Cornell University Library The original of this book is in the Cornell University Library. There are no known copyright restrictions in the United States on the use of the text. http://www.archive.org/details/cu31924002975757 Cornell University Library QD 81.P95 1907 Qualitative chemical a na| y sis ;, a 9" i l f) e ll ! n 3 1924 002 975 757 QUALITATIVE CHEMICAL ANALYSIS A GUIDE IN QUALITATIVE WORK, WITH DATA FOR ANALYTICAL OPERATIONS AND LABORATORY METHODS IN INORGANIC CHEMISTRY. The D. Van Nostrand Company intend this booh to be sold to the Public at the advertised price, and supply it to the Trade on terms which will not allow of discount. *, NG NEW YORK: D. VAN NOSTRAND COMPANY 23 Murray and 27 Warren Sts. ■ 1907 QUALITATIVE CHEMICAL ANALYSIS: A GUIDE IN QUALITATIVE WORK, WITH DATA FOR ANALYTICAL' OPERATIONS AND LABORATORY METHODS IN INORGANIC CHEMISTRY. BY ALBERT B. gRESCOTT, AND OTIS C. JOHNSON, PROFESSORS IN THE UNIVERSITY OF MICHIGAN. SIXTH REVISED AND ENLARGED EDITION, ENTIRELY REWRITTEN. WITH AN APPENDIX BY H. H. WILLARD CONTAINING A FEW IMPROVED METHODS OF ANALYSIS. NEW YORK: D. VAN NOSTRAND COMPANY 23 Murray and 27 Warren Sts. • 1907 Copyrighted 1901, by D. VAN NOSTRAND COMPANY. PREFACE. In this, the fifth full revision of this manual, the text has been rewritten and the order of statement in good part recast. The subject- matter is enlarged by fully one-half, though but one hundred pages have been added to the book. It has been our aim to bring the varied resources of analysis within reach, placing in order before the worker the leading characteristics of elements, upon the relations of which every scheme of separation de- pends. This is desired for the working chemist, and no less for the working student. However limited may be the range of his work, we would not contract his view to a single routine. It is while in the course of qualitative analysis especially that the student is forming his personal acquaintance with the facts of chemical change, and it is not well that his outlook should be cut off by narrow routine at this time. The introductory pages upon Operations of Analysis, setting forth some of the foundations of qualitative chemistry, consist of matter restored and revised from the editions of 1874 and 1880. This subject- matter, omitted in 1888, is now desired by teachers. For the portion upon Solution and Ionization, we are indebted to Dr. Eugene C. Sulli- van, a pupil of Professor Ostwald, now teaching qualitative analysis. The pages upon the Periodic System have been added to afford a more connected comparison of the elements than that undertaken in each group by itself, in previous editions, and referred to in the preface in 1874. The use of notation with negative bonds, in balancing equations for changes of oxidation, introduced by one of the authors in 1880, has been retained substantially as in the last edition. Other authors adopt the same notation with various modifications. For the present revision there has been a general search of literature, and authorities are given for what is less commonly known or more deserving of further IV PREFACE. inquiry. The number of citations is so large that to save room special abbreviation is resorted to. For convenient reference, on the part of teachers, students and analysts using the book, the section for each element and each acid is arranged in uniform divisions. For instance, in each section, solu- bilities are given in paragraph 5, the action of alkalis in paragraph 6a, the action of sulphur compounds in paragraph 6e, etc. In the para- graph (9) for estimations it should be said, nothing more than a general statement of methods is given, for the benefit of qualitative study, with- out directions and specifications for quantitative work, in which, of course, other books must be used. * The authors desire to say with the fullest appreciation that Perry F. Trowbridge, instructor in Organic Chemistry in this University, has performed a large amount of labor in this revision, collecting data from original authorities, confirming their conclusions by his own experi- ments, elaborating material, and making researches upon questions as they have arisen. University of Michigan, April, 1 90 1. CONTENTS. PART I.— THE PRINCIPLES OF ANALYTICAL CHEMISTRY. PAGE The Chemical Elements and their Atomic Weights 1 Table op the Periodic System of the Chemical Elements 3 Discussion op the Periodic System 3 Classification op the Metals as Bases 10 Commonly Occurring Acids 13 The Operations op Analysis 13 Solution and Ionization SO Order op Laboratory Study 24 PART II THE METALS. THE SILVER AND TIN AND COPPER GROUPS. (FIRST AND SECOND GROUPS). General Discussion 27 THE SILVER GROUP (FIRST GROUP). ^Iiead 29 >Mercury 37 Silver 45 Comparison of Certain Reactions of the Metals of the Silver Group 51 Table por Analysis op the Silver or First Group 52 Directions for Analysis with Notes 53 THE TIN AND COPPER GROUP (SECOND GROUP). THE TIN GROUP, OR SECOND GROUP, DIVISION A. Arsenic 56 ,. Antimony 72 THn 82 Comparison of Certain Reactions of Arsenic, Antimony and Tin. 90 Gold 91 Platinum 93 Molybdenum 97 THE COPPER GROUP, OR GROUP II, DIVISION B. \ Bismuth 100 . Copper 104 MJadmium 110 Comparison of Certain Reactions of Bismuth, Copper and Cad- mium 112 VI CONTENTS. PAGE The Precipitation op the Metals op the Second Group 113 Table por the Analysis of the Tin Group (Second Group, Division A). 116 Directions for Analysis with. Notes 118 Table por Analysis op the Copper Group (Second- Group, Division B). . 124 Directions for Analysis with. Notes 126 RARER METALS OP THE TIN AND COPPER GROUP. Ruthenium 129 Rhodium 130 Palladium 131 Iridium 132 Osmium 133 Tungsten 134 Vanadium 13 5 Germanium 136 Tellurium 137 Selenium 138 THE IRON AND ZINC GROUPS (Third and Fourth Groups) 140 THE IRON GROUP (THIRD GROUP). * Aluminum 142 Chromium 1 47 Iron 151 Table for Analysis op the Iron Group (Third Group) 160 Directions for Analysis with Notes 161 the zinc group (fourth group). Cobalt 163 Nickel 168 ■O Manganese 172 Zinc 178 Comparison of Some Reactions of the Iron and Zinc Group Bases 182 Table for the Analysis of the Zinc Group (Fourth Group) 183 Directions for Analysis with Notes 1&4 Analysis of Iron and Zinc Groups after Precipitation by Ammonium Sulphide 186 Iron and Zinc Groups in Presence of Phosphates 188 Iron and Zinc Groups in Presence of Oxalates 189 Table of Separation of Iron, Zinc and Calcium Group Metals and Phosphoric Acid by Means of Alkali Acetate and Ferric Chloride 191 Table of Separation of Iron, Zinc and Calcium Group Metals and Phosphoric Acid by Means of Ferric Chloride and Barium Carbonate 192 the rarer metals op the iron and zinc groups. Cerium 193 Columbium (Niobium) 193 Didymium 194 Erbium 195 V CONTENTS. vil PAGE Gallium 195 Oluoinum (Beryllium) 195 Indium 196 Lanthanum 197 Neodymium 197 Praseodymium 197 Samarium 197 Scandium 198 Tantalum , . 198 Terbium 198 Thallium 199 Thorium 199 Titanium 200 Uranium 201 Ytterbium 202 Yttrium 202 Zirconium 202 The Calcium Group (Fifth Group). (The Alkaline Earth Metals) 203 Barium 205 Strontium 208 Calcium 210 Magnesium 214 Table for the Analysis of the Calcium Group (Fifth Group) 217 Direction for Analysis with Notes 218 Separation of Barium, Strontium, and Calcium by the Use of Alcohol 220 Alkaline Earth Metals as Phosphates 220 Alkaline Earth Metals as Oxalates 220 The Alkali Group (Sixth Group) 221 Potassium 222 Sodium 226 . Ammonium 229 ^ Caesium 233 Rubidium 234 Iiithium 234 Directions for Analysis with Notes 236 PAET III.— THE NON-METALS. Balancing of Equations 238 Hydrogen 243 Boron 245 Boric Acid 245 Carbon 247 Acel ic Acid 249 Citric Acid 251 Tartaric Acid 252 Carbon Monoxide 254 Oxalic Acid 255 Carbon Dioxide (Carbonates) 259 Till CONTENTS. PASS. Cyanogen . ,. 263 Hydrocyanic Acid 263 Hydroferrocyanic Acid 267 Hydroferricyanic Acid 269 Cyanic Acid , 271 Thiocyanic Acid 272- Nitrogen 273 Hydronitric Acid 274 .Nitrous Oxide 275. Nitric Oxide 275- Nitrons Acid 276 Nitrogen Peroxide 277' Nitric Acid 277 Oxygen 282 Ozone 284 Hydrogen Peroxide 285 Fluorine , • 288- Hydrofluoric Acid 289 Fluosilicic Acid 289 Silicon 290 Silicic Acid 290' Phosphorus 292 Phosphine 295- Hypophosphorous Acid 295 Phosphorous Acid • 297 Hypophosphoric Acid 298- Phosphoric Acid 298 Sulphur 304 Hydrosulphuric Acid 306 Thiosulphuric Acid 312 Hyposulphurous Acid 314 Dithionic Acid 314 Trithionic Acid ' 315- Tetrathionic Acid 315 Pentathionic Acid 316- Tablr of Thionic Acids 317 Sulphurous Acid 318 Sulphuric Acid 321 Persulphuric Acid 326 Chlorine 327 Hydrochloric Acid 330' Hypochlorous Acid 337 Chlorous Acid 337 Chlorine Peroxide 338 Chloric Acid 339. Perchloric Acid 341 Bromine 342- Hydrobromic Acid 345. Hypobromous Acid 348. CONTENTS. ir PAGE Bromic Acid 348 Iodine 350' Hydriodic Acid 353 Iodic Acid 357 Periodic Acid 360 Comparative Reactions or the Halogen Compounds 361 PART IT.— SYSTEMATIC EXAMINATIONS. Removal op Orqanic Substances 36S- Preliminary Examination of Solids 363 Conversion op Solids into Liquids 366 Conversion op Solutions into Solids 367 Treatment of a Metal or an Allot 367 Separation of Acids from Bases 368 Table for Preliminary Examination of Solids 370 Behavior op Substances Before the Blow-Pipe 374 Table of ^the Grouping of the Metals 375 Table for the Separation op the Metals 376 Acids — First Table 378 Acids — Second Table 386- Acids— Third Table 387 Acids— Fourth Table 388. Notes on the Detection of Acids 389' Principles 393 Equations 396- Problems in Synthesis 397 Table of Solubilities 398 Reagents ■ 403- ABBREVIATIONS. A A. Ch. Am. Am. S. Arcb. Pharm. Am, Chem. B. Bl. B. J. Comey. C. N. Ch. Z. C. r. C. C. Dingl. D. Fehling. Fresenius. G. O. Gazzetta. Gill). Gmelin-Krant. J. J. O. J. pr. J. Soc. Ind. JT. Anal. J. Am. Soc. J. Pharm. Ladenburg, M. Phil. Mag. Pogg. Proc. Roy. Soc. Pharm. J. Trans. Ph. O. Tr. "Watt's. * Indicates continuance to the present time. Liebig's Annalen. 1832* Annales de Cbimie et de Physique. 1789* American Chemical Journal. 1879* American Journal of Science. 1818* Analyst. 1876* Archives der Pharmacie. 1822* American Chemist. 1870-77. Berichte der Deutschen Chemischen Gesellschaft. 1868* Bulletin de la Societe Chimique. 1859* Berzelius Jahresbericht. 1822-51. Comey's Dictionary of Solubilities. 1896. Chemical News. 1860* Chemiker Zeitung. 1877* Comptes Kendus des Seances de l'Academie des Sciences. 1835* Chemisches Centralblatt. 1830* Dingier' s Poly technische. Journal. 1820* Dammer's Anorganische Chemie. 1892* Fehling's Handbuch der Chemie. 1871* Fresenius: Qualitative Chemical Analysis. Graham-Otto: Lehrbuch der anorganischen Chemie. Gazzetta chimica italiana. 1871* Gilbert's Annalen der Physik und Chemie. 1799-1824. Gmelin-Kraut: Handbuch der anorganischen Chemie. 1877. Jahresbericht fiber die Fortschritte der Chemie. 1847* Journal of the Chemical Society. 1849* Journal fur praktische Chemie. 1834* Journal of the Society of Chemical Industry. 1882* Journal of Analytical Chemistry. 1887-1893. Journal of the American Chemical Society. 1876* Journal de Pharmacie et de Chimie. 1809* Handworterbuch der Chemie. 1882-1895. Monatshefte fiir Chemie. 1880* Menschutkin. Locke's Translation. 1895. Philosophical Magazine. 1798* Poggendorff's ADnalen der Physik und Chemie. 1824-1877. Proceedings of the Royal Society of London. 1832* Pharmaceutical Journal and Transactions. 1841* Pharmaceutische Centralhalle. 1859* Transactions of the Royal Society. 1665* Watt's Dictionary of Chemistry. 1888. Wells' Trans., 1897. 1885. ABBREVIATIONS. Xl W. A. Wiedemann's Annalen. 1877* W. A. (Beibl.) Wiedemann's Annalen Beiblatter. 1877* "Wormley. Wormley's Microchemistry of Poisons. 1867. Wurtz. Dictionnaire de Chimie. 1868. Z. Zeitschrift fur analytische Chemie. 1862.* Z. Ch. Zeitschrift fiir Chemie. 1865-1871. Z. anorg. Zeitschrift fiir anorganische Chemie. 1891* Z. angew. Zeitschrift fiir angewandte Chemie. 1888* Z. phys. Ch. Zeitschrift fiir physicalische Chemie. 1887* PAET I. THE PRINCIPLES OF ANALYTICAL CHEMISTRY. §1. The Chemical Elements and their Atomic Weights. Name. Sym- bol. 0=16. H=l. Name. Sym- bol. 0=16. H=l. Aluminum Al 27.1 26.9 Neodymium. . . Nd 143.6 142.5 Sb A 120.2 39.9 119.3 39.6 Neon Ne Ni 20. 58.7 19.9 58.3 As 75.0 74.4 N 14.04 ■ 13.93 Ba 137.4 136.4 Os 19L 189.6 Bl 208.5 206.9 O 16.00 15.88 B 11.0 10.9 Pd 106.5 105.7 Br 79.96 79.36 Phosphorus P 31.0 30.77 Cadmium Cd 112.4 111.6 Pt 194.8 193.3 •Caesium Cs 132.9 131.9 Potassium K 39.15 38.85 Ca 40.1 39.7 Praseodymium. Pr 140.5 139.4 C 12.00 11.91 Ra 225. 223.3 Ce 140.25 139.2 Rh 103.0 102.2 Chlorine CI 35.45 35.18 Rb 85.5 84.9 Chromium Cr 52.1 51.7 Ruthenium Ru 101.7 100.9 Cobalt Co 59.0 58.55 Sm 150.3 149.2 Columbium Cb 94. 93.3 Sc 44.1 43.8 Copper Cu 63.6 63.1 Se 79.2 78.6 Er 166. 164.8 Si 28.4 28.2 Fluorine F Gd 19. 156. 18.9 154.8 Silver Ag Na 107.93 23.05 107.11 ■Gadolinium Sodium 22.88 Ga 70. 69.5 Strontium Sr 87.6 86.94 Germanium . . . Ge 72.5 72. Sulphur S 32.06 31.82 Gl 9.1 9.03 Ta 183. 181.6 Gold Au He 197.2 4. 195.7 4. Tellurium Terbium Te Tb 127.6 160. 126.6 158.8 H 1.008 1.000 Thallium Tl 204.1 202.6 In 115. 114.1 Thorium Th 232.5 230.8 1 126.97 126.01 Thulium Tm 171. 169.7 Iridium Ir Fe Kr 193.0 55.9 81.8 191.5 55.5 81.2 Tin Sn Ti W 119.0 48.1 184. 118.1 Iron Titanium 47.7 Krypton 182.6 Lanthanum La 138.9 137.9 U 238.5 236.7 Lead Pb Li 206.9 7.03 205.35 6.98 Xenon V Xe 51.2 128. 50.8 127. Magnesium Mg 24.36 24.18 Ytterbium Yb 173.0 171.7 Manganese Mn 55.0 54.6 Yt 89.0 88.3 Hg 200.0 198.5 Zn 65.4 64.9 Molybdenum. . . Mo 96.0 95.3 Zr 90.6 89.9 2 TABLE OF THE PERIODIC SYSTEM OF CHEMICAL ELEMENTS. §2. S i— I W 013 o rd 0> 13 o •c Ph C4 coo M > o a. Fe= Co= Ni= 55.9 59.0 58.7 Ru= Rh= Pd= 101.7 103 106.5 Os= lr= Pt= 191 193.0 194.8 1—1 xd a la co o >-i ■* en tU || id o en cm II CQ ^ £- rH "■ II 1? 11 II s t-1 r> xd KOC o to eq CO <=> °. 0» t^ co cq t- cq m rH CO r-1 ,_( ^ Q,5 II ii s i - ii a 1 O CO || CO H » || || o S 5 3 > x"d Tt* o o n la P. l-i ui o 06 •*p CO t- O r- i— i xd tcoc P od cq* as cd ifl 3 " rH «? . s H s « ii i! 5 ^ s ii r s ii a ° " H ° II 5 ^ £ " 1- N (j h O rH © lfl rH II II 3 " 2 N s m 01 < II ° ^ II F o J 1 « CO >■ -1 1-4 1— 1 bra (-J •* ^ o CO 1£> CN) _4j © » H IO H ~ O 5 * - I % II 3 II a II " II N || g || - II — °> is t ° re X ra C3 £ o co m ce l-J ttor. CO O (M lO CO 00 _, t-^ „oio«OinO°; o> II " || o II 5 1 = -1 Z * cc o © £C ,« o en °°. go ^ oq ^ rH cq ill" J ^ z < ¥ X i 1> O § 6 4 S 0) P. H io" oi H a 3 ca to" a 3 a 10 CD H a a a o ■e H Q „ § °? CO r t- © P a *6 a* CD £ CD £ H a lS| o'° 2-3 H a ■° 2 ^5 CD" g - rH 4J cc ° to - 9^ a I »o o " .S h p, . a as|| ffl . 2 o „ CO "*h bii h 2 ■ S rt . i. fl 2 10 p *• 3 IIS* . S.c2i ■o o §3. DISCUSSION OF THE PERIODIC SYSTEM. 3 §3. In this system of the chemical elements certain regular gradations of chemical character are to be studied and held in view, to simplify the multitude of facts observed in analysis. Passing from Li 7.03 to F 19.05 in the first Series of this system, the elements are successively less and less of the nature to constitute bases and more and more of the nature to form acids, as their atomic weights increase. The acid-forming elements are electro-negative to the elements which form bases.* But in passing from 19.05 to the next higher atomic weight, Na 23.05, we return from the acid extreme to the basal extreme and begin another period, in gradation through the seven Groups. There is a like return from one extreme to the other in the steps between chlorine and potassium * Bases are the oxygen compounds of the metals. Acids are compounds of elements for the most part not metals. In the chemical union of sodium with chlorine, for example, these two elements differ widely from each other in their various properties. The chlorine is the opposite of the sodium in that very power by virtue of which the one combines with the other in the making of sodium chloride, a distinct product. In the polarity of electro- lysis the sodium is the positive element, while the chlorine is the negative element. The power of opposite action exercised by the one element upon the other, in their combination together, is represented by the opposite polarity of the one in relation to the other during electrolysis. Electrolysis is an exercise of the same energy that is otherwise manifested in chemical union or in a chemical change. Strictly speaking, it may be said that it is only in electrical results that a positive or a negative polarity appears. But the term positive polarity, applied to sodium because it goes to the negative pole of a battery, is a term which well designates the oppositeness of the chemical action of sodium in its union with chlorine. That is to say, the metals are in general " positive," the not-metals in general " negative," in the relation of the former to the latter, and this relation may be termed one of " polarity," whether it appear in electrolysis, in chemical combination, or in a chemical change. In chemical combination, the atoms of each element act with a " polarity," the extent of which may be expressed in terms of hydrogen equivalence or " valence." The valence of an element, when in combination with another element, may be counted as relatively " positive " or " negative " to the latter. For example, in the compound known as hydro- sulphuric acid, the sulphur is negative, the hydrogen positive, in the relation of one to the other, as represented by the diagram, H +- H+~ S ' in which the plus and minus signs of mathematics are used to represent the " positive " and " negative " activities of chemical elements. That is, the sulphur acts with two units of valence, both in negative polarity. In sulphuric acid the sulphur is positive in relation to both the oxygen and the hydroxyl, as indicated in the diagram .(HO)-+ J - +-° (HOI-+" j +- That is, the sulphur acts with six units of valence, all in positive polarity. In respect to oxidation and reduction, the difference between the action of sulphur in hydrosulphuric acid on the one hand, and in sulphuric acid on the other hand, is a difference equivalent to eight units of valence, the combining extent of eight atoms of hydrogen. This value is in agreement with the factors of oxidizing agents in volumetric analysis. In the same sense there is a change of " polarity " equivalent to the extent of eight units of valence, in reducing periodic acid to hydriodic acid, in reducing arsenic acid to arsine, or in reducing carbon tetrachloride to methane. That is, in any of the groups from IV. to VII. there is a difference, equivalent to the combining extent of eight hydrogen units, be- tween the negative polarity of the element in its regular combination with hydrogen, such as NH S , and its positive polarity in its highest combination with oxygen, such as N0 2 (OH). 4 DISCUSSION OF THE PERIODIC SYSTEM. §*. ■and in those between bromine and rubidium. This fact of a periodic return in the gradation of the properties of the elements, as their atomic weights ascend, constitutes a periodic system. A period is termed a Series. A Group in this system consists of the corresponding members of all the Series, which members are found to agree in valence, so that the number •of the groups, from I. to VII. (not in VIII.), expresses the typical valence of the element's as grouped. Further inquiry shows that all the properties of the elements are in relation to their atomic weights, as they appear in the periodic system. But this system is not to be depended upon to give information of the facts; it is rather to be used as a compact simpli- fication of facts found independently, by the student and by the author- ities on whom the student must depend. A full account of the Periodic ■System, as far as it is understood, is left to works on General Chemistry. §4. The remarkable position of Group VIII., made up of three series, •each of three elements near each other in atomic weight, respectively in Series 4, 6, and 10, is in central relation to the entire system. In this group there is something of a return, from negative to positive polarity, from higher to lower valence. Group VIII. lies between Group VII. and Group I., that is to say in this group there is a return from negative to positive nature, and from higher to lower valence. Moreover, the newly discovered elements related to argon, destitute of combining value as they are, appear to constitute a Group 0. The latest results render this position of the argon group of elements so probable that it has been placed in the ■chart for convenience of study, subject to further conclusions. (W. Eamsay. £r. Assoc. Adv. Sci., 1897, 598-601; B. 1898, 31, 3111. J. L. Howe, G. N., 1899, 80, 74; 1900, 82, 15, 52. Ostwald, Grundr. Allg. Chem., 3te Auf., 1899, S. 45.) In comparison with the members of Group VII. those of 'Group VIII. certainly have a diminished negative polarity, and a lower valence, the latter being easily variable. Some of the particulars are given below under the head, " Metals in Eelation to Iron." The most remark- able thing about Group VIII. is the fact that the return to Group I. from Group VIII. is less complete than the return from Group VII. That is to say, the character of copper is divided between Group VIII. and Group I.. and the same is true of silver and of gold. This relation to Group VIII. ■can be traced, in some particulars, to zinc and cadmium and mercury in Group II. For these reasons Series 4 and 5 may be studied as one long period of seventeen members, Series 6 and 7 as another long period and Series 10 and 11 as a third and final long period. §5. It is to be observed that each one of the Groups, from I. to VII., falls in two columns, a column consisting of the alternate elements in the group. Thus, H, Li, K, Rb and Cs make up the first column of Group I. It is among the alternate members of a group that the closer grade-relations of §9. DISCUSSION OF THE PERIODIC SYSTEM. a the elements are found. The gradations represented under one column are distinct from those under the other in the same group. The well known alternate elements of a Group, so far as found clearly graded together in respect to given properties, are to be studied as a Family of elements. Again a number of elements next each other in a Series are to be studied together, either by themselves or with an adjoining half-group. For the studies of analytical chemistry the following given are the more strongly marked of the families of the well known elements. §6. The Alkali Metals.— -Li 7.03, (Na 23.05), K 39.11, Rb 85.4, Cs 132.9. The first part and sodium of the second part of Group I. In the grada- tion of these elements the basal power increases qualitatively with the rise in atomic weight. The hydroxides and nearly all salts of these metals are freely soluble in water, wherein they are unlike the ordinary metals of all the other groups. For the most part, however, these solubilities increase) with the atomic weight of the metal, and the carbonate and orthophosphata of lithium are but slightly soluble. §7. The Alkaline Earth Metals.— (Mg 24.3), Ca 40.1, Sr 87.60, Ba 137.40. These metals, like those of the alkalis, form stronger bases as they hava higher atomic weights. Both in Group I. and in Group II. the member in Series 3 (Na, Mg), though in the second set of alternate members, agrees in many ways with the next three of the first set of alternates. The hydroxides of these metals are not freely soluble in water but are regularly more soluble as the atomic weight of the metal is higher. The sulphides are freely soluble; the carbonates and orthophosphates quite insoluble. The sulphates have a graded solubility, decreasing as the atomic weight is higher, an order of gradation the reverse of that of the hydroxides and of wider range. That is, at one extreme the magnesium sulphate is freely soluble, at the other barium sulphate is insoluble. §8. The Zinc Family.— Kg 24.3, (Al 27.1), Zn 65.4, Cd 112.4, , Hg 200.0. These metals, save aluminum, belong to the second alternates of Group II., and, like those of the corresponding half of Group I., in their gradation they are in general less strongly basal as they rise in their atomic weights. Aluminum, here drawn in from Group III. second half, has the valence of the third group, and differs from the others in not forming a sulphide. The sulphide of magnesium is soluble, the sulphides of zinc, cadmium and mercury insoluble in water, and these three show this grada- tion, that the zinc sulphide is the one dissolved by dilute acid, while the mercury sulphide is the one requiring a special strong acid to dissolve it, both these differences being depended upon in analysis. Mercury, sepa- rated from cadmium by two removes in the periodic order, is but a distant member of this family. §9. Metals in Relation to Iron.—Cr 52.1, Mn 55.0, Fe 55.9, Ni 58.70, 6 DISCUSSION OF THE PERIODIC SYSTEM. §10. Co 59.00. The atomic weights of these metals lie nearly together. They all belong to one Series, the fourth, representing Groups VI. and VII., and make the first of the instances of three members together in one series in Group VIII. Chromium, being in the first division of its group, could not be expected to grade with sulphur and selenium, nor would manganese be expected to grade with chlorine and bromine, but the disparity is strik- ing in both cases, especially in the comparison of melting points. The valence of both chromium and manganese appears partly exceptional to their positions in the system but the maximum valence of each is regular. That all of these five elements, neighbors to chlorine and bromine, are counted as metals, is not contrary to the periodic order. Group VIII. binds Group I. to Group VII. After Co 59.00 follow Cu 63.6 and then Zn 65.4. Indeed each of " the well-known metals related to iron " is capable of serv- ing as either a base or an acid, by change of valence. These metals are the special subjects of oxidation and reduction. So far they resemble their non-metallic neighbors, the halogens. Of the five, chromium and man- ganese (nearest the halogens) form the best known acids. Nickel and cobalt, like copper, have a narrower range of valence, a more limited extent of oxidation and reduction, within which they as readily act. These valences, in capacity of combination with other elements, not including the most unusual valences, may be written in symbols as follows: 2-3-6 2-3-4-6-7 2-3-6 2-3 2-3 1-2 2 Cr , Hn , Pe . , Hi , Co , Cu , Zn On reaching zinc, 65.4, in this gradation, the capacity of oxidation and reduction disappears. Sulphides are formed by such of these metals as act with a valence of two (all except chromium), and these sulphides are insolu- ble in water. In the conditions of precipitation sulphides are not formed with the metal in any valence other than two. Chromium acting as a base with a valence of three, like aluminum whose only valence is three, refuses to unite with sulphur. Trivalent iron in precipitation by sulphides is mainly reduced to ferrous sulphide (FeS). In chromates the chromium valence is reduced from six to three by hydrogen sulphide acting in solu- tion. A carbonate is not formed by chromium, this being another agree- ment with aluminum, and the same is true of trivalent iron. §10. The'Metals not Alkalis in Group I., Second Part, and their Relatives in Group 7771.— Cu 63.6, Ag 107.92, , Au 197.2. In gradation these metals are less strongly basal, and more easily reduced from their com- pounds to the metallic state, as their atomic weights rise. This is in agree- ment with the gradation among the second set of alternates in Group II.. the Zinc Family. It likewise agrees with second part of Group VII., the halogens. These elements of Group I. are to be studied with those of Group VIII., especially with those respectively nearest them in atomic §12. DISCUSSION OF THE PERIODIC SYSTEM. 7 weight: Cu 63.6 with Ni 58.70 and Co 59.00, Ag 107.92 with Pd 107.0, and Au 197.2 with Pt 194.9. Those with atomic weights above that of copper rank as "noble metals," from their resistance to oxidation and other qualities, so ranking in higher degree as their atomic weights increase. Their melting points (those of Pd, Ag, An, Pt) rise in the same gradation. By action of ammonium hydroxide upon solutions of their salts these (seven) metals form metal ammonium compounds, all of which are soluble in water except the compounds of platinum and gold (highest in atomic weight). All of the seven named form sulphides insoluble in water, in condition of precipitation. For the most part their sulphides are relatively more stable than their oxides. Silver differs from the others in the insolu- bility of its chloride, and agrees irregularly in this fact, one prominent in analysis, with mercury in its lower valence, and partly with lead. §11. The Nitrogen Family of Elements.— N 14.04, P 31.0, As 75.0, Sb 120.4, , Bi 208.1. The entire second part of Group V., and from the first part the Leading Element of the group. Nitrogen and phosphorus count as non-metals, antimony and bismuth as metals, arsenic as inter- mediate, the polarity being more positive as the atomic weight increases. In combinations with hydrogen, like ammonia and ammonium compounds, phosphine and phosphonium salts, and also like analogous organic bases where carbo-hydrogen takes the place of a part or all of the hydrogen, there is a remarkable unity of type in this family. The same is true of the com- binations with oxygen, like nitric acid. It is in Group V. that the group valence for oxygen begins to diverge in gradation from the group valence for hydrogen. In ammonium compounds nitrogen exercises a valence of five, it doubtless is true, but this total of five units is always limited in polarity to a balance of three negative units at most. In ammonia: N - 3 . HHH. In ammonium chloride : N - 4 + 1 = - 3 . HHHHC1. Bismuth is a distant member, a vacancy falling between it and antimony. Phosphorus, arsenic and antimony are in gradation with each other as to their indifference to chemical combination and readiness of reduction to the elemental state, these qualities intensifying with the rise in atomic weight. In this gradation nitrogen, belonging among the other alternate members, has no part. In its chemical indifference it stands in extreme contrast to phosphorus. §12. Relation of Tin and Lead to the Nitrogen Family. — These metals are in Group IV., each combining both as dyad and tetrad, a valence dis- tinctly unlike the valence of the nitrogen family, which is entirely regular for Group V. In Series 7: Sn 119.0, Sb 120.4. In Series 11: Pb 206.92, Bi 208.1. The metals in the first named pair are two removes from those in the second pair, all being among the second alternate members. In their salts tin and antimony are more easily subject to changes of valence than 8 DISCUSSION OF THE PERIODIC SYSTEM. §12. are lead and bismuth. In further comparison, arsenic, in its deportment as a metal, may be included, making the list: As 75.0, Sb 120.4 (Sn 119.0), Bi 208.1, (Pb 206.92). Of these,' only arsenic forms a higher oxide soluble in water (separation after treatment with nitric acid and evaporation). Arsenic and antimony form gaseous hydrides, in this agreeing with phosphorus and nitrogen, the others do not. The stability of the hydrides of N, P, As, Sb, all in the type of ammonia, is in the ratio inverse to that of the atomic weight. All of these metals are precipitable as hydroxides save arsenic, all are precipitated as sulphides, and these have chemical solubilities some- what in gradation with atomic weights, the arsenic sulphide being most fully separable by chemical solvents. The sparing solubility of the chloride of lead, referred to in description of silver, is approached by the insolu- bility of the oxy-chlorides of bismuth, tin, and antimony, and this fact must be borne in mind, when precipitation by hydrochloric acid is employed for separation of silver and univalent mercury in analysis. Nitrogen in its trivalent union with hydrogen, the leading element of the group of alkali metals, constitutes an active alkali. In its prevalent union with oxygen, the leading element of Group VI., that is with oxygen and hydroxyl, nitrogen forms an acid which is very active though not very stable, its decomposition being represented by its gunpowder salt. The degree of negative polarity of nitrogen, or its capacity for acid formation, in accordance with its place next to oxygen among the atomic weights, is shown in that singular instable body, hydronitric acid, H N 3 , of decided acid power, constituting well marked salts, such as Na N 3 , in which a ring of nitrogen alone acts as an acid radical. The first four members of the nitrogen family agree with each other in forming trivalent and pentavalent anhydrides and acids, the pentavalent ones being the more stable. The pentavalent acids are of especial interest. In nitric acid the five units of positive valence of an atom of nitrogen are met by two atoms of oxygen with two units each of negative valence and a unit of negative valence of hydroxyl: H — — NZJq . The same constitution is found in metaphos- phoric acid HO P 2 , meta-arsenic acid HO As 2 , and in antimonic acid HO Sb 2 . The so-called ortho acids, phosphoric and arsenic, have the constitution (HQ) 3 P and (H0) 3 As , respectively. Phosphoric and arsenic acids have a remarkable likeness to each other in nearly all the properties of all their salts, behaving alike in analysis so long as preserved from action of reducing agents. These sharply separate arsenic, usually in one of its trivalent forms, AsH 3 or As 2 S 3 . Antimony is reduced from its acid .even more readily than is arsenic, in accordance with the gradation stated above. In the solubility of its metal salts the acid of nitrogen is, again, in §14. DISCUSSION OF THE PERIODIC SYSTEM. 9 strong contrast with the acids of the elements of the second part, phos- phoric and arsenic acids. Metal nitrates are generally all soluble in water. Of the metal phosphates and arsenates, that is the full metallic salts of phosphoric and arsenic acids, in their several forms, only those of the alkali metals dissolve in water. §13. The Halogens.— F 19.05, CI 35.45, Br 79.95, I 126.85. The lead- ing element of Group VII., one of its first set of alternate members. and the three known members of the second alternates. In the halogen family fluorine has a relation like that of nitrogen in its family, taking part in the group gradation as to polarity, solubility of compounds and other qualities, but standing quite by itself in respect to certain properties. It is the most strongly electro-negative of the known elements, a fact in accord with the relation of its atomic weight. For the common work of analysis we may confine our study of the halogens to chlorine, bromine, and iodine. In the order of their atomic weights, these elements appear, respectively, in gaseous, liquid, and solid state, under common conditions. Their hydrogen acids, HC1, HBr, and HI, show a stability in proportion to the electro-negative polarity of the halogen, hydriodic acid being so unstable as to suffer decomposition in the air. In the solubility of their metal salts these acids are nearly alike, all being soluble except the silver, univalent mercury, and lead salts, but the iodides of divalent mercury, bismuth and divalent palladium are sparingly soluble. Each of these halogens, most especially iodine, forms a class of salts each containing two metals, one of the united metals being that of an alkali, such as (KI) 2 Hgl, and K 2 Pt Cl a . The periodides show that iodine atoms have a power of uniting with each other, in the molecules of salts, a power partly shared by bromine and chlorine and probably exercised in many complex halogen compounds. By this means two atoms of a halogen may serve the same as one atom of oxygen, in the linkings of molecular structure. Of the oxygen acids of chlorine, bromine and iodine, those in which the halogen has a valence of five are more stable than the others. These acids are chloric, HO CI 2 ; bromic, HO Br 2 ; and iodic, HO I 2 . Chloric acid agrees with nitric acid, HO N 2 , in the fact that it forms soluble salts with all the metals. Chlorates decompose more violently than nitrates; iodates for the most part less readily than the latter. Of the oxygen acids with a halogen valence of seven, periodic acid, HO I 3 , also (H0) 5 1 , is pre- served intact without difficulty. §14. The Belations of Sulphur. — S 32.07. Sulphur is the first member of a family including selenium and tellurium. It differs from oxygen almost as much as phosphorus differs from nitrogen, and we may say more than silicon differs from carbon. The higher valence of Group VI., exer- 10 THE CLASSIFICATION OF THE METALS AS BASES. §15. cised toward oxygen, cannot be met by oxygen itself. Of the acids of sulphur, H 2 S , in which sulphur has two electro-negative units of valence, is quite unstable, while (H0) 2 S 2 , in which the sulphur has six electro- positive units of valence, is the most stable. The sulphides (salts of H 2 S) of the heavier metals quite generally are insoluble in water, an important means of separation in analysis. The sulphates (salts of H 2 S0 4 ) of the larger number of the metals are soluble in water, the exceptions being important to observe, those of Pb 206.92, Ba 137.40, Sr 87.60, and (with sparing solubility) Ca 40.1. Of these sulphates, that of barium (least solu- ble), is the one usually employed in analytical separation. §15. The Relations of Carbon. — C 12.0. Carbon, in a central position in respect to polarity, stands alone in its capacity for a multitude of dis- tinct compounds with hydrogen and oxygen, with and without nitrogen, these being the so-called organic compounds. This capacity goes with the power of carbon atoms to unite with each other in the same mole- cule. It appears in acetylene C 2 H 2 (HCeC H), also in oxalic aeidj (HO) C — C (0 H). The same capacity of union of the atoms of an element with each other, in the molecules of compounds, is exercised by other elements in fewer instances, as by nitrogen in hydronitric acid, by oxygen in ozone, by sulphur in thiosulphurie acid, and by iodine in periodides. In carbon, nitrogen, and oxygen we see a decreasing grada- tion of this capacity, as the atomic weights ascend. Silicon, next to carbon in Group IV., but in the opposite set of alternates, agrees with carbon in the formation of many corresponding compounds, while it is entirely desti- tute of the capacity of uniting its atoms to each other in building up combinations. §16. The Classification of the Metals as Bases. The grouping of all the elements, both metals and not metals, according to their properties as related to their atomic weights, is the object of The Periodic System, briefly given in the foregoing pages for studies bearing especially upon the main methods of analysis. . The ordinary grouping of the bases in the work of analysis, outlined in the next paragraph, is done by the action of a few chemical agents, termed "group reagents," which have been chosen from a large number of re- agents, as being more satisfactory than others, for the use of the greater number of analysts. This ordinary grouping, therefore, is not the only •way in which the metals can be separated, in the practice of analytical chemistry, nor is any one scheme of separation adopted throughout by all , authorities. The principal separations of analysis can be well understood by gaining an acquaintance with the properties of the leading bases and acids. $16. THE CLASSIFICATION OF THE METALS AS BASES. 11 in their action upon each other. Without this acquaintance, the analyst is the servant of routine, and his results liahle to fallacy. The following named are the bases of more common occurrence. The Alkali Bases. The sixth group* Potassium (Kalium), K?.\ Sodium (Natrium), Na 1 . Ammonium, (NH 4 )'- The Alkaline Earth Bases. The fifth group. Magnesium, Mg n . Calcium, Strontium, Barium, Ca 11 , Sr 11 , Ba 11 . Not precipitated from their salts by any of the group reagents. Potas- sium and sodium are found after re- moving all the following named groups. Ammonium is Hound by tests of the original, this base being added in the " group reagents." In combination in potassium hy- droxide, KOH , and in potassium salts, such as the chloride KC1 , and the nitrate, KN0 3 . In the base, sodium hydroxide and its salts. Forms ammonium hydroxide, NH 4 OH , representing ammonia, NH 3 , and water, and serving as the base of ammonium salts, such as (NH 4 ) 2 S0 4 , ammonium sulphate. (Precipitated by carbonates, which fact alone does not separate them from the following named groups.) Separated by precipitation as a phosphate after removing all the fol- lowing named bases. Forms magne- sium hydroxide, Mg(0H) 2 , and mag- nesium salts, such as MgS0 4 . Separated by precipitation with Ammonium Carbonate, adding NH 4 C1 to Tceep magnesium from pre- cipitation. Calcium carbonate, a normal salt, CaCO., . * The sixth division of the bases, in the order in which they are separated from each other toy precipitation with the group reagents. tThe Roman numerals (as i) express units of valence, each equivalent to an atom of hydrogen, in the formation of salts and other combinations. 12 CLASSIFICATION OF THE METALS AS BASES. 516. The Zinc and Iron Groups. The Zinc Group. The fourth group. Zn 11 : zinc salts. Mn 11 : manganous salts. Mn 111 : manganic salts. Mn": salts unstable. Mn VI : salts of manganic acid. Mn vn : salts of permanganic acid. Ni n : nickel salts. Co 11 : cobaltous salts. Co 111 : eobaltic salts. The Iron Group. The third group. Fe n : ferrous salts. Fe 111 : ferric salts. Cr 111 : chromic salts. C.r VI : chromates. Al m : aluminum salts. Metals falling with Copper and Tin, The second group. ■ The Copper Group. Division B, second group. Mercury (Hydrargyrum). Hg n : mercuric salts. Hg 1 : mercurous salts. Silver (Argentum). Ag 1 : silver salts. Lead (Plumbum). Pb 11 : lead salts. Bi m : bismuth salts. Cu 11 : copper or cupric salts. Cu 1 : cuprous salts. Cd": cadmium salts. (Precipitated by sulphides, this-, being a separation from the fore- going, not from the following named groups of bases.) by precipitation with Ammonium Sulphide, after removal of all the following named bases as directed below. (The precipitates- are all sulphides.) Separated by precipitation with Ammonium Hydroxide, in presence of NH 4 C1 , after the removal of the groups named following. (The pre- cipitates are all hydroxides.) Precipitated by H 2 S in acidulated solution. (The precipitates are sul- phides.) Separated by the insolubility of the precipitated sulphides in treat- ment with Ammonium Sulphide. §18. THE OPERATIONS OF ANALYSIS. 13 The Tin Group. Division A, second group. Sn 11 : stannous salts. Sn IV : stannic salts and stannates. Sb 111 : antimonous compounds. Sb v : antimonic compounds. As 111 : arsenous compounds. As v : arsenic compounds and arsen- ates. Metals Precipitated as Chlorides. The Silver Group. The first group. Separated by dissolving the pre- cipitated sulphides with Ammonium Sulphide. The silver, lead, and univalent mercury, grouped in the division last above given. Silver and the mer- cury of mercurous salts can be re- moved, as chlorides, by precipitation with hydrochloric acid. The precip- itate of lead is not insoluble enough to remove this metal entirely, in sep- aration from other groups. §17. The Acids of Certain Commonly Occurring Salts. Name of Acid. Name of Salt. Formula. Showing Hydroxyl. Anhydrii Carbonic Carbonate H 2 C0 3 (HO) 2 CivO C0 2 Oxalic Oxalate H 2 C 2 4 (HO) 2 C 2 iv0 2 C 2 a Nitric Nitrate HN0 3 (HO)NV0 2 N 2 5 Nitrous Nitrite HN0 2 (HO) Nino N 2 8 Phosphoric (ortho) Phosphate H 3 P0 4 (H0)„PV0 .P 2 O s Metaphosphoric Metaphosphate HP0 3 (H0)PV0 2 P 2 5 Pyrophosphoric Pyrophosphate H 4 P 2 7 (HO)4PV 2 0, p 2 o 5 Sulphuric Sulphate H 2 S0 4 (HO) 2 Svi"0 2 so„ Sulphurous Sulphite H 2 S0 3 so 2 Hydrosulphuric Sulphide H 2 S Hydrochloric Chloride HC1 Hydrobromic Bromide HBr Hydriodic Iodide HI Chloric Chlorate HCIO3 (H0)C1V0 2 ci 2 o. Iodic Iodate HIO a (H0)IV0 2 i 2 o B The Operations of Analysis. §18. Chemical analysis is the determination of any or all of the compo- nents of a given portion of matter, whether this be solid, liquid or gaseous. A portion of matter is made up of one or more definite and distinct sub- stances, or' chemical individuals, each of which is either a " compound " or 14 TEE OPERATIONS OF ANALYSIS. §19- an " element " and js always and everywhere the same. It is required of analysis to determine a chemical compound as a hody distinct Urom the chemical elements that have formed it. For example, the analyst may have in hand a mixture containing sodium sulphate, Na 2 S0 4 ; sodium sul- phite, Na 2 S0 3 , and sodium thiosulphate, Na 2 S 2 3 , but not containing any sodium or sulphur or oxygen as these bodies are severally known to the world and described in chemistry. In this instance the analyst in his ordinary work does not separate the sulphur or the sodium, as elements uncombined with oxygen, either in qualitative or in quantitative oper- ations. Each one of the compounds of the sulphur with the oxygen is usually sought for and found and weighed as a chemical individual. Cer- tain of the chemical elements, however, are frequently separated free from all combination, as a method of determination of their compounds. §19. The analysis of gaseous material is termed Gas Analysis; that of mixtures of the complex compounds of carbon, Organic Analysis. An examination of organic matter, when limited to a determination of its ulti- mate chemical elements is styled Ultimate Organic Analysis. When it is Undertaken to determine individual carbon compounds actually existing in organic matter, it has been spoken of as Proximate Organic Analysis. If the same distinction were to be applied to inorganic analysis, we should have to say that it is mostly " proximate " but is sometimes " ultimate " in its methods of operation. §20. The term Qualitative Chemical Analysis as commonly used is con- fined to a chemical examination of material, chiefly inorganic, in the solid or liquid state, the inquiry being limited for the most part to well known substances. §21. In the methods of analysis of a mixture, it is often required to separate individual substances from each other, but sometimes a distinct compound can be identified and sometimes its quantity can be estimated while it is in the presence of other bodies. Both the identification and separation are accomplished, nearly always, by effecting changes, physical and chemical. Methods of analysis are as numerous as are the ways of bringing into action the physical and chemical forces by which chemical changes are wrought. The characteristics of any chemical individual, by which it is distinguished and removed from others, lie in its responses to the physical and chemical forces, including especially the chemical action of certain well known compounds called reagents. §22. The response toward heat and pressure fixes the melting and boiling points, its ordinary solid or liquid or gaseous state. The operations "in the dry way " are done over a flame or in a furnace, with or without solid "reagents" and with regard to oxidation. They represent some of the §27. THE OPERATIONS OF ANALYSIS. 15 methods of metallurgical manufacture. The liquid, state, whether by fusing or by solution, is the state commonly necessary or favorable to chem- ical change and its control. §23. The deportment of a solid substance toward light comprises its color and that of its solutions, as well as that of its vapor, in ordinary light, and the bands and primary colors it exhibits in the uses of the spectroscope (Crookes, J. C, 1889, 55, 255; Welsbach, M.. 1885, 6, 47). §24. The conduct of a chemical compound in electrolysis is, in various cases, a means both of identification and of separation. Electric conduc- tivity methods are used for establishing the presence or absence of minute traces of substances (Kohlrausch Whitney, Z. phys. Ch., 1896, 20, 44). Again, traces of dissolved matters too minute for other means of detection can be revealed by the difference of electric potential between electrode and solution (Ostwald, Lehrb., 2 Aufl., II, 1, 881; Behrend, Z. phys. Ch., 1893, 11, 466; Hulett, Z. phys. Ch., 1900, 33, 611). §25. By far the most extensive of the resources of analysis lie in the chemical reaction of one definite and distinct substance with another, ac- cording to the character of each, giving rise to a chemical product having peculiarities of its own in evidence of its origin. In this way the com- pounds are bound in regular relations to each other. Therefore it belongs to the analyst to gain personal acquaintance with the behavior of the repre- sentative constituent bases and acids toward each other. §26. Operations for chemical change are commonly conducted in solu- tion. The material for analysis is dissolved, and is treated with reagents that are in solution. A solid or a gas is dissolved in a liquid in making a solution. "When the dissolved substance is converted into one that will not dissolve a precipitate is formed. It is necessary therefore to under- stand the nature of solution and to give heed to its obvious limitations. Certain facts and conclusions as to the chemical state of dissolved com- pounds are presented under the head next following, " Solution and Ioniza- tion." But it must first be observed that the universal solvent, water, is always understood to be present in somewhat indefinite proportion in opera- tions " in the wet way." It serves as a vehicle, as such not being included in any statement of the substances operated upon, nor formulated in equa- tions, any more than is the material of the test tube, but often some portion of it enters into combination or suffers decomposition, and then it must be placed among the substances engaged in chemical change. §27. No other property of substances has so great importance in analysis and in all chemical operations, as their solubility in water. It must never be forgotten that there are degrees of solubility, but there is hardly such a fact as absolute solubility, or insolubility, regardless of the proportion •of the solvent. There are liquids which are miscible with each other 16 TEE OPERATIONS OF ANALYSIS. §28. in all proportions, but solids seldom dissolve in all proportions of the sol- vent, neither do gases. For every solid or gas, there is a least quantity of solvent which can dissolve it. One part of potassium hydroxide is soluble in one-half part of water (or in any greater quantity), but not in a less quantity of the solvent. One part of sodium chloride requires at least two and a half parts of water to dissolve it. One part of mercuric chloride will dissolve in two parts of water at 100 degrees, but when cooled to 15 degrees so much of the salt recrystallizes from the solution, that it needs twelve parts more of water at the latter temperature to keep a perfect solution. Lead chloride dissolves in about twenty parts of hot water, about half of the salt separating from the solution when cold. Calcium sulphate dis- solves in about 500 times its weight of water — this dilute solution forming one of the ordinary reagents. Barium sulphate is one of the least soluble precipitates obtained, requiring about 430,000 parts of water for its solution at ordinary temperature (Hollemann, Z. phys. Ch., 1893, 12, 131). In ordi- nary reactions it is not appreciably soluble in water. Lead sulphate dis- solves in about 21,000 parts of water: in many operations this solubility may be disregarded, but in quantitative analysis the precipitate is washed with alcohol instead of water, losing less weight with the former solvent. These examples indicate the necessity of discriminating between degrees of solubility. Also the solubility of a particular compound is dependent upon the physical form of that compound (§69, 5 V); e. g., amorphous magnesium ammonium phosphate is quite soluble in water, the crystalline salt being almost insoluble. When a solvent has dissolved all of a substance that it can at a particular temperature, in contact with the solid, the solution is said to be saturated at that temperature. It frequently happens that a. saturated solution of a substance at a higher temperature may be cooled without separation of the solid. Such a solution (at the lower temperature) is said to be supersaturated and precipitation frequently is induced by jarring the solution, more surely by adding a crystal of the dissolved sub- stance. §28. The ordinary liquid reagents are solutions in water — sulphuric acid and carbon disulphide being exceptions. Hydrochloric acid, liquid hydro- sulphuric acid, and ammonium hydroxide (reagents) are solutions of gases in water; on exposure to the air these gases gradually separate from their solutions. All these gases escape much more rapidly when their solutions are warmed. The majority of liquid reagents are solids in aqueous solu- tion. (See the list of Eeagents.) §29. Substances are said to dissolve in acids, or in alkalis, and this is termed chemical solution; more definitively it is chemical action and solu- tion, the solution being counted as a physical change. We say that cal- cium oxide dissolves (chemically) in hydrochloric acid; that is, in the §33. THE OPERATIONS OF ANALYSIS. 17 reagent named hydrochloric acid, a mixture of that acid and water. The acid unites with the calcium oxide, forming a soluble solid, which the water dissolves. Absolute hydrochloric acid cannot dissolve calcium oxide. §30. Solids can be obtained, without chemical change, from their aqueous solutions: Firstly, by evaporation of the water. This is done by a careful application of heat. Secondly, solids can be removed from solution, with- out chemical change, by (physical) precipitation — accomplished by modify- ing the solvent. If a solution of potassium carbonate, or of ferrous sul- phate, be dropped into alcohol, a precipitate is obtained, because the salts will not dissolve, or remain dissolved, in the mixture of alcohol and water. But, in analysis, precipitation is more often effected by changing the dis- solved substance instead of the solvent. §31. Solids can be separated from their solution by precipitation due to chemical change, to the extent that the product is insoluble in the quantity of the solvent present. Calcium can be in part precipitated from not too dilute solutions of its salts, by addition of sulphuric acid; but there still remains not precipitated the amount of calcium sulphate soluble in the water and acid present, which is enough to give an abundant precipitate with ammonium oxalate, the precipitated sulphate being previously re- moved by nitration. Time and heat are required for the completion of most precipita- tions. If it is necessary to remove a substance, by precipitation, before testing for another substance, the mixture should be warmed and allowed to stand for some time, before filtration. Neglect of these precautions often occasions a double failure; the true indication is lost, and a false indication is obtained. §32. Eeagents should be added in very small portions, generally drop by drop. Often the first drop is enough. Sometimes the precipitate redis- solves in the reagent that produced it, and this is ascertained if the reagent be added in small portions, with observation of the result of each addition. If it is a final test, a quantity of precipitate which is clearly visible is suffi- cient, but if the precipitate is to be filtered out and dissolved, a considerable quantity should be formed. If the precipitate is to be removed and the filtrate tested further, the precipitation must be completed — by adding the reagent as long as the precipitate increases, with the warmth and time requisite in the operation; and a drop of the same reagent should be added to the filtrate to obtain assurance that the precipitation has been completed. It will be found, with a little experience, that some reagents must be used in relatively large quantities. On the contrary, the acids, sulphuric, hydro- chloric and nitric, are required in a volume relatively very small. §33. Certain very exact methods of identification can be conducted by drop tests upon a black or white ground, or upon a glass slide and especially 18 THE OPERATIONS OF ANALYSIS. §34. with help of a microscope and with studies of crystalline form. Further see Behrens, Z. 1891, 30, 125; and Herrnschmidt and Capelle, Z. 1893, 32, 608. §34. Precipitates are removed — usually by filtration, sometimes by decan- tation. If they are to be dissolved, they must be first washed till free from all the substances in solution. For complete precipitation some excess of the reagent must have been used. Beside the reagent there are other dis- solved matters, after precipitations, some of which are indicated by the equation written for the change. All these dissolved substances permeate and adhere to the porous precipitate with greater or less tenacity. If they are not wholly washed away, some portion of them will be mixed with the dissolved precipitate. Then, the separation of substances, the only object of the precipitation is not accomplished, while the operator, proceeding just as though it was accomplished, undertakes to identify the members of a group by reactions on a mixture of groups. The washing, on the filter, is best completed by repeated additions of small portions of water — around the filter border, from the wash bottle — allowing each portion to pass through before another is added. The washings should be tested, from time to time, until they are free from dissolved substances. §35. In dissolving precipitates — by aid of acids or other agents — use the least possible excess of the solvent. Endeavor to obtain a solution nearly or quite saturated, chemically. If a large excess of acid is carried into the solution to be operated upon, it usually has to be neutralized, and the solution then becomes so greatly encumbered and diluted that reactions become faint or inappreciable. Precipitates may be dissolved on the filter, without excess of solvent, by passing the same portion of the (diluted) solvent repeatedly through the filter, following it once or twice with a few drops of water. The mineral acids should be diluted to the extent required in each case. For solution of small quantities of carbonates and some other easily soluble precipitates the acids may be diluted with fifty times their weight of water. Washed precipitates may also be dissolved in the test-tube, by rinsing them from the filter, through a puncture made in its point, with a very little water. If the filter be wetted before filtration, the precipitate will not adhere to it so closely. §36. "When the addition of a reagent is to cause a change in the acid, alkaline or neutral condition of the solution, the addition of sufficient reagent to cause the desired change should always be governed by testing a drop of the solution, on a glass rod, with a piece of litmus paper. §37. "When substances in separate solution are brought together, an evidence of the formation of a new substance is the appearance of a solid in the mixture, a precipitate. A chemical change between dissolved sub- stances — salts, acids, and bases — will be practically complete when one or §40. THE OPERATIONS OF ANALYSIS. 19 more of the products of such change is a solid or a gas, not soluble in the mixture. As an example, Calcium carbonate +. Hydrochloric acid = Cal- cium chloride + Water + Carbon dioxide (gas). §38. In the practice of qualitative analysis, the student necessarily refers to authority for the composition of precipitates and other products. For example, when the solution of a carbonate is added to the solution of a calcium salt, a precipitate is obtained; and it has been ascertained by quanti- tative analysis that this precipitate is normal calcium carbonate, CaCO, , invariably. "Were there no authorized statement of the composition of this precipitate, the student would be unable, without making a quantitative analysis, to declare its formula or to write the equation for its production. When the results of analytical operations are substances of unknown, uncer- tain, or variable composition, equations cannot be given for them. §39. The written equation represents only the substances, and the quan- tity of each, which actually undergo the chemical change that is to be expressed. Thus, if a reagent is used to effect complete precipitation, an excess of it must be employed, beyond the ratio of its combining weight in the equation. That is, if magnesium sulphate be employed to precipitate barium chloride, the exact relative amount of magnesium sulphate indicated by the equation : BaCl 2 + MgS0 4 = BaSO 4 + MgCl 2 , fails to precipitate all of the barium. The soluble sulphate must be in a slight excess. On the other hand, to effect complete precipitation of the sulphate the barium must be in a slight excess. §40. By translating chemical equations into statements of proportional parts by weight, they are prepared to serve as standard data of absolutely pure materials, and applicable in operations of manufacture, with large or small quantities, after making due allowance for moisture and other im- purities, necessary excess, etc. In quantitative analysis the equation is the constant reliance. For example, in dissolving iron fey the aid of hydro- chloric acid, we have the equation: Pe + 3HC1 = FeCl 2 + H 2 . 56 + 72.9 = 126.9 + 2 . Also in precipitating ferrous chloride by sodium phosphate, we have the equation: FeCl 2 + Na 2 HP0 4 ,12H 2 = FeHP0 4 + 2NaCl + 12H s O . 126.9 + (142.1 + 216) = 152 + 117 . Suppose it is desired to determine from the above: (1) How much hydrochloric acid, strength 32 per cent, is required to dissolve 100 parts of iron wire. (2) What quantities of 32 per cent hydrochloric acid and iron wire are necessary to use in preparing 100 parts of absolute ferrous chloride. 20 SOLUTION AND IONIZATION. §41. (3) What materials and what quantities of them, may be used in prepar- ing 100 parts of ferrous phosphate. In practice allowance must be made for the facts that the iron wire will not be quite pure, and that a considerable excess of the hydrochloric acid would be necessary to the complete solution of the iron. Also that some excess of the phosphate would be necessary to the full precipitation of the iron. Irrespective of impurities, oxidation product and excess, the re- quired quantities are found by the combining weights as follows: ■. J 56/72.9 = 100/x = parts of absolute HC1 for 100 parts of iron wire. ' \ 32/100 = x/y = parts of 32 per cent HC1 for 100 parts of iron wire. f 126.9/72.9 = 100/x 2. -j 32/100 = x/y = parts of 32 per cent HC1 for 100 parts of FeCl 2 , absolute. (. 126.9/56 = 100/z = parts of iron wire for 100 parts of FeCl 2 . 3.- 152/72.9 = 100/x 32/100 = x/y = parts of 32 per cent HC1 for 100 parts of TeHP0 4 . 152/56 = 100/z = parts of metallic iron for 100 parts of T"eHP0 4 . . 152/358.1 = 100/u = parts of Na 2 Bn?0 4 ,12H 2 for 100 parts of FeHP0 4 Practice in reducing the combining numbers of the terms in an equation to simple parts by weight, is a very instructive exercise, even in the early part of qualitative chemistry. It enforces correct and clear ideas of the significance of formulas and equations, and refers all chemical expressions to the facts of quantitative work. §41. The chief requirement in qualitative practice is an experimental acquaintance with the chemical relations of substances, rather than the identification of one after the other by routine methods. The acids and bases, the oxidizing and reducing agents, are all linked together in a net- work of relations, and the ability to identify one, as it may be presented in any combination or mixture, depends upon acquaintance with the entire fraternity. §42. The full text of the book, rather than the analytical tables, should be taken as the guide in qualitative operations, especially in those upon known material. The tabular comparisons are commended to attention, especially for review. In actual analysis, the tables serve mainly as an index to the body of the work. Solution and Ionization. §43. The Theory of Electrolytic Dissociation, proposed by Arrhenius in 1887 (Z. phys. Ch., 1887, 1, 631), assumes that salts, acids, and bases in water solution are present not as the intact molecule but split up into certain components, and that the characteristics of the dissolved substance result very largely from the extent to which this breaking down of the §43. SOLUTION AND IONIZATION. 21 molecule has taken place. The facts upon which the theory is based are in a word the parallelism between osmotic pressure,* electric conductivity, :and chemical activity of substances in solution. The gas-laws (Boyle's, Gay-Lussac's, Henry's, and Dalton's) are found to hold for dissolved substances, osmotic pressure being substituted for gas-pressure (van 't Hoff, Z. phys. Ch., 1887, 1, 481). Avogadro's Hypoth- esis is therefore applicable to solutions as well as to gases, and as abnormal .gas-pressure points to dissociation in the gas (NH 4 C1, PC1 5 ) so excessive ■osmotic pressure is taken as indicating dissociation of the dissolved sub- stance. The osmotic pressure is a measure of this dissociation. Faraday gave the name ions to the components of a substance conducting ihe electric current in solution. It is an observed fact that transmission •of the current by a solution is always accompanied by movement of the ions in opposite directions (Hittorf, Pogg. 1853, 89, 177). This is quite independent of any separations taking place at the electrodes. From this it is concluded that the ions carry the electricity from one pole to the other through the solution. If the ions are the carriers of electricity then the power of a solution to conduct the current will be in proportion to their number, that is, to the extent of dissociation of the dissolved substance. And experiment shows that the dissociation calculated from the osmotic pressure is identical with the dissociation calculated from the electric conductivity. Further, if in analysis of a substance in solution we are dealing not with the substance in its integrity but with certain ions, then our ordinary analytical reactions are reactions of the ions, and we may expect that where the substance for some reason is transformed from the ionized condition to the undivided molecule these reactions will fail. Here again the chemi- ■cal activity will be proportional to the number of ions; and experiment shows that unquestioned quantitative parallelism exists, to take the case of acids, between (1) the characteristic acid activity — the dissolving of metals, the influence as catalyzer on such changes as the inversion of cane- sugar and the saponification of esters; (2) the extent of dissociation as indicated by osmotic pressure, and (3) the extent of dissociation as indicated by electric conductivity. The same parallelism holds for other bodies in solution. The very active acids and bases and the neutral salts undergo wide dissociation in watsr solution, while weak acids and bases retain almost entirely the non-dissociated condition. The Electrolytic Dissociation Theory in its assumption of a separation * The pressure by virtue of which a soluble substance in contact with the solvent, as common ■salt in water, is enabled to rise against the force of gravity and distribute itself uniformly "throughout the solvent, just as a gas by virtue of the gas-pressure occupies the entire space at its disposal. 22 SOLUTION AND IONIZATION. §44. into ions groups together and gives system and meaning to these three classes of facts, experimentally absolutely independent and up to Arrhenius' time without any suspected relationship. In each case the results calculated on the assumption of such a dissociation are in quantitative agreement with those obtained by measurement. Corresponding in actual experience to the view that the common analyti- cal reactions are due to the ions rather than to, the molecule as a whole, is the analyst's practice of testing for acid radicle or basic radicle without regard to the other component; and on the other hand, to take a specific case, the fact that the sulphur in H 2 S does not give the same precipitation reactions as that in K 2 S or H 2 S0 4 or H 2 S0 3 or'H 2 S 2 3 . Further, HgCl 2 in its chemical behavior is unlike other mercuric salts and unlike other chlorides. The mercury is not readily precipitated by alkali hydroxides nor is the chloride readily precipitated by silver salts. In agreement with this, its conductivity and osmotic pressure- are also unlike those of the great majority of neutral salts, both pointing to very slight dissociation into the ions. CdCl 2 is another neutral salt anomalous in that its conductivity and osmotic pressure are both low. And here also for precipitation of the chloride a considerable concentration of the reagent is necessary. Similar- instances of the parallelism referred to are numberless. §44. The Law of Mass-Action embodies the familiar principle that the chemical activity of a substance is proportional to its concentration. It was first recognized, although imperfectly, by Berthollet and was given mathematical expression by Guldberg and "Waage in 1867. The latter investigators found it to accord well with the observed facts in some cases ^ in others there were wide discrepancies which were later shown by Ar- rhenius to disappear when the concentration, not of the reacting body as a whole but only of that part present in the ionized condition, was taken into consideration. We must assume that every chemical reaction is rever- sible, that is, that none of them proceed until the reacting substances are completely transformed. Then by a simple process of reasoning it is found that when equilibrium sets in the product obtained by multiplying together' the concentrations of the reacting substances will be in a certain definite ratio to the product of the concentrations of the substances formed, con- centration being defined as the quantity in unit volume.* For example, in the reaction indicated by the equation CH 3 C0 2 H + C 2 H 5 0H = CH 3 C0 2 C 2 H 5 + H 2 , when equilibrium sets in ab = kcd , in which a and b are the concentrations of acid and alcohol respectively, c and d those of ester and water, while k is a constant peculiar to the reaction. Where the * The unit of quantity is the molecular weight taken in grams (the " mol "). Where there are 18.23 grama HC1 in a liter either in solution or as gas the concentration is }6, where there are 72.92 grams in the same volume the concentration is 2. and so on. §45. SOLUTION AND IONIZATION. 23 reaction is a dissociation, as with gaseous NH 4 C1 , we have ab = k'c , a and b representing the concentrations of NH 3 and HC1 respectively, c that of the undecomposed NH 4 C1, and k' the constant characteristic of this change. Dissociation into ions must follow the same laws, and for the electrolytic, dissociation of acetic acid a similar equation holds, a and b in this case standing for concentration of H and acetic ions, c for concentration of non- dissociated acetic acid, while the constant is one governing only this par- ticular dissociation. It is apparent from each of these equations that, if we add one of the products of the reaction and thus increase its concentra- tion, the concentration of the other product must decrease in the same- proportion — the extent of the reaction will be decreased; while, on the other hand, removing either or both of the products will tend to make the transformation complete. This deduction is of great significance. In making ethyl acetate from the acid and alcohol, in order to use the materials- as completely as possible, the ester is distilled off as rapidly as produced while the water is taken up by some absorbent. Introducing gaseous NHj or HC1 diminishes the dissociation of NH 4 C1 by heat, and similarly adding either H ions or acetic ions will diminish the dissociation of acetic acid. Acetic acid is much weakened by the presence of a neutral acetate. A ferrous solution moderately acidified with acetic acid gives no precipitate on saturation with H 2 S , but on addition of sodium acetate the black FeS is brought down. Similarly a weak base, as NH 4 0H, is made still less effective by the presence of its strongly-dissociated neutral salt, as NH 4 C1 . Quantitative agreement is obtained between observed effect of NH 4 C1 on NH 4 0H as saponifying agent and that calculated from the equation: °NH • C 0H' = k °NH OH (Arrhenius, Z. phys. Ch., 1887, 1, 110). §45. The Solubility-Product. — In the saturated solution which always remains after precipitation we have the usual dissociation equilibrium, as: C Ae ' C C1' °AeCl • ^ ow ^ e quantity of non-dissociated substance in a saturated solution is invariable and the right side of this equation is therefore constant. That is, in saturated solution the product of the con- centrations of the ions is always the same for a given substance (Nernst). This Ostwald has called the Solubility-Product. Where the saturated solu- tion is made by bringing the salt into contact with the solvent c ^ • C ny • From stich a solution precipitation will take place on addition of either a silver salt or a chloride, for such addition largely increases the concentration of one ion and, to restore equilibrium, the concentration of the other ion must decrease in the same proportion, which is possible only by precipita- tion. From this follows the old empirical rule to add an excess of the reagent in making a precipitation. Experiments on this point give quanti- 24 ORDER OF LABORATORY STUDY. §46. tative agreement with the theory (Nernst, Z. phys. Ch., 1889, 4, 372; Noyes, Z. phys. Ch., 1890, 6, 241; 1892, 9, 603). The Solubility-Product of the alkaline-earth carbonates is °M " °C0 " = I n t ne solution of a neutral salt, as CaCl 2 , Ca ions are present in large concentration. When a substance containing C0 3 ions in large concentration is added, as Na 2 C0 3 , the solubility-product is exceeded and precipitation takes place. Carbonic acid, however, is shown by con- ductivity and osmotic pressure measurements to be but slightly disso- ciated, that is, it contains few C0 3 ions, and in accord with this is the familiar fact that the alkaline earths are not precipitated by carbonic acid. Similarly the fixed alkali hydroxides, strongly dissociated, will precipitate alkaline-earth hydroxides, while ammonium hydroxide, shown by other measurements to contain but few hydroxyl ions, will not. For the metallic sulphides the solubility-product is c jj •• C g" • 'The alkali sulphides as normal salts contain the S ion in large concentra- tion and so produce precipitation even of the more soluble sulphides of ~the Iron and Zinc Groups. The slightly dissociated H 2 S contains sufficient S ions to reach the solubility-product of the sulphides of the Silver, Tin, and Copper Groups, but not enough to attain to the larger solubility- product of the Iron and Zinc Group sulphides. A strong acid, as HC1 . •containing as it does H ions, one of the dissociation products of H 2 S , drives back the dissociation of the H 2 S, so decreasing the concentration of the ;S ions and making precipitation of the sulphide more difficult. For the application of the dissociation theory to the details of analytical ~work we are indebted chiefly to Ostwald. See his " Scientific Foundations of Analytical Chemistry " and " Outlines of General Chemistry." Okdek op Laboratory Study. §46. The following is a suggestive outline to be modified by the teacher to suit the ability of the students, and the amount of time to be given to the study: a. A review of chemical notation and the writing of salts. I. A study of the action of the Fixed Alkalis upon solutions of the salts 'of the metals in the order of their groupings; including the action of an •excess of the reagent. The fact of the reaction should be stated; e. g., lead acetate + potassium hydroxide = a white precipitate readily soluble in excess of the reagent. The text should then be consulted for the products ■of the reaction (6a), and the reactions expressed in the form of equations: 2Pb(C 2 H: a 2 ) 2 + 4KOH = Pb 2 0(0H) 2 * (white) + 4KC 2 H 8 2 + H 2 Pb 2 0(OH) 2 + 4KOH (excess) = 2K 2 Fb0 2 + 3H 2 or Pb(C 2 H a 2 ) 2 + 4KOH (excess) = K 2 Fb0 2 + 2KC 2 H 3 2 + 2HVO . * It has been found helpful to require students to underscore all precipitates. §46. ORDER OF LABORATORY STUDY. 25 The results should all he tabulated and then summarized in form of a carefully worded generalization (§205, 6a). c. Action of Ammonium Hydroxide (volatile alkali) upon solutions of the salts of the metals, etc., as in (b) above ; e. g., lead nitrate + ammonium hydroxide = a white precipitate not dissolving in excess. Consult text (§57, 6a) and write the equation: 3Pb(2T0 3 ) 2 + 41JH 4 0H = 2PbO.I'b(M , Os) ;! + 42TH: i NO a + 2H 2 . After the work has heen completed in the laboratory and the results discussed in the class room, summarize in the form of a generalized state- ment (§207, 6a). d. A study of the action of the Fixed Alkali Carbonates, and generaliza- tion of the results (§205, 6a). e. A study of the action of Ammonium Carbonate. Summarize the re- sults (§207, 6a). f. A study of the solvent action of acids, HC1 , HN0 3 , and H 2 S0 4 , upon the Hydroxides and Carbonates obtained by precipitation. g. Action of Hydrosulphuric Acid as a precipitating agent upon salts of the metals in neutral and acid solutions. h. The use of Ammonium Sulphide as a reagent. i. The solvent action of acids, HC1 , HN0 3 , and HC 2 H 3 2 , upon the sulphides obtained by precipitation. ;. Action of Hydrochloric Acid and Soluble Chlorides. Action of Hydrobromic Acid and Soluble Bromides. Action of Hydriodic Acid and Soluble Iodides. fc. Precipitation by Soluble Sulphates, Phosphates, and Oxalates. I. The solvent action of Hydrochloric and Acetic Acids upon the Phos- phates obtained by precipitation. m. The reverse of certain of the above reactions as illustrating the precipitation of Acids; e. g., Ammonium oxalate + calcium chloride = a white precipitate. Consult the text (§227, 8), and write the equation r. (NH 4 ) 2 C 2 4 + CaCl 2 = CaC 2 4 + 2NH 4 C1 . n. Application of the above reactions to the Grouping of the Metals, for Analysis. o. A study of the limit of visible precipitation with several reagents; upon a particular metal, or upon a number of metals. p. A study of the analysis of the individual metals and acids; combining them, and effecting their separation and detection. The new work of each day to be followed by the analysis of " unknown " mixtures prepared by the teacher to illustrate the new work and to give an instructive review of the preceding work. The order of the study of the metals and acids may be varied greatly. In no case should the metals of a whole group be studied without considering the relations to the other groups. 26 ORDER OF LABORATORY STUDY. §46. q. The study in the class room of Oxidation and Reduction, with work in the laboratory to illustrate. r. The study of problems in Synthesis involving analytical separations, accompanied by laboratory experiments. s. The analysis of a series of Dry " Unknown " Mixtures. t. A special study of the analysis of Phosphates, Oxalates, Borates, Silicates, etc., and certain of the Rarer Metals. u. The analysis of mixtures in solution, illustrating Oxidation and Reduction. v. A study of Electrolysis as a means of detection in qualitative analysis. PAET.II.-THE METALS. THE SILVEE AND TIN AND COPPER GROUPS. (First and Second Gboups.) §47. The Silver group (first group) includes the metals whose chlorides are insoluble in water and which are precipitated from solutions upon the -addition of hydrochloric acid or soluble chlorides : Pb, Hg', Ag . The Tin and Copper group (second group) includes those metals whose sulphides are precipitated by hydrosulphuric acid from solutions acid with dilute hydrochloric acid, and whose chlorides (soluble in water for the most part) are not precipitated by hydrochloric acid or soluble chlorides. Lead* Pb 206.9S Germanium Ge 72.5 Mercury Hg 200.0 Iridium Ir 193.1 Silver . Ag 107.92 Osmium Os 191.0 Arsenic As 75.0 Palladium Pd 107.0 Antimony Sb 120.4 Khodium Rh 103.0 Tin Sn 119.0 Ruthenium PvU 101.7 Gold Au 197.2 Selenium Se 79.2 Platinum Pt 194,9 Tellurium Te 127.5? Molybdenum Mo 96.0 Tungsten ■w 184. Bismuth Bi 208.1 Vanadium V 51.4 Copper Cu 63.6 Cadmium Cd 112.4 §48. Owing to the partial solubility of lead chloride in water, it is never completely precipitated in the first group; hence it must also be tested ior in the second group. Monovalent mercury belongs to the first group and divalent mercury- to the second. Silver, then, is the only exclusively first-group metal. §49. The metals included in these groups are less strongly electro- positive than those of the other groups. Only bismuth, antimony, tin, and molybdenum decompose water, and these only slowly and at high temperatures. The oxides of silver, mercury, gold, platinum, and palla- dium are decomposed below a red heat. Copper, lead, and tin tarnish by * In this list of the metals of the Silver, Tin and Copper Groups the more common, those in "the first column, are arranged in the order of their discussion and separation in analysis. The rare metals are arranged in alphabetic order, but are discussed in order of their relations to each other, beginning at § 104. 28 GENERAL DISCUSSION. §50* oxidation in the air. In general, these metals do not dissolve in acid* with evolution of hydrogen, or do so with difficulty. Nitric acid is the best solvent for all, except antimony and tin, which are rapidly oxidized by it. Concerning the separation and detection of the metals of these groups by electrolysis, see Schmucker, Z. anorg., 1894, 5, 199, and Cohen,. J., Soc. Ind., 1891, 10, 327 (§12). §50. Mercury, arsenic, antimony, and tin form, each two stable classes- of salts. Therefore, the lower oxides, chlorides, etc., of these metals act as reducing agents; and their higher oxides, chlorides, etc., as oxidizing agents, each to the extent of its chemical force. Arsenic, antimony, tin, molybdenum, and several of the rare metals of these groups enter into- acidulous radicles, which form stable salts. Arsenic, selenium and tellu- rium are metalloids rather than metals. Arsenic, antimony, and bismuth belong to the Nitrogen Series of Elements. §51. A large proportion of the compounds of these' metals are insoluble in water. Of the oxides or hydroxides, only the acids of arsenic are soluble in water. The only insoluble chlorides, bromides, and iodides are in these groups. The sulphides, carbonates, oxalates, phosphates, borates, and cyanogen compounds are insoluble. Most of the so-called soluble compounds of bismuth, antimony, and tin, and some of those of mercury, dissolve only in acidulated water, being decomposed by pure water, with formation of insoluble basic salts. §52. Among the many soluble double salts of the metals of these groups are especially to be mentioned the double iodides with KI and the iodides of Pb , Hg , Ag , Bi and Cd . Platinum forms a large number of stable double chlorides, soluble and insoluble; and gold forms double chlorides, cyanides, etc. §53. The oxides of arsenic act as acid anhydrides and form soluble salts- with the alkalis ; oxides of antimony, tin, and lead, are soluble in the fixed alkalis; oxides of silver, copper, and cadmium, in ammonium Irydroxide. Metallic lead, like zinc, dissolves in the fixed alkalis with evolution of hydrogen. §54. The solubility of certain sulphides in the alkali sulphides forming sulpho salts or double sulphides, separates the metals of the second group into two divisions. A (tin group) — As , Sb , Sn , Ge , Au , Ir , Mo , Pt , Se , Te , W . and V ; sulphides soluble in yellow ammonium sulphide; and B (copper group) — Hg , Pb , Bi , Cu , Cd , 0s , Pd , Rh , and Ru ; sulphides not soluble in yellow ammonium sulphide. §55. Mercury, antimony, silver, and gold do not form hydroxides. The oxides of gold are very unstable. §56. The metals of these groups are all easily reduced to the metallic state by ignition on charcoal. Except mercury and arsenic, which vaporize §57, 4. LEAD. 29 readily, and certain rarer metals difficultly fusible, the reduced metals melt to metallic grains on the charcoal. The Silver Group (First Group). Lead, Mercury (Mercurosum), Silver. §57. Lead (Plumbum) Pb = 206.92 . Valence two and four. 1. Properties. — Specific gravity, 11.37 (Reich, J. pr., 1859, 78, 328). Melting point, 327.69° (Callendar and Griffiths, C. N., 1891, 63, 2). It begins to vaporize at a red heat and boils at a white heat. Vaporization is said to take place at 360° (Demarcay, C. r., 1882, 95, 183). It can be distilled in vacuo (Schuller, B., 1883, 16, 1312). Pure lead is almost white, soft, malleable, very slightly ductile, tarnishes in the air from formation of a film of oxide. The presence of traces of most of the other metals makes the lead sensibly harder. It is a poor conductor of heat and electricity. It forms alloys with most metals; lead and tin in various pro- portions form solder and pewter; lead and arsenic form shot metal; lead and antimony, type metal; lead, bismuth, tin and silver form a fusible alloy melting as low as 45° ; bell metal consists of tin, copper, lead and zinc. 2. Occurrence. — It is rarely found native (Chapman, Phil. Mag., 1866, (4), 31,, 176); its most abundant ore is galena, PbS; it also occurs as cerussite, FbC0 3 ; anglesite, PbSO„; pyromorphite, 3Pb 3 P 2 O s + PbCl 2 ; krokoite, PbCr0 4 ; and also in many minerals in combination with arsenic, antimony, etc. The United States produces more lead than any other country. Spain produces about one-fourth the world's supply. 3. Preparation. — From galena (a) It is roasted in the air, forming variable quantities of PbS0 4 , PbO , and PbS ; then the air is excluded and the tempera- ture raised, and the sulphur of the sulphide reduces both the PbO and the PbS0 4 , S0 2 being formed: PbS0 4 + PbS = 2Pb + 2S0 2 . 2PbO + PbS = 3Pb + S0 2 . (6) Similar to the first except that some form of carbon is used to aid in the reduction, (c) It is reduced by fusing with metallic iron: PbS + Fe — Pb + FeS . Frequently these methods are combined or varied according to. the other ingredients of the ore. 4. Oxides.— Lead forms four oxides, Pb 2 , PbO , Pb0 2 , and Pb„0 4 . Lead, suboxide (Pb 2 0) is little known; it is the black powder formed when PbC 2 4 is. heated to 300°, air being excluded. Lead oxide (litharge, or massicot) is formed by intensely igniting in the air Pb , Pb 2 , Pb0 2 , Pb 3 4 , Pb(OH) 2 , PbC0 3 , PbC 2 4 , or Pb(NO a ) 2 . It has a yellowish-white color, melts at a red heat, and is volatile at a white heat. Trilead tetroxide (red lead or minium), Pb 3 4 , is formed by heating PbO- to a dull-red heat with full access of air for several hours. Strong, non-reduc- ing acids, such as HN0 3 , H 2 S0 4 , HC10, , etc., convert it into a lead salt and Pb0 2 (a). But concentrated hot H 2 S0 4 converts the whole into PbS0 4 , oxygen being evolved (6). But with the dilute acid and reducing agents, such as. C 3 H 5 (OH) 3 , C H 12 O , H 2 C 2 4 , H 2 C 4 H 4 O , Zn , Al , Cd , Mg, As , Pb , etc., it is all reduced to the dyad lead without evolution of oxygen (c), (d), and (e). Hydracids usually reduce the lead and are themselves oxidized (f). (a) Pb 3 4 + 2H 2 S0 4 (dilute) = Pb0 2 + 2PbS0 4 + 2H 2 (6) 2Pb 3 4 + 6H 2 S0 4 (concentrated and hot) = 6PbS0 4 + 6H 2 + O a (c) Ph 3 4 + H 2 C 2 4 + 6H2T0 3 = 3Pb(N0 3 ) 2 + 4H 2 + 2C0 2 (d) 10Pb 3 O 4 + As 4 + 30H 2 SO 4 = 30PbSO 4 + 4H 3 As0 4 + 24H a (e) Pb 3 4 + Zn + 4H 2 S0 4 = 3PbS0 4 + ZnS0 4 + 4H 2 (f) Pb 3 4 + 8HC1 = 3PbCl 2 + Cl 2 + 4H 2 The valence of Pb 3 4 is best explained by the theory that it is a union of the? dyad and tetrad (Pb" and Pbiv) , Pb„0 4 = 3PbO + PbrvO s . 30 LEAD. §57, 00. Lead dioxide or peroxide, Pb0 2 , is formed: (1) by fusion of PbO with KC10, or KNO s ; (2) by fusing Pb s 4 with KOH ; (3) by treating any compound of Pb" with CI , Br , K a re(CN) 6 , EMnO, , or H 2 2 in presence of KOH; (4) by treating Pb s 4 with non-reducing acids: Pb 8 4 + 4HN0 3 = Pb0 2 + 2Pb(NO s ) 2 + 2H 2 0. Ignition forms first Pb 8 4 and above a, red heat PbO, oxygen being given off. It dissolves in acids on same conditions as Pb„0 4 . Very strong solution of potassium hydroxide, in large excess, dissolves it, with formation of " potassium plumbate," K 2 PbO s . Lead dioxide is a powerful oxidizing agent, one of the strongest known. Digested with ammonium hydroxide, it forms lead nitrate and water. Triturated with one-sixth of sulphur, or tartaric acid, or sugar, it takes fire; with phosphorus, it detonates. 5. Solubilities. — a. — Metnl. — Nitric acid is the proper solvent for metallic lead, "the lead nitrate formed is readily soluble in water but insoluble in concentrated nitric acid *; hence if the concentrated acid be used to dissolve the lead, a white residue of lead nitrate will be left which dissolves on the addition of water. Dilute sulphuric acid is without action, the concentrated acid is almost without action in the cold (Calvert and Johnson, J. C, 1863, 16, 66), but the hot concentrated acid slowly changes the metal to the sulphate with evolution of sulphur dioxide, a portion of the salt being dissolved in the acid, precipitating •on the addition of water. Hydrochloric acid very slowly dissolves the metal (more rapidly when warmed), evolving hydrogen; the chloride formed dissolves in the acid in quantities depending upon eonditions of temperature and con- centration (c). The halogens readily attack the metal forming the correspond- ing haloid salts. Alloys of lead are best dissolved by first treating with nitric acid, if a white residue is left it is washed with water and, if not dissolved, it is then treated with hydrochloric acid, in which it will usually be soluble. Water used for drinking or cooking purposes should not be allowed to stand in lead pipes. Pure water free from air is without action upon pure lead, but water containing air and carbon dioxide very slowly attacks lead, forming' the hydroxide and basic carbonate. This action is promoted by the presence of salts, as ammonium nitrate, nitrite, chloride, etc.; the action seems to be hindered by the presence of sulphates. 6. — Oxides. — Lead oxide, litharge, PbO , and the hydroxides, 2Pb0.H 2 0; -3Pb0.H 2 O , are readily dissolved or transposed by acids forming the correspond- ing salts, i. e., PbO + H 2 S0 4 = PbS0 4 + H 2 . The oxide and hydroxide are soluble in about 7000 parts of water, to which they impart an alkaline reaction. They are soluble in the fixed alkalis forming plumbites; soluble in certain salts as NH 4 C1, OaCl 2 , and SrCl 2 (Andre, C. r., 1883, 96, 435; 1887, 104, 359); very ^soluble in lead acetate, forming basic lead acetate. Lead dioxide, Pb0 2 , lead peroxide, is insoluble in water or nitric acid; it is •dissolved by the halogen hydracids with liberation of the halogen and reduction •of the lead forming a dyad salt: Pb0 2 + 4HC1 = PbCl 2 + Cl 2 + 2H 2 0; it is attacked by hot concentrated sulphuric acid, forming the sulphate and liberat- ing oxygen; it is soluble in glacial acetic acid forming Pb(C 2 H s 2 ) 4 , unstable (Hutchinson and Pollard, J. C, 1896, 69, 212). Some of the salts of the tetrad lead seem to be formed when the peroxide is treated with certain acids in the cold. They are, however, very unstable, being decomposed to the dyad salt upon warming (Fischer, J. C, 1879, 35, 282; Nickels, A. Cii., 1867, (4), 10, 328). The peroxide is slowly soluble in the fixed alkali hydroxides forming plum- bates, i. e., PbO, + 2KOH = K 2 Pb0 3 + H 2 . Trilead tetroxide, Pb„0 4 , red lead, minium, is insoluble in water, is at- tacked by nearly all acids in the cold forming the corresponding dyad lead salt and lead peroxide, Pb0 2 . Upon further treatment with the acids Using heat the lead peroxide is decomposed as described above. The presence of many reducing agents, as alcohol, oxalic acid, hydrogen peroxide, etc., greatly * The solubility of a salt' is lessened by the presence of another substance having an ion in ■common with it (§45). In some cases, as with Ptol s and EI, this is offset in concentrated solution t>y formation of a complex compound. §57, 5c. LEAD. 31 facilitates the solution of red lead or lead peroxide in acids, i.e., nitric acid ■does not dissolve lead peroxide, but if a few drops of alcohol be added the solution is readily obtained upon warming, leaving the lead as the soluble nitrate, which greatly facilitates the further analysis. c. — Salts. — The carbonate, borate, cyanide, ferrocyanide, phosphate, sul- phide, sulphite, iodate, chromate, and tannate are insoluble in -water. The sulphate is soluble in about 21,000 parts of water at 18° (Kohlrausch and Eose, Z. phys. Gh., 1893, 12, 241), the presence of HN0 3 or HC1 in- creases its solubility in water; it is insoluble in alcohol even when quite dilute ; sparingly soluble in concentrated H 2 S0 4 , from which solution it is precipitated by the addition of water or alcohol; less soluble in dilute H 2 S0 4 than in water; soluble in 682 parts 10 per cent HC1, in 35 parts 31.5 per cent (Eodwell, J. C, 1862, 15, 59); transposed and dissolved by excess of HC1, HBr ; or HI forming the corresponding haloid salt; insoluble in HF (Ditte, A. Gh., 1878, (5), 14, 190); soluble in ammonium sulphate, nitrate, acetate, tartrate and citrate, and from these solutions not readily precipitated by ammonium hydroxide or sulphate (Fleischer, J. C, 1876, 29, 190; Woehler, A., 1840, 34, 235). The sulphate is almost completely "transposed to the nitrate by standing several days with cold concentrated nitric acid (Eodwell, I. c). The oxalate is sparingly soluble in water, insol- uble in alcohol; the ferricyanide is very slightly soluble in cold water, more soluble in hot water; the chloride is soluble in 85 parts water at 20° and in 32 parts at 80° (Ditte, C. r., 1881, 92, 718); the bromide is soluble in 16C parts water at 10°, in about 45 parts at 80°; the iodide is soluble in 1235 parts water at ordinary temperature, and in 194 parts at 100° (Denot, J. pr., 1834, 1, 425). The chloride is less soluble in dilute HC1 or H 2 S0 4 than in water, but is more soluble in the concentrated acids (Ditte, I. c.) ; HN0 3 increases the solubility of the chloride more and more as the HN0 3 is stronger. The chloride is less soluble in a solution of NaCl than in water (Field, J. C, 1873, 26, 575); soluble in NH 4 C1 —90 grams dissolving in 200 grams NH 4 C1 with 200 cc. water (Andre, G. r., 1893, 96, 435). The chloride, bromide, and iodide are insoluble in alcohol. The iodide is moderately soluble in solutions of alkali iodides; it is decomposed by ether. The basic acetates are permanently soluble if carbonic acid is strictly excluded. The basic nitrates are but slightly soluble in water, and are precipitated on adding solutions of KN0 3 to a solution of basic lead acetate. The relative insolubility of PbCl 2 in cold water or in dilute HC1 makes it possible to precipitate the most of the lead (by means of HC1) from solutions containing also the other metals of the Silver Group; while its solubility in hot water is the means of its separation from the other chlorides of that group (§61). The lead is separated and identified in the second group as the insoluble sulphate. (§95). 32 LEAD. §57, 6. 6. Reactions, a. — Fixed alkali hydroxides precipitate, from solutions of lead salts, basic lead hydroxide (1), Pb 2 0(0H) 2 (Schaffner, A., 1844, 51, 175),. white, soluble * in excess of the reagent as plumbite (2) (distinction from silver, mercury, bismuth, copper, and cadmium). The normal lead hy- droxide, Pb(0H) 2 , may be formed by adding a solution of a lead salt to. a solution of a fixed alkali hydroxide. (i) 2Pb(N0 3 ) 2 + 4KOH = Pb 2 0(OH) 2 + 4XN0 3 + H 2 (2) Pb 2 0(OH) 2 + 4KOH = 2K 2 Pb0 2 + 3H 2 . Ammonium hydroxide precipitates white basic salts, insoluble in water and in excess of the reagent (distinction from silver, copper, and cad- mium); with the chloride the precipitate, insoluble in water, is PbCl 2 .PbO.H 2 (Wood and Bordeu, C. N., 1885, 52, 43); with the nitrate- 2PbO.Pb(N0 3 ) 2 (D., 2, 2, 558). "With the acetate, in solutions of ordinary strength, excess of ammonium hydroxide (free from carbonate) gives no> precipitate, the soluble tribasic acetate being formed. Alkali carbonates precipitate lead basic carbonate, white, the composition varying with the conditions of precipitation. With excess of the reagent and in hot concentrated solutions the precipitate consists chiefly of Pb 3 (0H) 2 (C0 3 ) 2 . Precipitation in the cold approaches more nearly to the normal carbonate (Lefort, Pharm. J., 1885, (3), 15, 26). Solutions of lead salts when boiled with freshly precipitated barium carbonate are com- pletely precipitated. Carbon dioxide precipitates the basic acetate but not completely. b. — Oxalic acid and alkali oxalates precipitate lead oxalate, PbC 2 4 , white,. from solutions of lead salts, soluble in nitric acid, insoluble in acetic acid. A solution of lead acetate precipitates a large number — and a solution of lead subacetate a still larger number — of organic acids, color substances, resins, gums, and neutral principles. Indeed it is a rule, with few excep- tions, that lead subacetate removes organic acids (not formic, acetic, butyric, valeric, or lactic). Tannic acid precipitates solutions of lead acetate, and of the nitrate incompletely, as yellow-gray lead tannate,. soluble in acids. Soluble cyanides precipitate lead cyanide, Pb(CN) 2 , white, sparingly soluble in a large excess of the reagent and reprecipitated on boiling. Potassium f erro- cyanide precipitates lead ferrocyanide, Pb 2 Fe(CN') 6 , white, insoluble in water or dilute acids. Potassium ferricyanide precipitates from solutions not too dilute lead ferricyanide, Pb 3 (Pe(CN) ) 2 , white, sparingly soluble in water, soluble in nitric acid. Solutions of lead salts are precipitated by potassium sulphocyanate as lead sulphocyanate, Pb(CNS), , white, soluble in excess of the reagent and in nitric acid. c. — Lead nitrate is readily soluble in water, and dissolves the oxide to form the basic nitrate, which may also be formed by precipitating lead acetate with * Nearly all the salts of lead are soluble in the fixed alkali hydroxides, PbS forming almost the only notable exception. §57, Qe. LEAD. 33 potassium nitrate. The solubility of lead nitrate is greatly increased by the presence of the nitrates of the alkalis and of the alkaline earths, a complex compound being formed (Le Blanc and Noyes, Z. phys. Ch., 1890, 6, 385). d. — The higher oxides of lead are all reduced by hypophosphorous acid, lead phosphate being formed. Lead phosphite, PbHPO s , white, is formed by nearly neutralizing phosphorous acid with lead carbonate or precipitating Ha 2 HPO„ with Pb(NO„) 2 (Amat, C. r., 1890, 110, 901). Sodium phosphate, Na,HP0 4 , precipitates from solutions of lead acetate the tribasic lead phosphate, Pb 3 (P0 4 ) 2 , white, insoluble in the acetic acid which is set free (D., 2, 2, 562): 3Pb(C 2 H 3 2 ) 2 + 2Na 2 HP0 4 = Pb 3 (P0 4 ) 2 -f- 4NaC 2 H 3 2 + 2HC 2 H 3 2 . The same precipitate is formed when sodium phosphate is added to lead nitrate, soluble in nitric acid, insoluble in acetic acid. Lead phosphate is also precipitated upon the addition of phosphoric acid to solutions of lead acetate or lead nitrate. The pyrophosphate, Pb 2 P 2 0, , white, amorphous, is formed by precipitating a lead solution with Na 4 P 2 7 , soluble in excess of the precipitant, in nitric acid, and in potassium hydroxide; insoluble in ammonium hydroxide and in acetic acid (Gerhardt, A. Ch., 1849, (3), 25, 305). The metaphosphate, Pb(PO„) 2 , white, crystalline, is obtained by the action of NaP0 3 upon Pb(NO„) 2 in excess. e. — Hydros ulphuric acid and the soluble sulphides precipitate — from neutral, acid, or alkaline solutions of lead salts-^-lead sulphide, PbS, brownish black, insoluble in dilute acids, in alkali hydroxides, carbonates, or sulphides. Freshly precipitated CdS, MnS, FeS, CoS, and NiS also give the same precipitate. Hydrosulphuric acid and the soluble sulphides transpose all freshly precipitated lead salts to lead sulphide.* Moder- ately dilute nitric acid — 15 to 20 per cent — dissolves lead sulphide with, separation of sulphur (1), some of the sulphur, especially if the nitric acid be concentrated, is oxidized to sulphuric acid, which precipitates a portion ■of the lead {2), unless the nitric acid be sufficiently concentrated to hold that amount of lead sulphate in solution. The oxidation of sulphur always ■occurs when nitric acid acts upon sulphides, and in degree proportional to the strength of acid, temperature, and duration of contact. (1) 6PbS + 16HN0 3 = 6Pb(NO„) 2 + 3S 2 + 4NO + 8H 2 <2) 3PbS + 8HN0 3 = 3PbS0 4 + 8NO + 4H 2 In solutions too strongly acidulated, especially with hydrochloric acid, 'either no precipitation takes place, or a brick-red double salt, Pb 2 SCl 2 , * The condition for equilibrium is that a certain ratio of concentration exist between the ions, in the case of PbS0 4 between the S ions and the S0 4 ions. These concentrations are the same as those in a solution obtained by digesting the two salts, PbS0 4 and PbS, together in water. PbS0 4 dissolves more freely than PbS, and for equilibrium therefore c so ;, must be corres- pondingly greater than c g „. But adding H 2 S or a soluble sulphide to PbS0 4 gives just the opposite of this condition, and transformation accordingly results, increasing the S0 4 " con- centration by formation of soluble sulphate aDd decreasing the S" concentration by precipita- tion of PbS, until the equilibrium-ratio is produced or, if the quantity of PbSO) 4 present is in- sufficient for this, until all the PbS0> 4 has been transformed to sulphide. On the other hand, treatment of PbS with a very large excess of H 2 S0 4 will cause the reverse action, S ions going into solution until the same equilibrium results as before. The general principle is then that unless a constituent of the more soluble substance is in great preponderance in the solution the least soluble of two or more possible products will .always be formed. This principle determines the direction in which a reaction takes place ; AgCl + KI = Agl + KC1 ; CaSO„ + JI»,CO t = CaCO a + Na 2 S0 4 (§44). 34 LEAD. §57, 6/v is formed, the precipitation being incomplete. In neutral solutions con- taining 100,000 parts of water lead is revealed as the sulphide; a test ■which is much more delicate than the formation of the sulphate. Ferric chloride decomposes lead sulphide, forming lead chloride, ferrous chloride and sulphur. The reaction takes place in the cold and rapidly when warmed (Gabba, C. €., 1889, 667). When galena, PbS , is pulverized with fused KHS0 4 , H 2 S is evolved (Jan- nettaz, J. C, 1874, 27, 188). Lead thiosulphate, PbS 2 O a , white, is precipitated by adding sodium thiosul- phate to solutions of lead salts; the precipitate is readily dissolved in an excess of the reagent, forming the double salt, PbS 2 3 ,2Na 2 S 2 O a (Lenz, A., 1841, 40, 94) ; on boiling, all the lead is slowly precipitated as sulphide (Vohl, A., 1855, 96, 237). Sodium sulphite precipitates lead sulphite, PbS0 3 , white, less soluble in water than the sulphate, slightly soluble in sulphurous acid; decomposed by sulphuric, nitric, hydrochloric, and hydrosulphuric acids and by alkali sulphides; not decomposed by cold phosphoric and acetic acids. Sulphuric acid and soluble sulphates precipitate from neutral or acid solutions, lead sulphate, PbS0 4 7 white, not readily changed or permanently dissolved by acids, except hydrosulphuric acid, yet slightly soluble in strong acids (5c). Soluble in the fixed alkalis and in most ammonium salts, especially the acetate, tartrate, and citrate (Woehler, A., 1840, 34, 235). Soluble in warm sodium thiosulphate solution, in hot solution decomposed, lead sulphide, insoluble in thiosulphate, being formed (dis- tinction and separation from barium sulphate, which does not dissolve in thiosulphates). The test for lead as a sulphate is from five to ten times less delicate than that with hydrosulphuric acid; but lead is quantitatively separated as a sulphate, by precipitation with sulphuric acid in the presence of alcohol, and washing with alcohol. When heated with potassium chromate transposition takes place and yellow lead chromate is formed (h). Excess of potassium iodide transposes lead sulphate (f), a distinction of lead from barium. Eepeated washing of lead sulphate with a solution of sodium chloride completely transposes the lead to the chloride (Matthey, J. C.,. 