MEDICAL ollege of Pharmacy California College of Pharmacy QUALITATIVE CHEMICAL ANALYSIS A GUIDE IN QUALITATIVE WORK, WITH DATA FOR ANALYTICAL OPERATIONS AND 1 LABORATORY METHODS BY ALBERT B. PRESCOTT AND OTIS C. JOHNSON PROFESSORS IN THE UNIVERSITY OF MICHIGAN SEVENTH EDITION, THOROUGHLY REVISED BY JOHN C. OLSEN, A.M., PH.D. Professor of Chemicai Engineering, Polytechnic Institute, Brooklyn, N. Y. Author of "Quantitative Chemical Analysis" Editor, Van Nostrand's "Chemical Annual" California College of Pharmacy NEW YORK D. VAN NOSTRAND COMPANY Eight Warren Street 1920 Copyright, 1916, 1917 BY D. VAN NOSTRAND COMPANY Lo PREFACE TO THE SEVENTH EDITION. Since the last revision of this very comprehensive text on Qualitative Analysis was issued, a considerable number of new methods of analysis have been published, and some of these have been found to be valuable and have come into more or less general use. Many of the new methods have been tried out and some have been found to be unreliable unless extreme care is taken or only pure solu- tions used. The attempt has been made to include in this revision only those methods which have been found to be reliable. A very marked change has also taken place in the method of presen- tation of chemical reactions. Ionic reactions are given in many texts to the exclusion of molecular reactions. Physical chemical theories are also frequently presented and discussed at great length. The attempt has been made in the revision of this text to present only briefly the modern conceptions of solution, leaving a fuller presenta- tion for the lecturer or separate texts on Physical Chemistry. Mole- cular reactions have also been largely retained in the belief that the material to be analyzed by the chemist in practical work is quite as often a molecular compound as an ion. All molecular weights, solubilities and other constants of the elements and their principal compounds have been brought up to date. The principal minerals, methods of preparation and determinations of ele- ments, as well as the reactions, have been revised, references being given to the literature as heretofore. In general, the attempt has been made to retain the excellent features of this text which have given it such an extended use in the past, both as a class room and as a reference text, while adding the valuable results of recent progress in the science. Acknowledgment is made with pleasure of valuable assistance ren- dered the undersigned in preparing the revision to Professor R. J. PREFACE. Colony of Cooper Union, who revised the sections on the properties, oc- currence and preparation of the metals ; to Mr. Wm. H. Ulrich, who tried out many of the new methods; to Mr. N. F. Borg, who contributed much of the revision of the acids ; and to Mr. M. P. Matthias, who has revised the Index. J. C. OLSEN October 2, 1916. PREFACE TO THE FIFTH EDITION. In this, the fifth full revision of this manual, the text has been re- written 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 sub- ject-matter, omitted in 1888, is now desired by teachers. For the portion upon Solution and lonization, 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 v VI 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 estimation 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, 1901. CONTENTS. PART I. THE PRINCIPLES OF ANALYTICAL CHEMISTRY. PAQH THE CHEMICAL ELEMENTS AND THEIR ATOMIC WEIGHTS 1 TABLE OF THE PERIODIC SYSTEM OP THE CHEMICAL ELEMENTS 2 DISCUSSION OF THE PERIODIC SYSTEM 3 CLASSIFICATION OF THE METALS AS BASEB 10 COMMONLY OCCURRING ACIDS 13 THE OPERATIONS OF ANALYSIS 13 SOLUTION AND IONIZATION 20 ORDER OF LABORATORY STUDY 25 PART II. THE METALS. THE SILVER AND TIN AND COPPER GROUPS. (FIRST AND SECOND GROUPS). GENERAL DISCUSSION 27 THE SILVER GROUP (FIRST GROUP). Lead 29 Mercury 37 Silver 45 Comparison of Certain Reactions of the Metals of the Silver Group " 51 TABLE FOR ANALYSIS OF 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 Tin 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 Cadmium ^ 110 Comparison of Certain Reactions of Bismuth, Copper and Cad- mium 113 Vii vin CONTEXTS. PAGB THE PRECIPITATION OF THE METALS OF THE SECOND GROUP 113 TABLE FOR THE ANALYSIS OF THE TIN GROUP (SECOND GROUP, DIVISION A). 116 Directions for Analysis with Notes 118 TABLE FOR ANALYSIS or THE COPPER GROUP (SECOND GROUP, DIVISION B). . 126 Directions for Analysis with. Notes -125 RARER METALS OF THE TIN AND COPPER GROUP. Ruthenium 131 Rhodium 132 Palladium 133 Iridium 134 Osmium 135 Tungsten 136 Germanium 137 Tellurium 13g Selenium 139 THE IRON AND ZINC GROUPS (THIRD AND FOURTH GROUPS) 141 THE IRON GROUP (THIRD GROUP). Aluminum 144 Chromium 148 Iron 153 TABLE FOR ANALYSIS OF THE IRON GROUP (THIRD GROUP) 163 DIRECTIONS FOR ANALYSIS WITH NOTES 164 THE ZINC GROUP (FOURTH GROUP). Cobalt 167 Nickel 173 Manganese 177 Zinc 183 Comparison of Some Reactions of the Iron and Zinc Group Bases 187 TABLE FOR THE ANALYSIS OF THE ZINC GROUP (FOURTH GROUP) i8g DIRECTIONS FOR ANALYSIS WITH NOTES 189 ANALYSIS OF IRON AND ZINC GROUPS AFTER PRECIPITATION BY AMMONIUM SULPHIDE 191 IRON AND ZINC GROUPS IN PRESENCE OF PHOSPHATES 193 IRON AND ZINC GROUPS IN PRESENCE OF OXALATES 194 Table of Separation of Iron, Zinc and Calcium Group Metals and Phosphoric Acid by Means of Alkali Acetate and Ferric Chloride 196 Table of Separation of Iron, Zinc and Calcium Group Metals and Phosphoric Acid by Means of Ferric Chloride and Barium Carbonate 197 THE RARER METALS OF THE IRON AND ZINC GROUPS. Cerium 198 Columbium (Niobium) 198 Didymium 199 Erbium 200 Gallium.. . 200 CONTENTS. iX PAGE Glucinum (Beryllium) 200 Indium 201 Lanthanum 202 Neodymium 202 Praseodymium 202 Samarium 202 Scandium 202 Tantalum 203 Terbium 203 Thalliu in 204 Thorium 204 Titan ium 205 Uranium 206 Vanadium 207 Ytterbium 208 Yttrium 208 Zirconium 209 THE CALCIUM GROUP (FIFTH GROUP). (THE ALKALINE EARTH METALS) 209 Barium 211 Strontium 214 Calcium 216 Magnesium 220 TABLE FOR THE ANALYSIS OF THE CALCIUM GROUP (FIFTH GROUP). ...... 223 DIRECTION TOR ANALYSIS WITH NOTES 224 SEPARATION OF BARIUM, STRONTIUM, AND CALCIUM BY THE USE OF ALCOHOL 226 ALKALINE EARTH METALS AS PHOSPHATES 226 ALKALINE EARTH METALS AS OXALATES 226 THE ALKALI GROUP (SIXTH GROUP) 227 Potassium 228 Sodium 232 Ammonium 235 Caesium 239 Rubidium 240 Lithium 24') DIRECTIONS FOR ANALYSIS WITH NOTES 242 PAET III. THE NON-METALS. BALANCING OF EQUATIONS 246 Hydrogen 250 Boron 252 Boric Acid 252 Carbon 254 Acetic Acid 256 Citric Acid 258 Tartaric Acid ." 259 Carbon Monoxide 262 Oxalic Acid 263 Carbon Dioxide (Carbonates) 267 CONTENTS. PAGE Cyanogen 271 Hydrocyanic Acid 271 Hydroferrocyaiiic Acid 275 Hydroferricyanic Acid 277 Cyanic Acid 279 Thiocyanic Acid 280 Nitrogen 281 Hydrazoic Acid .282 Nitrous Oxide 283 Nitric Oxide 283 Nitrous Acid 284 Nitrogen Peroxide 285 Nitric Acid 285 Oxygen 291 Ozone 293 Hydrogen Peroxide 294 Fluorine 297 Hydrofluoric Acid 298 Fluosilicic Acid 298 Silicon 299 Silicic Acid 299 Phosph orus 301 Phosphine 304 Hypophosphorous Acid 304 Phosphorous Acid 306 Hypophosphoric Acid 307 Phosphoric Acid 308 Sulphur 313 Hydrosulphuric Acid 315 Thiosulphuric Acid 321 Hyposulphurous Acid 323 Dithionic Acid 324 Trithionic Acid 324 Tetrathionic Acid 325 Pentathionic Acid 325 TABLE OF THIONIC ACIDS 326 Sulphurous Acid 327 Sulphuric Acid 331 Persulphuric Acid 336 Chlorine 337 Hydrochloric Acid 341 Hypochlorous Acid 348 Chlorous Acid 349 Chlorine Peroxide 353 Chloric Acid 350 Perchloric Acid 353 Bromine 354 Hydrobromic Acid 357 Hypobromous Acid 360 CONTENTS. XI PAGE Bromic Acid 360 Iodine '. 362 Hydriodic Acid 365 lodic Acid 369 Periodic Acid 372 COMPARATIVE REACTIONS OF THE HALOGEN COMPOUNDS 373 PART IT. SYSTEMATIC EXAMINATIONS. REMOVAL OF ORGANIC SUBSTANCES 374 PRELIMINARY EXAMINATION OF SOLIDS 375 CONVERSION OF SOLIDS INTO LIQUIDS 378 TREATMENT OF A METAL OR AN ALLOY 379 SEPARATION OF ACIDS FROM BASES 381 TABLE FOR PRELIMINARY EXAMINATION OF SOLIDS 382 BEHAVIOR OF SUBSTANCES BEFORE THE BLOW-PIPE 386 TABLE OF THE GROUPING OF THE METALS 387 TABLE FOR THE SEPARATION OF THE METALS 388 ACIDS FIRST TABLE 390 ACIDS SECOND TABLE 398 ACIDS THIRD TABLE t . . 399 ACIDS FOURTH TABLE 400 NOTES ON THE DETECTION OF ACIDS 401 PRINCIPLES 4Q5 EQUATIONS 409 PROBLEMS IN SYNTHESIS 410 TABLE OF SOLUBILITIES 412 REAGENTS 415 ABBREVIATIONS. A. A. Ch. Am. Am. S. Arch. Pharm. Am. Chem. B. Bl. B. J. Comey. C. N. Ch. Z. C. r. C. C. Dingl. D. Fehling. Fresenius. G. O. Gazzetta. Gilb. Gmelin-Kraut. J. J. C. J. pr. J. Soc. Ind. J. Anal. J. Am. Soc. J. Pharm. Laden burg. M. Phil. Mag. Pogg. Proc. Roy. Soc. Pharm. J. Trans. Ph. C. Tr. Watt's. 1868* * Indicates continuance to the present time. Liebig's Annalen. 1832* Annales de Chimie 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. Bulletin de la Societe Chimique. 1859* Berzelius Jahresbericht. 1822-51. Comey's Dictionary of Solubilities. 1896. Chemical News. 1860* Chemiker Zeitung. 1877* Comptes Rendus des Seances de 1' Academic des Sciences. Chemisches Centralblatt. 1830* - Dingler's Polytechnische 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 iiber die Fortschritte der Chemie. 1847* Journal of the Chemical Society. 1849* Journal fiir 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* PoggendorfTs Annalen 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. 1835* Wells' Trans., 1897. 1885. COLLgat xiv ABBREVIATIONS. 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 fiir analytische Chemie. 1862.* Z. Ch. Zeitschrift fiir Chemie. 1865-1871. Z. anorg. Zeitscbrift fiir anorganische Chemie. 1891* Z. angew. Zeitscbrift fiir angewandte Chemie. 1888* Z. phys. Ch. Zeitschrift fiir physicaliscbe Chemie. 1887* 1. INTERNATIONAL ATOMIC WEIGHTS FOR 1916. Compiled by the International Committee on Atomic Weights, consisting of F. W. Clarke, W. Ostwald, T. E. Thorpe, and G. Urbain, Aluminum Al O=16 27 1 Molybdenum O Mo = .16. 96 Antimony . Sb 120.2 Neodymium -Nd 144 3 Argon A 39.88 Neon Ne 20 2 Arsenic As 74.96 Nickel Ni 58 6;5 Barium Bismuth Boron Ba Bi B 137.37 208.0 11 Niton (radium emanation) Nitrogen Nt N 222.4 14 01 Bromine Br 79 92 Osmium. Os 190 9 Cadmium Cd 112 40 Oxygen o 16 00 CSBSI im Cs 132 81 Palladium Pd 106 7 Calcium Ca 40 07 Phosphorus p 31 04 Carbon c 12.00 Platinum Pt 195 2 Cerium Chlorine Ce Cl 140.25 35.46 Potassium Praseodymium K Pr 39.10 140 9 Chrom 1 um Cr 5? Radium Ra 226 Cobalt Co 58 97 Rhodium Rh 102 9 Columbium Ch 93 5 Rubidium Rb 85 45 Copper Cn 63 57 Ruthenium Ru 101 7 Dysprosium Erbium Europium % Eu 162.5 167.7 152 Samarium Scandium Selenium . . . Sa Sc Se 150.4 44.1 79 2 Fluorine F 19 Silicon . . Si 28.3 Gadolinium. Gd 157 3 Silver Ag 107 88 Gallium Gn 69 9 Sodium Nn 23.00 Germanium GP 72 5 Strontium Sr 87.63 Glucinum Gl 9 1 Sulfur. S 32.06 Gold Helium Au HP 197.2 4.00 Tantalum Tellurium Ta Te 181.5 127.5 Holmium Hrt 163.5 Terbium Tb 159.2 Hvdrogen H 1.008 Thallium 71 204.0; Indium In 114 8 Thorium Th 232.4 Iodine I 126.92 Thulium Tm 168.5 Iridium. Ir 193.1 Tin Sn 118.7 Iron FP 55.84 Titanium. Ti 48.1 Krypton Kr 82.92 Tungsten W 184.0 Lanthanum. . . T* 139.0 Uranium u 238.2 Lead Pb 207.20 Vanadium V 51.0 Lithium Li 6.94 Xenon Xe 130.2 Lutecium Lu 175 Ytterbium (Neoytter- Magnesium Me 24 32 bium) . . Yb 173 5 Mangansse Mn 54 93 Yttrium .... Yt 88 7 Mercury Hr 200 6 Zinc .... Zn 65.37 Zirconium Zr 90.6 Jour. Am. Chem. Soc., 1915, 37, 2449. Th 00 t^ OOCOO5 THOCO II II II O5rH :CO Oi O5 8 CO ^O C^HH a H -Ji -si U q o so r-| uj O r-t O PH II 10 g 7 o ii 2 O a o S O rH CO O O CO 8 . (M t& PS 00 OC 00 C5 3. DISCUSSION OF THE PERIODIC SYSTEM. 3 3. In the periodic system of the chemical elements certain regular gradations of chemical character are to be studied and held in view, to sim- plify the multitude of facts observed in analysis. Passing from Li G.94 to F 19.0 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 F 19.0 to the next higher atomic weight, Na 23.0, 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 (H0)-+ I +- 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 ? , and its positive polarity in its highest combination with oxygen, such as NO a (OH). 4 DISCUSSION OF THE PERIODIC SYSTEM. 4. 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 elements 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. Ramsay. Br. Assoc. Adv. Sci., 1897, 598-601; B. 1898, 31, 3111. J. L. Howe, C. N. 9 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 Relation 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., falJa 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. 5 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 are the more strongly marked of the families of the well known elements. 6. The Alkali Metals. 116.94, (Na 23.0), K 39.10, Rb 85.45, Csl32.81. 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 orthophosphate of lithium are but slightly soluble. 7. The Alkaline Earth Metals. (Mg 24.32), Ca 40.07, Sr87.63 ; Ba 137.37. These metals, like those of the alkalis, form stronger bases as they have 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 arq 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. Mg 24.32, (Al 27.1), Zn 65.37, Cd 112.4, , H<* 200.6. 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 52.0, Mn 54.93, Fe 55.84, Ni 58.68, 6 DISCUSSION OF THE PERIODIC SYSTEM. 10, Co 58.97. 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 seriei 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 58.97 follow Cu 63.57 and then Zn 65.37. 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 , Mn , JFe , Ni , Co , Cu , Zn On reaching zinc, 65.37, 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 VIII. Cu 63.57, Ag 107.88, , 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 * These metals form unstable hydrated basic carbonates. 12. DISCUSSION OF THE PERIODIC SYSTEM. 7 weight: Cu 63.57 with Ni 58.68 and Co 58.97, Ag 107.88 with Pd 106. 7, and Au 197.2 with Pt 195.2. 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, Au, 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 14.01, P 31.04, As 74.96, Sb 120.2, , Bi 208.0. These elements include the leading element of Group V., and the entire second part of the group. Nitrogen and phos- phorus act as non-metals, antimony and bismuth as metals, while arsenic is intermediate, 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. While in ammonium compounds nitrogen exercises a valence of five, 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 . HHHHG1. 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., having valences of four and five, differing from the valence of the nitrogen family. In Series 7: Sn 118.7 is closely related to Sb 120.2. In Series 11: Pb 207.20 is closely related to Bi 208.0. 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 are lead and bismuth. In riti irnDMiA ftftl 1 EfiE 8 DISCUSSION OF THE PERIODIC SYSTEM. 12. further comarison, arsenic, in its deportment as a metal, may be included making the list: As 74.96, Sb 120.2 (Sn 118.7), Bi 208.0 (Pb 207.20). 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, while 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 (As, Sb, Sn, Bi, Pb) 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 insolubility of the oxychlorides of bismuth, tin, and antimony, and this fact must be borne in mind, when precipitation by hydrochloric acid is employed for the 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 decom- position being represented by gunpowder. 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 unstable body, hydrazoic acid, HN 3 (also called azimide), of decided acid power, constitut- ing well marked salts, such as sodium azoimide, 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 NZlQ . 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 orthq acids, phosphoric and arsenic, have the constitution (H0) 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 the 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 salts with metals, 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. TJie Halogens.* 19.00, Cl 35.46, Br 79.92, I 120 92.' These metals constitute the leading elements of Group VI f., 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 trie 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 HgI 2 and K 2 Pt C1 6 . The periodides show that iodine atoms have the 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 Cl 2 ; bromic, HO Br 2 ; and iodic, HO I 2 . Chloric acid resembles 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 oxygon acids with a halogen valence of seven, periodic acid, HO I 3 , also (H0) 5 1 , is pre- served intact without difficulty. Perchloric acid is more stable than chloric acid. 14. The Relations of Sulphur. $ 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, namely those of Pb 207.20, Ba i:>7.:>7, Sr 87.63, and (with sparing solubility) Ca 40.07. Of these sulphates, that of barium (least soluble), 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 (H C EEC H), also in oxalic acid, (HO) OC CO (OH), 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 hydrazoic acid, by oxygen in ozone, by sulphur in thiosulphuric acid, and by iodine in perio- dides. In carbon, nitrogen, and oxygen we see a decreasing gradation 01 this capacity, as the atomic weights ascend. Silicon, next to caroon in Group IV., but in the opposite set of alternates, agrees with carbon m tne formation of many corresponding compounds, and also exhibits to some extent the capacity of uniting its atoms to each otuer in ounding up com- binations. 16. The Classification of the Metals as liases. The object of the Periodic System is to group .ill the elements, both metallic and non-metallic, according to their general properties as related to their atomic weights. This has been briefly given in the foregoing pages for study 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 witli the properties of the leading bases and acids. JJ16. 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 liable to fallacy. The following named are the bases of more common occurrence. Metals Precipitated as Chlorides, ine Silver Group. The first group.* Silver (Argentum). f Ag 1 : silver salts. Mercury (Hydrargyrum) , Hff 1 : merer rous salts. Lead (Plumbum). Ph 11 ; lead salts, Silver and the mercury of mereur- ons salts can be removed as chlorides by precipitation with hydrochloric acid. The precipitate of lead is not insoluble enough to remove this metal entirely in separation from other groups. Metals falling with Copper and Tin. Precipitated ~by H.^S in acidulated The second group. solution. (The precipitates are sul- phides.) The Tin Group. Division A, second grout). bn xi stannous salts. Sn iV : stannic salts and stannates. Sb m : antimonous compounds. Sb v : antimonic compounds. As 111 : arsenous compounds. As v : arsenic compounds and arsen- ates. Separated by dissolving the precip- itated sulphides with Ammonium Sulphide. Separated by the insolubility of the precipitated sulphides on treatment with Ammonium Sulphide. The Copper Group. Division B, second groun. Hg n : mercuric salts. Pb 11 : lead salts. Bi m : bismuth salts. Cu 11 : copper or cupric salts. Cu 1 : cuprous salts. Cd 11 : cadmium salts. * The first division of the bases, in the ordar in which they are separated from each other by precipitation with the group reagents. t The Roman numerals (as J ) express units o* valence, each equivalent to an atom of hydrogen, in the formation of salts and other combinations. CLASSIFICATION OF THE METALS AS BASES. 16. The Iron Group. The third group. Fe 11 : ferrous salts. Fe m : ferric salts. Cr m : chromic salts. Cr vl : chrornates. Al m : aluminium salts. The Zinc Group. The fourth group. Zn 11 : zinc salts. Mn IT : manganous salts. manganic salts. unstable salts. salts of manganic acid. salts of permanganic acid Ni 11 : nickel salts. Co 11 : cobaltous salts. Co 111 : cobaltic salts. The Alkaline Earth Bases, The fifth group. Magnesium, Mg 11 . Mn 111 Mn IV Calcium, Ca lr . Strontium, Sr n . Barium, Ba n . The Alkali Bases. Tlie sixth group. Separated by precipitation ivith Ammonium Hydroxide, in presence of NH,C1, after the removal of previ- ously named groups. (The precip- itates are all hydroxides.) Separated by precipitation with Ammonium Sulphide, after remoi.u of all previously named oases, as di- rected above. (The precipuates are all sulphides.) (Precipitated by carbonates, wh'ch fact alone does not separate tnem from the following named groups.) Separated by precipitation as ti phosphate after removing all the pre- viously named bases. Forms magne- sium hydroxide, Mg(OH),, and mag- nesium salts, such as MgS0 4 . Separated by precipitation with Ammonium Carbonate, adding NH^Cl to keep magnesium from precipitation. Calcium carbonate, a normal salt, CaCO 3 . Not precipitated from their salts by any of the group reagents. Potassium and sodium are found after removing aa the previously named groups. Ammonium is found by tests of the original, this base being added in the "group reagents." 18. Potassium (Kaliurn), K 1 . THE OPERATIONS OF ANALYSIS. 13 In combination in potassium hy- droxide, KOH, and in potassium salts, such as the chloride KC1, and the ni- trate KNO, . Sodium (Natrium), In the base, sodium hydroxide and its salts. Ammonium Forms ammonium hydroxide^ NH 4 OH, representing ammonia, NHj, and water, and serving as the base of ammonium salts, such as ammonium sulphate. 17. THE ACIDS or CERTAIN COMMONLY OCCURRING SALTS. Name of Acid. Name of Salt. Formula. Showing Hydroxyl. Anhydride. 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 3 Nitric Nitrate HN0 3 (HO)NV0 2 N 2 5 Nitrous Nitrite HN0 2 (HO)NniO N 2 3 Phosphoric (ortho) Phosphate H 3 P0 4 (HO) 3 PvO P 2 5 Metaphosphoric Metaphosphate HP0 3 (HO)Pv0 2 P 2 0s Pyrophosphoric Pyrophosphate H 4 P 2 T (HO) 4 PV 2 3 P 2 5 Sulphuric Sulphate H 2 S0 4 (HO) 2 Svi"0 2 S0 3 Sulphurous Sulphite H 2 S0 3 S0 2 Hydrosulphuric Sulphide H 2 S Hydrochloric Chloride HC1 Hydrobromic Bromide HBr Hydriodic Iodide HI Chloric Chlorate HC10 3 (HO)C1V0 2 C1 2 5 locdc lodate HI0 3 (HO)IVO 2 I 2 O 6 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 an "element " and is always and everywhere the same. It is required in analysis to detect a chemical compound as a body distinct from the chemical elements that have formed it. For example, the analyst may have in hand a mixture containing sodium sulphate, Na SO, ; sodium sul- phite, Na,SO i , and sodium thiosulphate, Na 2 S.0 3 , but not containing any 14 THE OPERATIONS OF ANALYSIS. 19. 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 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, 27. THE OPEHATIOXS OF ANALYMX. 15 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- lirity methods are used for establishing the presence or absence of minute traces of substances (Kohlrausch Whitney, Z. pJiys. (77?.., 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, Lelirb., 2 Aufl., II, 1, 881; Behrend, Z. phys. Cli., 1893, 11, 466; Hulett, Z. phys. Cli., 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 loniza- fcion." 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 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 16 THE OPERATIONS OF ANALYSIS. 28. 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. pliys. 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 &); e. g., amorphous magnesium ammonium phosphate is quite soluble in water, the crystalline salt being almost insoluble. The solubility of a solid is ,-ilso dependent upon the size of the particles of the solid, a finely divided solid being more soluble than large particles of the same substance (Hullett and Allen, J. Am. Soc., 24, 667, 1902). In analysis it is customary to heat and then allow a precipitate to stand with the solution in which it was formed in order to obtain com- plete precipitation. 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 substance. 28. The ordinary liquid reagents are water solutions concentrated sul- phuric acid and carbon disulphide being exceptions. Hydrochloric acid, liquid hydrogen sulphide, and ammonium hydroxide 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 Reagents.) 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 filtration. 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 THE OPERATION* OF ANALYSIS. 34. with the help of a microscope and with studies of crystalline form. Further see Behrens, Z. 1891, 30, 1^5; and Herrnschmidt and Capelle, Z. 189; . 32, 608. 34. Precipitates are removed usually hy filtration, sometimes by d ^can- tation. If they are to be dissolved, they must he 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 othe dis- solved matters, after precipitations, some of which are indicated Ir the equation written for the change. All these dissolved substances pemeate and adhere to the porous precipitate with greater or less tenacity. If they are not wholly washed away, some portion of them will he 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 :ilter. 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. Fndeavor to obtain a solation 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 iho panic 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 ?omo other easily soluble precipitates the acids may be diluted with fifty times their weight of water. Washed precipitates may also be dissolved ri 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 a reagent is added in order to produce a change in the acid, alkaline or neutral condition of the solution, the addition of suff.cient 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. AY hen 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, i.e. aprecipitate. A chemical change between dissolvec 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 mixtiTC. As an example, Calcium carbonate + Hydrochloric acid Cal- * in in chloride + Water -(- Carbon- dioxide (gas). 3'. Tn the practice of qualitative analysis, the student necessarily refers to authority for the composition of precipitates and other products. For exam tie, when the solution of a carbonate is added to the solution of a calcivm sail, a precipitate is obtained; and it has been ascertained by quanti- tativ analysis that this precipitate is normal calcium carbonate, CaCO, , invar 'ably. Were there no authorized statement of the composition of this precipitate, the student would be unable, without making a quantitative anah ?is, to declare its formula, or to write the equation for its production. Whe:. the results of analytical operations are substances of unknown, uncer- tain, or variable composition, equations cannot be given for them. g3 l \ The written equation represents only the substances, and the quan- tity (>f each, which actually undergo the chemical change that is to be expressed. Thus, if a reagent is used to effect complete precipitation, an exces -; of it must be employed, beyond the ratio of its combining weight in the e 'illation. That is, if magnesium sulphate be employed to precipitate bariu ii chloride, the exact relative amount of magnesium sulphate indicated by th 3 equation: BaCL + MgS0 4 = BaS0 4 -f- MgCl, , fails to precipitate all of th ? 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 by the aid of hydro- chloric acid, we have the equation: Fe + 2HC1 = FeCL + H 2 . 55.8 + 72.9 = 126.8 + 2 . Also in precipitating ferrous chloride by sodium phosphate, we have the equation: FeCL + BTa 2 HP0 4 ,12H 2 = FeHPO 4 + 2NaCl + 12H 2 . 126.8 + (142.2 + 216) = 151.8 + 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 for 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 55.8/72.9 = 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 126.8/72.9= 100/x. 2. -j 32/100 =x/y = parts of 32 per cent HC1 for 100 parts of FeCl 2 , absolute, ( 129.8/55.8 = 100/z = parts of iron wire for 100 parts of FeCl 2 . , 151.8/72.9 = 100/x. j 32/100 =x/y = parts of 32 per cent HC1 for 100 parts of FeHPO 4 . 1 151.8/55 8 =100/z = parts of metallic iron for 100 parts of FeHPO 4 . ' 151.8/358.2 = 100/u = parts of Na>HPO 4 , 12H : O for 100 parts of FeHPO 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 formula 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 IOKIZATION. 43. The Theory of Electrolytic Dissociation, proposed by Arrhenius in 1887 (Z. pliys. Ch., 1887, 1, 631), assumes that acids, bases and salts in water solution are present not as the intact molecule but split up into two or more parts called ions which are charged with negative or positive elec- tricity. The facts upon which the theory is based are the osmotic pressure,* * 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 through- out the solvent, just as gas by the virtue of the gas-pressure occupies the entire space at its disposal. 43. SOLUTION AND IONIZAT1ON. 21 lowering of the freezing point, raising of the boiling point and electrical conductivity of water solutions of acids, bases and sails. Because such water solutions conduct the electric current, acids, bases and salts are culled electrolytes. The osmotic pressure of a solution is believed to be proportional to the number of particles of the dissolved substance present in unit volume of the solution. In the case of non-electrolytes the osmotic pressure is pro- portional to the molecular weight; but the osmotic pressure of electrolytes is greater than corresponds to their molecular weight. This is readily ex- plained by the assumption that some of the molecules are split into two or more parts. In a similar manner the freezing point of a liquid is lowered by the pres- ence of a dissolved substance, and the amount of the lowering is propor- tional to the number of dissolved particles. In the case of non-electrolytes the lowering is proportional to the molecular weight, but dissolved electro- lytes depress the freezing point to a greater extent, indicating dissociation of the molecules. The boiling point of a liquid is raised by the presence of a dissolved substance, but this effect is greater in the case of electrolytes because the molecule is dissociated. This reasoning is similar to that applied to gas pressures. The gas-laws (Boyle's, Guy-Lussac'a, 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; Morse and Frazer, Am. 9 34, 1 (1905); 37, 324, 425, 558; 38, 175 (1907); Lewis, J. Am. Soc.,ZQ, 6G8 (1908) ). Avogadro's Hypothesis 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 6 ) so excessive osmotic pressure, lowering of freezing point and raising of boiling point is taken as indicating dissociation of the dissolved substance. The osmotic pressure as well as the abnormal freezing and boiling point may be taken as a measure of this dissociation. The fact that solutions of non- electrolytes do not conduct the electric current while solutions of electrolytes do conduct the electric current indi- cates that molecules are incapable of carrying the current, but that the component parts into which the molecule is split carry the current. Faraday gave the name ions to the components of a substance conducting the 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 2 SOLUTION AND JONIZAT10N. 43. 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 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 saponi- fication of esters; (2) the extent of dissociation as indicated by osmotic pressure, and (3) the extent of dissociation as indicated by electric conduct- ivity. The same parallelism holds for other bodies in solution. A water solution of an acid such as hydrochloric or sulphuric should not be regirded as consisting of molecules of the acid and water but as a solution of mole- cules of HC1 and the ions H and Cl or molecules of H,S0 4 and ions H ai:il S0 4 respectively, the hydrogen ion being charged with positive electricity ai d the acid ion with negative electricity. The very active acids and bases ai.d the neutral salts undergo wide dissociation in water solution, while weak acids and bases retain almost entirely the non-dissociated condition, the strength of the acids being proportional to the concentration of the H ion. The Electrolytic Dissociation Theory in its assumption of a separation into ions groups together and gives system and meaning to these three classes of facts, experimentally absolutely independent and up to Arrhonius' 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. For instance, HJS or K 2 S will produce precipitates of metallic sulphides because the sulphur is present in solu- tion as an ion. On the other hand HSO ; , H 2 S0 3 , or H,S_0 3 wi 1 not precipitate metals as sulphides, because in these acids the ion is com- posed of the sulphur and the oxygen present. Further, HgCL in its chemical behavior is unlike other mercuric salts and unlike other chlorides. The mercury is not readily precipitated by alkali hydroxides n-n* is the chloride readily precipitated by silver salts. In agreement with this, its conductivity and osmotic pressure are also unlike those of the great 44. SOLUTION AND lONtZATIONn 23 majority of neutral salts, both pointing to very slight dissociation into the ions. CdClo is another neutral salt anomalous in that its conductivity and osmc:ic pressure are both low. And here also for precipitation of the chlor'.de a considerable concentration of the reagent is necessary. Similar instances of the parallelism referred to are numberless. 40:. The Law of Mass-Action embodies the familiar principle that the chemical activity of a substance is proportional to its concentration. It was irst recognized, although imperfectly, by Berthollet and was given math 3inatical expression by Guldberg and Waagc 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- rhenuis 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 \ r hen 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 fie reaction indicated by the equation CH,CO,H -\- C 2 H 5 OH CH 3 C 0,0,11, -f H 2 , when equilibrium sets in ab = kcd , in which a and b are t-ie 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 reaction is a dissociation, as with gaseous NH 4 C1 , we have ab = k'c , a and b repre renting the concentrations of NH 3 and HC1 respectively, c that of the uncle jomposed 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 aod 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 * Tl 3 unit of quantity is the molecular weight taken in grams (the " mol ") Where there are 18.23 | rams HCI in a liter either in solution or as gas the concentration is M, where there. are 72.92 ^;ams in the same volume the corcentratiori is 2 and so on, 24 SOLUTION AND IONIZATION. 45. \vhile the water is taken up by some absorbent. Introducing gaseous NH :j 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. Aeetic 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 OH , 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 OH as saponifying agent and that calculated from the equation: NH ' C OH' ~ k NH OH (Arrhenius, Z. phys. Cli., 1887, 1, 110). In general every acid is weakened by the addition of the neutral salt of the acid to its solution. Similarly bases are weakened by the addition of the neutral salt of the base to its solution. 45. The Solubility-Product. In the saturated solution which always remains after precipitation we have the usual dissociation equilibrium, as: C Ae ' C C1' ~~ AffCl ' ^ ow ^ ne q uaT1 tity ^ 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 ^ ff - ~ " C Q^ From such 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- tative agreement with the theory (Nernst, Z. phys. Ch., 1889, 4, 372; Noyes, Z. phys. Ch., 1890, 6, 241; 1892, 9, 603; 1898, 26, 152). The Solubility-Product of the alkaline-earth carbonates is M " C CO " ^ n ^ e so l u ti n 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 46. ORDER OF LABORATORY STUDY. 25 alkaline-earth hydroxides, while ammonium hydroxide, shown by othei measurements to contain but few hydroxyl ions, will not. For the metallic sulphides the solubility-product is j^- 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/' ORDER OF LABORATORY STUDY. 846. 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 the formulas of salts. &. 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. O or Pb(C 2 H 3 O 2 ) il + 4KOH (excess) = K 2 PbO 2 + 2KC 2 H 3 O 2 + 2H 2 O . The results should all be tabulated and then summarized in the form of a carefully worded generalization (205, 6a). c. Action of Ammonium Hydroxide (volatile alkali) upon solutions of the suits of the metals, etc., as in (b) above; e.g., the addition of ammonium hydroxide to lead nitrate produces a white precipitate not dissolving in ex* cess. Consult text (57, Go) and write the equation: 3Pb(N0 3 ) 2 + 4NH 4 OH = 2PbO Pb(NO 3 ) 2 + 4NH 4 NO 3 + 2H 2 O . After the work has been completed in the laboratory and the results * It Las been found helpful to require students to underscore alt precipitates. 26 ORDER OF LABORATORY STUDY. 46. discussed in the class room., summarize in the form of a generalized state- ment (207, 6ff). 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 , HNO S , 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. t. The solvent action of acids, HC1 , HN0 3 , and HC 2 H,0 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. Jc. 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., the addition of calcium chloride to ammonium oxidate produces a white precipitate. Consult the text (227, 8) and write the equation: (NH 4 )C 2 4 + CaCL = CaC 2 4 + 2NH.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 should be followed by the analysis of "unknown 5 ' 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. 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, Eorates, Silicates, etc., and certain of the Rarer Metals. u. The analysis of mixtures in solution, illustrating Oxidation and Reduction, PART II.-THE METALS. THE SILVER AND TIN AND COPPER GROUPS. (FIRST AND SECOND GROUPS.) 47. The Silver group (first group) includes the metals whose chlorides are insoluble in water and which are precipitated from solutions upon tlK addition of hydrochloric acid or soluble chlorides : Pb, Hg', Ag . The Tin and Copper group (second group) includes those metals whos sulphides are precipitated by hydrosulphuric acid from solutions acid witl dilute hydrochloric acid, and whose chlorides (soluble in water for the most part) are not precipitated by hydrochloric acid or soluble chlorides. Lead* Pb 207.20 Germanium Ge 72.5 Mercury Hg 200.6 Iridium Ir 193.1 Silver Ag> 107.88 Osmium Os 190.9 Arsenic As 74.96 Palladium Pd 106.7 .Antimony Sb 120.2 Ehodium Bh 102.9 Tin Sn H8.7 Kuthenium Bu 101.7 Gold Au 197.2 Selenium Se 79.2 Platinum Pt ?95.2 Tellurium Te 127.5 Molybdenum Mo 96.0 Tungsten "W 184.0 Bismuth Bi 208.0 Vanadium V 51.0 Copper Cu 63.57 Cadmium Cd U2.4 48. Owing to the partial solubility of lead chloride in water, it is nevei completely precipitated in the first group; hence it must also be tested for 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 1O4. 28 GENERAL DISCUSSION. 50. oxidation in the air. In general, these metals do not dissolve in acids with evolution of hydrogen, or do so with difficulty. Nitric acid is the best solvent for all, except for antimony and tin, which are rapidly oxidized by it. Antimony may be dissolved by treatment with a little strong nitric acid and tartaric acid. The best solvent for tin is hot strong hydrochloric acid. Concerning the separation and detection of the metals of these groups by electrolysis, see Schmucker, Z. anorg., 1891, 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 acid radicles, which form stable salts. Arsenic, selenium and tellurium 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 1 , Ag , Bi and Cd . Platinum forms a large number of stable double chlorides, soluble and insoluble; antl 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 hydroxide. Me- tallic 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 tho metals of the second group into two divisions. A (copper group) Kg, Pb, Bi, Cu, Cd, Os, Pd, Rh, and Ru; sulphides not soluble in yellow ammonium sulphide; and B (tin group) As, Sb, Sn, Ge, An, Ir, Mo, Pt, Se, Te, W, and V; sulphides soluble in yellow ammonium sulphide. 55. Mercury, antimony, silver, and gold do not form hydroxides. The oxides of gold a.re 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 GKOUP (FIRST GROUP). Lead, Mercury (Mercurosiun), Silver. 57. Lead (Plumbum Pb = 207.20. Valence two and foul. 1. Properties. Specific gravity, 11.34 (Reich, J. pr., 1859,78, 328). Melting point, 327.4. B. S., circular No. 35, 2nd ed., 1915. Vaporization is said to take place at 360 (Demarcay, C. r., 1882, 95, 183). Boiling point about 1525 (H. C. Greenwood, C. N., 39, 49). It can be distilled in vacuo, (Schuller, B., 1883, 16, 1312). Pure lead is almost white, soft, malleable, very slightly ductile; freshly cut surfaces tarnish in the air from formation of a film of oxide. The presence of traces of most of the other metals makes lead sensibly harder. It is a poor con- ductor of heat and electricity, and forms alloys with most metals; lead and tin in various proportions form solder and pewter; lead and arsenic form shot metal; lead and antimony, type metal; lead, bismuth, tin $ nd cadmium form easily fusible alloys of low melting points (minimum 55.5; Ch. Z., 30, 1139-1143; J. Soc. Ind., 25, 1221).; bell metal consists of tin, copper, lead and zinc. 2. Occurrence. It is rarely found native (Chapman, Phil. Mag., 1886, (4), 31, 176); the most abundant lead mineral is galena, PbS; it also occurs as cerussite, PbCO 3 ; anglesite, PbSO 4 ; pyrqmorphite, Pb 6 Cl(Pp 4 )5; crocoite, PbCrO 4 ; and in many other minerals in combination with arsenic, antimony, etc. The United States produces more lead than any other country. Spain produces about one- fourth of the world's supply. 3. Preparation. (a) From argentiferous lead ores; after roasting, if neces- sary, the ore is smelted in a rectagular bl st furnace with a properly propor- tioned mxiture of coke and limestone. The lead (base bullion) produced is de- silverized and refined by the Betts (electrolytic) process, or desilverized by the Parkes process and subsequently refined in a reverberatory furnace. (6) From galena, by the roast-reaction process in reverberatory furnaces, and ore hearths; the ore is roasted with access of air, forming variable quantities of PbSO 4 , PbO, and PbS. Air is then excluded and the temperature raised, the sulphur of the sulphide then reduces both oxide and sulphate with formation of SO2^ PbSO 4 + PbO + TPbS - 5Pb + 3SO 2 In variations of this process carb n is used to aid in the reduction. 4. Oxides. Lead forms four oxides, Pb 2 O , PbO , Pb0 2 , and Pb 3 O 4 . Lead suboxlde (Pb 2 O) is little known: it is the black powder formed when PbC 2 O 4 is heated to 300, air being- excluded. Lead oxide (litharge, or massicot) is formed by intensely igniting- in the air Pb , Pb 2 O , PbO 2 , Pb 3 O 4 , Pb(OH) 2 , PbCO ri , PbC,0 4 , or'Pb(NO :! ) 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 O 4 , is formed by heating 1 PbO to a dull-red heat with full access of air for several hours. Strong, non-reduc- ing- acids, such as HNO 3 , H,SO 4 , HC1O 3 , etc., convert it into a lead salt and Pb0 2 (a). But concentrated hot ELSO 4 converts the whole into PbSO 4 , oxygen being- evolved (&). But with the dilute acid and reducing agents, such as C 3 H 5 (OH) 3 , C G H 12 O,, , H 2 C 2 4 , H 2 C 4 H 4 O 6 , Zn , Al , Cd , Mg , As, Pb , etc., it is all reduced to the dyad lead without evolution of oxygen (c), (d), and (e). Hydracids usualty reduce the lead and are themselves oxidized (f). (a) Pb 3 4 + 2H 2 S0 4 (dilute) = PbO 2 + 2PbS0 4 + 2H 2 (5) 2Pb 3 4 -f 6H 2 SO 4 (concentrated and hot) = 6PbS0 4 + 6H 2 O + 3 (c) Pb 3 O 4 + H 2 C 2 4 + 6HN0 3 = 3Pb(NO 8 ) a + 4H 2 + 2C0 2 (d) 10Pb 3 4 + As 4 + 30H,S0 4 = 30PbS0 4 + 4H 3 AsO 4 + 24H 2 (e) Pb 3 4 + Zn + 4H 2 SO 4 = 3PbSO 4 + ZnSO 4 + 4H 2 (f) Pb 3 4 + 8HC1 = 3PbCl 8 + Cl, + 4H 2 The valence of Pb 3 O 4 is best explained by the theory that it is a union of the dyad and tetrad (Pb" and Pbrv) , Pb 8 O, = 3PbO + 30 LEAD. 57, 5a. Lead dioxide or peroride, PbO, , is formed: (1) by fusion of PbO with KC1O, or KNO, ; (2) by fusing Pb 3 4 with KOH : (3) by treating- any compound of Pb" with Cl , Br , K 3 Fe(CN) 6 , KMn0 4 , or H 2 2 in presence of KOH; (4) by treating Pb 3 4 with non-reducing acids: Pb 3 4 + 4HN0 3 = PbO 2 + 2Pb(N0 3 ) 2 + 2H 2 O. Ignition forms first Pb 3 4 and above a red heat PbO, oxygen being given oft'. It dissolves in acids on same conditions as Pb 3 4 . Very strong solution of potassium hydroxide, in large excess, dissolves it, with formation of " potassium plumbate," K 2 Pb0 3 . 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. Meto 1 * 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 w r hite residue of lead nitrate will be left which dissolves on the addition ef water. If concentrated and hot, the nitric acid is reduced to NO which, on contact with the oxygen of the air, becomes N 2 Oa (241, 6). The reactions are as follows: 3Pb + 8HNO 3 = 3Pb(NO 3 ) 2 +2NO + 4H 2 O 4NO + O 2 = N 2 O 3 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 concen- trated 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. The following reaction takes place (266, 6 A): Pb+2H 2 SO 4 = PbSO 4 + SO 2 + 2H 2 O. Hydrochloric acid very slowly dissolves the metal (more rapidly when warmed),, evolving hydrogen; the chloride formed dissolves in the acid in quantities depend- ing upon conditions of temperature and concentration (c) . The halogens readily attack the metal forming the corresponding 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 Avithout 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, I i Marge, PbO , and the hydroxides, 2PbO.H 2 O; 3PbO.H 2 O, are readily dissolved or transposed by acids forming the correspond- ing salts, i. ., PbO + H 2 S0 4 = PbS0 4 -j- H.,O ." The oxide and hydroxide are soluble in about 7000 parts of water, to which they impart arr alkaline reaction. They are soluble in the fixed alkalis forming plumbites; soluble in certain salts as NH 4 C1 , CaClo , and SrCL (Andre, C. r., 1883, 96, 435; 1887, 104, 359); very soluble in lead acetate, forming a strongly alkaline solution of basic lead acetate. Lead dioxide, Pb0 2 , lead pero-ride, 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: PbO, + 4HC1 = PbCL + C1 2 + 2ELO; 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 3 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. Ch., 1867, (4), 10, 328). The peroxide is slowly soluble in the fixed alkali hydroxides forming plum- bates, i. e., PbO 2 + 2KOH = K 2 PbO 3 + H,O . Trilead tetroxide, Pb 3 4 , red lead, nihunm, is insoluble in water, is at- tacked by nearly all acids in the cold forming the corresponding dyad lead * 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 Pbla and KI, this Is offset In concentrated solution by the formation of a complex compound. 