AlBeFI; meliphanite, Si,0,...FBe,Ca,Na; and leucophanite, Si,0), FBeCa,Na. Beryllium is a bivalent metal, and forms a white oxide, BeO, which is soluble in acids. Beryllium salts react acid in aqueous solu- tion and possess a sweetish, astringent taste. Frequently the ele- ment is called Glucinum, Gl. REACTIONS IN THE WET WAY. A solution of BeSO,-4H.O should be used. 1. Ammonia and Ammonium Sulphide produce a white pre- cipitate of Be(OH), similar in appearance to Al(OH), insoluble in an excess of the precipitant, but readily soluble in HCl, forming a colorless solution. The yellow color often obtained in dissolving the hydroxide in hydrochloric acid is due to traces of ferric chloride.* 2. KOH precipitates white, gelatinous beryllium hydroxide, readily soluble in an excess of the reagent, forming potassium beryl- late: Be(OH)2+2KOHz−2 Be(OK)2+2H2O.f The alkali beryllates are decomposed hydrolytically on boiling their dilute aqueous Solutions, all of the beryllium being pragipi- * * Haber and v. Ordt, Zeitschr. f. anorgan. Chem., 38 (1904), p. 382. † According to Hantzsch, the beryllium hydroxide is present in solution in a colloidal condition as in the case of zinc (see page 139). j ZIRCONIUM. 4 I 3 tated as hydroxide. The precipitate thus obtained is, according to Haber and v. Ordt, much denser than that thrown down by am- monia, and differs from the latter by being insoluble in potassium carbonate and difficultly soluble in ammonium carbonate solutions; it is also much more difficultly soluble in dilute acids. A solution containing a considerable excess of potassium alkali hydroxide is not precipitated by boiling. 3. Ammonium Carbonate produces a white precipitate of beryl- lium carbonate, readily soluble in an excess of the reagent (difference from aluminium); by boiling the solution, the beryllium is pre- cipitated as white basic carbonate. This property enables us to separate beryllium from iron and aluminium. The separation is not sharp, however, which can also be said of the separation by means of caustic potash.* 4. BaCO3 precipitates beryllium completely in the cold as hydroxide. 5. Oxalic Acid and Ammonium Oxalate cause no precipitation (difference from thorium, zirconium, erbium, yttrium, cerium, lanthanum, and didymium). 6. K2SO4 gives with beryllium sulphate a beautifully-crystalline SO4—K double salt, > Be +2H2O, which is soluble in a concentrated SO4—K solution of K2SO4 (difference from Ce, La, and Di). 7. BeCl2 is soluble in a mixture of equai volumes of saturated aqueous and ethereaſ hydrochloric acid, while the hydrous alumin. ium chloride is not (good method for separating Be and Al).f There are no characteristic dry reactions for beryllium. ZIRCONIUM, Zr. At. Wt. 90.6. Occurrence.—Zircon, ZrO2,SiO2, tetragonal, isomorphous with rutile, Ti2O4, thorite (orangite), ThC2,SiO2, cassiterite, SnzO4, polianite, JIn2O4, and plattierite, Pb2O4; baddeleyite, ZrO2, mono- clinic. , ir Zirco ium forms two oxides: Zirconium dioxide, ZrO2, and zirconium pentoxide, Zr2O5. The former is the more important order to make a quantitative separation, the beryllium hydroxide or carbonaº must be redissolved and the precipitation repeated several times. t F. S. Havens, Zeitschr. f. anorg. Ch., 18 (1898), p. 147. 4 I4. REACTIONS OF SOME OF THE RARER METALS. and can be dissolved by heating for a long time with a mixture of two parts of concentrated H2SO4 and One part of water and after- wards diluting. The mineral zircon, ZrSiO4, cannot be decom- posed by such treatment. It must be finely pulverized and fused with four times as much sodium carbonate at a high heat in a platinum crucible; SOdium silicate, NaaSiO4, and sodium zirconate, Na4ZrO4, are formed. On treating the melt with water, the former salt dissolves, while the latter is decomposed hydrolytically, form- ing sodium hydroxide and Sandy, insoluble zirconium hydroxide; the latter retains some of the caustic soda very persistently. After washing the residue, it is heated without previous drying with con- centrated Sulphuric acid at a temperature just near boiling-point, and in this way the anhydrous Sulphate Zr(SO4)2 is obtained. By pouring water over the latter the Salt Zr(SO4)2.4H2O is formed, which dissolves slowly in cold water, but more readily in hot-water, forming a solution with an acid reaction. f REACTIONS IN THE WET WAY. A solution of zirconium nitrate or a freshly-prepared one of zirconium oxychloride may be used for the following reactions: 1. NH4OH and (NH4)2S produce a white gelatinous precipitate of Zr(OH)4, insoluble in an excess of reagent. 2. KOH and NaOH likewise produce the same precipitate insoluble in an excess of reagent (difference from Al and Be). When the zirconium hydroxide is produced in the cold it is readily soluble in dilute acids; but when thrown down from a boiling solution it is very difficultly soluble in dilute acids, though it will dissolve even then in concentrated acids without difficulty. 3. (NH4)2CO3 produces a white, flocculent precipitate of basic carbonate, readily soluble in an excess of the reagent, out reprecipi- tated by boiling. *- | 4. K2CO3 and Na2CO3 produce white precipitates somewhat soluble in an excess, but reprecipitated by ammonia. 5. BaCO3 causes incomplete precipitation, even on boiling. 6. Oxalic Acid gives a white, flocculent precipitate of zirconium oxalate; readily soluble in an excess of Oxalic acid, difficultly ble in dilute hydrochloric acid, and readily’soluble in º: late. From the solution in (NH4)2C2O4 the zirconium is hºt pre- cipitated by the addition of dilute HCl (difference from Th). ZIRCONIUM t 4I4(l 7. Ammonium Oxalate behaves the same as Oxalic acid. From the solution in ammonium oxalate zirconium is not pre- cipitated on the addition of hydrochloric acid (difference from thorium). Remark.-A solution of zirconium sulphate behaves quite differ- ently from that of the nitrate and oxychloride towards oxalic acid and ammonium oxalate, a fact which, although published by Ber- Zelius and also by Pfaff, had been entirely forgotten by most chem- ists until their attention was recently called to it by R. Ruer.* On treating an aqueous solution of zirconium sulphate with oxalic acid or ammonium oxalate there is no precipitation; in fact, precipitation will not take place from nitrate or chloride solu- tions when these contain sufficient Sulphuric acid, sodium or am- monium sulphate. The cause of this different behavior lies in the fact that zirconium forms complex compounds with sulphuric acid and alkali sulphates. Thus the solution of zirconium sulphate contains the acid H2|ZrO(SO4)2], and on treating a solution of the oxychloride or nitrate with sodium or ammonium sulphate (but not the potassium salt) the sodium or ammonium salt of this complex acid is formed: ZrOCl2 +2Na2SO4–2NaCl-i-Na2|ZrO(SO4)2]. These compounds, however, are electrolytically dissociated in aqueous solution as follows: H2(ZrO(SO4)]= H'-i-H +[ZrO(SO4)2]”. As the zirconium is present in the anion it cannot react with oxalic acid. 8. HF causes no precipitation (difference from Th and Y). 9. K2SO4.—A concentrated, cold solution of K2SO4 precipitates little by little all of the zirconium as zirconium potassium sulphate, insoluble in an excess of the reagent (difference from Al and Be). The precipitate, when broduced in the cold, dissolves readily in con- siderable dilute HCl. If it is produced from a boiling solution, basic zirconium sulphate is formed by hydrolysis, which is quite insoluble in dilute HCl (difference from Th and Ce). 10. Na2SO4 produces no precipitation, even on boiling the solu- tion, which is slightly acid with sulphuric acid (difference from Ti). * * Zeitschr. f. anorg. Ch., 42 (1904), p. 85. 4I4b REACTIONS OF SOME OF THE RARER METALS. 11. H2O2 precipitates white, voluminous zirconium peroxide, Zr2O5, from slightly-acid solutions, which evolves chlorine on being warmed with concentrated HCl. 12. Na2S2O3 precipitates zirconium completely as the hydrox- ide, the precipitate being always contaminated with sulphur. 13. Turmeric Paper, after being moistened with the hydrochloric acid solution of a zirconium salt and dried, is colored reddish-brown (difference from Th). 14. HC1. Ruer * recommends the following reaction for the identification of zirconium. Zirconium is precipitated in the cold by ammonia, filtered, washed, and separated from the filter as completely as possible. It is dissolved in hydrochloric acid (or if the precipitate is small in amount the paper and precipitate are treated together with not-too-strong HCl and filtered). The hydrochloric acid solution is evaporated to dryness on the water- bath and the residue taken up in a little water. To the cold, sat- urated Solution hydrochloric acid is added drop by drop, when the presence of zirconium will be evident by the formation of a volumi- nous precipitate of zirconium oxychloride. The precipitate is re- dissolved by heating the solution, and the liquid is allowed to cool. After some time fine, silky needles of ZrOCl2.8H2O will precipitate. In the somewhat unusual case that zirconium is present in the form of the insoluble meta-zirconium acid, the latter is transformed into zirconium sulphate by heating with concentrated sulphuric acid (2:1), this is dissolved in water, the zirconium precipitated by ammonia, and the above process carried out. REACTIONS IN THE DRY WAY. ZrO2 is infusible in the oxyhydrogen flame (difference from the other earths), but glows brightly. THORIUM, Th. At. Wt. 232.5. Occurrence.—Thorite (orangite), Th9iO4, with 50–58 per cent. ThC)2; thorianite, a mineral recently discovered in Ceylon, with 72–76 per cent. Th92 and 11–12 per cent. UO2; f gadolinite, SiO4Beſy,Ce,La, Di, Th,O)2Pe; monazite, PO4(Ce,NLa, Di, Th), with * Zeitschr. f. anorg. Ch., 42 (1904), p. 85. f Chem.-Ztg. Rcp. 1905, p. 91. THORIUM. 415 2–8 per cent. Thos; and in the rare niobates, Samarskite, pyro- chlore, euxenite, etc. Thorite, monazite, and gadolinite are decom- posed by acids, preferably sulphuric acid. REACTIONS IN THE WET WAY. A solution of Th(SO,), should be used. 1. (NH)OH, (NH),S, or KOH produces a white precipitate of Th(OH)4, insoluble in an excess of the reagent, but readily Soluble in dilute acids. By igniting the hydroxide, Th(O, is obtained, which is only soluble in concentrated sulphuric acid after long digestion. 2. K.CO, or Na,CO, precipitates the white carbonate, soluble in an excess of the reagent, and is not reprecipitated by the addition of ammonia. On boiling, the solution becomes turbid, but clears again on cooling. 3. (NH9CO, precipitates the white carbonate, readily soluble in an excess; on warming to 50° a basic carbonate is precipitated, which redissolves on cooling the Solution. Ammonia causes no pre- cipitation in this solution. 4. BaCO, completely precipitates thorium salts in the cold. 5. K,SO, precipitates Th(SO.), K,4-2H,0, difficultly soluble in water and insoluble in concentrated K.SO, solution (difference from Y). The corresponding sodium compound is readily soluble in water. 6. Oxalic Acid precipitates, from Solutions which are not too. acid, all of the thorium as white, crystalline oxalate, practically in- soluble in oxalic and dilute mineral acids. 7. Ammonium Oxalate likewise precipitates thorium oxalate, which dissolves on boiling with a large excess of this reagent. The solution remains clear after cooling, provided the original solution did not contain too much free Sulphuric acid, and enough ammo- nium oxalate was used. From the boiling Solution of the ammo- nium double oxalate, HCl precipitates practically all of the thorium as oxalate (difference from Zr). In the presence of ammonium acetate, ammonium oxalate pro- duces no precipitation; but by the addition of HCl almost all of the thorium will be precipitated as Oxalate. 8. HF produces a white, gelatinous precipitate, which soon changes to a heavy powder. KF causes the same reaction. 416 REACTIONS OF SOME OF THE RARER METALS. 9. Na,S.O, precipitates all of the thorium on boiling. There are no characteristic dry reactions. THE GADOLINITE METALS. YTTRIUM, Y (At. Wt. 89), and ERBIUM,” Er (At. Wt. 166). Occurrence.—Yttrium is an important constituent of gadolinite, SiO, Be(Y,Ce,La,Di,Th, Er,0), Fe, and of yttrotantalite, (Nb,Ta)O.Y, an isomorphous mixture of yttrium tantalate and yttrium niobate. The two elements Y and Er are also found in cerite, thorite, and monazite. REACTIONS IN THE WET WAY. A solution of Y(NO.), and one of Er(NO), should be used. Yttrium. Erbium. 1. NH, OH and (NH,),S pre- Behaves like yttrium. cipitate the white hydroxide, insoluble in an excess. 2. KOH and NaOH precipi- Behaves like yttrium. tate the white hydroxide, insol- uble in an excess; the presence of tartaric acid does not prevent precipitation; but in this case: yttrium tartrate is precipitated (difference from Al, Be, Th, and Zr). On igniting the precipi- tate, the oxide is obtained, which is readily soluble in acids. 3. (NHA),CO, produces a Behaves like yttrium, ex- white precipitate of the carbo- cept that the Solution does not nate, readily soluble in an ex- become turbid on standing. cess of the reagent; after stand- ing some time the solution be- comes turbid, owing to the dep- osition of a double salt, Y2(CO3)3·2(NH4)2CO3 •2H2O. * Erbium is not an element itself, but consists of at least three elements— holium, thulium, and dysprosium. The separation of these elements is ex- ceedingly difficult; So we will consider simply the reactions of the mixture. YTTRIUM.–ERBIUM. 4I 7 Yttrium. 4. K.CO, and Na,CO, pre- cipitate the white carbonate, readily soluble in excess; after standing Some time an insoluble double Salt separates out. 5. BaCO, does not precipi- tate yttrium in the cold, and only incompletely on warming. 6. Oxalic acid precipitates white yttrium oxalate, insolu- ble in an excess, difficultly solu- ble in HCl, and perceptibly solu- ble in ammonium oxalate. 7. K,SO, forms a double salt which is soluble in K.SO, solution. (difference from Zr, Th, Ce, La and Di). 8. HF produces white amor- phous YF, which becomes pul- verulent on warming, and is in- soluble in water and in HF (difference from Al, Be, Ur, and Ti). Yttrium solutions do not give an absorption spectrum. Erbium. Behaves like yttrium, only the solution does not become turbid on standing. Erbium is not precipitated at all, even on warming. In erbium solutions, oxalic acid produces a light-red, pul— verulent precipitate; otherwise the reaction is the Same as with yttrium. Behaves like yttrium. Erbium solutions give a characteristic absorption spec- trum. REACTIONS IN THE DRY WAY. Yttrium oxide is strongly luminous on being heated; oth- erwise there is no reaction. Erbium oxide, on being heated on a platinum wire, colors the flame distinctly green. If the light is viewed through a spectroscope, a number of bright lines will be seen in the dark green which coincide with the dark lines obtained in the absorption spectrum. 418 REACTIONS OF SOME OF THE RARER METALS. THE CERITE METALS. CERIUM, Ce. At. Wt. 140.3. LANTHANUM, La. At. Wt. 138.9. PRASEODYMIUM, Pr. At. Wt. 140.5. NEODYMIUM, Nd. At. Wt. 143.6. Occurrence.—These three metals are important constituents of cerite, (SiO,) (Ce,Al), Ca,(OH), and of orthite (allanite), (SiO4)6CeCa,(OH), besides being usually found associated with the gadolinite earths in gadolinite, etc. DIDYMIUM: CERIUM. Cerium forms two oxides, Ce,O, and CeO2; both are basic an- hydrides, from which salts are derived. The cerous salts are white, while the ceric Salts are orange red. REACTIONS IN THE WET WAY. *- 1. Cerous Salts. A solution of cerous nitrate, Ce(NO), should be used. 1. NH, OH and (NH),S each produce a white precipitate of Ce(OH), insoluble in an excess of the reagent, but readily soluble in acids. In the presence of tartaric and citric acids, etc., the above reagents cause no precipitation (difference from Y). 2. NaOH or KOH also precipitate white Ce(OH), even in the presence of tartaric acid, etc. The white Ce(OH), becomes yellow gradually on standing in the air, on account of being oxidized to Ce(OH). 3. K.CO, and (NHA),CO, each produce a white precipitate insol- ſuble in an excess of the reagent. W 4. Oxalic Acid or Ammonium Oxalate precipitate white cerous oxalate, insoluble in an excess of the reagent, and in dilute mineral acids. On ignition, white, insoluble CeO, is formed. If the oxa- late is contaminated with praseodymium oxalate, a cinnamon- * Only when the cerous oxalate is pure. If it contains traces of prase- odymium, the CeO2 is obtained as a bright-yellow powder. CERIUM. 4 IQ colored oxide is obtained, which is completely soluble in dilute acids. º 5. BaGO, slowly precipitates all of the cerium in the cold. 6. KSO, produces a white, crystalline precipitate of the double salt Ce SO.), 3K,SO, difficultly soluble in cold water, but readily soluble in iſot water. It is insoluble in a saturated solution of KSO, (difference from Y and Er). ,2. Ceric Salts. A solution of either ceric nitrate, Ce(NO3)4, or of ceric ammonium nitrate, Ce(NO.), :2NH, NO,--H,O, should be used. The beautiful orange-red color of these solutions is characteristic of all ceric salts, as is also their tendency to form difficultly soluble basic salts. Preparation of Ceric Compounds.-AS has been already stated, cerous hydroxide on standing in the air gradually changes to yellow, on account of the formation of ceric hydroxide. This oxidation takes place immediately on the addition of chlorine or hypochlo- rites. If the solution of a cerous Salt is treated with caustic potash solution and chlorine is conducted into it, the white cerous hydroxide which was at first formed is quickly changed to light-yellow ceric hydroxide. The latter compound dissolves in dilute acids, forming orange solutions. It dissolves in concentrated hydrochloric acid with evolution of chlorine, forming cerous chloride. If white cerous hydroxide is heated in the air, it loses water and is changed into CeO, which is almost white when cold, dark orange when hot, and is almost entirely insoluble in concentrated hydrochloric and nitric acids. In the presence of reducing substances (such as KI, FeSO4, etc.) it dissolves in acids, forming cerous salts: 2CeO2+8HCl·H2KI=2KCl +4H2O+I2+2CeCls. CeO, also dissolves in concentrated sulphuric acid with evolution of oxygen and formation of cerous Sulphate. It can also be readily brought into Solution by fusing with potassium pyrosulphate and dissolving the melt in considerable hot water to which a little acid is added. If a mixture of cerous and praseodymium hydroxides is ignited in the air, a cinnamon-colored mass is obtained, which contains all 42O REACTIONS OF SOME OF THE RARER METALS. of the cerium as dioxide and is readily soluble in dilute acid's, form- ing ceric salts. If concentrated HCl is used, there is an evolution of chlorine, the ceric salt being reduced to cerous chloride. Con- centrated nitric acid dissolves it, forming cerous and ceric salts; a distinct evolution of oxygen can always be detected. 3. The reason why the brown mass containing a little praseody- mium dissolves although the pure oxide does not, is probably the following: CeO, like MnO, and PbO, (see pages 115 and 158), plays the part of an acid anhydride, so that the brown mass contains the praseodymium as the salt of ceric acid. On treating this salt with a stronger acid, the praseodymium salt of the latter is formed, set- ting free ceric acid (ceric hydroxide), which in the hydrated form is readily soluble in acids, forming ceric salts. Cerous salts can be oxidized in acid solutions by (a) Heating with PbO, and HNO, (1 part of acid to 2 parts of water). (b) Treatment with persulphuric acid. (c) Electrolysis. If a cerous Salt is treated with hydrogen peroxide in acid solution a yellow coloration is at first formed, which soon disappears with evolution of hydrogen. Ceric salts, therefore, are reduced by hydro- gen peroxide. If, however, a cerous salt is treated with H.O, and the solution is made barely alkaline with ammonia, a dark-orange precipitate is formed which looks something like ferric hydroxide, and probably consists of CeO,-- H.O. This is the most sensitive reaction for cerium and is especially suited for detecting the presence of cerium in Solutions of lanthanum and didymium salts. Basic Ceric Salts.—If a solution of ceric nitrate is evaporated on the water-bath to a consistency of syrup, the mass dissolves readily in water after it has become cold, and the solution can be boiled without becoming turbid. If, however, a little nitric acid is added, a yellow precipitate is immediately formed, consisting of basic ceric nitrate; on the addition of more acid the precipitate re- dissolves. This can be explained as follows: By treating the Solu- tion of ceric nitrate with considerable water it becomes hydrolyzed considerably, but the basic salt produced is present in the hydrosol state and is changed by the acid into the hydrogel form. As lanthanum and didymium salts do not yield basic salts under LANTHANUM. 42 I these conditions, this property can be used for separating cerium from these metals. It is characteristic of cerium to form with ammonium nitrate an easily crystallizable salt, ceric ammonium nitrate: Ce(NO3)4.2NH4NO3.H2O. All ceric salts may be readily reduced by the ordinary reducing agents (alcohol, HI, SO, H.S, HNO, H.O., etc.) to cerous salts. REACTIONS IN THE DRY WAY. The borax bead is colored dark brown when hot and light yellow to colorless when cold, after being heated in the oxidizing flame. In the reducing flame the bead becomes colorless, although strongly ignited CeO, will remain suspended in the bead, giving it a turbid yellowish appearance. W LANTHANUM, La. At. Wt. 138.9. Lanthanum forms only one oxide,” La,O, which, even after being strongly ignited, dissolves readily in acids. Its salts are colorless and yield no absorption spectrum, so that lanthanum may be distinguished in this way from didymium and erbium. REACTIONS IN THE WET WAY. A solution of lanthanum nitrate La(NO3)3 should be used. 1. NH, OH and (NHA),S precipitate a white basic salt which is difficult to filter. 2. KOH and NaOH precipitate the white hydroxide La(OH),. There is no change to be noticed on treating with oxidizing agents (difference from Ce).f La(OH), is soluble enough in water to turn red litmus-paper blue, and it decomposes ammonium salts on warm- ing with evolution of ammonia. The fused oxide is readily soluble in acids. 3. (NH,),CO, produces a white precipitate, insoluble in an excess of the reagent. 4. Oxalic Acid produces a white crystalline precipitate, insolu- ble in dilute mineral acids and in ammonium oxalate solution. * H2O, is said to cause the formation of La,Oo(?). f Ibid. * SUMMARY OF THE REACTIONS OF THE RARE EARTHS. (NH4)2COs Na2S2O3 White precipitate on ble in excess. White precipitate in- White precipitate in- White precipitate in- * * * * Al2O3 h. acid prevents soluble in excess. soluble in excess. soluble in excess. *::::::::::, almost neu- the precipitation. ------ ~ ––– * bl White precipitate solu- il W. Fº º: e in excess, reprecipi- * g gº & White precipitate very | . soluble in excess an BeO tated on boiling or dilut- White precipitate in- difficultly soluble in a reprecipitated on boil- No precipitate. 1Dig. . soluble in excess. 1Ing. Tartaric acid prevents large excess. the precipitation. * , White precipitate in- W h i t e precipitate * * * * * tº º ZrO soluble in excess. White precipitate in- | slightly soluble in excess, tºº.º.º. *. gº' With ...; 2 hº acid prevents soluble in excess. but #precipitated by reprecipit . ed on boiling. boiling p OD. the precipitation. 4\-W.M.L - s sº * * g. g White precipitate solu- Wºº precipitate in- Wºº precipitate in- ble in º ğ. re- l soluble in excess. soluble in excess. precipitate addition g Almost completely pre- T 110, Tartaric acid prevents It is precipitated be- of NH4OH. The solution Behaves as with K2CO3. cipitated on boiling. the precipitation. fore the cerite metals. becomes turbid on boil- ing, but clears on cooling. ** White precipitate in- * tº e 3- solº in *:::::: wh While precipitate #. " . artaric acid does not ite precipitate in- |uble in excess. er tº k Qº ſº. Y.Os * soluble in excess. je.ºu.”insoºj; Behaves as with K2CO3. No precipitate. |KOH or NaOH NH,OH or (NH4)2S or Na2CO3 + White precipitate solu- Fººt the precipitation, ut yttrium tartrate is then formed. double salt is formed. # * Er,0s Pink precipitate insol- uble in excess. Tartaric acid prevents the precipitation. * Pink precipitate insol- uble in excess. Tartaric acid prevents the precipitation. White precipitate solu- ble in excess. On boil- ing, all of the erbium is reprecipitated. Behaves as with KoCO3, except that the precipi- tate is more soluble in an eXCeSS. No precipitate. ~~ White precipitate in- soluble in excess, which is gradually colored yel- low on standing in the air. Tartaric acid prevents the precipitation. & Behaves as with KOFX. The precipitate becomes orange when treated with H2O2. NW hite, , , precipitate slightly solubie is excess. Behaves as with K2CO3. # CeO2 Yellow Ce(OH)4 is pre- cipitated, soluble in HCl with evolution of Cl and form.tion of cerous chlo- ride. After ignition it be- comes insoluble in acids. Rehaves as with KOH. Yellow precipitate al- most entirely insoluble in €XCeSS. Yellow precipitate sol- uble in considerable ex- cess, but reprecipitated on boiling. No precipitate. No precipitate. Di2O3 Violet precipitate in- soluble in excess. Tartaric acid prevents the precipitation. Behaves as with KOH. LazOs White precipitate in- soluble in excess, Tartaric acid prevents the precipitation. Behaves as with KOH. The precipitate produced from an acetic acid solu- tion with sqilute ama- monia is a basic salt, which becomes blue on treatment with iodine. White precipitate in- soluble in excess. Behaves as with K2CO3. No precipitate. .Violet precipitate of didymium carbonate, in- soluble in excess. Behaves as with K2COs. No precipitate. + Use a double normal solution: . concentrated solutions give precipitates, in nearly every case, soluble in excess of the reagent. f This reaction is used to separate aluminium from beryllium; the separation is not quantitative. s *. § SUMMARY OF THE REACTIONS OF THE RARE EARTHS.