1879, 36, 124). See footnote on previous page. f. — Hydrochloric acid and soluble chlorides precipitate, from solutions not too dilute, lead chloride, PbCl 2 , white. This reaction constitutes lead a member of the FIRST GROUP— as it also is of the second. The solu- bility of the precipitate is such (5c) that the filtrate obtained in the cold gives marked reactions with hydrosulphuric acid, sulphuric acid, chro- mates, etc.; and that it can be quite accurately separated from silver chloride and mercurous chloride by much hot water. Also, small propor- tions of lead escape detection in the first group, while its removal is necessarily accomplished in the second group. §57, 7. LEAD. 35 Hydrobromic acid and soluble bromides precipitate lead Iromide, PbBr 2 , white, somewhat less soluble in water than the chloride (5c); soluble in excess of concentrated potassium bromide, as 2KBr.PbBr 2 , which is decom- posed and PbBr 2 precipitated by dilution with water. Hydriodie acid and soluble iodides precipitate lead iodide, Pbl 2 , bright yellow and crystalline, much less soluble in water than the chloride or bromide (5c); soluble in hot moderately concentrated nitric acid and in solution of the fixed alkalis; soluble in excess of the alkali iodides, by forming double iodides, KIPbI 2 with small excess of KI , and 4KI.Pbt 2 with greater excess of KI ; these double iodides are decomposed by addi- tion of water with precipitation of the lead iodide. Lead iodide is not precipitated in presence of sodium citrate; alkali acetates also hold it in solution to some extent, so that it is less perfectly precipitated from the acetate than from the nitrate. Freshly precipitated lead peroxide, Pb0 2 , gives free iodine when treated with potassium iodide (Ditte, C. r., 1881, 93, 64 and 67). In detecting lead as an iodide in solutions of the chloride by precipita- tion with potassium iodide and recrystallization of the yellow precipitate from hot water, cafe must be taken that the potassium iodide be not added in excess to form the soluble double iodides. g. — Arsenous acid does not precipitate neutral solutions of lead salts; from alkaline solutions or with soluble arsenites a bulky white precipitate of lead arsenite is formed, insoluble in water, but readily soluble in all acids and in the fixed alkali hydroxides. Arsenic acid and soluble arsenates precipitate lead arsenate, white, from neutral or alkaline solutions of lead salts, soluble in the fixed alkali hydroxides and in nitric acid, insoluble in acetic acid. For the composition of the arsenites and arsenates of lead see (D., 2, 2, 565). Hot potassium stannite (SnCl 2 in solution by KOH) gives with lead salts or lead hydroxide a black precipitate of metallic lead. h. — Chromic acid and soluble chromates — both K 2 Cr0 4 and K 2 Cr 2 7 — ■ precipitate lead chromate, PbCr0 4 , yellow, soluble in the fixed alkali hydroxides (distinction from bismuth), insoluble in excess of chromic acid (distinction from barium), insoluble in ammonium hydroxide (distinction from silver), decomposed by moderately concentrated nitric and hydro- chloric acids, insoluble in acetic acid. 7. Ignition. — Lead salts when fused in a porcelain crucible with sodium carbonate are converted into lead oxide, PbO (a). After fusion and diges- tion with warm water, the aqueous solution is tested for acids, and the residue for bases after dissolving in nitric or acetic acid. If charcoal (or some organic compounds as sugar, tartrates, etc.) be present, metallic lead is formed (6); and with excess of charcoal the acid radicle may also be ehanged (c). If the fusion with sodium carbonate is made on a piece of charcoal, instead of in a crucible, using the reducing flame of the blow- J& 36 LEAD. §57, 8. pipe, globules of metallic lead are produced and at the same time the charcoal is covered with a yellow incrustation of lead oxide, PbO . (a) PbCl 2 + Na 2 C0 3 = 2NaCl + PbO + C0 2 (6) 2PbS0 4 + 2Na 2 C0 3 + C = 2Pb + 2Na 2 S0 4 + 3CO, (c) 2PbS0 4 + 2Na 2 C0 3 + 5C = 2Pb + 2Na 2 S + 7C0 2 8. Detection. — Lead is precipitated, incompletely, from its solutions by HC1 as PbCl 2 ; separated from AgCl and HgCl by hot water, and confirmed by H 2 S , H 2 S0 4 , K 2 Cr0 4 , and KI . It is separated (in the second group) from As , Sb , Sn , etc., by non-solubility of the sulphide in (NH 4 ) 2 S X ; from HgS by HN0 3 ; from Bi , Cu , and Cd by precipitation with dilute sulphuric acid. Insoluble compounds are transposed by an alkali sulphide, being then treated as lead in the second group, or they are examined by ignition as described in (7). 9. Estimation. — (a) As an oxide into which it is converted by ignition (if a carbonate or nitrate), or by precipitation and subsequent ignition. (6) As a sulphate. Add to the solution twice its volume of alcohol, precipitate with H 2 S0 4 , and after washing with alcohol ignite and weigh, (c) It is converted into an acetate, or sodium acetate is added to the solution, then precipitated With K 2 Cr 2 7 , and after drying at 100°, weighed as Pb0rO 4 . (d) It is con- verted into PbS , free sulphur added, and after ignition in hydrogen gas weighed as PbS . (e) The lead is precipitated with standardized sodium iodate and the excess of iodate is determined by retitration. Lead iodate is less soluble in water than lead sulphate (Cameron, J. 0., 1879, 36, 484). (f) In presence of bismuth, ignite the halogen compound, or convert into a sulphide and ignite in a current of bromine. The haloid salts of bismuth sublime upon ignition (Steen, Z. angew., 1895, 530). (g) Gas volumetric method. Precipitate as a chromate, filter, wash and transfer to an azotometer with dilute sulphuric acid and estimate the amount of chromium by the volume of oxygen set free by hydrogen peroxide (Baumann, Z. angew., 1891, 329), 10. Oxidation. — Metallic lead precipitates the free metals from solutions of Hg , Ag , Au , Pt , Bi , and Cu . Lead as a dyad is oxidized to the tetrad as stated in (4), also electrolytically in separation from Cu (Nissen- son, Z. angew., 1893, 646). Pb IV is reduced to Pb° in presence of dilute H 2 S0 4 by nascent hydrogen, and by all metals capable of producing nascent hydrogen (such as Al , Zn , Sn , Mg , Fe), and to Pb" by soluble compounds of Hg', Sn", Sb'", As'", (AsH 3 gas), Cu', Fe", Cr'", Mn", Mn'", Mn IV , Mn VI . Also by H 2 C 2 4 , HN0 2 , H 3 P0 2 , H 3 P0 3 , P , S0 2 , H 2 S , HC1 , HBr , HI , HCN , HCNS , H 4 Fe(CN) 6 , glycerine, tartaric acid, sugar, urea, and very many other organic compounds. In many cases the reduction to Pb" or to Pb° takes place in presence of K0H . The freshly precipitated peroxide oxidizes ammonia, NH 3 , to nitrite and nitrate in the course of a few hours' (Pollaeci, Arch. Pharm., 1886, 224, 176). From lead solutions Zn , Mg , Al , Co , and Cd precipitate metallic lead. §58, 5a. MERCURY. 37 §58. Mercury (Hydrargyrum) Hg = 200.0 . Valence one and two. 1. Properties. — Specific gravity, liquid, 13.5953 (Volkmann, W. A., 1881, 13, 209) ; solid, 14.1932 (Mallet, Proc. R. Soc, 1877, 26, 71). Melting (freezing) point, ^-38.85°, (Mallet, Phil. Mag., 1877, (5), 4, 145). Boiling point, 357.33° at 760 mm. (Ramsay and Young-, J. C, 1885, 47, 657). It is the only metal which is a liquid at ordinary temperatures, white when pure, with a slightly bluish tinge, and having- a brilliant silvery lustre. The precipitated or finely divided mercury- appears as a dark gray powder. Mercury may be " extinguished " or " dead- ened," i. e., reduced to the finely divided state, by shaking with sugar, grease, chalk, turpentine, ether, etc. It is slightly volatile even at — 13° (Regnault, C. r., 1881, 93, 308); is not oxidized by air or oxygen at ordinary temperature (Shenstone and Cundall, J. C, 1887, 51, 619). The solid metal is composed of octahedral and needle-shaped crystals, is very ductile and is easily cut with a knife. Owing to its very strong cohesive property it forms a convex surface with glass, etc. It is a good conductor of electricity, and forms amalgams with Al, Ba, Bi, Cd, Cs, Ca , Cr , Co, Cu , Au , Fe , Pb , Mg , Mn , Ni , Os , Pd , Pt , K , Ag , Na , Tl , Sn , and Zn . An amalgam containing about 30 per cent of copper is used for filling teeth (Dudley, Proc. Am. Assc. for Adv. of Sci., 1889, 145). 2. Occurrence. — The principal ore of mercury is cinnabar, HgS , red, found in California, Illyria, Spain, China, the Ural, and some other localities. The free metal is sometimes found in small globules in rocks containing the ore. It is also found amalgamated with gold and silver, and as mercuric iodide and mercurous chloride. 3. Preparation. — (a) The ore is roasted with regulated supply of air: HgS + 2 = Hg + S0 2 . (6) Lime is added to the ore, which is then distilled: 4HgS + 4Ca0 = 3CaS + CaS0 4 + 4Hg . (c) The ore is heated with iron (smithy scales) : Hg , FeS , and S0 2 are produced. The mercury is usually con- densed in a trough of water. Commercial mercury is freed from dirt and other impurities by pressing through leather or by passing through a cone of writ- ing paper having a small hole in the apex. For the separation of mercury from small quantities of Pb , Sn , Zn , and Ag without distilling, see Briihl (B., 1879, 12, 204), Meyer (B., 1879, 12, 437), and Crafts (Bl., 1888, (2), 49, 856). 4. Oxides. — Mercury forms two oxides, Hg 2 and HgO . Mercurous oxide, Hg 2 , is a black powder formed by the action of fixed alkalis on mercurous salts. It is converted by gentle heat into Hg and HgO and by a higher (red) heat, to Hg and . Mercuric oxide, HgO , is made (1) by keeping Hg at its boiling point for a month or longer in a flask filled with air; (2) by heating HgNO a or Hg(NO s ) 2 with about an equal weight of metallic mercury: Hg(NO s ) 2 + 3Hg = 4HgO + 2NO; (8) by precipitating mercuric salts with KOH or NaOH . Made by (1) and (2) it is red, by (3) yellow. On heating it changes to vermillion red, then black, and on cooling regains its original color. A red heat decomposes it completely into Hg and O . Mercury forms no hydroxides. 5. Solubilities. — a. — Metal. — Unaffected by treatment with alkalis. The most effective solvent of mercury is nitric acid. It dissolves readily in the dilute acid hot or cold; with the strong acid, heat is soon generated; and with con- siderable quantities of material, the action acquires an explosive violence. At ordinary temperatures, nitric acid, when applied in excess, produces normal mercuric nitrate, but when the mercury is in excess, mercurous nitrate is formed; in all cases, chiefly nitric oxide gas is generated. Both mercurous and mercuric nitrates require a little free nitric acid to hold them in solution. This free nitric acid gradually oxidizes mercurosum to mercuricum, making a clear solution of Hg(WO„) 2 , if there is sufficient HUO, present, otherwise a basic mercuric nitrate may precipitate. A solution of mercurous nitrate may be kept free from mercuric nitrate by placing some metallic mercury in the bottle containing it; still after standing some weeks a basic mercurous nitrate crystallizes out, which a fresh supply of nitric acid will dissolve. Sulphur attacks mercury even in the barometric vacuum, forming HgS (Schrotter, J. C, 1873, 26, 476). H 2 S0 4 concentrated at 25° has no action on Hg (Pitman, 38 MERCURY. §58, 56.. J. Am. Soc, 1898, 20, 100). With the hot concentrated acid S0 2 is evolved and. Hg 2 S0 4 is formed if Hg be in great excess; HgS0 4 if the H 2 S0 4 be in excess. Hydrochloric acid gas at 200° is without action (Berthelot, A. Oft., 1856, (3), 46, 492) ; also the acid sp. gr., 1.20. Bailey and Fowler (J. C, 1888, 53, 759) say that dry hydrochloric acid gas in presence of oxygen and mercury, at ordinary tem- perature for three weeks, forms' Hg.OCl, without evolution of hydrogen: 2Hg + 2HC1 + 2 = Hg 2 OCl 2 ,H 2 . Hydrobromic and hydripdic acids, gases, both attack mercury, evolve H , and form respectively HgBr and Hgl (Ber- thelot, I. a). Hydrosulphuric acid, dry gas, at 100° does not attack dry Hg (Berthelot, Uc). Hydrosulphuric acid, in solution, and alkali sulphides form HgS . Chlorine, bromine and iodine, dry or moist, attack the metal: mercurous salts are formed if the mercury be in excess,' mercuric salts if the halogens be- in excess. 6. — Oxides. — Mercurous oxide is insoluble in water or alkalis. Hydrochloric- acid forms HgCl; sulphuric acid forms Hg,SO, , changed by boiling with excess of acid to HgS0 4 ; nitric acid forms HgNO a , changed by excess of acid to Hg(NO„)j . Mercuric oxide is soluble in acids, insoluble in alkalis, soluble in 20,000 to 30,000 parts water (Bineau, C. r., 1855, 41, 509). It is decomposed by alkali chlorides forming HgCl 2 * (Mialhe, A. Ch., 1842, (3), 5, 177), soluble in NH 4 C1, from which solution NH 4 OH precipitates NH 4 Cl,lTHgH: 2 Cl + NH 2 HgCl (Ditte, C. r., 1891, 112, 859), soluble in KI 2 forming 2KI,HgI 2: (Jehn, J. C„ 1872, 25, 987). c. — Salts. — Mercury forms two well marked classes of salts — mercurous. monovalent, and mercuric, divalent — most mercurous compounds are per- manent in the air, but are changed by powerful oxidizing agents to mercuric compounds. The latter are somewhat more stable, but are changed by many reducing agents, first to mercurous compounds and then to metallic mercury (10). Solutions of mercury salts redden litmus. Many of the salts of mercury are either insoluble in water, or require the presence of free acid to keep them in solution, being decomposed by water at a certain degree of dilution, precipitating a basic salt and leaving an acid salt in solution. Mercurous chloride, bromide, and iodide are insolu- ble in water; the sulphate is soluble in 500 parts cold and 300 parts hot water, soluble in dilute nitric acid (Wackenroder, A., 1842, 41, 319). The acetate has about the same solubilities as the sulphate. Mercurous nitrate is completely soluble in water. On standing it gradually changes to- mercuric nitrate, prevented by the presence of free mercury, but if free mercury be present a precipitate of basic mercurous nitrate gradually forms. Mercuric chloride is soluble in 16 parts of cold water and 3 parts * The Law of Mass-Action requires that where the constituents of a slightly-ionized substance are present that substance shall form at the expense of those more strongly ionized. Such a slightly-ionized body is HgCl 2 . When.HgO is brought into contact with KC1 solution Hg and CI combine to form the non-dissociated HgCl.,, leaving K and O, which unite with water, im- parting to the solution a strong alkaline reaction. KBr and KI act even more strongly. HgO, although from the ready decomposition of its salts by water and from its easy reducibility a weak base, yet will replace the alkali metals where a little-dissociated Hg compound results. An excess of Hg(NO,), dissolves chloride, bromide, and iodide of Hg and Ag owing to the same cause, the Hg" ions of the stroDgly dissociated nitrate decreasing the already slight dissociation of the mercuric haloids i§44). The failure of HgCl, to give many of the pre- cipitation-reactions obtainable with other soluble mercuric salts is of course due to the sama fact— the slight concentration of Hg" ions (§ 45). §58, 6a. MERCURY. . 39 warm water; the bromide is soluble in 94 parts water at 9° and 4-5 parts at 100°, decomposed by warm nitric or sulphuric acids; the iodide is soluble in about 25,000 parts water (Bourgoin, A. Ch., 1884 (6), 3, 429), soluble in Na 2 S 2 3 (Eder and Ulen, J. C, 1882, 42, 806), and in many alkali salts, forming double salts. Normal mercuric sulphate is decom- posed by water into a soluble acid sulphate and the basic sulphate, HgS0 4 , 2HgO , which is practically insoluble (soluble in 43,478 parts water at 16°, Cameron, Analyst, 1880, 144). The normal nitrate is deliquescent, very soluble in a small amount of water, but more water precipitates the nearly insoluble basic nitrate, 3HgO.N 2 5 , changed by repeated washing into the oxide, HgO (Millon, A. Ch., 1846 (3), 18, 361). The basic nitrate is soluble in dilute nitric acid. The cyanide is soluble in eight parts water at 15°. The acetate is readily soluble, the chromate and citrate sparingly, and the sulphide, iodide, iodate, basic carbonate, oxalate, phosphate, arse- nate, arsenite, ferrocyanide, and tartrate are insoluble in water. 6. Reactions, a. — Fixed alkali hydroxides precipitate, from solutions of mercurous salts, mercurous oxide, Hg 2 , black, insoluble in alkalis, readily transposed by acids; from solutions of mercuric salts, the alkali, added short of saturation, precipitates reddish-brown basic salts, when added in excess, the orange-yellow mercuric oxide, HgO , is precipitated. If the solution of mercuric salt be strongly acid no precipitate will be obtained owing to the solubility of the mercuric oxide in the alkali salt formed; or, in the language of the Dissociation Theory, owing to the slight dissocia- tion of the soluble mercuric salt (§45). Ammonium hydroxide and car- bonate precipitate from solutions of mercurous salts mixtures of mercury and mercuric ammonium compounds. The same is true of the action of ammonium hydroxide on insoluble mercurous salts : 2HgCl + 2NH 4 0H = Hg + NH 2 HgCl + 2H 2 + NH 4 C1 ; 6HgN0 3 + 6NH 4 0H = 3Hg + (NH 2 HgN0 3 ) 2 HgO + 4NH 4 N0 3 + 5H 2 ; 4Hg 2 S0 4 + 8NH 4 0H = 4Hg + (HgH 2 N) 2 S0 4 .2HgO + 3(NH 4 ) 2 S0 4 + 6H 2 ; or uniting the salt in dif- ferent manner, 4HgCl + 4NH 4 0H = 2Hg + Hg 2 NCl.NH 4 Cl + 2NH 4 C1 + 4H 2 . Examination with a microscope reveals the presence of Hg° . The mercuric ammonium precipitate dissolves in a saturated solution of (NH 4 ) 2 S0 4 containing ammonium hydroxide and can thus be separated from the Hg (Francois, J. Pharm., 1897 (6), 5, 388; Turi, Oazzetta, 1893, 23, ii, 231; Pesci, Gazzetta, 1891, 21, ii, 569; Barfoed, J. pr., 1889, (2), 39, 201). With mercuric salts ammonium hydroxide produces " white precipi- tate," recognizable in very dilute solutions; that with cold neutral solu- tions of mercuric chloride being mercurammonium chloride, (NH 2 Hg)Cl , also called nitrogen dihydrogen mercuric chloride (a); if the solution be hot and excess of ammonium hydroxide be added, dimercurammonium chloride, also called nitrogen dimereuric chloride (b) is formed. Treat- 40 MERCURY. §58, 66. ing with fixed alkali hydroxide until no' more ammonia is evolved changes the former compound to the latter (Pesei, I. c). The precipitates are easily soluble in hydrochloric acid, slightly soluble in strong ammonium hydroxide, and more or less soluble in ammonium salts, especially am- monium nitrate and carbonate (Johnson, C. N., 1889, 59, 234). A soluble combination of ammonium chloride with mercuric chloride, 2NH 4 C1. HgCl 2 , or ammonium mercuric chloride, called " sal alembroth," is not precipitated by ammonium hydroxide, but potassium hydroxide precipi- tates therefrom the white mercurammonium chloride, (NH 3 ) 2 HgCl 2 (c) : (a) HgCl 2 + 2NH 4 OH = NH 2 HgCl + NH 4 C1 + 2H 2 (6) 2HgCl 2 + 4NH 4 OH = NHg 2 Cl + 3NH 4 C1 + 4H 2 (C) 2ITH 1 Cl.HgCl 2 + 2KOH = (NH 3 ) 2 HgCl 2 + 2KC1 + 2H 2 A solution of HgCl 2 in KI with an excess of KOH (Nessler's Reagent) is precipitated by NH 4 0H (or by ammonium salts), as NHg 2 I (§207,. 6k). Fixed alkali carbonates precipitate from mercurous salts an unstable mer- curous carbonate, Hg 2 C0 3 , gray, blackening to basic carbonate and oxide when heated. Carbonates of barium, strontium, calcium and magnesium precipitate mercurous carbonate in the cold. Mercuric salts are precipitated as red-brown basic salts, which, by excess of the reagent with heat, are converted into the yellow mercuric oxide. The basic salt formed with mercuric chloride is an oxy- chloride, HgCl 2 .(HgO) 2 , s , or 4 ; with mercuric nitrate, a basic carbonate, (Hg0) s HgC0 3 . Barium carbonate precipitates a basic salt in the cold, from the nitrate, but not from the chloride. 6. — Oxalic acid and soluble oxalates precipitate from solutions of mercurous salts mercurous oxalate, Hg 2 C 2 4 , white, slightly soluble in nitric acid; from solutions of mercuric salts, except HgCl 2 , mercuric oxalate, HgC 2 4 , white, easily soluble in hydrochloric acid, difficultly soluble in nitric acid. A solution of HgCl 2 boiled in the sunlight with (1TH 4 ) 2 C 2 4 gives HgCl and C0 2 . Hydrocyanic acid and alkali cyanides decompose mercurous salts into me- tallic mercury, a gray precipitate, and mercuric cyanide, which remains in solution. Mercuric salts are not precipitated, since the cyanide is readily soluble in water. Soluble ferrocyanides form with mercurosum a white gela- tinous precipitate, soon turning bluish green; with mercuricum a white precipi- tate, becoming blue on standing. Soluble ferricyanides form with mercurous salts a yellowish green precipitate; with mercuric salts a green precipitate, soluble in hydrochloric acid. Potassium thiocyanate precipitates mercurous thiocyanate, HgCNS , white, from solutions of mercurous salts (Claus, J. pr., 1838, ' 15, 406) ; from solutions of mercuric salts, mercuric thiocyanate, Hg(CNS) 2 , soluble in hot water (Philipp, Z. Oft., 1867, 553). c. — Nitric acid never acts as a precipitant of mercury salts, the salts being more soluble in strong nitric acid than in water or the dilute acid; also nitric acid dissolves all insoluble salts of mercury except HgS , which is insoluble in the hot acid (sp. gr. 1.42) (Howe, Am., 1887, 8, 75). HgCl is slowly dissolved by xiitric acid on boiling. All mercurous salts are oxidized to mercuric salts by ■excess of nitric acid. d. — Hypophosphorous acid reduces mercuric salts to Hg°, but the presence of hydrogen peroxide causes the formation of HgCl from HgCl, and is of value as a quantitative method for estimation of mercury (Vanino and Treubert, B., 1897,30, 1999). Phosphoric acid and alkali phosphates precipitate, from mercurous salts, mercurous phosphate, Hg 3 P0 4 , white, if the reagent be in excess; but if HgNO., "be in excess, Hg 3 P0 4 .HgN0 3 , white, with a yellowish tinge. Mercurous phos- jhate is soluble in dilute HNO a , insoluble in H a P0 4 . From mercuric nitrate, §58, 6e. MERCURY. 41 mercuric phosphate, Hg a (P0 4 ) 2 , white, is precipitated, soluble in HNO„ , HC1 , and ammonium salts, insoluble in H 3 P0 4 . Phosphoric acid does not precipitate HgCl 2 , and Na 2 HP0 4 does not precipitate the white Hg 3 (P0 4 ) 2 from HgCl 2 , but on standing a portion of the mercury separates as a dark brown pre- cipitate (Haaek, J. C, 1891, 60, 400; 1892, 62, 530). e.— Hydrosulphuric acid and soluble . sulphides, precipitate from mer- curous salts, mercuric sulphide, HgS , black, and mercury, gray. Mercurous sulphide, Hg 2 S , does not exist at ordinary temperatures. According- ,to Antony and Sestini (Gazzetta, 1894, 24, i, 193), it is formed at — 10° by the action of H 2 S on HgCl,. decomposing at 0° into HgS and Hg . From mercuric salts there is formed, first, a white precipitate, soluble in acids and excess of the mercuric salts, on further additions of the reagent, the precipitate becomes yellow T orange, then brown, and finally black. This progressive variation of color is characteristic of mercury. The final and stable black precipitate is mercuric sulphide, HgS ; the lighter colored, precipitates , consist of unions of the original mercuric salt with mercuric sulphide, as HgCl 2 .HgS , the proportion of HgS being greater with the- darker precipitates. When sublimed and triturated, the black mercuric sulphide is converted to the red (vermillion), without chemical change. Mercuric sulphide is insoluble in dilute HN0 3 (distinction from all other metallic sulphides) ; insoluble inHCl (Field, J. C, 1860, 12, 158); soluble in chlorine (nitro-hydrochloric acid); insoluble in (NH 4 ) 2 S except when K0H or NaOH be present (Volhard, A., 1891, 255, 252); soluble in K 2 S (Ditte,. C. r., 1884, 98, 1271), more readily if K0H be present (separation from Pb , Ag , Bi , and Cu) (Polstorff and Biilow, Arch. Pharm., 1 891, 229, 292). It is soluble in K 2 CS 3 (one part S , two parts CS 2 , and 23 parts K0H , sp. gr. 1.13) (separation from Pb , Cu , and Bi); reprecipitated as sulphide by HC1 (Eosenbladt, Z., 1887, 26, 15), Mercurous nitrate forms with sodium thiosulphate a grayish black precipi- tate, part of the mercury remaining in solution. Mercurous chloride forms metallic mercury and some mercury salt in solution as double salt (Schnauss, J. C, 1876, 29, 342). Mercuric chloride added to sodium thiosulphate forms a white precipitate, which blackens on standing; if the mercuric chloride be- added in excess a bright yellow precipitate is formed, which blackens when boiled with water, nitric acid or sulphuric acid, but does not dissolve or blacken on boiling with hydrochloric acid. Sodium thiosulphate added to- mercuric chloride forms a white precipitate, which blackens on standing or on adding excess of thiosulphate, but if excess of thiosulphate be rapidly added to- HgCl 2 no precipitate is formed; boiling or long standing produces the black precipitate. Mercuric salts are not completely precipitated by sodium thio- sulphate. The black precipitate is HgS. Sulphurous acid and soluble sulphites form from mercurous solutions a black precipitate (Divers and Shimidzu, /. C, 1886, 49, 567). Mercuric nitrate with sulphurous acid forms slowly a flocculent white precipitate soluble in nitric acid. The precipitate and solution contain mercurosum as evidenced by HC1 . Mercuric nitrate with soluble sulphites forms a voluminous white pre- cipitate, soluble in HNO, and containing mercurosum. Mercuric chloride is not precipitated by sulphurous acid or sulphites in the cold, but is reduced, by boiling with sulphurous acid, to HgCl and then to Hg° - 42 MERCURY. §58, 6f. Sulphuric acid and soluble sulphates precipitate from mercurous solu- tions not too dilute, mercurous sulphate, Hg 2 S0 4 , white, decomposed by boiling water, sparingly soluble in cold water (5c), soluble in nitric acid and blackened by alkalis. Mercuric salts are not precipitated by sulphuric acid or sulphates. For action of H 2 S0 4 on HgCL see next paragraph and (§269, 8, footnote). f. — Hydrochloric acid and soluble chlorides precipitate from solutions of mercurous salts, mercurous chloride, HgCl , " Calomel," white, insoluble in water, slowly soluble in hot concentrated HC1 . Boiling nitric acid slowly dissolves it, forming Hg(N0 3 ) 2 and HgCl 2 ; dissolved by chlorine or nitro- hydrochloric acid to HgCl 2 ; soluble in Hg(N0 3 ) 2 (5& footnote) (Dreschsel, J. C, 1882, 42, 18). This precipitation of mercurous salts by hydro- chloric acid is a sharp separation from mercuric salts and places mer- curous mercury in the Fiest (Silvee) Group of Metals. Mercuric salts are not precipitated by hydrochloric acid or soluble chlorides, unless the mercuric solution is more concentrated than possible for a mercuric chloride solution under the same conditions, i. e., a strong solution of Hg(N0 3 ) 2 gives a precipitate of HgCl 2 on addition of HC1, soluble on addition of water. Mercuric chloride is not decomposed by sulphuric acid. A compound HgCl 2 .H 2 S0 4 is formed which sublimes undecom- posed. The same compound is formed when HgS0 4 is treated with HC1 and distilled (Ditte, A. Ch., 1879, (5), 17, 120). Hydrobromic acid and soluble bromides precipitate, from solutions of mercurous salts, mercurous bromide, HgBr, yellowish white, insoluble in water, alcohol, and dilute nitric acid; from concentrated solutions of mercuric salts, mercuric bromide, HgBr 2 , white, decomposed by concen- trated nitric acid. Mercuric bromide is soluble in excess of mercuric salts (56 footnote), or in excess of the precipitant; hence, unless added in suitable proportions, no precipitate will be produced. Sulphuric acid does not transpose HgBr, but forms compounds exactly analogous to those with HgCl 2 . Excess of concentrated H 2 S0 4 gives some Br with HgBr 2 . Hydriodic acid and soluble iodides precipitate from solutions of mer- curous salts, mercurous iodide, Hgl, greenish yellow — "the green iodide of mercury " — nearly insoluble in water, insoluble in alcohol (distinction from mercuric iodide), soluble in mercurous and mercuric nitrates ; decom- posed by soluble iodides with formation of Hg and Hgl 2 , the latter being dissolved as a double salt with the soluble iodide : 2HgI + 2KI = Hg -f- HgI 2 .2KI . Mercurous chloride is transposed by HI or KI to form Hgl , excess of the reagent reacts according to the above equation (D., 2, 2, 867). Ammonium hydroxide in the cold decomposes Hgl into Hg and Hgl 2 (Francois, J. Pharm., 1897, (6), 5, 388). Mercuric salts are precipitated as mercuric iodide, Hgl 2 , first reddish- •§58, 7. MEBCVRY. 43 jellow then red, soluble in 24,814 parts of water at 17.5° (Bourgoin, A. Ch., 1884, (6), 3, 429), soluble in concentrated nitric and hydrochloric acids; ■quickly soluble in solutions of the iodides of all the more positive metals, -i. e. in excess of its precipitant, by formation of soluble double iodides; as (KI) 2 HgI 2 variable to KIHgI 2 . A hot concentrated solution of potas- sium iodide dissolves 3HgI 2 for every 2KI. The first crystals from this •solution are KIHgI 2 . These are decomposed by pure water, and require a little alkali iodide for perfect solution, but they are soluble in alcohol -and ether. A solution of dipotassium mercuric tetraiodide, K 2 HgI 4 = (KI) 2 HgI 2 (sometimes designated the iodo-hydrargyrate of potassium), is ^precipitated by ammonium hydroxide as mercurammonium iodide, NHg 2 I ^Nessler's test), and by the alkaloids (Mayer's reagent). Potassium bromate precipitates, from solutions of mercurous nitrate, mer- curous bromate, HgBr0 3 , white, soluble in excess of mercurous nitrate and in nitric acid; from solutions of mercuric nitrate, mercuric bromate, Hg(BrO s ) 2 , soluble in nitric acid, hydrochloric acid, and in excess of mercuric nitrate, soluble in 650 parts of cold and 64 parts of hot water (Rammelsberg, Pogg., 1842, 55, 79). No precipitate is formed when potassium bromate is added to mercuric chloride (56, footnote). Iodic acid and soluble iodates precipitate solutions of mercurous salts as mercurous iodate, HgI0 3 , white with yellowish tint, solu- ble with difficulty in dilute nitric acid, readily soluble in HC1 by oxidation to mercuric salt. Mercuric nitrate is precipitated as mercuric iodate, Hg(IO a ) 2 , white, soluble in HC1 , insoluble in HNO,' and H,S0 4 , soluble in NHjCl , trans- posed and then dissolved by KI . Mercuric chloride is not precipitated by JEIO a (56, footnote) (Cameron, C. N., 1876, 33, 253). g. — Arsenous acid or arsenites form a white precipitate with mercurous nitrate, soluble in HNO, (Simon, Pogg., 1837, 40, 442). Mercuric nitrate is precipitated by a solution of arsenous acid; the precipitate is soluble in HNO, (D., 2, 2, 920). Arsenic acid or Na,HAs0 4 precipitates from mercurous nitrate 3Hg a As0 4 .HgN0 3 .H 2 , light yellow if the HgN0 3 be in excess (D., 2, 2, 921); ■dark red Hg 3 As0 4 if the arsenate be in excess. Hg 3 As0 4 is changed by cold HC1 to HgCl and H 3 As0 4 , by boiling with HC1 to Hg° , HgCl 2 , and H 3 As0 4 ; and is soluble unchanged in cold HNO s , insoluble in water and acetic acid (Simon, Pogg., 1837, 41, 424). Arsenic acid and soluble arsenates precipitate from mercuric nitrate, Hg 3 (As0 4 ) 2 , white, soluble in H1T0 S and HC1 , slightly .soluble in water. , Arsenic acid and potassium arsenate do not precipitate mercuric chloride from its solutions. Stannous chloride precipitates solutions of mercuric salts (by reduction), as mercurous chloride, white; or if the stannous chloride be in excess, as metallic mercury (a valuable final test for mercuric salts) (10). ji. — Soluble chromates precipitate from mercurous solutions mercurous -chromate, Hg 2 Cr0 4 , brick-red, insoluble in water, readily transposed by HC1 to HgCl and H 2 Cr0 4 , soluble with difficulty in HNO, without oxidation (Kichter, B., 1882, 15, 1489). Mercuric nitrate is precipitated by soluble chromates as a light yellow precipitate, rapidly turning dark brown, easily soluble in dilute • acids and in HgCl,. Mercuric chloride forms a precipitate with normal chro- mates, but not with K 2 Cr,0 7 . 7. Ignition. — Mercury from all its compounds is volatilized by heat as vthe undecomposed salt or as the free metal. Mercurous chloride (Debray, 44 MERCURY. §58, 8. J. C, 1877, 31, 47) and bromide and mercuric chloride and iodide sublime (in glass tubes) undecomposed — the sublimate condensing (in the cold part of the tube) without change. Most other compounds of mercury are decomposed by vaporization, and give a sublimate of metallic mercury (mixed with sulphur, if from the sulphide, etc.). All compounds of mer- cury, dry and intimately mixed with dry sodium carbonate, and heated in a glass tube closed at one end, give a sublimate of metallic mercury as a gray mirror coat on the inner surface of the cold part of the tube. Under the magnifier, the coating is seen to consist of globules, and by gently rubbing with a glass rod or a wire, globules visible to the unaided eye are obtained. 8. Detection. — Mercury in the mercurous condition belongs to the first geoup (silver group), and is completely precipitated by HC1 . It is iden- tified by the action of ammonium hydroxide, changing the white precipi- .tate of mercurous chloride to the black precipitate of metallic mercury 'and nitrogen dihydrogen mercuric chloride (a delicate and characteristic . test for Hg')- Mercury in the mercuric condition belongs to the seconi> group (tin and copper group), and is separated from all other metals of that group by the non-solubility of the sulphide in (NH 4 ) 2 S X and in dilute HN0 3 . The sulphide is dissolved in nitrohydrochloric acid, and the pres- ence of mercury confirmed by the precipitation of Hg° on a copper wire, or by the reduction to HgCl or Hg c by SnCl 2> . 9. Estimation. — (a) As metallic mercury. The mercury is reduced by means, of CaO in a combustion-tube at a red heat in a current of C0 2 . The sublimed mercury is condensed in a flask of water, and, after decanting 1 the water, dried in a bell-jar over sulphuric acid without application of heat. The mercury may also be reduced from its solution by SnCl 2 (or H 3 P0 3 at 100°) and dried as above. (6) As mercurous chloride. It is first reduced to Hg' by H 3 PO s (Uslar,. £.,1895, 34, 391), which must not be heated above 60°, otherwise metallic mer- cury will be formed; and after precipitation by HC1 and drying- on a weighed filter at 100°, it is weighed as HgCl . Or enough HC1 is added to combine with the mercury, then the Hg" is reduced to Hg' by FeS0 4 in presence of MaOH : 2HgO + 2PeO + 3H 2 = Hg 2 + Si'e(OH) s . H 2 S0 4 is added, which causes the formation of HgCl , which is dried on a weighed filter at 100°. (c) As HgS . It is precipitated by H 2 S, and weighed in same manner as the chloride. Any free sulphur mixed with the precipitate should be removed by CS 2 . (° >> h;g o" toiS to W £ W fe oo •a 00 * ** '0 W 4S CO in *™ PU aj b d • 3 cat; P o o g-co" P o «» ■p o--- - "lj oj-dH *IB H ^ c £ to O K to-dfi H p,3» 8'S p *-^ toj>> P^ « -fl ID o SS o £a a p to i>3 - o Q »p-"3 fl 0Q §p? - •MO OS CO io Cj =os .•S fn" W ^-p w i— r w •|o ^ a "3 so "o" . CO &7o; :s -d'W 2°. W in" SOS >>^ "3 PI ft-" « a) h -a to^ 2 3 fl os^ > cam -t",^ »,_, OS OS W o ^0««S^ --^ to-p ^ in m '-' to to o ;coea»OS>Oh sosH ftp1«"W p«»o p ^^^Q CQ -P W ,Q 'P ^-^.Q *P Pi Pi Ph xs fi 60S T) TJ 5 5 3 cd A » 3 So §63, 6«. DIRECTIONS FOB ANALYSIS WITH NOTES. 53 DIRECTIONS FOE THE ANALYSIS OF THE METALS OF THE FlRST GbOTJP. §62. Manipulation. — To the solution add hydrochloric acid (whenever •directions call for the addition of a reagent it is to be used reagent istrength unless otherwise stated) drop by drop (§32) until no further precipitate is formed and the solution is distinctly acid to litmus (§36). The precipitate will consist of the chlorides of Pb , Hg', and Ag , e. g., :Pb(N0 3 ) 2 + 2HC1 = PbCl 2 + 2HN0 3 . Shake thoroughly and allow to istand a few moments before filtering; if the solution is warm it should be cooled to the temperature of the room. Decant the solution and precipitate upon a filter paper previously wetted (§35) with water and wash two or three times with cold water or until the filtrate is not strongly acid to litmus. The washings with cold water should be added to the first filtrate and the whole marked and set aside to be tested for the metals of the remaining groups (§16). §63. Notes. — 1. Failure to obtain a precipitate upon the addition of HC1 to an acid reaction is proof of the absence of Hg 7 and Ag , but a solution of a lead salt may be present, of such a degree of dilution that the lead chloride formed will be soluble in the dilute acid (§57, 5c). 2. The solution should not be strongly acid with nitric acid, as it forms nitrohydrochloric acid with the hydrochloric acid, causing oxidation of the Hg' (§58, 5c). Lead chloride is also more soluble in nitric acid than in dilute hydrochloric acid (§57, 5c). By a study of the solubilities of the silver group metals it will be seen that H = S0 4 , HC1 , HBr or HI cannot be used in prepar- ing a solution for analysis when these metals are present. 3. A great excess of acid is to be avoided, as it may interfere with the reac- tion in Group II. (§57, 6e). Complete precipitation should be assured by testing the nitrate with a drop of HC1 , when no further precipitation should occur (§32). If a white precipitate is formed by adding a drop of HC1 to the filtrate it is evident that the precipitation was not complete and more HC1 should be added and the group separation repeated (§11). 4. The presence of a slight excess of dilute acid does not aid or hinder the precipitation of the Hg' or Ag , but as PbCl 2 is less soluble in dilute HC1 than in water, a moderate excess of the acid causes a more complete precipita- tion of that metal in the first group. 5. Concentrated HC1 dissolves the chlorides of the first group quite appre- ciably (§59, 5c). 6. Hydrochloric acid added to certain solutions may cause a precipitate when none of the first group metals are present. Some of the more important conditions are mentioned: a. A concentrated solution of BaCL is precipitated without change by the addition of HC1 , readily soluble in water (§186, 5c). b. An acid solution of Sb , Bi , or Sn , with some other acid than HC1 , ;and saturated with water as far as possible without precipitation, on the addition of HC1 , precipitates the oxychloride of the corresponding metal (§76, 6f). These precipitates are readily soluble in an excess of the HC1 . It must, however, be remembered that a trace of AgCl will also be dissolved by an excess of HC1 (§59, 5c). c. Solutions of metallic oxides in the alkali hydroxides are precipitated when neutralized with acids, e. g., K 2 Zn0 2 + 2HC1 = Zn(OH) 2 + 2KC1 . d. The sulphides of As , Sb , Sn , Au , Pt , Mo (Ir , W , Ge , V , Se and Te) in solution with the alkali polysulphides are reprecipitated together with ■sulphur on the addition of HC1 (§69, 6e). e. Soluble polysulphides and thiosulphates give a precipitate of sulphur, -white, with HC1 (§256, 3a). 54 DIRECTIONS FOR ANALYSIS WITH NOTES. §63, 6/L f. Certain soluble double cyanides, as Ni(CN) 2 .2KCM' , are precipitated as insoluble cyanides, Ni(CN') 2 , on the addition of HC1 (§133, 66). g. Solutions of silicates (§249, 4), borates, tungstates, molybdates; also- benzoates, salicylates, urates, and certain other organic salts, are precipitated, by acidulation with HC1 , many of the precipitates being soluble on further- addition of the acid. h. Acidulation with HC1 may induce changes of oxidation or reduction, which in certain mixtures may result in precipitation: for example, Cu" salts. with KCNS in ammoniacal solution (§77, 66) ; mixture of solutions of KI and KI0 3 (§280, 6,B,7), etc. 7. If the precipitate, obtained by the addition of HC1 to the solution, is. colored or does not give further reactions which are conclusive and perfectly satisfactory in every respect, it should be separated' by filtration, and treated as a solid substance taken for examination (see conversion of solids into- liquids, §301). 8. Compounds of the first group metals insoluble in water or acids are trans- posed to sulphides by digestion with an alkali sulphide. The lead and silver sulphides thus formed are readily soluble in hot dilute nitric acid. The mer- curous compounds are changed to mercuric sulphide (§58, 5a and 6e), a, second group mercury compound insoluble in HN0 3 . 9. If but one metal of the first group be present, the action of NH t 0E determines which it is; PbCl 2 does not change color or dissolve; HgCl blackens^ and AgCl dissolves (§60). §64. Manipulation. — The precipitate (white) on the filter should now be washed once or twice with hot water. The first hot water should be poured upon the precipitate a second time. This hot filtrate is divided into four portions and each portion tested separately for lead with the- following reagents, H 2 S0 4 , H 2 S , X 2 CrtOy> and KI (§57, 6 e, h, and /) : PbCl 2 + H 2 S0 4 = PbS0 4 (white) + 2HC1 PbCl 2 + H 2 S = PbS (black) + 2HC1 2PbCl 2 + K 2 Cr 2 7 + H 2 = 2PbCr0 4 (yellow) + 2KC1 + 2HC1 PbCl 2 + 2KI = Pbl 2 (yellow) + 2KC1 The yellow precipitate with potassium iodide (the KI must not be used in great excess (§57, 5c)) should be allowed to settle, the liquid decanted,, and the precipitate redissolved in hot water, to a colorless solution which upon cooling deposits beautiful yellow crystalline scales of Pbl 3 (charac- teristic of lead). §65. Notes. — 1. Lead is never completely precipitated in the first group (§57, 6f). The presence of a, moderate excess of dilute HC1 and the cooling of the solution both favor the precipitation. 2. Lead can be completely separated from the second group metals by sul- phuric acid applied to the original solution (§57, 6e, §95 and §98), but that would necessitate a regrouping of the metals; as, Ba , Sr , and Ca would also* be precipitated (Zettnow, Z., 1867, 6, 438). 3. In order to precipitate the lead chloride, not removed in the first group, in the second group with H 2 S , the solutions must not be strongly acid, either the excess of HC1 should be removed by evaporation or the solution should be diluted (§57, 6e, and §81, 3, 5 and 9). 4. If the lead chloride is not all washed out with hot water it is changed to an insoluble basic salt (white) by the NH,0H , part remaining on the filter and part carried through mechanically which causes turbidity to the am- monium hydroxide solution of the AgCl and makes necessary the filtration of that solution before the addition of HU0 3 , otherwise it does not interfere. 5. The precipitation of lead as the sulphide while not characteristic of lead. §68, 3. DIRECTIONS FOB ANALYSIS WITH NOTES. 55 is exceedingly delicate, much more so than the formation of the white PbS0 4 (§57, 5c). In extremely dilute solutions no precipitate occurs, merely a brown coloration to the solution. The presence of free acid lessens the delicacy of the test. 6. PbCr0 4 is blackened by alkali sulphides and dissolved by the fixed alkalis (important distinction from BaCr0 4 ) ; the solubility in the fixed alkalis is also- an important distinction from bismuth chromate (§76, 6ft). 7. Other tests for lead by reduction on charcoal before the blow-pipe, or in the wet way by Zn, should not be omitted (§57, 7 and 10). If to a solution of lead salt nearly neutral a strip of zinc be added, the lead will soon be deposited on the zinc as a spongy mass. §66. Manipulation. — The white precipitate remaining on the filter after washing with hot water consists of HgCl and AgCl , with usually some PbCl 2 which was not removed. To this precipitate NH 4 0H , one or two cc. is added and allowed to pass through the filter into a clean test-tube. An instantaneous blackening of the precipitate is conclusive evidence of the presence of mercurosum; 2HgCl + 2NH 4 0H = Hg + NH 2 HgCl + NH 4 C1 + 2H 2 . The AgCl is dissolved by the NH 4 0H : AgCl + 2NH 4 0H = 2NH 3 . AgCl + 2H 2 0, and is found in the filtrate; its presence being confirmed by its reprecipitation on rendering the solution acid with HN0 3 : 21TH 3 . AgCl + 2HN0 3 = AgCl + 2NH 4 N0 3 . §67. Notes.— Mercury. — 1. The black precipitate on the filter, caused by the- addition of NH,OH to the HgCl may be examined under the microscope for the detection of globules of Hg°, or the precipitate may be digested with concentrated solution of (NH 4 ) 2 S0 4 , which dissolves the MH 2 HgCl , leaving theHg" unattacked (§58, 6«). 2. If the original solution contains no interfering metals, the distinctive reactions of mercurous salts with iodides, chromates and phosphates should be obtained (§58, 6e, ft and d). 3. The precipitation with HC1 and blackening with NH 4 OH is conclusive evi- dence of the presence of mercury in the mercurous condition; should further confirmation be desired, the black precipitate may be dissolved in nitro- hydroehloric acid, the excess of acid removed by evaporation and the free metal obtained as a coating on a copper wire, by immersing the freshly polished wire in the solution of HgCl,, (§58, 10). 4. Mercury has but few soluble mercurous compounds, and in preparing solutions of the insoluble compounds for analysis, oxidizing agents are usually employed and the mercury is then found entirely in the second group as a sulphide (§96 and §97). 5. Additional proof may be obtained by mixing a portion of the black residue with sodium carbonate, drying and heating in a glass tube (read §58, 7, also §97,7). §68. Silver. — 1. The presence of a large excess of Hg(NO s ) 2 prevents the precipitation of AgCl from solutions of silver salts by HC1 (§59, 5c). In this case the metals should be precipitated by H 2 S and the well-washed precipitate digested with hot dilute HUO, . The silver is dissolved as AglT0 3 , while the mercury is unattacked: 6Ag 2 S + 16HN0 3 = 12AgN0 3 + 3S 2 + 4NO + 8H 2 . After evaporation of the excess of HN0 3 the solution may be treated with HC1 as an original solution. 2. A small amount of AgCl with a large amount of HgCl is not dissolved by NH,OH , but is reduced to Ag° by the Hg° formed by the addition of the MH 4 OH to the HgCl (§58, 60, §59, 10 and §60). 3. If Hg 7 . be present and Ag is not detected, the black precipitate on the 56 ARSENIC. §68, 4. filter should be digested for some time with (NE,),S , washed, and boiled with hot dilute nitric acid. The Ag , if any be present, is dissolved and separated from the HgS : NH 2 HgCl + (HH,),S + 2H 2 = HgS + NH 4 C1 + 2HH.0H Hg + (NH 4 ) 2 S X = HgS + (NHO.S.-! 4. If only a trace of silver be present, its detection by adding HN0 3 to the HH 4 OH solution of the chloride may fail, unless the excess of the NH 4 OH be first removed by evaporation (because of the solubility of the AgCl in the ammonium salt, §59, 5c). 5. As a further test for silver, the chloride, precipitated by the nitric acid, may be reduced to the metal by zinc; by adding to the ammoniacal solution a few drops of potassium stannite (§71, 6o and 8); by warming with grape sugar in alkaline mixture. In all cases the well-washed grayish black metal may be dissolved in nitric acid as AgN0 3 . 6. To identify the acid of silver salts which are insoluble in HNO, (AgCl , AgBr , Agl), (l) Add metallic zinc and a drop of H.S0 4 ; when the silver is all reduced test for the acid in the filtrate. (2) Fuse with Na,C0 3 , add water, and test the filtrate for acids. (3) Add H 2 S , or an alkali sulphide, digest warm for a few minutes, filter and test filtrate for acids. (-4) Boil with KOH or NaOH (free from HC1), and test the filtrate in the same manner. It must not be overlooked that by the first three methods, and not by the last, hromates and iodates are reduced to bromides and iodides (§257, 6JS). The Tin and Copper Group (Second Group). Arsenic, Antimony, Tin, Gold, Platinum, Molybdenum, Mercury, lead, Uismuth, Copper, Cadmium (Euthenium, Khodium, Palladium, Iridium, Osmium, Tungsten, Vanadium, Germanium, Tellurium, Selenium). The Tin Group (Second Group, Division A). Arsenic, Antimony, Tin, Gold, Platinum, Molybdenum (Iridium, Tungs- ten, Vanadium, Germanium, Selenium, Tellurium). §69. Arsenic. As = 75.0. Valence three and five. 1. Properties. — Specific gravity, pure crystalline 5.727 at 14°; amorphous 4.716 e size of the tube). 64 ARSENIC. §69, 67>. a Bunsen burner provided with a flame spreader. The flame should be- applied to the tube between the calcium chloride tube and the constricted portion. The tube should be supported to prevent sagging in case the- glass softens, and it is customary to wrap a few turns of wire gauze around the portion of the tube receiving the heat. The heat of the flame decom- poses the arsine and a mirror of metallic arsenic is deposited in the con- stricted portion of the tube just beyond the heated portion. This may be tested as described under c 1. When a sufficient mirror has been obtained the flame is withdrawn, and, removing the rubber tube, the- escaping gas * is ignited. 6. Arsenous Hydride (arsine), AsH, , burns when a stream of it is ignited, where it enters the air, and explodes when its mixture with air is ignited- It burns with a somewhat luminous and slightly bluish flame (distinction from hydrogen); the hydrogen being first oxidized, and the liberated arsenic becoming incandescent, and then undergoing oxidation; the vapors, of water and arsenous anhydride passing into the air: 2AsH 3 + 30 2 = As 2 3 -f- 3H 2 . If present in considerable quantity a white powder may be observed settling on a piece of black paper placed beneath the flame. If the cold surface of a porcelain dish be brought in contact with the flame the oxidation is prevented and lustrous black or brownish-black spots of metallic arsenic are deposited on the porcelain surface; 4AsH 3 -f- 30 2 = As 4 -4- 6H 2 . A number of spots should be obtained and all the tests for metallic arsenic applied. The arsenic in the silver nitrate solu- tion is present as arsenous acid and can be detected by the usual tests (6e) by first removing the excess of silver nitrate with ^dilute hydrochloric acid or calcium chloride. To generate arsine, magnesium or iron t may be used, instead of zinc, and hydrochloric acid instead of sulphuric acid. Arsine cannot be formed in the presence of oxidizing agents as the halogens, nitric acid, chlorates, hypo- chlorites, etc. Arsinuretted hydrogen (arsine) may also be produced from arsenous compovmds by nascent hydrogen generated in alkaline solution. Sodium, amalgam, J zinc (or zinc and magnesium) and potassium hydroxide or alumi- num and potassium hydroxide may be used as the reducing agent. There is no reaction with AsV , or with compounds of antimony (§70, 6/) ; hence when * Arsine is an exceedingly poisonous gas, the inhalation of the unmixed gas being quickly" fatal. Its dissemination in the air of the laboratory, even in the small portions which are not appreciably poisonous, should be avoided. Furthermore, as it is recognized or determined, in its various analytical reactions, only by its decomposition, to permit it to escape undecomposed is so far to fail in the object of its production. The evolved gas should be constantly run into silver nitrate solution, or kept burning. + According to Thiele (C. C, 1890, 1, 877) arsenic may be separated from antimony in the Marsh test by using electrolytically deposited iron instead of zinc. Stibine is not evolved. According to Sautermeist'er [Analyst, 1891, 218) arsine is not produced when hydrochloric acid acts upon iron containing arsenic, but if several grams of zinc be added a very small amount of arsenic in the iron may be detected. ± Sodium amalgam is conveniently prepared by adding (in small pieces at a time) one part of sodium to eight parts (by weight) of dry mercury warmed on the water bath. When cold the amalgam becomes solid and is easily broken. It should be preserved in well stoppered bottles. §69, 6'c. ARSENIC. 65 the arsenic is present in the triad condition (Asv may be reduced to As'" by J50 2 ) the use of one of the above reagents serves admirably for the detection of arsenic in the presence of antimony. This experiment may be made in a test-tube, the arsenic being detected by covering the tube with a piece of filter paper moistened with silver nitrate. It is very difficult to drive over the last traces of the arsenic and therefore the method is not satisfactory for quanti- tative work (Hager, J. C, 1885, 48, 838; Johnson, 0. N., 1878, 38, 301; and Clark, ■J. C, 1893, 63, 884). If ferrous sulphide contains metallic iron and arsenic, arsine may be gen- erated with the hydrogen sulphide. It cannot be removed by washing the gases with hydrochloric acid (Otto, B., 1883, 16, 2947). Arsine does not combine with hydrogen sulphide until heated to 230°, while stibine, SbH 3 , combines at ordinary temperature (method of separation) '(Brunn, B., 1889, 22, 3202; Myers, J. C, 1871, 24, 889). As dry hydrogen sul- phide is without action upon dry iodine, it may be freed from arsine by passing the mixture of the dried gases through a tube filled with glass wool inter- spersed with dry iodine. AsH a + 3I 2 = Asl 3 + 3HI (Jacobson, B., 1887, 20, 1999). Arsenous hydride is decomposed by passing through a tube heated to redness (mirror in Marsh test) 4AsH 3 = As 4 + 6H 2 . Nitric acid oxidizes it to arsenic acid, 3AsH 3 + 8HNO, = 3H 3 As0 4 -f 8NO + 4H 2 0; and may be used instead of silver nitrate to effect a separation of arsine and stibine in the Marsh test. The nitric acid solution is evaporated to dryness and the residue "thoroughly washed with water. Test the solution for arsenic with silver nitrate and ammonium hydroxide (Ag 3 As0 4 , reddish brown precipitate, 6;'). Dissolve the residue in hydrochloric or nitrohydrochloric acid and test for antimony with hydrogen sulphide (Ansell, J. C, 1853, 5, 210). c. — Comparison of the mirrors and spots obtained with arsenic and anti- mony. — 1. Both the mirror and spots obtained in the Marsh test exhibit the properties of elemental arsenic (5a). The reactions of these deposits having analytical interest are snch as distinguish arsenic from antimony. Arsenic Mirror. Antimony Mirror. Deposited beyond the flame; the Deposited before or on both sides ;gas not being decomposed much be- of the flame ; the gas being decom- low a red heat. posed considerably below a red heat. Volatilizes in absence of air at The mirror melts to minute glob- 450° (1), allowing the mirror to be ules at 432°, and is then driven at driven along the tube; it does not a red heat, melt. By vaporization in the stream of The vapor has no odor, gas, escapes with a garlic odor. By slow vaporization in a cur- By vaporization in a current of rent of air a deposit of octahedral air, a white amorphous coating is and tetrahedral crystals is obtained, obtained; insoluble in water, soluble forming a white coating soluble in in hydrochloric acid, and giving re- water and giving the reactions for actions for antimonous oxide, arsenous oxide. 66 ARSENIC. §69, 6'c The heated mirror combines with hydrogen sulphide, forming the lemon-yellow arsenous sulphide, which, being volatile, is driven to the cooler portion of the tube. The dry sulphide is. not readily attacked by dry hydrochloric acid gas (6/). Arsenic Spots. Of a steel gray to black lustre, Volatile by oxidation to arsenous oxide at 218°. Dissolve in hypochlorite.* Warmed with a drop of ammon- ium sulphide form yellow spots, soluble in ammonium carbonate, in- soluble in hydrochloric acid (6«). With a drop of hot nitric acid, dissolve clear. The clear solution, with a drop of solution of silver nitrate, when treated with vapor of ammonia, gives a brick-red precipi^ tate. The solution gives a yellow pre- cipitate when warmed with a drop of ammonium molybdate. With vapor of iodine, color yel- low, by formation of arsenous iodide, readily volatile when heated. The heated mirror combines with hydrogen sulphide forming the orange antimonous sulphide, which is not readily volatile. The sulphide is readily decom- posed by dry hydrochloric acid gas, forming antimonous chloride which is volatile, and may be driven over the unattacked arsenous sulphide. Antimony Spots. Of a velvety brown to black sur- , face. Volatile, by oxidation to anti- monous oxide, at a red heat. Do not dissolve in hypochlorite. Warmed with ammonium sul- phide, form orange-yellow spots, in- soluble in ammonium carbonate., soluble in hydrochloric acid (§70. 6e). With a drop of hot dilute nitric acid, turn white. The white fleck, by- action of nitric acid treated with silver nitrate and vapor of ammo- nia, gives no color until warmed with a drop of ammonium hydrox- ide, then gives a black precipitate. With the white fleck no further action on addition of ammonium molybdate. With vapor of iodine, color more or less carmine-red, by formation of antimonous iodide, not readily volatile by heat. * The hypochlorite reagent, usually NaCIO, decomposes in the air and light on standing. It should instantly and perfectly bleach litmus paper (not redden it). It dissolves arsenic by oxidation to arsenic acid. As, + lOJVaCIO + 6H,0 = 4H a AsO, + lONaCl. §69, Q'd. arsenic. 67 , 2. To the spot obtained on the porcelain surface, add a few drops of nitric acid and heat; then add a drop of ammonium molybdate. A yellow precipitate indicates arsenic. Antimony may give a white precipitate with the nitric acid, but gives no further change with the ammonium molybdate (Deniges, C. r., 1890, 111, 824). 3. Oxidize the arsenic spot with nitric acid and evaporate to dryness. Add a drop of silver nitrate or ammonio-silver nitrate (6/). A reddish- brown precipitate indicates arsenic. 4. After the formation of the mirror in Marsh's test the generating flask may be disconnected and a stream of dry hydrogen sulphide passed over the heated mirror. If the mirror consists of both arsenic and anti- mony, the sulphides of both these metals will be formed, and as the arsenous sulphide is volatile when heated, it will be deposited in the cooler portion of the tube. The sulphides being thus separated can readily be distinguished by the color. If now a current of dry hydrochloric acid gas be substituted for the hydrogen sulphide the antimonous sulphide will be decomposed to the white antimonous chloride which volatilizes and may be driven past the unchanged arsenous sulphide (5c). 5. The tube containing the mirror is cut so as to leave about two inches on each side of the mirror and left open at both ends. Incline the tube and beginning at the lower edge of the mirror gently heat, driving the mirror along the tube. The mirror will disappear and if much arsenic be present a white powder will be seen forming a ring just above the heated portion of the tube. This powder consists of crystals of arsenous oxide, and should be carefully examined under the microscope and iden- tified by their crystalline form (Wormley, 270). 6. The crystals of arsenous oxide obtained above are dissolved in water and treated with ammonio-silver nitrate forming the yellow silver arse- nite (6/) : or with ammonio-copper sulphate forming the green copper x arsenite (6fc) (Wormley, 259). Any other test for arsenous oxide may be applied as desired. 7. Magnesia mixture (6t) is added to the solution of the mirror or spots in nitric acid. A white crystalline precipitate of magnesium ammonium arsenate, MgNH 4 As0 4 , is formed (Wormley, 316). d. — Reinsch's Test. — If a solution of arsenic be boiled with hydrochloric acid and a strip of bright copper foil, the arsenic is deposited on the copper as a gray film. Hager (O. C, 1886, 680) recommends the use of brass foil instead of copper foil. When a large amount of arsenic is present the coating of arsenic separates from the copper in scales. The film does not consist of pure metallic arsenic, but appears to be an alloy of arsenic and copper. Arsenous compounds are reduced much more readily than arsenic compounds. The hydrochloric acid should compose at least one-tenth the volume of the solution. The arsenic is not deposited if the acid is not present. This serves as one of the most satisfactory methods of determining the presence or absence of arsenic in €8 ARSENIC. §69, 6'e. hydrochloric acid. Dilute the concentrated acid with five parts of water and boil with a thin strip of bright copper foil. A trace of arsenic if present will soon appear on the foil. For further identification of the deposit, wash the foil with distilled water, dry, and heat in a hard glass tube, as for the oxida- tion of the arsenic mirror (6'c, 5). The crystals may be identified by the mic- roscope and by any other tests for arsenous oxide. It is important that the surface of the copper should be bright. This is obtained by rubbing the sur- face of the foil with a file, a piece of pumice or sand-paper just before using. The copper should not contain arsenic, but if it does contain a small amount no film will be deposited due to its presence unless agents are present which cause partial solution of the foil. If a strip of the foil, upon boiling with hydrochloric acid for ten minutes, shows no dimming of the brightness of the copper surf ape; the purity of both acid and copper may be relied upon for the most exact work. Antimony, mercury, silver, bismuth, platinum, palladium and gold are deposited upon copper when boiled with hydrochloric acid. Under certain conditions most of these deposits may closely resemble that of arsenic. Of these metals mercury is the only one that forms a sublimate when heated in the reduction tube (7), and this is readily distinguished from arsenic by examination under the microscope. Antimony may be volatilized as an amor- phous powder at a very high heat. Organic material may sometimes give a deposit on the copper which also yields a sublimate, but this is amorphous and does not show the octahedral crystals when examined under the microscope (Wormley, 269 and ff.; Clark, ,/. C, 1893, 63, 886). e. — Detection in Case of Poisoning. — Arsenic in its various compounds is largely used as a poison for bugs, rodents, etc., and frequently cases arise of accidental arsenical poisoning. It is also used for intentional poisoning, chiefly suicidal. It is usually taken in the form of arsenous oxide (white arsenic), or " Fowler's Solution " (a solution of the oxide in alkali carbonate). One hun- dred fifty to two h,undred milligrams (two to three grains) are usually sufficient to produce death. Violent vomiting is a usual symptom and death occurs in from three to six hours. In cases of suspected poisoning vomiting should be induced as soon as possible by using an emetic followed by demulcent drinks, or the stomach should be emptied by a stomach pump. Freshly prepared ferric hydroxide is the usual antidote, of which twenty-five to fifty grams (one to two ounces) may be given. The antidote may be prepared by adding magnesia (magnesium oxide), ammonium hydroxide, or cooking soda (sodium bicarbo- nate) to ferric chloride or muriate tincture of iron: straining in a clean piece of muslin, and washing several times. If magnesia be used it is not necessary to wash, as the magnesium chloride formed is helpful rather than injurious. A portion of the ferric hydroxide oxidizes some of the arsenous compound, being itself reduced to the ferrous condition, and forming an insoluble ferrous arsenate. When the ferric oxide is in excess the ferrous arsenate does not appear to be acted upon by the acids of the stomach. Of course it will be seen that the ferric hydroxide will have no effect upon the arsenic which has entered into the circulation. It frequently becomes necessary for the chemist to analyze portions of sus- pected food, contents of the stomach, urine; or, if death has ensued, portions of the stomach, intestines, liver, or other parts of the body. At first a careful examination should be made of the material at hand for solid white particles, that would indicate arsenous oxide. If particles be found they can at once be identified by the usual tests. Liquid food or liquid contents of the stomach should be boiled with dilute hydrochloric acid, filtered and washed and the filtrate precipitated with hydrogen sulphide, etc. When solid food or portions of tissue are to be analyzed, it is necessary first to destroy the organic material. ■Several methods have been proposed: (1) Method of Presenilis and Babo. — The tissue is cut in small pieces and about an equal weight of pure hydrochloric acid added to this, enough water should be added to form a thin paste and dilute the hydrochloric acid five or six times. The mass is heated on the water bath and crystals of potassium chlorate added in small amounts at a time with stirring until a clear yellow liquid is obtained containing a very small amount of solid particles. The heating is continued until there is no odor of chlorine, but concentration should ^69, 7. ARSENIC. 69 ~be avoided by the addition of water. The solution should be cooled and filtered; "the arsenic now being present in the filtrate as arsenic acid. This solution should be treated with sodium bisulphite or sulphur dioxide to reduce the arsenic acid to arsenous acid and then the arsenic may be precipitated with hydrogen sulphide. It is advisable to pass the hydrogen sulphide through the -warm liquid for twenty-four hours to insure complete precipitation. A yel- lowish precipitate of organic matter will usually be obtained even if arsenic be absent. The precipitate should be filtered, washed, and then dissolved in dilute ammonium hydroxide, which separates it from other sulphides of the silver, tin and copper groups, that may be present. A portion at least of the precipitated organic matter will dissolve in the ammonium hydroxide. The filtrate should be acidulated with hydrochloric acid, filtered and washed. Dissolve the precipitate in concentrated nitric acid and evaporate to dryness. Redissolve in a small amount of water, add a drop of nitric acid, filter and test the filtrate by Marsh's test or any of the other tests for arsenic. (2) Hydrochloric acid diluted alone may be used for the disintegration of the soft animal tissues. The solution will usually be dark colored and viscous and not at all suited for further treatment with hydrogen sulphide; but may be at once subjected to the Reinsch test (6'd). (3) Method of Danger and Flandin. — The tissue may be destroyed by heat- ing in a porcelain dish with about one-fourth its weight of concentrated sul- phuric acid. When the mass becomes dry and carbonaceous it is cooled, treated with concentrated nitric acid and evaporated to dryness. Moisten with waiter, add nitric acid, and again evaporate to dryness; and repeat until the mass is colorless. Dissolve in a small amount of 'water and test for arsenic by the usual tests. This method is objectionable if chlorides are present as the volatile arsenous chloride will be formed. (4) Method by distillation with hydrochloric acid. The finely divided tissue is treated, in a retort, with its own weight of concentrated hydrochloric acid and distilled on the sand bath. Salt and sulphuric acid may be used instead of hydrochloric acid. A receiver containing a small amount of water is connected to the retort and the mass distilled nearly to dryness. If preferred, gaseous hydrochloric acid may be conducted into the retort during the process of dis- tillation, in which case all the arsenic (even from arsenous sulphide (So)) will be carried over in the first 100 cc. of the distillate. The receiver contains the arsenic, a great excess of hydrochloric acid and a small amount of organic matter. To a portion of this solution the Eeinsch test may be applied at once and other portions may be diluted and tested with hydrogen sulphide or the solution may at once be tested in the Marsh apparatus. For more detailed instructions concerning the detection and estimation of arsenic in organic matter, special works on Toxicology and Legal Medicine must be consulted. The following are valuable works on this subject: Micro- -Chemistry of Poisons, Wormley; Medical Jurisprudence, Taylor; A System of Legal Medicine, Hamilton; Ermittelung von Giften, DragendorfE; Poisons, Taylor; etc. 7. Ignition. — Metallic arsenic is obtained by igniting any compound containing arsenic with potassium carbonate and charcoal,* or 'with potas- sium cyanide: 2As 2 3 + 6KCN = As 4 + 6KCNO 2As 2 S 3 + 6KCN = As 4 + 6KCNS 2As 2 S„ + 6TiTa 2 C0 3 + 6KCN = As 4 + 6Na 2 S + 6KCNO + 6C0 2 . 4H 3 As0 4 + 5C = As 4 + 5C0 2 + 6H 2 * A very suitable carbon for the reduction of arsenic is obtained by igniting an alkali tartrate in absence of air to complete carbonization. 70 ARSENIC. §69, 8- If this ignition be performed in a small reduction-tube * (a hard glass tub& about 7 mm. in diameter, drawn out and sealed at one end), the reduced, arsenic sublimes and condenses as a mirror in the cool part of the tube. The test may be performed in the presence of mercury compounds, but. more conveniently after their removal; in presence of organic material, it is altogether unreliable. If much free sulphur be present the arsenic should be removed by oxidation to arsenic acid by nitric acid or hydro- chloric acid and potassium chlorate, then precipitation after addition of ammonium hydroxide by magnesium mixture and thoroughly drying before- mixing with the cyanide or other reducing agent. 8. Detection. — Arsenic is precipitated, from the solution acidulated with hydrochloric acid, in the second group by hydrosulphuric acid as the- sulphide (6e). By its solution in (yellow) ammonium sulphide it is sepa- rated from Hg , Pb , Bi , Cu , and Cd . By reduction to arsine in the Marsh apparatus it is separated with antimony from the remaining second group metals. The decomposition of the arsine and stibine with silver- nitrate precipitates the antimony, thus effecting a separation from the- arsenic, which passes into solution as arsenous acid. The excess of AgN0 r is removed by HC1 or CaCl 2 and the presence of arsenic confirmed by its- precipitation with H 2 S . For other methods of detection consult the text: (6, 6' and 7). For distinction between As v and As'" see (& and §88, 4). 9. Estimation. — (1). As lead arsenate, Pb 3 (As0 4 ) 2 . To a weighed por- tion of the solution containing arsenic acid, a weighed amount of PbO is- added, after evaporation and ignition at a dull red heat is weighed as- Pb 3 (As0 4 ) 2 . The weight of the added PbO is subtracted from the residue,, and the difference shows the amount of arsenic present reckoned as As 2 5 . (2). It is precipitated by MgS0 4 in presence of NH 4 0H and NH 4 C1 , and after drying at 103°, weighed as MgNH 4 As0 4 .H 2 ; antimony is not precipitated if a tartrate be present (Lesser, Z., 1888, 27, 218). (S). The MgNH 4 As0 4 is converted by ignition into Mg,As 2 7 , and weighed. (4). The solution of arsenous acid containing HC1 is precipitated by H 2 S .. * As much of the reduction-glass tubing contains arsenic (?) Fresenius (Z., 20, 531 and 22, 391> recommends the folio-wing modification of the above method : A piece of reduction tubing about. 10 mm. diameter and 15 cm. long is drawn out to a narrow tube at one end. The other end of the- tube is connected with a suitable apparatus for generating and drying carbon dioxide. The sample to be tested is thoroughly dried and mixed with the dry cyanide (or charcoal) and car- bonate, placed in a small porcelain combustion boat and put in the middle of the reduction tube. The air is then driven from the tube by the dry carbon dioxide and the whole heated gently until all moisture is expelled. The tube is then heated to redness near the point of con- striction and when this is done the boat is heated, gently at first to avoid spattering of the fus- ing mass, then to a full redness till all the arsenic has been driven out. During the whole of the experiment a gentle stream of carbon dioxide is passed through the tube. The arsenic collects as a mirror in the narrow part of the tube just beyond the heated portion. The small end of the tube may now be sealed, the mirror collected by a gentle flame, driven to any desired portion of, the tube and tested with the usual tests (6' c5). Compounds of antimony when treated in this, way do not give a mirror. As small an amount as 0.00001 gram of As 2 3 will give a distinct mir- ror by this method. §69, 10. ARSENIC. 71 The precipitate is separated from free sulphur by solution in NH 4 0H and reprecipitated with HC1 . It is then dried and weighed as As 2 S 3 . (5). By precipitation as in (4) and removal of sulphur by washing the precipitate with CS 2 . Dry and weigh as As 2 S s . (6). ITranyl acetate, in presence of ammonium salts, precipitates NH 4 TT0 2 As0 4 ; by ignition this is converted into uranyl pyroarsenate (U0 2 ) 2 As 2 7 , and weighed as such. (7). Small amounts may be converted into the metallic arsenic mirror by the Marsh apparatus and weighed or compared with standard mirrors (Gooch and Moseley, C. N., 1894, 70, 207). (8). As'" is converted into As v by a graduated solution of iodine in presence of NaHC0 3 . The end of the reaction is shown by the blue color imparted to starch. (9). As'" is oxi- dized to As v by a graduated solution of K 2 Cr 2 7 , and the excess of K 2 Cr 2 7 determined by a graduated solution of FeS0 4 . (10). As'" is con- verted to As v by a weighed quantity of K 2 Cr 2 7 with HC1 , and the excess of chlorine is determined by KI and Na 2 S 2 3 . (11). As'" is oxidized to As v by a graduated solution of KMn0 4 . The end of the reaction is indi- cated by the color of the KMn0 4 . (12). As v is reduced to As'" by a grad- uated solution of HI . The action takes place in acid solutions. (IS). In neutral solution, as arsenate, add an excess of standard AgN0 3 , and in an aliquot part estimate the excess of AgN0 3 with standard NaCl . (H). Dis- tillation as AsCl 3 (Piloty and Stock, B., 1897, 30, 1649; see also 6'e 4). (15). The arsenic compound is converted into AsH 3 and this passed into a solution of standard silver nitrate, the excess of which is estimated with standard NaCl or the excess of AgNO, is removed and the arsenous acid titrated as in methods (9) or (11). Many other methods have been recommended. 10. Oxidation.— As-'"H 3 is oxidized to As'" by AgN0 3 , H 2 S0 3 , H 2 S0 4 , and HI0 3 ; and to As v by KMn0 4 (Tivoli, Gazzetta, 1889, 19, 630), HN0 2 , HN0 3 , CI and Br (Parsons, C. N., 1877, 35, 235). As° is oxidized to As'" by H 2 2 (Clark, J. C, 1893, 63, 886), HN0 3 , H 2 S0 4 hot, CI , HC10 , HC10 3 , Br , HBr0 3 , HI0 3 , Ag' (Senderens, C. r., 1887, 104, 175), and to As v by the same reagents in excess except H 2 S0 4 and Ag', which oxidize to As'" only. As'" is also oxidized to As v in presence of acid by Pb0 2 , Cr VI ; by compounds of Co, Ni, and Mn, with more than two bonds; and in alkaline mixture by Pb0 2 , Hg 2 , HgO , CuO , K 2 Cr0 4 , K 3 Fe(CN) c , etc. (Mayer, J. pr., 1880 (2), 22, 103). Arsine is oxidized to metallic arsenic by HgCl 2 (Magencon and Bergeret, J. C, 1874, 27, 1008), and by As'", the As'" also becoming As° (Tivoli, C. C, 1887, 1097). As v and As'" are reduced to metallic arsenic by fusion with CO , with free carbon, or with carbon com- bined, as H 2 C 2 4 , KCN , etc. (7). By SnCl 2 (Gg) and H 3 P0 2 (6d) in strong HC1 solution; also with greater or less completeness by some free metals, such as Cu , Cd , Zn , Mg , etc. Eideal (C. N., 1885, 51, 292) recommends 72 ANTIMONY. §70, 1. the use of the copper-iron wire couple for the detection of small quantities of arsenic by reduction to the elemental state. 0.0000075 grams may be detected. In solution As v is reduced to As'" by H 3 P0 2 , H 2 S , H 2 S0, , Na 2 S 2 3 (6e), HC1 , HBr , HI (6/), HCNS , etc. As v and As'" are reduced to As~"'H 3 by nascent hydrogen generated by the action of Zn and dilute H 2 S0 4 , or, in general, by any metal and acid which will give a ready generation of hydrogen, as Zn , Sn , Fe , Mg , etc., and H 2 S0 4 and HC1 (Draper, Dingl, 1872, 204, 320). As'" is reduced to As-"'H 3 by nascent hydrogen generated in alkaline solution as, Al and KOH, Zn and KOH, sodium amalgam, etc. (separation from antimony) (Davy, Ph. C, 1876, 17, 275; Johnson, C. N., 1878, 38, 301). §70. Antimony (Stibium) Sb = 120.4. Valence three and five (§11). 1. Properties. — Specific gravity, 6.697 (Schroeder, /., 1859, 12). Melting point, 432° (Ledebur, Wied. Beibl., 1881, 650). Boiling point, between 1090° and 1450° (Carnelley and Williams, J. C, 1879, 35, 566). Its molecular weight. is unknown, as its vapor densitjr has not been taken. Antimony is a lustrous, silver white, brittle and readily pulverizable metal. It is but little tarnished in dry air and oxidizes slowly in moist air, forming a blackish gray mixture of antimony and antimonous oxide. At a red heat it burns in the air or in oxygen with incan- descence, forming white inodorous (distinction from arsenic) vapors of anti- monous oxide. 2. Occurrence. — Native in considerable quantities in northern Queensland, Australia (Mac Ivor, C. N., 1888, 57, 64); as stibnite, Sb 2 S 3 ; as valentinite, Sb 2 O s ; in very many minerals usually combined with other metals as a double sulphide (Campbell, Phil. Mag., 1860, (4), 20, 304; 21, 318). 3. Preparation. — (a) The sulphide is converted into the oxide by roasting in the air, and then reduced by fusion with coal or charcoal. (6) The sulphide is fused with charcoal and sodium carbonate: 2Sb,S 3 + 6Na 2 C0 3 + 3C = 4Sb + 6Na 2 S + 9C0 2 . (o) It is reduced by metallic iron: Sb 2 S„ + 3Pe = 2Sb + 3PeS . (d) To separate it from other metals with which it is frequently combined requires a special process according to the nature of the ore (Dexter, J. pr., 1839, 18, 449; Pfeifer, A., 1881, 209, 161). 4. Oxides. — Antimony forms three oxides, Sb 2 O s , Sb 2 4 , and Sb 2 O s . (o) Antimonous oxide, Sb 2 O s , is formed (1) by the action of dilute nitric acid upon Sb°; (2) by precipitating SbCL, with Na 2 CO s or NH,OH; (3) by dissolving Sb° in concentrated H 2 S0 4 and precipitating with lTa 2 CO s ; (4) by burning antimony at a red heat in air or oxygen; (5) by heating Sb 2 4 or Sb,0 5 to 800° (Baubigny, G. r., 1897, 124, 499, and 560). It is a white powder, turning yellow upon heat- ing and white again upon cooling; melts at a full red heat, becoming crystalline upon cooling; slightly soluble in water, fairly soluble in glycerine (56). Anti- monous oxide sometimes acts as an acid, Sb 2 3 + 2NaOH = 2NaSbO„ + H 2 0; but more commonly as a base. Ortho and pyro antimonous acids are known in the free state. The meta compound exists only in its salts (D., 2, 1, 198). (6) Diantimony tetroxide, Sb 2 4 , is formed by heating Sb° , Sb 2 S 3 , Sb 2 O s , or Sb 2 O in the air at a dull red heat for a long time. The antimony in this compound is probably not a tetrad, but a chemical union of the triad and pentad: 2Sb 2 4 = 2Sb"'Sbv0 4 = Sb 2 8 .Sb 2 5 . It is found native as antimony ochre, (c) Antimonic oxide, Sb 2 5 , is formed by treating Sb° , Sb 2 O s or Sb 2 4 with concentrated nitric acid. When heated to 300° it loses oxygen, forming Sb 2 4 (Geuther, J. pr., 1871, (2), 4, 438). It is a citron-yellow powder, insoluble in water but reddening moist blue litmus paper. Antimonic acid ■exists in the three * forms, analogous to the arsenic and phosphoric acids, *BeilsteinandBlaeae(C. C, 1889, 8031 have prepared a number of antimonates and conclude that the acid is always the meta, H SbO, . §70, 56. ANTIMONT. 73 i. e., ortho, meta and pyro (Geuther, I. c, and Conrad, C. N., 1879, 40, 198). The ortho acid, H„Sb0 4 is formed by the decomposition of the pentaehloride with water and washing- until the chloride is all removed (Conrad, I. c, and Dau- brawa, A., 1877, 186, 110). The most of the antimonates formed in the wet way by precipitation from the acid solution of antimonic chloride are the ortho antimonates. By heating the ortho acid to 200° the meta acid, HSb0 3 , is formed. Strong ignition of Sb 2 O s with potassium nitrate and extraction with water gives the potassium metantimonate, KSb0 3 , and by adding nitric acid to a solution of this salt the free acid is formed. The ortho acid dried at 100* gives the pyro acid: 2H 3 Sb0 4 = H 4 Sb 2 7 + H 2 (Conrad, I. a), which upon further heating to 200° gives the meta acid. The pyroantimonic acid forms two series of salts, M 4 Sb 2 7 and M 2 H 2 Sb 2 0, . The sodium salt Na 2 H 2 Sb 2 T is insoluble in water and is formed in the quantitative estimation of antimony (9), and also in a method for the detection of sodium (§206, 6g). For the latter the soluble potassium salt K 2 H 2 Sb 2 7 is used as the reagent. It is prepared by fusing antimonic acid with a large excess of potassium hydroxide; then dissolving, filtering, evaporating and digesting hot, in syrupy solution, with a, large excess of potassium hydroxide, best in a silver dish, decanting the alkaline liquor, and stirring the residue to granulate, dry. This reagent must be kept dry, and dissolved when required for use: inasmuch as, in solution, it changes to the tetrapotassium pyroantimonate, K 4 Sb 2 7 , which does not precipitate sodium. The reagent is, of course, not applicable in acid solutions. The reaction is as follows: K 2 H 2 Sb 2 7 + 2NaCl = Na 2 H 2 Sb 2 7 + 2KC1 (§11). The ortho acid, H 3 Sb0 4 , is sparingly soluble in water, easily soluble in KOH, but insoluble in NaOH. The meta acid, HSbO„ , is sparingly soluble in water, easily soluble in both the fixed alkalis; the pyro acid, H 4 Sb 2 7 , is sparingly (more easily than the meta) soluble in water; the normal fixed alkali salts, K 4 Sb 2 7 , are soluble in water, also the acid potassium salt, K 2 H 2 Sb 2 0, , but not the corresponding sodium salt, Na 2 H 2 Sb 2 7 . 5. Solubilities. — a. — Metal. — Antimony is attacked but not dissolved by nitric acid, forming Sb 2 O s (a) or Sb 2 0„ (1>), depending upon the amount and degree of concentration of the acid; it is slowly dissolved by hot concentrated sulphuric acid, evolving S0 2 and forming Sb 2 (S0 4 ) 3 (c); it is insoluble in HC1 out of con- tact with the air, but the presence of moist air causes the oxidation of a small amount of the metal to Sb,0 3 , which is dissolved in the acid without evolution of hydrogen (Ditte and Metzner, A. Ch., 1896, (6), 29, 389). The best solvent for antimony is nitric acid, followed by hydrochloric acid or nitrohydrochloric acid containing only a small amount of nitric acid. Anti- monous chloride, SbCl 3 , is at first formed (d), but if sufficient nitric acid be present this is rapidly changed to antimonic chloride, SbCl 5 (e). If, however, too much nitric acid be present, the corresponding oxides (not readily soluble in nitric acid) are precipitated (6c). The halogens readily attack the metal foMning at first the corresponding trihalogen compounds (d). Chlorine and bramine (gas) unite with the production of light, and if the halogen be in excess, the pentad chloride (e) or bromide is formed (Berthelot and Petit, A. Ch., 189i, (6), 18, 65). The pentiodide, Sbl 5 , does not appear to exist (Mac Ivor, J. C, 1876, 29, 328). (o) 2Sb + 2HNO„ = Sb 2 O a + 2NO + H 2 (6) 6Sb + 10HNO 3 = 3Sb 2 5 + 10NO + 5H 2 (C) 2Sb + 6H 2 S0 4 = Sb 2 (S0 4 ) 3 + 3S0 2 + 6H 2 ( CO o «» -—.--; ca P T3«» -

— O »w o » ;.vr «2 6Dco csB "- 1 43 ^3 00 cos . - cd COO — > h > *-* o^C-3 a i — goafe-'d 3 §p Ph wo "1 a 01 CO COO ca .So 13 g2-^ W .So|2i' «1 CO rA a <1 I <1 -a 6o ffi - Cm 1^ td °Pa < in < a w < O fc Stsw ticoa qJ P. Scot .-a j3AO« ,h cd -P p ^ p. gflOr,,. J« P» •H PhC ca .-h S a Pj a -3 w o W 8 a o a B cp s ft H W §73, 5a. GOLD. 91 §73. Gold (Aurum) Au = 197.2. Valence one and three. 1. Properties. — Specific gravity, 19.30 to 19.34 (Kose, Pogg., 1848, 75, 403). Melt- ing point, 1061.7° (Heycock and Neville, J. C, 1895, 67, 189). It is a yellow metal, -that from different parts of the world varying slightly in color; the presence of very small traces of other metals also affects the color. It is softer than silver and harder than tin; possesses but little elasticity or metallic ring. It is the most malleable and ductile of all metals; one gram can be drawn into a wire 2000 metres long. The presence of other metals diminishes the ductility. It may be rolled into sheets 0.0001 mm. thick. At a very high heat it vaporizes (Deville and Debray, A. Ch., 1859, (3), 56, 429). It is a good conductor of electricity, equal to copper, not so good as silver. It has a high coefficient of expansion and cannot be moulded into forms but must be stamped. On account of its softness, gold is seldom used absolutely pure, but is hardened by being alloyed with other metals, as Ag , Cu , etc. 2. Occurrence. — Gold is usually found native, but never perfectly pure, being always alloyed with silver, and occasionally also with other metals. It is found as gold-dust in alluvial sand, sometimes in nuggets, and sometimes disseminated in veins of quartz. 3. Preparation. — (1) Washing. Which consists in treating the well-powdered ■ore with a stream of water, the heavy gold settling to the bottom. (2) Amalga- mation. Which consists in dissolving the gold in mercury and then separating it from the latter by distillation. (3) By fusing with metallic lead, which dis- solves the gold, the liquid alloy settling to the bottom of the slag. The gold is afterward separated from the lead by cupellation. The silver is separated from the gold by dissolving it in nitric or sulphuric acid. Or the whole is dissolved in nitrohydrochloric acid, and the gold precipitated in the metallic state by some reducing agent; ferrous sulphate being usually employed. Another method is to pass chlorine into the melted alloy. The silver chloride rises to the surface, while the chlorides of Zn , Bi , Sb , and As (if present) are vola- tilized, and the pure gold remains beneath. A layer of fused borax upon the surface prevents the silver chloride from volatilizing. (4) By treatment with a solution of KCN . (5) By amalgamation with mercury and electrolysis at the same time. 4. Oxides and Hydroxides. — Aurous oxide, Au,0 , is very unstable, heating to about 250° decomposes it into the metal and oxygen. The hydroxide is pre- pared by reducing the double bromide with S0 2 in ice-cold solution; heating to 200° changes it to the oxide (Kriiss, A., 1886, 237, 274). Aurie hydroxide, Au(0H) a , is prepared by precipitation from the chloride solution with MgO (Kriiss, I. a). It is a yellow to brown powder, changing to the oxide upon dry- ing at 100°. Heating to 250° gives the mental and oxygen (§10). 5. Solubilities. — a. — Metal. — Gold is not at all tarmsnea or m any way acted upon by water at any temperature, or by hydrosulphuric acid. Neither nitric nor hydrochloric acid attacks it under any conditions; but it is rapidly attacked by chlorine (as gas or in water solution), dissolving promptly in nitrohydro- chloric acid, as auric chloride, AuCl 3 ; by bromine, dissolving in bromine water, as auric bromide, AuBr 3 ; and by iodine; dissolving when finely divided in hydri- odic acid by aid of the air and potassium iodide, as potassium auric iodide, KIAuI s : 4Au + 12HI + 4KI + 30 2 = 4KIAuI 3 + 6H 2 . Potassium cyanide solution, with aid of the air, dissolves precipitated gold as potassium auro- ■cyanide, KAu(CBr) 2 : 4Au + 8KCN + 2 + 2H 2 = 4KAu(CN) 2 + 4KOH . Gold is separated, from its alloys with silver and base metals, by solution in nitric acid; the gold being left as a black-brown powder — together with platinum and oxides of antimony and tin. When the gold-silver or gold-copper has not over 20 per cent gold, nitric acid of 20 per cent disintegrates the alloy, and effects the separation; when the gold is over 25 per cent, silver or lead (three parts) must be added, by fusion, to the alloy before solution. (If gold- ;silver alloy contains 60 per cent or more of silver, it is silver color; if 30 per •cent silver, a light brass color; if 2 per cent silver, it is brass color.) If gold and other metals are obtained in solution by nitrohydrochloric acid, leaving most of the silver as a residue, the noble metals can be precipitated by :zinc or ferrous sulphate, and the precipitate of gold, silver, etc., treated with 92 GOLD. §73, 5L nitric acid, which will now dissolve out any proportion of silver not less than, 15 per cent, to 85 per cent of gold, and dissolve the baser metals. Concentrated, sulphuric acid dissolves silver, and leaves gold. 6. — The oxides and hydroxides of gold are insoluble in water, soluble in acids. c. — The salts of the oxyacids are not stable, being decomposed by hot water.. Gold sulphide is insoluble in water or acids, except nitrohydrochloric acid, soluble in alkali sulphides. Aurous salts are decomposed by water, forming Au° and Au'" . Auric chloride is deliquescent; both the chloride and bromide- are readily soluble in water. The iodide is decomposed by water, forming aurous iodide. The double chlorides, bromides, iodides and cyanides are soluble in water. 6. Reactions, a. The fixed alkali hydroxides and carbonates in excess do not precipitate AuCl 3 solutions, as a soluble aurate, KAu0 2 , readily forms; but upon boiling and neutralizing the excess of alkali, Au(0H) :j is precipitated. Ammonium hydroxide precipitates from concentrated solutions a reddish-yellow ammonium aurate, (NH 3 ) 2 Au 2 3 , " fulminating- gold." b. Oxalic acid reduces gold chloride from solutions, slowly (nitric: acid should be absent and the presence of ammonium oxalate is advan- tageous), but completely. The gold separates in metallic flakes or forms a mirror on the side of the test-tube. 2AuCl 3 + 3H 2 C 2 4 = 2Au + 6C0 2 + 6HC1 . As platinum, palladium, and other second group metals are not reduced by oxalic acid, this method of removal of gold should be? employed upon the original solution before the precipitation of the second, group metals as sulphides. Potassium gold cyanide, KCN.Au(CN) 3 , is. formed when a neutral solution of AuCl 3 is added to a hot saturated- solution of KCN . It is very soluble in water and by heating above 200" it is decomposed into CN and KCN.AuCN , which latter product is formed, when gold is dissolved in KCN in the presence of air (5a). c. A solution of AuCl 3 is precipitated as Au° by a solution of KN0 2 . d. Sodium. pyrophosphate forms with AuCl 3 a double salt which has found application in gold plating, e. Hydrosulphuric acid precipitates from gold chloride solution, hot or cold, gold sulphide, variable from An 2 S to Au 2 S 2 , brown,, insoluble in acids, hot or cold, except in nitrohydrochloric acid, in which it readily dissolves ; soluble in alkali sulphides to a thio-salt. Alkali- sulphites precipitate gold chloride solution as double sulphite, i. e. Au 2 (S0 3 ) 3 .(NH 4 ) 2 S0 3 .6NH 3 + 3H 2 . Upon boiling the sulphite acts as a reducing agent, giving metallic gold. f. Potassium iodide, added in small portions to solution of auric chloride (so that the latter is constantly in excess where the two salts are in contact), and when equivalent proportions have been reached, gives a yel- low precipitate of aurous iodide, AnI , insoluble in water, soluble in large excess of the reagent; the precipitate accompanied with separation of hee iodine, brown, which is quickly soluble in small excess of the reagent as a. colored solution: AuCl 3 + 4KI = Aul + 3KC1 + I 2 with KI . But, on gradually adding auric chloride to solution of potassium iodide, so that the* §74, 1. PLATINUM. 93 latter is in excess at the point of chemical change, there is first a dark- green solution of potassio-auric iodide, KIAuL, ; then a dark-green precipi- tate of auric iodide, Aul 3 , very unstable, decomposed in pure water, more quickly by boiling; changed in the air to the yellow aurous iodide. g. Stannous chloride gives a purple precipitate containing the oxides of tin with the gold, " purple of Cassius " insoluble in acids. h. Ferrous sulphate is the most common reagent for the detection of gold, reducing all gold salts to the metallic -state; AuCl 3 -)- 3FeS0 4 = An + Fe 2 (S0 4 ) 3 + FeCl 3 . 7. Ignition. — Gold is reduced from many of its compounds by light, and from all of them by heat — its separation in the dry way being readily effected by fusion with such reagents as will make the ma.terial fusible. Very small pro- portions are collected in alloy with lead, by fusion; after which the lead is. vaporized in " cupellation " (§59, 7). 8. Detection. — In the dry way gold is detected by fusion of the mineral, matter with lead, to the formation of a " button " which is then ignited to drive off the lead, leaving the gold and silver behind as the metals.. In the wet way the material, if not in solution, is digested with nitro- hydrochloric acid which dissolves all the gold. The excess of acid is re- moved by evaporation and the gold is precipitated by oxalic acid or ferrous sulphate, and identified by its color and insolubility in acids. If the- gold be not removed from the original solution it is precipitated in Group II. by H 2 S , passes into Division A (tin group) by (NH 4 ) 2 S , and may be detected in the flask of the Marsh apparatus by the usual methods. 9. Estimation. — Gold is always weighed in the metallic state, to which form it is reduced: (1) By ignition alone if it is a salt containing no fixed acid; if in an ore, by mixing with lead and fusion to an alloy, and final removal of the lead by ignition at a white heat in presence of air. (2) By adding to the solu- tion some reducing agent, usually FeS0 4 , H 2 C 2 4 , chloral hydrate, or some- easily oxidized metal, such as Zn , Cd , or Mg . (3) Gold is also estimated volu- metrically by H 2 C 2 4 and the excess of H 2 C 2 4 used, determined by KMnO, . 10. Oxidation. — Gold is reduced to the metallic state by very many reducing agents, among which may be mentioned the following : Pb , Ag ,. Hg, Hg', Sn, Sn", As, As'", AsH 3 , Sb, Sb'", SbH 3 , Bi, Cu, CuV Pd, Pt, Te, Fe, Fe", Al, Co, Ni, Cr"', Zn, Mg, H 2 C 2 4 , HN0 2 , P, H 3 P0 2 , H 3 P0 3 , PH 3 , H 2 S0 3 , and a great number of organic substances. §74. Platinum. Pt = 194.9 . Valence two and four. 1. Properties. — Specific gravity at 17.6°, 21.48 (Deville and Debray, C. r., 1860, 50, 1038). Melting point, 1775° (Violle, C. r., 1879, 89, 702). Pure platinum is a tin-white metal, softer than silver, hardened by the presence of other metals, especially iridium, which it frequently contains. It is surpassed in ductility and malleability only by Au aDd Ag . Platinum black is the finely divided metal, a black powder, obtained by reducing an alkaline solution of the platinous. salt with alcohol (Low, B., 1S90, S3, 389) ; platinum sponge, a gray spongy mass,. 94 PLATINUM. §74, 2 by ignition of the platinum ammonium double chloride; platinized asbestos (usually 10 per cent Pt), the metal in finely divided form deposited by reduction, from the salt upon asbestos. These finely divided forms of platinum have great power of condensation of gases, and by their presence alone bring about a num- ber of important chemical reactions (catalytic reaction); e.g., a current of hydrogen mixed with air ignites when passed over platinum black, also hydrogen and chlorine unite. SO z unites with O to form S0 3 ; alcohol is oxi- dized to acetic acid, formic and oxalic acids to C0 2 , As'" to Asv , etc. 2. Occurrence. — Found in nature only in the metallic state, generally alloyed with palladium, iridium, osmium, rhodium, ruthenium, etc. The Ural Moun- tains furnish the largest supply of platinum. 3. Preparation. — Usually by the wet method. The finely divided ore is treated with nitrohydroehloric acid until the platinum is all dissolved. The filtrate is then treated with lime water to a slightly acid reaction; this removes the greater part of the Fe , Cu , Ir , Eh , and a portion of the Pd . The filtrate is now evaporated to dryness, ignited and washed with water and hydrochloric acid. This gives a commercial platinum which is melted with six times its weight of lead and the finely divided alloy digested with dilute HNO, , which dissolves out the Pb , Cu , Pd , and Eh . The black powder which remains is dissolved in nitrohydroehloric acid, the Pb remaining, removed with H 2 S0 4 , and the Pt precipitated with MH 4 C1 . The precipitate contains a little rhodium, which is removed by gently igniting' the mass with potassium and ammonium di-sulphate, and exhausting with water, which dissolves out the rhodium sulphate (§105, 7). In the laboratory the platinum residues are boiled with KOH or K 2 CO a and reduced with alcohol. The fine black powder is filtered, washed with water and hydrochloric acid and ignited. 4. Oxides and Hydroxide's. — Platinum forms two oxides, PtO and PtO, . Platinous hydroxide is formed by treating a dilute solution of platinous potas- sium chloride with NaOH and boiling (Jorgensen, /. pr., 1877, (2), 16, 344). A black powder easily soluble in HC1 or HBr , reduced by formic acid to Pt° , gentle heating changes it to the oxide PtO . Platinic hydroxide, Pt(OH) 4 , is formed by treating a solution of H 2 PtCl„ with Na 2 C0 3 in excess, evaporating to dryness, washing with water and then with acetic acid. It is a red-brown powder, soluble in NaOH, HC1 , HNO„ , and H 2 S0 4 ; insoluble in HC 2 H 8 0, . Gentle heating changes it to the oxide Pt0 2 (Topsoe, B., 1870, 3, 462). 5. Solubilities. — a — Metal. — Platinifm is not affected by air or water, at any temperature; is not sensibly tarnished by hydrosulphuric acid gas or solution; and is not attacked at any temperature by nitric acid, hydrochloric acid or sulphuric acid, but dissolves in nitrohydroehloric acid (to platinic chloride) less readily than gold. 6. — Oxides and hydroxides. — See 4. c. — Salts. — Platinum forms two classes of salts (both haloid and oxy), platinous and platinic. The oxysalts are not stable. None of the platinous salts are permanently soluble in pure water. The chloride is soluble in dilute hydrochloric acid and the sul- phate in dilute sulphuric acid. Platinic chloride, PtCl 4 , and bromide, all the platinicyanides (as PbPt(CN)„), and the platinocyanides of the metals of the alkalis and alkaline earths (as K 2 Pt(CN) 4 ), are soluble in water. The platinous and platinic nitrates are soluble in water, but easily decomposed by it, with the precipitation of basic salts. The larger number of the metallo-platinic chlorides or " chloroplatinates " are soluble in water, including those with sodium [Na-jPtClu or (N'aCl) 2 PtCl 4 ], barium, strontium, magnesium, zinc, aluminum, copper; and those with potassium, and ammonium, are sparingly soluble in water, and owe their analytical importance as complete precipitates to their insolubility in alcohol. Of the metallo-platinous chlorides (the "chloroplatinites") — those with sodium [Na 2 PtCl 4 ], and barium, are soluble; zinc, potassium and ammonium, sparingly soluble; lead and silver, insoluble in water. Platinic sulphate, Pt(S0 4 ) 2 , is soluble in water (§10). 6. Eeactions. — a. — Platinous chloride, PtCl 2 , is precipitated by KOH as Pt(0H) 2 , soluble in excess of the reagent to K 2 Pt0 2 , potassium platinite, which solution is reduced by alcohol to "platinum black" (1). Platinic chloride, PtCl 4 , a brown-red solid, soluble in alcohol and water, forms with KOH or HH 4 OH , not too dilute, a yellow crystalline precipitate of an alkali (K or NH 4 ) platinum chloride, e. g., K 2 PtCl,, , sparingly soluble in water, soluble in excess §74, 7, i. PLATINUM. 95 of the alkalis and repreeipitated by hydrochloric acid. K 2 CO„ and (NH 4 )jCO, give the same precipitate, insoluble in excess of the reagent. A more complete precipitation of the K or NE, is obtained by the use of the chlorides. The sodium platinum chloride, Na 2 PtCl„ , is very soluble in water and is not formed by precipitation with sodium salts. 6. — Oxalic acid does not reduce platinum salts (distinction from gold). A solution of chloral hydrate precipitates pla- tinum from its solutions. Platinous and platinic salts form with cyanides a great number of double salts, c. — See 5c. 4. — Hypophosphorous acid reduces platinum salts to metallic platinum. Phosphates do not precipitate platinum salts. e. Hydrosulphuric acid precipitates solutions of the platinous salts as the hlack sulphide, PtS , insoluble in acids, sparingly soluble in water and in alkali sulphides; platinic salts are precipitated as platinic sulphide, PtS 2 , black; slowly soluble in alkali sulphides (Eibau, C. r., 1877, 85, 283), insoluble in acids except nitrohydrochloric. Sulphur dioxide decolors a solution of platinum chloride giving a compound which does not respond to the usual reagents for platinum and requires long boiling with HC1 for the removal of the S0 2 (Birnbaum, A., 1871, 159, 116). f. The chlorides of potassium and ammonium are estimated quantita- tively by precipitation from their concentrated solutions with a solution of platinic chloride. Potassium iodide colors a solution of platinum chloride brown-red and precipitates the black platinic iodide, Ptl 4 , excess of the KI forming K 2 PtI e , brown, sparingly soluble (5c). g. Stannous chloride does not precipitate the platinum from platinic chloride (distinc- tion from gold), but reduces it to platinous chloride. h. Ferrous sulphate solution on boiling with a platinum chloride solu- tion precipitates the platinum as the metal, the presence of acids hinders the reduction. 7. Ignition. — All platinum compounds upon ignition are reduced to the metal. Owing to the high point of fusibility of the metal and to the difficulty with which it is attacked by most chemicals, platinum has an extended use in the chemical laboratory for evaporating dishes, cruci- bles, foil, wire, etc. In the use of platinum apparatus without UNNECESSARY INJURY IT SHOULD BE REMEMBERED : (1) That free chlorine and or amine, attack platinum at ordinary tem- peratures (forming platinic chloride, bromide); and free sulphur, phos- phorus, arsenic, selenium, and iodine, attack ignited platinum (forming platinous sulphide, platinic phosphide, platinum-arsenic alloy, platinic selenide, iodide). Hence, the fusion of sulphides, sulphates, and phos- phates, with reducing agents, is detrimental or fatal to platinum crucibles. The ignition of organic substances containing phosphates acts as free phosphorus, in a slight degree. The heating of ferric chloride, and the fusion of bromides, and iodides, act to some extent on platinum. 