57, 5a. LEAD. 31 salt and ead peroxide, PbO 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 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, the lead being reduced and then converted into the soluble nitrate. 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 Rose, Z. phys. Ch. f 1893, 12, 241), the presence of HN0 3 or HC1 in- creasing 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 i precipitated by the addition of water or alcohol; less soluble in dilute 11,80^ than in water; soluble in 682 parts 10 per cent HC1 , in 35 parts 31.5 per cent (Eodwcll, 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. (77?,., 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 (Rodwell, 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 166 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 . ;//., 1834, 1, 425). The chloride is less soluble in dilute HC1 or H.SO, than in water, but is more soluble in the concentrated acids (Ditte, I. c.) ; HNO :>> increases the solubility of the chloride more and more as the HNO. { is stronger. The chloride is less soluble in a solution of NaCl than in water (Field, J. 0., 1873, 26, 575); soluble in NH 4 C1 90 grams dissolving in 200 grams NH 4 C1 with 200 cc. water (Andre, (7. 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 KNO n 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). 3% LEAD. 57, 6. 6. Reactions, a. Fixed alkali hydroxides precipitate, from solutions of lead salts, basic lead hydroxide (1), Pb,0(OH) 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(OH) 2 , may be formed by adding a solution of a lead salt to a solution of a fixed alkali hydroxide. (1) 2Pb(N0 3 ) 2 + 4KOH = Pb,0(OH) 2 + 4KN0 3 + H 2 (2) Pb a O(OH), + 4KOH = 2K,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 PbClo.PbO.H L> (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 basic lead 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 ;j (OH)o(CO.,) 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), , white, sparingly soluble hi a large excess of the reagent and reprecipitated on boiling. Potassium ferro- cvanide 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 (Fe(CN) 6 ) 2 , white, sparingly soluble in water, soluble in nitric acid. Solutions of lead salts are precipitated by potassium sulpho- cyanate as lead sulphocyanate, Pb(CNS)2, white, soluble in excess of the reagent and in nitric acid. c. Lead nitrate is very soluble in water, the solution dissolving the oxide to form a basic nitrate, which may also be formed by precipitating lead acetate with * Nearly all the salts are soluble in the fixed alkali hydroxides, PbS forming almost the only* notable exception. 57, Ge. 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 beiny formed ( Le I '.lane and A'oyes, Z. pliy*. ('//., 1890, 6, 385). . But adding H 2 S or a soluble sulphide to PbS<> 4 gives just the opposite of this condition, and transformation accordingly results, increasing the SO 4 " con- ce itrati'Mi by formation of soluble sulphate and decreasing the S" concentration by precipita- tion of PbS, until the equilibrium-ratio is produced or, if the quantity of Pl>SO 4 present is in- s'lTu-ient for this, until all the PbSO 4 has been transformed to sulphide. On tlr> other hand, treatment of PbS with a very large excess of H 2 SO 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 + KCl ; CaSO 4 + Na 2 COi = CaCOa + Na 2 SO4 (U4). 34 LEAD. 57, 6/. 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. (7., 1889, 667). When galena, PbS , is pulverized with fused KHSO 4 , H,S is evolved (Jan- nettaz, J. C., 1874, 27, 188). Lert ,3 at 760 mm. (Ramsay and Young 1 , /. 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," /. 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 knilV. 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. Amalgams with special alloys of gold, silver, tin, and zinc are used for filling teeth. 2. Occurrence. Occasionally found native in small globules associated with cinnabar, in the containing gangue, and as amalgam (Ag 2 Hg 3 to Ag 36 Hg); the principal mercury mineral is cinnabar, HgS. It occurs also as calomel, HgCl , generally associated with cinnabar. Found in Austria, Spain, Peru, China, Russia, California, Texas and Oregon. 3. Preparation. The extraction of mercury from cinnabar, which may be considered as practically the only ore of this metal, is effected: (a) by oxidation with a regulated supply of air, and volatilization of the liberated metal, which distils over and is condensed: HgS + O 2 = Hg + SO 2 ; (b) by mixing the ore with lime, and distilling: 4HgS + 4CaO = 4Hg + 3CaS + CaSO 4 . (c) The ore is heated with iron (smithy scales): Hg, FeS , and SO 2 are produced. The mercury is usually condensed 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 writing or filter paper having a small pin-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). A. Oxides. Mercury forms two oxides, Hg.O 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 O . Mercuric o,mde, 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 3 or Hg(N0 3 ) 2 with about an equal weight of metallic mercury: Hg(N0 3 ), -f 3Hg = 4 HgO + 2NO; (3) by precipitating mercuric salts with KOH or NaOH . Made by (1) and (2) it is red, by (3) yellow. On heating it changes tp 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 movst 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, and the acid is cold and dilute, mercurous nitrate is formed; in a 1 ! 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 mercurous to mercuric, making a clear solution of Hg(NO 3 )2 , if there is sufficient HNO a present, other- wise 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, 38 MERCURY. 58, 5J. J. C., 1873,26, 476). H 2 SO 4 , concentrated at 25 has no action on Hg (Pitman. J. Am. Soc., 1898, 20, 100). With the hot concentrated acid S0 2 is evolved and Hg,SO 4 is formed if Hg be in great excess; HgSO 4 if the H 2 SO 4 be in excess. Hydrochloric acid gas at 200 is without action (Berthelot, A. Ch., 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 2 OCl 2 without evolution of hydrogen: ^Kg- -f 2HC1 + (X = Hg 2 OCL,H 2 O . Hydrobromic and hydriodic acids, gases. both attack mercury, evolve H , and form respectively HgBr and Hgl (Ber- thelot, I. c.). Hydrosnlphnric acid, dry gas, at 100 does not attack dry Hg (Berthelot, I.e.). H.vdrosulphnric 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 palts if the halogen be in excess, b. Oxides. Mercurous oxide, Hg 2 O , is a black powder insoluble in water or alkalis. Hydrochloric acid forms HgCl ; sulphuric acid forms Hg 2 SO 4 , changed by boiling with excess of acid to HgSO 4 ; nitric acid forms HgNO 3 , changed by excess of acid to Hg(NO 3 )2 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 red when produced by heating in the dry way and orange yellow when formed by precipitation with alkalis. 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 produces the white precipitate NH 4 Cl,NHgH 2 Cl + NHoHgCl (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 merctirous compounds and tben 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 Cl combine to form the non-dissociated HgCl 2 , 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 3 i 2 dissolves chloride, bromide, and iodide of Hg and Ag owing to the same cause, the Hg" ions of the strongly dissociated nitrate decreasing the already slight dissociatioSikof the mercuric haloids (44). The failure of HgCl a to give many of the pre- cipitation-re3||k>ns obtainable with other soluble mercuric salts is of course due to the same fact the sligTWoncentration 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. Cli., 1884 (6), 3, 429), soluble in Na 2 S,0, (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 combination of -the mercuric salt with the alkali salt formed, producing a double salt in which the mercury is present in the acid ion un- affected by the hydroxyl ion. Ammonium hydroxide and carbonate pre- cipitate from solutions of. mercurous salts, black mixtures of mercury and mercuric ammonium compounds. The same is true of the action of ammonium h} r droxide on insoluble mercurous salts : 2HgCl -|- 2NH 4 OH = Hg + NH,HgCl + 2H 2 -f HH 4 C1 ; 6HgN0 3 + 6tfH 4 OH = 3Hg + (NH,HgNO ; ) 2 HgO -f 4NH 4 tfO ! + 5H 2 ; 4Hg 2 S0 4 + SNH 4 OH = 4Hg + (HgH 2 N)oS0 4 .2HgO + 3(NH 4 ).,S0 4 + 6H,0 ; or uniting the salt in dif- ferent manner, 4HgCl + 4NH 4 OH == 2Hg -f Hg 2 NCl.NH 4 Cl + 2NH 4 C1 -f- 4H 2 . Examination with a microscope reveals the presence of Hg , The mercuric ammonium precipitate dissolves in a saturated solution of (NH 4 ).,S0 4 containing ammonium hydroxide and can thus be separated from the Hg (Francois, J. Pliarm., 1897 (6), 5, 388; Turi, Gazzctta. 1S9.S, 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, (NHoHg)Cl , also called nitrogen dihydrogen mercuric chloride (a) ; with hot solution and excess of ammonium hydroxide, dimercurammomum chloride, NHg 3 Cl, also called nitrogen dimercuric chloride (6) is formed, Treat- 40 MERCURY. 58, 6&. ing with fixed alkali hydroxide until no more ammonia is evolved changes the former compound to the latter (Pesci, /. 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, (7. N., 1889, 59, 23-t). A soluble combination of ammonium chloride with mercuric chloride, 2NH 4 C1. HgClo , 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 ),HgCL (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 + 4ELO (c) 2NH 4 Cl.HgCl 2 + 2KOH = (NH 3 ) 2 HgCL + 2KC1 + 2H 2 A solution of HgCl 2 in KI with an excess of KOH (Xessler's Reagent) is precipitated by NH 4 OH (or by ammonium salts), as NHg 2 I (207, 6k). Fixed alkali carbonates precipitate from mercurous salts an unstable )ncr- curous carbonate, Hg 2 CO 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-broicn 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 , 3 , or 4 ; with mercuric nitrate, a basic carbonate, (HgO) 3 HgCO 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 oxalatc, Hg 2 C 2 O 4 , white, slightly soluble in nitric acid: from solutions of mercuric salts, except HgCl 2 , '-mercuric 0). e. Hydrosulphuric acid and soluble sulphides, precipitate from mer- curous salts, mercuric sulphide, HgS , black, and nie-rcury, gray. Mercurous sulphide, Hg.JS , (loos 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 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-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 HNO, (distinction from all other metallic sulphides); insoluble in HC1 (Field, J. (7., 1860, 12, 158); soluble in chlorine (nitro-hydrochloric acid); insoluble in (NH 4 ) 2 S except when KOH or NaOH be present (Volhard, A., 1891, 255, 252); soluble in K 2 S (Ditte, C. r., 1884, 98, 1271), more readily if KOH be present (separation from Pb , Ag , Bi , and Cu) (Polstorff and Billow, Arch. Pharm., 3891, 229, 292). A little HgS (0.5-1.0 mg.) may dissolve in ammonium polysulphide when a large amount of mercury is present (A. A. Noyes, J., Am. Cliem. Soc., 29, 170). It is soluble in K 2 CS 3 (one part S , two parts CS 2 , and 23 parts KOH, sp. gr. 1.13) (separation from Pb , Cu , and Bi) ; reprecipitated as sulphide by HC1 (Rosenbladt, 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. (7., 1876, 29, 342). Mercuric chloride added to sodium thiosulphate forms a white precipitate, which blackens on standing 1 ; 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 rapidty added to HgCl, 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 of complex sulphite (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 mer- curosum as evidenced by HC1 . Mercuric nitrate with soluble sulphites forms a voluminous white precipitate, soluble in HNO S 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, 6/. 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 2 see next paragraph and (269, 8, footnote). /. 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(N"0 3 ) 2 and HgCL ; dissolved by chlorine or nitro- hydrochloric acid to HgCl 2 ; soluble in Hg(NO.,) 2 (57; footnote) (Dresehspi, 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 FIRST (SILVER) 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 (5& footnote), or in excess of the precipitant; hence, unless added in suitable proportions, no precipitate will be produced. Sulphuric acid does not transpose HgBr 2 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 HgI 2 , the latter being dissolved as a double salt with the soluble iodide: 2HgI 4- 2KI Hg + 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 HgI 2 (Francois, J. Pharm., 1897, (6), 5, 388). Mercuric salts are precipitated as mercuric iodide, HgI 2 , first reddish- 1)8, 7. MERCURY. 43 yellow then red, soluble in 24,814 parts of water at 17.5 (Bourgoin, A. CU., 1884, (6), 3, 439), 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).,HgI., 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 KIHgl., . 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),HgI, (sometimes designated the iodo-hydrargyrate of potassium), is precipitated by ammonium hydroxide as mercurammonium iodide, NHg 2 I ( Xcssler'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(Br0 3 ) 2 , white, soluble in nitric acid, hydrochloric acid, and in excess of mercuric nitrate, 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 (5&, footnote). lodic acid and soluble iodates precipitate solutions of mercurous salts as mwcurous 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(I0 3 ) 2 , white, soluble in HC1 , insoluble in HN0 3 and H 2 S0 4 , soluble in NH 4 C1 , trans- posed and then dissolved by KI . Mercuric chloride is not precipitated by KIO 3 (56, footnote) (Cameron, C. N., 1876, 33, 253). g. Arsenous acid or arsenites form a white precipitate with mercurous nitrate, soluble in HN0 3 (Simon, Pogg., 1837, 40, 442). Mercuric nitrate is precipitated by a solution of arsenous acid; the precipitate is soluble in HNO 3 (/>., 2, 2, 920). Arsenic acid or Na,HAsO 4 precipitates from mercurous nitrate :{Hg 3 As0 4 .HgNO 3 .H 2 , lig-ht yellow if the HgNO 3 be in excess (D., 2, 2, 921); dark red Hg 3 As0 4 if the arsenate be in excess. Hg 3 AsO 4 is changed by cold HC1 to HgCl and H 3 As0 4 , by boiling- with HC1 to Hgo , HgCl, , and H 3 AsO 4 ; and is soluble unchanged in cold HNO 3 , insoluble in water and acetic acid (Simon, Pogg., 1837, 41, 424). Arsenic acid and soluble arsenates precipitate from mercuric nitrate, Hg 3 (AsO 4 ) 2 , white, soluble in HNO 3 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, gray to black (a valuable final test for mercuric salts) (10). ft. Soluble chromates precipitate from mercurous solutions mercurous tfu'oniate, Hg,CrO 4 , brick-red, insoluble in water, readily transposed by HC1 to HgCl and H 2 Cr0 4 , soluble with difficulty in HNO 3 without oxidation (Richter, 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 KoCr.,O 7 . 7. Ignition. Mercury from all its compounds is volatilized by heat a* the 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 1 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 GROUP (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 SECOND 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 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 CO 2 . The sublimed mercury is condensed in a flask of water, and, after decanting- 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 (or H 3 P0 3 at 100) and dried as above, (ft) As mercurous chloride. It is first reduced to Hg' by H 3 PO 3 (Uslar, Z., 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 FeSO 4 in presence of NaOH : 2HgO + 2FeO + 3H 2 O = Hg,,0 + 2Fe(OH) s . H 2 SO 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 . (d) As HgO , by heating the nitrate in a bulb-tube in a current of dry air not hot enough to decompose the HgO. (e) Volumetrically, by Na 2 S 2 3 ; from the nitrate the precipitate is yelloiL-. from the chloride it is white: 3Hg(N0 3 ) 2 + 2Na 2 S 2 3 + 2H 2 O = Hg 3 S 2 (NO 3 ) 2 + 2Na,SO 4 + 4HNO 3 SHgCl, -f 2Na 2 S 2 3 + 2H 2 Hg 3 S 2 CL -f 2N;a 2 S0 4 -f 4HC1 . (f) Volumetrically, HgCL is reduced to Hg 2 by FeSO 4 in presence of KOH , and after acidulating- with H 2 SO 4 the excess of FeS0 4 is determined by K 2 Cr 2 O- or KMnO 4 (Jliptner, C. C., 1882, 727). (#) By iodine. It is converted into HgCl and then dissolved in a graduated solution of I dissolved in KI : 2HgCl + OKI + I a = 2K 2 HgI 4 + 2KC1 . The excess of iodine is determined by Na,,S,O 3 . (7i) The measured solution of HgCL, is added to a graduated solution of KI: 4KI -f- HgCL = XoHgl. + 2KC1 . The instant- the amount of HgCL shown in the equation is exceeded a red precipitate of HgI 2 appears, (t) Volumetric, 59, 2. SILVER. 45 by adding a few drops of ammonium hydroxide to HgCl, and then titrating with standard KCN , the ammonium hydroxide precipitate disappears when the mercury becomes Hg(CN) 2 (Haimay, J. C., 187H, 26, 570; Tuson, J. C 1 ., 1877, 32, 679). (;) Electrolytwally, by obtaining the mercury as HgNO 3 , Hg(NO 3 ) 2 , or Hg 2 SO 4 and precipitating as Hg on platinum by the electric current. Mercuric chloride cannot be used, as it is partly reduced to HgCl , and that is not readily reduced to Hg by the electric current (Hannay, I.e.). 10. Oxidation. Free mercury (Hg) precipitates Ag, An , and Pt from their solutions, and reduces mercuric salts to mercurous salts (Hada, J. C., 1896, 69, 1667). Potassium permanganate in the cold oxidizes the metal to Hg 2 , when hot to HgO (Kirchmann, J. C., 1873, 26, 476). Mercury and mercurous salts are oxidized to mercuric salts by Br , Cl , I , HN0 3 , H 2 S0 4 (concentrated and hot), and HC10 ;5 . Reducing agents, as Pb , Sn , Sn", Bi , Cu , Cu', Cd , Al , Fe , Co , Zn , Th 1 , Mg, H 3 P0 2 , H 3 P0 3 and H 2 S0 3 , precipitate, from the solutions of mercuric and mercurous nitrates, dark-gray Hg ; from solution of mer- curic chloride, or in presence of chlorides, first the white., HgCl , then gray Hg. Strong acidulation with nitric acid interferes with the reduction, and heating promotes it. The reducing agent most frequently employed is stannous chloride: 2HgCL + SnCl 2 = 2HgCl + SnCl, 2HgCl -f SnCL = 2Hg + SnCl 4 or HgCL + SnCl 2 = Hg + SnCl 4 also 2Hg(N0 3 ) 2 + SnCl 2 2HgCl + Sn(N0 3 ) 4 A clean strip of copper, placed in a slightly acid solution of a salt of mer- cury, becomes coated with metallic mercury, and when gently rubbed with cloth or paper presents the tin-white lustre of the metal; the coating being driven off by heat; 2HgN0 3 + Cu = 2Hg + Cu(N0 3 ) 2 . Formic acid reduces mercuric to mercurous chloride, and in the cold does not affect further reduction. Dry mercuric chloride, moistened with alcohol, is reduced by metallic iron, a bright strip of which is corroded soon after immersion into the powder tested (a delicate distinction from mercurous chloride). 59. Silver (Argentum) Ag = 107.88. Monovalent. 1. Properties. Specific (/rarity 10.512 heated in vacuo (Dumas, C. N., 1878, 37, 82). Melting point, 960.7 (Heycock and Neville, /. C., 1895, 67, 1024). Does not appreciably vaporize at 1567 (V. and C. Meyer, B., 1879, 12, 1428). It is the whitest of metals, harder than gold and softer than copper. Silver is hardened by copper; United States silver coin contains 90 per cent silver and 10 per cent copper. In malleability and ductility it is inferior only to gold; and as a con- ductor of heat and electricity it exceeds all other metals. 2. Occurrence. Found in a free state in United States, Mexico, Peru, Siberia, etc.; alone, and with gold as a component of other minerals, e. g., galena, pyrite, phalcopyrite, and many other ores. The most important silver minerals are iRejd, c, N,, 1865, 12, 242; Neumann, J. .., J875, 38, J32, CALIFORNIA COLLEGE 46 SILVEE. 59, 3. argentite, Ag 2 S, stephanite, Ag 5 SbS 4 , pyrargyrite, Ag 3 SbS 3 , proustite, Ag 3 AsS 3 , cerargyrite, AgCl . 3. Preparation. (a) Argentiferous ores are smelted with lead ores, coke and limestone in a blast furnace; silver (and gold) alloys with the reduced lead, and is subsequently separated from it by Parkes' or Betts' process. (6) It is amalgamated with mercury and the mercury separated by distillation, (c) It is brought into solution and the metal precipitated by copper, (d) It is very easily reduced from the oxide or carbonate by heat alone, and from all its [com- pounds by ignition with hydrogen, carbon, carbon monoxide and organic compounds. 4. Oxides. Silrer oxide, Ag 2 O , argentic oxide, is formed by the action of alkali hydroxides on silver salts or by heating 1 the carbonate to 200. It is a brown powder, a strong oxidizing agent, decomposed at 300 into metallic silver and oxygen. Concerning the existence of argentous oxide, Ag 4 O , and silver peroxide, Ag*.>O 2 , and their properties, see Muthmann (J5., 1887, 20, 983) ; Pford- ten (B., 1887, 20, 1458) and Bailey (C. N., 1887, 55, 263). 5. Solubilities. a. Metal. The fixed alkali-s do not act upon silver, hence silver crucibles are used instead of platinum for fusion with caustic alkalis. Ammonium hydroxide dissolves finely divided silver, no action if air be excluded. Acetic acid is without action (Lea, Am. ., 1892, 144, 444). Nitric acid is the ordinary solvent for silver, the 50 per cent acid being most effective, while the dilute acid free from nitrous acid has little or no action (Lea, I. c.); silver nitrate is formed, nitrogen peroxide being the chief product of the reduction of the nitric acid (Higley and Davis, Am., 1897, 18, 587). Silver is not oxidized by water or air at any temperature; it is attacked by phosphorus or by sub- stances easily liberating phosphorus; it is tarnished in contact with hydrosul- phuric acid, soluble sulphides, and many organic compounds containing sulphur; except that pure dry hydrosulphuric acid is without action upon pure dry silver (Cabell, C. N., 1884, 50, 208). Dilute sulphuric acid slowly dissolves finely divided silver (Lea, 1. c.), a sulphate being formed while, with the hot concen- trated acid, sulphur dioxide is evolved. Hydrochloric acid, sp. gr., 1.20, is without action upon pure silver, but the metal is readily attacked by chlorine, bromine or iodine, b. Oxide. Silver oxide, Ag 2 O , soluble in 3000 parts of water, com- bines with nearly all acids, except CO 2 , forming the corresponding salts. The hydroxide is not known. c. Sails. Silver forms a greater number of insoluble salts than any other known metal, though in this respect mercury and lead are quite similar. The nitrate is very soluble in water, 100 parts water dissolv- ing 227.3 parts at 19.5, soluble in glycerol, and sparingly soluble in alcohol and ether. The chlorate dissolves in about ten parts cold water; the acetate in 100 parts; the sulphate in about 200 parts cold water and 88 parts at 10G, and is more soluble in nitric or sulphuric acid than in water; the borate, thiosulphate, and citrate are sparingly soluble in water. The oxalate, tartrate, carbonate, cyanide, ferrocyanide, ferricyanide, phos- phate, sulphide, sulphite, chloride, bromide, iodide, iodate, arsenite, arse- nate, and chromate are insoluble in water. The chloride is soluble in 244 parts HC1 , but its solubility is very much lessened by the presence of mercurous chloride (Ruyssen and Varenne, Bl. 3 1881, 36, 5). If a solution of silver nitrate be dropped into concentrated hydrochloric acid no precipitate appears until one half per cent of the HC1 becomes AgCl (Pierre, J. C., 1872, 25, 123). Concentrated nitric acid upon long continued boiling scarcely attacks AgCl (Thorpe, J. C., 1872, 25, 453); sulphuric acid, sp. gr. 1.84, completely transposes even the fused ;<59, 66. SILVER. 47 chloride on long boiling (Sauer, J. C. } 1874, 27, 335). Silver chloride is also soluble in ammonium hydroxide and carbonate; in sodium chloride forming a double salt; in a concentrated solution of mercuric nitrate (68, 1; 58, 56 footnote); and in many other metallic chlorides and alkali salts to a greater or less extent. All the salts of silver which are insoluble in water are soluble in ammonium hydroxide, except the sulphide and iodide; in ammonium carbonate, except the bromide, iodide, and sulphide, the bromide very slightly soluble; in cold dilute nitric acid, except the chloride, bromide, bromate, iodide, iodate, cyanide, and thio- c van ate; in a solution of potassium cyanide (and by many other cyanides) except the sulphide; and in alkali thiosulphates almost without exception. 6. Reactions, a. The fixed alkali hydroxides precipitate from solu- tions of silver salts (in absence of citrates), silver oxide, Ag 2 , grayish brown, insoluble in excess of the reagents; soluble in acids, alkali cyanides, and thiosulphates; somewhat soluble in ammonium salts. Most silver salts are transposed 011 boiling with the fixed alkalis, except the iodide, which is not thus transposed (Vogel, J. C., 1871, 24, 313). Ammonium hydroxide, in neutral solutions of silver salts, forms the same precipitate, Ag,0 , very easily dissolving in excess, by formation of silver ammonium hydroxide, NH 3 AgOH : AgN0 3 + 2NH 4 OH = NH 8 AgOH + NH 4 N0 3 + H 2 (Prescott, J. Am. Soc., 1880, 2, 32). In solutions con- taining much free acid, all precipitation is prevented by the ammonium salt formed with the formation of silver ammonium nitrate, NH 3 AgN0 3 or in the presence of excess of ammonia as (NH 3 ).,AgN0 3 . Alkali carbonates precipitate silver carbonate, Ag 2 C0 3 , white or yellow- ish white, very slightly soluble in water and in the fixed alkali carbonates, readily soluble in ammonium hydroxide -and carbonate, transposed by inorganic acids forming the corresponding salts. Carbon dioxide does not transpose silver salts. &. Oxalic acid and soluble oxalates precipitate silver oxalatc, Ag,,C 2 O 4 , white, slightly soluble in water, soluble with difficulty in dilute nitric or sulphuric acids, readily soluble in ammonium hydroxide. When heated it decomposes with detonation, forming- metallic silver. Potassium cyanide precipitates from neutral or slightly acid solutions jilrrr cyanide., AgCN , white, quickly soluble in excess of the reagent as silver potassium cyanide, AgCN.KCN . Hydrocyanic acid precipitates solutions of silver salts but the precipitate does not dissolve in excess of the reagent. Silver cyanide is transposed by H L ,S0 4 or HC1 and is soluble in ammonium hydroxide and carbonate (Schneider, J. pr., 18(58, 104, 83). The ready solubility of nearly all silver compounds in potassium cyanide (5r) affords a means of separating silver from many minerals. Potassium ferrocyanide precipitates silver ferrooyanide, Ag 4 Fe(CN) e , yellow- ish white, soluble with difficulty in ammonium hydroxide and carbonate; 48 SILVER. 59, 6c. metallic silver separates on boiling" and a ferricyanide is formed. The ferro- cyanide is not decomposed by hydrochloric acid, but it is changed to the ferricyanide by nitric acid. Exposure to the air gives it a blue tinge. Potas- sium ferricyanide precipitates xiln-r fcrrici/anide, Ag 3 Fe(CN) 6 , reddish yellow, readily soluble in ammonium hydroxide and carbonate. Potassium thlocyanate gives silver thiocyanate, AgCNS , white, soluble in ammonium hydroxide and carbonate, insoluble in dilute acids. Concentrated sulphuric acid with the aid of heat dissolves silver thiocyanate when some free silver nitrate is present. This may be used as a separation from silver chloride, which is transposed by hot concentrated sulphuric acid only on long-continued boiling (5c). To effect this separation a little silver nitrate should be added to the silver precipitates and then concentrated sulphuric acid and heat. To avoid danger of decomposition of the chloride the mixture should not be heated above 200. The pure silver thiocyanate (silver nitrate being absent) is decomposed by hot concentrated sulphuric acid with formation of a black precipitate containing- silver. c. Silver nitrate is soluble in 500 parts of concentrated nitric acid (Schultz, Z. Ch., 1869, 531), and is precipitated from its concentrated water solutions by the addition of concentrated nitric acid. <1. Disodium phosphate precipitates silver phosphate, Ag- 3 P0 4 , yellow, soluble in dilute nitric acid, in phosphoric acid, and in ammonium hydroxide and carbonate; but little soluble in dilute acetic acid. Sodium pyrophosphate precipitates silver /)//ro/>/f06'y)7/e, white, same solubilities as the orthophosphate. e. Hydrosulphuric acid and soluble sulphides precipitate from neutral, acid or alkaline solutions silver sulphide., Ag.,8 , black, soluble in moderately strong nitric acid (distinction from mercur}*), slightly soluble in potassium cyanide (distinction from copper), insoluble in alkali sulphides (distinction from arsenic, antimony, and tin). Certain insoluble sulphides form silver sulphide from solutions of silver nitrate,* e. g., cupric sulphide gives silver sulphide, cuprous sulphide gives silver sulphide and metallic silver, in both cases cupric nitrate resulting (Schneider, J. C. t 1875, 28, 133 and 612). Thiosulphates precipitate silver thiosulphate, Ag ? S 2 O 3 , white, unstable, readily soluble in excess of the precipitant, by formation of double thiosulpahtes; with excess of sodium thiosulphate Na4Ag 2 (S 2 O 3 ) 3 is formed (Cohen, J. C., 1896, 70, ii, 167). Silver thiosulphate turns black on standing or heating; Ag ; S L O + H 2 O = Ag 2 S + H 2 SO 4 . Sulphurous acid and soluble sulphites precipitate silver sulphite, Ag 2 SO 3 , white, readily soluble in excess of alkali sulphite or in dilute nitric acid; on boiling precipitated as metallic silver with formation of sulphuric acid. Sulphuric acid and soluble sulpht'ies precipitate silver sul- phate, Ag.jSO.1 , white, from concentrated solutions of the nitrate or chlorate; sparingly soluble in water, quite soluble in concentrated sulphuric acid. /. Hydrochloric acid and soluble chlorides precipitate silver chloride, AgCl , white, curdy ; separated on shaking the solution ; turning violet to brown on exposure to the light; fusible without decomposition; very easily soluble in ammonium hydroxide as ammonia silver chloride, (NH 3 ) 3 (AgCl) 2 (Jarry, C. r., 1897, 124, 288), according to the following equation : f 2AgCl + 3NH 4 OH - 3NH 3 .2AgCl + 3H 2 O. On acidifying the solution with nitric acid the silver chloride is repre- cipitated as follows : 3NH 3 .2AgCl + 3HNO 3 = 2AgCl + 3NH 4 NO 3 . If mercurous chloride be present with silver chloride the solubility in ammo- nium hydroxide is greatly lessened, in fact a great excess of mercurous * Ag2$ is one of the least soluble of the sulphides. See 57, 6e, footnote. 59, 7. SILVE1L 49 chloride may entirely prevent the solution of silver chloride in ammonium hydroxide by forming metallic silver. AgCl + :$HgCl + 4NH 4 OH = Ag + 2Hg> + 2NH 2 HgCl + 2NH 4 C1 + 4H 2 0. Silver chloride is quite soluble in a solution of mercuric nitrate, which, if present in large excess, may entirely prevent the precipitation of the silver chloride by hydrochloric acid. The precipitation by hydrochloric acid (in absence of a great excess of Hg(NO ;i ) 2 ) is the most delicate of the ordinary tests for silver, being recognized in 250,000 parts of water. As mercuric salts are not at all pre- cipitated by HC1 and lead salts only imperfectly, silver is the only metal which belongs exclusively to the FIRST OR SILVER GROUP OF BASES (16). Hydrobromic acid and soluble bromides precipitate silver bromide, AgBr , white, with a slight yellowish tint; but slightly soluble in excess of alkali bromides, and much less easily soluble in ammonium hydroxide than silver chloride. If silver nitrate be added to a bromide containing an excess of am- monium hydroxide, the precipitate which first forms readily dissolves on shak- ing", no solution is obtained with the iodide. Hydriodic acid and soluble iodides precipitate silver iodide, Ag-I , pale yellow, soluble in excess of the concentrated reagents by formation of double iodides, as KIAgT , which are decomposed by dilution with much water. The precipi- tate dissolves in 26,000 parts of ten per cent ammonium hydroxide; not at all in a five per cent solution (Longi, Gazzetta, 1883, 13, 87). It is insoluble in dilute acids, but is decomposed by hot concentrated nitric or sulphuric acids. Silver bro-mate formed by adding- potassium bromate to silver nitrate is soluble in about 600 parts water and in 320.4 parts nitric acid (sp. gr. t 1.21) at 25, and readily soluble in ammonium hydroxide. Silver iodate formed in manner simi- lar to the bromate is soluble in about 28,000 parts water and in 1044.3 parts nitric acid (sp. gr., 1.21) at 25, and readily soluble in ammonium hydroxide (Longi, I.e.). 0. Soluble arsenites precipitate silver arsenite, Ag 3 As0 3 , yellow, very readily, soluble in dilute acids and in ammonium hydroxide. Soluble arsenates precip.- tate silver (irtsntatc, Ag 3 As0 4 , red-brown, soluble in ammonium hydroxide, nitric acid, arsenic acid, and almost insoluble in acetic acid. A solution of alkali stannite as K 2 Sn0 2 precipitates metallic silver from solutions of silver salts. A solution of silver nitrate in a great excess of ammonium hydroxide constitutes a very delicate reagent to detect the presence of tin in the stannous condition in the presence of fixed alkalis; antimony does not interfere if a great excess of ammonium hy- droxide be present. li. Chromates and dichromates, as K 2 Cr0 4 and K 2 Cr 2 O 7 , precipitate silver chromate, Ag 2 Cr0 4 , dull-red, sparingly soluble in water and in dilute nitric acid, soluble in ammonium hydroxide. 7. Ignition. Silver nitrate melts undecomposed at 218, at a red heat it is decomposed into Ag , O , N , and NO (Fischer, Pogg., 1848, 74, 120). Silver chloride fuses at 451, the bromide at 427, and the iodide at 527. On charcoal with sodium carbonate, silver is reduced from all its compounds by the blow- pipe, attested by a bright malleable globule. Lead and zinc, and elements more volatile, may be separated from silver by their gradual volatilization under the blow-pipe, or in the assay furnace (see Cupellation in works on the assay of the precious metals). 50 SILVER. 59, 8. 8. Detection. Silver is identified by its precipitation with hydrochloric acid, the insolubility of the precipitate in hot water, and its solubility in ammonium hydroxide, with reprecipitation on rendering acid with nitric acid (61). 9. Estimation. (a) As metallic silver, into which it is converted by direct ignition if it is the oxide or carbonate, or by ignition in hydrogen if the chloride, bromide, iodide or sulphide (Vogel, J. C., 1871, 24, 1009). (b) It is precipitated as Ag<31 , and after igniting- to incipient fusion, weighed, (c) It is converted into Ag,S by H 2 S , and weighed after drying- at 100; inadmissible in case of an acid that might liberate free sulphur, (d) Add KCN until a solution of KAg(CN), is formed, precipitate with HNO 3 , and after drying at 100, weigh as AgCN . (e) Volumetrically, by adding a graduated solution of NaCl until a precipitate is no longer formed. This may be varied by adding the measured silver solution to the graduated NaCl solution, containing a few drops of K 2 Cr0 4 , until the red precipitate begins to form, (f) Volumetrically, add a graduated solution of ammonium thiocyanate, containing ferric sulphate, until the red color ceases to disappear, (g) Add the measured silver solution to a standard solution of KCN until a permanent white precipitate is formed. 10. Oxidation. Metallic silver precipitates gold and platinum from their solutions, reduces cupric chloride to cuprous chloride, 1 mercuric chloride to mercurous chloride, and permanganates to manganese dioxide 2 . Silver is precipitated from its solutions by: Pb , PbS 3 . Kg , As 4 , AsH 3 , Sb , SbH 3 , Sn , Sn", Bi , Cu , Cu' 8 , Cd , Te , 2'e , FeS 1 , Al , Mn , Zn , Mg , P 4 , PH 3 , H,P0 2 , H 2 S0 3 , SiH 4 , H 2 2 6 , and H (very slowly) 7 . In alkaline mixture silver is also reduced by Hg', As"', Sb'", Bi'", and Mn". An amalgam of mercury and tin reduces insoluble compounds of silver in the wet way, the silver amalgamates with the mercury and the tin becomes Sn IV (Laur, C. r., 1882, 95, 38). Ferrous sulphate in the cold incompletely reduces silver salts; on boiling, the ferric salt formed is reduced and the silver dissolved (Lea, 7. c.). In the gradual reduction of silver by certain organic reagents, the metal is obtained as a bright silver coating or mirror upon the inner surface of the test tube or other glass vessel. Usually a slight^ ammoniacal solution of silver nitrate is used and allowed to stand some time with the reagent; such as alcoholic solution of oil of cloves or cassia, formic acid, aldehyde, chloral, tartaric acid, etc. (ienlle warming facilitates the result. If a good mirror is desired, great care must be taken to free the inner surface of the glass from all organic impurities by careful washing with ether, chloroform, etc. In these deoxidations, generally the nitric acid radical of the silver nitrate is not decomposed, but nitric acid is left: 4AgNO 3 + 2HoO = 4Ag- -f 4HNO 3 + 0, . Light acts upon nearly all salts of silver when mixed with gelatine or other organic substances used in preparing photographic plates, etc. These plates contain various silver salts, frequently the bromide or iodide, or both together. The nature of the chemical change is not fully understood. It has been shown, however, that silver chloride on exposure to the light loses chlorine, and there is considerable evidence to prove that when the silver halides are acted upon by light, a subhalide such as Ag 2 Cl , Ag 2 Br or Ag 2 I is formed. When the pi ate which has been exposed to light is treated with a reducing agent, the reduc- tion of the silver is carried to the metallic state, the black silver producing the image. The nitrate in crystal or pure water solution, the phosphate, bromide, 1 Lea, Am. S., 1892, 144, 444. " /)., 2, 2, 759. s Skey, C. N., 1871, 23, 232. 4 Senderens, C. r., 1887, 1O4. 175. 6 r>.. 2. 1, 4515. Rietfler, J. C., 1896, 7O, ii, 471. 7 Pellet, B., 1874, 7, 656 ; Sehwarzenbach and K ritsvlK'u-sky, ;.., ISSU. ar,, 1574 ; < 'ooke, C. N.. 1888. 58. 103. Millon. Am. .S., 18i, . 417. ;60. COMPARISON OF REACTIONS OF METALS OF THE SILVER GROUP. 51 8 -I o "^^ '43 ^ I I be M fc W CO W w w oo = * 0> od T in w w &c 6fl WW W - 5 W W W &flr5 be be WWW Vo W LI ft :i :B T3 " ( W M f 2 ti gj s TABLE FOR ANALYSIS OF THE SILVER OR FIRST GROUP. 61. 8 | m c 5 CD m 02 is H Sll gto ^ s ^ EH 2 a srii r^T3 111 - : I !?* 5 IP & ~ 8& ajlJ U 1 in iaf-1 P g |* o-gS^J 3 * - -all s.l-= 5 f -g"^'^ 8 l^i - l^3l MEaJfa'P 3 x.-Sl?| ! 5 I fe -T1 -" 1 "*'! S QJ e3 CQ =;} -c ft ^ PQ 2 S g tin e from dissolv 3 il.l Si! ill 63, 60. DIRECTIONS FOE ANALYSIS WITH NOTES. 53 iodide and cyanide are not decomposed by light alone; but light greatly hastens their decomposition by organic substances, or other reducing agents, as of solution of silver nitrate in rain water, or \\ritten as an ink upon fabrics. Silver is the base of most indelible inks. DIRECTIONS FOR THE ANALYSIS OF THE METALS OF THE FIRST GROUP. 62. Manipulation. To the solution acid with nitric acid add hydro- chloric acid (whenever directions call for the addition of a reagent it is to be used reagent strength 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 , Hgf, and Ag . Shake thoroughly and allow to stand 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' 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 2 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 filtrate 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. It will be noted that a number of these are alkaline solutions and will give precipitates with other acids than HC1. It is advisable in these cases to acidify with HNO 3 before adding HC1 in order to avoid error from this source. 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, G/ 1 ). 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). 54 DIRECTIONS FOR ANALYSIS WITH NOTES. 63, 6/. c. Solutions of metallic oxides in the alkali hydroxides are precipitated when neutralized with acids, e. g., K 2 ZnO 2 + 2HC1 = Zn(OH) 2 + 2KC1 . d. r \ he 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, 3o). f. Certain soluble double cyanides, as Ni(CN) 3 .2KCN' , are precipitated as insoluble cyanides, Ni(CN) 2 , on the addition of HC1 (133, (>7>). g-. Solutions of silicates (249, 4), borates, tuiigstates, molybdates; also benzoates, salicylates, urates, and certain other organic salts, are precipitated by acidulation with HC1 , many of the precipitates being- soluble 011 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, M) ; mixture of solutions of KI and KI0 3 (280, 6,5,7), 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, 5tt and (te), a second group mercury compound insoluble in HNO 3 . 9. If but one metal of the first group be present, the action of NH 4 OH 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 , K 2 Cr 2 7 , and KI (57, 6 e, h, and /) giving white, black, and yellow precipitates : The yellow precipitate with potassium iodide (the KI must not be used in great excess (57, 5c)) should he 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 PbI 2 (charac- teristic of lead). 65. Notes. 1. Lead is never completely precipitated in the first group (57, Of). 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 4 OH , part remaining on the filter and part carried through mechanically which causes turbidity to the am- 68, 3. DIRECTIONS FOR ANALYSIS WITH NOTES. 55 monium hydroxide solution of the AgCl and makes necessary the filtration of that solution before the addition of HNO 3 , otherwise it does not interfere. 5. The precipitation of lead as the sulphide while not characteristic of lead, is exceedingly delicate, much more so than the formation of the white PbSO 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. PbCrO 4 is blackened by alkali sulphides and dissolved by the fixed alkalis (important dislinclioii from BaCr0 4 ); the solubility in the fixed alkalis is also an important distinction from bismuth chromate (76, Gh). 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 , and some PbCl 2 if it has not been well washed. To this precipitate NH^OH , 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 evidence of the presence of mercurous mercury. The AgCl is dissolved by the NH 4 OH , and is found in the filtrate ; its presence being confirmed by its reprecipitation on rendering the solution acid with HN0 3 . 67. Notes. Mercury. 1. The black precipitate on the filter, caused by the addition of NH 4 OH to the HgCl may be examined under the microscope for the detection of globules of Hg, or the precipitate may be digested with con- centrated solution of (NH 4 ) 2 SO 4 , which dissolves the NH 2 HgCl , leaving the Hg unattacked (58, 60). 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, h and d). 3. 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). 4. 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(N0 3 ) 2 prevents the precipitation of AgCl from solutions of silver salts by HC1 (59, 5<0- In this case the metals should be precipitated by H,S and the well-washed precipitate digested with hot dilute HNO 3 . The silver "is dissolved as AgNO,, . while the mercury is unattacked: 6Ag,S -f 16HNO 8 = 12AgNO 3 + 3S 2 + 4NO + 8H 2 O . After evaporation of the excess of HNO 3 the solution may be treated with HC1 as an original solution. . A small amount of AgCl with a large amount of HgCl is not dissolved by NH 4 OH , but is reduced to Ag by the Hg formed bv the addition of the 1TH 4 OH to the HgCl (58, 0- Oxides. Arscnous oxide exists in two forme, crystalline and amorphous, the solubilities of which differ considerably (27). At ordinary temperature 100 parts of water dissolve 3.7 parts of the amorphous and 1.7 parts of the crystal- line, several hours being necessary to effect the solution. 100 parts of boiling water dissolve 11.46 parts of the amorphous and 10.14 parts of the crystalline oxide in three hours (Winkler, J. pr., 1885, (2), 31, 247). The presence of acids greatly increases the solubility in water (Schultz-Sellac, B., 1871, 4, 109). Arsenous oxide is readily soluble in alkali hydroxides or carbonates to arsenites 58 ARSENIC. 69, 5c. (Clayton, C. N., 1891, 64, 27), Arsenic 'pcnto,ri(lc, As 2 O n , is deliquescent, soluble in water forming H 3 As0 4 . The meta and pyro acids are easily soluble in water forming the ortho acid (Kopp, A. Ch., 1856, (3), 48, 106). c. Salts. Arsenic does not act as a base with oxyacids, but its oxides combine with the metallic oxides to form tw r o classes of salts, arsenites and arsenates. Arsenites of the alkalis are soluble in water, all others are insoluble or only partially so; all are easily soluble in acids. Alkali arsenates, and acid arsenates of the alkaline earths, are soluble in water; all are soluble in mineral acids, including H 3 AsO 4 (LeFevre, C. r., 1889, 108, 1058). See also under the respec- tive metals. Arsenous sulphide, As 2 S 3 , is insoluble in water when prepared in the dry way ; when prepared in the moist way it may be transformed into the soluble colloidal * form by treatment with pure water, from which solutions it is precipitated by solutions of most inorganic salts or acids (Schulze, J. pr., 1882 (2), 25, 431). The presence of acids or solutions of salts prevents the solubility of As 2 S 3 in water. Boiling water slowly decomposes the sulphide forming As 2 3 and H 2 S (Field, C. N., 1861, 3, 115.; Wand, Arch. Phar., 1873, 203, 296). It is completely decomposed by gaseous HC1 form- ing AsCl 3 (Piloty and Stock, R., 1897, 30, 1649), very slightly decomposed and the arsenic dissolved by hot concentrated acid (Field, I. c.). Chlorine water and nitric acid decompose it readily with formation of H 3 As0 4 ; with sulphuric acid As 2 3 and S0 2 are formed (Rose, Pogg., 1837, 42, 536). The alkali hydroxides or carbonates dissolve it readily with formation of RAs0 2 and EAsS 2 (E = K , Na and NHJ V., 2, 1, 183) ; soluble in alkali sulphides and poly-sulphides forming R 4 As 2 S 5 , and RAsS 2 (Berzelius, Pogff.