—(Continued.) º º Absorption Oxalic Acid (NH4)2C2O, HF PaCOs Spectrum K.SO, 2’ t sº- is - Al2O3 No precipitate. No precipitate. No precipitate. sºil. None. No precipitate. BeO No precipitate. No precipitate. No precipitate. sºil. Nono. No precipitate. T. S. L. A concentrated so- > lºſiº of K2SO4 slow- y precipitates zir- º • º - e. * * ſº conium potassium White Yº: Nº. º No precipitate in sulphate, soluble in precipitate soluble in soluble in excess, an th cold Incom- - ble dilut ZrO excess, difficultly sol- not reprecipitated by No precipitate. e • .-- d None. considerable dilute 2 uble in very dilute addition of consider- pletely precipitate HCl, unless, it was HC1. able HCl On Warmlng. preci.itated from a • 'Loll in solution, - when it is almost in- soluble in HCl. White precipitate soluble in excess on ‘2 e * - boiling, and not re- *ſ º Y. precipitate º * º: - Precipita t e s , a insoluble in excess, ing provided that in- double salt insoluble ThC), : in § mineral º: original solution No precipitate. tºº None. in a concentrated acids; soluble in am- | did not ecºaun too w K2SO4 solution. In Onium acetate. much free acid. The 2SO4 w solution is precipi- tated by considera- ble HCl g. White, amorphous v e s a tº White precipitate ſº - = recipitate insoluble No precipitation in w No precipitation. Y.O, ºliº.º.º.º. s º, º: É. HF, and insoluble the cold. Incom: None. The double s a lit is and difficultly solu- ble in dilute HC1. CeSS- in . , dilute , mineral acids after being ig- nited. pletely precipitated On Warmlng. soluble in K2SO4 so- lution. , * § * § Er,0s . Pink precipitate insoluble in , excess, difficultly soluble in dilute HC1. Pink precipitate on boiling, scarcely soluble at all. No precipitation. No precipitation in the cold, incom- pletely precipitated on warming. Characteristic. No precipitation; the double salt is sol- uble in K2SO4 solu- tion. Ce2Os White precipitate insoluble in oxalic acids and, in dilute mineral acids. . White precipitate insoluble in excess. No precipitation. Slowly but c.o.m - pletely precipitated in the cold. None. A double salt is precipitated, insolu- ble in K2SO4 solu- tion. CeO2 Brown ceric oxa- late is at first precip- itated, which on the addition of more ox- alic acid is changed to white cerous oxa- late with loss of CO2. Behaves as with oxalic acid. No precipitation. Completely pre- cipitated in the cold. None. A double salt is precipitated, insolu- ble in K2SO4 solu- tion. La2Os . White precipitate insoluble in excess. Same as with ox- alic acid. The pre- cipitate is insoluble in (NH4)2C2O4. No precipitation. Completely pre - cipitated in the cold. None. A double salt is precipitated, insolu- ble in R2SO4 solu- tion. Di,Os Violet precipitate insoluble in excess. The solid salt shows in reflected light the well-known absorp- tion spectrum. Behaves as with oxalic acid. No precipitation, Completely pre- cipitated in the cold. Characteristic. A double salt is precipitated, insolu- ble in KeSO4 solu- tion. NotE.-These tables were compiled with the help of those published by C. Glaser, Chem. Zeitung, 1896, page 612. ANALYSIS OF GADOLINITE (CERITE). Gadolinite contains SiO, Y.O., Er,0s, Ce,0s, Di,Os, La,O, Fe,0s, Fe0, BaO, MgO, Na2O, H.O, and often some of the metals of the H2S group. About 2 gms, of the finely-divided mineral are treated repeatedly with conc. HCl, with the addition of a little HNO3, and evaporated to dryness on the water-bath; the mineral is completely decomposed by this treatment. The dry mass is moist- ened with 2 c.c. of conc. HCl (the acid is allowed to act upon the mass for fifteen minutes), 100 c.c. of water are added, and the silicic acid filtered off. The metals will be found in the filtrate, which is now saturated with H2S, and any metals of the H.S group filtered off. This precipitate is examined according to Table IX, page 386. The filtrate is freed from H.S by boiling and then treated with a boiling solution of oxalic acid, added little by little, while the solution is constantly stirred. The precipitate is filtered off. PRECIPITATE (Ce, La, Di, Y, and Er as oxalates). SoLUTION (Be, Al, Fe, Ca, Mg, and Na). The precipitate is washed with water, dried, and ignited in a porcelain dish, whereby a cinnamon-colored mass, consisting of a mixture of the oxides of the rare metals, is obtained. . It is dissolved in as little HNO, as possible, a little alcohol is added (to reduce ceric salt to cerous), evaporated on the water-bath to dryness, dissolved in a little water, treated with an excess of solid K.SO, and allowed to stand over night. In the morning the precipi- tate is filtered and washed with K2SO, solution. PRECIPITATE (Ce, La, and Di as SoLUTION (Y and Er R2(SO,)3' 3K2SO4.) as R2(SO4)3' 3R.S.O.). The precipitate is dissolved in very dilute HCl, pre- cipitated with oxalic acid, washed, dried, and ignited. The oxides formed are then dissolved in a little HCl, carefully precipitated with NaOH (avoiding an ex- cess) and the solution is then saturated with chlorine and filtered. The solution is precipitated with oxalic acid, filtered, ignited, and the oxides dissolved in a very little HNO3, evaporated to dry- NH,OH + (NH,),S are added and the filtrate tested in the usual way for Mg and Na. The precipitate (consisting of Be(OH)2, Al(OH), FeS, and CaC.O.) is filtered off, washed, dried, gently ignited, and then di- gested with conc. HCl and 1 c.c. HNO3 on the water-bath until completely dissolved. The excess of acid is evaporated off, the residue dissolved in water, ammonia is added, and the precipitate is filtered off. . Only calcium now remains in the filtrate, which is precipi- tated as carbonate and confirmed as usual. The precipitate (consisting of Be(OH)2, Al(OH)4, and Fe(OH)2) is dissolved in a little HCl, evaporated almost to dryness, and then treated with an excess of (NH4)2CO3 solution, allowed to stand 15 minutes, and filtered. In case a precipitate of Fe(OH)3 and Al(OH)3%. is obtained, the filtrate is treated, with a few... -*. * … …- - § § PRECIPITATE (Ce(OH),). Solution (LaCl, + DiCl3). After being washed, the precipitate is dis- solved in HNO3, and the yellowish-orange solution is treated with H2O2, un- til decolorized, when a little dilute ammonia is added. A brown precip- itate shows the presence of Ce. It is acidified with HCl, boiled, precipitated with o x a lic a cid, filtered, washed, dried, and ignited. The oxides are dissolved in as little HNO3 as pos- sible and examined in the spectroscope. The pres– ence of Di is shown by the formation of dark absorp- tion bands. To test for La, the solu- tion is treated with am- monium acetate, precipi- tated with very dilute ammonia, filtered through asbestos, and washed three times with H2O. The pre- cipitate ... is then treated with solid iodine, and if La is present a blue color will be noticed in a few minutes. ness, dissolved in water and the solu- tion examined with the spectroscope; black absorption bands show the presence of Er. The solution is treated with HF, and the formation of a white, amor- phous precipitate insoluble in F shows the presence of Y. drops of (NHA).S. (to remove the last traces of iron) until no further precipitation of FeS is formed. This is filtered off, and the filtrate is boiled for some time and filtered. PRECIPITATE. SoLUTION. Basic beryllium carbo- nate. It is dissolved in HCl, an excess of KOH is added, whereby the precipitate at first formed redissolves. The solution is diluted with water and boiled. A white precipitate shows the presence of Be. *Iron and aluminium are tested for in the usual way. Rejected. 428 REACTIONS OF SOME OF THE RARER METALS. 5. K,SO, precipitates white, crystalline La,(SO), .3K,SO, in- soluble in a concentrated K.SO, solution. º 6. Lanthanum Sulphate is only soluble in ice-cold water; on , warming the saturated solution to 30° the salt separates out thickly (difference from cerium). 7. Iodine.—If ammonia is added to a cold, dilute acetic acid solution of a lanthanum salt, and the slimy precipitate is washed with water and then treated with solid iodine, the whole mass grad- ually assumes a blue color which is similar to that produced by the action of iodine upon starch (this property is peculiar to lanthanum). The blue color is destroyed by the addition of acids or alkalies. NEoDYMIUM, Nd. At. Wt. 143.6. DIDYMIUM łº, Pr. At. Wt. 140.5. It is very difficult to separate these two metals from one another. It is accomplished only by repeated fractional crystallization of the ammonium double nitrate. Neodymium apparently forms only one oxide, Nd,0s; it appears bluish after being ignited, and is readily soluble in acids, forming violet salts, which afford a characteristic absorption spectrum. Praseodymium, on the other hand, forms a greenish-white oxide, Pr,O, which on being ignited is changed into dark-brown peroxide, Pr/O. On being heated in a stream of hydrogen, the latter is re- duced back to Pr,0,. The peroxide dissolves in acids with loss of oxygen, forming greenish salts corresponding to the lower oxide and yielding a characteristic absorption spectrum. The Didymium Reactions take place with a mixture of the two elements. A solution of didymium nitrate, Di(NO), is used. Didymium salts are violet and show a characteristic absorption spectrum (difference from Ce and La). The behavior toward NH,0H, (NH,),S, KOH, (NH,),CO, and K.SO, is exactly the same as with lanthanum. Oxalic acid precipitates the reddish oxalate, which in other respects is like lanthanum oxalate. TANTALUM, Ta. (At. Wt. 183), and NIOBIUM, Nb (At. Wt. 94). These two elements belong to the nitrogen group. Both tantalic and niobic acids, however, dissolve in strong acids under certain conditions and are precipitated then by ammonia and ammonium sulphide. | TANTALUM. 429 Occurrence.—The chief minerals are tantalite, Fe(Taſ)a)2, nio- e. Ol' columbite, Fe(NbO3)2, and yttrotantalite, an isomorphous lxture of (Ta2O7)3 Y4 and (Nb2O7)3 Y4. Tantalum, and to some extent niobium also, replaces the phosphorus in monazite. Cassit- erite and wolframite usually carry small amounts of both niobic and tantalic acids. TANTALUM, Ta. At. Wt. 183. Metallic tantalum * is ductile, although the presence of a little impurity makes it harder than tool steel. Its specific gravity is 16.5 and melting point 2250° C. On ignition in the air it assumes a yellow to blue tinge caused by a thin coating of oxide. Tanta- lum is not attacked by boiling H2SO4, HCl, HNO3, or even aqua regia, but it is slowly dissolved by hydrofluoric acid with evolution of hydrogen; any metal remaining undissolved is then brittle on account of absorbed hydrogen. The concentrated solution forms with concentrated KOH insoluble, crystalline potassium tantalum fluoride, K2(TaF7). By evaporating the solution in HF with concentrated H2SO4 until the former acid is all expelled, the residue dissolves in a little cold water, but the solution becomes turbid on dilution or especially by boiling. Tantalum forms two oxides, Ta2O4 and Ta2O5, the former being indifferent chemically and the latter an acid anhydride. After ignition the pentoxide is insoluble in acid and is not rendered soluble by fusion with pyrosulphate although it is volatilized by heating with ammonium fluoride. Fusion of the oxide with caustic alkali in a silver crucible gives rise to alkali tantalates, both the meta- and hexatantalates being known; only the former are soluble in water. Potassium hexatantalate, KSTag019 + 16H2O, is soluble in water and caustic potash solution, while the sodium altº in water, but not in caustic soda. There aimi. insoluble. " - *E**** ºne ºr way. A solution of tassium hexatantalate loud be used. 1. Mineral Aids-(a) H2SO4 precipitates tantalic acid from cold dilute solutiºns, and the precipitation becomes almost quantita- tive only on boºing. Hot concentrated H2SO4 dissolves the precipi- tate produced º, the dilute acid. On diluting the solution with water * W. ºn Bolton, z, Electrokem ki (1905), Mºſt 6, p. 45. 43o REACTIONS OF SOME OF THE RARER METALS. after it has become cold, the tantalic acid is reprecipitated (differ- ence from niobium). (b) HCl at first produces a precipitate from a concentrated solu- tion, which dissolves in an excess of the acid, forming an opalºcent solution. From this solution sulphuric acid precipitates tantalic acid in the cold, but the precipitation is not quantitative even on boiling. (c) HNO, has the same action as HCl, 2. NH, OH and (NHA),S precipitate from the hydrochloric acid solution either tantalic acid itself or an acid ammonium tantalate; tartaric acid prevents the precipitation. 3. Tannic Acid produces in acid solutions a light-brown precipi- tate (difference from niobic acid). 4. K. Fe(CN), produces in acid solutions a light-yellow precipi- tate, which becomes brown on the addition of a little ammonia. 5, HF, KF.—If a concentrated solution of tantalic acid in hydrofluoric acid is treated with KF, the difficultly soluble K.Taf, is formed, which separates from the solution in the form of ortho- rhombic needles (200 parts of water dissolve 1 part of salt) (differ- ence from niobium). On boiling the solution of tantalic potas- sium fluoride, the very difficultly soluble oxyfluoride precipitates (KTa,0s Fla.). By means of this reaction the merest trace of tantalic acid can be detected in the presence of niobic acid. 6. Zn and HCl do not produce colored solutions (difference from niobium). REACTIONS IN THE DRY WAY. Ta,0s is infusible. The bead of salt of phosphorus remains colorless in both oxidizing and reducing flames. The addition of FeSO, does not cause the formation of a blood red color (difference from Ti and Nb). ... Niobium forms three oxide Nº. 0. Nºb, , and Nb,0s, of which the last is an acid anhydride. Nb,0. Ta,0s, is insoluble in acids after it has been ignited and is not rendered soluble by fusing with potassium pyrosulphate. By fusing with KOH or K.CO, potassium hexaniobate is formed, which is soluble in water. The sodium salt is insoluble in caustic soda, but soluble in water. - - THALLIUM. 43I REACTIONS IN THE WET WAY. A solution of potassium hexaniobate should be used. 1. Mineral Acids produce in alkali niobate solutions a white, amorphous precipitate of niobic acid, which is only slightly soluble in an excess of the acid. Concentrated sulphuric acid, however, dissolves the niobic acid on warming, and the solution remains clear after being diluted with cold water (difference from tantalum). If the niobic acid is treated with boiling hydrochloric acid, it only dissolves slightly, but on pouring off the acid, the residue is soluble in water.” 2. NH, OH and, (NHA),S precipitate niobic acid from the sul- phuric acid solution, and the precipitate is soluble in HF. 3. Tannic Acid Tincture produces an Orange-red precipitate. 4. K. Fe(CN)6 produces a light-yellow precipitate. 5. H(KF2).-If niobic acid is dissolved in an excess of HF and KF is then added, readily soluble niobic potassium fluoride is formed (12.5 parts of water dissolve 1 part of the salt). By boiling the dilute aqueous Solution, Soluble potassium niobic oxyfluoride is formed, which is even more soluble (difference from tantalum). 6. Zinc produces a beautiful blue coloration when added to an acid solution of niobic acid (difference from tantalum). REACTIONS IN THE DRY WAY. The bead of salt of phosphorus is blue, violet, or brown in the reducing flame (according to the amount of niobic acid which is present): the bead becomes red on the addition of FeSO,. METALS OF THE H.S GROUP. THALLIUM, VANADIUM, MOLYBDENUM, TUNGSTEN, SELENIUM, TELLURIUM, RHODIUM, PALLADIUM, OSMIUM, IRIDIUM, RUTHENIUM. THALLIUM, Tl. At. Wt. 204.1. Occurrence.—Thallium is found in nature very sparingly; in small amount in many varieties of pyrite, and accompanying potas- * This behavior reminds º of metastannic acid. (Cf. page 218.) 432 REACTIONS OF SOME OF THE RARER METALS. sium in carnallite and Sylvite, in many lithium micas and in many mineral waters. It replaces the silver to a considerable extent in copper-silver selenide, in crookesite, (AgIlCu),Se, and in ber- zelianite, (CuAgTl),Se. There are no characteristic thallium mine- rals. The principal sources of our thallium is the dust from sul- phuric acid plants where pyrite containing thallium is used. Metallic thallium reminds one of lead in its color, softness, high specific gravity, and low melting-point. Lead. Thallium. Specific gravity . . . . . . 11.24 to 11. 388 11.8 Melting-point......... 326°C. 290° C. It is a monovalent element, dissolves readily in nitric and sul- phuric acids, but not in hydrochloric acid. It forms two oxides: thallous oxide, Tl,O and thallic oxide, Tl,O. Both oxides are anhydrides of bases and give rise to thallous and thallic salts. REACTIONS IN THE WET WAY. A. Thallous Compounds. Thallous compounds are colorless and soluble in water as a rule. The sulphide, chloride, bromide, iodide, and chromate are insoluble in water. Thallous oxide is a colorless powder, whose aqueous solution reacts alkaline and absorbs carbon dioxide with avidity. A solution of thallous sulphate should be used for the following reactions: 1. H.S causes no precipitation from solutions which contain mineral acids; in neutral solutions, thallium is incompletely pre- cipitated as black thallous Sulphide, Tl,S. T.S is readily soluble in mineral acids, but insoluble in acetic acid and alkaline sulphides. It is oxidized readily on standing in the air to thallous sulphate. 2. (NH),S precipitates all of the thallium as Tl,S. 3. KOH, NaOH, or NH, OH produce no precipitation. 4. Alkali Carbonates cause precipitation only in very con- centrated solutions, for thallous carbonate is fairly soluble (100 parts of water dissolve 5 parts of the salt). 5. HCl produces a heavy, white precipitate of thallous chloride, very slightly soluble in water, and still less so in water containing a little hydrochloric acid. | THALLIUM. - 433 6. KI precipitates yellow thallous iodide, TII, from even the most dilute solutions; this is the most sensitive reaction for thal- lium. 7. Alkali Chromates precipitate yellow thallous chromate, in- Soluble in cold nitric or sulphuric acids. 8. H.PtC1, precipitates light-yellow thallium chlorplatinate, which is almost entirely insoluble in water; 1 part dissolves in 15,585 parts of water at 15° C. and in 1948 parts of water at 100° C. 9. Al,(SOI), If a solution of thallous sulphate is treated with aluminium sulphate and the solution is then allowed to crystallize, glistening, colorless octahedrons are obtained of thallium alum, TIAl(SO,),4-12H,O. Thallium is like lead in respect to its specific gravity and to the solubility of its halogen compounds; but, on the other hand, it is similar to the alkalies with regard to the solubility and alkaline reaction of the hydroxide and carbonate, and with regard to its forming an insoluble chlorplatinate and an alum. B. Thallic Compounds. Thallic compounds cannot as a rule be prepared by the oxida- tion of thallous compounds (with the exception of thallic chloride, which is readily obtained by the action of chlorine water upon thallous chloride). They are obtained by the solution of thallic oxide * in acids, and can be distinguished from thallous compounds by the ease with which they suffer decomposition in aqueous solu- tion. Thus thallic sulphate is decomposed on boiling its aqueous solution into thallic hydroxide and sulphuric acid; the nitrate be- haves similarly. The chloride, TICl, is a hygroscopic and not very stable sub- stance; on being heated to 100° C. chlorine is evolved with the formation of thallous chloride. e KOH, NaOH, and NH, OH precipitate brown thallic hydroxide, Tl(OH), from solutions of thallic salts, which changes to TIO(OH) on standing in the air, is difficultly soluble in acids and insoluble in an excess of alkali. * Tl2O3 is not attacked in the cold by concentrated sulphuric acid, but is dissolved on warming. The hydroxide, TIO(OH), is much more soluble. 434 REACTIONS OF SOME OF THE RARER METALS. KCl and alkali chromates do not cause precipitation. KI precipitates thallous iodide with deposition of iodine. REACTIONS IN THE DRY WAY. Thallium salts color the non-luminous gas-flame a beautiful emerald green. The thallium spectrum consists of a green line. WANADIUM, W. At. Wt. 51.2. Occurrence.—Vanadinite, Pb,(VO,),Cl; carnotite; * mottramite, (CuIPb);V.O.o-H2H,O; many clays and in almost all granites. Vanadium forms, like nitrogen, five oxides: V.O, V,0, V.O., V.O., W.O. | V & \!. The Ét three are basic anhydrides, while the last two oxides possess completely the character of acid anhydrides. V.O., is the anhydride of hypovanadic acid, W.O.(OH),. It is a blue powder, soluble in concentrated acids, forming blue divanadyl salts: V.O,--2H,SO, = V2O,(SO.), +2H,O. If the solution of divanadyl Sulphate is treated with sodium car- bonate or ammonia (avoiding excess), hypovanadic acid separates out as a grayish-white precipitate, which, like the anhydride, is soluble in acids with a blue color and in alkalies with a brown color. The divanadyl compounds are readily formed by reducing solutions of the pentoxide in mineral acids with sulphurous acid (cf. page 435), and serve, on account of their blue color, for the detection of vanadium. V.O., is the anhydride of vanadic acid and is an orange-red crys- talline mass, which is readily fusible byºff non-volatile. It is orly slightly soluble in water, forming a yºlow slightly-acid solution, but readily soluble in concentrated 3%lutions of caustic alkaliès, forming vanadates. , ſº Like phosphoric acid, vanadic ačd exists in the form of meta-, pyro-, ortho-, and poly-compounds, of which the meta-compounds * According to Friedel and Cumenge, carnotite contains 18 per cent. V.O. and 55 per cent. VOs, as well as (K, Ca, Ba, H, As, and P). (Cf. Hille- brand and Ransome, Am. Jr. of Science, X, 138.) VANADIUM. " 435 {< are the most stable and the ortho-compounds the least so. Thus an aqueous Solution of potassium or sodium orthovanadate is hydro- yzed, even in the cold, into the pyro-salt and alkali hydroxide, 2Na,VO,--H,Oze Na,V.O.--2NaOH, and on boiling the meta-salt is formed: Na,V.O.--H,Oze 2NaVO,--2NaOH. The meta-, pyro-, and ortho-salts of the alkalies are colorless or slightly yellow, while the polyvanadates, e.g., the tetra- and hexa- vanadates, are intensely orange or reddish. Thus the colorless or light-yellow solutions of the Ortho-, meta-, and pyrovanadates are colored intensely orange on the addition of acid. * REACTIONS IN THE WET WAY. 1. NHC1.—If a piece of solid ammonium chloride is added to a solution of an alkali vanadate, colorless ammonium metavanadate eparates out, Na,V.0,4-4NH,Cl=2NH.VO,--2NH,--H,O+4NaCl, ifficultly soluble in a concentrated Solution of ammonium chloride. 2. (NH),S produces no precipitation, but causes the solution to turn brown owing to the formation of sulpho-salts, from which brown W.S., is precipitated on the addition of acid; * the precipitate is soluble in alkalies, alkali carbonates, and in alkali Sulphides, forming a brown solution. '3. H.S gives no precipitation in acid solution, but reduces com- pounds of vanadic acid to divanadyl compounds, so that the Solu- tion is colored blue. 4. Reducing Agents (SO, H.S., HBr, alcohol, oxalic and tartaric acids, sugar, etc.) reduce acid solutions containing vanadates to blue vanadyl Salts: V.O.--SO,-SO,--V.O. HI reduces vanadic acid to green V.O. : V.O.--4H[=2H,O-i-4I+V.O. *The precipitation of V.Ss is not quantitative; the filtrate is always col- ored blue and contains detectable amounts of vanadyl salts. * ~ * * * 436 REACTIONS OF SOME OF THE RARER METALS. The green color only appears after the iodine has been removed by continued boiling of the solution. Metals, such as Zn, Al, and Có, cause still further reduction of vanadic acid, so that the Solution turns at first blue, then green, and finally violet. 5. H.O,-If an acid solution of a vanadate is treated with a few drops of H.O, and shaken, the solution becomes colored red- dish brown. This is a very delicate reaction. 6. Mercurous Nitrate precipitates white mercurous vanadate from neutral solutions; the precipitate is soluble in nitric acid. 7. Lead Acetate precipitates yellow lead vanadate, Soluble in nitric acid. Detection of Vanadium in Rocks (Hillebrand).