96 PLATINUM. §74, 1, &. (#) The alkali hydroxides (not their carbonates) and the alkaline earths,, especially baryta and lithia, with ignited platinum in the air, gradually corrode platinum (by formation of platinites: 2Pt + 2BaO + 2 = 2BaPt0 2 . Silver crucibles are recommended for fusion ' with alkali hydroxides. (S) All metals which may he reduced in the fusion — especially compounds of lead, bismuth, tin, and other metals easily reduced and melted — and all metallic compounds with reducing agents (including even alkalis and earths)' form fusible alloys with ignited platinum. Mercury, lead, bismuth, tin,, antimony, zinc, etc., are liable to be rapidly reduced, and immediately to- melt away platium in contact with them. {4) Silica with charcoal (by formation of silicide of platinum) corrodes ignited platinum, though very slowly. Therefore, platinum crucibles should not be supported on charcoal in the furnace, but in a bed of mag- nesia, in an outer crucible of clay. Over the flame, the best support is the triangle of platinum wire. (5) The tarnish of the gas-flame increases far more rapidly upon the already tarnished surface of platinum — going on to corrosion and crack- ing. The surface should be kept polished — preferably by gentle rubbing with moist sea-sand (the grains of which are perfectly rounded, and do not scratch the metal). Platinum surfaces are also cleansed by fusing borax upon them, and by digestion with nitric acid. 8. Detection. — Platinum is identified by the appearance of the reduced metal ; by its insolubility in HC1 or HN0 3 and solubility in HN0 3 -f- HC1 ; and by its formation of precipitates with ammonium and potassium chlorides. It is separated from gold by boiling with oxalic acid and am- monium oxalate, which precipitate the gold, leaving the platinum in solu- tion. The filtrate from the gold should be evaporated, ignited, and the residue examined and after proving insolubility in HC1 or HN0 3 , dissolved in nitrohydrochloric acid and the presence of platinum confirmed with. NH 4 C1 . If the gold and platinum have been precipitated in the second group with H 2 S and dissolved with (NH 4 ) 2 S X they may be separated from As , Sb , and Sn by dissolving the reprecipitated sulphides in HC1 + KC10 3 , evaporating to remove the chlorine and boiling after adding KOH in ex- cess, with chloral hydrate, which precipitates the Au and Pt , leaving the As , Sb , and Sn in solution. The An and Pt may then be dissolved in HN0 3 -f- HC1 and separated as directed above. FeS0 4 may be use to pre- cipitate Au and Pt , separating them from As , Sb , and Sn . 9. Estimation. — Platinum is invariably weighed in the metallic state. It is brought to this condition: (1) By simple ignition; (2) by precipitation as (NH 4 ),PtCl„ , K 2 PtCl , or PtS 2 and ignition; (3) by reduction, using Zn , Mg, or FeS0 4 . 10. Oxidation. — Solutions of platinum are reduced to the metallic state by the §75, 6c. MOLYBDENUM. 97 following metals: Pb , Ag , Hg , Sn (Sn" to Pt" only), Bi , Cu , Cd , Zn , Pe , Pe" , Co , and Mi . Very many organic substances reduce platinum ■compounds to the metallic state. §75. Molybdenum. Mo = 96.0 . Valence two, three, four and six. 1. Properties.— Specific gravity, 8.56 (Loughlien, Am. 8., 1868, (2), 45, 131). Pure molybdenum appears not to have been melted; when heated to a very high heat in a graphite crucible it takes up carbon and melts. It is a silver- white, hard, brittle metal, not oxidized in the air or water at ordinary tem- peratures. Upon heating in the air it becomes brown, then blue, and finally burns to the white MoO s . Heated to a red heat in contact with steam, it forms first a blue oxide, then Mo0 3 . 2. Occurrence. — Not found native, but occurs chiefly as molybdenite, MoS,; as an oxide in molybdenum ochre, MoO s ; and as wulf enite, PbMo0 4 . 3. Preparation. — (i) By heating the oxide, sulphide or chloride in a current •of oxygen free hydrogen (von der Pfordten, B., 1884, 17, 732; Rogers and Mitchell, J. Am. Soc, 1900, 22, 350) ; (2) by heating with C and Na 2 CO„ ; (3) by heating Mo0 3 with KCN (Loughlien, I.e.). 4. Oxides and Hydroxides. — Molybdous hydroxide, Mo0.xH 2 O , is formed when molybdous chloride or nitrate is precipitated with alkali hydroxides or carbon- ates, dark brown becoming blue in the air by oxidation. Mo(OH), , black, turning red-brown by oxidation in the air, is formed by treating MoCl 3 with KOH; also by electrolysis of ammonium molybdate (Smith, B., 1880, 13, 751). By heating the hydroxide in a vacuum Mo.O : is obtained as a black mass, insoluble in acids. MoO., , a dark bluish mass, insoluble in KOH or HC1 , is formed by igniting a. mixture of ammonium molybdate, potassium carbonate and boric acid, and exhausting the fused mass with water (Muthmann, A., 1887, 238, 114). Molybdic anhydride (acid), Mo0 3 , white, occurs in nature; it is obtained by the ignition of the lower oxidized compounds in the air or in the presence of oxidizing agents. 5. Solubilities. — Molybdenum is readily soluble in nitric acid with oxidation to Mo0 3 , evolving NO; in hot concentrated sulphuric acid, evolving S0 2 . The various lower oxides of molybdenum are soluble in acids forming corresponding salts, not very stable, oxidizing on exposure, to molybdic acid and molybdates; on the other hand, reducing agents reduce molybdates to the lower forms of molybdenum salts, nearly all of which are colored brown to reddish brown or "violet. The salts of molybdenum are nearly all soluble in water. Molybdic anhydride, Mo0 3 , white, is sparingly soluble in water and possesses basic properties towards stronger acids, dissolving in them to form salts. The •chlorides and the sulphates are soluble in water (Sehulz-Sellack, B., 1871, 4, 14) ; the nitrates in dilute nitric acid. The anhydride MoO a combines with the alkalis to form molybdates, soluble in water. Molybdates of the other metals are insoluble in water. Solutions of the alkali molybdates are decomposed by acids forming, MoO B , which dissolves in excess of the acids. 6. Reactions. — a. — The dyad, triad and tetrad molybdenum salts are precipi- tated by the alkali hydroxides and carbonates, forming the corresponding hydroxides, insoluble in excess of the precipitant. These hydroxides oxidize in the air to a blue molybdenum molybdate. 6. — A solution of a molybdate acidulated with hydrochloric acid gives no red color with KCNS (distinction from Pe'") ; but if Zn be added, reduction to a lower oxide of molybdenum takes place and an intense red color is produced. Phosphoric acid does not destroy the color (difference from ferric thiocyanate) . Upon shaking with ether the sulphocyanate is dissolved in the ether, transferring the red color to the ether layer. In molybdic acid solutions, acidulated with hydrochloric acid, potassium ferrocyanide gives a reddish brown precipitate. An alkaline solution of molybdates is colored a deep red to brown by a solution, of tannic acid. c. — See 5, 98 MOLYBDENUM. §75, Qd. d. — Tribasic phosphoric acid and its salts precipitate, from strong nitric acid solutions of ammonium molybdate,* somewhat slowly, more rapidly on warming, ammonium phospho-molybdate, yellow, of variable composition, soluble in ammonium hydroxide and other alkalis, sparingly soluble in excess of the phosphate. Hydrochloric acid may be used instead of nitric. The sodium phospho-molybdate is soluble in water, and precipitates am- monium from its salts; also, it precipitates the alkaloids — for whicli reac- tion it has some importance as a reagent, f Arsenic acid and arsenates give the same reaction; ammonium arseno-molybdate being formed (g). e. — Neutral or alkaline solutions of molybdates are colored yellow to brown by hydrosulphuric acid but are not precipitated. From the acid solutions a small amount of the hydrogen sulphide gives no precipitate but colors the solution blue; with more hydrosulphuric acid the brown or red-brown precipitate, MoS., , molybdenum trisulphide, is obtained after some time. The precipitate is soluble in ammonium sulphide, better when hot and not too concentrated, as ammonium thiomolybdate, (NH 4 ) 2 MoS 4 , from which acids precipitate the trisulphide (Berzelius, Pogg., 1826, 7, 429), soluble in nitric acid, insoluble in boiling solution of oxalic acid (separation from stannic sulphide). If Na. 2 S 2 3 be added to a solution of ammonium molybdate, slightly acid, a, blue precipitate and blue-colored solution is obtained. If the solution be more strongly acid, a red brown precipitate is obtained. An acid solution of a molybdate treated with hypophosphorous and sulphurous acids gives an in- tense bluish green precipitate or color, depending upon the amount of molyb- denum present. f. — Halogen compounds not important in analysis of molybdenum. g. — Arsenic acid and arsenates form, with a nitric acid solution of ammonium molybdate, a yellow precipitate of ammoniwm arseno-molybdate, in appearance and reactions not to be distinguished from the ammonium phospho-molybdate; except the precipitation does not take place until the solutions are slightly warmed, while with phosphates the precipitation begins even in the cold. Stannous salts give with (NH 4 ) 2 Mo0 4 a blue solution of the lower oxides of molybdenum (a delicate test for Sn") (Longstaff, G. N., 1899, 79, 282). h. — Solutions of the alkali molybdates are soluble in water and precipitate solutions of nearly all other metallic salts, forming molybdates of the corre- sponding metals, insoluble in water* e.g., K 2 ]Mo0 4 + Pb(N0 3 ) 2 = PbMoO, + 2KNO, . * The reagent ammonium molybdate, (NH,!, MoO„ is prepared by dissolving molybdic acid, MoO, (100 grams), in ammonium hydroxide (250 cc. sp. gr. 0.90 with 250 cc. water) cooling, and slowly pouring this solution into well cooled fairly concentrated nitric acid (750 cc. sp. gr. 1.42 with 750 co. water) with constant stirring. t Sodium. Phospho-molybdate— Sonnenschein's reagent for acid solutions of alkaloids— is pre- pared as follows : The yellow precipitate formed on mixing acid solutions of ammonium molyb- date and sodium phosphate— the ammonium phospho-molybdate— is well washed, suspended in water, and heated with sodium carbonate until completely dissolved. The solution is evapor- ated to dryness, and the residue gently ignited till all ammonia is expelled, Bodium being sub- stituted for ammonium. If blackening occurs, from reduction of molybdenum, the residue is moistened with nitric acid, and heated again. It is then dissolved with water and nitric acid to strong acidulation ; the solution being made ten parts to one part of residue. It must be kept from contact with vapor of ammonia, both during the preparation and when preserved for use. §75, 10. MOLYBDENUM. 99 7. Ignition. — With microcosmic salt, in the outer blow-pipe flame, all com- pounds of molybdenum give a bead which is greenish while hot, and colorless on cooling; in the inner flame, a clear green bead. With borax, in the outer flame, a bead, yellow while hot, and colorless on cooling; in the inner flame, a brown bead, opaque if strongly saturated (molybdous oxide). On charcoal, in the outer flame, molybdic anhydride is vaporized as a white incrustation; in the inner flame (better with sodium carbonate), metallic molybdenum is obtained as a gray powder^ separated from the mass by lixiviation. Dry molyb- dates, heated on platinum foil with concentrated sulphuric acid to vaporiza- tion of the latter form, on cooling in the air, a blue mass. 8. Detection. — In the ordinary process of analysis, molybdenum appears in Division A (tin group) of the second group with As , Sb , Sn , Au , and Pt . The solution remaining in the Marsh apparatus is decanted from the residue (Sn, Sb, Au, Pt and excess of Zn) and heated with concen- trated HN0 3 , the molybdenum is oxidized to molybdic acid. This solution, evaporated to dryness, dissolved in ammonium hydroxide and poured into moderately concentrated HC1 forms a solution of ammonium molybdate which may be identified by the many precipitation and reduction tests (6 b, c, d, e, i, etc., 7, and 9). If the molybdenum be present as a molybdate it may be precipitated from its nitric acid solution by Na 2 HP0 4 ? washed, dissolved in ammonium hydroxide, the phosphate removed by magnesia mixture (§189, 6a), and the filtrate evaporated to crystallization (Maschke, Z., 1873, 12, 380). The crystals may be tested by the various reduction tests for molybdenum. 9. Estimation. — (1) Molybdic anhydride and ammonium molybdate may be reduced to the dioxide by heating in a current of hydrogen gas. The heat must not be permitted to rise above dull redness. Or the temperature may rise to a white heat, which reduces it to the metallic state, in which form it is weighed. (2) Lead acetate is added to the alkali molybdate, the precipitate washed in hot water, and after ignition weighed as PbJiloO,, . (3) Volumet- rically. The molybdic acid is treated with zinc and HC1 , which converts it into M0CI3 . This is converted into molybdic acid again by standard solution of potassium permanganate. 10. Oxidation. — Reducing agents convert molybdic acid either into the blue intermediate oxides, or, by further deoxidation, into the black molybdous oxide, MoO . In the (hydrochloric) a - cid solutions of molybdic acid, the blue or black oxide formed by reduction, will be held in solution with a blue or brown color. Nitric acidulation is, of course, incompatible with the reduction. Certain reducing agents act as follows: Cane sugar in the feebly acid boiling solution, forms the blue color — seen better after dilution; a delicate test. Stannous chloride forms first the blue, then the brown, or the greenish brown to black-brown, solution of both the intermediate oxide and the molybdous oxide. Zinc, with HC1 or H 2 S0 4 , gives the blue, then green, then brown color, by progressive reduction. Formic and oxalic acids do not react. A solution of 1 milligram of sodium (or ammonium) molybdate in 1 cc. of concentrated sulphuric acid (about 1 part to 1840 parts) is in use as Frcehde's Reagent for alkaloids. The molybdenum in this solution, which must be freshly prepared for use each time, is reduced by very many organic substances; and with a large number of alkaloids, it gives distinctive colors, blue, red, brown and yellow. 100 BISMUTH. §76, 1. The Copper Group (Second Group, Division B). Mercury (Mercuricum), Lead, Bismuth, Copper, Cadmium (Ruthenium, Rhodium, Palladium, Osmium). §76. Bismuth, Bi = 208.1 . Valence three and five. 1. Properties. — Specific gravity, 9.7474 (Classen, B., 1890, 23, 938) ; melting point, 269.22 (Callendar and Griffiths, C. N., 1891, 63, 2) ; it vaporizes at 1700° and the density of the vapor shows that the molecule Bi has begun to dissociate (Biltz and V. Meyer, B., 1889, 22, 725). It is a hard, brittle, reddish-white, lustrous metal; forming' beautiful rhombohaedral crystals when a partially cooled mass is broken into and the still molten mass decanted. Alloys of bismuth with other metals give compounds of remarkably low melting points, e. g., an alloy of: Bi two, Sr± one, and Pb one part by weight melts at 93.7°; and an alloy of: Bi fifteen, Pb eight, Sn, four, and Cd three parts by weight melts at 68° " Wood's Metal." 2. Occurrence. — It is a comparatively rare metal, not very widely distributed, usually found native. It is found in greatest quantities in Saxony; also found in Bohemia, France, England and South America. As mineralogical varieties it occurs as bismuth ochre (Bi 2 3 ), bismuthite (4Bi 2 O a .3C0 2 .4H 2 0), bismuth glance (Bi 2 S 3 ), etc. 3. Preparation. — The rock containing bismuth, usually with large amounts of cobalt, etc., is roasted to remove sulphur and arsenic, which is nearly always present. The mass is then fused with charcoal. The molten bismuth settles to the bottom below the layer of cobalt. The cobalt becomes solid •while the bismuth is still molten, and the two are separated mechanically. The metal is further purified by melting with KNO, or KC1T . 4. Oxides. — Bismuth trioxide, Bi 2 3 , is formed by heating the metal in the presence of air, or by igniting the hydroxide; it is a pale citron-yellow powder. The hydroxide, Bi(OH) 3 , white, is formed by precipitating a solution of a salt of bismuth with an alkali hydroxide. If bismuth chloride is used the hydroxide formed always contains some oxychloride, BiOCl (Strohmeyer, Pogg., 1832, 26, 549). The meta hydroxide, BiO(OH) , is formedupon drying the orthohydroxide at 100° (Arppe, Pogg., 1845, 64, 237). Bismuth pentoxide, Bi 2 5 , is formed by igniting Bi(OH) 3 with excess of KOH or NaOH in presence of the air, and washing the cooled mass repeatedly with cold dilute nitric acid (Strohmeyer, I. c); or by treating Bi(OH) 3 with three per cent H 2 2 in strong alkaline solu- tion (Hasebrock, B., 1887, 20, 213). It is a heavy dark brown powder. At 150° it gives off O, and at the temperature of boiling mercury becomes Bi 2 3 . It is decomposed in the cold by HC1 with evolution of chlorine. BismutMc acid, HBiO, , or more probably Bi 2 5 .BLO , is formed upon conducting a rapid current of chlorine into Bi(OH), suspended in concentrated KOH solution. It is a beautiful scarlet red powder which at 120° gives off its water, becoming Bi.Oji (Muir, J. 0., 1876, 29, 144; Muir and Carnegie, J. C, 1887, 51, 86). It is doubtful if any alkali salt of bismuthic acid exists, although mixtures of KBi0 3 and HBi0 3 are claimed by Hoffmann (A., 1884, 223, 110), and Andre (0. r., 1891, 113, 860). The so-called oisrrmih tetroxide, Bi 2 4 , is probably a mixture of the trioxide and pentoxide (§12). 5. Solubilities. — a. — Metal. — Metallic bismuth is insoluble in hydrochloric acid *; soluble in warm concentrated sulphuric acid with evolution of sulphur dioxide; readily soluble in nitric acid and in nitrohydrochloric acid. It burns in chlorine with production of light; it combines with bromine, but more slowly than antimony; it combines readily upon fusing together with I , S , Se , Te , As , and Sb , besides the many metals with which it combines to form com- * A trace of bismuth can always be found in solution when the metal Is boiled with hydro- ■ehloric acid, but no more than whea the metal has been boiled with pure water (Ditte and Uetzner,' A. Ch., 1896, (6), 39, 389), §76, 6a. BISMUTH. 101 mercial alloys (1). The halogen derivatives of pentad bismuth are not known (Muir, J. C, 1876, 29, 144). 6 — Oxides and hydroxides. — Bismuth oxide, Bi 2 0„ , and the hydroxides, Bi(OH) 3 and BiO(OH), are soluble in hydrochloric, nitric and sulphuric acids; insoluble in water and the alkali hydroxides or carbonates. The presence of glycerol prevents the precipitation of bismuth hydroxides "from solutions of its salts by the alkalis.* Bismuth pentoxide, Bi 2 6 , is solu- ble in HC1 , HBr , and HI with evolution of the corresponding halogen and formation of the triad salt. Nitric and sulphuric acids in the cold have but little or no action; when hot the triad bismuth salt is formed with evolution of oxygen. c. — Salts. — Most of the salts of bismuth are insoluble in water. The ■chloride, bromide, iodide, nitrate,- and sulphate are soluble in water acidu- lated with their respective acid, or with other acids forming " soluble " "bismuth salts. ■ Pure water decomposes the most of the solutions of bis- muth salts forming corresponding oxy-salts (§70, 5d footnote). The chloride, bromide and sulphate are deliquescent. d. — Water. — A solution of bismuth chloride in water acidulated with lydrochloric acid is precipitated on further dilution with water, bismuth oxy-chloride, BiOCl being formed; e. g., BiCl 3 + H 2 = BiOCl + 2HC1, insoluble in tartaric acid (distinction from antimony, §70, 5d). The hydro- chloric acid set free serves to hold a portion of the bismuth in solution. The presence of acetic, citric, and other organic acids prevents the pre- cipitation of solutions of bismuth salts upon further dilution with water. The washing of the precipitated oxy-salt with pure water removes more of the acid forming a salt still more basic. Bi(Br0 3 ) 3 -(- H 2 =BiONO a + 2HN0 3 12BiON0 3 + H 2 = 6Bi 2 O s ,51^05 + 2HN0 3 'This is prevented by the presence of one part ammonium nitrate to five hundred parts water (Lowe, J. pr., 1858, 74, 341). Bismuth nitrate crystallizes with ten molecules of water, Bi(N0 3 ) 3 . 10H 2 . It is decomposed by a small amount of water forming the basic nitrate, BiON0 3 ; this is soluble in dilute nitric acid, when further dilution with water to any extent is possible without precipitation of the basic «alt, but a drop of hydrochloric acid or a chloride causes a precipitate of -the oxychloride in the diluted solution. The bromide is readily decom- posed by water to BiOBr ; the iodide is stable to cold water, but is decom- posed by hot water to BiOI (Schneider, A. Ch., 1857 (3), 50, 488); the normal sulphate very readily absorbs water to form Bi 2 (S0 4 ) 3 .3H 2 , which is decomposed by more water to Bi 2 3 .S0 3 . 6. Reactions, a. — The alkali hydroxides precipitate from solutions of bismuth salts bismuth hydroxide, Bi(0H) 3 , white; insoluble in excess of the fixed alkalis (distinction from Sb and Sn), insoluble in ammonium * Lowe (O. N., 1883, 45, 296) dissolves the hydroxides of copper and bismuth in glycerol, adds jglucose and gently warms. The copper is completely precipitated and separated from the bis- muth. Upon boiling the filtrate for some time the bismuth is completely precipitated as the .metal. 102 BISMUTH. §76, 65. hydroxide (distinction from Cu and Cd). The hydroxide is converted by boiling into the oxide, Bi 2 3 , yellowish white. The precipitation is pre- vented by the presence of tartaric acid, citric acid, glycerol, and certain other organic substances (Kohler, J. C, 1886, 50, 428). The alkali carbonates precipitate basic bismuth, carbonate, Bi 2 3 .CO = , white, insoluble in excess of the reagent. Freshly precipitated barium carbonate forms the same precipitate without heating. 6. — Oxalic acid and soluble oxalates precipitate bismuth oxalate, Bi 2 (C 2 4 ) a ,. white, soluble in moderately dilute acids. Potassium cyanide forms a white crystalline precipitate insoluble in excess of the reagent but soluble in nitric or hydrochloric acid. Potassium ferrocyanide forms a yellowish white pre- cipitate, potassium ferricyanide a brownish yellow, both soluble in hydrochloric acid. c. — The action of nitric acid upon bismuthi and its salts is fully explained under (5). d. — Metallic bismuth is precipitated when bismuth salts are warmed with hypophosphorous acid (separation from Zn and Cd) (Muthmann and Mawron, Z., 1874, 13, 209). From solutions of bismuth nitrate (5). §77. Copper (Cuprum) Cu = 63.6 . Valence one and two. 1. Properties. — Specific gravity, electrolytic, 8.914; melted, 8.921; natural crys- tals, 8.94; rolled and hammered sheet, 8.952 to 8.958 (Marchand and Scheerer, J. pr., 1866, 97, 193). Melting point, 1080.5 (Heycock and Neville, J. C, 1895, 67, 190). A red metal, but thin sheets transmit a greenish-blue light, and it also shows the same greenish-blue tint when in a molten condition. Of the metals in ordinary use, only gold and silver exceed it in malleability. In ductility it is inferior to iron and cannot be so readily drawn into exceedingly fine wire. Although it ranks next to iron in tenacity, its wire bears about half the weight which an iron wire of the same size would support. As a conductor of heat it is surpassed only by gold. Next to silver it is the best conductor of electricity. Dry air has no action upon it; in moist air it becomes coated with a film of oxide which protects it from further action of air or of water. It forms a number of very important alloys with other metals; bronze (copper and tin), brass (copper and zinc with sometimes small amounts of lead or tin), German silver (copper, nickel and zinc). 2. Occurrence. — Copper is found native in various parts of the world, and especially in the region of Lake Superior. It is found chiefly as sulphides in enormous quantities in Montana, Colorado, Chili and Spain; as a carbonate in Arizona. It is very widely distributed and occurs in various other forms. Copper pyrites is Cu3?eS 2 ; copper glance, Cu 2 S; green malachite, Cu 2 (OH) 2 C0 8 ; blue malachite, Cu 8 (OH) 2 (C0 3 ) 2 ; red copper ore, Cu 2 0; and tenorite, CuO . 3. Preparation. — For the details of the various methods of copper-smelting and refining, the works on metallurgy should be consulted. In the laboratory pure copper may be produced (J?) by electrolysis; (2) reduction by ignition in hydrogen gas; (3) reduction of the oxide by ignition with carbon, carbon monoxide, illuminating gas, or other forms of carbon; (4) reduction of the oxide by K or Ma at a temperature a little above the melting point of these metals; (5) reduction by fusion with potassium cyanide: CuO + KCN = Cu + KCNO . For its reduction in the wet way, see 10. 4. Oxides and Hydroxides. — Cuprous oxide (Cu 2 0), red, is found native; it is prepared: (i) by reducing CuO by means of grape-sugar in alkaline mixture; (2) by igniting CuO with metallic copper; (3) by treating an ammoniacal cupric ■solution with metallic copper; then adding KOH and drying. Cuprous hydrox- ide, CuOH , brownish yellow, is formed by precipitating cuprous salts with KOH or NaOH . Cupric oxide, CuO , black, is formed by igniting the hydroxide, §77, 5c. COPPER. 105 carbonate, sulphate, nitrate and some other cupric salts in the air; or by- heating the metal in a current of air. Cupric hydroxide, Cu(OH) 2 , is formed by precipitating cupric salts with KOH or NaOH . It is stated by Rose (Pogg.,. 1863, 120, 1) that tetracupric monoxide, (Cu 4 , is formed by treating a cupric salt with KOH and a quantity of K SnO ., insufficient to reduce it to the metallic state. A peroxide of copper, CuO, , is supposed to be formed by treating Cu(OH) 2 with H 2 2 at 0° (Kriiss, B., 1884, 17, S593). (§10.) 5. Solubilities. — a. — Metal. — Copper does not readily dissolve in acids with evolution of hydrogen; it dissolves most readily in nitric acid chiefly with evolution of nitric oxide- 3Cu + 8HN0 3 = 3Cu(N0 s ) 2 + 4H 2 + 2NO (Freer and Higley, Am., 1899, 21, 37-7) ; also in hot concentrated sulphuric acid, with evolution of sulphurous anhydride: Cu + 2H 2 S0 4 = CuS0 4 + 2H 2 + S0 2 . If dry hydrochloric acid gas be passed over heated copper, CuCl is formed with evolution of hydrogen (Weltzien, A. CK, 1865, (4), 6, 487). A saturated solution, of hydrochloric acid at 15° dissolves copper as CuCl with evolution of hydrogen. The action is very rapid if the copper be first immersed in a platinum chloride solution. Heat favors the reaction and the presence of 10H.O to one HC1 pre- vents the action (Engel, C. r., 1895, 121, 528). Hydrobromic acid concentrated acts slowly in the cold and rapidly when warmed, forming CuBr 2 , with evolu- tion of hydrogen. Cold hydriodic acid, in absence of iodine, is without action (Mensel, B., 1870, 3, 123). Ammonium sulphide, (NH t ),S , colorless, acts upon copper turnings with evolution of hydrogen, forming Cu 2 S (Heumann, /. C.,, 1873, 26, 1105). o. — Oxides. — Cuprous oxide and hydroxide are insoluble in water, soluble- in hydrochloric acid with formation of cuprous chloride, white, unstable,, readily oxidized by the air to colored cupric salts. Cupric oxide, black, and hydroxide, blue,- are insoluble in water, soluble in dilute acids; in a mixture of equal parts glycerine and sodium hydroxide, sp. gr. 1.20 (sepa- ration from Cd) (Donath, J. C, 1879, 36, 178), in a mixture of tartrates and fixed alkalis (but precipitated as Cu 2 by heating with glucose) -{sepa- ration from Cd and Zn) (Warren, C. N., 1891, 63, 193); insoluble in ammonium hydroxide in absence of ammonium salts (Maumene, J. C. r 1882, 42, 1266). c. — Salts. — All salts of copper, except the sulphides, are soluble in am- monium hydroxide. All cuprous salts are insoluble in water, soluble in hydrochloric acid and reprecipitated upon addition of water. They are- readily oxidized to cupric salts on exposure to moist air. Cuprous chloride and bromide are soluble in ammonium chloride solution (Mohr, J. C, 1874, 27, 1099). Cupric salts, in crystals or solution, have a green or blue color; the chloride (2 aq.) in solution is emerald-green when concen- trated, light blue when dilute; the sulphate (5 dq.) is "blue vitriol." Anhydrous cupric salts are white. The crystallized chloride and chlorate are deliquescent; the sulphate, permanent; the acetate, efflorescent. Cupric basic carbonate, oxalate, phosphate, borate, arsenite, sulphide,, cyanide, ferrocyanide, ferricyanide, and tartrate are insoluble in water. The ammonio salts, the potassium and sodium cyanides, and the potassium and sodium tartrate, are soluble in water. In alcohol the sulphate and acetate are insoluble.; the chloride and nitrate, soluble. Ether dissolves, the chloride. 106 COPPER. §77, 6a. 6. Reactions. — a. — Fixed alkali hydroxides precipitate acid solutions ai cuprous chloride, first as the white cuprous chloride, changing with more of the alkali to the yellow cuprous hydroxide, insoluble in excess. Ammonium hydroxide and carbonate precipitate and redissolve the hydroxide to a color- less solution, which turns blue on exposure. The colorless ammoniacal solution is precipitated by potassium hydroxide. Fixed alkali carbonates precipitate the yellow' cuprous carbonate, Cu 2 CO s . Fixed alkalis — KOH — added to saturation in solutions of cupric salts, 'precipitate cupric hydroxide, Cu(0H) 2 , deep blue, insoluble in excess unless ■concentrated (Loew, Z., 1870, 9, 463), soluble in ammonium hydroxide (if too much fixed alkali is not present), very soluble in acids, and changed, by standing, to the black compound, Cu 3 2 (0H) 2 ; by boiling, to CuO . If tartaric acid, citric acid, grape-sugar, milk-sugar, or certain other organic substances are present, the precipitate either does not form at all, -or redissolves in excess of the fixed alkali to a blue solution. The alkaline tartrate solution may be boiled without change; in presence of glucose, the application of heat causes the precipitation of the yellow cuprous oxide. Alkali hydroxides, short of saturation, form insoluble basic salts, -of a lighter blue than the hydroxide. Ammonium hydroxide added short of saturation precipitates the pale •blue basic salts; added just to saturation, the deep blue hydroxide (in both ■cases like the fixed alkalis); added to supersaturation, the precipitate dis- solves to an intensely deep blue solution (separation from bismuth). The iblue solution is a cuprammonium compound, not formed unless ammonium :salts be present. It has been isolated as CuS0 4 .(NH 3 ) 4 (§77, 5b). The deep blue solution probably consists of this compound in a hydrated condition, i. e. Cu(0H) 2 .2NH 4 0H.(NH 4 ) 2 S0 4 ; or (NH 4 ) 4 Cu(0H) 4 S0 4 . Other salts than the sulphate form the corresponding compounds: CuCl 2 -4- 4NH 4 0H = Cu(0H) 2 .2NH 4 0H.2NH 4 Cl . The blue color with ammonium hydroxide is a good test for the presence of copper in all but traces (one to 25,000), its sensitiveness is diminished by the presence of iron (Wagner, Z., 1881, 20, 351). Ammonium carbonate, like ammonium hydroxide, precipitates and redissolves to a blue solution. Carbonates of fixed alkali, metals — as X 2 C0 3 — precipitate the greenish-blue, basic carbonate, Cu 2 (0H) 2 C0 3 , of variable composition, according to conditions, and converted by boiling to the black, basic hydroxide and finally to the black oxide. Barium carbon- ate precipitates completely, on boiling, a basic carbonate. From the blue ammoniacal solutions a concentrated solution of a fixed alkali precipitates the blue hydroxide, changed on boiling to the black oxide, CuO . 6. — Oxalates, cyanides, ferrocyanides, ferricyanides and thiocyanates pre- cipitate their respective cuprous salts from cuprous solutions not too strongly acid. The ferricyanide is brownish-red, the others are white. The thiocyanate is used to separate copper from palladium (Wohler, A. Ch., 1867, (4), 10, 510); and -also from cadmium. In solutions of cupric salts, oxalates precipitate cupric §77, 6e. COPPER. 107 oxalate, CuC 2 4 , bluish-white, insoluble in acetic acid, and formed from mineral acid salts of copper by oxalic acid added with alkali acetates. Potassium cyanide forms the yellowish-green cupric cyanide, Cu(CN) 2 , soluble in excess of the reagent with formation of the double cyanide, 2KCN.Cu(CN) 2 , unstable, changing in whole or in part to cuprous cyanide. The potassium cyanide also dissolves cupric oxide, hydroxide, carbonate, sulphide, etc., changing rapidly to cuprous cyanide in solution in the alkali cyanide. This explains why hydrogen sulphide does not precipitate solutions of copper salts in potassium cyanide, used as a separation from cadmium. Potassium ferrooyanide precipitates cupric ferrocyanide, Cu 2 Fe(CN) e , reddish-brown, insoluble in acids, decomposed by alkalis ; a very delicate test for copper (1 to 200,000) ; forming in highly dilute solu- tions a reddish coloration ("Wagner, Z., 1881, 20, 351). Potassium ferri- cyanide precipitates cupric ferricyanide, Cu 3 (Fe(CN) 6 ) 2 , yellowish-green, insoluble in hydrochloric acid. Potassium thiocyanate, with cupric salts, forms a mixed precipitate of cuprous thiocyanate, white, and a black precipitate of cupric thiocyanate, which gradually changes to the white cuprous compound, soluble in NH,OH; in the presence of hypophosphorous or sulphurous acid the cuprous thiocyanate is precipitated at once (distinction from cadmium and zinc) (Hutchinson, J. C, 1880, 38, 748). Ammonium benzoate (10 per cent solution) precipitates copper salts completely from solutions slightly acidified (separation from cadmium) (Gucci, B., 1884, 17, 2659). If to a solution of cupric salt slightly acidulated with hydrochloric acid, an excess of a solution of nitroso-B-naphthol in 50 per cent acetic acid be added, the copper will be completely precipitated on allowing to stand a short time (separation from Pb , Cd , Hg , Mn , and Zn) (Knorre, B., 1887, 20, 283). Potassium xanthate gives with very dilute solutions of copper salt a yellow- coloration; according to Wagner (I. c.) one part copper in 900,000 parts water may be detected. c. — Nitric acid rapidly oxidizes cuprous salts to cupric salts, d. — A solution of cupric sulphate slightly acidulated with hydrochloric acid is precipitated as cuprous chloride by sodium hypophosphite (Cavazzi, Gazzetta, 1886, 16, 167) ; if the slightly acidulated copper salt solution be boiled with an excess of the hypophosphite the copper is completely precipitated as the metal. Sodium phosphate, Na,HP0 4 , gives a bluish-white precipitate of copper phosphate, CuHP0 4 , if the reagent be in excess and Cu 3 (P0 4 ) 2 if the copper salt be in excess. Sodium pyrophosphate precipitates cupric salts, but not if tartrates or thiosulphates be present (separation from cadmium) (Vortmann, B., 1888, 21, 1103). c. — Cuprous salts (obtained by treating cupric salts with SnCl 2 ) when boiled with precipitated sulphur deposit the copper as Cu 2 S (separation from cad- mium) (Orlowski, J. C, 1882,42, 1232). Cuprous salts are precipitated or trans- posed by hydrosulphuric acid or soluble sulphides, forming cuprous sulphide,* Cu 2 S , black, possessing the same solubilities as cupric sulphide. With cupric salts H 2 S gives CuS, black (with some Cu 2 S), produced alike in acid solutions (distinction from iron, manganese, cobalt, nickel) * Freshly precipitated cuprous sulphide transposes silver nitrate forming silver sulphide, metallic silver and cupric nitrate ; with cupric sulphide, silver sulphide and cupric nitrate are formed (Schneider, Pogg., 1874, 152, 471). Freshly precipitated sulphides of Fe, Co, Zn, Cd, Pb, Bi, Sm", and Sn IT , when boiled with CuCl in presence of NaCl give Cu 2 S and chloride of the metal: with CuCI„, CuS and a chloride of the metal are formed, except that SnS gives Cu a S, CuCl and Sn" (Baschig, B., 1884, 17, 697). 108 COPPER. §77, 6f. and in alkaline solutions (distinction from arsenic, antimony, tin). — Solu- tions containing only the one-hundred-thousandth of copper salt are- colored hrownish by the reagent. The precipitate, CuS , is easily soluble- by nitric acid (distinction from mercuric sulphide); with difficulty soluble by strong hydrochloric acid (distinction from antimony) ; insoluble in hot dilute sulphuric acid (distinction from cadmium) ; insoluble in fixed alkali sulphides, and but slightly soluble in ammonium sulphide (distinction from arsenic, antimony, tin); soluble in solution of potassium cyanide- (distinction from lead, bismuth, cadmium, mercury). Concerning the formation of a colloidal cuprio sulphide, see "Spring- (B., 1883,, 16, 1142). According- to Brauner (C. N., 1896, 74, 99) cupric salts witk excess- of hydrogen sulphide always yield a very appreciable amount of cuprous- sulphide. See also Ditte (0. r., 1884, 98, 1492). Solutions of cupric salts are reduced to cuprous salts by boiling with sulphurous acid (Kohner, C. C, 1886, 813). Sodium, thiosulphate added to hot solutions of copper salts gives a black precipitate of cuprous sulphide. In solutions acidulated with hydrochloric: acid, this is a separation from cadmium (Vortmann, M., 1888, 9, 165). /. — Hydrobromic acid added to cupric solutions and concentrated by evaporation gives a rose-red color. Delicate to 0.001 m. g. (Endemann and Prochazka, C. N., 1880, 42, 8). Of the common metals only iron interferes. Potassium bromide and sulphuric acid may be used instead of hydrobromic acid. Hydriodic acid and soluble iodides precipitate, from concentrated solu- tions of copper salts, cuprous iodide, Cul , white, colored dark brown by the iodine separated in the reaction * (a). The iodine dissolves with color in excess of the reagent, or dissolves colorless on adding ferrous sulphate or soluble sulphites, by entering into combination. Cuprous iodide' dissolves in thiosulphates (with combination). The cuprous iodide is precipitated, free from iodine, and more com- pletely, by adding reducing agents with iodides; as, Na 2 S0 3 , H 2 S0 3 , FeSO, (&). (a) 2CuS0 4 + 4KI = 2CuI + I 2 + 2K 2 S0 4 (6) 2CuS0 4 + 2KI + 2FeS0 4 = 2CuI + K 2 S0 4 + Fe 2 (S0 4 ) 3 2CuS0 4 + 4KI + H 2 SO s + H 2 = 2CuI + 2K 2 S0 4 + H 2 S0 4 + 2HI g— Arsenites, as KAs0 2 , or arsenous acid with just sufficient alkali hydrox- ide to neutralize it, precipitate from solutions of cupric salts (not the acetate) the green copper arsemte, chiefly CuHAs0 3 (Scheete's green, "Paris green"), readily soluble in acids and in ammonium hydroxide, decomposed by strong potassium hydroxide solution. From cupric acetate, arsenites precipitate, on boiling, copper aceto-arsenite, (CuOAs 2 O s ) s Cu(C 2 H s 2 ) 2 , Schweinfurt green or Imperial green, " Paris green," dissolved by ammonium hydroxide and by acids, decomposed by fixed alkalis. Soluble arsenates precipitate from solutions of cupric salts cupric arsenate, bluish-green, readily soluble in acids and in ammonium hydroxide, h. — Potassium bichromate does not precipitate solutions of cupric salts; * The precipitation is incomplete unless the free iodine, one of the products of the reaction, is removed by means of a reducing agent (§44). §77, 10. COPPER. 109 normal potassium chromate forms a brownish-red precipitate, soluble in am- monium hydroxide to a green solution, soluble in dilute acids. 7. Ignition. — Ignition with sodium carbonate on charcoal leaves metallic copper in finely divided grains. The particles are gathered by triturating the charcoal mass in a. small mortar, with the repeated addition and decantation of water until the copper subsides clean. It is recognized by its color, and its softness under the knife. Copper readily dissolves, from its compounds in beads of borax and of microcosmic salt, in the outer flame of the blow-pipe. The beads are green while hot, and' blue when cold. In the inner flame the borax bead becomes colorless when hot; the microcosmic salt turns dark green when hot, both having a reddish-brown tint when cold (Cu 2 0) (helped by add- ing tin). Compounds of copper, heated in the inner flame, color the outer flame green. Addition of hydrochloric acid increases the delicacy of the reaction, giving a greenish-blue color to the flame. 8. Detection. — Copper is precipitated from its solutions by H 2 S , form- ing CuS . By its insolubility in (NH 4 ) 2 S x and solubility in hot dilute HN0 3 it is separated with Pb , Bi , and Cd from the remaining metals of the tin and copper group. Dilute H 2 S0 4 with C 2 H 5 0H removes the lead and ammonium hydroxide precipitates the bismuth as Bi(0H) 3 , leaving the Cu and Cd in solution. The presence of the Cu is indicated by the blue color of the ammoniacal solution, by its precipitation as the brown ferro- cyanide after acidulation with HC1 (6&) ; and by its reduction to Cu° with Fe°, from its neutral or acidulated solutions (10). Study the text on reactions (6) and §102 and §103. 9. Estimation. — (1) It is precipitated on platinum by the electric current or by means of zinc, the excess of zinc may be dissolved by dilute hydrochloric acid. (2) It is converted into CuO and weighed after ignition, or the oxide is reduced to the metal in an atmosphere of hydrogen and weighed as such. (3) It may be precipitated either by H 2 S or Ha 2 S 2 03 , and, after adding free sulphur and igniting in hydrogen gas, weighed as cuprous sulphide, or it may be precipitated by KCNS in presence of H 2 S0 3 or H s P0 2 , and, after adding S , ignited in H and weighed as Cu 2 S . Cu 2 , CuO, Cu(2TO s ) 2 ,CuCO B , CuS0 4 , and many other cupric salts, are converted into Cu 2 S by adding S and igniting in hydrogen gas. (Jf) By adding KI to the cupric salt and titrating the liber- ated I by Na 2 S 2 8 ; not permissible with acid radicals which oxidize HI . (5) By precipitation as Cul and weighing after drying at 150° (Browning, Am. 8., 1893 [3], 46, 280). (6) By titrating in concentrated HBr , using a solution of SnCL in concentrated HC1; the end reaction is sharper than with SnCl 2 alone (Etard and Lebeau, C. r., 1890, 110, 408). (7) By titration with Ma 2 S. Zinc does not interfere (Borntrager, Z. angew., 1893, 517). (8) By reduction with S0 2 and precipitation with excess of standard NH 4 C1TS; dilu- tion to definite volume and titration of the excess of NH^CNS in an aliquot part, with AgNO, (Volhard, A., 1878, 190, 51). (9) Small amounts are treated with an excess of NH,OH and estimated colorimetrically by comparing with standard tubes. \ 10. Oxidation. — Solutions of Cu" and Cu' are reduced to the metallic state by Zn , Cd , Sn , Al , Pb , Fe , Co , Ni , Bi , Mg *, P , and in presence of K0H by K 2 Sn0 2 . A bright strip of iron in solution of cupric salts acidulated with hydrochloric acid, receives a bright copper coating, recog- nizable from solutions in 120,000 parts of water. With a zinc-platinum * Warren, 0. N., 1895, 71, 93. HO CADMIUM. §78, 1. couple the copper is precipitated on the platinum and its presence can be confirmed by the use of H 2 S0 4 , concentrated, and KBr - an intense violet color is obtained (Creste, J. C, 1877, 31, 803). Cu" is reduced to Cu' by Cu° (Boettger, J. C, 1878, 34, 113), by SnCl 2 in presence of HC1, in presence of KOH by As 2 3 and grape sugar, by HI , and by S0 2 . Metallic copper is oxidized to Cu" by solutions of Hg", Hg', Ag', Pt IV , and Au'", these salts being reduced to the metallic state. Ferric iron is reduced to the ferrous condition (Hunt, Am. 8., 1870, 99, 153). Copper is also oxi- dized by many acids. §78. Cadmium. Cd = 112.4 . Valence two. 1. Properties.— Speci/te gravity, liquid, 7.989; copied, 8.67; hammered, 8.6944. Melting point, 320.68° (Callendar and Griffiths, C. N., 1891, 63, 3). Boiling point 763° to 772° (Carnelley and Williams, J. C, 1878,' 33, 284). Specific heat is 0.0567! Yapor density (H = 1), 55.8 (Deville and Troost, A. Gh., 1860, (3), 58, 257). Prom these data the gaseous molecule of cadmium is seen to consist of one atom (Richter, Anorg. Chem., 1893, 363). It is a white crystalline metal, soft, but harder than tin or zinc; more tenacious than tin; malleable and very ductile, can easily be rolled out into foil or drawn into fine wire, but at 80° it is brittle. Upon bending- it gives the same creaking sound as tin. It may be completely distilled in a current of hydrogen above 800°, forming silver white crystals (Kammerer, B., 1874, 7, 1724). Only slightly tarnished by air and water at ordinary temperatures. When ignited burns to CdO . When heated it com- bines directly with CI , Br , I , P , S , Se , and Te . It forms many useful alloys having low melting-points. 2. Occurrence.— Found as greenockite (CdS) in Greenland, Scotland and Penn- sylvania; also to the extent of one to three per cent in many zinc ores. 3. Preparation. — Reduced by carbon and separated from zinc (approximately) hy distillation, the cadmium being more volatile. It may be reduced by fusion with H , CO , or coal gas. 4. Oxide and Hydroxide. — Cadmium forms but one oxide, CdO , either by burning the metal in air or by ignition of the hydroxide, carbonate, nitrate, oxalate, etc. It is a brownish-yellow powder, absorbs C0 2 from the air, becom- ing white (Gmelin-Kraut, 3, 64). The hydroxide -Cd(OH) 2 is formed by the action of the fixed alkalis upon the soluble cadmium salts; it absorbs C0 2 from the air. 5. Solubilities. — a. — Metal. — Cadmium dissolves slowly in hot, moderately dilute hydrochloric or sulphuric acid with evolution of hydrogen; much more readily in nitric acid with generation of nitrogen oxides. It is soluble in ammonium nitrate without evolution of gas; cadmium nitrate and ammonium nitrite are formed (Morin, C. r., 1885, 100, 1497). 6. — The oxide and hydroxide are insoluble in water and the fixed alkalis, soluble in ammonium hydroxide, readily soluble in acids forming salts; soluble in a cold mixture of fixed alkali and alkali tartrate, reprecipitated upon boiling (distinction from copper) (Behal, /. Pharm., 1885, (5), 11, 553). c— Salts.— The sulphide, carbonate, oxalate, phosphate, cyanide, ferrocyanide and ferricyanide are insoluble (§27) in water, soluble in hydrochloric and nitric acids, and soluble in NH,OH , except CdS . The chloride and bromide are deliquescent, the iodide is perma- nent; they are soluble in water and alcohol. 6. Reactions, a. — The fixed alkali hydroxides — in absence of tartaric a-nd citric acids, and certain other organic substances — precipitate, from solutions of cadmium salts, cadmium hydroxide, Cd(0H) 2 , white, insoluble §78, 6i. CADMIUM. Ill in excess of the reagents (distinction from tin and zinc). Ammonium hydroxide forms the same precipitate which dissolves in excess. If the concentrated cadmium salts be dissolved in excess of ammonium hydroxide with gentle heat and the solution then cooled, crystals of the salt, with variable amounts of ammonia, are obtained; e. g., CdCl 2 (NH 3 ) 3 , CdS0 4 (NH 3 ) 4 , Cd(N0 3 ) 2 (NH 3 ) 6 (Andre, C. r., 1887, 104, 908 and 987; Kwasnik, Arch. Pharm., 1891, 229, 569). The fixed alkali carbonates pre- cipitate cadmium carbonate, CdC0 3 , white, insoluble in excess of the reagent, ammonium carbonate forms the same precipitate dissolving in excess. Barium carbonate, in the cold, completely precipitates cadmium salts as the carbonate. 6. — Oxalic acid and oxalates precipitate cadmium oxalate, white, soluble in mineral acids and ammonium hydroxide. Potassium cyanide precipitates cadmium cyanide, white, soluble in excess of the reagent as Cd(CM')4.2KClT; ferrocyanides form a white precipitate; ferricyanides a yellow precipitate, both soluble in hydrochloric acid, and in ammonium hydroxide. Potassium sulphocyanate does not precipitate cadmium salts (distinction from copper). Cadmium salts in presence of tartaric acid are not precipitated by fixed alkali hydroxides in the cold; on boiling-, cadmium oxide is precipitated (separation from copper and zinc) (Aubel and Ramdohr, A. Ch., 1858, (3), 52, 109). c. — Nitric acid dissolves all the known compounds of cadmium, d. — Soluble phosphates precipitate cadmium phosphate, white, readily soluble in acids. Sodium pyrophosphate precipitates cadmium salts, soluble in excess and in mineral acids, not in dilute acetic. The reaction is not hindered by the pres- ence of tartrates or of thiosulphates (separation from Cu) (Vortmann, B., 1888, 21, 1104). , e. — Hydrogen sulphide and soluble sulphides precipitate, from solutions neutral, alkaline, or not too strongly acid, cadmium sulphide, yellow; insoluble in excess of the precipitant (Fresenius, Z., 1881, 20, 236), in ammonium hydroxide, or in cyanides (distinction from copper) ; soluble in hot dilute sulphuric acid and in a saturated solution of sodium chloride * (distinction from copper) (Cushman, Am., 1896, 17, 379). Sodium thiosulphate, Na 2 S 2 3 , does not precipitate solutions of cadmium salts (Follenius, Z., 1874, 13, 438), but in excess of this reagent, ammonium, salts being absent, sodium carbonate completely precipitates the cadmium as carbonate (distinction from copper) (Wells, 0. N., 1891, 64, 294). Cadmium salts with excess of sodium thiosulphate are not precipitated upon boiling with hydrochloric acid (distinction from copper) (Orlowski, /. C, 1882, 42, 1232). /. — The non-precipitation by iodides is a distinction from copper, g. — Soluble arsenites and arsenates precipitate the corresponding cadmium salts, readily soluble in acids and in ammonium hydroxide, h. — Alkali chromates precipitate yellow cadmium chromate from concentrated solutions only, and soluble or* addition of water. • i. — A solution of copper and cadmium salts, very dilute, when allowed to spread upon a filter paper or porous porcelain plate, gives a ring of the cad- mium salt beyond that of the copper salt, easily detected by hydrogen sulphide (Bagley, J. C, 1878, 33, 304). * Owing to the formation of incompletely-dissociated CdCl 2 . Cdl, is still less dissociated and accordingly CdS dissolves more readily in HI than in HC1 and much more readily than in HNO, of the same concentration. On the other hand, of course, precipitation of the sulphide -takes place with more difficulty from the iodide than from the other salts. 112 REACTIONS OF BISMVTE, COPPER AND CADMIUM. §78,7. 7. Ignition. — On charcoal, with sodium carbonate, cadmium salts are reduced before the blow-pipe to the metal, and usually vaporized and reoxidized nearly as fast as reduced, thereby forming a characteristic brown incrustation (CdO). This is volatile by reduction only, being driven with the reducing flame. Cad- mium oxide colors the borax bead yellowish while hot, colorless when cold; microcosmic salt, the same. If fused with a bead of K 2 S, a yellow precipitate of CdS is obtained (distinction from zinc) (Chapman, J. C, 1877, 31, 490). 8. Detection. — Cadmium is precipitated from its solutions by H 2 S form- ing CdS. By its insolubility in (NH^S^. and solubility in hot dilute HNOj it is separated with Pb , Bi , and Cu from the remaining metals of the second group. Dilute H 2 S0 4 with C 2 H 5 0H removes the lead and NH 4 0H precipitates the bismuth as Bi(0H) 3 , leaving the Cu and Cd in solution. If copper be present, KCN is added until the solution becomes colorless, when the Cd is detected by the formation of the yellow CdS with H 2 S . If Cu be absent the yellow CdS is obtained at once from the ammoniacal solution with H 2 S . See also 6i. 9. Estimation. — (1) It is converted into, and after ignition weighed as an oxide. (2) Converted into, and after drying at 100°, weighed as CdS. (3) Pre- cipitated as CdC 2 4 and titrated by XMn0 4 . (4) Electrolytically from a slightly ammoniacal solution of the sulphate or from the oxalate rendered acid with oxalic acid. (5) Separated from copper by XI; the I removed by heating; the excess of XI removed by XN0 3 and H..SO,; the cadmium precipitated by Na 2 CO s and ignited to CdO (Browning, Am. S., 1893, 146, 280). (6) By adding a slight excess of H 2 S0 4 to the oxide or salt, and evaporation first on the water bath and then on the sand bath, weighed as CdS0 4 (Follenius, Z., 1874, 13, 277). 10. Oxidation. — Metallic cadmium precipitates the free metals from solutions of Au , Pt , Ag , Hg , Bi , Cu , Pb , Sn , and Co ; and is itself reduced by Zn , Mg , and Al . §79. Comparison of Certain Reactions of Bismuth, Copper, and Cadmium. Taken in Solutions of their Chlorides, Nitrates, Sulphates, or Acetates. Bi Cu Cd XOH or NaOH, in Bi(OH)„, white. Bi(OH) 3 , white. BiOCl, white (§76, Sd). Partial precipita- tion in solutions not very strongly acid (§76, 6f). Bi,S 3 , black, in- soluble in XCN. Bi, spongy precipi- tate. Bi, black. Bi, black. Cu(OH) 2 , dark blue. Blue solution. Cd(OH) 2 , white. NH,OH, in excess Dilution of satu- Colorless solution.. Precipitation of Cul, with libera- tion of iodine (§77, en. Cu 2 S and CuS, black, soluble in XCN. Cu, bright coating (§77, 10). Cu 2 0, yellow (§77, 56). Cu, precipitated metal. Glucose, XOH, and CdS, yellow, insol- uble in XCN. Cd, gray sponge with zinc, no ac- tion with iron. X 2 Sn0 2 + XOH.. — *■ §81, 4- PRECIPITATION OF METALS OF SECOND GROUP. 113 Systematic Analysis of the Metals of the Tin and Copper Group. The precipitation of the metals of the second group (Tin and Copper Group) by hydrosulphuric acid, and their separation into Division A (Tin Group) and Division B (Copper Group). See §312. §80. Manipulation. — The nitrate from Group I. (§62), or the original solution, if the metals of the silver group be absent, is rendered acid with a few drops of HC1 , warmed and saturated with hydrosulphuric acid gas. 2H 3 As0 4 + xHCl + 5H 2 S = As 2 S 5 + xHCl + 8H 2 or 2H 3 As0 4 + xHCl + 5H,S ==As 2 S 3 + xHCl + S 2 + 8H 2 SnCl 4 + 2H 2 S = SnS 2 + 4HC1 SnCl 2 + H 2 S = SnS + 2HC1 2Bi(NO s ) 3 + 3H 2 S = Bi 2 S 3 + 6HNO s CdS0 4 + H 2 S = CdS + H 2 S0 4 The precipitate, after being allowed to settle a few minutes, is filtered and thoroughly washed with hot water containing a little HC1 . A portion of the filtrate diluted with water is again tested with H 2 S to insure complete precipitation (§81, 2), and if necessary the whole of the filtrate is diluted and again precipitated. The filtrate containing no metals of the second group is set aside to be tested for the remaining metals (§128). §81. Notes. — 1. Hydrosulphuric acid gas should be used in precipitating the metals of the second group. It should be generated in a Kipp apparatus, using ferrous sulphide, ITeS , and dilute commercial sulphuric acid (1-12). Commercial hydrochloric acid may be used instead of sulphuric. The gas should be passed through a wash bottle containing water to remove any acid that may be carried over mechanically. It should always be conducted through a capillary tube into the solution to be analyzed. Less gas is required and the solution is less liable to be thrown from the test tube by the excess of unab- sorbed gas. S. In testing the nitrate for complete precipitation, instead of the gas, a cold saturated water solution of the gas may well be employed. This dilutes the solution at the same time. In treating the unknown solution with H 2 S or in making a saturated water solution of the gas, it should be passed into the liquid until, upon shaking the test tube or bottle capped with the thumb, there is no formation of a partial vacuum due to the further absorption of the gas by "the liquid. 3. H 2 S is decomposed by HN0 3 or HN0 3 + HC1 (nitrohydrochloric acid) (§257, 6B), hence these acids must not be present in excess. If these acids were used in preparing the solutions for analysis, they must be removed by evaporation. Sulphuric acidulation is not objectionable to precipitation with H 2 S , but could not be used until absence of the metals of the calcium group (Group V.) had been assured. Ji. The precipitation of the silver group has left the solution acid with HC1 and prepares the solution for precipitation with H 2 S , if other acids are not present in excess and if too much HC1 was not employed. The presence of a great excess of HC1 does not prevent the precipitation of arsenic (§69, 6e), but does hinder or entirely prevent the precipitation of the other metals of this group, especially tin, lead (§57, 6e), cadmium and bismuth. The solution must be acid or traces of Co , Ni and Zn (§135, 6e) will be precipitated. No instruc- tions can be given as to the exact amount of HC1 to be employed. About one part of HC1 to 25 of the solution should be present to prevent the precipitation 114 PRECIPITATION OF METALS OF SECOND GROUP. §81, 5. of Zn , and it is seldom advisable to use more than one part of HC1 to ten of the solution * (this refers to the reagent HC1 , §334). 5. The precipitation takes place better from the warm solutions than from the cold (§31); hence it is directed to warm the solution before passing in the H 2 S , and before filtering heat again nearly to boiling. If arsenic be present,, the solution should be kept at nearly the boiling point, and the gas passed into the solution for several minutes (§69, 6e). 6. The precipitated sulphides of the metals of the tin and copper group- second group) present a variety of colors, which aid materially in the further analysis of the group. CdS , SnS 2 , As^Sj and As 2 S 5 are lemon-yellow; Sb 2 S s and Sb 2 S„ are orange; SnS , HgS , PbS , Bi 2 S 8 , Cu 2 S and CuS are black to- brownish-black. If too much HC1 be present, lead salts frequently precipitate a red double salt of lead chloride and lead sulphide (§57, 6e). Mercuric chloride at first forms a white precipitate of HgCl 2 .2HgS , changing from yellow to red r and finally to black with more H 2 S , due to the gradual conversion to HgS (§58, 6e). 7. Addition of water to the solution before passing in H 2 S may cause the precipitation of the oxychlorides of Sb , Sn or Bi (5d; §70, §71 and §76). These should not be redissolved by the addition of more HC1 , as they are readily- transposed to the corresponding sulphides by H 2 S , and the excess of acid necessary to their resolution may prevent the precipitation of cadmium or cause the formation of the red precipitate with lead chloride. S. Arsenic when present as arsenic acid is precipitated exceedingly slowly from its cold solutions, and tardily even from the hot solutions. Frequently the other metals of the group may be completely precipitated and removed by filtration, when a further treatment with H 2 S causes a precipitation of the arsenic as As 2 S 5 from the hot solution. This slow formation of a yellow pre- cipitate is often a very sure indication of the presence of pentad arsenic (§69,. 6'e,i). °. The presence of a strong oxidizing agent as H1T0, , K 2 Cr 2 7 , 3?eCl s , etc.,. causes with H 2 S the formation of a white precipitate of sulphur (§125, 6e),. which is often mistaken as indicating the presence of a second group metal. If the original solution be dark colored, it is advisable to warm with hydro- chloric acid and alcohol (§125, 6f and 10) to effect reduction of a possible higher oxidized form of Cr or Mn before the precipitation with H 2 S , thus avoiding the unnecessary precipitation of sulphur. 10. Complete precipitation of the metals of the second group with H 2 S may fail: (1) from incomplete saturation with the gas (§81, 2); (2) from the pres- ence of too much HC1 (§81, 4) ; (3) from the presence of much pentad arsenic (§69, 6«). The first cause of error may be avoided by careful observance of the directions in note (2). To prevent the second cause of error a. portion of the filtrate, after the removal of the precipitate by filtration, should be largely diluted with water (10 volumes) and H 2 S (gas or saturated water solution) again added. In case a further precipitate is obtained, the whole of the filtrate should be diluted and again precipitated with H 2 S . This should be repeated until the absence of second group metals is assured. If a slow formation of a yellow precipitate indicating Asv is observed, H 2 S should be passed into the * Addition of a strong acid, containing H ions in large quantity, diminishes the already slight dissociation of the H 2 s (§44), thus decreasing in number the S ions, whose concentration multi- plied by that of the metal ions must equal the solubility-product of the sulphide in question, before precipitation can take place. Precipitation of some of the sulphides of the Tin and Copper Group may be entirely prevented in this way. It frequently happens that addition of water alone will cause precipitation of these sulphides from a strongly acid solution which has been saturated with H 2 S. This appears strange in view of the fact that the acid which prevented precipitation and the acid which finally produced it were both diluted by the added water in the same proportion. But as a matter of fact dilution does not have the same effect on a strong acid as on a weak one. Dissociation is always in- creased by dilution, but in much greater ratio in the case of a weakly-dissociated body as H 2 S than where the dissociation of the substance is already practically complete, as in the case of the strong acid. Dilution in the case mentioned increases the relative concentration of the » ions and so the solubility-product is reached and precipitation results. §83, 4. PRECIPITATION OF METALS OF SECOND GROUP. 115- hot solution for fully 30 minutes (Note 5) or the solution should be treated with S0 2 or some other agent for the reduction of Asv to As'" (§69, 10). §82. Manipulation.— After the precipitate has been well washed with hot water the point of the filter is pierced with a small stirring rod and the precipitate washed into a test-tube, using as small an amount of water as possible. Yellow ammonium sulphide (NH 4 ) 2 S X (§83, 2) is then added and the precipitate digested for several minutes with warming : As 2 S s + 2(NH,),S, = (NH 4 ) 4 As 2 S B + S 2 SnS + (NH 4 ) 2 S 2 = (NH 4 ) 2 SnS 3 2SnS 2 + 2(NH 4 ) 2 S 2 = 2(NH 4 ) 2 SnS 3 + S 2 2Sb 2 S 3 + 6(NH 4 ) 2 S 2 = 4(lTH 4 ) 3 SbS 4 + S 2 2MoS„ + 2(NH 4 ) 2 S 2 = 2(NH 4 ) 2 MoS 4 + S 2 The precipitate is then filtered and washed once or twice with a small amount of (NH 4 ) 2 S X , and then with hot water. The filtrate consisting of solutions of the sulphides of As, Sb, Sn, An, Pt, Mo (Gr, Ir, Se, Te,. W, V), constitutes the Tin Group (Division A of the second group). The precipitate remaining upon the filter, consisting of the sulphides of Hg; Pb, Bi, Cu, Cd (Os, Pd, Rh, and Ru), constitutes the Copper Group (Division B of the second group, §95). §83. Notes. — 1. The precipitate of the sulphides of the tin and copper group- must be thoroughly washed with hot water (preferably containing" H 2 S and about one per cent of reagent HC1 to prevent the formation of soluble colloidal sulphides (§69, 5c), to insure the removal of the metals of the iron and zinc groups, which would be precipitated on the addition of the ammonium sulphide (§144). 2. Yellow ammonium sulphide, (NH 4 ) 2 S X , forms upon allowing the normal sulphide, (NH 4 ) 2 S , to stand for sometime, or it may be prepared for imme- diate use by adding sulphur to the freshly prepared normal sulphide (§257, 4).. For arsenic sulphides the normal ammonium sulphide may be employed, but the sulphides of antimony are soluble with difficulty, and stannous sulphide is scarcely at all soluble in that reagent; while they are all readily soluble in the yellow polysulphide (6e; §69, §70 and §71). 3. Cupric sulphide, CuS , is sparingly soluble in the yellow ammonium sul- phide and will give a grayish-black precipitate upon acidulation with HC1 . The sulphides of the tin group are soluble in the fixed alkali sulphides, K 2 S and Na 2 S; cupric sulphide is insoluble in these sulphides. Mercuric sulphide, however, is much more soluble in fixed alkali sulphides than cupric sulphide is in the (NH 4 ) 2 S X . If copper be present and mercury be absent, it is recom- mended to use K 2 S or Na 2 S instead of (NH 4 ) 2 S X for the separation of the second group of sulphides into divisions A (tin group) and B (copper group). But if Hg" be present, the (NH 4 ) 2 Sjc should be used, and the presence or absence of traces of copper be determined from a portion of the filtrate from the silver group before the addition of H 2 S (§103). 4. The sulphides dissolve more readily in the (HH 4 ) 2 S X when the solution is; warmed. An excess of the reagent is to be avoided, as the acidulation of the solution causes the precipitation of sulphur (§256, 3), which may obscure the. precipitates of the sulphides present. 116 TABLE FOR THE ANALYSIS OF THE TIN GROVP. m. * A °a a += cd ^ o ■5 ojiS t^ Sis - •2 2 3 a.g - a « a & - ■3 P-i - >» » So; ■+j .a £ "* g s « -9 £n3 £ ^ a; .3 s «r K -C-s +» _ -rt co " ft 4> o> .!? "d - H a p 4) T3 fl-a co ,d « -dp? 03 ,s >oo -d i p £ -d o 3 ." h ^ " o ,d C3 M ca ca h co o O o ft) £ a> o - ca Ti ft » o J - o o 4) co 4> hn 4> ■ CO " .M CO ft o h O S3 S > — - ft P O K a d CD ,sa -t- ° f-* CO *3 is ^rd u %2 i— I w CO -d . PI 05 C3 CO 3 o >) M o a o pi ^ ai ca k^ li S n u ** o 2 £ > § "g "S -d 8 "3 ca Ph ■d 0) 10 o ft a o w SQ PI CO ^ 3J5 ° fc s ■d >,S O ..W pl« a bo o CO ^3 ^3 K. P rf3 P pQ t3 /-s .13 ■" 9 > « b* ca p u>s 'H ij cow co S a) S« ■ cacn,a S 5 bfi 09 ft •H o CD U Ph J< ° " o ^ 33 g f-i o > ca o 3 -^ P-d bo ca eo o -f-»0 q d cat- o CO 'ft *" 3 hB4J T3 ft ^ .Up R d^^s "so a) R +3 oj h — ■» -* ft -o d"d g-53 > s ° R co.S o w =M CD ti 2 ".SO -g^ g,^ gw „ d d 75 'P a S g Sd 5- 01 S P o d So 13 Pti Pi 5< » o -^ B h P """OS mi B -d -^ 0) g 2--S ^3 P fi •w B ca w ^ a g - g < d fill h d ^"2 CJ) 3 +J r. g Mft^i d ? U 3 Bfi )53 O Ph P<0 P4 < CD rt co .^ ,|H "5 ^ v d d a 2 tc § « S.3 - ^ _ S3 a §i la « »g«^ ftg-H-d § gd d^ peg S o ca - S P _ M t3 4) "o ."K B x| A o TABLE FOB THE ANALYSIS OF, THE TIN GROUP. 117 S'i 3 * o *> A s» A' «H O ^ ^-^ CD O K ^ •° S « o sT.S « ft lg|.s 3 © a .9 ® o d ^ 2 >■ o a , o« b, O [j O Is .9 ° u o 3 -i ■gS 3 « £« 118 DIRECTIONS FOR ANALYSIS WITH NOTES. §85. §85. Manipulation. — The solution of the sulphides in (NHJ-jS* is care- fully acidulated with hydrochloric acid: 2(BTH 4 ) 2 S 2 + 4HC1 = 4NH 4 C1 + S 2 + 2H 2 S (lTH 4 ) 4 As 2 S 5 + 4HC1 = As 2 S 2 + 4NH 4 C1 + 2H 2 S 2(NH 4 ) a SbS 4 + 6HC1 = Sb 2 S 5 + 6NH 4 C1 + 3H 2 S (NH 4 ) 2 SnS s + 2HC1 = SnS 2 + 2NH 4 C1 + H 2 S The precipitate obtained when the metals of the tin group are present, is usually yellow or orange-yellow and is easily distinguished from a pre- cipitate of sulphur alone (SnS and MoS 3 are brownish-black). It should be well, washed with hot water and then dissolved in hot HC1 using small fragments of ZC10 3 (§69, 6e) to aid in the solution: 2As 2 S„ -I- 10C1 2 + 16H 2 = 4H 3 As0 4 + 20HC1 + 3S 2 SnS 2 + 4HC1 = SnCl 4 + 2H,S PtS 2 -f £C1 2 = PtCl 4 + S 2 The solution h boiled (to insure removal of the chlorine (§69, 10) until it no longer bleaches litmus paper. §86. Notes. — 1. If the precipitate obtained is white, it probably consists of sulphur alone and indicates absence of more than traces of the metals belong- ing to this group (GeS 2 is white, §111, 6). 2. Care should be taken not to use too much HC1 in precipitating the sul- phides from the (NH 4 ) 2 S X solution, as some of the sulphides (especially SnS 2 ) are quite soluble in concentrated HC1 . 3. It will be noticed (§85) that the lower sulphides of Sb and Sn are oxidized by the (NH 4 ) 2 S X , and are precipitated by the HC1 as the higher sulphides Sb 2 S 5 and SnS 2 respectively. This fact may be most readily observed by the precipitation of a solution of SnCl 2 with H 2 S , giving a brown precipitate of SnS , then dissolving this precipitate in (NH 4 ) 2 S X and reprecipitating with HC1 as the orange-colored SnS 2 . 4. Hot reagent HC1 (§334) dissolves the sulphides of tin quite readily without reduction; the sulphides of antimony, slowly forming SbCl 3 only; and the sulphides of arsenic practically not at all, or at most only traces. The , sulphides of Au and Pt are not soluble in HC1 . MoS 3 is soluble in hot con- centrated HC1 . The relative solubility of these sulphides in HC1 is used by some chemists as the basis of a separation of As from Sb and Sn (§69, 6e, also bottom of next note, 5). 5. The sulphides of arsenic are readily soluble in ammonium carbonate (§69, 5c) and are thus separated from the sulphides of Sb and Sn, which are prac- tically insoluble. The following table suggests a method of analysis based upon this property of these sulphides, §86, 6. DIRECTIONS FOR ANAL7SIS WITE NOTES. Digest with solution of ammonium carbonate and filter. 119 Residue: SnS 2 , Sb 2 S,, , (S) . Dissolve in hot hydrochloric acid (5c, §70 and §71). Solution: SnCl 4 , SbCl 3 . Treat with zinc and hydrochloric acid in Marsh's apparatus (§69, 6'o). Deposit: Sn , (Sh) . Dissolve by hydro- chloric acid. Solution: SnCl 2 . (Residue, Sh .) Test by ammoniacal silver nitrate and by mercuric chlo- ride (§71, 6i and j). Gas: SbH a . (Test the spots, §69, 6'c, 1.) Receive the gas in solution of silver nitrate. Dissolve the precipitate (SbAg 3 ) (§70, 6;'), and test by H 2 S (§87 and Solution: (NH 4 ) 3 AsS 4 + (NH 4 ),As0 1 , and (NH 4 ) 4 As 2 S 5 + (NH 4 ) 4 As 2 5 . Precipitate by hydrochloric acidj filter; wash the precipitate and dissolve it by chlorine gener- ated from a minute fragment of potassium chlorate and a little hydrochloric acid (§69, 5c). Expel all free chlorine (note 9, and §69, 10). Solution: H 3 As0 4 . Apply Marsh's Test, as directed in §69, 6'a, testing the spots (§69, 6'c) ; receiving the gas in solu- tion of silver nitrate, and test- ing the resulting solution (§87). Examine the original solution, as indicated in §88, 1. The plan above given may be varied by separating antimony and tin by ammo- nium carbonate in fully oxidized solution, as follows: The Sb 2 S 5 and SnS 2 are dissolved by nitrohydrochloric acid, to obtain the antimony as pyroantimonic acid. The solution is then treated with excess of ammonium carbonate, in » vessel wide enough to allow the earrbonic acid to escape without waste of the solution. The soluble diammonium dihydrogen pyroantimonate, (NH 4 ) 2 H 2 Sb 2 7 , is formed. Meanwhile the SnCl 4 is fully precipitated as H 2 Sn0 3 (§71, 6a), and may be filtered out from the solution of pyroantimonate. The liability of failure, in this mode of separating antimony and tin, lies in the non-formation of pyroantimonic acid by nitrohydrochloric acid. The ordi- nary antimonic acid forms a less soluble ammonium salt, but this acid is not so likely to occur in obtaining the solution with nitrohydrochloric as anti- monous chloride, SbCl 3 . Excess of ammonium carbonate does not redissolve the Sb 2 3 which it precipitates from SbCl 3 , as stated in §70, 6«. The above plan may also be varied as follows: After removal of the arsenic sulphide with (NHi^CO,, , the residue is dissolved in strong HC1 , not using KC10 3 or HN0 3 . The solution consists of SnCl 4 and SbCl 3 . Divide in two portions: (1) Add Sn on platinum foil. A black precipitate indicates Sb° . (2) Add iron wire, obtaining Sb° and Sn"; filter and test the filtrate for Sn by HgCl 2 (Pieszczek, Arch. Pharm., 3891, 229, 667). 6. The sulphides of As , Sb and Sn are all decomposed by concentrated nitric acid, which furnishes a basis of an excellent separation of the arsenic from the antimony and tin (Vaughan, American Chemist, 1B75, 6, 41). The sulphides reprecipitated from the (NH 4 ) 2 S X solution by HC1 are well washed, transferred to an evaporating dish, heated with concentrated HNO, until brown fumes are no longer evolved, and then evaporated to dryness, using sufficient heat to expel the HN0 3 and the H 2 S0 4 formed by the action of the HNO, upon the S . The heating should be done on the sand bath. The cooled residue is digested for a few minutes with hot water, the arsenic passing into solution as H 3 As0 4 , and the antimony and tin remaining as residue of Sb z 5 and Sn0 2 . The pres- ence of arsenic may be confirmed by the reactions with AgN0 3 (§69, 6/), CuS0 4 '"i, 6fc) by the Marsh test (§69, 6'a), or by precipitation with magnesia mix- 120 DIRECTIONS FOB. ANALYSIS WITH NOTES. §86, 7. ture (§69, 6i). A portion of the residue may be tested in the Marsh apparatus for the Sb (§70, 6;), another portion may be reduced and dissolved in an open dish with Zn and HC1 (not allowable if As be present, §71, 10), and the result- ing SnCl 2 identified by the reaction with HgCl 2 (§71, 6i). 7. The precipitated sulphides must be thoroughly washed to insure the removal of the ammonium salts, since in their presence the dangerously ex- plosive nitrogen chloride (§268, 1) could be formed when the sulphides were dissolved in HC1 with the aid of KC10 3 . 8. Instead of chlorine (HC1 + KC10 3 ), nitrohydroehloric acid may be em- ployed, but it is liable to cause the formation of a white precipitate of Sb 2 5 and Sn0 2 . 9. The chlorine should all be removed, as the metals cannot be reduced by the Zn and H.SO, in the Marsh apparatus in the presence of powerful oxidizing agents as CI . This would also require evaporation to expel the HN0 3 , if nitrohydroehloric acid were used to effect solution. iO. Hydrogen peroxide, H.O, , decomposes the sulphides of arsenic and anti- mony with oxidation. The arsenic will appear in the solution, the antimony remaining as a white precipitate of the oxide (a sharp separation) (Luzzato, Arch. Pharm., 1886, 224, 772). §87. Manipulation. — The solution of the metals of the tin group is "then ready to be transferred to the Marsh apparatus (the directions for the use of the Marsh apparatus are given under arsenic (§69, 6'a), and, should be carefully studied and observed. They will not be repeated here). Only a portion of the solution should be used in the Marsh appar- atus, the remainder being reserved for other tests. The gas evolved from the Marsh apparatus is passed into a solution of silver nitrate, which by its oxidizing action effects a good separation between the arsenic and ■antimony (§89, 2) : AsH-3 + 6AgNO s + 3H 2 = H 3 AsO s + 6Ag + 6HN0 3 SbH 3 + 3AgN0 3 = SbAg 3 + 3HN0 3 The hard glass tube of the Marsh apparatus is heated while the gas is being generated, a mirror of arsenic and antimony being deposited, due to the decomposition of the gases (§69, 6'c) : 2SbH 3 = 2Sb + 3H 2 . The ignited gas is brought in contact with a cold porcelain surface for the production of the arsenic and antimony spots (§69, 6'6). Failure to obtain mirror, spots, or a black precipitate in the AgN0 s is proof of the absence of both arsenic and antimony. The black precipitate obtained in the silver nitrate solution is separated by filtration, washed and reserved to be tested for antimony. The filtrate is treated with HC1, or a metallic chloride, as CaCl 2 or NaCl , to remove the excess of silver and, after evapor- ation to a small volume, is precipitated with H 2 S . A lemon-yellow pre- cipitate indicates arsenic. The black precipitate from the silver nitrate solution is dissolved in hot reagent HC1 : SbAg 3 -j- 6HC1 = SbCl 3 + 3AgCl . The excess of acid is removed by evaporation, a little water is added (§70, 5d and §59, 5c) and the AgCl removed by filtration. The filtrate is divided into two portions. To one portion H 2 S is added; an orange precipitate indicates antimony. The H 2 S may give a black precipi- tate of Ag 2 S from the AgCl held in solution by the HC1 . If this be the §89,4. DIRECTIONS FOB ANALYSIS WITH NOTES. 121 case, to the other portion one or two drops of ( KI are added and the solution filtered. This filtrate is now tested for the orange precipitate with H 2 S . The mirror ohtained in the hard glass tube should he examined as directed in the text, especially by oxidation and microscopic examination (§69, 6'c 5). The spots should he tested with NaCIO and by the other tests as given in the text (§69, 6'c 1). §88. Notes. — Arsenic. — 1. All compounds of arsenic are reduced to arsine by the Zn and H 2 SO« in the Marsh apparatus. Hence if strong oxidizing agents are absent, the original solution or powder may be used directly in the Marsh apparatus for the detection of arsenic; but sulphides should not be present. 2. The burning arsine forms As 2 3 , which may be collected as a heavy white powder on a piece of black paper placed under the flame. Antimony will also deposit a similar heavy white powder. 3. The arsine evolved is not decomposed (faint traces decomposed) upon passing through a drying tube containing soda lime or through a solution of KOH (distinction and separation from antimony). 4. Arsenites and arsenates are distinguished from each other by the following reactions: (a) Arsenous acid solution acidulated with HC1 is precipitated in the cold instantly by H 2 S; arsenic acid under similar conditions is precipitated exceedingly slowly (§69, 6e). (6) Neutral solutions of arsenites give a yellow precipitate with AgN0 3 ; neutral solutions of arsenates give a brick-red pre- cipitate. Both precipitates are soluble in acids or in ammonium hydroxide (§59, Sg). (e) Magnesia mixture precipitates arsenic acid as white magnesium ammonium arsenate, Mg2TH 4 As0 4 ; no precipitate with arsenous acid (§189, 6g). (d) HI gives free iodine with arsenic acid; not with arsenous acid (§69, 6f). (e) Alkaline solutions of arsenous acid are immediately oxidized to the pentad arsenic compounds by iodine (§69, 10). (f) Potassium permanganate is imme- diately decolored by solutions of arsenous acid or arsenites; no reaction with arsenates (§69, 10). §89. Notes. — Antimony. — 1. If antimony be present in considerable amount, it (in the form of the sulphide) is most readily separated from arsenic by boiling with strong HC1 (solution of the antimony sulphide, (§70, 6e)); or by digesting with (NH 1 ) 2 CO B or NH 4 OH (solution of the arsenic (§69, 5c)). 2. For the detection of traces of antimony, the most certain test is in its volatilization as stibine in the Marsh apparatus and precipitation as SbAg 3 , antimony argentide, with AgN0 3 ; this is a good separation from arsenic -and ' tin, and after filtration it remains to dissolve the SbAg 3 in concentrated HC1 and identify the Sb as the orange precipitate of Sb 2 S 3 . The formation of the black precipitate in the AgN0 3 solution must not be taken as evidence of the presence of antimony, as arsine gives a black precipitate of metallic silver with AglTOj . A trace of antimony may be found in the filtrate from the SbAg 8 , hence a slight yellow-orange precipitate from this solution must not be taken as evidence of arsenic without further examination (§69, 7). 5. Sb 2 S 3 is precipitated from solutions quite strongly acid with HC1 , i. e., in the presence of equal parts of the concentrated acid (sp. gr. 1.20). Tin is not. precipitated as sulphide if there be present more than one part of the con- centrated acid to three of the solution (§70, 6e). This is a convenient method Of separation. The addition of one volume of concentrated HC1 to two volumes of the solution under examination before passing in the H 2 S will prevent the precipitation of the tin while allowing the complete precipitation of the anti- mony. 4- If the sulphides of As , Sb and Sn are . evaporated to dryness with con- centrated HN0 3 ; the residue strongly fused with Na 2 CO a and NaOH; and the cooled mass disintegrated with cold water, the filtrate will contain the arsenic as sodium arsenate, Na 3 As0 4 , and the tin as sodium stannate, Na 2 SnO„; while the antimony remains as a residue of sodium pyroantimonate, 2Ta 2 H 2 Sb 2 T (§70, 7). 122 DIRECTIONS FOB ANALYSIS WITH NOTES. §89, ,5. 5. Stibine is evolved much more slowly than arsine in the Marsh apparatus, and some metallic antimony will nearly always be found in the flask with the tin (§70, 6;). 6. If organic acids, as tartaric or citric, be present, they should be removed by careful ignition with K,.C0 ; , as preliminary to the preparation of the sub- stance for analysis, since they hinder the complete precipitation of the anti- mony with H ? S (§70, 6e). 7. Antimonic compounds are reduced to the antimonous condition by HI with liberation of iodine (§70, 6f and 10). Chromates oxidize antimonous salts to antimonic salts with formation of green chromic salts (§70, 6ft). KMnO, also oxidizes antimonous salts to antimonic salts, a manganous salt being formed in acid solution (§70, 6ft). No reaction with antimonic salts. Antimonous salts reduce gold chloride; antimonic salts do not (§73, 10). §90. Manipulation. — The contents of the generator of the Marsh appar- atus should be filtered and washed. The filtrate, if colorless, may be rejected (absence of Mo). A colored filtrate, blue to green-brown or black, indicates the probable presence of some of the lower forms of molybdenum. The solution should be evaporated to dryness with an excess of HNO . , which oxidizes the molybdenum to molybdic acid, Mo0 3 . The residue is dissolved in NH 4 0H (the zinc salt present does not interfere) and poured into moderately concentrated nitric or hydrochloric acid (§75, 6d footnote). This solution is tested for molybdenum by Na 2 HP0 4 . The original solu- tion should also be examined for the presence of molybdenum as molybdic acid or molybdate (§75, 6d). The residue from the generator of the Marsh apparatus may contain Sb , Sn , Au , and Pt with an excess of Zn . It should be dissolved as much as possible in HC1 . Sb , Au , and Pt are insoluble (§70, 5a). The Sn passes into solution as SnCl 2 and gives a gray or white precipitate with HgCl 2 , depending on amount of the latter present (§71, 6/) : SnCl 2 + HgCl 2 = SnCl, + Hg SnCl 2 + 2HgCl 2 = 2HgCl + SnCl 4 The presence of Sn" should always be confirmed by its action in fixed alkali solution upon an ammoniacal solution of AgN0 3 , giving Ag° (§71, 6t). Au and Pt may be detected in the residue, but it is preferable to precipi- tate them from a portion of the original solution by boiling with ferrous sulphate (6h, §§73 and 74). Both metals are precipitated. They are then dissolved in nitro-hydrochloric acid and evaporated to dryness with am- monium chloride on the water bath. The residue is treated with alcohol which dissolves the double chloride of gold and ammonium, leaving the platinum double salt as a precipitate, which is changed to 'the metal upon ignition. The alcoholic solution is evaporated, taken up with water and the gold precipitated by treating with FeS0 4 (§73, 6h), by boiling with oxalic acid (§73, 66), or by treating with a mixture of SnCl 2 and SnCl 4 (Cassius' purple) (§73, 6g). If a portion of the original solution, free from HNO., , be boiled with §94,5. DIRECTIONS FOR ANALYSIS WITH NOTES. 123 oxalic acid the gold is completely precipitated as the metal, separation from the platinum which is not precipitated (§74, 66). §91. — Notes. — Molybdenum. — 1. In the regular course of analysis, molyb- denum remains in the flask of the Marsh apparatus as a dark colored solution, the Zn and H 2 S0 4 acting as a redtfcing agent upon the molybdic acid. 2. If the molybdenum be present in solution as molybdic acid or a molybdate, it may be separated in the acid solution from the other metals by phosphoric acid in presence of ammonium salts, forming the ammonium phosphomolyb- date; insoluble in acids, but soluble in ammonium hydroxide (§75, 6(2). 3. In ammoniacal solution of a phosphomolybdate, magnesium salts precipi- tate the phosphoric acid, leaving the molybdenum as ammonium molybdate in solution, which may be evaporated to crystallization (method of recovering ammonium molybdate from the ammonium phosphomolybdate residues). §92. Tin. — 1. Tin requires the presence of much less HC1 to prevent its pre- cipitation by H 2 S than arsenic or antimony (§89, 3). 2. The yellow ammonium sulphide (NH„) 2 S X must be used to effect solution if, tin (Sn") be present, SnS being practically insoluble in the normal am- monium sulphide (§71, 5c). 3. Tin in the stannous condition, dissolved in the fixed alkalis (stannites), Teadily precipitates metallic silver black from solutions of silver salts. An arsenite (hot) or ari antimonite in solution of the fixed alkalis produces the same result, but not if the silver salt be dissolved in a great excess of ammo- nium hydroxide (§70, 6i). This reaction also detects stannous salts in the presence of stannic salts. .{. Tin in the Marsh apparatus is reduced to the metal, and then by solution of the residue in HC1 , forms SnCl, , which may be detected by the reduction of HgCl 2 to HgCl or Hg° (§71, 6/), and by the action in fixed alkali solution upon the strong ammoniacal solution of silver oxide (§71, 6i). 5. If the Zn in the Marsh apparatus is completely dissolved, the Sn must be looked for in the solution, which in this case must not be rejected. The tin Temains as the metal as long as zinc is present (§135, 10). 6. The presence of the tin may be confirmed by its action as a powerful reducing agent (§71, 10). If it be present as Sniv , these tests must be made after reduction in the Marsh apparatus or in an open dish with zinc and HC1 . §93. Gold. — 1. Gold will usually be met with in combination with other metals as alloys, and is separated from most other metals by its insolubility in all :acids except nitrohydrochloric acid. 2. If more than 25 per cent of gold be present in an alloy, as with silver, the other metal is not removed by nitric acid (§73, 5a). Either nitrohydro- chloric acid must be used or the alloy fused with about ten times its weight of silver or lead, and this alloy dissolved in nitric acid when the gold remains behind. 3. If the presence of gold is suspected in the solution, it should be precipi- tated with I?eS0 4 before proceeding with the usual method of analysis. b. If gold be present (in the usual method of analysis) it will remain as a metallic residue in the Marsh apparatus, insoluble in HC1 and may be identi- fied by the reactions for Au° . 5. The reactions of gold chloride with the chlorides of tin forming Cassius' purple (§73, 6g) is one of the most characteristic tests for ,gold. §94. Platinum. — 1. Notes 1 to 4 under gold apply equally well for platinum, -except that it is necessary to toil with FeS0 4 to insure complete precipitation of the platinum. 2. Oxalic acid is the best reagent for the separation of gold from platinum (§73, 66). 3. The most important problems in the analysis of platinum consist in its separation from the other metals of the platinum ores (§74, 3). 124 TABLE FOR ANALYSIS OF THE COPPER GROUP. 395. PQ o 02 t> l-H P 1= o o Q o o H CQ Ph t= O M P5 W Ph P-i o O M Eh Ph O (M pi <1 M o W B <1 EH to 5 j * w CO A ft l-H CO =fl -p i — i -a S CO bfi w * 2 S p « EG p 1 i — I OJ 'A -p *a 03 ,_ -p ,5" 'ft OJ •h -p CJ 03 oj u u *H *£ a"" tp.a -ag o3 - n o 3w o o *■ W a) to 2 ft /— n ft" o > 2 T) ^H o fl On •— N O o '-g S t- & P gin X §'t!0 W o Ph P O CO 0) (DO •p °j;q So** «W p £°* pat O p 1 OJ ft H-~H " "+J OJ * rt -P o a » to p 'p H Wig «H CS U ° S-2 o ™ I -3 -P 41 -P Olgg "** P ^ OJ 13 o CO N n o CO a D n 03 O !*, m s 3 o O .S n at _? "A O 03 S. O o •o o s .2 e o I ■d -3 £^2 1« T3 i-s A •e-S rf ci M « O h ■d-d ■l-H P t» pc* BiA d u *j p<-fl o P u o CD d w * h "-d'3" o . sh "pli 2 ! 3£^o ft ft rt ca ^ a! "-i . ? « S*fU ra ■" p" « -,"a =h ft 9 03 o ft° O -p p nn'S oI P 'fe m OS'S p « a 3 ■-J5-P o h og^ft ">A o ft ffl -p -p ts ft adding 7 color- a traces may be recipi- des ft3" ni.'S„ p o not snip and ? rfD P OJ t > « ^ffl 0£L >% o 0} C- - p J T S ^ O T3 l- ■sW P* M I S • H U ft H •" 2. Pi OJ urj ^ ^ M ft n ^ fa oj Ojfl^ 5 5 o ■H M.S o >a a 03 .g^ o 6 ft 0J EQ I I p S o oj 3.2-h P 03 P „ «H 'O i> o n: m m "S ft-p P Ph ^ oj n -a oM o CO I s 'a o 03 <— ' P OJ ft.3 « is W ^3 i-H OJ O ■S A g ft-S «h"S o oj -p .3 oj . .2 ftfl ej +3 ,Q o p °° o »rt C oj 595. TABLE FOR ANALYSIS OF TEE COPPER GROUP. 13S a T3 C ft-p .2 is > o -d a >d >" » -g S--S-S — CO „ I M:S< " •**12 bniu hfr. rO £ .g ft ft^- u cS u ' fl.2 o -P o .23 £-*eb 5£ g O h -la 1)5* £ o ■* ftft£ 0) bo H m o q <£> .pH f" ft o"g fl^ftg o « m ■B ID B ■ ■-H +j h* <— '<£ 2 -p rS *h «» OOOCDO q coaeoaoao q eooeoaeos •ft n a "h B S* 126 DIRECTIONS FOB ANALYSIS WITH NOTES. §96. §96. Manipulation. — The well washed residue after digesting the pre- cipitated sulphides of the second group (the Tin and Copper Group) in (NH 4 ) 2 S X may contain any of the metals of the Copper Group, and in •addition frequently contains sulphur, formed by the action of the H 2 S upon oxidizing agents: 4FeCl 3 + 2H 2 S = 4FeCl 2 + 4HC1 + S 2 . Pierce the point of the filter with a small stirring rod and, with as little water as possible, wash the precipitate into a test-tube, beaker, or small casserole. Sufficient reagent nitric acid (§324) should be added to make about one part of the acid to two parts of water and the mixture boiled vigor- ously for two or three minutes : * 2Bi 2 S 3 + 16HN0 3 = 4Bi(N0 3 ) 3 + 4NO + 8H 2 + 3S 2 6CdS + 16HN0 3 = 6Cd(N0 3 ) 2 + 4NO + 8H 2 + 3S 2 Mercuric sulphide is unattacked (§58, 6e) and remains as a black pre- cipitate together with some sulphur as a yellow to brown-black precipitate. The precipitate is filtered and washed with a small amount of hot water. The filtrate is set aside to be tested later, and the black residue on the filter is dissolved in nitro-hydrochloric acid : 2HgS + 2C1 2 = 2HgCl 2 + S 2 . This solution is boiled to expel all chlorine and the presence of mercury determined by reduction to HgCl or Hg° by means of SnCl 2 (§58, 6g) : HgCl 2 + SnCl 2 = Hg + SnCl 4 , 2HgCl 2 + SnCl 2 = 2HgCl + SnCi;; or by the deposition of a mercury film on a strip of bright copper wire (§50, 10) : HgCl 2 + Cu = Hg + CuCl 2 . Confirm further by bringing in •contact with iodine in a covered dish: Hg + I 2 = Hgl 2 (Jannaesch, Z. anorg., 1896, 12, 143). The mercury may also be detected by using HH 4 0H and KI as the reverse of the Nessler's test (§207, 61c) (delicate 1 to 31,000) (Klein, Arch. Pharm., 1889, 227, 73). §97. Notes.— 1. The concentration of HN0 3 (1-2) is necessary for the solution of the sulphides of Pta , Bi , Cu and Cd , and may also dissolve traces of HgS . However, the concentrated HNO, (sp. gr., 1.42) dissolves scarcely more than traces of HgS (§58, 6e). Long-continued boiling of HgS with concentrated HNO, changes a portion of the HgS to Hg(N"0 3 ) 2 .HgS , a white precipitate, insoluble in HNO, . %. In the use of nitrohydrochloric acid to dissolve the HgS , the HC1 should be used in excess to insure the decomposition of the nitric acid, which would interfere with the reduction tests with SnCl 2 and Cu° . One part of HN0 3 to three parts HC1 gives about sufficient HC1 to decompose all the HNO, , hence in this reaction -a little more than that proportion of HC1 should be used. 3. A small amount of black residue left after boiling the sulphides with HNO, may consist entirely of sulphur, which can best be determined by burning the residue on a platinum foil and noting the appearance of the flame, the odor, and the disappearance of the residue. The residue of sulphur frequently possesses the property of elasticity (§256, 1). 4. Boiling the sulphides of the copper group with HN0 3 will always oxidize * If preferred the precipitate on the filter may be washed with the boiling hot nitric acid of the above mentioned strength, pouring the same acid back upon the precipitate, reheating eacn time, until no further action takes place. §99,5. DIRECTIONS FOR ANALYSIS WITH NOTES. 127 a trace at least of sulphur to H 2 S0 4 (§256, 6B, 2), which will form PbS0 4 if any- lead be present: S 2 + 4HN0 a = 2H 2 S0 4 + 4N0 3PbS + 8HN0„ = 3PbS0 4 + 4H 2 + 82TO If the boiling be not continued too persistently, the amount of PbS0 4 formed is soluble in the HITO,, present (§57, 5c), and does not at all remain behind with the HgS . 5. If the Sb and Sn are not removed, through an insufficiency of (NH 4 ) 2 S X they will appear as a white precipitate mixed with the black precipitate of HgS , due to the fact that HN0 3 decomposes the sulphides of Sb and Sn , forming the insoluble Sb 2 5 and SnO, : 6Sb 2 S a + 2OHNO3 = 6Sb 2 0„ + 9S 2 + 20110 + 10H 2 O 6. Traces of mercury may be "detected by using a tin-gold voltaic couple. The Hg deposits on the Au , and can be sublimed and identified with iodine vapor. Arsenic gives similar results (Lefort, G. r., 1880, 90, 141). 7. Mercury may quickly be detected from all of its compounds by ignition in a hard glass tube with fusion mixture (Na 2 C0 a + K 2 C0 S ) (§58, 7), and then adding a few drops of HNO s (concentrated) and a small crystal of KI . Upon warming the iodine subljmes and combines with the sublimate of Hg , forming the scarlet red Hgl 2 . As and Sb both give colored compounds with iodine, de- composed by H1T0 8 (Johnstone, G. N., 1889, 59, 221). §98. Manipulation. — To the filtrate containing the nitric acid solution of the sulphides of Pb , Bi , Cu , and Cd , should be added about two cc. of concentrated H 2 S0 4 and the mixture evaporated on a sand bath or over the naked flame in a casserole or evaporating dish until the fumes of H„S0, are given off: Pb(lT0 3 ) 2 + H 2 S0 4 = PbS0 4 + 2H2TO a Cu(N0 8 ) 2 + H 2 S0 4 = CuS0 4 + 2HNO3 About 20 cc. of 50 per cent alcohol should be added to the well cooled mixture and the whole transferred to a small glass beaker. Upon giving the beaker a rotatory motion the heavy precipitate of PbS0 4 will collect in the center of the beaker, and its presence even in very small amounts may be observed. The filtrate from the PbS0 4 should be decanted through a wet filter, and the PbS0 4 in the beaker may be further identified by its transference into the yellow chromate with K 2 Cr0 4 or into the yellow iodide with KI (57, 6/ and *). §99. Notes. — 1. In analysis, if lead was absent in the silver group, it is advantageous to test only a portion of the nitric acid solution with H 2 S0 4 for lead, and if that metal be not present, the above step may be omitted with the remainder of the solution and the student may proceed at once to look for Bi , Cu and Cd . If, however, lead is present, the whole of the solution must be treated with H 2 S0 4 . 2. The nitric acid should be removed by the evaporation, as PbS0 4 is quite appreciably soluble in HNO, (§57, 5c). 3. The H 2 S0 4 should be present in some excess, as PbS0 4 is less soluble in dilute H 2 S0 4 than in pure water (§57, 5c). 4. Alcohol should be present, as it greatly decreases the solubility of PbS0 4 in water or in dilute E 2 S0 4 (§57, 5c, 6e). 5. Too much alcohol must not be added, as sulphates of the other metals present are, also less soluble in alcohol than in water (§77, 5c). These sul- phates, if precipitated by the alcohol, are readily dissolved on dilution with water. 128 DIRECTIONS FOR ANALYSIS WITH NOTES. §99, g. 6. If the (N'H 4 ) 2 S X had not been well removed by washing, ammonium sul- phate would be present at this point, greatly increasing the solubility of PbSO, (§57, 5c). §100. Manipulation.— The filtrate from the PbS0 4 should be boiled to expel the alcohol (or if Pb be absent evaporate the nitric acid solution of division B) and then carefully neutralized with NH 4 0H . An excess of NH 4 0H should be added to dissolve the precipitates of Cu(0H) 2 and Cd(0H) 2 , leaving the Bi(0H) 3 as a white precipitate. The solution should be filtered, the precipitate thoroughly washed, and then treated upon the filter with a hot solution of potassium stannite, K 2 Sn0 2 . A black pre- cipitate is evidence of the presence of Bi (§76, 6*7). §101. Notes.— 1. If the precipitate of the sulphides of the second group was not well washed, the hydroxides of the metals of the iron group (Al , Cr and Fe) may be present at this point. The precipitate of Al(OH) 3 would be white, but would not give a black precipitate with K 2 Sn 4 , thus maintaining the equilibrium for H,PO ( , but disturbing that between solid and dissolved PePO, , which requires a certain concentration of PO, ions. To restore the latter more FeP0 4 dissolves, only to react with the H ions as before, and this process continues until the H ions of the hydrochloric acid are reduced to such small quantity as to be in equilibrium with the P0 4 ions or, if the HC1 is in excess, until the PePO, is entirely dissolved. This process takes place whenever a strong acid dissolves the salt of a weak one. It is analogous to the solution of a base in an acid, forming non-dissociated water. §126, 6e. IRON. - 157' excess of sodium acetate and ferric chloride is added drop by drop, until a red color indicates complete precipitation of the phosphate and forma- tion of ferric acetate. The mixture is then boiled and filtered hot. Evidently another portion of the solution must be tested for iron. All of the phosphoric acid present is thus precipitated and separated from the metals of the remaining groups. • Care should be taken to avoid an excess of the ferric chloride as the ferric phosphate is somewhat soluble- in ferric acetate solution. The alkali hydroxides transpose ferric phos- phate (freshly precipitated), forming ferric hydroxide and alkali phosphate,. The transposition is not complete in the cold. With fixed alkali hydroxide aluminum phosphate is dissolved, thus effecting a separation from chrom- ium and iron. Ferric phosphate warmed with ammonium sulphide forms ferrous sulphide, ammonium phosphate and sulphur: 4FeP0 4 + 6(NH 4 ) 2 S ='4FeS + 4(NH 4 ) 3 P0 4 + S 2 . e. — Hydrosulphuric acid is without action upon ferrous salts in acid or neutral solutions, except a slight precipitate is formed with neutral fer- rous acetate. Alkali sulphides and H 2 S in alkaline mixture, form ferrous- sulphide, FeS, black, insoluble in excess of the reagent, readily soluble in dilute acids with evolution of hydrogen sulphide. The moist precipitate is slowly converted, in the air, to ferrous sulphate and finally to basic ferric sulphate, Fe 2 0(S0 4 ) 2 ■ Ferric salts are reduced to ferrous salts with liberation of sulphur by H 2 S (1), or soluble sulphides, the latter at once reacting to precipitate ferrous sulphide {2) : (1) 4FeCI 3 + 2H 2 S = 4Fe€l 2 + 4HC1 + S 2 (2) 4FeCl 3 + GCBmjS = 4FeS + 12NH 4 C1 + S 2 After the removal of the metals of the .second group by H 2 S , the iron present will always be in the ferrous condition (it will therefore be neces- sary to test the original solution to find the condition of the iron at the- beginning of the analysis). The excess of H 2 S should be removed by boiling and the iron oxidized by carefully adding nitric acid drop by drop and boiling until the solution assumes a pale straw color (6o). If this be- done the iron will be completely precipitated in the third group by the- ammonium hydroxide (6a). ■ Ferrous sulphite is but little soluble in pure water, easily soluble in excess of sulphurous acid, to a colorless solution. The moist salt oxidizes rapidly on exposure to the air (Fordos and Gelis, J. Pharm., 1843, (3), 4, 333). Ferric sulphite is only known as a red solution formed by the action of SO. upon freshly precipitated Fe(OH) 3 , rapidly reduced to the ferrous condition accord- ing to the following' equation: Fe 2 (S0 3 ) 3 = FeS0 3 + FeS 2 O (Gelis, G. C, 1862, 896). Ferrous thiosulphate, FeS 2 3 , is formed, together with some FeS and FeS0 3 , by the action of S0 2 upon Fe° (Fordos and Gelis, I. c). Ferric salts are reduced by sodium thiosulphate to ferrous salts in neutral solutions with formation of sodium tetrathionate: 2FeCl 3 + 2Na 2 S,,0 3 = 2FeCl 2 + 2MaCl + Na 2 S t O„ (Fordos and Gelis, C.r-., 1842, 15, 920); in acid solutions sulphuric acid and sulphur are- formed: 4FeCl 3 + 2M"a 2 S 2 O s + 2H 2 = 4FeCl 2 + 4NaCl + aHJSO, + S 2 (Men- 158 IRON. §126, 6/. schutkin, 78). Ferric iron is precipitated as basic nitrate by the addition of a solution of ammonium sulphate to a solution of iron in HNO, evaporated to dryness and taken up with water (separation from aluminum) (Beilstein and Luther, C. C, 1891, i, 809). /. — Chlorides and bromides of both ferrous and ferric iron are formed hut only ferrous iodide exists. Ferric salts are reduced to ferrous salts by hydriodic acid with liberation of iodine. g. — Soluble arsenites and arsenates precipitate solutions of ferrous and ferric salts, forming 1 the corresponding- arsenites and arsenates. Basic ferric arsenite, 4Fe 2 O 3 .As 3 , s + 5H 2 , is formed when an excess of ferric hydroxide is added to arsenous acid. It is insoluble in acetic acid. It is formed when moist ferric hydroxide is given as an antidote in case of arsenic poisoning (§69, U and 6'e; D., 3, 352). h. Ferrous salts are rapidly oxidized to ferric salts by solutions of chro- mates, the chromium being reduced to the triad condition (9 and 10). With ferric salts potassium chromate forms a reddish-brown precipitate. i. — Zinc oxide precipitates solutions of Fe'" , Al , Cr'" and Cu completely and Pb partially, effecting a separation of these metals from Idi , Co and Mi (Meineke, Z. angew., 1888, 358). 7. Ignition. — The larger number of iron salts are decomposed, as solids, by heat; FeCl 3 vaporizes partly decomposed, at » very little above 100° Igni- tion in the air changes ferrous compounds, and ignition on charcoal or by reducing flame changes ferric compounds to the magnetic oxide, which is attracted to the magnet. Ferrous oxalate ignited in absence of air gives FeO . Ferric oxide ignited in a current of hydrogen gives Fe s 4 from 330° to 440°, FeO irom 500° to 600°, and Fe° above 600° (Moissan, A. Ch., 1880, (5), 21, 199). In the outer flame, the borax bead, when moderately saturated with any compound of iron,, acquires a reddish color while hot, fading and becoming light yellow when cold, or colorless, if feebly saturated. The same bead, held persistently in the reducing flame, becomes colorless unless strongly saturated, when it shows the pale green color of ferrous compounds. The reactions with microcosmic salt are less distinct, but similar. Cobalt, nickel, chromium and ■copper conceal the reaction of iron in the bead. Ferric compounds, heated briefly in a blue borax bead holding a very little cupric oxide, leave the bead blue; ferrous compounds so treated change the blue bead to red — the color of cuprous oxide. 8. Detection. — After removal of the first two groups the iron (now in the ferrous condition) is oxidized by HN0 3 and then precipitated in pres- ence of NH 4 C1 with Al and Cr'" by an excess of NH 4 0H . The Al is re- moved by boiling with excess of KOH . If more than traces of Fe be present it is detected in presence of the Cr(0H) 3 , by dissolving in HC1 and obtaining the blood-red solution with KCNS . In case Cr be present in great excess the Cr(0H) 3 and Fe(0H) 3 are fused on a platinum foil with Na 2 C0 3 and KN0 3 , oxidizing the Cr to a chromate soluble in water. After filtering, the precipitate of Fe 2 3 is dissolved in HC1 and tested with KCNS. The original solution must be tested to determine whether the iron was present in the ferrous or ferric condition. A portion of the original isolation acidified with HC1 gives blood red color with KCNS- if Fe'" is §126, 10. IRON. 159 present, no color for the Fe". Another portion gives a blue precipitate with K 3 Fe(CN) if Fe" is present, only a brown or green color for the Fe'" (66). 9. Estimation. — (1) After oxidation to Fe'" , if necessary, it is precipitated with NHiOH , dried, ignited to a dull-red heat and weighed as Fe 2 3 . (2) By- precipitation with nitroso-/?-naphthol in slightly acid solution (Knorre, B., 1887, 20, 283). Volumetrically: (3) As ferrous iron, by titration with a standard solution of KMn0 4 : 10FeSO 4 + 2KM:n0 4 + SH^SO, = 5Fe 2 (SO„) 3 + K 2 S0 4 + 2MnS0 4 + 8H 2 . (4) By titration with a standard solution of K 2 Cr 2 7 , using a solution of K a 'Fe(Cli!) 6 as an external indicator: 6FeS0 4 + K 2 Cr 2 7 + 7H 2 S0 4 = 3Fe 2 (S0 4 ) 3 + K 2 S0 4 + Cr 2 (S0 4 ) s + 7H 2 . (5) As ferric iron, by titration with a standard solution of Bra 2 S 2 O s , using KCNS as an indicator: 2FeCl 3 + 2Na 2 S 2 3 = 2FeCl 2 + Ua 2 S 4 0„ + 2NaCl . A few drops of a solution of CuS0 4 are added, which seems to hasten the reaction and gives more accurate results; or use excess of the M"a 2 S 2 3 and titrate back with standard iodine (Crafts, J. 0., 1873, 26, 1162). (6) The iron as ferric salt is treated with an excess of a standard SnCl 2 solution, the excess of the SnCl 2 being determined by a standard solution of iodine in potassium iodide: 2FeCl 3 + SnCl 2 = 2FeCl 2 + SnCl 4 . (7) Potas- sium iodide is added to the nearly neutral ferric chloride; the flask is stoppered and warmed to 40°. The iodine set free is titrated by standard Na^Oa (very accurate for small amounts of iron). (8) When present in traces it is determined colorimetrically as Fe(CNS) 3 in etherial solution (Lunge, Z. angew., 1894, 669). 10. Oxidation. — Metallic iron precipitates the free metals from solu- tions of Au , Pt , Ag , Hg , Bi , and Cu (separation from Cd). Solutions of Fe" are changed to Fe'" solutions by treating with solutions of Au , Ag , Cr VI , Mn vn , Mn VI , and H 2 2 . In presence of some dilute acid, such as H 2 S0 4 or H 3 P0 4 by Pb0 2 , Pb 3 4 , Mn 3 4 , Mn0 2 , Mn 2 3 , Co 2 3 , Ni 2 3 . The following acids also oxidize Fe" to Fe'", HN0 2 , HN0 3 , HC10 , HC10 2 , HC10 3 , H 2 S0 4 (if concentrated and hot), HBrO , HBr0 3 HI0 3 , also Br , CI . Br and CI in presence of KOH changes Fe" and Fe'" to K 2 Fe0 4 . Barium ferrate is the most stable of the ferrates ; they are strong oxidizers, acting upon nitrites, tartrates, glycerol, alcohol, etherj, ammonia, etc. (Eosell, J. Am. Soc, 1895, 17, 760). Fe'" is reduced to Fe" by solutions of Sn", Cu', H 3 P0 2 , H 3 P0 3 , H 2 S , H 2 S0 3 , Na 2 S 2 3 , and HI . Also by nascent hydrogen, or by any of the metals which produce hydrogen when treated with acids, including Pb , As , Sb , Sn , Bi , Cu *, Cd , Fe , Al , Co , Ni , Zn , and Mg f. * Carnegie, J. C, 1888, 53, 468. t Warren, C. N„ 1889, 60, 187. 