-, 1826, 7, 137 ; Nilsson, J. 0. , 1872, 25, 599). Whether the ortho, meta or pyro salt is formed, depends upon the amount of alkali sulphide present Arsenic sulphide, As 2 S 5 , is insoluble in water ; soluble in HC1 gas, as AsCl 3 ; insoluble in dilute HC1 , soluble in HN0 3 or chlorine water, as H 3 As0 4 ; soluble in alkali hydroxides and carbonates, as R 3 AsS 4 and R 3 As0 3 S : As 2 S 5 -f 6NH 4 OH = (NH 4 ) 3 AsS 4 + (NH 4 ) 3 As0 3 S + 3H 2 (Mc- Cay, Ch. Z., 1891, 15, 476); soluble in alkali sulphides, as K 3 AsS 4 (Nilsson, J. pr., 1876 (2), 14, 171). Arsenous chloride, bromide and iodide (AsCl 3 , AsBr, , Asl,) are decomposed by small amounts of water into the corresponding oxyhalogen compounds, AsOCl , etc. A further addition of water decomposes these compounds into arsenous oxide and the halogen acids. 6. Reactions. a. The alkali hydroxides and carbonates unite with arsenous and arsenic oxides (acids), the latter with evolution of carbon dioxide, forming- soluble alkali arsenites and arsenates. These alkali salts are chiefly meta arse- nites and ortho arsenates (Bloxam, J. (7., 1862, 15, 281; Graham, Pogff., 1834, 32, 47). * Colloids is a name given by Graham to a class of glue-like bodies in distinction to the crystal- loids, which have a weU-defined solid form. The colloids are indefinitely soluble in water, giving the little-understood " pseudo-solutions," which stand midway between the mechanical suspension or emulsion and the true solution. Gelatine, starch, the metallic sulphides, silicic acid, and the hydroxides of iron and aluminum are some of the substances that may take on the colloid torn, TBe colloid lolutioni are ai a rule broken up by addition of an ei4 or a neutral H4t, 69, 6e. ARSENIC. 59 ft. Oxalic acid does not reduce arsenic acid* (Nay lor and Braithwaite, Pharm. ./. Trans., 1883, (3), 13, 464). Potassium ferricyaiiide in alkaline solution oxi- dizes arsenous compounds to arsenic compounds, very rapidly when gently warmed, c. Nitric acid readily oxidizes all other compounds of arsenic to arsenic acid. d. Hypophosphites in presence of concentrated hydrochloric acid reduce all oxycompounds of arsenic to the metallic state. 0.00001 gram oi arsenic may be detected by boiling with 10 cc. strong hydrochloric acid and 0.2 gram calcium hypophosphite (Engel and Bernard, C. r., 1896, 122, 390; Thiele and Loof, C. C., 1890, 1, 877 and 1078; and Hager, J. C., 1874, 27, 868). e. Hydrosulphuric acid precipitates the lemon-yellow arsenous sulphide, As 2 S 3 , from acidulated solutions of arsenous acid. The precipitate forms in presence of concentrated hydrochloric acid. Citric acid and other organic compounds hinder the formation of the precipitate, but do not wholly prevent it if strong hydrochloric acid be present. Nitric acid should not be present in strong excess as it decomposes hydrosulphuric acid, with precipitation of sulphur. In aqueous solutions of arsenous acid the sulphide forms more as a yellow color than as a precipitate, being soluble to quite an extent in pure water, especially when boiled (5c) : As,S 3 + 3H,0 = As 2 3 + 3H 2 S . This has been given as a method of separating arsenous sulphide from all other heavy metal sulphides (Clermont and Frommel, /. C., 1879, 36, 13). The precipitate is not formed in solutions of the arsenites except upon acidu- lation. The hydrogen sulphide converts the oxy salts of arsenic into the thio salt, which is decomposed by acid with precipitation of the sulphide of arsenic : Na.iAsOa + 3H 2 S = Na 3 AsS 3 + 3H 2 O 2Na 3 AsS 3 + 6HC1 = As 2 S 3 + 3H 2 S Alkali sulphides produce and, by further addition, dissolve the precipi- tate (5c) : 2AsCl 3 + 3(NH 4 ) 2 S = AS 2 S 3 + NH 4 C1 As 2 S 3 + 2(NH 4 )S = (NH 4 ) 4 AS2S5 or A 2 S 3 + (NH 4 ) 2 S = 2NH 4 AsS 2 Arsenous sulphide is also soluble in alkali hydroxides and carbonates, with evolution of C0 2 , forming arsenites and thioarsenites (5c). The thioarse- nites are precipitated by acids forming As 2 S 3 : (NH 4 ) 4^8085 + 4HC1 = As 2 S 3 + 2H 2 S + 4NH 4 C1 or 2NH 4 AsS t + 2HC1 = As 2 S 3 + H 2 S + The solubility of the sulphides of arsenic in yellow ammonium sulphide separates arsenic with antimony and tin from the other more common metals of the second group; and the solubility in ammonium carbonate effects an approximate separation from antimony and tin (Hager, J. C., 1885, 48, 838). Arsenous sulphide is soluble in solutions of alkali sul- phites containing free sulphurous acid (separation from antimony and tin): 4As 2 S 3 + 32KHSO,, = 8KAs0 2 + 12K,S,0, + 3S 2 + US0 2 + 16H 2 0. It may also be separated from antimony and tin by boiling with strong hydrochloric acid, the As 2 S.< remaining practically insoluble; the sulphides of antimony and tin being dissolved. It is easily dissolved by strong 60 AltsmiC. 69, 6/1 hydrochloric acid, the As S ; remaining practically insoluble; the sulphides of antimony and tin being dissolved. It is easily dissolved by strong nitric acid, and by free chlorine or nitrohydrochloric acid, as arsenic acid: 6As 2 S 3 + 20HN0 3 + SH 2 = 12H 3 As0 4 + 9S 2 + 20NO ; 2As 2 S 3 + 10CL, + 16H 2 =3 4H 3 As0 4 + 3S 2 + 20HC1 . Usually a portion of the sulphur is oxidized to sulphuric acid, completely if the nitric acid or chlorine be in great excess and heat be applied: As 2 S 3 + 14C1 2 + 20H 2 = 2H 3 As0 4 + 3H 2 S0 4 + 28HC1 . Arsenic pentasulpliide, As 2 S 5 , is formed by passing H 2 S for a long time into a solution of alkali arsenate and then adding acid (McCay, Am., 1891, 12, 547); by saturating a solution of arsenic acid with H 2 S and placing, in stoppered bottle, in boiling water for one hour; or by passing a rapid stream of H 2 S into an HC1 solution of H 3 As0 4 (Bunsen, A., 1878, 192, 305 ; Brauner and Tomicek, J. C., 1888, 53, 146); 2H 3 As0 4 + 5H 2 S + xHCl = As,S. -f- 8H 2 -f- xHCl . Carbon disulphide extracts no sulphur from the precipitate, indicating the absence of free sulphur. The presence of FeCl 3 or heating the solution does not reduce the As 2 S 5 to As 2 S 3 . If there be a small amount of HC1 and the H 2 S be passed in slowly about 15 per cent of As 2 S 3 is formed: 2H 3 As0 4 + 5H 2 S + xHCl = As 2 S 3 + S 2 -f 8H 2 + xHCl . If NH 4 C1 be present more As 2 S 3 is formed. According to Thiele (C. C., 1890, 1, 877), arsenic acid cold treated with a slow stream of H 2 S gives arsenous sulphide, while the hot acid with a rapid stream of the gas gives the pentasulphide. Arsenic sulphide has the same solubili- ties as arsenous sulphide. When distilled with hydrochloric acid gas and a reducing agent, arsenous chloride is formed (AsCl 5 is not known to exist). The solutions in the alkali hydroxides, carbonates and sul- phides form arsenates and thioarsenates (5c). Ammonium sulphide added to a neutral or alkaline solution of arsenic acid forms arsenic sulphide which remains in solution as ammonium thioarsenate (5c). The addition of acid at once forms arsenic sulphide, not arsenous sulphide and sulphur. The reaction is much more rapid than with hydrosulplmric acid and is facilitated by warming. Arsine, AsH . , does not combine with hydrosulphuric acid until heated to 230, while stibine, SbH 3 , combines at the ordinary temperature (Brunn, B., 1889, 22, 3202). Acidulated solutions of arsenic boiled with thiosulphates form arsenous sulphide (separation f rom Sb and Sn) (Lesser, Z., 1888, 27, 218). Arsenic may be removed from sulphuric acid by boiling with barium thiosulphate and no foreign material is introduced into the acid: As 2 3 -f- 3BaS,0 : , = As 2 S 3 + 3BaS0 4 ; 2H 3 As0 4 + 5Na 2 S 2 3 = As 2 S 3 + 5Na 2 S0 4 + S 2 + 3H 2 0. (Thorn, J. C., 1876, 29, 517; Wagner, Dingl, 1875, 218, 321). Sulphurous acid readily reduces arsenic acid to arsenous acid : H 3 As0 4 + H 2 S0 3 = H 3 As0 3 + H 2 S0 4 (Woehler, A., 1839, 30, 224). 69, Qi. ARSENIC- 61 f. The arsenic from all arsenical compounds treated with concentrated hydrochloric acid and then distilled in a current of hydrochloric acid gas, passes into the distillate as arsenous chloride, AsCl 3 . Nearly all of the arsenic will be carried" over in the first 50 cc. of the distillate. This is a very accurate quantitative separation of arsenic from antimony and tin and from other non-volatile organic and inorganic material. The AsCl 3 passes over at 132, condenses with HC1 and may be tested with SnCl 2 (g), or, after decomposition with water (5c) by the usual tests for arsenous acid (Huf schmidt, B., 1884, 17, 2245; Beckurts, Arch. Pharm., 1884, 222, 684; Piloty and Stock, B., 1897, 30, 1649). Hydrobromic acid in dilute solutions is without action upon the acids of arsenic. The concentrated acid reduces arsenic acid to arsenous acid: H 3 As0 4 + 2HBr - ~ H,As0 3 -f Br, + H 2 . Hydriodic acid reduces arsenic acid to arsenous acid with liberation of iodine. This is a method of detecting As v in the presence of As'". 0.0001 gram of H 3 As0 4 may be detected in the presence of one gram of As 2 3 : 2H 3 As0 4 + 4HI = As 2 3 + 2I 2 + 5H 2 (Naylor, J. C., 1880, 38, 421). Chloric and bromic acids oxidize arsenous compounds to arsenic acid with formation of the corresponding hydracid: 3As 2 3 + 2HBr0 3 -f~ 9H 2 = 6H S AsO 4 + 2HBr . lodic acid oxidizes arsenous compounds to arsenic acid with liberation of iodine: 5As 2 3 + 4HI0 3 + 13H 2 = 10H 3 As0 4 + 2I 2 . g. Stannous chloride, SnCL , reduces all compounds of arsenic from their hot concentrated hydrochloric acid solutions, as flocculent, black-brown, metal- loidal arsenic, containing three or four per cent of tin. The arsenic, in solution with the concentrated hydrochloric acid, acts as arsenous chloride: 4AsCl 3 -f- GSnCl, As 4 + 6SnCl 4 .' The hydrochloric acid should be 25 to 33 per cent; if not over 15 to 20 per cent, the reaction is slow and imperfect. In a wide test-tube place 0.1 to 0.2 gram of the (oxidized) solid or solution to be tested, add about 1 gram of sodium chloride, and 2 or 3 cc. of sulphuric acid, then about 1 gram of crystallized stannous chloride', agitate, and heat to boiling several times, and set aside for a few minutes. Traces of arsenic give only a brown color; notable proportions give the flocculent precipitate. A dark gray precipitate may be due to mercury (58, 60), capable of being gath- ered into globules. If a precipitate or a darkening occurs, obtain conclusive evidence whether it contains arsenic or not, as follows: Dilute the mixture with ten to fifteen volumes of about 12 per cent hydrochloric acid; set aside, decant; gather the precipitate in a wet filter, wash it with a mixture of hydro- chloric acid and alcohol, then with alcohol, then with a little ether, and dry in a warm place. A portion of this dry precipitate is now dropped into a small hard-glass tube, drawn out and closed at one end, and heated in the flame; arsenic is identified by its mirror (7), easily distinguished from mercury (58, 7). Antimony is not reduced by stannous chloride; other reducible metals give no mirror in the reduction-tube. Small proportions of organic material impair the delicacy of this reaction, but do not prevent it. It is especially applicable to the hydrochloric acid distillate, obtained in separation of arsenic, according to f. h. Chromates boiled with arsenites and sodium bicarbonate give chromium arsenatc (Tarugi, J. (7., 1896, 70, ii, 340 and 390). i. Magnesium salts with ammonium chloride and ammonium hydroxide precipitate from solutions of arsenates, magnesium ammonium arsenate, MgNH,AsO 4 , white, easily soluble in acids. The reagents should be first mixed together, and used in a clear solution (" magnesia mixture ") to make sure that enough ammonium salt is present to prevent the precipitation of magnesium hydroxide by the ammonium hydroxide. The crystalline precipi- 62 ARSENIC. 69, 6;. tate forms slowly but completely. Compare with the corresponding magnesium ammonium phosphate (189, Qd). Maync-xium urxcnite is insoluble in water, but is soluble in ammonium hydroxide and in ammonium chloride (distinction from arsenates). j. Silver nitrate solution precipitates from neutral solutions of arsenites, or ammonio-silver nitrate * precipitates from a water solution of arsenous oxide, silver arsenite, Ag 3 AsO" 3 , yellow, readily soluble in dilute acids or in ammonium hydroxide (59, 6#). Neutral solutions of arsf. nates are precipitated as silver arsenate, Ag 3 As0 4 , reddish brown, having the same solubilities as the arsenite. k. Copper sulphate solution precipitates from neutral solutions of arsenites, or ammonio-copper sulphate (prepared in the same manner as the ammonio- silver oxide described above) precipitates from water solutions of arsenous oxide, the ffreen copper arsenitc, CuHAsO 3 (Scheele's green), soluble in ammo- nium hydroxide and in dilute acids. Copper acetate, in boiling solution, pre- cipitates the green copper aceto-arsenite (CuOAs 2 O 3 ) 3 Cu(C 2 H 3 O a ) 2 (Scliweinfurt green), soluble in ammonium hydroxide and in acids. Both these salts are often designated as Paris green (77, 6#). Copper sulphate with excess of free alkali is reduced to cuprous oxide with formation of alkali arsenate (10). K 3 As0 3 + 2CuS0 4 + 4KOH = K 3 As0 4 + 2K 2 S0 4 + Cu 2 O + 2H 2 O . Solutions of arsenates are precipitated by copper sulphate as copper arsenate, CuHAsO, , greenish blue, the solubilities and conditions of precipitation being the same as for the arsenites. 1. Ferric salts precipitate from arsenites, and freshly precipitated ferric hydroxide (used as an antidote, Wormley, 246), forms with arsenous oxide, variable basic ferric arsenites, scarcely soluble in acetic acid, soluble in hydro- chloric acid. Water slowly and sparingly dissolves from the precipitate the arsenous anhydride; but a large excess of the ferric hydroxide holds nearly all the arsenic insoluble. To some extent the basic ferric arsenites are trans- posed into basic ferrous arsenates, insoluble in water, in accordance with the reducing power of arsenous oxide. In the presence of alkali acetates, arsenic acid, or acidulated solutions of arsenates, are precipitated by ferric salts as ferric arsenate, FeAsO 4 , yellowish white, insoluble in acetic acid (compare 126, 6ef). m. Ammonium molybdate, (NH 4 ) 2 MoO 4 , in nitric acid solution, when slightly warmed with a solution of arsenic acid or of arsenates gives a yellow precipitate of ammonium arseno-molybdale, of variable composition. No precipitate is formed with As'". This precipate is very similar in appearance and properties to the ammonium phospho-molybdate; except the latter precipitates completely in the cold. 6'. Special Reactions, a. Marsh's Test. This is an extremely deli- cate test for arsenic, especially adapted for the detection of this ele- ment when present in small quantities, even when large quantities of other elements are present. Arsenic, from all of its soluble com- pounds, is reduced by the action of dilute sulphuric or hydrochloric acid on zinc, forming at first metallic arsenic and then arsenous hydride, AsH,, gaseous: As 2 3 + 6Zn + 6H 2 S0 4 = 2AsH 3 + 6ZnS0 4 + 3H 2 ; H 3 As0 4 + 4Zn + 4H 2 S0 4 = AsH 3 + 4ZnS0 4 + 4H 2 . The arsenic is precipitated with the other metals of the second group by hydrogen sulphide, separated with antimony, tin (gold, platinum and molybdenum) by yellow ammonium sulphide. This solution is precipitated by dilute hydrochloric acid and the mixed sulphides, well washed^ are dissolved in hydrochloric acid using as small an amount of potassium chlorate crystals as possible. The solution is boiled (till it does not bleach . litmus paper) * Prepared by adding ammonium hydroxide to a solution of silver nitrate till the precipitate at first produced is marly all redissoived. T If the ammonium salts are not thoroughly removed by washing there is danger of the for- mation of the very explosive chloride of nitrogen (268, 1) when the precipitate is treated with hydrochloric acid and potassium chlorate. CALIFORWW COUEfil u**ta. * P "ARMACY 63 to remove excess of chlorine and is then ready for the Marsh apparatus. This apparatus consists of a strong Erlenmeyer flask of about 125 cc. capacity fitted with a two hole rubber stopper. Through one hole is passed a thistle (safety) tube, reaching nearly to the bottom of the flask; in the other is fitted a three-inch Marchand calcium chloride tube, which projects just through the stopper and is filled with glass-wool and granular calcium chloride to dry the gases generated in the flask. To the other end of the Marchand tube is fitted, with a small cork or rubber stopper, a piece of hard glass tubing of six mm. diameter and one foot long. This tube should be constricted one-half, for about two inches, beginning at the middle of the tube and extending toward the end not fastened to the calcium chloride tube. The outer end of the tube should also be con- stricted to about one mm. inner diameter. A short piece of rubber tubing should connect this constricted end with a piece of ordinary glass tubing, dipping into a test tube about two-thirds filled with a two per cent solu- tion of silver nitrate. The rubber tubing should make a close joint with the constricted end of the hard glass tube, and yet not fit so snug but that it can be easily removed. From 10 to 20 grams of granulated zinc * are placed in the flask with sufficient water to cover the end of the thistle tube. Four or five cubic centimeters of reagent sodium carbonate are added and the stopper tightly fitted to the flask. Dilute sulphuric acid (one of acid to three of water) should now be added, very carefully at first, f until a moderate evolution of hydrogen is obtained,. The hydrogen should be allowed to bubble through the silver nitrate for about five minutes. There should be no appreciable blackening of the solution (59, 10), thus proving the absence of arsenic from the zinc and the sulphuric acid. The purity of the reagents having been estab- lished the solution containing the arsenic may be added in small amounts at a time through the thistle tube. If arsenic be present there will be almost immediate blackening of the silver nitrate solution. 6AgN0 3 + AsH 3 + 3H,0 = 6Ag + H 3 As0 3 + GHN0 3 The hard glass tube should now be heated J to redness by a flame from * The zinc and all the reagents should be absolutely free from arsenic. If the zinc be strictly chemically pure it will be but slowly attacked by the acid. It should be platinized (349, 4a)or should contain traces of iron. Hote ' A. Ch., 1884, (6>, 3, 141) removes arsenic from zinc by adding anhydrous MgCl 4 to the molten metal, AC1 3 being evolved. The zinc purified in this way is readily attacked by acids. t The acid first added decomposes the alkali carbonate forming carbon dioxide which rapidly displaces the air and greatly lessens the danger of explosion when the gas is ignited. If too much acid be added before the carbonate is decomposed violent frothing may take place and the liquid contents of the flask forced into the calcium chloride tube. \ Before heating the tube or igniting the gas, a towel should be wrapped around the flask to insure safety in case of an explosion due to the imperfect removal of the air ; or the tube con- necting the hard glass tube with the Marchand tube should be of larger size and provided with u plug of wire gauze (made of 10 or 20 circles of gauze the size of the tube). A flame cannot pass such a plug of wire gauze. 64 AlttiKXic. 69, 6'fe. 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. As small a quantity of arsenic as 0.002 mg. will produce a visible mirror and if 20 g. of material is used for analysis, this would represent one part of arsenic in 10,000,000 parts. 1). Arsenous Hydride (arsine), AsH 3 , 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 + 30 2 = As 4 + GH.,0 . 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, or by cautiously neutralizing with ammonia the arsenic may be precipitated as the yellow silvery arsenite (6/). To generate arsine, magnesium or iron f 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 compounds by nascent hydrogen generated in alkaline solution. Sodium amalgam, t zinc (or zinc and magnesium) and potassium hydroxide or alumi- num and potassium hydroxide may be used as the reducing agent. There is * 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. t According to Thiele (C. C., 1890, 1, 877) arsenic may be separated from antimony in the Marsh test by using electrolytirally deposited iron instead of zinc. Stibine is not evolved. According to Sautermeister (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 tiroo) 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. 6*5 no reaction with As v , or with compounds of antimony (70, 6;'); hence when the arsenic is present in the triad condition (Asv may be reduced to As"' by SO.,) the use of one of the above reagents serves aclinirably i'or 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, /. C., 1885, 48, 838; Johnson, C. 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 .v/M/w, SbH 3 , combines at ordinary temperature (method of separation) (Hrunn, B., 1889, 22, 3202; Myers, /. 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 s + 3l a = Asl, -f SHI (Jacobson, B., 1887, 20, 1999). Arsenous hydride is decomposed by passing through a tube heated to redness (mirror in Marsh test) 4AsH g = As 4 + 6H 2 . Nitric acid oxidizes it to arsenic acid, 3AsH, -f 8HNO 3 = 3H 8 AsO 4 -f 8NO + 4H,O; 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 8 AsO 4 , reddish brown precipitate, &j). Dissolve the residue in hydrochloric or nitrohydrochloric acid and test for antimony with hydrogen sulphide (Ansell, J. 0., 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 such as distinguish arsenic from antimony. ARSENIC MIRROR. ANTIMONY MIRROR. Deposited beyond the flame; ar- Deposited before or on both sides sene not being decomposed much be- of the flame; stibine being decom- low a red heat. poied 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 630, 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. GG ARSENIC. 69, 6> 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 (6e). 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 Of a velvety brown to black sur- face. Volatile, by oxidation to monous oxide, at a red heat. anti- Do not dissolve in hypochlorite. Warmed with ammonium sul- phide, form orange-yellow spots, in- soluble in ammonium carbonate, soluble in hydrochloric acid (70. Qe). 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 4 -f lOXaCIO + 6H 2 O = 4H 3 AO 4 + lONaCl. 69, 6'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- nil c (C)/): or with ammonio-copper sulphate forming the green copper ai-smite (6fc) (Wormley, 259). Any other test for arsenous oxide may be applied as desired. 7. Magnesia mixture (Gi) is added to the solution of the mirror or spots in nitric acid. The solution must be strongly alkaline. A white crystalline precipitate of magnesium ammonium arsenate, MgNH 4 As0 4 , is formed (Wormley, 316). d. Remsch'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 (C. 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 no* present. This serves as one of the most satisfactory methods of determining the presence or absence of arsenic jn 68 AR8ENIC. 69, 6'g. 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 surface; 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 ma} r 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 w r hich also yields a sublimate, but this is amorphous and does not show the octahedral crystals w r hen examined under the microscope (Wormley, 269 and ff.; Clark, J. (7., 1893, 63, 886). . Detection in Case of Poisoning 1 . 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 hundred 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 trailed 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 Fresenius and Babo. The tissue is cut into 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. Th^ heating is continued until there is no odor of chlorine, but concentration should be avoided by the addition of water. The sclu.ion should be cooled and filtered, 69, 7. ARSENIC. 69 the arsenic now being 1 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. Kedissolve 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 Reinseh 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 water, add nitric acid, and again evaporate to dryness; and repeat until the mass is colorle'ss. 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 IT ay be conducted into the retort during the process of dis- tillation, in which case all the arsenic (even from arsenous sulphide (5c)) will be carried over in the first 100 cc. of the distillate. The receiver contains the arsenic, a great excess of hj'drochloric acid and a small amount of organic matter. To a portion of this solution the Reinsch 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: Medical Jurisprudence. Forensic Medicine and Toxicology, Witthaus and Becker, Vol. iv, 1911; Laboratory Manual for the Detection of Poisons and Powerful Drugs, Dr. Wilhelm Autenreith (translated by Wm. H. Warren), 1915; Allen's Commer- cial Organic Analysis, Vol. vi (4th edition, 1912); The Qualitative Analysis of Medicinal Preparations, H. C Fuller, 1912; Elementary Chemical Micros- copy, Emile M. Chamot, 1915; Biochemisches Handlexikon, V. Band, Dr. Emil Abderhalden, 1911; Mikrochemische Analyse, 1 and 11 Teil, P. D. C. Kley, 1915; Medical Jurisprudence, Taylor; Ermittelung von Giften, Dragendorff. 7. Ignition. Metallic arsenic is obtained by igniting any compound containing arsenic with potassium carbonate and charcoal,* or with potas- sium cyanide : 2As 2 0, + 6KCN = As, + 6KCNO 2As 2 S 3 + 6KCN = As 4 + 6KCNS 2As 2 S 8 + 6Na 2 C0 3 + 6KCN = As 4 + 6Na 3 S + 6KCNO + 6CO, . 4H,AsO< + 50 = As 4 + 5C0 2 + 6H 2 O * A very suitable carbon for the reduction of arsenic is obtained by igniting an alkali tartrate In absence of air to complete carbonization. 70 A&8ENM. 69, 8. If this ignition be performed in a small reduction-tube * (a hard glass tube 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). B} r 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 3 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 (6 and 88, 4). 9. Estimation. (1). As lead arsenate, Pb,(As0 4 ) 2 . To a weighed portion of the solution containing arsenic acid, a weighed amount of PbO is added. After evaporation and ignition at a dull red heat the residue is weighed as Pb 3 (As0 4 ) 2 from which the weight of the added PbO is subtracted. 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 OH 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). (-5). The MgNH 4 As0 4 is converted by ignition into Mg 2 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., 2O, 531 and 22, 397) recommends the following modification of the above method : A piece of reduction tubing about 16 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, gentlj r 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 tin? way do not give a mirror. As small an amount as 0.00001 gram of As 3 O 3 will give a distinct mir- ror by this method. 69, 10. ARSENIC. 71 The precipitate is separated from free sulphur by solution in NHjOH and ivprecipitated with HC1 . It is then dried at 100 and weighed as As 2 S 3 . (-5). By precipitation as in (4) and removal of sulphur by washing with CS 2 . Dry at 100 and weigh as As 2 S 3 . (6). Uranyl acetate, in presence of ammonium salts, precipitates N"H 4 U0 2 As0 4 ; by ignition this is converted into uranyl pyroarsenate (U0 2 ) 2 As.,0 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,Cr,0 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. (13). 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 . (14). 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, and this passed into a, solution of standard silver nitrate, the excess of which is estimated with standard NaCl or the excess of AgN0 3 is removed and the arsenous acid titrated as in methods (9) or (11). (16). Small amounts are determined by conversion to AsH 3 and the stain produced on mercury bromide paper compared with the stain produced by known amounts of arsenic. (Seeker and Smith, U. 8. Bull. No. 147, p. 212-214). Many other methods have been recommended. 10. Oxidation. As~"'H 3 is oxidized to As'" by AgN0 3 , H 2 SO, , H 2 S0 4 , and HIO S ; and to As v by KMn0 4 (Tivoli, Qazzetta, 1889, 19, 630), HN0 2 , HN0 3 , Cl and Br (Parsons, C. N., 1877, 35, 235). As is oxidized to As"' by H 2 6 2 (Clark, J. C., 1893, 63, 886), HN0 3 , H 2 S0 4 hot, Cl , HC10 , HC10 3 , Br, HBrO ;? , 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) , etc. (Mayer, J. pr., 1880 (2), 22, 103). Arsine is oxidized to metallic arsenic by HgCl 2 (Magencon and Bergeret, J. (7., 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 , KCT, etc. (7). By SnCl 2 (6*7) and H 3 P0 2 (Gd) in strong HC1 solution; also with greater or less completeness b such as Cu , Cd , Zn , Ms , etc. Rideal (C. N.. IfiftUf. 72 ANTIMONY. 70, 1. HC1 solution; also with greater or less completeness by some free metals, such as Cu , Cd , Zn , Mg , etc. Rideal, (C. N., 1885, 51, 292) recommends the use of the copper-iron wire couple for the detection of small quantities of arsenic by reduction to the elemental state. 0.00000?5 grams may be detected. In solution As v is reduced to As'" by H 3 PO, , H S , H S0 3 , Na 2 S 2 3 (6e), HC1 , HBr , HI (G/), 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 = 120.2. Valence three and five (11). 1. Properties. Specific Gravity, 6.62 (Z. anorgan. Chem., 1902, 177). Melting point, 630 (Cir. B. S., 35, 1915). Boiling point, between 1500 and 1700 (B., 1889, 725). Its molecular weight is unknown, as its vapor density 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 incandescence, forming white inodorous (dis- tinction from arsenic) vapors of antimonous 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 3 ; in very many minerals usually combined with other metals as a double sulphide (Campbell, Pliil. 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 2 S 3 + 6Na 2 CO 3 +30 = 4Sb + 6Na 2 S + 9CO, . (c) It is reduced by metallic iron: Sb 2 S 3 + 3Fe = 2Sb + 3FeS . (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 3 , Sb 2 4 , and Sb,O 5 . (a) Antimonous oxide, Sb 2 O s , is formed (1) by the action of dilute nitric acid upon Sb; (2) by precipitating SbCl s with Na 2 C0 3 or NH 4 OH; (3) by dissolving Sb in concentrated H 2 S0 4 and precipitating with Na 2 C0 3 ; (4) by burning antimony at a red heat in air or oxygen; (5) by heating Sb 2 O 4 or Sb 2 5 to 800 (Baubigny, C. 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 (5&). Anti- monous oxide sometimes acts as an acid, Sb 2 O 3 + 2Na.OH = 2NaSb0 2 + H,O; 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 O 4 , is formed by heating Sb , Sb 2 S 3 , Sb,0, , or Sb 2 5 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 3 O 4 = 2Sb"'SbvO 4 = Sb 2 O 3 .Sb 2 O 5 . It is found native as antimony ochre, (c) Antimonic oxide, Sb,O s , is formed by treating Sb , Sb,O 3 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, *Beilstein and Blaese (C. C., 1889, 803' have prepared a number of antimonates and conclude that the acid is always the meta, H SbO, . 70, oft. ANTIMONT. 73 j. c., ortho, meta and pyro (Geuther, I. c., and Conrad, C. N., 1879, 40, 198). The ortho acid, H 3 SbO 4 is formed by the decomposition of the pentachloride 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,O 3 with potassium nitrate and extraction with water gives the potassium metantimonate, KSbO s , 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 SbO 4 = H 4 Sb 2 O 7 -f H,O (Conrad, L c.), 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 7 . The sodium salt Na 2 H 2 Sb 2 O 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, 6*7). For the latter the soluble potassium salt K 2 :Er,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 O 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 O 7 + 2NaCl = Na 2 H 2 Sb 2 7 + 2KC1 (11). The ortho acid, H s Sb0 4 , is sparingly soluble in water, easily soluble in KOH, but insoluble in NaOH. The meta acid, HSbO 8 , 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, R 4 SbnO 7 , are soluble in water, also the acid potassium salt, K 2 H 2 Sb 2 7 , but not the corresponding sodium salt, Na 2 H._,Sb 2 O 7 . 5. Solubilities. a. Metal. Antimony is attacked but not dissolved by nitric acid, forming Sb 2 O 3 (a) or Sb 2 O 5 (&), 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,O 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 forming at first the corresponding trihalogen compounds (d). Chlorine and bromine (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., 1891, (6), 18, 65). The pentiodide, SbI 5 , does not appear to exist (Mac Ivor, J. C., 1876, 29, 328). (a) 2Sb + 2HN0 3 = Sb 2 3 + 2ND + H 2 O (6) 6Sb + 10HNO 3 = 3Sb 2 O 5 -f 10NO + 5H Z (c) 2Sb + f,H 2 SO 4 = Sb 2 (SO 4 ) 3 -f 3S0 2 + 6H 2 O (d) 2Sb + 3CL = 2SbCl s (e) SbCl 3 + C1 2 = SbCl 5 ft. Oaridett. Antimonous oxide, Sb 2 O 3 , is soluble in 55,000 parts of water at 15 and in 10,000 parts at 100 (Schulze, J. Pr., 1883, (2), 27, 320); insoluble in alcohol; soluble in hydrochloric (a), sulphuric and tartaric (&) acids with formation of the corresponding salts. The dry ignited oxide is scarcely at all soluble in nitric acid; the moist, freshly precipitated oxide, on the other hand, dissolves readily in the dilute or concentrated acid, be it hot or cold. Under certain conditions of concentration a portion of the antimony precipitates out upon standing as a white crystalline precipitate. It is soluble in the fixed 74 A\TLMOXY. 70, 5c. alkali hydroxides with formation of metantimonites (c) (Terreil, A. Cli., 1866, (4), 7, 350). Fixed alkali carbonates dissolve a small amount of the oxide with the probable formation of some antimonite (rf) (Schneider, Poyy., 1859, 108, 407). It is fairly soluble in glycerine (Kohler. DingL, 18S5. 258. 520). (a) Sb 2 O 3 + 6HC1 = 2SbCl 3 + 3H 2 O (6) Sb 2 3 + H 2 C 4 H 4 G = (SbO) 2 C 4 H 4 .+ H 2 (c) Sb 2 3 + 2KOH = 2KSb0 2 + H 2 (d) Sb 2 3 + NasCOs = 2NaSb0 2 + CO 2 Antimony tetroxide, Sb 2 4 , is insoluble in water, slowly dissolved by hot concentrated hydrochloric acid. Antimonic oxide, Sb 2 O, , is insoluble in water; soluble in hydrochloric and tartaric acids without reduction; hydriodic acid dissolves it as antimonous iodide with liberation of iodine (Of) ; slowly soluble in concentrated fixed alkalis; soluble in alkaline solution of glycerine (Kohler, J. C., 1886, 50, 428). The hydrated oxides of antimony (acids) "have essentiallv the same solubilities as the oxides (4). c. Salts. Antimonous chloride, SbCL, , is very (lelniucwent. decomposed by pure water, forming a basic salt; soluble in water strongly acidulated with an inorganic acid, or tartaric, citric, or oxalic acids (6&), but not when acidulated with acetic acid; it is also soluble in concentrated solutions of the chlorides of the alkalis and of the alkaline earths (Atkinson, (''. .V., 1883, 47. 175). The bromide and iodide are dcl'uiucsccnt and require moderately concentrated acid lo keep them in solution. The sulphate, Sb, (S0 4 ) 3 , dissolves in moderately con- centrated sulphuric acid. Antimonous tartrate and the potassium antimonous tartrate (tartar-emetic) are soluble in water without acidulation; the latter is soluble in glycerine and insoluble in alcohol. The trichloride, bromide and iodide are soluble in hot CS 2 ; the chloride and bromide are soluble in alcohol without decomposition, but the iodide is partially decomposed by alcohol or ether (Mac Ivor, J. (7., 1876, 99, 328). The pentachloride, SbCl 5 , is a liquid, very readily combining with a small amount of w^ater to form crystals containing one or four molecules of water. The addition of more water decomposes the salt forming the basic salt; if, however, a few drops of HC1 have been added first, any desired amount of water (if added at one time) may be added without causing a precipitation of the basic salt. If after acidulation water be added slowly, the basic salt will soon be precipitated. Antimonous sulphide, Sb 2 S 3 , is readily soluble in K 2 S , and on evapora- tion large yellow transparent crystals of K 4 Sb 2 S. are^obtained (a) (Ditte, C. r., 1886, 102, 1G8 and 212). It is soluble in moderately concentrated HC1 with evolution of ELS (5); slowly decomposed by boiling with water into Sb 2 3 and H 2 S (c): and on boiling with NH 4 C1 into SbCl, and (NH 4 ),S (de Clermont, C. r., 1879, 88, 972). Dilute H 2 S0 4 is almost without action, dilute HNO., gives Sb 2 3 (d). Sparingly soluble in hot NH 4 OH solution, soluble in the fixed alkalis (on fusion or boiling) (e); insoluble in (NH 4 ) 2 C0$ (distinction from arsenic); insoluble in the fixed alkali carbonates in the cold but on warming they effect complete solution (f) (distinction from tin); very sparingly soluble in normal ammonium sulphide; readily soluble in yellow ammonium sulphide with oxidation (g) (6e). The pentasulphide, Sb^S. , is insoluble in water; soluble in the alkali sulphides (Ji), and in the fixed alkali carbonates and hydroxides; insoluble in ammonium carbonate and sparingly soluble in ammonium hydroxide, more readily when warmed (D., 2, 1, 217). On boiling with water it slowly decomposes into Sb,0,., ?0, 5d. ANTIMONV. 75 H 2 S and S (Mitscherlich, J. pr., 1840, 19, 455). Hydrochloric acid on warming dissolves it a? SbCl, (/): (a) Sb 2 S 3 + 2K 2 S = K,Sb,S 5 (1)) Sb 2 S 3 + C.EC1 = 2SbCL + 3H 2 S (c) Sb 2 S 3 + 3H,0 = Sb 2 3 + 3H 2 S (d) 2Sb 2 S 3 + 4HNO 3 = 2Sb,0 3 + 3S 2 + 4NO + 2H 2 (e) 2Sb,S 3 + 4KOH = 3KSbS 2 + KSbO 2 + 2H 2 (/) 2Sb,S 3 -f 2Nsi,CO, = :;NaSbS 2 + NaSbO, + 2CO 2 (y) 2Sb,S 3 + 6(NH 4 ) 2 S 2 = 4(NH 4 ) 3 SbS 4 + S 2 (70 Sb a S 5 + 3(NH 4 ) 2 S = 2(NH 4 ) 3 SbS 4 (i) Sb 2 S 5 + 6HC1 = 2SbCl 3 + 3H 2 S + S 2 d. Water.* With the exception of the compounds of antimony with some organic acids, as tartaric and citric, all salts of antimony are decom- posed by pure WATER. For this reason it will be seen that water is a very important reagent in the analysis of antimony salts. The salts with inorganic acids all require the presence of some free acid (not acetic) to keep them in solution. If the acid be tartaric the further addition of water causes no precipitation of the antimony salt. Water decomposes the inorganic acid solutions precipitating the basic salt, setting more acid free which dissolves a portion of the basic salt. The addition of more water causes a further precipitation and at the same time dilutes the acid so that upon the addition of a sufficient amount of water a nearly com- plete precipitation may be obtained. If the precipitate of the basic salt be washed with water the acid is gradually displaced, leaving finally the anti- mony as oxide. With solutions of antimonous chloride the basic salt precipitated is white antimonous oxychloride, Sb 4 Cl 2 5 , " Powder of Algaroth," soluble in tartaric acid (distinction from bismuth, 76, 5d) (Mac Ivor, C. N., 1875, 32, 229), 4SbCl 3 + 5H,0 = Sb 4 Cl,0 5 + 10HC1 . The basic salt repeatedly washed with water is slowly (rapidly if alkali carbonate be used) changed to the oxide, Sb 2 3 (Malaguti, J. pr., 1835, 6, 253), Sb 4 Cl 2 5 + H 2 = 2Sb,0 3 + 2HC1. With antimonic chloride, SbCl 5 , the basic salt is SbOCl, ; SbCl, + H 2 = SbOCl 3 + 2HC1 (Williams, (7. N., 1871, 24, 224). Solutions of the tartrates of antimony and of antimony and potassium are not precipitated on the addition of water; and antimonous chloride *The acidity of water solutions of certain salts having- a weak base and the alkalinity of others containing a weak acid is due to a partial decomposition (hydrolysis) of the salt by the ions of the water, H* and OH', forming- again the original acid and base. ]Va 2 CO 3 , for instance, is split up into the weak non-dissociated H 2 CO 3 and the strongly-dissociated NaOH, whose OH ions give the "alkaline reaction." FeCl 3 in water forms soluble colloidal Fe(OH)a, which may be separated by dialysis from the free HC1 resulting or precipitated by addition of a neutral salt, as NaCl, to the dilute solution; KCN gives alkaline KOII and non-dissociated HCW, readily detected by its odor. In other cases precipitation is caused, as in the treatment of bismuth or antimony solutions with water or on heating "VaoZnOa solution, hydrolysis in general being increased by raising the temperature. The action of water on soap belongs to this class. 76 ANTIMONY. 70, 60. dissolved in excess of tartaric or citric acid solution is not precipitated on addition of water. G. Reactions. a. The alkali hydroxides and carbonates precipitate from acidulated solutions of inorganic antimonous salts, antimonous oxide * Sb (a) (Rose, Pogg., 1825, 3, 441), white, bulky, readily becoming- crystalline on boiling; sparingly soluble in water (56), readily soluble in excess of the fixed alkalis, forming a metantimonite (6) (Terreil, A. Ch., 1866, (4), 7, 350); slowly soluble in a strong excess of a hot solution of the fixed alkali carbonate (c) (distinction from tin); insoluble in ammonium hydroxide or ammonium car- bonate. The freshly precipitated oxide is readily soluble in acids (not in acetic acid). If the alkaline solution of the antimony be carefully neutralized with an acid (not tartaric or citric) the oxide is precipitated (d) and at once dissolved by further addition of acid. The presence of tartaric or citric acids prevents the precipitation of the oxide by means of the alkalis or alkali carbonates. Antimonous oxide acts as a feebly acidic anyhdride toward alkalis, with which it combines, dissolving in their solutions and forming antimonites, which are found to be monobasic, so far as capable of isolation. Sodium ar.timonite, NaSbO, , is the most stable and the least soluble in water; potassium anti- monite, KSbO 2 , is freely soluble in dilute potassium hydroxide solution, but decomposed by pure water. By long standing (24 hours), a portion of the antimonous oxide deposits from the alkaline solution, and the presence of alkali hydrogen carbonates causes a nearly complete separation of that oxide (e). (a) 2SbCl 3 + 6KOH = Sb 2 3 + 6KC1 + 3H 2 2SbCl s + 3Na 2 C0 3 = Sb 2 3 + 6NaCl + 3CO 2 (6) Sb 2 0, + 2KOH = 2KSb0 2 + H 2 or SbCl 3 + 4KOH = KSb0 2 + 3KC1 + 2H 2 O (c) Sb 2 O 3 + Na,CO 3 = 2NaSb0 2 + C0 2 (d) 2KSb0 2 + 2HC1 = Sb 2 3 + 2KC1 + H 2 O (e) 2NaSb0 2 + 2NaHC0 3 = Sb 2 8 + 2Na 2 C0 3 + H 2 Antimonic salts are precipitated under the same conditions as the antimonous salts. The freshly formed precipitate is the orthoantimonic acid, H s Sb0 4 = SbO(OH) 8 = Sb 2 O 5 ,3H 2 O (a) (Conrad, C. N., 1879, 40, 198); insoluble in am- monium hydroxide or carbonate; soluble, more readily upon warming, in excess of the fixed alkali hydroxides and carbonates as metantimonate (6). (a) SbCl 5 + 5KOH = Sbo'(OH), + 5KC1 + H 2 O (6) SbO(OH) 8 + KOH = KSb0 3 + 2H 2 O 6. The freshly precipitated antimonous oxide is soluble in oxalic acid, but (in absence of tartaric acid) the antimony soon slowly but completely separates out as a white crystalline precipitate; unless an alkali oxalate be present, when the soluble double oxalate is formed. The precipitate of antimony oxalate dissolves upon the further addition of hj^drochloric acid. Freshly precipitated antimonic oxide dissolves readily in oxalic acid and does not separate out upon standing. Acetic acid precipitates the solutions of antimony salts if tartaric acid be absent. Potassium cyanide gives a white precipitate with antimonous salts soluble in excess of the cyanides. With potassium ferrocyanide antimonous chloride (not tartrate) gives a white precipitate, soluble in hydrochloric acid (distinction from tin), or fixed alkali hydroxides (Warren, C. N., 1888, 57, 124). Potassium ferricyanide is reduced to ferrocyanide by antimonous salts in alkaline solution (Baumann, Z. angew., 1892, 117). c. From the solutions of the fixed alkali antimonites or antimonatee the oxides or hydrated oxides (acids) are precipitated upon neutralization with nitric acid (or other inorganic acids) ; the freshly formed precipitates readily * Menschutkin (page 186) says the precipitate formed by the action of alkalis upon antimonous salts is the meta acid, HSbO a . 70, 6. ANTIMONY. 77 dissolving- in an excess of the acid. Antimonous nitrate is very unstable and the antimonic nitrate is not known to exist. It is quite probable that these solutions in nitric acid are merely solutions of some of the hydrated oxides (acids). d. Compounds of antimony with the acids of phosphorus are not known, (Na 2 HP0 4 does not precipitate antimony salts, separation from tin, 71, Qd). c. Hydrogen sulphide precipitates, from acid * solutions of antimonous salts, antimonous sulphide (a), Sb 2 S 3 , orange-red; in neutral solutions (tartrates) the precipitation is incomplete. In strong fixed alkali solu- tions (6a) the precipitation is prevented, or rather the sulphide first formed (&) is at once dissolved in the excess of the fixed alkali (c), sparingly in NH 4 OH . The alkali sulphides give the same precipitate sparingly soluble in normal ammonium sulphide, readily soluble in the fixed alkali sulphides (d) and in yellow ammonium sulphide (e). Antimonous sulphide is slowly decomposed by boiling water (f) ; insoluble in ammonium carbon- ate (distinction from As); slowly soluble in boiling solution of the fixed alkali carbonates (g) (distinction from Sn) ; soluble in hot moderately con- centrated hydrochloric acid (h) (distinction from arsenic). The alkaline solutions of antimonous sulphide are oxidized upon standing by the oxygen of the air or rapidly in the presence of sulphur (e) ; from the alkaline solu- tions hydrochloric acid precipitates the antimony as trisulphide, penta sulphide or a mixture of these, depending upon the degree of .oxidation (i). (a) 2SbCl s + 3H 2 S = Sb 2 S 3 + 6HC1 (6) 2KSb0 2 + 3H 2 S = Sb 2 S 3 + 2KOH + 2H 2 O (c) 2Sb 2 S 3 + 4KOH = 3KSbS 2 + KSbO 2 + 2H 2 O (d) Sb 2 S 3 + K 2 S = 2KSbS 2 (e) 2Sb 2 S 3 + 6(NH 4 ) 2 S 2 = 4(NH 4 ) 3 SbS 4 + S a (f) Sb 2 S 3 -f 3H 2 = Sb 2 s + 3H 2 S (g) 2Sb 2 S 3 + 2K 2 C0 3 = 3KSbS 2 -f- KSbO 2 + 2C0 2 (ft) Sb 2 S 3 + 6HC1 = 2SbCl s + 3H 2 S (i) 3KSbS 2 -f KSb0 2 + 4HC1 = 2Sb 2 S 3 + 4KC1 + 2H 2 or 2(NH 4 ) 3 SbS 4 + 6HC1 = Sb 2 S 6 + 6NH 4 C1 + 3H 2 S Hydrosulphuric acid f and alkali sulphides precipitate (under like condi- tions as for antimonous salts), from solutions of antimonic salts, antimonic sulphide, Sb 2 S 5 , orange, having the same solubilities as the tri-sulphide. The alkaline solution of the sulphide consists chiefly of the ortho-thioanti- monate instead of^the meta, as in antiinpnous compounds. Sb 2 S 5 -f 3K 2 S = 2K 3 SbS 4 ; 4Sb 2 S 5 + 18KOH = 5K 3 SbS 4 + 3KSb0 3 + 9H 2 . When dissolved in HC1 the penta-sulphide is reduced to SbCl 3 with liberation of sulphur, Sb 2 S 5 + 6HC1 = 2SbCl 3 + 3H 2 S + S 2 . * According to Loviton (J. C., 1888, 54, 993) the precipitation takes place in the presence of quite strong hydrochloric acid tone to one) separation from tin, which is precipitated only when three or more parts of water are present to one of the acid. See also Noyes and Bray, J. Am. Soc. t 29, 137 (1917). t In order to precipitate pure antimonic sulphide, the solution of the antimonic salt must be cold, and the hydrogen sulphide added rapidly. If the solution be warmed or the hydrogen sul- phide added slowly more or less antimonous sulphide is precipitated (B6sek,.7. C., 1895, 67,515). 