* Five gms. of the finely-powdered rock are fused over the blast: lamp with 20 gms. of Na,CO, and 3 gms. of NaNO,. The product of the fusion is extracted with water, the manganate formed is re- duced by the addition of a little alcohol, and the solution is filtered. The aqueous solution can contain As, P, Mo, Cr, W, W. It is nearly neutralized with nitric acid (the amount necessary having been determined by a blank test), almost evaporated to dryness, taken up in water, and filtered. The alkaline solution is then treated with mercurous nitrate, whereby mercurous phosphate, arseniate, chro- mate, molybdate, and tungstate with some basic mercurous carbo- nate are precipitated. The Solution is boiled and filtered, the pre- cipitate dried and separated from the filter, ignited in a platinum crucible, and then fused with a little sodium carbonate. The prod- uct of the fusion is extracted with water, when a yellow color shows that chromium is present. The solution is acidified with sulphuric acid, and traces of Pt, Mo, and As are precipitated by means of H.S (best in a small suction flask). This precipitate is filtered off and the excess of H.S is removed by boiling, while a stream of carbonic acid gas is being passed through it, evaporated to dryness, and the excess of sulphuric acid removed by carefully heating in an air-bath. The residue is now dissolved in 2 or 3 c.c. of water and shaken with a few drops of H.O.; a brownish-yellow color shows the presence of vanadium. If chromium is present, on adding H2O2 and ether to the sulphuric acid solution and shaking, the ether will be colored blue by chromium and the aqueous solution yellow by vanadium.f * Amer. Journ. of Science, 1898, p. 209. t E Champapne, Chem. Centralbl., 1904, II, p. 1167. MOLYBDENUM. 437 REACTIONS IN THE DRY WAY. The borax bead is colorless in the oxidizing flame if slightly săturated with the vanadium compound, yellow if strongly satu- rated, and green in the reducing flame. MOLYBDENUM, Mo. At. Wt. 96.0. Occurrence.—Molybdenite, MoS,; wulfenite, PbMoQ,; powel- lite, CaMoQ. Molybdenum has a valence of 2, 3, 4, and 6, and forms the following oxides: MoC), Mo.O, MoC), and MoQ. The first three are basic anhydrides, while the last oxide, MoQ, is an acid anhydride, forming a white mass (yellow when warm) which is readily fusible, but very difficultly volatile. In the cooler part of he crucible, colorless, transparent, thin, orthorhombic plates of oO, are deposited. MoQ, is only very slightly soluble in water, but issolves readily in alkalies and in ammonia, forming molybdates. olybdic acid itself can be readily obtained as a solid mass by acidi- ying the solution of an alkali molybdate; it is soluble in an excess f the acid (difference from tungstic acid). The most important 3ommercial molybdate is the acid ammonium molybdate, corre- ponding to the formula (NH,), Mo.O,--4H,O. REACTIONS IN THE WET WAY. A solution of ammonium molybdate should be used. The alkali molybdates are soluble in water; the remaining salts 'e mostly insoluble in water but soluble in acids. 1. Dilute Acids precipitate from concentrated alkali molybdate solutions white H.MoO, soluble in an excess of acid. Concentrated Sulphuric Acid.—If a trace of a molybdenum. compound is evaporated with a drop of concentrated sulphuric acid almost to dryness in a porcelain dish, the mass is colored intensely blue. This is an exceedingly delicate reaction. 2. H.S at first colors acid molybdenum solutions blue, and pre- cipitates little by little the molybdenum as brown molybdenum tri- sulphide, MoS, soluble in ammonium sulphide, forming a brown solution and reprecipitated by the addition of acids. Molybdenum 4.38 REACTIONS OF SOME OF THE RARER METALS. *3. sulphide is oxidized by warm sulphuric acid to white molybdic acid, and is changed by roasting in the air into MoQ. 3. Zinc.—If a molybdate Solution which is acid with hydro- chloric or sulphuric acid is treated with zinc, the solution is colored at first blue, then green, and finally brown. Other reducing agents such as SnCl, Hg,(NO), etc., cause the same reaction. 4. KCNS causes no change when added to acid molybdenum solutions, but if the solution is then treated with zinc or stannous chloride, a blood-red coloration is produced on account of the forma- tion of molybdenum Sulphocyanide; the reaction also takes place in the presence of phosphoric acid (difference from iron). If the solu- tion is shaken with ether, the colored compound is dissolved in the latter. 5. Sodium Phosphate.—If a few drops of a solution of sodium phosphate are added to a molybdate Solution strongly acid with nitric acid, a yellow crystalline precipitate of ammonium phospho- molybdate is formed, slowly in the cold, but much more quickly on warming the solution (cf. Phosphoric Acid, page 328). Arsenic acid causes the precipitation of a similar compound. 6. Mercurous Nitrate precipitates white mercurous molybdate from neutral solutions; the precipitate is soluble in nitric acid. 7. Lead Acetate precipitates white lead molybdate, soluble in nitric acid. REACTIONS IN THE DEY WAY. Alkali molybdates, alone or with sodium carbonate, are reduced on charcoal to gray molybdenum, a white incrustation of MoQ, being formed at the same time. Borax Bead.—All molybdenum compounds impart a yellow color to the warm bead, after it has been heated in the oxidizing flame, becoming colorless when cold; in the reducing flame a dark- brown colored bead is produced, which is turbid if considerable MoC), was used. TUNGSTEN, W. At. Wt. 184. Occurrence.—Tungsten is not very often found in nature, but there are a number of well-crystallizing tungsten minerals, such as the minerals of the Scheelite group, ** TUNGSTEN. 439 Scheelite, CaWO4; cuproscheelite, (CaCu)WO,; reinite, FeWO,; stolzite, Pb WO. These minerals all crystallize in the tetragonal system and form with powellite, CaMoC), and wulfenite, PbMoO, a very interesting isomorphous group. Another isomorphous group, which consists of minerals crystallizing in the monoclinic system, is formed by hiibnerite, MnWO,; wolframite, (MnFe)WO, and fer- berite, FeWO. The most important tungsten mineral is wolfram- ite, which is usually contaminated with small amounts of silicic, tantalic, and niobic acids. Tungsten forms two oxides, WO, and WO, WO, is a brown powder, readily obtained by heating WO, to a dull redness in a stream of hydrogen. It is pyrophoric and must, therefore, be cooled in a stream of hydrogen before it is allowed to come into contact with the air. By igniting strongly in a stream of hydrogen, metallic tungsten is obtained, which is stable in the air. This behavior is important and is taken advantage of in the quantitative determination of tungsten. WO, is an acid anhydride obtained by the ignition of tungstic acid, of ammonium or mercurous tungstates, or by the oxidation of the dioxide, on heating in the air. The trioxide is a canary-yellow powder, insoluble in water and dilute acids, and only slightly soluble in concentrated hydrochloric and hydrofluoric acids. It dissolves readily by warming with po- tassium or sodium hydroxides, and less readily in ammonia. It is most easily dissolved by fusing with Sodium carbonate, sodium tungstate being formed: WO,--Na,CO,-Na, WO,--CO,. It is changed to potassium tungstate by fusing with potassium pyrosulphate: WO,--K.S.O. =K.WO,4-2SO,. If the product of this last fusion is treated with water, usually none of the tungsten goes into Solution, because if an excess of potassium pyrosulphate is present (which is usually the case) it reacts with the potassium tungstate, forming free tungstic acid: K.WO,--K.S.O,4-H,0=2K,SO,--H.WO,. 44O.” REACTIONS OF SOME OF THE RARER METALS. If not enough pyrosulphate remains to complete the above de- composition, some of the tungsten will be dissolved, but never all of it. If a little sulphuric acid is added to the water, none of the tung- Sten will go into Solution. This property enables one to separate tungsten from titanium. If ammonium carbonate were added, all of the tungsten dissolves, which enables us to separate tungstic from silicic acid. REACTIONS IN THE WET WAY. A Solution of sodium tungstate should be used. 1. Mineral Acids, HCl, HNO, H.SO, produce, in the cold, a white, amorphous precipitate of hydrated tungstic acid, H.WO,--H,O.* By boiling the solution, the yellow anhydrous acid H.WO, is obtained, insoluble in dilute acids, but soluble to an appreciable extent in concentrated hydrochloric acid. Tungstic acid must always be washed with water which contains acid or a dissolved Salt, as otherwise tungstic acid will form a pseudo- solution with pure water, so that a turbid filtrate will be obtained (cf. pages 70, 174). Phosphoric acid behaves differently toward solutions of the alkali tungstates than do the other mineral acids; it produces a white precipitate soluble in an excess of phosphoric acid; a complex phospho-tungstic acid is formed, e.g., OP(WO.),(ONa). If the solution of an alkali tungstate is boiled with free tungstic acid, the latter gradually goes into Solution, forming a metatungstate: Na,WO,--3WO,-Na,W.O. Mineral acids cause no precipitation in Solutions of metatung- states. If the solution is boiled with an excess of acid, the soluble metatungstic acid is gradually changed to insoluble, ordinary tung- stic acid, which is then precipitated. w 2. H.S produces no precipitation in acid solutions. 3. (NH,),S gives no precipitation in a solution of an alkali tungstate, but if the solution is afterward acidified, light-brown tungsten trisulphide is precipitated, which has the property of form- ing pseudo-solutions with pure water, but is insoluble in hydrochloric acid. The precipitate redissolves in ammonium sulphide. * The presence of tartaric acid prevents the precipitation. SELENIUM. 44 I. 4. Reducing Agents.-If the solution of an alkali tungstate is treated with HCl and zinc, the tungstic acid at first precipitated by the HCl is soon turned to a beautiful blue color, owing to the forma- tion of W.O. SnCl, produces a yellow coloration at first, but on adding HCl and warming, a beautiful blue precipitate is obtained. This is one of the most sensitive reactions for tungstic acid. - 5. Mercurous Nitrate precipitates white mercurous tungstate from neutral Solutions. 6. Lead Acetate precipitates white lead tungstate from neutral Solutions. REACTIONS IN THE DRY WAY. The salt of phosphorus bead is colorless in the oxidizing flame, and blue in the reducing flame, becoming blood red on the addition of a little FeSO4. SELENIUM, Se.. At. Wt. 79.2. Occurrence.—Although selenium is quite widely distributed in nature, it is invariably found in very Small amounts, usually re- placing sulphur, forming isomorphous compounds with lead, silver, copper, and mercury; claustalite, PbSe; berzelianite, (CuAgTl),Se; naumannite, (Ag, Pb)Se; tiemannite, HgSe; lehrbachite, (Pb, Hg)Se; onofrite, Hg(SeS); eucairite, (Ag,Cu),Se. It is also found in small amounts in many varieties of pyrite and chalcopyrite, and indeed the small amounts which are found in these minerals form the chief source of the Selenium of commerce. By roasting these minerals (as in the manufacture of Sulphuric acid) all of the Selenium is volatilized, and is consequently deposited in the lead chambers as a mud from which it is extracted with a Solution of potassium cyanide and afterwards precipitated with acid: KCN-I-Se-KCNSe and KCNSe--HCl–HCN-H KCl-H Sc. Selenium, like sulphur, exists in two allotropic forms. The modification soluble in carbon disulphide is obtained by reducing selenious acid in the cold with Sulphurous acid; it is a brick-red powder. After heating this red Selenium with hot water for some 442 REACTIONS OF SOME OF THE RARER METALS. time it is changed into black Selenium, and is then insoluble in car- bon disulphide. Selenium melts at 200° C. On heating in the air, Selenium burns with a bluish flame (giving off an odor similar to that of rotten radishes) forming white crys- talline selenium dioxide, SeO, which will sublime on being heated in a stream of oxygen. Selenium forms one oxide, SeO, and two acids: Selenious acid, H.SeO, and Selenic acid, H.SeO,. Selenious acid, H. SeO, is obtained in the form of long colorless needles by oxidizing Selenium with nitric acid, or by the solution of its anhydride, SeO, in water. Unlike sulphurous acid, it is not changed on standing in the air into Selenic acid; but, on the con- trary, is reduced by dust, etc., to red Selenium. The acid is dibasic, and forms salts in which either one or both of the hydrogen atoms are replaced by metal. . The acid salts are all soluble in water, but the neutral salts are all insoluble with the exception of those of the alkalies. Selenic acid, H.SeO4, is obtained in solution by conducting chlo- rine into water which contains either suspended Selenium or dis- solved selenious acid: Se-H3Cl,--4H,O = H.SeO,--6HC1. Sodium seleniate is obtained by fusing Selenium with sodium carbonate and potassium nitrate. Selenic acid is a dibasic acid, and behaves similar to a peroxide, evolving chlorine when boiled with concentrated hydrochloric acid, being reduced to selenious acid: * H.SeO,--2HCl=H.O.--H,SeO,--Cl,. " REACTIONS IN THE WET WAY, (a) Selenious Acid. A solution of either potassium Selenite or of free selenious acid should be used. 1. H.S produces a lemon-yellow precipitate, consisting of sele- nium and sulphur, from Solutions in water or in dilute hydrochloric acid: H.SeO,--2H.S=3H,0+Se+S, ~~~~ The precipitate is soluble in ammonium sulphide. SELENIUM. 443 2. Reducing Agents (SO2, SnCl2, FeSO4, Zn, Fe, HI, NH2. NH2-H2SO4, NH2OH.HCl) reduce the solution of selenious acid in dilute hydrochloric acid, causing the precipitation of red selenium, which changes black after the Solution has been heated for some time. H3PO3 precipitates selenium from hot concen- trated solutions, but does not from a dilute solution slightly acid with HCl. 3. Barium Chloride precipitates from neutral solutions white barium selenite, BaSeO, soluble in dilute acids. 4. CuSO, produces a greenish-blue, crystalline precipitate (difference from Selenic acid). (b) Selenic Acid. A solution of potassium seleniate should be used. 1. H.S causes no precipitation unless the solution is boiled with hydrochloric acid. In the latter case the selenic acid is first reduced to selenious acid, and is then reduced to selenium by the hydrogen sulphide. 2. BaCl, gives a white precipitate of barium seleniate, BaSeO, insoluble in water and in dilute acids; Soluble, with evolution of chlorine, on being boiled with hydrochloric acid: Base0,--2HCl=BaCl, -H,SeO,--Cl,--H,0. 3. CuSO, produces no precipitation. 4. SO, does not reduce selenic acid. Method for Testing Sulphuric Acid for Selenium.* Five or six drops of the acid to be tested are added to the solution of a little codein in Sulphuric acid, when, if Selenium is present, a green coloration will be apparent. The test is a very delicate one. REACTIONS IN THE DRY WAY. All selenium compounds emit the odor of decayed radishes on being mixed with sodium carbonate and heated on charcoal before the blowpipe. If a selenium compound is heated at the end of a thread of asbes- tos in the upper reducing flame of the Bunsen burner, it will be re- * Dragendorff, Chem. Centralbl., 1900, 944. 444 REACTIONS OF SOME OF THE RARER METALS. duced to selenium; and if a test-tube filled with water is held above the flame, a red coating of selenium will be deposited upon the glass.” If a few drops of concentrated Sulphuric acid f are placed in a larger test-tube (large enough to hold the Smaller test-tube) and the tube on which the Selenium is deposited is emptied and placed within the larger tube, the selenium will dissolve : in the sulphuric acid, form- ing a green solution; but on the addition of water, red selenium will be reprecipitated (difference from tellurium): SeSO,--H,0=Se-i-H.S.O. Green TELLURIUM, Te. At. Wt. 127.6. Occurrence.—Tellurium is a rarer metal than selenium, always occurring in the form of a telluride, and usually combined with the noble metals: calaverite, (Au,Ag)Te2; krennerite, (Au,Ag)Te2; Sylvanite, (Au,Ag)Tea; nagyagite, AugSb2Pb10TecS15; coloradoite, HgTe; silver telluride, Ag3Te; and often in small amounts in galena and copper ores. Emmonsite of Cripple Creek, Colorado, is a ferric telluride with 70.71 per cent. Te02 and 22.76 per cent. Fe2O3. Tellurium itself is a bluish-white, brittle substance, which melts at 500° C. and can be distilled in a stream of hydrogen. It burns in the air with a bluish-green flame, forming tellurium dioxide, Te02. It is insoluble in carbon disulphide, and can be oxidized by means of nitric acid to tellurous acid. On being fused with potassium cyanide, out of contact with the air, it is changed to potassium telluride, 2KCN +Te=K2Te + (CN)2, which dissolves in water, forming a cherry-red solution. If air is conducted through this solution, the tellurium is precipitated in the form of a black powder (difference from selenium): K2Te + H2O+O=2|KOH +Te. Tellurium may be separated from Selenium by means of this last reaction. The two metals are fused with potassium cyanide, the melt is treated with water, and the tellurium precipitated by passing * Cf. page 27. t The sulphuric acid should be freed from water by heating in a platinum : crucible to a temperature just below the boiling-point, and the crucible with its contents allowed to cool in a desiccator. f Slowly in the cold, readily on warming. TELLURIUM. 445 a current of air through the solution; the selenium is precipitated from the filtrate by acidifying with hydrochloric acid. Tellurium forms two oxides: TeO, and Te0. Tellurium dioxide (the anhydride of tellurous acid) is usually obtained in the form of a white mass, which melts on gentle heating, forming a yellow liquid. Tellurium dioxide does not sublime (dif- ference from selenium). It is scarcely soluble at all in water, is slightly soluble in ammonia and in dilute acids, but readily Soluble in concentrated acids or in caustic potash Solutions. TeO, dissolves in fairly concentrated sulphuric acid, forming the basic Sulphate, Te,0,...SO, while with nitric acid it forms the basic nitrate, Te2O3(OH)NO3. Both of these compounds are hydrolyzed readily, forming insoluble tellurous acid; the latter in turn loses water and forms the anhydride. On dissolving Te02 in caustic potash, potassium tellurite, I(2TeO3, is obtained. Only the alkali tellurites are soluble in water. Tellurium trioxide (telluric anhydride) is formed by heating telluric acid. It is a yellow powder, insoluble in water and nitric acids, scarcely affected by boiling with concentrated hydrochloric acid, but is readily dissolved by boiling with a concentrated solution of potassium hydroxide (but not by sodium hydroxide), forming potassium tellurate. Telluric acid, H.TeC),4-2H,0, is a very weak acid, obtained by oxidizing tellurous acid with chromic acid, and precipitating the telluric acid by the addition of concentrated nitric acid. The acid forms a colorless crystalline mass, is readily soluble in water, and is converted by means of concentrated hydrochloric acid into tellurous acid, with evolution of chlorine. The acid dissolves readily in caustic potash (or Soda) Solution, forming the readily soluble alkali tellurate, which reacts strongly alkaline in aqueous solution. By gently heating the hydrated telluric acid, the anhydrous acid, H.TeC), is obtained in the form of a white powder and is totally different from the hydrated acid. The latter is soluble in water and in caustic alkalies, and is completely reduced by boiling with concen- trated hydrochloric acid; but the anhydrous acid is insoluble in water and in concentrated sodium hydroxide solution, and is only very slightly attacked by boiling concentrated hydrochloric acid, although readily Soluble in warm potassium hydroxide solution. Only the tellurates of the alkalies are soluble in water; the others 446 REACTIONS OF SOME OF THE RARER METALS. are usually obtained in the form of amorphous precipitates soluble in acids. REACTIONS IN THE WET WAY. (a) Tellurows Acid. A solution of potassium tellurite, K.Te0, should be used. 1. H.S precipitates brown TeS, from acid solutions, which is readily soluble in ammonium sulphide. 2. Reducing Agents.-SO, precipitates tellurium completely from dilute hydrochloric acid solutions in the form of a black pow- der, even in the presence of tartaric acid; but from a strongly-acid Solution no tellurium is precipitated even on boiling (difference from Selenium). In the latter case about 200 c.c. of hydrochloric acid (sp. gr. 1.175 are used) (cf. Vol. 2). SnCl, or Zn causes black tellurium to precipitate from solutions which are not too acid. H.PO, precipitates the tellurium only from concentrated solu- tions, not at all from cold dilute solutions. FeSO, reduces neither tellurous nor telluric acids (difference from Selenium). 3. HC1 produces a white precipitate of H.Te0,. (b) Telluric Acid. A solution of potassium tellurate should be used. 1. HCl causes no precipitation; but if the solution is then boiled chlorine is evolved, and on dilution with water tellurous acid is precipitated. 2. H.S and reducing agents have the same effect upon hot solu- tions of tellurates as upon tellurites. 3. Lead Salts precipitate difficultly-soluble lead tellurate. REACTIONS IN THE DEY WAY. Metallic tellurium is formed by heating any telluride in the upper reducing flame, and can be collected on the lower surface of a test-tube filled with water in the form of a black film, soluble in concentrated Sulphuric acid. The latter Solution is of a carmine red color (difference from Selenium); and on the addition of water black tellurium is deposited: TeSO,--H.O=Te-H H.SO,. Carmine red PALLADIUM. 447 THE PLATINUM METALS. PLATINUM, PALLADIUM, RHODIUM, OSMIUM, RUTHENIUM AND IRIDIUM. Platinum has been already described on page 232. PALLADIUM, Pd. At. Wt. 106.5. Sp. Gr. 11.9. Occurrence.—The platinum metals form an isodimorphous group, but only in the case of palladium are both forms known—the regu- lar and the hexagonal: (a) Regular System. (b) Hexagonal System. Platinum (Pt, Fe). Iridosmium (Sysserskit) (Ir, Os). Iridium (Ir, Pt). OSmiridium (Newjanskit) (Ir, Os, Platinum iridium (Pt, Ir, Rh). Pt, Rh, Ru) or (Os, Ir, Rh). Palladium (Pd, Pt, Ir). Palladium (Pd, Pt, Ir). Properties.—Rolled, hammered, or cast palladium possesses an almost silver-white color, but when precipitated from Solutions it is in the form of a black powder. If it is suspended in water when in the finely-divided form, it is transparent with a reddish color. Palladium has the lowest melting-point of all the platinum metals, viz., 1500° C. On being heated in the air, it appears bluish, owing to the formation of Pd,0; the latter, however, is decomposed by stronger heating. Behavior towards acids: While the remaining platinum metals are attacked by no acid except aqua regia, palladium is dissolved slowly by warm nitric acid (also in the cold when it is alloyed with other metals such as CuAg, etc.), forming a brown solution of Pd(NO), Finely-divided, precipitated palladium is soluble in hydrochloric acid when exposed to the action of air at the same time, and less readily soluble in sulphuric acid. It is readily attacked by fusing with potassium pyrosulphate, forming soluble palladium sulphate, PdSO4. The best solvent for palladium is aqua regia. Finely-divided palladium has the very characteristic property 448 REACTIONS OF SOME OF THE RARER METALS. of being able to adsorb almost 700 times its own volume of hydro- gen, and possesses consequently a very strong catalytic action. If hydrogen and oxygen (air) are conducted at the same time over Some gently-ignited, finely-divided, metallic palladium, the hydro- gen is burnt to water, and in the same way carbon monoxide may be changed to carbon dioxide. Methane, however, is only decom- posed by igniting the palladium more strongly, so that this gives us a method for separating methane from a mixture of H and CO (cf. Vol. 2, Gas Analysis). Palladium forms two oxides, both of which possess strongly basic properties: PdC) and PdC). From the former the palladous, and from the latter the palladic, compounds are derived. The palladous compounds are much more stable than the palladic compounds, and the latter constantly exhibit the tendency to change into the former. $ By dissolving finely-divided palladium in hydrochloric acid, palladous chloride is obtained; or, better, by dissolving the metal in aqua regia, in which case a mixture of palladous and palladic chlorides is at first obtained. If this solution, however, is evapo- rated to dryness, palladic chloride loses chlorine and is completely changed into palladous chloride, so that on treating the residue with water a solution of palladous chloride is obtained. Since palladic chloride is decomposed completely by evaporation, it is evident that palladic chloride cannot exist in hot solutions. ^. REACTIONS IN THE WET WAY. (a) Palladous Compounds. A solution of palladous chloride, PdCl, should be used. 1. H.S precipitates black palladous sulphide from acid and neu- tral solutions. The precipitate is insoluble in ammonium sulphide, PdS, but soluble in boiling hydrochloric acid, or more readily in aqua regia. 2. KOH or NaOH precipitates a brown basic salt, soluble in an excess of the reagent. If the solution is acidified with HCl, KOH produces no precipitate (difference from platinum). 3. Na,CO, produces a brown precipitate of palladous hydroxide, soluble in excess but reprecipitated on boiling. PALLADIUM. 449 4. NH4OH gives a flesh-colored precipitate of [Pd(NH3)2Cl2].” soluble in an excess of ammonia, forming a colorless Solution (con- faining palladodiamine chloride, Pd(NH3)4Cl2), from which yellow crystalline palladosamine chloride, Pd(NH3)2Cl2, is precipitated on the addition of hydrochloric acid. The latter compound is diffi- cultly soluble in dilute hydrochloric acid and is used for the prepara- tion of pure palladium. In a solution of palladous nitrate, ammonia causes no precipita- tion, but forms colorless palladodiamine nitrate, Pd(NH3)4(NO3)2. 