160 ANALYSIS OF THE IRON GROUP. §127. §127. Table for Analysis op the Iron or Third Group (Phosphates, and Oxalates being absent). See §312. To the clear filtrate from the Second Group, in which H 2 S will cause no nre- cipitate (§80), and freed from H 2 S by boiling, add a few drops of Nitric Acid and boil an instant (to oxidize ferrosum*). Immediately arM Ammonium Chloride (§134, 56; §189, 56) and an excess (§135, 6a) of Ammonium Hydroxide (§116). If there is a precipitate, filter and wash. Precipitate: Al(OH) 8 , Cr(OH)„ , Fe(OH)„ . Pierce the point of the filter, and with a little water wash the precipitate into a casserole or evaporating dish; add a few drops of Potassium or Sodium Hydroxide and boil for several minutes. If a residue remains filter and wash. ' Residue: Cr(OH) 3 , Fe(OH) s . Fuse a portion of the residue on a platinum foil with potassium nitrate and sodium carbonate, cool, digest in warm water and filter (§125, 7). Besidue: Fe 2 O a . Dissolve the residue in HC1 and test for iron with potassium thio- cyanate (§126, 66). If the residue after re- moval of the aluminum does not indicate an ex- cess of Cr by its green color, it may be dis- solved in HC1 and test- ed for the blood-red color with KCNS . Iron being found, to de- termine whether it is ferric or ferrous, or bothf, in the original solution, test the latter, after acidulating with hydrochloric acid, with KCNS for ferricum, and with ^Ee(CN), for ferrosum (§126, 66). Solution: Na 2 Cr0 4 , K 2 Cr0 4 (Na.COa) . Acidify with HCH s 2 and precipitate the chro- mium as lead chromate (yellow) with a solu- tion of lead acetate (§57, 6ft). If the original solution contains a chromate it will be yellow (normal chromate), or red (acid chromate), and will give the reactions for chromates with Pb(C 2 H 3 2 ) 2 , BaCl 2 , etc. (§125, 6h). If the chromium is present as a chromic salt, Cr 2 (S0 4 ) s , the solution will have a green or bluish-green color and will give the general reactions as de- scribed at §125, 6. Chromates should be re- duced by boiling with HC1 and CH 5 OH be- fore proceeding with the regular course of analysis (§125, 6f). Study §136, §128, §129, §130 and §131. Study §136, §128, §129, §130, §131. Solution : KA10 2 . Make the solution slight- ly acid with hydro- chloric acid, and then- add ammonium car- bonate. A precipitate is Al(OH)„ . The same result is ob- tained with nearly equal certainty by add- ing an excess of NH 4 C1 to the alkaline solution (§124, 6a; §130). Lead and antimony give similar results if (through carelessness) they have not been, removed (§131, 6). Study §136, §128, §129, §131, 6, and §124, 6. * In the filtrate from the Second Group iron Is necessarily in the ferrous condition (126 6e). + Ferrous salts, which have been kept in the air, are never wholly free from ferric compounds. §129, 8. DIRECTIONS FOR ANALYSIS WITH NOTES. 161 Directions foe the Analysis of the Metals of the Third Group. §128. Manipulation. — Boil the filtrate from the second group (§80) to •expel the H 2 S and then oxidize any ferrous iron that may be present by the addition of a few drops of HN0 3 , continuing the boiling to a clear .straw-colored solution (§126, 6c): 3FeS0 4 + 4HST0 3 = Fe 2 (S0 4 ) 3 + Fe(N0 3 ) 3 + NO + 2H 2 Add to the solution about one-half its volume of NH 4 C1 (5&, §§134 and 189) and warm and then add NH 4 0H in a decided excess (§135, 6a) : MgCl 2 + NH 4 C1 + NH.OH = NH 4 MgCl 3 + NH 4 0H Fe 2 (S0 4 ) 3 + 6NH 4 OH = 2Fe(OH) 3 + 3(HH 4 ) 2 S0 4 ZnS0 4 + 4STH 4 OH = (NH 4 ) 2 Zn0 2 + (NH 4 ) 2 S0 4 + 2H 2 Heat nearly to boiling for a moment, filter, and wash with hot water. Notice that the filtrate has a strong odor of ammonium hydroxide and set aside to be tested for the metals of the succeeding groups (§138). §129. Notes. — (i) If the H 2 S is not all expelled, it becomes oxidized by the HN0 3 with deposition of a milky precipitate of sulphur (§257, 6B), which tends to obscure the reactions following: 6H 2 S + 4HW0 3 = 3S 2 + 4NO + 8H 2 0. Also any H 2 S not decomposed by the HN0 3 would cause a precipitate of the sulphides of the fourth group upon the addition of the NH 4 OH: H 2 S + NiCl 2 + 2lTH 4 OH = NiS + 2NH 4 C1 + 2H 2 . (2) Any iron that may have been present in the original solution in the ferric condition is reduced to the ferrous condition by the H 2 S (§126, 6e) : 4FeCl s -)- 2H 2 S = 4FeCl 2 + S 2 + 4HC1 . The ferrous hydroxide is not com- pletely insoluble in the ammonium salts present (§117), and hence unless the •oxidation with the HNO, be complete, some of the iron will be found in the next group. (3) If considerable iron be present the solution becomes nearly black upon -addition of nitric acid, due to the combination of the nitric oxide with the ferrous iron (§241, 8a). Therefore the boiling, and addition of HN0 3 , a drop ■or two at a time, must be continued until the solution assumes a bright straw color. (4) If nitric acid be added in excess there is danger that Mn will be oxid- ized to the triad or tetrad condition then it is precipitated with iron in the third group (§134, 6a). The careful addition of the nitric acid (avoiding an excess) prevents this oxidation of the manganese. (5) Ammonium hydroxide precipitates a portion of Mn (§134, 60) and Mg (§189, 6a), but these hydroxides are soluble in NH 4 C1 (5c, §§134 and 189); hence if that reagent be added in excess the Mn (§134, 6a) and Mg are not at all precipitated by the NH 4 OH: 2MnCl 2 + 2NH 4 OH = Mn(OH) 2 + (NH 4 ) 2 MnCl 4 Mn(OH) 2 + 4NH 4 C1= (NH 4 ) 2 MnCl 4 + 2NH 4 OH 2MgCl 2 + 2NH 4 OH = Mg(OH) 2 + NH 4 MgCl a + NH 4 C1 Mg(OH) 2 + 3NH 4 C1 = NH 4 MgCl„ + 2NH 4 OH (6) Ammonium chloride lessens the solubility of Al(OH) a in the NH,OH solution and effects an almost quantitative precipitation of that metal (§117). (7) NH 4 0H precipitates solutions of Co , Ni and Zn , but these precipitates are readily soluble in an excess of the NH.OH (§116). ■ To insure the presence of an excess of NE,OH the odor should be noted after shaking the test tube and after the solution has been heated. (S) The precipitates of the hydroxides of Al , Cr and Fe'" filter much more Tapidly if the precipitation takes place from a hot solution (§124, 4 and 6a). 162 DIRECTIONS FOR ANALYSIS WITH NOTES. §129, 9. (9) In the presence of chromium the filtrate from the third group is usually of a slight violet color, due to the solution of a trace of chromium hydroxide in the NH 4 OH (§125, 6a). Boiling the solution to remove excess of ammonia prevents this. (10) A small portion of the nitrate of the second group after the removal of the H 2 S by boiling should be tested for the presence of phosphates by am- monium molybdate (§75, 6d). If phosphates are found to be present, the method of analysis of the succeeding groups must be considerably modified. These modifications are fully discussed under §145 to §153. §130. Manipulation. — The well washed precipitates of Al , Cr , and Fe'" hydroxides are transferred to a small casserole or evaporating dish by piercing the point of the filter and washing the precipitate from the filter with as small an amount of water as possible; and then boiled for a minute or two with an excess of NaOH : Al(OH) 8 + NaOH = NaA10 2 + 2H 2 Cr(OH) 3 + NaOH = NaCr0 2 + 3H 2 (in the cold) NaCr0 2 + 2H 2 = Cr(OH)„ + NaOH (upon boiling) The alkaline liquid is filtered (§131, 1) (the filtrate is reserved to be tested for aluminum), and the remaining precipitate fused on a platinum foil with a mixture of equal parts of KN0 3 and Na 2 C0 3 : 2Cr(0H) 3 + 2KN0 3 + Na 2 C0 3 = K 2 Cr0 4 + Na 2 Cr0 4 + 2N0 + C0 2 + 3H 2 (§125, 7). The fused mass is then dissolved in water, filtered, rendered acid with acetic acid and tested for chromium with Pb(C 2 H 3 2 ) 2 , a yellow precipitate at this point being sufficient evidence of the presence of chromium: Na 2 Cr0 4 + K 2 Cr0 4 + 2Pb(C 2 H 3 2 ) 2 = 2PbCr0 4 + 2NaC 2 H 3 2 + 2KC 2 H 3 2 (§57, 6ft). The residue of the fused mass not soluble in water should be washed with hot water and then dissolved in HC1 : Fe 2 3 + 6HC1 = 2FeCl 3 + 3H 2 , and tested for iron with KCNS : FeCl 3 + 3KCNS = Fe(CNS) 3 -f 3KC1. If iron has been found to be present, the original solution acidulated with HC1 (or a few drops of the filtrate from the first group) should be tested with KCNS for the presence of ferric iron (§126, 6b) and with K 3 Fe(CN) 6 for the dark blue precipitate of Fe 3 (Fe(CN) 6 ) 2 indicating the presence of ferrous iron (§126, 6b): 3FeS0 4 + 2K 3 Fe(CN) 6 = Fe 3 (Fe(CN) 6 ) 2 + 3K 2 S0 4 . The alkaline filtrate obtained after boiling the precipitated hydrox- ides with NaOH , is slightly acidulated with HC1 : KA10 2 + 4HC1 = A1C1 3 + KC1 + 2H 2 , and then precipitated with (NH 4 ) 2 C0 3 , a white gelatinous precipitate being evidence of the presence of aluminum: 2A1C1 3 + 3(NH 4 ) 2 C0 3 + 3H 2 = 2A1(0H) 3 + 6NH 4 C1 + 3C0 2 . Or an excess of NH 4 C1 may be added directly to the alkaline filtrate, giving the white gelatinous precipitate of aluminum oxide-hydroxide: 2KA10 2 -f- 2NH 4 C1 + H 2 = A1 2 0(0H) 4 + 2KC1 + 2NH 3 (§124, 6a). §132, 1. COBALT. 163 §131. Notes. — (1) Chromium hydroxide when precipitated from solutions of pure chromic salts by NaOH is readily soluble in an excess of the cold reagent (§125, 6a); but in presence of ammonium salts or of ferric hydroxide the chromium hydroxide is not completely soluble in a cold solution of the fixed alkali. This prevents the use of the cold fixed alkali as a means of separation of Cr and Al from Fe'" . The student is therefore directed to boil the mixture of these three hydroxides with NaOH , thus precipitating the whole of the chromium and effecting a quantitative separation of Cr and Fe'" from Al . If the alkaline liquid is too concentrated to filter, it must be diluted -with water. (2) Unless the precipitate of the hydroxides is a very dark green, due to the presence of a large amount of chromium, a portion of the precipitate should be dissolved in HC1 and tested with KCNS for the presence of iron. The presence of a moderate amount of chromium does not interfere. (S) In the absence of chromium the presence of more than traces of iron gives a brown color to the ammonium hydroxide precipitate (§126, 6a), alu- minum hydroxide being a white gelatinous precipitate. (It) If the fused mass has a green color, manganese (§134, 7) is evidently present in large quantities and was not completely separated by the ITH 4 C1 and NH 4 0H (§134, 6a). By dissolving the fused mass in water and carefully warming with HC1 , the manganate, K 2 Mn0 4 , may be reduced (a) (§134, 5c) without effecting a reduction of the chromate, which may be precipitated as BaCr0 4 by BaCL after neutralization with 1TH 4 0H . Or the fused mass may be warmed with hydrochloric acid and alcohol, effecting complete reduction (B), and this solution again precipitated with lTH.,OH , which will prevent more than traces of the manganese from being precipitated with the third group hydroxides. If again upon fusion with KNO, and K.C0 3 a green mass is obtained, the operation should be repeated: (a) K 2 Mn0 4 + 8HC1 = MnCl 2 + 2KC1 + 2C1 2 + 4H 2 (6) 2K 2 Cr0 4 + 10HC1 + 3C 2 H O = 2CrCl s + 4KC1 + 3C 2 H 4 + 8H 2 (5) The presence of chromium as chromic salts is usually indicated by the green or bluish-green color of the original solution. Chromium as chromates (red or yellow) should be reduced to chromic salts by boiling with HC1 and" C,H„0 before proceeding with the regular group separations (§125, 66 and f). H 2 S will effect this reduction but gives also a precipitate of sulphur which should be avoided when convenient to do so: 2K 2 Cr 2 7 + 16HC1 + 6H 2 S = 4CrCl 3 + 4KC1 + 3S 2 + 14H 2 . (6) Too much stress cannot be laid upon the necessity for removing all the metals of one group before testing the filtrate for the metals of the next succeeding group. If through lack of sufficient H 2 S or too much HC1 , lead or antimony are not completely removed in the second group, they will give all the reactions for aluminum (§57, 6a, and §70, 6a) ; hence as a safeguard it is advised to test the white precipitate, indicating aluminum, with H 2 S . A black or orange precipitate is evidence of unsatisfactory- work and the student should repeat his analysis. (7) The presence of a trace of white precipitate in the final test for aluminum may be due to the presence of that metal in the fixed alkali (§124, 6a, footnote), or it may be caused by the use of too concentrated fixed alkali, which may dissolve silica from the glass of the test tubes or remove it from the filter paper (§249, 5). The Zixc Group (Fourth Group). Cobalt, Nickel, Manganese, and Zinc. §132. Cobalt. Co = 59.00 . Usual valence two and three. 1. Properties. — Specific gravity, powder from the oxide reduced by hydrogen, mean of five samples, 8.957 (Eammelsberg, Pogg., 1849, 78, 93); melting point, 1500° (Pictet, 0. r., 1879, 88, 1317). Cobalt is similar to iron in appearance, is 164 COBALT. §132, 2. harder than Fe or Ni . It is malleable, very ductile and most tenacious of any metal, the wire being about twice as strong as iron wire (Deville, A. Ch., 1856 (3), 46, 202). The fine powder oxidizes in the air quite rapidly and may even take fire spontaneously; in a compact mass it is but little tarnished in moist air. At a white heat it burns rapidly to Co 3 4 . It is attracted by the magnet and can be made magnetic, retaining (unlike steel) its magnetism at a white heat. 2. Occurrence. — Cobalt does not occur in a free state, except in meteoric iron. It is found in linnaeite (Co s S 4 ); skutterudite (CoAs s ); speiss cobalt (CoNiFeAs 2 ) ; glance cobalt (CoFeAsS 2 ); wad (Co.MnO,2Mn0 2 + 4H 2 0); etc. 3. Preparation. — (1) By electrolysis of the chloride. (2) By heating with potassium or sodium. (3) By heating any of the oxides, hydroxides or the chloride in hydrogen gas. (4) By fusion of the oxalate under powdered glass. (5) Also reduced by carbon in various ways. 4. Oxides and Hydroxides. — Cobaltous oxide, CoO , is made (i) by heating any of its oxides or hydroxides in hydrogen to (not above) 350°; (2) by ignition of Co(OH) 2 or CoC0 3 , air being excluded; (3) by heating Co 3 4 to redness in a stream of C0 2 (Bussell, /. C, 1863, 16, 51) ; (4) by heating any of the higher oxides to a white heat (Moissan, A. Ch., 1880, (5), 21, 242). Cobaltous hydroxide is made from cobaltous salts by precipitation with fixed alkalis; oxidizes if exposed to the air (6a). The most stable oxide is the cobaltoso-cobaltic (Co„0 4 ) tricobalt tetroxide; it is made by heating any of the oxides or hydroxides, the carbonate, oxalate or nitrate to a dull-red heat in the air or in oxygen gas. Several oxide-hydroxides are known, e.g., Co a 2 (OH) 4 , Co 3 O(OH) , Co 3 O 3 (0H) 2 . ■Cobaltic oxide, Co 2 0„ , is made by heating the nitrate just hot enough for de- composition, but not hot enough to form Co 3 4 . Cobaltic hydroxide, Co(OH) 3 , is made by treating any cobaltous salt with CI , HCIO , Br or I in presence of a fixed alkali or alkali carbonate. It dissolves in HC1 with evolution of chlo- rine, in H 2 S0 4 with evolution of oxygen, forming a cobaltous salt. Co0 2 has not yet been isolated, but McConnell and Hanes (J. C, 1897, 71, 584) have shown that it exists as H 2 CoO s and in certain cobaltites. 5. Solubilities. — a. — Metal. — Slowly soluble on warming in dilute HC1 or H 2 S0 4 , more rapidly in HH0 3 , not oxidized on exposure to the air or when heated in contact with alkalis. Like iron, it may exist in a passive form (Niekles, /. pr., 1854, 61, 168; St. Edme, C. r., 1889, 109, 304). With the halogens it forms cobaltous compounds (Hartley, J. C, 1874, 27, 501). b. — Oxides and hydroxides. — Cobaltous oxide (gray-green) and hydroxide (rose-red) are in- soluble in water; soluble in acids, in ammonium hydroxide, and in concentrated solutions of the fixed alkalis when heated (Zimmerman, A., 1886, 232, 324); the various higher oxides and hydroxides are insoluble in ammonium hydroxide or chloride (separation from nickelous hydroxide after treating with iodine in alkaline mixture) (Donath, Z., 1881, 20, 386), and are decomposed by acids, evolving oxygen with non-reducing acids, or a halogen from the halogen acids, and forming cobaltous salts. Co 3 4 is said to be soluble in acids with great diffi- culty (Gibbs and Geuth, Am. S., 1857, (2), 23, 257). c.—Salte — Cobalt forms two classes of salts: cobaltous, derived from CoO , and cobaltic, from Co 2 3 . The latter salts are quite unstable, decomposing in most cases at ordinary tem- peratures, forming cobaltous salts. The cobaltous salts show a remarkable variation of color. The crystallized salts with their water of crystallization are pink; the anhydrous salts are lilac-blue. In dilute solution the salts are pink, but most of them are blue when concentrated or in presence of strong acid. A dilute solution of the chloride spreads colorless upon white paper, turning blue upon heating and colorless again upon cooling, used as " sympa- thetic ink." Cobaltous nitrate and acetate are deliquescent; chloride, hygroscopic; sulphate, efflorescent. The chloride vaporizes, undecomposed, at a high temperature. The carbonate, sulphide, phosphate, borate, oxalate, cyanide, ferrocyanide and ferricyanide are insoluble in water. The potassium-cobaltous oxide is in- soluble; the ammonio-cobaltous oxide, and the double cyanides of cobalt and the alkali metals, soluble in water. Alcohol dissolves the chloride and nitrate; ether dissolves the chloride, sparingly, more so if the ether be saturated with HC1 gas (separation from Ni) (Pin'erua, C. r., 1897, 124, 862). Most of the salts insoluble in water form soluble compounds with ammonium hydroxide. §132, 66. COBALT. 165 6. Reactions, a. — The fixed alkali hydroxides precipitate, from solu- tions of cobaltous salts, blue basic salts, which absorb oxygen from the air and turn olive green, as cobaltoso-cobaltic hydroxide; or if boiled before oxidation in the air, become rose-red, as cobaltous hydroxide, Co(0H) 2 . The cobaltous hydroxide is not soluble in excess of the reagent, "but is somewhat soluble in a hot concentrated solution of KOH (distinction from Ni) (Eeichel, Z., 1880, 19, 468). Freshly precipitated Pb(0H) 2 , Zn(0H) 2 , and HgO precipitate Co(OH) 2 from solutions of cobaltous salts at 100°. Ammonium hydroxide causes the same precipitate as the fixed alkalis; incomplete, even at first, because of the ammonium salt formed in the reaction, and soluble in excess of the reagent to a solution which turns brown in the air by combination with oxygen, and is not precipitated by potassium hydroxide. The reaction of the precipitate with ammonium salts forms soluble double salts (as with magnesium); the reaction of the precipitate with ammonium hydroxide produces, in different conditions, different soluble compounds noted for their bright colors, as (NH 3 ) 4 CoCl 2 , used in plating other metals, in making coins of small denominations, in hardening armor plate, projectiles, etc. The presence of small amounts of phosphorus or arsenic renders it much more fusible, without destroying its ductility; a larger amount makes it brittle. 2. Occurrence. —Nickel almost always occurs in nature together with cobalt. It is found as millerite, NiS 2 ; as nickel blende, NiS; as iron nickel blendej NilTeS; as cobalt nickel pyrites, (NiCoFe) 3 S 4 , etc. 3. Preparation.— (1) By electrolysis. (2) By heating in a stream of hydrogen. The oxide is reduced in this manner at 270° (W. Miiller, Pogg., 1869, 136, 51). (3) By fusing the oxalate under powdered glass (C0 2 being given off). (If) Reduction by igniting in CO . (5) Reduction by fusing with carbon in a variety of methods. (6) By heating the carbonyl,* Ni(C0) 4 to 200°. 4. Oxides and Hydroxides. — Nickelous oxide is formed when the carbonate, nitrate, or any of its oxides or hydroxides are strongly ignited. Nickelous hydroxide is formed by precipitation of nickelous salts with fixed alkalis. NicTcelic oxide, Ni 2 O a , is made from NiC0 3 , Ui(N'0 3 ) 2 or NiO by heating in the air not quite to redness, with constant stirring. It is changed to NiO at a red heat. Nickelic hydroxide, Ni(OH)„ , is formed by treating nickelous' salts first with a fixed alkali hydroxide or carbonate and then with CI , NaCIO , Br or NaBrO (not formed by iodine), a black powder forming no corresponding salts (Campbell and Trowbridge, /. Anal., 1893, 7, 301). A trinickelic tetroxide, KTiaOj , magnetic (corresponding to Co 3 4 , l?e s 4 , Mn 3 4 and Pb s 4 ), is formed, according to Baubigny (C r., 1S78, 87, 1082), by heating NiCl 3 in oxygen gas at from 350° to 440°; and by heating Ni 2 3 in hydrogen at 190° (Moissan, A. (Jh„ 1880, (5), 21, 199). 5. Solubilities. — a. — Metal. — Hydrochloric or sulphuric acid, dilute or con- centrated, attacks nickel but slowly (Tissier, 0. r., 1860, 50, 106) ; dilute nitric acid dissolves it readily, while towards concentrated nitric acid it acts very similar to passive iron (Deville, G. r., 1854, 38, 284). It is not attacked when heated in contact with the alkali hydroxides or carbonates. 6. — Oxides and hydroxides. — Nickelous oxide and hydroxide are insoluble in water or fixed alkalis, soluble in ammonium hydroxide and in acids. Nickelic oxides and hydroxides are dissolved by acids with reduction to nickelous salts, with halogen acids the corresponding halogens are liberated. The moist nickelic hydroxide, formed by the action of CI , Br , etc., in alkaline solution, after washing with hot water liberates free iodine from potassium iodide (distinction from cobalt) . Nickelic hydroxide when treated with dilute sulphuric acid forms NiS0 4 , oxygen being evolved. With nitric acid the action is similar, distinction from cobaltic hydroxide, which requires a more concentrated acid to effect a similar reduction, c. — Salts. — The salts of nickel have a delicate green color in crystals and in solution; when anhydrous, they are yellow. The nitrate and chloride are deliquescent or efflorescent, according to the hygrometric state of the atmosphere; the acetate is efflorescent. The chloride vaporizes at high tem- peratures. The carbonate, sulphide, phosphate, borate, oxalate, cyanide, ferrocyanide and ferricyanide are insoluble; the double cyanides of nickel and alkali metals, soluble in water. The chloride is soluble in alcohol, and the nitrate in dilute alcohol. Most salts of nickel form soluble compounds by action of ammonium hydroxide. 6. Reactions, a. — Alkali hydroxides precipitate solutions of nickel salts as nickel hydroxide, Ni(0H) 2 , pale green, not oxidized by exposure to the air (§132, 6a), insoluble in excess of the fixed alkalis (distinction from zinc), soluble in ammonium hydroxide or ammonium salts, forming a greenish-blue to violet-blue solution. Excess of fixed alkali hydroxide *Nlokel carbonyl is prepared by heating the nickel ore in a current of CO. It is a liquid, sp.. gr., 1.3186, boiling at 43° and freezing at —25°. When heated to 200° it is decomposed into Nl and, CO (Berthelot, C. r., 1891, 112, 1343 ; 113, 679 ; Mond, J. Soc. Ind., 1892, 1.1, 750). 170 nickel. §133, 66. will slowly precipitate nickel hydroxide from the ammoniacal solutions (distinction from cobalt). Alkali carbonates precipitate green basic nickelous carbonate, Ni 5 (0H) 6 (C0 3 ) 2 (composition not constant), soluble in ammonium hydroxide or ammonium salts, with blue or greenish-blue color. Carbonates of Ba, Sr, Ca, and Mg are without action on nickelous chloride or nitrate in the cold (distinction from Fe"', Al , and Cr"'), but on boiling precipitate the whole of the nickel. o. — Oxalic acid and oxalates precipitate, very slowly but almost completely, rafter twenty-four hours, nickel oxalate, green. Alkali cyanides, as KCN , pre- cipitate nickel cyanide, N'i(CH') 2 , yellowish-green, insoluble in hydrocyanic ■acid, and in cold dilute hydrochloric acid; dissolving in excess of the cyanide, by formation of soluble double cyanides, as potassium nickel cyanide (KCN) 2 lTi(ClJ') 2 . The equation of the change corresponds exactly to that for ■cobalt (§132, 66); and the solution of double cyanide is reprecipitated as Ni(CN) 2 by a careful addition of acids (like cobalt); but hot digestion, with the- liberated hydrocyanic acid, forms no compound corresponding to cobalti- -cyanides, and does not prevent precipitation by acids (distinction from cobalt). It will be observed that excess of hydrochloric or sulphuric acid will dissolve the precipitate of Ni(CET) 2 . Ferrocyanides, as K 4 Fe(CN)„ , precipitate a .greenish-white nickel ferrocyanide, Ni 2 T , e(CN) ll , insoluble in acids, soluble in ammonium hydroxide, decomposed by fixed alkalis. Ferricyanides precipitate greenish-yellow nickel ferricyanide, insoluble in acids, soluble in ammonium hydroxide to a green solution (§132, 6b). A solution of nitroferricyanide precipitates solutions of cobalt and nickel salts, the latter being soluble in dilute ammonium hydroxide (Cavalli, Gazsetta, 1897, 27, ii, 95). A solution of potassium xanthate precipitates neutral solutions of nickel and -cobalt, the former being soluble in ammonium hydroxide (distinction), from which solution it is precipitated by (NH 4 ) 2 S (Phipson, 0. N., 1877, 36, 150). 'The xanthate also precipitates nickel in alkaline solution in presence of Na 4 P 2 7 (a separation from Fe"') (Campbell and Andrews, J. Am. Soc, 1895, 17, 125). Nickel salts are not precipitated by an acetic acid solution of nitro'so-/3- naphthol (separation from cobalt) (Knorre, B., 1885, 18, 702). c. — Potassium nitrite in presence of acetic acid does not oxidize nickelous compounds (distinction from cobalt), d. — Sodium phosphate, Na 2 HP0 4 , pre- cipitates nickel phosphate, Mi 8 (P0 4 ) 2 , greenish- white. ei 1$ CD a! 09 a « o 6 • J=i CO S <* ca E o el a ft ° «> a A Pi (1 a :a ft a iz; O O +» ■*> CD r& S -2 ■a the Ir Nitra E 9 s. Sif w o ft •a <5 s o a CD ft po S, m « o » ® A to* « ° -8 u. to 8 • ® Ph R to o O H PI o 03 ++ CO O i-a E 3 ~ -a 4-+ CD P3 1 ° w o . z -fa 02 » O * h -a o : b Comparison of some K< Jw Tfafer Solution, a u. to <0 to O * e 3 ^ E 1 s ° .c o n w o - - - - cd* ft JO ++ ++ T3 O cd ,£) O O CD T3 o £ O Jzi « , si « ^ g CD® P. +3 •o-a-Seiig as 1 saaa£.a||e »3™S « e § s; &.g 83 3 « .2 .h o « o . e 2 *j a s 3 a. a;< 35 gas s as a a §«s f« - u . C2 B d' ■ «"3 £ Q Q. O O DO 3 * fl «H - C8 O o MS !ja+'i;H H «.-. n • OStBubS 1 ; s S © " CD _ Son •a $ If ■S-S. p, a CD © Js -a C § » SIM 5 a s « a 5 ^ 2 2 g s a ■3 = 1^5?' ill o o Q +3 o 1 fl fl ++ ^ # ■*- ++ * §137. TABLE FOR ANALYSIS OF THE ZINO GROUP. 183 §137. Table for Analysis of the Zinc Group (Fourth Group) (Phosphates and Oxalates being absent). Into the clear ammoniacal filtrate from the Third Group pass HYDROSUL- PHURIC ACID GAS, and if a precipitate appears, warm until it subsides. Filter and wash with a one per cent solution of NH 4 C1 . (Test filtrate, in " which H 2 S gives no precipitate for the Fifth Group.) Precipitate: CoS , NiS , MnS , ZnS . Treat on the filter with cold dilute Hydrochloric Acid. Residue: CoS, NiS* (black). For Cobalt: Dissolve in nitro- hydrochloric acid, -evaporate and add NaHC0 3 and H 2 2 ; warm gently and filter. A green color to the filtrate indi- cates cobalt (§140). Test the black resi- due with the borax bead (blue color characteris- tic of cobalt, §132, 7). If sufficient nickel be present to ob- scure the blue bead (§133, 7), dissolve the sul- phides in nitro- hydrochloric acid, evaporate and add an excess of ni- troso-/? -naphthol in acetic acid so- lution (§132, 66); filter, wash and test the brick-red precipitate with the borax bead. Study §132, 6c, §136, §138, §139, §140, §141, §144, §145 and ff. For Nickel : Dissolve the sul- phides in nitro- hydrochloric acid, evaporate and add an ex- cess of nitroso-/?- naphthol in acet- ic solution to re- move the cobalt §132, 66). Filter and add. to fil- trate ammonium hydroxide till al- kaline, filter and to the filtrate addH 2 S. A black precipitate, NiS, indicates nickel. Or: Dissolve the CoS and NiS, add excess of hot KOH and Br, boil, filter, wash (until fil- trate gives no precipitate with AgNO a ), add so- lution of hot KI and test the fil- trate with CS 2 . If free iodine ap- pears, nickel is present(§133,6f). Study the text at §133, 6a, ft, e and f ; §132, 66 and c; §136, §138, §139, §140, §141, §144, §145 and ff. Solution: MnCl 2 , ZnCl 2 (H 2 S,HCl). Boil the solution thoroughly to remove the H 2 S , cool, and add a decided excess of potassium or sodium hydroxide and digest without warming (§135, 6a). Filter and wash. Precipitate: Mn(OH) 2 * Dissolve in nitric acid and boil with an excess of Pb0 2 and HNO s . Violet solution (HMn0 4 ) indi- cates manganese (characteristic reaction, §134, 6c). Dark-colored orig- inal solutions in- dicating an alka- li salt of manga- nese should be reduced by warming with HC1 before pro- ceeding with the analysis (§134, 5c and 6f). Confirm by study of the text, §134, 7, §136, §138, §139, §142, §143, §144, §145 and if- Solution: K,ZnO : . Test for zinc by adding H 2 S. A white precipitate (ZnS) indicates zinc. Study the text at §135, 6a and e, §136, §138, §139, §142, §143, §144, §145 and if. • Small portions of cobalt and nickel sulphides may be dissolved by the cold dilute HC1, and will be precipitated with the Hn(OH|, . These traces will not interfere with the further tests. for manganese. 184 DIRECTIONS FOB ANALYSIS WITH NOTES. §138. DlKECTIONS FOB THE ANALYSIS OF THE METALS OF THE FOUKTH GeOTJP. §138. manipulation. — Into the warm strongly ammoniacal filtrate from the third group (§128), H 2 S gas is passed until complete precipitation is ohtained : MnCl 2 .2NH 4 Cl + 2NH 4 OH + H 2 S = MnS + 4NH 4 C1 + 2H 2 (ira 4 ) 2 Zn.0 2 + 2H 2 S = ZnS + (NH 4 ) 2 S + 2H 2 The solution is warmed until the precipitate subsides, allowed to stand for a few minutes, and is then filtered and the precipitate washed with hot water containing about one per cent of NH 4 C1 (§139, 2). The filtrate should be again tested with H 2 S and if complete precipitation has been obtained it is set aside to be tested for the metals of the succeeding groups (§191). The well washed precipitate of the sulphides of Co , Ni , Mn , and Zn is digested on the filter or in a test-tube with cold dilute HC1 (one part of reagent HC1 to four of water) : MnS -f 2HC1 = MnCl 2 + H 2 S . The black precipitate remaining undissolved contains the sulphides of Co and Hi , the filtrate contains Mn and Zn as chlorides with an excess of HC1 ■and the H 2 S which has not escaped as the gas. §139. Notes. — (1) Instead of passing the H 2 S into the ammoniacal solution, a "freshly prepared solution of ammonium sulphide may be used. The yellow ammonium sulphide, (NH 4 ) 2 Sx , should not be employed to precipitate the metals of the fourth group, as nickel sulphide is quite appreciably soluble in -that reagent (§133, 6e). (2) The sulphides of the fourth group, especially MnS and ZnS , should not be washed with pure water, as they may be changed to the colloidal sulphides, soluble in water. The presence of a small amount of NH 4 C1 prevents this, and ■does not in any way interfere with the analysis of the succeeding groups. (3) If the precipitates are to be treated on the filter with the dilute HC1, the acid solution should be poured on the precipitate three or four times. For digestion in a test tube, the point of the filter is pierced and the precipitate washed into the test tube with as little water as possible. (4) The sulphides of Co and Ni are not entirely insoluble in the cold dilute SCI , and traces of them may usually be detected in the precipitate for Mn (§137, footnote). (5) Dilute acetic acid readily dissolves MnS but scarcely attacks ZnS (§135, 6e). If desired, dilute acetic may be used, first removing the Mn and then adding dilute HC1 to dissolve the Zn . (6) If large amounts of iron are present, a portion of the Mn will always appear in the third group (§134, 6a), and is detected by the green color of the fused mass when testing for Cr: 3Mn(OH) 2 -f- 4KNO s + Wa 2 CO s = 2K 2 MnO, + Ua 2 Mn0 4 + 4NO + C0 2 + 3H 2 . Too much HNO, in the oxidation of the iron favors this precipitation of Mn with Fe'" due to the oxidation of the Mn to the triad or tetrad combination. §140. Manipulation. — The black precipitate of cobalt and nickel sul- phides should first be tested with the borax bead (§141, 8) for the blue bead of cobalt (delicate and characteristic but obscured by the presence of an excess of nickel (§132, 7)). The sulphides are then dissolved in hot HC1 , using a few drops of HN0 3 (§141, 1), and boiled to expel excess of HN0 3 : 6CoS + 12HC1 + 4HN0 3 = 6CoCl 2 + 3S 2 + 4N0 + 8H 2 0. Divide the solution into three portions: To one portion of the solution §142. DIRECTIONS FOB ANALYSIS WITH NOTES. 185 add an excess (§142, 2) of nitroso- ^-Naphthol, filter, and wash with hot water and then with hot HC1 (§132, 6b). Test the red precipitate with the borax head for cobalt. Eender the filtrate ammoniacal, filter again and test this last nitrate with H 2 S for the black precipitate of NiS (§133, 6& and e). To another portion of the solution add NaHC0 3 in excess, then add H 2 2 , warm and filter, a green color to the filtrate indicates cobalt (§132, 10). The third portion of the solution is boiled with an excess of NaOH, bromine water (10, §§132 and 133) is added and the solu- tion is again boiled. The black precipitate of the higher hydroxides (§141, -4) of Co and Ni is thoroughly washed with hot water and then treated on the filter with hot solution of KI (§133, 6/), catching this last filtrate in a test-tube containing CS 2 (§141, 6). Free iodine is evidence of the presence of nickel. §141. Notes.— (i) HNO, interferes with the nitroso- /3-naphthol reaction that follows the solution of the sulphides of Co and Ni , hence an excess is to be avoided. A crystal of KC10 3 may be used instead of HN0 3 . (2) If an insufficient amount of nitroso- /3-naphthol has been used a portion of the cobalt may be in the filtrate and will give the black precipitate for nickel. The filtrate must be tested with the reagent to insure complete removal of the cobalt. (3) Test with the borax bead as follows: Make a small loop on the end of a platinum wire, dip this loop when hot into powdered borax, and heat the adhering mass in the flame until a uniform transparent glassy bead is obtained. Repeat until a bead the size of a kernel of wheat has been made. Bring this hot bead into contact with the precipitate or solution to be tested and fuse again in the burner flame. Allow the bead to cool and notice the appearance. A deep blue indicates cobalt, obscured, however, by a large excess of nickel. (4) The nickel and cobalt may also be oxidized for the KI test as follows: Add five or ten drops of bromine to the solution to be tested in a beaker, warm on a water bath under the hood until the bromine is nearly all expelled, then add rapidly an excess of a hot saturated solution of Na 2 CO a . The black precipitate so obtained will filter rapidly. (5) The test for nickel by adding KI to the mixed higher oxides of cobalt arid nickel is characteristic of nickel and is also a very delicate test. Fully nine-tenths of the cobalt salts sold for chemically pure, show the presence of nickel by this test. (6) In the reaction of nickelic hydroxide with potassium iodide some potas- sium iodate is formed and a greater amount of free iodine will be obtained if a drop of hydrochloric acid be added to the filtrate: KI0 3 + 5KI + 6HC1 = 3l 2 + 6KC1 + 3H 2 (7) If the sulphides of Ni and Co be digested with yellow ammonium sul- phide, a portion of the NiS will be dissolved (§133, 6e) and may be reprecipi- tated as a gray precipitate (black with free sulphur) upon acidulating the- filtrate with acetic acid. It is not a delicate test. §142. Manipulation.— The solution of the sulphides of manganese and zinc in cold dilute hydrochloric acid is boiled thoroughly to insure the removal of the hydrosulphuric acid (§143, 1), cooled (§135, 6a), and then treated with an excess of sodium hydroxide. The zinc forms the soluble zincate, Na 2 Zn0 2 , while the manganese is precipitated as the hydroxide,, white, rapidly turning brown by oxidation : MnCl 2 + 2NaOH = Mn(OH), + 2KC1 ZnCl 2 + 4NaOH = Na 2 Zn0 2 + 2NaCl + 2H 2 186 ANALYSIS OF IRON AND ZINC OROUPS. §143, 1. Filter and test the filtrate with H 2 S > a white or grayish-white precipitate indicates zinc (characteristic). Dissolve the well washed precipitate of Mn(0H) 2 in nitric acid and boil with an excess of lead peroxide, adding more nitric acid. A violet color to the nitric acid solution indicates the presence of manganese (very delicate and characteristic) : 2Mn(6H) 2 + 5Pb0 2 + IOHITO3 = 2HMnO, + 5Pb(N0 3 ) 2 + 6H 2 §143. Notes. — 1. If the H 2 S is not completely removed the Zn will be pre- cipitated as the sulphide upon adding- the WaOH , and will not be separated from the manganese: ZnCl 2 + H 2 S + 2NaOH = ZnS + 2NaCl + 2H 2 . 2. Frequently the precipitate of zinc sulphide is dark gray or almost black. This is usually due to the presence of traces of other sulphides. If iron has not been all removed, through failure to oxidize completely with the nitric acid, it may appear as a precipitate with the manganese, and also as a black precipi- tate with the zinc sulphide. 3. Small amounts of Co and Ni are frequently dissolved by the cold dilute HC1 and will appear with the precipitate o{ Mn(OH) 2 . They do not interfere with the final test for manganese. 4. The, precipitate of Mn(OH) 2 must be washed to remove all the chloride, as the manganese will not be oxidized to permanganic acid until the chloride is completely oxidized to chlorine. 5. Instead of Pb0 2 , red lead, Pb 3 4 , is frequently employed with the nitric acid to oxidize the manganese to permanganic acid: 2Mn(OH) 2 + 5Pb 3 4 + 30HNO 3 = 2HMn0 4 + 15Pb(N0 3 ) 2 + 16H 2 6. It is very difficult to procuie Pb0 2 or Pb 3 4 which does not contain traces of manganese. The student should always boil the lead oxides with nitric acid, and if a violet-colored solution is formed, this should be decanted and the operation repeated until the solution is perfectly colorless after the black precipitate of Pb0 2 has subsided. Then the unknown solution in HN0 3 may be acl^ed and the boiling repeated to test for the manganese. 7. The student is not advised to apply the permanganate test to the original substances. All reducing agents interfere, and Mn0 2 frequently fails to give permanganic acid when boiled with Pb0 2 and HN0 3 until after reduction (§134, 6c). Analysis of Ikon and Zinc Geoups aftek Precipitation bt Ammonium Sulphide. §144. It is preferred by some to precipitate the metals of the third and fourth groups together, by means of ammonium sulphide; using ammonium chloride to prevent the precipitation of magnesium (§189, 5i and 6a), and to insure the complete precipitation of the aluminum as the hydroxide §124, 6a). In the manipulation for this method of separation, the H 2 S is not removed from the second group-filtrate, nor is nitric acid used to oxidize any iron that may be present. To the second group filtrate (§80), warmed, an excess of NH 4 C1 is added (§189, 5c), then NH 4 0H till strongly alkaline, and, paying no attention to any precipitate that may be formed (6a, §§124, 125 and 126), normal ammonium sulphide is added (or what is equivalent H 2 S is passed into the alkaline mixture). Aluminum and chromium are precipitated as the hydroxides, the remaining metals as the sulphides. The following table illustrates a plan of separation of the ammonium sulphide precipitates of the third and fourth group metals, phosphates being absent: 5144. ANALYSIS OF IRON AND ZINC GROUP. 187 Co is •3 a co o 'H •d I? o o o 3 S fl P< 0) fl O n O >M •« 111 a B o« cfl 60P(+J -d fl CO 5 a *s -. B w o w o o O M o a W o w o o -d fl cd w to -a T3 m to o u S * g w o •-* p* - § W-SH S^-fl Is-dS . ft«So t3 B CO co CO .a o fl-d Sod •5 o u ■ + 1 -*> .. o rt !K m CO CO B aw fl — ' .0!ra P-»co' W 3 •SO^flcoco-iSnS CD » •3 H ■>» 2, CD 01 flO « CS^ ^r-," <° Sfc .2" cd -■a > o X X o a w o O ho fl o .-g § <8 £ -d fl a. - " fl c; 4j 'd cd •2,2=1^2 Jfl > =l2 , S 188 IRON AND ZINC GROUPS. §145. §145. The presence of phosphates greatly complicates the work of the analysis of the metals of the third, fourth, and fifth groups. The phos- phates of the alkali metals are soluble, those of the other metals insoluble in water. As the solutions for precipitation of first and second group metals are acid; phosphates remain in solution and do not in any way interfere with the analysis for the metals of those groups; i. e., silver phosphate in nitric acid solution is readily transposed by HC1 ; copper phosphate in acid solution is readily transposed by H 2 S ; etc. §146. When the filtrate from the second group is rendered strongly ammoniacal (§128) the phosphates of all the metals present, except those of the alkalis, are precipitated. Phosphates of cobalt, nickel and zinc are redissolved by an excess of ammonium hydroxide. Freshly precipitated ferric phosphate is transposed by the alkali hydroxides (incompletely in the cold). The phosphates of Al, Cr^ and Zn are soluble in the fixed alkalis, the solution of chromium phosphate is decomposed by boiling,, precipitating Cr(0H) 3 and leaving the alkali phosphates in solution. §147. In analysis a portion of the filtrate from the second group (after the removal of the H 2 S) (§128) should be tested for phosphoric acid with ammonium molybdate (§75, 6d). If phosphates are present the usual methods of analysis for third, fourth, and fifth groups must be modified. Several methods have been recommended: §148. First. — To the filtrate from the second group, H 2 S, being re- moved (§128), an excess of the reagent ammonium molybdate is added, the mixture set aside in a warm place for several hours, until the yellow ammonium phospho-molybdate has completely formed and settled (§75, Gd). Filter and evaporate nearly to dryness to remove the nitric acid. Take up with water and a little hydrochloric acid if necessary to obtains clear solution, and remove the excess of molybdenum with H 2 S (§75, 6e). From this point proceed by the usual methods of analysis (§§127, 128 and ff.). §149. Second. — Precipitation of the phosphate as ferric phosphate in acetic acid solution. This method of separation rests upon the fact that the phosphates of the fourth group and of the alkaline earths are soluble,. and the phosphates of Al , Cr"' and Fe'", insoluble in acetic acid. To the filtrate from the second group, freed from H 2 S by boiling (128), and nearly neutralized with Na 2 C0 3 , an excess of NaC 2 H 3 2 is added and then FeCl 3 solution, drop by drop, as long as a precipitate is formed.. Care must be taken to avoid an excess of FeCl 3 , as the ferric phosphate is soluble in a solution of ferric acetate. As soon as the phosphate is all precipitated the blood-red ferric acetate is formed at once, indicating the presence of a sufficient amount of FeCl 3 . The mixture should be boiled §151. IRON AN$) ZINC GROUPS. 189> to precipitate the ferric acetate as basic ferric acetate (§126, 66) and at once filtered. Upon the addition of the sodium acetate the aluminum and chromium are precipitated as phosphates, provided there be sufficient phosphate present to combine with them; if not the whole of the phosphate will be precipitated and the first, drop of FeCl 3 will give a red solution showing the addition of that reagent to be unnecessary. By the above method of manipulation any iron present in the original solution is in the ferrous condition and does not react to precipitate the phosphate, as ferrous phosphate is soluble in acetic acid. If the iron has been previously oxidized with nitric acid it will react with the phosphate- upon the addition of the sodium acetate ; but if there be more iron present than necessary to combine with the phosphate, the red ferric acetate solu- tion will be formed with the excess of the iron and render the precipita- tion of the phosphate incomplete. In this case the previous oxidation of the iron is detrimental. If alkaline earth salts are present in quantity more than sufficient to- combine with the phosphoric acid radical, not all of these metals will be precipitated with the third group metals upon the addition of ammonium hydroxide. The table (§152) illustrates the separation of the metals In presence of the phosphates by the use of FeCl 3 in acetic acid solution. §150. Third. — A method of separation of the third group metals with phosphates from the remaining metals is based upon the action of freshly- precipitated barium carbonate. Solutions of Al , Cr'", and Fe"' are pre- cipitated as the hydroxides by digestion in the cold with freshly precipi- tated BaC0 3 (6a, §§124, 125 and 126): 2A1C1 3 + 3BaC0 3 + 3H 2 = 2A1(0H) 3 + 3BaCl 2 + 3C0 2 . Solutions of the chlorides or nitrates of' the fourth group and of the alkaline earths are not transposed by cold digestion with BaC0 3 . Sulphates of the fourth group are transposed by freshly precipitated BaC0 3 in the cold: CoS0 4 + BaC0 3 = BaS0 4 + C0CO3 , etc.; and must not be present in this method of separation (§126, 6a). ' If an excess of ferric chloride be present the phosphates will all be precipitated as ferric phosphate and the Al, Cr'" and excess of Fe'" as. the hydroxides upon the digestion with BaCO, . The table (§153) gives an illustration of the use of the BaC0 3 in effecting the separation. It should be observed that presence or absence of FeCl 3 or of BaC0 3 in the sample must be fully determined before their addition as reagents. §151. Oxalates do not interfere with the usual course of analysis of the- first two groups of metals; with the other metals oxalates interfere very much the same as phosphates. They, however, with other interfering 190 IRON AND ZINC GROUPS. §151. organic matter, can readily be removed by ignition. If the presence of an oxalate has been established (§§188, 6& and 227, 8), the second group nitrate should be evaporated to dryness, moistened with concentrated HN0 3 and gently ignited. The residue, dissolved in HC1 , is then ready for the usual process of analysis. For the analysis in presence of silicates and borates the student is referred to the text under those elements (§§249, 8 and 221, 8). $152. IRON, ZINC AND CALCIUM GROUP METALS. 191 § 3 Sh * "5 •5 g o •a a oJ t? N ff o .3 P. »8 W £ r>-f ^ ^ £ ® £ P u 5 Kg o g ° a 2« a fl;g lite fn . . cs "O -; Sag* ■ra-al >sl o 8 * « s ° s _■• fl O d td us * 3 'ft d o ftsos t « gO rtfl 9 3 S _. 03 CO (4 bo o § 3 9 I fl 9-*" 3-^-g . •I'd H O-P «? O ft 3^»2 " -S ±i d « fl » * ° a 198 IRON, ZINC AND CALCIUM GROUP METALS. 5153. o O CO ns 'd p* coo CO • "S3 «*2 CO »rt 55 ■dfr P< . - a o o pa- l«o ""3 8 d ',3 0" 3 " I «L « ft co Pit- 's -a £ "- S d S » 13 3 9 ft ,4 CO m ^ - * g o A Peg d p ^j CO d .s O) XI O ^ CO boE .a co,a 'fls^ff: 3 ga & -" J o o d^ ba *d « * " ft O CO O ; a a II p" 1 h o .2 "> -£ O M c 23d <|! w IU MJ " -t-> CO > o S u a h a a •2 S .So ala w *H ^ d S' ffl w J) d " -H CO CO +J SO .0) .S rt M BOS l S3 W o CO ^ V ft CO u a S CO B O «H -p O EC OJ Pl'rd o • |H CO M » CO a CO Sc? §.2 ^^ 8^ 2 ^^ . O cO • « «h r^i-i tH o oO «»'rt _° ■ rt u fd ^— '-"^v W ^H co "2 .fcl >> 3^3 1-* i— i o ^ gco a nr £h o ra a? * rrl 1* ■g|-l Blo :a ISM*' ft*£« ' rt cO « . ft .a oj u ••* cO oS -d » h fl cap-So H P4 « 3^ PI c3 d '^iIh ■S ^ a .2 .S "5 d ~ S O _ c3 0* " m d *S «» ^-coft- -d j, .,« fe CO f^G-^ 1 o o . - P 1 '■§ •a ; S3 -*J PH O "-^ O a. ; -73 .S 5/3 >>Sr t> gc0 lu .-, ^ S -^^ .-^ o cs: 1 « 5 1 = S o § V m {§155. CERIUM— COWMBIUM. 193 The Barer Metals of the Iron and Zinc Groups. Cerium, Columbium (Niobium), Didymium, Erbium, Gallium, Glucinmm (Beryllium), Indium, Lanthanum, Neodymium, Praseodymium, Sama- rium, Scandium, Tantalum, Terbium, Thallium, Thorium, Titanium, Uranium, Ytterbium, Yttrium, Zirconium. §154. Cerium. Ce = 139.0 . Valence three and four. Specific gravity, 6.628. Melts higher than Sb and lower than Ag (Hilljebrandt and Norton, Pogg., 1875, 156 ; 466). Cerium is a comparatively rare metal, never, found native; it is found in many minerals in Sweden, especially in cerite,' -which is chiefly a silicate of Ce , La , Ne , Pr , Al and Fe; also found in a brick-making clay near Frankfurt, Germany (Strohecker, J. pr., 1886, (3), 33, 133 and 260). It was first described in 1803 by Klaproth, but in 1839 Mosander showed the supposedly pure cerium oxide to consist of oxides of at least three metals: Ce, La, D (Ne and Pr) (Pogg., 1842, 56, 503). The metal is obtained from the chloride, CeCl 3 , by electrolysis or by heating with sodium. It is a steel-gray, lustrous, malleable, ductile metal; fairly stable in air under ordinary conditions. When heated in air it burns with incandescence. It burns in CI , Br and in vapor of I , S and P . Soluble in acids. Two oxides are known, Ce 2 3 and Ce0 2 , forming two classes of salts, cerous and eerie, the latter being less stable. Ignition in air or oxygen changes Ce 2 3 to Ce0 2 . Ce 2 3 is white or grayish-white, soluble in acids and formed by igniting Ce 2 (C0 3 ) 3 , Ce 2 (C 2 4 ) 3 or Ce0 2 in an atmosphere of hydrogen. Cerous salts are white and form color- less solutions in water. Ceric oxide, Ce0 2 , is yellowish-white, orange-yellow -when hot, soluble in acids with difficulty; the hydroxide dissolves readily. •Ceric salts are yellow or red, forming yellow solutions. Ceric hydroxide, Ce(OH) 4 , dissolves in HC1 with evolution of chlorine, forming colorless cerous ■chloride. Sulphurous acid decolorizes solutions of ceric salts, forming cerous salts. Fixed alkali hydroxides and ammonium sulphide precipitate, from .solutions of cerous salts, the white cerous hydroxide, turning yellow by absorp- tion of oxygen, with formation of ceric hydroxide. The precipitate is in- soluble in excess of the fixed alkalis (distinction from Al and Gl). The pre- cipitation is hindered by the presence of tartaric acid (distinction from yttrium). Ammonium hydroxide precipitates a basic salt. Alkali carbonates precipitate cerous carbonate, soluble in excess of the fixed alkali carbonates. Oxalic acid forms cerous oxalate, white, from moderately acid solutions, soluble in hot (NH,)jCjO ( , but reprecipitated on dilution with cold water. A con- centrated solution of K„SO, forms the double sulphate, K 3 Ce(S0 4 ) 3 , white, sparingly soluble in water, insoluble in K„S0 4 solution (distinction from Gl). Na 2 S 2 0„ does not precipitate cerium salts. BaC0 3 does not precipitate cerous salts in the cold, but precipitates them completely on boiling. Ceric salts are completely precipitated by BaC0 3 in the cold. Alkali hypochlorites precipitate cerous salts as the yellow ceric hydroxide. If cerous nitrate be boiled with Pb0 2 and HN0 3 , eerie nitrate, a deep yellow solution is formed (delicate test for cerium). Cerium givas no absorption spectrum, but the spark spectrum •shows several brilliant lines. §155. Columbium (Niobium). Cb = 93.7 . Valence five. Columbium usually occurs with tantalum in such minerals as columbite and tantalite; it is also found in tantalum free minerals as euxenite, pyrochlor, etc. The metal is prepared by passing the penta-chloride mixed with hydrogen repeatedly through a hot tube. It is a steel-gray lustrous metal, specifie gravity, 7.06 at 15.5°. By ignition in the air it burns readily to the pentoxide. Not attacked by chlorine in the cold, but when warmed combines readily, forming CbCl 5 . The metal is not soluble in hydrochloric, nitric or nitrohydro- 194 DIDYMIUM. gigg chloric acids, but is readily soluble in hot concentrated sulphuric acid, forming a colorless solution (Roscoe, G. N., 1878, 37, 25). It forms several oxides, CbO Cb0 2 and Cb 2 5 . Columbic acid (anhydride) Ct^O, , is a white powder, yellow when hot (distinction from tantalum); it is obtained by ignition of the lower oxides, or by decomposition of solutions of the salts by water or alkalis and igniting. Cb0 2 , black, is prepared by strongly igniting Cb 2 B in a current of hydrogen. Cb 2 5 , not too strongly ignited, is soluble in acids, from which solutions NHjOH and (NH,),S precipitate Columbia acid containing some am- monia. By mixing Cb 2 5 with charcoal and heating in a current of chlorine, a mixture of CbOCl 3 and CbCl„ is obtained. CbCl 6 is a yellow crystalline solid (needles), melting at 194° and distilling at 240.5° (Deville and Troost, C. r., 1867, 64, 294). Upon treating the chloride with water, it is partially decomposed to columbic acid, a large portion remaining in solution and not precipitated by H 2 S0 4 (distinction from tantalum). Cb 2 5 not- previously ignited dissolves in HE; which solution when mixed with KF , the HF being in excess, gives a double fluoride, 2KF.CbF B ; if the HF be not in excess, a double oxy-fluoride is obtained, 2KF.CbOF 3 (Kruess and Nilson, B., 1887, 20, 1676). The potassium columbium fluoride is much more soluble than either the corresponding tita- nium or tantalum compounds. Fusion of columbic acid with the alkalis gives the columbates, the potassium salt being quite soluble in water and in potas- sium hydroxide; the sodium salt is only soluble in water after removal of the excess of the ' sodium hydroxide. From a solution of potassium columbate, sodium hydroxide precipitates, almost completely, sodium columbate. Carbon dioxide precipitates columbic acid from solutions of columbates. Soluble salts of Ba , Ca and Mg form white bulky precipitates with a solution of potassium columbate. AgN0 3 gives a yellowish-white precipitate, CuS0 4 ^ green pre- cipitate. Cb 2 5 in presence of HC1 or H 2 S0 4 gives a blue to brown, color with Sn or Zn, due to partial reduction of the Cb (distinction from tantalum). Fused with sodium meta-phosphate, columbic acid gives in the inner flame a violet to blue bead; a red bead by addition of FeS0 4 . oicn t>-j • f Neodymium. Nd = 143.6 . Valence three. 8156. Didymmm = < _ J , . _ _„„ __ , ,, I Praseodymium. Pr = 140.5. Valence three. Specific gravity, 6.544. Melts with greater difficulty than Ce or La . Present in cerite in Sweden and in monazite sand from Brazil. Didymium was reported about 1840 by Mosander, having been separated from cerium and lanthanum. In 1885 Welsbach (M., 1885, 6, 477) separated didymium salts into two distinct salts, neodymium and praseodymium. By the absorption spectrum bands other chemists are of the opinion that the so-called didymium consists of a group of elements, nine or more (Kruess and Nilson, B., 1887, 20, 2166; Kreuss, A., 1892, 265, 1). Concerning the separation of didymium compounds, see Dennis and Chamot (J. Am. Soc, 1897, 19, 799). By repeated fractionation of the nitrate (several thousand times) Welsbach obtained a pale green salt and a rose-colored salt, which gave different spectra but which, united, gave the spectrum of didymium. Didymium oxide absorbs water to form the hydroxide, which absorbs C0 2 from the air, but does not react alkaline to litmus. The salts are soluble in water to a reddish solution. The saturated sulphate solu- tion does not deposit crystals until heated to boiling; while lanthanum sulphate precipitates from the saturated solution at 30°. Fixed alkalis precipitate the hydroxide; NH 4 OH , a basic salt; insoluble in excess of the reagents. Alkali carbonates form a bulky precipitate, insoluble in excess of the reagent, barium carbonate precipitates slowly but completely. Precipitation by alkalis is pre- vented by tartaric acid. Oxalic acid precipitates didymium salts completely, soluble with difficulty in HC1 . The double potassium sulphate forms much more slowly and less completely than with cerium. The salts give a distinct and characteristic absorption spectrum. Consult Jones (Am., 1898, 20, 345), Scheie (Z. anorg., 1898, 17, 319), Boudard (C. r., 1898, 126, 900), Demarcay (C. r., 1898, 126, 1039), and Brauner (C. N„ 1898, 77, 161). §159. ERBIUM— -GALLIUM— GLUCINUM. 195 §157. Erbium. Er = 166.0 . Valence three. Erbium metal has not been prepared. As oxide or earth it is described by Cleve (0. r., 1880, 91, 381) as that yttrium earth the most beautiful rose colored. It forms a characteristic absorption spectrum, and a spark spectrum with sharp lines in the orange and green. This earth has not been thoroughly studied and quite probably consists of the oxides of several metals (Boisbau- dran, C. r., 1886, 102, 1003; Soret, 0. r., 1880, 91, 378; Crookes, G. N., 1886, 54, 13). The oxide gives upon ignition an intense green light; it is not fusible or volatile. §158. Gallium. Ga = 70.0 . Valence three. Specific gravity, the solid, at 23° to 24.5°, 5.935 to 5.956; the melted, at 24.7°, 6.069. Melting point, 30.15°; frequently may be cooled to 0° without again be- coming solid. It is a grayish-white metal, crystallizing in octahaedra or in broad plates. It is quite brittle and gives a bluish-gray mark on paper. It gives a very weak and fugitive flame spectrum; the spark spectrum shows two beautiful violet lines. When heated in the air or in oxygen it is but slightly oxidized; does not vaporize at a white heat; soluble in acids and alkalis; attacked by the halogens (with iodine only upon warming). In the Periodic System it is the Ekaaluminum of Mendelejeff, who described the general prop- erties before the metal was discovered (0. r., 1875, 81, 969). It occurs in zinc blende (black) from Bensberg on the Rhine; in brown blende from the Pyrenees; and in some American zinc blendes (Cornwall, Ch. Z., 1880, 4, 443). It is prepared by electrolysis after previous purification of the ore by chemical methods. 4300 kilos of the Bensberg ore gave 55 kilos of pure gallium (Bois- baudran and Jungfleisch, G. r., 1878, 86, 475). The oxide, Ga 2 3 , is a white powder obtained by igniting the nitrate. After strong ignition it is insoluble in acids or alkalis. It is easily attacked on fusion with KOH or KHSO, . The alkalis and the alkali carbonates precipitate the salts as the hydroxide, perceptibly soluble in fixed alkali carbonates, more easily in ammonium hydroxide and in ammonium carbonate, and very readily in the fixed alkalis. Tartrates hinder the precipitation of the hydroxide. The salts of gallium are colorless and for the most part soluble in water. The neutral solutions upon warming precipitate a basic salt, dissolving again upon cooling. Excess of zinc forms a basic zinc salt which precipitates the gallium as oxide or basic salt. BaCO., precipitates gallium salts in the cold. K 4 Fe(CN') 9 gives a precipitate, insoluble in SCI , noticeable in very dilute solutions (1-175,000). H 2 S does not precipitate gallium salts from solutions acid with mineral acids; from the acetate or in presence of ammonium acetate the white sulphide, Ga 2 S a , is precipitated; (NH,),S precipitates the sulphide. Gallium chloride, GaCl s , is a colorless salt, melting at 75° and volatilizing at 215° to 220°. The vapor density indicates the molecule to be Ga 2 Cl„ , which decomposes to GaCl, at about 400° (Friedel and Kraft, G. r., 1888, 107, 306). Upon evaporat- ing a solution of the chloride on a water bath the salt is perceptibly volatil- ized, not so if HjSOj be present. Gallium sulphate forms with ammonium sulphate an alum. For separation from other metals, see Boisbaudran, C. r., 1882, 95, 410, 503, 1192, 1332. §159. Glucinum (Beryllium). Gl = 9.1 . Valence two. Specific gravity, 1.85 (Humpidge, Proc. Roy. Soc, 1871, 39, 1). Melting point, below 1000 (Debray, A. Ch., 1855, (3), 44, 5). It is a white malleable metal, obtainable in hexagonable crystals (Mlson and Pettersson, B., 1878, 11, 381 and 906). It was first discovered in 1797 by Vauquelin from beryl. It is stable in the air, does not decompose steam at a red heat, and at red heat is scarcely attacked by oxygen or sulphur. It is a strongly positive element, 198 SCANDIUM— TANTALUM— TERBIUM. §165„ §165. Scandium. So = 44.1 . Valence three. It is found in euxenite and gadolinite with, yttrium. Its name comes from Scandinavia, where it was first found. It is separated from ytterbium, with which it is always closely associated, by heating the nitrates; the basic scan- dium nitrate being precipitated before the ytterbium basic nitrate, or by precipitating as the double potassium sulphate, the corresponding ytterbium, salt remaining in solution. The oxide, Sc 2 O s , is a white flocculent infusible- powder, readily soluble in warm acids. The solutions of the salts show no absorption bands in the spectrum. The spark spectrum of the chloride gives over 100 bright lines (Thalen, C. r., 1880, 91, 45). Solutions of the salts taste sweet and have an astringent action. The alkalis precipitate the hydroxide a white bulky precipitate, insoluble in excess of the precipitant. Tartrates, hinder the precipitation in the cold, but not upon heating. Na 2 C0 3 gives a bulky white precipitate, soluble in excess of the reagent. HjS is without action, but (NH,) 2 S precipitates the hydroxide. K,S0 4 precipitates the double scandium sulphate, 3K2S0 4 .Sc 2 (S0 4 ) 3 , soluble in water but not in a saturated KS0 4 solution. §166. Tantalum. Ta = 182.8 . Valence five. Tantalum occurs in tantalite and columbite, silicates, nearly always ac- companied by columbium. It is prepared by heating the tantalum alkali fluoride with K or Ma in a well-covered crucible (Rose, Pogg., 1856, 99, 65). It is a black or iron-gray powder with a metallic lustre. Specific gravity, 10.78. Heated in the air it burns with incandescence to form TajO,, . It is insoluble in acids except HF , in which it dissolves with evolution of H . Upon ignition in a current of chlorine, TaCl 5 , volatile, is formed. Solution of alkalis has; no action, upon fusion with the fixed alkalis an alkali tantalate is formed. Ta 2 5 is a white infusible powder, specific gravity, 8.01 (Marignac, A. Ch., 1866, (4), 9, 254). The oxide fused with fixejd alkalis gives also an alkali tantalate, M'Ta0 3 . When KOH is used, the fused mass is soluble in water. When NaOH is used, water removes the excess of alkali, leaving the MaTa0 3 as a white residue, which dissolves in pure water, but not in NaOH solution. Tantalum. chloride is a yellow solid, melting at 211.3° and boiling at 241.6", with 753 mm. atmospheric pressure (Deville and Troost, G. r., 1867, 64, 294). It is com- pletely decomposed by water, forming the hydrated acid, 2H i TaO 3 .H 2 = H^TajO, . The freshly precipitated acid is soluble in acids and reprecipitated by MH 4 OH . The acid is readily soluble in HF , which solution with KF forms a characteristic double salt, 2KF.TaF 5 , crystallizing in fine needles, insoluble in water slightly acidulated with HIP (distinction and separation from colum- bium). A solution of alkali tantalate gives with HC1 a precipitate of tantalic acid, soluble in excess of the HC1 . From this solution NH 4 OH or (NH,) 2 S precipitates tantalic acid; H-SO, precipitates tantalic sulphate. Tartaric acid prevents the precipitation with MH 4 OH and (TTHj)^ . A solution of tantalic acid gives no coloration with zinc (distinction from Cb). Solutions of alkali tantalates form tantalic acid with C0 2 . The acid fused with sodium meta- phosphate gives a colorless bead (distinction from Si0 2 ) , which does not become blood-red upon adding PeS0 4 and heating in the inner flame (distinction from titanium). §167. Terbium. Tr = 160. Valence three. The terbium compounds are very similar to the yttrium compounds. The salts are colorless and give no absorption spectrum. The double potassium terbium sulphate has about the same solubilities as the corresponding cerium compound, and so the terbium is frequently precipitated with cerium com- pounds. Terbia, Tr 2 3 , is the darkest colored of the yttrium earths, soluble= §169. THALLIUM— THORIUM. 199> in acids and sets BH, free from ammonium salts. The hydroxide is a, gelatinous precipitate which absorbs C0 2 from the air. It is quite probable- that terbia is a mixture of rare earths (Boisbaudran, C. r., 1886, 102, 153, 395 v 483 and 899). §168. Thallium. Tl = 204.15 . Valence one and three. Thallium was discovered by Crookes by means of the spectroscope in 1861,, in selenium residues of the H 2 S0 4 . factory at Tilkerode in the Hartz Mountains, Germany (G. N., 1861, 3, 193, 303; 1863, 7, 290; 1863, 8, 159, 195, 219, 231, 243, 255 and 279). It is found widely distributed in many varieties of iron and copper pyrites, but in large proportions it is only found in Crookesite in Sweden. This mineral contains as high as 18.55 per cent Tl (Nordenskjoeld, A., 1867, 144, 127). It is prepared by reduction from its solutions with Zn or Al; by electrolysis; by precipitation with KI , and then reduction by Zn or Al or by electrolysis. Specific gravity, 11.777 to 11.9 (Werther, J. pr., 1863, 89', 189). Melting point, 290° (Lamy, C. r., 1862, 54, 1255). It is a bluish-white metal, softer than lead, malleable and ductile; tarnishes rapidly in the air; may be preserved under water, which it does not decompose below a red heat; soluble in H 2 S0 4 and HUO, , in HC1 with great difficulty; combines directly with CI , Br , I , P , S , Se , arid precipitates from their solutions Cu , Ag , Hg , Au and Fb in the metallic state. As a monad its compounds are stable, and not easily oxidized; as a triad it is easily reduced to the univalent condition. Thallious oxide, T1 2 , is black; on contact with water it forms an hydroxide, T10H , freely soluble in water and in alcohol, to colorless solutions. The car- bonate is soluble in about 20 parts of water; the sulphate and phosphate are soluble; the chloride very sparingly soluble; the iodide insoluble in water. Hydrochloric acid precipitates, from solutions not very dilute, thallious chloride, T1C1 , white, and unalterable in the air. As a silver-group precipitate, thallious chloride dissolves enough in hot water to give the light yellow pre- cipitate of iodide, Til , on adding a drop of potassium iodide solution, the precipitate being slightly soluble in excess of the reagent. H,S precipitates the acetate, but not the acidified solutions of its other salts. (TJH.,) 2 S pre- cipitates T1 2 S , which, on exposing to the air, soon oxidizes to sulphate. Eerrocyanides give a yellow precipitate, Tl 4 l?e(CN) ; phosphomolybdic acid a yellow precipitate; and potassium permanganate a red-brown precipitate, con- sisting in part of T1 2 3 . Chromates precipitate yellow normal chromate; and platinic chloride, pale orange, thallious platinic chloride, Tl 2 PtCl,, . Thallium compounds readily impart an intense green color to the flame, and one emerald- green line to the spectrum (the most delicate test). The flame-color and' spectrum, from small quantities, are somewhat evanescent, owing to rapid vaporization. Thallic oxide, Tl 2 O s , dark violet, is insoluble in water; the hydroxide, an oxyhydroxide, TIO(OH), is brown and gelatinous. This hydrox- ide is precipitated from thallic salts by the caustic alkalis, and not dissolved be excess. Chlorides and bromides do not precipitate thallic solutions; iodides precipitate Til with I. Sulphides and HzS precipitate thallious sulphide, with sulphur. Thallic oxide, suspended in solution of potassium hydroxide, and treated with chlorine, develops an intense violet-red color. Thallic chloride and sulphate are reduced to thallious salts by boiling their water solutions. §169. Thorium. Th = 232.6 . Valence four. Thorium is a rare element found in thorite (a silicate), orangite and some other minerals. It was described by Berzelius in 1828 (Pogg., 1829, 16, 385), who also prepared the metal by reduction of the potassium thorium fluoride with potassium. The metal is a gray powder; specific gravity, 11.000; stable in air at ordinary temperature, but igniting when heated; attacked by vapors of CI , Br , I and S . Sparingly soluble in dilute acids, easily soluble in concen- trated acids; insoluble in the alkalis (Nilson, B., 1882, 15, 2519 and 2537; Kruess 200 TITANIUM. §170. and Nilson, B., 1887, 20, 1665). Thorium forms one oxide, Th0 2 , upon ignition of the oxalate. It is a snow-white powder, not easily soluble in acids if highly ignited (Cleve, J., 1874, 261). The hydroxide, Th(OH) 4 , is formed by precipita- tion of the salts by the alkalis. It is a white, heavy, gelatinous precipitate drying to a hard glassy mass. The chloride, ThCl 4 , and the nitrate, Th(M'0,) 4 , are deliquescent. The chloride is a white bodj r melting at a white heat and then subliming in beautiful white needles (Kruess and Nilson, I. c). The sulphate is soluble in five parts of cold water. The carbonate, oxalate and phosphate are insoluble in water; the oxalate is scarcely soluble in dilute mineral acids. Alkali hydroxides or sulphides precipitate thorium hydroxide, Th(OH), , insoluble in excess of the reagent. Tartaric and citric acids hinder the pre- cipitation. Alkali carbonates precipitate the basic carbonate, soluble in ex- cess, if the reagent be concentrated. The solution in (NH 4 ) 2 CO = readily repre- cipitates upon warming. BaCO , precipitates thorium salts completely. Oxalic acid and oxalates form a white precipitate (distinction from Al and Gl), not soluble in oxalic acid or in dilute mineral acids; soluble in hot concentrated (NH 4 ) 2 C 2 4 and not reprecipitated on cooling and diluting (distinction from Ce and La). A saturated solution of K 2 S0 4 slowly but completely precipitates a solution of Th(S0 4 ) 2 , forming potassium thorium sulphate; insoluble in a saturated K,S0 4 solution, sparingly soluble in cold water, readily soluble in hot water. HP precipitates ThF 4 , insoluble in excess, gelatinous, becoming crystalline on standing. Boiling freshly precipitated Th(0H) 4 with KF in presence of HF forms K 2 ThP„.4H 2 , a heavy fine white precipitate almost insoluble in water. The distinguishing reactions of thorium are the precipitation with oxalates and with K,S0 4 , and failure to form a soluble compound on fusion with Na 2 C0 3 (distinction from Si0 2 and Ti0 2 ). §170. Titanium. Ti = 48.15 . Valence three and four. Titanium is found quite widely distributed as rutile, brookite, anatase, titanite, titaniferous iron, PeTiO a , and in many soils and clays. Never found native. It is prepared by heating the fluoride or chloride with K or Na . It is a dark gray powder, which shows distinctly metallic when magnified. Heated in the air it burns with an unusually brilliant incandescence; sifted into the flame it burns with a blinding brilliance. Chlorine in the cold is without action, when heated it combines with vivid incandescence. It decomposes water at 100°. It is soluble in acids, with evolution of hydrogen, forming titanous salts. At a higher temperature it combines directly with Br and I. It is almost the only metal that combines directly with nitrogen when heated in the air (Woehler and Deville, A., 1857, 103, 230; Merz, J. pr., 1866, 99, 157). The most common oxide of titanium is the dioxide, Ti0 2 , analogous to C0 2 and Si0 2 . It occurs more or less pure in nature as rutile, brookite and anatase; it is formed by ignition of the hydrated titanic acid or of ammonium titanate (Woehler, J., 1849, 268). Ignition of Ti0 2 in dry hydrogen gives Ti 2 0, , an amorphous black powder, dissolving in H 2 S0 4 to a violet-colored solution (Ebel- men, A. Ch., 1847, (3), 20, 392). TiO is formed when Ti0 2 is ignited with Mg: 2Ti0 2 + Mg = TiO + MgTi0 3 (Winkler, B., 1890, 23, 2660). Other oxides have been reported. Titanic acid, Ti0 2 , is a white powder, melts somewhat easier than Si0 2 , soluble in the alkalis unless previously strongly ignited. Mixed with charcoal and heated in a current of chlorine! TiCl 4 is formed. The bromide is formed in a similar manner. Ti0 2 acts as a base, forming a series •of stable salts; also as an acid, forming titanates. TiCl 4 is a colorless liquid, fuming in the air; it boils at 136.41° (Thorpe, J. C, 1880, 37, 329); it is de- composed by water, forming titanic acid, which remains in solution in the HC1 present. Solutions of most of the titanic salts, when boiled, deposit the insoluble meta-titanic acid. HF dissolves all forms of titanic acid; if the ;solution be evaporated in presence of H 2 S0 4 no TiP 4 is volatilized (distinction irom SiF 4 ). When evaporated with HP alone, TiP 4 is volatilized. The double ■potassium titanium fluoride, K 2 TiF„ , formed by fusing Ti0 2 with acid KF , is •sparingly soluble in water (96 parts), readily soluble in HC1 . Solutions of titanic salts in water or acid solutions of titanic acid are precipitated by §171. URANIUM. 201 alkali hydroxides, carbonates and sulphides as the hydrated titanic acid, insolu- ble in excess of the preeipitants and in ammonium salts. BaC0 3 gives the same- precipitate. K 1 Fe(ClT) 6 gives a reddish-yellow precipitate; K 3 Fe(CN) a yellow precipitate. Na^PO, precipitates the titanium almost completely, even in the- presence of strong HC1 . An acid solution of Ti0 2 when treated with Sn or Zn. gives a pale blue to violet coloration to the solution, due to a partial reduction of the titanium to the triad condition. These colored solutions are precipitated by alkali hydroxides, carbonates and sulphides, HS is without action. The solution reduces Fe'" to Fe" , Cu" to Cu' , and salts of Hg , Ag and Au to the metallic state; the titanium becoming again the tetrad. The reduction by Sn or Zn takes place in presence of HF (distinction from columbic acid). Titanium compounds fused in the flame with microcosmic salt give in the reducing flame a yellow bead when hot, cooling to reddish and violet (reduction of the tita- nium). With FeS0 4 in the reducing flame a blood-red bead is obtained. §171. Uranium. U = 239.6 . Valence four and six. Specific gravity, 18.685 (Zimmermann, A., 1882, 213, 285). Melts at a bright red heat (Peligot, A. Gh., 1869, (4), 17, 368). Found in various minerals; its chief ore is pitch-blende, which contains from 40 to 90 per cent of Us0 8 . Prepared by fusing TTC1 4 with K or Na (Zimmermann, A., 1883, 216, 1; 1886,. 232, 273). It has the color of nickel, hard, but softer than steel, malleable,, permanent in the air and water at ordinary temperatures; when ignited burns, with incandescence to XT s O s ; unites directly with CI , Br , I and S when heated;, soluble in HC1 , H-SO^ and slowly in HUO, . Uranous oxide, TJO : , formed by igniting the higher oxides in carbon or hydrogen, is a brown powder, soon turning yellow by absorption of oxygen from the air. Uranous hydroxide is. formed by precipitating uranous salts with alkalis. Uranic oxide, UO s , is. formed by heating uranic nitrate cautiously to 25°, and upon ignition in the- air both this and other uranium oxides, hydroxides and uranium oxysalts with volatile acids are converted into TJ 3 O s = U0 2 2TJ0 3 . Uranium acts as a base in two classes of salts, uranous and uranyl salts. Uranous salts are green and give- green solutions, from which alkalis precipitate uranous hydroxide, insoluble in excess of the alkali; alkali carbonates precipitate tr(OH) 4 , soluble in (MH 4 ) 2 C0 3 ; with BaCO, the precipitation is complete even in the cold. H 2 S is. without action; (NH 4 ) 2 S gives a dark-brown precipitate; K 4 Fe(CN) gives a. reddish-brown precipitate. In their action toward oxidizing and reducing agents uranous and uranyl (uranic) salts resemble closely ferrous and ferric salts; uranous salts are even more easily oxidized than ferrous salts, e. g., by exposure to the air, by HN0 3 , CI, HC10 3 , Br, KMnO, , etc. Gold, silver and platinum salts are reduced to the free metal. The hexad uranium (TTvi) acts- as a base, but usually forms basic salts, never normal: we have TTO : .(M'0 3 ) 2 , not TJ(Kr0 3 ) 6 ; TTO^SC^ , not TT(S0 4 ) 3 . These basic salts were formerly called uranic salts, but at present CO"0 2 )" is regarded as a basic radical and called uranyl, and its salts are called uranyl salts, e. g., XT0 2 C1 2 uranyl chloride, (tT0 2 ) s (P0 4 ) z uranyl orthophosphate. Solutions of uranyl salts are yellow; KOH and NaOH give a yellow precipitate, uranates, K 2 TX 2 7 and TH^TIjOt ,. insoluble in excess. Alkali carbonates give a yellow precipitate, soluble in excess; BaCO, and CaCO, give TT0 3 . HzS does not precipitate the uranium, but slowly reduces uranyl salts to uranous salts (Formanek, A., 1890, 257, 115). (NH 4 ) 2 S gives a dark-brown precipitate. K 4 Fe(ClT) 6 gives a reddish-brown precipitate. Used in the analysis and separation of uranium compounds (Fresenius and Hintz, Z. angew., 1895, 502). Sodium phosphate gives a yellow precipitate. The hexad uranium acts as an acid toward some stronger bases. Thus we have K 2 TT 2 7 and Na 2 TJ 2 7 , formed by precipitating uranyl salts with KOH and NaOH; compare the similar salts bf the hexad chromium, K 2 Cr 2 Oy and ITajjCrjOv • Other oxides of uranium are described, but are doubtless com- binations of U0 2 and tJ0 3 . Zn , Cd , Sn , Pb , Co , Cu , Fe , and ferrous salts reduce uranyl salts to uranous salts. Solutions of Sn , Pt , Au , Cu , Hg and Ag are reduced to the metal by metallic uranium (Zimmermann, I. c). For method of recovery of waste uranium compounds, see Laube (Z. angew., 1889, 575). '202 YTTERBIUM— YTTRIUM— ZIRCONIUM. §172, §172. Ytterbium. Yb = 173.2 . Valence three. Obtained as an earth by Marignac (C. r., 1878, 87, 578) from a gadolinite earth; by Delafontaine (C. r., 1878, 87, 933) from sipylite found at Amherst, Va. Nilson (B., 1879, 12, 550; 1880, 13, 1433) describes its preparation from euxenite and its separation from So . It has the lowest bacisity of the yttrium earths. The double potassium ytterbium sulphate is easily soluble in water and in potassium sulphate. The oxalate forms a white crystalline precipitate, in- soluble in water and in dilute acids. The salts are colorless and give no absorption spectrum. For the spark spectrum see Welsbach (M., 1884, 5, 1). The oxide, Yb 2 8 , is a white powder, slowly soluble in cold acids, readily upon warming'. The hydroxide forms a gelatinous precipitate, insoluble in NHjOH but soluble in KOH . It absorbs C0 2 from the air. The nitrate melts in its "water of crystallization and is very soluble in water. §173. Yttrium. Y = 89.0 . Valence three. Yttrium is one of the numerous rare metals found in the gadolinite mineral •at Ytterby, near Stockholm, Sweden; also found in Colorado (Hidden and Mackintosh, Am. 6'., 1889, 38, 474). The metal has been prepared by electro- lysis of the chloride; also by heating the oxide, Y 2 O s , with Mg (Winkler, B., 1890, 23, 787). The study of these rare earths is by no means complete. It is ■also claimed that they have not yet been obtained pure, but that the so-called pure oxides really consist of a mixture of oxides of from five to twenty ele- ments (Crookes, C. N., 1887, 55, 107, 119 and 131). The most of these rare earths do not give an absorption spectrum, but give characteristic spark spectra; and it is largely by this means that the supposedly pure oxides have been shown to be mixtures of the oxides of several closely related elements (Wels- bach, M., 1883, 4, 641; Dennis and Chamot, J. Am. Soc, 1897, 19, 799). Yttrium salts are precipitated by the alkalis and by the alkali sulphides as the ■hydroxide, Y(OH), , a white bulky precipitate, insoluble in the excess of the reagents (distinction from Gl). The oxide and hydroxide are readily soluble in acids; boiling with NH4CI causes solution of the hydroxide as the chloride. 'The alkali carbonates precipitate the carbonate Y 2 (C0 3 ) s , soluble in a large excess of the reagents. If the solution in ammonium carbonate be boiled, the hydroxide is precipitated. Soluble oxalates precipitate yttrium salts as the white oxalate (distinction from Al and Gl) ; soluble with some difficulty in HC1 . The double sulphate with potassium is soluble in water and in potassium sulphate (distinction from thorium, zirconium and the cerite metals). BaCO s forms no precipitate in the cold (distinction from Al , Fe'" , Cr"' , Th , Ce , la, Nd and Pr). Hydrofluoric acid precipitates the gelatinous fluoride, YF 3 , insoluble in water and in HF . The precipitation of yttrium salts is not hindered by the presence of tartaric acid (distinction from Al , Gl , Th and Zr). The analysis of yttrium usually consists in its detection and separation in gadolinite (silicate of Y , Gl , Fe , Mn , Ce and La). Fuse with alkali car- bonate, decompose with HC1 , and filter from the Si0 2 . .Neutralize the filtrate and precipitate the Y, La and Ce as oxalates with (NH,),C,0 ( . Ignite the precipitate and dissolve in HC1 . Precipitate the La and Ce as the double potassium sulphates, and from the filtrate precipitate the yttrium as the 'hydroxide with NE ( OH . Ignite and weigh as the oxide. In order to effect icomplete separations the operations should be repeated several times. §174. Zirconium. Zr = 90.4 . Valence four . "Zirconium is a rare metal found in various minerals, chiefly in zircon, a silicate; never found native. The metal was first prepared by Berzelius in 1824 by fusion of the potassium zirconium fluoride with potassium (Pogg., 1825, 4, 117). Also prepared by electrolysis of the chloride (Becquerel, A. Gh., 1831, 48, 337). The metal exists in three modifications: crystalline, graphitoidal and amorphous- The amorphous zirconium is a velvet-black powder, burning when §175. TEE CALCIUM CROUP. 203 heated in the air. Acids attack it slowly even when hot, except HF , which •dissolves it in the cold. It forms but one oxide, Zr0 2 , analogous to Si0 2 and Ti0 2 . Zr0 2 is prepared from the mineral zircon by fusion with a fixed Alkali. Digestion in water removes the most of the silicate, leaving the alkali zirconate as a sandy powder. Digestion with HC1 precipitates the last of the Si0 2 and dissolves the zirconate. The solution is neutralized, strongly •diluted and boiled; whereupon the zirconium precipitates as the basic chloride free from iron. Or the zirconium may be precipitated by a saturated solution ■of KSO, , and after resolution in acids precipitated by NH,OH and ignited to Zr~O s (Berlin, J. pr., 1853, 58, 145; Roerdam., C. C, 1889, 533). Zr0 2 is a white infusible powder, giving out an intense •white light when heated ; it shows no lines in the spectrum. It is much used with other rare earths, La,0 :! , Y 2 0, , etc., to form the mantles used in the Welsbach gas-burners (Drossbach, C. C, 1891, 772; Welsbach, J., 1887, 2670; C. N., 1887, 55, 192). The oxide (or hydroxide precipitated hot) dissolves with difficulty in acids to form salts. The hydroxide, ZrO(OH).. , precipitated in the cold dissolves readily in acids. As an acid, .zirconium hydroxide, ZrO(OH) 2 = H 2 Zr0 3 , forms zirconates, decomposed by acids. As a base it forms zirconium salts with acids. The sulphate is easily ■soluble in water, crystallizing from solution with 4H 2 . The phosphate is insoluble in water, formed by precipitation of zirconium salts by Na.HPO, or H a P0 4 . The silicate, Zr0 2 .SiO_ , is found in nature as the mineral zircon, usually containing traces of iron. Zirconium chloride is formed when a current of chlorine is passed over heated Zr0 2 , mixed with charcoal. It is a white solid, may be sublimed, is soluble in water. Solutions of zirconium salts are precipitated as the hydroxide, ZrO(OH) 2 , by alkali hydroxides and sulphides, a white flocculent precipitate, insoluble in excess of the reagents, insoluble in UHjCl solution (difference from Gl). Tartaric acid prevents the precipitation. Alkali carbonates precipitate basic zirconium carbonate, white, soluble in excess of KHC0 3 or (11114)2003; boiling precipitates a gelatinous hydroxide from the latter solution. BaCO, does not precipitate zirconium salts com- pletely, even on boiling. The precipitates of the hydroxide and carbonate are soluble in acids. Oxalic acid and oxalates precipitate zirconium oxalate, solu- ble in excess of oxalic acid on warming, and soluble in the cold in (NHJ.CjO, (difference from thorium) ; soluble in HC1 . A saturated solution of K 2 S0 4 precipitates the double potassium' zirconium sulphate, white, insoluble in excess of the reagent if precipitated cold, soluble in excess of HC1; if precipitated hot, almost absolutely insoluble in water or HC1 (distinction from Th and Ce). Zirconium salts are precipitated on warming with Na 2 S 2 8 (separation from Y , Nd and Pr)'. Solution of H.0 2 completely precipitates zirconium salts. Tumeric paper moistened with a solution of zirconium salt and HC1 is colored orange upon drying (boric acid gives the same reaction) (Brush, J. pr., 1854, 62, 7). HF does not precipitate zirconium solutions, as zirconium fluoride, ZrF, , is soluble in water and in HF (distinction from Th and Y). The Calcium Gbottp (Fifth Geotjp). (The Alkaline Eaeth Metals.) Barium. Ba = 137.40 . Calcium. Ca = 40.1 . Strontium. Sr = 87.60 . Magnesium. Mg = 24.3 . §175. Like the alkali metals, Ba , Sr - and Ca oxidize rapidly in the air at ordinary temperatures — forming alkaline earths — and decompose water, forming hydroxides with evolution of heat. Mg oxidizes rapidly in the air when ignited, decomposes water at 100°, and its oxide — in physical proper- ties farther removed from Ba , Sr , and Ca than these oxides are from each 204 TEE CALCIUM CROUP. §176. other — slowly unites with water without sensible production of heat. As compounds, these metals are not easily oxidized beyond their quantivalence as dyads, and they require very strong reducing agents to restore them to the elemental state. §176. In basic power, Ba is the strongest of the four, Sr somewhat stronger than Ca, and Mg much weaker than the other three. It will be observed that the solubility of their hydroxides varies in the same decreas- ing gradation, which is also that of their atomic weights; while the solubility of their sulphates varies in a reverse order, as follows: (§7) : §177. The hydroxide of Ba dissolves in about 30 parts of water; that of Sr , in 100 parts; of Ca , in 800 parts; and of Mg , in 100,000 parts. The sulphate of Ba is not appreciably soluble in water (429,700 parts at 18.4°;, Hollemann, Z. phys. Gh., 1893, 12, 131); that of Sr 'dissolves in 10,000 parts; of Ca , in 500 parts; of Mg , in 3 parts. To the extent in which they dissolve in water, alkaline earths render their solutions caustic to the taste and touch, and alkaline to test-papers and phenolphthalein. §178. The carbonates of the alkaline earths are not entirely insoluble- in pure water: BaC0 3 is soluble in 45,566 parts at 24.2° (Hollemann,. Zeit. phys. Gh., 1893, 12, 125); SrC0 3 in 90,909 parts at 18° (Kohlrausch and Eose, Zeit. phys. Oh., 1893, 12, 241); CaC0 3 in 80,040 parts at 33.8°" (Hollemann, I. c); MgC0 3 in 9,434 parts (Chevalet, Z., 1869, 8, 91). The presence of NH 4 0H and (NH 4 ) 2 C0 3 lessens the solubility of the carbonates of Ba , Sr , and Ca , while their solubility is increased by the presence of NH 4 C1 . MgC0 3 is soluble in ammonium carbonate and in ammonium chloride, so much so that in presence of an abundance of the latter it is not at all precipitated by the former, i. e. (NH 4 ) 2 C0 3 does not precipitate a solution of MgCl 2 as the NH 4 C1 formed holds the Mg in solution. §179. These metals may be all precipitated as phosphates in presence of ammonium salts, but their further separation for identification or esti- mation would be attended with difficulty (§145 and ff.). §180. The oxalates of Ba, Sr, and Mg are sparingly soluble in water, calcium oxalate insoluble. Barium chromate is insoluble in water (§§27 and 186, 5c), strontium chromate sparingly soluble, and calcium and mag- nesium chromates freely soluble. §181. In qualitative analysis, the group-separation of the fifth-group metals is effected, after removal of the first four groups of bases, by precipitation with carbonate in presence of ammonium chloride, after which magnesium is precipitated from the filtrate, as phosphate. §182. The hydroxides of Ba, Sr, and Ca, in their saturated solutions, necessarily dilute, precipitate solutions of salts of the metals of the first four groups and of Mg , as hydroxides. In turn, the fixed alkalis precipi- tate, from solutions of Ba , Sr , Ca , and Mg , so much of the hydroxides §186, 4. BARIUM 205 of these metals as does not dissolve in the water present *; hut ammonium hydroxide precipitates only Mg , and this but in part, owing to the solubility of Mg(0H) 2 in ammonium salts. §183. Solutions containing Ba , Sr . Ca , and Mg , with phosphoric, oxalic, boric, or arsenic acid, necessarily have the acid reaction, as occurs in dis- solving phosphates, oxalates, etc., with acids; such solutions are precipi- tated by ammonium hydroxide or by any agent which neutralizes the solu- tion, and, consequently, we have precipitates of this kind in the third group (§145 and//.): CaCl 2 + H„P0 4 + 2NH 4 0H = CaHP0 4 + 2NH 4 C1 + 2H 2 CaH 4 (P0 4 ) 2 + 2NH 4 OH = CaHP0 4 + (NH 4 ) 2 HP0 4 + 2H 2 . If excess of the ammonium hydroxide be added the precipitate is Ca 3 (P0 4 ) 2 . In the case of a magnesium salt the precipitate is MgNH 4 P0 4 . §184. The carbonates of the alkaline earth metals are dissociated by heat, leaving metallic oxides and carbonic anhydride. This occurs with difficulty in the ease of Ba . §185. Compounds of Ba , Sr , and Ca (preferably with HC1) impart char- acteristic colors to the non-luminous flame, and readily present well-defined spectra. §186. Barium. Ba = 137.40 . Valence two. 1. Properties. — Specific gravity, 3.75 (Kern, G. N., 1875, 31, 243); melting point, above that of cast iron (Frey, A., 1876, 183, 368). It is a white metal, stable in dry air, but readily oxidized in moist air or in water at ordinary temperature, hydrogen being' evolved and barium hydroxide formed. It is malleable and ductile (Kern, I. a). 2. Occurrence. — Barium can never occur in nature as the metal or oxide, or hydroxide near the earth's surface, as the metal oxidizes so readily, and the oxide and hydroxide are so basic, absorbing acids readily from the air. Its most common forms of occurrence are heavy spar, BaS0 4 , and witherite, BaCO s . 3. Preparation. — (i) By electrolysis of the chloride fused or moistened with strong HC1 . (2) By electrolysis of the carbonate, sulphate, etc., mixed with Hg and HgO , and then distilling the amalgam. (3) By heating the oxide or various salts with sodium or potassium and extracting the metal formed with mercury, then separating by distillation of the amalgam. 4. Oxides and Hydroxides.— The oxide, BaO , is formed by the action of heat upon the hydroxide, carbonate, nitrate, oxalate, and all its organic salts. The corresponding hydroxide, Ba(OH) 2 , is made by treating the oxide with water. The peroxide, Ba0 2 , is made by heating the oxide almost to redness in oxygen, or air which has been freed from carbon dioxide; by heating the oxide with potassium chlorate (Liebig, Pogg., 1832, 26, 172) or cupric oxide (Wanklyn, B., 1874, 7, 1029). It is used as a source of oxygen, which it gives off at a white heat, BaO remaining; also in the manufacture of hydrogen peroxide, H 2 2 , which is formed by treating it with dilute acids: Ba0 2 + 2HC1 = BaCl 2 + H 2 2 . *The presence of an excess of fixed alkali renders these hydroxides much less soluble, the high concentration of the hydroxyl ions, one of the factors of the solubility product, diminish- ing the other factor. (§45). 206 BARIUM. §186, 5a. 5. Solubilities. — a. — Metal. — Metallic barium is readily soluble in acids with evolution of hydrogen. 6. — Oxides and hydroxides. — Barium oxide is acted upon by water with evolution of heat and formation of the hydroxide, which is soluble in about 30 parts of cold water and in its own weight of hot water (Rosenstheil and Ruehlmann, J., 1870, 314). Barium peroxide, BaO, , is very sparingly soluble in water (Schone, A., 1877, 192, 257); soluble in acids with formation of H»0 2 . c. — Salts. — Most of the soluble salts of barium are permanent; the acetate is efflorescent. The chloride, bromide, bromate, iodide, sulphide, ferrocyanide, nitrate, Irypophosphite, chlorate, acetate, and phenylsul- phate, are freely soluble in water; the carbonate, sulphate, sulphite, chr ornate,* phosphite, phosphate, oxalate, iodate, and siUco-fluoride, are insoluble in water. The sulphate is perceptibly soluble in strong HC1 . The chloride is almost insoluble in strong hydrochloric acid (separation from Ca and Mg) (Mar, Am. 8., 1892, 143, 521); likewise the nitrate in strong hydrochloric and nitric acids. The chloride and nitrate are insolu- ble in alcohol. 6. Reactions, a. — The fixed alkali hydroxides precipitate only con- centrated solutions of barium salts (5&). No precipitate is formed with ammonium hydroxide (§45). The alkali carbonates precipitate barium carbonate, BaC0 3 , white. The precipitation is promoted by heat and by ammonium hydroxide, but is made slightly incomplete by the presence of ammonium salts (Vogel, J. pr., 1836, 7, 455). Barium Carbonate — BaC0 3 — is a valuable reagent for special purposes, chiefly for separation of third and fourth group metals. It is used in the form of the moist precipitate, which must be thoroughly washed. It is best precipitated from boiling solutions of barium chloride and sodium or ammonium carbonate, washed once or twice by decantation, then by filtra- tion, till the washings no longer precipitate solution of silver nitrate. Mixed with water to consistence of cream, it may be preserved for some time in stoppered bottles, being shaken whenever required for use. When dissolved in hydrochloric acid, and fully precipitated by sulphuric acid, the filtrate must yield no fixed residue. This reagent removes sulphuric acid (radical) from all sulphates in solution to which it is added (e) : Na 2 S0 4 + BaCO, = BaS0 4 -(- Na 2 C0 3 . When salts of non-alkali metals are so decomposed, of course, they are left insoluble, as carbonates or hydroxides, nothing remaining in solution: FeSd .+ BaCO s = BaS0 4 + !FeC0 3 Fe 2 (S0 4 ) 8 + 3BaC0 3 + 3H 2 = 3BaS0 4 + 2Fe(OH) a + 3C0 2 The chlorides of the third group, except Fe" , are decomposed by barium carbonate; while the metals of the fourth group (zinc, manganese, cobalt, nickel), are not precipitated from their chlorides by this reagent. Tartaric *Kohlrausch and Rose, Z. phys. Ch., 1893, 12, 241 ; Schweitzer, Z., 1890, 89, 414. §186, 7. BARIUM. 207 ucid, citric acid, sugar, and other organic substances, hinder or prevent the decomposition by barium carbonate. 6. — Ammonium oxalate precipitates barium oxalate, BaC,0 + , from solutions f rom Sr (and of course from Ca) ; but precipitation by the (very dilute §187, 5c) solution of strontium sulphate is a more certain test between Ba and Sr . BaS0 4 is not transposed by solutions of alkali carbonates (distinction from Sr and Ca, §188, 6a foot- note). /. — Solutions of iodates, as NaI0 3 , precipitate, from barium solutions not very dilute, barium iodate, Ba(IO a ) 2 , white, soluble in 600 parts of hot or 1746 parts of cold water (distinction from the other alkaline earth metals). g. — Neutral or ammoniacal solutions of arsenous acid do not precipitate barium salts (distinction from calcium). Soluble arsenates precipitate solutions of barium salts, soluble in acids, including arsenic acid. h. — Soluble chromates, as K 2 Cr0 4 , precipitate solutions of barium salts as barium chromate, BaCr0 4 , yellow; almost insoluble in water (separa- tion from calcium and from strontium except in concentrated solutions), sparingly soluble in acetic acid, moderately soluble in chromic acid and readily soluble in hydrochloric and nitric acids. Bichromates, as K 2 Cr 2 7 , precipitate solutions of barium salts (better from the acetate) as the normal chromate (very accurate separation from strontium and calcium) (Grittner, Z. angew., 1892, 73). i. — Fluosilicic acid, BuSiP„ , precipitates white, crystalline barium fluo- silicate, BaSiF 6 , slightly soluble in water (1-4000), not soluble in alcohol (distinction from strontium and calcium). If an equal volume of alcohol be added the precipitation is complete, sulphuric acid not giving a precipitate in the nitrate (Fresenius, Z., 1890, 29, 143). 7. Ignition. — The volatile salts of barium as the chloride or nitrate impart a yellowish-green color to the flame of the Bunsen burner, appearing blue when viewed through a green glass. The spectrum of barium is readily distinguished from the spectra of other metals by the green bands Baa, /3 and y . Barium carbonate is very stable when heated, requiring a very high heat to decompose it into BaO and CO, . 208 STRONTIUM. §186, 8. 8. Detection. — In the filtrate from the fourth group, barium is precipi- tated with strontium and calcium as the, carbonate by ammonium car- bonate. The white precipitate (well washed) is dissolved in acetic acid and the barium precipitated with K 2 Cr 2 7 as BaCr0 4 which separates it from strontium and calcium. The barium is further identified by the non-solubility of the chromate in acetic acid, the solubility in hydrochloric acid, and precipitation from this solution by sulphuric acid. It may also be confirmed by the color of the flame with any of the volatile salts (7) (not the sulphate). 9. Estimation. — Barium is weighed as a sulphate (Fresenius and Hurtz, Z. angew., 1896, 253), carbonate or fluosilicate (BaSiIT ). It is separated from strontium and calcium: (1) By digesting the mixed sulphates at ordinary tem- peratures for 12 hours with ammonium carbonate. The calcium and strontium, are thus converted into carbonates, which are separated from the barium sulphate by dissolving in hydrochloric acid. (2) By hydrofluosilicic acid. (3) By repeated precipitation as the chromate in an acetate solution. It is separated from calcium by the solution of the nitrate of the latter in amyl alcohol (§188, 9). The hydroxide and carbonates are also determined by alkalimetry. Volumetrically it is precipitated as the chromate, thoroughly washed, dissolved in dilute HC1 and the Crvi determined by H 2 2 (Baumann, Z. angew., 1891, 331). 10. Oxidation. — Barium compounds are reduced to the metal when heated with Ha or K (3). Ba0 2 oxidizes MnCl 2 to Mn 2 O a (Spring and Lucion, Bl., 1890* (3), 3, 4). §187. Strontium. Sr = 87.60 . Valence two. 1. Properties. — Specific gravity, 2.4 (Franz, J. pr., 1869, 107, 254). Melts at a. moderate red heat and is not volatile when heated to a full red. It is a " brass- yellow " metal, malleable and ductile. It oxidizes rapidly when exposed to- the air, and when heated in the air burns, as does barium, with intense illumination (Franz, I.e.). 2. Occurrence. — Strontium occurs chiefly in strontianite, SrCO s , and in celestine, SrS0 4 . 3. Preparation. — First isolated in 1808 by Davy by electrolysis of the hydrox- ide (Trans. Royal Soc, 345). It is made by electrolysis of the chloride (Frey,. A., 1876, 183, 367); by heating a saturated solution of SrCl 2 with sodium amalgam and distilling off the mercury (Franz, I. c.) ; by heating the oxide with powdered magnesium the metal is obtained mixed with MgO (Winkler, B., 1890,. 23, 125). 4. Oxides and Hydroxides. — Strontium oxide, SrO , is formed by igniting the hydroxide, carbonate (greater heat required than with calcium carbonate),, nitrate and all organic strontium salts. The hydroxide, Sr(OH) 2 , is formed by the action of water on the oxide. The peroxide, Sr0 2 .8H 2 , is made by pre cipitating the hydroxide with H 2 2 ; at 100° this' loses water and becomes Sr0 2 . a white powder, melting at a red heat, used in bleaching works (Conroy /. Soc. Ind., 1892, 11, 812). 5. Solubilities. — o. — Metal. — Strontium decomposes water at ordinary tem- perature (Winkler, I. a), it is soluble in acids with evolution of hydrogen. 6. — Oxides and hydroxides. — The oxide, SrO , is soluble in about 100 parts water at ordinary temperature, and in about five parts of boiling water forming the hydroxide (Scheibler, Neue Zeitschrift fur Ruebenzucker, 1881, 49, 257). The: peroxide is scarcely soluble in water or in ammonium hydroxide, soluble in- acids and in ammonium chloride. §187, 6h. STRONTIUM, 209 c. — Salts. — The chloride is slightly deliquescent; crystals of the nitrate and acetate effloresce. The chloride is soluble, the nitrate insoluble in absolute alcohol. The nitrate is insoluble in boiling amyl alcohol (§188, 5c). The sulphate is very sparingly soluble in water (1-10,090 at 20.1°) (Hollemann, Z. phys. Ch., 1893, 12, 131); yet sufficiently soluble to allow its use as a reagent to detect the presence of traces of barium. Less soluble in water containing ammonium salts, sodium sulphate, or sulphuric acid than in pure water; quite appreciably soluble in HC1 or HN0 3 ; insoluble in alcohol. Strontium fluosilicate is soluble in water (distinction from barium). The chromate is soluble in 831.8 parts water at 15° (Fresenius, Z., 1890, 29, 419); soluble in many acids including chromic acid; and more soluble in water containing ammonium salts than in pure water. 6. Reactions, a. — The fixed alkalis precipitate strontium salts when not too dilute, as the hydroxide, Sr(0H) 2 , less soluble than the barium hydroxide. No precipitate with ammonium hydroxide. The alkali car- bonates precipitate solutions of strontium salts as the carbonate. Stron- tium sulphate is completely transposed on boiling with a fixed alkali car- bonate (distinction from barium, §188, 6a footnote). 6. — Oxalic acid and oxalates precipitate strontium oxalate, insoluble in ■water, soluble in hydrochloric acid (Souchay and Lenssen, A., 1857, 102, 35). c. — The solubility of strontium salts is diminished by the presence of con- centrated nitric acid, but less so than barium salts, d. — In deportment with phosphates, strontium is not to be distinguished from barium. e. — See 6e, §§186 and 188. Sulphuric acid and sulphates (including CaS0 4 ) precipitate solutions of strontium salts as the sulphate, unless the solution is diluted beyond the limit of the solubility of the precipitate (5c). A solution of strontium sulphate is used to detect the presence of traces of barium (distinction from strontium and calcium). In dilute solutions the precipitate of strontium sulphate forms very slowly, aided by boiling or by the presence of alcohol, prevented by the presence of hydrochloric or nitric acids (5c). It is almost insoluble in a solution of ammonium sulphate (separation from calcium). f. — The halides of strontium are all soluble in water and have no application in the analysis of strontium salts. Strong hydrochloric acid dissolves stron- tium sulphate, but in general diminishes the solubility of strontium salts in "water, g. — Neutral solutions of arsenites do not precipitate strontium salts, the addition of ammonium hydroxide causes a precipitation of a portion of the strontium. Arsenate of strontium resembles the corresponding barium salt. Alkaline arsenates do not precipitate strontium from solution of the sulphate (distinction from calcium, §188, 6g). h. — Normal chromates precipitate strontium chromate from solutions not too dilute (5c), soluble in acids. In absence of barium, strontium may be separated from calcium by adding to the nearly neutral solutions a solution of K 2 Cr0 4 plus one-third volume of alcohol. The calcium 210 CALCIUM. §187, 6t. chromate is about 100 times as soluble as the strontium chromate (Fre- senius and Eubbert, Z., 1891, 30, 672). No precipitate is formed with potassium bichromate (separation from barium). i. — Fluosilicic acid does not precipitate strontium salts even from quite concentrated solutions, as the strontium fluosilicate is fairly soluble in cold water and more so in the presence of hydrochloric acid (Fresenius, Z. 1890 29,143). 7. Ignition. — Volatile strontium compounds color the flame crimson. In pres- ence of barium the crimson color appears at the moment when the substance (moistened with hydrochloric acid, if a non-volatile compound) is first brought into the flame. The paler, yellowish-red flame of calcium is liable to be mis- taken for the strontium flame. The spectrum of strontium is characterized by eight bright bands; namely, six red, one orange and one blue. The orange line Sr a, at the red end of the spectrum; the two red lines, Sr /3 and Sr y t and the blue line, Sr S , are the most important. 