78 ANTIMONY. * 70, 6/. All salts of antimony when warmed with sodium, thiosulphate, Na 2 S 2 O 3 , are precipitated as the sulphide (separation of arsenic and antimony). 2SbCl 3 + 3Na 2 S 2 3 + 3H 2 O == Sb 2 S 3 + 3Na,,SO 4 + GHC1 . Sulphurous acid reduces antimonic salts to antimonous salts (Knorre, Z. angeic., 1888, 155). Sulphates of antimony are not prepared by precipitation, but by boiling the oxides with strong sulphuric acid. They dissolve only in very strongly acidulated water. /. Antimony occurs most frequently for analysis as the chlorides; it is therefore important that the student familiarize himself with the deport- ment of these salts with the various reagents, used in qualitative analysis. The most important of the properties have been discussed under 5a, &. c, d. Hydrochloric acid, or any other inorganic acid, carefully added to a solu- tion of antimony salts in the fixed alkalis will precipitate the correspond- ing oxide or hydrated oxide, soluble upon further addition of the acid. Potassium iodide added to antimonous chloride solution, not too strongly acid, gives a yellow precipitate of antimonous iodide, soluble in hydro- chloric acid. The precipitation does not take place in the presence of tartaric or oxalic acids. Hydriodic acid (or potassium iodide in acidu- lated solutions) added to solutions of antimonic salts causes a reduction of the antimony to an antimonous salt with liberation of iodine (distinc- tion from Sn IV : SbCl 5 + 2HI = SbCl 3 + 2HC1 + I 2 . The iodine may be detected by heating and obtaining the violet vapors, or by adding carbon disulphide and shaking. It should be remembered that the solution to be tested must be acid, for in alkaline solutions the reverse action takes place, iodine oxidizing antimonous salts to antimonic salts: SbCl 3 -)- 8KOH + I 2 K 3 Sb0 4 + 2KI + 3KC1 + 4H 2 (Weller, A., 1882, 213, 364). Also the absence of other oxidizing agents which liberate iodine from hydriodic acid must be assured. g. If antimony and arsenic compounds occurring together are strongly oxidized with nitric acid there is danger that the insoluble precipitate of anti- monic oxide may contain arsenic, as antimonic arsenate, insoluble (Menschut- kin). Stannous chloride reduces antimonic compounds to the antimonous condition, but in no case causes a precipitation of the metal (distinction from arsenic). Ji. Antimonous salts in acid, neutral or alkaline solution, rapidly reduce solutions of chromates to chromic compounds. Acid solutions of antimonous salts reduce solutions of manganates and permanganates to manganous salts; with alkaline solutions to manganese dioxide. These reactions are capable of quantitative application in absence of other reducing agents. The antimony is oxidized to the antimonic condition (9 and 10). i. An antimonous compound when evaporated on a water bath with an ammoniacal solution of silver nitrate gives a black precipitate (Bun sen, A., 1855, 106, 1). A solution of an antimonous compound in fixed alkali when treated with a solution of silver nitrate gives a heavy black precipitate of metallic silver, insoluble in ammonium hydroxide, and thus separated from the precipitated silver oxide. If instead of 'a water solution of silver nitrate, a solution with great excess of ammonium hydroxide (one to sixteen) be added, no precipitation occurs in the cold (distinction from Sn") ; nor upon heating until the excess of ammonia has been driven off. Antimonates with silver nitrate give a white precipitate of silver antimonate, soluble in ammonium hydroxide. 70, 6/. ANTIMONY. 79 j. Stibine. By the action of zinc and sulphuric or hydrochloric acid all compounds of antimony are first reduced to the metallic state. The formation of stibine is a secondary reaction and requires the moderately rapid generation of hydrogen in acid solution. If a few drops of a solu- tion of an antimony salt, acidulated with hydrochloric acid, be placed upon a platinum foil and a small piece of zinc be added, the antimony is immediately deposited as a black stain or coating adhering firmly to the platinum; 2SbCl 3 + 3Zn -- 2Sb -f 3ZnCl, . In this test tin, if present, deposits as a loose spongy mass, while arsenic, if present, does not adhere so firmly to the platinum as the antimony. In the presence of arsenic this test should be applied with caution under a hood as a portion of the arsenic is almost immediately evolved as arsine (69, 6'&). If hydrogen be generated more abundantly than in the operation above mentioned, by zinc and dilute sulphuric or hydrochloric acid, the gaseous antimony hydride, stibine, SbH 3 , is obtained for examination. For com- parison with arsine and details of manipulation see " Marsh's Test " under arsenic (69, 6' a) : Sb,O 3 + GZn + 6H 2 S0 4 = GZnSO, + 3H 2 + 2SbH, SbCl 3 + 3Zn + 3HC1 = SZnCL + SbH 3 Stibine is a colorless, odorless gas, not nearly so jpoisonous as arsine. It burns with a luminous and faintly bluish-green flame, dissipating vapors of antimonous oxide and of water (a); or depositing antimony on cold porcelain held in the flame, as a lusterless brownish-black spot (&). The gas is also decomposed by passing through a small glass tube heated to low redness (c), forming a lustrous ring or mirror in the tube. The stibine is decomposed more readily by heat than the arsine and the mirror is deposited on both sides of the heated portion of the glass tube. The spots and mirror of antimony are compared with those of arsenic in 69, 6'c. The antimony in stibine is deposited as the metal when the gas is passed into a concentrated solution of fixed alkali hydroxide or when it is passed through a IT tube filled with solid caustic potash or soda-lime (distinction and separation from arsenic). (a) 2SbH 3 + 30 2 = Sb,0 3 + 3H 2 O (6) 4SbH 3 -f 30 2 = 4Sb + 6H 2 O (c) 2SbH 3 = 2Sb + 3H 2 When the antimony hydride (stibine) is passed into a solution of silver nitrate, the silver is reduced, leaving the antimony with the silver, as (tiiHntnuonx urniide 9 SbAg 3 , a black precipitate, distinction front arsenic, which enters into solution (69, 6'a and &); SbH 3 -f 3AgN0 3 SbAg 3 + 3HNO, . The precipitate should be filtered and washed free from unde- composed silver salt (and arsenous acid, if that be present), and dissolved \\ith dilute hydrochloric acid (HC1 does not dissolve uncombined anti- 80 ANTiuoyr. 70, 7. mony, 5o) : SbAg 3 + 6HC1 = SbCL + 3AgCl + 3H, . The solution con- sists of antimonous chloride,, leaving silver chloride as a precipitate. However, in the excess of hydrochloric acid used a small portion of the silver chloride may be dissolved (59, oc), interfering with the final test for the antimony. If this be the case the silver should be removed by a drop of potassium iodide (8). Stibine is not evolved by the action of strong- KOH upon zinc or aluminum, nor by sodium amalgam in neutral or alkaline solution (distinction from triad arsenic); the antimony is precipitated as the metal (Fleitmann, -/. C., 18! 329). Stibine is slowly oxidized by sulphur to Sb 2 S 3 in the sunlight at ordinary temperature and rapidlj' when the sulphur (in a U tube mixed with glass wool) is heated to 100. The reaction takes place according to the following equation: 2SbH 3 + 3S 2 = Sb,S 3 + 3H 2 S (Jones, J. C., 1876, 29, 645). 7. Ignition. By ignition in the absence of reducing agents, antimonic acid and anhydride are reduced to antimonous antimonate. Sb,O 3 .Sb.O 5 or Sb,O 4 (Sb'"SbvOJ, a compound unchanged at a dull red heat, but when heated to 800 this oxide is further reduced to antimonous oxide (ib). The antimonates of the fixed alkali metals are noi vaporized or decomposed when ignited in the absence of reducing agents; hence, by fusion in the crucible "with sodium carbonate and oxidizing agents, i. e., with sodium nitrate and car- bonate, the compounds of antimony are converted into non-volatile sodium pyroantimonate, Na 4 Sb 2 O 7 , and arsenic compounds if present are at the same time changed to sodium orthoarsenate, Na 3 AsO 4 . If now the fused mass be digested and disintegrated in cold water and filtered, the antimonate is sepa- rated as a residue, Na;,H 2 Sb,O 7 (4c), while the arsenate remains in solution with the excess of alkali. The operation is much more satisfactory when the arsenic and antimony are previously fully oxidized as by digestion with nitric acid as the oxidation by fusion in the crucible is not effected soon enough to retain all the arsenic or antimony which may be in the state of lower oxides, sulphides, etc. If compounds of tin are present in the operation and it' the fusion is not done with excess of heat, so as to convert sodium nitrite to caustic soda and form the soluble sodium stannate the tin will be left as stannic oxide, SnO, , in the residue with the Na^H-Sb.O, . But if sodium hydroxide is added in the operation, the tin is separated as stannate in solution with the arsenic (Meyer, ./. C.. 1849, 1, 388). All compounds of antimony are completely reduced in the dry way on char- coal with sodium carbonate, more rapidly with potassium cyanide; the metal fusing to a brittle globule. The reduced metal rapidly oxidizes, the white antimonous oxide rising in fumes, and making a crystalline deposit on the support. If now ammonium sulphide be added to this white sublimate, an orange precipitate is a sure indication of the presence of antimony (Johnstone, C. y., is$3. 58, 296). The same white oxide is formed on heating antimony or its sulphides in a glass tube, through which air is allowed to pass. 8. Detection. Antimony 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 ammoniuin sulphide * it i? sepa- rated from Hg , Pb , Bi , Cu , and Cd . In the Marsh apparatus the anti- mony is precipitated on the Zn as the metal, a portion being still further reduced to stibine. By passing the gases, stibine and arsine, into AgNO, solution, the antimony is precipitated as SbAg, , antimony argent ide, sepa- * Antimony as sulphide solution in potassium sulphide may be detected electrolytically, being deposited as Sb. Delicate to one part in 1,500,000 (Kohn, J. Soc. Iwl., 1891. 1 0, 327). 70, 10. ANTIMONY. 31 rating it from the arsenic which is oxidized and passes into solution as arsenous acid. The SbAg 3 is dissolved in HC1 and the presence of the antimony is confirmed by the precipitation of the orange colored sulphide with H 2 S . Study text at 6 and 84 to 89. For distinction between Sb v and Sb'" see 89, 7. 9. Estimation. (1) Tartaric acid and water are added to SbCl 3 , which is then precipitated by H 2 S as Sb 2 S , and after washing on a weighed Gooch filter, it is heated to 230 in a stream of CO 2 , in order to exclude oxygen, and weighed. (2) Antimonous oxide, sulphide, or any oxysalt of antimony is first boiled with fuming nitric acid, which converts it into SbjOs , and then by ignition it is reduced to Sb 2 O 4 , and weighed as such. (3) The trichloride is precipitated by gallic acid, and weighed after drying at 100. (4) In the presence of tin and lead oxidize the hydrochloric acid solution of the salts with KC1O (the tin must be present as Sn IV ) and distil in a current of HC1 . The stannic and anti- mony chlorides are volatile (separation from lead). To the distillate add metallic iron, obtaining stannous chloride and metallic antimony; filter and wash (sep- aration from tin). Fuse the precipitate with sodium nitrate and sodium car- bonate, digest the fused mass with cold water, filter, wash, dry and weigh as Na. : H 2 Sb 2 O 7 (7) (Tookey, J. C., 1862, 15, 462; and Thiele, A., 1894, 263, 361). (5) For estimation of antimony and separation from arsenic and tin by the use of oxalic acid, see Lessen (Z., 1888, 27, 218) and Clarke (C. N., 1870, 21, 124). (6) Volumetrically. The antimony compound is converted into stibine (6j) and the gas passed into standard silver nitrate solution. The solution is filtered and the excess of silver nitrate is titrated with standard sodium chloride. If arsenic be present it must also be estimated (69, 9 (15)), and the true amount of antimony present computed from the two determinations (Houzeau, J. C., 1873, 26, 407). (7) Sb'" is oxidized to Sb v in presence of NaHCO 3 by a standard solution of iodine. The end of the reaction is shown by the blue color given to starch. (8) Sb"' is oxidized by KC1O 3 in strong HC1 solution to SbCls . KI is added, which reduces the Sb? to Sb'" with the liberation of I 2 , which is titrated with Na-jS-jOs solution. (9) Sb"' is oxidized to Sb v in presence of H 2 C 4 H4O e by KMnO 4 . (10) Sb'" is oxidized to Sb v by K 2 CrO 7 , and the excess of K 2 Cr 2 O 7 used is determined by a standard solution of FeSO 4 , K 3 Fe(CN) 6 being used to show the end of the reaction. (11) The antimony as the triad salt is treated with an excess of standard K 3 Fe(CN) 6 ; the excess of which is estimated in a gas apparatus with H 2 2 (Baumann, Z, angew,, 1892, 117), 10. Oxidation. Stibine, SbH 3 , is decomposed by heat alone into anti- mony and hydrogen (6;). By burning in the air it is oxidized to Sb.,0 ; and H 2 . Passed into a solution of silver nitrate, SbAg 3 is produced, or passed into a solution of antimonous chloride or potassium hydroxide, sp. gr. 1.25, metallic antimony is produced. Excess of chlorine, bromine, or nitric acid in presence of water oxidizes it to Sb v ; but if the SbH 3 be in excess metallic antimony is precipitated. With excess of iodine in pres- ence of water Sb'" is produced; if the stibine be in excess metallic anti- mony. Metallic antimony is oxidized by nitric acid, chlorine or bromine to Sb"' or Sb v , depending upon the amount of these reagents and the temperature. Iodine oxidizes the metal to Sb'" only, except in alkaline mixtures when Sb v is formed. Antimonous compounds are oxidized to antimonic compounds by Cl , Br , HN0 3 , K 2 Cr 2 7 , and KMn0 4 ; by silver oxide in presence of the fixed alkalis (6t); by gold chloride in hydrochloric acid solution, gold being 82 TIN. 71, 1. deposited as a yellow precipitate (73, 10). The antimony is precipitated as Sb 2 5 unless sufficient' acid be present to dissolve the oxide: 4AuCl 3 -f- 3Sb 2 3 -f 6H 2 = 4Au + 3Sb 2 O g + 12HC1 . Antimonic compounds are reduced to antimonous compounds by HI (6/) and by SnCL (69 and 71, 10); the antimony not being further recuci'd (distinction from As). Antimonic and antimonous compounds are reduced to the metallic state by Pb , Sn , Bi , Cu , Cd , Fe , Zn , and Mg ; but in the presence of dilute acids and metals which evolve hydrogen the antimony is still further reduced to stibine. Iron alone or in the presence of plat- inum (iron platinum wire couple) precipitates the antimony from acid solu- tions as Sb ; 0.000012 grams can be detected (Rideal, C. N., 1885, 51, 292). Sodium amalgam with dilute sulphuric acid evolves stibine from all antimony solutions (Van Bylert, B., 1890, 23, 2968) but the generation of hydrogen in alkaline solution, i. e., Zn -f- KOH , causes the reduction of the antimony salt to the metal only, in no case evolving stibine. 71. Tin (Stannum). Sn = 118.7. Valence two and four. 1. Properties. Specific gravity, 7.2984 (Rammelsberg, B., 1870, 3, 724); welting point, 231.68 (Callendar and Griffiths, C. N., 1891, 63, 2). Boils at 2275 (Greenwood, Proc. Roy. Soc., 82, 396, 1908). Does not distil in a vacui m at a red heat (Schuller, J., 1884, 1550). Tin is a silver white metal, does not t; rnish readily in pure air. At a red heat it decomposes steam with evolution of hydrogen; at a white heat it burns in the air with a dazzling white light, fo:ming ShOo . It is softer than gold and harder than lead, can readily be hammered or rolled into thin sheets (tinfoil); at 100 it can be drawn into wire a id at 200 can be pulverized. Tin possesses a strong tendency to crystalline struc- ture, and when bar or block tin is bent a marked decrepitation "Zinngesc hrei" (Levol, A. Ch., 1859, (3), 66, 110) is noticed, due to the friction of the en stals. Block tin exposed to severe cold (winter of 1867-68, at St. Petersburg, --39) crumbles to a grayish powder (Fritsche, B., 1869, 2, 112), considered to be an allotropic modification. This same property of crumbling is noticed in samples of tin that have been preserved several hundred years (Schertel, J. pr., 1879, 2, 19, 322). The grayish powder is an allotropic modification of tin, the tran- sition temperature being 20 C. Tin forms alloys with many metals. Eronze consists of copper and tin, brass frequently contains from two to five per cent of tin, solder consists of lead and tin. All the easily fusible metals as Wood's metal, etc., contain tin. For many references concerning tin alloys, see Watts (IV, 720). 2. Occurrence. The chief ore of tin is cassiterite or tinstone, a nearly pure crystallized dioxide, SnO 2 ; found in England, Australia, Malay Peninsula, Bolivia, Mexico, and to a very limited extent in the United States; (Z)., 2, 1, 643). Stannite, Cu 2 FeSnS 4 , is found in small quantities in various tin veins. 3. Preparation. The reducing- agent employed is carbon. The impure ore, SnO, , is first roasted, which removes some of the arsenic as As 2 O 3 , and some of the sulphur as SO, . Then, by washing-, the soluble and some of t-.eju- soluble impurities are washed away, the heavier SnO 2 remaining-. It is then fused with powdered coal, lime being 1 introduced to form a fusible slag with the earthy impurities. It is refined by repeated fusion. Strictly pure ,in is best made by treating- the refined tin with HN0 3 , and then reducing- the oxide thus formed by fusion with charcoal; or by reducing- the purified chloride 4. Oxides and Hydroxides. Tin forms 'two stable oxides and cor respo id ing- classes of salts: starmous oxide, SnO, black or blue black, and stannic c xide. SnO 2 , white; the latter acts both as a base, in stannic salts, and as an Anhy- dride, in stannates. Stan-nous o.ridc is formed (1) by precipitating SnCl 2 with K 2 CO 8 , washing- with boiled water in absence of air, drying- at 80 or 1 >wer; then dehydrating by heating- in an atmosphere of hydrogen or carbon dioxide 71, o&. TIN. 83 (Loni?e, r. r., jssci, ::i); (2) by melting- a mixture of SnCl 2 and Na 2 C0 3 with .stirri.ig- until it becomes black, and removing the NaCl by washing (Sandal, rhil. May., 1838, (3), 12, 216; Bottger, /!., 18.39, 29, 87). titan-nous hi/dro.ridc, Sn(C I) 2 *, white to yellowish white, is formed by adding- alkalis or alkali carbc nites to stannous chloride, washing and drying at a low temperature (Ditt ;, A. Ch., 1882, (5), 27, 145). (12.) &(< iiic oxide exists in two forms, crystalline and amorphous. The native tinst ne is nearly pure crystalline SnO 2 . For preparation see Bourgeois (C. r., 1.887, -04, 231) and Levy and Bourgeois (('*. r., 1882, 94, 1305). Amorphous SnO 2 is fo tned (1) by heating tin in the air to a white heat; (2) stannic salts are preci 'Hated by alkali carbonates, the precipitate washed and ignited; (3) tin is DJ dized by nitric acid; (4) tin filings are ignited in a retort with HgO (/>., - , 1, 647). Stannic lii/dro.i-idc or stannic acid exists in two forms: (1) Nor- mal i "annic acid, SnO(OH) 2 = H 2 SnO :{ , is formed when a solution of stannic chloi de is precipitated by barium or calcium carbonate (Freing, Poyg., 1842, 55, 519); if an alkali carbonate be used some alkali stannate is also formed. (2) Meta tannic acid, H^Sn-.Oj,. , is formed by decomposition of tin with nitric acid Hay, C. N., 1870, 22, 298; Scott, C. N., 1870, 22, 322); insoluble in acids but chan ed on standing with acids to normal stannic acid, which is readily soluble in a< ds (56). It is also formed when stannic chloride is boiled in concen- trate I solution with most of the alkali salts: 5SnCl 4 + 20Na 2 S0 4 -4- 15H 2 = H 10 S.i 5 16 + 20NaCl + 2()NaHS0 4 , or according to Fresenius (16th edition), 271: SnCl 4 + 4Na 2 SO, + 4H 2 O = Sn(OH) 4 + 4NaCl + 4NaHSO 4 . It is also form- d together with hydrochloric acid when stannic chloride is boiled with a large excess of water. 5. Solubilities. a. Metal. Tin dissolves in hydrochloric acid slowly when the i "Ad is dilute and cold, but rapidly when hot and concentrated, stannous chlor de and hydrogen being produced (a); in dilute sulphuric acid slowly, with separation of hydrogen (6), (not at all even in hot acid if more dilute than H 2 SC 1 .6H 2 p (Ditte, A. Ch., (5), 27, 145); in hot concentrated sulphuric acid rapid y, with separation of sulphurous anhydride and sulphur (c); nitric acid, rapid, y converts it into metastannic acid, insoluble in acids (d)' } very dilute nitric acid dissolves it without evolution of gas as stannous nitrate and ammo- nium nitrate (e] (Maumene, BL, (2), 36, 598); nitro-hydrochloric acid dis- solves tin easily as stannic chloride (/), potassium hydroxide solution dissolves it ve.:y slowly, and by atmospheric oxidation (p); or, at high temperatures, with evolution of hydrogen (h). Bromine vapors readily attack melted tin with formation of SnBr< , colorless crystals, melting point 30 (Carnelley and O'Shca, J. C., 1878, 33, 55). Dry chlorine gas attacks tin readily in the cold, producing stannic chloride as vapor or colorless liquid. The action is vigorous enough in strong chlorine to produce a flame. Sn + 2HC1 = SnCL + H 2 Sn + H 2 S0 4 SnS0 4 + H 2 Sn + 2H 2 S0 4 = SnSO, + 2H 2 O + S0 2 and then 4SriS0 4 + 2SO, + 4H 2 S0 4 = 4Sn(S0 4 ) 2 + S 2 + 4H 2 O (d) 15Sn + 20HN0 3 + 5H 2 = 3H 10 Sn 5 O 15 + 20NO (e) 4Sn + 10HN0 3 = 4Sn(NO 3 ) 2 + 3H.O + NH 4 N0 8 (0 Sn + 2C1 2 = SnCl t (g) 2Sn + 4KOH + 2 = 2K 2 SnO 2 + 2H 2 O (7t) Sn -f 2KOH = K 2 Sn0 2 + H 2 6. Oxides. Stannous oxide is insoluble in water, soluble in acids (Ditte, A. Ch., 1882, (5), 27, 145; Weber, /. C., 1882, 42, 1266), oxidized by nitric acid when heat?d, forming the insoluble metastannic acid. Sidiinous hydroxide is readily soluble in all the solvents of the oxide, and is also readily soluble in fixed alkali hydroxides. Sia-nnic o.i-ide, Sn0 2 , is insoluble in water; soluble with difficulty in alkalis; insoluble in acids except in concentrated H 2 SO 4 (D., 2, 1, 648). Sulphur forms SnS 2 and SO-,; chlorine forms SnCl 4 (Weber, Pogg., 18(51, 112, 619). Normal stinuiic acid, BLSnOj , freshly precipitated, is soluble in According to other authorities Sn(OH) 8 does not exist, hut a liydratccl oxide is formed, SnO . Sii(OH) a (Graham-Otto, a, 2, 1207 ; D., 2, 1, 657; Graelin-Kraut, , 107). 84 TIN. 71, 5c. fixed alkali hydroxides and in acids (Ditte, C. r., 1887,104, 172); insoluble in water and changed by hot nitric acid to the insoluble metastannic acid. Mela- stannic acid, H 10 Sn 5 Oio, is insoluble in water and acids, HC1 changes it to metastannic chloride insoluble in the acid, but soluble in water after removal of the acid; soluble in the fixed alkalis as metastannates, which are soluble in water and precipitated by acids. Metastannic acid in contact with HC1 is gradually changed to stannic acid (Barfoed, J. pr., 1867, 101, 368). C. Salts. The sulphides and phosphates of tin are insoluble in water, also stannous oxychloride; stannous sulphate,* bromide and iodide; and stannic chloride and bromide dissolve in pure water with little or no decomposition (Personne, C. r., 1862, 54, 216; and Carnelley and O'Shea, J. C., 1878, 33, 55). Stannous chloride is soluble in less than two parts of water (Engel, A. Ch., 1891, (6), 17, 347); but more water decomposes it, unless a strong excess of acid be present: 2SnCl 2 -f- H 2 = SnO.SnCl 2 + 2HC1 . Pure stannic chloride is a liquid; sp. gr., 2.2; boiling point, 114 (Walden, Z. ph. Ch., 43); solidifies at 33 (Besson, C. r., 1889, 109, 940). The liquid combines with water, liberating heat to form crystals of SnCl.).3H 2 O , which are readily soluble in excess of water (D., 2, 1, 662). Stannic chloride is completely decomposed by boiling water. The nitrates of tin are very easily decomposed by water and require free acid to keep them in solution (Weber, J. pr. y 1882, (2), 26, 121; Montemartini, Gazzetta, 1892, 22, 384). Stannic iodide is readily soluble in water (Schneider, Pogg., 1866, 127, 624). Stannic sulphate is easily soluble in water, but is de- composed by a large excess (Ditte, C. r., 1887, 104, 171). Stannous and stannic chloride, and stannic iodide are soluble in alcohol. Stannous nitrate and stannic sulphate, and bromide are deliquescent. Stannous sulphide is insoluble in water, soluble in HC1 with formation of H2S; decomposed by HNO 3 with oxidation to metastannic acid; insoluble in solution of the normal alkali sulphides, but soluble in the polysulphides with oxidation to a stannic compound (6e). Stannic sulphide is soluble in HC1 , with evolution of H 2 S ; and in solutions of the alkali sulphides. 6. Reaction's. a. Alkali hydroxides and carbonates precipitate from solu- tions of stannous salts, stannous hydroxide, Sn(OH)2 (4) , white, readily soluble in excess of the fixed alkali hydroxides, insoluble in water, ammonium hy- droxide and the alkali carbonates (distinction from antimony). It is also precipitated by barium carbonate in the cold (Schaffner, A., 1844, 51, 174). SnCl 2 + 2KOH Sn(OH) 2 + 2KC1 Sn(OH) 2 + 2KOH = K 2 Sn0 2 + 2H 2 SnCl 2 + 4KOH K 2 SnO, + 2KC1 + 2H 2 SnCl 2 + Na 2 CO s + H 2 O = Sn(OH) 2 + 2NaCl + C0 2 By gently heating the solution of potassium stannite, K 2 Sn0 2 , crystalline stannous oxide, SnO ,. is formed. By rapid boiling of a strong potassium hydroxide solution of stannous hydroxide part of the tin is oxidized and the remainder precipitated as metallic tin; 2K 2 Sn0 2 + ^ 2 = Sn -f- K 2 Sn0 3 + 2KOH . The reaction proceeds more rapidly upon the addition of a little tartaric acid. Stannic salts are precipitated by alkali hydroxides and carbonates as stannic acid, H 2 Sn0 3 soluble in excess of the fixed alkali hydroxides, insoluble in ammonium hydroxide and the alkali carbonates (Ditte, A. Ch., 1897 (6), 30, 282). SnCl< + 4KOH = H 2 SnO 8 + 4KC1 + H 2 H 2 Sn0 8 + 2KOH = K 2 SnO, + 2H 2 O SnCl 4 + 6KOH = K 2 Sn0 8 + 4KC1 + 3H 2 O SnCl, + 2Na 2 C0 8 + H 2 = H 2 Sn0 8 + 4NaCl + 2CO 2 * Stannous sulphate is decomposed by an excess of cold water forming 2Sn8O 4 .4SnO.8H a O| and by a small amount of hot water forming SnSO^.gSnO (Ditte, A. Ch., 1883, (5), 27, 161). |71, 6. TIN. 85 Metastannic salts are precipitated as metastannic acid soluble in potassium hydroxide not too concentrated, not readily soluble in sodium hydroxide, insoluble in ammonium hydroxide excepting when freshly precipitated in the- cold, and the alkali carbonates. 6. Oxalic acid forms a white crystalline precipitate with a nearly neutral solution of staiinous chloride, soluble in hydrochloric acid, not readily soluble in ammonium chloride. If a nearly neutral solution of stannous chloride be added drop by drop to a solution of ammonium oxalate, the white precipitate which forms at once dissolves in the excess of the ammonium oxalate. Stannic chloride is not precipitated by oxalic acid or ammonium oxalate (Hausmann and Loewenthal, A., 1854, 89, 104). Potassium cyanide precipitates both stannous and stannic salts, white, in- soluble in excess of the cyanides. Potassium ferrocyanide precipitates from stannous chloride solution stannous ferrocyanide, Sn 2 Fe(CN) 6 , white, insoluble in water, solublp in hot concentrated hydrochloric acid. From stannic chloride is precipitated a greenish white gelatinous precipitate, soluble in hot hydro- chloric acid, but reprecipitated upon cooling (distinction from antimony) (Wyrouboff, A. Ch., 1876, (5), 8, 458). Potassium ferricyanide precipitates from solutions of stannous chloride, stannous ferricyanide, Sn 3 (Fe(CN) ) 2 , white, readily soluble in hydrochloric acid. On warming, the ferricyanide is reduced to ferrocyanide with oxidation of the tin. No precipitate is formed by the ferricyanide with stannic chloride. c. The nitrates of tin are not stable. Stannous nitrate is deliquescent and soon decomposes on standing exposed to the air. Stannous salts when heated with nitric acid are precipitated as SnO 3 ; but if stannous chloride be warmed with a mixture of equal parts of nitric and hydrochloric acids, stannic chloride and ammonium chloride are formed (Kestner, A. Ch. t 1860, (3), 58, 471). df. Hypophosphorous acid does not form a precipitate with stannous or stannic chlorides, nor are these salts reduced when boiled with the acid. Sodium hypophosphite forms a white precipitate with stannous chloride, soluble in excess of hydrochloric acid; no precipitate is formed with stannic chloride. Phosphoric acid and soluble phosphates precipitate from solutions of stannous salts, not too strongly acid, stannous phosphate, white, of variable composition, soluble in some acids and KOH; insoluble in water (Lenssen, A., 1860, 114, 113). With stannic chloride a white gelatinous precipitate is formed, soluble in HC1 and KOH , insoluble in HNO S and HC,H 8 2 . If the stannic chloride be dissolved in excess of NaOH before the addition of Na 2 HPO 4 and the mixture then acidulated with nitric acid, the tin is completely precipitated as stannic phosphate (separation from antimony). However, the precipitate always car- ries a little antimony (Bornemann, Z. angew., 1899, 635). e. Hydrosulphuric acid and soluble sulphides precipitate from solutions of stannous salts dark brown hyd rated stannous sulphide, SnS (a), insol- uble in dilute, soluble in moderately concentrated HC1 (b). It is readily dissolved with oxidation by alkali supersulphides, the yellow sulphides, forming thiostannates (c); from which acids precipitate the yellow stannic sulphide (d). The normal, colorless alkali sulphides scarcely dissolve any stannous sulphide at ordinary temperature, compare (69, 6e and 70, 6e), but hot concentrated K 2 S dissolves SnS forming K 2 SnS 3 and Sn (e) (Ditte, C. r., 1882, 94, 1419; Baubigny, J. C., 1883, 44, 22). Potassium and sodium hydroxides dissolve it as stannites and thiostannites (f), from which acids precipitate again the brown stannous sulphide (g). Am- monium hydroxide and the alkali carbonates do not dissolve it (distinction from arsenic, 69, 6e). The insolubility in fixed alkali carbonates is a 86 2W. 71, 6/. distinction from antimony (70, 6e). Nitrohydrochloric acid (free chlorine) dissolves it as stannic chloride, with residual sulphur (h). Nitric acid oxidizes it to raetastannic acid without solution (t) (separation from arsenic, 69, 6e). (a) SnCl 2 + H 2 S = SnS + 2HC1 (&) SnS + 2HC1 = SnCL + ELS (c) SnS + (NH 4 ) 2 S, = (NH. t ) 2 "snS 3 (d) (NH 4 ) 2 SnS 8 + 2HC1 = SnS 2 + 2NH 4 C1 + H 2 S (e) 2SnS + K 2 S = K 2 SnS 3 + Sn (f) 2SnS + 4KOH = K 2 Sn0 2 + K 2 SnS 2 + 2H 2 O (0) (K,Sn0 2 + K 2 SnS 2 ) + 4HC1 = 2SnS + 4KC1 + 2H 2 (7i) 2SnS + 4C1 2 = 2SnCl 4 -f S 2 (?) SOSnS + 40HN0 3 + 10H 2 = GH 10 Sn 5 15 + 40NO + 15S 2 Solutions of stannic salts are precipitated as stannic sulphide, SnS 2 . hydrated, yellow, having much the same solubilities as those given for stannous sulphide, with this difference, that stannic sulphide is moderately soluble in normal, colorless, alkali sulphides. The following equations illustrate the most important reactions: SnCl 4 + 2H 2 S = SnS 2 + 4HC1 SnS 2 + 4HC1 = SnCl 4 + 2H 2 S SnS 2 + (NH 4 ) 2 S = (NH 4 ) 2 SnS 3 2SnS 2 + 2(NH 4 ) 2 S 2 = 2(NH 4 ) 2 SnS 3 + S 2 3SnS 2 + 6KOH = K 2 SnO 3 + 2K 2 SnS 3 + 3H 2 O (K 2 Sn0 3 + 2K 2 SnS 3 ) + 6HC1 = 3SnS 3 + GKC1 + 3H 2 O SnS 2 + 2C1 2 = SnCl 4 + S 2 15SnS 2 + 20HN0 3 + 5H 2 = 3H 10 Sn 5 15 + 15S 2 + 20NO Sodium thiosulphate does not form a precipitate with the chlorides of tin (separation from As and Sb) (Lesser, Z., 1888, 27, 218). Sulphurous acid and sodium sulphite precipitate from stannous chloride solution not too strongly acid, sta-nnous sulphite, SnS0 3 , white, readily soluble in HC1 . When wariaed in the presence of hydrochloric acid, si^lphur dioxide acts as an oxidizing agent upon the stannous salt. A precipitate of Sn 6 O 10 S 2 or SnS 2 is formed, or H 2 S is evolved and SnCl 4 formed, depending upon the amount of HC1 present. 6SnCl 2 + 2S0 2 + GH 2 = Sn 6 O 10 S 2 + 12HC1 6SnCl 2 + 2S0 2 + 8HC1 = SnS 2 + 5SnCl 4 + 4H 2 O 3SnCl 2 + S0 2 + 6HC1 = 3SnCl 4 + H 2 S + 2H 2 O Stannic chloride does not give a precipitate with sulphurous acid or sodium sulphite. The sulphates of tin are formed by dissolving the freshly precipitated hydroxides in sulphuric acid and evaporating at a gentle heat. They cannot be formed by precipitation and are decomposed by water (Ditte, A. CJi., 1882, (5), 27, 145). f. Potassium iodide added to a concentrated water solution of stannous chlo- ride forms first a yellow precipitate soluble in excess of the SnCl, . Further addition of KI gives a yellow precipitate rapidly turning to dark orange needle- like crystals, often forming in rosette-like clusters. If a drop of the stannous chloride solution be added to an excess of potassium iodide the yellow p^ecipi- tate is formed, which remains permanent unless a further quantity of sts nnous chloride be added when the orange precipitate is formed. The orange p -ecipi- tate is probably SnI, , and is soluble in HC1 , KOH , and C,H 5 OH , soluble in large excess of KI and sparingly soluble in H 2 O with some decomposition. 71, 7. TIN. 87 The yellow precipitate is probably a double salt of stannous iodide and potas- sium iodide, and has about the same solubilities as the orange precipitate (Personue, J., 1862, 171; Boullay, A. (7/,, 1K27, (2), 34, 372). Potassium iodide in concentrated solution precipitates xtumnc iodide, yellow, from very concentrated \vate> solutions of stannic chloride. The precipitate is readily soluble in water to a i olorless solution (Schneider, T e i 3 & 03 O " to M IB ts T 8 "13 1-T N S S * IP 5 C co ~ M - rt a> < ^ n ^3 p liiuji sgs*/ w P ^ o-S ~ c oq c W .r rib. eg f* !i & III LI lils 18 03 +j y ^^ ^ ^ D03 S^ s " ^^s ^^ s ccO^T- GS^rt o" X n 03 ja a;^ ^b- 03 O.* W W iii i , '02 O M . "3 K ^-O ho: ,-o O ^ 03 ^'3M .S p;z ^ p 0) 71 ^ 03 v gj -r^ _ ^ ^fc"? 9 ^ V ^' r- -^11 '~-o ^o^^o'b- o g S 03 03 o O o ^ b- >rH -H v - > rQ C ,0 x- ^ r j- C^ > ||^a " ^ "S rr ^^g.2 ;&H ^ 0,0 c fc M -t- K (D lip <*g w W ^ TS^'^S C 2 'o^ 111? swl 5.1 ^.2 ^co 2i W r* jfrfJ ^'S 5^1 1 M ,a.st w o f-t 02 O 02 0) WS S" c X s gS!s* ^ o o 13 02 fl S3 be ^ _r b . fid^ t O^i3 . err IgSgV' S C K t P W 73, 5a. GOLD. 91 73. Gold (Aurnm) Au= 197.2. Valence one and three. 1. Properties. Specific gravity, 19.30 to 19.34 (Rose, Pogg., 1848, 76, 403). Melting point, 1063 (Cr. B. 8., 35, 1915). It is a yellow metal, that from dif- ferent 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 (DeviUe 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 incl.-ils, 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. This consists in treating the well-powdered ore with a stream of water, the heavy gold settling to the bottom. (2) Amalga- mation. This 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 1 . The gold is afterward separated from the lend 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 . () By amalgamation with mercury and electrolysis at the same time. 4. Oxides and Hydroxides. Aurous oxide, Au 2 O , is very unstable, heating to about 250, decomposing it into the metal and oxygen. The hydroxide is pre- pared by reducing the double bromide with SO 2 in ice-cold solution; heating to 200, changing it to the oxide (Kriiss, A., 1886, 237, 274). Auric hydroxide, Au(OH)., , is prepared by precipitation from the chloride solution with MgO (Kriiss, 7. e.). It is a yellow to brown powder, changing to the oxide upon dry- ing at 100. Heating to 230 gives the metal and oxygen (^10). 5. Solubilities. a. Metal. Gold is not at all tarmsnea or in 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 aurif 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 3 : 4Au + 12HI -f 4KI + 30 2 = 4KIAuI 3 + 6H 2 . Potassium cyanide solution, with aid of the air, dissolves precipitated gold as potassium auro- cyowide, XAu readily dissolved on addition of more hydrochloric or nitric acid (distinction from the silver group chlorides). Hydrobromic acid and soluble bromides do not precipitate solutions of bis- muth chloride, but do precipitate solutions of the nitrate, forming the oxy- bromide, BiOBr , white. The presence of potassium bromide prevents the pre- cipitation of a bismuth chloride solution by water and also dissolves the oxy- chloride which has been precipitated by the addition of water. Hydriodic acid and soluble iodides precipitate from solutions of bismuth salts, unless strongly acid, bismuth iodide, black or brownish gray crystals, quite readily soluble in excess of the reagent * or in strong HC1 without warm- * Bismuth iodide dissolves in solution of potassium iodide with an intense yellow color, deli- cate to one-millionth (Stone J. Sac. Chem. Jwl., 1887, 6, 416). The potassium iodide solution of bismuth iodide is used as Drag-endorff's reagent to detect the presence of an alkaloid. Leger (Bf,, 1888, 50, 91) uses cmchonine and potassium iodide to prove the presence of bismuth. Del- icate to one-five hundred thousandth. Other metals must be removed. 76, 9. BISMUTH. 103 ing 1 . It is reprecipitated on diluting the solution with water. Bismuth iodide is scarcely at all decomposed l>.v washing with cold water, but on boiling with water it is deeom posed into bismuth o\y-iodide, BiOI , red, insoluble in KI , soluble in HC1 , and in HI (Gott and Muir, J. C'., 1888, 53, 137). Chloric acid dissolves bismuth hydroxide, but the compound deeomposes upon evaporation (\Vachter, A., 1844, 52, 233). Potassium bromate and iodate both precipitate solutions of bismuth nitrate. The iodate formed is scarcely soluble, the bromate easily soluble in HN0 3 . g. Potassium or sodium stannite hot, when added in excess to bismuth solutions, cause a black precipitate, from reduction to metallic bismuth, a very delicate reaction.* The stannite is made, when wanted, by adding to a stun nous chloride solution, in a test-tube, enough sodium or potas- sium hydroxide to redissolve the precipitate at first formed: 2BiCl 3 -f- 3K 2 Snol + (iKOH = 2Bi + 6KC1 + 3K,Sn0 3 + 3H 2 (Vanino and Treu- l.H>rt, B., 1898, 31, 1113). h. Solutions of bismuth salts, nearly neutral, poured into a hot solution of potassium bichromate precipitates the orange red chromate, (BiO) 2 Cr 2 O 7 ; but if poured into a cold solution of the neutral chromate a citron-yellow precipi- tate, .iBLCK.L'OO, , is formed. These precipitates are soluble in moderately concentrated acids, insoluble in fixed alkalis (distinction from Pb). The pre- cipitate with KvCro0 7 is used in the quantitative determination of bismuth (9). 7. Ignition. On charcoal, with sodium carbonate, before the blow-pipe, bis- muth is readily reduced from all its compounds. The (/lobule is easily fusible, brittle (distinction from lead), and gradually oxidizable under the flame, form- ing an Incrustation (Bi 2 O 3 ), orange-yellow while ho V "lemon-yellow when cold, the edges bluish-white when cold. The incrustation disappears, or is driven by the reducing flame, without giving color to the outer flame. With borax or microcosmic salt, bismuth gives beads, faintly yellowish when hot, colorless when cold. A mixture of equal parts cuprous iodide and sulphur forms an excellent reagent for the detection of bismuth in minerals by the use of the blow-pipe. The reagent mixed with the unknown is fused on charcoal or on a piece of aluminum sheet. A red sublimate indicates bismuth. Mercury gives a mix- ture of red and yellow sublimates (Hutchings, C. N., 1877, 36, 249). Bismuth chloride may be sublimed at the temperature of boiling sulphur; recommended as a separation from lead (Remmler, B., 1891, 24, 3554). 8. Detection. Bismuth is precipitated from its solutions by H 2 S form- ing Bi 2 S 3 . By its insolubility in (NH 4 ) 2 S or (NHJJS., and solubility in hot dilute HN0 3 it is separated with Pb , Cu , and Cd from the other metals of the tin and copper group. Dilute H 2 S0 4 removes the lead and NH,OH precipitates the bismuth as Bi(OH) 3 , leaving the Cu and Cd in solution. The presence of the bismuth is confirmed by the action of a hot solution of K 2 Sn0 2 or NaOH and formaldehyde f on the white precipitate of Bi(OH) 3 , giving metallic bismuth (6*7) or by dissolving the Bi(OH) 3 in HC1 and its precipitation as BiOCl upon dilution with water (5d). 9. Estimation. (1) As metallic bismuth formed by fusion with potassium cyanide. (2) As Bi 2 O 3 formed by ignition of bismuth salts of organic acids, or of the salts of volatile inorganic oxyacids. (3} By precipitation by HsS , and *For a modification of this test se3 Muir (J. C ., 1S77, :\", 43). t Sodium stannite reduces lead hydroxide while formaldehyde does not. Traces of lead may be present owing to imperfect separation. 104 COPPER. 76,10. after drying at 100, weighing- as Bi 2 S s . (4) By precipitation by K 2 Cr 2 7 , and after drying at 120, weighing as (BiO) 2 Cr 2 O 7 . (.5) VolumetricaUy. By" precipi- tation with K 2 Cr 2 7 . Dissolve the chromate in dilute acid, transfer to an azotometer and reduce the chromate with hydrogen peroxide (Baumann, Z. angew., 1891, 331). (6) By precipitation as a phosphate with standard sodium phosphate; dilution to definite volume and determination of the excess of phosphate in an aliquot part with uranium acetate (Muir, J. (7., 1877, 32, 674). 10. Oxidation. Metallic bismuth reduces salts of Hg ? Ag, Pt , and An to the metallic state. Bismuth is precipitated as free metal from its solutions by Pb , Sn , Cu , Cd , Fe , Al , Zn , Mg , and HH 2 P0 2 (6d). All salts of bismuth are oxidized to Bi 2 5 by Cl or H 2 2 in strong alkaline mixture (Hasebrock, B., 1887, 20, 213; Scruff, A. Ch., 1861 (3), 63, 474). All compounds of bismuth are reduced to the metal by potassium stannite K 2 Sn0 2 (60). Bismuth chloride or bromide heated in a current of hydro- gen is partially reduced to the free metal (Muir, J. C., 1876, 29, 144). It is precipitated as free metal upon warming in alkaline mixture with grape sugar (56). 77. Copper (Cuprum) Cu = 63.57 . 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, 1083 (Cir. B. S., 35, 1915). 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 pro- tec l s it from further action of air or of water. It forms a number of very im- portant 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. Chalcopyrite (CuFeS 2 ); chalcocite (Cu 2 S); green malachite (Cu 2 (OH) 2 CO,); blue malachite (Cu 3 (OH) 2 (CO)j); cuprite (Cu 2 O); and tenorite (melanconite) (CuO). 3. Preparation. For the details of the various methods of copper-owning- and refining, the works on metallurgy should be consulted. In the laboratory pure copper may be produced (1) by electrolysis; (2) reduction by ignition in hydrogen gas; (5) reduction of the oxide by ignition with carbon, carbon monoxide, illuminating gas, or other forms, of carbon; (4) reduction of the oxide Toy K or Na at a temperature a little above the melting point of these metals; (5) reduction by fusion with potassium cyanide: CuO + KCN = Cu -f KCNO . For its reduction in the wet way, see 10. 4. Oxides and Hydroxides. Cuprous oxide (Cu 2 O), red, is found native; it is prepared: (1) 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 KOBE or NaOH . It is stated by Rose (Pogg., 1863, 120, 1) that tetracupric monoxide, (Cu 4 O , is formed by treating 1 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 (Kriiss, #., 1884, 17, 2593). (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 3 ) 2 + iH 2 O -f- 2NO (Freer and ILigley, Am.. 1899, 21, 377); also in hot concentrated sulphuric acid, with evolution of sulphurous anhydride: Cu + 2H 2 S0 4 = CuS0 4 + 2H 2 O -f SO 2 . If dry hydrochloric acid gas be passed over heated copper, CuCl is formed with evolution of hydrogen (Weltzien, A. Ch., 18G5, (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 w r hen 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 4 ) 2 S , colorless, acts upon copper turnings with evolution of hydrogen, forming Cu,S (Heumann, J. C. y 1873, 26, 1105). &. 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., 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 aq.) 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. n. Fixed alkali hydroxides precipitate acid solutions o 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 C0 3 . Fixed alkalis KOH added to saturation in solutions of cupric salts, precipitate cupric hydroxide, Cu(OH) 2 , deep blue, insoluble in excess unless concentrated (Loe\v, 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 (OH) 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 nqt 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 glurnso, 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 blue solution is a cuprammonium compound, not formed unless ammonium salts be present. It has been isolated as CuS0 4 .