5. NH,C1 forms a soluble complex salt: [PdCl,](NH,),. 6. KCI, when added to a concentrated solution, causes the pre- cipitation of difficultly-soluble, reddish-brown PdCl,K, (octahe- drons). 7. HI or KI produces a black precipitate of palladous iodide, even in very dilute solutions. The precipitate is insoluble in water, alcohol, ether, and HI, but soluble in KI and NH. (This and the following reaction are characteristic for palladium.) 8. Hg(CN), produces a yellowish-white gelatinous precipitate of palladous cyanide, Pd(CN), difficultly soluble in HCl, readily soluble in KCN and NHA. On being ignited, the spongy metal remains. 9. Reducing Agents.-H..SO, formic acid (HCOOH), Zn, Fe, TeSO, Cu,Cl,f alcohol, and CO Í reduce palladium salts to the metal itself. In the presence of HCl, stannous chloride forms at first a red, then a brown, and finally a green, Solution; but in the absence of the acid, SnCl, causes a partial reduction to the metal, while the solution turns green. * This compound is an isomer of palladosamine chloride, and is often written thus: PdCl2, Pd(NH3)4C12. f In the presence of considerable NaCl or HCl there is no reduction with Cu2Cl2. † PdCl, + CO + H2O = 2HCl + CO2 + Pd. This reaction enables us to detect small amounts of CO in gas mixtures; e.g., in the air. For this purpose the gas is led by means of a small drawn-out glass tube into 10 c.c. of a solution which contains 1 mg. of PdCl2 and 2 drops of dilute HCl. If CO is present, black På will be deposited, and the solution will become decolorized little by little. (Potain and Drouin, Compt, rend., 126, 938.) 45O REACTIONS OF SOME OF THE RARER METALS. (b) Palladic Compounds. These give the same reactions as palladous compounds, on ac- count of their being readily changed into the latter. The principal difference, however, is the insolubility of the ammonium salt of hydrochlorpalladic acid. If a concentrated, cold solution of palla- dous chloride is shaken with chlorine water and then treated with ammonium chloride, a red crystalline precipitate of [PdCl](NH.), is soon formed. REACTIONS IN THE DRY WAY. All palladium compounds are decomposed on ignition, leaving behind the metal, which is soluble in nitric acid or in aqua regia, and the solution thus obtained can be tested by the above reactions. RHODIUM, Rh. At. Wt. 103.0. Sp. Gr. = 12.6. Properties.—Rhodium possesses the color and lustre of alumin- ium; it is more infusible than platinum, its melting-point lying above 2000° C.; on cooling the hot metal it sputters and appears bluish, owing to oxidation. The solubility of rhodium depends en- tirely upon the fineness of the material. When precipitated from a solution of its chloride by means of formic acid or other reducing agents at a temperature not exceeding 100° C., it exists in an extremely finely-divided state (rhodium- black) and dissolves readily in boiling concentrated Sulphuric acid, or more readily in aqua regia. If, however, the finely-divided metal is ignited strongly, it becomes (like the compact metal) almost in- soluble in aqua regia. If rhodium is alloyed with other metals (Pb, Zn, Bi, Cu, etc.), it is left in a finely-divided condition after treatment of the alloy with acids, and is consequently soluble in aqua regia. When it is alloyed with much platinum or palladium a considerable amount of rho- dium will dissolve in aqua regia; but when it is alloyed with a little platinum, most of the rhodium and a part of the platinum remain undissolved. On being fused with potassium pyrosulphate, potassium rhodium RHODIUM. 45 I sulphate is formed, which dissolves in water, forming a yellow solu- tion, but becomes red on the addition of HCl. Rhodium forms three oxides: RhC), Rh,0, and RhC,; all pos- sess a well-defined basic nature. The sesquioxide, Rh,0, alone * forms a series of salts, of which sodium-rhodium chloride is the most important for the analytical chemist; for when in this form it is easiest to bring rhodium into solution. This Salt is prepared by mixing the finely-divided metal very intimately with twice as much dry sodium chloride, placing it in a porcelain boat and gently ignit- ing it in a current of moist chlorine gas. The salt thus formed has the composition [RhCls]Nas, and is soluble in water (45 parts of water dissolve 1 part of the salt). From this solution large, dark-red, glistening triclinic prisms of [RhCl]Naa-H 9H,0 can be crystallized out. REACTIONS IN THE WET WAY. A solution of sodium rhodium chloride should be used. 1. H.S precipitates (very slowly in the cold, but much more quickly on warming) black rhodium sulphide, Rh,S, ; insoluble in (NH,),S, soluble in boiling nitric acid. 2. KOH and NaOH produce at first no precipitate; but after standing some time a yellow precipitate of rhodium hydroxide, Rh(OH,)+H,O, separates out. The precipitate is soluble in an excess of the reagent, but is reprecipitated on boiling in the form of brownish-black Rh(OH)2. If a solution of potassium rhodicSulphate were used, KOH pre- cipitates the yellow compound immediately. * On adding KOH to a solution of rhodium chloride, at first no precipitate is produced; but on the addition of a little alcohol brownish-black rhodium hydroxide is deposited. 3. NH4OH produces (in concentrated solutions and after stand- ing some time) a yellow precipitate of chlorpurpureorhodium chlo- ride, Rh(NHA):Cl, insoluble in hydrochloric acid. 4. KNO, on being warmed with sodium rhodium chloride solu- tion, causes the precipitation of difficultly—Soluble, orange-yellow Rh(NO.), Ks, soluble in HCl. * A sodium-rhodium sulphite of the formula 4Rh(SOs),6Na,SO, + 9H,O has been prepared by Bunsen. 452 REACTIONS OF SOME OF THE RARER METALS. * 5. Reducing Agents.-Formic acid in the presence of ammo- nium acetate precipitates the black metal, as does zinc in the presence of acids. REACTIONS IN THE DRY WAY. All rhodium compounds are reduced to metal on being heated in a stream of hydrogen, or by heating on charcoal with sodium car- bonate before the blowpipe. The metal is easily recognized by its insolubility in aqua regia, its being brought into solution by fusing with potassium pyrosulphate and then treating with water, and by the formation of the brown hydroxide when KOH and a little alcohol are added to the solution thus obtained. OSMIUM, Os. At. Wt. 191. Sp. Gr. = 21.3–22.477. Osmium and ruthenium are distinguished from the other plati- num metals by their forming volatile oxides. Properties.—The compact metal possesses a bluish-white color, very similar to zinc, and is the heaviest of all metals. It can be melted by heating in an electric furnace.* Very finely divided osmium is oxidized by the air i at ordinary temperatures, and at about 400° C. it ignites and burns rapidly to OsO4, which is volatile at 100° C. The denser the metal the higher the temperature necessary to effect the oxidation. Behavior towards Acids.-In the compact condition osmium is insoluble in all acids; but in the finely-divided state (as obtained by treating its zinc alloy with nitric acid) it is soluble in nitric acid, more soluble in aqua regia, and most soluble in fuming nitric acid, forming osmium tetroxide; the latter can be separated from the solution by distillation. Compact osmium is brought into solution by fusing with NaOH and either KNO, or KClO. The melt contains a salt of perosmic acid (OsO4). * F. Mylius and R. Dietz, B. B., 1898, 3, 187. † Cf. Ot. Sulc, Zeit. für anorgan. Chemie, XIX, 332. OSMIUM. 453 Osmium forms five oxides: OsO, OS,Oa, OSO y Osmious oxide, grayish-black, Osmium sesquioxide, black, Osmic oxide, Élack-gray, insoluble in acids insoluble in acids insoluble in acids [OsO.), OsO4. Osmic acid, known only in the Perosmic acid, colorless needles form of its derivatives soluble in water Osmium tetroxide, OsO4, (the anhydride of perosmic acid,) is the most important osmium compound in the eyes of the analytical chemist, and is obtained by the oxidation of the substance in the air by dissolving the finely-divided metal in fuming nitric acid or in aqua régia, or by fusing with NaOH and KNO, or KClO, treat- ing the melt with nitric acid and distilling. Osmium tetroxide is a colorless, crystalline mass which sublimes at comparatively a low temperature and melts, forming colorless vapors at 100° C. The vapor has a chlorine-like odor, attacks the mucous membrane, and is poisonous. The chlorides of osmium can only be obtained in the dry way; OsCl, OsCl, OsCl, are known. The potassium salt of the hypo- thetical hydrochlorosmic acid, H.OsCls, forms dark-red octahedrons, soluble in water and decomposed by boiling the Solution. By heating finely-divided osmium with KCl in a current of chlorine, K.OsCl is formed; it dissolves in cold water, forming a red solu- tion. REACTIONS IN THE WET WAY. A solution of K.OsCl should be used. 1. If a solution of osmium chloride is treated with dilute nitric acid, the mixture distilled from a small retort, and the vapors re- ceived in caustic soda solution, the latter will be colored yellow, owing to the formation of potassium osmiate. If this solution is now acidified, osmium tetroxide is set free, and can be recognized by its very penetrating odor. On adding a little Sodium thiosul- phate to the acid solution and warming, a brown precipitate of os- mium sulphide is formed. 2. H.S precipitates brownish-black osmium sulphide, insoluble in ammonium sulphide. 3. KOH, NH, OH or K.CO, precipitate reddish-brown osmium hydroxide. 454 REACTIONS OF SOME OF THE RARER METALS. 4. Reducing Agents.-If the solution of the chloride is treated with tannic acid or alcohol and a little hydrochloric acid is added, it is colored dark-blue, owing to the formation of osmium dichloride, OsCl,; KI colors the solution a deep reddish-purple. Indigo is decolorized by solutions containing OsO4. Ferrous sul- phate precipitates black OSmium dioxide; stannous chloride pro- duces a brown precipitate soluble in HCl, forming a brown solution. REACTIONS IN THE DRY WAY. All osmium compounds are reduced to metal on being heated in a stream of hydrogen. RUTHENIUM, Ru. At. Wt. 101.7. Sp. Gr. = 12.261, crystallized; 11.0, fused. Properties.—Ruthenium exists in the form of a dark gray or black powder, and in the form of bright porous sticks; it is brittle, can be powdered, and is melted in the oxyhydrogen flame. On being melted, a part of the ruthenium is oxidized to ruthe- nium tetroxide, a volatile Substance having a penetrating odor simi- lar to that of OsO4. The molten metal spurts on cooling. Behavior towards Acids.—Ruthenium is almost completely in- soluble in all acids, even aqua regia. By fusing with KOH and RNO, (or KClO,) it is oxidized to potassium rutheniate, K.RuO,. On heating with NaCl in a current of chlorine, soluble K.RuCl, is formed. The solution in water of the greenish-black melt is of an orange-yellow color, and colors the human skin black. Ruthenium is unaffected by fusing with potassium pyrosulphate. It forms the following oxides: RuO, Ru,0, RuO, [RuO,], [Ru,O.], RuO,. The most important of the oxides is RuO4. It is formed: (a) By roasting the metal itself, or its oxide, above 1000° C. (osmium formed the volatile tetroxide at 400°C.). (b) By fusing the metal with KOH and KNO, in a silver cruci- ble, dissolving the melt in water, Saturating the cold Solution with chlorine gas, and distilling the solution from a small retort: KRuO,--Cl-KCl-i-RuO. RUTHENIUM. 455 treating the solution of potassium-ruthenium chloride and Cl, and subsequently distilling. By distilling a dilute solution after the addition of nitric acid no RuO, will be evolved (difference from osmium). Ruthenium tetroxide forms gold-yellow, glistening orthorhom- bic needles that are very volatile and emit a characteristic odor; it boils at 100° C. and is only slightly soluble in water. It is changed by the addition of alcohol and HCl into ruthenium trichloride, RuOl, (or sesquichloride, Ru,Cl). If the solution of the latter salt is made ammoniacal, treated with sodium thiosulphate and warmed, an intense reddish-violet coloration will be produced. (This is a very sensitive and characteristic reaction.) On treating a solution of potassium rutheniate with nitric acid, black Ru(OH), is precipitated; it dissolves in hydrochloric acid, forming a yellow solution of Rucla. REACTIONS IN THE WET WAY. A solution of Ruci, should be used. 1. H.S produces no precipitation at first, but after some time the solution becomes azure-blue, and brown ruthenium sulphide is pre- cipitated (very sensitive and characteristic). 2. (NH,),S precipitates the brownish-black sulphide, difficultly soluble in an excess of the reagent. 3. KOH and NaOH precipitate black ruthenium hydroxide, Ru(OH), soluble in acids but insoluble in alkalies. 4. KCNS, in the absence of other platinum metals, produces gradually a red, then purple, and on warming a violet, coloration (very characteristic). 5. KNO, imparts an orange-yellow color to the solution, owing to the formation of KaRu(NO.), becoming a beautiful dark red on the addition of a little colorless ammonium sulphide; on adding more ammonium sulphide, brown ruthenium sulphide is precip- itated. 6. Zinc at first colors the solution of the chloride azure blue, but subsequently the solution is decolorized and ruthenium itself is precipitated. 456 REACTIONS OF SOME OF THE RARER Merak { 7. Hydroxylamine reduces ruthenium tetrachloride bººtle. nium trichloride (difference from platinum). Ss. IRIDIUM, Ir. At. Wt. 193. Sp. Gr. 22.4. Properties.—When produced by the ignition of iridium ammo- nium chloride, it is obtained in the form of a gray, spongy mass, very difficultly soluble in aqua regia. After being strongly ignited, it is almost completely insoluble in aqua regia. It is more soluble in aqua regia when it is precipitated from solu- tions in a very finely-divided form by means of formic acid, or when it is alloyed with other metals (Au, Ag). The metal is unaffected by fusing with potassium pyrosulphate (difference from rhodium). It is oxidized by fusing with NaOH and KNO, in a silver crucible, but the compound formed (Ir,O, combined with sodium) is only partly soluble in water. If the melt is treated with aqua regia, however, a dark-red solution of Na,IrCls will be obtained. By heating the metal with NaCl in a current of chlorine, Na,IrCl, iš readily obtained. Iridium forms the following oxides: 3& IrO, Ir,Os, IrO, and the hydroxide Ir(OH), Black Blaii, black Needles with a Indigo-blue metallic lustre powder The dark-green color of the chlorides is very characteristic: IrCl, and IrCls, IrCl. B Green lack REACTIONS IN THE WET WAY. A solution of Na,IrCl, should be used. k 1. H.S at first decplorizes the solution, owing to the reduction of the tetrachloride to the trichloride, accompanied by the deposition of sulphur; Subsequently brown Ir,S, is precipitated, readily soluble in (NHA),S. 2. (NH,),S precipitates the same compound. 3. NaOH, on being added to the Solution, changes the color from dark red to green; on warming the Solution it is at first colored reddish and finally azure blue: 2IrCl,--2NaOH=2IrCl,--NaCl-i-H,0+NaOCl. # 457 KIrCl, formed is readily sºluble in water and in KCl solution (dif- ference from platinum). 4. KCl precipitateſ, brownish-black potassium iridium chloride, K.IrClº, insoluble º and ºn-alcohol, difficultly soluble in water. 5. NHC1 precipitates dark-red ammonium iridium chloride, (NH4), IrCl3, in uple in a saturated solution of NH,Cl. 6. Reducing Agents usually change the solution to a greenish color, owing tº the reduction of the tetrachloride to trichloride; or the solutiºn º and the black, finely-divided metal is de l, and then some KCI, there º ºe no precipitation, because the 3K,Ir(NO), K.IrCls. | Oxalic acid, ferrous sulphate, stannous chloride, and hydroxyl- 'mine reduce the tetrachloride to trichloride. Zinc reduces it to al, and so does formic acid on warming in the presence of am- }nium acetate. If considerable mineral acid is present, the re- ction takes place less readily. REACTIONS IN THE DRY WAY. On being fused with soda, the gray, brittle metal can be ob- |ained by the action of the upper reducing flame. It is insoluble in Separation of the Platinum Metals. The separation of the platinum metals is one of the most difficult tasks met with in analytical chemistry. If the metals are already in solution, the following table can be used to advantage. If, how- * The alcohol reduces NaOCl to NaCl. The six platinum metals, in the presence of gold and mercury, are assumed to be present in solution as chordes. The 3. solution is placed in a small retort, treated with dilute HNO, heated to boiling, and the vapors conducted into NaOH solution. The solution remaining in the retort contains the remaining metals. - After cooling a few cubic centimeters of ether are added. The Tiquid shaken, and the ether layer removed by means of a separ Ammonium acetate and fºrmic acid are added and The sºlution boiled for several hours, being connected, with a The solution is then filtered and the precipitate washed with ammonium acetate and water. | I with NH3Cl. NH4OH, a little Na2S2O3. A black pre- || and warming. An inten- c i p i ta t e lºsive violet coloration shows the shows Ru – presence of | _- r - preci ed by H. when di- lute, shows the presence of Rh. the presence of Pa. ANALYSIS OF THE PLATINUM METALS.* The solution is repeatedly shaken with ether until the ether extract appears colorless. Aqueous Solution (contains the remaining metals). Distillate. If the NaOH becomes yel- low, Os may be present. In this case. the odor of OsO4 will be detected on acidifying with HCl; and if a little Na2S2O3 is then added and the so- lution warm- ed, a brown precipitate of osmium sulphide will formed showing the presence of Os. ratory funnel. ETHEREAL Solution. If the ethe- real solution is yellow, return-flow condenser. PRECIPITATE (Pt. Pd, Ir, Rh, Ru, Hg). Solution. AuClais pro- bably pres- ent.t The et her is evaporated off, the resi- due dissolvi ed in a little water, and the resulting solution tested with SnCl2: the formation of The dried precipitate is heated in a porcelain boat in a stream of hydrogen, which causes any Hg to be volatilized, and is subsequently condensed on the colder portions of the tube in the form of a gray coating. The remaining metals after becoming cold again are boiled with hot HCl, in order to remºve any Sn or other foreign metal which may have been precipitated with this group on the addition of fºr- mic acid. The dried residue is then mixed with common salt and heated to dull redness in a moist chlo- Fine stream (the mass must not melt). After cooling, the residue, is treated with waterf and ammo- nium chloride added to the solution so long as a precipitate is formed, which is filtered off PRECIPITATE [. The annline lowed to cool. g (NH4)2PtCls. (NH4)2TrCl3, (NH4)2Rucial. recipitate is dissolved in hot water, hydroxyl- ydrochloride is added, and the solution al- NH2Cl is again added, and the precip- itate filtered off. the purple of Cassius shows the presence of Aw. PRECIPITATE. Solution (IrCl3(NH4)2Rucle(NH4)3. (NH4)2PtCl6. This is evaporated to dryness, ignited A yellow || and fused with KOH and KNO3 in a sil- precipitate || ver crucible. The melt is dissolved in consisting of re r octa- he dºr on s shows the presence of water and the bluish-black residue fil- tered off. SolºtroN (PdCl4Maz. Rhôlanas). The solution is evaporated slowly to dryness, with an excess of NH4OH, and the residue allowed to crystallize from a little warm, dilute NH4OH (insoluble, dark flocks are filtered off; they usually consist chiefly of Ru). . The Rh crys 5 as chloropurpureo chloride in yellow crystals. wºeſ are filtered off. CRYstals: Rh(NH3)3Cls Moth ER Liceo º * Mylius and Dietz, B. B. :898. 3187. Residue. Solution. (Ir(OH)3). (K2Ru();). It is mixed It is saturated in the with NaCl || cold with Cl. distilled and heated | from a retort. and the vapors received in alco- hol that has been acidi- fied with HCl. yel- lowish-brown precipitate shows the presence of Ru. It is confirmed by adding in a current of chlorine. The product is dissolved in water and precipitated the product dis- This contains palladious di- amine chloride, Pd(NH3)2Cl2. It is saturated with conc. HCl, when crystalline Pd(NH3)2Cl2 will be precipitated. This is filtered-off, dried in a stream of H. ignited, and the residue dysolved in conc. HNQa. Af- *ér cautiously evaporating the solved in water. A 'brown solution and taking up rose-colored sºu- || the residue in water, the solu- These are solu- ble in hot concen- trated H2SO4 with a yellow color. A small portion is mixed with NaCl. and heated in a stream of C1, and tion, which is not |tion is treated with Hg(CN)2. - - yellow precipitate shows ----- They recommend this method for the qualitative examination only. f Chromic acid wiiſ impart a yellow or blue color to the ether and FeCls a yellow color. f_If a residue remains after extracting with water, it is filtered off and the process repeated; if this does not suffice, it is examined for Ir + Ru after fusing with KOH and KNO3. : º C) 7. s º O 5. - : F- º tº S. S St. tº: > s § s S. t- º IRIDIUM. 459 ever, the metals are present in a more compact form, it is quite difficult to bring them into solution. In the latter case the metal is handled in as finely-divided con- dition as possible (filings, etc.) and treated with aqua regia; this serves to dissolve the greater part of the platinum and palladium, as well as Small amounts of rhodium and iridium. The residue (osmium, ruthenium, rhodium, iridium, and small amounts of plati- num and palladium) is dried, placed in a porcelain crucible and fused for some time with ten times as much zinc (or lead) in a cur- rent of illuminating-gas.” In this way the platinum metals are alloyed with the zinc. The mass is allowed to cool in the stream of illuminating gas, and is then treated with hydrochloric acid to dis- Solve out the zinc.; this leaves the platinum metals behind in a finely- divided state. They are filtered off from the acid solution, dried, introduced into a porcelain boat, placed in a tube made of diffi- cultly fusible glass, and heated to dark redness in a stream of Oxygen. The greater part of the osmium escapes as osmium tetroxide; it is absorbed by the caustic soda solution and tested according to the table. The residue is intimately mixed with sodium chloride and heated in a moist stream of chlorine. The mass is then dissolved in water and examined according to the table. If lead were used instead of zinc in the above procedure, the alloy should be treated with dilute nitric acid, which dissolves the lead and the greater part of the palladium. The lead is precipi- tated with the calculated amount of Sulphuric acid, and the filtrate is tested for palladium by transforming it into palladosamine chlo- ride, and then into palladium cyanide. The residue from the nitric acid treatment is treated in the same way as when zinc was used.