8. Detection. — Strontium is precipitated with barium and calcium from the filtrate of the fourth group by ammonium carbonate. The well washed precipitate of the carbonates is dissolved in acetic acid and the barium removed by K 2 Cr 2 7 . The strontium and calcium are separated from the excess of chromate by reprecipitation with (NH 4 ) 2 C0 3 . The precipitate is again dissolved in HC 2 H 3 2 and from a portion of the solution the stron- tium is detected by a solution of CaS0 4 (6e). The name test (7) is of value in the identification of strontium. 9. Estimation. — Strontium is weighed as a sulphate or a carbonate. The hydroxide and carbonate may be determined by alkalimetry. It is separated from calcium: (1) By the insolubility of its sulphate in ammonium sulphate. (2) By boiling the nitrates with amyl alcohol (§188, 9). (S) By treating the nitrates with equal volume of absolute alcohol, and ether (§188, 9). For separation from barium see §186, 9. §188. Calcium. Ca = 40.1 . Valence two. 1. Properties.— Specie gravity, 1.6 to 1.8 (Caron, C. r., 1860, 50, 547). Meltmp point, at red heat (Matthiessen, A., 1855, 93, 284). A white metal having very much the appearance of aluminum, is neither ductile nor malleable (Frey, A., 1876, 183, 367). In dry air it is quite stable, in moist air it burns with incandescence, as it does also with the halogens. It dissolves in mercury, form- ing an amalgam. 2. Occurrence. — Found in the mineral kingdom as a carbonate in marble* limestone, chalk and arragonite; as a sulphate in gypsum, selenite, alabaster, etc.; as a fluoride in fluor-spar; as a phosphate in apatite, phosphorite, etc. It is found as a phosphate in bones; in egg-shells and oyster-shells as a car- bonate. It is found in nearly all spring and river waters. 3. Preparation. — (1) By ignition of the iodide with sodium in closed retorts (Dumas, C. r., 1858, 47, 575). (2) By fusion of a mixture of 300 parts fused CaCl 2 , 400 parts granulated zinc and 100 parts Wa until zinc vapor is given off. From the CaZn alloy thus obtained the zinc is removed by distillation in a graphite crucible (Caron, I. c). (3) By electrolysis of the chloride (Frey, I. c). (4) By reducing the oxide, hydroxide or carbonate with magnesium (Winkler, B., 1890, 23, 122 and 2642). 4. Oxides and Hydroxides. — The oxide, CaO , is a strong base, non-fusible, non-volatile; it is formed by oxidation of the metal in air; by ignition of the §188, 5c. CALCIUM. 211 hydroxide, the carbonate (limestone), nitrate, and all organic calcium salts* The corresponding hydroxide, Ca(OH) 2 (slaked lime), is made by treating the oxide with water. Its usefulness when combined with sand, making mortar, is too well known to need any description here. The peroxide, Ca0 2 .8H„0 , is. made by adding hydrogen peroxide or sodium peroxide to the hydroxide: Ca(OH) 2 + H 2 2 = Ca0 2 + 2H 2 (Conroy, /. Soc. Ind., 1892, 11, 808). Drying at 130° removes all the water, leaving a white powder, C.a0 2 , which at a red heat loses half its oxygen (Schoene, A., 1877, 192, 257). It cannot be made by- heating the oxide in oxygen or .with potassium chlorate (§186, 4). 5. Solubilities. — a. — Metal. — Calcium is soluble in acids with evolution of hydrogen; it decomposes water, evolving hydrogen and forming Ca(OH) 2 . &. — Oxide and hydroxide. — CaO combines with dilute acids forming cor- responding salts, it absorbs C0 2 from the air becoming CaC0 3 .* In moist air it becomes Ca(0H) 2 , the reaction takes place rapidly and with increase of volume and generation of much heat in presence of abundance of" water. The hydroxide, Ca(0H) 2 , is soluble in acids, being capable of titration with standard acids. It is much less soluble in water than barium or strontium hydroxides (Lamy, C. r., 1878, 86, 333); in 806 parts, at 19.5° (Paresi and Eotondi, B., 1874, 7, 817); and in 1712 parts at 100° (Lamy, I. c). The solubility decreases with increase of temperature. In saturated solutions one part of the oxide is found in 744 parts of water at 15° (Lamy, I. c). A clear solution of the hydroxide in water is lime water (absorbs C0 2 forming CaC0 3 ), the hydroxide in suspension to a greater or less creamy consistency is milk of lime. c. — Salts.- — The chloride, bromide, iodide, nitrate, and chlorate are deliquescent; the acetate is efflorescent. The carbonate, oxalate, and phosphate are insoluble in water. The chloride, iodide, and nitrate are soluble in alcohol. The nitrate is soluble in 1.87 parts of equal volumes of ether and alcohol (Fresenius, Z., 1893, 32, 191); readily soluble in boiling amyl alcohol (Browning, Am. 8., 1892, 143, 53 and 314) (separation from, barium and strontium). The carbonate is soluble in water saturated with carbonic acid (as also are barium, stron- tium, and magnesium carbonates), giving, hardness to water. The oxalate is insoluble in acetic acid, soluble in hydrochloric and nitric acids. The sulphate is soluble in about 500 parts of water f at ordinary temperature,, the solubility not varying much in hot water until above 100° when the solubility rapidly decreases. Its solubility in most alkali salts is greater than in pure water. Ammonium sulphate (1-4) requires 287 parts for the solution of one part of CaS0 4 (Fresenius, Z., 1891, 30, 593) (separation from Ba and Sr). Eeadily soluble in a solution of Na 2 S 2 3 (separation from barium sulphate) (Diehl, J. pr., 1860, 79, 430). It is soluble in 60 parts hydrochloric acid, 6.12 per cent at 25°, and in 21 parts of the same * Dry CaO does not absorb dry CO, or S0 2 below 350°. (Veley, J. C, 1893, 63, 831). t Goldhammer, C. C, 1888, 708; Droeze, B., 1877, 10, 330; Boisbaudran, J.. Ch., 1874, (5), 3, HT Kohlrausoh and Bose, Z. phys. Ch., 1893, 12, 241 ; Raupenstraueh, M., 1885, 6, 563). 212 CALCIUM. §188, 6a. acid at 103° (Lunge, J. Soc. Ind., 1895 14, 31). The chromate is soluble in 214.3 parts water at 14° (Siewert, J., 1862, 149) ; in dilute alcohol it is rather more soluble (Fresenius, I. c, page 672); very readily soluble in acids including chromic acid. 6. Reactions, a. — The fixed alkali hydroxides precipitate solutions of calcium salts not having a degree of dilution beyond the solubility of the calcium hydroxide formed (5b), i. e. potassium hydroxide will form a precipitate with calcium sulphate since the sulphate requires less water for its solution than the hydroxide (5b and c) ; also the calcium hydroxide is less soluble in the alkaline solution than in pure water. Ammonium hydroxide does not precipitate calcium salts. The alkali carbonates pre- cipitate calcium carbonate, CaC0 3 , insoluble in water free from carbon dioxide, decomposed by acids. Calcium sulphate is completely trans- posed upon digestion with an alkali carbonate * (distinction from barium). Calcium hydroxide, Ca(0H) 2 , is used as a reagent for the detection of carbon dioxide (56 and §228, 8). b. — Alkali oxalates, as (NH 4 ) 2 C 2 4 , precipitate calcium oxalate, CaC 2 4 , from even dilute solutions of calcium salts. The precipitate is scarcely at all soluble in acetic or oxalic acids (separation of oxalic from phosphoric acid (§315), but is soluble in hydrochloric and nitric acids. The pre- cipitation is hastened by presence of ammonium hydroxide. Formed slowly, from very dilute solutions, the precipitate is crystalline, octahedral. If Sr or Ba are possibly present in the solution to be tested (qualitatively), an alkali sulphate must first be added, and after digesting a few minutes, if a precipitate appears, SrS0 4 , BaS0 4 , or, if the solution was concentrated, perhaps CaS0 4 , it is filtered out, and the oxalate then added to the filtrate. If a mixture of the salts of barium, strontium, and calcium in neutral or alkaline solution be treated with a mixture of (NH 4 ) 2 S0 4 and (NH 4 ) 2 C 2 4 , the barium and strontium are precipitated as sulphates and the calcium as the oxalate; separated from the barium and strontium on addition of hydrochloric acid (Sidersk}', Z., 1883, 22, 10; Bozomoletz, B., 1884, 17, 1058). A solution of calcium chloride is used as a reagent for the detec- tion of oxalic acid (§227, 8). In solutions of calcium salts containing a strong excess of ammonium chloride, potassium ferrocyanide precipitates the calcium (distinction from barium and strontium) (Baubigny, Bl., 1895, (3), 13, 326). * Here experiment shows that for equilibrium the S0 4 ions must be present in solution in large excess of CO, ions. With strontium also an excess of S0 4 ions is required, although not ao .great as in the case of calcium. For barium, however, equilibrium demands that the concen- tration of CO s ions exceed that of S0 4 . This condition is already fulfilled when an alkali car- bonate is added to BaSO« and therefore no change takes place in this case, while in the others the sulphate is transformed into carbonate. It is important to notice that the relative or ab- solute quantities of solid carbonate and sulphate present do not affect the equilibrium, which Is determined solely by the substances in solution (§57, 6e, footnote). §188, 9. CALCIUM. 213 c. — See 5c. d. — By the action of alkali phosphates, solutions of calcium are not distinguished from solutions of barium or strontium. e. — Pure sodium sulphide, Na. 2 S , gives an abundant precipitate with calcium salts; even with CaS0 4 . The precipitate is Ca(OH) 2 : CaCl 2 + 2Na2S + 2H 2 = Ca(OH).. + 2NaCl + 2NaHS . The acid sulphide, NaHS , does not precipitate calcium salts (Pelouze, A. Ch., 1866, (4), 7, 172). Alkali sulphites precipitate calcium sulphite, nearly insoluble in water, soluble in hydrochloric, nitric or sulphurous acid; barium and strontium salts act similarly. Sulphuric acid and soluble sulphates precipitate calcium salts as CaS0 4 , distinguished from barium by its solubility in water and in hydrochloric^ .acid; from barium and strontium by its solubility in ammonium sulphate (5c). A water solution of calcium sulphate is used to detect strontium after barium has been removed as a chromate. Obviously a solution of strontium sulphate will not precipitate calcium salts. f. — Calcium chloride, fused, is much used as a drying agent for solids, liquids and gases. Chlorinated lime, calcium hypochlorite, Ca(C10) 2 (Kingzett, J. C, 1875, 28, 404), is much used as a bleaching agent and as a disinfectant, g. — Neutral or ammoniacai solutions of arsenites form a precipitate with calcium «alts (distinction from barium). A solution of calcium salts including solu- tions of calcium sulphate in ammoniacai solution is precipitated by arsenic acid as CaMTHjAsO., (distinction from strontium after the addition of sulphuric acid) (Bloxam, C. N., 1886, 54, 16). li. — Normal chromates, as K 2 Cr0 4 , precipitate solutions of calcium salts as calcium chromate, CaCr0 4 , yellow, provided the solution be not too dilute (5c). The precipitate is readily soluble in acids and is not formed with acid chro- mates as K 2 Cr 2 7 (separation from barium), i. — Fluosilicic acid does not precipitate calcium salts even in the presence of equal parts of alcohol (separa- tion from barium). 7. Ignition. — Calcium sulphate, CaS0 4 .2H 2 , gypsum , loses its water of crystallization at 80° and becomes the anhydrous sulphate, CaSO t , plaster of Paris; which on being moistened forms the crystalline CaS0 4 .2H 2 , expands and " sets." Calcium carbonate, limestone, when heated (burned) loses carbon dioxide and becomes lime, CaO . Compounds of calcium, preferably the chloride, render the flame yellowish red,. The presence of strontium or barium obscures this reaction, but a mixture containing calcium and barium, moistened with hydrochloric acid, gives the calcium color on its first introduction to the flame. The spectrum of calcium is distinguished by the bright green line, Ca /3, and the intensely bright •orange line, Ca a, near the red end of the spectrum. 8. Detection. — Calcium is separated in analysis from the metals of the •other groups and from barium, with strontium, as described at §187, 8. A portion of the solution of strontium and calcium acetate is boiled with potassium sulphate ; after standing for some time (ten minutes), the filtrate is tested with ammonium oxalate. A white precipitate insoluble in the acetic acid present, but soluble in hydrochloric acid is evidence of the presence of calcium. The flame test (7) is confirmatory. 9. Estimation.— Calcium is weighed as an oxide, carbonate, or sulphate. The ■carbonate is obtained by precipitating as oxalate, and gently igniting the dried precipitate; higher ignition changes the carbonate to the oxide. The sulphate is precipitated in a mixture of two parts of alcohol to one of the solution. The hydroxide and carbonate may be determined by alkalimetry. Calcium may be separated from barium and strontium by the solution of its nitrate in amy! g 14 MAGNESIUM. §189, 1. alcohol (5c). The best method of separation from strontium is to treat the nitrates with a mixture of equal volumes of alcohol and ether. The calcium nitrate dissolves, but not more than one part in 60,000 of the strontium is found in the solution (§195). In the presence of iron, aluminum and phos- phoric acid, calcium is best precipitated as an oxalate in the presence of citric acid (Passon, Z. angew., 1898, 776). See also 9, §186 and §187. §189. Magnesium. Mg = 24.3 . Valence two. 1. Properties.— Specific gravity, 1.75 (Deville and Caron, A. Ch., 1863, (3), 67, 346) ; melting point, a little below 800°, does not appear to be volatile (Meyer, B., 1887, 20, 497). A white, hard, malleable and ductile metal; not acted upon by water or alkalis at ordinary temperature and only slightly at 100° (Ballo, B., 1883, 16, 694). When heated in air or in oxygen it burns with incandescence to MgO . It combines directly when heated in contact with TSS , P , As , S and CI . It forms alloys with Hg and Sn , forming compounds which decom- pose water. 2. Occurrence. — Magnesite, MgC0 8 ; dolomite, CaMg(C0 8 ) 2 ; brucite, ]ffig(6H) 2 ; epsom salts, MgS0 4 .7H 2 0; and combined with other metals in a great variety of minerals. 3. Preparation. — (i) By electrolysis of the chloride or sulphate (Bunsen, A., 1852, 82, 137). (2) By ignition of the chloride with sodium or potassium (WShler, A., 1857, 101, 563). (3) Mg 2 Pe(CN) 6 is ignited with Na 2 C0 8 , and this product ignited with zinc (Lanterbronn, German Patent No. 39,915). 4. Oxide and Hydroxide. — Only one oxide of magnesium, MgO , is known with certainty. Formed by burning the metal in the air, and by action of heat upon the hydroxide, carbonate, nitrate, sulphate, oxalate and other mag- nesium salts decomposed by heat. The corresponding hydroxide, Mg(0H) 2 , is formed by precipitating magnesium salts with the fixed alkalis. 5. Solubilities. — a. — Metal. — Magnesium is soluble in acids including carbonic acid, evolving hydrogen: Mg + C0 2 + H 2 = MgC0 3 + H 2 (Ballo, B., 1882, 15, 3003) : it is also attacked by the acid alkali carbonates, as NaHC0 3 , to form MgC0 3 , Na 2 C0 3 and H (Ballo, I. c). Soluble in ammonium salts: Mg + 3NH 4 C1 = NH 4 MgCl 3 + 2NH 3 + H 2 . With the halogens it acts tardily (Wanklyn and Chapman, J. C, 1866, 19, 141). b. — Oxide and hydroxide. — Insoluble in water, soluble in acids. Mg(0H) 2 is soluble in 111,111 parts of water at 18° (Kohlrausch and Rose, Zeit. phys. Ch., 1893, 12, 241). In contact with water the oxide is slowly changed to the hydroxide, Mg(0H) 2 , and absorbs C0 2 from the air. Sol- uble in ammonium salts:* Mg(0H) 2 + 3NH 4 C1 = NH 4 MgCl 3 + 2NH 4 0H. c. — Salts. — The chloride, bromide, iodide, chlorate, nitrate, and acetate (4 aq) are deliquescent', the sulphate (7 aq) slightly efflorescent. The carbonate, phosphate, borate, arsenite, and arsenate are insoluble in water; the sulphite, oxalate, and chromate soluble; the tartrate sparingly soluble. The carbonate is soluble; the phosphate, arsenite, and arsenate are insoluble in excess of ammonium salts. 6. Keactions. a. — The fixed alkali hydroxides and the hydroxides of barium, strontium and calcium precipitate magnesium hydroxide, Mg(0H) 2 , * The conditions here are the same as in the ease of Mn(OH)„ §134, 6a, footnote. §189, 7. MAGNESIUM. 215 white, gelatinous, from solutions of magnesium salts; insoluble in excess of the reagent but readily soluble in ammonium salts : Mg(0H) 2 -f- 3NH 4 C1 = MgCl 2 .NH 4 Cl + 2NH 4 0H . With ammonium hydroxide but half of the magnesium is precipitated, the remainder being held in solution by the ammonium salt formed in the reaction: 2MgS0 4 -4- 2NH 4 0H = Mg(0H) 2 + (NH 4 ) 2 Mg(S0 4 ) 2 (Eheineck, Dingl, 1871, 202, 268). The fixed alkali carbonates precipitate basic magnesium carbonate,- Mg 4 (0H) 2 r (C0 3 ) 3 , variable to Mg 6 (0H) 2 (C0 3 ) 4 : 4MgS0 4 + 4Na 2 C0 3 + H 2 = Mg 4 (0H) 2 (C0 3 ) 3 + 4Na 2 S0 4 + C0 2 . If the above reaction takes place in the cold the carbon dioxide combines with a portion of the magnesium carbonate to form a soluble acid magnesium carbonate: 5MgS0 4 -f- 5Na 2 C0 3 + 2H 2 = Mg 4 (0H) 2 (C0 3 ) 3 + MgH 2 (C0 3 ) 2 + 5Na 2 S0 4 . On boiling, the acid carbonate is decomposed with escape of C0 2 . Ammonium carbonate does not precipitate magnesium salts, as a soluble double salt is at once formed. Acid fixed alkali carbonates, as NaHC0 3 , do not precipi- tate magnesium salts in the cold; but upon boiling, C0 2 is evolved and the carbonate is precipitated (Engel, A. Ch., 1886, (6), 7, 260). 6. — Soluble oxalates do not precipitate solutions of magnesium salts, as they form soluble double oxalates. If to the solution of double oxalates, preferably magnesium ammonium oxalate, an equal volume of 80 per cent acetic acid be added, the magnesium is precipitated as the oxalate (separation from potas- sium or sodium (Classen, Z., 1879, 18, 373). d. — Alkali phosphates — as Na 2 HP0 4 — precipitate magnesium phosphate, MgHP0 4 , if the solution be not very dilute. But even in very dilute solutions, by the further addition of ammonium hydroxide (and NH 4 C1), a crystalline precipitate is slowly formed, magnesium ammonium phosphate — MgNH 4 P0 4 . Stirring with a glass rod against the side of the test-tube promotes the precipitation. The addition of ammonium chloride, in this test, prevents formation of any precipitate of magnesium hydroxide (56). The precipitate dissolves in 13,497 parts of water at 23° (Ebermayer, J. pr., 1853, 60, 41); almost absolutely insoluble in water containing ammonium hydroxide and ammonium chloride (Kubel, Z., 1869, 8, 125). e. — Magnesium sulphide is decomposed by water, and magnesium salts are not precipitated by hydrosulphuric acid or ammonium sulphide; but MgO + H 2 (1-10) absorbs H 2 S , forming in solution MgH 2 S 2 , which readily gives off H 2 S upon boiling (a very satisfactory method of preparing H 2 S absolutely arsenic free) (Divers and Shmidzu, J. C, 1884, 45, 699). Normal sodium or potassium sulphide precipitates solutions of magnesium salts as the hydroxide with formation of an acid alkali sulphide: MgS0 4 + 2Na 2 S + 2H 2 = Mg(OH), + Na 2 S0 4 + 2NaHS (Pelouze, A. Ch., 1866, (4), 7, 172). Sulphuric acid and soluble sulphates do not precipitate solutions of magnesium salts (distinction from Ba , Sr and Ca). f. — Magnesium chloride, in solution, evaporated on the water bath evolves hydrochloric acid (7). g. — Soluble arsenates precipitate magnesium salts in deportment similar to the corresponding phosphates. 7. Ignition. — Magnesium ammonium phosphate when ignited loses ammonia 216 MAGNESIUM. §189, 8. and water, and becomes the pyrophosphate: 2Mg\N , H 4 P0 1 = Mg 2 P 2 0, + H 2 + 2NH 3 . The carbonate loses C0 2 and becomes MgO . In dry air magnesium chloride may be ignited without decomposition, but in the presence of steam MgO and HC1 are formed: MgCl 2 + H 2 = MgO + 2HCI; a technical method for preparing HC1 (Heumann, A., 1877, 184, 227). 8. Detection. — If sufficient ammonium salts have been used, the mag- nesium will be in the filtrate from the precipitated carbonates of barium, strontium and calcium. From a portion of this filtrate the magnesium is precipitated as the white magnesium ammonium-phosphate, MgHH 4 P0 4 , byNa 2 HP0 4 . 9. Estimation.- — After removal of other non-alkali metals, magnesium is pre- cipitated as MgNH 4 P0 4 , then changed by ignition to Mg 2 P 2 7 (magnesium pyrophosphate) and weighed as such. Separated as MgCl 2 from KC1 and NaCl by solution in amyl alcohol, evaporated with H 2 S0 4 and weighed as MgS0 4 (Riggs, Am. 8., 1892, 44, 103). It is estimated volumetrically by precipitation as MgNH 4 P0 4 , drying at about 50° until all free NH,OH is removed. An excess of standard acid is then added and at once titrated back with standard fixed alkali, using methyl orange as an indicator (Handy, J. Am. 8oe., 1900, 22, 31). 10. Oxidation. — Magnesium is a powerful reducer; ignited with the oxides or carbonates of the following elements magnesium oxide is formed and the corresponding element is liberated : Ag , Hg , Pt , Sn *, B , Al , Th, CJ, Si, Pb, PJ, As, Sb, Bi, Cr, Mo, Mn, Fe, Co, Hi, Cu, Cd , Zn , Grl , Ba , Sr , Ca , Eh , K , Ha , and Li . In some cases the reaction takes place with explosive violence. Prom their corresponding salts in neutral solution Mg precipitates Se , Te , As , Sb , Bi , Sn , Zn f , Cd , Pb , Tl , Th , Cu , Ag , Mn f, Fe f, Co , Hi , An , Pt , and Pd (Scheibler, B., 1870, 3, 295; Villiers and Borg, C. r., 1893, 116, 1534). * Winkler, B., 1890, 23, 44, 130 and 172 ; 1891, 24, 893. + Kern, C. JW., 1876, 33, 113 and 336. t Seubert and Schmidt, A., 1893, 367, 318. §190 ANALYSIS OF THE CALCIUM GROUP. 217 2 ^ ax> T2 Co C5 03 9 CO 03 Pj s ^ I CO o 2 O rH 1 Is, rH eg W coo rQ B pi 8 H 03 o CO o ° H c8 ^ J3 Ph 03 <3 03 Jh p «h bO 9 ^ B 1-3 o *-t 2 -§ § Ho g M O . « O cos H M o s* -d >> rB a 3 • iH a o a a a *a 13 a /— \ oo w § H — i d o O ca *• to 5 •= -d i 3 43 a.-; 2 -til ft W 'S cm - W 03 B d w "ft*K bd "S _ ^ P4-H •HrH O oSfl« tp « - fj cd „43-h - ■d ^Qj ~2 5 a **H -P. „-r= .. o> Sao ,2"" jfl 10 93 2 "8 2 -t- o -3 «H +J ** ° 03 n u o ft -J3 •« 5 rB r/3 C3 l-H "3 03 II a.& o o U 03 si ^ .2 ft 3 C3 a 2 'xk be TO O ft 2 CO ^ - 13 B ca _o ^+3 B£ m O o>fe ■f -^ ft fl ftfefl S 0) g * ri a c S 5 a ca^ .2ra 5-2t- -H "q m r* *& CO 5 ^ a .s ™ S ■- a ^ 03 03 'O-S*' 03O .1 m S "§ a -e a ° g .s ft.ri 03 a C3 |» O cars ca u 53 h "> ftca"^ h a oT o ^ -So **- §'-sa» ^■2 a-H ■So,SS t-1 T3 a ca t-Oi 505,0, >?-d a ca a ca 1O0J 1-03 rH t-I C09C09 ^T3 a 03 -l-> ca .-a .& C3 0) rH a.S >> " rt 5 o 5 s. 02 T ftca [0 •a Tj-r tttf O ft ej ca o ft ca r-l .„_, .rH 03 rH ftr^ .S 0) mil fH £ +J rB r5 P "3r2-2 < S oJa'gS.S >* a. "lap »Or . rB to BCD IIS ca ft^< 01 a ca mot t-0> IS 218 DIRECTIONS FOB ANALYSIS WITH NOTES. §191. DIRECTIONS FOR ANALYSIS OF THE METALS OF THE CALCIUM GeOTJP. (The Alkaline Earths.) §191. Manipulation. — To the filtrate from the fourth group in which H 2 S (§192, 1) gives no .precipitate (§138) add NH 4 0H and ammonium carbonate as long as a precipitate is formed : BaCl 2 + (11114)2003 = BaC0 3 + 2NH 4 C1 . Digest with warming, filter and wash. The filtrate should he tested again with ammonium carbonate and if no precipitate is formed it is set aside to be tested for magnesium and the alkali metals (§§193 and 211). The well washed white precipitate is dissolved in acetic acid, using as little as possible : SrC0 3 + 2HC 2 H 3 2 = Sr(C 2 H 3 2 ) 2 + C0 2 + H 2 . To a small portion of the acetic acid solution add a drop of K 2 Cr 2 7 ; if a precipitate — BaCr0 4 — is obtained, the K 2 Cr 2 7 must be added to the whole solution: 2Ba(C 2 H 3 2 ) + K 2 Cr 2 7 + H 2 = 2BaCr0 4 + 2KC 2 H 3 2 + 2HC 2 H 3 2 . Filter, wash the precipitate, dissolve it in HC1 and pre- cipitate the barium as barium sulphate, with a drop of sulphuric acid. To the filtrate from the barium chromate add NH 4 0H and (NH 4 ) 2 C0 3 , warm, filter, and wash. Dissolve the white precipitates of SrC0 3 and CaC0 3 in acetic acid and divide the solution into two portions. Portion 1. — For Strontium. — With a platinum wire obtain the flame test, crimson for strontium; calcium interferes (7, §§187, 188 and 205). Add a solution of calcium sulphate and boil; set aside for about ten min- utes. A precipitate — SrS0 4 — indicates strontium. This SrS0 4 may be moistened with HC1 and the crimson flame test obtained. Portion 2. — For Calcium. — Add a solution of potassium sulphate, boil, and set aside for ten minutes. Filter (to remove any strontium that may be present; also a portion of the calcium may be precipitated, §188, 6e) and add ammonium oxalate to the filtrate. Dissolve the precipitate in HC1 . A white precipitate — CaC 2 4 — insoluble in acetic acid by its forma- tion in that solution, and soluble in HC1 is proof of the presence of calcium. §192. Notes.— 1. The failure of (NH,) 2 S (or H 2 S in presence of NH 4 OH) to form a precipitate with solutions of the alkaline earths and of the alkalis, marks a sharp separation of these metals from the metals of the preceding groups. 2. Do not boil after the addition of ammonium carbonate, as this will drive off ammonium hydroxide and carbonate, increasing the solubility of the CaCO a (note 3 and §178). 8. The precipitation, of barium, strontium and calcium by ammonium car- bonate in the presence of ammonium chloride, is not as complete as would be desirable in very delicate analyses. The carbonates of barium, strontium and calcium are all slightly soluble in ammonium chloride solution; and while the prescribed addition of ammonium hydroxide, and excess of ammonium car- bonate, greatly reduces the solubility of the precipitated carbonates, yet even with these the precipitation is not absolute, though more nearly so with strontium than with barium and calcium. Thus, in quantitative analyses, if §194,4- DIRECTIONS FOR ANALYSIS WITH NOTES. 219 barium and calcium are precipitated as carbonates, it must be done in the absence of ammonium chloride or sulphate, and the precipitate washed with water containing ammonium hydroxide. 4. If barium be absent, as evidenced by the failure to obtain a precipitate with K 2 Cr 2 0, , the solution may at once be divided into two portions to test for strontium and calcium. 5. With care the reprecipitation by ammonium carbonate, for the separa- tion from the excess of K 2 Cr 2 7 , may be neglected and the nitrate from the barium, yellow, at once divided into two portions and tested for Sr and Ca . Eeprecipitation always causes the loss of some of the metals, due to the solu- bility of the carbonates in the ammonium acetate formed. Qn the other hand, traces may escape observation in the yellow chromate solution. 6. Before reprecipitation with (NH 4 ) 2 CO s , an excess of ammonium hydroxide should be added to prevent the liberation of C0 2 when the ammonium car- bonate is added. 7. Strontium sulphate is so sparingly soluble in water (§187, 5c) that its precipitation by CaS0 4 (or other sulphates in absence of Ca) is sufficiently delicate to detect very small amounts of that metal. However, it is sufficiently soluble in water to serve as a valuable reagent to detect the presence of traces of barium. Obviously SrSO, will not precipitate solutions of calcium salts. Solutions of strontium and barium salts (except SrS0 4 ) are all precipitated by CaS0 4 . The presence of excess of calcium salts lessens the delicacy of the precipitation of strontium salts by calcium sulphate. 8. — In very dilute solutions the sulphates of the alkaline earths are not precipitated rapidly. Time should be allowed for the complete precipitation. Boiling and evaporation facilitates the reaction. 9. It should be noticed that the test for calcium as an oxalate is made upon that portion of the calcium not removed by K 2 SO,; or in other words upon a solution of CaS0 4 (1-500). A solution of SrS0 4 (1-10,000) may be present but is not precipitated by (1TH 4 ) 2 C 2 4 . The presence of a great excess of (N , H 4 ) 2 S0 4 prevents the precipitation of traces of calcium salts by (NH 4 ) 2 C 2 4 . §193. Manipulation. — To a portion of the filtrate from the carbonates of Ba , Sr , and Ca add a drop or two of (NH 4 ) 2 S0 4 and then a few drops of (NH 4 ) 2 C 2 4 ; filter if a precipitate is obtained and test the nitrate for Mg with Na 2 HP0 4 . A white precipitate — MgNH 4 P0 4 — is evidence of the presence of magnesium. The other portion of the nitrate from the car- bonates of Ba , Sr , and Ca is reserved to be tested for the alkali metals (§211). §194. Notes. — 1. By some, magnesium is classed in the last or alkali group instead of in the alkaline earth group. It is not precipitated by the (NH 4 ) 2 C0 3 , yet in the general properties of its salts it is so closely related to Ba , Sr and Ca , that it is much better regarded as a subdivision of that group than as belonging to the alkali group (§175 and ff.). 2. Traces of Ba, Sr and Ca may remain in solution after adding (NH 4 ) 2 CO a and warming; due to the solvent action of the ammonium salts present. To prevent these traces giving a test for magnesium with Na 2 HP0 4 , a drop or two of (HH 4 ) 2 S0 4 is added to remove barium or strontium and a few drops of (NH 4 ) 2 C 2 4 to remove calcium. The precipitate (if any forms) is removed by filtration, before the Na 2 HP0 4 is added. 3. The precipitate of MglTH 4 P0 4 does not always form rapidly if only small amounts of Mg are present, and the solution should be allowed to stand. Bubbing the sides of the test tube with a glass stirringfrod promotes the pre- cipitation. 4. The precipitation of Mg as MgNH 4 P0 4 is fairly delicate (1-71,492) (Kissel, Z., 1869, 8, 173) ; but not at all characteristic, as the phosphates of nearly all the metals are white and insoluble in water. Hence the reliability of this test for 218 DIRECTIONS FOR ANALYSIS WITH NOTES. §191. dlbections foe analysis of the metals of the calcium group. (The Alkaline Eaeths.) §191. Manipulation.— To the filtrate from the fourth group in which H 2 S (§192, 1) gives no .precipitate (§138) add NH 4 0H and ammonium carbonate as long as a precipitate is formed: BaCl 2 + (NH 4 ) 2 C0 3 = BaC0 3 + 2NH 4 C1 . Digest with warming, filter and wash. The filtrate should be tested again with ammonium carbonate and if no precipitate is formed it is set aside to be tested for magnesium and the alkali metals (§§193 and 211). The well washed white precipitate is dissolved in acetic acid, using as little as possible: SrC0 3 + 2HC 2 H 3 2 = Sr(C 2 H 3 2 ) 2 + C0 2 + H 2 . To a small portion of the acetic acid solution add a drop of K 2 Cr 2 7 ; if a precipitate — BaCr0 4 — is obtained, the K 2 Cr 2 7 must be added to the whole solution: 2Ba(C 2 H 3 2 ) + K 2 Cr 2 7 + H 2 = 2BaCr0 4 + 2KC 2 H 3 2 -f- 2HC 2 H 3 2 . Filter, wash the precipitate, dissolve it in HC1 and pre- cipitate the barium as barium sulphate, with a drop of sulphuric acid. To the filtrate from the barium chromate add NH 4 0H and (NH 4 ) 2 C0 3 , warm, filter, and wash. Dissolve the white precipitates of SrC0 3 and CaC0 3 in acetic acid and divide the solution into two portions. Portion 1. — For Strontium. — With a platinum wire obtain the flame test, crimson for strontium; calcium interferes (7, §§187, 188 and 205). Add a solution of calcium sulphate and boil; set aside for about ten min- utes. A precipitate — SrS0 4 — indicates strontium. This SrS0 4 may be moistened with HC1 and the crimson flame test obtained. Portion 2. — For Calcium. — Add a solution of potassium sulphate, boil, and set aside for ten minutes. Filter (to remove any strontium that may be present; also a portion of the calcium may be precipitated, §188, 6e) and add ammonium oxalate to the filtrate. Dissolve the precipitate in HC1 . A white precipitate — CaC 2 4 — insoluble in acetic acid by its forma- tion in that solution, and soluble in HC1 is proof of the presence of calcium. §192. Notes.— 1. The failure of (NH,),S (or H 2 S in presence of NH 4 OH) to form a precipitate with solutions of the alkaline earths and of the alkalis, marks a sharp separation of these metals from the metals of the preceding groups. 2. Do not boil after the addition of ammonium carbonate, as this will drive off ammonium hydroxide and carbonate, increasing' the solubility of the CaCO, (note 3 and §178). 3. The precipitation, of barium, strontium and calcium by ammonium car- bonate in the presence of ammonium chloride, is not as complete as would be desirable in very delicate analyses. The carbonates of barium, strontium and calcium are all slightly soluble in ammonium chloride solution; and while the prescribed addition of ammonium hydroxide, and excess of ammonium car- bonate, greatly reduces the solubility of the precipitated carbonates, yet even with these the precipitation is not absolute, though more nearly so with strontium than with barium and calcium. Thus, in quantitative analyses, if §194,4. DIRECTIONS FOR ANALYSIS WITH NOTES. 219 barium and calcium are precipitated as carbonates, it must be done in the absence of ammonium chloride or sulphate, and the precipitate washed with water containing ammonium hydroxide. 4. If barium be absent, as evidenced by the failure to obtain a precipitate with K 2 Cr 2 T , the solution may at once be divided into two portions to test for strontium and calcium. 5. With care the reprecipitation by ammonium carbonate, for the separa- tion from the excess of K 2 Cr 2 7 , may be neglected and the filtrate from the barium, yellow, at once divided into two portions and tested for Sr and Ca . Eeprecipitation always causes the loss of some of the metals, due to the solu- bility of the carbonates in the ammonium acetate formed. On the other hand, traces may escape observation in the yellow chromate solution. 6. Before reprecipitation with (NH 4 ) 2 C0 3 , an excess of ammonium hydroxide should be added to prevent the liberation of C0 2 when the ammonium car- bonate is added. 7. Strontium sulphate is so sparingly soluble in water (§187, 5c) that its precipitation by CaSO, (or other sulphates in absence of Ca) is sufficiently delicate to detect very small amounts of that metal. However, it is sufficiently soluble in water to serve as a valuable reagent to detect the presence of traces of barium. Obviously SrS0 4 will not precipitate solutions of calcium salts. Solutions of strontium and barium salts (except SrS0 4 ) are all precipitated hy CaS0 4 . The presence of excess of calcium salts lessens the delicacy of the precipitation of strontium salts by calcium sulphate. 8. — In very dilute solutions the sulphates of the alkaline earths are not precipitated rapidly. Time should be allowed for the complete precipitation. Boiling and evaporation facilitates the reaction. 9. It should be noticed that the test for calcium as an oxalate is made upon that portion of the calcium not removed by K 2 S0 4 ; or in other words upon a solution of CaS0 4 (1-500). A solution of SrS0 4 (1-10,000) may be present but is not precipitated by (NH 4 ) 2 C 2 4 . The presence of a great excess of <1TH 4 ) 2 S0 4 prevents the precipitation of traces of calcium salts by (NH 4 ) 2 C 2 4 . §193. Manipulation. — To a portion of the filtrate from the carbonates of Ba , Sr , and Ca add a drop or two of (NH 4 ) 2 S0 4 and then a few drops of (NH 4 ) 2 C 2 4 ; filter if a precipitate is obtained and test the filtrate for Mg with Na 2 HP0 4 . A white precipitate — MgNH 4 P0 4 — is evidence of the presence of magnesium. The other portion of the filtrate from the car- bonates of Ba , Sr , and Ca is reserved to be tested for the alkali metals (§211). §194. Notes. — 1. By some, magnesium is classed in the last or alkali group instead of in the alkaline earth group. It is not precipitated by the (ini 4 ) 2 COs , yet in the general properties of its salts it is so closely related to Ba , Sr and Ca , that it is much better regarded as a subdivision of that group than as belonging to the alkali group (§175 and ff.). 2. Traces of Ba, Sr and Ca may remain in solution after adding (1TH 4 ) 2 C0 3 and warming; due to the solvent action of the ammonium salts present. To prevent these traces giving a test for magnesium with Na 2 HP0 4 , a drop or two of (NH 4 ) 2 S0 4 is added to remove barium or strontium and a few drops of (NH 4 ) 2 C 2 4 to remove calcium. The precipitate (if any forms) is removed by nitration, before the Ma 2 HP0 4 is added. 3. The precipitate of MglTH 4 P0 4 does not always form rapidly if only small amounts of Mg are present, and the solution should be allowed to stand. Bubbing the sides of the test tube with a glass stirring 'rod promotes the pre- cipitation. Jf. The precipitation of Mg as MgNH 4 P0 4 is fairly delicate (1-71,492) (Kissel, Z., 1869, 8, 173) ; but not at all characteristic, as the phosphates of nearly all the metals are white and insoluble in water. Hence the reliability of this test for 220 SEPARATION OF BARIUM, STRONTIUM AND CALCIUM. §194, 5. magnesium depends upon the rigid exclusion of the other metals (not alkalis) by the previous processes of analysis. 5. Lithium phosphate is not readily soluble in water or ammonium salts and may give a test for magnesium. See §210, 6d. §195. The unlike solubilities in alcohol, of the chlorides and nitrates of barium, strontium and calcium enable us to separate them quite closely by absolute alcohol, and approximately by " strong alcohol," as follows : Dissolve the carbonate precipitate in HC1 , evaporate to dryness on the water-bath, rub the residue to a fine powder in the evaporating dish, and digest it with alcohol. Filter through a small filter, and wash with alcohol (5c, §§186, 187 and 188). Residue: BaCL . Dissolve in water, test with CaSOj , SrS0 4 , X 2 Cr 2 7 , etc. Filtrate: SrCl 2 and CaCl 2 . Evaporate to dryness, dissolve in water, change to nitrates by precipitating with (1TH4) 2 C0 S , wash- ing and "dissolving in H2TO s . Evaporate the nitrates to dryness, powder, digest with alcohol* filter and wash with alcohol (or digest and wash with equal volumes of alcohol and ether). Residue: Sr(NO s ) 2 . Precipitation by CaS0 4 in water solution; flame test, etc. nitrate: Ca(lTO,) 2 . Precipitation by H 2 S0 4 in alcohol solution, by (NHJjCjO, , etc. Or, the alcoholic filtrate of SrCl 2 and CaCl 2 may be precipitated with (a drop of) sulphuric acid, the precipitate filtered out and digested with solution of (NH 4 ) 2 S0 4 and a little NH 4 0H. Residue, SrS0 4 . Solution contains CaS0 4 , precipitable by oxalates. §196. If the alkaline earth metals are present in the original material as phosphates, or in mixtures such that the treatment for solution will bring them in contact with phosphoric acid; the process of analysis must be modified. One of the methods given under analysis of third and fourth group metals in presence of phosphates (§145 and //.) must be employed. §197. The presence of oxalates will also interfere, necessitating the evaporation and ignition to decompose the oxalic acid (§151). * Instead of alcohol the residue of the nitrates may be boiled with amyl alcohol. Calcium nitrate is dissolved making a complete separation from the strontium nitrate (8188, 6c), §200. THE ALKALI GROUP. 221 The Alkali Geoup (Sixth Gkoup). Potassium. K = 39.11. Caesium. Cs = 132.9. Sodium. Na = 23.05. Rubidium. Rb = 85.4. Ammonium. (NH 4 )'. Lithium. Li = 7.03. §198. The metals of the alkalis are highly combustible, oxidizing quickly in the air, displacing the hydrogen of water even more rapidly than zinc or iron displaces the hydrogen of acids, and displacing non-alkali metals from their oxides and salts. As elements they are very strong reducing agents, while their compounds are very stable, and not liable to either re- duction or oxidation by ordinary means. The five metals, Cs , Rb , K , Na, li, present a gradation of electro-positive or basic power, caBsium being strongest, and the others decreasing in the order of their atomic weights, lithium decomposing water with less violence than the others. Their specific gravities decrease,* their fusing points rise, and as carbon- ates their solubilities lessen, in the same order. In solubility of the phos- phate, also, lithium approaches the character of an alkaline earth (§6). Ammonium is the basal radical of ammonium salts, and as such has many of the characteristics of an alkali metal. The water solution of the gas ammonia, NH 3 (an anhydride), from analogy is supposed to contain ammonium hydroxide, NH 4 0H, known as the volatile alkali. Potassium and sodium hydroxides are the fixed alkalis in common use. §199. The alkalis are very soluble in water, and all the important salts of the alkali metals (including NH 4 ) are soluble in water, not excepting their carbonates, phosphates (except lithium), and silicates; while all other metals form hydroxides or oxides, either insoluble or sparingly soluble, and carbonates, phosphates, silicates, and certain other salts quite insoluble in water. Their compounds being nearly all soluble, the alkali metals are not pre* cipitated by ordinary reagents, and, with few exceptions, their salts do not precipitate each other. In analysis, they are mostly separated from other metals by non-precipitation. §200. In accordance with the insolubility in water of the non-alkali hydroxides and oxides, the alkali hydroxides precipitate all non-alkali metals, except that ammonium hydroxide does not precipitate barium, strontium, and calcium. These precipitates are hydroxides, except those of mercury, silver, and antimony. But certain of the non-alkali hydroxides and oxides, though insoluble in water, dissolve in solutions of alkalis; hence* when added in excess, the alkalis redissolve the precipitates they at first pro- duce with salts of certain metals, viz. : the hydroxides of Pb , Sn , Sb (oxide)* * Except those of potassium (0.875) and sodium (0.9735). 222 POTASSIUM. §201. Zn, Al, and Cr dissolve in the fixed alkalis; and oxide of Ag and hy- droxides of Cu , Cd , Zn , Co , and Ni dissolve in the volatile alkali. §201. Solutions of the alkalis are caustic to the taste and touch, and turn red litmus blue; also, the carbonates, acid carbonates, normal and dibasic phosphates, and some other salts of the alkali metals, give the "alkaline reaction" with test papers. Sodium nitroferricyanide, with hydrogen sulphide, gives a delicate reaction for the alkali hydroxides (§207, 66). §202. The hydroxides and normal carbonates of the alkali metals are not decomposed by heat alone (as are those of other metals), and these metals form the only acid carbonates obtained in the solid state. §203. The fixed alkalis, likewise many of their salts, melt on platinum foil in the flame, and slowly vaporize at a bright red heat. All salts of ammonium, by a careful evaporation of their solutions on platinum foil, may be obtained in a solid residue, which rapidly vaporizes, wholly or partly, below a red heat (distinction from fixed alkali metals). §204. The hydroxides of the fixed alkali metals, and those of their salts most volatile at a red heat, preferably their chlorides, impart strongly characteristic colors to a non-luminous flame, and give well-defined spectra with the spectroscope. §205. Potassium. K = 39.11 . Valence one. 1. Properties. — Specific gravity, 0.875 at 13° (Baumhauer, B., 1873, 6, 655). Meltmg point, 62.1° (Hagen, C. C, 1883, 129). Boiling point, 719° to 731° (Car- nelley and Williams, B., 1879, 12, 1360); 667° (Perman, J. C, 1889, 55, 328). Silver-white metal- with a bluish tinge. At ordinary temperature of a wax-like consistency, ductile and malleable; at 0° it is brittle. It is harder than Na and is scratched by Li , Pb , Ca and Sr . The glowing vapor is a very beautiful intense violet (Dudley, Am., 1892, 14, 185). It is next to caesium and rubidium, the most electro-positive of all metals, remains unchanged in dry air, oxidizes rapidly in moist air, and decomposes water with great violence, evolving hydrogen, burning with a violet flame. At a red heat CO and C0 2 are decomposed, at a white heat the reverse action takes place. Liquid chlorine does not attack dry potassium (Gautier and Charpy, C. r., 1891, 113, 597). Acids attack it violently, evolving hydrogen. 2. Occurrence. — Very widely distributed as a portion of many silicates. In sea water in small amount as KC1 . In numerous combinations in the large salt deposits, especially at Stassfurt; e. g., carnallite, KCl.MgCl., + 6H 2 0; kainite, K 2 S0 4 .M:gS0 4 .:MgCI 2 + 6H 2 , etc. As an important constituent of many plants — grape, potato, sugar-beet, tobacco, fumaria, rumex, oxalis, etc. 3. Preparation. — (i) By reduction of the carbonate with carbon. (2) By electrolysis of the hydroxide (Horning and Kasemeyer, B., 1889, 22, 277c; Castner, B., 1892, 25, 179c). (3) By reduction of K 2 CO g or KOH with iron car- bide: 6KOH + 2FeC 2 = 6K + 2Fe + 2CO + 2C0 2 + 3H 2 (Castner, C. N., 1886, 54, 218). (4) By reduction of the carbonate or hydroxide with Pe or Mg '(Winkler, B., 1890, 23, 44). 4. Oxides and Hydroxide. — Potassium owide,* K 2 , is prepared by carefully * The existence of the oxides M' a O of K, Na and Rb is disputed (Erdmann and Koethner, A., 1896, 294, 55). §205, 6b. POTASSIUM. 223 heating potassium with the necessary amount of oxygen (air) (Kuhnemann, C. C, 1863, 491); also by heating K 2 4 with a mixture of K and Ag (Beketoff, C. C, 1881, 643). It is a hard, gray mass, melting above a red heat. Water changes it to KOH with generation of much heat. Potassium hydroxide, KOH, is formed by treating K or K 2 with water; by boiling a solution of K 2 C0 8 with Ba, Sr or Ca oxides; by heating K.CO, with Pe 2 O s to a red heat and decomposing the potassium ferrate with water (Ellershausen, G. C, 1891, (1), 1047; (2), 399). Pure water-free KOH is a white, hard, brittle mass, melting at a red heat. It dissolves in water with generation of much heat. Potassium superoxide, K 2 4 , is formed when K is heated in contact with abundance of air (Harcourt, J. C, 1862, 14, 267); also by bringing K in contact with KN0 3 heated until it begins to evolve (Bolton, C. N., 1886, 53, 289). It is an amor- phous powder of the color of lead chromate. Upon ignition in a silver dish oxygen is evolved and K 2 and Ag 2 formed (Harcourt, I. a). Moist air or water decomposes it with evolution of oxygen. It is a powerful oxidizing agent, oxidizing S° to Svi , P° to Pv , K , As , Sb , Sn , Zn , Cu , Pe , Ag and Pt to the oxides (Bolton, I. c; Brodie, Proc. Boy. Hoc., 1863, 12, 209). 5. Solubilities. — K and K 2 dissolve in water with violent action, forming KOH , which reacts with all acids forming soluble salts. Potassium dissolves in alcohol, forming potassium alcoholate and hydrogen. Potassium platinum chloride, acid tartrate, silico-fluoride, picrate, phos- phomolybdate, perchlorate, and chlorate are only sparingly soluble in cold water, and nearly insoluble in alcohol. The carbonate arid sulphate are insoluble in alcohol. 6. Reactions, a. — Potassium and sodium hydroxides are very strong bases, fixed alkalis, and precipitate solutions of the salts of all the other metals (except Cs , Rb , and Li), as oxides or hydroxides. These precipi- tates are quite insoluble in water, except the hydroxides of Ba, Sr, and Ca . Excess of the reagent causes a resolution with the precipitates of Pb , Sb , Sn , Al , Cr , and Zn , forming double oxides as, K 2 Pb0 2 , potas- sium plumbite, etc. Potassium carbonate is deliquescent, strongly alkaline, and precipitates solutions of the salts of the metals (except Cs , Rb , Na , and li), forming normal carbonates with Ag , Hg', Cd , Fe", Mn , Ba , Sr , and Ca ; oxide with Sb ; hydroxide with Sn , Fe'", Al , Cr'" and Co'"; basic salt with Hg", and a basic carbonate with the other metals. b.— The potassium salts of HCN, H 4 Fe(CN) 8 , H 3 Fe(CN) 8 , and HCNS find extended application in the detection and estimation of many of the heavy metals. Tartaric acid, H 2 C 4 H^0 6 , or more readily sodium hydrogen tartrate, NaHC 4 H 4 6 , precipitates, from solutions sufficiently concentrated, potas- sium hydrogen tartrate, KHC 4 H 4 e , granular-crystalline. If the solution be alkaline, tartaric acid should be added to strong acid reaction. The test must be made in absence of non-alkali bases. The precipitate is in- creased by agitation, and by addition of alcohol. It is dissolved by fifteen parts of boiling water or eighty-nine parts water at 25°, by mineral acids, by solution of borax, and by alkalis, which form the more soluble normal tartrate, K 2 C 4 H 4 6 , but not by acetic acid, or at all by alcohol of fifty per cent. 224 POTASSIUM. §205, 6c. Picric acid, C 6 H 2 (N0 2 ) 3 0H , precipitates, from solutions not very dilute, the yellow, crystalline potassium picrate, C„H 2 (N0 2 ) 3 0K , insoluble in alco- hol, by help of which it is formed in dilute solutions. The dried precipi- tate detonates strongly when heated. c. — If a neutral solution of a potassium salt be added to a solution of cobaltic nitrite,* a precipitate of the double salt potassium cobaltic nitrite, K s Co(N , 2 )j , will be formed. In concentrated solutions the precipitate forms immediately, dilute solutions should be allowed to stand for some time; sparingly soluble in water, insoluble in alcohol and in a solution of potassium salts, hence the precipitation is more valuable as a separation of cobalt from nickel than as a test for potassium (§132, 6c). Potassium nitrate is not found abundantly in nature, but is formed by the decomposition of nitrogenous organic substances in contact with potassium salts, " saltpeter plantations "; or by treating a hot solution of NaNO, with KC1 (Z)., 2, 2, 72). It finds extended application in the manufacture of gun- powder. . d— See §206, 6d. e. — Potassium sulphide may be taken as a type of the soluble sulphides which precipitates solutions of the metals of the first four groups as sulphides except: Hg' becomes HgS and Hg°, Fe"' becomes FeS and S r and Al and Cr form hydroxides. The sulphides of arsenic, antimony and 1 tin dissolve in an excess of the reagent, more rapidly if the alkali sulphide contain an excess of sulphur. For the general action of H 2 S or soluble sulphides as a reducing agent see the respective metals. Potassium sul- phate is used to precipitate barium, strontium, and lead. It almost always occurs in nature as double salt with magnesium, K 2 S0 4 .MgS0 4 .MgCl 2 -f- 6H 2 , kainite, and is used in the manufacture of KA1(S0 4 ) 2 , K 2 C0 3 and KOH . As a type of a soluble sulphate it precipitates solutions of lead,, mercurosum, barium, strontium, and calcium; calcium and mereurosum incompletely. f. — Potassium chloride precipitates the metals of the first group, acting thus as a type of the soluble chlorides. It is much used with sodium nitrate in the preparation of potassium nitrate for the manufacture of gunpowder, in the preparation of K 2 C0 3 , KOH , and also as a fertilizer. Potassium bromide as a type of the soluble bromides precipitates solutions of Pb , Ag , and Hg (Hg" incompletely). Potassium iodide finds extended use in analytical chemistry in that it forms many soluble double iodides;, it is also extensively used in medicine. As a type of a soluble iodide it precipitates solutions of the salts. of Pb, Ag, Hg, and Cu'. Cu" salts are precipitated as Cul with liberation of iodine. Fe'" salts are merely reduced to Fe" salts with liberation of iodine. Arsenic acid is merely reduced to arsenous acid with liberation of iodine. * One co. of cobaltous nitrate solution and three co. of acetic acid are added to five cc. of a ten per cent solution of sodium nitrite. This gives a yellowish solution having an odor of nitrous acid. . §205, 7. POTASSIUM. 225 Potassium chlorate is used as a source of oxygen and as an oxidizing' agent in acid solutions. Sodium perchlorate, NadO* , precipitates from solutions of potassium salts potassium perchlorate, KCIO, , sparingly soluble in water and almost insoluble in strong alcohol (Kreider, Z. anorg., 1895, 9, 342). Potassium iodate is iised as a reagent in the detection of barium as Ba(IO s ) 2 . g. — The oxides of arsenic act as acid anhydrides toward KOH and form stable soluble potassium salts, arsenites and arsenates, which react with the salts of nearly all the heavy metals, h. — Potassium chromate and dichromate are both exten- sively used as reagents, especially in the analysis of Ag , Pb and Ba salts. i. — Fluosilicic acid, H 2 SiF 6 , precipitates from a neutral or slightly acid solution of potassium salts, potassium fluosilicate (silico-fluoride), X 2 SiF 6 , soluble in 833.1 parts of water at 17.5°; in 104.8 parts at 100°; and in 327 parts of 9.6 per cent HC1 at 14° (Stolba, /. pr., 1868, 103, 396\ The precipitate is white, very nearly transparent. /. — Platinic Chloride, PtCl 4 , added to neutral or acid solutions not too dilute, with hydrochloric acid if the compound be not a chloride, precipi- tates potassium platinic chloride, (KCl) 2 PtCl 4 , crystalline, yellow. Non- alkali bases also precipitate this reagent, and if present must be removed before this test. The precipitate is soluble in 19 parts of boiling water, ■or 111 parts of water at 10°. Minute proportions are detected by evapor- ating the solution with the reagent nearly to dryness, on the water-bath, and then dissolving in alcohol; the yellow crystalline precipitate, octahe- dral, remains undissolved, and may be identified under the microscope. k. — An alcoholic solution of BiCl 8 in excess of Ha 2 S 2 8 gives a yellow pre- cipitate with solutions of potassium salts (Pauly, O. C, 1887, 553). I. — Gold chloride added to sodium and potassium chloride forms double salts, e. #., KCl.AuClg + 2H 2 . If these salts are dried at 100° to 110° to remove water .and acids, the sodium salt is soluble in ether (separation from potassium) (Fasbender, C. C, 1894, 1, 409). 7. Ignition. — Ignited potassium hydroxide or potassium carbonate is a valuable desiccating agent for use in desiccators or in liquids. A mixture -of molecular proportions of K 2 C0 3 and Na 2 C0 3 melts at a lower tempera- ture than either of the constituents, and is frequently employed in fusion for the transposition of insoluble metallic compounds : BaSO t + K 2 C0 3 = BaC0 s + K 2 S0 4 . Potassium compounds color the flame violet. A little of the solid substance, or residue by evaporation, moistened with hydrochloric acid, is brought on a platinum wire into a non-luminous flame. The wire should be previously washed with HC1, and held in the flame to insure the absence of potassium. The presence of very small quantities of sodium enables its yellow flame completely to obscure the violet of potas- sium; but owing to the greater volatility of the latter metal, flashes of violet are sometimes seen on the first introduction of the wire, or at the border of the flame, or in its base, even when enough sodium is present to conceal the violet at full heat. The interposition of a blue glass, or 226 SODIUM. §205, 8. prism filled with indigo solution, sufficiently thick, entirely cuts off the yellow light of sodium, and enables the potassium flame to be seen. The red rays of the lithium flame are also intercepted by the blue glass or indigo prism, a thicker stratum being required than for sodium. It organic substances are present, giving luminosity to the flame, they must be removed by ignition. Certain non-alkali bases interfere with the examination. Silicates may be fused with pure gypsum, giving vapor of potassium sulphate. Bloxam (J. C, 1865, 18, 229) recommends to fuse insoluble alkali compounds with a mixture of sulphur, one part, and barium nitrate, six parts; cool, dissolve in water, remove the barium with NH 4 QH and (NH 4 ) 2 C0 3 and test for the alkalis as usual. The volatile potassium compounds, when placed in the flame, give a widely-extended continuous spectrum, containing two characteristic lines; one line, K «, situated in the outermost red, and a second line, K /J, far in the violet rays at the other end of the spectrum. 8. Detection. — Potassium is usually identified by the violet blue color which most of its salts impart to the Bunsen flame (7). Sodium inter- feres but the intervention of a cobalt glass (§132, 7) or a solution of indigo cuts out the yellow color of the sodium flame and allows the violet of the potassium to be seen. Some of the heavy metals interfere, hence the test should be made after the removal of the heavy metals (§§211 and 212). Potassium may be precipitated as the platinichloride (6;); as the per- chlorate (6/); as the silico-fluoride (6i); as the acid tartrate (6&); etc. Certain of these reactions are much used for the quantitative estimation (9) of potassium but are seldom used for its detection qualitatively. 9. Estimation. — (1) Potassium is converted into the sulphate or phosphate and weighed as sueh. (2) It is precipitated and weighed as the double chloride with platinum. (3) If present as KOH or K 2 CO„ it is titrated with standard acid (Kippenberger, Z. angew., 1894, 495). (4) It is precipitated with H 2 SiF and strong- alcohol. (5) Indirectly when mixed with sodium, by converting into the chlorides and weighing as such; then determining the amount of chlorine and calculating the relative amounts of the alkalis. (6) It is pre- cipitated as the bitartrate in presence of alcohol and, after nitration and solution in hot water, titrated with deci-normal KOH. (7) By precipitation as the perchlorate, XC10 4 (Wense, Z. angew., 1892, 233; Caspari, Z. angew., 1893, 68). 10. Oxidation. — Potassium is a very powerful reducing agent, its affinity for oxygen at temperatures not too high is greater than that of any other element except Cs and Eb . For oxidizing action of K 2 4 see 4. §206. Sodium. Na = 23.05 . Valence one. 1. Properties.— Specific gravity, 0.9735 at 13.5° (Baumhauer, B., 1873, 6, 665); 0.7414 at the boiling point (Ramsay, B., 1880, 13, 2145). Melting point, 97.6° (Hagen, B., 1883, 16, 1668). Boiling point, 742° (Perman, C. N., 1889, 59, 237). §206, 6d. SODIUM. 227 A silver-white metal with a strong metallic lustre. At ordinary temperatures it is softer than Li or Pb, and can be pressed together between the fingers; at — 20° it is quite hard; at 0° very ductile. It oxidizes rapidly in moist air and must be kept under benzol or kerosene. It decomposes water violently even at ordinary temperatures, evolving hydrogen, which frequently ignites from the heat of the reaction: 3Na + 2H 2 = 2NaOH + H 2 . It burns, when heated to a red heat, with a yellow flame. Pure dry Na is scarcely at all attacked by dry HC1 (Cohen, O. N., 1886, 54, 17). 2. Occurrence. — Never occurs free in nature, but in its various combinations one of the most widely diffused metals. There is no mineral known in which its presence has not been detected. It occurs in all waters mostly as the chloride from traces in drinking waters to a nearly saturated solution in some mineral waters and in the sea water. It is found in enormous deposits as rock salt, NaCl; as Chili saltpeter, NalTO,; in lesser quantities as carbonate, borate, sulphate, etc. 3. Preparation. — (1) By igniting the carbonate or hydroxide with carbon; (2) by igniting the hydroxide with metallic iron; (3) by electrolysis of the hydroxide; (4) by gently heating the carbonate with Mg . 4. Oxides and Hydroxides. — Sodium oxide, 2Ta 2 , is formed by burning sodium in oxygen or in air and heating again with Na to decompose the Na 2 O a (§205, 4, footnote). Sodium hydroxide, NaOH , is formed by dissolving the metal or the oxide in water (Eosenfeld, J. pr., 1893, (2), 48, 599); by treating a solution of sodium carbonate with lime; by fusion of NaNO B with CaCO a , CaO and Na^COa are formed and the mass is then exhausted with water; by igniting Na 2 C0 3 with Fe 2 0„ , forming sodium ferrate, which is then decom- posed with hot water into NaOH and Pe(OH) 8 (Solvay, 0. 0., 1887, 829). It is a white, opaque, brittle crystalline body, melting under a red heat. The fused mass has a sp. gr. of 2.13 (Filhol, A. GJi., 1847, (3), 21, 415). It has a very powerful affinity for water, gradually absorbing water from CaCL (Muller- Erzbach, B., 1878, 11, 409). It is soluble in about 0.47 part of water according to Bineau (0. r., 1855, 41, 509). Sodium peroxide, Na 2 2 , is formed by heating sodium in C0 2 free air or oxygen (Prud'homme, G. 0., 1893, (1), 199). It reacts as H 2 2 , partly reducing and partly oxidizing. It may be fused without decomposition. Water decom- poses it partially into NaOH and H 2 2 . 5. Solubilities. — Sodium and sodium oxide dissolve in water, forming the hydroxide, the former with evolution of hydrogen. In acids the corresponding sodium salts are formed, all soluble in water except sodium pyroantimonate, which is almost insoluble in water, and the fluosilicate sparingly soluble. The nitrate and chlorate are deliquescent. The carbonate (10 aq), sul- phate (10 aq), sulphite (8 aq), phosphate (12 aq), and the acetate (3 aq) are efflorescent. 6. Reactions, a. — As reagents sodium hydroxide and carbonates act in all respects like the corresponding potassium compounds, which see. 6. — By the greater solubility of the picrate and acid tartrate of sodium, that metal is separated from potassium (§205, 66). c. — Sodium nitrate occurs in nature in large quantities as Chili saltpeter, used as a fertilizer, for the manu- facture of nitric acid, with KC1 for making KKO, , etc. d. — Sodium phosphate, Na 2 HP0 4 , is much used as a reagent in the precipitation and estimation of Pb, Mn, Ba, Sr, Ca, and Mg. The phosphates of all metals except the alkalis are insoluble in water (lithium phosphate is only sparingly soluble (§210, he), soluble in acids). Solu- 228 SODIUM. §206, 6e. tions of alkali phosphates precipitate solutions of all other metallic salts as phosphates (secondary, tertiary or basic) except: HgCl 2 precipitates as a basic chloride (§58, 6d), and antimony as oxide or oxychloride (§70, 6d). e, f, g, ft. — As reagents the sodium salts react similar to the corresponding potassium salts, which see. i. — Sodium fluosilicate is soluble in 153.3 parts H 2 at 17.5° and in 40.66 parts at 100° (Stolba, Z., 1872, 11, 199); hence is not precipitated by fluosilicic acid except from very concentrated solutions (separation from K). /. — Sodium platinic chloride, (NaCl) 2 PtCl t , crystallizes from its concentrated solutions in red prisms, or prismatic needles (distinction from potassium or ammonium). A drop of the solution to be tested is slightly acidified with hydrochloric acid from the point of a glass rod on a slip of glass, treated with two drops of solution of platinic chloride, left a short time for spontaneous evaporation and crystallization, and observed under the micro- scope. h. — Solution of potassium pyroantimonate, K 2 H 2 Sb 2 7 , produces in neutral or alkaline solutions of sodium salts a slow-forming, white, crystal- line precipitate, Na 2 H 2 Sb 2 7 , almost insoluble in cold water. The reagent must be carefully prepared and dissolved when required, as it is not per- manent in solution (§70, 4c). 7. Ignition. — Sodium bicarbonate, NaHC0 3 , loses H 2 and C0 2 at 125° becoming Na 2 C0 3 , no further decomposition till 400° when a very small amount of NaOH is formed (Kirsling, Z. angew., 1889, 332). Sodium compounds color the flame intensely yellow, the color being scarcely affected by potassium (at full heat), but modified to orange-red by much lithium, and readily intercepted by blue glass. Infusible com- pounds may be ignited with calcium sulphate. The test is interfered with by some non-alkali bases, which should be removed (§§211 and 212). The spectrum of sodium consists of a single broad band at the D line in the yellow of the solar spectrum separable into two bands, D ; and D„, by prisms of higher refractive power. The amount of sodium in the atmosphere, and in the larger number of substances designed to be " chemically pure " is sufficient to give a dis- tinct but evanescent yellow color to the flame and spectrum. 8. Detection. — Sodium is usually detected by the color of the flame, yellow, in absence of the heavy metals. In the usual process of analysis the presence or absence of sodium is determined in the presence of magnesium (as Na 2 HP0 4 is the usual reagent for the detection of mag- nesium, it is evident that the presence or absence of the sodium must be determined before the addition of that reagent); and as that metal gives a yellowish color to the flame it mUst be removed if small quantities of sodium are to be detected. For this purpose the filtrate from Ba , Sr and Ca is evaporated to dryness and gently ignited to expel all ammonium salts; then taken up with a small amount of water and the magnesium precipitated as the hydroxide with a solution of barium hydroxide. After §207, 5. AMMONIUM. p 229 filtration the barium is removed by (NH 4 ) 2 C0 3 or H 2 S0 4 and the filtrate tested for sodium by the flame or by the pyroantimonate test (6fc). 9. Estimation. — (i) If present as hydroxide or carbonate, by titration with standard acid (Lunge, Z. angew., 1897, 41). (2) By converting into the chloride or sulphate and weighing as such. (3) In presence of potassium by converting into the chloride, weighing as such, then estimating the amount of chlorine with AgN0 3 and computing the amounts of K and Na. (4 It is precipitated by K 2 H 2 Sb 2 0, and dried and weighed as Na 2 H 2 Sb 2 7 . 10. Oxidation.- — Sodium ranks with potassium as a very powerful re- ducing agent. It is not quite so violent in its reaction and being much cheaper is almost universally used instead of potassium. Sodium peroxide may act both as a reducing and oxidizing agent. The action is similar to H 2 2 in alkaline solution, which see (§244, 6). §207. Ammonium. (NH 4 )'. Valence one. 1. Properties. — Specific gravity of NH, gas, 0.589 (Fehling, 1, 384); of the liquid, 0.6234 at 0° (Jolly, A., 1861, 117, 181). The liquid boils at —33.7°, at ■0° the liquid has a tension of 4.8 atmospheres (Bunsen, Pogg., 1839, 46, 95). Xiquid ammonia is a colorless mobile liquid, burns in air when heated or in oxygen without being previously heated. At ordinary temperature it is a gas with very penetrating odor. It burns with a greenish-yellow flame, and com- bines energetically with acids to form salts, the radical NH, being monovalent and acting in many respects similar to K and Na . At 0° one volume of water absorbs 1049.6 volumes of the gas; at 15°, 727.22 volumes (Carius, A., 1856, 99, 144). One gram of water, pressure 760 mm. and temperature 0°, absorbs 0.899 gram of NH,; with temperature 16°, 0.578 gram (Sims, A., 1861, 118, 345). 2. Occurrence. — Free ammonia does not occur in nature. Various ammonium salts occur widely distributed: in rain water, in many mineral waters, in almost all plants, among the products of the decay or decomposition of nitrogenous organic bodies, etc. 3. Preparation. — It is obtained from the reduction of nitrates or nitrites by nascent hydrogen in alkaline solution, e. g., 8A1 + 5KOH + 3KN0 3 + 2H 2 = 8KA10, -f- 3NH 3 ; by the reduction with the hydrogen of the zinc-copper couple; "by boiling organic compounds containing nitrogen with KMnO, in strong alkaline solution (as in water analysis) ; also by the oxidation of nitrogen in organic bodies with strong sulphuric (Kjeldahl method of nitrogen determina- tion). It is prepared on a larger scale by heating an ammonium salt with lime (or some other strong base). Nearly all the ammonium hydroxide and am- monium salts of commerce are obtained as a by-product in the production of illuminating gas by the destructive distillation of coal. 4. Hydroxide. — Ammonium hydroxide, NH 4 0H, is made by passing ammonia, NH 3 , into water. The gas is absorbed by the water with great avidity, and a strongly alkaline solution is produced. A solution having a sp. gr. of 0.90 at 15° contains 28.33 per cent of NH 3 (Lunge and Wiernik, Z. angew., 1889, 183). 5. Solubilities. — Ammonia, NH 3 , and all ammonium salts are soluble in water. Ammonia dissolves less readily in a strong solution of potassium hydroxide than in water. The carbonate (acid), and phosphate are efflores- cent. The nitrate and acetate are deliquescent, the sulphate slightly deli- quescent. 230 AMMONIUM. §207, 6a. 6. Reactions, a. — The fixed alkali hydroxides and carbonates liberate ammonia, NH 3 , from all ammonium salts, in the cold and more rapidly upon heating. Ammonium hydroxide, volatile alkali, colors litmus blue, neutralizes acids, forming salts, and precipitates solutions of the metals of the first four groups, manganese and magnesium salts imperfectly; due to the solubility of the hydroxide formed, in the ammonium salt produced by the reaction, and with these metals if excess of ammonium salts be present no precipitate will be formed by the NH„0H . The precipitate is a hydroxide except: with Ag and Sb it is an oxide, with mercury a sub- stituted ammonium salt and with lead a basic salt (see below, h and I). With salts of Ag , .Cu , Cd , Co , Ni , and Zn the precipitate redissolves in excess of the reagent. Ammonium carbonate, (NN 4 ) 2 C0 3 , is unstable and used only in solution. It is formed by adding ammonium hydroxide to a solution of the acid carbonate of commerce. It precipitates solutions of all the non-alkali metals, chiefly as carbonates except magnesium salts which are not at all precipitated, as a soluble double salt is at once formed (separation of Ba , Sr , and Ca from Mg). With salts of Ag , Cu , Cd , Co , Ni , and Zn , the precipitate is redissolved by an excess of the ammonium carbonate. o. — Dilute solutions of picric acid with ammonium hydroxide form in- tensely colored yellow solutions, a precipitate of ammonium picrate is formed if the solutions are quite concentrated. Tartaric acid precipitates ammonium salts very closely resembling the precipitate of potassium acid tartrate. The ammonium salt is more soluble in water than the potas- sium salt and does not leave K 2 C0 3 upon ignition. Sodium nitroferri- cyanide, Na 2 Fe(N0)(CN) 5 , added to a mixture of NH 4 0H and H 2 S [(NH 4 ) 2 S] gives a very intense purple color, characteristic of alkali sulphides and the manipulation may be modified so as to give a very deli- cate test for the presence of an alkali hydroxide or of hydrosulphurie acid. In no case, however, can the H 2 S be directly added to the sodium nitro- ferricyanide as it causes oxidation of the sulphur. To test for ammonia the gas should be liberated by KOH and distilled into a solution of H 2 S ; and this solution added to the Na 2 Fe(N0)(CN) 5 . c. — Ammonium nitrite, NH,NO, , is used in the preparation of nitrogen (§235, 3) ; ammonium nitrate in the preparation of nitrous oxide, N 2 , " laughing gas " (§237). d. — Ammonium phosphate, as a reagent, acts similarly to sodium phosphate. When sodium phosphate, Na.HPO, , is used to precipitate metals in the presence of ammonium hydroxide, a double phosphate of the metal and ammonium is frequently formed as MnWH 4 P0 4 , MgNH,,]^ , etc. By some chemists microcosmic salt, NaNH 4 HP0 4 , is preferred to sodium phosphate, Na 2 HP0 4 , as a reagent. e. — When ammonium hydroxide is saturated with H 2 S , ammonium sul- phide, (NH 4 ) 2 S , is formed. Complete saturation is indicated by the failure §207, 6h. AMMONIUM. 231 to precipitate magnesium salts, that is, NH 4 OH precipitates magnesium salts while (NH 4 ) 2 S does not. Freshly prepared ammonium sulphide is colorless, but upon standing becomes yellow with loss of ammonia and formation of the poly-sulphides, (NH 4 ) 2 S X . The yellow poly-sulphide may also be formed by dissolving sulphur in the normal ammonium sul- phide. As a precipitant ammonium sulphide acts similarly to the fixed alkali sulphides. The sulphides of Sb'" and Sn" are with great difficulty soluble in the normal ammonium sulphide, but readily soluble in the poly-sulphide. Nickel sulphide, NiS , is insoluble in normal ammonium sulphide but is sparingly soluble in the yellow poly-sulphide (distinction from cobalt). (NH 4 ) 2 S gives a rich purple color with sodium nitroferri- cyanide (&). Ammonium sulphate as a precipitating reagent acts similar to all soluble sulphates (§205, 6e). A 25 per cent solution of (NH 4 ) 2 S0 4 is used to dissolve CaS0 4 (§188, 5c) (distinction from Ba and Sr). f. — Ammonium chloride is much used as a reagent. It prevents pre- cipitation of the salts of Mn by the NH 4 0H , and is of special value in the precipitation of the third group as hydroxides and the fourth group as sulphides by preventing the formation of soluble colloidal compounds. The solubility of the precipitates of the carbonates of the fifth group is slightly increased by the presence of ammonium chloride ; i. e., very dilute solutions of barium chloride are not precipitated by ammonium carbonate in presence of a large excess of ammonium chloride. The salts of mag- nesium are not precipitated by the alkalis or by the alkali carbonates in presence of ammonium chloride. The solubility of A1(0H) 3 is diminished by the presence of NH 4 C1 (§124, 6a, and §117). g, Ik — Similar as reagents to the corresponding 1 potassium salts, i. — ITruo- silicic acid, H,SiP, , does not precipitate ammonium salts, the ammonium fluosilicate being' very soluble in water (distinction from potassium). /. — Plat- inum chloride, PtCl 4 , forms with ammonium salts the yellow double ammonium platinum chloride, (NH 4 ) 2 PtCl 6 , very closely resembling the potassium salt with the same reagent, but upon ignition only the spongy metallic platinum is left, i. e., no chloride of the alkali metal, as KC1 . h. — A solution of potassium mercuric iodide, K 2 HgI 4 , containing also potassium hydroxide — Nessler's test * — produces a brown precipitate of nitrogen dimercuric iodide, NHg 2 I , dimercur-ammonium iodide (§58, 6a), soluble by excess of KI and by HC1 ; not soluble by KBr (distinction from HgO): NH, + 2HgI 2 = NHgJ + 3HI WH.OH + 2K 2 HgI 4 + 3KOH = HHg.I + 7KI + 4H 2 * This reagent may be prepared as follows : To a solution of mercuric chloride add solution of potassium iodide till the precipitate is nearly all redissolved ; then add solution of potassium hydroxide sufficient to liberate ammonia from ammonium salts ; leave until the liquid becomes clear, and decant from any remaining sediment. 232 AMMONIUM. §207^61, This very delicate test is applicable to ammonium hydroxide or salts; traces forming only a yellow to brown coloration. The potassium mercuric iodide, "Meyers Reagent," alone, precipitates the alkaloids from neutral or acid solutions, but does not precipitate ammonium salts from neutral or acid solutions. Ammonium hydroxide in alcoholic solution does not give a precipitate with Nessler's reagent, but from this solution a precipi- tate is formed with HgCl 2 (De Koninck, Z., 1893, 32, 188). I. — Mercuric chloride, HgCl 2 , forms, in solutions of ammonium hy- droxide or ammonium carbonate, the " white precipitate " of nitrogen dihydrogen mercuric chloride, NH 2 HgCl , or mercur-ammonium chloride. If the ammonium is in a salt, not carbonate, it is changed to the carbonate and precipitated, by addition of mercuric chloride and potassium carbonata previously mixed in solutions (with pure water), so dilute as not to precipi- tate each other (yellow). This test is intensely delicate, revealing the presence of ammonia derived from the air by water and many substances (Wittstein, Arch. Pharm., 1873, 203, 327). m. — Add a small quantity of recently precipitated and well-washed silver chloride, and, if it does not dissolve after agitation, then add a little potassium hydroxide solution. The solution of the AgCl , before the addition of the fixed alkali, indicates free ammonia; after the addition of the fixed alkali, ammonium salt. (Applicable in absence of thiosulphates, iodides, bromides and sulpho- cyanates.) m. — Sodium phosphomolybdate (§75, 6d) precipitates ammonium from neutral or acid solutions; also precipitates the alkaloids, even from very dilute solu- tions, and, from concentrated solutions, likewise precipitates K , Eb and Cs (all the fixed alkalis except Na and Li). 7. Ignition. — Heat vaporizes the carbonate, and the haloid salts of am- monium, undecomposed (dissociated but reuniting- upon cooling); decomposes the nitrate with formation of nitrous oxide and water, and the phosphate and borate with evolution of ammonia. NH, heated to 780° or higher is dissociated into N and H (Ramsay and Young, J. C, 1884, 45, 88). 8. Detection. — As ammonium hydroxide and chloride are used in the regular process of analysis, the original solution must be tested for the presence or absence of ammonium compounds. The hydroxide or the carbonate may be detected by the odor (1) ; the action on red litmus paper suspended in the test-tube above the heated solution; the blue color im- parted to paper wet with copper sulphate; the blackening of mercurous nitrate paper; and if in considerable quantity, the white vapors when brought into contact with the vapors of volatile acids. In combination as salts the gas is liberated by the fixed alkali hydroxides or carbonates (oxides or hydroxides of Ba, Sr ; or Ca may be used) and distilled into Nessler's reagent, or collected in water and the test with HgCl 2 (6Z) applied or any of the tests for ammonium hydroxide. , 9. Estimation. — Ammonium salts are usually estimated by distillation into a standard acid, from a solution made alkaline with KOH. , and titration of the excess of the acid with a standard NH,OH solution, using tincture of cochineal §208, 5. CAESIUM. 233 as an indicator. It may be converted into the chloride and precipitated by PtCl 4 and weighed as the double' platinum salt. 10. Oxidation. — Ammonium salts in solution, treated with chlorine gas, gen- erate the unstable and violently explosive "nitrogen chloride" (NCl a ?) (a). The same product is liable to arise from solid ammonium salts treated with chlorine. Gaseous ammonia, and ammonium hydroxide, with chlorine gas, generate free nitrogen (B), a little ammonium chlorate being formed if the ammonia is in excess. Hypochlorites or hypobromites (or chlorine or bromine dissolved in aqueous alkali, so as to leave an alkaline reaction) liberate, from dissolved ammonium salts, all of their nitrogen (as shown in the second equa- tion of 6) ; the measure of the nitrogen gas being a means of quantitative estimation of ammonium. With iodine, ammonium iodide and the explosive iodamides (c) are produced; or under certain conditions an iodate (d). Ammo- nium hydroxide is liable to atmospheric oxidation to ammonium nitrite and nitrate. Permanganates oxidize to nitrate (e) (Wanklyn and Gamgee, J. 0.,. 1868, 21, 29). In presence of Cu the O of the air oxidizes the nitrogen of ammonia to a nitrite (f) (Berthelot and Saint-Gilles, A. Gh., 1864, (4), 1, 381). Ammonia is somewhat readily produced from nitric acid by strong reducing agents (g). It is formed with carbonic anhydride, in a water solution of cyanic acid, and, more slowly, in a water solution of hydrocyanic acid. It is. generated, by fixed alkalis, in boiling solution of cyanicfes (ft) ; also in boiling solutions of albuminoids and other nitrogenous organic compounds, this forma- tion being hastened and increased by addition of permanganate (Wanklyn's process). Fusion with fixed alkalis transforms all the nitrogen of organic bodies into ammonia. (a) NH 4 C1 + 3C1 2 = NCI, + 4HC1 (6) 8HTH 3 + 3C1 2 = 6BTH 4 C1 + ET 2 2NH 4 C1 + 3C1 2 = 8HC1 + N 2 (c) 2NH, + I 2 = NH.I + NH,I (d) 6NH 4 OH + 3I 2 = 5NH 4 I + NH 4 IO a + 3H 2 (e) 6NH 4 OH + 8H]Mn0 1 = 3NH 4 N0 3 + 8MnO(OH) 2 + 5H 2 (f) 12Cu + 2NH 3 + 90 2 = 12CuO + 2HN0 2 + 2H 2 (g) 3HNO a + 8AI + 8KOH = 8KA10 2 + 3NH„ + H 2 (ft) HCN + KOH + H 2 = NH, + KCH0 2 (formate). §208. Caesium. Cs = 132.9 . Valence one. 1. Properties.— Specific gravity, 1.88 at 15° (Setterberg, A., 1882, 211, 100). Melting point, between 26° and 27°. It is quite similar to the other alkali metals; silver-white, ductile, very soft at ordinary temperature. It burns rapidly when heated in the air, and takes fire when thrown on water. It may be kept under petroleum. It is the most strongly electro-positive of all metals. 2. Occurrence. — Widely distributed but in small quantities; as caesium aluminum silicate (mineral castor and pollux) (Pisani, 0. r., 1864, 58, 715); in many mineral springs (Miller, C. N., 1864, '10, 181) ; in the ash of certain plants, tobacco, tea, etc. 3. Preparation. — By electrolysis of a mixture of CsCN with Ba(ClT) 2 ; by ignition of CsOH with Al in a nickel retort (Beketoff, C. C, 1891, (2), 450). 4. Oxide and Hydroxide. — An oxide has not yet been prepared. The hydroxide, CsOH , is a grayish-white solid, very deliquescent, absorbs C0 2 from the air; dissolves in water with generation of much heat, forming a strongly caustic solution. 5. Solubilities. — Caesium dissolves with great energy in water, acids or alcohol, liberating hydrogen and forming the hydroxide, salts or alcoholate respectively. The hydroxide is soluble in water and alcohol. The salts are all quite readily soluble. The double platinum chloride, Cs 2 PtCl 4 , aud the acid tartrate, CsHC,H,O , being least soluble and used in preparation of the salts free from the other alkali metals. 234 RUBIDIUM— LITHIUM. §208, 6. 6. Reactons. — In all its reactions similar to the other fixed alkalis'. 7. Ignition. — Caesium salts color the non-luminous flame violet. The spec- trum gives two sharply defined lines, Cs a and Cs /S, in the blue and a third faint line in the orange-red Cs y, also several faint lines in the yellow and green. With the spectroscope three parts of CsCl may be detected in presence of 300,000 .to 400,000 parts KC1 or NaCl; and one part in presence of 1,500,000 parts LiCl (Bunsen, Pogg., 1875, 155, 633). 8. Detection. — By the spectroscope (7 and §210, 7). 9. Estimation.—^) As the double platinum chloride; (2) as the chloride with RbCl , estimation of the amount of CI and calculation of the relative amounts of the metals; (3) as the sulphate obtained from ignition of the acid tartrate and treatment with H 2 S0 4 (Bunsen, Pogg., 1863, 119, 1). §209. Rubidium. Rb = 85.4 . Valence one. 1. Properties. — Specific gravity, 1.52 (Bunsen, A., 1863, 125, 367). point, 38.5°; at — 10° soft as wax. A lustrous silver-white metal with a tinge of yellow, oxidizes rapidly in the air, developing much heat and soon igniting. Volatile as a blue vapor below a red heat. The metal does not keep well under petroleum, but is best preserved in, an atmosphere of hydrogen. Next to caesium it is the most electro-positive of all metals. 2. Occurrence. — Widely distributed in small quantities, usually with caesium, and frequently with the other alkali metals, always in combination. None of the alkali metals can occur free in nature. 3. Preparation. — From the mother liquor obtained in the preparation of Li salts (Heintz, J. pr., 1862, 87, 310): (1) By ignition of the acid tartrate with charcoal; (2) electrolysis of the chloride; (3) by ignition with Big 1 or Al (Winkler, B., 1890, 23, 51; Beketoff, B., 1888, 21, c, 424). 4. Oxide and Hydroxide. — The oxide Rb 2 has not been with certainty pre- pared. The hydroxide, RbOH , is formed when the metal is decomposed by water; also through the action of Ba(0H) 2 upon Rb 2 S0 4 . It is a gray-white, brittle mass, melting under a red heat. 5. Solubilities. — The metal dissolves in cold water, in acids and in alcohol with great energy, evolving hydrogen. The hydroxide is readily soluble in water with generation of heat. The salts are all quite readily soluble. The acid tartrate is about eight times less soluble than the corresponding Cs salt. Among the less soluble salts are to be mentioned the perchlorate, the fluosili- cate, the double platinum chloride, the silicotungstate, the picrate, and the phosphomolybdate. The alum is less soluble than the corresponding potassium alum. 6. Reactions. — Similar to the other fixed alkalis. 7. Ignition. — The salts give a violet color to the flame. The spectrum gives two characteristic lines in the violet, Rb a and Rb (3 ; two less intensive in the outer red, Rb y and Rb 6; a fifth Rb s in the orange ; and many faint lines in the orange, yellow and green. As small a, quantity as 0.0000002 gram of RbCl can be detected (Bunsen, I.e.). 8. Detection. — By the spectroscope (7 and §210, 7). 9. Estimation.— (1) By weighing with CsCl as the chlorides, determining the amount of CI and calculating the proportion of the metals; (2) as the double platinum chloride. §210. Lithium. li = 7.03 . Valence one. 1. Properties.— Specific gravity, 0.5936, the lightest of all known solid bodies (Bunsen and Matthiessen, A., 1855, 04, 107). Melting poimt, 180°; does not vaporize at a red heat. It is a silver-white metal with a grayish tinge; harder than K or Ma but softer than Pb , Ca or Sr; it is tough and may be drawn into wire and rolled into sheets. It is more electro-positive than the alkaline earth metals but less electro-positive than K or Na . The pure metal is quite similar §210, 8. LITHIUM. 235 in appearance and in its chemical properties to K and Na , but does not react so violently as those metals. It does not ignite in the air until heated to 200°, and then burns quietly with a very intense white light. It also burns with vivid incandescence in CI , Br , I , O , S and dry C0 2 . It decomposes water readily, forming LiOH and H , but not with combustion of the hydrogen or ignition of the metal. 2. Occurrence. — It is a sparingly but widely distributed metal. Usually pre- pared from lepidolite, triphylene or petalite. Traces are found in a great many minerals, in mineral springs, and in the leaves and ashes of many plants; e. g., coffee, tobacco and sugar-cane. 3. Preparation. — It is prepared pure only by electrolysis, usually of the chloride. A larger yield is obtained by mixing the LiCl with NHUCl or KC1 (Giintz, G. r., 1893, 117, 732). The metal is also obtained by ignition of the carbonate with Mg , but the metal is at once vaporized and oxidized. 4. Oxide and Hydroxide. — It forms one oxide, Li 2 , by heating the metal in oxygen or dry air; cheaper by the action of heat upon the nitrate. The corresponding hydroxide, LiOH , is made by the action of water upon the metal or its oxide; cheaper by heating the carbonate with calcium hydroxide. 5. Solubilities. — The metal is readily soluble in water with evolution of hydrogen, forming the hydroxide; soluble in acids with formation of salts. The oxide, LLO , dissolves in water, forming the hydroxide. The most of the lithium salts are soluble in water. A number of the salts, including the chloride and chlorate, are very deliquescent. The hydroxide, carbonate and phosphate are less soluble in water than the corresponding compounds of the other alkali metals. In this respect lithium shows an approach to the alkaline earth metals. LiOH is soluble in 14.5 parts water at 20° (Dittmar, J. Soc. Ind., 1888, 7, 730); Li 2 C0 3 in 75 parts at 20°; Li 3 P0 4 in 2539 parts pure water and 3920 parts ammoniacal water, more soluble in a solution of NH.,C1 than in pure water (Mayer, A., 1856, 98, 193). 6. Reactions. — Lithium salts in general react similar to the corresponding potassium and sodium salts. They are as a rule more fusible and more easily decomposed upon fusion." Soluble phosphates precipitate lithium phosphate, more soluble in NH,C1 solution than in pure water (distinction from mag- nesium). In dilute solutions the phosphate is not precipitated until the solu- tion is boiled. The delicacy of the" test is increased by the addition of NaOH, forming a double phosphate of Na and Li (Eammelsberg, A. Ch., 1818, (2), 7, 157). The phosphate dissolved in HC1 is not at once precipitated by neutraliz- ing with NH,OH (distinction from the alkaline earth metals). Nitrophenic acid forms a yellow precipitate, not easily soluble in water. 7. Ignition. — Compounds of lithium impart to the flame a carmine-red color, obscured by sodium, but not by small quantities of potassium compounds. Blue glass, just thick enough to cut off the yellow light of sodium, transmits the red light of lithium; but the latter is intercepted by a thicker part of the blue prism, or by several plates of blue glass. The spectrum of lithium con- sists of a bright red band, Li a , and a faint orange line, Li /3 . The color tests have an intensity intermediate between those of sodium and potassium. 8. Detection.— By the spectroscope. — To the dry chlorides of the alkali metals a few drops of HC1 are added and the mass extracted with 90 per cent alcohol. The solution contains all the rare alkalis and some Na and K . Evaporate to dryness, dissolve in a small amount of water and precipitate with platinum chloride. The double platinum and potassium chloride is more soluble than the corresponding salt of Kb and Cs'. Boil repeatedly with small portions of water to remove the potassium, and frequently examine the residue by the spectroscope as follows: Wrap a small amount of the precipitate in a moistened filter paper, then in a platinum wire and carefully char. After charring is complete, ignite before the spectroscope. The K spectrum grows fainter, that of Eh and Cs appear. Evaporate to dryness the filtrate from the precipitate of the platinum double salts, add oxalic acid and ignite, moisten with HCl, evaporate and extract with absolute alcohol and ether. Upon evaporation of the extract LiCl is obtained, almost pure. Test with the spectroscope and by forming the insoluble phos- phate. 236 DIRECTIONS FOB ANALYSIS WITH NOTES. §210, 9".. 9. Estimation. — After separation from other elements it may be weighed asr a sulphate, carbonate or phosphate, Li s P0 4 . It may also be estimated by the- comparative intensity of theWes in the spectroscope (Bell, Am., 1886, 7, 35). directions for the analysis of the metals of the alkali group (Sixth Group). §211. If the material is found not to contain magnesium, the clear filtrate from the carbonates of Ba , Sr , and Ca , after testing for traces, with (NH 4 ) 2 S0 4 and (NH 4 ) 2 C 2 4 (§193), may at once be tested for the pres- ence of potassium and sodium. If magnesium be present it should be- removed in order to test for small amounts of sodium. Potassium and large amounts of sodium may be readily detected in the presence of mag- nesium. It is evident that the magnesium must not be removed by the usual reagent used to detect the presence of that element, i. e. Na 2 HP0 4 . It is recommended by many to use ammonium phosphate, (NH 4 ) 2 HP0 4 . This reagent removes the magnesium, and permits the application of th& flame test for the fixed alkalis ; but the presence of the phosphate obstructs the gravimetric determination of the alkalis. The phosphate may be removed by lead acetate and the excess of the lead by hydrogen sulphide. §212. As a better method it is directed to evaporate the nitrate con- taining the magnesium and the alkalis to dryness, ignite gently to remove the ammonium salts. Dissolve the residue in water and add Ba(0H) 2 to precipitate the magnesium as Mg(0H) 2 (§§177 and 182). After filtration,, the excess of barium in the nitrate is removed by H 2 S0 4 , and the filtrate from the barium sulphate is ready to be tested for the fixed alkalis by the- flame test or by gravimetric methods as may be desired. The presence of sodium obscures the flame reaction for potassium, but the introduction of a cobalt glass (§132, 7) or an indigo prism cuts out the sodium flame- and allows the violet potassium flame to be seen. Study 6, 7, 8, and 9 of §§205 and 206. §213. The free use of ammonium salts during the process of analysis, makes it necessary that the testing for ammonium be done in the original, solution or in the filtrate from the Tin and Copper Group. Add an excess of KOH or NaOH to the solution and warm gently. Notice the odor (§207, 1). Suspend a piece of moistened red litmus paper in the test-tube; in the presence of ammonia it will be changed from red to blue color. To detect the presence of small amounts of ammonium salts, heat the strongly alkaline mixture nearly to boiling and pass the evolved gas into water. Test this solution (ammonium hydroxide) with Nessler's Eeagent (§207, 6k) or by the precipitation with HgCl 2 (§207, 6Z). Study §207, 6, 7, 8, and 9. §214. The rare metals of the Alkali Group: lithium, rubidium, and §2] 5. DIRECTIONS FOR ANALYSIS WITE NOTES. 237 caesium, are rarely met with in the ordinary analyses. If their presence is suspected they are tested for and detected by the spectroscope (7, §§208, 209 and 210). §215. Lithium, because of the insolubility of its phosphate (§210, 5c), interferes with the detection of magnesium. If the filtrate after the removal of barium, strontium, and calcium be evaporated to dryness and gently ignited to remove all ammonium salts; the residue, dissolved in water and treated with an excess of barium hydroxide, will give a precipi- tate of the magnesium as the hydroxide, leaving the lithium in solution. The barium hydroxide precipitate may be tested for magnesium and from the nitrate the excess of barium hydroxide may be removed by sulphuric acid before testing for the alkali metals. PAET IIL-THE NON-METALS. §216. Balancing Equations in Oxidation and Reduction. Statement of Bonds in Plus and Minns Numbers,* according to chemical polarity, positive and negative (see §3 footnote). In the terms of this notation the plus bond is the unit of Oxidation and the minus bond is the unit of Reduction. A bond, that is a unit of active valence, is either a plus one or a minus one. The formula of a molecule of hydrochloric acid is stated, H +I C1 _I . That of water, (H +I ) 2 0~ n . (The plus sign is understood when no sign is written before the valence number.) H +/ q£. / Plus and minus bonds are represented as positive and negative quan- tities. In the formula of hydrochloric acid, as above, the difference between the polarity of the hydrogen atom and that of the chlorine atom is stated as a difference of two. In any compound the sum of the plus bonds and the minus bonds of the atoms forming a molecule is zero. Free elements, not having active valence, have zero bonds in this notation, f The Oxidation of any element is shown by an increase, and its Eeduetion by a decrease, in the sum of its bonds. When one substance reduces another the element which is reduced loses as many bonds as are gained by the element which is oxidized. It is evident that, changes in valence being reciprocal in oxidation and reduction, there is no gain or loss in the sum of the bonds of two elements which act upon each other. The use of this notation is illustrated in the following equations : 3SnCl 2 + H 2 S0 8 + 6HC1 = 3SnCl 4 + H 2 S + 3H 2 In this equation the three atoms of tin gain six bonds; the bonds of the sulphur in the H 2 S0 3 have then been diminished by six; that is, it has given up six bonds to the tin, and having only four in the first place must now have minus two (4 -6 = -2). • O C. Johnson, C. JV., 1880, 43, 61. See also Ostwald, GruvOr. allg. Chem., 3te Aufl., 1899, S. 439. •Hf there is polarity in the union of like atoms with each other in forming an elemental Molecule, the sum must he zero, as in the formation of the molecules of compounds. §217, f. BALANCING OF EQUATIONS. 239 3SnCl 2 + HIO a + 6HC1 = 3SnCl f + HI + 3H 2 Here also the three atoms of tin gain six bonds, and these are furnished by the iodine of the HI0 3 . It has five in the first place, and being diminished by six, has one negative bond remaining (5 -6 = -] ). [In other words, unless we deny that iodine has five bonds in HI0 3 , we must admit that it has one negative bond in HI (written HT - ').] 8HMCn.0 4 + 5AsH, + 8H 2 S0 4 = 5H„As0 4 + 8MnS0 4 + 12H 2 In this equation eight atoms of manganese in the first member have 56 bonds, and a like amount in the second member has only 16, losing 40, and this 40 has been gained by the five atoms of arsenic. They now have 25, after gaining 40. They must then have had — 15 in the first place (25 — 40 = -15). That is, the atom of arsenic in arssnous hydride has -3 bonds (As-"'H 3 ). SnClj + HgCl 2 = Hg + SnCl 4 This equation illustrates the statement that free elements have no bonds. The tin gains two bonds, and these two bonds are taken from the mercury in the HgCl 2 . §217. Rule for Balancing Equations. The" number of oxidation bonds which any element has is determined by the following rules : a. Hydrogen has always one positive bond. ~f— I . "2. — b. Oxygen has always two negative bonds. j_ A Q c. Free elements have no bonds. d. The sum of the bonds of any compound is zero. e. In salts the bond of the metal is always positive. /. In acids and in salts the acid radical has always negative bonds. Thus, the bond of free Pb is zero, but in PbCl 2 the lead has two posi- tive bonds, and each atom of chlorine has one negative bond. In Bi 2 S s , each atom of Bi has three positive bonds (e), and each atom of S has two negative bonds (/). In ammonium nitrite, NH 4 N0 2 , or H 4 = N — — N = , the nitrogen of the NH 4 has four negative bonds and one positive bond. The other nitrogen, that of the acid radical N0 2 , has three positive bonds. Each atom of hydrogen has one positive bond and each atom of oxygen two negative bonds, the sum being zero : +4 — 4+1 + 3 — 4 = 0. In the following salts, etc., the bond of each element is marked above, with its proper sign, plus being understood if no sign is given. Then f ol- 240 BALANCING OF EQUATIONS. §218,7. lows the equation in full, the bonds of each atom being multiplied by the number of atoms, and all being added, the sum is seen to be aero. Hg"(Nvo-" a ) 2 .2 + 10 — 12 = o Bi'" 2 (Sviacid sulphate of potassium, and a fluo- ride, fused to a bead on the loop of platinum wire, in the clear flame of the Bunsen gas-lamp, an evanescent yellowish-green color is imparted to the flame. Borates fused in the inner blow-pipe flame with potassium acid sulphate give the green color to the outer flame. If a crystal of boric acid, or a solid residue of borate previously treated with sulphuric acid, on a porcelain surface, is played upon by the flame of Bunsen's Burner, the green flame of boron is obtained. §222, 1. CARBON. 247 If a powdered borate (previously calcined), is moistened with sulphuric acid and heated on platinum wire to expel the acid, then moistened with glycerine and burned, the green flame appears with great distinctness. The glycerine is only ignited, then allowed to burn by itself. Barium does not interfere (being held as sulphate, non-volatile) ; copper should be previously removed in the wet way. The glycerine flame gives the spec- trum. But in all flame tests, boric acid must be liberated. Borates (fused on platinum wire with sodium carbonate) give a char- acteristic spectrum of four lines, equidistant from each other, and extend- ing from Ba y in the green to Sr i 2 5 becomes bismuth oxalate and C0 2 . 5. — Mn" +n becomes Mn". (That is, all compounds of manganese having more than two bonds are reduced to the dyad.) In absence of other free acid, MnC 2 4 is formed, and C0 2 is given off. If some non-reducing acid be present, such as H 2 S0 4 , it unites with the manganese, and all of the oxalic acid is converted into C0 2 . 6. — Co 2 3 and Co(0H) 3 form cobaltous oxalate, and C0 2 is evolved. 7. — Ni 2 8 and Ni(0H) 3 become nickelous oxalate, and C0 2 is evolved. ■8. — H 2 Cr0 4 is reduced to chromic oxalate, and C0 2 is evolved. As a rule, reducing agents have no action on oxalic acid at ordinary "temperatures. By fusion, however, a few metals, K , Na , Mg , etc., reduce it to free carbon. B. — With non-metals and their compounds. 1.— HCW , HCNS , H 4 Fe(CN) 6 , and H 3 Fe(CN) 6 seem to be without action upon oxalic acid. 2. — HN0 2 seems to have no action upon H 2 C 2 4 . With HN0 3 , C0 2 , NO , and H 2 are formed. The nitric acid should be concentrated. Test for the C0 2 by passing the gases into a solution of BaCl 2 containing KOH . 3. — H 3 P0 2 , H3PO3 , and H 3 P0 4 do not act upon oxalic acid. If.. — Concentrated sulphuric acid, with a gentle heat, decomposes oxalic acid, by removing the elements of water from it, with effervescence of ■carbon dioxide and carbon monoxide: H 2 C 2 4 -f- H 2 S0 4 = H 2 S0 4 .H 2 + C0 2 -)- CO . With oxalates, the decomposition generates the same gases. Other strong dehydrating agents produce the same result. The effervescing gases, C0 2 and CO , give the reactions for carbonic anhy- dride; also, if in a sufficient quantity, the CO will burn with a blue flame, when ignited. 5. — With chlorine, hydrochloric acid is formed and the oxalic acid becomes C0 2 (Gmelin's Hand-booh, 9, 116). This reaction takes place more readily in the presence of KOH , forming KC1 and K 2 C0 3 . HC10 forms C0 2 and CI . If the oxalic be in excess HC1 is formed. The action 258 OXALIC ACID. §227, B6, is more rapid in the presence of a fixed alkali, an alkali chloride and carbonate being formed. HC10 3 forms C0 2 and varying proportions of CI and HC1 . A high degree of heat and excess of oxalic acid favoring the production of HC1 (Calvert and'Davies, J. C, 1850, 2, 193). 6. — Bromine decomposes oxalic acid, in alkaline mixture, forming a bromide and a carbonate. In acid mixture a similar reaction takes place if a hot saturated solution of oxalic acid be used in excess. With HBrO., , bromine and C0 2 are formed; with excess of oxalic acid and heat hydro- bromic acid is formed. 7. — HI0 3 forms C0 2 and I . With mixtures of chlorates, bromates, and iodates, the chlorate is first decomposed, then the bromate, and finally the iodate (Guyard, J. C, 1879, 36, 593). 7. Ignition. — The oxalates are all dissociated on ignition. Those of the metals of the alkalis and alkaline earths are resolved at an incipient red heat, into carbonates and carbon monoxide (a) — a higher temperature decomposing the alkaline earth carbonates. The oxalates of metals, whose carbonates are easily decomposed, but whose oxides are stable, are re- solved into oxides, carbonic anhydride, and carbon monoxide (&). The oxalates of metals, whose oxides are decomposed by heat, leave the metal, and give off carbonic anhydride (c). As an example of the latter class, silver oxalate, when heated before the blow-pipe, decomposes explosively,, with a sudden puffing sound — a test for oxalates : (a) CaC 2 4 = CaC0 3 + CO (6) ZnC 2 4 = ZnO + C0 2 + CO (c) Ag 2 C 2 4 = 2Ag + 2C0 2 8. Detection. — (a) By warming with concentrated sulphuric acid after decomposition of carbonates with dilute sulphuric acid; showing the pres- ence of C0 2 by absorption in Ca(0H) 2 or in a solution of BaCI, alkaline with KOH ; and showing the presence of CO by its combustibility. (6) In solution by precipitation in neutral, alkaline, or acetic acid solution by calcium chloride, and solubility of the precipitate in dilute hydrochloric acid. Frey {Z., 1894, 33, 533), recommends the formation of a zone of precipitation. To the HC1 solution containing BaCl 2 and CaCl 2 he adds carefully a solution of NaC 2 H 3 2 and watches the zone of contact. 