(NH 3 ) 4 (77, 56). The (loop blue solution probably consists of this compound in a hydrated condition, t. e. Cu(OH) 2 .2NH 4 OH.(NH 4 ) 2 S0 4 ; or (NH 4 ) 4 Cu(OH) 4 S0 4 . Other salts than the sulphate form the corresponding compounds: CuCl 2 + 4NH 4 OH = Cu(OH),.2NH 4 OH.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 K 2 C0 3 precipitate the greenish-blue, basic carbonate, Cu 2 (OH) 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, f err o cyanides, 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 thiocyamite 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, 60. COPPER. 10? oxalate, CuC2O 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 a brownish precipitate of cupric cyanide, Cu(CN) 2 , which immediately changes to the yellowish green cupric cuprous cyanide with the evolution of cyanogen. On warming, this precipitate changes to the white cuprous cyanide, Cu 2 (CN) 2 , with further evolution of cyanogen. Excess of potassium cyanide dissolves the precipitate with the formation of the double cuprous salt, K 3 Cu(CN) 4? or, in the presence of a smaller excess of potassium cyanide, K,Cu(CN) 3 . (Kunschert, Z. Anorg., 41, 359, 1904). The potassium cyanide also dissolves cupric oxide, hydroxide, carbonate, sulphide, etc., changing rapidly to the double cuprous cyanide in solution in the alkali cyanide. Hydrogen sulphide does not precipitate the copper from solutions of copper salts in potassium cyanide on account of the very slight concentration of the copper ions in such solutions (5.10~ 20 . Kunschert, ibid.). This serves as a separation from Cadmium which is precipitated by hydrogen sulphide from a potas- sium cyanide solution. Potassium ferrocyanide precipitates cupric ferro- cyanide, Cu 2 Fe(CN) (! , reddish-brown, insoluble in acids, decomposed by alkalis; a very delicate test for copper (1 to 200,000) ; forming in highly dilute solutions a reddish coloration (Wagner, Z., 1881, 20, 351). Potassium ferricyanide 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 4 OH; in the presence of hypophosphorous or sulphurous acid the cuprous thiocyanate is precipitated at once (distinction from cadmium and zinc) (Hutchinson, J. (7., 1880, 38, 748). Ammonium benzoate (10 per cent solution) precipitates copper salts completely from solutions slightly acidified (separation from cadmium) (Gucci, B., 1884, IT, 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.e.} 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, Gazzctta, 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_,HPO 4 , gives a bluish-white precipitate of copper acid phosphate, CuHPO 4 , if the reagent be in excess and Cu,(PO 4 ) a 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). e. Cuprous salts (obtained by treating cupric salts with SnCL) when boiled with precipitated sulphur deposit the copper as Cu.S (separation from cad- mium) (Orlowski, ./. ('., 1882, 42, 12:52). Cuprous salts "are precipitated or trans- 108 COPPER. 77, 6/. 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 CiuS), produced alike in acid solutions (distinction from iron, manganese, cobalt, nickel) and in alkaline solutions (distinction from arsenic, antimony, tin). Solu- tions containing only the one-hundred-thousandth of copper salt are colored brownish 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 ammonium monosulphide, but slightly soluble in ammonium polysulphide and fixed alkali. (Bossing, Z., 41, 1) { (distinction from arsenic, antimony, tin) ; soluble in solution of potassium c} r anide (distinc- tion from lead, bismuth, cadmium, mercury). According to Noyes (J. Am. Soc. 29, 170) 5 to 10 mg. of copper may dissolve in ammonium polysulphide when a large amount of copper is present. Concerning the formation of a colloidal cupric sulphide, see Spring (B., 1883, 16, 1142). According to Brauner (C. AT., 1896, 74, 99) cupric salts with excess of hydrogen sulphide always yield a ver.y appreciable amount of cuprous sul- phide. See also Ditte (C. 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, If., 1888, 9, 165); in acetic acid solution, separation from cadmium and zinc (Dovath, Z., 40, 141), /. Hydrobromic acid added to cupric solutions and concentrated by evaporation gives a rose-red color. Delicate to 0.001 m. g. (Endemann and Prochazka, (7. A 7 ., 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 J (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- * Freshly precipitated cuprous sulphide transposes silver nitrate forming 1 silver sulphide, metallic silver und cupric nitrate ; with cupric sulphide, silver sulphide and cupric nitrate are formed (Schneider, Pogg., 1874, 152, 471). Freshly precipitated sulphides of Fe Co, Zii, Cd. Pb, Bi, Sn", and Sn lv , when boiled with CuCl in presence of NaCl give Cu.,S and chloride of the metal: with CuCl 2 , CuS and a chloride of the metal are formed, except that SnS gives Cu,S, CuCl and S IV (Raschig, B., 1884, 17, 697). t Thio salts having the formulas (NH02 CiiaS? and NaiCujS? are formed (Rossing, Z. anorg. 25, 407.) t 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, 0. COPPER. 109 pletely, bj adding reducing agents with iodides; as, Na 2 SO. } , H S0 3 , FeS0 4 (6). (a) 2CuS0 4 + 4KI = 2CuI + I 2 -f 2K 2 S0 4 (6) 2CuS0 4 + 2KI + 2FeSO, 2CuI + K 2 S0 4 + l ; e 2 (S0 4 ), 2CuSO 4 + 4KI -f H 2 S0 3 + H 2 = 2CuI + 2K 2 SO 4 + H 2 S0 4 + 2HI ff. Arsenites, as KAsO 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 arsenite, chiefly CuHAs0 3 (Scheele'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,O g ),Chl(C,H,Oj) a , 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. * Potassium bichromate does not precipitate solutions of cupric salts: 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 microco&mic 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 OH removes the lead and ammonium hydroxide precipitates the bismuth as Bi(OH) 3 , leaving the Cn 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 (66) ; 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 and weighed as the metal, or by means of zinc, the excess of zinc being 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 Na 2 S 2 O s , and, after adding free sulphur and igniting in hydrogen gas, weighed as cuprous sul- phide, or it may be precipitated by KCNS in presence of H ? SOs or H PO 2 , and, after adding S , ignited in H and weighed as Cu 2 S . Cu ? O , CuO , Cu(NO 3 ) 2 , CuCO 3 , CuSO 4 and many other cupric salts, are converted into Ou 2 S by adding S and igniting in hydrogen gas. (4) By adding KI to the cupric salt and titrating the liberated I by Na 2 S 2 O 3 ; not permissible with acid radicals which oxidize HI . (5) By precipitation as Cul and weighing after drying at 150 (Browning, 110 CADMIUM. 77, 10. Am. S., 1893 [3], 46, 280). (6) By titrating- in concentrated HBr , using a solution of SnClo in concentrated HC1; the end reaction is sharper than Avith SnCL, alone (Etard and Lebeau, C. r., 1890, 110, 408). (7) By titration with Na,S. Zinc does not interfere (Borntrager, Z. angeio., 1893, 517). (8) By reduction with SO, and precipitation with excess of standard NH 4 CNS; dilu- tion to definite volume and titration of the excess of NH 4 CNS in an aliquot part, with AgNO 3 (Volhard, A., 1878, 190, 51). (9) Small amounts are treated with an excess of NH 4 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 KOH 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 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. (7., 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 An'", these salts being reduced to the metallic state. Ferric iron is reduced to the ferrous condition (Hunt, Am. $., 1870, 99, 153). Copper is also oxi- dized by many acids. 78. Cadmium. Cd 112.4. Valence two. 1. Properties. '-pecific gr vtiy, liquid, 7.989; cooled, 8.67; hammered, 8.6944. Melting point, 320.9 (Cir. B. S., 36, 1915). Boiling point, 763 to 772 (Car- nelley and Williams, /. C., 1878, 33, 284). Specific heat is 0.0567. Vapor density (H = 1), 55.8 (Deville and Troost, A. Ch., 1860, (3), 68, 257). From these data the gaseous molecule of cadmium is seen to consist of one atom (Richter, Anonj. Chan., 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 SO 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 (Kanimerer, 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 Cl , 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) by 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 CO, from the air, becom- ing white (Gmelin-Kraut, 3, 64). The Jiydroxule -Cd(OH) 2 is formed by the action of the fixed alkalis upon the soluble cadmium salts; it absorbs CO 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 *Warren, C, N., 1895, 71, 92, 78,6 CADMIUM. Ill 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). ft. The oxide and Jiiidrn.ridc 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, J. Plianii., 1885, (5), 11, 553). C. Salts. The sulphide, carbonate, oxalate, phosphate, cyanide, ferrqcy_anide and ferricyanide are insoluble (27) in Avater, soluble in hydrochloric and nitric acids, and soluble in NH 4 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 and citric acids, and certain other organic substances precipitate, from solutions of cadmium salts, cadmium hydroxide, Cd(OH) 2 , white, insoluble 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(CN).,.2KCN; ffirrocyanides 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. Cli., 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). * Owing to the formation of incompletely-dissociated CdCl 3 . CdI 2 is still less dissociated and accordingly CclS dissolves more readily in HI than in HC1 and much more readily than in HNO 3 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 CADMIUM. 78, 6/. Sodium thiosulphate, Na 2 S s O 8 , does not precipitate solutions of cadmiuir 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, C. N., 1891, 64, 294). Cadmium salts with excess of sodium thiosulphate are not precipitated upon boiling 1 with hydrochloric acid (distinction from copper) (Orlowski, J. C., 1882, 42, 1232). f> 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 oiv addition of water. * A solution of copper and cadmium salts, very dilute, when allowed to spread upon a filter paper or porous porcelain ple.te, gives a ring of the cad- mium salt beyond that of the copper salt, easily detected by hydrogen sulphide (Bag-ley, J. C., 1878, 33, 304). 7. Ignition. On charcoal, with sodium carbonate, cadmium salts are reduced by the blow-pipe to the metal, which is 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 forming CdS. By its insolubility in (NH 4 ) 2 S or (NH 4 ) 2 S Z and solubility in hot dilute HN0 3 it is separated with Pb , Bi , and Cu from the other metals of the second group. Dilute H 2 S0 4 with C 2 H 5 OH removes the lead and NH 4 OH precipitates the bismuth as Bi(OH) 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 Gi. 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 O 4 and titrated by KMnO 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 KI; the I removed by heating; the excess of KI removed by KNO 2 and H 2 SO 4 ; the cadmium precipitated by Na 2 CO 8 and ignited to CdO (Browning, Am. 8., 1893, 146, 280). (6) By adding a slight excess of H 2 SO 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 . , 1. PRECIPITATION OF METALS OF SECOND GROUP. 113 79. Comparison of Certain Reactions of Bismuth, Copper, and Cadmium. Taken in Sohilio-ns of their Chlorides, Nitrates, Sulphates, or Acetates. Bi Cu Cd KOH . r NaOH, in Bi(OH) 8 , white. Cu(OH) ;i , dark blue Cd(OH) a , white. NH 4 OH, in excess Dilution of satu- rated solutions Bi(OH) 3 , white. BiOCl, white (76, :,'/) Blue solution. Colorless solution. Iodides Partial precipita- Precipitation of tion in solutions Cul with libera- Sulphides Iron or zinc not very strongly acid (76, 6f). Bi. 2 S 3 , black, in- soluble in KCN. Bi spongy precipi- tion of iodine (77, 6/). Cu 2 S and CuS, black, soluble in KCN. Cu bright coating CdS, yellow, insol- uble in KCN. Cd gray sponge Glucose, KOH, and Vtonf tate. Bi, black. (77, 10). Cu 2 0, yellow (77, K.h\ with zinc, no ac- tion with iron. K 2 8nO, -f KOH.. Bi, black. 50). Cu, precipitated SYSTEMATIC ANALYSIS OF THE METALS OF THE TIN AND COPPER GROUT. The precipitation of the metals of the second group (Tin and Copper Group) by hydrosulphuric acid, and their separation into Division A (Copper Group) and Division B (Tin Group). See 312. 80. Manipulation. The filtrate from Group 1 (62), or the original solution, if the metals of the silver group be absent is warmed nearly to boiling and saturated with H 2 S gas. The volume of the solution should be about 50 c.c. and it should contain about 6 per cent of concentrated HC1 by volume. After passing H 2 S for about 15 minutes through the hot solution, allow it to cool, dilute with an equal volume of water and again pass H 2 S for some time through the cold solution. Shake well to coagulate the precipitate and filter. Pass H 2 S again through the filtrate, filter and repeat until no further H 2 S precipitate is obtained. 2H 3 AsO 4 + xHCl + 5H,S = As 2 S 6 + xHCl + 8H 2 O or 2H 3 AsO 4 + xHCl + 5H 2 S = As 2 S, + xHCl + S, + 8H a O SnCl 4 + 2H 2 S = SnS 2 -f- 4HC1 SnCl 2 + H 2 S - SnS + 2HC1 2Bi(NO 3 ) 3 + 3H 2 S = Bi 2 S 3 + 6HNO, CdSO, + H 2 S = CdS + H,SO 4 81. Notes. /. Hydrosulphuric acid gas should be used in precipitating the metals of the second group. It may be generated in a Kipp apparatus, using ferrous sulphide, FeS , and dilute commercial sulphuric acid (1-12). Com- mercial hydrochloric acid may be used instead of sulphuric. The gas should 114 PRECIPITATION OF METALS OF SECOND GROUP. 81, #. 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 capil- lary tube into the solution to be analyzed contained in a flask. Less gas is required and the solution is less liable to be thrown from the test tube by the excess of unabsorbed gas. 2. In treating the unknown solution with H 2 S, it should be passed into the liquid until, upon shaking the flask, capped with the thumb, there is no forma- tion of a partial vacuum due to the further absorption of the gas by the liquid. 8. H 2 S is decomposed by HNO 3 or HNO 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.) has been assured. If this group is present strontium and especially, barium, will invariably be present in the H 2 S precipitate on account of the oxi- dation of the sulphur to sulphuric acid. For this reason, oxidizing agents must be removed from the solution so far as possible. If ferric chloride is present, 15 milligrams of barium may be present in this precipitate as sulphate. Curt man and Frankel (J. Am. Soc., 33, 724, 1911.) For detection of the barium see 301, 5. 4- 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. A moderate excess of HC1 is necessary to insure the precipi- tation of arsenic if present in the arsenic condition. For this purpose the solution must be hot and must contain at least 6 per cent by volume of con- centrated HC1. Under these conditions the arsenic precipitates slowly (69). The strong acid, especially when hot, prevents the precipitation of the other metals, especially tin, lead and cadmium. For this reason, the solution must be cooled and diluted and again saturated with H 2 S in order to precipitate these metals. The solution must not be too largely diluted or traces of Co , Ni and Zn will be precipitated. About one part of HC1 to 25 of the solution must be present to prevent the precipitation 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 , 324). 5. 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 2 S 3 , and As 2 S 5 are lemon-yellow; Sb-S 5 and SboS? are orange; SnS , HgS , PbS, Bi 2 S 3 , 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 (67, 6e}. Mercuric chloride at first forms a white precipitate of HgCl 2 .2HgS, changing from yellow to red, and finally to black with more H 2 S , due to the gradual conversion to HgS (58, 6e). 6. 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 nec- essary to their resolution may prevent the precipitation of cadmium or cause the formation of the red precipitate with lead chloride. 7. The presence of a strong oxidizing agent as HNO 3 , K L >Cr 2 O 7 , FeCl 3 , 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. * 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 sulphidos 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. Hut 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 S ions and so the solubility-product is reached and precipitation results. 83,2. PRECIPITATION OF METALS OF SECOND GROUP. 115 If the original solution be dark colored, it is advisable to warm with hydro- chloric acid and alcohol (126, 6/ 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. 8. 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, 6e). The first cause of error may be avoided by careful observance of the directions in note (2). Too much acid may be present because excess of acid had been used in dissolving the unknown. After precipitating the first group, excess of nitric should be removed by evaporating the solution nearly to dryness then diluting and adding the required amount of HC1. As a further precaution a portion of the filtrate from the H 2 S precipitate should be diluted with several volumes of water and H^S passed. If a precipitate is obtained, the entire solu- tion should be diluted and saturated with H 2 S. As v must be precipitated by passing H 2 S rapidly through the hot moderately acid solution before dilution as long as the slow formation of the arsenic precipitate continues. 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 beaker, using as small an amount of water as possible. If As, Sb and Sn are present,* ammonium sulphide (NH 4 ) 2 S (38, 2) is then added and the precipitate digested for several minutes with warming: As 2 S 8 + 2(NH 4 ) 2 S 2 = (NH 4 ) 4 As 2 S 5 + S 2 SnS + (NH 4 ) 2 S 2 = (NH 4 ),SnS 3 2SnS 2 + 2(NH 4 ) 2 S, = 2(NH 4 ) 2 SnS 3 + S 2 2Sb 2 S 3 + 6(NH 4 ) 2 S 2 =:4(NH 4 ) 3 SbS 4 + S 2 2MoS 3 + 2(NH 4 ) 2 S 3 2(NHJ 2 MoS 4 + S 2 The precipitate is then filtered and washed once or twice with a small amount of (NH 4 ) 2 S, and then with hot water. The filtrate consisting of solutions of the sulphides of As, Sb , Sn , Au, Ft, Mo (Gr, Ir, Se, Te, IV, !'), 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, Eli, and Ru), constitutes the Copper Group (Division B of the second group,, 95). 83. IVo/rs. 1. The precipitate of the sulphides of the tin and copper group must be thoroughly washed with hot water (preferably containing 1 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). :.'. Yellow ammonium sulphide, (NH 4 ) 2 S X , forms upon allowing- Iho 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). * 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 warming with 1 or 2 cc. (NH-OaSx . If it all dissolves, only As, Sb, Sn can be present; if nothing dissolves, none of these can be present; if part dissolves, then the whole 2nd group precipitate must be so treated. To see if anything has dissolved in the (NH 4 ) 2 Sx it is acidified slightly with HC1 (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 s treated with (NHOzSx , the solution is filtered and acidified just as the test portion was. 116 TABLE FOR THE ANALYSIS OF THE TIN GROUP. >84. Pi .g 3 PH W M p g* K *" s^,3 I -^ S ^ . 5 * 5 H -P V S f i|i|ilJ5 i 03 W "isl 4 ft^n3 S l)0 . w fi C6 >84. TABLE FOR THE ANALYSIS OF THE TIN GROUP. 117 *l5 fl ^w H 4 " H -- O c3 limit i lit ilif. O p Ci s ^ !|I o B a --- f]!j 3 * ft OC3 il . - 1 . lii s 118 DIRECTIONS FOR ANALYSIS WITH NOTES. 83, 3. 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,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 (N"H 4 ),S X . If copper be present arid mercury be absent, it is recom- mended to use K L ,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 S X 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 (NH 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. 85. Manipulation. The solution of the sulphides in (NH 4 ) 2 S X - is care- fully acidulated with hydrochloric acid: 2(NH 4 ) 2 S 2 + 4HC1 = 4NH 4 C1 + S 2 + 2ELS (NH 4 ) 4 As 2 S 3 + 4HC1 = As 2 S, + 4NH 4 C1 + 2H,S 2(NH 4 ) 3 SbS 4 + 6HC1 = Sb 2 S 5 + 6NH 4 C1 + 3H 2 S (NH 4 ) 2 SnS 3 + 2HC1 = SnS 2 + 2NH 4 C1 + H 2 S The precipitate obtained when the metals of the tin group are present, is usually yellovr 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 vita hot water and then dissolved in hot HC1 using small fragments o2 LC10 3 (69, 6e) to aid in the solution: 2As 2 S 3 -:- 10C1 2 + 16H,,0 = 4H 3 As0 4 + 20HC1 + 3S 2 SnS 2 + 4HC1 = SnCl 4 + rts 2 -:- ^ci 2 The solution iz 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,,) are quite soluble in concentrated HC1 . .?. 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 l>\ the precipitation of a solution of SnCL with H 2 S , giving a brown precipitate of SnS , then dissolving this precipitate in (NH 4 ) L .S X and reprecipitating with HC1 as the orange-colored SnS,. . //. Hot reagent HC1 (324) 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- DIRECTIONS FOR ANALYSIS WITH NOTES. 119 eentrated HC1 . The relative solubility of these sulphides in HC1 is used as the basis of the following separation of As from Sb and Sn (69, 6e, also bottom of next note, 5). Free. : As 2 S 5 , SboS 5 , SnS 2 . (sp. gr. 1.2). Expel all H 2 S . Heat for a few moments with concentrated HC1 Dilute a little and filter. Residue :As 2 S ft , (S 2 ). Apply either of the following tests: (1) Pour warm NH 4 OH over the precipi- tate. Add H 2 O 2 to this solution and boil. Cool and add a few cubic centi- meters of NH 4 C1 and a little MgCl 2 and obtain a white crystalline precipitate of MgNH 4 AsO 4 . (2) Dissolve As 2 S 5 in HC1 + crystal of KC1O 3 . Boil. Make alkaline with NH 4 OH , and add NH 4 C1 and MgCl 2 as in (1). Filtrate: SbCl 3 , SnCl 4 . Boil to be sure of complete expulsion of H 2 S . Test for Sb: Put drop of solution on silver coin. Bend piece of tin in form of i 1. Touch one end to drop and other to coin outside the drop. Allow to stand for a few moments. Brown or black spot on coin is due to metallic Sb . Test for Sn: Heat solution with iron wire until reduction is complete. Filter and to filtrate add HgCL A white precipitate of HgCl or a gray precip- itate of Hg shows tin. The precipitated sulphides of As , Sb , Sn are 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 HC1 (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 SbCl 3 and SnCl 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 O 2 , 2 to 3 cc. NH 4 C1 , and 2 to 3 cc. "magnesia mixture," which is MgCl 2 -f- NH 4 C1 4- NH 4 OH . Cool, and let stand for a time. The As v is precipitated as NH. ; MgAsO 4 , a white, crystalline precipitate exactly like NH,MgPO 4 in ap- pearance. As 2 S 6 + 16NH 4 OH + 20H 2 O 2 = 2(NH 4 ) 3 AsO 4 + 5(NH 4 ) 2 SO 4 + 28H 2 O . (NH 4 ) 3 As0 4 + MgCl 2 + [NH 4 OH + NH 4 C1] = MgNH 4 AsO 4 + 2NH 4 C1 . 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 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 filtrate slowly (a few drops at a time), to a few cc. of ammonium molybdate, (NH 4 ) 2 MoO 4 solution. A deep blue color or precipitate will appear if Sn * is present, due to the reduction of the MoO 3 to a lower exide. Or, instead of adding this filtrate to molybdate solution, it may be treated with HgCl 2 , a white precip- itate of HgCl being formed if Sn * is present. Note that this is reversing the test for Hg with SnCl 2 . The HgCl 2 test is most characteristic. 120 DIRECTIONS FOR ANALYSIS WITH NOTES. \,5. The precipitation of Ag 2 S 3 , unlike that of the other sulphides, is not prevented by the presence of any amount of HC1 , however large, but, on the contrary, is aided. It may, therefore, be necessary, after removing all other sulphites in the N/5 HC1 solution, to add several cc. of concentrated HC1 , heat to boiling, and pass in H 2 S for some time to precipitate the rest of the As . In the cold, H 3 Asp 4 is very slow- ly precipitated by H 2 S , but strong HC1 and heat accelerate the reaction very much. It is essential that the sulphides be thoroughly washed before treatment with HC1. CuS is slightly soluble in (NH 4 ) 2 S X and may give a coloration when the solution is acidified. (NH 4 ) 2 S , which is colorless, gives no precipitate of S upon addition of excess of acid; (NH 4 ) 2 S X , yellow, always gives more or less S , white and difficult to filter. 2 S 2(NH 4 ) 2 S X + 4HC1 (NH 4 ) 2 S + 2HC1 = Make a blank test on the Fe wire used, to see that its solution in HC1 gives no test for Sn with molybdate. 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 practically insoluble. The following table suggests a method of analysis based upon this property of these sulphides. Digest the mixed sulphides with solution of ammonium carbonate and filter. Residue: SnS 2 , Sb a S 5 , (S) . Dissolve in hot hydrochloric acid (5c, 70 and 71). Solution: SnCl 4 , SbCl 8 . Treat with zinc and hydrochloric acid in Marsh's apparatus (69, 6'a). Deposit: Sn , (Sb) . Dissolve by hydro- chloric acid. Solution: SnCl 2 . (Residue, Sb .) Test by ammoniacal silver nitrate and by mercuric chlo- ride (71, 6i and ;). Gas: SbH 3 . (Test the spots, 69, 6'c, 1.) Receive the gas in solution of silver nitrate. Dissolve the precipitate (Sb Ag s ) (70, 6j), and test by H 2 S (87 and 89). Solution: (NH 4 ),AsS 4 + (NH 4 ) 3 As0 4 and (NH 4 ) 4 As 2 S 5 + (NH 4 ) 4 As 2 5 . Precipitate by hydrochloric acid; 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 8 As0 4 . Apply Marsh's Test, as directed in 69, fi'rt, 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 arsenic may also be identified by adding HC1 to the ammonium carbonate solution, passing H 2 S and dissolving the precipitate in a small amount of concentrated HN0 3 . The arsenic will be oxidized to H 3 As0 4 . Divide the solution into two parts. Cautiously neutralize one portion with 86, 10. DIRECTIONS FOR ANALYSIS WITH NOTES, 121 ammonia. When the solution is nearly neutral, add AgN0 3 arid a drop or two of ammonia without shaking the solution. A reddish brown ring of Ag 3 As0 4 will form at the neutral zone of the solution. To the other por- tion, add magnesia mixture and ammonia until the solution is alkaline. A crystaline precipitate of MgNH 4 As0 4 will form on standing. 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 a vessel wide enough to allow the carbonic acid to escape without waste of the solution. The soluble diammonium dihydrogen pyroantimonate, (NH 4 ) 2 H 2 Sb 2 O 7 , is formed. Meanwhile the SnCl 4 is fully precipitated as H 2 SnO 8 (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- tnonous chloride, SbCl 3 . Excess of ammonium carbonate does not redissolve the Sb 2 3 which it precipitates from SbCl s , as stated in 70, 6a. The above plan may also be varied as follows: After removal of the arsenic sulphide with (NH 4 ) 2 CO 8 , the residue is dissolved in strong HC1 , not using KC1O 3 or HNO 3 . The solution consists of SnCl 4 and SbCl, . 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 Hg-Cl 2 (Pieszczek, Arch. Pharm., 1891, 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 Chemtet, 1875, 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 HN0 3 until brown fumes are no longer evolved, and then evaporated to dryness, using sufficient heat to expel the HN0 8 and the H 2 S0 4 formed by the action of the HNO 3 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 AsO 4 , and the antimony and tin remaining as residue of Sb 2 O fl and Sn0 2 . The pres- ence of arsenic may be confirmed by the reactions with AgN0 8 (69, 6;), CuSO 4 (69, Gfc) by the Marsh test (69, 6'a), or by precipitation with magnesia mix- 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, (71, 61). 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 KC1O S . 8. Instead of chlorine (HC1 + KC1O 3 ), nitrohydrochloric acid may be em- ployed, but it is liable to cause the formation of a white precipitate of Sb,0 5 and Sn0 2 . 9. The chlorine should all be removed, as the metals cannot be reduced by the Zn and H S SO 4 in the Marsh apparatus in the presence of powerful oxidizing agents as Cl . This would also require evaporation to expel the HNO 8 , i nitrohydrochloric acid were used to effect solution. 10. Hydrogen peroxide, H 2 O 3 , 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). 122 DIRECTIONS FOR ANALYSIS WITH NOTES. 87. 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 + 6AgN0 3 + 3H 2 = H 3 As0 3 + 6Ag + 6HNO 3 SbH 3 + 3AgN0 3 = SbAg 3 + 3HNO 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'&). Failure to obtain mirror, spots, or a black precipitate in the AgN0 3 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 nitrate is treated with HC1 , or a metallic chloride, as CaCl 2 or Nad , 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, + 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 nitration. The nitrate 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 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 obtained in the hard glass tube should be examined as directed in the text, especially by oxidation and microscopic examination (69, 6'c 5). The spots should be tested with NaClO 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 4 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.,0 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, 90. DIRECTIONS FOR ANALYSIS WITH NOTES. 123 3. The arsine evolved is not decomposed (faint traces decomposed) upon passing- through a drying 1 tube containing- soda lime or through a solution of KOH (distinction and separation from antimony). 4. Arsenites and arsenafes 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). (b) 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, 6#). (c) Magnesia, mixture precipitates arsenic acid as white magnesium ammonium arsenate, MgNH 4 AsO 4 ; no precipitate with arsenous acid (189, 60). (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). (/) 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 dig-esting with (NH 4 ) 2 C0 3 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 AgNO 3 solution must not be taken as evidence of the presence of antimony, as arsine gives a black precipitate of metallic silver with AgNO 3 . A trace of antimony may be found in the filtrate from the SbAg s , hence a slight yellow-orange precipitate from this solution must not be taken as evidence of arsenic without further examination (69, 7). 3. Sb,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, 6). This is a convenient method 0f 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 C0 3 and NaOH; and the cooled mass disintegrated with cold water, the nitrate will contain the arsenic as sodium arsenate, Na 3 As0 4 , and the tin as sodium stannate, Na 2 SnO 3 ; while the antimony remains as a residue of sodium pyroantimonate. Na 2 H 2 Sbo0 7 (70, 7). 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 ig-nition with K 2 CO 3 as preliminary to the preparation of the sub- stance for analysis, since they hinder the complete precipitation of the anti- mony with H 2 S (70, 6e). 7. Antimonic compounds are reduced to the antimonous condition by HI with liberation of iodine (70, Qf and 10). Chromates oxidize antimonous salts to antimonic salts with formation of green chromic salts (70, 6ft). KMnO 4 also oxidizes antimonous salts to antimonic salts, a mang-anous 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 nitrate, if colorless, may be 124 DIRECTIONS FOB ANALYSIS WITH NOTES. 91,^. 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 HN0 3 , which oxidizes the molybdenum to molybdic acid, Mo0 3 . The residue is dissolved in NH 4 OH (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 , Ail , 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 a SnCl 4 -f Hg SnCl, + 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, 6i). 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 (6/&, 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, 6&), or by treating with a mixture of SnCl 2 and SnCl 4 (Cassius' purple) (73, 6#). If a portion of the original solution, free from HN0 3 , be boiled with oxalic acid the gold is completely precipitated as the metal, separation from the platinum which is not precipitated (74, 6&). 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 SO 4 acting- as a reducing 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, 6d). 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 cr3 r stallization (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 HoS than arsenic or antimony (89, 3). 96. DIRECTIONS FOR ANALYSIS WITH NOTES. 125 2. The yellow ammonium sulphide (NH 4 ).,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), readily precipitates metallic silver black from solutions of silver salts. An arsenite (hot) or an 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 (7O, 6i). This reaction also detects stannous salts in the presence of stannic salts. 4- 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 Hg-Cl 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 remains 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 FeS0 4 before proceeding with the usual method of analysis. 4. 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, G time, until no further acflon takes place, 126 TABLE FOR ANALYSIS OF THE COPPER GROUP. ^ EO t -a w HI g s ii 43 0>iO ' a 5* -d e CJ *2 _. 3*0 M IS: S o 8 nc of f f H 2 SO 4 ,t , until fumes precipitation - %/ '43 &,- ^ S 3 O' 'I 13! n-g s a e P g w s 2 s^ 1 ! ' 95. TABLE FOlt ANALYSIS OF THE COPPER GROUP. fcJDJL C 1 t> rn TJ > P-l+3 o> d a) ' mi .. a; oj TH ^ 003 Isl "So 1 -+J &0! TH rS ...TH O 00 CD O -pO OJTH &O5 r 00 m Is" 00 oo ^~ X CO 05 | 00 CO ID -d -a c ** 5 a ^ ^8 2 Is o s ^ d w ? S H 'd i-t sS r2 O h p o 2 a si ,H S -o d 3 Si Isll ^ a ca w ?& 2&^i S^Q! fl >, - ilfl og a, 3 Mil >% o 5^ -^ C d 3 8 o ^ ^i^^ -S DIRECTIONS FOR ANALYSIS WITH NOTES. 97, 1. Mercuric sulphide is unattacked (58, Qe) 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 bo tested later, and the black residue on the filter is dissolved in nitro-hydrochloric acid : 2HgS + 2C1 2 = 2HgCl 2 + 3 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, (58, 60) : HgCL + SnCl 2 == Hg + SnCl 4P 2HgCl 2 + SnCl 2 == 2HgCl + SnCl,'; or by the deposition of a mercury film on a strip of bright copper wire (5C, 10) : HgCl 2 + Cu = Hg + CuCl 2 . Confirm further by bringing in contact with iodine in a covered dish: Hg + I 2 = HgI 2 (Jannaesch, Z. anorg., 1896, 12, 143). The mercury may also be detected by using NH 4 OH and KI as the reverse of the Nessler's test (207, 6fc) (delicate 1 tj> 31,000) (Klein, Arch. PJiarm., 1889, 227, 73). 97. Notes. 1. The concentration of HNO 3 (1-2) is necessary for the solution of the sulphides of Pb , Bi , Cu and Cd , and may also dissolve traces of HgS . However, the concentrated HNO :i (sp. gr., 1.42) dissolves scarcely more than traces of HgS (58, 6e). Long-continued boiling of HgS with concentrated HNO 3 changes a portion of the HgS to Hg(N0 3 ) t .HgS , a white precipitate, insoluble in HNO 3 . 2. 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, and Cu . One part of HNO 3 to three parts HC1 gives about sufficient HC1 to decompose all the HN0 3 , hence in this reaction a little more than that proportion of HC1 should be ued. 3. A small amount of black residue left after boiling the sulphides with HNO 3 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 HNO S will always oxidize a trace at least of sulphur to H 2 S0 4 (256, 6B, 2), which will form PbSO 4 if any lead be present: S 2 + 4HNC-3 = 2H 2 S0 4 + 4NO 3PbS + 8HNO 3 = 3PbS0 4 + 4H 2 -f- 8NO If the boiling ,be not continued too persistently, the amount of PbS0 4 formed is soluble in the HN0 3 present (57, 5o), and does not at all remain behind with the HgS . 6. Even if only 1 to 2 mg. of As or Sb are present with a large quantity (500 ing.) of an element of the copper group enough is dissolved by either (NH 4 ) 2 S or (NH 4 ) 2 S X for the detection of these metals. If only 3 to 5 mg. of Sn are present with a large quantity of elements of the copper group, all of the tin may remain undissolved. When Cd is present and the tin is in the stannous state as much as 15 mg. of Sn may remain undissolved even in the polysulphide (A. A. Noyes, J. Am. Chem. Soc. 29, 170). The Sn or Sb (present on account of an insuffi- ciency of (NH 4 ) 2 S X ) will appear as a white precipitate mixed with the black precipitate of HgS , due to the fact that HNO 3 decomposes the sulphides of Sb and Sn , forming the insoluble Sb>O 5 and SnO. : 6Sb 2 S 3 -f- 20HNO 3 = 6Sb 2 O 6 + 9S 2 -f 20NO + 10H 2 O If these metals have not been detected this precipitate must be tested. Aftei testing for mercury in a portion of the precipitate, the paper may be burned 100. DIRECTIONS FOR ANALYSIS WITH NOTES. 129 in a porcelain crucible and the residue fused with sulphur and sodium carbonate in the covered crucible. The tin and antimony will be converted into soluble thio salts and tested for according to 84. 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, C. r., 1880, 90, 141). 7. Merciiry may quickly be detected from all of its compounds by ignition in a hard glass tube with fusion mixture (Na,CO s + K,C0 3 ) (58, 7), and then adding- a few drops of HN0 3 (concentrated) and a small crystal of KI . Upon warming- the iodine sublimes and combines with the sublimate of Hg , forming the scarlet red HgI 2 . As and Sb both give colored compounds with iodine, de- composed by HNO 3 '(Johnstone, C. 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 2 S0 4 are given off: Pb(NO s ) 2 + H,,S0 4 = PbS0 4 + 2HNO 3 Cu(N0 3 ), + H 2 S0 4 = CuS0 4 + 2HN0 3 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, G/ and h). 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^SO 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,S0 4 . 2. The nitric acid should be removed by the evaporation, as PbSO 4 is quite appreciably soluble in HNO 3 (57, 5c). 3. The H,SO 4 should be present in some excess, as PbSO 4 is less soluble in dilute H;,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 H 2 SO 4 (57, 5c, Ge). 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. 6. If the (NH 4 ) 2 S X had not been well removed by w r ashing, ammonium sul- phate would be present at this point, greatly increasing the solubility of PbS0 4 (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 OH . An excess of 130 DIRECTIONS FOR ANALYSIS WITH NOTES. 101, 1. NH 4 OH should be added to dissolve the precipitates of Cu(OH) 2 and Cd(OH) 2 , leaving the Bi(OH) 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, Qg). 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 A1(OH) 3 would be white, but would not give a black precipitate with K 2 SnO, . 2. If an insufficient quantity of (NH 4 ) 2 S X was used, Sb and Sn would be present and give a white precipitate with the NH 4 OH . 3. If the lead had not been removed it would appear as a white precipitate with the NH 4 OH , and would give a brownish-black precipitate with the hot K 2 SnO 2 (57, 6#). The presence of a permanent white precipitate with NH 4 OH must never be taken as final evidence of the presence of Bi . 4. As a confirmatory test for the presence of Bi , a portion of the white precipitate with NH 4 OH should be dissolved in HC1 and the solution evapo- rated nearly to dryness to remove the excess of HC1 . Now upon adding water, a white precipitate of BiOCl , bismuth oxychloride, will be obtained if Bi is present (76, 5(7). 102. Manipulation. If the ammoniacal filtrate from the Bi(OH) 3 is of a blue color, that is sufficient evidence of the presence of Cu unless nickel was precipitated in the second group. In absence of a blue color a portion of the solution should be acidulated with acetic acid and then to this solu- tion a few drops of potassium ferrocyanide, K 4 Fe(CN) c , should be added. A brick-red precipitate is evidence of copper. Or to the acidulated solu- tion a bright nail or piece of iron wire may be added, obtaining a film of metallic copper. If sufficient copper be present to give a blue color to the solution, before testing for cadmium a solution of KCN should be added until the blue color disappears. Then the addition of H 2 S will give a yellow precipitate for cadmium. 103. Notes. 1. The precipitate of the brick-red Cu 2 Fe(CN) 6 is a much more delicate test for copper than the blue color to the ammoniacal solution (77, 66). Cd gives a white precipitate, insoluble in the acid. 2. The student should not forget that in the regular course of analysis a trace of copper may be lost by the solubility of the sulphide in (NH 4 ),S X . If mercury has been shown to be absent, the sulphides of the tin group (second group, division A) should be dissolved by the addition of a fixed alkali sulphide (71, 6e), K 2 S or Na,S , which does not dissolve CuS. In case mercury be present, the presence or absence of small amounts of copper must be deter- mined by the usual reactions for copper upon the original solution, having due regard for the possible interference of metals which the analysis has shown to be present. 3. Potassium cyanide, KCN , in excess changes cupric salts to the soluble double salt of cuprous cyanide and potassium cyanide, K 3 Cu(CN) 4 , which is colorless and not precipitated by sulphides. With cadmium salts the soluble double cyanide. K,Cd(CN) 4 , is formed, which is decomposed by sulphides forming CdS , yellow. 4- If preferred, the sulphides of Cu and Cd may be precipitated from the ammoniacal solution by H 2 S and then the black CuS dissolved with KCN , leaving a yellow precipitate of CdS . 104, 6. RUTHENIUM. 131 5. Copper and cadmium may be separated from each other by reduction of the copper (from the ammoniacal solution acidulated with HC1) with SnCL (77, 10): 2CuCL + SnCl, = 2CuCl -f SnCl 4 , and its precipitation with milk of sulphur (77, Ge), forming- Cu.S , removal of the tin with NH 4 OH and the precipitation of the cadmium with H 2 S . 6. From Ihe solutions of copper and cadmium acidulated with HC1 , a hot solution of Na 2 S 2 O 3 precipitates the copper as Cu,S (77, Ge), while the cadmium remains in solution. From this solution the cadmium is detected as the sul- phide by neutralization with NH.OH and precipitation with H,S or (NH 4 ),S. 7. The ammoniacal solution of Cu and Cd may be precipitated with H,S , and the resulting- sulphides, after filtering and washing-, boiled with hot dilute H 2 SO 4 (one of acid to five of water). In this solution the CuS (77, 5c) is unattacked while the CdS is dissolved. The filtrate upon dilution with water gives the yellow CdS with H 2 S or (NH 4 ) 2 S (78, 6e). BARER METALS OF THE TIN AND COPPER GROUP. (Second Group.) Ruthenium, Rhodium, Palladium, Iridium, Osmium, Tungsten, German- ium, Tellurium, Selenium. 104. Ruthenium. Ru = 101.7. Valence two to eight. 1. Properties. Specific gravity, 11.0 to 11.4 (Deville andDebray, C.r., 1876, 83, 926). Melting-point 2450 ? (Cir. B. S. t 35, 1915). Next to osmium it is the most difficultly fusible of all the platinum metals. A black powder or a grayish-white crystalline brittle metal. 2. Occurrence. In small quantities in platinum ores. 3. Preparation. Ignite the Pt residues in a stream of chlorine in presence of NaCl . Dissolve the fused mass in H 2 O , add KN0 2 , neutralize with Na 2 C0 3 , evaporate to dryness and extract the double nitrites with absolute alcohol (separation from rhodium). Add water to the solution, distill off the alcohol, add HC1 and obtain a red solution of potassium ruthenium chloride. This is changed to the double ammonium salt and then precipitated with HgCl 2 , which upon recrystallization and ignition gives pure Ru (Gibbs, Am. 8., 1862, (2), 34, 349 and 355). 4. Oxides and Hydroxides. The hydroxides, Ru(OH) 2 , Ru(OH) 3 , and Ru(OH) 4 , are precipitated from the respective chlorides by KOH . They are dark brown to black. Perruthenic anhydride or acid, RuO 4 , is a golden yellow crystalline powder, volatile even at ordinary temperatures. It has a peculiar odor, somewhat like ozone, is sparingly soluble in water, melts at 50 and boils at a little over 100 (Deville and Debray, B., 1875, 8, 339). It is pre- pared by heating K 2 RuCl 5 with KOH into which a current of chlorine is passed or by distillation of a Ru salt with KC10 3 and HC1 . The vapor is yellow and is strongly irritating to the membrane of the throat. 5. Solubilities. Ru is soluble with difficulty in nitrohydrochloric acid, in- soluble by fusion witk KHSO 4 , but is soluble by fusion with KOH , especially in presence of KNO 3 . Soluble in chlorine, forming' a mixture of RuCl 2 , RuCl 3 , and RuCl 4 . The double nitrites are soluble in water and alcohol (sepa- ration from rhodium). 6. Reactions. The alkalis precipitate from ruthenic chloride the dark yellow hydroxide, soluble in acids, insoluble in the fixed alkalis, soluble in NH 4 OH with a greenish-brown color. H 2 S precipitates slowly the black sulphide (formed at onc by (NH 4 ) 2 S), the solution becoming blue. The sulphide is insoluble in alkali ^sulphides. KI gives with hot solutions a black precipitate 132 RHODIUM. 104, 7. of ruthenic iodide. KCNS forms, after some time in the cold, a red coloration, which upon heating* assumes a beautiful violet color (characteristic). The double nitrites are soluble, and if to the solution (NH^nS be added, a char- actetristic crimson red liquid is obtained. Upon standing the solution becomes brown, or a brown precipitate is caused by excess of the (NH 4 ) 2 S . 7. Ignition. If Ru0 4 be heated to a dull-red heat the violet-blue dioxide is formed (Debray and Joly, C. r., 1888, 106, 328). 8. Detection. By oxidation and distillation as Ru0 4 . 9. Estimation. Keduced to the metal and weighed as such. 10. Oxidation. Ru0 4 heated with HC1 forms RuCl 3 , evolving chlorine. B, U 8 x is changed to Bu0 4 by distilling with KC10 3 and HC1 . Zn reduces Ru solutions to the metal, with an indigo-blue color during transition from Ruiv to Run . 105. Rhodium. Rh = 102.9 . Valence two, three and four. 1. Properties. Specific grarity, 12.1 (Deville and Debray, C. r., 1874, 78, 1782). Melting point, 1950 (Cir. B. S., 35, 1915). It is a white metal, nearly as ductile and malleable as Ag. The metal precipitated by alcohol or formic acid appears as a black spongy mass (Wilm, B., 1881, 14, 629). 