| *The operation can be very conveniently performed in a common clay pipe. The gas is conducted through the stem and ignited at the bowl. In this way the gas continually streams through the molten alloy, keeps it well stirred, and thereby yields a uniform alloy. † For an accurate separation of the platinum metals, consult the work of Sainte-Claire-Deville, Debray, and Stas: “Procés verbaux du comité internat. des poids et mesures,” 1877–1878 and 1879. * ** * INDEX. A. Absorption Spectrum, 68 Acetates, Reactions of, 296 Acetic Acid, 295 , , Acetone, 39, 296 Acid Reaction, 394, 403 Acids, Dissociation of, 14 Acids, Division into Groups, 243 Acids, Examination for, 394 Acids, Suitable Concentrations, 31 Acids and Bases, Detection of, 194 Acids and Bases, Pseudo, 194 Agate, 355 Albite, 40 Alcohol, Solubility of Salts in, 57 Alkalies, Concentrations of, 31 Alkalies, Reactions of, 36,410 Alkalies, Separation from Magne- sium, 51 Alkalies, Separation from Group III, 142 Alkaline Earths, 53 Alkaline Earths, Separation from Alkalies, 51 Alkaline Earths, Separation from Group III, 142 Alkaline Earths, Sulphites, 303 Alkaline Earths, Sulphites and Thio- sulphates, 337 Alkaline Reaction of Solutions, 394, 400 Alkaline Solutions, Note Concern- ing, 392 Alloys, Analysis of, 399 Aluminium, 69 Aluminium, Detection in the Pres- ence of Organic Substances, 74 Aluminium, Separation from Fe, Cr, and U, 112 Alunite, 69 Ammonia, see Ammonium Ammonia in Drinking Water, 46 Ammonium, Reactions of, 44 Ammonium Carbamate, 55 * * 46I Ammonium Chloride, Use of, 16, 50 Ammonium Molybdate, Reagent, 189 Ammonium Salts, Ignition of, 48 Ammonium Sulphide Group, 69,412 Ammonium Sulphide Group, Analy- sis of, 144, 145 Amygdaline, 269 Analysis, Course of, 364 Anatase, 108 Anglesite, 54, 156 Å. 54 Anions, 12 Anions, Examination for, 380, 394 Anions, Preliminary Examination for, 369 Anions, Reactions of, 243 Annabergite, 126, 132 Antimonic Compounds, 205 Antimony, 199 Antimony, Separation from Other Metals of H2S Group, 226, 386 Antimony Trioxide, 200 Apatite, 54, 180 Aqua Regia, 247 Aragonite, 54 Argentite, 237 Arrhenius, Ion Theory, 11 Arsenic, 180 Arsenic, Detection in Wall Papers, 195 Arsenic Acid, Reactions, 187 Arsenic Pentoxide, 186 Arsenic Trioxide, 181 Arsenious Acid, Reactions, 183 Arsenolite, 180 Arsenopyrite, 180 Arsine (Arseniuretted Hydrogen), 191 Asbestos, 49 Atacamite, 168 Augite, 49 Autenrieth and Windaus, Solubility of the Sulphites and Thiosulphates of the Alkaline Earths, 337 Azurite, 168 462 INDEX. B. Barite, 59 Barium, 59 Barium, Separation from Calcium and Strontium, 62, 391 Barium Chloride Test, 397 Barium Sulphate, Decomposition of, Bases, see Alkalies and Anions Bases, Dissociation of, 14 Bead Tests, 23, 366 Benzaldehyde, 269 Beryl, 412 Peryllium, 412 Berzelianite, 432, 441 Berzelius-Marsh Test, 190 Bismite, 163 Bismuth, Reactions, 163 Bismuth, Separation from Other Metals of H2S Group, 179, 226, 386 Bismuthenite, 163 Bismuth-sodium. Thiosulphate Re- action, 39 Bismutite, 163 Blockmann, Strength of Reagents, 30 Blowpipe Reactions, 28, 368 Bog Ore, 88 Boracic Acid, see Boric Acid Borates, Solubility of, 309 Borax, 40, 309 Borax Beads, 23, 312 Boric Acid, 309 Braunite, 113 Breithauptite, 126 Bromine (free), Reactions, 261 Bromine, Detection in Non-electro- lytes, 261 Bromine, Detection in Presence of Cl and I, 268 * Brookite, 108 Brucine Reaction, 289, 341 Brucine, Reagent, 341 Bunsenite, 126 C. Cacodyl Oxide, 296 Cadmium, Reactions, 175 Cadmium, Separation from Other Metals of H2S Group, 179, 226, 386 Cadmium, Separation from Am- monium Sulphide Group, 178 Caesium, Reactions, 407 Calamine, 138 Calaverite, 444 Calcite, 54 Calcium, Reactions, 54. Calcium, Separation from Ba and Sr., 62, 391 Calcium Carbide, 57 Calcium Nitride, 57 Calcium Phosphide, 57 Calomel, 154 Carbamate of Ammonium, 55 Carbides, Solution of, 379 Carbon Disulphide, Detection of, 174 Carbonates, Behavior on Ignition, 308 Carbonates, Solubility of,306 Carbonic Acid, Dissociation of, 14 Carbonic Acid, Reactions, 305 Carnallite, 49 Carnallite, Dissociation of, 10 Carnot, Detection of Potassium, 39 Caro’s Acid, 347 Cassiterite, 210 Cations, Reactions of, 36 Caustic Alkali, Detection in Presence of Carbonates, 403 Celestite, 58 Ceric Compounds, Reactions, 419 Cerite, Analysis of, 426 Cerite Metals, 418 Cerium, 418 Cerous Compounds, Reactions, 418 Cerussite, 156 Chalcedony, 355 Chalcocite, 168 “Chamber Acid,” 157 Charcoal Reductions Blowpipe, 28, 368 Charcoal Stick Reactions, 24 Chili Saltpetre, 40 Chloanthite, 126 Chlorates, 343 Chloric Acid, Reactions, 343 Chloric Acid, Detection in the Pres- ence of HNOs and HC1, 344 Chlorides, Behavior on Ignition, 251 Chlorides, Solubility of, 248 Chlorine (free), Reactions, 253 Chlorine, Detection in Non-electro- lytes, 251 Chlorine, Detection in Presence of Brand I, 268 Chromates, 79 Chromic Acid, Reactions, 85 Chromic Acid, Oxidation by Means of, 6, 82 Chromic Compounds, Reactions, 77 Chromite, 76 A g Chromite, Analysis of, 378 Chromium, 76 Chromium, Separation from Fe, Al, and U, 112, 388 Before the INDEX. 463 ./ Chromium Peroxide, 76 Chromous Compounds, Reactions, 76 Chrysoberyl, 412 f Cinnabar, 146 ~ Citrates, Solubility of, 318 Citric Acid, 318 Claudetite, 180 Claustalite, 441 Clay, 69 Closed-tube Test, 26, 365 _Cobalt, Reactions, 132 Cobalt, Separation from Mn, Ni, and Zn, 143, 388 Cobaltite, 132 Cod in Test for Selenium, 443 Colloidal Substances, 70 Color Imparted to the Flame, 23, 368 Coloradoute, 444 Complex Ions, 17 Concentration of Reagents, 30 Confirmatory Test, 2 Coppe-, 168 Copper, Separation from Other Metals of H.S Group; 179, 226, 386 Copper, Separation from (NHA),S Group, 178, 382 Corunaun, 69 Course of Analysis, 364 Cream of Tartar, 314, 317 Crocoite, 76 Crookesite, 432 Cryolite, 349 ?_riº Cupric Compounds, Reactº 171 Cupric Xanthogenate, 173*. Cuprite, 168 Cuproscheelite, 439 Cuprous Compounds, Reactions, 170 Cuprous Xanthogenate, 174 Cyanates, Solubility of, 297 Cyanic Acid, 297 Cyanides, see Cyanogen Compounds Cyanides, Behavior on Ignition, 274 Cyanides, Complex Compounds, 93 Cyanides, Decomposition of, 378 Cyanogen, Complex Compounds of, 93, 94, 95 Cyanogen Compounds, Decomposi- tion of, 102, 378 D. Denigé's Reagent, 319 Devarda's Alloy, 6 Déville, St. Claire, 9 Diaspore, 69, 88 Dicyanogen, 275 Didymium, 428 Diphenylamine, Reagent, 288, 340 Dissociation, 9 Dissociation, Influence of Concen- tration upon, 15 Dissociation of Carbonic Acid, 14 Dissociation of Carnallite, 10 Dissociation of Electrolytes, 14 Dissociation of Hydrogen Sulphide, 14 Dissociation of Phosphorus Penta- chloride, 9 Dissociation of Sodium Carbonate, 14 Dissociation of Sodium Sulphide, 14 Dissociation of Water, 190. Dithionate of Manganese, 124 Division of Metals into Groups, 29 Dolomite, 49 Drinking Water, Ammonia in, 46 Dry Reactions, 19ſ, 365 E. Electrolytes, 12 Electrolytes, Degree of Dissociation, 14 Electrolytic Dissociation, 11 Emery, 69 Emplectite, 163 Epsomite, 49 Erbium, Reactions, 416 Erythrite, 132 Ethyl Acetate, 296 Ethyl Xanthogen Disulphide, 174 Ethylene Platinous Chloride, 236 Eucairite, 441 Euclase, 412 Euxenite, 415 ...” 2 F, ; Feldspar, 36 Ferberite, 439 Ferric Compounds, Reactions, 97 Ferric Compounds, Reduction of, 11 Ferricyanides, 279 Ferricyanides, Decomposition on Ig- nition, 102, 281 Ferrocyanides, 276 Ferrocyanides, De composition on Ig- nition, 102, 279 Ferrous Compounds, 90 Ferrous Compounds, Oxidation of, 4 Ferrous Oxide, Detection in Presence of Iron, 96 Filters, Size of, 19 Filtration, 19 464 INDEX. Fischer, Emil, Detection of Hydro- gen Sulphide, 292 Flame, Parts of, 20 Flame Reactions, 23, 368 Flame Spectra, 63 Fluorides, Analysis of, 353, 378 Fluorides, Solubility of, 350 Fluorine, Detection in Insoluble Fluorides, 352 Fluorite, 349 Formic Acid, 270 Franklinite, 138 Free Alkali in Presence of Carbon- ates, 403 Free Chlorine, 253 Fulgurator,67 Fusibility, Test of, 22 G. Gadolinite, 416 Gadolinite, Analysis of, 426 Gadolinite Metals, 416 Galena, 156 Garnierite, 126 Gas, Composition of Illuminating, 20 “Gas—water,” 269 Gersdorffite, 126 Gold, Detection of Small Amounts in Alloys and Ores, 230, 231 Gold, Reactions, 224 Gold, Separation from Platinum, 237 – Göthite, 88 Greenochite, 175 Greiss, Peter, Detection of Nitrous Acid, 287 Groups, Division of Acids into, 243 Groups, Division of Metals into, 29 Guldberg and Waage, “Law of Chemical Mass-action,” 9 Gutzeit Test for Arsenic, 197 & H. Halite, 40 Halogens, Detection of HCI, HBr, and HI in the Presence of One Another, 268 Halogens, Oxidation by, 4 Hausmannite, 113 Hematite, 88 Hornblende, 49 Horn Silver, 237 Hübnerite, 439 Hyalite, 355 Hydriodic Acid, 262 Hydriodic Acid, Detection in the Presence of HCl and HBr, 268 Hydriodic Acid, Reduction by, 8, 11 Hydrobromic Acid, 258 Hydrobromic Acid, Detection in the Presence of HCl and HI, 268 Hydrochloric Acid, 245 Hydrochloric Acid, Detection in the Presence of HBr and HI, 268 Hydrochloric Acid, Detection in the Presence of HNO, and HClO4, 344 Hydrochloric Acid, Oxidation by Nitric Acid, 247 Hydrochloric Acid, Oxidation by Peroxides, 246 . . Hydrochlorplatinic Acid, Reagent, 37, 234 Hydrocyanic Acid, 269 Hydrocyanic Acid, Detection in the Presence of Hydroferrocyanic Acid, 278 Hydroferricyanic Acid, 279 Hydroferrocyanic Acid, 276 Hydroferrocyanic Acid, Detection in the Presence of HCN, 278 Hydrofluoric Acid, 349 Hydrofluosilicic Acid, 354 Hydrogel, 70 Hydrogen, Reduction by, 6 Hydrogen Peroxide, 41 - Hydrogen Peroxide, Oxidation by, 5 - Hydrogen Peroxide, Reduction of Permanganate, 125 Hydrogen Sulphide, see Hydrosul- phuric Acid Hydrogen Sulphide, as Reducing Agent, 7 Hydrogen Sulphide, Detection in Presence of H2SOs and H.S.O.s, 336 Hydrogen Sulphide, Dissociation of, 14 Hydrolysis, 19 Hydrolysis of Ferric Salts, 97 Hydrolysis of Titanium, 110 Hydrosol, 70 Hydrosulphuric Acid, 289 Hydroxyl Ions, Detection in Pres- ence of Carbonates, 403 Hypochlorites, Solubility of, 256 Hypochlorous Acid, 255 Hypophosphites, Solubility of, 299 Hypophosphorous Acid, 299 I. Illuminating-gas, 20 Indicators, Theory of, 194 INDEX. 465 Indigo,254 Infusible Precipitate, 149 Insoluble Substances, Analysis of, 376 Iodates, Solubility of, 324 Iodic Acid, 324 Iodide of Starch, 267 Iodides, Solubility of, 263 Iodine, Detection in Non-electrolytes, 265 - Iodine, Detection in Presence of Bromine, 268 Iodine, Free, 265 Iodine, Reactions of, 265 Ion Theory, 11 Ions, Reactions of, 16 Iridium, 456 Iron, 88 Iron, Detection in the Presence of Ferrous Oxide, 96 Iron, Rusting of, 92 Iron, Separation from Al, Cr, and U, 112 Iron, Stäte of Oxidation in Original Substance, 388 Isatine, 254 } J. Jasper, 355 K. Raolin, 69 . Kermesite, 200 Rieserite, 49 Rlaproth, 104 Rrennerite, 444 L. Lacmoid, 19 Lanthanum, 421 Laxmannite, 76 Lead, Reactions, 156 Lead, Separation-from Ag and Hg, 242, 385 Lead, Separation from Other Metals of H2S Group, 226, 386, 179 Lead, Separation from Ammonium Sulphide Group, 178 Lead Acetate (Sugar of Lead), 295 Lead Sulphate, Solution of, 377 Lehrbachite, 441 Lepidolite, 408, 409 Leucophanite, 412 Liebigite, 104 Limonite, 88 Liquids, Analysis of, 400 Litharge, 158 Lithium, 409 Löllingite, 180 M. Magnesia Mixture, 327 Magnesite, 49 Magnesium, Reactions, 50 Magnesium, Reduction with, 26 Magnesium, Separation from Alka- lies, 51, 393 Magnetite, 49, 88 Malachite, 168 Manganese, 113 Manganese, Separation from Ni, Co, and Zn, 143, 388 Manganic Acid, 122 Manganite, 113 Manganites, 115 Manganous Compounds, 117 Marcasite, 88 Marsh Test for Arsenic, 190 Mass-action Law, 9 Mass-action Law, Applied to Hy- drolysis of Bismuth Solutions, 164 Mass-action Law, Applied to Precipi- tation by Ammonia, 50 Mass-action Law, Applied to Pre- cipitation by Hydrogen Sulphide, 91 Mass-action Law, Applied to Re- duction of Ferric Salts, 287 Meerschaum, 49 Meliphanite, 412 Melting-points, 22 Menaccanite, 108 Mercuric Compounds, 147 Mercuric Mercury, Separation from Other Metals Precipitated by H2S, 179, 226, 386 Mercuric Mercury, Separation from Ammonium Sulphide Group, 178 Mercurous Compounds, 153 Mercurous Mercury, Separation from Ag and Pb, 242, 385 Mercury, 146 Metalloids, 243 Metals, Division into Groups, 29 Metals, General Tables for the Ex- amination of, 382 Metaphosphates, Solubility of, 322 Metaphosphoric Acid, 321 Metastannic Acid, 218, 377 Methyl Orange, 19e Methylene Blue, 292 Mica, 69 466 INDEX. Mcrocosmic Salt, see Phosphorus, Salt of Millerite, 126 Mimetite, 156, 189 Minium, 158 Molybdate of Ammonium, 189 Molybdenum, Reactions, 437 Monazite, 414 Monticellite, 49 Mottramite, 434 Muscovite, 36, 69 N. Nagyagite, 224 Natrite 40 Naumannite, 441 Neodymium, 428 Nessler's Reagent, 47, 150 Newjanskite, 447 Niccolite, 126 Nickel, Reactions, 126 Nickel, Separation from Mn, Co, and Zn, 143, 388, 400 Niobium, 430 Nitrates, Solubility of, 339 Nitric Acid, Detection in the Presence of Nitrous Acid, 341 . Nitric Acid, Detection in the Presence of HClO, and HCl, 344 1Nitric Acid, Oxidation with, 4 Nitric Acid, Oxidation of Hydro- chloric Acid, 247 Nitric Acid, Reactions, 337 Nitrites, Solubility of, 286 Nitrose, 285 Nitrosoplatinic Chloride, 235 Nitrosyl Chloride, 247 Nitrosyl Sulphuric Acid, 285 Nitrous Acid, Detection in the Presence of HNOa, 341 Nitrous Acid, Reactions, 284 Noumeite, 126 O. Olivine, 49 Onofrite, 441 Organic Matter, Removal of, 73,125, 385 Organic Substances, Tests for Halo- gens, Sulphur, etc., 252, 261, 265, 293 Orpiment, 180 Orthite, 418 Orthoclase, 69 Osmic Acid, 453 | Osmium, Reactions, 452 Oxalates, Behavior on Ignition, 314 Oxalates, Solubility of, 313 Oxalic Acid, 312 Oxalic Acid, Detection in Preliminary Examination, 384 Oxalic Acid, Oxidation by Perman- ganate, 125 Oxidation, Définition of, 3 Oxidation by Chromic Acid, 6, 82 Oxidation by Halogens, 4 Oxidation by Hydrogen Peroxide, 5 Oxidation by Nitric Acid, 4 Oxidation by Permanganic Acid, 5, 123 Oxidation of One Molecule at Ex- pense of Another Similar One, 122 Oxidation of HCl by HNO, 246 Oxidation of HCl by Peroxides, 247 Oxidation and Reduction in Vitreous Fluxes, 23 ^ P. Palladic Compounds, 450 Palladium, 447 Palladodiamine Chloride, 449 Palladosamine Chloride, 449 Palladous Compounds, 448 Perchloric Acid, 345 Permanganates, see Acid Permanganic Acid, 122 Permanganic Acid, Oxidation by, 5, 124 Perowskit, 108 Peroxide of Hydrogen, 42 Peroxide of Sodium, 41 Peroxides, Solution of, 378 Peroxides, Test for, 400 Persulphuric Acid, 346 Petallite, 409 Phenolphthalein, 19e Phosphates, Solubility of, 326 Phosphides, Solution of, 379 Phosphites, 319 Phosphoric Acid, Detection in Pre- liminary Examination, 384 Phosphoric Acid, Reactions, 326 Phosphorous Acid, Reactions, 319 Phosphorus, 331 Phosphorus, Detection of, in Iron and Steel, 329 Phosphorus, Salt of, 331 Phosphorus Pentachloride, Dissocia- tion of, 9 Pitch-blende, 104 Permanganic | INDEX. 467. Platinic, Chloride, see Hydrochlor- platinic Acid Platinum, Reactions, 232 Platinum, Separation from Gold, 237 Platinum Metals, 447 Platinum Residues, Treatment of, 236 Plattnerite, 413 Plumbic Acid, 159 Polianite, 113, 413 Potash, 36 Pot ssium Acid Sulphate Fusion, 73 Pºm Chlorplatinate, Solubility of. 37 . f Potassium Percarbonate, 125 Potassium Persulphate, 346 Potassium Pyroantimonate, 208 Potassium Pyrosulphate, Fusion with, 73 ~Potassium, Reactions, 36 Potassium, Reduction with, 26 Powellite, 437 Praseodymium, 428 Precipitates, Washing of, 19 Preliminary. Examination, 365 Proustite, 180 Prussian Blue, 100 Prussiate of Potash, Yellow, 276 Prussiate of Potash, Red, 279 Prussic Acid, 269 Pseudo Acids and Bases, tºd Pyragryrite, 200 Pyrite, 88 Pyrochlore, 415 Pyrolusite, 113 Pyromorphite, 156 Pyrophosphates, Solubility of, 323 Pyrophosphoric Acid, Reactions, 323 Pyrosulphate of Potassium, Fusion with, 73 Q. Quartz, 355 R. Rare Earths, Reactions, 422 Rare Elements, 407 Rare Elements in Rocks, 436 Reaction, Definition, 1 Reactions, Sensitiveness of, 29, 33 Reactions in the Dry Way, 19 Reactions in the Wet Way, 1 Reactions Used for Qualitative Analysis, 1 Reagent, 1 Reagents, Concentration of, 30 Red Lead, 158 Reduction, Definition, 6 \ Reduction by Nascent Hydrogen, 6 Reduction by Devarda's Alloy, 7 Rººn by Hydrogen Sulphide, Reduction by Stannous Chloride, 8 Reduction in a Glass Tube, 26 Rºction in Upper Reducing Flame, Reduction on Charcoal, 28 Reduction on Charcoal Stick, 24 Rºduction with Sodium, Potassium, § Magnesium, 26 Reihite, 439 Reinsch Arsenic Test, 198 Rhodium, 450 Rochelle Salt, 314 Rock Salt, 40 Rubidium, Reactions, 408 Ruby, 69 Rºsting of Iron, 92 R tthenium, Reactions, 454 S. St. Claire Déville, 9,459 Salt of Phosphorus Beads, 23 Saltpetre, 36 Salts, Dissociation of, 14 Salts, Hydrolysis of, 19 Salts, Solubility Table of, 372 Samarskite, 104, 415 Sapphire, 69 Scheelite, 439 Selenic Acid, Reactions, 443 Selenious Acid, Reactions, 442 Selenites, 442 Selenium, 441 Senarmontite, 200 Sensitiveness of Reactions, 29, 33 Separation of the Five Metal Groups from One Another, 382 Separation of the Metals of Group I, 242, 385 Separation of the Metals of Group II, 179, 226, 386 Separation of the Metals of Group. III, 112, 144, 388, 390 Separation of the Metals of Group IV, 62, 391 Separation of the Metals of Group. V, 51,393 Separation of Acids, 396 Separation of Gold from Platinum, 237 Separation of the Platinum Metals, 458 Separation of the Cerite Metals, 426, 468 INDEX. Serpentine, 49 Siderite, 54, 88 Silicates, 357 Silicates, Analysis of, 377 Silicates, Decomposition with HF,361 Silicates, Soluble in Acids, 359 Silicates, Undecomposable by Acids, 360 Silicates, Water-soluble, 358, Silicic Acid, 355 Silicon, 362 Silver, Reactions, 237 Silver, Separation from Lead and Mercury, 212 Silver Telluride, 444 Skeleton Bead, 362 Skutterudite, 132 Smaltite, 132 Smithsonite, 138 Sodium, Reactions, 40 Sodium, Reduction with, 26 Sodium Carbonate, Dissociation of, 14 Sodium Peroxide, 41 Sodium Xanthogenate, 173 Solubility Product, 18 Solubility Table, 372 Solution of the Substance, 378 Solution of Insoluble Substances, 372, 376 Solutions, Analysis of, 400 Spectroscope, 64 Spectrum Analysis, 63 Sphalerite, 138 Spinel, 49 Spodumene, 408 Stannic Acid, 217 Stannic Compounds, 215 b-Stannic Chloride, 218 b-Stannic Compounds, 218 Stannous Chloride, Reduction with, 8 Stannous Compounds, 211 Stannyl Chloride, 218 Steatite, 49 Stibnite, 199 Stolzite, 439 Strength of Reagents, 30 Strontianite, 54, 58 Strontium, Reactions, 58 Strontium, Separation from Ba and Ca, 62, 391 Sulphates, Solubility of, 348 Sulphides, Behavior on Ignition, 293 Sulphides, Solubility of, 290 Sulphites, Solubility of, 301 Sulpho Acids, Separation Sulpho Bases, 226, 386 from Sulphocyanates, Behavior on Igni- tion, 283 Sulphocyanic Acid, 281 Sulphur, 293 Sulphuretted Hydrogen, see Hydro- sulphuric Aï Sulphuric Acid, Reactions, 348 Sulphurous Acid, Detection in the Presence of H2S2O, and H2S, Sulphurous Acid, Reactions, 300 sºurous Acid, Reduction by, 7, 381 Sylvanite, 224, 444 Sylvite, 36 Sysserskit, 447 T. Talc, 49 Tantalite, 429 Tantalum, Reactions, 429 Tartar Emetic, 202, 314 Tartaric Acid, 314 Tartrates, Solubility of, 314 Telluric Acid, Reactions, 446 Tellurium, 444 Tellurous Acid, Reactions, 446 Thallic Compounds, 432 Thallium, Reactions, 431 Thallous Compounds, 432 Thermonatrite, 40 Thiocyanic Acid, 281 Thiosulphates, Solubility of, 334 Thiosulphuric Acid, Reactions, 333 Thiosulphuric Acid, Detection in the Presence of Sulphurous Acid and Hydrogen Sulphide, 336 Thorium, 414f Tiemannite, 441 Tin, Reactions, 210 Tin, Separation from Copper Group, 226, 386 Tinkal, 40, 309 Tinstone, 210, 222, 378 Titanite, 108 Titanium, Reactions, 108 Titanium Dioxide, Analysis of, 378 Tremolite, 49 Trona, 40 Tungsten, Reactions, 438 Turmeric, 19, 311 Turnbull’s Blue, 95 U. Ullmannite, 126 Uranite, 104 Uranium, 104 INDEX 469 Uranium, Separation from Fe, Al, and Cr, 112, 388, 390 IJranyl Compounds, 106 Urea, Decomposition of, 44 V. Valentinite, 200 Vanadic Acid, 434 Wanadinite, 434 Wanadium, 434 Vanadium, Detection in Rocks, 436 Vivianite, 132 W. Wall-papers, Arsenic in, 196 Washing Precipitates, 19 Water, Ammonia in, 46 Water, Dissociation of, 190. Water Glasses, 4, 358 Water Opal, 355 Wave Lengths, 65 Witherite, 54, 59 Wolframite, 439 Wulfenite, 437, 439 Y. Xanthogen Compounds, 173 Y. Yttrium, 416 Yttrotantalite, 416 Z. Zinc, Reactions, 138 Zinc, Separation from Mn, Ni, Co, and Zn, 143, 388, 390 . Zincite, 138 Zircon, 413 Zirconium, 413 SHORT-TITLE CATALOGUE OF THE PUBLICATIONS JOHN WILEY & SONS NEW YORK LoNDoN: CHAPMAN & HALL, LIMITED ARRANGED UNDER SUBJECTS Descriptive circulars sent on application. Books marked with an asterisk (*) are sold at net prices only. All books are bound in cloth unless otherwise stated. AGRICULTURE–HORTICULTURE–FORESTRY. Armsby's Principles of Animal Nutrition..... . . . . . . . . . . . . . Budd and Hansen’s American Horticultural Manual: Part I. Propagation, Culture, and Improvement. . . . . . . . . . . . . . . 12mo, Part II. Systematic Pomology. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Elliott's Engineering for Land Drainage............ . . . . . . . . . . . . . . . . . 12mo, Practical Farm Drainage. (Second Edition, Rewritten.)... . . . . . 12mo, Graves's Forest Mensuration.............. . . . . . . . . . . . * . . 8vo, Green's Principles of American Forestry. . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Grotenfelt's Principles of Modern Dairy Practice. (Woll.).... . . . . . . . 12mo, * Herrick's Denatured or Industrial Alcohol . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Kemp and Waugh's Landscape Gardening. (New Edition, Rewritten. In Preparation.) * McKay and Larsen's Principles and Practice of Butter-making . . . . . . 8vo, Maynard's Landscape Gardening as Applied to Home Decoration...... 12mo, Sanderson's Insects Injurious to Staple Crops. . . . . . . . . . . . . . . . . . . . . . 12mo, Sanderson and Headlee's Insects Injurious to Garden Crops. (In Prepa- ration.) * Schwarz's Longleaf Pine in Virgin Forest. . . . . . . . . . . . . . . . . . . . . . . . 12mo, Stockbridge's Rocks and Soils. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Winton's Microscopy of Vegetable Foods. . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Woll's Handbook for Farmers and Dairymen.............. . . . . . . . . . . . 16mo, ARCHITECTURE. Baldwin's Steam Heating for Buildings. . . . . . . ‘. . . . . . . . . . . $ tº dº e º a º e a 12mo, Berg's Buildings and Structures of American Railroads. . . . . . . . . . . . . . . 4to, Birkmire's Architectural Iron and Steel........... . . . . . . . . . . . . . . . . . . . .8vo, Compound Riveted Girders as Applied in Buildings. . . . . . . . . . . . . . . . . 8vo, Planning and Construction of American Theatres.............. . . . . .8vo, Planning and Construction of High Office Buildings..... . . . . . . . . . .8vo, Skeleton Construction in Buildings. . . . . . . . . . . . . . . . . . . . l 1 :| : i . . . . . . . .8vo, $4 00 25 50 50 Briggs's Modern American School Buildings. . . . . . . . . . . . . . . . . . . . . . . . . 8vo, $4 00 Byrne's Inspection of Materials and Wormanship Employed in Construction. 16mo, Carpenter's Heating and Ventilating of Buildings.......... . . . . . . . . . . . .8vo, * Corthell's Allowable Pressure on Deep Foundations............... . . . 12mo, Freitag's Architectural Engineering. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .8vo, Fireproofing of Steel Buildings.... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Gerhard's Guide to Sanitary Inspections. (Fourth Edition, Entirely Re- vised and Enlarged.) . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, * Modern Baths and Bath Houses. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Sanitation of Public Buildings.............. . . . . . . . . . . . . . . . . . . . . 12mo, Theatre Fires and Panics..... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, * The Water Supply, Sewerage and Plumbing of Modern City Buildings. 8vo, Johnson's Statics by Algebraic and Graphic Methods...... . . . . . . . . . . . . 8vo, Kellaway's How to Lay Out Suburban Home Grounds. . . . . . . . . . . . . . . 8vo, Kidder's Architects’ and Builders' Pocket-book........... . . . . . . 16mo, mor., Merrill's Stones for Building and Decoration.... . . . . . . . . . . . . . . . . . . . . . 8vo, Monckton's Stair-building. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4to, Patton's Practical Treatise on Foundations.... . . . . . . . . . . . . . . . . . . . . . . 8vo, Peabody’s Naval Architecture. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8wc, Rice's Concrete-block Manufacture................ . . . . . . . . . . . . . . 8vo, Richey's Handbook for Superintendents of Construction . . . . . . . 16mo, mor. Building Foreman’s Pocket Book and Ready Reference. . 16mo, mor. * Building Mechanics' Ready Reference Series: * Carpenters' and Woodworkers' Edition. . . . . . . . . . . . 16mo, mor. * Cement Workers' and Plasterers' Edition........... 16mo, mor. * Plumbers’, Steam-Fitters', and Tinners' Edition. . . 16mo, mor. * Stone- and Brick-masons' Edition. . . . . . . . . . . . . . . . 16mo, mor. Sabin's House Painting.............. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Siebert and Biggin's Modern Stone-cutting and Masonry. . . . . . . . . . . . . . 8vo, Snow's Principal Species of Wood......... . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Towne's Locks and Builders' Hardware....... . . . . . . . . . . . . . . . . . 16mo, mor. Wait's Engineering and Architectural Jurisprudence........ . . . . . . . . . . .8vo, Sheep, Law of Contracts........ . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Law of Operations Preliminary to Construction in Engineering and Architecture......... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Sheep, Wilson's Air Conditioning.......... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Worcester and Atkinson's Small Hospitals, Establishment and Maintenance, Suggestions for Hospital Architecture, with Plans for a Small Hospital. . . . . . . . . . . . . . . . . . . e g º ſº tº e º & © tº º ſº, º is tº º e º g e g º e 12mo, ARMY AND NAVY. Bernadou's Smokeless Powder, Nitro-cellulose, and the Theory of the Cellu- lose Molecule. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Chase's Art of Pattern Making. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Screw Propellers and Marine Propulsion. . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Cloke's Enlisted Specialists' Examiner. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Gunner's Examiner... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Craig's Azimuth... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4to, Crehore and Squier's Polarizing Photo-ch1.onograph.............. . . . . . . . 8vo, * Davis's Elements of Law........ . . . . . . . . . . . . . . . .8vo, * Treatise on the Military Law of United States. . . . . . . . . . . . . . . . . 8vo, * Dudley's Military Law and the Procedure of Courts-martial...Large 12mo, Durand's Resistance and Propulsion of Ships................ . . . . . . . . . . .8vo, * Dyer's Handbook of Light Artillery................ . . . . . . . . . . . . . . . . 12mo, Eissler's Modern High Explosives. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Fiebeger's Text-book on Field Fortification........ . . . . . . . . . . Large 12mo, Hamilton and Bond's The Gunner's Catechism................ . . . . . . . . 18mo, * Hoff's Elementary Naval Tactics......... . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, i 2 5 :i 0 0 5 O 2 Ingalks's Handbook of Problems in Direct Fire, ... . . . . . . . . . . . . . . . . . . .8vo, $4 * Lissak’s Ordnance and Gunnery............ . . . . . . . . . . . . . . . . . . . . . . . .8vo, * Ludlow's Logarithmic and Trigonometric Tables.............. . . . . . . . .8vo, * Lyons's Treatise on Electromagnetic Phenomena. Vols. I. and II.8vo, each, * Mahan's Permanent Fortifications. (Mercur.)