9. Estimation. — (a) It is precipitated as CaC 2 4 ; alter washing, the Ca is determined by §188, 9, from which the oxalic acid is calculated. (6) By the amount of KMeO, which it will reduce, (c) By measuring the amount of C0 2 evolved when it is oxidized in any convenient manner, usually by Mn0 2 . ( a yellow-white " K 2 MgFe(CEr)„ (only in concentrated solu- tion). Manganese, a white " Mn 2 Fe(CN) (soluble in HC1). Mercury (HgO, a white " Hg 4 Fe(CN)„ (gelatinous). Mercury (Hg"), a white " Hg 2 Fe(CN)„ , turning to Hg(CN) 2 and Fe a (Fe(CN)a) 2 , blue. Molybdenum, a brown " Nickel, a greenish-white " Mi 2 Fe(CN) . Silver, a white " Ag 4 Fe(ClI) 6 , (slowly turning blue). Tin (Sn" and Sniv), white " (gelatinous). ■Uranium (uranous), brown " XTFe(CIT) . Uranium (uranyl), red-brown " (tr0 2 ) 2 Fe(CN) 6 . Zinc, a white, gelatinous " Zn 2 Fe(CN)„ . See Wyrouboff (A. Ch., 1876 (5), 8, 444; and 1877, (5), 10, 409). Insoluble ferrocyanides are transposed by alkalis (§330, 6, Class 11.) It will be observed (§230, 6) that ferrocyanides are ferrous combinations, while ferricyanides are ferric combinations. And, although ferrocyanides are far less easily oxidized than simple ferrous salts, being stable in the air, they are §232. EYDR0FERRIC7ANIC ACID. 269 nevertheless reducing agents, of moderate power: 2K 4 Fe(CN'),, + Cl 2 = 2K 3 Fe(CN) + 2KC1 . Pb0 2 with sulphuric acid forms Pb" and H 3 Fe(CN) e . Ag' with fixed alkali forms an alkali ferricyanide and metallic silver. Crvi with phosphoric acid, gives Cr'" and H 3 Fe(CN) (Schonbein, J. pr., 1840^ 20, 145). Co'" with phosphoric acid forms Co" and H 8 Fe(Cl]') l , . Ni'" with acetic acid gives Ni" and H a Fe(CN') . Mn0 2 with phosphoric acid gives Mn" and H 3 Fe(CN) a . ■Mnvn forms with potassium hydroxide MnO : and potassium ferricyanide. With sulphuric acid, manganous sulphate and hydroferricyanic acid. Ferricyanides when boiled with NH,OH give ferroeyanides (Playfair, J. C> 1857,9,128). HN0 2 forms first hydroferricyanic acid, then hydronitroferricyanic acid and NO. HN0 3 forms hydroferricyanic acid, and then hydronitroferricyanic acid, NO* being evolved. CI forms first hydroferricyanic and hydrochloric acids. Excess of chlorine to be avoided in preparation of ferricyanides. HC10 a forms hydroferricyanic and hydrochloric acids. Br forms hydroferricyanic and hydrobromic acids. HBrO, forms hydroferricyanic and hydrobromic acids. X , iodine is decolored by potassium f errocyanide, and some potassium ferri- cyanide and potassium iodide are formed. The action is slow and never complete (Gmelin's Hand-book, 7, 459). HI0 3 forms hydroferricyanic acid and free iodine. In analysis, soluble ferroeyanides are recognized by their reactions with ferrous and ferric salts and cepper salts (see 66, §126 and §77). Separated from ferricyanide, by insolubility of alkali salt in alcohol. Ferroeyanides are estimated in solution with sulphuric acid by titrating with standard KMn0 4 . Also by precipitation with CuS0 4 either for gravimetric de- termination or volumetrically, using a ferric salt as an external indicator. §232. Hydroferricyanic acid. H 3 Fe(CN) 6 = 215.164 . H' 3 Fe'"(CN)-' 6 . Absolute hydroferricyanic acid, H 3 Fe(CN) 6 , is a non-volatile, crystallizable 6olid, readily soluble in water, with a brownish color, and an acid reaction to test-paper. It is decomposed by a slight elevation of temperature. In the transposition of most ferricyanides, by sulphuric or other acid, the hydro- ferricyanic acid radical is broken up. Potassium ferricyanide is the usual starting point in the preparation of most ferricyanides. It is prepared by passing chlorine into a cold solution of K 4 Fe(CN) until a few drops of the liquid give a brownish color, but no pre- cipitate with a ferric salt. The solution is evaporated to crystallization and the salt repeatedly recrystallized from water. Large red prismatic crystals, very soluble in water, freely soluble in alcohol (distinction from K,F6(CH') ). The free acid is made by adding to a cold saturated solution of K 3 Fe(CN)„ three volumes of concentrated HC1 and drying the precipitate which forms, in a vacuum (Joannis, C. r., 1882, 94, 449, 541 and 725). Lustrous, brownish- green needles, very soluble in waier and alcohol, insoluble in ether. The ferricyanides of the metals of the alkalis and alkaline earths are soluble in water; those of most of the other metals are insoluble or sparingly soluble. The soluble ferricyanides have a red color, both in crystals and solution; those insoluble have different, strongly marked colors. Potassium and sodium ferri- cyanides are but slightly, or not at all, precipitated from their water solutions by alcohol (separation from ferroeyanides). _ Ferricyanides are not easily decomposed by dilute acids; but alkali hydrox- ides, either transpose them or decompose their radicals (§230, 6). 270 BTDROFERRICTANIC ACID. §232. Solutions of metallic ferricyanides give, with soluble salts of: Aluminum, no precipitate. Antimony, no precipitate. Bismuth, light-brown precipitate, BiFe(CN)„ , insoluble in HC1 . Cadmium, yellow precipitate, Cd a [Fe(CN) 6 ] 2 , soluble in acids and in ammo- nium hydroxide. Chromium, no precipitate. Cobalt, brown-red precipitate, Co 3 [Fe(CN) e ] 2 , insoluble in acids. With ammo- nium chloride and hydroxide, excess of ferricyanide gives a blood-red solution, a distinction of cobalt, from nickel, manganese and zinc. Copper, a yellow-green precipitate, Cu„[Fe(CN) 6 ] 2 , insoluble in HC1. Gold, no precipitate. Iron (ferrous), dark Hue precipitate, Fe3[Fe(CN) e ] 2 , insoluble in acids. Iron (ferric), no precipitate, a darkening of the liquid. Lead, no precipitate, except in concentrated solutions (dark brown). Manganese, brown precipitate, Mn„[Fe(CN) 6 ] 2 , insoluble in acids. Mercury (mercurous), red-brown precipitate, turning white on standing. Mercury (mercuric), no precipitate. Nickel, yellow-green precipitate, Ni s [Fe(CN),,] 2 , insoluble in hydrochloric acid. With ammonium chloride and hydroxide, excess of ferricyanide gives a copper-red precipitate. Silver, a red-brown precipitate, Ag 3 Fe(CN) a , soluble in NHjOH . Tin (stannous), white precipitate, Sn a [Fe(ClT) 6 ] 2 , soluble in hydrochloric acid. Tin (stannic), no precipitate. Uranium (uranous), no precipitate. Zinc, orange precipitate, Zn 3 [Fe(CN) ] 2 , soluble in HC1 and in NH 4 OH . Ferricyanides, ferric combinations, are capable of acting as oxidizing agents, becoming ferrocyanides, ferrous combinations. 4K s Fe(ClSr) 8 + 2H 2 S = 3K 4 Fe(CN)„ + H 4 Fe(CN), + S 2 2K 3 Fe(CN)„ + 2KI = 2K 4 Fe(CN)„ + I 2 . Nitric acid, or acidulated nitrite, by continued digestion in hot solution, efEects a still higher oxidation, of ferricyanides, with the production, among other products, of nitroferricyanides or nitroprussides (Playfair, Phil. Mag., 1845, (3), 26, 197, 271 and 348). These salts are generally held to have the composi- tion represented by the acid H 2 Fe(NO) (CN) B . Sodium nitroprusside is used as a reagent for soluble sulphides — that is, in presence of alkali hydroxides, a test for hydrosulphuric acid; in presence of hydrosulphuric acid, a test for alkali hydroxides (§207, 66). K 3 Fe(CN) 6 is reduced to K 4 Fe(CN)„ by Pd , Th , Mg and As, but not by Pb , Hg , Ag , Sb , Sn , Au , Pt , Bi , Cu , Cd , Te , Al , Fe , Co , Mn , Zn and In . When a sheet of any metal except Au and Pt is placed in contact with a solution of K„Fe(CN)„ and FeCl 3 , a coating of Prussian blue is soon formed (Boettger, J. C, 1873, 26, 473). Pb" with potassium hydroxide forms Pb0 2 and potassium ferrocyanide (Watts' Dictionary, 1889, 2, 340). Sn" with potassium hydroxide forms potassium stannate, K 2 Sn0 3 and potas- sium ferrocyanide (Watts' Dictionary, I. c.). Cr'" forms in alkaline mixture a chromate and a ferrocyanide (Bloxam, C. N., 1885, 52, 109). Mn" with potassium hydroxide forms Mn0 2 and potassium ferrocyanide (Boudault, J. pr., 1845, 36, 23). Co" and Ni" are not oxidized. In alkaline solutions K 8 Fe(CN) oxidizes sugar, starch, alcohol, oxalic acid and indigo (Wallace, J. C, 1855, 7, 77; Mercer, Phil. Mag., 1847, (3), 31, 126). HNOj and HN0 3 both form hydronitroferricyanic acid, H 2 Fe(NO)(CN) B . NO in alkaline solution becomes a nitrate (Wallace, I. a). P in alkaline solution becomes a phosphate (Wallace, I. a). §233. CYANIC ACID. 271 HH 2 P0 2 forms H 4 Fe(CN)„ and H 3 P0 4 . H 2 S forms S , then H 2 S0 4 and H 4 Pe(CN)„ (Wallace, I. c). SO a forms H 2 S0 4 and H 4 Fe(CN)„ . CI decomposes f erricyanides. HC10„ acts upon K a Fe(CN)„ , forming potassium superferricyanide, X 2 Fe(CN)» (Skraup, A., 1877, 189, 368). V ' HI forms H 4 Fe(CN) 6 and I . Ferricyanides in solution are detected by the reactions with ferrous and ferric salts (§126, 66). Insoluble compounds are ignited (under a hood) with a fixed alkali, giving an alkali cyanide, ferric oxide, and an oxide of the metal in combination. Detect the alkali cyanide as directed (§230, 8). A ferri- cyanide is estimated by reduction to f errocyanide with KI in presence of con- centrated HC1; the liberated iodine being titrated with standard Na 2 S 2 0„ . Or it is reduced to ferrocyanide by boiling with KOH and FeS0 4 , filtering, acidulating with H 2 SO, and titrating with KMn0 1 . §233. Cyanic acid. HCNO = 43.048 . H — — C=N. The cyanates of the alkalis and of the fourth-group metals may be made by- passing cyanogen gas into the hydroxides. The cyanates of the alkalis are easily prepared by fusion of the cyanide with some easily reducible oxide. C 2 N 2 + 2KOH = KCNO + KCN + H 2 KCN + PbO = KCNO + Pb 4KCN + Pb„0 4 = 4KCNO + 3Pb The free acid may be obtained by heating cyanuric acid, H 3 C 3 N a 3 , to Tedness, better in an atmosphere of C0 2 . Cyanic acid is found in the dis- tillate. H 3 C 3 N 3 3 = 3HCNO . Absolute cyanic acid, HCNO , is a colorless liquid, giving off pungent, irri- tating vapor, and only preserved at very low temperatures. It cannot be formed by transposing metallic cyanates with the stronger acids in the pres- ence of water, by which it is changed into carbonic anhydride and ammonia: HCNO ,+ H 2 = NH 3 + C0 2 . The cyanates, therefore, when treated with hydrochloric or sulphuric acid, effervesce with the escape of carbonic anhydride (distinction from cyanides), the pungent odor of cyanic acid being perceptible: 2KCNO + 2H 2 S0 4 + 2H 2 = K 2 S0 4 + (NH 4 ) 2 S0 4 + 2C0 2 . The ammonia remains in the liquid as ammonium salt, and may be detected by addition of potassium hydroxide, with heat. The cyanates of the metals of the alkalis and of calcium are soluble in water; most of the others are insoluble or sparingly soluble. All the sohitions gradually decompose, with evolution of ammonia. Silver cyanate is sparingly soluble in hot water, readily soluble in ammonia; soluble, with decomposition, in dilute nitric acid (distinction from cyanide). Copper cyanate is precipitated greenish-yellow. Ammonium cyanate in solution changes gradually, or immediately when boiled, to urea, or carbamide, with which it is isomeric: NH 4 CNO = C0(NH 2 ) 2 . The latter is recognized by the characteristic crystalline laminse of its nitrate, when a few drops of the solution, on glass, are treated with a drop of nitric acid. Also, solution of urea with solution of mercuric nitrate, forms a white precipitate, CH 4 N 2 0(HgO) 2 , not turned yellow (decomposed) by solution of sodium carbonate (no excess of mercuric nitrate being taken). Solution of urea, on boiling, is resolved into ammonium carbonate, which slowly vaporizes: CH 4 N 2 + 2H 2 = (NH 4 ) 2 C0 3 . Cyanates, in the dry way, are reduced by strong deoxidizing agents to cyanides. For detection of a cyanate in presence of cyanides, see Schneider, B., 1895, -28, 1540. 272 THIOCYANIC ACID. §234. §234. Thiocyanic acid. HCNS = 59.118 . H — S — C = N. An aqueous solution of HCNS may be obtained by treating lead thiocyanate suspended in water with H 2 S , also by treating barium thiocyanate with H.SO^ in molecular proportions. The anhydrous acid is obtained by treating dry 35(0118)2 with H 2 S . Potassium thiocyanate is formed by fusing KCN with S. Or two parts of K 4 Fe(CN) 6 with one part of sulphur. Also by fusing the cyanide or ferrocyanide of potassium with potassium thiosulphate, K 2 S 2 3 : 2KCN + S 2 = 2KCNS K t Fe(CN), + 3S 2 = 4KCNS + Fe(CNS) 2 4KCN + 4K 2 S 2 3 = 4KCNS -f 3K 2 S0 4 + K 2 S 2K 4 Fe(CN) + 12K 2 S 2 0„ = 12KCNS + 9K 2 S0 4 + K 2 S + 2FeS Thiocyanic acid is quite as frequently called sulphocyanic acid, and its salts either thiocyanates or sulphocyanates. It corresponds to cyanic acid, HCNO , oxygen being substituted for sulphur. Absolute thiocyanic acid, HCNS , is a colorless liquid, crystallizing at 12° and boiling at 85°. It has a pungent, acetous odor, and reddens litmus. It is soluble in water. The absolute acid decomposes quite rapidly at ordinary temperatures; the dilute solution slowly; with evolution of carbonic anhydride, carbon disulphide, hydrosulphuric acid, hydrocyanic acid, ammonia, and other prodiicts. The same products result, in greater or less degree, from transposing soluble thiocyanates with strong acids; in greater degree as the acid is stronger and heat applied; while in dilute cold solution, the most of the thiocyanic acid remains undecomposed, giving the acetous odor. The thiocyanates, insoluble in water, are not all readily transposed. Thiocyanates of metals, whose sul- phides are insoluble in certain acids, resist the action of the same acids. The thiocyanates of the metals of the alkalis, alkaline earths; also, those of iron (ferrous and ferric), manganese, zinc, cobalt and copper — are soluble in water. Mercuric thiocyanate, sparingly soluble; potassium mercuric thiocyanate, more soluble. Silver thiocyanate is insoluble in water, insoluble in dilute nitric acid, slowly soluble in ammonium hydroxide. Solutions of metallic thiocyanates give, with soluble salts of: Cobalt, very concentrated, a blue color, Co(CNS) 2 , crystallizable in blue needles, soluble in alcohol, not in carbon disulphide. The coloration is promoted by warming, and the test is best made in an evaporating dish. In strictly neutral solutions, iron, nickel, zinc and manganese, do not interfere. Copper, if concentrated, a black crystalline precipitate, Cu(ClTS) 2 , soluble in thiocyanate. With sulphurous acid, a white precipitate, CuCNS; also with hydrosulphuric acid (used to separate a. thiocyanate from a chloride) (Mann, Z., 1889, 28, 668). Iron (ferrous), no precipitate or color. Iron (ferric), an intensely blood-red solution of Fe(CNS) s , decolored by solu- tion of mercuric chloride (§126, 66, distinction from acetic acid); decolored by phosphoric, arsenic, oxalic and iodic acids, etc., unless with excess of ferric salt; decolored by alkalis and by nitric acid, not by dilute hydro- chloric acid. On introduction of metallic zinc, it evolves hydrosulphuric acid. Ferric thiocyanate is soluble in ether, which extracts traces of it. from aqueous mixtures, rendering its color much more evident by the concentration in the ether layer. Z>ead, gradually, a yellowish crystalline precipitate, Pb(CNS) 2 , changed by boiling to white basic salt. Mercury (mercurous), a white precipitate, HgCNS , resolved by boiling into Hg and Hg(CNS) 2 . The mercurous thiocyanate, HgCNS , swells greatly on ignition (being used in "Pharaoh's serpents "), with evolution of mer- cury, nitrogen, thiocyanogen, cyanogen and sulphur dioxide. §235, 1. NITROGEN. 273 Mercury (mercuric), in solutions not very dilute, a white precipitate, Hg(CNS) 2 , somewhat soluble in excess of the thiocyanates, sparingly soluble in water, moderately soluble in alcohol. On ignition, it swells like the mercurous precipitate. Platinum. Platinic chloride, gradually added to a hot, concentrated solution of potassium thiocyanate, forms a deep-red solution of double thiocyanate of potassium and platinum (KCNS) 2 Pt(CN'S) 4 , or more properly, K 2 Pt(CNS) , potassium thiocyanoplatinate. The latter salt gives bright-colored precipi- tates with metallic salts. The thiocyanoplatinate of lead (so formed) is golden-colored; that of silver, orange-red. Silver, a white precipitate, AgCNS , insoluble in water, insoluble in dilute nitric acid, slowly soluble in ammonium hydroxide, readily soluble in excess of potassium thiocyanate; blackens in the light; soluble in hot concentrated H 2 S0, (separation from AgCl) (Volhard, A., 1877, 190, 1). Certain active oxidizing agents, viz., nascent chlorine, and nitric acid contain- ing nitrogen oxides, acting in hot, concentrated solution of thiocyanates, pre- cipitate perthiocyanogen, H(CNS) 3 , of a yellow-red to rose-red color, even blue sometimes. It may be formed in the test for iodine, and mistaken for that element, in starch or carbon disulphide. If boiled with solution of potassium hydroxide, it forms thiocyanate. Concentrated hydrochloric acid, or sulphuric acid, added in excess to water solution of thiocyanates, causes the gradual formation of a yellow precipitate, pertMooyanie acid, (HCN) 2 S„ , slightly soluble in hot water, from which it crystallizes in yellow needles. It dissolves in alcohol and in ether. Potassium thiocyanate can be fused in closed vessels, without decomposition; but with free access of air, it is resolved into sulphate and cyanate, with evolution of sulphurous acid. When thiocyanic acid is oxidized, the final product, as far as the sulphur is concerned, is always sulphuric acid or a sulphate. In many cases (in acid mix- ture) it has been proven that the cyanogen is evolved as hydrocyanic acid. In other cases the game reaction is assumed as probable. Pb0 2 and Pb 8 4 form Pb" and sulphuric acid, in acid mixture only (Hardow, J. C, 1859, 11, 174). H 3 As0 4 forms H a AsO, , hydrocyanic and sulphuric acids. Co'" forms Co" , hydrocyanic and sulphuric acids. Hi'" forms Ni" , hydrocyanic and sulphuric acids. CrVl forms Cr'" , hydrocyanic and sulphuric acids. Mn"+n forms Mn" , hydrocyanic and sulphuric acids. In alkaline mixture, a cyanate and sulphate are formed (Wurtz's Diet. Chim., 3, 95). HNO, forms sulphuric acid and nitric oxide. HN0 S forms sulphuric acid and nitric oxide. CI forms at first a red compound of unknown composition, then HC1 , H 2 S0 4 and HCN are produced. In alkaline mixture a chloride and sulphate are formed. HC10 same as with CI . HC10 3 forms sulphuric, hydrochloric and hydrocyanic acids. Br forms HBr and H 2 S0 4 ; but with alkalis, a bromide and sulphate. HBr0 8 forms HBr and H 2 S0 4 . HIO a forms H,S0, and free iodine. §235. Nitrogen. K = 14.04. Valence one to five (§11). 1. Properties. — Weight of molecule, N 2 , 28.08. Vapor density, 14 (Jolly, W. A., 1879, 6, 536). At — 123.8°, under pressure of 42.1 atmospheres, it condenses to a liquid (Sarrau, C. r., 1882, 94, 718). Boiling point, —194.4° (Olszewski, W. A., 1897, 31, 58). Liquid nitrogen is colorless and transparent. The gas is taste- less, odorless and colorless. Not poisonous, but kills by excluding air from the lungs. Does not burn or support combusion. It is very inert, not attacking other free elements. Its simplest combinations are the following: N— "'H', , ~N 2 , NO , N,0 ;l , 1T0 2 and N 2 0., . The number of organic compounds contain- ing nitrogen is very large. The nitrogen in all compounds that are the 274 HYDRONITRIC ACID. §235, 2. • immediate products of vegetable growth has a valence of minus three and may without change of bonds be converted into N-'"H' S . This statement is made with a limited knowledge of the facts and without, at present, having conclusive proof; and merely predicting that future research will verify it. 2. Occurrence.— It constitutes about four-fifths of the volume of the atmos- phere. It occurs as a nitrate in various salts and in various forms as a con- stituent of animal and vegetable growths. 3. Formation. — (a) From the air, the oxygen being removed by red-hot copper, the C0 2 by potassium hydroxide, the ammonia and water by passing through H 2 S0 4 . (6) Ignition of ammonium dichromate, (NH t ) 2 Cr 2 0, = N 2 + Cr 2 O s -f- 4H 2 . (c) By heating ammonium nitrate and peroxide of manganese to about 200° (Gatehouse, G. N., 1877, 35, 118). (d) Ignition of 1TH 4 C1 and X 2 Cr 2 7 : 2NH 4 C1 + K 2 Cr 2 7 = 2KC1 + N 2 + Cr 2 5 + 4H 2 . Unless the temperature be carefully guarded traces of NO are formed, which may be removed by passing the gases through FeS0 4 . (e) Action of chlorine upon 2ra: s : 8NH 3 + 3C1 2 = 6NH 4 C1 + N 2 . The HE, must be kept in excess to avoid the formation of the dangerously explosive chloride of nitrogen, NC1 3 . (f) Bemoving the oxygen from the air by shaking with NH 4 OH and copper turnings, (g) Burning phosphorus in air over water, (ft) By passing air through a mixture of FeS and sawdust; then through a pyrogallate solution, and finally through concentrated H 2 S0 4 . (i) By shaking air with Fe(0H) 2 and Mn(OH) 2 . (;') By passing air through an alkaline pyrogallate. (fe) By passing air, from which C0 2 has been removed, mixed with hydrogen over heated platinum black, the hydrogen having been added in just sufficient quantity to form water with all the oxygen (Damoulin, /., 1851, 321). (I) By warming a concentrated solution of NH 4 N0 2 or a mixture of KN0 2 and NH 4 Clr KH,N0 2 = N 2 + 2H 2 . Potassium dichromate is added to oxidize to nitric acid any of the oxides of nitrogen that may be formed (Gibbs, B., 1877, 1387). (m) By action of potassium or sodium hypobromite upon ammonium chloride : 3NaBrO -f- 2NH 4 C1 = N 2 "+ 3NaBr + 2HC1 + 3H 2 . 4. Preparation. — Nitrogen has been economically produced by most of the above methods. 5. Solubilities. — Nitrogen is nearly insoluble in all known liquids. 6. Reactions. — At ordinary temperatures nitrogen is not acted upon by other compounds. Nodules growing on the roots of leguminous plants absorb nitro- gen and build up nitrogenous compounds therewith. 7. Ignition. — Under electric influence it combines slowly with hydrogen; also with B , Cr , Mg , Si and V . 8. Detection. — Nitrogen is more easily detected by the nature of its com- pounds than by the properties of the liberated element. 9. Estimation. — (o) As free nitrogen by measuring the volume of the gas. (6) By oxidation of the organic substance with hot concentrated H 2 S0 4 , which also converts the nitrogen into ammonium sulphate. For details, see works on organic analysis, (c) By decomposition of the organic material with potas- sium permanganate in strong alkaline solution, forming ammonia, (d) By combustion of the organic compound in' presence of CuO and Cu° , absorbing the C0 2 by KOH and determining the nitrogen by volume. (For Hydroxylamine, see foot-note, page 278.) §236. Hydronitric acid Azoimide). N S H = 43.128. Constitution, || ^)NH N Curtius, B., 1890, 23, 3023. A clear mobile liquid of a penetrating odor, a very irritative effect upon the nostrils and the skin, and readily exploding with exceeding violence. Boiling point, about 37°. Soluble in water and alcohol. An acid of marked activity, dissolving a number of metals with evolution of hydrogen. Its salts, the trinitrides of the metals of the alkalis §238,6. NITROUS OXIDE— NITRIC OXIDE. 275 and the alkaline earths, are soluble in water and erystallizable (Dennis, J. Am. Soc, 1898, 20, 225). Potassium trinitride precipitates from thorium salts, the hydroxide of this metal in quantitative separation from cerium, lanthanum, neodymium and praseodymium (Dennis, J. Am. Soc, 1896, 18, 947). Hydro- nitric acid is formed by treating ammonia with sodium, and the resulting sodamide, NaNH 2 , with nitrous oxide: 2NaNH 2 + N 2 = NaN 3 + NaOH + NH„ (Wislicenus, B., 1892, 25, 2084). §237. Nitrous oxide. N 2 = 44.08 . N' 2 0- ", N — — N . Nitrous oxide becomes a colorless liquid at 0° under pressure of three atmospheres (Farady, A., 1845, 56, 157). Melts at — 99° and boils at — 92° (Wills, J. 0., 1874, 27, 21). It is a colorless gas with slight sweetish smell and taste. Supports combustion. When breathed acts as an anaesthetic of short duration; and is used in dentistry for that purpose. Decomposed by heat completely at 900° into N and O (Meyer, Pyrochemisch. Untermcli., 1885). Passed over red-hot iron N and Fe 2 3 are formed. K and Na burn in nitrous oxide, liberating the nitrogen. As a rule both gases and solids that burn in air burn also in nitrous oxide. It is formed: (a) By heating ammonium nitrate in a retort from 170° to 260°: NH 4 N0 3 = N 2 + 2H 2 . (6) By passing NO through solution of S0 2 . (c) By action of HNO, ; sp. gr., 1.42, diluted with an equal volume of water, upon metallic zinc. (. It is a base with an alkaline reaction and a strong reducing agent. When pure it is a crystalline solid, odorless, melting at 33.05°, boiling at 68° at 23 mm. pressure ; oxidized by oxygen to HN0 2 (Lobry de Bruyn, B., 1892, 25, 3, 190 and 684). It is a good antiseptic and preservative. It combines with acids to form salts : SfH a OH + HC1= NH 2 OH . HC1. Hydroxylamine hydrochloride is decomposed by alkalis forming the free base, which is decomposed by the halogens, KJInO,, K 2 Cr.O„ BaO a and PbO a . Its solution in ethsr reacts with sodium forming a white precipitate of If H 2 ONa. §241, 7. NITRIC ACID. 279 1. Pb0 2 is not changed. Pb 3 4 is changed thus: Pb 3 4 -f 4HN0 3 = Pb0 2 + 2Pb(N0 3 ) 2 + 2H 2 . 2. Hg' becomes Hg". 3. Sn" becomes Sn Iv . Stannous chloride and hydrochloric acid, heated with a nitrate, form stannic chloride, and convert nitric acid to ammonia (which remains as ammonium salt). See §71, 6c. 4. Sb'" becomes Sb v , forming Sb 2 B , insoluble. 5. As'" becomes As v , forming H 3 As0 4 . 6. Cu' becomes Cu". 7. Fe" becomes Fe"'. B. — With non-metals and their compounds. 1. Carbon (ordinary, not graphite) becomes C0 2 if the nitric acid be hot and concentrated. H 2 C 2 4 becomes C0 2 , in hot concentrated acid. H 4 Fe(CN)„ becomes first H 3 Fe(CN) 6 and then hydronitroferricyanic acid. HCNS is oxidized, the sulphur becoming H 2 S0 4 . 2. Nitrites are all decomposed, nitrates being formed, the nitric acid not being reduced. The nitrous acid liberated immediately dissociates: 3HN0 2 = 2N0 + HN0 3 + H 2 . S. P°, PH 3 , HH 2 P0 2 and H 3 P0 3 become H 3 P0 4 . That is P v - n becomes P v . h- S becomes H 2 S0 4 . H 2 S becomes first S° and then H 2 S0 4 . H„S0 3 becomes H 2 S0 4 ; and in general S^f 1 becomes S'". 5. HC1, nitrohydrochloric acid: 2HN0 3 + 6HC1 = 2N0 + 4H 2 + 3C1 2 (Koninck and Nihoul, Z. anorg., 1890, 477). See §269, 6B2. HC10 3 is not reduced. Chlorates are all transposed but not decom- posed until the temperature and degree of concentration is reached that would dissociate the HC10 3 if the nitric acid were absent. 6. Br" is not oxidized. HBr becomes Br° and is not further oxidized. All bromates are transposed but the HBr0 3 is not decomposed until a tem- perature and degree of concentration is reached that would cause the dissociation of the HBr0 3 if the nitric acid were absent. 7. 1° becomes HI0 3 . Very slowly unless the fuming nitric be used. HI become first 1° ; then as above. 8. In general organic compounds are oxidized. Straw, hay, cotton, etc., are inflamed by the strong acid {Kraut, B., 1881, 14, 301). For action on starch, see Lunge, B., 1878, 11, 1229, 1641. With many organic bodies substitr.tion products are formed, the oxides of nitrogen taking the place of the hydrogen. 7. Ignition.— Nitric acid is dissociated by heat: 4HNO s = 4N0 2 + 2H 2 + O a , complete if at 256° (Carius, B., 1871, 4, 828). No nitrates are volatile as such; 280 NITRIC ACID. §241,8. ammonium nitrate is dissociated: WH 4 W0 3 = N 2 + 2H 2 0. Some nitrates, e.g. those of X and Na , are first changed to nitrites with evolution of oxygen 'only' and at an intense white heat further changed to oxides with evolution of ~N 2 6 as well as oxygen. As a final result of ignition the nitrates of all ordinary metals are left as oxides, except that those of Hg , Ag , Au and Pt are reduced to the free metal. A mixture of potassium nitrate and sodium carbonate in a state of fusion is a powerful oxidizer; e. g., changing Sn" to Smy , As'" to Asv , Sb"' to Sbv Te" to Pe'" , Cr'" to Crvi , jinvi-n to Mnvi , svi-n to Svi , etc. Seated on charcoal, or with potassium cyanide, or sugar, sulphur or other easily oxidizable substance (as in gunpowder), nitrates are reduced with deflagration or explosion, more or less violent. With potassium cyanide, on platinum foil, the deflagration is especially vivid. In this reaction free nitrogen, is evolved. Strongly heated with excess of potassium hydroxide and sugar or other carbonaceous compound, in a dry mixture, nitrates are reduced to ammonia, which is evolved, and may be detected. In this carbdnaceous mixture, the nitrogen of nitrates reacts with alkalis, like the unoxidized nitrogen in car- bonaceous compounds. 8. Detection. — Most of the tests for the identification of nitric acid are made by its deoxidation^ disengaging a lower oxide of nitrogen, or even, by complete deoxidation, forming ammonia. If, with concentrated sulphuric acid, a bit of copper turning, or a crystal of ferrous sulphate, is added to a concentrated solution or residue of nitrate, the mixture gives off abundant brown vapors; the colorless nitric oxide, NO , which is set free from the mixture, oxidizing immediately in the air to nitrogen peroxide, N0 2 : 2KlTO a + 4H 2 S0 4 + 3Cu = K 2 S0 4 + 3CuS0 4 + 4H 2 + 2NO 2KNO s + 4H 2 S0 4 + 6FeS0 4 = K 2 S0 4 + 3Fe 2 (S0 4 ) 3 + 4H 2 + 2NO The three atoms of oxygen furnished by two molecules of nitrate suffice to oxidize three atoms of copper; so that 3Cu0 with 3H 2 S0 4 , may form 3CuS0 4 and 3H 2 . The same three atoms of oxygen (having six bonds) suffice to oxidize six molecules of ferrous salt into three molecules of ferric salt; so that 6FeS0 4 with 3H 2 S0 4 , can form 3Fe 2 (S0 4 ) 3 and 3H 2 . Now if, by the last-named reaction, the nitric oxide is disengaged in cold solution, with excess of ferrous salt and of sulphuric acid, instead of passing off, the nitric oxide combines with the ferrous salt, forming a black-brown liquid, (FeS0 4 ) 2 N0 , decomposed by heat and otherwise un- stable: 2KN0 3 + 4H 2 S0 4 + 10FeS0 4 = K 2 S0 4 + 3Fe 2 (S0 4 ) 3 + 4H 2 + 2(FeS0 4 ) 2 N0 . a. — This exceedingly delicate " Brown ring " test for nitric acid or nitrates in solution may be conducted as follows: If the solution of a nitrate is mixed with an equal volume of concentrated H 2 S0 4 , the mixture allowed to cool and a concentrated solution of FeS0 4 then cautiously added to it, so that the fluids do not mix, the junction shows at first a purple, afterwards a brown color (Fresenius, Qual. Anal., 16th ed., 387). A second method of obtaining the same brown ring is: Take sulphuric acid to a §241, 8h. NITRIC ACID. 281 quarter of an inch in depth in the test-tube ; add without shaking a nearly- equal bulk of a solution of ferrous sulphate, cool; then add slowly of the solution to be tested for nitric acid, slightly tapping the test-tube on the side but not shaking it. The brown ring forms between the two layers of the liquid. A third method often preferred is: Take ferrous sulphate solution to half an inch in depth in the test-tube; add two or three drops of the liquid under examination and mix thoroughly; incline the test-tube and add an equal volume of concentrated H 2 S0 4 in such a way that it will pass to the bottom and form a separate layer. Cool and let it stand a few minutes without shaking. I. — Indigo solution. — In presence of HC1 heat moderately and blue color is destroyed. Interfering substances, HC10 3 , HI0 3 , HBr0 3 , Fe'", Cr", Mn™ and all that convert HC1 into CI . c. — Sodium salicylate is added to the solution, H 2 S0 4 is slowly added, test-tube being inclined. Avoid shaking, keep cool for five minutes. A yellow ring indicates HN0 3 . To increase the brilliancy of the color, shake* cool and add to HN 4 0H . d. — Ammonium test. — Treat the solution with KOH and Al wire, warm until gas is evolved. Pass the gas into water containing a few drops of Nessler's reagent. A yellowish-brown precipitate indicates HN0 3 : 3HN0 3 + 8A1 + 8K0H = 3NH 3 + 8KA10 2 + H 2 . Nothing interferes with this test, but action is delayed by Cl v , I v and many other oxidisers. e. — Nitrite test. — Eeduce the nitrate to nitrite by warming with Al and KOH . At short intervals decant a portion of the solution, add a drop of XI, acidify with HC 2 H 3 2 and test for I with CS 2 . This test should always be made in connection with (d). Other oxidisers including Cl v , Br v , I 9 , and As v are reduced before the reduction of the HN0 3 begins : 3HNO a + 2A1 + 5KOH = 3K2T0 2 + 2KAJ0 2 + 4H 2 2KN0 2 + 2KI + 4HC 2 H S 2 = I, + 4KC 2 H 3 2 + 2H 2 + 2NO Other means of making the nascent hydrogen are sometimes preferred; e. g., sodium amalgam, a mixture of Zn and Fe both finely divided and used with excess of hot KOH , or finely divided Mg in presence of H 3 P0 4 . /. — Add three drops of the solution to- be tested to two drops of diphenylamine, (C 6 H 5 ) 2 NH , dissolved in H 2 S0 4 . A blue color indicates a nitrate, Cl°, Cl v , Br v , I v , Mn vn , Cr VI , Se lv , and Fe'" interfere with this test. g. — Brucine, dissolved in concentrated sulphuric acid, treated (on a porcelain surface) with even traces of nitrates, gives a fine deep-red color, soon paling to reddish-yellow. If now stannous chloride, dilute solution, be added, a fine red- violet color appears. (Chloric acid gives the same reaction.) h. — Phenol, C H OH , gives a deep red-brown color with nitric acid, by for- mation of nitrophenol (mono, di or tri), C 6 H 4 (N0 2 )OH to C 6 H 2 (N0 2 ) ? OH , " picric acid " or nitrophenic acid. A mixture of one part of phenol (cryst. carbolic acid), four parts of strong sulphuric acid, and two parts of water, 282 OXYGEN. §241, 8i. constitutes a reagent for a very delicate test for nitrates (or nitrites), a few drops being sufficient. With unmixed nitrates the action is explosive, unless upon very small quantities, , The addition of potassium hydroxide deepens and brightens the color. According to Sprengel (/. C, 1863, 16, 396), the some- what similar color given by compounds of chlorine, bromine, iodine and by organic matter may be removed by adding ammonium hydroxide without diminishing the brightness of the color formed by the nitrates. i. — According to Lindo (C. N., 1888, 58, 176), resorcinal is five times more delicate a test than phenol. Ten grammes of resorcinol are dissolved in 100 cc. of water; one drop of this solution with one drop of a. 15 per cent solution of HC1 and two drops of concentrated H 2 S0 4 are added to 0.5 cc. of the nitrate to be tested. Nitrous acid gives the same purple color. ;'. — A little pyrogallol is dissolved in the liquid to be tested (less than one mg. to one cc.) and ten drops of concentrated H 2 'S0 4 are dropped down the side of the test tube so as to form two layers; at the surface of contact a brown or yellow coloration appears if nitric acid is present. One mg. of nitric acid in one litre of potable water can thus be detected (Curtman, Arch. PTiarm., 1886, 223, 711). 9. Estimation. — (a) If the base is one capable of readily forming a silicate, the nitrate is fused with Si0 2 and estimated by the difference in weight. (6) By treating with hot sulphuric acid, passing the distillate into BaC0 3 and esti- mating the nitric acid by the amount of barium dissolved, (c) Treating with Al and KOH and estimating the distillate as NH S . (d) Neutralizing the free acid with ammonium hydroxide, and after evaporation and drying at 115°, weighing as ammonium nitrate, (e) In presence of free H,SO, a ferrous solu- tion of known strength is added in excess to the nitrate aitd the amount of ferrous salt remaining is determined by a, standard solution of potassium permanganate, (f) The volume of hydrogen generated by the action of potas- sium hydroxide upon a known quantity of aluminum is measured; and the test is then repeated under the same conditions, but in presence of the nitrate. The difference in the volume of the hydrogen obtained represents the quantity of NH, that has been formed. i §242. Oxygen. = 16.000 . "Usual valence two. 1. Properties.— A colorless, odorless gas; specific gravity, 1.10562 (Crafts, C. r., 1888, 106, 1662). When heated it diffuses through silver tubing quite rapidly (Troost, G. r., 1884, 98, 1427). It liquifies by cooling the gas under great pres- sure and then suddenly allowing it to expand under reduced pressure. It hoils at — 113° under 50 atmospheres pressure; and at — 184° under one atmosphere pressure (Wroblewski, C. r., 1884, 98, 304 and 982). Its critical temperature is about — 118°, and the critical pressure 50 atmospheres. Specific gravity of the liquid at —181.4°, 1.124 (Olszewski, M., 1887, 8, 73). Oxygen is sparingly soluble in water with a slight increase in the volume (Winkler, B., 1889, 22, 1764). Slightly soluble in alcohol (Carius, A., 1855, 94, 134). Molten silver absorbs about ten volumes of oxygen, giving it up upon cooling (blossoming of silver beads) (Levol, G. r., 1852, 35, 63). It transmits sound better than air (Bender, B., 1873, 6, 665). It is not combustible, but supports combustion much better than air. In an atmosphere of oxj'gen, a glowing splinter bursts into a flame; phosphorus burns with vivid incandescence; also an iron watch spring heated with burning sulphur. It is the most negative of all the elements except fluorine; it combines directly or indirectly with all the elements except fluorine; with the alkali snetals rapidly at ordinary temperature. The combination of oxygen with elements or compounds is termed combustion or oxidation. The temperature at which the combination takes place varies greatly: Phosphorus at 60°; hydrogen in air at 552°; in pure oxygen at 530° (Mallard and Le Chate- lier, Bl., 1883, (2), 39, 2); carbon disulphide at 149°; carbon at a red heat; while the halogens do not combine by heat alone. 2. Occurrence. — The rocks, clay and sand constituting the main part of the earth's crust contain from 44 to 48 per cent of oxygen; and as water contains §242,4/. OXYGEN. ri . 283 88.81 per cent, it has been estimated that one-half of the crust is oxygen. Except in atmospheric air, which contains about 23 per cent of uncombined oxygen, it is' always found combined. 3. Formation.— (a) By igniting HgO . (6) By heating KC10 3 to 350°, KC10 4 is produced and oxygen is evolved; at a higher temperature the KC10 4 becomes KC1 . In the presence of MnO. the KCIOiYis completely changed to KC1 at 200°, without forming KC10 4 , the Mn0 2 not being changed. Spongy platinum, CuO , Fe 2 3 , Pb0 2 , etc., may be substituted for MnO, (Mills and Donald, J. C, 1882, 41, 18; Baudrimont, Am. S., 1872, 103, 370). Spongy platinum, ruthenium, rhodium and indium with chlorine water or with hydrogen peroxide evolve oxygen. The spongy ruthenium acts most energetically (Schoenbein, A. Ch., 1866, (4), 7, 103). (c) Action of heat on similar salts furnishes oxygen; e. g., KC10 and KC10 2 form KC1 , KBr0 3 forms KBr , KI0 3 and KI0 4 form KI , and KN0 8 forms KNO, (at a white heat K 2 , NO and are formed), (d) By the action of heat on metallic oxides as shown in the equations below, (e) By, heating higher oxides or their salts with sulphuric acid. Crvi iS changed to Cr"' , Co'" to Co" , Ni"' to Ni" , Biv to Bi'" , Fevi to Fe"' , Pbiv to Pb" , and ]ffln"+n to Mn"; in each case a sulphate is formed and oxygen given off: ' a. 2HgO (at 500°) = 2Hg + 2 6. 10KClO 3 (at 350°) = 6KC10 4 + 4KC1 + 30 2 (Teed, /. C, 1887, 51, 283) 2KC10 3 (at red heat) = 2KC1 + 30 2 2KC10, + nMmO, (at 200°) = nMiiO, + 2KC1 + 30, c. KC10 2 = KC1 + 2 2KBrO„ = 2KBr + 30 2 2KIO a = 2KI + 30 2 XI0 4 = KI + 20 2 2KN0 3 '= 2KN0 2 + 2 4KN0 2 (white heat) = 2K 2 + 4N0 + 0, 4. 2Pb 3 4 (white heat) = 6PbO + 2 SSb 2 B (red heat) = 2Sb 2 4 + 2 ' Bi 2 O s (red heat) = Bi 2 3 + 2 :4CrO s (about 200°) =2Cr a O, + 30 2 4K 2 Cr 2 7 (red heat) = 2Cr 2 3 + 4K 2 Cr0 4 + 30, 6Fe 2 3 (white heat) = 4Fe 3 4 + 3 3Mn0 2 (white heat) = Mn.O, + 2 /!' 6Co 2 3 (dull-red heat) = 4Co s 4 + 2 2Mi 2 8 (dull-red heat) = 4MiO + 2 2Ag 2 (300°) =4Ag + 2 2Ba0 2 (800°) = 2BaO + 2 e. 2K 2 Cr 2 7 + 8H 2 S0 4 = 4KCr(S0 4 ) 2 + 30 2 + 8H 2 4KMn0 4 + 6H 2 S0 4 = 2K 2 S0 4 + 4MnS0 4 -f- 50 2 + 6H a O 2Pb 3 4 + 6H 2 S0 4 = 6PbS0 4 + 6H 2 + 2 4. Preparation.— (a) By heating KCIO s to 200° in closed retorts in the pres- ence of Mn0 2 or Fe 2 3 . If KC10 3 be heated alone, higher heat (350°) is required, and the gas is given off with explosive violence. About equal parts of the metallic oxide and KC10 3 should be taken. (&) BaO heated in the air to 550° becomes Ba0 2 , and at 800° is decomposed into BaO and O , making theoretically a cheap process, (c) By heating calcium plumbate. The calcium plumbate is regenerated by heating in the air (Kassner, J. C, 1894, 66, 11, 89). (d) By passing sulphuric acid over red-hot bricks: 2H 2 S0 4 = 2S0 2 + 2H ,0 + 2 ; the S0 2 is separated by water, and after conversion into H 2 S0 4 (§266, 4) is used over again, (e) By warming a saturated solution of chloride of lime with a small amount of cobaltic oxide, freshly prepared and moist. The cobaltic oxide seems to play the same role as NO in making H 2 S0 4 (Fleitmann, A. Oft., 1865, (4), 5, 507). (f) The following cheap process is now employed on a large scale. Steam is passed over sodium manganate at a dull-red heat; Mn,0, and 284 OZONE. §242, 5. oxygen are formed. Then, without change of apparatus or temperature, air- instead of steam is passed over the mixture of Mi.O, and NaOH . The Mn 2 8 is thus again oxidized to Na 2 Mn0 4 , and free nitrogen is liberated: 4Na 2 Mn0 4 + 4H 2 (dull-red heat) = 8K"aOH + 2Mn 2 O a + 30 2 8NaOH + 2Mn 2 8 + air, 3(0 2 + 4N 2 ) = 4Na 2 Mn0 1 + 4H 2 + 12N 2 " 5. Solubilities. — See 1. 6. Reactions. — Pure oxygen may be breathed for a short time without injury.. A rabbit placed in pure oxygen at 24° lived for three weeks, eating voraciously all the time, but nevertheless becoming thin. The action of oxygen at 7.2° is. to produce narcotism and eventually death. When oxygen is cooled by a. freezing mixture it induces so intense a narcotism that operations may be performed under its influence. Compressed oxygen is " the most fearful poison known." The pure gas at a pressure of 3.5 atmospheres, or air at a pressure of 22 atmospheres, produces violent convulsions, simulating those of strychnia, poisoning, ultimately causing death. The arterial blood in these cases is found, to contain about twice the quantity of its normal oxygen. Further, compressed, oxygen stops fermentation, and permanently destroys the power of yeast. At varying temperatures oxygen combines directly with all metals except silver, gold and platinum, and with these it may be made to combine by pre- cipitation. It combines with all non-metals except fluorine; the combination, occurring directly, at high temperatures, except with CI , Br and I , which require the intervention of a third body. 7. Ignition. — Most elements ' when ignited with oxygen combine readily... Some lower oxides combine with oxygen to form higher oxides, and certain other oxides evolve oxygen, forming elements or lower oxides. Oxides of gold,, platinum and silver cannot be formed by igniting the metals in oxygen; thejr must be formed by precipitation. 8. Detection.— Uncombined oxygen is detected by its absorption by an alka- line solution of pyrogallol; by the combination with indigo white to form indigo blue; by its combination with colorless NO to form the brown N0 2 ; by its combination with phosphorus, etc. It is separated from other gases by its absorption by a solution of chromous chloride, pyrogallol or by phosphorus. In combination in certain compounds it is liberated in whole or in part by simple ignition; as with KC10„ , KMnO, , HgO , Au 2 0„ , Pt0 2 , Ag 2 , Sb 2 5 ,. etc. In other combinations by ignition with hydrogen, forming water. 9. Estimation. — Free oxygen is usually estimated by bringing the gases in contact with phosphorus or with an alkaline solution of pyrogallol (C0 2 having- been previously removed), and noting the dimunition in volume. Oxygen in_ combination is usually estimated by difference. §243. Ozone. 3 = 48.000. — \/ Ozone was first noticed by Van Marum in 1785 as a peculiar smelling gas formed during the electric discharge; and which destroyed the lustre of mercury. Schoenbein (Pogg., 1840, 50, 616) named the gas ozone and noticed its powerful oxidizing properties. It is said to be an ever-present constituent of the air, giving to the sky its blue color; present much more in the country and near the seashore than in the air of cities (Hartley, J. C, 1881, 39, 57 and 111; Houzeau, C. r., 1872, 74, 712). Ozone is always mixed with ordinary oxygen, partly due to dissociation of the ozone molecule, which is stable only at low temperatures (Hautefeuille and Chappuis, C. r., 1880, 91, 522 and 815). It is- prepared by the action of the electric discharge upon oxygen (Bichat and Guntz, 0. r., 1888, 107, 344; Wills, B., 1873, 6, 769). By the oxidation of moist phosphorus at ordinary temperature (Leeds, A., 1879, 198, 30; Marignac, C. r. r 1845, 20, 808). By electrolysis of dilute sulphuric acid, using lead electrodes- §244, 1. HYDROGEN PEROXIDE. 285. (Planti, C. r., 1866, 63, 181). By the action of concentrated sulphuric acid on potassium permanganate (Schoenbein, J. pr., 1862, 86, 70 and 377). Many readily oxidized organic substances form some ozone in the process of oxida- tion (Belluci, B., 1879, 12, 1699). Ozone is a gas, the blue color of which can be plainly noticed in tubes one metre long. Its odor reminds one somewhat of chlorine and nitrogen peroxide, noticeable in one part in 500,000. It acts, upon the respiratory organs, making breathing difficult. When somewhat concentrated it attacks the mucous membrane. It caused death to small animals which have been made to breathe it. For further concerning the physiological action, see Binz, C. C, 1873, 72. Its specific gravity is 1.658 (Soret, A.. 1866, 138, 4). It has been liquified to a deep-blue liquid, boiling at — 106° (Olszewski, M., 1887, 8, 230). The gas is sparingly soluble in water (Carius, B. y 1873, 6, 806). It decomposes somewhat into inactive oxygen at ordinary tem- perature, and completely when heated above 300°, with increase of volume. A number of substances decompose ozone without themselves being changed; e. g., platinum black, platinum sponge, oxides of gold, silver, iron and copper,, peroxides of lead and manganese, potassium hydroxide, etc. It is one of the most active oxidizing agents known, the presence of water being necessary. When ozone acts as an oxidizing agent there is no change in volume; but one-- third of the oxygen entering into the reaction, inactive oxygen remaining. Moist ozone oxidizes all metals except gold and platinum to the highest pos- sible oxides. Pb" becomes Fb0 2 Sn" becomes SnO, Hg' becomes Hg" Bi'" becomes Bi 2 5 Pd" becomes Pd0 2 Cr"' becomes Crvi Fe" becomes Fe 2 O s ; in presence of KOH , K 2 Fe0 4 Sin" becomes MnO,; in presence of H 2 SO, or HNO, , HMnO, is formed. Co" becomes Co'" Ni" becomes MI'" . With the salts of nickel and cobalt the action is slow,, rapid with the moist hydroxides. K 4 Fe(CM) e becomes K 3 Pe(CN) 8 NjOj becomes HN0 3 , in absence of water M0 2 is formed S0 2 becomes H 2 SO, H 2 S becomes S and H 2 , the sulphur is then oxidized to H,S0 4 (Pollacci,. C. C, 1884, 484) P and PH 3 become H 3 P0 4 HC1 becomes CI and H 2 HBr becomes Br and H 2 I becomes HI0 3 and HIO, (Ogier, C. r„ 1878, 86, 722) HI and KI become I and H 2 , then Iv Most organic substances are decomposed; indigo is bleached much more- rapidly than by chlorine (Houzeau, C. r., 1872, 75, 349). Alcohol and ether are rapidly oxidized to aldehyde and acetic acid. Ozone is usually detected by the liberation of iodine from potassium iodide, potassium iodide starch paper being used. Because HNO, and many other- substances give the same reactiou, thallium hydroxide paper is preferred by Schoene (B., 1880, 13, 1508). The paper is colored brown, but the reaction is. much less delicate than with potassium iodide starch paper. It is estimated quantitatively by passing the gas through a solution of KI rendered acid with H 2 S0 4 , and titration of the liberated iodine: 3 + 2HI = 2 + I 2 + H 2 . §244. Hydrogen peroxide. H 2 2 = 34.016 . H — — — H. 1. Properties. — Pure hydrogen peroxide (99.1 per cent) is a colorless syrupy- liquid, boiling at 84° to 85° at 68 mm. pressure. It does not readily moisten the containing vessel. It is volatile in the air, irritating to the skin, and 286 HYDROGEN PEROXIDE. §244,2. reacts strongly acid to litmus. The ordinary three per cent solution can be evaporated on the water bath until it contains about 60 per cent H,0 2 , losing about one-half by volatilization. The presence of impurities causes its decom- position with explosive violence. Before final concentration under reduced pressure it should be extracted with ether (WolfEenstein, B., 1894, 27, 3307). The dilute solutions are valuable in surgery in oxidizing putrid flesh of wounds, etc.; they are quite stable and may be preserved a long time especially if acid (Hanriott, 0. r., 1885, 100, 57). The presence of alkalis decreases the stability. Concentrated solutions evolve oxygen at 20°, and frequently explode when heated to nearly 100°. It contains the most oxygen of any known compound; one-half of the oxygen being available, the other half combining with the hydrogen to form water. 2. Occurrence. — In rain water and in snow (Houzeau, C. r., 1870, 70, 519). It is also said to occur in the juices of certain plants. 3. Formation. — (a) By the electrolysis of 70 per cent H 2 S0 4 (Bicharz, W. A., 1887, 31, 912). (6) By the action of ozone upon ether and water (Berthelot, €. r., 1878, 86, 71). (c) By the action of ozone upon dilute ammonium hydroxide (Carius, B., 1874, 7, 1481). (d) By the decomposition of various peroxides with acids, (e) By the action of oxygen and water on palladium sponge saturated with hydrogen (Traube, B., 1883, 16, 1201). (f) By the action of moist air on phosphorus partly immersed in water (Kingzett, J. C, 1880, 38, 3). 4. Preparation.— Ba0 2 is decomposed by dilute H 2 S0 4 , the BaS0 4 being removed by filtration. The.Ba0 2 is obtained by heating BaO in air or oxygen to low redness. At a higher heat the Ba0 2 is decomposed into BaO and O (Thomsen, B., 1874, 7, 73). Sodium peroxide, Na 2 2 , is formed by heating sodium in air or oxygen (Harcourt, J. C, 1862, 14, 267); by adding H 2 2 to UaOH solution and precipitating with alcohol. Prepared by the latter method it contains water. 5. Solubilities It is soluble in water in all proportions; also in alcohol, which solvent it slowly attacks. Ba0 2 is insoluble in water, decomposed by acids, including C0 2 and H 2 SiP 8 with formation of H 2 2 . Na^O, is soluble in water with generation of much heat. It is a powerful oxidizing agent. 6. Reactions. A.— With metals and their compounds.— Hydrogen peroxide usually acts as a powerful oxidizing agent to the extent of one- half its oxygen. Under certain conditions, however, it acts as a strqng reducing agent. Some substances decompose it into H 2 and without changing the substance employed, e. g., gold, silver, platinum, manganese ■dioxide, charcoal, etc. (Kwasnik, B., 1892, 25, 67). Many metals are oxidized to the highest oxides, e. g., Al , Fe , Mg , Tl , As , etc. Gold and platinum are not attacked. 1. Pb" becomes Pb0 2 (Schoenbein, J. pr., 1862, 86, 129; Jannasch and Lesinsky, B., 1893, 26, 2334). 2. Ag 2 becomes Ag and . 3. HgO becomes Hg and . J,.. Au 2 3 becomes Au and . 5. As'" becomes As v . 6. Sn" becomes Sn Iv . 7. Bi'" becomes Bi v . 8. Cu" in alkaline solution (Fehling's solution) becomes Cu 2 (Hanriott, Bl., 1886, (2), 46, 468). 9. Fe" becomes Fe'" (Traube, B., 1884, 17, 1062). 10. Tl' becomes T1 2 3 (Schoene, A., 1879, 196, 98). §244,96. HYDROGEN PEROXIDE. 287 11. Cr'" becomes Cr^ in alkaline mixture (Lenssen, J. pr., 1860, 81, 278). 12. Cr" with H 2 S0 4 gives a blue color, HCr0 4 , perchromic acid, soon changing to green by reduction to Cr'". By passing the air or vapor through a chromic acid solution, ozone is separated from hydrogen perox- ide, the latter being decomposed (Engler and Wild, B., 1896, 29, 1940). IS. Mn" in alkaline mixture becomes Mn0 2 . In presence of KCN a' separation from Zn (Jannasch and Niederhofheim, B., 1891, 24, 3945; Jannasch, Z. anorg., 1896, 12, 124 and 134). Mn" +n with H 2 S0 4 forms MnS0 4 , oxygen being evolved both from the H 2 2 and from the Mn compound (Brodie, J. C, 185'5, 7, 304; Lunge, Z. angew., 1890, 6). H. BaO , SrO , and CaO become the peroxides. 15. NaOH becomes Na 2 2 .8H 2 . 16. NH 4 0H becomes NH 4 N0 2 (Weith and Webber, B., 1874, 7, 17 and 45). ' B. — With non-metals and their compounds. 1. K 4 Fe(CN) 6 becomes K 3 Fe(CN) 6 (Weltzien, A., 1866, 138, 129); in alkaline solution the reverse action takes place : 2K 3 Fe(CN) e + 2K0H + H 2 2 = 2K 4 Fe(CN) 6 + 2H 2 + 2 (Baumann, Z. angew., 1892, 113). 2. 3 becomes 2 (Schoene, I. c, page 239). .'• 8, H 3 P0 2 becomes H 3 P0 4 . ' k- H 2 S and sulphides, and S0 2 and sulphites, become H 2 S0 4 or sulphates (Classen and Bauer, B., 1883, 16, 1061). 5. CI becomes HC1 (Schoene, I. c, page 254). It is a valuable reagent for the estimation of chloride of lnne : CaOCl 2 -+-• H 2 2 = CaCl 2 + H 2 + 2 (Lunge, Z. angew., 1890, 6). 6. I becomes HI (Baumann, Z. angew., 1891, 203 and 328). KC1 , Or , and KI liberate oxygen from H 2 2 but no halogen is set free; except that with commercial H 2 2 free , iodine may always be obtained from KI (Schoene, A., 1879, 195, 228; Kimgzett, J. C, 1880, 37, 805). 7. Ingition. — The peroxide of barium is formed by igniting BaO to dull red- ness; strong ignition causes decomposition of the BaO, into BaO and O . The peroxide of calcium cannot be formed by ignition of lime in air or oxygen. 8. Detection. — In a dilute solution of tincture of guaiac mixed with malt infusion, a blue color is obtained when H 2 2 is added. To the solution sup- posed to contain H,0„ add a few drops of lead acetate; then KI , starch, and a little acetic acid; with H 2 2 a blue color is produced (Schoenbein, I. c; Struve, Z., 1869, 8, 274). As confirmatory, its action on KJInO, and on K 2 Cr 2 7 should be observed. A ten per cent solution of ammonium molybdate with equal parts of concentrated sulphuric acid gives a characteristic deep yellow color with H 2 2 (Deniges, C. r., 1890, 110, 1007; Crismer, Bl., 1891, (3), 6, 23). H 2 2 gives some extremely delicate color tests with the aniline bases (Ilosvay, B., 1895, 28, 2029; Deniges, J. Pharm., 1892, (5), 25, 591). 9. Estimation. — (a) By measuring the amount of oxygen liberated with Mn0 2 (Hanriott, BL, 1885, (2), 43, 468). (6) By the amount of standard KMnO, 288 FLUORINE. §245. reduced, or by measuring the volume of oxygen set free: 2KMnO, + 3H 2 S0 4 -+- 5H 2 2 = K 2 S0 4 + 2MnS0 4 + 8H 2 + 50 2 . (c) By decomposition of KI in. presence of an excess of dilute H„S0 4 ; and titration of the liberated iodine with standard Na 2 S„0 3 . (d) Dissolve a weighed sample of Ba0 2 in dilute HC1 , add. KaFeCC^e; transfer to an azotometer and add KOH . The volume of oxygen is a measure of the amount of H 2 2 (Baumann, I. a). §245. Fluorine. F = 19.05 . Valence one. Since Davy's experiments in 1813, many others have attempted the isolation of fluorine. In his zeal the unfortunate Louyet fell a victim to the poisonous, fumes which he inhaled. Faraday, Gore, Fremy, and others took up the prob- lem in succession, but it was not ultimately solved until H. Moissan, in 1886, produced a gas which the chemical section of the French Academy of Sciences, decided to be fluorine. Many ingenious experiments had been made in order- to obtain fluorine in a separate state, but it was found that it invariably combined with some portion of the material of the vessel in which the opera- tion was conducted. The most successful of the early attempts to isolate fluorine appears to have been made, at the suggestion of Davy, in a vessel of fluor-spar itself, which could not, of course, be supposed to be in any way affected by it. Moissan's method was as follows: The hydrofluoric acid having been very carefully obtained pure, a little potassium hydrofluoride was dis- solved in it to improve its conducting power, and it was subjected to the action of the electric current in a U tube of platinum, down the limbs of which the electrodes were inserted; the negative electrode was a rod of platinum, and the positive was made of an alloy of platinum with 10 per cent of iridium. The U tube was provided with stoppers of fluor-spar, and platinum delivery tubes. for the gases, and was cooled to — 23°. The gaseous fluorine, which was extri- cated at the positive electrode, was colorless, and possessed the properties of chlorine, but much more strongly marked. It decomposed water immediately,, seizing upon its hydrogen, and liberating oxygen in the ozonized condition; it exploded with hydrogen, even in the dark, and combined, with combustion,, with most metals and non-metals, even with boron and silicon in their crystal- lized modifications. As , Sb , S , I , alcohol, ether, benzol and petroleum took fire in the gas. Carbon was not attacked by it (Moissan, 1886, G. r., 103, 202 and 256; J. C, 50, 1886, 849 and 976; A. Ch., 1891, (6), 24, 224). Fluorine, in several characteristics, appears as the first member of the Chlorine Series of Elements. It cannot be preserved in the elemental state, as it combines with the materials of vessels (except fluor-spar), and instantly decomposes water, forming hydrofluoric acid, HF , an acid prepared by acting on calcium fluoride with sulphuric acid (a). Fluorine also combines with silicon as SiF 4 , silicon fluoride, a gaseous compound, prepared by acting on calcium fluoride and silicic anhydride with sulphuric acid (6). On passing silicon fluoride into water, a part of it is transposed by the water, forming silicic and hydrofluoric acids (c); but this hydrofluoric acid does not at all remain free, but combines with the other part of the fluoride of silicon, as fluosilicic acid (hydrofluosilicic acid), (HF) 2 SiF 4 or H 2 SiF„ (d) (Offermann, Z. angew., 1890, 617). This acid is used as a reagent; forming metallic fluo- silicates (silicofluorides) , soluble and insoluble (§246). a. CaF 2 + H 2 S0 4 = CaS0 4 + 2HF 6. 2CaF 2 + Si0 2 + 2H 2 S0 4 = 2CaS0 4 + 2H 2 + SiF 4 c. SiF 4 + 2H 2 = Si0 2 + 4HF (not remaining free) d. 2HF + SiF 4 = H 2 SiF a c and d. 3SiF 4 + 2H 2 = Si0 2 + 2H 2 SiF, §247. HYDROFLUORIC ACID—FLUOSILICIC ACID. 289 §246. Hydrofluoric acid. HF = 20.058 . H'F-', H — F . A colorless, intensely corrosive gas, soluble in water to a liquid that reddens litmus, rapidly corrodes glass, porcelain, and the metals, except platinum and gold (lead but slightly). Both the solution and its vapor act on the flesh as an insidious and virulent caustic, giving little warning, and causing obstinate ulcers. The anhydrous acid at 25° has a vapor density of 20, indicating that the molecule at this temperature is H 2 F 2 . But at 100° it is only 10, indicating that at that temperature the molecule is HF . The anhydrous liquid acid boils at 19.44° and does not solidify at —34.5°. The fluorides of the alkali metals are freely soluble in water, the solutions alkaline to litmus and slightly corrosive to glass; the fluorides of the alkaline earth metals are insoluble in water; of copper, lead, zinc and ferricum, spar- ingly soluble; of silver and mercury readily soluble. Fluorides are identified by the action of the acid upon glass. Calcium chloride solution forms, in solution of fluorides or of hydrofluoric acid, a gelatinous and transparent precipitate of calcium fluoride, CaF 2 , slightly soluble in cold hydrochloric or nitric acid and in ammonium chloride solution. Barium chloride precipitates, from free hydrofluoric acid less perfectly than from fluorides, the voluminous, white, barium fluoride, BaF 2 . Silver nitrate gives no precipitate. Sulphuric acid transposes fluorides, forming hydrofluoric acid, HF (§245, a). The gas is distinguished from other substances by etching hard glass — previously prepared by coating imperviously with (melted) wax, and writing through the •coat. The operation may be conducted in a small leaden tray, or cup formed •of sheet lead; the pulverized fluoride being mixed with sulphuric acid to the •consistence of paste. If the fluoride be mixed with silicic acid, we have, instead of hydrofluoric acid, silicon fluoride, SiF 4 (§245, 6); a gas which does not attack glass, but when passed into water produces fluosilicic acid, H 2 SiF a (§245, c and d). See below. Also, heated with acid sulphate of potassium, in the dry way, fluorides dis- engage hydrofluoric acid. If this operation be performed in a small test-tube, the surface of the glass above the material is corroded and roughened: CaF 2 + 2KHS0 4 = CaSO, + K 2 S0 4 + 2HF . By heating a mixture of borax, acid sulphate of potassium, and a fluoride, fused to a bead on the loop of platinum wire, in the clear flame of the Bunsen gas-lamp, an evanescent ■color is imparted to the flame. §247. Fluosilicic acid. H 2 SiF 6 = 144.716 . Fluosilicic acid* (hydrofluosilicic acid), (HF) 2 SiF 4 , or H 2 SiF„ , is soluble in water and forms metallic fluosilicates (silicofluorides) , mostly soluble in water; those of barium (§186, 6i), sodium and potassium, being only slightly soluble in water, and made quite insoluble by addition of alcohol. Potassium fluosilicate is precipitated translucent and gelatinous. Ammonium hydroxide precipitates silicic acid with formation of ammonium fluoride. With concentrated sulphuric acid, they disengage hydrofluoric acid, HF . By heat, they are resolved into fluorides and silicon fluoride: BaSiF„ = BaF 2 + SiF 4 . * Fluosilicic acid is directed to be prepared by taking one part each of fine sand and finely pow- dered fluor-spar, with six to eight parts of concentrated sulphuric acid, in a small stoneware bottle or a glass flask, provided with a wide delivery-tube, dipping into a little mercury in a small porcelain capsule, which is set in a large beaker containing six or eight parts of water. The stoneware bottle or flask is set in a small sand-bath, with the Band piled about it, as high as the material, and gently heated from a lamp. Each bubble of gas decomposes with deposition ■of gelatinous silicic acid. When the water is filled with this deposit, it may be separated by straining through cloth and again treating with the gas for greater concentration. The strained liquid is finally filtered and preserved for use. 290 SILICON— SILICON DIOXIDE. §248. §248. Silicon. Si = 28.4. Valence four (§15). There are three modifications of silicon: (a) Amorphous. — A dark brown powder; specific gravity, 2.0; non-volatile; infusible; burns in the air, forming- Si0 2 , and in chlorine, forming SiCl 4 . It is not attacked by acids except HP- Si + 6HF = H 2 SiF 6 + 2H 2 . It is dissolved by KOH with evolution of hydrogen. (6) Graphitoidal. — May be fused, but is not oxidized upon ignition in air or in oxygen. It is not attacked by HP , but is dissolved by a mixture of HP and HNO a , forming H 2 SiP„ . It is attacked slowly by fused KOH . (c) Adamantine silicon, crystalline silicon. — Grayish-black, lustrous, octahedral crystals, formed by fusing the graphitoidal form. Specific gravity, 2.49 at 10° (Woehler, A., 1856, 97, 261). It scratches glass but not topaz. It melts between the melting points of pig iron and steel, 1100° to 1300° In chemical properties it is very similar to the graphitoidal form, being attacked with even greater difficulty. Silicon is never found free in nature, but always in combination as silica, Si0 2 , or as silicates. Amorphous silicon is formed by passing vapor of SiCl 4 over heated potassium; by heating magnesium in SiP 4 vapor; by heating a mixture of Mg and Si0 2 ; by electrolysis of a fused silicate. It is readily prepared by heating a mixture of magnesium, one part, with sand, four parts, in a wide test-tube of hard glass (Gattermann, B., 1889, 22, 186). The graphitoidal form is crystalline and by many is said to be the same as the adamantine form. Method of preparation essentially the same (Warren, C. N., 1891, 63, 46). The crystalline form is made by fusing a silicate or K 2 SiF 6 with Al; by passing vapors of SiCl 4 over heated Ma or AJ in 9, carbon crucible (Deville, A. Ch., 1857, (3), 49, 62; Deville and Caron, A. Ch., 1863, (3), 67, 435; Woehler, I. a). §249. Silicon dioxide. ■ Si0 2 = 60.4 . (Silicic anhydride; silica.) Silicic acid. H 2 Si0 3 = 78.416 . SFO-*, and H'jffi^O-", , = Si = and H — — Si — — H. 1. Properties. — Silica, silicic anhydride, Si0 2 , is a white, stable, infusible solid; insoluble in water or acids; soluble in fixed alkalis with formation of silicates. Specific gravity of quartz, 2.647 to 2.652; of amorphous silica, 2.20 at 15.6°. Silicic acid, silicon hydroxide, H,SiO, t is a white, gelatinous solid, generally insoluble in water, and soluble in mineral acids. A dilute solution in water is. obtained by dialysis of the fixed alkali silicate with an excess of HC1 until the chlorides are all removed. It may be boiled for some time before the acid precipit'ates out. Upon standing silicic acid soon separates. 2. Occurrence. — Silicon is never found free in nature; it is always combined with oxygen in the form of silicon dioxide, Si0 2 , as quartz, opal, flint, sand,, etc.; or the silicon dioxide is in combination with bases as silicates; asbestos, soapstone, mica, cement, glass, etc. All geological formations except chalk contain silicon as the dioxide or as a silicate. 3. Formation. — Crystalline silica is formed by passing silicon fluoride into water, forming silicic acid and fluosilicic acid: 3SiF 4 + 3H 2 = H 2 SiO s + 2H 2 SiP 6 . The precipitate of silicic acid is dissolved in boiling NaOH and then heated in sealed tubes. Below 180° crystals of tridymite are formed, and above 180° crystals of quartz (Maschke, Pogg., 1872, 145, 549). 4. Preparation.— Pure amorphous silica is prepared by fusing finely pow- dered quartz with six parts of sodium carbonate, dissolving the cooled mass in water, and pouring into fairly concentrated hydrochloric acid. The precipitate is filtered, well washed and ignited. Or SiP 4 vapors are passed into water (§246) and the gelatinous precipitate washed, dried and ignited. Crystalline- §249, 7. SILICON DIOXIDE. 291 silica is prepared by fusing silicates with microcosmic salt or with borax (Eose, J. pr., 1867, 101, 228). Silicic acid. — The various hydroxides of silica act as weak acids. Metasilicio acid, H;SiO ; , , has been isolated; it is formed by decomposing silicon ethoxide, SUOCjHb^ , with moist air (Ebelmen, J. pr., 1846, 37, 359). Also by dialysis of a mixture of sodium silicate with an excess of hydrochloric acid until the chlorides are all removed, concentrating, allowing to gelatinize, and drying over sulphuric acid. Other hydroxides, acids, have been isolated, but there is some question as to their exact composition. 5. Solubilities. — Silica, Si0 2 , is insoluble in water or acids except HF , which dissolves it with formation of gaseous silicon fluoride, SiF 4 (§246). Of the silicates only those of the fixed alkalis are soluble in water, water glass. These silicates in solution are readily decomposed by acids, in- cluding carbonic acid, forming silicic acid, gelatinous. While anhydrous silicic anhydride, Si0 2 , is insoluble in mineral acids, the freshly precipi- tated hydroxide, silicic acid, is soluble in those acids. Silicic acid is decomposed by evaporation to dryness in presence of mineral acids, with separation of the anhydrous Si0 2 ; which is insoluble in more of the same acids, which previously had effected its solution. The most of the silicates found in nature are of complex composition. They are combinations of Si0 2 with bases. They are, as a rule, insoluble in water or acids. 6. Reactions. — Solutions of the alkali silicates precipitate solutions of all other metallic salts with formation of insoluble silicates; they are decomposed by acids with separation of silicic acid, a gelatinous precipi- tate, soluble in hydrochloric acid. Evaporation decomposes silicic acid with separation of insoluble silicic anhydride, Si0 2 . Ammonium salts precipitate gelatinous silicic acid from solutions of potassium Or sodium silicate. Therefore in the process of analysis the silicic acid, not removed in the first group by hydrochloric acid, will be precipitated in the third group on the addition of ammonium chloride. Silica, Si0 2 , is soluble in hot fixed alkalis forming silicates; it is not soluble in ammonium hydroxide, nor are solutions of alkali silicates pre- cipitated on addition of ammonium hydroxide as they are on the addition of ammonium salts. Boiling Si0 2 with the fixed alkali carbonates forms soluble silicates with greater or less readiness. Nearly all silicates are decomposed by heating in sealed tubes to 200° with concentrated HC1 or H 2 S0 4 . 7. Ignition. — Silicates fused with the alkalis form soluble alkali sili- cates, and oxides of the metal previously in combination. If alkali car- bonates are employed the same products are formed with evolution of C0 2 . Preferably a mixture (in molecular proportions) of potassium and sodium carbonates, four parts, should be used to one part of the insoluble silicate. Silica, Si0 2 , is also changed to a soluble silicate by fusing with fixed alkali hydroxides or carbonates. 392 PHOSPHORUS. §249, 8. Si0 2 does not react with K 2 S0 4 or Na 2 S0 4 , even when fused at a very high temperature (Mills and Meanwell, J. C, 1881, 39, 533). In the fused bead of microcosmic salt particles of silica swim undissolved. If a silicate be taken, its base will, in most cases, be dissolved out, leaving a " skeleton of silica " un- dissolved in the liquid bead. But with a bead of sodium carbonate, silica (and most silicates) fuse to a clear glass of silicate. Silica is separated from the fixed alkalis in natural silicates, by mixing the latter in fine powder with three parts of precipitated calcium carbonate, and one-half part of ammonium chloride, and heating in a platinum crucible to redness for half an hour, avoiding too high a heat. On digesting in hot water, "the solution contains all the alkali metals, as chlorides, with calcium chloride -and hydroxide. 8. Detection. — Silicates are detected by conversion into the anhydride, Si0 2 . The silicate is fused with about four parts of a mixture of potas- sium and sodium carbonates, digested with warm water, filtered, and evaporated to dryness with an excess of hydrochloric acid. The dry resi- due is moistened with concentrated HC1 and thoroughly pulverized; water is added and the precipitate of Si0 2 is thoroughly washed. Further con- firmation may be obtained by warming the precipitate of Si0 2 with calcium fluoride and sulphuric acid (in lead or platinum dishes), forming the gaseous silicon fluoride, SiF 4 . This is passed into water where it is decomposed into gelatinous silicic acid and fluosilicic acid: 3SiF 4 -\- 3H 2 = H 2 Si0 3 + 2H 2 SiF 6 (§246). Silica, Si0 2 , is usually treated as directed for silicates, but may be at once warmed with calcium fluoride and sul- phuric acid. 9. Estimation. — The compound containing a silicate or silica is fused with fixed alkali carbonates as directed under detection, and the amount of well- ■washed Si0 2 determined by weighing after ignition. §250. Phosphorus. P = 31.0. Usual valence three or five (§11). 1. Properties. — Phosphorus is prepared in several allotropic modifications. Specific gravity of the yellow, solid, at 20°, 1.82321; liquid, at 40°, 1.74924; solid, at 44°, 1.80681 (Pisati and de Franehis, B., 1875, 8, 70). At ordinary tempera- tures it is brittle and easily pulverized. At about 45° it melts, but may be cooled in some instances (under an alkaline liquid) as low as +4° without solidifying. When it solidifies from these lower temperatures, as it does by stirring with a solid substance, the temperature immediately rises to about 45°. Boiling point, 287.3° at 762 mm. pressure (Schroetter, A., 1848, 68, 247; Kopp, A., 1855, 93, 129). The density of the vapor at 1040° is 4.50 (Deville and Troost, C. r., 1863, 56, 891). The computed density for. the molecule P 4 is 4.294. At a white heat the density, 3.632, indicates dissociation of the molecule to P 2 (Meyer and Biltz, B., 1889, 22, 725). Specific gravity of the red amorphous modification at 10°, 1.964. Ordinary crystalline yellow stick phosphorus is a nearly colorless, trans- parent solid; when cooled slowly it is nearly as clear as water. In water con- taining air it becomes coated with a thin whitish film. If melted in fairly large quantities and cooled slowly it forms dodecahedral and octahedral crys- tals (Whewell, C. N., 1879, 39, 144). Heated in absence of air above the boiling point it sublimes as a colorless gas, depositing lustrous transparent crystals (Blondlot, G. r., 1866, 63, 397). At low temperatures phosphorus oxidizes slowly •in the air with a characteristic odor, probably due to the formation of ozone §250, 4. PHOSPHORUS. 293 -and phosphorous oxide, P 2 3 (Thorpe and Tutton, J. C, 1890, 57, 573). It ignites spontaneously in the air at 60°, burning with a bright yellowish white light ^producing much heat. From the finely divided state, as from the evaporation •of its solution in carbon disulphide, it ignites spontaneously at temperatures -at which the compact phosphorus may be kept for days. It must be preserved under water. Great precaution should be taken in working with the ordinary or yellow phosphorus. Burns caused by it are very painful and heal with great difficulty. Ordinary phosphorus is luminous in the dark, but it has been shown 'that the presence of at least small amounts of oxygen are neces- sary. The presence of H 2 S , S0 2 , CS 2 , Br , CI , etc., prevent the glowing (Schroetter, J. pr., 1853, 58, 158; Thorpe, Nature, 1890, 41, 523). Upon heating in absence of air, better in sealed tubes, to 300° it is changed to the red modi- fication (Meyer, B., 1882, 15, 297). Red phosphorus is a dull carmine-red tasteless powder. It is not poisonous, while the ordinary yellow variety is intensely poisonous, 200 to 500 milligrams being sufficient to cause death. While the yellow modification is so readily and dangerously combustible when exposed to the air even at ordinary tem- peratures, the red variety needs no special precautions for its preservation. It does not melt when heated to redness in sealed tubes, but is partially -changed to the yellow crystalline form (Hittorf, Pogg., 1865, 126, 193). If amorphous phosphorus be distilled in the absence of air, it is changed to the crystalline form, action beginning at 260°. Heated in the air from 250° to 260° it takes fire (Schroetter, I. a). A black crystalline metallic variety of phos- phorus is described by Hittorf (I. a); also Bemsen and Kaiser (Am., 1882, 4, 459) •describe a light plastic modification. Phosphorus is largely used in match- making. Yellow phosphorus is used in the ordinary match, and the red (amorphous) in the safety matches, the phosphorus being on a separate surface. 2. Occurrence. — It is never found free in nature. It is found in the primitive rocks as calcium phosphate, occasionally as aluminum, iron, or lead phosphate, •etc. Plants extract it from the soil, and animals from the plants. Hence traces of it are found in nearly all animal and vegetable tissues; more abundantly in the seeds of plants and in the bones of animals. 3. Formation. — Ordinary phosphorus is formed by heating calcium or lead phosphates with charcoal. The yield is increased by mixing the charcoal with sand or by passing HC1 gas over the heated mixture. By igniting an alkali -or alkaline earth phosphate with aluminum (Bossel and Frank, B., 1894, 27, 52). Bed phosphorus is formed by the action of light, heat or electricity on ordinary phosphorus (Meyer, B., 1882, 15, 297). By heating ordinary phosphorus with a small amount of iodine (Brodie, J. pr., 1853, 58, 171). 4. Preparation. — Ordinary phosphorus is prepared from bones. They are first burned, which leaves a residue, consisting chiefly of Ca 3 (P0 4 ) 2 ; then H 2 S0 4 is added, producing soluble calcium tetrahydrogen diphosphate (a). After filtering from the insoluble calcium sulphate the solution is evaporated and ignited, leaving calcium metaphosphate (6). Then fused with charcoal, reducing two-thirds of the phosphorus to the free state (c). The mixture of sand, Si0 2 , with the charcoal is preferred, in which case the whole of the phosphorus is reduced (d). Hydrochloric acid passed over red-hot calcium phosphate and charcoal reduces the whole of the phosphorus. This process works well in the laboratory, and has also been successfully employed on a larger scale. Either of the calcium phosphates may be used (e) and (f). (a) Ca„(P0 4 ) 2 + 2H 2 S0 4 = 2CaS0 4 + CaH 4 (P0 4 ) 2 (6) CaH 4 (P0 4 ) 2 + ignition = Ca(P0 3 ) 2 + 2H 2 (c) 3Ca(P0 3 ) 2 + 10C = Ca 3 (P0 4 ) 2 + 10CO + P 4 (d) 2Ca(P0 8 ) 2 + 10C + 2Si0 2 = 2CaSi0 3 + P 4 + 10CO (e) 2Ca 8 (P0 4 ) 2 + 16C + 12HC1 = 6CaCl 2 + P 4 + 16CO + 6H 2 (f) 2Ca(P0 3 ) 2 + 12C + 4HC1 = 2CaCl 3 + P 4 + 12CO + 2H 2 Bed or amorphous phosphorus is prepared by heating ordinary phosphorus ior a long time (40 hours) at 240° to 250° in absence of air. At 260° the reverse -change takes place. If the heating is under pressure and at 300°, the change -to the red phosphorus is almost immediate. It is washed with CS 2 to remove ■all traces of yellow phosphorus and is dried at 100°- 294 PHOSPHORUS. §250, 5.. 5. Solubilities. — A trace of phosphorus dissolves in water. Alcohol dissolves 0.4, ether 0.9, olive oil 1.0, and turpentine 2.5 per cent of it,, while carbon disulphide dissolves 10 to 15 times its own weight. Bed phosphorus is insoluble in water, ether, or carbon disulphide. 6. Reactions. — When phosphorus is boiled with a fixed alkali or alkaline earth hydroxide, phosphorus hydride, phosphine (§249), PH 3 , and a hypophosphite (§250) are formed. Phosphorus, when warmed in an atmosphere of N or C0 2 , combines directly with many metals to form phosphides. These phosphides are usually brittle solids decomposing with water or dilute acids with formation of phosphoretted hydrogen, PH 3 . In nearly all the reactions of phosphorus both varieties react the- same, the red variety with much less intensity, and frequently requiring the aid of heat. It is ignited when brought in contact with Pb0 2 , Pb 3 4 _ HgO , Ag 2 , Cr0 3 , K 2 Cr 2 7 and when heated with CuO or Mn0 2 . Solu- tions of platinum, gold, silver, and copper salts are decomposed by phos- phorus with separation of the corresponding metal (Boettger, /. C, 1874,. 27,1060). With HN0 3 , H 3 P0 4 and NO are formed; when heated with KN0 3 a rapid oxidation takes place. It combines with oxygen, forming P 2 3 or P 2 5 . With yellow phos- phorus the reaction begins at ordinary temperature; with the red variety not till heated to 250° to 260° (Baker, J. C, 1885, 47, 349). Water is decomposed at 250°, forming PH 3 and H 3 P0 4 (Schroetter, I. c). Combination with red phosphorus and sulphur takes place at ordinary temperatures, forming P 2 S 3 or P 2 S 5 , depending upon the proportion of each employed (Kekule, A., 1854, 90, 310). With ordinary phosphorus, the action is explosive. CI or Br react with incandescence at ordinary temperatures, forming trihalogen or pentahalogen compounds, depending upon the amount of halogen employed. With iodine, PI 3 is formed. The halogen compounds of phosphorus are decomposed by water with formation of the corresponding hydracids and phosphorous or phosphoric acids, depending upon the degree of oxidation of the phosphorus. In the presence of water phosphorus is oxidized to H 3 P0 4 by CI, Br, I, HC10 3 , HBr0 3 , or HI0 3 with formation of the corresponding hydracid : P 4 + 10C1 2 + 16H 2 = 4H 3 P0 4 + 20HC1 . 7. Ignition. — When sodium carbonate is heated to redness with phosphorus, the carbonic anhydride is reduced and carbon is set free. Phosphorus heated with magnesium in a vapor of carbon dioxide forms P 2 Mg s , which can be heated to redness in absence of air without decomposition. Heated in the air it becomes oxidized (Blunt, A. Gh., 1865, (4), 5, 487). Phosphorus also combines with Cu , Ag , Cd , Zn and Sn when it is heated with these elements in sealed tubes. It does not combine with Al and but slightly with Fe (Emmerling,, J. C, 1879, 36, 508). §252, 3. PHOSPHINE— HYPOPBOSPEOROUS ACID. 295 8. Detection. — By its phosphorescence; by formation of PH 3 when boiled with KOH (Hofmann, B., 1871, 4, 200); by oxidation to H 3 P0 4 and detection as such (§75, 6d). 9. Estimation. — Oxidation to H 3 P0 4 , precipitation with magnesia mixture as MgNH 4 P0 4 , ignition to, and weighing as Mg 2 P 2 O t (§189, 9). §251. Phosphine. PH 3 = 34.024. P-"'H' 3 ,H — P — H. I H Phosphine, PH 3 , is a colorless gas having a very disagreeable odor. As usually prepared, it is spontaneously inflammable, burning in the air with formation of metaphosphoric acid: 2PH 8 + 40 2 = 2HPO„ + 2H 2 . It is liquified and frozen at very low temperatures; boiling point, about — 85°; ■melting point, — 132.5° (Olszewski, M., 1886, 7, 371). It is very poisonous, spar- ingly soluble in water, which solution has the peculiar odor of the gas and has an exceedingly bitter taste. It is formed by boiling phosphorus with a fixed alkali or alkaline earth hydroxide (a); by ignition of H 3 PO, or H s P0 8 (6); by ignition of hypophosphites (c) ; ( by the decomposition of the alkaline earth phosphides with water or dilute acids (d) : (a) P 4 _ + 3KOH + 3H 2 = 3KH 2 F0 2 + PH S (6) 2H 3 P0 2 = HPO„ + PH„ + H 2 4H s PO» = 3HPO„ + PH„ + 3H 2 (c) 4NaH 2 P0 2 = Na 4 P 2 7 + 2PH 3 + H 2 (d) Ca 3 P 2 + 6H 2 = 3Ca(OH) 2 + 2PH 3 Ca B P 2 + 6HC1 = 3CaCl 2 + 2PH„ It is a strong reducing agent; transposes many metallic solutions: 3CuSO, + 2PH S = Cu 8 P 2 + 3H 2 S0 4 ; reduces solutions of silver and gold to the metallic state: 8AglT0 8 + PH, + 4H 2 = H s P0 4 + 8HNO a + 8Ag; is oxidized to H 3 P0 4 by hot H 2 S0 4 , CI , HCIO , HN0 2 , HN0 3 , H, AsO, , etc. A liquid phosphorus hydride, P 2 H 4 , and a solid, P 4 H 2 , are known (Besson, C. r., 1890, 111, 972; Gattermann and Hausknecht, B., 1890, 23, 1174). §252. Hypophosphorous acid. H 3 P0 2 = 66.024 . H I H' 3 P'0-" 2 . H — — P = 0. I H 1. Properties. — Hypophosphorous acid was discovered in 1816 by Dulong (A. Ch., 1816, 2, 141). It is a colorless syrupy liquid; specific gravity, 1.493 at 18.8°. At 17.4° it becomes a white crystalline solid (Thomsen, B., 1874, 7, 994). Although containing three hydrogen atoms it forms but one series of salts, e. g., NaH 2 P0 2 , Ba(H 2 P0 2 ) 2 , etc. 2. Occurrence. — Not found in nature. 3. Formation. — All ordinary metals form hypophosphites except tin, copper and mercurosum. Silver and ferric hypophosphites are very unstable. (1) A 296 HYPOPHOSPHOROUS ACID. §252,4. few metals, such as zinc and iron, dissolve in H a P0 2 , giving off hydrogen and forming a hypophosphite. (2) The alkali and alkaline earth salts may be formed by boiling phosphorus with the hydroxides (Mawrow and Muthmann, Z. angew., 1896, ii, 268). (3) As all hypophosphites are soluble, none can be formed by precipitation. All may be formed from their sulphates by trans-, position with barium hypophosphite. (4) All may be made by adding H 3 PO, to the carbonates or hydroxides of the metals. 4. Preparation. — To prepare pure H 3 P0 2 , BaO and P (in small pieces) are warmed in an open dish with water until FE, ceases to be evolved. The liquid is filtered and excess of BaO is removed by passing in C0 2 . After again filtering, the liquid is evaporated to crystallization of the barium salt. This is dissolved in water and decomposed by the calculated quantity of H 2 S0 4 . The solution is filtered and evaporated in an open dish, care being taken not to heat above 110° . Upon cooling the white crystalline tablets are obtained. 5. Solubilities. — The free acid is readily miscible in water in all proportions. The salts are all soluble in water, a number of them are soluble in alcohol. 6. Reactions. — A. — "With metals and their compounds. Hypophosphorous acid is a very powerful reducing agent, being oxidized to phosphoric acid or a phosphate. 1. Pbiv becomes Pb" in acid or alkaline mixture. 2. Ag' becomes Ag° in acid or alkaline mixture. 3. Hg" becomes Hg' and then Hg° in acid or alkaline mixture. 4- AsV and As'" become As° in presence of HC1 . 5. Bi'" becomes Bi° in presence of alkalis or acetic acid. 6. Cu" becomes Cu 2 H 2 and on boiling Cu° (separation from Cd). 7. Pe'" becomes Pe" , no action in alkaline mixture. 8. CrVi becomes Cr'" , no action in alkaline mixture. 9. Co'" becomes Co" , no action in alkaline mixture. 10. Hi'" becomes Ni" , no action in alkaline mixture. 11. Mn"+n becomes Mn" in acid solution. 12. Mniv+n becomes Mniv in alkaline mixture. B. — With non-metals and their compounds. 1. H 3 Pe(CN) 6 becomes H 4 Pe(CN) . 2. HN0 3 and HNO, become NO . S. H 3 P0 2 on heating becomes H 3 P0 4 and PH 3 . 4. H.,S0 3 becomes free sulphur with formation of some H 2 S (Ponndorf, J. C, 1877,31,275). H 2 S0 4 becomes first H 2 S0 3 then S . See above. 5. CI becomes HC1 in acid mixture, a chloride with alkalis. HCIO and HC10 8 form same products as CI . 6". Br becomes HBr in acid mixture, a bromide with alkalis. HBr0 3 forms HBr . 7. I forms HI , in alkaline mixtures an iodide. HI, dry, reacts violently, forming H 3 P0 3 and PH 4 I (Ponndorf, I.e.). HI0 3 forms HI . 7. Ignition.— On ignition hypophosphites leave pyrophosphates, evolving PE S . The acid decomposes on heating to PH 3 and H„P0 4 (or HP0 8 if at a red heat). 8. Detection. — Hypophosphorous acid may be' known from phosphorous acid by adding cupric sulphate to the free acid and heating the solution to 55°. "With hypophosphorous acid a reddish-black precipitate of copper hydride (Cu 2 H 2 ) is thrown down, which, when heated in the liquid to 100°, is decomposed with the deposition of the metal and the evolution of hydrogen, whilst with phosphorous acid the metal is precipitated and hydrogen evolved, but no Cu 2 H 2 is formed. Further, hypophosphorous acid reduces the permanganates immediately, but phosphorous acid only after some time. Phosphites precipitate barium, strontium, and calcium §253,6. PHOSPHOROUS ACID. 297 salts, while hypophosphites do not. When hypophosphorous acid is treated with zinc and sulphuric acid it is converted into phosphoretted hydrogen. On boiling hypophosphorous acid with excess of alkali hydrox- ide, first a phosphite then a phosphate is formed, with evolution of hydrogen. 9. Estimation. — (1) By oxidation with nitric acid and then proceeding- as with phosphoric acid. (2) By mercuric chloride acidulated with HC1; the temperature must not rise above 60°, otherwise metallic mercury will be formed. The precipitated HgCl, after washing and drying at 100°, is weighed. NaH 2 P0 2 + 4HgCl 2 + 2H 2 = 4HgCl + H s P0 4 + NaCl + 3HC1 §253. Phosphorous acid. H 3 P0 3 = 82.024 , H I H' 3 F"0-" 3 ,H — — P — — H. 1. Properties. — Phosphorous anhydride, P 2 8 , is a snow-white solid, melting at 22.5°, and boiling- at 173.1° (Thorpe and Tutton, J. C, 1890, 57, 545). The vapor density of the gaseous oxide indicates the molecule to be P 4 8 . Specific gravity of the liquid at 21°, 1.9431; of the solid at the same temperature, 2.135. It has an odor resembling that of phosphorus. Heated in a sealed tube at 200° it decomposes into P 2 4 and P (T. and T., J. C, 1891, 59, 1019). It reacts slowly with cold water, forming H 3 PO„; with hot water the reaction is violent and PH 3 is evolved. Upon exposure to the air it oxidizes to P 2 5 . The acid, H 8 PO a , is a crystalline solid, very deliquescent, melting at 74° (Hurtzig and Geuther, A., 1859, 111, 171). It is a dibasic acid, forming no tribasic salts (Amat, C. r., 1889, 108, 403). One or two of the hydrogen atoms are replaceable by metals forming acid or normal salts. The third hydrogen is never replaced by a metal, but may be replaced by organic radicles (Eailton, J. C, 1855, 7, 216; Michaelis, J. C., 1875, 28, 1160). Neither meta nor pyro- phosphorous acids are known, but a number of pyrophosphites have been pre- pared (Amat, G. r., 1888, 106, 1400; 1889, 108, 1056; 1890, 110, 1191 and 901; A. Ch., 1891, (6), 24, 289). 2. Occurrence. — Does not occur in nature. 3. Formation. — P 2 0\, is formed together with P 2 5 when phosphorus is ignited in the air. H s PO„ is formed together with H s P0 4 when phosphorus is oxidized with HN0 3 ; by the oxidation of H 3 P0 2 ; by the action of P upon a concentrated solution of CuS0 4 in absence of air: 3CuSO, -4- P 4 + 6H 2 = Cu„P 2 + 2H 3 PO„ + 3H 2 S0 4 (Schiff, A., 1860, 114, 200). 4. Preparation To prepare phosphorous anhydride, P 2 3 , phosphorus is burned in a tube with an insufficient supply of air (Thorpe and Tutton, I. c.). The acid, H 3 P0 S , is prepared by dissolving the anhydride in cold water; by decomposing PC1 3 with water (Hurtzig and Geuther, I.e.). 5. Solubilities. — The acid is miscible in water in all proportions. Alkali phosphites are soluble in water, most others are insoluble (distinction from hypophosphites) . 6. Keactions.— Phosphorous acid is a strong reducing agent, oxidizing to phosphoric acid when exposed to the air. It reduces salts of silver and gold to the metallic state and is changed to phosphoric acid by most of the strong oxidizing acids and by many of the higher metallic oxides. HgCl 2 becomes HgCl and then Hg° , CuCl 2 becomes CuCl then Cu° (Kammelsberg, /. C, 1873, 298 HYPOPHOSPHORIC ACID— PHOSPHORIC ACID. §253, 7. 26, 13). Concentrated H 2 S0 4 with heat forms H a P0 4 and S0 3 (Adie, J. C, 1891, 59, 230). H 2 SO a forms H 2 S and H a P0 4 (Woehler, A., 1841, 39, 252). Nascent hydrogen (Zn and H 2 S0 4 ) produce PH B (Dusart, C. r., 1856, 43, 1126). 7. Ignition. — The acid is decomposed by ignition, forming HP0 3 and P or PH a (Vigier, Bl., 1869, (2), 11, 125; Hurtzig and Genther, I. &). Phosphites are decomposed by heat, leaving a pyrophosphate and a phosphide and evolving PH a or H (Kammelsberg, B., 1876, 9, 1577; and Kraut, A., 1875, 177, 274). 8. Detection. — By oxidation to H s P0 4 and detection as such. It is distin- guished from hypophosphorous acid by reducing CuS0 4 to Cu°, while the latter forms Cu 2 H 2 ; also by the solubilities of the salts (§252, 8). Its reactions with oxidizing agents distinguish it with hypophosphorous acid from phos- phoric acid. 9. Estimation. — By oxidation to H a P0 4 and estimation as such. §254. Hypophosphoric acid. H 4 P 2 6 = 162.032 . II II H' 4 F v ,0- , ',,H — — P — P — — H. I I I I H H Hypophosphoric acid is formed together with phosphorous and phosphoric acids by slowly oxidizing phosphorus in moist air (Salzer, A., 1885, 232, 114 and 271); also by oxidizing phosphorus with dilute ENO, in presence of silver nitrate (Philipp, B., 1885, 18, 749). It consists of small colorless hygroscopic crystals which melt at 55°. It decomposes when heated to 70° into H a PO s and BP0 3 , and at 120° gives H 4 P 2 T and PH, (Joly, C. r., 1886, 102, 110 and 760). It is oxidized to H s P0 4 by warm HN0 3 , slowly by EMnO, in the cold, rapidly when heated. It is not oxidized by H 2 2 , chlorine water or H 2 CrO„; HgCl, becomes HgCl (Amat, C. r., 1890, 111, 676). It is not reduced by Zn and H 2 S0 4 (distinction from H 8 P0 2 and H 3 P0„). With a solution of silver nitrate it gives a white precipitate which does not blacken in the light (distinction from H a P0 2 and H3PO3). It forms four series of salts, all four hydrogen atoms being replaceable by a metal. The hypophosphates are much more stable towards, oxidizing agents than hypophosphites or phosphites. §255. Phosphoric acid. H 3 P0 4 = 98.024 . II H' 3 P v 0- \ ,H — — P — — H. I I H 1. Properties. — Phosphoric anhydride, P 2 O *, is a white, flakey, very delique- scent solid, fusible, subliming undecomposed at » red heat. It is very soluble in water, forming three varieties of phosphoric acid: ortho, H 3 P0 4 ; meta, HPO a ; * According to Tilden and Barnett (J. C, 1896, 69, 154) the molecule is P 4 O 10 not P a O a ; P t O, not P a O a (Thorpe and Tutton, J. C, 1891, 59, 1022) ; and P 4 S, not P 2 S S (Isambert, C.r., 1886, 108, §255, 3. PHOSPHORIC ACID. 299 .and pyro, H 4 P a 7 . Orthophosphoric acid is a translucent, feebly crystallizable and very deliquescent soft solid. Specific gravity, 1.88 (Schiff, A., 1860, 113, 183); ^melting point, 41.75° (Berthelot, Bh, 1878, (2), 29, 3). It is changed by heat, first to pyrophosphoric acid, then to metaphosphoric acid. Orthophosphoric acid forms three classes of salts: M'H 2 P0 4 , primary, monobasic or mono- metallic phosphates; M' 2 HP0 4 , secondary, dibasic or dimetallic phosphates; and M' 3 F0 4 , tertiary, tribasic, trimetallic or normal phosphates. The first two are acid salts, but Na^HPO, is alkaline to test paper. Metaphosphoric acid, HPO„ ,H — O — P = 0,isa white waxy solid, volatile at a red heat O (ordinary glacial phosphoric acid owes its hardness to the universal presence of sodium metaphosphate). It is a monobasic acid, but there are various poly- meric modifications, distinguished from each other chiefly by physical differ- ences of the acids and their salts (Tammann, Z. phys. Ch., 1890, 6, 122). O II II Pyrophosphoric acid, H 4 P 2 7 , H — O — P — O — P — O — H , is a glass-like O I I H H solid (Peligot, A. Ch., 1840, (2), 73, 286), very soluble in, but unchanged by, water at ordinary temperature; changed by boiling water to H s P0 4 . Heated ■to redness HP0 3 is formed. It forms two classes of salts: M' 2 H 2 P 2 7 and HI' 4 P 2 7 . 2. Occurrence. — Phosphates of Al , Ca , Mg and Pb are widely distributed in minerals. Guano consists quite largely of calcium phosphate. Calcium and magnesium phosphates are found in the bones of animals and in the ashes of plants. The free acids are not found in nature. 3. Formation. — Phosphoric anhydride, P 2 5 , is formed by burning phosphorus in great excess of air; also by burning phosphorus in NO , NO. , or C10 2 . Orthophosphoric acid, H 3 P0 4 , is formed by long exposure of phosphorus to moist air, or by oxidation with HNO, ; by oxidation of H a P0 2 or H„P0 3 with the halogens, HNO, , HC10 a , etc.; by treating P,0 5 , HP0 3 , or H 4 P 2 7 with boiling water; by combustion of PH S in moist air; and by action of water on PC1 5 . It is also formed from metallic phosphates by transposition with acids in cases where a precipitate results, as a lead or barium phosphate with sul- phuric acid, or silver phosphate with hydrochloric acid. But when the pro- ducts are all soluble, as calcium phosphate with acetic acid or sodium phosphate with sulphuric acid, the transposition is only partial; so that unmixed phos- phoric acid is not obtained. A non-volatile acid, like phosphoric, is not sepa- rated from liquid mixtures, as the volatile acids are, like hydrochloric. The change represented by equation (a) can be verified, that is, pure phosphoric acid can be separated; but the changes shown in equations (6) and (c) do not comprise the whole of the material taken. In the operation (6) some sodium phosphate and some nitric acid will be left, and in (c) some trihydrogen phosphate will no doubt be made. a. CaH 4 (P0 4 ) 2 + H 2 C 2 4 = CaC 2 4 + 2H 3 P0 4 B. NajHPO, + 2HN0 3 = 2NaN0 3 + H 3 P0 4 and Na 2 HP0 4 + HUO, = NaN0 3 + NaH 2 P0 4 c. 2CaHP0 4 + 2HC1 = CaCl 2 + CaH 4 (P0 4 ) 2 Metaphosphoric acid is formed by treating P 2 5 with cold water; by decom- position of lead metaphosphate with H 2 S or of the barium salt with H 2 S0 4 ; by ignition to dull redness of phosphorus or any of its acids in the presence of air and moisture. Pyrophosphoric acid, H 4 P 2 7 , is formed by the decomposition of lead pyro- phosphate, Pb 2 P 2 7 , with H 2 S or of the corresponding barium salt with H 2 S0 4 ; or by heating H 3 P0 4 to a little above 200° until no yellow silver phosphate, Ag,P0 4 , is obtained on dissolving in water and treatment with silver nitrate after neutralization with NH 4 OH . 300 PHOSPHORIC ACID. §255,4. 4. Preparation. — To prepare P 2 5 , phosphorus is burned in a slow cur- rent of dry oxygen heating to about 300°, then in a more rapid current of the gas, and finally the P 2 5 is distilled in an atmosphere of oxygen (Shenstone, Watts' Die, 1894, IV, 141). H 3 P0 4 is prepared by warming- phosphorus, one part, with nitric acid, sp. gr. 1.20, ten to twelve parts,, with addition of 300 to 600 milligrams of iodine to 100 grams of phos- phorus, until the phosphorus is completely dissolved. The excess of HN0 3 is removed by evaporation, water is added and the solution is sat- urated with H 2 S to remove any arsenic that may be present. The solution is then evaporated to a syrupy consistency at temperatures not above 150° (Krauthausen, Arch. Pharm., 1877, 210, 410; Huskisson, B., 1884,. 17, 161). Many orthophosphates are formed by the action of H 3 P0 4 upon metallic oxides or carbonates ; by the reaction between an alkali phosphate and a soluble salt of the heavy metal; by fusion of a metaphosphate with the corresponding metallic oxide or hydroxide; also by long continued boiling of meta or pyrophosphates. Metaphosphates are formed by double- decomposition with NaP0 3 or by fusion of a monobasic phosphate or any phosphate having but one hydrogen equivalent substituted for a metal,, the oxide of which is non-volatile, e. g., NaNH 4 HP0 4 . Pyrophosphates are formed by double decomposition with Na 4 P 2 7 ; by action of H 4 P 2 7 on certain oxides or hydroxides; also by ignition of dibasic orthophos- phates, e. g., Na 2 HP0 4 . Na 2 H 2 P 2 7 may be prepared by titrating a sat- urated solution of Na 4 P 2 7 with HN0 3 until the solution gives a red color with methyl orange. Upon standing the salt separates in large crystals (Knorre, Z. angew., 1892, 639). 5. Solubilities. — All the phosphoric acids are readily soluble in water,, as are all alkali phosphates. Alkali primary orthophosphates have an acid reaction in their solutions; alkali secondary and tertiary phosphates are alkaline in their solutions; the latter is easily decomposed, even by C0 2 , forming the secondary salt. A number of non-alkali primary ortho- phosphates are soluble in water, e. g., CaH 4 (P0 4 ) 2 . All normal and di- metallic orthophosphates are insoluble except those of the alkalis. The normal and dimetallic phosphates of the alkalis precipitate solutions of all other salts. The precipitate is a normal, dimetallic, or basic phos- phate, except that with the chlorides of mercury and antimony it is not a phosphate but an oxide or an oxyehloride. All phosphates are dissolved or transposed by HN0 3 , HC1 , or H 2 S0 4 , and all are dissolved by HC 2 H 3 2 except those of Pb , Al and Fe'" . All are soluble in H 3 P0 4 except those of lead, tin, mercury, and bismuth. The non-alkali meta and pyrophosphates are generally insoluble in water. The pyrophosphates of the alkaline earth metals are difficultly solu- ble in acetic acid. The most of the pyrophosphates of the heavy metals, §255, 6A. PHOSPHORIC ACID. 301 except silver, are soluble in solutions of alkali pyrophosphates, as double pyrophosphates soluble in water (distinction from orthophosphates). Ferric iron as a double pyrophosphate loses the characteristic properties of that metal (Persoz, J. C, 1849, 1, 183). Phosphates are insoluble in alcohol. 6. Reactions. — A. — With metals and their compounds. — Phosphoric acid dis- solves some metals, e. g., Fe , Zn and Mg with evolution of hydrogen. It unites with the oxides and hydroxides of the alkalis and alkaline earths and with other freshly precipitated oxides and hydroxides except perhaps antimonous oxide. It also decomposes all carbonates evolving' C0 2 . Phosphates are formed in the above reactions, the composition of which depends upon the conditions of the experiment. Free orthophosphoric acid is not precipitated by ordinary salts of third, fourth and fifth group metals (in instance of ferric chloride, a distinction from pyrophosphoric acid and metaphosphoric acid),* but is precipitated in part by silver nitrate, and lead nitrate and acetate. Ammoniacal solution of calcium chloride or of barium chloride precipitates the normal phosphate. Free metaphosphoric acid precipitates solutions of silver nitrate, lead nitrate, and lead acetate, the precipitates being insoluble in excess of metaphosphoric acid, and soluble in moderately dilute nitric acid. Barium, calcium and ferrous chlorides, and magnesium, aluminum, and ferrous sulphates, are not precipi- tated by free metaphosphoric acid. Ferric chloride is precipitated, a distinc- tion from orthophosphoric acid. Free pyrophosphoric acid gives precipitates with solutions of silver nitrate, lead nitrate or acetate, ar.d ferric chloride; no precipitates with barium or calcium chloride, or with magnesium or ferrous sulphate. Orthophosphoric acid — or an orthophosphate with acetic acid — does not coagu- late egg albumen or gelatine. This is a distinction of both orthophosphoric acid and pyrophosphoric acid from metaphosphoric acid*. With silver nitrate soluble orthophosphates form silver orthophosphate, Ag 3 P0 4 , yellow; with metaphosphates, silver metaphosphate, AgP0 3 , white; and with pyrophosphates, silver pyrophosphate, Ag 4 P 2 7 , white, all soluble in ammonium hydroxide. Silver metaphosphate is soluble in excess of an alkali metaphosphate (distinction from pyrophosphates). If a disodium or dipotassium orthophosphate is added to solution of silver nitrate, free acid is formed, and an acid reaction to test-paper is induced (a). But with a trisodium or tripotassium phosphate, the solution remains neutral (8) — a means of distinguishing the acid, phosphates from the normal. (a) Na^HPO, + 3AgN0 3 = Ag s P0 4 + 2NaN0 3 + HUO, (6) Na 3 P0 4 + 3AgNO„ = Ag 3 PO, + SNaNO, Free orthophosphoric acid forms no precipitate with reagent silver nitrate. With lead acetate or nitrate, Na 2 HP0 4 forms Pb 3 P0 4 , white, insoluble in acetic acid, as are also the phosphates of aluminum and f erricum. With * A solution containing 5 p. c. ferric chloride, mixed with one-fourth its volume of a 10 p. o. solution of orthophosphoric acid, requires that near half of the latter he neutralized (so that phosphate is to phosphoric acid as 1.114 is to 1.000) before precipitation occurs. On the other hand, 4 cc. of a 5 p. c. solution of ferric chloride, mixed with 1 cc. of a 6 p. u. solution of meta- phosphoric acid, form a precipitate, to dissolve which, 20 cc. of the same metaphosphoric acid solution or 6 cc. of a 24 p. c. solution of hydrochloric acid are required. Four cc. of a 5 p. c. solution of silver nitrate with 1 cc. of a 10 p. c. solution of orthophosphoric acid give a precipi- tate, to dissolve which requires 7 cc. of the same orthophosphoric acid solution. [The Author's report of work by Mr. Morgan, Am. Jour. Phar., 1876, 48, 534. Kratschmer and Sztankovansky, Z., 1883, 21, 520.] 