2. Occurrence. Found in platinum ores. 3. Preparation. Fusion of the Pt residues with Pb , digestion with HN0 3 and then Cl , converting the Rh into the chloride, from which solution it is precipitated as the double ammonium chloride by fractional precipitation. See Gibbs (/. pr., 1865, 94, 10) and Wilm (5., 1883, 16, 3033). 4. Oxides and Hydroxides. Rh(OH) 3 is precipitated from a solution of sodium rhodium chloride by an excess of KOH . It is a black gelatinous pre- cipitate, forming the oxide upon ignition. Rhodium fused with KOH and KNO 3 gives Rh0 2 , a brown powder, insoluble in acids or alkalis. 5. Solubilities. The pure metal or the alloy with Au or Ag is almost in- soluble in acids; alloyed with Bi , Pb , Cu or Pt , it is soluble in HNO 3 (Deville and Debray, I.e.). Attacked by chlorine the most easily of all the Pt metals. The precipitated metal, a gray powder, is soluble in HC1 in presence of air to a cherry-red color. 6. Reactions. Alkali hydroxides and carbonates precipitate solutions of Rh salts as Rh(OH) 3 , yellow, insoluble in acids, soluble in excess of NH 4 OH , forming a rhodium ammonium base, precipitated by HC1 as a bright yellow crystalline salt, chloro-purpureo-rhodium chloride, Rh(NH 3 ) 5 Cl 3 . Alkali nitrites precipitate alcoholic solutions of rhodium chloride as alkali-rhodium nitrite (Gibbs, Am. S., 1862, (2), 34, 341) (separation from ruthenium). From a hot solution of Rh salt, H 2 S precipitates the sulphide, insoluble in the alkali sulphides: the sulphide precipitated from the cold solution is soluble in alkali sulphides. KI precipitates from hot solutions a black-brown rhodium iodide. 7. Ignition. When the metal or its compounds are repeatedly fused with HPO 3 or KHS0 4 , the corresponding Rh salts are formed. The mass fused with KHS0 4 is soluble in water to a yellow color, turning red with HC1 . 8. Detection. By ignition as given above. Also to the concentrated neutral solution add fresh NaCIO solution. To the yellow precipitate add a small amount of HC 2 H 3 2 and shake till an orange-yellow solution is obtained. After a short time the solution becomes colorless, then a gray precipitate separates out and the solution assumes a sky-blue color (Demarcay, C. r., 1885, 101, 951). 9. Estimation. It is reduced to the metal and weighed as such. 10. Oxidation. Solutions of rhodium salts are reduced to the metal by Zn . All Rh compounds are reduced to the metal by heating in a current of hydrogen. 106, 6. PALLADIUM. 133 10. Oxidation. Solutions of rhodium salts are reduced to the metal by Zn. All Rh compounds are reduced to the metal by heating in a current of hydrogen. 106. Palladium. Pd = 100.7. Valence two and four. 1. Properties. Specific gravity, 11.4 (Deville and Debray, C. r., 1857, 44, 1101)' Melting point 1549 (Cir. B. S., 36, 1915). It conducts electricity about one-eighth as well as silver (Matthiessen, Pogg., 1858, 103, 428). Palladium has about the color and lustre of silver. The metal when only slightly heated assumes a rainbow tint from green to violet. Because of its general properties, it is to be classed with the platinum metals, yet in its reaction with acids it is markedly different. In the air at ordinary temperature it is but slightly tarnished, but at a red heat it becomes covered with a coating of the oxide. The finely divided metal, palladium sponge, absorbs many times its volume of hydrogen, retaining the most of the hydrogen even at 100. At a high heat the hydrogen is all driven off. It is much used in gas analysis for the separation of hydrogen from other gases (Hempel, B., 1879, 12, 636, 1006). Also used for scale graduations of the best scientific instruments. 2. Occurrence. It is a never-failing element in the platinum ores, native or alloyed with Pt , Au or Ag . 3. Preparation. The obtaining of pure palladium involves its separation from the other platinum metals, i. e., platinum, iridium, osmium, rhodium and ruthenium. The student is referred to the various works on metallurgy; also to the following: Bunsen, A., 1868, 146, 265; Wilm, B., 1885, 18, 2536; and Mylius and Forster, B., 1892, 25, 665. 4. Oxides and Hydroxides. Palladium monoxide, PdO , is the most stable of the oxides of Pd . It is formed by the gentle ignition of Pd(N0 3 ) 2 or the precipitation of PdCl 2 with Na 2 CO 3 , forming Pd(OH) 2 , and then igniting. Palladic oxide, PdO 2 , when gently ignited loses half its oxygen, becoming PdO . 5. Solubilities. a. Metal. It is slowly dissolved by boiling with HC1 or H 2 SO 4 ; HN0 3 dissolves it, even in the cold, forming Pd(NO 3 ) a . It is more readily soluble in nitrohydrochloric acid, forming PdCl 4 . It is not at all attacked by H 2 S . An alcoholic solution of iodine blackens it, and when fused with KHSO 4 it becomes the sulphate (distinction from platinum). 6. Oxides. Pd0 2 is soluble in HC1 with evolution of Cl , forming PdCl 2 . Pd(OH) 2 is readily soluble in acids forming palladous salts, c. Suits. Palladic chloride, PdCl 4 , the most stable of the palladic salts is decomposed by boiling with water or by much dilution with cold water, forming PdCl 2 . It forms double chlorides with other metals, as calcium palladic chloride, CaPdCl , which for the most part are stable, and soluble in water and alcohol. Potassium palladic chloride, K 2 PdCl e , is but sparingly soluble in water, insoluble in alcohol; par- tially decomposed by both solvents. Palladous chloride is readily soluble in water with a brownish-red color; with metallic chlorides, it forms double chlorides, as potassium-palladous chloride, K.PdCl 4 , all of which are soluble in water. Palladous iodide is insoluble in water, alcohol or ether; insoluble in dilute hydrochloric acid or hydrioclic acid; slightly soluble by iodides and by chlorides". Palladous nitrate, Pd(N0 3 ) 2 , is soluble in water with free nitric acid; the solution being decomposed by dilu- tion, evaporation, or by standing, with precipitation of variable basic nitrates. Palladous sulphate, PdSO 4 , dissolves in water, but decomposes in solution on standing. 6. Reactions. Palladous chloride is precipitated by potassium hydroxide or sodium hydroxide; as brown basic salt or as brown palladous hydroxide, Pd(OH), , soluble in excess of the hot reagents. Ammonium hydroxide gives a flesh-red precipitate of palladio-diammonium chloride, (NH 3 ) 2 PdCl 2 . The flesh- red precipitate is soluble in excess of the ammonia, and from this solution reprecipitated by hydrochloric acid, with a yellow color. The fixed alkali carbonates precipitate the hydroxide; ammonium carbonate acts like the hydroxide. Potassium cyanide precipitates palladous cyanide, Pd(CN) 2 , white, soluble in excess of the reagent. Phosphates give a brown precipitate. Hydrosulphuric acid and sulphides precipitate the dark-brown palladous sulphide, PdS , insoluble in the ammonium sulphides, soluble in nitrohydro- chloric acid. Potassium iodide precipitates palladous iodide, PdI 2 , black, visible in 500,000 parts of the solution, with the slight solubilities stated in 5c, an important separation of iodine from bromine. In very dilute solutions, only a iRiDWM. 106, 7. color is produced, or the precipitate separates after warming. At a red heat, the precipitate is decomposed. Palladous nitrate gives most of the above reactions; no precipitate with ammonia, and a less complete precipitate with iodides. 7. Ignition. Nearly all the palladium compounds are reduced by heat, before the blow-pipe, to a " sponge." If this be held in the inner flame of an alcohol lamp, it absorbs carbon at a heat below redness; if then removed from the flame, it glows vividly in the air, till the carbon is all burnt away (distinction from platinum). 8. Detection. Palladium is precipitated with the second group metals by H 2 S, not dissolved by (NH 4 ) 2 S X (separation from the tin group). It is distinguished from mercury by its precipitation as a cyanide with mercuric cyanide. It is precipitated from quite dilute solutions by KI (distinction from Bi and Cd) ; an excess of the KI dissolves the black palladous iodide, PdI 2 , to a dark brown solution. KCNS does not precipitate palladium salts, not even after the addi- tion of SO 2 (separation from Cu). The addition of H,SO 4 and alcohol separates lead from palladium. The presence of the metal should be further confirmed by reduction and study of the properties of the " sponge " obtained. 9. Estimation. (Jf) As metallic palladium, to which state it is reduced by mercuric cyanide or potassium formate, and ignition, first in the air and then in hydrogen gas. (2) As K 2 PdCl G . Evaporate the solution of palladic chloride with potassium chloride and nitric acid to dryness, and treat the mass when cold with alcohol, in which the double salt is insoluble. Collect on a weighed filter, dry at 100, and weigh. 10. Oxidation. Palladium is reduced as a dark-colored precipitate, from all compounds in solution, by sulphurous acid, stannous chloride, phosphorus, and all the metals which precipitate silver (59, 10). Ferrous sulphate reduces palladium from its nitrate, not from its chloride. Alcohol, at boiling heat, reduces it; oxalic acid does not (distinction from gold 73, 6&). 107. Iridium. Ir = 193.1 . Usual valence three and four. 1. Properties. Specific gravity, 22.421 (Deville and Debray, C. r., 1875, 81, 839). Melting point 2350 ? (Cir. B. S., 35, 1915). When reduced by hydrogen it is a gray powder, which by pressing and igniting at a white heat changes to a metallic mass capable of takir.g a polish. It is used mostly as an alloy with platinum, forming a very hard, durable material for standard weights and measures. A platinum-iridium dish containing 25 to 30 per cent iridium is not attacked by nitrohydrochloric acid. 2. Occurrence. Found in platinum ores, usually as an alloy with platinum or osmium. 3. Preparation. The platinum residues are mixed with Pb and PbO and heated at a red heat for one-half hour, then treated with acids. The residue contains the iridium as osmium-iridium or platinum-iridium with other plat- inum metals. This residue is mixed with NaCl in a glass tube and heated to a red heat in a current of chlorine. Much of the osmium passes over as the volatile perosmic acid, and is condensed. The double sodium chlorides of Ir , Os , Rh , Pt , Pd and Ru are dissolved in water, filtered and, when boiling hot, decomposed by H 2 S . The iridium is reduced from the tetrad to the triad, but is not precipitated until after all the other metals. By stopping the current of H 2 S just as the brown iridium sulphide begins to form, a complete separation can be made by filtration. By recrystallization the pure sodium double salt, 6NaC1.2lrCl 3 + 24H 2 , is obtained, which is changed to the tetrad ammonium double salt, (NH 4 ) 2 IrCl 6 , by the addition of NH 4 C1 and oxidation with chlorine (Wohler, Pogg., 1834, 31, 161). This upon ignition gives the pure metal as iridium sponge. Or, the double sodium salt is ignited with sodium carbonate, exhausted with water and reduced by ignition in a current of hydrogen, leav- ing the metal as a fine gray powder. (See also 106, 3). 4. Oxides and Hydroxides. Iridium forms two series of oxides and hydrox- ides, the metal acting as a triad and tetrad respectively. IrO 3 is formed by 108, 5. OSMIUM. 135 igniting the metal in the air at a bright red heat, henoe the scaling of platinum dishes which contain iridium. The hydroxide, Ir(OH) 4 , is formed by boiling- a solution of the trichloride, IrCl s , in a fixed alkali hydroxide or' carbonate. Careful addition of KOH to IrCl 3 in a vessel full of liquid and closed to exclude air gives Ir(OH) 3 , easily oxidi/ed to Ir(OH) 4 (Clans, J. pr., 1846, 39, 104). 5. Solubilities. Freshly precipitated iridium may be dissolved in nitrohydro- chloric acid. The ignited metal is insoluble in all acids. Its proper solvent is chlorine. Iridium trichloride, IrCl 3 , is soluble in water and forms with the alkali chlorides double chlorides, soluble in water, insoluble in alcohol. The tetrachloride with sodium chloride, Na 2 IrCl , is formed when the platinum residues mixed with NaCl are heated in a current of chlorine. It is soluble in water. The corresponding ammonium salt may be formed from the sodium salt by precipitation from the concentrated solution with NH 4 C1 , a reddish- brown precipitate, soluble in 20 parts of water (Vauquelin, A. Ch., 1806, 59, 150 and 225). The potassium double salt is sparingly soluble in water. 6. Reactions. Fixed alkali hydroxides or carbonates precipitate from toil- ing solutions of iridium chloride, IrCl 3 or IrCl 4 , iridiuju hydroxide, Ir(OH) 4 , dark blue, insoluble in all acids except HC1 . Potassium nitrite added to a hot solution of iridium salts gives, first a yellow color and finally a yellow precipi- tate, insoluble in water or acids. Hydrogen sulphide reduces IrCl 4 to IrCl 3 , and then precipitates the trisulphide, Ir 2 S, , brown, soluble in alkali sulphides. 7. Ignition. When iridium is fused with potassium acid sulphate it is oxid- ized, but does not go into solution (difference from rhodium, 105, 7). Ignition on charcoal reduces all iridium compounds to the metal. Fusion in the air with sodium hydroxide or with sodium nitrate causes oxidation of the metal, the iridiiim oxide formed being partially soluble in the fixed alkali. 8. Detection. See 3 and 6. 9. Estimation. It is converted into the oxide by igniting with KN0 3 and then reduced by ignition in an atmosphere of hydrogen. 10. Oxidation. Formic acid (from hot solution), zinc and H 2 S0 4 or HC1 reduce iridium compounds to the metal. SnCl 2 , FeSO 4 and H 2 C 2 O 4 reduce tetrad iridium to triad, but do not further reduce (separation from gold, 73, 60, h and &). 108. Osmium. Os = 190.9. Valence two to eight. 1. Properties. Specific gravity, 22,477, the heaviest of all bodies (Deville and Debray, C. r., 1876, 82, 1076). Melting point 2700? (Cir. B. S., 36, 1915). In the absence of air it may be heated above the vaporization point of Pt without melting or oxidizing. In presence of air, when heated a little above the melting point of Zn, it burns to the volatile poisonous perosmic acid, OsO 4 . In com- pact form it is very hard, cutting glass, and possesses a metallic lustre, with a bluish color resembling Zn. 2. Occurrence. Always present in the residues of the platinum ores, m com- bination with iridium. 3. Preparation. The iridium osmium alloy or other Os containing material is finely divided and distilled in a current of chlorine or with nitrohydrochloric acid, the osmium passing into a receiver containing KOH. By repeated addi- tions of HNO 3 and further distillation, the osmium may all be driven into the receiver. The distillate is treated with HC1 and Hg and the amalgam ignited in a current of hydrogen (Berzelius, Pogg., 1829, 15, 208). 4. Oxides. Osmium forms five different oxides, OsO , Os 2 3 , OsO, , OsO., , Os0 4 . The first three are bases, the salts of which have been but little studied; OsO 3 forms salts with bases, and OsO 4 acts rather as an indifferent peroxide. Perosmic acid, OsO 4 , exists as white glistening needles, melting under 100, sparingly soluble in water, its solution having a very penetrating odor, resembling that of chlorine. The fumes of the acid are very poisonous, and cause inflammation of the eyes. H 2 S is recommended as an antidote (Clauss, A., 1847, 63, 355). 5. Solubilities. The metal in compact condition is not at all attacked by any acid. The precipitated metal is slowly dissolved by nitrohydrochloric or funv 136 TUNGSTEN. 108, 6. ing nitric acid. By heating the metal in a current of chlorine a mixture of OsCl 2 and OsCl 4 is formed. They are both unstable. 6. Reactions. Perosmic acid, OsO 4 , when boiled with alkalis, is reduced to osmates, as K 2 OsO 4 . A solution of perosmic acid decolors indigo, oxidizes alcohol to aldehyde, and liberates iodine from potassium iodide. In the pres- ence of a strong mineral acid, H 2 S precipitates osmium sulphide, OsS 4 , brown- ish black (Claus, /. pr., 1860, 79, 28); insoluble in alkali hydroxides, carbonates or sulphides. 7. Ignition. Osmium when heated on a piece of platinum foil gives an in- tensely luminous flame of short duration. By holding the foil in the reducing flame and then again in the oxidizing flame, the luminosity may be repeated. If a mixture of the metal or of the sulphide and potassium chloride be heated in a current of chlorine, a double salt of potassium osmic chloride is formed, sparingly soluble in cold \vater, more readily in hot water. Alcohol precipitates it from its solutions as a red crystalline powder. 8. Detection. By the intensely luminous flame when ignited on a platinum foil; by oxidation and distillation as perosmic acid and identification by odor, action on indigo and on potassium iodide. 9. Estimation. It is weighed as the metal (see 3). 10. Oxidation. Os0 4 is reduced to Os0 2 by ferrous sulphate. Zn and many other metals in presence of strong acids precipitate the metal. The metal is also obtained from all osmium compounds by ignition in a current of hydrogen. 109. Tungsten (Wolframium). W = 184. Valence two to six. 1. Properties. Specific gravity, 18.71-18.74 (Z. Anorg, 1909, 1910). Melting point, 3000 (Cir. B. S., 35, 1915). A tin-white or steel-gray metal, brittle, harder than agate. That precipitated from acid solutions is a velvet-black powder. Non-magnetic. Stable in the air at ordinary temperature; burning at a high temperature, it decomposes steam at a red heat. Dry metallic tungsten powder, compressed into a bar in a hydraulic press, may, by repeated heatings and swaging or rolling, be converted into a ductile form, and drawn into wire (Met. and Cham. Eng., 3, XII, 1914). 2. Occurrence. Tungsten does not occur in nature in large amounts, nor is it widely disseminated. The most common tungsten minerals are scheelite, (CaWO 4 ), and wolframite (FeWO 4 and MnWO 4 , in variable proportions). It never occurs native. 3. Preparation. By reduction of WO 3 in H at a red heat (Zettnow, Pogg., 1860, 111, 16); by ignition of W0 3 and Na under NaCl . Tungstic acid of commerce is prepared by igniting for several hours: 100 parts Na 2 CO 3 , ignited; 150 parts finely ground wolframite; and 15 parts NaNO 3 . The cooled mass is exhausted with water and the filtrate poured into hot, moderately concentrated HC1 (Franz, J. pr., 1871, (2), 4, 238). 4. Oxides. WO, is obtained as a brown powder by decomposing WC1 4 with water (Roscoe, 1. c.). WO 3 is a lemon-3 r ellow, soft powder, insoluble in water or acids. It is formed by ignition of the metal, lower oxides or decomposable salts in the air. The blue tungsten oxides are compounds between WO 2 and W0 8 . 5. Solubilities. The metal is scarcely at all attacked by HC1 or H 2 S0 4 , slowly by HNO 3 or nitrohydrochloric, slowly soluble in alkalis. The halogens com- bine directly upon heating. WO 2 is readily soluble on heating with HC1 and H 2 SO 4 to a red color. It is also soluble in KOH with red color, evolving hydrogen. Both the acid and alkaline solutions deposit the blue oxide on standing (von der Pfordten, A., 1884, 222, 158). WO 3 is insoluble in water or acids, not even soluble in hot concentrated H 2 S0 4 . Soluble in KOH , K 2 CO, and NH 4 OH . In an atmosphere of C0 2 it reacts with the chlorides of Ca , Mg, Co, Ni and Fe (not with those of Pb , Ag , K and Na), e.g., MC1 2 + 2W0 3 = MW0 4 + WO 2 C1 2 . Heated with chlorine, W0 2 C1 2 is formed, and also WC1 4 , decomposed by water. S , H 2 S or HgS form WS 3 on heating with W0 3 . Soluble alkali tungstates are formed by fusion of the acid, WO 3 , with the alkali metal carbonates, more slowly by boiling with the carbonates. Acids form, from solutions of the alkali tungstates, a white precipitate of the hydrated acid turning yellow on boiling, insoluble in excess of the acids (dis- 111, 5. GERMANIUM. 137 tinction from MoO 3 ), soluble in NH 4 OH . Phosphoric acid changes tungstic acid to the metatungst ic acid, which is soluble in water arid not precipitated by other acids. Long 1 boiling 1 of the solution of metatungsl ic acid causes 1 lie- precipitation of tungstic acid. Fusion of W0 3 with KHSO, gives a compound of potassium tungstate and tungstic acid, uo1 readily soluble in water but very readily soluble in (NH 4 ),CO 3 (distinction from silica, 249, 5). 6. Reactions. Solutions of salts of Ba , Ca , Pb , Ag and Hg produce white precipitates with solutions of alkali tungstates. H 2 S precipitates WS ; . from acid solutions, the sulphide dissolving readily in (NH 4 ) 2 S , forming a thiotungstate (NH 4 ) 2 WS 4 . The tungstates, like the molybdates, form complex compounds with phosphoric acid, i.e., phosphomolybdates and phospholuiig- states, which react very similarly with ammonium salts and with organic bases (75, Grf). K 4 Fe(CN) gives with tungstates (in presence of acids) a deep brownish-red fluid, forming after some time a precipitate of the same color. Solution of taniiic acid gives a bro\vii color or precipitate. 7. Ignition. With NaPO 3 , W0 3 dissolves, on fusion, to a clear or yellowish bead in the oxidizing flame; in the reducing flame it has a blue color, changing to red on addition of FeSO 4 . Heated on charcoal in presence of Na 2 C0 3 with the blow-pipe, using the reducing flame, the metal is obtained. 8. Detection. If a tungstate be fused with Na 2 CO 3 , the mass warmed with water and the water then absorbed with strips of filter paper, the tungsten may be detected by moistening the strip with HC1 and warming, obtaining the yellow color of WO 3 ; and the blue color of a lower oxide by moistening with SnCl 2 and warming. (NH 4 ) 2 S does not color the paper, even after adding HC1 , but on warming a blue or green color is obtained. 9. Estimation. It is converted into WO 3 and weighed as such after ignition. 10. Oxidation. WO 3 gives with SnCl, , or Zn in presence of HC1 or H 2 SO 4 , a beautifid blue color, due to the formation of oxides between W0 2 and W0 3 , blue oxides of tungsten (delicate and characteristic). 111 Germanium. Ge = 72.5. Valence two and four. 1. Properties. Specific gravity, 5.469 at 20.4 (Winkler, /. pr., 1886, (2), 34, 177); melting point, 958 (Cir. B. S., 35, 1915). A gray-white crystalline metal. Fused under borax it gives a grayish-white regulus with a metallic lustre. It is stable in the air, volatilized at a high heat (Meyer, B., 1887, 20, 497), and is easily pulverized. It burns in oxygen to form germanic oxide, GeO 2 . 2. Occurrence. It is found in small quantity in argyrodite, a sulphide of silver and germanium, (3Ag 2 S.GeS 2 ) , a silver ore from Freiburg, Saxony; eux^ni e (msbate and titanate of yttrium, erbium, cerium and uranium) from Sweden (Kriiss, C. C., 1888, 75); also contains small amounts of germanium. 3. Preparation. It is formed by reduction of the oxide, GeO 2 , with H , C or Mg (Winkler, B., 1891, 24, 891) ; also by reduction of the sulphide in H . 4. Oxides. It forms two oxides, GeO and GeO 2 . To prepare pure Ge0 2 , the mineral argyrodite is pulverized and intimately mixed with equal weights of Na 2 C0 3 and S and heated to a good full ignition. The mass must be added carefully to prevent foaming. The fused mass is exhausted with H 2 O , the germanium going into solution as a thiosalt. With a decided excess of H 2 S0 4 , the sulphide is completety precipitated. The precipitate is now dissolved in KOH , the sulphides of A.g , Cu and Pb remaining undissolved. By adding to the KOH solution H 2 S0 4 not quite to neutralization, the As and Sb sulphides are precipitated on boiling, while the GeS remains in solution with some As 2 S 3 ; H 2 S is carefully added to the solution until the As 2 S 3 is all precipitated, then the filtrate is made strongly acid with H 2 SO 4 , and the solution evaporated till SO 3 fumes escape. The mass is dissolved in hot water, and upon cooling Ge0 2 crystallizes out (Winkler, I. c.). 5. Solubilities. Germanium is insoluble in HC1 , soluble in nitrohydrochloric acid as GeCl 4 , and oxidized with HNO 3 to Ge0 2 . Hot concentrated H 2 S0 4 evolves S0 2 and forms Ge(S0 4 ) 2 . Insoluble in KOH solution but dissolves with incandescence in fused KOH . It unites directly with 01 , Br and I (Winkler, I. c.). Germanic oxide, GeO 2 , is a white powder, very sparingly soluble in water or acids. Fused with fixed alkali hydroxides or carbonates it is converted into compounds soluble in water. GeCl 4 is a liquid, boiling at 84; 138 TELLURIUM. 112, 6. it is decomposed by water. If a solution of the oxide in excess of HC1 be evaporated to dryness the Ge is all volatilized. GeS 2 is soluble in 222 parts water, in alkali sulphides and hydroxides; insoluble in HC1 or H 2 SO 4 , which precipitate it from its solutions; soluble in nitrohydrochlorie acid with separa- tion of sulphur. Nitric oxide changes it to Ge0 2 with separation of sulphur. 6. Reactions. Germanium salts give almost no characteristic reactions with the various reagents. H 2 S precipitates germanic sulphide, GeS 2 , white, from solutions of the salts quite strongly acid. The sulphide is soluble in ammonium sulphide, forming a thio salt, thus placing Ge in division A of the second group. 7. Ignition. Heated before the blow-pipe in the reducing flame without an alkaline flux the metal is formed, and at the same time a Avhite coating of the oxide. It forms a colorless bead with borax. 8. Detection. In the mineral, argyrodite, by heating in an atmosphere of H 2 S or illuminating gas, an orange-yellow sublimate is obtained, which may be examined under the microscope and in the wet way (Haushofer, C. C., 1888, 867). 9. Estimation. It is converted into the sulphide, GeS 2 , and then heated with HNO 3 and weighed as GeO 2 . 10. Oxidation. Zn in acid solutions of Ge salts precipitates the metal as a dark brown slime. If GeS-, is heated in a current of H , GeS is at first formed with H 2 S, finally Ge. 112. Tellurium. Te = 127.5. Valence two, four and possibly six. 1. Properties. Specific gravity, 6.2445 (Berzelius, Pogg., 1834, 32, 1 and 577). Melting point, 452 (Cir. B. S., 35, 1915). Te is crystalline, silver white, brittle, stable in the air and in boiling water; heated in the air, it burns with a greenish flame. In its general properties and reactions it stands closely related to S and Se (2). 2. Occurrence. In few places and in small quantities in Germany, Mexico, Bolivia, United States and Japan. Some of the minerals are: tellurite, (TeO 2 ) ; tetradymite (Bi 2 (TeS) 3 ) ; ferrotellurite, (FeTeO.,) , sylvanite, (AgAuTe 4 ) ; calaverite, (AuTe 2 ). It also occurs native. n. Preparation. (1) Fusion with alkali carbonate and C , which converts it into a telluride, as Na 2 Te; then solution in (air free) water, the air being excluded as much as possible, and the filtrate precipitated by passing air through the solution. The Te is precipitated as a gray metallic powder, con- taining what Se may have been present. (2) Conversion into TeCl 4 by distilla- tion in a current of chlorine, decomposition of the chloride with water to H 2 TeO 3 and precipitation of the Te with KHSO 3 . (,?) From lead chamber scale by digestion with Na 2 C0 3 and KCN , forming KCNTe . The decanted solution is acidified with HN0 3 and the Te precipitated with H 2 S (Schimose, C. N., 1884, 49, 157). (4) For purification of the commercial Te , see Brauner (If., 1889, 10, 411) and Schimose (C. N., 1884, 49, 26, and 1885, 51, 199). 4. Oxides and Hydroxides. TeO is said to be formed by heating TeS0 3 in a vacuum above 180: TeSO., = TeO + S0 2 (Divers and Schimose, C. N., 1*8::, 47, 221). TeO 2 forms when Te is burned in the air, and when TeCl 4 is decomposed by boiling water. It is a white crystalline solid, sparingly soluble in H 2 O , more soluble in acids from which solutions water causes a white precipitate of Te0 2 or H 2 TeO., . H 2 TeO 3 is formed when a HNO 3 solution of Te is immediately poured into cold water, warming to 40 changes it to TeO 2 . H 2 TeO 4 is made by fusing TeO, with KN0 3 , treating the K 2 Te0 4 so *btained with soluble lead or barium salt and decomposing this salt with H 2 SO 4 or H 2 S , colorless crystals, insoluble in alcohol or ether-alcohol (separation from H 2 SO 4 ). It can be recrystallized from water and upon heating forms TeO 3 (Clarke, Am. 8., 1877, 114, 281; 1878, 116, 401). 5. Solubilities. Te is insoluble in HC1; HN0 3 and nitrohydrochloric acids oxidize it to H,TeO 4 ; in H 2 SO 4 it becomes H 2 Te0 3 with evolution of SO, (Hilger, A., 1874. 171, 211): soluble in warm concentrated solution of KCN, from which solution HC1 precipitates all the Te . H 2 TeO 3 is fairly soluble in water, red- c|ens moist litmus paper and easily decomposes into TeO 2 and H 2 . Acid solu- **!." rwifl * PHARMACY 113, *. SELENIUM. 139 tions of Te0 2 are precipitated upon addition of water or upon standing. Te0 2 and ILTeOj, form soluble alkali salts with :lis alkalis from which solutions of the other im-laJlic sails precipitate the respective teliUrites. H 2 Te0 4 is soluble in water, acids and alkalis; alkali carbonates form acid teimratoc, less soluble than the corresponding normal salts. Solutions of the alkali tellurates loiixi insoluble tellurates with soluble salts of the other metals, e. a. K,TeO 4- BaCL = BaTe0 4 + 2KC1 . G. Reactions. Tellurium is classed with second group metals because of its precipitation from solutions of tellurites and tellurates by H,S . The precipi- tate is not a sulphide, but is Te mixed with varying proportions of S , for CS 2 removes nearly all the sulphur (Becker, A., 1876, 180, 257). In appearance the precipitate of Te with H.>S very much resembles SnS , and is very soluble in (NH 4 ) a S. At a high temperature Te and H unite directly, forming H 2 Te (Brauner, M., 1889, 10, 446). H 2 Te is best prepared by heating together Te and Fe or Zn and decomposing these tellurides with HC1 (analogous to the corresponding reac- tions with sulphur, 257, 4). A colorless gas, odor similar to H 2 SJ , burns with a blue name, fairly soluble in water and is precipitated as Te from its solution by the oxj r gen of the air. H 2 Te precipitates solutions of metallic salts very similarly to H 2 S and H 2 Se . 7. Ignition. Te combines on ignition with most metals to form tellurides. TeO 3 ignited, decomposes into Te0 2 and O . All lower Te compounds ignited with KNO 3 give K 2 TeO 4 . All Te compounds give on charcoal with the blow- pipe a white powder, which colors the reduction flame green and disappears. Heated in an open glass tube, Te compounds give a sublimate of TeO 2 . which melts upon heating. Te compounds fused with KCN in a current of hydrogen form potassium tellurocyanate, KCNTe; soluble in water but precipitated by a current of air as Te (distinction and separation from Se). Heated with Na 2 C0 3 on charcoal Te compounds give Na,Te , which blackens silver with formation of Ag 2 Te . 8. Detection. By reduction to Te and solution in cold concentrated H 2 S0 4 to a purplish-red solution (characteristic). Separated from Se by fusion with KCN in a current of hydrogen and precipitation from the solution by a current of air. 9. Estimation. The Te compound is heated in a current of Cl , TeCl 4 being sublimed. This is decomposed by water to Te0 2 , which is reduced to Te by S0 2 and weighed as such after drying at 100 . 10. Oxidation. Hydrogen at a high temperature reduces Te compounds to H 2 Te . H,S reduces Te compounds to Te mixed with S . Fusion with KNO 3 oxidizes all Te compounds to K 2 TeO 4 . S0 2 reduces Te compounds to Te . SnCL and Zn in acid solutions give with Te compounds a black precipitate of Te . Te compounds warmed with dextrose in alkaline solution are reduced to Te . Tellurates boiled with HC1 evolve chlorine and are reduced to H 2 TeO 3 , which precipitates as Te0 2 on adding water if too much HC1 be not present (distinction from Se). 113. Selenium. Se = 79.2 . Valence two and four, possibly six. 1. Properties. Specific gravity, red, cryst., 4.47; gray metallic 4.80; (/. phys. Chem. 4, 491; 1900). Melting point 217-220 (Cir. B. S., 35, 1915). The molten Se does not become completely solid until cooled to 50. Selenium with tellurium is closely related to sulphur, and like sulphur exists in amorphous forms (266, 1). The precipitated Se is red. The brown or brown-black powder obtained by quickly cooling from the molten state is insoluble in 82. Boiling point, 676 to 683 (Carnelley and Williams, C. N., 1879, 39, 286). 2. Occurrence. In no place abundantly; never native. It is found in com- bination with minerals in the Hartz Mountains, Sweden, Argentine Republic and Mexico (Billandot, C. N., 1882, 46, 60). It occurs in very small quantities with some sulphides of Fe , Cu and Zn . 140 SELENIUM. 113, 3. 3. Preparation. In the lead chambers of the H 2 S0 4 works it is found as a red deposit with some S , As 2 O 3 , Sb,O 3 , PbSO 4 , etc. The scale is washed with water and digested with KCN solution at 80 to 100, until the red color entirely disappears. The filtrate is then treated with HC1 , which precipitates the e . It is further purified by oxidation to SeO 9 . sublimed and then reduced with S0 2 (Nilson, B., 1874, 7, 1719). 4. Oxides and Hydroxides. H 2 SeO 3 is prepared by oxidizing Se with HNO 3 , or nitrohydrochloric acid. H,Se0 3 evaporated to dryness gives H 2 O and Se0 2 , crystalline. SeO, is also formed by burning Se in air or oxygen; it has an odor similar to decaying radish. It sublimes at 250-280 as a yellow vapor, condensing to white needles on cooling. SeO 3 is not known. H.Se0 4 , pure, is a white crystalline mass, melting at 58. H 2 Se0 4 .H,0 is crystalline at 38, and if recrystallized melts at 25. The selenic acid usually obtained is a thick oily liquid, resembling H 2 S0 4 and containing about 95 per cent H 2 Se0 4 . It is obtained by fusing Se or SeO 2 with KN0 3 and precipitation of the K,SeO 4 with soluble salts of Ba , Pb , Ca or Cu and decomposing the washed precipitates, suspended in water, with H 2 S0 4 or H 2 S . 5. Solubilities. Se dissolves in cold concentrated H 2 S0 4 to a green colored solution without oxidation (dilution with water precipitates the Se) ; if the solution be warmed SO, is evolved and the green color disappears (dilution with water gives precipitate), the Se being oxidized to SeO, . HNO 3 and nitro- hydrochloric acid oxidize it to SeO 2 . Selenous oxide, SeO 2 , is soluble in water in all proportions, forming H 2 Se0 3 . The selenites and selenates of the alkaline earths are insoluble and may be formed by adding a solution of the metal to an alkali selenite or selenate, e. g., Na 2 Se0 3 + BaCl 2 = BaSe0 3 + 2NaCl . Many of the selenites are soluble in excess of H 2 Se0 3 . Selenates are less stable than selenites. BaSeO 4 is soluble in HC1 (distinction and separation from BaS0 4 ) and upon long-continued boiling is reduced to BaSeO, . 6. Reactions. Selenous acid precipitates with H 2 S a mixture of Se and S , lemon yellow, bright red upon heating (Divers and Shimose, C. N., 1885, 51, 199). This mixture is soluble in (NH 4 ) 2 S , hence in qualitative analysis Se is classed among the metals of division A, second group, while because of its general properties it belongs with sulphur. "When Se and H are heated to- gether they begin to combine directly at 250, forming H_,Se (Ditte, C. r,, 1872, 74, 980); which in practically all its reactions is similar to H^S . H.,Se is also formed by treating K 2 Se , FeSe , etc., with dilute HC1 or H 3 SO; HN0 3 gives H,SeO 3 with selenides. H 2 Se is a colorless gas, odor similar to HoS but more penetrating. It is more poisonous than H 2 S , burns when ignited, combines slowly but completely with Hg , evolving hydrogen. 100 cc. of water dissolves 331 cc. of the gas at 13, the solution reacting acid and depositing red flakes of Se on standing. It precipitates the selenides of the metals having almost the same solubilities as the corresponding sulphides (von Reeb, J. Pharm., 1869, (4), 9 173). With soluble sulphites H 2 Se gives a precipitate of a mixture of Se and S'. 7. Ignition. When Se or compounds of Se are fused with KCN in a current of hydrogen, potassium selenocyanate, KCNSe , is formed. Long boiling with HC1 separates the Se , but this does not take place on exposure of the solution to the air (separation from tellurium). Selenium compounds heated on char- coal with Na 2 C0 3 are changed to Na 2 Se , which yields a black stain with Ag and H 2 Se with dilute acids. 8. Detection. If in solution as selenites it is precipitated with H 2 S (soluble in (NH 4 ) 2 S); oxidized to SeO, and obtained as the white needles by sublima- tion, and reduced from its solution in water to the red Se by S0 2 . If present as selenides, decomposed by HC1 or H 2 SO 4 , forming H 2 Se , which is conducted into water and the Se precipitated by passing air or oxygen through the solu- tion. 9. Estimation. Oxidized to selenic acid and precipitated as BaSe0 4 and weighed as such. If BaS0 4 be present the precipitate is reduced in H , and the resulting BaSeO., separated by solution in HC1 . Selenides are heated in a current of chlorine in a hard glass tube, being converted into SeCl 4 , which vaporizes and is decomposed in water; continued chlorination of the water solution forms H 2 SeO 4 . 10. Oxidation. Se is oxidized to SeO 2 by HN0 8 , nitrohydrochloric acid, 114. THE IRON AND ZINC GROUPS. 141 H 2 SO 4 hot concentrated, by heating in air or oxygen, etc. H 2 SeO 3 is oxidized to H 2 SeO 4 by continued chlorination, and by fusion with KN0 3 . H 2 SeO 4 is reduced to H 2 Se0 3 by boiling- with HC1 . SO, reduces selenous compounds to the red Se , even in H 2 SO 4 solutions (distinction from tellurium) (Keller, J. Am. Soc., 1900, 22, 241)." H 2 S forms a precipitate of Se mixed with S . SnCl 2 precipitates Se from HC1 or H 2 SO 4 solutions of selenous compounds. THE IRON AND ZINC GROUPS (THIRD AND FOURTH GROUPS). 114. The Metals of the Earths and the more Electro-Positive of the Heavy Metals. Aluminum.,, Al = 27.1 Chromium Cr = 52.0 Iron Fe '= 55.84 Cobalt Co = 58.97 Nickel Ni = 58.68 Manganese Mn = 54.93 Zinc Zn = 65.37 Cerium Ce = 140.25 Columbium Cb = 93.5 Erbium E = 167.7 Gallium Ga = 69.9 Glucinum Gl = 9.1 Indium In = 114.8 Lanthanum La = 139.0 Neodymium Nd = 144.3 Praseodymium Pr = 140.9 Samarium Sa = 150.4 Scandium Sc = 44.1 Tantalum Ta = 181.5 Terbium Tb = 159.2 Thallium .Tl = 204.00 Thorium Th = 232.4 Titanium Ti = 48.10 Uranium U = 238.2 Vanadium V = 51 .0 Ytterbium Yb = 173.5 Yttrium Y = 188.7 Zirconium Zr = 90.6 115. The metals above named gradually oxidize at their surfaces in the air, and their oxides are not decomposed by heat alone. Zinc, iron, cobalt, nickel, and, with more difficulty, manganese, chromium, and most of the other metals of the groups, are reduced from their oxides by igni- tion at white heat with charcoal. They are all reduced from oxides by the alkali metals. Iron is gradually changed from ferrous to ferric combinations by contact with the air. Chromium and manganese are oxidized from bases to acid radicals by ignition with an active supply of oxygen in presence of alkalis; these acid radicals acting as strong oxidizing agents (8, 9). 116. The oxides and hydroxides of these metals are insoluble in water and they are precipitated from all their salts by alkalis. In the case of zinc, the precipitate redissolves in all the alkalis; the aluminum hydroxide redissolves in the fixed alkalis, but very slightly in ammonium hydroxide; the precipitate of chromium redissolves in cold solution of fixed alkalis, precipitating again on dilution and boiling; the hydroxides of cobalt and nickel dissolve in ammonium hydroxide. The oxide of chromium after ignition is insoluble in acids; the oxides of aluminum and iron are soluble with difficulty. The presence of tartaric acid, citric acid, sugar, and some other organic substances, prevents the precipitation of bases of these groups by alkalis. 142 THE IRON AND ZINC GROUPS. 117. 117. Ammonium salts, as NH 4 C1 , dissolve moderate quantities of the hydroxides of manganese, zinc, cobalt, nickel, and ferrous hydroxide; but, so far from, dissolving the hydroxide of aluminum, they lessen its slight solubility in ammonium hydroxide. 118. It thus appears that ammonium hydroxide, with ammonium chloride, the latter necessary on account of magnesium (189, 6a), man- ganese (134, 6a), and aluminum, will fully precipitate only aluminum, chromium, and iron of the important metals above named. These metals therefore constitute the THIRD GROUP (127), and the re- agent of this group is AMMONIUM HYDROXIDE in the presence of AMMONIUM CHLORIDE. Since aluminum, chromium, and iron are precipitated by ammonium hydroxide in the presence of ammonium chloride (Fe" by its previous oxidation with HN0 3 is present as Fe'") constituting the THIRD GROUP; the remaining of the most important metals cobalt, nickel, manganese, and zinc constitute the FOURTH GROUP (137). They are precipitated by the group reagent, AMMON- IUM SULPHIDE or HYDROSULPHURIC ACID in an AMMONIACAL SOLUTION. Some chemists do not make this classification of these metals, but precipitate them all as one group with ammonium sulphide (144), from neutral or ammoniacal solutions. The sulphides of Fe, Co, Ni , Mn , and Zn are not formed in presence of dilute acids, which acids keep them in solution during the second group precipitation; but are insoluble .in water, which enables them to be precipitated by alkali sulphides, and separated from the fifth and sixth groups. The other two metals, Al and Cr, do not form sulphides, in the wet way, but are precipitated as hy- droxides by the alkali sulphides. 119. Hydrosulphuric acid scarcely precipitates the metals of these groups, unless it be from some of their acetates (135, 6e), owing to the solubility of the sulphides in the acids, which would be set free in their formation. Thus, this change cannot occur FeCl 2 + H 2 S = FeS -|- 2HC1 because the two products would decompose each other. Hydrogen sulphide does not precipitate the metals of these groups in acid solution unless the acid is very weak (acetic acid 135, 6e). The hydrogen ions of strong acids, which are largely dissociated, reduce the concentration of the sulphur ion of the hydrogen sulphide below the point where it can pre- cipitate the sulphides of these metals. For the same reason these sul- phides are dissolved by strong acids and the reaction FeCL + H 2 S <=> FeS -f 2HC1 cannot proceed from right to left. As acetic acid and other weak acids are only slightly dissociated, the concentration of the hydrogen ion is very nmcli less and the decrease in the concentration of the sulphur ion of the hydrogen sulphide is slight. The soluble sulphides axe dissociate^ to a muqh greater extent giving a con- THE IRON AND ZINC GROUPS. 143 centration of the sulphur ion sufficient to precipitate the sulphides of these metals. (See 45.) Therefore when ii is desired i<> precipitate the metals as sulphides, neutralized hydrosulphuric acid an alkali suljihitle is used in neutral or alkaline solution; or, what is equivalent, hydrosulphuric acid gas is passed into the strongly ammonifical xolnlion. 120. As most of the normal chemically salts of heavy metals are hydro- lyzed, in water, giving free acids, so that their solutions have an acid reaction to test-paper, we can onlv assure ourselves of the requisite neutrality by adding sulTieient ammonium hydroxide, which, itself precipitates the larger number of the bases, as we have just seen (116). But the resulting. precipitate of hydroxide, as Fe(OH) 2 , is immediately changed to sul- phide, FeS , by subsequent addition of ammonium sulphide; as the student may observe, by the change in the color of the precipitate. Ferric and manganic salts are reduced to ferrous and mauganous salts, by hydrosulphuric acid, in solution, with a precipitation of sulphur, and the corresponding reaction occurs with chromates. 121. Soluble carbonates precipitate all the metals of these groups, in accordance with the general statement for bases not alkali (205, 6a). With aluminum and chromium, the precipitates dissolve sparingly in ex- cess of potassium or sodium carbonate; with Co, Ni and Zn , the precipitate dissolves in excess of (NH 4 ) 2 C0 3 . In the case of ferrous and manganous salts, the precipitates are normal carbonates; with zinc, cobalt, and nickel salts, they are basic carbonates; while with ferric, aluminum, and chrom- ium salts, the precipitates are hydroxides. Barium carbonate precipitates Al , Cr'" and Fe'", which, in the cold and from salts not sulphates, is a separation from the fourth group metals. 122. Soluble phosphates precipitate these as they do other non-alkali bases. The acid solutions of phosphates of the metals of the third and fourth groups are precipitated by neutralization. Phosphates of Co , Ni , and Zn are redissolved by excess of NH 4 OH , and those of Al , Cr , and Zn by excess of the fixed alkalis. The recently precipitated phosphates of all the metals of these groups which form sulphides, are transformed to sul- phides by ammonium sulphide, due to the fact that the sulphide is less soluble than the phosphate: FeHP0 4 -f (NH 4 ) 2 S = FeS + (NH 4 ) 2 HP0 4 . Hsnce, the only phosphates which may occur in a sulphide precipitate are those of Al , Cr , Ba , Sr , Ca , and Mg . 123. The rnetals of the third and fourth groups are not easily reduced from their compounds to the metallic state by ignition before the blow- pipe, even on charcoal, except zinc, which then vaporizes. Three of them, however iron, cobalt, and nickel are reducible to magnetic oxides. The larger number of them give characteristic colors to beads of borax and of microcosmic salt, fused on a loop of platinum wire before the blow-pipe. 144 ALUMINUM. 124, 1. None of them color the flame or give spectra, unless vaporized by a higher temperature than that of a Biinsen burner (spark spectra). THE IROX GROUP (THIRD GROUP). Aluminum, Chromium, Iron. 124, Aluminum. Al 27.1. Valence three. 