........ . . . . . .8vo, half mor. Manual for Courts-martial................ . . . . . . . . . . . . . . . . . . . . . . 16mo, mor. * Mercur's Attack of Fortified Places. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, * Elements of the Art of War. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Nixon's Adjutants' Manual. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 24mo, Peabody's Naval Architecture. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Phelps's Practical Marine Surveying...... . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Putnam's Nautical Charts. ... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Rust’s Ex-meridian Altitude, Azimuth and Star-Finding Tables. . . . . . . . 8vo, Selkirk's Catechism of Manual of Guard Duty. (In Press.) Sharpe's Art of Subsisting Armies in War........ . . . . . . . . . . . . . . 18mo, mor. Taylor's Speed and Power of Ships. (In Press.) * Tupes and Poole's Manual of Bayonet Exercises and Musketry Fencing. 24mo, leather, * Weaver's Military Explosives..... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Woodhull's Military Hygiene for Officers of the Line. . . . . . . . . Large 12mo, ASSAYING. 13etts's Lead Refining by Electrolysis........ . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Fletcher's Practical Instructions in Quantitative Assaying with the Blowpipe. 16mo, mor. Furman and Pardoe’s Manual of Practical Assaying. (Sixth Edition, Re- vised and Enlarged.). . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Lodge's Notes on Assaying and Metallurgical Laboratory Experiments..8vo, Low's Technical Methods of Ore Analysis. . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Miller's Cyanide Process... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Manual of Assaying. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Minet’s Production of Aluminum and its Industrial Use. (Waldo.)...12mo, Ricketts and Miller's Notes on Assaying. . . . . . . . . . . . . . . . . . . . . . . . . .8vo, Robine and Lenglen's Cyanide Industry. (Le Clerc.). . . . . . . . . . . . . . . . 8vo, Seamon's Manual for Assayers and Chemists. (In Press.) Ulke's Modern Electrolytic Copper Refining. . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Wilson's Chlorination Process. . . . . . . . . . . . . . . . . . . . . . º 12mo, Cyanide Processes. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, ASTRONOMY. Comstock's Field Astronomy for Engineers............ . . . . . . . . . . . . . . . .8vo, Craig's Azimuth. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4to, Crandall's Text-book on Geodesy and Least Squares. . . . . . . . . . . . . . . . . . 8vo, Doolittle's Treatise on Practical Astronomy.... . . . . . . . . . . . . . . . . . . . . . . . .8vo, Hayford's Text-book of Geodetic Astronomy...... . . . . . . . . . . . . . . . . . . . 8vo, Hosmer's Azimuth............ . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16mo, mor. Merriman's Elements of Precise Surveying and Geodesy. . . . . . . . . . . . . . 8vo, * Michie and Harlow's Practical Astronomy. . . . . . . . . . . . . . . . . . . . . . . . .8vo, Rust's Ex-meridian Altitude, Azimuth and Star-Finding Tables. . . . . . .8vo, * White's Elements of Theoretical and Descriptive Astronomy........ 12mo, CHEMISTRY. * Abderhalden's Physiological Chemistry in Thirty Lectures. (Hall and Defren). . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Abegg's Theory of Electrolytic Dissociation. (von Ende.). . . . . . . . . 12mo, Alexeyeff's General Principles of Organic Syntheses. (Matthews.).......8vo, Allen's Tables for Iron Analysis. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Armsby's Principles of Animal Nutrition........ . . . . . . . . . . . . . . . . º 8vo Arnold's Compendium of Chemistry. (Mandel.)............ . . . .i.arge 12mo, 1 1. i : 00 00 00 00 50 3 Association of State and National Food and Dairy Departments, Hartford Meeting, 1906. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Jamestown Meeting, 1907................ . . . . . . . . . . . . . . . . . . . . . . . .8vo, Austen's Notes for Chemical Students........ . . . . . . . . . . . . . . . . . . . . . 12mo, Baskerville's Chemical Elements. (In Preparation.) Bernadou's Smokeless Powder.—Nitro-cellulose, and Theory of the Cellulose Molecule. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, * Biltz's Introduction to Inorganic Chemistry. (Hall and Phelan.). . . 12mo, Laboratory Methods of Inorganic Chemistry. (Hall and Blanchard.) 8vo, * Blanchard's Synthetic Inorganic Chemistry. . . . . . . . . . . . . . . . . . . . . . . . . 12mo, * Browning's Introduction to the Rarer Elements...... . . . . . . 8vo, * Claassen's Beet-sugar Manufacture. (Hall and Rolfe). º : ſ . . . . . . 8vo, Classen's Quantitative Chemical Analysis by Electrolysis. (Boltwood.).8vo, Cohn's Indicators and Test-papers. . . . . . . . . . . . . . . . . . 12mo, Tests and Reagents.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Danneel's Electrochemistry. (Merriam ). ... . . . . . . . . . . . . . . . . . . . . . . 12mo, Dannerth's Methods of Textile Chemistry. . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Duhem's Thermodynamics and Chemistry, (Burgess.). . . . . . . . . . . . . . . . 8vo, Effront's Enzymes and their Applications. (Prescott.). . . . . . . . . . . . . . . 8vo, Eissler's Modern High Explosives. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Erdmann's Introduction to Chemical Preparations. (Dunlap.).... . . . 12mo, * Fischer's Physiology of Alimentation. . . . . . . . . . . . Large 12mo, Fletcher's Practical Instructions in Quantitative Assaying with the Blowpipe. 16mo, mor. Fowler's Sewage Works Analyses. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Fresenius's Manual of Qualitative Chemical Analysis. (Wells.).......... 8vo, ; : Manual of Qualitative Chemical Analysis. Part I. Descriptive. (Wells.)8vo, 3 Quantitative Chemical Analysis. (Cohn.) 2 vols. . . . . . . . . . . . . . . 8vo, When Sold Separately, Vol. I, $6. Vol. II, $8. Fuertes's Water and Public Health. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Furman and Pardoe's Manual of Practical Assaying. (Sixth Edition, Revised and En larged.). . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Getman's Exercises in Physical Chemistry. . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Gill's Gas and Fuel Analysis for Engineers. . . . . . . . . . . . . . . . . . . . . . . . . 12mo, * Gooch and Browning's Outlines of Qualitative Chemical Analysis, Large 12mo, Grotenfelt's Principles of Modern Dairy Practice. (Woll.). . . . . . . . . . . 12mo, Groth's Introduction to Chemical Crystallography (Marshall). . . . . . . . . 12mo, Hammarsten's Text-book of Physiological Chemistry. (Mandel.). . . . . . 8vo, Hanausek's Microscopy of Technical Products. (Winton.)............... 8vo, * Haskins and Macleod's Organic Chemistry. . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Hering's Ready Réference Tables (Conversion Factors)........... 16mo, mor. * Herrick's Denatured or Industrial Alcohol......... . . . . . . . . . . . . . . 8vo, Hinds's Inorganic Chemistry. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Laboratory Manual for Students. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, * Holleman's Laboratory Manual of Organic Chemistry for Beginners. (Walker.) . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Text-book of Inorganic Chemistry. (Cooper.). . . . . . . . . . . . . . . . . . 8vo, Text-book of Organic Chemistry. (Walker and Mott.). . . . . . . . . . . . . 8vo, * Holley's Lead and Zinc Pigments. . . . . . . . . . . . . . . . . . . . . . . . . Large 12mo, Holley and Ladd's Analysis of Mixed Paints, Color Pigments, and Varnishes. Large 12mo, Hopkins's Oil-chemists' Handbook. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Jackson's Directions for Laboratory Work in Physiological Chemistry. . 8vo, Johnson's Rapid Methods for the Chemical Analysis of Special Steels, Steel- making Alloys and Graphite. . . . . . . . . . . . . . . . . . . . . . . Large 12mo, Landauer's Spectrum Analysis. (Tingle ) . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Lassar-Cohn's Application of Some General Reactions to Investigations in Organic Chemistry. (Tingle.). . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Leach's Inspection and Analysis of Food with 'Special Reference to State Control. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .8vo, Lob's Electrochemistry of Organic Compounds (Lorenz.). . . . . . . . . . . . . . 8vo, Lodge's Notes on Assaying and Metallurgical Laboratory Experiments...8vo, Low's Technical Method of Ore Analysis. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Lowe's Paint for Steel Structures. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Lunge's Techno-chemical Analysis. (Cohn.). . . . . . . . . . . . . . . . . . . . . . . . 12mo, 12 1. ;::: 1 : 4 * * McKay and Larsen's Principles and Practice of Butter-making........ 8vo, Maire's Modern Pigments and their Vehicles. . . . . . . . . . . . . . . . . . . . . . . 12mo, Mandel's Handbook for Bio-chemical Laboratory............ . . . . . . . . . 12mo, * Martin's Laboratory Guide to Qualitative Analysis with the Blowpipe 12mo, Mason's Examination of Water. (Chemical and Bacteriological.)...... 12mo, Water-supply. (Considered Principally from a Sanitary Standpoint.) 8vo, * Mathewson's First Principles of Chemical Theory......... . . . . . . . . . . . .8vo, Matthews's Laboratory Manual of Dyeing and Textile Chemistry. . . . . .8vo, Textile Fibres. 2d Edition, Rewritten... . . . . . . . . . . . . . . . . . . . . . . .8vo, * Meyer's Determination of Radicles in Carbon Compounds. (Tingle.) Third Edition. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . g ſº º e g g g º º 12mo, Miller's Cyanide Process....... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Manual of Assaying. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Minet's Production of Aluminum and its Industrial Use. (Waldo.)...12mo, * Mittelstaedt's Technical Calculations for Sugar Works. (Bourbakis.) 12mo, Mixter's Elementary Text-book of Chemistry. . . . . . . . . . . . . . . . . . . . . . . . 12mo, Morgan's Elements of Physical Chemistry........ . . . . . . . . . . . . . . . . . . . 12mo, Outline of the Theory of Solutions and its Results.............. . 12mo, * Physical Chemistry for Electrical Engineers..................... . 12mo, * Moore's Outlines of Organic Chemistry. . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Morse's Calculations used in Cane-sugar Factories. . . . . . . . . . . . . . . 16mo, mor. * Muir's History of Chemical Theories and Laws. . . . . . . . . . . . . . . . . . . . . 8vo, Mulliken's General Method for the Identification of Pure Organic Compounds. Vol. I. Compounds of Carbon with Hydrogen and Oxygen. Large 8vo, Vol. II. Nitrogenous Compounds. (In Preparation.) Vol. III. The Commercial Dyestuffs. . . . . . . . . . . . . . . . . . . Large 8vo, * Nelson's Analysis of Drugs and Medicines. . . . . . . . . . . . . . . . . . . . . . . 12mo, O'Driscoll's Notes on the Treatment of Gold Ores................ . . . . . . .8vo, Ostwald's Conversations on Chemistry. Part One. (Ramsey.). . . . . . 12mo, { { & 4 tº º 4 & Part Two. (Turnbull.). . . . . 12mo, Introduction to Chemistry. (Hall and Williams.) (In Preparation.) Owen and Standage's Dyeing and Cleaning of Textile Fabrics.......... 12mo, * Palmer's Practical Test Book of Chemistry. . . . . . . . . . . . . . . . . . . . . . . . . 12mo, * Pauli's Physical Chemistry in the Service of Medicine. (Fischer ) . . 12mo, Penfield's Tables of Minerals, Including the Use of Minerals and Statistics of Domestic Production.... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Pictet's Alkaloids and their Chemical Constitution. (Biddle.). . . . . . . . . 8vo, Poole's Calorific Power of Fuels. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Prescott and Winslow's Elements of Water Bacteriology, with Special Refer- ence to Sanitary Water Analysis. . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, * Reisig's Guide to Piece-Dyeing. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Richards and Woodman's Air, Water, and Food from a Sanitary Stand- point. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo. Ricketts and Miller's Notes on Assaying..... . . . . . . . . . . . . . . . . . . . . . 8vo, Rideal's Disinfection and the Preservation of Food. . . . . . . . . . . . . . . . . . . . .8vo, Sewage and the Bacterial Purification of Sewage. . . . . . . . . . . . . . . . . . . .8vo, Riggs's Elementary Manual for the Chemical Laboratory. . . . . . . . . . . . . . . 8vo, Robine and Lenglen's Cyanide Industry. (Le Clerc.). . . . . . . . . . . . . . . . 8vo, Ruddiman's Incompatibilities in Prescriptions. . . . . . . . . . . . . . . . . . . . . . . 8vo, Whys in Pharmacy............... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, * Ruer's Elements of Metallography. (Mathewson.) . . . . . . . . . . . . . . . . 8vo, Sabin's Industrial and Artistic Technology of Paint and Varnish. . . . . . .8vo, Salkowski's Physiological and Pathological Chemistry. (Orndorff.)..... 8vo, Schimpf's Essentials of Volumetric Analysis. . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Manual of Volumetric Analysis. (Fifth Edition, Rewritten) . . . . . 8vo, * Qualitative Chemical Analysis. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Seamon's Manual for Assayers and Chemists. (In Press.) Smith's Lecture Notes on Chemistry for Dental Students. . . . . . . . . . . . . 8vo, Spencer's Handbook for Cane Sugar Manufacturers.............. . 16mo, mor. Handbook for Chemists of Beet-sugar Houses. . . . . . . . . . . . . . . I6mo, mor. Stockbridge's Rocks and Soils. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Stone's Practical Testing of Gas and Gas Meters. . . . . . . . . . . . . . . . . . . . . . 8vo, * Tillman's Descriptive General Chemistry. . . . . . . . . . . . . . . . . . . . . . . . . . . .8vo, * Elementary Lessons in Heat . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Treadwell's Qualitative Analysis. (Hall.).... . . . . . . . . . . . . . . . . . . . . . . . 8vo, 5 º :; 2 ; Treadwell's Quantitative Analysis. (Hall.).......... . . . . . . . . . . . . . . . . . 8vo, $4 00 Turneaure and Russell's Public Water-supplies.............. . . . . . . . . . . .8vo, Van Deventer's Physical Chemistry for Beginners. (Boltwood.)...... 12mo, Venable's Methods and Devices for Bacterial Treatment of Sewage...... 8vo, Ward and Whipple's Freshwater Biology. (In Press.) Ware's Beet-sugar Manufacture and Refining. Vol. I... . . . . . . . . . . . . . . 8vo, Vol. II...... . . . . . . . . . .8vo, Washington's Manual of the Chemical Analysis of Rocks............ . . . , 8vo, * Weaver's Military Explosives. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Wells's Laboratory Guide in Qualitative Chemical Analysis.... . . . . . . . Short Course in Inorganic Qualitative Chemical Analysis for Engineering Students. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Text-book of Chemical Arithmetic............ . . . . . . . . . . . . . . . . . , 12mo, Whipple's Microscopy of Drinking-water............... . . . . . . . . . . . . . . . . 8vo, Wilson's Chlorination Process................ . . . . . . . . . . . . . . . . . . . . . . . 12mo, Cyanide Processes. . . . . . . . . . . . . . . . . . . . . . . . . tº e º 'º e º is e º a º e º ſº dº e e 12mo, Winton's Microscopy of Vegetable Foods... . . . . . . . . . . . . . . . . . . . . . . . . . 8vo Zsigmondy's Colloids and the Ultramicroscope. (Alexander.), . Large 12mo, CIVIL ENGINEERING. ii 0 0 5 O BRIDGES AND ROOFS. HYDRAULICS. MATERIALS OF ENGINEER- ING. RAILWAY ENGINEERING. Baker's Engineers' Surveying Instruments............ . . . . . . . . . . . . . . . 12mo, Bixby's Graphical Computing Table . . . . . . . . . . . . . . . Paper 19% X24% inches. Breed and Hosmer's Principles and Practice of Surveying, Vol. I. Elemen- tary Surveying. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Vol. II. Higher Surveying.............. . . . . . . . . . . . . . . . . . . . . . . . .8vo, * Burr's Ancient and Modern Engineering and the Isthmian Canal...... 8vo, Comstock's Field Astronomy for Engineers. . . . . . . . . . . . . . . . . . . . . . 8vo, * Corthell's Allowable Pressure on Deep Foundations . . . . . . . . . . . . . . . 12mo, Crandall's Text-book on Geodesy and Least Squares... . . . . . . . . . . . . . . . 8vo, Davis's Elevation and Stadia Tables. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Elliott's Engineering for Land Drainage........... . . . . . . . . . . . . . . . . . . 12mo, Practical Farm Drainage. (Second Edition Rewritten.). . . . . . . . 12mo * Fiebeger's Treatise on Civil Engineering............ . . . . . . . . . . . . 8vo, Flemer's Photographic Methods and Instruments............. . . . . . . . . . .8vo, Folwell's Sewerage. (Designing and Maintenance.)......... . . . . . . . . . . .8vo, Freitag's Architectural Engineering.............. . . . . . . . . . . . . . . . . . . . . .8vo, Goodhue's Municipal Improvements....... . . . . 12mo, O O * Hauch and Rice's Tables of Quantities for Preliminary Estimates... 12mo, Hayford's Text-book of Geodetic Astronomy...... . . . . . . . . . . . . . . . . . . . 8vo, Hering's Ready Reference Tables (Conversion Factors.).......... 16mo, mor. Hosmer's Azimuth. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16mo, mor, Howe' Retaining Walls for Earth. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, * Ives's Adjustments of the Engineer's Transit and Level. . . . . . . 16mo, bds. Johnson's (J. B.) Theory and Practice of Surveying. . . . . . . . . . Large 12mo, Johnson's (L. J.) Statics by Algebraic and Graphic Methods........... . . 8vo, Kinnicutt, Winslow and Pratt's Purification of Sewage. (In Preparation.) * Mahan's Descriptive Geometry......... . . . . . . . . . . . . . . . . . . . . . . . . . . . .8vo, Merriman's Elements of Precise Surveying and Geodesy. . . . . . . . . . . . . . 8vo, Merriman and Brooks's Handbook for Surveyors. . . . . . . . . . . . . . 16mo, mor. TNugent's Plane Surveying. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Ogden's Sewer Construction............... . . . . . . . . . . . . . . . . . . . . . . . . . . , 8vo, Sewer Design. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Parsons's Disposal of Municipal Refuse... . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Batton's Treatise on Civil Engineering. . . . . . . . . . . . . . . . . . 8vo, half leather, Reed's Topographical Drawing and Sketching. . . . . . . . . . . . . . . . . . . . . . . . 4to, Rideal's Sewage and the Bacterial Purification of Sewage.... . . . . . . . . . . 8vo, Riemer's Shaft-sinking under Difficult Conditions. (Corning and Peele.).8vo, Siebert and Biggin's Modern Stone-cutting and Masonry. . . . . . . . . . . . . . 8vo, Smith's Manual of Topographical Drawing. (McMillan.)............... . 8vo, : O O O 0 6 Soper's Air and Ventilation of Subways. . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, * Tracy's Exercises in Surveying. . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, mor. Tracy's Plane Surveying. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . , 16mo, mor. * Trautwine's Civil Engineer's Pocket-book.......... g ºf e º a tº sº e º s is 16mo, mor. Venable's Garbage Crematories in America. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Methods and Devices for Bacterial Treatment of Sewage. . . . . . . . . . . . 8vo, Wait's Engineering and Architectural Jurisprudence..... . . . . . . . . . . . . . . 8vo, Sheep, Law of Contracts. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Law of Operations Preliminary to Construction in Engineering and Architecture. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Sheep, Warren's Stereotomy—Problems in Stone-cutting. . . . . . . . . . . . . . . . . . . . . . . 8vo, * Waterbury's Vest-Pocket Hand-book of Mathematics for Engineers. 2% X5% inches, mor. * Enlarged Edition, Including Tables . . . . . . . . . . . . . . . . . . . . . . . . . . II) Of. Webb's Problems in the Use and Adjustment of Engineering Instruments. 16mo, mor. Wilson's Topographic Surveying. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, BRIDGES AND ROOFS. Boller's Practical Treatise on the Construction of Iron Highway Bridges...8vo, * Thames River Bridge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Oblong paper, Burr and Falk's Design and Construction of Metallic Bridges. . . . . . . . . . 8vo, Influence Lines for Bridge and Roof Computations........ . . . . . . . . . . 8vo, Du Bois’s Mechanics of Engineering. Vol. II. . . . . . . . . . . . . . . . . . . Small 4to, Foster's Treatise on Wooden Trestle Bridges. . . . . . . . . . . . . . . . . . . . . . . . . . 4to, Fowler's Ordinary Foundations...... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .8vo, Greene's Arches in Wood, Iron, and Stone.......... . . . . . . . . . . . . . . . . . .8vo, Bridge Trusses. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Roof Trusses. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .8vo, Grimm's Secondary Stresses in Bridge Trusses. . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Heller's Stresses in Structures and the Accompanying Deformations. . . . 8vo, Howe's Design of Simple Roof-trusses in Wood and Steel. . . . . . . . . . . . . . . 8vo. Symmetrical Masonry Arches. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Treatise on Arches. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Jacoby's Structural Details, or Elements of Design in Heavy Framing, 8vo, Johnson, Bryan and Turneaure's Theory and Practice in the Designing of Modern Framed Structures. . . . . . . . . . . . . . . . . . . . . . . . . . Small 4to, * Johnson, Bryan and Turneaure's Theory and Practice in the Designing of Modern Framed Structures. New Edition. Part I. . . . . . . . . 8vo, Merriman and Jacoby's Text-book on Roofs and Bridges: Part I. Stresses in Simple Trusses..... . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Part II. Graphic Statics. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Part III. Bridge Design. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Part IV. Higher Structures.......... . . . . . . . . . . . . . . . . . . . . . . . . . . .8vo, Morison's Memphis Bridge....... . . . . . . . . . . . . . . . . . . . . . . . . . . . . Oblong 4to, Sondericker's Graphic Statics, with Applications to Trusses, Beams, and Arches. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Waddell's De Pontibus, Pocket-book for Bridge Engineers...... . 16mo, mor. * Specifications for Steel Bridges... . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Waddell and Harrington's Bridge Engineering. (In Preparation.) Wright's Designing of Draw-spans. Two parts in one volume........... 8vo, HYDRAULICS. Barnes's Ice Formation. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Bazin's Experiments upon the Contraction of the Liquid Vein Issuing from an Orifice. (Trautwine.)... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Bovey's Treatise on Hydraulics. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .8vo, Church's Diagrams of Mean Velocity of Water in Open Channels. Oblong 4to, paper, ;i . 1 I O 1 i3 : 3 3 :: 00 Hydraulic Motors....... e e s is e g º º e º is a e s tº * * * * * * * * * * * * * * * * * * * * * * * * * 8vo, 00 00 50 00 Coffin's Graphical Solution of Hydraulic Problems. . . . . . . . . . . . . 16mo, mor, $2 50 Flather's Dynamometers, and the Measurement of Power.... . . . . . . . . 12mo, Folwell's Water-supply Engineering. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Frizell's Water-power. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . tº e º us tº º q & d e º ſº a 8vo, Fuertes's Water and Public Health. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Water-filtration Works........ . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Ganguillet and Kutter's General Formula for the Uniform Flow of Water in Rivers and Other Channels. (Hering and Trautwine.). . . . . . . 8vo, Hazen's Clean Water and How to Get It. . . . . . . . . . . . . . . . . . . . Large 12mo, Filtration of Public Water-supplies........ . . . . . . . . . . . . . . . . . . 8vo, IHazelhurst's Towers and Tanks for Water-works. . . . . . . . . . . . . . . . . . . . . , 8vo, Herschel's 115 Experiments on the Carrying Capacity of Large, Riveted, Metal Conduits. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Hoyt and Grover's River Discharge. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Hubbard and Kiersted's Water-works Management and Maintenance. 8vo, * Lyndon's Development and Electrical Distribution of Water Power. 8vo, Mason's Water-supply. (Considered Principally from a Sanitary Stand- point.) . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Merriman's Treatise on Hydraulics. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ... 8vo, * Molitor's Hydraulics of Rivers, Weirs and Sluices. . . . . . . . . . . . . . . . . . . 8vo, Morrison and Brodie's High Masonry Dam Design. (In Press.) * Richards's Laboratory Notes on Industrial Water Analysis. . . . . . . . . . 8vo, Schuyler's Reservoirs for Irrigation, Water-power, and Domestic Water- supply. Second Edition, Revised and Enlarged. . . . . . . Large 8vo, * Thomas and Watt's Improvement of Rivers. . . . . . . . . . . . . . . . . . . . . . . . 