302 PHOSPHORIC ACID. §255,64. PbCl 2 the precipitate always contains a chloride. Free phosphoric acid, H 3 P0 4 , forms an acid phosphate, PbHP0 4 (Heintz, Pogg., 1848, 73, 119). Lead salts also form white precipitates with soluble pyro and metaphos- phates ; the pyro salt, Pb 2 P 2 7 , is soluble in an excess of Na 4 P 2 7 . Bis- muth salts form BiP0 4 , insoluble in dilute HN0 3 . Solutions of orthophosphates give, with soluble ferric, chromic, and aluminum salts, mostly the normal phosphates, FeP0 4 , etc. The ferric phosphate is but slightly soluble in acetic acid, and for this reason it is made the means of separating phosphoric acid from metals of the earths and alkaline earths (§152). Solution of sodium or potassium acetate is added; and if the reaction is not markedly acid, it is made so by addition of acetic acid. Ferric chloride (if not present) is now added, drop by drop, avoiding an excess. The precipitate, ferric phosphate, is brownish- white. With zinc and manganous salts, the precipitate is dimetallic or normal — ZnHP0 4 , or Zn 3 (P0 4 ) 2 — according to the conditions of precipitation. When a manganic compound is mixed with aqueous phosphoric acid, the solution evaporated to dryness and gently ignited, a violet or deep blue mass is obtained, from which water dissolves a purple-red manganic hydrogen phosphate, a distinction from manganous compounds. With salts of nickel, a light green normal phosphate is formed; with cobalt, a reddish normal phosphate. Soluble salts of the alkaline earth metals, with dimetallic alkali phos- phates, as Na 2 HP0 4 , form white precipitates of phosphates, two-thirds metallic, as CaHP0 4 ; with trimetallic alkali phosphates, white precipitates of phosphates, normal or full metallic, as Ca 3 (P0 4 ) 2 . The precipitates are soluble in acetic acid, and in the stronger acids. Concerning the am- monium magnesium phosphate, see §189, 6d. Magnesium salts with ammonium hydroxide give a precipitate of double pyrophosphate, soluble in alkali pyrophosphate solution. Magnesium salts with ammonium hydroxide are not precipitated by soluble metaphosphates unless very concentrated. Ammonium molybdate, in its nitric acid solution (§75, 6d), furnishes an exceedingly delicate test for phosphoric acid, giving the pale yellow pre- cipitate, termed ammonium phosphomolybdate. The molybdate should be in excess, therefore it is better to add a little of the solution tested (which must be neutral or acid) to the reagent, taking a half to one cc. of the latter in a test-tube. For the full delicacy of the test, it should be set' aside, at 30° to 40°, for several hours. Ammonium molybdate reacts but slowly with meta or pyrophosphate solutions — and not until orthophosphoric acid is formed by digestion with the nitric acid of the reagent solution. §255, 8. PHOSPHORIC ACID. 303 B. — With non-metals and their compounds. — Phosphoric acid is not reduced by any of the reducing acids. Phosphates of the first two groups are transposed by H 2 S, and of the first four groups by alkali sulphides with formation of a sulphide of the metal, except Al and Cr , which form .a hydroxide; phosphoric acid or an alkali phosphate is also formed. HC1 , HN0 3 , and H 2 S0 4 transpose all phosphates and all are transposed by acetic acid except those of Pb , Al and Fe"' phosphates. Sulphurous acid transposes the phosphates of Ca, Mg, Mn, Ag, Pb, and Ba, also the ;arsenite and arsenate of calcium (Gerland, J. C, 1872, 25, 39). Excess of phosphoric acid completely displaces the acid of all nitrates, chlorides, and sulphates upon evaporation and long-continued heating on the sand bath. 7. Ignition with metallic magnesium (or sodium) reduces phosphorus from phosphates to magnesium phosphide, P,Mg, , recognized by odor of PH g , formed on contact of the phosphide with water. A bit of magnesium wire (or of sodium) is covered with the previously ignited and powdered substance in a glass tube of the thickness of a straw, and heated. If any combination of phosphoric acid is present, vivid incandescence will occur, and a black mass will be left. The latter, crushed and wet with water, gives the odor of phos- phorus hydride. Orthophosphoric acid heated to 213° forms pyrophosphoric acid; when heated to dull redness the meta acid is obtained, which sublimes upon further heating without change. Phosphoric anhydride, P 2 5 , cannot be prepared by ignition of phosphoric acid. Tribasic orthophosphates, normal pyrophosphates, and metaphosphates of metals whose oxides are not volatile and not decomposed by heat alone are unchanged upon ignition. Dimetallic orthophosphates, 3I' 2 HP0 4 , are changed to normal pyrophosphates upon ignition; also tribasic •orthophosphates when one-third of the base is volatile, e. g., MgNH 4 P0 4 . Mono-metallic or primary orthophosphates, M'H 2 P0 4 , become metaphosphates; also secondary or tertiary orthophosphates when only one atom of hydrogen is displaced by a metal whose oxide is non-volatile, e. g., NaN'H 4 HP0 4 . Acid pyrophosphates, M' 2 H 2 P 2 7 , form metaphosphates. When meta or pyro- phosphates are fused with an excess of a non-volatile oxide, hydroxide or carbonate the tertiary orthophosphate is formed (Watts', 1894, IV, 106). Phosphates of Al , Cr , Pe , Cu , Co , Ni , IHn , Gl and V when heated to a white heat with an alkali sulphate form oxides of the metals and an alkali tribasic orthophosphate; phosphates of Ba , Sr , Ca , Mg, Zn and Cd form double phosphates, partial transposition taking place {Derome, C. r., 1879, 89, 952; Grandeau, A. Ch., 1886, (6), 8, 193). 8. Detection. — The presence of orthophosphoric acid in neutral or acid solutions is detected by the use of an excess of an ammonium molybdate solution (§75, 6d). With pyro and metaphosphoric acids no reaction is obtained except as they are changed to the ortho acid by the reagents used. Disodium phosphate, Na 2 HP0 4 , after precipitation with silver nitrate, reacts acid to test papers. With trisodium phosphate the solu- tion is neutral (distinction). Orthophosphates are distinguished from pyro and metaphosphates by the color of the precipitate with silver nitrate : Ag 3 P0 4 is yellow, Ag 4 P 2 7 and AgP0 3 are white. Also by the fact that only the ortho acid is precipitated by ammonium molybdate. Nearly all pyrophosphates are soluble in sodium pyrophosphate, Na 4 P 2 7 (distinc- 304 SULPHUR. §255, 9. lion from orthophosphates). Hager (J. C, 1873, 26, 940) gives a method for detecting the presence of H 3 P0 3 , H 3 As0 3 , or HN0 S in H 3 P0 4 . Sodium metaphosphate does not give a precipitate with ZnS0 4 cold and in excess; with Na 4 P 2 7 and Na 2 H 2 P 2 7 a white precipitate of Zn 2 P 2 7 is obtained (Knorre, Z. angew., 1892, 639). 9. Estimation. — (a) By precipitation as magnesium ammonium phosphate, MgNH 4 P0 4 , and ignition to the pyrophosphate. (6) By precipitation and weighing as lead phosphate, Pb 8 (P0 4 ) 2 . (c) By precipitation from neutral or acid solution by ammonium molybdate and after drying at 140° weighing as ammonium phosphomolybdate. Consult Janovsky (J. C, 1873, 26, 91) for a review of all the old methods. §256. Sulphur. S = 32.07. Usual valence two, four and six (§14). 1. Properties. — Sulphur is a solid, in yellow, brittle, friable masses (from melting); or in yellowish, gritty powder (from sublimation) or in nearly white, slightly cohering, finely crystalline powder (by precipitation from its com- pounds). At — 50° it is white (Schoenbein, J. pr., 1852, 55, 161). The specific gravity of native sulphur is 2.0348 (Pisati, B., 1874, 7, 361). Melting point, 111 (Quincke, >b« 3 a "" 1 3 5 O 5§ = s 5 3 Q go o3 3 3 3 5 3 53112 1^3 © 3 +3 a n M m ■a" §.§ ■a ~5 C 'SI'S >> si +3 ft 3 m ■a 03 43 3 i '3 08 Is -o MS' ©©■3 3 3 h s2 IS*! Sjaax 3 ©bo .S'Sp .?3 a ■3 8 II. s ^ ■rtrjfl is* Sot a © *3„ 3 fe a 3 ^ 11 & © P £ © g Id 1 a © PL, o © 03 til 3 3 £ S«« gb«a ©Sa ».S.S MO 3 3 +3 fe+3 fa. © © m a pjsi a 3 3^ ^la 5 . +=i«jaB 1 ft '3 2 a 1 3 13 g fto © O *>-S 3 CO nj a ft ) ! ! =s 0^ .& ftBJ ft9 go 3 3 *| 3 *,§ fctSS a » Sdfl ©■a ©.a 03 -00 8| 03 a F "3 © S > S 0© ? '"Xi o.S £a feS» m 03 g « 3 I s > 43 +3 © s £ a £ 3 bj a ! 03 a a O s ft '3 2 a $ 3 d 03 +3 3 ¥ CD CD CJtH 0) Se; fe S* N !5 ■< ^ fi « n 'S '3 © a •a § © . A II 1 B-3 » ft| § a do ace 3 be . ■as H ^2 -O £ = ft© 33 k a O CO h © ©0, s" '3 OS 43 IS -2 3 I | 53 £•§ ■§1 g 3 a© S£ +3 03 oO Is ftx •3© 1.S ■3 2 © 3 i i?2 '8 o 5 5 -J 3. "3 £ a o O 03 c ti . tat) 111 £5 33 ©fe £|5 -si o >■*• 1" £ 3.3 O COa a I p. +3 2 2 ft c . 2 p. © a fe ©£ S c afl fc fc fc lz Z a O s 3 '3 OS © © ■a 3 a CO 05 o5 S bo § M o •3 >> s) O © i '3 * OB J 3 is £ ■£ 3 t 1 ^ 1 i © •a •a 3 a . 03 a «h a |i 3« l. c). HI and iodides form I and HBr , but in alkaline mixture an iodate and a bromide are produced. 7. Ignition. — Warming drives off all the bromine from its solutions in water or other solvents. Heat favors all reactions with bromine. 8. Detection. — Bromine is usually detected by shaking its solution in water with CS 2 , which dissolves it with a reddish-yellow color; if present in large quantities the color is brown to brownish black. In this case a large excess of CS 2 must be used or a very small portion of the unknown "taken, in order that the solution be dilute enough for the reddish-yellow bromine color to be distinguished from the violet color of iodine. Tether or chloroform may be used instead of carbon disulphide, but the solution is of a paler yellow. Starch solution gives a yellow color with bromine, but the reaction is less delicate than with CS 2 . 9. Estimation. — (a) The bromine is made to act upon KI , and the iodine •which is liberated is estimated by standard solution of Na 2 S 2 3 . (6) It is estimated by the amount of As 2 3 which it oxidizes in alkaline solution, (c) It is converted into HBr by H 2 S or H 2 SO s , and then precipitated by AgN0 3 . and weighed as AgvBr . §276, 6A. HYDROBROMIC ACID. 345, §276. Hydrobromic acid. HBr = 80.958. H'Br-', H — Br. 1. Properties.— Molecular weight, 149.9. Vapor density, 39.1. A colorless gas„ condenses to a liquid at — 69° and solidifies at — 73° (Faraday, A., 1845, 56, 155)* Its aqueous solution is colorless and is not decomposed by exposure to the air. The specific gravity of the saturated solution at 0° is 1.78; containing 82.02 per cent HBi, or very nearly HBr.H 2 . If a saturated solution is boiled,, chiefly HBr is given off, and if a dilute solution is boiled, chiefly H 2 is given off, until in both cases the remaining liquid contains 47.38 to 47.86 per cent of HBr , its sp. gr. 1.485, its boiling point constant at 126°, and its composition almost exactly HBr.5H 2 , which distils over unchanged. Its vapor density of 14.1 agrees with the calculated vapor density of HBr.5H 2 . 2. Occurrence. — Not found free in nature, in combination as bromides in sea. water and in some minerals. 3. Formation. — (a) By action of bromine upon phosphorus immersed in water, the amorphous phosphorus is preferred: P 4 + 10Br 2 + 16H 2 = 4H„P0 4 + 20HBr . (6) By action of H s P0 4 or H 2 S0 4 on KBr (Bertrand, J. C, 1876, 29,. 877). (c) By transposition of BaBr 2 by cold dilute H 2 S0 4 added in molecular proportions, ((f) By passing a mixture of Br and H over platinum sponge, (e) By action of Br on H 3 P0 2 . (f) By adding Br to Na 2 SO s . Metallic bromides are formed: (i) By direct union of the elements, but in a. few cases heat is required to effect the combination. (2) By action of HBr upon the metallic oxides, hydroxides and carbonates. (3) Many bromides are formed by action of HBr on the free metal, ous salts and not ic being formed. (4) Bromides of the first group are best made by precipitation. (5) Bromides of K , Ma , Ba , Sr and Ca are made by the action of bromine on their hydrox- ides and subsequent fusion: 6KOH + 3Br 2 = KBrO„ + 5KBr + 3H 2 2XBrO„ (ignited) = 2KBr + 30 2 4. Preparation. — (a) H 2 S is added to a solution of bromine in water until the yellow color disappears; the solution is then distilled. The first portion of the distillate is rejected if it contains H 2 S, and the latter portion if it con- tains H 2 S0 4 (Recoura, C. r., 1890, 110, 784). (6) H 2 S0 4 is added to a concen- trated solution of KBr; after twenty-four hours the greater portion of the KHS0 4 has crystallized out. The remaining liquor is then distilled. The product usually contains traces of H 2 S0 4 . (c) By passing bromine into hot paraffine (Crismer, B., 1884, 17, 649). 5. Solubilities. — Silver and mercurous bromide are insoluble in water,, lead bromide is sparingly soluble ; all other bromides are soluble. Hydro- bromic acid and soluble bromides precipitate solutions of the metals of the first group, lead salts incompletely. Lead bromide is less soluble than the corresponding chloride. The presence of soluble bromides increases the solubility of lead bromide. A small amount of hydrobromic acid, decreases its solubility, but a larger excess increases it (Ditte, C. r., 1881^ 92, 718). In alcohol, the alkali bromides are sparingly or slightly soluble : calcium bromide, soluble; mercuric bromide, soluble; mercurous bromide, insolu- ble. Silver bromide is soluble in NH 4 0H . 6. Reactions. — A. — With metals and their compounds. — Hydrobromic acid dissolves many metals with the formation of bromides and evolution of hydrogen, e. g., Pb , Sn , Fe , Al , Co , Ni , Zn , and the metals of the 346 SYDROBROMIG ACID.. % §276, 6A, 1. calcium and the alkali groups. It unites with salt forming oxides and hydroxides to produce bromides without change of valence: PbO 4- 2HBr = PbBr 2 + H 2 . But if the valence of the metal in the oxide or hydroxide is such that no corresponding bromide can be formed, then reduction takes place as follows : 1. Pb"+ n becomes PbBr 2 and Br . 2. As v becomes As'" and Br . The HBr must be concentrated and in excess, and the As v compound merely moistened with water: H,As0 4 + 2HBr = H,As0 3 + Br 2 + H 2 . In presence of much water the reverse action takes place : H 3 As0 3 + Br 2 + H 2 = H 3 As0 4 + 2HBr . 3. Sb v becomes Sb'" and Br . 4. Bi v becomes BiBr 3 and Br . 5. Fe VI becomes Fe'" and uot Fe" , and Br . 6. Cr^ 1 becomes CrBr 3 and Br (a separation from a chloride if the solu- tion be dilute) (Friedheim and Meyer, Z. anorg., 1891, 1, 407). KBr is not decomposed by a boiling concentrated solution of K 2 Cr 2 7 (separation from KI) (Deehan, /. C, 1887, 51, 690). 7. Co" +n becomes CoBr 2 and Br . 8. Ni"+ n becomes NiBr 2 and Br . 9. Mn"+ n becomes MnBr 2 and Br (§269, 8; Jannasch and Aschoff, Z. anorg., 1891, 1, 144 and 245). KMn0 4 liberates all the bromine from KBr in presence of CuS0 4 (a separation of bromide from chloride (Baubigny and Eivals, C. r., 1897, 124, 859 and 954). Silver nitrate solution precipitates, from solutions of bromides, silver oromide, AgBr, yellowish-white in the light, slowly becoming gray to black. The precipitate is insoluble in, and not decomposed by, nitric acid, soluble in concentrated aqueous ammonia, nearly insoluble in concentrated solution of ammonium carbonate, slightly soluble in excess of alkali bromides, soluble in solutions of alkali cyanides and thiosulphates. It is slowly decomposed by chlorine. Solution of mercurous nitrate precipitates mercurous oromide, HgBr, yellowish-white, soluble in excess of alkali bromides. Solutions of lead salts precipitate, from solutions not very dilute, lead oromide, PbBr 2 , white. B. — With non-metals and their compounds. 1. H 3 Fe(CN) 6 becomes H 4 Fe(CN) 6 and Br . The HBr must be in excess and concentrated, also the ferricyanide should be merely moistened with water, as in the presence of much water the reverse action takes place: 2X 4 Fe(CN) e + Br 2 = 2K 3 Fe(CN) e + 2KBr . 2. HN0 2 , in dilute solutions, no action (distinction from HI) (Gooch and Ensign, Am. S., 1890, 140, 145 and 283). HN0 3 becomes NO and Br . §276, 8. HYDROBROMIC ACID. 347 8. Phosphorus compounds are not reduced. k. H 2 S0 4 becomes S0 2 and Br . Both acids must be concentrated and hot, otherwise the reverse action takes place: S0 2 + Br 2 -4- 2H 2 = H 2 S0 4 + 2HBr . With H 2 S0 4 , sp. gr. 1.41, no bromine is set free even when solution is boiled (Feit and Kubierschky, J. Pharm., 1891, (5), 24, 159). The bromine of bromides is all liberated when warmed to 70° or 80° with ammonium persulphate (separation from a chloride) (Engel, C. r., 1894, 118, 1263). 5. Chlorine liberates bromine from all bromides, even from fused silver bromide (Nihoul, Z. angew., 1891, 441). HC10 3 becomes HC1 and Br . If the HC10 3 be concentrated other pro- ducts may appear. 6. HBrO liberates Br fro:$i both acids; the same with HBr0 3 . 7. HIO3 becomes I and Br . 8. Hydrogen peroxide liberates the bromine from hydrobromic acid at 100° (a distinction and separation from chloride). The bromine can best be removed by aspiration (Cavazzi, Gazzetta, 1883, 13, 174). 7. Ignition. — Some bromides can be sublimed undecomposed in presence of air; e. g., AsBr, , SbBr 3 , HgBr and HgBr 2 . Some can be sublimed only by exclusion of air and moisture; e. g., AlBr 3 and NiBr 3 . Bromides of sodium and potassium are not changed by heat. Silver bromide melts undecomposed. Many bromides, however, are more or less decomposed when ignited in pres- ence of air and moisture: CuBr 2 becomes CuBr and Br . 8. Detection. — Bromides are usually oxidized to free bromine, which is detected by its physical properties and by its color when dissolved in CS 2 (§275, 5). The oxidizing agent used to liberate the bromine varies according to the conditions. Chlorine is more commonly employed and acts when cold (6B5). A large excess of chlorine is to be avoided, as it decolorizes bromine solutions with formation of a chlorbromide. Nitric acid when dilute acts slowly unless hot. H 2 S0 4 , dilute, fails to oxidize the HBr even when hot; but when concentrated and hot is sometimes preferred. If chlorine be used, the mixture if alkaline must first be acidified; otherwise a colorless bromate will be formed, free bromine not being a visible intermediate step in the oxidation: KBr + 6K0H + 3C1 2 = KBr0 3 + 6KC1 + 3H 2 . If an iodide be present: (a) In absence of a chloride precipitate with silver nitrate, and digest the precipitate with HB^OH , which will dissolve the AgBr and none of the Agl . The nitrate may be treated with H 2 S , which precipitates the silver as Ag 2 S , leaving the bromine in the nitrate as NH 4 Br , which may be detected in the usual way. (J) To the acid mixture add chlorine water and carbon disulphide, shake and continue the addition of the chlorine water until the violet color of the iodine solution disappears, when the brown color due to the bromine may be observed: 2KI + SKBr + 7C1 2 + 6H 2 = 2HI0 3 + Br 2 348 B7P0BR0M0V8 ACID—BROMW ACID. §276, 9„ -f- 4ZC1 -4- 10HC1 . (c) To the solution from which the bases have been, removed add a cold saturated solution of potassium chlorate and dilute- sulphuric acid (one of acid to four of water); warm until the solution is. of a pale straw color, or colorless if only iodides are present. It may be- necessary to add more of the solution of potassium chlorate to complete- the oxidation of the iodine. Dilute the solution with water, cool, and shake with carbon disulphide. See also §269, 8. 6X1 + 6KBr + 2KC10 3 + 7H 2 S0 4 = 3l 2 + 3Br 2 + 7K 2 S0 4 + 2HC1 + 6H 2 6I 2 + xBr 2 + 10KCIO 3 + 5H 2 S0 4 + 6H 2 = 12HIO„ + xBr 2 + 5K 2 S0 4 + 10HCI 9. Estimation. — (a) It is converted into AgBr , and after gentle ignition weighed as such. (6) The bromide is oxidized to free bromine, which is passed into a solution of KI and the liberated iodine titrated with standard Ma 2 S 2 O s . (c) The bromide is oxidized to bromine, which is passed into an alkaline solution of arsenous acid. The excess of the arsenous acid is titrated with a standard solution of KMnO, . §277. Hypobromous acid. HBrO = 96.958 . H'Br'O-", H — — Br. The anhydride, Br 2 , has not been isolated. The acid, HBrO , is a very- unstable yellow liquid, a strong oxidizing and bleaching agent. The hypo- bromites are less stable than the corresponding hypochlorites. The calcium, and the alkali group hypobromites may be prepared by adding bromine to the respective hydroxides in the cold. The free acid is obtained by the action of bromine upon mercuric oxide: 2HgO + 2Br 2 + H 2 = Hg 2 OBr 2 + 2HBrO; also by the action of bromine upon silver nitrate: AgN0 3 + Br 2 -f- Hj-0 = AgBr + HBrO + HNO, (Dancer and Spiller, C. N., 1860, 1, 38; 1862, 6, 249). The free acid as an oxidizing agent reacts in many cases similar to free bromine. With HBr free Br is obtained from both acids (Schoenbein, J. pr.> 1863, 88, 475). §278. Bromic acid. HBr0 3 = 128.958 . H'Br v 0-" 3 , H — — Br ~ jj 1. Properties. — The anhydride, Br 2 5 , has not been isolated; and the acid, HBr0 3 , is known only in solution. It is a colorless liquid, smelling like bro- mine. It is a strong oxidizing agent. The solution of HBr0 3 is decomposed upon boiling, but by evaporating in a vacuum, a solution containing about 50 per cent of the acid may be obtained. 2. Occurrence. — Neither the acid nor its salts are found in nature. 3 Formation.— (a) By the electrolysis of HBr (Eiche, G. r., 1858, 46, 348). (6) By the decomposition of AgBr0 3 by Br: 5AgBr0 3 + 3Br 2 + 3H 2 = 5 AgBr + 6HBr0 3 . (c) An alkali bromate is made by adding bromine to a solution of chlorine in sodium ca.-bonate (Kaemmerer, J. pr., 1862, 85, 452). 4. Preparation.— Bromates of Ba , Sr , Ca , K and Na are made by the action of bromine upon the respective hydroxides at 100° : 6KOH + 3Br 2 = 5KBr ■+- KBrO s + 3H 2 . The free acid is prepared by adding dilute HjSO,, in slight excess to Ba(Br0 3 ) 2 ; the slight excess of H 2 S0 4 being removed by the cautious addition of Ba(OH) 2 . §278,8. BROMIC ACID. 349 5. Solubilities. — AgBr0 3 is soluble in 123 parts of water at 24.5° (Noyes, Z. phys. Ch., 1890, 6, 246). Ba(Br0 3 ) 2 is soluble in 124 parts of water at ordinary temperature and in 24 parts at 100° (Kammelsberg,. Pogg., 1841, 52, 81 and 86). Witb the exception of some basic bromates,. all other bromates are soluble in water. 6. Eeactions. — A.— With metals and their compounds. — Bromic acid is. a powerful oxidizing agent, acting in most respects like free bromine. It is usually reduced to hydrobromic acid, sometimes only to free bromine :. 1. Hg' becomes Hg" and a bromide. 2. As'" becomes As v and a bromide. S. Sb'" becomes Sb v and a bromide. 4. Sn" becomes Sn IV and a bromide. 5. Cu' becomes Cu" and a bromide. 6. Fe" becomes Fe'" and a bromide. 7. Mn" becomes Mn0 2 and bromine. 8. Cr'" becomes H 2 Cr0 4 and bromine. Silver nitrate precipitates in solutions not very dilute, silver Inmate* AgBr0 3 , white, sparingly soluble in water, soluble in ammonium hydroxide,, easily soluble by nitric acid, its color and solubility in ammonium hydroxide differing a little from the bromide (§276, 5). It is decomposed by hydro- chloric acid with evolution of bromine — a distinction from bromides and from other argentic precipitates. B. — With non-metals and theix compounds. 1. H 2 C 2 4 becomes C0 2 and Br. An excess of hot H 2 C 2 4 changes the- Br to HBr (Guyard, Bl., 1879, (2), 31, 299). HCNS becomes H 2 S0 4 , HBr and other products. H 4 Fe(CN) 8 becomes H 3 Fe(CN) 6 and HBr . An excess of HBr0 3 carries, the oxidation farther. 2. HN0 2 reduces HBr0 3 , forming HN0 3 and Br . 3. PH 3 , HH 2 P0 2 and H 3 P0 3 become H 3 P0 4 and HBr . 4. S and S0 2 become H 2 S0 4 and HBr . H 2 S forms first S then H 2 S0 4 . 5. HC1 becomes CI and Br . 6. HBr forms Br from both acids. 7. HI becomes I and Br . With an excess of HBr0 3 the products are HI0 3 and Br (Kaemmerer, I. c, Wittstein, Z., 1876, 15, 61). 7. Ignition. — All bromates are decomposed upon heating. KBr0 3 , NaBr0 3 and Ca(Br0 3 ) 2 evolve oxygen and leave the bromides. Co(Br0 3 ) 2 , Zn(Br0 3 ) 2 and other bromates evolve oxygen and bromine, leaving an oxide. 8. Detection. — The bromine is first liberated by some reducing agent that does not carry the reduction to the formation of HBr. H 2 C 2 4 is a 350 IODINE. §278, 9. very suitable agent for this purpose, since it does not change Br to HBr except when hot and concentrated. The Br is detected by CS 2 (§275, 8). Sulphuric and nitric acids liberate bromic acid from metallic bromates, the HBr0 3 remaining for some time intact, and the solution colorless. The gradual decomposition of the HBr0 3 is first a resolution into HBr and 0, and as fast as HBr is formed it acts with HBr0 3 , so as to liberate the bromine of both acids. Now, if the solution contained bromide as well as bromate, an abundance of free bromine is obtained immediately upon the addition of dilute sulphuric acid in the cold. Hence, if dilute sulphuric acid in the dilute cold solution does not color the carbon disulphide, and if the addition of solution of pure potassium bromide immediately develops the yellow color, while it is found that no other oxidizing. agent is present, we have corroborative evidence of the presence of a bromate. And, if we treat a solution known to contain bromide with dilute sulphuric acid and carbon disulphide, and obtain no color, we have conclusive evidence of the absence of bromates. Hydrochloric acid transposes bromates and quickly decomposes the bromic acid, liberating both bromine and chlorine. A mixture of bromate and iodate, treated with hydrochloric acid, fur- nishes bromine without iodine, coloring carbon disulphide yellow. The ignited residue of bromates, in all cases if the ignition be done with sodium carbonate, will give the tests for bromides. 9. Estimation. — The bromate is reduced to free bromine or to a bromide and determined as such. §279. Iodine. I = 126.85. Usual valence one, five and seven (§12). 1. Properties. — Specific gravity, 4.948 at 17° (Gay-Lu,ssac). Melting point, 114.2°. Boiling point, 184.35° at 760 mm. pressure (Ramsay and Young, J. C, 1886, 49, 453). At ordinary temperature iodine is a soft gray -black crystalline solid with a metallic lustre. The thin crystals have a brownish-red appear- ance. Precipitated iodine is a brownish-black powder. It vaporizes very appreciably at ordinary room temperature with a characteristic odor, and may be distilled with steam. The molecule of iodine vapor under about 800° is I 2 ; above that temperature dissociation takes place, until at 1700° it is complete and the molecule consists of single atoms (Biltz and Meyer, B., 1889, 22, 725). The vapor of iodine unmixed with other, gases is deep blue, mixed with air or other gases it is a beautiful violet. It is sparingly soluble in water to a "brown or yellowish-brown solution, which slowly bleaches litmus paper. It stains the skin yellow-brown. The solution gradually decomposes in the sun- light with formation of HI. It reacts similarly to bromine and chlorine, but With much less intensity. The free element combines with starch,* forming a compound of an intense blue color. This colored body is quite stable in the cold; decolors upon warming, the color returning upon cooling. The reaction of iodine with starch constitutes a very delicate reaction for the detection of the presence of iodine. It also serves as an indicator in the volumetric estima- tion of iodine, as all reducing agents destroy the color by taking the iodine into combination. Combined iodine does not react with starch. * The compound formed when iodine unites with starch is regarded by Bondonneau {Bl., 1877, <2), 28, 452) as an addition compound of the composition (C,H 10 O 5 ) 2 I . §279, 6A, 5. IODINE. 351 Colorless solutions are formed by all the alkali hydroxides with iodine; the fixed alkali hydroxides forming- iodides and iodates. With ammonia in water solution it dissolves more slowly, becoming colorless; the solution contains the most of the iodine as ammonium iodide, and deposits a dark-brown powder, termed " iodide of nitrogen," very easily and violently explosive when dry. According to Chattaway (Am., 1900, 24, 138) this compound has the composi- tion N a H 8 I 3 . The anhydride of iodic acid, I 2 5 , is the only stable compound of iodine and oxygen. The chief acids of iodine are: Hydriodic acid, HI; iodic acid, HIO s ; periodic acid, HI0 4 . Hypoiodous acid is said to be formed by the action of alcoholic iodine upon freshly precipitated mercuric oxide (Lippmann, G. r., 1866, 63, 968). Lunge and Schoche (B., 1882, 15, 1883) prepared iodide of lime which seemed to contain calcium hypoiodite, Ca(IO) 2 . 2. Occurrence Found free in some mineral waters (Wanklyn, 0. N., 1886, 54, 300). As iodides and iodates in sea water (Sonstadt, C. N., 1872, 25, 196, 231 and 241). In the ashes of sea plants. In small quantities in several minerals, especially in Chili saltpeter as sodium iodate. 3. Formation. — From iodides by nearly all oxidizing agents: 2KI + Br 2 = 2KBr + I,; and from iodates by nearly all reducing agents: 2HIO, + 5H = C,0, = I„ + 10CO 2 + 6H 2 . 4. Preparation. — (a) The ashes of the sea plants are digested in hot water and from the filtrate most of the salts removed by evaporation and crystalliza- tion. The iodides remain in the mother liquor and from this the iodine is obtained by treatment with Mn0 2 and H 2 SO t . (6) The sodium iodate in the mother liquor of the Chili saltpeter is reduced with S0 2 , the iodine precipitated as Cul with CuS0 4 . From the precipitate the iodine is recovered by distilla- tion with Mn0 2 and H 2 S0 4 . By far_the greatest portion of the iodine and iodides of commerce is obtained from the Chili saltpeter deposits. 5. Solubilities. — It is soluble in about 5500 parts water at 10° to 12° (Wittstein, J., 1857, 123), differing from CI or Br in that it forms no hydrate. It is much more soluble in water containing hydriodic acid or soluble iodides. From a concentrated solution in KI the compound KX, has been obtained. Iodine dissolves in very many organic solvents as alcohol, ether, chloroform, glycerol, benzol, carbon disulphide, etc. Car- bon disulphide readily removes the iodine from its solution or suspension in water; with small amounts of iodine imparting to the carbon disulphide a beautiful violet color, with large amounts the CS 2 solution is almost black. 6. Reactions. — A. — With metals and their compounds. — It unites slowly by the aid of heat with Pb and Ag; more rapidly with Hg, As, Sb, Sn, Bi , Cu , Cd , Al , Cr , Fe , Co , Ni , Mn , Zn , Ba , Sr ; Ca , Mg , K and Na . In oxidizing metallic compounds the iodine invariably becomes HI or an iodide, depending upon whether the mixture be acid or alkaline. It may, however, with certain substances act as a reducing agent, becoming oxidized to iodate or periodate. 1. Hg' becomes Hg" in acid and in alkaline mixture. 2. As'" becomes As v in presence of alkalis only. 3. Sb'" becomes Sb v in presence of alkalis only. •4. Sn" becomes Sn 17 in acid or in alkaline mixture. 5. Cr'" becomes Cr" in presence of alkalis only. 352 IODINE. §279, GA, 6.. 6. Fe" becomes Pe'" in presence of alkalis only. 7. Co" becomes Co'" in presence of alkalis only. 8. Ni" is not oxidized. 9. Mn" becomes Mn IV in presence of alkalis only. B. — With non-metals and their compounds. 1. K 4 Fe(CN)„ is oxidized, forming K 3 Fe(CN) 6 and KI, action slow and incomplete. 2. HN0 3 forms HI0 3 and NO . Strong HN0 3 must be used (at least. sp. gr. 1.42). Action is slow. A very good method of making HI0 3 . 3. HH 2 P0 2 becomes H 3 P0 4 with acids and with alkalis. 4. H 2 S becomes S and HI; no action if both substances be perfectly dry (Skraup, C. C, 1896, i, 469) (separation of H 2 S from AsH 3 ). According to Saint-Gilles (A. Ch., 1859, (3), 57, 221), in alkaline mixture from six to seven per cent of the sulphur is oxidized to a sulphate. H 2 S0 3 becomes H 2 S0 4 and HI. With a thiosulphate a tetrathionate is. formed: 2Na 2 S 2 3 + I 2 = Na 2 S 4 6 + 2NaI (Pickering, J. C, 1880, 37, 128). 5. CI becomes IC1 or IC1 3 , depending upon the amount of chlorine present, water should be absent. In the presence of water HC1 and HI0 3 are formed; in alkaline mixture a chloride and a periodate: I 2 + 7C1 2 -+- 16NaOH = 14NaCl + 2NaI0 4 + 8H 2 . HC10 3 forms HI0 3 and HClr 5HC10 3 + 3I 2 + 3H 2 = 6HI0 3 + 5HC1 . 6. Br becomes IBr, decomposed by water (Bornemann, A., 1877, 189,. 183). In alkaline mixture with an excess of Br a bromide and an iodate: I 2 + 5Br 2 + 12K0H = 2KI0 3 + lOKBr + 6H 2 . HBr0 3 becomes Br and HI0 3 . 7. Iodine combines with KI in concentrated solution to form ZI 3 (KII 2 ) . 7. Ignition. — See I. 8. Detection. — Iodine is recognized by the yellow to black color when mixed with water; the violet color when dissolved in carbon disulphide; the reddish color when dissolved in chloroform or ether; the blue color when added to a cold solution of starch; the violet color of the vapors, etc. The presence of tannin interferes with the usual tests for iodine unless a drop or two of ferric chloride solution be added (Tessier, Z., 1874, 11, 313). 9. Estimation. — (a) It is reduced to an iodide, precipitated with AgN0 3 , and after drying at 150°, weighed as Agl . It is estimated volumetrieally with a standard solution of N , a 2 S 2 3 , using starch as an indicator. (6) The iodine dissolved in potassium iodide is treated with an alkaline solution of hydrogen peroxide in an azotometer, the oxygen liberated being a measure of the amount of iodine present (Baumann, Z. angew., 1891, 204). §280, 5. HYDRIODIC ACID. 353 §280. Hydriodic acid. HI = 127.858. H'l-' , H — I . 1. Properties. — Molecular weight, 127.858. Vapor density, 63.927. A colorless incombustible gas. At atmospheric pressure it solidifies at — 51°. At 0° it liquefies under a pressure of 3.97 atmospheres (Faraday, A. Ch., 1845, (3), 15, 266). The constant boiling point of the aqueous solution of the gas is 127°, ■which solution contains 57 per cent of HI and has a specific gravity of 1.694 (Eoscoe, J. C, 1861, 13, 160). Gaseous HI is dissociated by heat, slowly at 260°; rapidly at 240° (Lemoine, A. Ch., 1877, (5), 12, 145). Iodine separates from the water solution of the acid when exposed to the air. 2. Occurrence. — Not found free in nature, but in combination as iodide or iodate. 3. Formation. — (a) By direct union of the elements at a full red heat (Merz and Holzmann, B., 1889, 22, 869). (6) By direct union of the elements in pres- ence of platinum black at 300° to 400° (Lemoine, C. r., 1877, 85, 34). (c) From J3al 2 by adding H,SO„ in molecular proportions, (d) By the action of iodine upon Na 2 SO„ or Na 2 S 2 3 (Mene, C. r., 1849, 28, 478). (e) By the action of iodine upon moist calcium hypophosphite: Ca(H 2 P0 2 ) 2 -\- 4I 2 + 4H 2 = CaH<(P0 4 ) a -4- 8HI (Mene, I. a). Iodides are formed by the direct action of iodine upon the metals; or better, hy'the action of HI upon the oxides, hydroxides or carbonates of those metals ■whose iodides are soluble in water. Iodides of lead, silver and mercury are formed by precipitation. 4. Preparation. — (a) By passing H 2 S into a mixture of finely divided iodine suspended in water, adding more iodine as fast as the color disappears: 2I 2 + SH 2 S = 4HE + S 2 (Pellagri, Gazeetta, 1875, 5, 423). (6) By bringing moist red phosphorus in contact with iodine: P 4 + 10I 2 + 16H 2 = 4H 3 P0 4 + 20HI (Meyer,B., 1887, 20, 3381). (c) By passing vapors of iodine into hot liquid paraffine (Crismer, B., 1884, 17, 649). (d) By heating iodine with copaiba oil (Bruylants, B., 1879, 12, 2059). It cannot be prepared by adding H 2 SO, to an iodide and distilling (5). 5. Solubilities. — Iodides of lead, silver, mercury and cuprosum are in- .soluble. Iodides of other ordinary * metals are soluble, those of bismuth, tin and antimony requiring a little free acid to hold them in solution. Lead iodide is sparingly soluble in water (§57, 5c). Mercuric iodide is readily soluble in excess of potassium iodide, forming a double iodide, X,HgI 4 ; most other iodides are more soluble in a solution of potassium iodide than in pure water. The iodides of the alkalis, Ba , Ca and Hg" are soluble in alcohol; Hgl and Agl are insoluble. All iodides in solution are transposed by HC1 or by dilute H 2 S0 4 . Hot concentrated H 2 S0 4 decomposes all iodides, those of Pb, Ag and Hg slowly but completely, S0 2 and I being produced: 2KI + 2H 2 S0 4 = K 2 S0, + I 2 + S0 2 + 2H 2 . HNOg in excess first transposes then decomposes soluble iodides: 6KI -\- 8HN0 3 = 6KNO3 + 3I 2 + 2N0 + 4H 2 . If the HN0 3 be concentrated the iodine is further oxidized: 3I 2 + 10HN0 3 == 6HI0 3 + 10NO + 2H 2 . Long-continued boiling with HN0 3 , sp. gr. 1.42, decomposes the insoluble iodides. Chlorine in the cold decomposes all soluble iodides, by heating with chlorine the insoluble iodides are also decomposed: 2KI -f~ Cl 2 = * Thallium Iodide, Tl I, is perfectly insoluble in cold water, a distinction and separation from bromides and chlorides (Huebner, Z., 1872, 11, 397). Palladous iodide is insoluble in water. 354 ETDRIODIO ACID. §280, 6A. 2KC1 -+- I 2 . With an excess of chlorine the iodine is further oxidized: I 2 + 5C1 2 + 6H 2 = 2HI0 3 + 10HC1 . Silver iodide is almost insoluble in ammonium hydroxide or ammonium carbonate (distinction from silver chloride). It is soluble in KCN . Agl and Pbl 2 are soluble by decomposi- tion in solution of alkali thiosulphates : Agl + Na 2 S 2 3 = Nal -j- NaAgS 2 3 . Lead iodide is soluble in a solution of the fixed alkalis. 6. Reactions. — A. — With metals and their compounds. — Silver nitrate solution in excess precipitates, from solutions of iodides, silver iodide, Agl , yellow-white, blackening in the light without appreciable separation of iodine. For solubilities see paragraph above. Solution of mercuric chloride precipitates the bright, yellowish-red to red, mercuric iodide, Hgl 2 . The precipitate redissolves on stirring, after slight additions of the mercuric salt, until equivalent proportions are reached, when its color deepens. For the solubilities of the precipitate see §58, 6/. Solution of mercurous nitrate precipitates mercurous iodide, Hgl , yellow to green (§58, 6f). Solution of lead nitrate or acetate precipitates, from solutions of iodides not very dilute, lead iodide, Pbl 2 , bright-yellow — soluble, as stated in full in §57, 5c. Palladous chloride, PdCl 2 , precipitates, from solutions of iodides, pal- ladous iodide, Pdl 2 , black, insoluble in water, alcohol or dilute acids, and visible in 500,000 parts of solution. The reagent does not precipitate bromine at all in moderately dilute solutions, slightly acidulated with HC1 . Palladous iodide is slightly soluble in excess of the alkali iodides, and is soluble in ammonium hydroxide (§106). Copper salts precipitate from solutions of iodides cuprous iodide (white) mixed with iodine (black) : 2CuS0 4 + 4KI = 2CuI + 2K 2 S0 4 + I 2 . If sufficient reducing agents (as sulphurous acid) are present to reduce the liberated iodine to HI , only the white cuprous iodide will be precipitated (a distinction from bromides and chlorides). When metals are attacked by HI an iodide is formed and hydrogen is evolved. Hydriodic acid unites with all metallic oxides and hydroxides (expect ignited Cr 2 3 ) to form iodides; frequently, however, iodine is liberated and an iodide of lower metallic valence is formed: 1. Pb" +n becomes Pb" . 2. As v becomes As'" ; ZI has no action upon normal K 3 As0 4 (Friedheim and Meyer, Z. anorg., 1891, 1, 409). 8. Sb v becomes Sb'" . Jf. Bi v becomes Bi'" . 5. Cu" becomes Cu' . Soluble iodides reduce normal cupric salts, but have no reducing action in alkaline mixture or upon cupric hydroxide. With phenylhydrazine sulphate and cupric sulphate the iodine of iodides is §280, 65, 6. ETDBIODIC ACID. 355 completely precipitated (separation from chlorides) (Raikow, Ch. Z., 1894, 18,1661). 6. Fe'" becomes Fe" (§269, 8). 7. Cr VI becomes Cr'" . K 2 Cr0 4 is not reduced by KI even upon boiling the concentrated solutions. K 2 Cr 2 T with KI slowly gives I and Cr'" in the cold. When KI is boiled with a concentrated solution of K 2 Cr 2 7 the iodine is completely liberated (separation from bromides and chlorides which are unchanged): 6KI + 5K 2 Cr 2 7 = 8K 2 Cr0 4 + Cr 2 3 + 3I 2 (Dechan, J. C, 1886, 50, 682; 1887, 51, 690). When Agl is boiled with K 2 Cr 2 T and H 2 S0 4 no iodine is evolved, chromium is reduced and the iodide becomes silver iodate: K 2 Cr 2 7 + Agl + 5H 2 S0 4 = 2KHS0 4 + Cr 2 (S0 4 ) 3 + AgI0 3 + 4H 2 (Macnair, J. C, 1893, 63, 1051). 8. Co" +n becomes Co"; KI has no reducing action upon cobaltic hy- droxide. 9. Ni" +n becomes Ni" ; KI reduces Ni'" , liberating iodine. 10. Mn"+ n becomes Mn" . When KI is boiled with KMn0 4 the manga- nese becomes Mn0 2 and the iodide is oxidized to an iodate: 6KMn0 4 + 3KI + 3H 2 = 3KI0 3 + 6Mn0 2 + 6K0H (Groeger, Z. angew., 1894, 13 and 52) (distinction from bromides, which do not decolor permanganates). B. — With non-metals and their compounds. 1. H 3 Fe(CN) 6 forms H 4 Fe(CN) 6 and I; the reaction also takes place in neutral mixture. 2. HN0 2 forms NO and I (separation of iodide from bromide and chloride) (Jannasch and Aschoff, Z. anorg., 1891, 1, 144 and 245). HN0 3 forms NO and I, with further oxidations to HI0 3 with concen- trated HN0 3 . The HN0 2 acts much more rapidly than the HN0 3 . 8. Ko reduction with phosphorous compounds. Jf. H 2 S0 4 dilute no action; with the concentrated acid in excess, S0 2 and I are formed: 2KI + 3H 2 S0 4 = I 2 + S0 2 + 2KHS0 4 + 2H 2 ; if KI be added in excess to boiling H 2 S0 4 , H 2 S and I are formed: 8KI + 9H 2 S0 4 = 4I 2 + H 2 S + 8KHS0 4 + 4H 2 (Jackson, J. C, 1883, 43, 339). Ammo- nium persulphate liberates iodine from iodides at ordinary temperature (Engel, C. r., 1894, 118, 1263). 5. CI in excess forms HC1 and HI0 3 ; with excess of HI , HC1 and I are formed. In the presence of a fixed alkali a periodate and a chloride are formed: KI + 8K0H + 4C1 2 = 8KC1 + KI0 4 + 4H 2 . Hypochlorous acid oxidizes to iodine, then to iodic in acid solution; in alkaline solution to periodate. HC10 3 with excess of HI forms HC1 and I; with excess of HC10 S HC1 and HI0 3 . 6. Br forms I and HBr or a bromide. 356 HYDRIODIC ACID. §280, 65, 7. HBr0 3 with excess of HI forms HBr and I ; with excess of HBr0 3 , Br and HI0 3 . 7. HIO3, iodine is liberated from both acids: HI0 3 + 5HI = 3I 2 -j- 3H 2 . HI0 4 gives iodine. 8. H 2 2 becomes H 2 , and I (§244, 6B6) (Cook, J. C, 1885, 47, 471). 9. Ozone promptly liberates .iodine from soluble iodides. Atmospheric oxygen decomposes HI and ferrous and calcium iodides slowly, the alkali iodides not at all. 7. Ignition As a general rule iodides strongly ignited in presence of air and moisture evolve iodine, leaving the oxide of the metal. Ignited in absence of air or moisture the following iodides are not decomposed: KI , Nal , Bal 2 , Cal 2 , Srl 2 , Mnl 2 , AH, , Snl 4 , Pbl 2 , Agl and Hgl„ . See Mitscherlieh (Pogg., 1833, 29, 193), Personne (C. r., 1862, 54, 216) and Gustavson (A., 1873, 172, 173). 8. Detection. — The iodide is oxidized to free iodine by one of the. re- agents mentioned in (6) above. With a dry powder hot concentrated H 2 S0 4 is usually employed when the iodine is detected by the violet fumes evolved, condensing in the cooler portion of the test tube. With solu- tions the usual reagent is chlorine water. The iodine is recognized by the violet color when shaken with CS 2 , or the bright-red color with CHC1 3 . In case a large amount of iodine be present the CS 2 solution may be almost black. In this case large dilution with CS 2 is necessary to detect the violet color. If but a small amount of iodine be present the chlorine must be added very cautiously or the iodide will all be oxidized to the colorless iodic acid.* With small amounts of iodide, nitric acid is less liable to cause error as relatively much more nitric acid is required to oxidize the iodine to iodic acid. For the detection of small amounts of iodide a cuprie salt strongly acidulated with HC1 is an excellent reagent for the oxidation: 2CuCl 2 + SKI = 2CuCl + 2KC1 + I 2 . If insoluble iodides are present they should be transposed by H 2 S, the insoluble sulphide removed by nitration, the excess of H 2 S removed by boiling, and the solution then tested for hydriodic acid. Or the insoluble iodide should be reduced by Zn and H 2 S0 4 : 2AgI + Zn + H 2 S0 4 = 2Ag + ZnS0 4 + 2HI . The filtrate may then be tested for hydriodic acid. The insoluble iodide may also be fused with Na 2 C0 3 , and after digestion with water the filtrate acidulated and tested for hydriodic acid. That is, the solution must be acidulated before chlorine water is added, else the iodine will be oxidized to an iodate or periodate. 9. Estimation. — Gravimetrically by precipitation as Agl and weighing as such after gentle ignition. "Volumetrically by oxidation to iodine and titration with standard Na 2 S 2 3 (Groger, Z. angew., 1894, 52). * To test potassium bromide for traces of an iodide it is recommended to add CS„ and cuprie sulphate or a small amount of ferric alum. Or add chlorine water and then a few crystals of ferrous sulphate ; then shake with CS a (Brito, C. IV., 1884, 50, 210). §281, 6A. iodic ACID. 357 §281. Iodic acid. HI0 S = 175.858 . H'I v 0-" 3 , H — — 1 = ° 1. Properties.— Iodic acid is a white crystalline solid; its solution saturated at 14° contains 68.5 per cent HIO s , and has a specific gravity of 2.1629 (Kaem- merer, Pogg., 1869, 138, 390). At 170° it loses water, forming iodic anhydride, I 2 5 , a white crystalline solid, which, at 300°, dissociates into iodine and oxygen. See Ditte, A. Cll., 1870, (4), 21, 5. It is readily soluble in water and in alcohol; the solutions redden litmus and afterwards bleach it. 2. Occurrence. — The free acid is not found in nature. It is found as Ca(I0 3 ) 2 in sea water, and as sodium iodate in Chili saltpeter (Sonstadt, C. N., 1872 25 196, 231 and 241; Guyard, Bl., 1874, (2), 22, 60). 3. Formation.— (a) By electrolyzing a solution of I or HI (Riche, C. r., 1858, 46, 348). (6) By the action of chlorine on iodine in the presence of much water. The HC1 formed cannot be expelled by boiling without decomposing the HI0 3 . It must be removed by the careful addition of Ag 2 . (e) By adding water to IC1 3 and washing with alcohol: 2IC1 3 + 3H 2 = HI0 3 + 5HC1 + IC1 . () By reducing to an iodide and estimating as such. (c) By treating with KI acidulated with H 3 S0 4 , and titrating the iodine lib- erated with standard KTa 2 S 2 0, . 360 PERIODIC ACID. §282. §282. Periodic acid. HI0 4 = 191.858 . H H H \ I / [I SI/ H'l^O-"^ or S! 5 T n ^-\, H — — 1 = or H — — I — 0— H. The anhydride, I 2 0, , has not been isolated, and but one acid is known in the free condition, HI0 4 ,2H,0 or H,,IO„ . This acid exists in colorless monoclinic crystals, which do not lose water at 100°. It melts at 133°, and at a higher temperature it decomposes into iodic anhydride, water and oxygen (Kimmins, J. C, 1887, 51, 356; and 1889, 55, 148). Numerous periodates have been prepared as if derived from one or the other following named acids: HI0 4 , H 3 I0 5 , H 5 I0 6 , H 4 I 2 0„ , HsI.O^ , H 12 I 2 13 , H 10 I 4 O 19 , H 10 I a O 28 (Kammelsberg, Pogg., 1865, 134, 368, 499). The free periodic acid, H G IO, , is prepared: (a) By oxidizing iodine with per- chloric acid: 2HC10 4 + I 2 + 4H 2 = 2H B IO e -f Cl 2 (Kaemmerer, Pogg., 1869, 138, 406). (6) By heating iodine or barium iodide with a mixture of barium oxide and barium peroxide, digesting with water, and transposing the Ba s (IO e ) 2 thus obtained with the calculated amount of sulphuric acid (Kam- melsberg, Pogg., 1869, 137, 305). (c) By conducting chlorine into sodium iodate in presence of sodium hydroxide: NaI0 3 + 3NaOH + Cl 2 = Na 2 H s IO + 2NaCl . This acid periodate dissolved in water with a little nitric acid and then precipitated with silver nitrate, forms the silver salt, Ag 2 H 3 IO„ . This precipitate is dissolved in nitric acid and evaporated on the water-bath, when orange-colored crystals of silver meta periodate are formed according to the following: 2Ag 2 H 3 I0 6 + 2HN0 3 = 2AgI0 4 + aAgWO, + 4H 2 . Water decom- poses this precipitate: 2AgI0 4 + 4H 2 = H 5 I0 6 + Ag 2 H 3 IO B . Or the silver periodate, AgIO„ , is decomposed by CI or Br (Kaemmerer, I. c, p. 390). The silver salts vary in color: AgI0 4 is orange; Ag 2 HI0 5 , dark brown; Ag 4 I 2 0„ , chocolate colored; while silver iodate is white (a distinction). In the general reactions periodic acid and periodates resemble iodic acid and iodates. H 2 C 2 4 becomes C0 2 and I . H 3 P0 2 becomes H 3 P0 4 and HI . H 2 S becomes S and HI . H 2 S0 3 becomes H 2 S0 4 and HI0 3 without separation of iodine when the two acids are present in molecular proportions. The presence of a greater pro- portion of H„S0 3 causes, first, separation of iodine with final complete reduc- tion to HI (Selmous, B., 1888, 21, 230): HI0 4 + H 2 S0 3 = HI0 3 + H 2 S0 4 3HI0 4 + 8H 2 S0 3 = HI0 3 + I 2 + 8H 2 S0 4 + H 2 2HI0 4 + 7H 2 S0 3 = I 2 + 7H 2 S0 4 + H 2 HI0 4 + 4H 2 S0 3 = HI + 4H 2 S0 4 HC1 becomes CI and IC1 3 HI forms I from both acids. 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The removal of the latter can be effected, 1st, by combustion at a red or white heat, with or without oxidiz- ing reagents; 2d (in part), by oxidation with potassium chlorate and hydro- chloric acid on the water-bath (§69, 6'el); 3d, by oxidation with nitric acid in presence of sulphuric acid, at a final temperature of the boiling point of the latter (§79, 6'e3); 4th, by solvents of certain classes of organic substances; 5th, by dialysis. These operations are conducted as follows: §285. Combustion at a red or white heat, of course, excludes analysis for mer- cury, arsenous and antimonous bodies (except as provided in §70, 7), and ammonium. The last-named constituent can be identified from a portion of the material in presence of the organic matter (§207, 3). If chlorides are present some iron will be lost at temperatures above 100°, and potassium and sodium waste notably at a white heat, and slightly at a full red heat.. Certain acids will be expelled, and oxidizing agents reduced. The material is thoroughly dried and then heated in a porcelain or platinum crucible, at first gently. It will blacken, by separation of the carbon of the organic compounds. The ignition is continued until the black color of the carbon has disappeared. In special cases of analysis, it is only necessary to char the material; then pulverize it, digest with the suitable solvents, and filter; but this method does not give assurance of full separation of all sub- stances. Complete combustion, without use of oxidizing agents, is the way most secure against loss, and entailing least change of the material; it is, how- ever, sometimes very slow. The operation may be hastened, with oxidation of all materials, by addition of nitric acid, or of ammonium nitrate. The material is first fully charred; then allowed to cool till the finger can be held on the crucible; enough nitric acid to moisten the mass is dropped from a glass rod upon it, and the heat of the water-bath continued until the mass is dry, when it may be very gradually raised to full heat. This addition may be repeated as necessary. The ammonium nitrate may be added, as a, solid, in the same way. §286. Oxidation with potassium chlorate and hydrochloric acid on the water-oath does not wholly remove organic matter, but so far disintegrates and changes it that the filtrate will give the group precipitates, pure enough for most tests. It does not vaporize any bases but ammonium, but of course oxidizes or chlorinates all constituents. It is especially applicable to viscid liquids; it may be followed by evaporation to dryness and ignition, according to the paragraph above. The material with about an equal portion of hydrochloric acid is warmed on the water-bath, and a minute portion of potassium chlorate is added at short intervals, stirring with a glass rod. This is continued until the mixture is wholly decolored and dissolved. It is then evaporated to remove chlorine, diluted and filtered. If potassium and chlorine are to be tested for, another portion may be treated with nitric acid, on the water-bath. The .organic matter left from the action of the chlorine or the nitric acid may be sufficient to prevent the precipitation of aluminum and chromium in the third group of bases; so that a portion must be ignited. As to arsenic and antimony, see §70, 7. §292. PRELIMINARY EXAMINATIONS OF SOLIDS. 363 §287. The action of sulphuric icith nitric acid at a gradually increasing heat leaves behind all the metals (not ammonium), with some loss of mercury and arsenic (and iron?) if chlorides are present in considerable quantity. In this, as in the operations before mentioned, volatile acids are lost — sulphides partly ■oxidized to sulphates, etc The substance is placed in a tubulated retort, with about four parts of con- centrated sulphuric acid, and gently heated until dissolved or mixed. A funnel is now placed in the tubulure, and nitric acid added in small portions, gradu- ally raising the heat, for about half an hour— so as to expel the chlorine, and not vaporize chlorides. The material is now transferred to a platinum dish and heated until the sulphuric acid begins to vaporise. Then add small portions ■of nitric acid, at intervals, until the liquid ceases to darken by digestion, after a portion of nitric acid is expelled. Finally, evaporate off the sulphuric acid, using the lowest possible heat at the close. §288. The solvents used are chiefly ether for fatty matter, and alcohol or ether, ■or both successively, for resins. Instead of either of these, benzol may be -used; and many fats and some resins may be dissolved in petroleum ether. It will be observed that ether dissolves some metallic chlorides, and that alcohol dissolves various metallic salts. Before the use of either of these sol- vents upon solid material, it should be thoroughly dried and pulverized. Fatty matter suspended in water solutions may be approximately removed by filter- ing through wet, close filters; also by shaking with ether or benzol, and decant- ing the solvent after its" separation. §289. By Dialysis, the larger part of any ordinary inorganic substance can be extracted in approximate purity from the greater number of organic sub- stances in water solution. The degree of purity of the separated substance depends upon the kind of organic material. Thus albuminoid compounds are almost fully rejected; but saccharine compounds pass through the membrane •quite as freely as some metallic salts. (Consult Watts' Dictionary, 1894, IV, 173). PKELIMIJSTAKY EXAMINATION OP SOLIDS. §290. Before proceeding to the analysis of a substance in the wet way, a •careful study should usually be made of the reactions which the substance undergoes in the solid state, when subjected to a high heat, either alone or in the presence of certain reagents, before the blow-pipe, or in the flame of the Bunsen burner. This examination in the dry way precedes that in the wet, and should be carried on systematically, following the plan laid down in the tables, and noting carefully every change which the substance under investiga- tion undergoes, and if necessary making reference to some of the standard works on blow-pipe analysis. In order to understand fully the nature of these reactions, the student should first acquaint himself with the character of the different parts of the flame, and the use of the blow-pipe in producing the- reducing and oxidizing flames. §291. The flame of the candle, or of the gas-jet, burning under ordinary circum- stances, consists of three distinct parts: a dark nucleus or zone in the centre, •surrounding the wick, consisting of unburnt gas — a luminous cone surrounding this nucleus, consisting of the gases in a state of incomplete combustion. Ex- terior to this is a thin, non-luminous envelope, where, with a full supply of oxygen, complete combustion is taking place: here we find the hottest part of the flame. The non-luminous or outer part is called the oxidizing flame; the luminous part, consisting of carbon and unconsumed hydrocarbons, is called the reducing flame. §292. The flame produced by the blow-pipe (or Bunsen burner) is divided into two parts: the oxidizing flame, where there is an excess of oxygen, correspond- ing to the outer zone of the candle-flame; and the reducing flame, where there is an excess of carbon, corresponding to the inner zone of the candle-flame. Upon the student's skill in producing these flames depend very largely the results in the use of the blow-pipe. In order to produce a good oxidizing flame, the jet of the blow-pipe is placed .just within the flame, and a moderate blast applied — the air being thoroughly mixed with the gas, the inner blue flame, corresponding to the exterior part 364 PRELIMINARY EXAMINATIONS OF SOLIDS. §293. of the candle-flame, is produced: the hottest and most effective part is just before the apex of the blue cone, where combustion is most complete. The reducing flame is produced by placing the blow-pipe just at the edge of the flame, a little above the slit, and directing the blast of air a little higher than for the oxidizing flame. The flame assumes the shape of a luminous cone, surrounded by a pale-blue mantle; the most active part of the flame is some- what beyond the apex of the luminous cone. §293. The blast with the blow-pipe is not produced by the lungs, but by the action of the muscles of the cheek alone. In order to obtain a better knowledge of the management of the flame, and to practise in producing a good reducing flame, it is well to fuse a small grain of metallic tin upon charcoal, and raising- to a high heat endeavor to prevent its oxidation, and keep its surface bright; or better, perhaps, to dissolve a speck of manganese dioxide in the borax bead on platinum wire — the bead becoming amethyst-red in the outer flame and colorless in the reducing flame. The beginner should work only with sub- stances of a known composition, and not attempt the analysis of unknown complex substances, until he has made himself perfectly familiar with the reactions of at least the more frequently occurring elements. The amount of substance taken for analysis should not be too large; a quantity of about the bulk of a mustard-seed being, in most cases, quite- sufficient. The physical properties of the substance under examination are to be first noted; such as color, structure, odor, lustre, density, etc. Heat in Glass Tube Closed at One End. §294. The substance, in fragments or in the form of a powder, is introduced into a small glass tube, sealed at one end, or into a small matrass, and heat applied gently, gradually raising it to redness, if necessary with the aid of the blow-pipe. When the substance is in the form of a powder it is more easily introduced into the tube by placing the powder in a narrow strip of paper, folded lengthwise in the shape of a trough; the paper is now inserted into the- tube held horizontally, the whole brought to a vertical position, and the paper withdrawn; in this way the powder is all deposited at the bottom of the tube. By this treatment in the glass tube we are first to notice whether the sub- stance undergoes a change, and whether this change occurs with or without decomposition. The sublimates, which may be formed in the upper part of the tube, are especially to be noted. Escaping gases or vapors should be tested as to their alkalinity or acidity, by small strips of moist red and blue litmus paper inserted in the neck of the tube. Heat in Glass Tube Open at Both Ends. §295. The substance is inserted into a glass tube from two to three inches long, about one inch from the end, at which point a bend is sometimes made; heat is applied gently at first, the force of the air-current passing through the tube being regulated by inclining the tube at different angles. Many sub- stances undergoing no change in the closed tube absorb oxygen and yield volatile acids or metallic oxides. As in the previous case, the nature of the sublimate and the odor of the escaping gas are particularly to be noted. The reactions of sulphur, arsenic, antimony and selenium are very characteristic; these metals, if present, are generally easily detected in this way (§69, 7). Heat in Blow-pipe Flame on Charcoal. §296. For this test, a well-burned piece of charcoal is selected, and a small cavity made in that side of the coal showing the annular rings; a small frag- ment of the substance is placed in the cavity, and, if the substance be a powder, it may be moistened with a drop of water. The coal is held horizont- ally, and the flame made to play upon the assay at an angle of about twenty- five degrees. The substance is brought to a moderate heat, and finally to intense ignition. Any escaping gases are to be tested for their odor; the- §300. PRELIMINARY EXAMINATIONS OF SOLIDS. 365 change of color which the substance undergoes, and the nature and color of the coating which may form near the assay, are also to be carefully noted, tiome substances, as lead, may be detected at once by the nature of the coating. Ignition of the Substance previously Moistened with a Drop of Cobalt Nitrate. §297. This test may be effected either by heating on charcoal, in the loop of platinum wire, or in the platinum-pointed forceps. A portion of the substance is moistened with a drop of the reagent, and exposed to the action of the outer flame. When the substance is in fragments, and porous enough to absorb the cobalt solution, it may be held in the platinum-pointed forceps and ignited. The color is to be noted after fusion. This test is rather limited; aluminum, zinc and magnesium giving the most characteristic reactions. Fusion with Sodium Carbonate on Charcoal. §298. The powdered substance to be tested is mixed with sodium carbonate, moistened and placed in the cavity of the coal. Some substances form, with sodium carbonate at a high heat, fusible compounds; others infusible. Many bodies, as silicates, require fusion with alkali carbonate before they can be tested in the wet way. Many metallic oxides are reduced to metal, forming globules, which may be easily detected. When this test is applied for the detection of sulphates and sulphides, the flame of the alcohol lamp is to be substituted for that of the gas-flame, as the latter generally contains sulphur compounds. Examination of the Color which may be imparted to the Outer Flame. §299. In this way many substances may be definitely detected. The test may be applied either on charcoal or on the loop of platinum wire, preferably in the latter way. When the substance will admit a small fragment is placed in the loop of the platinum wire, or held in the platinum-pointed forceps, and the point of the blue flame directed upon it. If the substance is in a powder it may be made into a paste with a drop of water, and placed in the cavity of the charcoal, the flame being directed horizontally across the coal. The color which the substance imparts to the outer flame in either case is noted. In most cases the flame of the Bunsen burner alone will suffice; the substance being heated in the loop of platinum wire, which, in all cases, should be first dipped in hydrochloric acid and ignited, in order to secure against the presence of foreign substances. Those salts which are more volatile at the temperature of the flame, as a rule give the most intense coloration. When two or more substances are found together it is sometimes the case that one of them masks the color of all the others; the bright yellow flame of sodium, when present in excess, generally veiling the flame of the other elements. In order to obviate this, colored media, as cobalt-blue glass, indigo solution, etc., are interposed between the flame and the eye of the observer. The appearance of the flame of various bodies, when viewed through these media, enables us often to detect very small quantities of them in the presence of large quantities of other substances. Treatment of the Substance with Borax and Microcosmic Salt. §300. This is best effected in the loop of platinum wire. This is heated and dipped into the borax or microcosmic salt and heated to a colorless bead; a small quantity of the substance under examination is now brought in contact with the hot bead, and heated, in both the oxidizing and reducing flames. Any reaction which takes place during the heating must be noticed; most of the metallic oxides are dissolved in the bead, and form a colored glass, the color of which is to be observed, both while hot and cold. The color of the bead varies in intensity, according to the amount of the substance used; a very 366 CONVERSION OF SOLIDS INTO LIQUIDS. §301. small quantity will, in most cases, suffice. Certain bodies, as the alkaline earths, dissolve in borax, forming beads which, up to a certain degree of satura- tion, are clear. When these beads are brought into the reducing flame, and an intermittent blast used, they become opaque. This operation is called flaming. As reducing agents, certain metals are employed in the bead of borax or microcosmic salt. For this purpose tin is generally chosen, lead and silver being taken in some cases. These metals cannot be used in the loop of plat- inum wire, as they will alloy the platinum. The beads are first formed in the loop of wire; then, while hot, shaken off into a porcelain dish, several being so obtained. A number of these are now taken on charcoal and fused into a large bead, which is charged with the substance to be tested, and then with the tin or other metal. For this purpose tin foil (or lead foil) is previously cut in strips half an inch wide, and the strips rolled into rods. The end of the rod is touched to the hot bead to obtain as much of the metal as required. Lead may be added as precipitated lead (" proof -lead "), and silver as precipitated silver. By aid of tin in the bead, cuprous oxide, ferrous oxide and metallic antimony are obtained and other reductions effected, as directed in §77, 7, and elsewhere. CONVEKSION OF SOLIDS INTO LIQUIDS. §301. Before the fluid reagents can be applied, solids must be reduced to liquids. To obtain a complete solution, the following steps must be observed: First. The solid, reduced to a fine powder, is boiled in ten times its quantity of water. Should a residue remain, it is allowed to subside, and the clear liquid poured off or separated by filtration. A drop or two evaporated on glass, or clean and bright platinum foil, will give a residue, if any portion has dis- solved. If a solution is obtained, the residue, if any, is exhausted, and well washed with hot water. Second. The residue, insoluble in water; is digested some time with hot hydrochloric acid. (Observe §305.) The solid, if any remain, is separated by filtration and washed, first with a little of this acid, then with water. The solution, with the washings, is reserved. Third. The well-washed residue is next digested with hot nitric acid. Observe if there are vapors of nitrogen oxides, indicating that a metal or other body is being oxidized. Observe if sulphur separates. If any residue remains it is separated by filtration and washing, first with a little acid, then with water, and the solution reserved. Sometimes it does not matter which acid is used first. But if a first-group base be present, H1T0 3 should be added first, for HC1 would form an insoluble chloride. If the substance contain tin (especially an alloy of tin) HN0 3 would form insoluble metastannic acid, H 10 Sn s O ls , in which case HC1 should be used first. Fourth. Should a residue remain it is to be digested with nitrohydrochloric acid, as directed for the other solvents. The acid solutions are to be evaporated nearly to dryness, and then redis- solved in water, acidulating, if necessary, to keep the substance in solution. Fifth. Should the substance under examination prove insoluble in acids, it is likely to be either a sulphate (of barium, strontium or lead); a chloride, or bromide, of silver or lead; a silicate or fluoride — perhaps decomposed by sul- phuric acid — and it must be fused Kith a fixed alkali carbonate, when the con- stituents are transposed in such manner as to render them soluble. The water solution of the fused mass will be found to contain the acid; the residue, insoluble in water, the metal, now soluble in hydrochloric or nitric acids (compare §266, 7). If more than one solution is obtained, by the several trials with solvents, the material contains more than one compound, and the solutions, as sepa- rated by filtration,^ should be preserved separately, as above directed, and analyzed separately! The separate results, in many cases, indicate tlie original combination of each metal. §303. TREATMENT OF A METAL OB AN ALLOT. 367 CONVEESION OF SOLUTIONS INTO SOLIDS. §302. Before solids in solution can be subjected to preliminary examination, either for metals or for acids, they must be obtained in the solid state. This is done by evaporation. TEEATMENT OP A METAL OE AN ALLOY.* §303. On account of the different effect that nitric acid has upon the un- combined metals, it is used as a solvent in their detection. Thus: Gold and platinum are not attacked by nitric acid. Tin and antimony are oxidized and converted into compounds that are insolu- ble both in water and an excess of the acid. 6Sb + lOHlTOs = 3Sb 2 5 + 10NO + 5H 2 15Sn + 20HNO 3 + 5H 2 = 3H 10 Sn 6 O lB + 202TO All the other metals are oxidized and converted into compounds that dissolve either in water or an excess of the acid; e. g.: 3Pb + 8HITO3 = 3Pb(M-O a ) 2 + 2NO + 4H 2 Bi + 4HNO3 = Bi(IT0 3 )3 + NO + 2H 2 Method of Procedure, f Place a small quantity of the metal or alloy, about equal in bulk to a pea, having 1 previously obtained it in as finely divided a state as possible, in an evaporating-dish, or any suitable vessel, cover well with nitric acid, sp. gr. 1.20, and apply heat. Continue the application of heat, replacing from time to time the acid lost by evaporation, until the metal or alloy is dissolved or wholly disintegrated. If complete solution takes place immediately, pass on to A. If a residue remains, decant the liquid portion upon a filter; again add nitric acid to the residue, heat, and again decant upon the same filter. Then thor- oughly wash with hot water, either by boiling with water and decanting, or by transferring the whole to and pouring hot water through the filter. Add the first portions of the hot-water filtrate to the nitric acid filtrate already obtained, and treat the mixture as directed in A, after having first evaporated a drop or two on platinum foil, to ascertain whether anything has really been dissolved. Treat the residue as directed in B. A. — The Nitric Acid Solution. This solution may contain any of the metals, except those mentioned under B. If the nitric acid has effected a whole or partial solution of the original metal or alloy, evaporate almost to dryness to remove excess of acid, add about ten times its bulk of water, and proceed with the separation and detection of the metals in the regular way. Should the concentrated liquid become turbid when diluted with water, the presence of bismuth is indicated. In this case enough acid must be added to ■clear up the solution.J *This section is furnished by Dr. J. W. Baird, Dean of the Massachusetts College of Pharmacy. t When gold or platinum constitutes more than one-quarter of the alloy, nitric acid fails to extract the whole of the base metals that otherwise are readily soluble. In such a case the amount of gold or platinum must be reduced to at least 25 per cent, by fusing the alloy with the requisite amount of that base metal whose absence is surely known. t Arsenic, if present in the original alloy, now exists in the form of arsenic acid, the precipi- tation of which requires heat and long-continued passage of H 3 S (§69, 6' e 7). 368 SEPARATION OF ACIDS FROM BASES. §303, B. B. — The Residue Insoluble in Nitric Acid. This may contain gold and platinum in their metallic forms, and tin * and antimony * in the form of metastannic and antimonic acids. The separation of the two former from the two latter depends upon the fact that the meta- stannic and antimonic acids are soluble in hydrochloric acid, forming SnCl 4 and SbCl 5 . Digest, therefore, the well-washed residue in concentrated hydrochloric acid at a boiling temperature for from 5 to 10 minutes; then add at once an equal volume of water (to dissolve the stannic chloride), and bring to the boiling point. If gold or platinum existed in the original metal or alloy it will now be found in the form of a dark-brown or black powder or mass, insoluble in the hydrochloric acid. If such a residue exists, decant while hot, again add hydro- chloric acid, heat, and again decant. The Hydrochloric Acid Solution. This solution may have a turbid appearance, especially when cold, due to the action of the water upon the SbCl, ; but without filtering proceed with the separation and detection of the tin and antimony by the usual process.f The Bark-colored Residue. Add, after washing, two volumes of hydrochloric and one of nitric acid; evaporate almost or quite to dryness, dissolve in a small quantity of water (to obtain a concentrated solution), and divide into two portions. The gold and platinum have been dissolved by the aqua-regia formed, and now exist as auric and platinic chlorides. First Portion — Test for Gold. Dilute with at least ten times its bulk of water; add a drop or two of a mix- ture of stannous and stannic chlorides; a purple or brownish-red precipitate (or coloration), purple of Cassius, constitutes the test for gold. A convenient way of preparing this mixture of stannous and stannic chlorides is to (a) Add a few drops of chlorine-water to a solution of stannous chloride; or (6) Add to a small quantity of stannous chloride enough ferric chloride to produce a faint coloration. Second Portion — Test for Platinum. Add, without dilution, an equal volume of a strong solution of ammonium chloride. The formation, either at first or on standing, of a lemon-yellow crystalline precipitate, consisting of the double chloride of platinum and ammonium, (NH 4 Cl) 2 PtCl 4 , constitutes the test for platinum. Addition of alcohol favors the precipitation. If the proportion of platinum is very small, the mixture, after ammonium chloride has been added, should be evaporated to dryness on a water-bath and the residue treated with dilute alcohol. The ammonium platinic chloride remains behind as a yellow crystalline powder. SEPAKATION OF THE ACIDS FEOM THE BASES. §304. The preliminary examination of the solid material in the dry way will give indications drawing attention to certain acids. Solutions can be evapo- rated to obtain a residue for this examination. Thus, detonation (not the * Traces may sometimes be dissolved. t Arsenic must be looked for in this as well as In the nitric acid solution. For when the alioy contains arsenic, part of it will combine with the antimony and tin, and be held in the residue. §309. SEPARATION OF ACIDS FROM BASES. 369 decrepitation caused by water in crystals) indicates chlorates, nitrates, bro- mates, iodates. Explosion or deflagration will occur if these, or other oxygen- furnishing salts — as permanganates, chromates — are in mixture with easily combustible matter (§273, 7). Hypophosphites, heated alone, deflagrate in- tensely. A brownish-yellow vapor indicates nitrates or nitrites (§241, 7) ; a green flame, borates (§221, 7). The odor of burning sulphur: sulphides, sulphites, thiosulphates, or free sulphur. The separation of carbon black: an organic acid. The formation of a silver stain: a sulphur compound (§266, 7). §305. When dissolving a solid by acids for work in the wet way, indications of the more volatile acids will be obtained: Sudden effervescence: a carbonate (oxalate or cyanate, §228, 6). Greenish-yellow vapors: a chlorate (§272). Brownish-yellow, chlornitrous vapors on addition of hydrochloric acid: a nitrate. The characteristic odors: salts of hydrosulphuric acid, sulphurous acid, hydro- bromic acid, hydriodic acid, hydrocyanic acid, acetic acid. The separation of sulphur: a higher sulphide, etc. It will be remembered that chlorine results from action of manganese dioxide, and numerous oxidizing agents, upon hydrochloric acid. §306. If the material is in solution, the bases will be first determined. (Certain volatile acids will be detected in the first-group acidulation — by indica- tions mentioned in the preceding paragraph.) Now, it should first .be con- sidered, what acids can be present in solution with the bases found? Thus, if barium be among the bases, we need not look for sulphuric acid, nor, in a solution not acid, for phosphoric acid. §307. As a general rule, the non-alkali metals must be removed from a solution before testing it for acids, unless it can be clearly seen that they will not interfere with the tests to be made. Metals need to be removed: because, firstly, in the testing for acids by precipi- tation, a precipitate may be obtained from the action of the reagent on the base of the solution tested, thus: if the solution contain silver, we cannot test it for sulphuric acid by use of barium chloride (and we are restricted to use of barium nitrate). And, secondly, in testing for acids by transposition with a stronger acid — the preliminary examination for acids — certain bases do not permit transposition. Thus, chlorides, etc., of lead, silver, mercury, tin and antimony, and sulphide of arsenic, are not transposed by sulphuric acid, or not promptly. §308. If neither arsenic nor antimony is among the bases, they may all be removed by boiling with slight excess of sodium or potassium carbonate, and filtering. Arsenic and antimony, and all other bases of the second group, may he removed by warming with hydrosulphuric acid, and filtering. When the bases are removed by sodium or potassium carbonate, the filtrate must be exactly neutralized by nitric acid, with the expulsion of all carbonic acid by boil- ing. Then, for nitric acid, the original substance may be tested. §309. The separation of phosphoric acid from bases is a part of the work of the third group of metals, and is explained in §§152 and 153. For removal of boric acid, see §221; oxalic acid, §151; and silicic acid, §249, 6 and 8. The non-volatile cyanogen acids can be separated from bases by digesting with potassium or sodium hydroxide (not too strong, §§231 and 232), adding potassium or sodium carbonate and digesting, and then filtering. The residue is examined for bases, by the usual systematic process. The solution will contain the alkali salts of the cyanogen acids, and may contain metals whose hydroxides or carbonates are soluble in fixed alkali hydroxides. 370 PRELIMINARY EXAMINATION OF 80LIDS. §310. s i-SS T3 a ca a P o P< a o a ca 60 h O TO g IS a o d 60 p .5 * 1.3 •H o teg s p O p" o IS o § S 3 e a. v s s § * P § 8 a ca ■» m g » a >P o 5 ^ S3 I 5 & 3 J «* "I "3 .>? in 8 *> a e R S P fJ H TO 03 a 10 o "3 "O P R p ca °* ca .» 60 ° O A J? 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O S 1 d cu cu 0) -fl ft cS d j" co d cu - .2 ft *co " CO m CS p &0 „ d « -s y—N -H 3 +J -+J I* C3 JS cu .fl ' ~P 2 -a ^ -p • n a o CO a v * -3 o" ^ag .2 " a •s ° ^ 0) CO (U "* a ^ cs a XI a p o ft a o o -a d cs a <{ cs ca co g co O - I 8 p 13 13 P "3 ft - ?, cs co CO V o * ft 1-1 cu , jd "P - ■a d "S p O a ft co L, B s ft o CO d o £ ■S 1'8 w P cos cy O CO o CO CO y. cs d w a .-s d S d D 3 3 )■§ is a a R P CS p OJ d » CO CO ft 60 d O O W . m cu » a, fl • - ? t- cu fl Sao g O 'cu 2 M (H £ d rrt tl « h S a ,y to w n 1 o d ej jj O -P P W « IB ,rt B ,. d S J*5ft, a 2"^ '" E S ° S ■2 • S § p,» ■P Si S S o f< 374^ .SUBSTANCES BEFORE THE BLOW-PIPE. §311. < En <1 o I— I O o o « H Eh i— i tS PH I— I Ph I o O Ph PQ xfl o m PP l=> Pq O o O w pq Co s CO O ^3 • i-i o s » b ;£ s ^ s'f S ^ » §■ =s rf? .« ■2-. (=1 , o » ■- 2 s s ai ■S ^ JS ■ ■ I 8 M *** M CO ^§ J §^^^ s § ° -■ r» § «> & S -tf 1 B L, <» .^ P «£..§ ^ H « g ^ m ■8.S : § "3 S St sw o i£ 0) «j t » -t-a o § ^h o e to co H o a t ■8 O C CO d as 3 5SI e c B IV O «s a "13 d •i-T oo o d D d 00 W so s.s CO S w a w at in if Eh to £3 a! tsi t3 ^> o to •g H . . rt rj" S IB EH 1 •rH EH 5S d d W bo +rfc ^ xn< o rfi o o is; oi isi oi »rfS -»] £ ^5 ■rT m 13 Fl > . o rf •S 60 f - °° • a o is? f3 W J3 «• »< 0) O -d Q aJ go p g ft §318. GROUPING OF THE METALS. 375 CO am rfS o 3 ftS as bo Ph P e u eo 003 © 3 a ° * 2 a a 3 a IS Pi Ik- a« a a s ri a a « £ " a " © © ^ h bo T3 »0 08 ^* d ® - a" g s ^ a * 2 * S , A •a ■a < S si •a a s e ■8 a e ft -8 S O © fl ,•■9 § S ift§3 ?Sl1 • s »? .|s ft .a 0D ^j * • • li «D8 o*2 •S Sao "Bi -a a "o a 3« £?|2 a E»^ a* .S*« o«! 5B< oS «** ■ B ^ j tf : 3 ■ a 2"L 2» a 85 a a- fta Sxi ai .gpa &b io § « *• ^ i : oo : ■ : to • ft- S 3 B 3 o SJ -r-t S ^ 5 a .33. cm oo S • ss ** . ou ft ^ ^ oo § :• 2 tf B • g .5 ^ s ? - I § §> -6% TABLE FOR TEE SEPARATION OF THE METALS. §313. §313. Table foe Eeview Pb Hg' Ag As"' As 7 Sbv Sb'" Sniv Sn" Ail"' p t lv Mo" Hg" Pb Bi Cu' Cu' Cd ~ Al Cr'" Fe" Fe'" Co Ni Mn" Mn"i Zn Ba Sr Ca Mg K Na NH 4 PbCI 2 HgCI AgCI PbCI 2 HgCI f H 2 S0 4 = PbS0 4 White. H 2 S = PbS Black. K 2 Cr0 4 = PbCr0 4 Yellow. Kl =Pbl 2 Yellow. lS f NH-HgCI I Hg Black. \ AgCI j S |(NH,) a (AgCI) 2 }Add HN0 3 JAgCI White. o As 2 S, As 2 S„ Sb 2 S G Sb 2 S 3 SnS 2 SnS Au 2 S 3 PtS 2 MoS 3 HgS PbS BUS, Cu 2 S CuS CdS 3 3. CO a 3 U c m S a jj o a 1'° 5 ? ■g 3 p. 3 cc a 3 '3 o a a (NH„) 4 As 2 S, (NH 4 ) 3 AsS 4 (NH 4 ) 3 SbS 4 (NH 4 ) 2 SnS 3 Solution. Solution. (NH 4 ) 2 MoS., As 2 S 3 As 2 S B Sb 2 S B SnS. Au 2 S 3 PtS 2 MoS 3 .lo H 3 As0 4 SbCU SnCI 4 AuCI, PtCI 4 MoCU J3 CO is l| II u 0- +-P HgS PbS DI0S3 Cu 2 S CuS CdS { HgS Dissolve in nitrohydrochli CD 3 be— J 3* I Pb(N0 3 ) 2 Bi(N0 3 l 3 Cu(N0 3 ) 2 \ Cd(N0 3 ) 2 J Sso , c mo oS 1 "?, *Kx ■^•Occ ■OS'S a s a x s ro z ° •a Or'o >, ft 3 E ■§ T! £2! ■a CoS NiS MnS ZnS -P+3 CoS NiS MnCL ZnCI 2 iti §313. table for the separation of the metals. Separations of the Metals. 377 |H, andSb at J? oh \ i , 513 8 H 3 As0 3 \ Remove AgN0 3 with CaCI 2 and add H 2 S-Us 2 S 3 Lemon yellow. Tj-gg o.er as SbAg 3 J- Dissolve in hot HCI, dilute, filter and add H 2 S^Sb 2 S 3 Orange. / an 6'Pi'i ■a /SnCU } Test with.HgCI 2 . -j HgCI, White ; or Hg Gray. Sb Tsia'T SbCU "I ^ j S bC I E reject or test in Marsh apparatus. Au i-ISe-i AuCI 3 l^loiAu ) .DisfoivMnnitrohydro-^ AuCI 3 .NH 4 CI _ Evaporate Pt £3 fig Hi AuCU -PJ A » IdSK II L Ptcu j«s"-|pt Inh 4 ci alcohol. is with excess of HN0 3 . Dissolve. ...,„. , iOE?° k j of HCI - Test tDls solution with Na 2 HP0 4 ^ 'Ammonium phosphomolybdate, Yellow For traces add HC 2 H 3 2 and test with K 4 Fe(CN) 6 -J Cu 2 Fe(CN) s Red-brown. Add KCN till blue color disappears, then H 2 S-j CdS Lemon-yellow. acid, evaporate to -j an< J ignite to All", Yellow ■ with exceBB of and digest with j (NH 4 ) 2 PtCI 6 Ignite to Pt°, alcohol. I Gray, togreen- Evaporate to dryness with excess of HN0 3 . Dissolve residue in NH 4 0H and add to an excess ttorblP" 1 ' ntion land test with SnCI. and Cu wire, nation of Pbl 2 or PbCrO ,. 1), AddhotK a Sn0 2 [Bi° Black. I|,.2NH 4 OH.2NH 4 N0 3 ] Deep blue solution evidence of copper, i,2NrL0H.2NH 4 N0 3 HCI and precipitate with (NH 4 ) 2 C0 3 j AI(0H) 3 White, gelatinous. , . Ucidify with HC 2 H 3 2 and add Pb(C 2 H 3 2 ) 2 -j PbCr0 4 Lemon-yellow. [Dissolve in HCI and add KCNS j Fe(CNSi 3 Blood red. I Test original solution (acid) with KCNS for Fe'" and with Kjfe.'CN). for Fe"-j Fe 3 [Fe CN)„] 2 Blue. [a. Test with borax bead. Blue bead. "1 Lb. Add NaHC0 3 and H 2 2 , Green solution. J r J. Test with borax bead. Brown bead. -. lb. Heat with I N .. nll . /add Kl. ^ Brand NaOH I N| < 0H >3 \ Free I in CS 2 J M Boil with Pb0 2 and HN0 3 {■ HMn0 4 Purple. hjft precipitate. } Test with torax bead " Ml- lAddNH.,0H , Solution. ( filter and add H 2 S < NiS, Black. l,jAddH 2 s[ZnS White. [Dissolve in HCI and add H 2 S0 4 -j BaS0 4 White. °. o 3 * SrCO. CaC0 3 Sr(C 2 H 3 2 > 2 Ca(C 2 H 3 2 ) 2 ,J« g j, .. 1. Add CaSO,, set aside ten minutes -j SrS0 4t White. J a a Moisten SrS0 4 with HCI and apply flame test. ■SoS V £*> 2. Add K 2 S0 4 , boil, set aside ten minutes. 5 * F *{JhVc o add \ CaC 2 4 , White, solubl soluble in HCI. !A, White. 'ipply flame test using cobalt glass. "Violet. -After removal of Mg apply flame test, yellow. Ho the original solution add KOH in strong excess, warm (note odor; and rest with moist litmus Paper; pass gas into Nessler's reagent j NHg 2 l, Brown. S78 A.CID8. FIRST TABLE. 5314. co a M < o ft 15 o £ T3 s H -S W a, I * H W CO 000 be » d » • rH cd as pq Pi O CO © a) a) 8^> o « a) i fq V 2 o o & g g § o 03 -4-a -*> rd •i-H M -, « nd o s * ^H FH ^ ^ ^« 0> H-3 CS f-i (H CD d g O pi d o -^ i — I ° ^ -d to pi cS 02 H 03 J^J CQ , — [ O OS CO P Pi o 1 fe fn d S '-£ d '&- ■2 «H ^ Ph o o 02 O o cp rd rO o !>^ i — I CO Kj CO _i ™ .3 U ■-H C3 rQ 03 d ■ — ' CO c3 rd FH O c3 03 03 -F-H cS 'r-H *^ 03 •H ti j h O pq d o "-+3 o5 ■^H rH m p4h O "^3 Ph 'tO ^ g ■rH O S3 03 *> d 3 • iH © T3 e_. •|S6 S o -S e «h •*■ ■+~. d ^ O o 03 . ••H -|J o d 0) » °° d ° 2 d S tap Fh o 03 ■rt o m Ph o 03 O rQ 03 -, Fh o '3 2 ^ rd 03 Ph -h; P Fh 02 H^> p-s 03 rQ 03 O 03 a ^ 03 Fh ^ rS •rH ^ 03 -rH 03 O 03 O • rH O d -c rd X Pi P d rcj 10 d rd 03 S-.s _H K O © -5 ^ O 03 O o CB r,-j !* ^ r^ 03 fO -+J 03 CO CO 03 O M ft j „d 02 P d d ■r-H r^ ■HH 02 03 03 Ph O o3 d IO rd '"' n 030 02 • — ' d o co d S o 03 .23 CO g w ^3 1— 1 t3 0) —H l *- a If 5-_2 t-~ 03 ; ° - M CO "c. « M S •3 -§ CO Q^ 03 « .5 § b0 03 d "o rQ > 03 CO^ CD 03 rS HJ 1^3 •Th O Pj ■cTg Ph 3 2 Fh P 03 d o rH » ^d 03 •rH 03 exj .y 003 a ^ d ■+» to 02 CB n;) FH rH H-^ 9 S S 03 O CJ Ph g a w 03 rd O •« 'd be rd Fh cp O 0) 03 rd -'H ni -, 03 rP Fh CO I B. o d * d r^ S o 03 CO ' ^ a© coo 03 Fh ■d d 03 co CO 03 03 08 H ^H tn ^H _03 a? ^ rfl J3 H3 * a fl ofl' , •rj © <0 Isg h ? fl fl S ■" ^i m ^ .2 «*o ° « "3 ago ° O «*H fl « fl O ° C3 f rM rd J 8-9 3 e © m ■o S 03 Q. ' 5 fl 2 5 fl fl ^ . to CJ u fl 03 C3 s a a? g336 ■a +■ e 2 O •" K B^rflS _- S o -^ fc a fl ^ 3* sag* 1 S3 « w "2 p 03.5 ^S|fl a » 5 * cS o J3 ft Si CO §314. ACIDS. FIRST TABLE. 37!). p§i s++" * £ ^ 8 „r » ~ ^sifi 2 pi ^ fH S.I'S-al'lll l§§ g i 8 ^|^g | s SO b d ~r a; K „, rtrnSai 1 -' o>io^ &o r c*s iHEO«i-HCSrE?^3 03 V rr| iH "H r^ i^ 2? „ Ej n cd s, s8 ns isi** *5-i! lis '-*> -5 " & fc ug-2 aT fa 5s S3 r^j rt ^ '3 * ^ BPh» 'S ° m ^ 2 " ^ §5 d "-SrS^OJ 3 S » S ° fluj"^ ^ «r*tj5o'^ S Pi "*" ^ ^ M S g bP S, & T3 § W - 1 -* ^aS°§-ti§1 S2«S: £ «, b g § § ^s^g •S«'3'5 , h2«'S S S S 2 e s-i^^fc!",^ S S ^ r i ^3 ' ^ K^ B rH • i-H 02 ^ O CT" °3 ■ r-H c3 ^ 2 hHO pi ■2 1^j£ "sS^S ;^-JspI l^l^l - b ^ S ^guTSS*-* «r S a Ji ei ! £ ° J* ^ § * ■-£■§§ «» O £j C! _H eOO.rt d rrtSB^BSR _d^o' rt rf ^BO ^5. a +|o:^| fill §<■? ^j^-if ^- ■ . P-I o ** - - - - . -^ ■ . - p 3 - I TS H tfi rA « v S > S. g. o"t3 a b m H o ^ (2 3 >» -9 « ^ 380 ACIDS. FIRST TABLE. 5314. 13 PI 03 03 1=1 O -8 e3 CJ 03 CJ bo =1 ■ .—1 13 13 03 b£ PI •rH O 03 13 CJ Pi -ij oj bo U, •9 'd OJ 02 r<= „, -a. s?» Pi 13 d * 03 £ s '3 03 a 0J 13 •i— I d 03 o 03 53l >* 13 bo d o d bD o d 03 cj CJ rd H bp oi d » Bh eS - 1 - 3 is 03 » rB ^» 13 cj o eo CO est o d o o OJ is CQ OJ -p> 03 d 03 ^1 O Ph OJ i3 13 •rH CJ 03 o ■s a rd ft a ^ I— I .£ s — 13 ■rH ^2 d S -s . o rd bo co 09 OJ o 0J S d o •rH t> OJ ■s OJ OJ 13 o SB M O d CQ CJ rd ' s ■■£ •13. r* 2* I * *■§ oj o ££ CJ "B ° d h OJ rr-j co fcC Oj £» rH rd Ph^ oj d J£ oj ^ bo CQ 03 -H OJ d h cj 03° "S bp £ d oj ■^ -*p CJ OJ ii-S * d ■H g I— I OJ 'S rO .a Ho 1=1 r§ rB <=> CO r§ ■g.s © d © o " rH J» +3 is d 02 « pj O o o cd '^ .—i •rH O COS o 03 CJ 13 oo d r- rd 3 ft ** 'd « d COS 02 v ~^t3 ±i oj o o i—i *- •a ^ s "3 03 ^ • c3 i — i CO pf pi O O £ > OJ d f^ 08 13 g S OJ OJ rH I * r<2 OJ 0J '^ re! OJ *d r* OQ g 03 OJ ,5 J 4^ 03 £ £ oj d d > ° •H J COD » Si s r« d 03 H ft ■ I— i 13 O 3 3 IS 13 rd *- *- &3 s 03 r° rQ «H +i -P -P + 1 -P tfJ CQ CO 02 02 Oj CJ OJ OJ OJ E-i cH H EH H OJ ■J=l OJ > rH OJ CQ rO o CJ •rH rH d rd CQ 13 03 fH d OJ CJ d o o rd CO •rH 13 --> rrH ^ s °* rS «» a ft rH CO 03 " O P 03 rS •rH CO ^c3 oj r» r— < O O CD ^ OJ « rd O ■^ d CQ O 03 O Pn=rH d ^ CO c« „ coo rrj s cs O OJ »fH 1/2 13 d 03 ^ 03 bo . 03 Ph d W ° • rH CJ 03 Qi r* d ^ -3 n- cj O -r" r»> d o n 03 " ."B 'K Ph rd OJ ft S - Ph CQ ^j ^ ■"£ •tn t« ^ 13 d o g =H 03 CQ r""a -rH OJ CQ CJ Ph -th o. CJ w 03 d CO 03 rH 13 •rH c3 §314. W Ti ACIDS. FIRST TABLE. 381 0) | S^sgS 3 j} »- s s * -g ~ JTg | ^ 8 » -3 9 C of E ^ *- -ti ft£< g :s S ° ^ d g « ^ s -b =H « * ^s« ggs « -g a „ ^ -S ■« * &■§■§ fe>£ §a s 8-i * * 5.1 a fc. SB 01,3??,-" — o o rd ^H oa 1 ' £DW3 ^ A rdEI tJ 10 CO S — I l> li m •£ ft qjf5« 'SOo ^ S -j3 o S^W ftuO eS Sv hn ^ 9,2^ fl ft'S • "3 p S -^ 'o « Sl^^fl Ifs S =3 f^ : . 1 2^1 d" s a s 1 -s -« e J -s H - -s ^■s inn f^t st^ii^s isls^i. ■sS l^' s l- ^H 2 « ^ IU -s 11 Jill 1 5 iwi^! c//j o O O) cd Ul ■j3 ca ^ rO CO cd rd •rH I 1 o ^J cd ti> cd pq a o d s. , , Si cd S5h«(H "3 ° 00 CO 2 s '« S CO s <» Ph 0) © Ph rd i ft OQ S ooo P ' o rH rd P en r=! Ph i-H P co 02 rH. _£ ^6 p ^ O rd rd Ph rP d ti h cp *$ r2 "g rQ ^ o .2 CO £ a s ■rH tH o H a o> o +? ■■£ 2 ca ^ 6 Pi O CO *H rf rd r c3 CO J3 S 00 ' .3 O rrH ■- CD U3 bOrH J ■" -3 ^ « <° co e* CD "* cd .-« .P «o te ft - P & ca CD coo CD r—H 0O0 rH rQ -^ !» r3 5R d> CD S s § Pi • ti oi cd ca +; cj cd ca 3 rP rP rt ft * 1 S "3 CO rd "3 CO £ be P CO ^ o ca «a ft ^i CD rR ^ E-l 03 ca CD S to ^3 £ ■g =a • I— I b0 0) 00 c» CO e« * oca IS 00 CO CM GOQ !> i — i CO CO CO CD £ o 0> rH V .2 M O rH as .go o W a d a -a .a o ca 5 _2 o •t-t Sft -S ft d rQ •3 ° ca h o o •rH ft f-* ca .-S > ^rd CD 53 H-i -rH r2 ^ d g os .a rS rD ^ w •fH rH o rC, o o !>• CD COS o CD d ca o t3 .d rd o rd CD d o o r- CD o Ph a o o CD P rH o o d o d rQ O C4 COO o bp • S rd CD C8 CD -3 r5 CD P rd o d o p p CD CO O Ph a o o rd CD o Ph W 3 Sr2G <" ftrd M 4- d ° " S °^ is§s m ca ^ f*1 cD ca ca « rH © o w ^ M CO rH rH «r^ ■■s-i --> d M « »S os ~|- CD «»,J-J o" S ?2-l M CO r^ .■S -"S ^ CD ca cd o .^i - C5 Sr2 OS s-. P O .2 ?£ ** |^^§ ^^^^ <« r _- ^P M a rH « 'P S, m 0) Mo „ -^ r2 | ^ CD S CD I P rT^ - cj -p d M g ca w ^ g S II ca eo g ..2 2 P H d w o ca rH o l— H rd O £ §314. A0ID8. FIRST TABLE. 383 CO I ■ p -\- >> + ~ -° & "3 W rd a> ■d o I — I A o d o -2 «° tt - W -^ s CO ft d *rH s + g o o J2 03 S>> . 4- <" r o O » !=* W 2 o - C3 * + 1 U r^J Si coo w 04 to cd O I O rd + -c ° ■* pi CJ H2 ft"* 3 CD, rH ft a o o HI rd a to a) O 60 o 60 + TOO + M to o3 rd ^ 0> H 'S o O S « nq coo .&-r> O 03 CD ft3 03 « Ul-H rS "S-g ■» ~ cd § 6pft d •£ S 5 bo g go o W o p) o -3 w 000 00 CO o ft s o o cu ^3 ft -a o pi pi Pi CO O „ CO 'j o a N - rjj COO ft a fe n .a pq $ a .^ o u m TO ^ tH co 60 W ■S g d ■s «a ra £ ox" h 'O -P «9 d o coo .a pg * coo .2 a> CQ S -a T3 O as •j-i 0) ^ 02 ft pi CO a> a> PI o h o PI PI _o %-3 CS PI •.—I a o o O TO CD ^3 pl rP! 0) Pi TO h" Ol TO & Ti PI ro s ^ a >? » o cJ- S .a "S a a 2 s O H3 00 «r !> 04 coo CD T3 !-< O CD t»>T3 Is a^ CO a o CD Pl CCi PI o ■8 TO CD Pl ■i-H a c -P! TO n3 CO o CD rt TO O " a o Ph rO CD o ^2 O -i-H P! ft CO CD ,P! TO a o u o ft CO PI TO l-l -M .° pi ft — i CD ^d > CO CD CQ o ft a g H3 CD * T3 p J CD 03 b0 g T3 CD Pl 60 co •73 CD ts Sh PI CD CD PI O O o TO T3 « nzi Ifl •rH O CD tn CD .a o t> -+J o p pl T3 CD 13 CD rd +" * .2 CD TO q; ft a CO TO CD > Al p TO ft 9; ■prcc CD CD fteo oo rd ess CD COO O » bo PI •rH CD rP rH r^ ! ft is rH pq T3 PI TO rP TO l>> rQ CO 'd ... T3 _o rC ecj CD T3 ■l—( T3 O w !>>-ri rO CD 13 i® CD £ CO ^ O ft «> O rd ftrd 'd CO CD -3 ^3 r»» rO CD CO O ft CO d TO rH CD t-t C3 CO CD ^3 . o . CO o O CS d f -1 3 CD IS d _o '-+3 o crj CD U CD rH TO IT o O CO CD ^d . O ft co ■3 ca CD rH d -r= .a * a CD Pl CD CO rH ° £ O CU rrj rO 02 d o -s TO O d TO CD d CD a ft d o rH 60 CD 'o > CD CD rH 03 «p) CD CO ^ rS +a CO ° d 60 co ^ V =? 00 o CQ ]R £ « "■h68 r ^ coo coo £ .2 CQ CD O J3 H-» q o >. a C3 a « «! t» 43 ■O g I o a T3 ■a o •l-t 8 o (HI m OS a © m 13 o 384 ACIDS. FIRST TABLE. §314. I*** .-s "Sl^^i lis *A\ S-SJ? ^3 si* 9 ! 111 s s w -^.^-^ o o 9 _S d 2 " ^ °s „, _£ lila g; -g.ei.1i I 1 - 8 ill .^ § «,.s S.iB.S'g, ■" 3 's j^ o Ph ■»-+_! £* 3 cs Ph _r <" _ _2 ^ P< o o fe ±? SS Sua a 3 Sit 1'atsU .SfS ^^ § M P h _H co Eh Eh < EH to §314. ACIDS. FIRST TABLE. 1 o u P (H o r>> oj 03 -2 n" « -2 1 Pi 03 -§ SB Ph rt 13 T3 ? 02 •5 ° Ph U s ■a s o a P •l-H Pi 3 • rH Is Pi O) 03 s ^-. bo^ P j. 03 ^ g rS © o d P o a P 02 -+J 03 gs o3 ^ Ph a O +3 rH >> t u c3 O '43 o3 fl o a> rP S <» 535 rH CJ c3 Pi CO (D rH Ph P •rH P o P r— H o 2 •* § S hC-H Pi ^3 II «H rH o -S 0J -rH p R o .2 ^ p T3 o O 02 o a> * rP O -^ OS o 03 bo Pi i 1 •rH O , n3 0) 60 PI cd rP T3 '3 o3 .2 •1— t Pi O •rH -(J 02 Pi a? CD rH be rH ^ P a> O N O o rS O S i=! OJ .9 • o ^ o pi o 02 cd _r rP 02 ep P! ai fl u o r— 1 o cj Is o 1 — 1 i-H T* CO ■i-H T3 02 a> -f^ ce m PI a> o P o 13 s rH o «4H bo •rH oj rQ 03 CO j — 1 rH '3 ■H-H Ol o SI o* r5 & _0 CQ J °° n3 p a> o be P P o3 p s 03 ^J bo's P to P a>. a> 5g rH o3 bo