1. Properties. Specific gravity, 2.708 (C. JV. f 105, 1912). The cast metal has specific gravity of 2.56. Melting point, 658.7 (Cir. B. S., 35, 1915). It is a tin-white metal (the powder is gray), odorless and tasteless, very ductile and malleable, about as hard as silver. Its boiling point is above 2200. Impurities increase the melting point. When molten it possesses great fluidity. As a con- ductor of heat it is about twice as good as tin and about one-third as good as silver. It conducts electricity about one-half as well as copper and silver (Dewar and Fleming, Phil. Mag., (5) 36, 271, 1893. Roy, Inst. Gt. Brit., June 5, 1896), and about three times better than iron. Commercial aluminum is never pure, containing small amounts of silicon and iron, and sometimes Cu and Pb , with 96 to 99.75 per cent aluminum. It is used for cooking utensils, canteens and other military equipments, boats, small weights, measures, articles of ornament and scientific instruments; as an alloy with copper (aluminum bronze) it finds exten- sive application. 2. Occurrence. Not found free in nature. Is found in corundum, ruby anc 1 sapphire, as nearly pure A1 2 O 3 ; in diaspore (A1(OH) 3 .A1 2 O3) ; in bauxite (ALOs.xHzO) ; in orthoclase (KlSi,jO s ) , and other feldspars ; in cryolite (Na^AlFe) . As a silicate in all clays and in very many minerals. It is widely distributed, constituting about one-twelfth of the earth's crust. 3 Preparation. (i) By electrolysis of the fused NaAlCl 4 . (2} By fusion of cryolite or the chloride with Na or K . (3) By heating NaAlCl 4 with zinc, with which it forms an alloy from which the zinc is driven off by a white heat. (4) By fusion of the chloride with potassium cyanide. (5} By fusing ALS 3 with iron. A great many new methods have been patented. (6) Aluminum is prepared commercially, by electrolysis of aluminum oxide dissolved in a bath of cryolite (Na-iAlF 6 ). The metal is deposited around the cathode, oxygen being evolved at the anode. See Dammer, 3, 79. 4. Oxide and Hydroxides. A1 2 O 3 is formed by heating- the hj-droxide, nitrate, acetate or other organic salt, difficultly soluble in acids after ignition, but may be dissolved after fusion with KHSO 4 or Na^CO., . A1(OH) 3 is formed when aluminum salts are precipitated with cold ammonium hydroxide. A1 2 0(OH) 4 is formed if the precipitation is made at 100. 5. Solubilities. a. Metal. Pure aluminum scarcely oxidizes at all in dry or moist artr; the electrolytically deposited powder oxidizes gradually in the air. Powdered or leaf aluminum when boiled with water evolves hydrogen, forming the hydroxide. It is -attacked by the halogens forming the corresponding halides (Gustavson, BL, 1881, (2), 36, 556). Dilute sulphuric acid attacks it slowly, evolving hydrogen (Ditte, C. r., 1890, 11O, 573); the hot concentrated acid dissolves it readily with evolution of S0 2 . Nitric acid, dilute or con- centrated, attacks it very slowly (Deville, A. CJi., 1855, (3), 43, 1-1: Montemartini, Gazzetta, 1892, 22, 397; Ditte, I.e., 782). Hydrochloric acid, dilute or concen- trated, dissolves it readily with evolution of hydrogen; also attacked readily by fixed alkalis, sparingly by NH 4 OH (Gottig, J5., 1896, 29, 1671), evolving hydrogen with formation of an aluminate: 2A1 + 2KOH + 2H 2 = 2KA10, -f 3H 2 . It is attacked by fixed alkali carbonates (D., 3, 87). When ignited with sodium carbonate, aluminum oxide is formed, sodium is vaporized and a small amount of aluminum nitride produced (Mallet, J. C., 1876, 30, 349). Fused KOH is decomposed by aluminum at very high temperature, the potassium being vaporized (Deville, J., 1857, 152). It is not at all attacked by cold four per cent acetic 'acid (vinegar) even in presence of NaCl , and when boiled for 124, 6a. ALUMINUM. 145 14 hours with the above mixture a square meter of surface (weighing 24.7426 grams) lost but 0.047 gram (one part in 526). b. Oxide and hydroxide. The oxide is insoluble in water, and when not too strongly ignited dissolves readily in dilute acids and in fixed alkalis. Corundum, crystallized A1 2 O 3 , is insoluble in acids, but is rendered soluble by fusion in fixed alkali carbonates or sulphates. The hydroxide A1(OH) 3 is insoluble in water, readily soluble in acids and in fixed alkalis, sparingly soluble in ammonium hydroxide, the solubility, however, being much decreased by the presence of ammonium salts. c. Salts. Aluminum phosphate is insoluble in water. The normal acetate is soluble, the basic acetate insoluble in water (separation from Cr and the fourth group).. The chloride is deliquescent. The double sulphates of aluminum and the alkali metals (alums) are soluble and readily melt in their water of crystallization, becoming anhydrous. Solutions of normal salts of aluminum have an acid reaction. 6. Reactions, a. The alkali hydroxides and carbonates * precipitate aluminum hydroxide (1), A1(OH) 3 (4), grayish -white, gelatinous insoluble in water, soluble in excess of the fixed alkali hydroxides f "() (Prescott, /. Am. Soc., 1880, 2, 27; Ditte, A. Ch., 1897 '(6), 30, 266), sparingly soluble in the fixed alkali carbonates and in ammonium hydroxide but much less so if ammonium salts be present. The solution of fixed alkali alnminate is precipitated as aluminum hydroxide by careful neutralization of the alkali with acids including hydrosulphuric (3), and carbonic, as basic hydroxide, by adding excess of ammonium chloride (4) (distinction from zinc which is precipitated by a small amount of NH 4 C1 , but redissolves on adding an excess) (Lowe, Z., 1865, 4, 350). The excess of potassium hydroxide liberates ammonia forming potassium chloride, thus reducing the amount of fixed alkali present. The precipitate is more compact and washes more readily than the gelatinous normal hydroxide. Barium car- bonate, on digestion in the cold for some time completely precipitates aluminum salts as the hydroxide (5) mixed with a little basic salt. (See 126, 6a.) The presence of citric, oxalic, or tartaric acid greatly hinders the precipitation of aluminum hydroxide, and an excess may entirely pre- vent its precipitation by the formation of a soluble double salt, e. g., KA1(C 4 H 4 6 ) L , . Other organic substances, as sugar, pieces of filter paper, etc., hinder the precipitation. To obtain complete precipitation all or- ganic substances should be decomposed. (1) A1C1, + 3KOH =-Al(OH), + 3KC1 2A1C1 3 + 3K 2 CO, + 3H 2 = 2A1(OH) 3 + 6KC1 + 3CO 2 (2) Al(OH), + KOH KA10, + 2H 2 O or A1C1, + 4KOH = KA1O 3 + 3KC1 + 2H 2 O (3) 2KA10 2 + H 2 S + 2H 2 = 2A1(OH) 3 + K 2 S (4) 2KA10 2 + 2NH 4 C1 + H 2 = A1 2 0(OH), + 2KC1 + 2NH S (5) 2A1C1 3 + 3BaC0 3 + 3H 2 = 2A1(OH) 8 + 3BaCl 2 + 3C0 2 * According to Langlota (A. Ch., 1856, (3), 48, 502) the precipitate with alkali carbonates always contains COi. He assigns the formula SCAhOsCCh) + SCAlsOs.SHaO). t A solution of barium hydroxide may be used to dissolve the Al(OH)s in separating from Fe(OH)3 and Cr(OH)s ; especially valuable in detecting the presence of small amounts of alu' minum when the reagents NaOH and KOH contain aluminum (Neumann, M., 1894, 15, 53). 146 ALUMINUM. 124, 6b. ft- Oxalates do not precipitate aluminum salts. The acetate of alum- inum is decomposed upon boiling, forming the insoluble basic acetate (separation of iron and aluminum from the fourth group) : A1(C 2 H 3 2 ) 3 + H 2 = A1(C 2 H 3 2 ) 2 OH + HC 2 H,0 2 . The basic acetate is best formed as follows: To the solution of aluminum salt add a little sodium or am- monium carbonate, as much as can be added without leaving a precipitate on stirring, then add excess of sodium or ammonium acetate, and boil for some time, when the precipitation at length becomes very nearly complete. Phenyl hydrazine, C 6 H-NHNH 2 , completely precipitates aluminum as the hydroxide from the neutral solution of its salts (complete separation of aluminum and chromium from iron which should be in the ferrous condition) (Hess and Campbe.ll, J. Am. Soc., 1899, 21, 776). c. Nitric acid is a very poor solvent for metallic aluminum, but a good solvent for the oxide and hydroxide. The metal dissolves in a solution of the normal aluminum nitrate, evolving hydrogen and forming the basic nitrate A1 4 O 6 (N0 8 ), (Ditte, C. r., 1890, 110, 782). d. Alkali phosphates precipitate aluminum phosphate, A1P0 4 , white, insoluble in water and acetic acid, soluble in mineral acids, and in the fixed alkalis (separation from FeP0 4 ) (Grueber, Z. angew., 1896, 741). A separation of Al and P0 4 may be effected by dissolving in hydrochloric acid adding tartaric acid and then ammonium hydroxide, and digesting some time with magnesia mixture (magnesium sulphate to which sufficient ammonium chloride has been added so that no precipitate is obtained when rendered strongly alkaline with ammonium hydroxide). The filtrate contains nearly all of the aluminum. The same method may be employed with Fe'" and*P0 4 . See also 7. e. The sulphide of aluminum cannot be prepared in the wet wa} r , that prepared in the dry way being decomposed by water (Curie, C. N. 9 1873, 28, 307). Hydrosulphuric acid does not precipitate aluminum from acid or neutral solutions; from its solutions in the fixed alkalis it is precipitated as the hydroxide on addition of sufficient hydrosulphuric acid to neutralize the fixed alkali (distinction from zinc which is rapidly precipitated from its alkaline solutions, as the sulphide). The alkali sulphides precipitate aluminum from its solutions, as the hydroxide; from acid or neutral solu- tion H 2 S is evolved: 2A1C1, + 3(NH 4 ) 2 S + 6H 2 = 2A1(OH) 3 + 6NH 4 C1 -f- 3H 2 S , from solutions in the fixed alkalis ammonia is evolved, fixed alkali sulphide being formed: 2KA10 + (NH 4 ) S + 2H = 2A1(OH) 3 + KS Sodium thio sulphate precipitates, from aluminum salts, in neutral solutions, aluminum hydroxide with free sulphur and liberation of sulphurous anhydride: 2A1 2 (S0 4 ) 3 + GNa,S 2 O 3 + 6H 2 O = 4A1(OH) 3 + 3S 2 + 6Na,SO 4 + 6SO, . A small amount of sodium tetrathionate is formed and also some hydrosulphuric acid (Yortmann, B., 1889, 22, 2307). Sodium sulphite also precipitates alu- 124, 0. ALUMINUM. 147 minum hydroxide, with liberation of sulphur dioxide: 2A1CL + 3Na 2 SO 3 + 3H 2 O = 2A1(OH) 3 + GNaCl + ?>SO, . Neither of 1he above reagents precipi- tate iron salts, thus effecting- a separation of aluminum (and chromium) from iron. Aluminum, chromium and ferric sulphates crvsialli/c wilh the sulphates of the alkali metals, forming a class of compounds, AI.TMS, of which th potassium aluminum compound is perhaps best known, KA1(S0 4 ) 2 .12H 2 , connnoii alum. These compounds melt in their water of crystallization, becoming anhydrous upon further heating. The freshly ignited alum is only sparingly soluble in cold water, but upon standing becomes readily soluble, dissolving in less than one part of hot wate 1 *. The alums are usu- ally less soluble than their constituent sulphates and may be precipitated by adding a saturated solution of alkali sulphate to a very concentrated so- lution of Al , Cr'" , or Fe'" sulphate. f. Aluminum chloride is a very powerful dehydrating- agent and is much used in organic chemistry as a halogen carrier. An impure aluminum chlorate, mixture of KC1O 3 and AL(S0 4 ) 3 , is much used in calico printing (Schlum- berger, Dingl., 1873, 207, 03). y. Aluminum salts are precipitated by solu- tions of alkali arsenites and arsenates, but not by arsenous or arsenic acids. 7i. Potassium chromate forms a yellow gelatinous precipitate, potassium bichromate gives no precipitate with aluminum salts, i. Solution of borax precipitates an acid aluminum borate, quickly changed to aluminum hydroxide. 7. Ignition. Compounds of aluminum are not reduced to the metal, but most of them are changed to the oxide, by ignition on charcoal. If now this residue is moistened with solution of cobaltous nitrate, and again strongly ignited, it assumes a blue color. This test is conclusive only with infusible compounds, and applies only in absence of colored oxides. Aluminum com- pounds ignited on charcoal in presence of sulphur are changed to A1 2 S 3 (Buch- erer, Z. angew., 1892, 483). To separate Al from PO 4 , fuse the precipitate or powdered substance with iy 2 parts finely divided silica and 6 parts dried sodium carbonate in a platinum crucible, for "half an hoiir. Digest the mass for some time in water; add ammonium carbonate in excess, filter and wash. The residue consists of aluminum sodium silicate; the solution contains the PO 4 , as sodium phosphate. The Al can be obtained from the residue by dissolving it in hydrochloric acid, evaporating to dryness to render the silica insoluble. Treat with hydrochloric acid and filter; the filtrate containing aluminum chloride. 8. Detection. After the removal of the first two groups it is precipi- tated with Cr and Fe'" as the hydroxide, A1(OH) 3 , by NH 4 OH in the pres- ence of NH 4 C1 . It is separated from Fe(OH) 3 and Cr(OH),, by boiling with KOH or NaOH or by fusion with an alkaline oxidizing agent such as NaCIO or Na 2 2 . From the nitrate acidulated with HC1 it is precipitated as hydroxide with (NH 4 ) 2 C0 3 ; or it is precipitated from the alkaline solution by an excess of NH 4 C1 (6a). 9. Estimation. Aluminum is usually weighed as the oxide, after ignition It is separated from zinc as a basic acetate; from chromium by oxidizing the latter to chromic acid, by boiling with potassium chlorate and nitric acid, or by fusing with KNO 3 and Na 2 CO 3 , or by action of Cl or Br in presence of KOH , or by NaaO 3 fused or in solution, and after acidulating with HC1 precipi- CHROMIUM. 124, 10. tating the aluminum with ammonium hydroxide. It may be separated from iron by boiling with KOH (6a), by Na 2 S 2 O 3 (6e), or by phenylhydrazine (6b). It is separated from iron by conversion into the oleate and dissolving the oleate of iron (Fe'" or Fe") in petroleum (Borntraeger, Z., 1893, 32, 187). It is some- times precipitated and weighed as the phosphate. 10. Oxidation. Aluminum reduces solutions of Pb , Ag , Hg *, Sn , Bi (incompletely), Cu f, Cd , Co , Ni , Zn J and Gl (in alkaline mixture only), Te, Se, An, and Pt , to the metallic state; ferric salts to ferrous salts; As and Sb with HC1 become respectively AsH 3 and SbH 3 ,with alkalis As'" is reduced to AsH 3 . As v is unchanged (69, 6'& and 10), and Sb"' and Sb v become Sb. Aluminum salts are not reduced to the metallic state by any other compounds at ordinary temperature; by fusion with K or Na metallic aluminum is obtained, much better, however, by the aid of the electric current. 125. Chromium. Cr = 52.0. Valence two, three and six. 1. Properties. Specific gravity, 6.92 (Moissan, C. r., 116). Melting point 1520 (Cir. B. ., 36, 1915). A grayish-white crystalline metal. The hardness of steel is greatly increased by the presence of less than one per cent of chromium. It is non-magnetic (Woehler, A., 1859, 111, 231). It burns to the oxide Cr 2 O 3 when heated to 200 to 300 in the air (Moissan, C. r., 1879, 88, 180). 2. Occurrence. Not found native. Chrome-ironstone or chromite (FeOCr 2 O 3 ) is the chief ore of chromium, and is usually employed in the manufacture of chro- mium compounds. Chromite and also Daubreelite (FeCr 2 S 4 ) , are frequently found in meteorites; it also occurs in crocoite (PbCrO 4 ) , and other rare chro- mates and sulphates. 3. Preparation. (1) By electrolysis of the chloride. (2) By fusing the chloride with potassium or sodium. (3) By ignition of the oxide with carbon. (4) By fusing CrCl 3 with Zn , Cd or Mg , using KC1 and NaCl as a flux, and removing the excess of the Zn , Cd or Mg by dissolving in nitric acid, which does not dissolve metallic chromium. (J) By ignition of the oxide with alu- minum (Goldschmidt, A., 1898, 301, 19). 4. Oxides and Hydroxides. CJiromous oxide, CrO , has not been isolated. The corresponding hydroxide, Cr(OH) 2 , is made by treating CrCL with KOH. Chromic oxide, Cr 2 O 3 , is made by a great variety of methods, among which are fusing the nitrate, or higher or low r er oxides and hydroxides in the air; heating mercurous chromate, or the dichromates of the alkalis: 4Hg 2 Cr0 4 = 2Cr 2 3 + SHg + 50 2 (NH 4 ) 2 Cr 2 7 = Cr 2 3 + N 2 + 4H 2 4K 2 Cr 2 O 7 = 2Cr,O 3 + 4K 2 Cr0 4 + 3O 2 In the last the K 2 CrO 4 may be separated by water. After heating to redness, Cr,O 3 is insoluble in acids. Chromic hydroxide, Cr(OH) 3 , is precipitated by adding NH 4 OH to chromic solutions. That formed by precipitating with KOH or NaOH retains traces of the alkali, not easily removed by washing. Chromium trioxide or chromic anhydride, Cr0 3 , is formed as brown-red needles upon addition of concentrated sulphuric acid to a concentrated solution of K 2 Cr 2 O 7 ; to be freed from sulphuric acid it must be recrystallized from water, in which it is readily soluble, or treated with the necessary amount of * Klandy, C. C., 1893, 201 ; Wislicenus, B. 1895, 28, 1323. t Tommasi, Bl., 1882, (2), 37. .{,:}. t Flavitsky, B., 1873, 6, 195 ; Zimmerman, Z., 1888, 27, 61 ? , 60. CHROMIUM. 149 BaCrO 4 (Moissan, A. Ch., 1885, (6), 6, 568). It is also prepared by transposi- tion of BaCrO 4 with HNO 3 or H 2 SO, ; PbCrO 4 with H 2 SO 4 ; and Ag2CrO 4 with HC1; etc. It melts at 196, decomposing at higher temperature into Cr 2 O 3 and O . It is used in dyeing silk and wool, but not cotton fabrics. It is a powerful oxidizing agent, being reduced to chromic oxide. The existence of chromic acid, H 2 CrO 4 , is disputed (Moissan, L c.; Field, C. N., 1892, 65, 153; and Ostwald, Zeit. phys. Ch., 1888, 2, 78). Two series of salts are formed as if derived from chromic acid, H 2 CrO 4 , and dichromic acid, H 2 Cr 2 O 7 . The salts are quite stable and find an extended application in analytical chemistry (Qh. 57, 59, 186, etc.). 5. Solubilities. a. Metal. Chromium is not at all oxidized by water or moist air at 100. Heated above 200 it is oxidized to Cr,O, , rapidly in pres- ence of KOH . It is soluble in HC1 or dilute H,S0 4 ; insoluble in concentrated HoSO 4 or in HNO ;i , dilute or concentrated. Chlorine or bromine attack it with formation of the corresponding- halides (Woehler, L c.; Ufer, A., 1859, 112, 302). &. Oxides and Hydroxides. Chromic o.ride, Cr,0 3 , is insoluble in water, slowly soluble in acids, but not at all if previously ignited (Traube, A., 1848, 66, 88); the hydroxide is insoluble in water, soluble in acids, sparingly soluble in ammonium hydroxide, soluble in fixed alkalis to chromites, reprecipitated again upon boiling-. The presence of other metallic hydroxides, as iron, etc., hinders the solution in fixed alkalis. Chromic anhydride, Cr0 3 , is very soluble in water, soluble in reducing acids to chromic salts. c. Salts. Chromic sulphide is not formed in the wet way, being decomposed by water; the phosphate is insoluble in water. The chloride exists in two modifications ; a deliquescent soluble chloride, which also forms a soluble basic chloride (Ordway, Am. S., 1858 (2), 26, 202); and a violet sublimed chromic chloride absolutely insoluble in water, hot or cold, or in dilute or concentrated acids, the presence of a very small amount of chromous or stannous chloride at once renders this modi- fication soluble in water (Peligot, A. Ch., 1846 (3), 16, 298); the bromide and sulphate also exist in soluble and insoluble modifications; the nitrate and also the basic nitrates are readily soluble in water (Ordway, 1. c.). There are many double salts, the sulphates of chromium and the alkali metals, chrome alum, forming salts similar to the corresponding aluminum compounds. There are two modifications of solutions of chromium salts, one having a green color and the other violet to red, the tints being modified somewhat by the degree of the concentration. All normal chromic salts in solution have an acid reaction, being partially hydrolized. 6. Reactions.* a. Alkali hydroxides and carbonates precipitate solu- tions of chromic salts, as chromium hydroxide, gelatinous, gray-green or gray-blue according to the variety of solution from which it is obtained (5c), insoluble in water, soluble in acids; soluble in excess of the fixed alkalis to chromites : Cr(OH) 3 + KOH <=> KCr0 2 + 2H 2 . This reaction * Chromous salts are very unstable, they are great reducing agents, oxidizing- rapidly when exposed to the air. They are almost nver met with in analysis. Chromous chloride, CrCl 2 , is formed when the metal is heated in contact with hydrochloric acid gas (Ufer, I. c ); also by re- duction of CrCI 3 with hydrogen in a heated tube (Moberg, J. pr., 1848, 44r, 322). Precipitates are formed in its solutions by the alkali hydroxides, carbonates, sulphides, etc. (Moissan, J37., 1HX2 (2), 37, 296). 150 CHROMIUM. 125, 6. is reversible. The KOH tends to form, while the water, especially when hot, tends to decompose the chromite. The chromium may therefore be reprecipitated from this solution if the excess of KOH is small and the solution is diluted and boiled. (Distinction from aluminum.) As ammo- nium chloride reacts with the caustic potash forming potassium chloride and the weak alkali, ammonium .hydroxide, the chromium may also be pre- cipitated by the addition of ammonium chloride and heating. The presence of ferric hydroxide and some other compounds greatly hinders the solution in fixed alkalis, hence chromium cannot be separated from iron by excess of fixed alkali. Chromium hydroxide is slightly soluble in excess of cold ammonium hydroxide to a violet solution, completely reprecipitated on boiling. The precipitate formed with the alkali carbonates is almost entirely free from carbonate: 2CrCl 3 + 3Na,C0 3 + 3H 2 = 2Cr(OH) 3 + GNaCl -(- 3C0 2 . Barium carbonate precipitates chromium from its solu- tions (better from the chloride) as a hydroxide with some basic salt, the precipitation being complete after long digestion in the cold (separation from the fourth group). For removal of excess of reagent, add H S0 4 and the filtrate will contain the chromium as a sulphate. Alkali dichromates are changed to normal chromates by alkali hydrox- ides or carbonates. 1). Chromium forms no basic acetate and remains in solution when the basic acetates of aluminum and ferric iron are formed (6&, 124 and 126). Potassium cyanide precipitates chromium hydroxide. Oxalates and ferro- cyanides cause no precipitate. H 2 Cr0 4 is reduced to chromic compounds by K 3 Fe(CN) 6 and KCNS. r. Nitrites or nitrates are without action upon chromium salts in the wet way, but upon fusion in presence of nitrites or nitrates and alkali carbonate a chromate is formed (separation from Fe and Al). d. Hypophosphorous acid reduces chromates to chromic salts. Soluble phosphates, as Na,HPO 4 , precipitate chromic phosphate, CrPO 4 , insoluble in acetic acid, decomposed by boiling 1 with KOH , leaving the phosphate in solu- tion (Kammerer, J. C., 1874, 27, 1005). e. Hydrosiilplmric acid is without action upon neutral or -acid solutions of chromic salts, chromites as KCr0 2 are precipitated as chromium hydroxide; 2KCr0 2 + HJ3 + 2H 2 =" 2Cr(OH) 3 + K 2 S . The hexad chromium of chromates is reduced to the triad condition with liberation of sulphur, in neutral or alkaline solutions, chromium hydroxide being formed: 2K 2 Cr,0- + 8H 2 S =: 4Cr(OH) 3 +2K 2 S + 3S 2 + 2H 2 ; in acid solutions a chromic salt is formed (10). Alkali sulphides precipitate chromic salts as the hydroxide liberating H 2 S : 2CrCL, + 3(NH 4 )S + GH 2 = 2Cr(OH) 3 + GNH 4 C1 + ?>H 2 S Chromates are reduced and precipitated as chromium hydroxide with sepa- ration of sulphur: 4K 2 Cr0 4 + 6(NH 4 ) 2 S + 4H 2 = 4Cr(OH) 3 + 8KOH 125, 8. CHROMIUM. + 3S 2 -f- 12NH 3 . Soluble sulphites and thiosulpliates reduce chromates in acid solution (Donath, /. C., 1879, 36, 401; Longi, Gazzetta, 1896, 26, ii, 119). f. Hydrochloric acid reduces chromates to chromic chloride on boiling,, with evolution of chlorine: 2K,Cr0 4 + 1GHC1 = 2CrCl 3 -f 4KC1 + 3C1 2 + SHoO ; more readily without evolution of chlorine in presence of other easily oxidized agents, as alcohol, oxalic acid, etc. : K 2 Cr 2 7 -f- 8HC1 -f- 3C 2 H 8 OH = 2KC1 -f 2CrCl 3 + 3C 2 H 4 (acetaldehyde)"+"?'H 2 . If the dry chromate be heated with sulphuric acid and a chloride (transposable by sulphuric acid) (269, 5), brown fumes of chromium dioxydichloride are evolved: K 2 Cr 2 7 -f 4NaCl -f- 3H 2 S0 4 *= 2Cr0 2 Cl 2 + K 2 S0 4 + 2Na 2 S0 4 + 3H 2 (269, Sd) (Moissan, Bl, 1885 (2), 43, 6). To obtain a quantity of Cr0 2 Cl 2 , Thorpe (J. C., 1868, 21, 514) recommends 10 parts of NaCl and 12 parts K 2 Cr,0 7 fused together and distilled with 30 parts of H 2 S0 4 . Hydrobromic acid reduces chromates to chromic bromide with evolution of bromine; hydriodic acid to chromic iodide with evolution of iodine. In the presence of hydrochloric or sulphuric acids all the bromine or iodine is set free. K 2 Cr 2 7 + 6HI + 4H 2 S0 4 = K 2 S0 4 + Cr 2 (S0 4 ) 3 4- 3I 2 + 7H 2 . Hydriodic acid acts most readily upon chromates, the hydrochloric least readily. Chromic hydroxide and chromic salts, when boiled with chloric or bromic acids, or potassiuni^hlorate or brornate and nitric, sulphuric or phosphoric acids, become chromic acid. g. Soluble arsenites and arsenates form corresponding' salts with chromic salts. Chromates in acid solution are instantly reduced to chromic salts by arsenites or arsenous acid. Chromic acid boiled with arsenous acid in excess gives CrAsO, (Neville, J. C., 1877, 31, 283). k. Potassium chromate colors an acid solution of chromic salt brown-yellow' on addition of ammonium hydroxide, a precipitate of the same color is obtained, chromic chromate (Maus, Po#H 3 2 ) 2 . These metals may also be separated by the methods given in 10". 9. Estimation. Chromium is usually estimated gravimetrically (1) as the oxide. It is brought into this form either by precipitation as a hydroxide (6a) and ignition or, in many cases, by simple ignition (4). (2) As chromate, it may be precipitated with barium chloride, dried and weighed as such; or in acetic acid solution it may be precipitated as PbCrO 4 by PbCC.HgO,), , dried and weighed. Volumetrically, as a chromate (if present as chromic salt it may be oxidized to a chromate). (3) By titration with a standard solution of ferrous sulphate. (//) By liberation of iodine from hydriodic acid (60) and measuring the amount of iodine liberated with standard sodium thiosulphate solution. 10. Oxidation. Chromous compounds are very strong reducing agents, changing HgCl 2 to HgCl , CuS0 4 to Cu, SnCl 2 to Sn, etc. Chromic com- pounds are oxidized to chromates by chlorates (Giacomelli, UOrosi, 1895, 18, 48; Storer, Am. S., 1869,98,190) (6/), Na,0 2 , Mn0 2 (Marchal and Wier- nick, Z. angew., 1891, 511), and Pb0 2 in acid solution; in alkaline mixture, by reducing Pb0 2 to PbO , Ag 2 to Ag, Hg 2 and HgO to Hg, CuO to CiuO , KMn0 4 and K 2 Mn0 4 to Mn0 2 (Donath and Jeller, C. C., 1887, 151); by Cl, Br, and I, forming the corresponding halide; and by H 2 2 * (Baumann, Z. angew., 1891, 139). The halogens in alkaline solution may be used to separate chromium from iron and aluminum. The halogens react with the alkali and oxidize the chromium to chromate according with the following reactions: 2NaOH + C1 2 = NaCl + NaCIO + H 2 O 2NaCrO 2 + SNaCIO + NaOH = 2Na2CrO 4 + 3NaCl + H 2 O. The ferric hydroxide is not acted upon, while the aluminum hydroxide dissolves in the fixed alkali and may be separated from the chromium by the addition of ammonium chloride and warming. Na 2 2 or H 2 2 in alka- line solution produce a similar reaction. 2NaCrO 2 + SNa^ + 2H 2 O = 2Na 2 CrO 4 + 4NaOH, the separation of iron and aluminum being effected by the method already given. * The use of H 2 O 2 in alkaline solution is proposed by Riggs (Am. , 1894, 148, 409) in the sepa- ration of Al, F'e and Cr. 100 cc. water, 10 cc. H 2 O 2 , and one gram of BfaOH are added to the freshly precipitated hydroxides and digested until effervescence ceases. Filter off the precipi- tate of ferric hydroxide, acidify tho filtrate with acetic acid and precipitate the aluminum witb ammonium hydroxide. Tfce chromium if present will be in the filtrate as sodium ohromatc. 126, 1. IRON. 153 A chromate is also formed when chromium compounds are fused with an alkali carbonate and an oxidizing agent (7). For this purpose sodium carbonate and potassium nitrate or sodium peroxide are frequently used. The following reactions take place: 2Cr(OH) 3 + 2Na 2 CO 3 + 3KNO 3 = 2Na 2 CrO 4 + 3KNO 2 -f 2CO 2 + 3H 2 O 2Cr(OH) 3 + 3Na 2 O 2 = 2Na 2 CrO 4 + 2NaOH + 2H 2 O. Ferric hydroxide is not acted upon while aluminium hydroxide is con- verted into aluminate. 2A1(OH) 3 + Na 2 CO 3 = 2NaAlO 2 + CO 2 + 3H 2 O . On dissolving the fused mass in water, the three metals may be separated by the method already given. Chromic oxide (not ignited) or chromic chloride at 440 in a current of chlorine become Cr0 2 Cl 2 (Moissan, El., 1880, (2), 34, 70). Chromic acid and chromates are reduced to chromic compounds by H 2 C 2 4 (Werner, /. C., 1888, 53, 602), K 4 Fe(CN) 6 , KCNS , H 2 S , (NH 4 ) 2 S , Na 2 S 2 0, , S0 2 , H 2 2 , etc. Of most common occurrence in qualitative analysis is the action of hydrosulphuric acid and alkali sulphides; at first sulphur is liberated, a part of which may be oxidized to sulphurous and sulphuric acids (Parsons, C. N., 1878, 38, 228). 2K 2 Cr 2 7 + 16HC1 + 6H 2 S = 4CrCl 3 + 4KC1 + 3S 2 + 14H 2 O 12H 2 Cr0 4 + 3S 2 = 4Cr 2 0,Cr0 4 + 6SO, -f 12H 2 O 2H 2 Cr0 4 + 3SO, = Cr 2 (S0 4 ) 3 + 2H 2 While H 2 2 in alkaline solution oxidizes Cr'" to Cr vl , in acid solution the reverse action takes place: 2H 2 Cr0 4 + 3H 2 S0 4 -f- 3H 2 2 = Cr 2 (S0 4 ) 3 + :>0 2 + 8H 2 (Baumann, 1. c.). With chromate in acid solution, the H,0 2 at first gives a deep blue solution (probably of the very unstable perchromic acid, HCr0 4 ), followed by the reduction to a chromic salt. The blue color gives a very delicate test for chromium. The test is rendered more deli- cate by the addition of a few c.c. of ether and shaking. The ether dis- solves the blue compound and forms a blue layer on standing. One part of chromic acid in 40,000 parts of water can be detected by this reaction. Vanadic acid interferes with the delicacy (Reichard, Z., 40, 577). 126. Iron, (Ferrum). Fe 55.84. Usual valence two and three. 1. Properties. Specific gravity, variable, depending upon the purity and methods of preparation. 7.85 at 16 (Caron, C. r., 1870, 70, 1263), 8.139 (Chandler-Roberts, C. N., 1875, 31, 137). Melting point, cast iron, 1130 to 1375; steel, 1375 to <1530; pure iron, 1530 (Cir. B. S., 36, 1915). Pure iron is silver-white, capable of taking a remarkably fine polish; it is among the most ductile of metals, in this property being approached by nickel and cobalt (73, 1); it is the hardest of the ductile metals (Calvert and Johnson, DingL, 1859, 162, 129), and in tenacity it is surpassed only by cobalt and nickel (132, 1). It softens at a red heat and may be welded at a white heat. Finely divided iron burns in the air when ignited; that made by reduction in hydrogen may ignite spontaneously when exposed to the air. When pure iron is heated, or cooled through certain ranges of temperature, polymorphic changes occur, which are 154 IRON. 126, 2. accompanied by an absorption (on heating) or an evolution (on cooling) of heat, and changes in the physical properties of the iron; these polymorphic modifica- tions are called Alpha, Beta and Gamma iron, respectively. Alpha iron is stable in all ranges of temperature up to 768 C. (Ac 2 ) ; * at 768 C. the change Alpha ^ Beta occurs; Beta iron is stable between 768 C. and 909 C. (Ac 3 ) ; * at 909 C. the change Beta - Gamma takes place (Burgess and Crowe, Reprint, 213, Bui. B. S., 10, 1913). Gamma iron is capable of dissolving, and retaining in solid solution, carbon as iron carbide (Fe?C) ; the presence of Fe 3 C in solid solution, however, progressively lowers the tem- perature of the A 3 * point, at which the change Gamma > Be t-i occurs, and causes the appearance of a third point (Ai) * at 700 C., the temperature of which is unaffected by varying percentages of carbon. Because of the pro- gressive lowering of the temperature at which the Gamma > Beta (Ar 3 ) * change takes place by increasing percentages of carbon, the (A 3 ) * point merges into the (A 2 ) * point at 768 C., and 0.33% C. At this temperature and concentration, and between it and 700 C. (Ai)* Gamma iron probably passes directly into the alpha modification, with a continued concentration of carbon in the solid solution still remaining, until the Ar 3 - 2 * point and ATI * point coincide (Ar 3 - 2 -i). * This point (700 C., and 0.835% C.) marks the lowest temperature at which carbon (as Fe 3 C) will remain in solid solution; if the temperature falls below Ari,* Carbon separates from the solid solution as Fe 3 C (cementite) and at the same instant the Gamma iron, which up to that moment had acted as a solvent for the carbon is transformed to Alpha iron, simultaneously with the precipitation of the cementite, forming that interstratified conglomerate of Alpha ferrite, and cementite, known as Pearlite (the eutectoid). Above 700 C. the solubility of Fe 3 C in Gamma iron increases with the tem- perature, reaching a maximum at a concentration of about 1.7% carbon and a temperature of 1130 C. Hence all slowly cooled iron-carbon alloys containing less than the eutectoid percentage of carbon (<0.835% C.; hypo-eutectoid steels) will consist of structurally free Alpha Ferrite (pure iron) and Pearlite; those with more than the eutectoid percentage of carbon (between 0.835% C. and 1.7% C.; hyper-eutectoid steels) will consist of structurally free Cementite (Fe 3 C) and Pearlite; those of eutectoid composition (0.835% C.; eutectoid steel), will consist of Pearlite only. When such steels, provided the carbon con- tent lies between 0.25% and 1.75%, are heated to a temperature higher than Ac 3 * (Ac cm * in the case of hyper-eutectoid steels) and then suddenly cooled (e.g., by quenching in water), the changes which would normally occur on slow cooling through Ar 3 , Ar 2 , Ari * are partially suppressed; the steel becomes hard and brittle, the carbon being retained in solid solution. This process is called "hardening." The hardened steel may be " tempered" by reheating to temperatures lower than Aci* (700 C.) which causes certain structural changes to take place leading to the formation of a series of transition products, the general effect of which is to soften the steel somewhat and lessen its brittleness to the extent desired. 2. Occurrence. Iron is rarely found native except in meteorites; the iron minerals used as ores are hematite (Fe 2 Oj), limonite (2Fe(OH) 3 .Fe 2 O s ) ,- mag- netite (Fe 3 O 4 ) ; to a less extent, siderite (FeCO 3 ), clay iron stone (FeCO.i with clay) black band (FeCO 3 with bituminous matter); it also occurs as pyrite (FeS 2 ), marcasite (FeS 2 ), pyrrhotite (FenSn+i) and widely distributed in many other min- erals and rocks. 3. Preparation. Pure iron is not usually found in the market, although some of the commercial products approach it very closely. Pig iron is produced by smelting iron ore mixed with coke and limestone in blast furnaces; the resulting product is subjected to remelting processes (prod- uct, gray and malleable iron castings), or to conversion processes (products, Bessemer steel; open- hearth steel, acid and basic; wrought iron); pure iron may be made by electrolysis, and by heating its purified salts with hydrogen. 4. Oxides and Hydroxides. Ferrous oxide, FeO , is made from Fe 2 O 3 by heat- ing it to 300 in an atmosphere of hydrogen; also by heating FeC 2 O4 to 160, * Ari-2-2 is derived from Arret (a halt, or pause) and refroidissement (cooling) ; hence, a halt in cooling at certain critical temperatures. Ar 3 =898C.; Ar 2 =768C.; An = 700C. Aci-:-3 (c = chauffant, heating) refers to corresponding points during heating. Acs is slightly higher (909 C.) than Ar. Accra or Arem refers to temperatures at which Cementite dissolves in, or is precipitated from Gamma iron during heating or cooling. 126, 51. IRQ!!. 155 air being excluded. It takes fire spontaneously in the air, oxidizing- to Fe 2 O 3 . Ferrous li!/.ri, Fe(OH), , is formed by precipitating- ferrous suits with KOH or NaOH , perfectly white when pure, but usually green from partial oxidation. /V/r/V o.r'nle, Fe,O ;i , is formed by heating- FeO , Fe(OH) 2 , or any ferrous salt consisting- of a volatile or organie acid in the air; more rapidly by heating Fe(OH) 3 , Fe(NO 8 ), , or Fe 2 (S0 4 ) 3 . Ferric hydroxide is formed by precipitat- ing- cold dilute ferric salts with alkalis or alkali carbonates, and drying at 100. If KOH or NaOH is used, the precipitate requires longer washing than when NH,OH is employed. By increasing the temperature and concentration of the solutions, the following definite compounds may be formed: FeO(OH) , Fe,0(OH) 4 , Fe 4 0,(OH) L ,, Fe 4 O 3 (OH),,, Fe 3 O 2 (OH) 5 . Fe 3 O 4 is slowly formed by heating- FeO or Fe,O 3 to a white heat. Its corresponding hydroxide may be made by precipitation: FeCL -f 2FeCl 3 + 8NH 4 OH = Fe 3 (OH) 8 + 8NH 4 C1 . Fe 3 (OH) s when heated to 90 forms Fe 3 4 . The black color and magnetic properties show that it is a chemical salt and not a mechanical mixture of FeO and Fe,O 3 . Fe'" acts as an acid towards the Fe"; this oxide, Fe 3 O 4 , or FeFe_,O 4 , may be called ferrous ferrite. Other ferrites have been formed, e. g., calcium ferrite, CaFe,O 4 ; MgFe 2 O 4 and BaFe 2 O 4 (List, #., 1878, 11, 1512); zinc ferrite, ZnFe,O 4 . Compare potassium aluminate, KA10, (124, 6a), and potas- sium chromite, KCrO, (125, 6a). Ferric acid, H L ,Fe6 4 , and its anhydride, FeO 3 , have not been isolated. Potassium ferrate, K 2 Fe0 4 , is made (/) by elec- trolysis; (2) by heating iron-filings, FeO or Fe 2 O 3 ," to a red heat with KN0 3 ; (3) by heating Fe(OH) 3 with potassium peroxide K 2 2 ; (4) by passing Cl or Br into a solution of 5 parts of KOH in 8 parts of water in which Fe(OH) 3 is suspended; the temperature should be not above 50. It has a purple color; is a strong oxidizing agent. It slowly decomposes on standing: 4K,FeO 4 + 10H 2 O = 8KOH + 4Fe(OH) 3 + 3O 2 . With barium salts it precipitates a stable barium ferrate, BaFeO 4 . 5. Solubilities. a. Metal. Iron dissolves, in hydrochloric acid and in dilute sulphuric acid, to ferrous salts, with liberation of hydrogen (a) ; concentrated cold H._>SO 4 has no action, but if hot, SO 2 is evolved and a ferric salt formed (6); in moderately dilute nitric acid, with heat, to ferric nitrate, liberating chiefly nitric oxide (c) ; in cold dilute nitric acid, forming ferrous nitrate with pro- duction of ammonium nitrate (d), of nitrous oxide (e), or of hydrogen f) (Langlois, A. Ch., 1856, [3], 48, 502). () Fe + H 2 S0 4 = FeS0 4 + H 2 (b) 2Fe + 6H 2 S0 4 = Fe 2 (S0 4 ) 3 -f 3S0 2 -f 6H 2 (c) Fe + 4HN0 3 = Fe(N0 3 ) 3 + NO + 2H 2 O (d) 4Fe -f 10HN0 3 = 4Fe(N0 3 ) 2 + NH 4 NO 3 + 3H 2 O (e) 4Fe + 10HN0 3 = 4Fe(N0 3 ) 2 + N 2 + 5H 2 O (0 Fe + 2HN0 3 = Fe(N0 3 ) 2 + H 2 In dissolving the iron of commerce in hydrochloric acid, the carbon which it always contains, so far as combined in the carbide of iron, will pass off in gaseous hydrocarbons (Campbell, Am., 1896, 18, 836), and so far as uncombined will remain undissolved, as graphitic carbon. The metal is attacked by moist air, forming chiefly 2Fe 2 O 3 .3H 2 O , iron rust. When hot iron is hammered, scale oxide, Fe 2 O 3 .6Fe6 , is formed. Cold concentrated HNO 3 or the action of the electric current forms passive iron. (Byers and Langdon, J. Am. Soc., 36, 759, 1913) (Byers, Ibid., 30, 1718, 1908). &. Oxides and hydroxides. Ferrous oxide and hydroxide unite with acids with rapid increase in temperature, forming ferrous salts, always mixed with more or less ferric salts. The ferrous salts are much more readily prepared by the action of dilute acids upon the metal, or upon FeC0 3 or FeS . Fe 3 O 4 , treated with an insufficient amount of HC1, forms FeCl 2 and Fe,O ;! : treated wit h HC1 sufficient for complete solution, a mixture of Fed,, and FeCL, is obtained^ which, when treated with excess of ammonium hydroxide and dried at 100 again exhibits the magnetic properties of the original. Ferric oxide, Fe,O 3 , dis- solves in acids, quite slowly if the temperature of preparation of the oxide has been high. Mitscherlich (J. pr., 1860, 81, 110) recommends warm digestion with tn parts of a mixture of sulphuric acid and water (8-3), If the oxide be 156 IRON. 126, be. heated with alkalis or alkali carbonates, it then dissolves much more readily in acids. Ferric hydroxide, Fe(OH) 3 , is insoluble in water (for a soluble colloidal ferric hydroxide, see Sabanejeff, C. C., 1891, i, 11), readily soluble in acids to ferric salts. Freshly precipitated ferric hydroxide readily dissolves in ferric chloride and in chromium chloride, not in aluminum chloride. A solution of ferric hydroxide in ferric chloride is soluble in water after evaporation to dry- ness if not more than ten parts of Fe,0 3 are present to one of the FeCl 3 (Be- champ, A. Ch., 1859, (3), 56, 306) c. Salts. Ferrous salts, in crystals and in solution, have a light green color. Solutions of the salts have a slight acid reaction toward litmus. The sulphate FeS0 4 .7H 2 , is efflorescent; the chloride, bromide, iodide, and citrate are deliquescent. Solutions of all ferrous salts are unstable, gradually changing to basic ferric salts, more or less insoluble in water. The carbonate, hydroxide, phosphate, borate, oxalate, cyanide, ferro- cyanide, ferricyanide, tartrate, and tannate are insoluble in water. Ferric salts in solution have a brownish-yellow color, redden litmus and color the skin yellow. The chloride, bromide, nitrate, and sulphate are deliquescent. The ferrocyanide, tannate, borate, phosphate, basic acetate, and sulphite are insoluble in water; the sulphate is soluble in alcohol (separation from ferrous sulphate). Ferric chloride is soluble in ether saturated with hydrochloric acid, separation from aluminum (Gooch and Havens, Am. 8., 1896, 152, 416). Solutions of ferric salts, when boiled, frequently precipitate a large portion of the iron as basic salt, especially if other soluble salts are present (Fritsche, Z. angeir., 1888, 227; Pickering, J. C., 1880, 37, 807) (70, 5d footnote). 6. Reactions, a. The alkali hydroxides precipitate ferrous hydroxide. Fe(OH) 2 , white if pure, but seldom obtained sufficiently free from ferric hydroxide to be clear white, and quickly changing, in the air, to ferroso- ferric hydroxide, of a dirty-green to black color, then to ferric hydroxide (4), otf a reddish-brown color. The fixed alkalis adhere to this precipitate. Ammonium chloride or sulphate, sugar, and many organic acids, to a slight extent, dissolve the ferrous hydroxide or prevent its formation (116 and 117). The soluble carbonates precipitate, from purely ferrous solutions, ferrous carbonate, FeC0 3 , white if pure, but soon changing, in the air, to the reddish-brown ferric hydroxide. Solutions of ferric salts are precipitated by the alkali hydroxides and carbonates as ferric hydroxide, Fe(OH) 3 , variable to Fe 2 3 .H 2 FeO(OH) reddish-brown insoluble in excess of the reagents (distinction from alumi- num and chromium which are soluble in excess of the fixed alkali hy- droxides and from cobalt, nickel and zinc which are soluble in ammonium hydroxide). Salts of the fixed alkalis adhere to. this precipitate with great tenacity and the precipitate obtained from the use of the fixed alkali carbonates invariably contains traces of a carbonate. Freshly precipitated barium carbonate completely precipitates ferric salts in the cold as ferric 126, 6*. IRON. 15? hydroxide (separation of ferric iron, with aluminum and chromium, from ferrous iron, cobalt, nickel, manganese, and zinc; 2FeCl ;{ + 3BaC0 3 + 3H 2 = 2Fe(OH) 3 + 3BaCl 2 + 3C0 2 ). The mixture should be allowed to stand several hours (chromium precipitates more slowly than aluminum or iron), and, sulphates must be absent, as freshly precipitated barium carbonate reacts with solutions of the sulphates of the fourth group; e. g., NiS0 4 + BaC0 3 = NiC0 3 + BaS0 4 . The reaction takes place most read- ily if the metals be present as chlorides. If the precipitate obtained be treated with an excess of dilute sulphuric acid the ferric hydroxide dis- solves, leaving the excess of barium as the insoluble sulphate. Freshly precipitated carbonates of Ca , Mg , Mn , Zn , and Cu react similar to the barium carbonate. I). Oxalic acid and soluble oxalates precipitate from solutions of ferrous salts, ferrous oxalate, FeC 2 O 4 , yellowish-white, crystalline, sparingly soluble in hot water, soluble in HC1 , HNO 3 and H,SO 4 ; ferric salts are not precipitated by oxalates except as reduction to ferrous oxalate takes place. The acetates, as NaC 2 H 3 2 , form in solutions of ferric salts a dull red * solution of ferric aceiate, Fe(C 2 H 3 2 ) 3 , which upon boiling is decomposed and precipitated as basic ferric acetate of variable composition (separation of iron and aluminum from phosphoric acid (d), chromium, and the metals of the fourth group). The red colored ferric acetate solution is not decolored by mercuric chloride (distinction from Fe(CNS) 3 ). The basic precipitates are soluble in HC1 , HN0 3 and H 2 S0 4 and are transposed by alkali hydroxides. Tannic acid precipitates concentrated solutions of ferrous salts: ferric salts are precipitated as blue-black ferric tannate (the basis of common ink), insoluble in water or acetic acid, very soluble in excess of tannic acid. Ferric salts are completely precipitated by ammonium succinate from hot solutions (Young, J. C., 1880, 37, 674). Both ferrous and ferric salts (not nitrates) slightly acid are completely precipitated by a solution of nitroso B. naphthql (separation from aluminum and chromium) (Knorre, B., 1887, 20, 283; Menicke, Z. uiu.n'ir., 1888, 5). If the Fe'" be in excess of the PO 4 the phosphate will all be pre- cipitated. Hydrochloric acid should be absent, i. e., excess of NaC,H :! 