4to, Turneaure and Russell's Public Water-supplies. . . . . . . . . . . . . . . . . . . . . 8vo, Wegmann's Design and Construction of Dams. 5th Ed , enlarged.... . .4to, Water-Supply of the City of New York from 1658 to 1895. . . . . . . . . . 4to, Whipple's Value of Pure Water... . . . . . . . . . . . . . . . . . . . . . . . . . . Large 12mo, Williams and Hazen's Hydraulic Tables. . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Wilson's Irrigation Engineering........... . . . . . . . . . . . . . . . . . . . . . . . . . . .8vo, Wood's Turbines. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, MATERIALS OF ENGINEERING. Baker's Roads and Pavements.... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Treatise on Masonry Construction........ . . . . . . . . . . . . . . . . . . . . . . . 8vo, Black's United States Public Works........ . . . . . . . . . . . . . . . Oblong 4to, Blanchard's Bituminous Roads. (In Press) tº Bleininger's Manufacture of Hydraulic Cement. (In Preparation.) * Bovey's Strength of Materials and Theory of Structures................ 8vo, Burr's Elasticity and Resistance of the Materials of Engineering. . . . . . . . 8vo, Byrne's Highway Construction..... .8vo, i4:ii I : Inspection of the Materials and Workmanship Employed in Construction. 16mo, Church's Mechanics of Engineering........... . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Du Bois's Mechanics of Engineering. Vol. I Kinematics, Statics, Kinetics. . . . . . . . . . . . . . . . . . . . . Small 4to, Vol. II. The Stresses in Framed Structures, Strength of Materials and Theory of Flexures. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Small 4to, * Eckel's Cements, Limes, and Plasters...... . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Stone and Clay Products used in Engineering. (In Preparation.) Fowler's Ordinary Foundations. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .8vo, * Greene's Structural Mechanics . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Holley's Lead and Zinc Pigments. . . . . . . . . . . . . . . . . . . . . . . . . Large 12mo, Holley and Ladd's Analysis of Mixed Paints, Color Pigments and Varnishes. Large 12mo, * Hubbard's Dust Preventives and Road Binders. . . . . . . . . . . . . . 8vo, Johnson's (C M.) Rapid Methods for the Chemical Analysis of Special Steels, Steel-making Alloys and Graphite. . . . . . . . . . . . . . . . .Large 12mo, Johnson's (J. B.) Materials of Construction. . . . . . . . . . . . . . . . . . . . Large 8vo, Keep's Cast Iron. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Lanza's Applied Mechanics. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Lowe's Paints for Steel Structures. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, :º7; Maire's Modern Pigments and their Vehicles...... is ſº tº gº ºn tº ſº gº & is ſº tº ſº tº gº & & is 12mo, $2 00 Maurer's Technical Mechanics. . . . . . . . . . . . . . . . . . tº e º 'º jº tº $ tº * * w ſº tº * * * * . . . . 8vo, Merrill's Stones for Building and Decoration.... . . . . . . . . # * tº dº ſº e 4 & © # * tº gº 8vo, Merriman's Mechanics of Materials. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Strength of Materials. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Metcalf's Steel. A Manual for Steel-users. . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Morrison's Highway Engineering... . . . . . . . . . . . . . . . . . . . . . . gº tº e s is s a s 8vo, Patton's Practical Treatise on Foundations. . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Rice's Concrete Block Manufacture.... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Richardson's Modern Asphalt Pavement. . . . . . . . . . 8vo, Richey's Building Foreman's Pocket Book and Ready Reference. 16mo, mor. * Cement Workers' and Plasterers' Edition (Building Mechanics' Ready Reference Series)...... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16mo, mor. Handbook for Superintendents of Construction. . . . . . . . . . . . . 16mo, mor. * Stone and Brick Masons' Edition (Building Mechanics' Ready Reference Series) . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16mo, mor. * Rieg's Clays: Their Occurrence, Properties, and Uses. . . . . . . . . . . . . . . . . 8vo, * Ries and Leighton's History of the Clay-working Industry of the United States. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo Sabin's Industrial and Artistic Technology of Paint and Varnish....... .8vo, * Smith's Strength of Material. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo Snow's Principal Species of Wood. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Spalding's Hydraulic Cement. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Text-book on Roads and Pavements. . . . . . . . . . . . . . . . . . . . . . . . . . Taylor and Thompson's Treatise on Concrete, Plain and Reinforced. . . . . 8vo, Thurston's Materials of Engineering. In Three Parts. . . . . . . . . . . . . . . . . . . 8vo, Part I. Non-metallic Materials of Engineering and Metallurgy. . . . 8vo, Part II. Iron and Steel. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .8vo, Part III. A Treatise on Brasses, Bronzes, and Other Alloys and their Constituents. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Tillson's Street Pavements and Paving Materials. . . . . . . . . . . . . . . . . . . . 8vo, * Trautwine's Concrete, Plain and Reinforced . . . . . . . . . . . . . . . . . . . . . 16mo, Turneaure and Maurer's Principles of Reinforced Concrete Construction. Second Edition, Revised and Enlarged . . . . . . . . . . . . . . . . . . . . . 8vo, Waterbury's Cement Laboratory Manual. . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Wood's (De V.) Treatise on the Resistance of Materials, and an Appendix on the Preservation of Timber. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Wood's (M. P.) Rustless Coatings: Corrosion and Electrolysis of Iron and Steel. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, RAILWAY ENGINEERING. Andrews's Handbook for Street Railway Engineers. . . . . .3 X5 inches, mor Berg's Buildings and Structures of American Railroads. . . . . . . . . . . . . . . 4to, |Brooks's Handbook of Street Railroad Location. . . . . . . . . . . . . . . . 16mc, mor. Butts's Civil Engineer's Field-book...... . . . . . . 16mo, mor. Crandall's Railway and Other Earthwork Tables. . . . . . . . . . . . . . . . . . . . . . .8vo, Transition Curve. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16mo, mor. * Crockett's Methods for Earthwork Computations. . . . . . . . . . . . . . . . . . . 8vo, Dredge's History of the Pennsylvania Railroad. (1879)... . . . . . . . . . . . . . Paper, Fisher's Table of Cubic Yards ...... . . . . Cardboard, Godwin's Railroad Engineers' Field book and Explorers' Guide . 16mo, mor. Hudson's Tables for Calculating the Cubic Contents of Excavations and Fºrm- bankments. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Ives and Hilts's Problems in Surveying, Railroad Surveying alıd Geodesy 16mo, mor. Molitor and Beard's Manual for Resident Engineers. . . . . . . . . . . . . . . 16mo, Nagle's Field Manual for Railroad Engineers. . . . . . . . . . . . . . . . . 16mº), mor. * Orrock's Railroad Structures and Estimates........ * > * is a & s is e a s ſº is 8vo, Philbrick's Field Manual for Engineers. . . . . . . . . . . . . . . . . . . . . . . . . 16mo, mor. 5 0 :;i O O00 : 5 O 0 0 5 O . . 50 i Raymond's Railroad Engineering. 3 volumes. Vol. I. Railroad Field Geometry. (In Preparation.) Vol. II. Elements of Railroad Engineering. . . . . . . . . . . . . . . . . . . . . 8vo, Vol. III. Railroad Engineer's Field Book. (In Preparation.) 00 00 00 3 50 9 Roberts' Track Formulae and Tables. (In Press.) Searles's Field Engineering. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16mo, mor, $3 00 Railroad Spiral. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16mo, mor. Taylor's Prismoidal Formulae and Earthwork. . . . . . . . . . . . . . . . . . . . . . . . . .8vo, * Trautwine's Field Practice of Laying Out Circular Curves for Railroads. 12mo, mor. * Method of Calculating the Cubic Contents of Excavations and Em- bankments by the Aid of Diagrams........... . . . . . . . . . . . . . . . .8vo, Webb's Economics of Railroad Construction............. . . . . . . . Large 12mo, Railroad Construction............ . . . . . . . . . . . . . . . . . . . . . . . . 16mo, mor. Wellington's Economic Theory of the Location of Railways..... Large 12mo, Wilson's Elements of Railroad-Track and Construction...... tº tº tº $ tº e º is e 12mo, DRAWING. Barr's Kinematics of Machinery.............. . . . . . . . . . . . . . . . tº e g º ſº tº º ºs e 8vo, * Bartlett's Mechanical Drawing.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, >k ſt & § { { { Abridged Ed.......... . . . . . . . . . . . . .8vo, Coolidge's Manual of Drawing. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, paper, Coolidge and Freeman's Elements of General Drafting for Mechanical Engi- TheerS. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Oblong 4to, Durley's Kinematics of Machines. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Emch’s Introduction to Projective Geometry and its Application. . . . . . 8vo, French and Ives' Stereotomy............ . . . . . . . . . . . . . . . . . . . . . . . . . . . .8vo, Hill's Text-book on Shades and Shadows, and Perspective ............. . 8vo, Jamison's Advanced Mechanical Drawing. . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Elements of Mechanical Drawing. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Jones's Machine Design: Part I. Kinematics of Machinery......... . . . . . . . . . . . . . . . . . . . . . . .8vo, Part II. Form, Strength, and Proportions of Parts.......... . . . . . . .8vo, * Kimball and Barr's Machine Design . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, MacCord's Elements of Descriptive Geometry. . . . . . . . . . . . . . . . . . . . . . . 8vo, Kinematics; or, Practical Mechanism....... . . . . . . . . . . . . . . . . . . . . .8vo, Mechanical Drawing........... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .4to, Volocity Diagrams. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, McLeod's Descriptive Geometry...... . . . . . . . . . . . . . . . . . . . . . . . Large 12mo, * Mahan's Descriptive Geometry and Stone-cutting. . . . . . . . . . . . . . . . . . 8vo, Industrial Drawing. (Thompson.)......... . . . . . . . . . . . . . . . . . . . . . .8vo, Moyer's Descriptive Geometry............ . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Reed's Topographical Drawing and Sketching... . . . . . . . . . . . . . . . . . . . . . 4to, Reid's Course in Mechanical Drawing........ . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Text-book of Mechanical Drawing and Elementary Machine Design..8vo, Robinson's Principles of Mechanism................ . . . . . . . . . . . . . .8vo, Schwamb and Merrill's Elements of Mechanism. . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Smith (A. W.) and Marx's Machine Design. . . . . . . . . . . . . . . . . . . . . . . . . . . .8vo, Smith's (R. S.) Manual of Topographical Drawing. (McMillan.)........ 8vo, * Titsworth's Elements of Mechanical Drawing. . . . . . . . . . . . . . . Oblong 8vo, Warren's Drafting Instruments and Operations. . . . . . . . . . . . . . . . . . . . 12mo, Elements of Descriptive Geometry, Shadows, and Perspective...... 8vo, Elements of Machine Construction and Drawing. . . . . . . . . . . . . . . . . . . 8vo, Elements of Plane and Solid Free-hand Geometrical Drawing. . . . 12mo, General Problems of Shades and Shadows. . . . . . . . . . . . . . . . . . . . . . . 8vo, Manual of Elementary Problems in the Linear Perspective of Forms and Shadow. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Manual of Elementary Projection Drawing. . . . . . . . . . . . . . . . . . . . 12mo, Plane Problems in Elementary Geometry. . . . . . . . . . . . . . . . . . . 12mo, Weisbach's Kinematics and Power of Transmission. (Hermann and Klein.). . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Wilson's (H. M.) Topographic Surveying. . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Wilson's (V. T.) Descriptive Geometry................ . . . . . . . . . . . . . . .8vo, Free-hand Lettering. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Free-hand Perspective. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Woolf’s Elementary Course in Descriptive Geometry. . . . . . . . . . . Large 8vo, 1 1 i2 i i i i 50 50 50 10 ELECTRICITY AND PEIYSICS. * Abegg's Theory of Electrolytic Dissociation. (von Ende.). . . . . . . . . 12mo, $1 25 Andrews's Hand-book for Street Railway Engineering. . . . .3 × 5 inches, mor. 1 25 Anthony and Brackett's Text-book of Physics. (Magie.) . . . . Large 12mo, 3 00 Anthony and Ball's Lecture-notes on the Theory of Electrical Measure- ments. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, 1 00 Benjamin's History of Electricity. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, 3 00 Voltaic Cell.......... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, 3 00 Betts's Lead Refining and Electrolysis....... . . . . . . . . . . . . . . . . . . . . . . . . 8vo, 4 00 Classen's Quantitative Chemical Analysis by Electrolysis. (Boltwood.).8vo, 3 00 * Collins's Manual of Wireless Telegraphy and Telephony. . . . . . . . . . . . 12mo, 1 50 Crehore and Squier's Polarizing Photo-chronograph.... . . . . . . . . . . . . . . . . . 8vo, 3 00 * Danneel's Electrochemistry. (Merriam.)... . . . . . . . . . . . . . . . . . . . . . . . 12mo, 1 25 Dawson’s “Engineering” and Electric Traction Pocket-book. ... 16mo, mor. 5 00 Dolezalek's Theory of the Lead Accumulator (Storage Battery). (von Ende.) 12mo, 2 50 Duhem's Thermodynamics and Chemistry. (Burgess.)...... . . . . . . . . . .8vo, 4 00 Flather's Dynamometers, and the Measurement of Power... . . . . . . . . . 12mo, 3 00 * Getman's Introduction to Physical Science.- . . . . . . . . . . . . . . . . . . . . . 12mo, 1 50 Gilbert's De Magnete. (Mottelay)............... . . . . . . . . . . . . . . . . . . . .8vo, 2 50 * Hanchett's Alternating Currents................ . . . . . . . . . . . . . . . . . . . 12mo, 1 00 Hering's Ready Reference Tables (Conversion Factors). . . . . . . . 16mo, mor. 2 50 * Hobart and Ellis's High-speed Dynamo Electric Machinery. . . . . . . . . . 8vo, 6 00 Holman's Precision of Measurements. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, 2 00 Telescopic Mirror-scale Method, Adjustments, and Tests. . . . Large 8vo, 75 * Karapetoff's Experimental Electrical Engineering.............. . . . . . . . 8vo, 6 00 Rinzbrunner's Testing of Continuous-current Machines. . . . . . . . . . . . . . . 8vo, 2 00 Landauer's Spectrum Analysis. (Tingle.)... . . . . . . . . . . . . . . . . . . . . . . . . 8vo, 3 00 Le Chatelier's High-temperature Measurements. (Boudouard—Burgess.) 12mo, 3 00. Löb's Electrochemistry of Organic Compounds. (Lorenz.)............ . .8vo, 3 00 * Lyndon's Development and Electrical Distribution of Water Power. .8vo, 3 00 * Lyons's Treatise on Electromagnetic Phenomena. Vols, I. and II. 8vo, each, 6 00 * Michie's Elements of Wave Motion Relating to Sound and Light. . . . . 8vo, 4 00 Morgan's Outline of the Theory of Solution and its Results. . . . . . . . . . 12mo, 1 00 * Physical Chemistry for Electrical Engineers. . . . . . . . . . . . . . . . . . 12mo, 1 50 * Norris's Introduction to the Study of Electrical Engineering. . . . . . . . . 8vo, 2 50 Norris and Dennison's Course of Problems on the Electrical Characteristics of Circuits and Machines. (In Press.) * Parshall and Hobart's Electric Machine Design. . . . . . . . . . . 4to, half mor, 12 50 Reagan's Locomotives: Simple, Compound, and Electric. New Edition. Large 12mo, 3 50 * Rosenberg's Electrical Engineering. (Haldane Gee—Kinzbrunner.). , 8vo, 2 00 Ryan, Norris, and Hoxie's Electrical Machinery. Vol. I. . . . . . . . . . . . . . . . 8vo, 2 50 Schapper's Laboratory Guide for Students in Physical Chemistry..... 12mo, 1 00 * Tillman's Elementary Lessons in Heat. . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, 1 50 Tory and Pitcher's Manual of Laboratory Physics. . . . . . . . . . . . Large 12mo, 2 00 Ulke's Modern Electrolytic Copper Refining.. . . . . . . . . . . . • * * * * * * * * * * * 8vo, 3 00 LAW. * Brennan's Hand-book of Useful Legal Information for Business Men. 16mo, mor. * Davis's Elements of Law............. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .8vo, * Treatise on the Military Law of United States. . . . . . . . . . . . . . . . . 8vo, * Dudley's Military Law and the Procedure of Courts-martial... Large 12mo, Manual for Courts-martial. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16mo, mor. Wait's Engineering and Architectural Jurisprudence........ . . . . . . . . . . , 8vo, Sheep, Law of Contracts. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Law of Operations Preliminary to Construction in Engineering and Architecture. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . tº º ºs e º ºs e º us tº 8vo, Sheep, 5 5 O MATHEMATICS. Baker's Elliptic Functions.............. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Briggs's Elements of Plane Analytic Geometry. (Bôcher.) . . . . . . . . . 12mo, * Buchanan's Plane and Spherical Trigonometry.... . . . . . . . . . . . . . . . . . 8vo, Byerley's Harmonic Functions. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Chandler's Elements of the Infinitesimal Calculus.... . . . . . . . . . . . . . . . 12mo, * Coffin's Vector Analysis. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Compton's Manual of Logarithmic Computations. . . . . . . . . . . . . . . . . . . 12mo, * Dickson's College Algebra. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Large 12mo, * Introduction to the Theory of Algebraic Equations. . . . . . Large 12mo, Emch’s Introduction to Projective Geometry and its Application. . . . . . 8vo, Fiske's Functions of a Complex Variable. . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Halsted's Elementary Synthetic Geometry. . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Elements of Geometry. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Rational Geometry. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Synthetic Projective Geometry. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Hancock's Lectures on the Theory of Elliptic Functions. (In Press.) Hyde's Grassmann's Space Analysis. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Johnson's (J. B.) Three-place Logarithmic Tables: Vest-pocket size, paper, * 100 copies, * Mounted on heavy cardboard, 8X10 inches, * 10 copies, Johnson's (W. W.) Abridged Editions of Differential and Integral Calculus. Large 12mo, 1 vol. Curve Tracing in Cartesian Co-ordinates. . . . . . . . . . . . . . . . . . . . . . 12mo, Differential Equations. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Elementary Treatise on Differential Calculus. . . . . . . . . . . . . . Large 12mo, Elementary Treatise on the Integral Calculus. . . . . . . . . . . . Large 12mo, * Theoretical Mechanics. ................ . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Theory of Errors and the Metliod of Least Squares.. . . . . . . . . . . . 12mo, Treatise on Differential Calculus. . . . . . . . . . . . . . . . . . . . . . . . Large 12mo, Treatise on the Integral Calculus. . . . . . . . . . . . . . . . . . . . . . . Large 12mo, Treatise on Ordinary and Partial Differential Equations. . . Large 12mo, $ 25 FCarapetoff's Engineering Applications of Higher Mathematics. (In Preparation.) 2 00 Laplace's Philosophical Essay on Probabilities. (Truscott and Emory.). 12mo, * Ludlow and Bass's Elements of Trigonometry and Logarithmic and Other Tables. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Trigonometry and Tables published separately. . . . . . . . . . . . . . . Each, * Ludlow's Logarithmic and Trigonometric Tables. . . . . . . . . . . . . . . . . . . 8vo, Macfarlane's Vector Analysis and Quaternions. . . . . . . . . . . . . . . . . . . . . . 8vo, McMahon's Hyperbolic Functions. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Manning's Irrational Numbers and their Representation by Sequences and €T16S. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . a s e e s s a In O, Mathematical Monographs. Edited by Mansfield Merriman and Robert S. Woodward. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Octavo, each No. 1. History of Modern Mathematics, by David Eugene Smith. No. 2. Synthetic Projective Geometry, by George Bruce Halsted. No. 3. Determinants, by Laenas Gifford Weld. No. 4. Hyper- bolic Functions, by James McMahon. No. 5. Harmonic Func- tions, by William E. Byerly. No. 6. Grassmann's Space Analysis, loy Edward W. Hyde. No. 7. Probability and Theory of Errors, lby Robert S. Woodward. No. 8. Vector Analysis and Quaternions, by Alexander Macfarlane. No. 9. Differential Equations, by William Woolsey Johnson. No. 10. The Solution of Equations, by Mansfield Merriman. No. 11. Functions of a Complex Variable, by Thomas S. Fiske. Maurer's Technical Mechanics. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Merriman's Method of Least Squares . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Solution of Equations. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Rice and Johnson's Differential and Integral Calculus, 2 vols, in one. Large 12mo, Elementary Treatise on the Differential Calculus. . . . . . . . . Large 12mo, Smith's History of Modern Mathematics. . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Veblen and Lennes's Introduction to the Real Infinitesimal Analysis of One Variable. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . * , s is tº a s e s is tº e º 8vo, 3 2 1 1 1. : : 00 00 50 00 00 00 * Waterbury's West Pocket Hand-book of Mathematics for Engineers. 2% X5% inches, mor, $1 00 * Enlarged Edition, Including Tables ... . . . g is g g g g º is is 6 & e s is e s e º is a II*C.r. Weld's Determinants. . . . . . . . . . . . . . . . . . . . . . . is g º ºs º º e º & e ſº e º is sº º e s e a º 8vo, Wood's Elements of Co-ordinate Geometry. . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Woodward's Probability and Theory of Errors. . . . . . . . . . . . . . g º e g & e s is 8vo, MECHANICAL ENGINEERING. 1. 1 2 1 50 00 00 00 MATERIALS OF ENGINEERING, STEAM-ENGINES AND BOILERS. Bacon's Forge Practice. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Baldwin’s Steam Heating for Buildings. . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Barr's Kinematics of Machinery. . . . . . . . . . . . . . . tº g tº º tº is e º ſº e º is tº º is º ºs º & 8vo, * Bartlett's Mechanical Drawing. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, 2k tº ſº $ & “ Abridged Ed. . . . . . . . . . . . . . . . . . . . . . . 8vo, * Burr's Ancient and Modern Engineering and the Isthmian Canal. . . . . 8vo, Carpenter's Experimental Engineering. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Heating and Ventilating Buildings. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Clerk's The Gas, Petrol and Oil Engine. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Compton's First Lessons in Metal Working. . . . . . . . . . . . . . . . . . . . . . . . 12mo, Compton and De Groodt's Speed Lathe. . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Coolidge's Manual of Drawing. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, paper, Coolidge and Freeman's Elements of General Drafting for Mechanical En- gineers. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Oblong 4to, Cromwell's Treatise on Belts and Pulleys. . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Treatise on Toothed Gearing. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Dingey's Machinery Pattern Making.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Durley's Kinematics of Machines. . . . . . . . * * tº e º 'º it tº e g g º is e º ºs e º & e s is e e s a 8vo, Flanders's Gear-cutting Machinery. . . . . . . . . . . . . . . . . . . . . . . . Large 12mo, Flather's Dynamometers and the Measurement of Power. . . . . . . . . . . . 12mo, Rope Driving. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Gill's Gas and Fuel Analysis for Engineers. . . . . . . . . . . . . . . . . . . . . . . . 12mo, Goss's Locomotive Sparks. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Greene's Pumping Machinery. (In Preparation.) Hering's Ready Reference Tables (Conversion Factors). . . . . . . . 16mo, mor. * Hobart and Ellis's High Speed Dynamo Electric Machinery. . . . . . . . . 8vo, Hutton's Gas Engine. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Jamison's Advanced Mechanical Drawing. . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Elements of Mechanical Drawing. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Jones's Gas Engine. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Machine Design: Part I. Kinematics of Machinery. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Part II. Form, Strength, and Proportions of Parts. . . . . . . . . . . . . 8vo, Rent's Mechanical Engineer's Pocket-Book. . . . . . . . . . . . . . . . . . . 16mo, mor. FCerr's Power and Power Transmission. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Kimball and Barr's Machine Design. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Levin's Gas Engine. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Leonard's Machine Shop Tools and Methods. . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Lorenz's Modern Refrigerating Machinery. (Pope, Haven, and Dean).. 8vo, MacCord's Kinematics; or, Practical Mechanism. . . . . . . . . . . . . . . . . . . . . 8vo, Mechanical Drawing. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4to, Velocity Diagrams. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, MacFarland's Standard Reduction Factors for Gases. . . . . . . . . . . . . . . . . 8vo, Mahan's Industrial Drawing. (Thompson.). . . . . . . . . . . . . . . . . . . . . . . 8vo, Mehrtens's Gas Engine Theory and Design. . . . . . . . . . . . . . . . . . . Large 12mo, Oberg’s Handbook of Small Tools. . . . . . . . . . . . . . . . . . . . . . . . . . Large 12mo, * Parshall and Hobart's Electric Machine Design. Small 4to, half leather, Peele's Compressed Air Plant for Mines. . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Poole's Calorific Power of Fuels. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Porter's Engineering Reminiscences, 1855 to 1882. . . . . * - - - - - - - - - - - - 8vo, Reid's Course in Mechanical Drawing. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Text-book of Mechanical Drawing and Elementary Machine Design,8vo, i 13 Richards's Compressed Air. . . . . . . . . . . . . . . . # * * * * * s tº t e s s a e s w w e s a w e 12mo, $ Robinson's Principles of Mechanism. . . . . . . . . . . . * * * * * g g º g g g º 'º ºn e º g º º 8vo, Schwamb and Merrill's Elements of Mechanism. . . . . . . . . . . . . . . . . . . . . 8vo, Smith (A. W.) and Marx's Machine Design. . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Smith's (O.) Press-working of Metals. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Sorel's Carbureting and Combustion in Alcohol Engines. (Woodward and Preston.). . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Large 12mo, Stone's Practical Testing of Gas and Gas Meters. . . . . . . . . . . . . . . . . . . . . 8vo, Thurston's Animal as a Machine and Prime Motor, and the Laws of Energetics. 12mo, Treatise on Friction and Lost Work in Machinery and Mill Work. . . 8vo, * Tillson's Complete Automobile Instructor. . . . . . . . . . . . . . . . . . . . . . . 16mo, * Titsworth's Elements of Mechanical Drawing. . . . . . . . . . . . . . . Oblong 8vo, Warren's Elements of Machine Construction and Drawing. . . . . . . . . . . . 8vo, * Waterbury's West Pocket Hand-book of Mathematics for Engineers. 23 X53 inches, mor. * Enlarged Edition, Including Tables. . . . . . . . . . . . . . . . . . . . . . . . . . ITY OT. Weisbach's Kinematics and the Power of Transmission. (Herrmann– Klein.). . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Machinery of Transmission and Governors. (Hermann–Klein.). , 8vo, Wood's Turbines. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, MATERIALS OF ENGINEERING. * Bovey's Strength of Materials and Theory of Structures. . . . . . . . . . . . 8vo, Burr's Elasticity and Resistance of the Materials of Engineering. . . . . . . 8vo, Church's Mechanics of Engineering. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Greene's Structural Mechanics. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Holley's Lead and Zinc Pigments. . . . . . . . . . . . . . . . . . . . . . . . . . Large 12mo Holley and Ladd's Analysis of Mixed Paints, Color Pigments, and Varnishes. Large 12mo, Johnson's (C. M.) Rapid Methods for the Chemical Analysis of Special Steels, Steel-Making Alloys and Graphite . . . . . . . . . . . Large 12mo, Johnson's (J. B.) Materials of Construction. . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Keep's Cast Iron. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Lanza's Applied Mechanics. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Maire's Modern Pigments and their Vehicles. . . . . . . . . . . . . . . . . . . . . . . 12mo, Maurer's Technical Mechanics. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Merriman's Mechanics of Materials. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Strength of Materials. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Metcalf's Steel. A Manual for Steel-users. . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Sabin's Industrial and Artistic Technology of Paint and Varnish. . . . . . 8vo, Smith's ((A. W.) Materials of Machines. . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, * Smith's (H. E.) Strength of Material... . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Thurston's Materials of Engineering. . . . . . . . . . . . . . . . . . . . . . . 3 vols., 8vo, Part I. Non-metallic Materials of Engineering, . . . . . . . . . . . . . . . . 8vo, Part II. Iron and Steel. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Part III. A Treatise on Brasses, Bronzes, and Other Alloys and their Constituents. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Wood's (De V.) Elements of Analytical Mechanics. . . . . . . . . . . . . . . . . . . 8vo, Treatise on the Resistance of Materials and an Appendix on the Preservation of Timber. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Wood's (M. P.) Rustless Coatings: Corrosion and Electrolysis of Iron and Steel. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, STEAM-ENGINES AND BOILERS. Berry's Temperature-entropy Diagram. . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Carnot's Reflections on the Motive Power of Heat. (Thurston.). . . . . 12mo, Chase's Art of Pattern Making. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, l b 0. :;i;: 0 0. :: 0 0. : O 0 Creighton's Steam-engine and other Heat Mötors. . . . . . . . . & & g g g g g g º is & 8vo, $5 00 Dawson's “Engineering” and Electric Traction Pocket-book. ... 16mo, mor. * Gebhardt's Steam Power Plant Engineering. . . . . . . . . . . . . . . . . . . . . . . 8vo, Goss's Locomotive Performance. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Hemenway's Indicator Practice and Steam-engine Economy. . . . . . . . . 12mo, Hutton's Heat and Heat-engines. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Mechanical Engineering of Power Plants. . . . . . . . . . . . . . . . . . . . . . . 8vo, Rent's Steam boiler Economy. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Rneass's Practice and Theory of the Injector. . . . . . . . . . . . . . . . . . . . . . . 8vo, MacCord's Slide-valves. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Meyer's Modern Locomotive Construction. . . . . . . . . . . . . . . . . . . . . . . . . . . 4to, Moyer's Steam Turbine. . . . . . tº º is tº s p & e s g º a 9 s = p s is e s & 6 s is tº & e º s e s e s tº a s 8vo, Peabody's Manual of the Steam-engine Indicator. . . . . . . . . . . . . . . . . . 12mo, Tables of the Properties of Steam and Other Vapors and Temperature- Entropy Table. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Thermodynamics of the Steam-engine and Other Heat-engines. . . . 8vo, Valve-gears for Steam-engines. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Peabody and Miller's Steam-boilers. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Pupin's Thermodynamics of Reversible Cycles in Gases and Saturated Vapors. (Osterberg.). . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Reagan's Locomotives: Simple, Compound, and Electric. New Edition. Large 12mo, Sinclair's Locomotive Engine Running and Management. . . . . . . . . . . . 12mo, Smart's Handbook of Engineering Laboratory Practice. . . . . . . . . . . . . 12mo, Snow's Steam-boiler Practice. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Spangler's Notes on Thermodynamics. . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Valve-gears. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Spangler, Greene, and Marshall's Elements of Steam-engineering. . . . . . 8vo, Thomas's Steam-turbines. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Thurston's Handbook of Engine and Boiler Trials, and the Use of the Indi- cator and the Prony Brake. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Handy Tables. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Manual of Steam-boilers, their Designs, Construction, and Operation 8vo, Manual of the Steam-engine. . . . . . . . . . . . . . . . . . . . . . . . . . . 2 vols., 8vo, Part I. History, Structure, and Theory . . . . . . . . . . . . . . . . . . 8vo, Part II. Design, Construction, and Operation . . . . . . . . . . . . . . 8vo, Wehrenfennig's Analysis and Softening of Boiler Feed-water. (Patterson ) 8vo, Weisbach's Heat, Steam, and Steam-engines. (Du Bois.). . . . . . . . . . . . 8vo, Whitham's Steam-engine Design. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Wood's Thermodynamics, Heat Motors, and Refrigerating Machines. . .8vo, MECHANICS PURE AND APPLIED. Church's Mechanics of Engineering. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Notes and Examples in Mechanics. . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Dana's Text-book of Elementary Mechanics for Colleges and Schools . 12mo, Du Bois's Elementary Principles of Mechanics: Vol. I. Kinematics. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo Vol. II. Statics. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Mechanics of Engineering, Vol. I. . . . . . . . . . . . . . . . . . . . . . . Small 4to, Vol. II. . . . . . . . . . . . . . . . . . . . . . . Small 4to, * Greene's Structural Mechanics. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo James's Kinematics of a Point and the Rational Mechanics of a Particle. Large 12mo, * Johnson's (W. W.) Theoretical Mechanics. . . . . . . . . . . . . . . . . . . . . . . 12mo, Lanza's Applied Mechanics. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Martin's Text Book on Mechanics, Vol. I, Statics. . . . . . . . . . . . . . . . . 12mo, * Vol. II, Kinematics and Kinetics. 12mo, Maurer's Technical Mechanics. . . . . . . . . . . . . . . . . . . . . . & s • * * * * * * * * * * * * VO, * Merriman's Elements of Mechanics. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo Mechanics of Materials. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Michie's Elements of Analytical Mechanics. . . . . . . . . . . . . . . . . . . . . . . . 8vo, 15 ./ y * sº 1 : l 1 i: : 1 i Robinson's Principles of Mechanism. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Sanborn's Mechanics Problems. . . . . . . . . . . . . . . . . . . . . . . . . . . . . Large 12mo, Schwamb and Merrill's Elements of Mechanism. . . . . . . . . . . . . . . . . . . . . 8vo, Wood's Elements of Analytical Mechanics. . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Principles of Elementary Mechanics. . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, MEDICAL. * Abderhalden's Physiological Chemistry in Thirty Lectures. (Hall and Defren). . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, von Behring's Suppression of Tuberculosis. (Bolduan.). . . . . . . . . . . . 12mo, Bolduan's Immune Sera. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Bordet's Studies in Immunity. (Gay.) . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Chapin's The Sources and Modes of Infection. (In Press.) Davenport's Statistical Methods with Special Reference to Biological Varia- tions. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16mo, mor. Ehrlich's Collected Studies on Immunity. (Bolduan.). . . . . . . . . . . . . . . 8vo, * Fischer's Physiology of Alimentation. . . . . . . . . . . . . . . . . . . . . . Large 12mo, de Fursac's Manual of Psychiatry. (Rosanoff and Collins.). . . . Large 12mo, Hammarsten's Text-book on Physiological Chemistry. (Mandel.). . . . . 8vo, Jackson's Directions for Laboratory Work in Physiological Chemistry. .8vo, Lassar-Cohn's Practical Urinary Analysis. (Lorenz.). . . . . . . . . . . . . . . 12mo, Mandel's Hand-book for the Bio-Chemical Laboratory. . . . . . . . . . . . . . 12mo, * Nelson's Analysis of Drugs and Medicines . . . . . . . . . . . . . . . 12mo, * Pauli’s Physical Chemistry in the Service of Medicine. (Fischer.) ..12mo, * Pozzi-Escot's Toxins and Venoms and their Antibodies. (Cohn.). . 12mo, Rostoski's Serum Diagnosis. (Bolduan.). . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Ruddiman's Incompatibilities in Prescriptions. . . . . . . . . . . . . . . . . . . . . . . 8vo, Whys in Pharmacy. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Salkowski's Physiological and Pathological Chemistry. (Orndorff.) . . . .8vo, * Satterlee's Outlines of Human Embryology. . . . . . . . . . . . . . . . . . . . . . 12mo, Smith's Lecture Notes on Chemistry for Dental Students. . . . . . . . . . . . . 8vo, * Whipple's Tyhpoid Fever. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Large 12mo, * Woodhull's Military Hygiene for Officers of the Line . . . . . . . . Large 12mo, * Personal Hygiene. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Worcester and Atkinson's Small Hospitals Establishment and Maintenance, and Suggestions for Hospital Architecture, with Plans for a Small Hospital. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, METALLURGY. Betts's Lead Refining by Electrolysis. . . . . . . . . . . . . . . . . . . . . . * * @ e º ſº e º s 8vo, Bolland's Encyclopedia of Founding and Dictionary of Foundry Terms used in the Practice of Moulding. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Iron Founder. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, & ºt 4 & Supplement. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Douglas's Untechnical Addresses on Technical Subjects. . . . . . . . . . . . . 12mo, Goesel's Minerals and Metals: A Reference Book. . . . . . . . . . . . . . 16mo, mor. * Iles's Lead-smelting. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Johnson's Rapid Methods for the Chemical Analysis of Special Steels, Steel-making Alloys and Graphite. . . . . . . . . . . . . . . . . . Large 12mo, Keep's Cast Iron. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Le Chatelier's High-temperature Measurements. (Boudouard–Burgess.) 12mo, Metcalf's Steel. A Manual for Steel-users. . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Minet's Production of Aluminum and its Industrial Use. (Waldo.). . 12mo, * Ruer's Elements of Metallography. (Mathewson.) . . . . . . . . . . . . . . . . . 8vo, Smith's Materials of Machines. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Tate and Stone's Foundry Practice. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Thurston's Materials of Engineering. In Three Parts. . . . . . . . . . . . . . . . 8vo, Part I. Non-metallic Materials of Engineering, see Civil Engineering, page 9. Part II. Iron and Steel. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Part III. A Treatise on Brasses, Bronzes, and Other Alloys and their Constituents. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, :; 4li 2 Ulke's Modern Electrolytic Copper Refining. . . . . . . . . . . . . . . . . . . . . . . . . 8vo, $3 00: West's American Foundry Practice. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Moulders' Text Book. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, MINERALOGY. Baskerville's Chemical Elements. (In Preparation.) * Browning's Introduction to the Rarer Elements. . . . . . . . . . . . . . . . . . . 8vo, Brush's Manual of Determinative Mineralogy. (Penfield.). . . . . . . . . . . . 8vo, Butler's Pocket Hand-book of Minerals. . . . . . . . . . . . . . . . . . . . . . 16mo, mor. Chester's Catalogue of Minerals. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, paper, Cloth, * Crane's Gold and Silver. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Dana’s First Appendix to Dana's New “System of Mineralogy”. . Large 8vo, Dana’s Second Appendix to Dana's New “System of Mineralogy.” Large 8vo, Manual of Mineralogy and Petrography. . . . . . . . . . . . . . . . . . . . . . . 12mo, Minerals and How to Study Them. . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, System of Mineralogy. . . . . . . . . . . . . . . . . . . . . . Large 8vo, half leather, 1 Text-book of Mineralogy. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Douglas's Untechnical Addresses on Technical Subjects. . . . . . . . . . . . . 12mo, Eakle's Mineral Tables. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Eckel's Stone and Clay Products Used in Engineering. (In Preparation.) Goesel's Minerals and Metals: A Reference Book. . . . . . . . . . . . . . 16mo, mor. Groth's The Optical Properties of Crystals. (Jackson.) (In Press.) Groth's Introduction to Chemical Crystallography (Marshall). . . . . . . . 12mo, * Hayes's Handbook for Field Geologists. . . . . . . . . . . . . . . . . . . . 16mo, mor. Iddings's Igneous Rocks. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Rock Minerals. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Johannsen's Determination of Rock-forming Minerals in Thin Sections. 8vo, With Thumb Index * Martin's Laboratory Guide to Qualitative Analysis with the Blow- pipe. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Merrill's Non-metallic Minerals: Their Occurrence and Uses. . . . . . . . . . . 8vo, Stones for Building and Decoration. . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Penfield's Notes on Determinative Mineralogy and Record of Mineral Tests. 8vo, paper, Tables of Minerals, Including the Use of Minerals and Statistics of Domestic Production. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Pirsson's Rocks and Rock Minerals. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, * Richards's Synopsis of Mineral Characters. . . . . . . . . . . . . . . . . . 12mo, mor. * Ries's Clays: Their Occurrence, Properties and Uses. . . . . . . . . . . . . . . . 8vo, * Ries and Leighton’s History of the Clay-working Industry of the United States. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Tillman's Text-book of Important Minerals and Rocks. . . . . . . . . . . . . . 8vo, Washington's Manual of the Chemical Analysis of Rocks. . . . . . . . . . . . . 8vo, MINING. * Beard's Mine Gases and Explosions. . . . . . . . . . . . . . . . . . . . . . . Large 12mo, * Crane's Gold and Silver. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Index of Mining Engineering Literature. . . . . . . . . . . . . . . . . . . . . . . 8vo, Ore Mining Methods. (In Press.) Douglas's Untechnical Addresses on Technical Subjects. . . . . . . . . . . . . 12mo, Eissler's Modern High Explosives. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Goesel's Minerals and Metals: A Reference Book. . . . . . . . . . . . . . 16mo, mor. Ihlseng's Manual of Mining. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Iles's Lead Smelting. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Peele's Compressed Air Plant for Mines. . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Riemer's Shaft Sinking Under Difficult Conditions. (Corning and Peele.)8vo, * Weaver's Military Explosives. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Wilson's Hydraulic and Placer Mining. 2d edition, rewritten. . . . . . . 12mo, Treatise on Practical and Theoretical Mine Ventilation . . . . . . . . 12mo, i O 0 O5 0.0 :5i OO 0.O i 5 O : O 0 : O O 5 O 17 SANITARY SCIENCE. Association of State and National Food and Dairy Departments, Hartford Meeting, 1906. . . . . . . . . . . . . . . . . . . . . . . . . . . . tº $ tº º ſº a n w is is º e º e & 8vo, Jamestown Meeting, 1907. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Bashore's Outlines of Practical Sanitation, . . . . . . . . . . . . . . . . . . . . . . 12mo, Sanitation of a Country House. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Sanitation of Recreation Camps and Parks. . . . . . . . . . . . . . . . . . . . 12mo, Chapin's The Sources and Modes of Infection. (In Press.) Folwell's Sewerage. (Designing, Construction, and Maintenance.). . . . . 8vo, Water-supply Engineering . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Fowler's Sewage Works Analyses. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Fuertes's Water-filtration Works. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Water and Public Health. . . . . . . . . . . . . . : - - - - - - - - - , , . . . . . . . . . 12mo, Gerhard's Guide to Sanitary Inspections. . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, * Modern Baths and Bath Houses. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Sanitation of Public Buildings. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, * The Water Supply, Sewerage, and Plumbing of Modern City Buildings. 8vo, |Hazen's Clean Water and How to Get It. . . . . . . . . . . . . . . . . . . Large 12mo, Filtration of Public Water-supplies. . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Rinnicut, Winslow and Pratt's Purification of Sewage. (In Preparation.) Leach's Inspection and Analysis of Food with Special Reference to State Control. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Mason's Examination of Water. (Chemical and Bacteriological), . . . . 12mo, Water-supply. (Considered principally from a Sanitary Standpoint). 8vo, * Merriman's Elements of Sanitary Engineering. . . . . . . . . . . . . . . . . . . . . 8vo, Ogden's Sewer Construction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Sewer Design. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Parsons's Disposal of Municipal Refuse. . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Prescott and Winslow's Elements of Water Bacteriology, with Special Refer- ence to Sanitary Water Analysis. . . . . . . . . . . . . . . . . . . . . . . . . 12mo, * Price's Handbook on Sanitation. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Richards's Cost of Cleanness. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Cost of Food. A Study in Dietaries. . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Cost of Living as Modified by Sanitary Science. . . . . . . . . . . . . . . . 12mo, Cost of Shelter. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, * Richards and Williams's Dietary Computer. . . . . . . . . . . . . . . . . . . . . . . 8vo, Richards and Woodman's Air, Water, and Food from a Sanitary Stand- Point. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Richey's Plumbers', Steam-fitters', and Tinners' Edition (Building Mechanics' Ready Reference Series). . . . . . . . . . . . . . . . . 16mo, mor. Rideal's Disinfection and the Preservation of Food. . . . . . . . . . . . . . . . . . 8vo, Sewage and Bacterial Purification of Sewage. . . . . . . . . . . . . . . . . . . . 8vo, Soper's Air and Ventilation of Subways. . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Turneaure and Russell's Public Water-supplies. . . . . . . . . . . . . . . . . . . . . . 8vo, Venable's Garbage Crematories in America. . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Method and Devices for Bacterial Treatment of Sewage. . . . . . . . . . 8vo, Ward and Whipple's Freshwater Biology. (In Press.) Whipple's Microscopy of Drinking-water. . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, * Typhoid Fever. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Large 12mo, Value of Pure Water. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Large 12mo, Winslow's Systematic Relationship of the Coccaceae. . . . . . . . . . . Large 12mo, MISCELLANEOUS. I’mmons's Geological Guide-book of the Rocky Mountain Excursion of the International Congress of Geologists. . . . . . . . . . . . . . . . . . Large 8vo Ferrel's Popular Treatise on the Winds. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Fitzgerald's Boston Machinist. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 18mo, Gannett's Statistical Abstract of the World. . . . . . . . . . . . . . . . . . . . . . . 24mo, Haines's American Railway Management. . . . . . . . . . . . . . . . . . . . . . . . . 12mo, Hanausek's The Microscopy of Technical Products. (Winton) . . . . . . . 8vo, 18 $ : iſ; 2 :i : : Jacobs's Betterment Briefs. A Collection of Published Papers on Or- ganized Industrial Efficiency. . . . . . . . . . . . . . . . . * * * * * * * * * * * * * 8vo, Metcalfe's Cost of Manufactures, and the Administration of Workshops..8vo, Putnam's Nautical Charts. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8vo, Ricketts's History of Rensselaer Polytechnic Institute 1824–1894. Large 12mo, Rotherham's Emphasised New Testament. . . . . tº e º e s tº a s is is a s s a e Large 8vo, Rust’s Ex-Meridian Altitude, Azimuth and Star-finding Tables. . . . . . . . 8vo, Standage's Decoration of Wood, Glass, Metal, etc. . . . . . . . . . . . . . . . . . 12mo, Thome's Structural and Physiological Botany. (Bennett). . . . . . . . . . 16mo, Westermaier's Compendium of General Botany. (Schneider). . . . . . . . . 8vo, Winslow's Elements of Applied Microscopy. . . . . . . . . . . . . . . . . . . . . . . . 12mo, HEBREW AND CHALDEE TEXT-BOOOKS. Gesenius's Hebrew and Chaldee Lexicon to the Old Testament Scriptures. (Tregelles.) . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Small 4to, half mor, Green's Elementary Hebrew Grammar. . . . . . . . . . . . . . . . . . . . . . . . . . . . 12mo, 19 $3 30 00 00 Oſ) 00 OO 00 25 ()0 50 803 5 t M i 4705 N ·Oſ= ~~ ~） - -| == co ！é -{-* "ż , õ=&5°、 ! ===）---- ~~~~** ***~ · :*=♥--- -===== -·№t=='uſº . ER: i. i ****<！--~--~~~~ ： ºrºso e º aedºſ e, e e - z ： ： ：Caeſa •· №* . ：：: * ·º a ſe ſſae ·g * ſ; º[X_' ）， 'ºſººººſº[…]ºſº.| ''****[z ºº º |-tº eſ، ، ، ſº º 8 ° CC• • • • • b ， º º © ： ， : e 4 × ± • • • * - | • ¡ ¿ g ∞∞∞ zo • • • • • • • • ¡ ¿ ºg � : ~） suit, *** ſj • • • • • • • • ſ ≡ - .ae » ， ， ， ， º º £© ºŒ ſ º.ſe s ſ Caesae º ·ſ º• ae aeſſºr • • • • • • • • • •C.， Cº ：： .ec • • •¿•••• •ú (Dºº * * Œ œ • • ſ) → Clſ- . Caeº，· _■ * : ºſ ºg í • • • •