S ft bO^ P ^ a) rH o 0J ^» o3 r=! 'p CO 03 «4H g b0 o3 =H CP O 02 ^ s S o3 .P % r4 CO "3 ■i-H if •rH S O P O ■rH H^> 03 rH CO o W 03 8 S 5-c 0) M K ef 03 JS o bo -m HH> 03 |S 13 03 3 m rP -ri >•> oj o3 03 P O T3 CJ CO o Ph o bo P •rH r— t O o 02 rP 'r* o P ?2 CD p bo=H bD-= 9 n3 T3 S ^ s r>, H rQ T3 ■8 & OD OJ ri^ ^ '3 OJ rH U -H P r« rS + 1 H3 -rH OJ fe rQ 02 2 9 03 cd H bO •F-t T3 s s o 0) CD o rH q-l g-g o CJ OJ r-l O CD M o .d ■M bJD a •M el $2 § s - <1 5 s 13 1 tg i T3 O "2 & ,S O ■+3 o° o « *» o n O '& " '8 £ co t5 M< «M 03 ft J3 'M CD M EH ^i ^3 Qo ■as •a > 1^ .H o> „ m J3 « « 8 T u5 -« * S r, 'Si h to U3 '8 0J ^2 N «» 1 § <* C CO c • •» "C :9 ■§ CD o « +3 -a a ^7 cS Ph-^ ■d a -a ^— ^ d ft ^2 r-H O CO S CO o 13 'St a W 3 l'§ CO «?0 oS o 3 .3 .S * g -S o © CO -T B ^ -g o a °3 '3 3 o Pi W a o ■f-i o g I cm O -, CO ■S .2 CO 'C ra o >> Ph 2 o h Ph ■s o ■8 s w IO S S 00 & iH i-H o CO P -^ • r3 HJ '-+J O ■+3 CQ CD o M CD CO S -e O CO 12 p p o p 10 10 S Pi t*n w COT CO g oj 13 'S CO % g «») ^ 02 . n h m 0J o ^ J3 -3 CO o u CW n >-, T3 t' p g « co CO « S3 |A g i Pi ** P. "3 s g ID -rH eO T3 60 fj O T3 ■S ? co 8 •o o « CO .2 is t) o FH ~-< 5^ p,t3 f 9 -(J CD § g g 9 CO gj O. rQ CQ O " .2 ■a CO ■§ 43 C© '3 to S 5 S , CQ CO 5 CO IS Ss 0) g cd 3 a « l£* lis 2 o rS 5 -g Z a S o o S s a lis s a hj 2 -a a ^ co §2 | X3 CD •S a S co co a A? cS ■• co go a 3 5 » .S S ° Si? 6e?» ■ 4+ * §316. ACIDS. THIRD TABLE. 387 © i o eo am am o H < M GO e8 * -"3 n d a D ft- ps Ah ft o ft O 02 Eh O !5 O M H •4 M ■4 M ■a -a * a o H M U) > o o _g-3-s,2 g t| g.K -JS oc/2 SS<» § * -gH3 .6 -i s "B -rt O C3 — ! eo to o «a w -3 .5 ja o '8 '5b5 fe i I I at o o tj-a tn •-- ce .is ~-"S "^ ci o-o ft t-i t> 03 "^ - — • ° ta - o a £ to 03 ' — ' H f* - v — * So » ca h o t.«*SO®.2 o.SbdoBN-g b003 cO . s la «>j j O r=! « « s " ^? ja p. p. o .2 d y ja ■» P k ii >, °iri si 5? >. H H 1 " 4S S s 3 » S a 2 T1 •=• g a "O a) ■R k o «B ft s -d 5"§ ^ 'R a o a> ■3 ft A 2 a 3 13 H O S ■a a o u ■a s W fl 2 o -=: ■3 fl o o .2 a) fan SO -d 53 o Tl 0J d ja GJ J5 -^ r3« g "W ^3 >> pa 3 * ^H ^- o j= 388 ACIDS. FOURTH TABLE. §317, 1. §317. Table for Identification and Separation of the Commonly Occurring Anions (Acids).* 1. Boil the material with dilute HN0 3 . There results: Effervescence; turbidity in a drop of lime-water. Effervescence, penetrating odor. Effervescence, red-brown fumes, odor. Odor, blackening of paper moistened with lead acetate, separa- tion of sulphur in the solution. Odor, | Often masked by the others; see special test* Vinegar odor j below. 2. Boil with concentrated Na 2 C0 3 solution; all cathions (bases) except, the alkalis are precipitated as carbonates or hydroxides, and removed by nitration. The filtrate contains all the anions (acids) and the excess of HN0 3 sets free C0 2 , and Si0 2 is precipitated; C0 a so 2 N 2 3 H 2 S HCN C 2 H 4 C0 3 " . Acidulation with identified in the microcosmic salt bead. 3. Ca(N0 3 ) 2 solution precipitates: as CaF 2 ") insoluble in v in acetic dilute as CaC 2 4 J acid; HC1 as Ca(CN) 2 F c 2 o 4 CN' \ The filtrate is made ammoniaeal. f insoluble; H 2 S0 4 liberates HF . ' soluble, reappearing with NH 3 ; I decolors KMn0 4 solution. OH' gives C 4 H 4 6 " as CaC 4 H 4 6 HAsO," as CaHAsO, HAs0 4 " as CaHAs0 4 HP0 4 " as CaHP0 4 o rO /heated with Fe" + Fe'" + Prussian blue on acidifying. with K* ions in concentrated solution po- tassium bitartrate precipitated. In the filtrate from the above, H 2 S precipitates As 2 S 3 at once in the cold. In the filtrate from the above, H 2 S slowly precipitates from hot solution S 2 + As 2 S 3 . In the filtrate from the above, ammonium molybdate gives yellow pre- cipitate; or Mg" + NH 4 " + OH' gives V MgNH 4 P0 4 . 4. In the filtrate from 3. Ba(N0 3 ) 2 precipitates: Cr0 4 "(Cr 2 7 ") as BaCr0 4 , yellow, soluble in HC1; the yellow color of the solution becoming green on boiling with alcohol. / * From Chem. Prakt. Abegg and Herz (1900), Breslau, Page 113 ; reviewed by Fresenius, Z., 1900, 39, 566. §318, 2. NOTES ON THE DETECTION OF ACIDS. 389 SO/' as BaS0 4 \ / unchanged, remains insoluble in HC1 . SiF/' as BaSiF„ \ insoluble on J gives off SiF 4 , which deposits in HC1 ; ignition \ Si0 2 in a drop of water; the ] residue, BaF 2 , is soluble in ^ HC1. 5. The filtrate from 4. is exactly neutralized with HN0 3 *; Zn(N0 3 ) 2 then precipitates : Fe(CN)/" as Zn 3 [Fe(CN) 6 ] 2 brownish-yellow \ dissolved ( brown I by OH' / coloration. Fe(CN)/'" as Zn 2 Fe(CN) 6 white gives with \ Prussian ) Fe'"' and H* ( blue. 6. A few drops of the filtrate from 5. are treated with as little Fe"" as possible : Eed J Fe(CNS) 3 ) on J permanent red color, coloration 1 Fe(C 2 H 3 2 ) 3 J heating 1 precipitate and colorless solution. In the absence of CNS' another drop is tested with Ag' for the halogens; if a precipitate results or if CNS' is present, one part of the solution is treated with CS 2 and a little Cl-water: I' violet coloration, disappears with , ( much Cl-water. Br' brown coloration, does not disappear with j The second portion is evaporated to dryness with K 2 Cr 2 7 , fused, and "the mass after cooling distilled with concentrated H 2 S0 4 ; appearance of oily brown drops of Cr0 2 Cl 2 , forming Cr0 4 '' with water : CI' . 7. A concentrated water-extract of the original substance is treated with concentrated H 2 S0 4 and solid FeS0 4 or Fe" solution, prepared cold; a brown coloration shows the presence of N0 3 '. The anions mentioned above to some extent exclude one another, being unstable when together in solution owing to their power of mutual oxida- tion and reduction, e.g., S0 3 " and S"; S0 3 " and N0 2 '; NO/ and CN'; NO/ and S"; NO/ and I'; NO/ and HAsO/'; S" and HAsO/' , etc. It is to be noticed that this always simplifies the analytical procedure. §318. Notes on the Detection of Acids. 1. The precipitation of tartrates by calcium salts is incomplete; from calcium sulphate solution a precipitate forms slowly or not at all. Calcium tartrate is soluble in the cold in a solution of KOH , precipitating- gelatinous on boiling; again soluble on cooUng (separation from citrate). Calcium tartrate -is soluble in acetic acid (separation from oxalate). 2. A number of basic carbonates give almost no effervescence when treated *In the original German text it is directed to use HC1 at this point. 390 NOTES ON THE DETECTION OF ACIDS. §318, 3. with acids. To detect the presence of small amounts of carbonate, it is recommended to place the dry powder in a test-tube and fill about three-fourths full of distilled water. Close the test-tube with a two-holed rubber stopper contain- ing a thistle tube reaching nearly to the bottom of the test tube, and a delivery tube reaching just through the stopper. Add dilute sulphuric acid and warm gently. The carbonate is decomposed, driven from the solution, and, owing to the limited air space, readily passes through the delivery tube into the solution of calcium hydroxide. 3. With the generation of an abundance of C0 2 , the precipitate first formed in the Ca(OH) 2 is redissolved (solution of lime in spring water). Boiling drives off the excess of C0 2 and causes the reprecipitation of the CaCOj . Barium hydroxide may be used instead of calcium hydroxide. 4. If compounds have been strongly ignited previous to solution for analysis, oxalates cannot be present. 5. In Table H (§315), if strong oxidizing agents are present, as KCIO, , K 2 Cr 2 T , KMnO, , etc., the oxalic acid will be decomposed on warming with hydrochloric acid. This may be avoided by adding calcium chloride to the solution, neutral or alkaline with ammonium hydroxide. The oxalate will be precipitated and thus separated from the oxidizing agents. After filtering, the precipitate is digested with dilute acetic acid, filtered and the filtrate tested for phosphate with ammonium molybdate. The residue is dissolved in hydrochloric acid, filtered if necessary (calcium sulphate does not dissolve readily), and the filtrate made alkaline with ammonium hydroxide. The pre- cipitate thus obtained is washed, dissolved in nitric acid and tested with potassium permanganate. The filtrate from the solution after the addition of calcium chloride is acidified with hydrochloric acid, heated to boiling and tested for sulphate by the addition of a few drops of barium chloride (§317). 6. In Table XT, if sulphites or thiosulphates are present, the solution in hydrochloric acid must be heated sufficiently to drive off all the sulphurous anhydride, or reactions for oxalates will be obtained, due to the sulphurous acid alone. If there be any doubt as to the complete removal of the sulphur- ous anhydride, the gas evolved by the reaction of the potassium perman- ganate should be passed into a solution of calcium hydroxide. A precipitate of calcium carbonate at this point is positive evidence of the previous presence of oxalic acid or oxalates. 7. Alkali ferro- and ferricyanides are separated from each other by the solubility of the latter in alcohol. 8. In testing for nitric acid the student must not be content with good results from one test. At least four tests should be made, and all of them should give positive results before final affirmative judgment is passed. Failure to bleach indigo solution in the presence of an excess of hydrochloric acid may be taken as conclusive ^evidence of the absence of nitrates. 9. In the analysis of minerals, silica or silicates will usually be present. The silica should be removed before proceeding with the analysis. Fuse the finely divided material with an excess of sodium carbonate, digest the cooled mass thoroughly in hot water, filter and evaporate the filtrate to dryness. Moisten the residue with concentrated hydrochloric acid, and again evaporate to dryness. Pulverize thoroughly, digest in water acidulated with hydro- chloric acid and filter. The residue, white, consists of the silica, Si0 2 . 10. Meta- or pyrophosphates do not react promptly with ammonium molyb- date. In the usual course of analysis they are changed to the orthophosphate (§255, 6A). 11. Phosphoric acid may be detected in the presence of arsenic acid by ammonium molybdate if the solution be kept cold; it is preferable to remove the arsenic before testing. In absence of interfering substances the color of the silver nitrate precipitate will indicate the presence or absence of arsenic acid (§69, 6/). See also note 26. 12. Sulphides which are transposed by hydrochloric acid are best detected by the odor of the evolved gas, and by passing the evolved gas into ammonium- hydroxide and testing with sodium nitroferricyanide. Other sulphides are decomposed by nitric acid or by nitrohydrochloric acid with separation of sulphur as, a leathery mass or as a yellow precipitate. Persistent heating of §318, 19. NOTES ON THE DETECTION OF ACIDS. 391 the sulphur with the reagent decomposing the sulphide will cause the oxida- tion of a portion of the sulplfur to a sulphate which may be detected in the usual manner. A portion of the precipitated sulphur should be burned on a platinum foil with careful observance of the odor of the evolved gas. 13. A sulphite (or other lower oxidized compound of sulphur) is readily- detected by adding barium chloride in excess to a portion of the solution, dissolving in HC1 , filtering if residue remains, and adding bromine or chlorine to the clear filtrate. A precipitation of barium sulphate indicates the oxidation of a lower compound of sulphur to a sulphate. lJf. It frequently becomes necessary to detect and estimate sulphides, thio- sulphates, sulphites and sulphates in mixtures containing two or more of the compounds. The method of procedure will vary according to the nature of the substance. The student will be aided by studying §§257, 8; 258, 8; and 265, 8. 15. The recognition of chlorides, bromides and iodides — by evolving their chlorine, bromine and iodine, in presence of each other — can be accomplished as follows — for the iodine the test being very easy; for chlorine, indirect but unmistakable; for bromine, dependent upon much care and discretion.* The iodine is liberated with dilute chlorine-water, added drop by drop, and is readily detected by starch, or carbon disulphide, according to §280, 8. (As to interference of thiocyanates, see §234.) The chlorine is vaporized (from another portion) as chlorochromic anhydride, and the latter identified by its color and its various products, as described in §269, 8d. Before the bromine is identified the iodine is to be either removed as free iodine, or oxidized to iodate (§276, 86). The- oxidation to iodic acid is effected as follows: Treat with chlorine-water till free iodine no longer shows its color; add a drop or two more of the chlorine-water, and dilute with water, keeping cool; then add the carbon disulphide, agitate and leave the solvent to settle, for the yellow color of bromine. The removal of free iodine may be accomplished as follows: Add chlorine-water, drop by drop, as long as the iodine tint seems to deepen by the addition; add the carbon disulphide, agitate, allow to subside, and remove the lower layer, either by taking it out with a pipette or by filtration through a wet filter. Repeat, if need be, till iodine color is no longer obtained; then continue, with dilute chlorine-water, in test for bromine. If iodide in large proportion is to be removed, it is well, first, to precipitate it out, as far as possible, by copper sulphate and a reducing agent (Note 17). The filtrate is then to be treated by either method above given. 16. The separation by ammonium hydroxide, as a solvent of the silver pre- cipitates — AgCl , AgBr , Agl — when conducted with dilute ammonium hydrox- ide, may be made complete between the chloride and the iodide, also between the bromide and the iodide, but it is very imperfect between the bromide and the chloride. The hot and strong solution of ammonium acid carbonate separates the chloride from the bromide (§269, 8a). 11. The direct removal of iodides by precipitation, leaving bromides and chlorides in solution, can be effected (approximately) by copper sulphate with, sulphurous acid (§77, 6f), or quite completely by palladous chloride (§106, 6). 18. Chloric acid is separated from hydrochloric and all other acids of chlorine, bromine and iodine (except from hypochlorous acid, and from traces of bromic acid), by remaining in solution during the precipitation by silver nitrate (§273, 5). 19. Chloric acid is separated from nitric acid— after finding that stiver nitrate gives no precipitate in another portion of the solution, acidulated — by evaporat- ing and igniting the residue, then dissolving and testing one portion of the solution by silver nitrate for the chloride formed from chlorate during igni- tion (§273, 7). The other portion of the solution is tested for nitric or nitrous acid. * In consequence of the relative com mercial values of bromine and iodine, and the medicinal relations of bromides and Iodides, it is of great importance to search commercial iodides for intentional and considerable mixtures of bromides— an impurity likely to escape cursory chemical examination. There are, however, very slight and usually unobjectionable propor- tions of bromides frequently to be found in the iodides of commerce, and occurring from the difficulty of exact separation in the manufacture of iodine from kelp. 392 NOTES ON THE DETECTION OF ACIDS. §318, 20 20. If we have to separate chloric acid both from nitric and hydrochloric acids, a solution of silver sulphate must be used instead of the nitrate, to precipitate out all the hydrochloric acid. The filtrate from this is evaporated, ignited, dissolved and tested for silver chloride, indicating chlorate in the original solution, and another portion is tested for nitric acid. Also, chlorates are distinguished (not separated) from nitrates, by oxidation of ferrous sulphate in solution with acetic acid on heating, and the consequent formation of the red solution of ferric acetate (§§126, 66; 152; 223, 6). The solution tested must contain no free acids, and no nitrites or other oxidizing agents beside the two in question, but may contain chlorides; and, of course, the ferrous sulphate must be pure enough not to color when heated alone with the acetic acid. Mix the ferrous sulphate solution with the acetic acid, boil, then add the solu- tion to be tested, and heat nearly to boiling, for some minutes. If no red color appears, chlorates are absent, and nitrates may be present. 21. Hypochlorites are separated with chlorates from chlorides (bromides), etc., by silver nitrate; and distinguished from chlorates (in the filtrate from AgCl , etc.) by bleaching litmus, and by their much more rapid decomposition and consequent precipitation of any silver in solution. They are also more active than chlorates, as oxidizing agents. 22. M. Dechan's method (§269, 8i) consists (1) in boiling the mixture with a solution of 40 grammes of K 2 Cr 2 7 , dissolved in 100 cc. of water, which lib- erates and expels all of the iodine without disturbing the bromine and chlorine. 5K 2 Cr 2 7 + 6KI = Cr 2 3 + 8K 2 Cr0 4 + 3I 2 (S) 8 cc. of a. dilute solution of sulphuric acid (consisting of equal volumes of H.S0 4 sp. gr. 1.84, and water) are added to 100 cc. of the dichromate solution, and on boiling, the bromine is distilled off without disturbing the chlorine; after which the chlorine is detected in the usual manner. For other methods of detecting chlorides in presence of bromides and iodides, see §269, 8. 23. For A. Longi's process for the analysis of a mixture of chlorides, bro- mides, iodides, chlorates, bromates, iodates, ferrocyanides and ferricyanides, see C. N., 1883, 47, 209. 2Jf. In the detection of chlorides in presence of cyanides and thiocyanates by the decomposition of the silver salts with concentrated sulphuric acid (§269, 8e), a drop or two of silver nitrate should be added to the precipitate before heating with the acid or a black precipitate will be obtained, apparently carbon. 25. For the detection of a bromide in the presence of an iodide, the most satisfactory method is by the use of potassium chlorate and dilute sulphuric acid as described in §276, 8c. The student should carefully note the change in color of the solution. The first very dark color is due to the liberation of iodine. There is usually a sudden change of color on the complete oxidation of the iodine, but if much bromine be present the solution will be quite dark straw color. This should be tested with carbon disulphide and the heating continued if free iodine is still present. 26. Arsenic "acid should not be present when testing for a phosphate. If the arsenic acid be reduced to arsenous acid by sulphurous acid it will not interfere with the ammonium molybdate test for a phosphate. The excess of sulphur- ous acid must be removed by boiling before testing for the phosphate. Arsenic is best removed by precipitation as sulphide in the second group. 27. Chromic acid is always identified by reduction and precipitation as chromic hydroxide, green, in the third group. The red or yellow color to the original substance usually gives evidence of the probable presence of the hexad chromium. The reduction is effected in the course of analysis by hydro- sulphuric acid with precipitation of sulphur. It is advisable to reduce all chromates by warming with hydrochloric acid and alcohol before proceeding with the analysis. Another portion of the substance may be reduced with sulphuric acid and alcohol and tested for chlorides. _ 28. Manganates are readily decomposed by water with formation of KMnO. and Mn0 2 . In the presence of a fixed alkali the manganate solution is green and does not rapidly change to permanganate. The manganates and perman- ganates in solution are all dark colored (green, purple-red) and should be reduced by warming with hydrochloric acid before proceeding with the analysis. &319, 7. PRINCIPLES. 393 §319. PRINCIPLES. In the following statements, the term salt includes only cases where the metal acts as 'a base, e. g., chromium salts include CrCl 3 , not K 2 Cr0 4 . Only salts of ordinary metals are included. 1. Hydroxides when brought in contact with acids form salts, provided they can be formed by any means in the presence of water. The same is true of oxides. But A1 2 S 3 and Cr 2 S 3 are not formed in presence of water. (Some oxides after ignition fail to unite with all acids, e. g., Sn0 2 , Fe 2 3 , A1 2 3 , but by long boiling unite with a few acids ; while ignited Cr 2 3 is insoluble in all acids). 2. All nitrates, chlorates and acetates are sohible, but salts of cuprosiim, bismuth, tin, antimony and the oxysalts of mercury require some free acid to hold them in solution. 3. All oxides and hydroxides are insoluble, except those of the alkalis, those of Ba, Sr and Ca slightly soluble. The fixed alkalis precipitate solutions of all other metallic salts, Ba, Sr and Ca incompletely. The precipitate with silver, antimonosum and mercury is an oxide, with Sn IV it is Sn0(0H) 2 , with Sb v , SbO(OH) 3 , in all other cases a normal hydroxide. [At boiling heat instead of normal hydroxides other hydroxides are some- times formed, e. g., Fe 4 3 (0H) 6 , and Cu 3 2 (0H) 2 ]. This precipitate re- dissolves in eight cases, forming, if potassium hydroxide be used . . . X 2 Pb0 2 , K 2 Sn0 2 , K 2 Sn0 3 , KSb0 2 , KSb0 3 , K 2 Zn0 2 , KA10 2 , KCr0 2 . The last precipitates on boiling. 4. Ammonium hydroxide precipitates solutions of the first four groups, manganese and magnesium imperfectly and not at all if ammonium ■chloride be present. The precipitate is a normal hydroxide, except that with Sn IV it is Sn0(0H) 2 , with Sb v , Sb0(0H) 3 , with Ag and Sb'" the •oxide, with Pb a basic salt, and with Hg a substituted mercuric ammonium compound, Hg' in addition forms Hg°. The precipitate redissolves in six cases, viz., silver, copper, cadmium, cobalt, nickel and zinc. With silver, NH 4 AgO is formed, with zinc (NH 4 ) 2 Zn0 2 . 5. The chlorides of the first group are insoluble, lead chloride slightly soluble. Hydrochloric acid and soluble chlorides precipitate solutions of salts of the first group, lead salts incompletely, and normal lead salts are not precipitated by mercuric chloride. (For action on higher oxides, etc., see §269, 6 A). 6. The bromides of the first group are insoluble, lead bromide slightly soluble (less than the chloride). Hydrobromic acid and soluble bromides precipitate solutions of the salts of the first group, lead salts incompletely. (For action on higher oxides, etc., see §276, 6A). 7. The iodides of lead, silver, mercury and cuprosum are insoluble. Hydriodic acid and soluble iodides precipitate solutions of lead, silver, 394 PRINCIPLES. §319, 8. mercury and cuprosum. Cupric salts are precipitated as cuprous iodide with liberation of iodine. Ferric salts are merely reduced to ferrous salts with liberation of iodine* Arsenic acid is merely reduced to arsen- ous acid with liberation of iodine. (Bismuth, stannous and antimonous iodides are really insoluble in water, and are readily formed by the action of hydriodic acid or soluble iodides on the dry or merely moistened salts. But the dissolved salts of these three metals fre- quently contain so much free acid that it prevents their precipitation by hydriodic acid or by soluble iodides. Arsenous iodide is decomposed by water. It may be formed from the chloride, best by adding hydriodic acid or a soluble iodide to a solution of arsenous acid in strong hydrochloric acid. Bismuth iodide is black; stannous, antimonous and arsenous iodides are yellowish red.) 8. The sulphates of lead, mercurosum, barium, strontium and calcium are insoluble, those of calcium and mercurosum slightly soluble. Sulphuric acid and soluble sulphates precipitate solutions of lead, mercurosum, barium, strontium and calcium; calcium and mercurosum incompletely. 9. (a) The sulphides of the first four groups are insoluble. Hydro- sulphuric acid transposes salts of the first two groups in acid, neutral, and alkaline mixtures, except arsenic, which is generally imperfectly precipitated unless some free acid or salt that is not alkaline to litmus paper be present. The result is a sulphide, but mercurosum forms mer- curic sulphide and mercury, and arsenic acid forms arsenous sulphide and free sulphur. Ferric solutions are reduced to ferrous with liberation of sulphur. In acid mixture other third and fourth group salts are not disturbed, but from solutions of their normal salts traces of cobalt, nickel, manganese, and zinc are precipitated. (For action on higher oxides, see §257, 6A). (6) Soluble sulphides transpose salts of the first four groups. The result is a sulphide, except that with aluminum and chromium salts it is a hydroxide, hydrosulphuric acid being evolved. With mercurous salts, mercuric sulphide and mercury are formed; with ferric salts, ferrous sulphide and sulphur. 10. The carbonates of the alkalis are soluble. Carbonates of the fixed alkalis precipitate solutions of all other metallic salts. The precipitate is : a. An oxide with antimonous salts. I. A normal hydroxide with Sn", Al , Cr'" and Fe"'; with Sn IV , Sn0(0H) 2 ; with Sb v , Sb0(0H) 3 . c. A normal carbonate with Ba , Sr and Ca salts and, if cold, with silver, mercurosum, cadmium, ferrosum and manganosum. d. A basic carbonate in other cases, except mercuric chloride, which forms an oxychloride. Carbonic is completely displaced by strong acids, for example, from all carbonates, by HC1 , HC10 3 HBr , HBr0 3 , HI , HI0 3 , H 2 C 2 4 , HN0 3 , H 3 P0 4 , H 2 S0 4 , and even by H 2 S, completely from carbonates of first four groups,« incompletely from those of the fifth and sixth groups (Nandin and Montholon, C. r., 1876, 83, 58). §319, 13e. PRINCIPLES. 395 11. All normal and di-metallic orthophosphates are insoluble except those of the alkalis. The normal and di-metallic phosphates of' the alkalis precipitate solutions of all other salts. The precipitate is a normal, di- metallic, or basic phosphate, except that with mercuric chloride and with the chlorides of antimony it is not a phosphate, but an oxide, or an oxy- chloride. All phosphates are dissolved, or transposed by nitric, hydrochloric and sulphuric acids, and all are dissolved by acetic acid except lead, aluminum and ferric phosphates. All are soluble in phosphoric acid except those of lead, tin, mercury and bismuth. 12. Ignition. — a. The oxides of lead and iron heated in the air to a red heat form Pb 3 4 and Fe 2 3 , but A at a white heat form PbO and Fe 3 4 . Other oxides, if ignited in the air to a white heat, when changed, either take up or lose oxygen and leave ultimately the following : Ag , Hg , Au , Pt , Sn0 2 , Sb 2 3 , As 2 3 , Bi 2 3 , CuO, CdO, Fe 3 4 , Cr 2 3 , A1 2 3 , Co 3 4 , NiO , Mn 3 4 , ZnO , BaO , SrO , CaO , MgO , K 2 , Na 2 . In a few cases ignition at a lower temperature gives other results, e. g., Co 2 3 , Ba0 2 , Na 2 2 , Sb 2 4 , etc. I. Alkali hydroxides ignited in air at a white heat are not changed. Other hydroxides give same as in (a). c. Alkali carbonates ignited in air at a white heat are not changed. Other carbonates evolve C0 2 and leave as in (a). d. Fixed alkali oxalates ignited at a white heat in absence of air are changed to carbonates, evolving CO . Ba , Sr and Ca oxalates and a few others at a red heat, in absence of air, form carbonates evolving CO , at a white heat these carbonates are changed to oxides evolving C0 2 . All oxalates ignited in presence of air at a white heat are changed as in (a), except the fixed alkali oxalates which are left as carbonates. In all cases when air is present the CO burns to C0 2 . e. All organic salts ignited at a white heat, in a current of air, leave residues as in (a), but forming carbonates if fixed alkalis are present. The products evolved depend upon the composition of the organic por- tion of the salt. 13. The following acids may be made by adding sulphuric acid in excess to their respective salts and distilling: a. Carbonic from all carbonates. 5. Nitric from all nitrates. d. Sulphurous from all sulphites. e. Hydrochloric from all chlorides except those of mercury. But sul- phuric acid transposes the chlorides of Ag, Sn and Sb with extreme difficulty, so that practically other methods are used to separate hydro- chloric acid from these metals. 396 EQUATIONS. §320. §320. Equations. It is recommended that in the class-room some attention be paid to the balancing of equations as representing the important analytical and synthetic operations, especially those involving oxidation and reduction. The work will be simplified by a careful study of §§216, 217 and 218 and application of the rule as ilKistrated there. When the time permits, the operations represented by the equations studied in the class-room should be performed by each student at his laboratory work-table. At first the teacher should select simpler equations illustrating analytical operations and the principles (§319). Later, the more difficult equations involving oxidation and reduction should be studied. The student should give the authority for every reaction. The accompanying list of equations is merely suggestive and may be expanded by the teacher as the time permits. In each equation the second substance is to be considered as in excess; that is, sufficient to produce the greatest possible change in the first substance. For description and methods of making the basic salts used in this list, see Dammer's Anorganishe Ghemie. J 1. Sb + HN0 3 J 2. As 4 + HN0 8 J 3. As 2 3 4- HNO3 4. Mn(OH) 2 + Pb0 2 + HNO„ 5. MnS0 4 + Pb 3 4 + H,S0 4 , dilute 6. M!n0 2 + KNO3 + K 2 CO s , fusion 7. S 2 + KBr0 3 + K 2 C0 3 , fusion 8. MnS + KNO s + K,C0 3 , fusion 9. Mn.0,, + Pb 3 4 + HNO s 10. Cr(0H) 3 + KKT0 3 + K 2 CO„ , fusion. 11. Pb 3 (AsO„) 2 +2n + H 2 S0 4 , dilute 12. Cu 2 As 2 7 + Zn + H 2 S0 4 , dilute 13. Pb(ET0 3 ) 2 + Al + KOH 14. Cu(NO s ) 2 + Al + KOH 15. Bi(lirO s )3 + Al + KOH 16. Hg 10 O 2 (NO 3 ),, + Al + KOH 17. MnS + Mn(NO s ) 2 + K 2 C0 3 , fus. 18. Mn 3 5 + Pb 3 4 + HBTOj 19. Fe + H,S0 4 , con., hot. 20. KI + KI0 3 + H 2 S0 4 , dilute 21. MnS0 4 + KMn0 4 + H 2 S0 4 , dilute 22. (NaCl + K 2 Cr0 4 + H 2 S0 4 ), dry, hot 23. KN0 3 + PeS0 4 + H 2 S0 4 , con., cold 24. K,Cr 2 0(CT0 4 ) 3 + HC1 , hot 25. Hg 8 0(N0 3 )„ + Al + KOH 26. Ag 3 As0 4 + SnCl 2 + HC1 , sp. gr. 1.18 27. Pb0 2 + K 2 C 2 4 + H 2 S0 4 , dilute 28. Pb 3 4 , white heat 29. NaH 2 P0 2 , ignition 30. Fe,O (AsO 3 ) 2 + PeS + HC1 31. PeBr, + HN0 8 32. Sn + HWO3 , hot 33. KOH + Br, , hot 34. Pel 2 + H„S0 4 , sp. gr. 1.84, hot 35. KBr + KBrO„ + H 2 S0 4 , dilute 36. PeS0 4 + KMn0 4 + H 2 S0 4 , dilute 37. K 2 Cr 2 0(Cr0 4 ) 3 + KOH + Br, 38. 4Hg 2 0,(N 2 0») 3 + Al + KOH 39. Ag 8 AsO s + SnCl 2 + HC1 , sp. gr. 1.18 40. Co 2 3 , ignition, white heat 41. H 2 S + HUO, , sp. gr. 1.42, hot 42. Hg s (As0 4 ) 2 + PeS + HC1 43. Pe s 3 (As0 3 ) 2 + KOH + Cl 2 44. Pel 2 + HNO, , sp. gr. 1.48, hot 45. Cr 2 (S0 4 ) 3 + Cr(NO„) 3 + K 2 C0 3 , fusion 46. Pb,(As0 4 ) 2 + Zn + H 2 S0 4 , dilute 47. KOH + CU , cold 48. KBr, + KI0 3 + H 2 S0 4 , dilute 49. (Cr 2 OHCl 5 + K 2 Cr0 4 + H 2 S0 4 ), dry, hot 5.0. Zn 4 3 (N0 3 ) 2 + PeS0 4 + H 2 S0 4 , concentrated, cold 51. Hg 3 (As0 4 ) 2 + SnCl 2 + HC1, sp.gr. 1.18 52. Mn 3 6 , ignition 53. Pe 2 2 S0 3 + Zn + H 2 S0 4 , dilute 54. Bi 2 S s + HNO3 , dilute, hot 55. Hg s As0 4 + PeS + HC1 56. Cr 2 (OH) 4 S0 4 + KOH + Cl 2 57. Pe(H 2 P0 2 ) 2 + HNO, 58. Cr 2 3 + KC10 3 + K 2 CO s , fusion 59. Cu 5 2 (As0 4 ) 2 + Zn + H 2 S0 4 , dil. 60. KOH + CL , hot 61. Mn.0,, + KC10 3 + K 2 C0 3 , fusion 62. HI0 3 + SnCl, + HC1 63. Bi 12 6 13 (NO3) 10 + PeS0 4 +H 2 S0 4 , con., cold 64. CrO a , ignition 65. EMnO, + H 2 C,0 4 + H 2 S0 4 , dilute 66. PeAs0 4 + SnCl. + HC1 , sp. gr. 1.18 67. Pe 8 Cl 8 + PeS + HC1 68. 5CuO.As 2 5 + Pe + HC1 69. HIO3 + H 2 C 2 4 , hot 70. (Cr 2 (OH) 5 Cl + K 2 Cr 2 7 + H,S0 4 ), dry, hot 71. Pe(N0 3 ) 2 + PeS0 4 + H 2 S0 4 , con., cold ' 72. Ag 2 S0 4 + Zn 73. H,S0 3 + HNO3 , sp. gr. 1.42 74. PeAs0 4 + FeS + HC1 75. Pb(As0 2 ) 2 + KOH + Cl 2 76. Fe(NO s ) 2 + HKTO s 77. Mn.O, + Mn(NO,), + K 2 C0 3 , fusion 78. Pe 8 0,(As0 3 ) 2 + KOH + Br 2 79. Pb 10 O 3 (OH) 8 (NO a ) a + Al + KOH §321. PROBLEMS IN SYNTHESIS. 39? §321. Problems in Synthesis. For the sake of a more thorough drill in the principles of oxidation and other reactions, a few problems are here given; a part of them the student should practically work at his table, but they are chiefly designed for class, exercises. Special care should be taken that a pure product be formed and that the ingredients be taken from the sources indicated. In each case the authority for every step in the process should be stated. 1. Silver oxide from metallic silver. 2. Mercuric bromide from mercurous chloride and sodium bromide. 3. Chromic chloride from potassium chromate and hydrochloric acid. 4. Arsenic acid from potassium arsenite. 5. Potassium arsenate from arsenous oxide and potassium hydroxide. 6. Lead nitrate from lead chloride and potassium nitrate. 7. Mercurous nitrate from mercuric chloride and sodium nitrate. 8. Mercurous oxide from mercuric oxide. 9. Mercuric bromide from metallic mercury and potassium bromide. 10. Lead nitrate from lead dioxide and potassium nitrate. 11. Lead chromate from lead hydroxide and chromium hydroxide. 12. Barium chromate from chrome alum and barium carbonate. 13. Mercuric chromate from mercuric sulphide and chromium hydroxide. 14. Chromium sulphate from potassium dichromate and zinc sulphide. 15. Phosphoric acid from sodium phosphate. 16. Phosphorus from calcium phosphate. 17. Lead iodate from sodium iodide and lead sulphide. 18. Silver iodate from silver chloride and iodine. 19. Ferric arsenate from ferrous sulphide and arsenous oxide. 20. Mercuric bromide from mercuric sulphide and sodium bromide. 21. Ammonium sulphate from ammonium chloride and sulphur. 22. Sodium chloride from commercial salt. 23. Phosphorus from sodium phosphate. 24. Lead sulphide from trilead tetroxide and ferrous sulphide. 25. Ferrous sulphate from ferric oxide and sulphuric acid. 26. Ammonium hydroxide from potassium nitrate. 27. Cadmium sulphate from cadmium phosphate and ferrous sulphide. 28. Mercurous nitrate from mercuric sulphide and nitric acid. 29. Barium sulphate from potassium thiocyanate and barium chloride. 30. Mercurous chloride from mercuric oxide and sodium chloride. 31. Sodium iodate from potassium iodate and sodium chloride. 32. Sodium phosphate from calcium phosphate and sodium chloride. 33. Strontium nitrate from sodium nitrate and strontium sulphate. 34. Potassium sulphate from potassium nitrate and sulphur. 35. Barium sulphate from barium chloride and zinc sulphide. 36. Potassium permanganate from manganese dioxide and potassium nitrate. 37. Arsenous chloride from lead arsenate and sodium chloride. 38. Potassium chromate from potassium nitrate and lead chromate. 39. Potassium iodide from potassium chloride and iodine. 40. Barium chlorate from sodium chloride and barium nitrate. 41. Arsenous sulphide from arsine and ferrous sulphide. 42. Copper sulphate from copper sulphide. 43. Silver nitrite from silver chloride and sodium nitrate. 44. Cuprous chloride from metallic copper and sodium chloride. 45. Manganous carbonate from manganese peroxide and sodium carbonate. 46. Manganous pyrophosphate from manganese peroxide and ammonium phos- phate. 47. Lead arsenate from lead sulphide and arsenous oxide. 48. Bismuth subnitrate from metallic bismuth and nitric acid. 49. Barium perchlorate from sodium chloride and barium hydroxide. 50. Lead iodate from metallic lead and iodine. 398 TABLE OF SOLUBILITIES. §322. §322. Table of Solubilities.* Showing the clasies to which the compounds of the commonly occurring elements belong in respect to their solubility in water, hydrochloric acid, nitric acid, or aqua regia. Preliminary Eemabks. For the sake of brevity, the classes to which the compounds belong are expressed in letters. These have the following signification: W or w, soluble in water. A or a, insoluble in water, but soluble in hydrochloric acid, nitric acid, or in aqua regia. I or i, insoluble in water, hydrochloric acid, or nitric acid. Further, substances standing on the border-lines are indicated as fol- lows: W — A or w — a, difficultly soluble in water, but soluble in hydrochloric acid or nitric acid. W — I or w — i, difficultly soluble in water, the solubility not being greatly increased by the addition of acids. A — I or a — i, insoluble in water, difficultly soluble in acids. If the behavior of a compound to hydrochloric and nitric acids is essen- tially different, this is stated in the notes. / Capital letters indicate common substances used in the arts and in medicine, while the small letters are used for those less commonly occur- ring. The salts are generally considered as normal, but basic and acid salts, as well as double salts, in case they are important in medicine or in the arts, are referred to in the notes. The small numbers in the table refer to the following notes. Notes to Table of Solubilities. (1) Potassium dichromate, W. (2) Potassium borotartrate, W. (3) Hydro- gen potassium oxalate, W. (4) Hydrogen potassium carbonate, W. (5) Hydro- gen potassium tartrate, W. (6) Ammonium potassium tartrate, W. (7) Sodium potassium tartrate, W. (8) Ammonium sodium phosphate, W. (9) Acid sodium borate W. CIO) Hydrogen sodium carbonate, W. (11) Tricalcium phosphate, A. (12) Ammonium magnesium phosphate, A. (13) Potassium aluminum sulphate, W. (14) Ammonium aluminum sulphate, W. (15) Potas- sium chromium sulphate, W. (16) Zinc sulphide, as sphalerite, soluble in nitric acid with separation of sulphur; in hydrochloric acid, only upon heating. (17) Manganese dioxide, easily soluble in hydrochloric acid; insoluble in nitric acid. (18) Nickel sulphide is rather easily decomposed by nitric acid; very difficulty by hydrochloric acid. (19) Cobalt sulphide, like nickel sulphide. (20) Ammonium ferrous sulphate, W. (21) Ammonium ferric chloride, W. * The following table of solubilities, is taken from Treseniua Qualitative Analysis, Well's translation of 16th German edition. §322. TABLE OF SOLUBILITIES. 399 (22) Potassium ferric tartrate, W. (23) Silver sulphide, only soluble in nitric acid. (24) Minium is converted by hydrochloric acid into lead chloride; by nitric acid, into soluble lead nitrate and brown lead peroxide which is insoluble in nitric acid. (25) Tribasic lead acetate, W. (26) Mercurius solubilis Hahne- manni, A. (27) Basic mercuric sulphate, A. (28) Mercuric chloride-amide, A. (29) Mercuric sulphide, not soluble in hydrochloric acid, nor in nitric acid, but soluble in aqua regia upon heating'. (30) Ammonium cupric sulphate, W. (31) Copper sulphide is decomposed with difficulty by hydrochloric acid, but easily by nitric acid. (32) Basic cupric acetate, partially soluble in water, and completely in acids. (33) Basic bismuth chloride, A. (34) Basic bismuth nitrate, A. (35) Sodium auric chloride, W. (36) Gold sulphide is not dissolved by hydrochloric acid, nor by nitric acid, but it is dissolved by hot aqua regia. (37) Potassium plantinic chloride, W — I. (38) Ammonium platinic chloride, W — I. (39) Platinum sulphide is not attacked by hydrochloric acid, is but slightly attacked by boiling nitric acid (if it has been precipitated hot), but is dissolved by hot aqua regia. (40) Ammonium stannic chloride, W. (41) Stannous sulphide and stannic sulphide are decomposed and dissolved by hot hydrochloric acid, and are converted by nitric acid into oxide which is insoluble in an excess of nitric acid. Sublimed stannic sulphide is dissolved only by hot aqua regia. (42) Antimonous oxide, soluble in hydrochloric acid, not in nitric acid. (43) Basic antimonous chloride, A. (44) Antimony sulphide is com- pletely dissolved by hydrochloric acid, especially upon heating; it is decom- posed by nitric acid, but dissolved only to a slight degree. (45) Calcium antimony sulphide, W— A. (46) Potassium antimony tartrate, W. (47) Hydro- gen calcium malate, W. 400 TABLE OF SOLUBILITIES. §322. SOLUBILITY I I & 9 a ■■a o 02 a a a o a a si B A a a a I < a a 1 o S3 o a 8 o3 « ID a s bn a a a M s o O w W w W W W-A A A&I A a 17 A A Chromate. w, w w tt w-a w-a w a w w u. a Sulphate.. ■W„., 5 W Wj4.jp. S0 I I W-I W W 13 .i4 W&I 16 w W W W Phosphate w W e W e ., a a a' A„ a ]3 a a a a a a Borate.... w 3 w„ W a a a w-a n » a a a a Oxalate • . . w, W W a u A a a w-a a w-a a a Fluoride.. w ' w W w-a w-a A-I a-i w w w-a a w-a w-a Carbonate w 4 W >0 W A A A A A A A A Silicate... w W a a a a a-1 a a a a a Chloride . . ■W a , W s6 W 21 . 8 8 W W W w w W&I W W W W Bromide . . TV w W w w w w w w&i w w w w W w W w w w w w w w w w w Cyanide. . . W w w w-a w w w » A a a-1 a-1 Ferrocy'de W w w w-a w w w A-I a i i Ferricy'de W w w w w a i i i Slphocy'de w w w w w w w w w w w w Sulphide. . w TV W W w W-A 46 a a a-i A„ A a 16 a, e Nitrate... w W , w W W w w W W w w w W Chlorate • ■ w w w w W w w w w w w w w Tartrate . . w 6 . 6 . 7 . a2 . 4S W 7 ■w. u u A w-a w w a w-a a w Citrate.... w w w a a w-a w w w w-a a w w Malate — W w w w&a w w-a 47 w w w w Succinate. w w w w-a w-a w-a w w-a w-a w w w-a Benzoate.. w w w w w w w Salicylate. w W w w-a w-a w-a w * ■ Acetate... W W w W w W w W w W w w w Formate.. w w w w w w w w w w w w w Arsenite.. w w w a a a a tt a a Arsenate.. w "W w a a a a a a a. a a a §322. TABLE. TABLE OF SOLUBILITIES. 401 CO 3 2 u © to _ci u CD to A I <53 ■a CD 3 8 3 o u CD s S cj •a a 3 J3 3 a CO s a 3 a ■a ca O a 2 1 B 3 .g s CO 3 1 ct) +J CO 6 d § CO CO 9 O a a 3 < a a A 2 4 A A A a a a a&i A41 Oxide w a A-I a w-a w a a a a Chromate ■w so W W-A A-I w-a W 27 W 30 w W w w a Sulphate a A a a a a a a a a a w-a Phosphate a a a a a a w-a a Borate a a a a a a a a a w a w . a Oxalate w-a w w a w-a a w w-a w w w Fluoride A a A a a A a a Carbonate a a a a a Silicate W W„ I W-I A-I "W, 8 W W-Aas W w 36 W aT . sa W ■w 40 W-A„ Chloride w w i w-i a-i w w w-a W w w w-a Bromide W w i W-A A A w a W a 1 w w w-a Iodide a-i I a W a a W w Cyanide i I i a i i i Ferrocy'de I w i w-a i Ferricy'de w w i a A w a w-a a w Sulphocy'de A a a 33 A A A29 a 31 a A a 38 a sa a4i a 41 A44.45 Sulphide w w W W w,, W W w„ w w Nitrate w w w w w w w w w W Chlorate w-a w„. a a w-a a w a w-a a a„ Tartrate w W a a a w-a w a Citrate w w-a w-a a w-a w w w Malate w-a a a a a w-a w W a Succinate W a w-a w-a a w-a a w-a a w W Benzoate Salicylate v W w W a6 w-a w W aa w W w w Acetate w W w w-a w w w w W w Formate a a a a a a A a Arsenlte a a a a a a a a a a Arsenate 402 REAGENTS. t §323. §323. Reagents.* During the past two years the reagents for use in qualitative chemical ■analysis at the University of Michigan have been made up on the basis of the normal solution; i. e., the quantity capable of combining with one gram of hydrogen or with its equivalent is taken in a litre for the normal solution. For example; Normal potassium hydroxide, KOH , requires 56.1 grams per litre of solution (not 56.1 grams to a litre of water), but the usual pure product contains about ten per cent of moisture, so it is directed to use 62.3 grams or 312 grams for a solution five times the normal strength, 5K Barium chloride, BaCl 2 .2H 2 , has a molecular weight of 244.2, but the hydrogen equivalent is (244.2 -+- 2) 122.1, so for a litre of half-normal solution, N/2, take 61 grams. In the following list of reagents, in the parenthesis immediately follow- ing the formula are given the grams per litre necessary for a solution of the strength indicated. Fresenius' standard follows the parenthesis. Acid, Acetic, HC 2 H 8 2 (300, 5N), sp. gr. 1.04, 30 per cent acid. Arsenic, H 3 As0 4 .y„ 11,0 (15, % H 3 As0 4 -4- 5). Fluosilicic, H 2 SiF 6 ~, §347. Hydrobromic, HBr (40, N/2). Hydriodic, HI (64, N/2). Hydrochloric, HC1 (182, 5N, sp. gr. 1.084), sp. gr. 1.12, 24 p. c. acid. Hydrosulpburic, H 2 S , saturated aqueous solution, §257, 4. Iodic, HIO„ (15, y.,, HIO„-r6). Nitric, HN0 3 (315, 5N, sp.gr. 1.165), sp. gr. 1.2, 32 p. c. acid. Nitrohydrochloric, about one part of concentrated HNO, to three parts HC1 . * Nitrophenic, C„H 2 (N0 2 ) s OH (picric acid). Oxalic, H,C 2 4 .2H 2 , crystals dissolved in 10 parts water. Phosphoric, H 3 P0 4 (16, N/2). Sulphuric, H .SO, , concentrated, sp. gr. 1.84. Sulphuric, dilute (245, 5N, sp. gr. 1.153), one part acid to five parts water. Sulphurous, H 2 S0 3 , saturated aqueous solution. Tartaric, H,C,H 4 0„ , crystals dissolved in three parts water. Alcohol, C 2 H e O , sp. gr. 0.815, about 95 p. c. Aluminum Chloride, A1C1 3 (22, N/2). Nitrate, Al(N0 3 ) 3 .7y.,H,0 (58, N/2). Sulphate, Al 2 (SO„) 3 .18H 2 (55, N/2). Ammonium Carbonate, (NH 4 ) 2 C0 3 (240, 5N), one part crystallized salt in four parts water, with one part ammonium hydroxide. Ammonium Chloride, NH 4 C1 (267, 5N), one part salt in eight parts water. Hydroxide, NH 4 OH (85NH 3 , 5N, sp. gr. 0.964), sp. gr. 0.96, 10 p. c. NH, . Ammonium Molybdate, (NH 4 ) 2 Mo0 4 (36MoO a , N/2, §75, 6d), 150 g. salt in one litre of NH 4 OH , pour this into one litre of HNO„ , sp. gr. 1.2. Ammonium Oxalate, (NH 4 ) 2 C 2 4 ., H 2 O»(40, N/2), one part crystallized salt in 24 parts water. Ammonium Sulphate, (NH 4 ) 2 S0 4 (33, N/2). Sulphide, (NH,),S , colorless, three parts NH 4 OH , saturate with H,S and add two parts of NH,OH . * In the greater number of cases, reagents should be "chemically pure." Different uses require different degrees of purity. An article of sodium hydroxide contaminated 'with ■chloride may be used in some operations ; not in others. ThoBe who have had training in analysis can do without specific directions, which cannot be made to cover all circumstances ; and the beginner must depend on others for the selection of reagents. §323. REAGENTS. 403 Ammonium Sulphide, (NH,) 2 Sj , yellow, allow the colorless to stand for some time or add sulphur. Antimonic Chloride, SbCl„ (30, N/2). Antimonous Chloride, SbCl, (38, N/2). Arsenous Oxide, As 3 3 (8, N/4), saturated aqueous solution. Barium Carbonate, BaC0 3 , freshly precipitated. Chloride, BaCl 2 .2H 2 (61, N/2), one part salt to 10 parts water. Hydroxide, Ba(OH) 2 .8H 2 (32, N/5), saturated aqueous solution. Nitrate, Ba(N0 3 ) 2 (65, N/2), one part to 15 of water. Bismuth Chloride, BiCl 3 (52, N/2, use HC1). Nitrate, Bi(NO s ) 3 .5H 2 (40, N/4, use HN0 3 ). Cadmium. Chloride, CdCl 2 (46, N/2). Nitrate, CdXN0 3 ) 2 .4H 2 (77, N/2). Sulphate, CdSO,.4ELO (70, N/2). Calcium Chloride, CaCl 2 .6H 2 (55, N/2), dissolve in 5 parts water. Hydroxide 2 Ca(OH) 2 , a saturated solution in water. Nitrate, Ca(N0 3 ) 2 .4H 2 (59, N/2). Sulphate, CaS0 4 .2H 2 , a saturated solution in water. Carbon Bisulphide, CS 2 , colorless. Chromic Chloride, CrCl 3 (26, N/2). Nitrate, Cr(NO a ) 3 (40, N/2). Sulphate, Cr 2 (S0 4 ) 3 .18H 2 (60, N/2). Cobaltous Nitrate, Co(NO a ) 2 .6H 2 (73, N/2), in 8 parts of water. Sulphate, CoS0 4 .7H 2 (70, N/2). Copper Chloride, CuCl 2 .2H 2 (43, N/2). Nitrate, Cu(N0 3 ) 2 .6H 2 (74, N/2). Sulphate, CuS0 4 .5H 2 (62, N/2), in 10 parts water. Cuprous Chloride, CuCl (50, N/2, use EC1). Ferric Chloride, FeCl 3 (27, N/2), 20 parts water to one part metal. Nitrate, Fe(N0 3 ) 3 .9H 2 (67, N/2). Ferrous Sulphate, FeS0 4 .7H 2 (80, N/3), use a few drops of H 2 S0 4 . Gold Chloride, HAuCl 4 .3H 2 , solution in 10 parts water. Hydrogen Peroxide, 3 p. c. solution. Indigo Solution, 6 parts fuming H 2 S0 4 to one part indigo, pulverize, stir and cool, allow to stand 48 hours and pour into 20 parts water. Xead Acetate, Pb(C 2 H s 2 ) 2 .3H 2 (95, N/2), dissolve in 10 parts of water. Chloride, PbCl 2 , saturated solution, N/7. Nitrate, Pb(NO s ) 2 (83, N/2). Magnesia Mixture: MgS0 4 , 100 g.; NH 4 C1 , 200 g.; NH 4 OH , 400 cc; H 2 , 800 cc. One cc. = 0.01 g. P. Magnesium Chloride, MgCl 2 .6H 2 (51, N/2). Nitrate, Mg(NO„) 2 .6H 2 (64, N/2). Sulphate, MgSQ 4 .7H 2 (62, N/2), in 10 parts of water. Manganous Chloride, MnCl 2 .4H 2 6 (50, N/2). Nitrate, Mn(N0 3 ) 2 .6H 2 (72, N/2). Sulphate, MnS0 4 .7H 2 (69, N/2). Mercuric Chloride, HgCl 2 (68, N/2), in 16 parts of water. Nitrate, Bg(N0 3 ) 2 (81, N/2). Sulphate, HgS0 4 (74, N/2). Mercurous Nitrate, HgN0 3 (131, N/2), one part salt, 20 parts water and one part HN0 3 . Nickel Chloride, NiCl 2 .6H 2 (59, N/2). Nitrate, Ni(N0 3 ) 2 .6H 2 (73, N/2). Sulphate, NiS0 4 .6H 2 (66, N/2). Palladous Sodium Chloride, Na,PdCl 4 , in 12 parts water. .Potassium Arsenate, K 3 As0 4 (26, y 3 K 3 As0 4 H- 5). Arsenite, KAs0 2 (24, % KAs0 2 -r- 3). Bromate, KBrO s (14, % KBr0 3 H- 6). Bromide, KBr (60, N/2). Carbonate, K 2 C0 3 (207, 3N). Chlorate, KC10 3 , the dry salt. Chloride, KC1 (37, N/2). 404 REAGENTS. §323. Potassium Chromate, K 2 Cr0 4 (49, N/2), in 10 parts water. Cyanide, KCN (33, N/2), in four parts water. Dichromate, K 2 Cr 2 7 (38, y 2 , K 2 Cr 2 7 -=- 4), in 10 parts water. Ferrocyanide, K 4 Fe(CN)„.3H 2 (53, N/2), 12 parts water. Ferricyanide, K 3 Fe(CN),, (55, N/2), in 10 parts water. Hydroxide, KOH (312 [90 p. c. KOH], 5N). Iodate, KIO s (18, y 2 KIO a -H 6). Iodide, KI (83, N/2), dissolve in 20 parts water. Mercuric Iodide, K 2 HgI 4 , Nessler's solution, §207, 6fc. Nitrate, KNO a (50, N/2), the crystallized salt. KTitrite, KNO s , the dry salt. Pyroantimonate, K 2 H 2 Sb 2 7 .6H 2 , see §70, 4c. Permanganate, KMn0 4 (16, y 3 KMn0 4 -r- 5). Thiocyanate, KCNS (49, N/2), in 10 parts water. Hydrogen Sulphate, KHS0 4 , fused salt. Sulphate, K 2 S0 4 (44, N/2), in 12 parts of water. Platinic Chloride, H 2 PtCl,.6H 2 , in 10 parts of water. Silver Nitrate, AgN0 3 (43, N/4), in 20 parts of water. Sulphate, Ag 2 S0 4 , saturated solution, N/13. Sodium Acetate, NaC 2 H 3 2 .3H 2 , in 10 parts of water. Carbonate, Na 2 C0 3 (159, 3N), one part anhydrous salt or 2.7 parts of the crystals, Na 2 CO 3 .10H 2 O , in 5 parts of water. Chloride, NaCl (29, N/2). Tetraborate, Na 2 B 4 O 7 .10H 2 O, horax, the crystallized salt. Hydroxide, NaOH (220 [90 p. c. NaOH], 5N), dissolve in 7 parts of water. Hypochlorite NaCIO, §270, 4. Nitrate, NaNO„ (43, N/2). Phosphate, Na 2 HP0 4 .12H 2 (60, N/2), dissolve in 10 parts of water. Phosphomolybdate, §75, 6d. Sulphate, (35, N/2). Sulphide, Na 2 S , one part NaOH saturated with H 2 S to one part of NaOH , unchanged. Acid Sulphite, the dry salt. Sulphite, Na 2 S0 3 .7H 2 (63, N/2), in 5 parts of water. Acid Tartrate, NaHC.H.O, , in 10 parts of water. Thiosulphate, Na 2 S 2 O s .5H 2 , in 40 parts of water. Stannic Chloride, SnCl 4 (33, N/2). Stannous Chloride, SnCl 2 .2H,0 (56, N/2), in 5 parts water strongly acid with HC1. Strontium Chloride, SrCl 2 .6H 2 (67, N/2). Nitrate, Sr(NO„) 2 (53, N/2). Sulphate, SrS0 4 , a saturated aqueous solution. Zinc Chloride, ZnCl 2 (34, N/2). Nitrate, Zn(NO s ) 2 .6H 2 (74, N/2). Sulphate, ZnS0 4 .7H 2 (72, N/2). INDEX. PAGE Acetates, detection of 251 ignition of 259 with ferric salts 154 Acetic acid 249-251 estimation of 251 formation of 250 glacial -., 250 occurrence of 249 preparation of 250 properties of 249 reactions of 250 solubilities of 250 Acids, detection of, notes on 389 displacement of weak by strong. 180 effect of concentrated sulphuric upon 378 list of " 13 precipitated by barium and cal- cium chlorides 386 preparation of 395 separation from bases 368 table of, precipitated by silver nitrate 387 table of separation of 388 Alkali carbonates, with third and fourth group salts 142 group 221 hydroxides, action on double cyanides 265 hydroxides, detection of in pres- ence of carbonates 262 hydroxides, reactions with 221 Alkalis, on third and fourth group metals 140 Alkali metals 5 Alkaline earth metals 5 earth metals in presence of phos- phates 220 earths, relative solubilities of . . . 204 Alkali sulphides, as reagents. 308, 309 action of, on stannic salts 86 action of, on stannous salts .... 85 Alloys, analysis of 367 with copper 104 Aluminum 142-146 acetate 144 compounds, ignition of 146 PAGB Aluminum, detection of 146, 162 distinction from chromium 148 estimation of 146 hydroxide, formation and prop- erties 144 hydroxide, solubility in ammo- nium chloride 161 occurrence of 143 oxidation of 146 oxide and hydroxides 143 phosphate, separation of ...145, 146 preparation of 143 properties of 142 reduction of 146 salts, reactions of 143 salts, with hydrosulphuric acid. 145 salts, with phenylhydrazin 144 separation of, from iron by ITa^Sa 3 and Wa 2 S0 3 145 separation of, from Cr and 4th group by basic acetates 143 separation of, from glucinum... 196 solubilities 143 Alums 145 Ammonia, occurrence 229 formation of, from nitric acid.. 278 preparation of 229 properties of 229 Ammonium 229-233 arsenomolybdate 62, 98 benzoate, in separation of Cu from Cd 107 carbonate, as a reagent 230 carbonate, in separation of As, Sb and Sn 119 chloride, as a reagent 231 chloride, in the third group.... 161 chloride, with PtCl 4 95 compounds, solubilities of 229 cyanate in formation of urea... 271 detection of 232 directions for detection 236 estimation of 232 hydroxide, as a reagent 230 hydroxide, as a distinguishing reagent for the first group. ... 54 406 INDEX. PAGE Ammonium hydroxide, detection by mercuric chloride 233 hydroxide, preparation and prop- erties of 229 molybdate, preparation of 98 molybdate, test for phosphates. 302 molybdate, with arsenic acid 67 oxidation of 233 phosphomolybdate 98 picrate, formation of 230 polysulphide, formation of 231 salts, detection by Nessler's re- agent 231 salts, ignition of 232 solution to be tested for 236 sulphate, in separation of stron- i tium and calcium 220 sulphide, as a reagent 231 sulphide, formation of 230 sulphide, preparation of 307 sulphide, on iron and zinc groups 180 sulphide, yellow, formation of. . 115 sulphide, yellow, in separation of cobalt and nickel 185 sulphide, yellow, in cupric salts. 115 test for nitric acid 281 thioacetate as a substitute for hydrosulphuric acid 307 Analysis of alkali group 236 proximate 14 operations of 13, 20 ultimate 14 Anions, table of separations of. . . 388 Antimonic acid 76 distinction from antimonous . 122 reduction to antimonous by stan- nous chloride .' 78 salts, action of hydriodic acid on 78 sulphide, precipitation of 77 Antimonites 74 Antimonous argentide 79 compounds with silver nitrate. . . 78 iodide, formation of 78 oxide, formation of 76 salts with permanganates 78 salts with chromates 78 sulphide 74 sulphide, precipitation of 77 Antimony 72-82 acids of 72 compounds, reduction with char- coal '. 80 detection of, in alloys 367 PAGE Antimony, detection of 80 detection of traces of 121 distinction from arsenic 78 estimation of 81 in the test for aluminum 163 metal with hydrosulphuric acid. 66 mirror 65 notes on analysis of 121 occurrence of 72 oxidation of 81 oxides of 72 pentachloride 74 preparation of 72 properties of 72 reduction of 81 reduction to metallic 79 salts 74 separation from arsenic by per- oxide of hydrogen 120 separation from arsenic 64 separation from tin by sodium thiosulphate 78 separation from tin 81 solubility of 73 spots 66 sulphide, separation from arsen- ous sulphide 121 sulphide, separation from stan- nous sulphide 121 with iodine 66 Argol, purification of 252 Arsenates, distinction from arsen- ites 70, 71 separation from phosphates 290 Arsenic 56-72 acid, precipitation by hydrosul- phuric acid 114 acid, reduction by hydrosul- phuric acid and hydriodic acid. 61 acid, reduction with sulphurous acid 60 acid, with ammonium molybdate 67 acid, with molybdates : 62 acid, with nitric acid 66 acid with silver nitrate 67 antidote for 62 compounds, ignition of 69 compounds, with concentrated hydrochloric acid «.. 61 compoTinds, with magnesium salts 61 compounds, with stannous chlor- ide 89 detection of 70 detection of, in poisoning 68 distinction from antimony 78 INDEX. 407 PAGE Arsenic, estimation of .. 70 in glass tubing 70 metal with hydrosulphuric acid. 66 method of Fresenius and Babo . . 68 mirror 64, 65 notes on analysis of 121 oxidation of 71 oxides of 57 occurrence of 57 pentasulphide, formation and properties of 60 preparation of 57 properties of 56 reaction with alkali sulphides... 59 reaction with hydrosulphuric acid 59 reduction of 71 reduction by stannous chloride. 61 separation from antimony 64 separation from antimony by peroxide of hydrogen 120 separation from Sb and Sn by use of thiosulphates 60 spots, formation of 64 spots, properties of 66 sulphide, separation of, from Sb„S 3 121 sulphides with ammonium car- bonate 118 trichloride, formation in analysis 61 with peroxide of hydrogen 71 with hydrosulphuric acid gas... 67 with iodine 66 with nitric acid 66 Arsenites, distinction from arsen- ates 121 Arsenous hydride 64 oxide, crystals, identification of. 67 sulphide, solubilities of 58 sulphide, with HC1 gas 67 Arsine 64 from alkaline mixtures 64 reactions with KOH 121 separation from stibine 65 with hydrosulphuric acid 60, 65 Atomic weights, table of 1 Azoimide (hydronitric acid) 274 Barium 205-208 carbonate, action on ferric salts. 154 carbonate, as a reagent 206 carbonate, as a reagent for third and fourth groups 142 Barium carbonate, as a reagent to precipitate chromium 148 carbonate, and ferric salts 153 carbonate, to separate phos- phates from third, fourth and fifth groups 189 chloride, separation of, from SrCl 2 and CaCl 2 by SCI 206 detection of 208 estimation of 208 hydroxide, formation of 205 iodide, properties 358 occurrence of 205 oxide, preparation of 205 peroxide, ignition of 287 peroxide, preparation 205 preparation of 205 properties of 205 salts, separation of sulphites from sulphates 207 salts, spectrum of 207 separation of, from Sr , Ca and Mg by sulphates 207 solubilities of 206 strontium and calcium, separa- tion of by alcohol 220 sulphate, separation 209 Bases, alkali 11 alkaline earth 11 copper, group of 12 definition of 3 fifth group of 11 first group of 13 fourth group of 12 iron group of 12 need for separation from acids. . 368, 369 second group of 12 silver group of 13 sixth group of 11 third group of 12 tin group of 13 zinc group of 12 Beryllium 195 Bismuth. 100-104 blowpipe, reactions of 103 chloride, sublimation of 103 detection of 103> detection in alloys 367 detection by cinchonine 102 detection as iodide 103 detection by alkaline stannite... 103 detection of traces of 102 dichromate 103 estimation of 103 408 INDEX. PAGE Bismuth hydroxide, solubility in glycerol 101 iodide, stability toward water. . . 103 nitrate, precipitation with HC1. 101 nitrate, reactions 101 notes on analysis of 128 occurrence of 100 oxidation of 104 oxides and hydroxides of 100 oxychloride, formation of 101 pentoxide, reaction with halogen acids 101 preparation of 100 properties of 100 reactions of, comparison with Cu and Cd 112 reduction by grape sugar 104 salts, reaction with the alkalis.. 101 separation from Cu by glycerol. 101 solubility of 100 sulphide, formation of 102 sulphide, separation of, from CuS 102 sulphide, separation of, from tin group 102 Blowpipe, examination of solids. . 374 Blue vitriol 105 Bonds, plus and minus 238 Borates, green flame by ignition of 246 in analysis 54 reactions of 246 Borax, bead, formation of 247 bead, test for Hn 184 bead, use of 365 Boric acid 245-247 estimation of 247 formation of 245 occurrence of 245 preparation of 245 properties of 245 solubility of 246 Boron 245 Bromates, detection of 349 estimation of 350 ignition of 349 preparation of 348 solubilities of 349 Bromic acid 348-350 properties of 348 reactions of 349 Bromides, detection of 347 detection in presence of iodides. 391, 392 estimation of 348 formation of 345 PAGB Bromides, ignition of 347 solubilities of 345 with first group metals 346 Bromine 342-344 detection of 344 estimation of 344 formation of 343 occurrence of 343 preparation of 343 properties of 342 reactions with 343 solubilities of 343 Brown ring, test for nitric acid. . . 280 Brucine, reactions with nitric acid 281 Cacodyl oxide, test for acetates... 250 Cadmium 110-112 detection of 112 estimation of 112 hydroxide 110 notes on analysis of 129 occurrence of 110 oxide 110 properties of 110 reactions of, comparison with Bi and Cu 112 Salts, absorption by gaseous sub- stances, separation from Cu. . Ill salts, fused with K 2 S 112 salts, with alkaline tartrates, separation from Cu Ill salts, with alkalis 110 salts, with ammonia Ill salts, with barium carbonate. . . Ill salts, with pyrophosphates, sepa- ration from Cu Ill salts, reactions with lSa^S 2 3 , separation from Cu Ill salts, reduction of by metals. . . . 112 salts, reduction of by ignition. . 112 separation of 110 separation from Cu by KCNS. . . Ill separation from Cu by glycerol. 105 separation from Cu by NazSaOj and Na 2 SO s Ill solubilities of 110 Caesium 233-234 Calcium 210-214 carbonate in spring water 211 carbonate, solubility of ........ 218 detection of 213 detection of by spectrum 213 estimation of 213 group 203 group, directions for analysis of. 218 INDEX. 409 PAGE Calcium hydroxide, formation and properties 211 hydroxide, formation by Na^S.. 213 hydroxide, to detect CO 212 oxide, formation and properties. 210 occurrence of 210 peroxide 211 preparation of 210 properties of 210 salts with UTa^S 213 salts, separation of oxalic from phosphoric acid by 212 separation from Ba and Sr by (NH 4 ) 2 S0 4 211 separation from Ba and Sr by amyl alcohol 211 solubilities of 211 sulphate, separation from stron- tium sulphate 209 sulphate, solubility in ammo- nium sulphate 220 sulphate, to detect strontium... 213 Carbon 247-249 detection of 249 preparation of 248 properties of 247 reactions of 248 reduction by ignition with 248 relations of 10 solubilities of 248 dioxide 7 258-263 dioxide, absorption by Ca(OH) 2 . 261 dioxide, detection in sodium car- bonate 262 dioxide, detection by calcium hy- droxide 212 dioxide, distinction from H 2 S , S0 2 , N" 2 O a , etc 261 dioxide, formation of 259 dioxide, occurrence of 259 dioxide, properties of 259 monoxide 254, 255 Carbonates, acid, decomposition of 230 decomposition of, by acids 262 detection of 262 detection of traces 390 estimation of 263 ignition of 262 occurrence of 259 preparation of 259 reactions with 260 Carbonic oxide, formic anhydride. 254 Cassius' purple 93 Cerium I 93 Chili saltpeter, occurrence of 277 PAGE Chloric acid 339-341 formation of 339 preparation of 339 properties of 339 separation of, from nitric acid.. 391 Chlorates, detection of 341 distinction from nitrates 392 estimation of 341 formation from chlorine 329 ignition of 340 oxidation by ignition of 341 preparation of 339 reactions with 339 solubilities of 339 Chlorides, detection of 149 detection of, in presence of bro- mides 335, 336, 391 detection of, in presence of cy- anides or thiocyanates....335, 392 formation of 331 ignition of 33 i Chloride of lime, formation of.... 337 estimation of, by H 2 287 Chlorine 327-330 action on metals 328 as an oxidizer 328 detection of .330 estimation of 330 formation of 327 occurrence of 327 peroxide, formation and proper- ties 338 properties of 327 solubilities of 328 Chlorochromic test for chlorides. . 335 anhydride 149 Chlorous acid, formation and de- tection 338 properties of 337 Chromates 150 in test for HC1 149 reduction of, by hydrochloric acid 149 reduction of by H 2 S 149 use in separation of barium. . . . 207 with antimonous salts 78 with As"' 149 with ferrous salts 7 . . . . 158 Chromic acid, detection of 150 formation of 149 identification of » 392 Chromium 147-151 distinction from aluminum 148 estimation of 150 410 INDEX. PAGE Chromium hydroxide, solubility in ammonium hydroxide 162 and manganese in third group separation 163 metal, solubility of 147 occurrence of 147 oxidation of 150 oxides and hydroxides 147 oxide, solubilities of 147 preparation of 147 properties of 147 reduction of 150 salts, solubilities of 147 salts, reaction of 148 separation from Al and Fe by H 2 2 150 separation from fourth group.. 148 separation from Fe by Na,S,O a and Ma 2 S0 145 Chromous salts 148 Cinchonine as a test for bismuth. 102 Citric acid 251-252 detection of oxalic acid in 251 distinction from tartaric 251 properties and reactions 251 Colloidal sulphides of the fourth group 184 Color, flame tests 365 Columbium, distinction from Ti. . 201 properties and reactions , 193 separation from tantalum 193 Cobalt .4 163-168 bead test 167 cobalticyanide separation from nickel 166 detection of 168 detection of in presence of Hi by H 2 O a . . . ; 185 estimation of 168 hydroxide 161 metal, solubilities of 164 nitrate, effect of ignition with. . 365 occurrence of 164 oxidation of 168 oxides and hydroxides 164 phosphate, a distinction from Ni 167 preparation of 164 properties of 163 reduction of 168 salts, solubilities of 164 salts, with alkalis 165 salts, with barium carbonate... 165 separation from nickel by ether 164 separation from nickel by KNO, 166 PAQB Cobalt, separation from nickel by KMn0 4 167 separation from nickel by ni- troso-/?-naphthol 166, 185 Colloidal sulphides of fourth group 184 Color, flame tests 365 distinction from Ti 201 Columbium, properties and reac- tions of 193 distinction from Ti 201 separation from tantalum 198 Copper 104-110 acetoarsenite 108 analysis of, notes 128 arsenite 108 compounds with cyanogen 107 detection of 109 detection of, in alloys 367 detection of traces of, with H 2 S 108 detection of, with HBr 108 electrical conductors 104 estimation of 109 f errocyanide, formation of 107 group, metals of 56, 100 hydroxide of 104 occurrence of 104 oxides of 104 precipitation of, by iron wire. . . 109 preparation of 104 properties of 104 reactions of, comparison with Bi and Cd 112 reduction by ignition 109 .-reduction of, by KCNS 107 salts, detection by potassium xanthate 107 salts, reaction with zinc-plati- num couple 109 salts, reduction of, with H 3 F0 2 . 107 salts, separation of, from Cd by Na 4 P 2 7 107 salts, solubilities of 105 separation of, from Bi by gly- cerol 101 separation of, from Cd by gly- cerol 105 separation of, from Cd by 2Ta 2 S 2 3 and Na 2 C0 3 Ill separation from Cd by nitroso- /3-naphthol 107 separation from Cd by ammo- nium benzoate 107 separation from Pd 106 traces, loss of 115 traces of, with K 4 Fe(CN) 107 INDEX. 411 PAGE Cream of tartar, formation of 252 Cuprammonium salts 106 Oupric hydroxide in NH.OH 105 hydroxide, effect of boiling 106 hydroxide, formation of 106 hydroxide, with glucose 106 hydroxide, with tartrates. . .105, 106 salts, reaction with glucose 105 salts, reaction with iodides 108 salts, reaction with lfa 2 S 2 3 108 salts, reduced by S0 2 108 sulphide colloidal 10S sulphide, formation of 107 sulphide, separation from Cd by H 2 S0 4 108 sulphide, solubility in (NHJjS,, 108 sulphide, solubility in KCN 108 sulphide, with K 2 S 115 sulphide, with (2raj 2 Sx 115 Cuprous iodide 108 oxide, formation of, by glucose. 105 salts, oxidation of, by As,0 3 110 salts, separation, fromCd by S. 107 salts, with metallic sulphides... 107 sulphide, formation by Na 2 S 2 3 . 108 thiocyanate, formation of 107 Cyanates, detection of, in presence of cyanides 271 Cyanic acid 271 Cyanide of silver, distinction from chloride 265 Cyanides, detection as thiocyan- ate 267 double, dissociated by acids 264 double, not dissociated by acids. 265 estimation of -„ 267 guaiaeum test 267 ignition of 266 preparation of 264 reactions with 264 simple, with mineral acids 265 solubility of 264 transposition by acids 267 Cyanogen properties and reac- tions 263 Danger and Flandin, detection of arsenic 69 Decomposition of organic mate- rial 362, 363 Dialysis, separation from organic material by 363 Didymium '. 194 D*Bhenylaniine test for nitric acid 281 PAGE: Dissociation, electrolytic 20> Dithionic acid, formation and properties 314 Dragendorff's reagent 102 Electrolytic dissociation 21 Epsom salts 304 Equations illustrating oxidation and reduction 396 rule for balancing 239 Erbium 195. Ethyl acetate, odor of 250' Everett's salt 154 Fatty material, removal of 363 Ferric acetate, formation of 250i acetate, separation of from chro- mium 154 basic nitrate, separation from aluminum 158 and ferrous compounds, distinc- tion 162 hydroxide, antidote for arsenic. 62 phosphate, formation of 156 salts, detection of traces 155 salts, with acetates 154 salts, with BaC0 3 153 salts, with HI and iodides 158 salts, with H 2 S 157 salts, with H 3 P0 2 156 salts, with K 3 Fe(CN) 6 155 salts, with K 4 Fe(CN) 6 155 salts, with KCNS 155 salts, with stannous chloride 89 salts, separation from ferrous sulphate 153 Ferric thiocyanate, distinction from ferric acetate 154 hindrance to reactions of 155 Ferricyanides, in distinction be- tween Co and Ml 166 reactions of 270 Ferrocyanides, detection of 269 detection and estimation 271 reactions of 268 Ferrous iron, detection of, in ferric salts 155 in the third group 161 in the third group with phos- phates 189 salts, traces in ferric salts 155 salts, with chromates 158 salts, with HNO3 156 salts, with KCN 154 412 INDEX. PAGE Ferrous salts, with KaPeCCN), 155 salts, with K 4 Fe(CN') e 154 sulphate, with gold salts 93 First group metals, table of 52 Fixed alkalis 221 alkali hydroxides on stibine 79 alkalis with salts of tin 84 Flame, blowpipe, production of. . . 364 or color tests 373 oxidizing and reducing 363 reactions with copper salts 109 Fluorides, solubilities of 289 Fluorine 288 Fluosilicates, formation of 289 Fluosilicic acid 247-248 in detection of potassium 225 in separation of Ba , Sr and Oa. 207 Formates, formation from cyan- ides 266 Fourth group, directions for anal- ysis 184 reagents 141 sulphides colloidal 184 table of :. 183 Fresenius and Babo, detection of arsenic 68 Froehde's reagent 99 Fulminating gold 92 Gallium (eka-aluminium) 195 Gases, absorption of by palladium 131 Germanium, properties and reac- tions 136 sulphide 118 Glass, etching by hydrofluoric acid 289 Glauber's salts 304 Glucinum (Beryllium) 195 distinction from yttrium 202 separation from aluminum 196 separation from cerium 193 Glucose, in formation of cuprous oxide 105 Gold 91-93 detection in alloys 367, 368 detection of 93 distinction from Fd 132 estimation of 93 fulminating 92 notes on analysis 123 occurrence, properties, etc 91 reduction by ferrous sulphate... 93 reduction with oxalic acid 92 salts with alkalis 92 salts with stannous chloride .... 89 PAGE Gold, separation from Ir 133 Greeuockite 101, 110 Gypsum 213 Halogens 9 as oxidizers 330 compounds, comparative table of 361 hydracids as reducers 330 Heat, upon substances in closed tubes 364, 370 upon substances in open tubes. 364, 371 Hydriodic acid 353-356 action on antimonic salts 78 action on arsenic salts 61 action on ferric salts 158 as a reducer . ; 354, 355 formation of 353 Hydrobromic acid .' . . . .344-348 detection of Cu with 108 formation of 345 occurrence of 345 preparation of 345 properties of 345 reactions of 345 Hydrochloric acid 330-336 action on Sb s S s 77 action on bismuth nitrate 102 effect of excess in second group. 113 formation of 331 formation from BIgCl 2 216 gas on arsenic sulphide 67 occurrence of 331 preparation of l . 331 properties of 330 reactions with 332 solubilities of 331 Hydrocyanic acid 263-267 formation of 264 occurrence of 264 on PbQ 2 264 preparation of 264 properties of 263 solubilities of 264 Hydrof erricyanic acid 269-271 Hydroferrocyanic acid 267-269 separation from hydroferri- cyanic acid 269 Hydrofluoric acid 289 Hydrofluosilicic acid (fluosilicic acid) 289 Hydrogen 243-244 absorption by Pd sponge 131 detection of 244 estimation of 244 INDEX. 413 PAGE Hydrogen, formation of 243 nascent 244 occluded 244 occurrence of 243 preparation of 243 properties of 243 reactions with 243 reducing 1 action of, with ignition 244 solubilities of 243 peroxide, detection of 287 peroxide, estimation of 287 peroxide, estimation of bismuth with 104 peroxide, formation oi 286 peroxide, occurrence of 286 peroxide, on sulphides of arsenic and antimony 120 peroxide, preparation of 286 peroxide, properties of 285 peroxide, reactions with 286 peroxide, reagent to separate Co from 2Ti 185 peroxide, separation from ozone 237 peroxide, separation of Al , 3Te and Cr with 150 peroxide, solubilities of 286 peroxide, with arsenic 71 Hydronitric acid 274-275 Hydrosulphuric acid 306-311 action on copper salts 107 action on ferric salts 157 aqueous solution 113 dissociation of 114 formation of 307 gas as a reagent 113 gas on antimony 67 gas on arsenic 67 occurrence of 307 on aluminum salts 145 on stannic salts 86 on stannous salts 85 on third and fourth group salts 141, 161 preparation of 307 properties of 306 uses as a reagent 308 with arsenic acid 114 with oxidizing agents 114 Hydrosulphurous acid 314 Hydroxylamine, formation and properties 278 Hypobromous acid, formation and properties 348 Hypochlorites, detection of 392 formation of 337 Hypochlorites, formation from chlorine 329 on arsenic 66 Hypochlorous acid 337 Hypoiodous acid, existence of 351 Hyposulphites, detection of 296 ignition of 296 Hypophosphites in formation of PH 3 296 Hypophosphoric acid 298 Hypophosphorous acid 295-297 estimation of 297 formation of 295 preparation of 296 properties of 295 reactions of 296 solubilities of 296 with bismuth salts 102 Hyposulphurous acid 314 Imperial green 108 Indigo test for nitric acid 281 Indium 196 Ink, sympathetic 154 Iodates, detection of 359 estimation of 359 formation of 359 ignition of 357 reactions of 358 Iodic acid 357-359 formation of 357 preparation of 357 properties of 357 reactions of 358 Iodide of nitrogen 351 I Iodides, decomposition by HN0 3 . 281 detection as Pdl 2 131 detection of 356 estimation of 356 formation of 353 ignition of 356 occurrence of 353 reactions of 354 separation of, from bromides and chlorides by KMn0 4 176 solubilities of 353 Iodine 350-352 detection of 352 estimation of '. 352 formation of 351 liberation by copper salts 108 occurrence of 351 on antimonous salts .' 78 on antimony 66 414 INDEX. PAGE Iodine, on arsenic 66 preparation of 352 properties of 350 reactions of 351 separation from Br by Pd 131 solubilities of 351 Ions 21 Ionization and solution 20, 24 Iridium 132-133 Iron 151-159 and zine groups 140 detection of 162, 163 detection of traces in copper... 154 detection of traces 154, 155 estimation of 159 group 142 group, separation from Co , Ni , and Mn by ZnO 158 hydroxides , 152 in relation to metals 6 occurrence of 151 oxidation of 159 oxides 152 preparation of 151 properties of 151 reduction 159 salts, ignition of 158 salts, solubilities of 153 salts, with alkalis 153 salts, with nitroso : j3-naphthol. . 154 salts, separation from Al as basic nitrate 158 separation from Al and Cr by nitroso- /?-naphthol 154 separation from Cr and Al 154 separation from Ni by xanthate 170 solubilities of 152 Lanthanum 197 Lead 29-36 acetate, properties of 32 chloride 34 chloride, precipitation of 53 chromate, formation of 35 compounds, ignition of 35 detection in alloys 367 detection of 36 estimation of 36 in the test for Al 163 iodide, formation and proper- ties 35 notes on analysis of 127 occurrence of 29 oxidation and reduction 36 oxides of ; 29 Lead oxides, solubilities of 30 preparation of 29 properties of 29 red 29 relation to nitrogen family 7 salts, reactions 32, 35 salts, solubilities of 31 solubilities of metallic 30 sulphate, formation and proper- ties of 34 sulphide, formation and proper- ties of 33 tests for 54 Leblanc-soda process 259 Lithium 234-236 Lime, slacked 211 stone (CaC0 3 ) 213 Light, action on silver salts 50 Magnesia mixture 145 Magnesium. 214-216 as a reducing agent 216 detection of 216 estimation of 216 hydroxide, formation 214 occurrence of 214 oxalate, separation of, from K and Na 215 oxide, formation of 214 preparation of 214 properties of 214 removal for detection of sodium. 236 salts, with ammonium salts 215 salts, with arsenic acid 61 salts, with Na 2 S 215 salts, solubilities of 214 Malachite 104 Manganates, identification 392 Manganese 172-177 detection of, 176, 186 estimation of 177 hydroxides of 172 hydroxides, solubilities of 173 ignition of .* 176 in third group 161, 163, 184 occurrence of 172 oxidation of 177 oxidation to permanganic acid. . 175 oxides 172 oxides, solubilities of 173 preparation and properties 172 reduction of 177 reduction by sulphites 175 salts, reactions with oxalic acid. 174 salts, solubilities of 173 INDEX. 415 PAOE Manganese salts, with alkalis 174 salts, with sulphides 175 separation from zinc with acetic acid 1S4 solubilities of 173 with KI 176 Manganic acid 173 Marsh's test 62 Mass action, law of 22, 38 Mayer's reagent 43, 232 Mercurammonium compounds ... 39 Mercuric chloride with stannous chloride 88 sulphide, formation and proper- ties 41 sulphide, with K 2 S 11 5 Mercury 37-45 chlorides 42 compounds, ignition of 43 detection and estimation of 44 iodides 42 metallic, analysis of ..>. 367 occurrence of 37 oxidation of 45 oxides 57 preparation and properties of... 37 salts, reactions 39, 43 salts, solubilities of 38 solubilities of 37 sulphide, analysis of 126 Metals, classification 10 grouping 375 table of separation 376 Metaphosphoric acid 299 Metastannic acid 83 Microcosmie salt 230 use in ignition 365 Milk of lime 211 Molybdates in analysis 54 with phosphates 98 Molybdenite 97 Molybdenum 97-99 deportment in second group. .. . 99 detection of 99, 122 estimation of 99 ignition tests 99 notes on analysis of 123 occurrence of 97 oxides and hydroxides 97 preparation and properties 97 reduction tests 99 solubilities of 97 Molybdic acid 97 Nascent hydrogen on nitric acid. . 278 PAGE 'Neodymium 194, 197 Nessler's reagent 43, 231 Nickel 168-172 detection of 171 detection of, in presence of Co by KI 185 distinction from cobalt 170 estimation of 171 hydroxides 169 hydroxides with KI 171 ignition of 171 occurrence of 169 oxidation of 171 oxides 169 properties and preparation 168 reduction 172 salts with alkalis 169 separation from Co , cyanide method 166 separation from Co , by nitroso- /3-naphthol 168 separation from Co , by KN0 2 . . 166, separation from Co , by sulphide 170 separation from Co , by xan- thate 170 solubilities of 169 solubility of NiS in ammonium sulphide 170 xanthate, separation from 3?e. . 170 Niobium (Columbium) 193-194 Nitrates, decomposition by igni- tion 280 distinction from chlorates 392 occurrence of 277 preparation of 277 proof of absence 390 solubilities of 278 Nitric acid 277-282 as an oxidizer 278 brown ring test 281 decomposition of, by HC1 279 detection of 280 detection by diphenylamine 281 detection by reduction to NH 3 . 278, 281 detection by reduction to nitrite. 281 dissociation, by heat 279 estimation of 282 formation of 277 indigo test 281 in separation of Sn , Sb and As. 119 sodium salicylate test 281 with phenol 281 with pyrogallol 282 with brucine 281 416 INDEX. PAGE Nitric acid, occurrence of . , : . 277 on antimony 66 on arsenic 66 preparation of 277 products of reduction 278 properties of 277 Nitric anhydride, formation of 278 oxide 104, 275, 215 Nitrites, decomposition by igni- tion 276 detection of 276 test for nitric acid 281 Nitrof erricyanides 270 Nitrogen 273-274 chloride 62, 120, 327 combination 'With elements 274 detection and estimation 274 family 7 formation, occurrence 274 peroxide 277 properties 273 Nitroso-/?-naphthol, separation of Co and Ni 166, 185 separation of Cu from Cd 107 ■with iron salts 154 Nitroprussides 270 Nitrous acid 276-277 as an oxidizer 276 as a reducer 276 formation of 276 occurrence of 276 properties of 276 reactions with 276 solubilities of 276 Noble metals, enumeration 7 Nordhausen sulphuric acid 322 Notes on detection of acids 389 on analysis of calcium group. 218, 219 on analysis of third group 161 Order of laboratory study 24 Organic substances, removal of. . 362, 363 Osmium 133 Osmotic pressure 21 Oxalates, decomposition by igni- tion of 390 decomposition by oxidation 390 detection of 258 distinction from tartrates. .253, 389 estimation of 258 ignition of 258 in 3d, 4th and 5th groups 189 reactions of 256 solubilities of 256 PAGB Oxalic acid m . 255-258 as a reducer 256 decomposition of by H 2 S0 4 257 formation of 255 in separation of gold 92 occurrence of 255 preparation and properties of... 255 solubility of 256 Oxidation, balancing equations in. 238 Oxidizing flame 363 Oxygen 282-284 as a poison 284 combinations with ignition 284 detection of 284 estimation of 284 formation of 283 occurrence of 282 preparation of 283 reactions with 284 Ozone 284 separation from 3£ 2 2 287 Palladium 131-132 distinction from gold and plati- num 131, 132 separation from copper 106 sponge 131 Palladous iodide in analysis 131 Paris green j. .62, 108 Pentathionic acid, formation and properties 316 Perchlorates, preparation and properties 341, 342 Perchromic acid 151 Periodic acid 360 system, table of 2 Permanganates, identification 392 action on antimonous salts 78 Permanganic acid , 173 Persulphuric acid 326 Phenol reaction for nitric acid 281 Phenylhydrazine, on aluminum salts 144 Phosgene, formation 254 Phosphates, changes by ignition. 303 detection 162, 303, 390 distinction between primary, secondary and tertiary 301 estimation of 304 in presence of third and fourth group metals. 142, 188, 189, 191, 192 occurrence of 299 reaction with ammonium molyb- date 188, 302 separation as ferric phosphate . . 188 INDEX. 417 PAGE Phosphates, solubilities of 300 Phosphides, formation of 303 Fhosphine 295 Phosphoric acid 298-304 preparation of 300 properties of ." 298 Phosphoric anhydride, formation of 299 Phosphorous acid 297-298 detection of 298 preparation and properties of... 297 Phosphorus 293-295 detection and estimation of 295 in combination with the halo- gens 294 occurrence and preparation of.. 293 properties of 292, 294 use in match-making 293 Phosphotungstates 135 Picric acid, in detection of potas- sium 224 Plaster of Paris (calcium sul- phate) 213 Platinized asbestos 94 Platinum 93-97 apparatus, care of 95 black 93 chloride, as a reagent 95 detection of 96,122,367 distinction from palladium. .131, 132 estimation of 96 iridium alloys, properties 132 notes on the analysis of 123 occurrence of 94 preparation and properties 93, 94 reduction of 95, 96 sponge 93 Polarity 3 Potassium 222-226 as a reducing agent 226 bichromate, in test for stron- tium and calcium 219 carbonate, as a reagent 223 chlorate, in preparation of oxy- gen 283 chloride with platinum chloride. 95 cyanide with copper salts 107 cyanide with ferrous salts 154 detection of 223, 22G estimation of 22G ferricyanide, formation of 269 f errocyanide, formation of . . 265, 267 hydroxide, as a, reagent 223 iodate, in separation of alkaline earths 207 TAGS Potassium iodide, as a reagent 224 iodide, in separation of AgCl from SbCl 3 120. iodide, in the test for nickel 185 iodide, on nickelic hydroxide.... 171 iodide, on permanganates 176 nitrite in separation of cobalt from nickel 166 occurrence, preparation and properties of 222 picrate 224 pyroantimonate 73, 228 salts, flame test 225 thiocyanate with copper salts . . . 107 thiocyanate with iron salts 155 xanthate, for detection of copper 107 Powder of algaroth 75. Praseodymium 194, 197 Precipitates, formation and re- moval of : 17, 18 Principles 393 Problems in molecular propor- tions 19 in synthesis 397 Prussian-blue, formation of.. 155, 266 Purple of Cassius 89, 93 Pyroantimonic acid 73 Pyrogallol, as a. test for nitric acid 282 Pyrophosphoric acid, formation.. 299 Pyrosulphuric acid, formation . . . 322 Reagents, care in the addition of. 17 list of 402 Reducing flame, description of... 363 Reduction, balancing equations in 238 with charcoal 364, 365, 371 Reinsch's test for arsenic 67 Rhodium, distinction from ruthe- nium 133 properties and reactions 130 Rochelle salts, composition of...- 253 Rosolic acid as a test for carbon dioxide 262 Rubidium, properties and reac- tions 234 Rule for balancing equations 239 Ruthenium, properties and reac- tions 129 Saltpeter, occurrence 277 Samarium, properties and reac- tions 197 Scandium, properties and reac- tions 198 418 INDEX. PAGE Scheele's green and Schweinfurt's green 62, 108 Eelenic acid, separation from sul- phuric acid 139 Selenium, properties and reac- tions 138, 139 Silica (silicon dioxide) 290 detection and estimation of 292 in the borax bead 292 in the third group 163 removal of 390 solubilities of 291 Silicates, decomposition by igni- tion 291 in analysis 54 Silicic acid 290-292 Silicon 290 distinction from tantalum 198 Silico-fluoride (fluosilicate) 289 Silicon fluoride, formation.... 288, 289 preparation and properties 290 separation from thorium 200 Silver 45-50 arsenate and arsenite, formation 62 bromate, properties of 349 chloride, formation and proper- ties 48 cyanate in distinction from chlo- rides 271 detection of 50, 367 estimation of 50 in presence of mercury salts 55 iodate, properties of 358 mirror, formation by tartrates. 253 nitrate, action on stibine 79 nitrate with stannous and anti- monous salts 78, 79, 88 occurrence and properties of.... 45 salts, action of light upon 50 solubilities of 46 thiocyanate, separation from silver chloride 272 Soda lime on stibine 79 process, Le Blanc's 259 process, Solvay's 260 Sodium 226-229 amalgam, action with arsenic... 64 as a reducing agent 229 detection of 73, 228 estimation of 229 flame test 228 hydroxide, formation of 227 nitroferricyanide as reagent. 230, 311 occurrence of 227 phosphate as reagent 227 PAGH Sodium phosphomolybdate as re- agent 98, 232 preparation and properties of.. 226, 227 pyroantimonate 73, 80 pyrophosphate with copper and cadmium 107 salicylate test for nitric acid ... 281 sulphide, preparation of 308 thiosulphate on cupric salts.... 108 thiosulphate with antimony salts 78 thiosulphate with third group metals 145 Solids, conversion into liquids.... 366 decomposition upon ignition. 370, 371 effect on ignition with cobalt nitrate 372 preliminary examination of 363 separation of 17 table for preliminary examina- tion 370 Solubility, degrees of 15, 16 Solubility-product 23 Solutions, conversion into solids.. 367 Solution and ionization 20-24 Solvay soda process 260 Sonnenschein's reagent 98 Stannic salts, solubilities 84 sulphide, formation and proper- ties of 86 Stannite, alkali, as a test for bis- muth 103 Stannous chloride on mercury salts 43 chloride as a reducing agent 88 chloride with gold salts 93 chloride with molybdic acid 99 salts, distinction from stannic salts 123 salts, solubilities 84 salts with silver nitrate 87 salts with sulphurous acid 86 sulphide, formation and proper- ties 85 Stibine, decomposition by soda lime 79 formation of 79 reaction with fixed alkali hy- droxides 79 reaction with silver nitrate 79 separation from arsine -65 Strontium 208-210 detection of 210, 213 estimation of 210 hydroxide, formation 208 occurrence of .......,.,,......•• 208 INDEX. 419 PAGE Strontium, preparation and prop- erties of 208 sulphate, distinction from 0aSO 4 209 sulphate, separation from BaS0 4 209 Sulphates, detection and estima- tion of 326 ignition of 325 preparation of 322 reduction by ignition with carbon 249 solubilities of 323 Sulphites, detection of 321 distinction from sulphates 321 estimation of 321 ignition of 321 interference in test for oxalates. 390 preparation of 318 separation from sulphates by Ba salts 207 solubilities of 319 Sulphides, detection and estima- tion of 311 formation of 307 ignition of 310 reactions of 309, 310 solubilities of 28, 308 Sulphur 304-306 combinations on ignition of 306 detection and estimation of 306 formation of 304 in the tin group 118 occurrence of 304 oxidation by reagents 305, 306 oxides 304 precipitation of 53, 114, 115 preparation and properties of . . . 304, 305 reactions in forming sulphides.. 305 relations of 9 separating copper from cadmium 107 solubilities of 305 Sulphuric acid 321-326 detection in presence of sul- ^^ phates 326 formation and occurrence of.... 322 properties of 321 reactions with 323, 324, 325 separation from Se 139 separation from Fe 137 anhydride, preparation of 322 Sulphurous acid 318-321 on arsenic acid 60 and sulphites as reducers 320 occurrence of 3 1 § preparation and properties of.. 318 formation of 3i8 PAGE Sulphurous acid, reduction of cupric salts 108 solubilities of 319 on stannous salts S6 Synthesis, problems in 397 Table for acids as precipitated by barium and calcium chlorides. 386 for acids precipitated by silver nitrate 387 for acids, preliminary 378 for analysis in presence of phos- phates by the use of alkali ace- tates and ferric chloride 191 for analysis in presence of phos- phates by use of ferric chloride and barium carbonate 192 for analysis of the Silver Group (first) 52 for analysis of the Copper Group (second) 124 for analysis of the Tin Group (second) 116 for analysis of the Iron Group (third) 160 for analysis of the Zinc Group (fourth) 183 for analysis of the Calcium Group (fifth) 217 of grouping of the metals 375 of separations of the metals.... 376 of separation of the ammonium sulphide precipitates of the Iron and Zinc Groups 187 of solubilities 398 Tannic acid with iron salts 154 Tantalum, distinction from silica. 198 distinction from titanium 198 properties and reactions of 198 separation from columbium 198 Tartar emetic, composition of 252 Tartaric acid 252-254 in detection of potassium 223 distinction from citric acid 251 formation and properties 252 Tartrate calcium, deportment with water 253 detection of 253 distinction from citrates 253 distinction from oxalates. ...253, 389 estimation of • • 254 Tartrates, ignition 253 reactions 253 solubilities 252 Tellurium 137-138 420 INDEX. PAGE Tellurium, distinction from sele- nium 138, 140 properties and reactions of 137 separation from sulphuric acid. . 137 Tenorite 104 Terbium 198-199 Tetrathionic acid, formation and properties 315 Thallious iodide 199 Thallium, properties and reac- tions 199 Thioacetate in formation of sul- phides 307 Thioeyanates, reactions with 272 Thiocyanic acid as a reducer 273 properties of 272 Thionic acids, table of compari- sons 317 Thiosulphates, detection of 313 distinction from sulphates and sulphites 314 estimation of 314 ignition of 313 formation and properties of. . . . 312 Thiosulphuric acid 312-314 Third group reagents 141 Thorium 199-200 Tin 82-89 creaking of 82 detection of 88, 122, 367 estimation of 88 Group, metals of 56 Group, separation from Copper Group '. 115 Group, sulphides with (NH 4 ) 2 S X 115 occurrence of 82 oxidation of r 88 oxides and hydroxides 82 preparation and properties of... 82 notes on the analysis of 123 relation to Nitrogen Family. ... 7 reduction by ignition 87 salts with the alkalis 84 salts with hydrosulphuric acid. . 85 separation from antimony 81 separation from antimony sul- phides 121 separation from arsenic 118 solubilities of 83 sulphides, colloidal 115 PAGE Tin with antimony and with arsenic 87 Titanium .200-201 distinction from columbium 201 distinction from tantalum 198 properties and reactions' of 200 separation from thorium 200 Trithionic acid, formation and properties 315 Tungsten, properties and reac- tions 134 TurnbulPs blue .,. '. 155 Unit of quantity 22 Uranium, properties and reactions 201 Urea, from ammonium cyanate. . . 271 Valence, negative 3 Vanadium 135-136 Volatile alkali (ammonium hy- droxide) 221 Water, action on bismuth salts... 101 action on antimonous salts 75 Welsbach burners 203 ■Wolframium (tungsten) 134 Wulfenite 97 Ytterbium, properties and reac tions 202 Yttrium 202 Zincates, formation of 179 Zinc 17S-181 detection and estimation of 180 Family 5 granulated 63, 178 Group, table for analysis 183 Group, comparative reactions... 182 hydroxide and oxide 178 ignition of 180 occurrence of 178 oxidation of 181 platinized 178, 243 preparation and properties 178 reduction of 181 salts, solubilities and reactions of 179 sulphide, formation in presence of acetic acid 179 Zirconium 202-203 APPENDIX. The methods described in this Appendix have been in use in the Uni- versity of Michigan for the past two years, and have proved to be thoroughly reliable. They are based fundamentally on older methods, which have, however, been so modified as to make them more satisfactory for qualitative tests. The modifications are due to Mr. Hobart H. Willard, Instructor in Qualitative Analysis in this University. APPENDIX. SEPAKATION OP AKSENIC, ANTIMONY, AND TIN. These metals are precipitated from a N/5 HCI solution by H 2 S and therefore belong to the 2nd group, but their sulphides, unlike those of Pb , Cu , Cd , Bi , are soluble in yellow ammonium sulphide, (NHJ.sS*, and are thus separated from them before treatment with 2NHN0 3 . This operation is necessary only when both divisions of the group are present, and is to be avoided when unnecessary. Hence a little of the 2nd group precipitate is tested by warm- ing with 1 or 2 cc. (NH 4 ) 2 S X . If it all dissolves, only As, Sb , Sn can be present; if nothing dissolves, none of these can be present; if pari dissolves, then the whole 2nd group precipitate must be so treated. To see if anything has dissolved in the (NH i ) ! S ll it is acidified slightly with HCI (test with litmus) ; a milky, white precipitate of S will always be formed, but if any sulphides are present they will appear as a flocculent, colored precipitate. If the whole 2nd group precipitate is treated with (NH 4 ) 2 S X , the solution is filtered and acidified just as the test portion was, the precipitated sulphides of As , Sb , Sn well washed with hot water and removed from the filter to a casserole by a spatula, or, if the amount is small, treated with the filter; a convenient amount of concentrated HCI (sp. gr. 1.2) is added and boiled a minute or two to expel H 2 S . The sulphides of Sb and Sn are dissolved to form the chlorides SbCI 3 and SnCI 4 while the As 2 S 3 is hardly attacked. Since the strong acid attacks the filter, the solution is diluted a little, which should cause no reprecipitation if all H 2 S was expelled, filtered, and the residue well washed. It "may be either As 2 S 3 and S, or S alone. A few cc. of warm NH 4 OH are poured over it, the solution being passed through again if necessary. The As 2 S 3 dissolves and the S remains. To the solution, which must be clear, add 1 or 2 cc. H 2 2 , 2 to 3 cc. NH 4 CI , and 2 to 3 cc. " magnesia mixture," which is MgCI 2 -f NH 4 CI -f NH.OH . Cool, and let stand for a time. The As is precipitated as NH 4 MgAs0 4 , a white, crystalline precipitate exactly like NH 4 MgP0 4 in appearance. This "magnesia mixture" is not intended to take the place of the "Marsh apparatus" but to confirm its results 2H 3 As0 3 + 3H 2 S + [HCI1 = As 2 S 3 + 6H 2 . 2H 3 As0 4 + 5H 2 S + [ HCI] = As 2 S 3 + S 2 + 8H 2 . 2As 2 S 3 + 6(NH 4 ) 2 S x = 4(NH 4 ) 3 AsS 4 + (3x — 5)S 2 . 4(NH 4 ) 3 AsS 4 + 6HCI = As 2 S 5 + 6NH 4 CI + 3H 2 S . As 2 S 6 + HCK12N, hot) =no action. As 2 S 5 + 16 N H 4 H + 20 H 2 2 = 2 ( N H 4 ) 3 As0 4 + 5 ( N H 4 ) 2 S0 4 + 28 H 2 . (NH 4 ) 3 As0 4 +MgCI 2 +[NH 4 OH + NH 4 CI] = MgNH 4 As0 4 + 2NH 4 CI. The filtrate from As 2 S 3 is to be tested for Sb and Sn . For the Sb , place a few drops on a clean silver coin; it should produce no discoloration. A piece of tin, bent into the shape of a broad U, is now placed on the coin so that one end is in the center of the drop and the other in contact with the silver outside Allow to stand about 5 minutes. If Sb is present it will be deposited as a brown spot on the silver covered by the drop, the Sn and Ag acting as a 424 APPENDIX. galvanic couple to reduce the Sb " " to metal. Another test consists in treating the solution with pure, fine Fe wire, the Sb being precipitated in black metallic form, while the Sn " ' " is merely reduced to Sn " " but not precipitated. Test the rest of the solution for Sn by heating with fine Fe wire until the solution is colorless or greenish, with no trace of yellow, to make sure that all the Sn " " is reduced to Sn " " . Ten minutes or more may be required. Filter and add the nitrate slowly (a few drops at a time) to a few cc. of ammo- nium molybdate, (NH 4 ) 2 Mo0 4 solution. A deep blue color or precipitate will appear if Sn " ' is present, due to the reduction of the MoO s to a lower oxide. Or, instead of adding this nitrate to molybdate solution, it may be treated with HgCI 2 , a white precipitate of HgCI being formed if Sn " is present. Note that this is reversing the test for Hg" with SnCI 2 . Remember that it is always necessary to reduce the Sn ' " " to Sn " " with Fe , since the former cannot reduce the HgCI 2 or (NH 4 ).,Mo0 4 . The HgCI 2 test is most characteristic. The precipi- tation of As 2 S 3 , unlike that of the other sulphides, is not prevented by the pres- ence of any amount of HCI , however large, but, on the contrary, is aided. It may, therefore, be necessary, after removing all other sulphides in the N/5 HCI solution, to add several cc. of concentrated HCI , heat to boiling, and" pass in H 2 S for some time to precipitate the rest of the As . In the cold, H 3 As0 4 is very slowly precipitated by H 2 S , but strong HCI and heat accelerate the reaction "very much. It is essential that the sulphides be thoroughly washed before treatment with HCI . CuS is slightly soluble in (NH 4 ) 2 S X and may give a coloration when the solu- tion is acidified. (NHJjS, which is colorless, gives no precipitate of S upon addition of excess of acid; (NHJ 2 S X , yellow, always gives more or less S, white and difficult to filter. 2(NH 4 ) 2 S„ + 4HCI = 4NH 4 CI + 2H 2 S+(x — 1)S 2 . (NH 4 ) 2 S + 2HCI = 2NH 4 CI + H 2 S. Make a blank test on the Fe wire used, to see that its solution in HCI gives no test for Sn with molybdate. Additional Tests for Bismuth. — First. — Dissolve the separated hydroxide in hydrochloric acid, then one drop of iodide of potassium will produce a black color and a larger quantity forms a yellow solution not decolorized by stannous chloride. A very delicate test. The yellow color of ferric salts and the green of copper salts are removed by stannous chloride. Second. — To the separated bismuth hydroxide add sodium hydroxide and for- maldehyde. Metallic bismuth is produced. 2Bi(OH) s + 3HCHO + 3NaOH = 2Bi -f-3HCOONa-]- 6H 2 0. Since sodium stannite reduces lead hydroxide to metallic lead and formaldehyde does not, it is preferred as a reducer. Owing to imperfect separation traces of lead might be present at this point. Separation of Third Group Cations, Al ' ", Fe ' ", Cr " '. Boil the nitrate from the second group to expel H 2 S ; any Fe present is in the ferrous condition. Since the precipitation of Fe " by NH 4 OH is not quite com- plete, it must be oxidized to Fe " '. To effect this, add 5 to 10 drops HNO s and boil till there is no further change of color, even upon adding another drop of acid. A slight, white precipitate here is sulphur from oxidation of H 2 S not APPENDIX. - 425 expelled by boiling. Add to the solution half its volume of NH 4 CI , heat to boil- ing, and add NH.OH, two or three drops at a time, till there is a slight excess. Boil a minute or two, and at once filter and wash. If precipitated in a boiling solution the hydroxides filter more rapidly. The precipitate consists of Fe(OH) 3 , AI(OH) 3 , Cr(OH) s , and more or less Mn(OH) 3 ; sometimes all the Mn is precipi- tated here. The more rapidly the precipitation and filtration are performed, the less Mn(OH), will be included; hence never delay at this point. Wash the precipitate with hot water, stirring it up with the jet from the wash-bottle. If the filtrate is amethyst colored and no Co " is found, too much NH 4 OH was added, and a trace of Cr(OH) s has dissolved. Further boiling may precipitate this, but often the color remains. If the third group precipitate is greenish or black and no Cr - " ' is present, it shows that not all the Fe " " was oxidized, and the precipitate must be dissolved in HCI and again treated with HN0 3 . The precipitate is transferred to a casserole or beaker, either by means of a spatula or with a fine jet of water, but any considerable dilution is to be avoided. Add 5 to 10 cc. NaOH , 10 to 15 cc. H 2 2 and boil for several minutes — until effervescence ceases. The H 2 2 oxidizes the Cr " ' ' to CrO/', while the Al ■ ■ " goes into solution as the ion AI0 2 '. The precipitate then consists of Fe(OH) 3 and some Mn(OH) 3 . Filter and wash with hot water. The filtrate, yellow if Cr is present, and containing NaAI0 2 and Na 2 Cr0 4 , is acidified with dilute H 2 S0 4 . (Test with litmus.) This converts the AI0 2 ' ion into Al ' " ' and Cr0 4 " into Cr 2 0/', the latter having an orange color. Divide the solution into two parts; cool one under running water and add a little H 2 2 ; a dark and beautiful blue color, due to perchromic acid, HCr0 4 , is formed if Cr is present; the color soon disappears, becoming green, and the more rapidly the warmer the solution. It is not formed at all in a warm solution. To the other portion add excess of (NH 4 ) 2 C0 3 , warm, and if no precipitate appears at once set aside a few minutes. A white, gelatinous, flocculent precipitate of AI(OH) 3 will be formed if Al " ' " is present. This precipitate is sometimes almost trans- parent; it is nearly always very light and liable to escape notice. The precipitate of Fe(OH) 8 and Mn(OH) 3 is dissolved in dilute H 2 S0 4 with the addition of a few drops of H 2 2 if solution does not readily occur, and a few drops of this solution are tested for Fe with KCNS or K 4 Fe(CN) . Test the solution for Win " by adding Pb s 4 and HN0 3 , boiling several minutes and allow- ing to settle without filtering. If it shows a reddish purple color, this is due to the presence of HMn0 4 formed by the oxidation of the Mn". Test the Pb 3 4 for Mn by boiling it with HN0 3 ; some samples give a slight color. If, in the original solution, Cr 2 7 " or Cr0 4 " is present, it must be reduced to Cr'" before precipitating the second group. Add considerable HCI and several cc. of alcohol, and boil till the color changes to green. Be sure that reduction is complete. Boil down to expel all the alcohol, then neutralize and proceed in the regular way for the second group. If not reduced by alcohol and HCI , H 2 S will at once give a green color, a large amount of sulphur being liberated, difficult to filter. Do not fail to test the reagent NaOH and H 2 2 for Al . The fact that a solution is green does not exclude the possibility of some Cr 2 0," being present; but a very small amount of Cr " ' ' will give a distinct color to the orange Cr 2 7 " solution. If further proof of the presence of Cr " ' is desired, the Cr 2 7 " may be precipitated by making alkaline with Na 2 CO„, adding excess of acetic acid (this is to avoid the presence of free mineral acid), and then BaCI 2 426 APPENDIX. or Pb(C 2 H s 2 ) 2 . Cr""" remains in the filtrate, the anion being precipitated as BaCr0 4 or PbCr0 4 . The reagent NaOH may contain A I or silica, which might give the test if no Al ' ' •" was present in the solution being analyzed. Test, therefore, about 10 cc. of it by acidifying with HCI and adding excess of (NH 4 ) 2 C0 3 . Note amount of precipitate formed, and do not report A I " ' unless a heavier precipitate than this is obtained in the analysis. This is known as a " blank test "; it is often neces- sary to apply such checks in analytical work, e. g., testing the Pb 3 4 for Mn . Strong solutions of NaOH or KOH attack and soften filter paper, interfering with filtration. If this difficulty is experienced, dilute with water. Remember that it is impossible to keep a solution of Fe ' ' free from Fe ' " ' unless all oxygen is carefully excluded, so that the latter is always associated with the former. NEVER FAIL to test the precipitate remaining after treatment with NaOH and H 2 O s for Mn ; it may all appear here and none in the fourth group. Analytical Equations. Cr'" Reduce Cr 2 0," or CrO," to Cr ' " ' thus: K 2 Cr 2 7 + 3C 2 H„0 + 8HCI = 2CrCI 3 + 3C 2 H 4 + 2KCI + 7H 2 . CrCI 3 + 3NH 4 OH=Cr(OH) 3 -f 3NH 4 CI. 2Cr(OH) 3 + 4NaOH-}-H 2 2 = 2Na 2 Cr0 4 + 8H 2 0. 2Na 2 Cr0 1 + 4H 2 S0 4 +H 2 2 = 2HCr0 4 + 4NaHS0 1 + 2H 2 0. (Cold solution.) 2Na 2 Cr0 4 -f 2HN0 3 = Na 2 Cr,0 7 + 2NaN0 3 -f. H 2 . Na 2 Cr 2 7 + 2NaOH = 2Na 2 Cr0 4 -f H.,0 . 2Na 2 Cr0 4 + 5H 2 S0 4 + 3H 2 2 (warm) =2Na 2 S0 4 + Cr 2 (SO.) 3 +30 2 + 8H 2 0. Separation of Nickel and Cobalt. Dissolve the CoS and NiS in HCI with a small crystal of KCI0 3 , boil, filter, add two cc. of NH 4 CI , and NH 4 OH in large excess, at least ten cc. more than enough to make the solution alkaline. Dilute to at least 25 cc. in a beaker or casserole, add .2 to .3 gram potassium persulphate, K 2 S 2 O s , and warm, with fre- quent stirring, until crystals dissolve. Boil for some time to expel most of the ammonia. The blue color of the (NH 3 ) 4 NiCl 2 does not change, but the cobaltous compound is oxidized to a complex cobaltic ammonium salt, the color changing to red. There should be no precipitate in the solution now (if there is, either the solution was not dilute enough or too little NH 4 OH was added. Dissolve in HCI and repeat process without adding more NH 4 CI). Add a few more crystals K 2 S 2 O s and boil again to make sure that oxidation is complete. "When no further change of color occurs, cool under the tap to room temperature, add 10 to 15 cc. NaOH and shake. If Ni is present it will be precipitated as Ni(OH) 3 , dark brown, turning black. The precipitate may form slowly, and requires some time for complete separation. Let stand at least 15 minutes; filter, and if Co is present, the filtrate will be pink or red. The amount of NaOH necessary to precipitate the Ni depends on the amount of NH 4 CI and NH,OH present; if a very large excess is present, more than 10 cc. NaOH may be required. No Co(OH), is precipitated unless the solution is warm. APPENDIX. 427 Separation of CI', Br', and I', by the Persulphate Method. To ten cc. of the original solution, add slight excess of Na 2 C0 3 , free from chlorine, and boil, to precipitate the heavy metals. The solution must react alkaline. Filter and add to the filtrate acetic acid, several cc. more than enough to neutralize it, dilute to 50-60 cc, add about one-half gram of K 2 S 2 8 , and heat. If an iodide is present, free iodine -will be liberated, and may be identified by shaking a few drops of the solution with CS 2 . Boil in a casserole until all iodine is expelled, which should require three to four minutes. If action is slow, more persulphate should be added. When the solution is colorless, add a few more crystals of persulphate and boil again, to make sure that no iodine remains. As the solution evaporates add distilled water to maintain the original volume. To remove Br' add two cc. of H 2 SO lr previously diluted with water, a little more K 2 S 2 8 , and heat to boiling point, but do not 'boil. A yellow or red coloration, if the separation of I has been properly conducted, indicates Br . Pour a little of the solution into a test tube, cool, and shake with CS 2 , which should be colored yellow or red but not violet, which would indicate that the I had not been com- pletely removed. If bromine is present, add one-half gram of K 2 S 2 O s to the main part of the solution, and boil until it is all expelled and the solution is colorless; then test with a little more K 2 S 2 O s and boil five minutes longer to make sure of the complete expulsion of the bromine. BE SURE THAT THE VOLUME OF THE SOLUTION DOES NOT FALL BELOW 50 TO 60 cc. Add distilled water from time to time to replace that lost by evaporation. When all bromine is removed, cool and add a few drops of silver nitrate; a white, curdy precipitate of silver chloride indicates the presence of CI . If too much silver nitrate is added, a white crystalline precipitate may be formed, but will dissolve upon dilution and warming. If ClO'j is present, the above procedure cannot be followed, for the I' would be oxidized to IO' 3 . In this case it is necessary to precipitate the CI', Br', and I' by adding to the original solution excess of silver nitrate and then nitric acid; this effects a separation, silver chlorate being soluble. Wash the precipitate of AgCI , AgBr , Agl , transfer to a test tube, add a piece of zinc, a little water, and a drop of sulphuric acid. Let it stand until it is perfectly black all the way through, showing complete reduction to metallic silver. Filter and treat the filtrate containing ZnCI 2 , ZnBr 2 , Znl 2 , according to the above method, starting at the beginning. Even if no heavy metals are present, Na 2 C0 3 should be added to neutralize any mineral acid that may be present and to form some sodium acetate when acetic acid is added. The persulphate method should be used only when the presence of I' or Br' has been proved by some short test (H 2 S0 4 , CI , HN0 2 , HN0 3 , or other oxidizer). In presence of a great excess of Br', CuSO,, KN0 2 , or HgCI is an excellent test for I'. Equations. 2KI + K 2 S 2 9 = 2K 2 S0 4 + l 2 . 2KBr + K 2 S 2 8 + H 2 SO, = 2K 2 S0 4 + Br 2 + H 2 SO< . KCI -f K 2 S 2 0„ + H 2 SO« (1.5— 2N) =No action. 2Agl + 2AgBr + 2AgCI + 3Zn = 6Ag + Znl 2 + ZnBr 2 + ZnCJ, . 428 APPENDIX. COKKECTIONS IN ATOMIC WEIGHTS, MADE BY THE INTEKNATIONAL COMMITTEE, JANUAEY, 1905. Journal of the American Cbemical Society, Vol. XXVII. PAGE. 72— Antimony 120.2 100— Bismuth , ., 208.5 342— Brpmine 79.96 193— Cerium ..' 140.25 193— Columbium 94. 288— Fluorine 19. 196— Indium 115. 350— Iodine 126.97 132— Iridium 193. 197 — Lanthanum .' 138.9 29 — Lead 206.9 214— Magnesium 24.36 131— Palladium 106.5 93— Platinum 194.8 PAGE. 222 — Potassium 39.15 234 — Rubidium 85.5 45— Silver 107.93 304— Sulphur 32.06 198 — Tantalum 183. 137— Tellurium *. 127.6 199— Thallium 204.1 199— Thorium 232.5 200 — Titanium 48.1 201 — Uranium 238.5 135— Vanadium 51.2 202— Ytterbium 173. 202 — Zirconium 90.6 Catalogue of Scientific Publications and Importations of the D. 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From the German of Prof. Ludwig Spangenburg, with a Preface by S. H. Shreve, A.M. No. 24. A PRACTICAL TREATISE ON THE TEETH OF WHEELS. By Prof. S. W. Robinson. 2nd edition, revised, with additions. No. 25. THEORY AND CALCULATION OF CANTILEVER BRIDGES. By R. M. Wilcox. No. 26. PRACTICAL TREATISE ON THE PROPERTIES OF CON- tinuous Bridges. By Charles Bender, C.E. No. 27. BOILER INCRUSTATION AND CORROSION. By F. J. Rowan. New edition. Revised and partly rewritten by F. E. Idell. No. 28. TRANSMISSION OF POWER BY WIRE ROPES. By Albert W. Stahl, U.S.N. Second edition, revised. No. 29. STEAM INJECTORS, THEIR THEORY AND USE. Trans- lated from the French of M. Leon Pochet. No. 30. MAGNETISM OF IRON VESSELS AND TERRESTRIAL Magnetism. By Prof. Fairman Rogers. D. VAN NOSTRAND COMPANY'S No. 31. THE SANITARY CONDITION OF CITY AND COUNTRY Dwelling-houses. By George E. Waring, Jr. Second edition, revised. No. 32. CABLE-MAKING FOR SUSPENSION BRIDGES. By W. Hildenbrand, C.E. No. 33. MECHANICS OF VENTILATION. By George W. Rafter, C.E. Second edition, revised. No. 34. FOUNDATIONS. By Prof. Jules Gaudard, C.E. Trans- lated from the French. Second edition. No. 35- THE ANEROID BAROMETER: ITS CONSTRUCTION AND Use. Compiled by George W. Plympton. Ninth edition, revised and enlarged. No. 36. MATTER AND MOTION. By J. Clerk Maxwell, M.A. Second American edition. No. 37. GEOGRAPHICAL SURVEYING: ITS USES, METHODS, and Results. By Frank De Yeaux Carpenter, C.E. No. 38. MAXIMUM STRESSES IN FRAMED BRIDGES. By Prof. William Cain, A.M., C.E. New and revised edition. No. 39. A HANDBOOK OF THE ELECTRO-MAGNETIC TELE- graph. By A. E. Loring. Fourth edition, revised. No. 40. TRANSMISSION OF POWER BY COMPRESSED AIR. By Robert Zahner, M.E. New edition, in press. No. 41. STRENGTH OF MATERIALS. By William Kent, C.E., Assoc. Editor "Engineering News." Second edition. No. 42. THEORY OF STEEL-CONCRETE ARCHES, AND OF Vaulted Structures. By Prof. Wm. Cain. Third edition, thoroughly revised. 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With diagrams and folding plates. By Thomas Nolan. Second edition, revised and enlarged. No. 52. IMAGINARY QUANTITIES: THEIR GEOMETRICAL IN- terpretation. Translated from the French of M. Argand bv Prof. A. S. Hardy. No. 53. INDUCTION COILS: HOW MADE AND HOW USED. Eleventh American edition. No. 54. KINEMATICS OF MACHINERY. By Prof. Alex. B. W. Kennedy. With an introduction by Prof. R. H. Thurston. No. ss. SEWER GASES: THEIR NATURE AND ORIGIN. By A. de Varona. Second edition, revised and enlarged. No. 56. THE ACTUAL LATERAL PRESSURE OF EARTHWORK. By Benj. Baker, M. Inst., C.E. No. 57. INCANDESCENT ELECTRIC LIGHTING. A Practical De- scription of the Edison System. By L. H. Latimer. To which is added the Design and Operation of Incandescent Sta- tions, by C. J. Field; and the Maximum Efficiency of Incandescent Lamps, by John W. Howell. No. 58. VENTILATION OF COAL MINES. By W. Fairley, M.E., and Geo. J. Andrg. No. 59. RAILROAD ECONOMICS; OR, NOTES WITH COMMENTS. By S. W. Robinson, C.E. 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EXPLOSIVE MATERIALS. By Lieut John P. Wisser. No. 71. DYNAMIC ELECTRICITY. By John Hopkinson, J. A. Shoolbred, and R. E. Day. No. 72. TOPOGRAPHICAL SURVEYING. By George J. Specht. Prof! A. S. Hardy, John B. McMaster, and H. F. Walling. Third edition, revised. No. 73. SYMBOLIC ALGEBRA; OR, THE ALGEBRA OF ALGE- braic Numbers. By Prof. William Cain. No. 74. TESTING MACHINES: THEIR HISTORY, CONSTRUC- tion and Use. By Arthur V. Abbott. No. 75. RECENT PROGRESS IN DYNAMO-ELECTRIC MACHINES. Being a Supplement to "Dynamo-electric Machinery." By Prof. Sylvanus P. Thompson. No. 76. MODERN REPRODUCTIVE GRAPHIC PROCESSES. By Lieut. James S. Pettit, U.S.A. No. 77. STADD\ SURVEYING. The Theory of Stadia Measure- ments. By Arthur Winslow. Sixth edition. No. 78. THE STEAM-ENGINE INDICATOR AND ITS USE. By W. B. Le Van. No. 79. THE FIGURE OF THE EARTH. By Frank C. Roberts, C.E. No. 80. HEALTHY FOUNDATIONS FOR HOUSES. By Glenn Brown. No. 81. WATER METERS: COMPARATIVE TESTS OF ACCURACY, Delivery, etc. Distinctive features of the Worthington, Ken- nedy, Siemens, and Hesse meters. By Ross E. Browne. No. 82. THE PRESERVATION OF TIMBER BY THE USE OF ANTI- septics. By Samuel Bagster Boulton, C.E. No. 83. MECHANICAL INTEGRATORS. By Prof. Henry S. H. Shaw, C.E. No. 84. FLOW OF WATER IN OPEN CHANNELS, PIPES, CON- duits, Sewers, etc. With Tables. By P. J. Flynn, C.E. No. 85. THE LUMINIFEROUS .ETHER. By Prof. De Volson Wood. No. 86. HANDBOOK OF MINERALOGY: DETERMINATION, DE- scription, and Classification of Minerals Found in the United States. By Prof. J. C. Foye. Fifth edition, revised. SCIENTIFIC PUBLICATIONS. No. 87. TREATISE ON THE THEORY OF THE CONSTRUCTION of Helicoidal Oblique Arches. By John L. Culley, C.E. No. 88. BEAMS AND GIRDERS. Practical Formulas for their Resist- ance. By P. H. Philbrick. No. 89. MODERN GUN COTTON: ITS MANUFACTURE, PROP- erties, and Analyses. By Lieut. John P. Wisser, U.S.A. No. 90. 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HOW TO BECOME AN ENGINEER; or. The Theoretical and Practical Training necessary in Fitting for the Duties of the Civil Engineer. By Prof. Geo. W. Plympton. No. 1 01. THE SEXTANT, and Other Reflecting Mathematical Instru- ments. With Practical Hints for their Adjustment and Use. By F. R. Brainard, U. S. Navy. No7io2. THE GALVANIC CIRCUIT INVESTIGATED MATHE- matically By Dr. G. S. Ohm, Berlin, 1827. Translated by William Francis. With Preface and Notes by the Editor, Thomas D. Lockwood, M.I.E.E. D. VAN NOSTRAND COMPANY'S No. 103. THE MICROSCOPICAL EXAMINATION OF POTABLE Water. With Diagrams. By Geo. W. Rafter. Second edition. No. 104. VAN NOSTRAND'S TABLE-BOOK FOR CIVIL AND ME- chanieal Engineers. Compiled by Prof. Geo. W. Plympton. No. 105. DETERMINANTS. An Introduction to the Study of, with Examples and Applications. By Prof. G. A. Miller. No. 106. COMPRESSED AIR. Experiments upon the Transmission of Power by Compressed Air in Paris. (Popp's System.) By Prof. A. B.. W. Kennedy. The Transmission and Distribution of Power from Central Stations by Compressed Air. By Prof. W. C. Unwin. Edited by F. E. Idell. Third edition. No. 107. A GRAPHICAL METHOD FOR SWING BRIDGES. A Rational and Easy Graphical Analysis of the Stresses in Ordinary Swing Bridges. With an Introduction on the General Theory of Graphical Statics, with Folding Plates. By Benjamin F. La Rue. No. 108. SLIDE-VALVE DIAGRAMS. A French Method for Con- structing Slide-valve Diagrams. By Lloyd Bankson, B.S., Assistant Naval Constructor, U. S. Navy. 8 Folding Plates. No. 109. THE MEASUREMENT OF ELECTRIC CURRENTS. Elec- trical Measuring Instruments. By James Swinburne. Meters for Electrical Energy. By C. H. Wordingham. Edited, with Preface, by T. Commerford Martin. With Folding Plate and Numerous Illustrations. No. no. TRANSITION CURVES. A Field-book for Engineers, Con- taining Rules and Tables for Laying out Transition Curves. By Walter G. Fox, C.E. No. in. GAS-LIGHTING AND GAS-FITTING. Specifications and Rules for Gas-piping. Notes on the Advantages of Gas for Cooking and Heating, and Useful Hints to Gas Consumers. Third edition. By Wm. Paul Gerhard, C.E. No. 112. A PRIMER ON THE CALCULUS. By E. Sherman Gould, M. Am. Soc. C. E. Third edition, revised and enlarged. No. 113. PHYSICAL PROBLEMS and Their Solution. By A. Bour- gougnon, formerly Assistant at Bellevue Hospital. Second ed. No. 114. MANUAL OF THE SLIDE RULE. By F. A. Halsey, of the "American Machinist." Third edition, corrected. No. 115. TRAVERSE TABLE. Showing the Difference of Latitude and Departure for Distances Between 1 and 100 and for Angles to Quarter Degrees Between 1 Degree and 90 Degrees. (Reprinted from Scribner's Pocket Table Book.) SCIENTIFIC PUBLICATIONS. No. 116. WORM AND SPIRAL GEARING. Reprinted from "Ameri- can Machinist." By F. A. Halsey. Second revised and enlarged edition. No. 117. PRACTICAL HYDROSTATICS, AND HYDROSTATIC FOR- mulas. With Numerous Illustrative Figures and Numerical Examples. By E. Sherman Gould. No. 118. TREATMENT OF SEPTIC SEWAGE, with Diagrams and Figures. By Geo. W. Rafter. No. 119. LAY-OUT OF CORLISS VALVE GEARS. With Folding Plates and Diagrams. By Sanford A. Moss, M.S , Ph.D Re- printed from "The American Machinist,'' with revisions and additions. Second edition. No. 120. ART OF GENERATING GEAR TEETH. By Howard A. Coombs. With Figures, Diagrams and Folding Plates. Re- printed from the "American Machinist." No. 121. ELEMENTS OF GAS ENGINE DESIGN. Reprint of a Set of Notes accompanying a Course of Lectures delivered at Cornell University in 1902. By Sanford A. Moss. Illustrated. No. 122. SHAFT GOVERNORS. By W. Trinks and C. Housum. Il- lustrated. No. 123. FURNACE DRAFT; ITS PRODUCTION BY MECHANICAL Methods. A Handy Reference Book, with figures and tables. By William Wallace Christie. Illustrated.