0, should be added (Knorre, Z. angew., 1893, 267). Potassium cyanide gives with solutions of ferrous salts a yellowisK-red pre- cipitate, which dissolves in excess of the reagent to potassium ferrocyanide, K 4 Fe(CN) 8 ; with solutions of ferric salts, ferric hydroxide is precipitated with evolution of hydrocyanic acid (equation (a), page 156). Potassium ferrocyanide precipitates ferrous salts as potassium ferrous ferrocyanide (b), K 2 FeFe(CN) 6 , (Everitt's salt), bluish-white, insoluble in *Meconic acid and formic acid form red solutions with ferric- salts : benzole acid gives a flesh colored precipitate; phenol, creosote, saligenin, and other hydroxy aromatic derivatives give a blue to violet color. Morphine gives a blue color. The following is recommended as a very satisfactory test for a trace of iron in copper sulphate. Dissolve one gram of the CuSO 4 in five cc. of water, add five cc. of a ten per cent. < therial solution of salicylic acid. Jf the layer of eontact assumes a violet color iron is present (Grigge, Z., 1895, 34, 450), 158 IRON. 126, 6*. acids, transposed by alkalis (c). This is converted into Prussian blue (see below), gradually by exposure to the air, immediately by oxidizing agents (d). With ferric salts, ferric ferrocyanide (e), Fe 4 (Fe(CN) 6 ) 3 , Prus- sian blue, is formed, insoluble in acids, decomposed by alkalis (/). If the reagent be added in strong excess the precipitate is partially dissolved to a blue liquid. Strong acids should not be present as they color the re- agent blue. In neutral solutions diluted to one in 500,000 the iron may be detected (Wagner, Z., 1881, 20, 350). The ferrocyanides are transposed by KOH and decomposed by fusion with NaNO ;s and Na,C0 3 , the iron being obtained as Fe 2 3 (Koningh, Z. angew., 1898, 463). Potassium ferri- cyanide precipitates from dilute solutions of ferrous salts ferrous ferri- cyanide (g), Fe 3 (Fe(CN) 6 ) 2 (Turnbull's blue), dark blue, insoluble in acids, transposed by alkali hydroxides (/?); with ferric salts no precipitate is obtained, but the solution is colored brown or green (i). This is a very important reagent for the detection of the presence of even traces of ferrous salts in the presence of ferric salts. As iron is so readily oxidized or reduced by various reagents the original solution should always Be tested. The solutions should also be sufficiently diluted to allow the detection of the precipitate of the ferrous ferricyanide in the presence of the dark colored liquid due to the presence of ferric salts. If no precipi- tate be obtained (indicating absence of ferrous iron) a drop of stannous chloride or some other strong reducing agent constitutes a delicate test for ferric salts and reconfirms the previous absence of ferrous salts. Potassium thiocyanate gives no reaction with ferrous salts; with ferric salts the Uood red ferric thiocyanate, Fe(CNS) 3 (solution),* is formed (/). This constitutes an exceedingly delicate test for iron in the ferric condi- tion (the original solution should always be tested). According to Wagner (Z., 1881, 20, 350) one part of iron" as ferric salt, may be detected in 1,600,000 parts of water. The red salt of ferric thiocyanate is freely soluble in water, alcohol, and ether; it is extracted by ether from aqueous solutions and thus concentrated, increasing the delicacy of the test (Natan- son, A., 1864, 130, 246). The red color of the liquid is destroyed by mercuric chloride (fc), also by phosphates, borates, acetates, oxalates, tar- trates, racemates, malates, citrates, succinates, and the acids of these salts. Nitric and chloric acids give red color with potassium thiocyanate, re- moved by heat. * The quantity of non-dissociated Fe(CMS) 3 , to which the color is due, is increased by an ex- cess of either of the products of the dissociation. The test for iron is therefore more delicate if considerable KCNS is added. The decoloration by HgCl 2 is due to the breaking- up of the Fe(CNS) 3 to form Hg(CNS) a which is even less dissociated in water solution than HgCl ? , 126, M. tltott. 159 (a) FeCl 3 + 3KCN + 3H 2 = Fe(OH) 3 + 3KC1 + 3HCBT (&) PeS0 4 + I^FefCN),. = K,FeFe(CN) + K 2 S0 4 (c) K,FeFe(CN) + 2KOH = Fe(OH), + K 4 Fe(CN) 6 ^K.FeFeCCN), + 0, + ^HCl = Fe 4 (Fe(CN) ) 3 + K 4 Fe(CN) 6 + 4KC1 + 2H 2 4FeCl 3 + 3K,Fe(CN) - Fe 4 (Fe(CN) u ) 3 + 12KC1 Fe 4 (Fe(CN) ) 3 + 12KOH = 4Fe(OH) 3 + 3K 4 Fe(CN) a 3FeS0 4 4- 2K 3 Fe(CN) (j =Fe s (Fe(CN) ), + 3K 2 S0 4 Fe 3 (Fe(CN) 6 ) 2 + GKOH = 3Fe(OH) 2 +'2K 3 Fe(CN) 6 FeCl 3 + K,Fe(CN) G = FeFe(CN) e + 3KC1 FeCl 3 + 3KCNS Fe(CNS)., + 3KC1 (fc) 2Fe(CNS) 3 + 3HgCl, = 3Hg(CNS) 2 + 2FeCl s c. Nitric acid readily oxidizes all ferrous salts to ferric salts, the reac- tion being hastened by the aid of heat. As the iron is reduced to the ferrous condition in the precipitation of the metals of the second group with hydrosulphuric acid, the oxidation with nitric acid is necessary to insure the precipitation of all the iron as hydroxide in the third group (Fe,(S0 4 ) 3 + K,SO 4 + 2MnSO 4 + 8H_,0 . (//) By titration with a standard solution of K,Cr,O 7 , using a solution of K 3 Fe(CN) G as an external indicator: 6FeSO 4 + K.CrA + 7H,SO 4 = 3Fe 2 (S0 4 ) 3 + K,S0 4 + Cr 2 (SO 4 ) 3 + 7H 2 O . (o) As ferric iron, by titration with a standard solution of Na,S,0 3 , using KCNS as an indicator: 2FeCl 3 + 2Na,S,0 3 = 2FeCl> + NaJS^Oe + 2NaCl . A few drops of a solution of CuSO 4 are added, which seems to hasten the reaction and gives more accurate results; or use excess of the N"a 2 SoO 3 and titrate back with standard iodine (Crafts, J. ., 1873, 26, 1162). (6) The iron as ferric salt is treated with an excess of a standard SnCl, solution, the excess of the SnCl, being determined by a standard solution of iodine in potassium iodide: 2FeCl 3 + SnCL = 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-,S 2 3 (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. angeiv., 1894, 669). 10. Oxidation. Metallic iron precipitates the free metals from solu- tions of An , 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 vl , 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, , ILS0 4 (if concentrated and hot), HBrO , HBr0 3 HIO., , also Br , Cl . Br and Cl 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, ether* ammonia, etc. (Resell, J. Am. Soc., 1895, 17, 760). Fe'" is reduced to Fe" by solutions of Sn", Cu', H 3 P0 2 , H 3 PO, , 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, 127. ANALYSIS OF THE IRON GROUP. 163 127. .TABLE FOR ANALYSIS or 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 pre- cipitate (80), and freed from H 2 S by boiling, add a few drops of Nitric Acid and boil an instant (to oxidize the ferrous iron *). Immediately add Ammonium Chloride (134, 56; 189, 56) and an excess (136, 6a) of Ammonium Hydroxide (116). If there is a precipitate, filter and wash. Precipitate: A1(OH) 3| Cr(OH) 3 , Fe(OH) 3 [Mn(OH),] . Pierce the point of the filter, and with a little water wash the precipitate into a casserole or evaporating dish. Add 5 to 10 cc. of NaOH and 10 to 15 cc. of H 2 O 2 or add a gram or two of sodium peroxide. Boil several minutes, dilute and filter. Residue: Fe (OH) 3 , Mn(OH) 3 Dissolve in dilute H 2 SO 4 , adding a few drops of Filtrate: Na 2 CrO 4 , NaAlO 2 Acidify with acetic acid and divide into two parts. H 2 O 2 if necessary. Di- (1) Test for chromium. (2) Test for aluminum. vide in two parts. Add a few drops of lead Add (NH 4 ) 2 CO 3 in ex- (1) Add KCNS. Blood- red color indicates iron /s -tno njLN acetate. The yellow lead chromate indi- cates chromium ( 57, cess and heat. A pre- cipitate is A1(OH) 3 . (, 1^0, DO). 6/i). The same result is ob- (2) Boil with HNO 3 and tained with nearly PbO 2 or PbjO 4 . Red- dish purple solution in- dicates manganese. 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 K 3 Fe(CN) 6 forferrosum ( 126, 66). 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 ;i 2 ) 2 , BaCl 2 , etc. (125, 6h). If the chromium is present as achromic salt,Cr 2 (SO 4 ) 3 , the solution will have a green or bluish-green color and will give the 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). general reactions as de- scribed at 125, 6. Chromates should be re- duced by boiling with HC1 and C 2 H 5 OH be- fore proceeding with the regular course of analysis ( 125, 6/). Study 136, 128, 129, 130, and 131. Study 136, 128, 129, 130, 131. 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). t Ferrous salts, which have been kept in the air, are never wholly free from ferric compounds. 164 DIRECTIONS FOR ANALYSIS WITH NOTES. 128. DIRECTIONS FOR 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 + 4HN0 3 = Fe 2 (SO 4 ) 3 + Fe(N0 3 ) 3 + NO + 2H 2 Add to the solution about one-half its volume of NH 4 C1 (56, 134 and 189) and warm and then add NH 4 OH in a decided excess (135, 6a): MgCl 2 + NH 4 C1 + NH 4 OH = NH 4 MgCl 3 + NH 4 OH Fe 2 (S0 4 ) 3 + 6NH 4 OH = 2Fe(OH), + 3(NH 4 ) 2 SO 4 ZnS0 4 + 4NH 4 OH = (NH 4 ) 2 Zn0 2 + (NH 4 ) 2 S0 4 + 2H 2 O 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. (1) 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, 6#), which tends to obscure the reactions following: 6H 2 S -4- 4HNO 3 = 3S 2 -4- 4NO + 8H 2 O. Also any H 2 S not decomposed by the HNO 3 would cause a precipitate of the sulphides of the fourth group upon the addition of the NH 4 OH: H 2 S -4- NiCl 2 + 2NH 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,S (126, 6e) : 4FeCl 3 + 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 HNO 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. (J) Ammonium hydroxide precipitates a portion of Mn (134, 6) and Mg (189, 60), but these hydroxides are soluble in NH 4 C1 (5c, 134 and 189): hence if that reagent be added in excess the Mg k not at all precipitated by the NH 4 OH, but the manganese is oxidized to Mn(OH) a and precipitated more or less completely ( 134, 6a): 2MnCl 2 + 2NH 4 OH = Mn(OH) 2 + (NH 4 ) 2 MnCl 4 Mn(OH) 2 + 4NH 4 C1 = (NH 4 ) 2 MnCl4 + 2NH 4 OH 2(NH 4 ) 2 MnCl 4 + 4NH 4 OH + O + H 2 O = 2Mn(OH) 3 + 8NH 4 C1 2MgCl 2 + 2NH 4 OH = Mg(OH) 2 + NH 4 MgCl 3 + HN 4 C1 Mg(OH) 2 + 3NH 4 C1 = NH 4 MgCl 3 + 2NH 4 OH Barium, calcium, and strontium are also precipitated as carbonate because the alkaline solution absorbs carbon dioxide from the air. As much as 15 mg. of barium may be present in this precipitate. (Curtman and Frankel, J. Am. Soc., 33, 724, 1911.) For detection of the barium see 186, 8. (6) Ammonium chloride lessens the solubility of A1(OH) 3 in the NH 4 OH solution .and effects &n almost quantitative precipitation of that metal (117), 130. DIRECTIONS FOR ANALYSIS WITH NOTES. 165 (7) NH 4 OH precipitates solutions of Co , Ni and Zn , but these precipitates are readily soluble in an excess of the NH 4 OH (116). To insure the presence of an excess of NH 4 OH the odor should be noted after shaking the test tube and after the solution has been heated. (8) The precipitates of the hydroxides of Al , Cr and Fe'" filter much more rapidly if the precipitation takes place from a hot solution (124, 4 and 6a). (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 filtrate 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 , Fe'" and Mn'" 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 a,n amount of water as possible. Add NaOH and H 2 2 and boil for a minute or two or add Na 2 2 in small portions and boil until the chromium is completely oxidized as indicated by the yellow color. 2Cr(OH) 3 + 3Na 2 O 2 = 2Na>Cr(>4 + 2H 2 O + 2NaOH A1(OH) 3 + NaOH = NaAlO 2 + 2H,O The alkaline solution is filtered (the filtrate is reserved to be tested for aluminum) and the remaining precipitate dissolved in dilute HC1 or H 2 S0 4 with the addition of a few drops of H 2 2 if necessary to dissolve the manganese hydroxide. Fe 2 O 3 + 3H 2 O 4 = Fe 2 (SO 4 ) 8 + 3H 2 O 2Mn(OH) 3 + 2H 2 SO 4 + H 2 O 2 = 2MnSO 4 + 6H 2 O + O 2 The iron is tested for by adding a few drops of ammonium or potassium sulphocyanate : Fe 2 (SO 4 ) 3 + 6KCNS = 2Fe(CNS) s + 3K 2 SO 4 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, 66) 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, 66) : 3FeSO 4 + 2K 3 Fe(CN) 6 - Fe 3 [Fe(CN) 6 ] 2 4- 3K 2 SO 4 Manganese is tested for by boiling a portion of the solution with nitric 166 DIRECTIONS FOR ANALYSIS WITH NOTES. 131, L acid and Pb0 2 or Pb 3 4 .* If manganese is present a solution of the reddish purple permanganic acid is obtained. 2MnSO 4 + 5PbO 2 + 6HNO S = 2HMnO 4 + 2PbSO 4 + 3Pb(NO 3 ) 2 + 2H 2 The alkaline filtrate obtained after boiling the precipitated hydroxides with NaOH and H 2 2 is acidified with acetic acid and one portion tested for chromium with Pb(C 2 H 3 2 ) 2 a yellow precipitate indicating chromium. Na2CrO 4 + Pb(C 2 K 3 O 2 ) 2 = PbCrO 4 + 2NaC 2 H 3 O 2 (57, 6h) Another portion is tested for aluminum by adding (NH 4 ) 2 C0 3 in excess and warming: NaA10 2 + 4HC 2 H 3 2 = A1(C 2 H 3 O 2 ) 3 + NaC 2 H 3 O 2 2A1(C 2 H 3 O 2 ) 3 -f 3(NH 4 ) 2 CO 3 + 3H 2 O = 2A1(OH) 3 + 6NH 4 (C 2 H 3 O 2 ) + 3CO 2 a white gelatinous precipitate being evidence of the presence of aluminum. A similar precipitate may also be obtained by adding excess of NH 4 C1 to the alkaline solution: 2KA1O 2 + 2NH4C1 + H 2 O = A1 2 O(OH) 4 + 2KC1 + 2NH 3 (124, 6a) 131. Notes. (1) Chromium hydroxide is oxidized by an alkaline oxidizing agent to a chromate. Sodium peroxide may be used or hydrogen peroxide and caustic soda which unite to form sodium peroxide. The same result may be obtained by fusing chromium hydroxide on a platinum foil with a mixture of equal parts of KNO 3 and 2Cr(OH) 3 + 3KN0 3 + 2Na 2 CO 3 = 2Na 2 CrO 4 + 3KNO 2 + 2CO 2 + 3H 2 O . Aluminum is converted into sodium aluminate 2A1(OH) 3 + Na 2 CO 3 = 2NaAlO 2 + 3H 2 O + CO 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. (3) 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. (4) Manganese remains undissolved when the mixed hydroxides are treated with sodium peroxide. When they are fused with KNOs and Na^COs the mangan- ese is converted into the green manganate. ( 134, 7). This dark green color is an excellent test for manganese which is almost invariably present with the hydroxides if it is in the solution of the unknown ( 134, 6a). On treating the fused mass with water a green solution is obtained which turns purple on stand- ing or cautiously acidifying with acetic acid. () K 2 Mn0 4 + 8HC1 = MnCl 2 + 2KC1 + 2C1 2 + 4H 2 O (6) 2K 2 CrO 4 + 10HC1 + 30^0 = 2CrCl 3 + 4KC1 + 3C 2 H 4 O + 8H 2 O (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 bailing with HC1 and * As PbOz and PbsCh frequently contain traces of manganese, the samples used when testing for manganese must first be tested as follows: Boil the oxide with dilute nitric acid and allow to settle. If the liquid is red, manganese is present. A pure sample of oxide must be obtained or the manganese removed by repeated boiling with nitric acid. 132, 4. COBALT. 167 C 2 H 6 O before proceeding with the regular group separations ( 126, 6e and O H 2 S will effect this reduction but gives also ;i precipitate of sulphur which should be avoided when convenient to do so: :2K Cr O 7 -f- lOHCl + r 'H,S = ) Also reduced by carbon in various ways. 4. Oxides and Hydroxides. Cobaltous oxide, CoO , is made (1) by heating any of its oxides or hydroxides in hydrogen to (not above) .350; (2) by ignition of Co(OH) 2 or CoCO 3 , air being excluded; (3) by heating Co 3 O 4 to redness in a, stream of CO, (Russell, J. C., 1863, 16, 51); (//) by heating any of the higher oxides to a white heat (Moissan, A. Ch., 1880, (5), 21, 242). CobaltoilS lii/ , Co 3 O 3 (OH)o. Cohdltic o.rifJc, Co 2 3 , is made by heating* the nitrate just hot enough for de- composition, but not hot enough to form Co 3 O 4 . CobaltlC lii/<]r<>.rid<\ Co(OH) 3 , is made by treating any cobaltous salt with Cl , HC1O , Br or I in presence of a fixed alkali or alkali carbonate. It dissolves in HC1 with evolution of chlo- rine, in H SO, with evolution of oxygen, forming a cobaltous salt. CoO 2 has not yet been isolated, but McConnell and Hanes (J. C., 1897, 71, 584) have -ho\\u that it exists as H a CoO, and in certain cobaltitee. 168 COBALT. 132, 5rt. 5. Solubilities. a. Metal. Slowly soluble on warming in dilute HC1 or H 2 SO 4 , more rapidly in HNO 3 , not oxidized on exposure to the air or when heated in contact with alkalis. Like iron, it may exist in a passive form (Nickles, J. pr., 1854, 61, 168; St. Edme, C. r., 1889, 109, 304). With the halogens it forms cobaltous compounds (Hartley, J. (7., 1874, 27, 501). 6. Oxides and 7i/dro,rM7es. 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 O 4 is said to be soluble in acids with great diffi- culty (Gibbs and Geuth, Am. 8., 1857, (2), 23, 257). cSnlts. Cobalt forms two classes of salts: coMltous, derived from CoO , and cobaltic, from Co 2 O 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) (Pinerua, C. r., 1897, 124, 862). Most of the salts insoluble in water form soluble compounds with ammonium hydroxide. 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(OH) 2 . The cobaltous hydroxide is not soluble in excess of the reagent, but is somewhat soluble in a hot concentrated solution of KOH or NaOH (dis- tinction from Ni) Reichel, Z., 1880, 19, 468). The cold solution is blue, the color being more intense when hot and especially if a little glycerine is present. On standing it turns green, then red, the changes being more rapid if a little hydrogen peroxide solution is added. Copper also gives a blue solution but does not give the other color changes. Donath, Z., 40, 137. Freshly precipitated Pb(OH) 2? Zn(OH) 2 , and HgO pre- cipitate Co (OH) 2 from solutions of cobaltous salts at 100. Ammo- nium hydroxide causes the same precipitate as the fixed alkalis; ncomplete, even at first, because of the ammonium snlt 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 8132, 6J. COBALT. 169 precipitate with ammonium hydroxide produces, in different conditions, different soluble compounds noted for their bright colors, as (NH 3 ) 4 CoCl 2 , (NH 3 ) n CoCl 2 , (NH s ) 4 CoCl 3 , etc. Alkali carbonates precipitate cobaltous basic-carbonate, Co 8 5 (C0 3 ) 3 , peach-red, which when boiled loses carbonic anhydride and acquires a violet, or, if the reagent be in excess, a blue color. The precipitate is soluble in ammonium carbonate and very slightly soluble in fixed alkali carbonates. Carbonates of Ba , Sr , Ca , or Mg do not precipitate cobaltous chloride or nitrate in the cold (separation from Fe'", Al, and Cr'"), but by prolonged boiling they precipitate them completely. However, if a solution of a cobaltous salt be treated with chlorine, a cobaltic salt is formed (5a), which is precipitated in the cold on digestion with BaC0 3 (distinction from Ni). 6. Oxalic acid and oxalates precipitate reddish-white cobaltous oxalate, CoC 2 O 4 , soluble in mineral acids and in ammonium hydroxide. Alkali cyanides as KCN precipitate the brownish-white cobaltous cyanide, Co(CN) 2 , soluble in hydrochloric acid, not in acetic or in hydro- cyanic acid, soluble in excess of the reagent, as double cyanides of cobalt and alkali metals (KCN) 2 Co(CN) 2 potassium cobaltous cyanide, the solu- tion having a brown color: CoCl 2 + 2KCN = Co(CN) 2 + 2KC1 . Then Co(CN) 2 + 2KCN = (KCN),Co(CN) 2 . Dilute acids, without digestion, reprecipitate cobaltous cyanide from this solution (the same as with Ni): (KCN) 2 Co(CN) 2 + 2HC1 = Co(CN) 2 + 2HCN + 2KC1 . But if the solu- tion, with excess of the alkali cyanide and with a drop or two of Jiydro- chloric acid,* insuring free HCN , be now digested hot for some time, the cobaltous cyanide is oxidized and converted into alkali cobalticyanide as K 3 Co(CN) 6 corresponding to ferricyanides, but having no corresponding nickel compound: 4Co(CN) 2 + 4HCN + O 2 = 4Co(CN) 3 (cobaltic cyanide) + 2H 2 O Co(CN) 3 + 3KCN = K 3 Co(CN) 6 (potassium cobalticyanide). In the latter solution neither alkalies nor acids precipitate the cobalt (important distinction from nickel, whose solution remains (KCN) 2 Ni(CN) 2 , and which, after digestion as above is precipitated). The potassium cobalticyanide solution, after removal of the Ni , may bo precipitated with HgN0 3 (Gibbs, J. C., 1874, 27, 92). The oxidation of the cobalt may be hastened by the presence of chromic acid, whirl i is * Moore (C. N., 1887, 56, 3) adds glacial phosphoric acid to the neutral solutions of cobalt and nickel, until the precipitate first formed begins to redissolve ; then he adds KCN and boils, continuing the boiling- and addition of KCN until K.OH fails to give a precipitate. He then warms with excess of bromine in presence of KOH, whereupon the nickel is completely pre- cipitated leaving the cobalt in solution. See also Hambly (C. N.< 1893, 65, 289). 170 COBALT. 132, 68. reduced to trivalent chromium compound: 6Co(CN) 2 + 24KCN + 2Cr0 3 + 3H 2 = = 6K 3 Co(CN) 6 + Cr 2 3 + C>KOH (McCulloch, C. N., 1889, 59, 51), also by means of NaCIO or NaBrO produced by passing chlorine into caustic soda or adding bromide to the alkali. Ferrocyanides, as K 4 Fe(CN) fi , precipitate colaltoits fcn'ocyuiiirtc, Co 2 Fe(CN) 6 , gray-green, insoluble in acids. Ferricyanides, as K 3 Fe(CN) 6 , precipitate cobdU- (W# ferricyanide, Co 3 (Fe(CN) ) 2 , brownish-red, insoluble in acids. But a more distinctive test is made by adding- ammonium cMwide and lij/diwidc, with the ferricyanide, when a blood-red color is obtained, in evidence of cobalt (distinc- tion from nickel). Potassium xanthate forms a green precipitate in neutral or slightly acid solutions of cobalt salts (133, 66). Nitroso-^-naphthol completely precipitates solutions of Cu , Fe , and Co ; Ag , Sn , and Bi salts are partially precipitated; and Pb , Hg , As , Sb , Cd . Al , Cr , Mn , Ni , Zn , Ca , Mg , and Gl remain in solution (Burgass, Z. angew.y 1896, 596). In analysis for the separation of cobalt and nickel it is recommended to proceed as follows : The solution of the metals preferably as sulphates or chlorides is acidulated with hydrochloric acid and treated with a hot solution of nitroso-/3-naphthol in 50 per cent acetic acid, until the whole of the cobalt is precipitated. The brick-red precipitate is then washed with cold HC1 , then with hot 12 per cent HC1 , and finally with water. The separation is quantitative. The precipitate may be ignited in air to the oxide or with oxalic acid in an atmosphere of hydrogen and weighed as the metal. For qualitative purposes the cobalt in the precipitate may be identified by the color of the borax bead (7). The nickel in the filtrate may be precipitated by hydrosulphuric acid and identified by the usual tests (Knorre, B., 1887, 20, 283 and Z. angeiv., 1893, 264). Ammonium TMocyanate in concentrated solutions of cobalt produces a brilliant blue color which disappears on dilution with water. On the ad- dition of amyl alcohol and ether in equal portions and shaking, the layer of alcohol and ether becomes blue. The test is very delicate especially if the cobalt solution is concentrated. Nickel salts produce no coloration of the amyl alcohol and ether, but iron interferes because the red Fe (CNS) 3 is dissolved by the alcohol. By the addition of 2 or 3 c.c. of concentrated* ammonium acetate solution and 2 or 3 drops of 50 per cent tartaric acid solution, the red color of the Fe (CNS) 3 may be removed. The blue color is probably due to the undissociated salt (NHJ 2 [Co(CNS) 4 ] which is soluble in ether and amyl alcohol (Vogel Ber. 12, 2314. Treadwell, Z. An. Hi. 26 (1901), 105). Ammonium thiocyanate may be employed for the separation of nickel and cobalt (Rosenheim. and Cohen, B., 33, 111, and Rosenheim and Huld- shinsky, B. f 34, 2050), 12 grams of NH 4 CNS are added to the nitric acid solution of the metals the volume being 15-20 c,c, Cobalt produces a deep 132, 6>. COBALT. 171 blue compound of the formula R 2 Co(CNS) 4 in which R is an alkali metal. This blue compound may be extracted by a mixture of 25 vol. ether and 1 vol. amyl alcohol. c. Potassium nitrite forms with both cobaltous and nickelous salts the double nitrites, Co(N0 2 ) 2 .2KNO, and Ni(N0 2 ),.2KN0 2 , soluble. The nickel compound is very stable, but if the cobalt compound, strongly acidulated with acetic acid, be warmed and allowed to stand for some time, preferably twenty-four hours; the cobalt is completely precipitated as the yellow crystalline potassium cobaltic nitrite, Co(N0 2 ) 3 .3KH0 2 (separation from Ni) : CoCl, + GKNO., + HC 2 H 3 2 + HN0 2 = Co(N0 2 ) 3 .3KN0 2 + 2KC1 + KC 2 H 3 6 2 + H 2 + NO. d. Phosphates, as Na 2 HP0 4 , precipitate cobaltous salts as the reddish cobalt on t> i>Ito^i>Jiatc, CoHPO 4 , soluble in acids and in ammonium hydroxide. Sodium pyrophosphate forms a gelatinous precipitate with solutions of cobalt salts, soluble in excess of the reagent. The addition of acetic acid causes a precipitation of the cobalt even in the presence of tartrates (separation from Ni, but not from Mn or Fe) (Vortmann, B., 1888, 21, 1103). If a solution of cobaltous salt be treated with a saturated solution of ammonium phosphate and hydrochloric acid, and when hot treated with an excess of ammonium hydroxide, a bluish precipitate of CoNH 4 PO 4 will appear on stirring (separa- tion from nickel*) (Clark, C. N., 1883, 48, 262; Hope, J. Soc. Ind., 1890, 9, 375). e. Hydrosulphuric acid, with normal cobaltous salts, gradually and imperfectly precipitates the black cobalt sulphide, CoS ; from cobalt acetate, the precipitation is more prompt, and is complete; but in presence of mineral acids, as in the second group precipitation, no precipitate is made. Immediate precipitation takes place with hydrosulphuric acid acting upon solutions of cobaltous salts in ammonium hydroxide. When formed, the precipitate is scarcely at all soluble in dilute hydrochloric acid or in acetic acid; slowly soluble in -moderately concentrated hydrochloric acid; readily soluble in nitric acid; and most easily in nitrohydrochloric acid. By exposure to the air, the recent cobaltous sulphide is gradually oxidized to cobalt sulphate, soluble, as occurs with iron sulphide (126, 60). Alkali sulphides precipitate immediately and perfectly the black cobaltous sul- phide, described above, insoluble in excess of the reagent. When cobaltous salts are boiled with sodium thiosulphate a portion of the cobalt is precipi- tated as the black sulphide. f. The higher oxides of cobalt and cobaltic salts are reduced by warming with halogen acids, liberating the corresponding halogens (HC1 does not reduce the cobalt in K,Co(CN) 8 ). g. Soluble arsenites and arsenates precipitate cobaltous salts, forming the corresponding cobaJt iirxenites or arsenates, bluish-white, soluble in ammonium hydroxide or in acids, including 1 arsenic acid. //. Soluble chromates precipi- tate colmltous eliminate, yellowish-brown, soluble in ammonium hydroxide and * Krauss (Z., 1891, 3O, 227) gives a good review of the most important methods for the separa- tion of cobalt and nickel. COBALT. 132, 6/. in acids, including chromic acid. No precipitate is formed with potassium dichromate. i. KMnO 4 added to an annnoniaeal solution of eobaltous salts oxidizes the cobalt and prevents its precipitation by KOH (separation from Ni) (Delvaux, G. r., 1881, 92, 723). ;'. Cobaltous salts in ammoniacal solution, warmed with H 2 O 2 and then rendered acid with acetic acid, are precipitated by ammonium molybdate (separation from Ni) (Carnot, C. r., 1889, 109, 109). 7. Ignition, In the bead of borax., and in that of microcosmic salt, with oxidizing and with reducing flames, cobalt gives an intense blue color. The blue bead of copper changes to brown in the reducing flame. If strongly saturated, the bead may appear black from intensity of color, but will give a blue powder. This important test is* most delicate with the borax bead. Manganese, copper, nickel, or iron interfere somewhat. By ignition, with sodium carbonate on charcoal or with the reducing flame, compounds of cobalt are reduced to the metal (magnetic). Cobaltous oxide dissolves in melted glass and in other vitreous substances, coloring the mass blue used to cut off the light of yellow flames (205, 7). The black cobaltoso-cobaltic oxide, Co 3 4 , as left by ignition of eobaltous oxide or nitrate, combines or mixes, by ignition, with zinc oxide from zinc com- pounds to form a green mass, with aluminum compounds a blue, and with magnesium compounds a pink mass. 8. Detection. After removal of the metals of the first three groups cobalt is precipitated by H 2 S in ammoniacal solution with Ni , Mn and Zn . The sulphides are digested with cold dilute HC1 which dissolves the Mn and Zn . The borax bead test (7) is now made upon the remaining black precipitate, and if Ni be not present in great excess * the characteristic blue bead is obtained. If the nickel be present in such quantities as to obscure the blue borax bead the sulphides are dissolved in hot cone. HC1, using a few drops of HN0 3 . The solution is heated to decompose all the nitric acid and, after dilution, the cobalt is precipitated with nitroso-/?-naphthol, according to directions given in 6&, and further identified by the bead test. The NH 4 CNS test (132, 6&) and the NaHC0 3 and H 2 2 test (140) may also be obtained. 9. Estimation. (1) As metallic cobalt, all compounds that may be reduced by ignition in hydrogen gas, e. g. CoCl 2 , Co(NO 3 )2, CoCO 3 , and all oxides and hydroxides. (2} As UoD , all soluble cobalt salts, all salts whose acids are expelled or destroyed by ignition, all oxides and hydroxides. The salt is converted into Co(OH) 2 by precipitation with a fixed alkali, and ignited in a stream of CO 2 . The carbonate and nitrate may be ignited directly in CO 2 , and organic salts are first igni ed in the air until the carbon is oxidized, and then again ignited in CO 2 . (3) After converting into a sulphate it is ignited at a dull-red heat and weighed as a sulphate. (4) After converting into the oxalate, titrated with KMnCV (5) In presence of nickel, it is oxidized in alkaline solution by H 2 O 2 , KI and HCl are added, and the liberated iodine titrated with sodium thiosulphate (Fischer, C. C., 1889, 116). (6) Electrolytically. (7) Separated * If more than ten parts of nickel are present to one part of cobalt, the characteristic blue bead is not obtained. 133, 5b. NICKEL. 173 from nickel by nitroso-/3-naphthol, and after ignition in hydrogen weighed as the metal (66). 10. Oxidation. Co" is oxidized to Co'" in presence of a fixed alkali by Pb0 2 , Cl , KC10 , Br , KBrO , I and H 2 2 * ; in presence of acetic acid by KN0 2 (Gc). Co'" is reduced to Co" by H 2 C 2 4 , H,P0 2 , H 2 S , H 2 S0 3 , HC1 , HBr , and HI . Metallic cobalt is precipitated from solution of CoCl 2 by Zn , Cd . and Mg . 133. Niekel. Ni = 58.70 . Usual valence two and three. 1. Properties. Specific gravity, 8.9 (Rchroeder, Pogff., 1850, 106, 226). Melting point, 1452 (Cir. B. S., 36, 1915). It is a hard white metal, capable of taking a high polish; malleable, ductile and very tenacious, forming wire stronger than iron but not quite so strong as cobalt ( 132, 1). It does not oxidize in dry or moist air at ordinary temperatures. It is magnetic but loses its magnetism like steel on heating to redness (Gangain, C. r., 1876, 83, 661). It burns with incandescence when heated in O , Cl , Br , or S . It is much used in plating other metals, in making coins of small denominations, in hardening armor plate, pro- jectiles, 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) ; pentlandite (FeNi)S ; niccolite (NiAs) ; garnierite (variable, perhaps H 2 (NiMg)SiO 4 + xH 2 O) ; frequently in pyrrhotite, (Fe n S n +i with Ni) and in numerous other rarer minerals. 3. Preparation. (1) By electrolysis. (2) By heating- jn a stream of hydrogen. The oxide is reduced in this manner a't 270 (W. Miiller, Pogg., 1869, 136, 51). (3) By fusing- the oxalate under powdered glass (C0 2 being given off). (4) Reduction by igniting in CO . (5) Reduction by fusing with carbon in a variety of methods. (6) By heating the carbonyl, f Ni(CO) 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. Nickclic oxide, Ni 2 3 , is made from NiC0 3 , Ni(NO 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) 3 , is formed by treating nickelous salts first with a fixed alkali hydroxide or carbonate and then with Cl , NaCIO , Br or NaBrO (not formed by iodine), a black powder forming no corresponding salts (Campbell and Trowbridge, J. Anal., 1893, 7, 301). A trinickelic tetroxide, Ni 3 O 4 , magnetic (corresponding to Co 3 O 4 , Fe 3 O 4 , Mn 3 O 4 and Pb 3 O 4 ), is formed, according to Baubigny (C. r., 1878, 87, 1082), by heating NiCL in oxygen gas at from 350 to 440; and by heating Ni 2 O 3 in hydrogen at 190 (Moissan, A. Ch., 1880, (5), 21, 199). 5. Solubilities. a. Metal. Hydrochloric or sulphuric acid, dilute or con- centrated, attacks nickel but slowly (Tissier, C. r., 1860, 50, 106); dilute nitric acid dissolves it readily, while towards concentrated nitric acid it acts very similar to passive iron (Deville, C. r., 1854, 38, 284). It is not attacked when heated in contact with the alkali hydroxides or carbonates, ft. Oxides and * Durant, C. N., 1897, 75, 43. t Nickel carbonyl is prepared by heating the nickel ore in a current of CO . It is a liquid, sp. gr. 1.3185, boiling at 43 and freezing at 25. When heated to 200 it is decomposed into Ni #nd CO (Berthelot, C, r,, 1891, 112, 1343; 113, 679; Mond, J. Soc, Ind., 1892. 11, 750). 174 NICKEL. 133, firt. hydroxides. Nickelous oxide and hydroxide are insoluble in water or fixed alkalis, soluble in ammonium hydroxide and in acids. Nickelie 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 Cl , 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 NiSO 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(OH) 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 will slowly precipitate nickel hydroxide from the ammoniacal solutions (distinction from cobalt). Alkali carbonates precipitate green 'basic nickelous carbonate, Ni 5 (OH) 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. ft. Oxalic acid and oxalates precipitate, very slowly but almost completely, after twenty-four hours, nickel oxalate, green. Alkali cyanides, as KCN , pre- cipitate nickel cyanide, Ni(CN) 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 Ni(CN) 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 ;he precipitate of Ni(CN) 2 . If an oxidizing agent such as NaCIO or NaBrO be added to the alkaline solution of the double cyanide, the nickel will be oxidized and precipitated (separation from cobalt), as Ni 2 O 3 .3H 2 O , according to the following equation: K 2 Ni(CN) 4 + NaBrO + NaOH = Ni 2 O 3 3H 2 O Ferrocyanides, as K 4 Fe(CN) 6 , precipitate a greenish-white nickel ferrocyanide, Ni-Fe(CN) 6 , 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), 133, 7. NICKEL. 175 A solution of nitre ferricyanide precipitates solutions of cobalt and nickel salts, the latter being soluble in dilute ammonium hydroxide (Cavalli, Gazzetta, 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 ),S (Phipson, C. N., 1877, 36, 150). The xanthate also precipitates nickel in alkaline solution in presence of Na 4 P,0 T (a separation from Fe'") (Campbell and Andrews, /. Am. Soc., 1895, 17, 125). Xirl-rl salts are not precipitated by an acetic acid solution of nitroso-/?- 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 HPO 4 , pre- cipitates nickel phosphate, Ni.,(P0 4 ) 2 , greenish-white. 0- Hydrosulphuric acid precipitates from neutral solutions of nickel salts a portion of the nickel as nickel sulphide, black (Baubigny, C.r.. 1882, 94, 1183; 95, 34). The precipitation takes place slowly, and from nickel- ous acetate is complete. In the presence of mineral acids no precipita- tion takes place. Alkali sulphides precipitate the whole of the nickel, as the black sulphide. Although precipitation is prevented by free acids, the precipitate, once formed, is nearly insoluble in acetic or in dilute hydrochloric acids; slowly dissolved by concentrated hydrochloric acid, readily hy nitric or nitro-hydrochloric. Nickel sulphide, NiS , is partially soluble in yellow ammonium sulphide,* from which brown-colored solution it is precipitated (gray, black mixed with sulphur) on addition of acetic acid (distinction from cobalt). It is insoluble in the monosulphide and is completely precipitated by passing hydrogen sulphide through an ammoniacal solution in the absence of air or oxidizing agents (Noyes, Bray and Spear, J. Am. Soc., 30, 497). Freshly precipitated nickel sulphide is soluble in KCN and reprecipitated as Ni(CN) 2 on adding HC1 or H 2 SO 4 (sep- aration from cobalt) (Guyard, Bl., 1876, (2), 25, 509). When nickel salts are boiled with a solution of NfeSzOs, a portion of the nickel is precipitated as the black sulphide. f- The halogen acids reduce the higher oxides of nickel to nickelous salts with liberation of the corresponding halogen. Potassium iodide added to freshly precipitated nickelic hydroxide gives free iodine (distinc- tion from cobalt). 9- Nickel salts are precipitated by arsenites and arsenates, white or green- ish-white, soluble in acids, including arsenic acid. h. Potassium chromate precipitates basic nickel chromate, yellow, soluble in acids, including chromic acid (Schmidt, A., 1870, 156, 19). K,Cr 2 O T forms no precipitate. 7. Ignition. Xickel compounds dissolve clear in the borax bead, giving 1 with the oxidizing flame a purple-red or violet color while hot, becoming yellowish- brown when cold; with the reducing 1 flame, fading- to a turbid gray, from reduced metallic nickel, and finally becoming- colorless. The addition of any potassium salt, as potassium nitrate, causes the borax bead to take a dark purple or blue color, clearest in the oxidizing- flame. With microcosmic salt, * Hare (J". Am. Soc., 1895, 17, 537) adds tartaric acid to the solutions of nickel and cobalt, and an excess of sodium hydroxide. He then passes in IF 2 S. The cobalt is completely precipitated while the nickel remains in solution, and can be precipitated upon acidulating- the nitrate. ft f*HI Lf-BK 1?6 X1CK&L. 133, 8. nickel gives a reddish-brown bead, cooling to a pale reddish-yellow, the rob- being alike in both flames. Hence, with this reagent, in the reducing flame, the color of nickel may be recognized in presence of iron and manganese, which are colorless in the reducing flame; but colxilt effectually obscures the bead test for nickel. The yellow-red of copper in the reducing flame, persisting in beads of microcosmic salt, also masks the bead test for nickel. By ignition with sodium carbonate on charcoal, compounds of nickel are reduced to the metal, slightly attracted by the magnet. 8. Detection. We proceed exactly as with cobalt for the nitroso-/3 naphthol precipitation. The Ni remains in the filtrate and can be precipi tated with, H 2 S (after neutralizing with NH 4 OH), and its presence con- firmed by the usual tests. Or dissolve the sulphides of Ni and Co io HN0 3 , evaporate nearly to dryness, add an excess of KOH or Na 2 C0 3 , boil, add bromine water and boil to complete oxidation of the Co and Ni , filter, wash thoroughly with hot water and add hot solution of KI to the precipitate on the filter paper. Free iodine (test with CS 2 ) is evidence of the presence of nickel. Nickel may also be detected in the presence of cobalt as follows : Dissolve the sulphides of Ni and Co in aqua regia. boll out the excess of chlorine, neutralizing with KOH and add KCN in slight excess. Add NaCIO or NaBrO and warm. The nickel is percipitated as the brown hvdrated oxide Ni,0 3 .3H.,0 . Dimethylglyoxime produces a red crystalline precipitate which forms the most sensitive test known for detecting nickel in the presence of cobalt. The reagent is prepared by dissolving 1 gram of the dimethyglyoxime in 100 c.c. of 98% alcohol. The solution should be made slightly alkaline with ammonia and boiled after adding a few drops of the reagent. The following reactions take place : 2(CH 3 ) 2 C 2 N 2 O 2 H 2 +NiCl 2 -f 2NH 4 OH = 2NH 4 C1 + [(CH 3 ) 2 C 2 N 2 O 2 H] 2 Ni + 2H 2 O If the amount of nickel is small the solution at first becomes yellow and on cooling deposits red needles. The test is sensitive to nickel when present in 400,000 parts of water (L. Tschugaeff Ber. 38 (1905), 2520). Ten times as much cobalt may be present but in the presence of larger quantities of cobalt the following procedure is followed. Excess of am- monia is added to the cobalt solution and then a few cubic centimeters of hydrogen peroxide. The solution is boiled to decompose the excess of hydrogen peroxide. Dime thy Iglyoxime is added and the solution again boiled. If a small amount of nickel is present a red scum is formed on the section and red crystals form on the sides of the beaker. On filtering the solution the red precipitate L; readily observed on the filter paper. 9. Estimation. * (1} Nickel hydroxide, oxide, carbonate or nitrate is ignited at a white heat and weighed as NiO . (2} It is converted into the sulphate and * Goulal (Z. angew., 1898, 177) gives a summary of the methods proposed for the volu- metric estimation of nickel. 134, 4. MANGANESE. 177 deposited on platinum as the free metal by the electric current. (-?) Volu- metrically. By lilration In a slightly alkaline solution with KCN , using- a small amount of freshly precipitated Agl as an indicator (Campbell and Andrews, J. Am. ftoc., 1895, 17, 127). 10. Oxidation. Ni" is changed to Hi'" in presence of fixed alkalis b} Cl , NaCIO , Br , and NaBrO (not by I , distinction from cobalt, Donatti, B., 1879, 12, 1868). Ni'" is reduced to Ni" by all non-reducing acids with evolution of oxygen; by reducing acids, H,C,,0 4 is oxidized to C0 2 , HNO ? to HN0 3 , H 3 P0 2 to H 3 P0 4 , H 2 S to S , H,S0 3 to H 2 S0 4 , HC1 to Cl , HBr to Br , HI to I , HCNS to HCN and H 2 S0 4 , H 4 Fe(CN) 6 to H 3 Fe(CN) 6 . Ni" is reduced to the metal by finely divided Zn , Cd , and Sn . 134. Manganese. Mn = 54.93. Valence two, three, four, six and seven. 1. Properties. Specific gravity, 7.392 (Glatzel, Ber., 1889). Melting point, 1260 (C'ir. B. S., 35, 1915). Boiling point, 1900 (v. d. Weyde, Ber. 44, 1879). It is a brittle metal, having the general appearance of cast iron, non-magnetic, takes a high polish. According to Deville it has a reddish appearance. It is readily oxidized, decomposing water at but little above the ordinary temperature (Deville, A. Ch., 1856, (3), 46, 199). It is used largely as ferromanganese in the manufacture of Bessemer steel. Oxides and hydroxides of manganese exist as dyad, triad, and tetrad; the salts exist most commonly as the dyad with some unstable triad and tetrad salts; as an acid it is a hexad in manganates and a heptad in permanganates. 2. Occurrence. Not found native. It accompanies nearly HI iron ores. Its chief ore is pyrolusite (MnO-2) . It is Hso found as braunite, (8Mn>O3.MnSiO ;i ) ; hausmannite, (Mn;iO 4 ) ; manganite, (Mn 2 O3.H 2 O) ; rhodocrosite, MnCO ;i ; ala- bandite, (MnS) ; and as a constituent of many other minerals. 3. Preparation. (/) By electrolysis of the chloride. (2) By reduction with metallic sodium or magnesium (Glatzel, B., 1889, 22, 2857). (3) By reduction with some form of carbon. It has not been reduced by hydrogen. (4) By ignition with aluminum (Goldschmidt, A., 1898, 301, 19). 4. Oxides and Hydroxides. (a) Mangunous oxide, MnO , represents the only base capable of forming stable manganese salts. It is formed (1) by simple ignition of Mn(OH) 2 , MnC0 3 or MnC,0 4 , air being excluded; (2) by ignition of any of the higher oxides of manganese with hydrogen in a closed tube (Moissan, A. Ch., 1880, (5), 21, 199). If prepared at as low a temperature as practicable, it is a dark gray or greenish-gray powder, and oxidizes quickly in the air to Mn 3 4 . If prepared at a higher heat it is more stable. Man- ganous hydroxide, Mn(OH) 2 , is formed from manganous salts by precipita- tion with alkalis." It quickly oxidizes in the air, forming MnO (OH), thus changing from white to brown. (?>) Manganic oxide, Mn 2 O 3 , is formed by heating any of the oxides or hydroxides to a red heat in oxygen gas or in air (Schnieder, Pogg., 1859, 107, 605). Mang-anic oxide-hydroxide, MnO (OH) , is formed (1) by oxidation of Mn(OH) 2 in the air; (2) by treating MnO, with concentrated H,SO 4 at a temperature of about 130, forming Mn 2 (S0 4 ) 3 and then adding water: Mn 2 (SO 4 ) 3 + 4H 2 = 2MnO(OH) + 3H 2 SO 4 (Carius, A., 1856, 98, 63). (c) TrimanganesQ tetroxide, Mn 3 O 4 , is formed when any of the higher or lower oxides of manganese or any manganese salts with a volatile acid are heated in the air to a white heat (Wright and Luff, B., 1878, 11, 2145). The corresponding hydroxide would be Mn 3 (OH) 8 ; this has not been isolated. A corresponding oxide-hydroxide is formed by adding freshly formed and moist Mn0 2 to an excess of MnCl 2 containing NH 4 C1 (Otto, A., 1855, 93, 372). 178 MANGANESE. 134, 5. (W) Manganese peroxide, MnO 2 , is formed (/) by heating Mn(NO.0 2 to 200 (Gorgeu, C. r., 1879, 88, 796): (2) by heating 1 MnCO., with KC10 3 to 300; (5) by boiling- any manganous salt with concentrated HNO 3 and KC1O 3 . A correspond- ing- hydroxide, Mn(OH) 4 , has not been isolated. Several other hydroxides, e. ff., MnO(OH) 2 , Mn,O 3 (OH) 2 , Mn : ,0 4 (OH) 4 etc., have been produced. The chief use of mang-anese dioxide is in the preparation of chlorine or bromine. (e) TVEang-anates. Manr/anic acid, H 2 MnO 4 , is not known in a free state. The corresponding salt. K 2 Mn0 4 . is formed when any form of manganese is fused with KOH or K,CO 3 (1) in the air, oxygen being absorbed: or (2) with KN0 3 or KC1O 3 , NO or KC1 being formed. A manganate of the alkali metals is soluble in water, with gradual decomposition into manganese dioxide and per- manganates: ?,K 2 Mn0 4 + 2H,O = 2KMnO. -f Mn0 2 + 4KOH . Free alkali retards, and free acids and boiling promote, .this ehang-e. Manganates have a green color, which turns to the red of permanganates during the decomposi- tion inevitable in solution. This is the usual method of manufacturing KMn0 4 . (/) Permanganic acid is not in use as an acid, but is represented by .the per- mang-anates, as KMnO 4 . The permanganic acid radical is at once decomposed by addition of hot H 2 SO 4 to a solid permanganate (1), but in water solution this decomposition does not at once take place, except by contact with oxidiz- able' substances. The oxidizing power of permanganates extends to a great number of substances, possesses different characteristics in acid and in alka- line solutions, and acts in many cases so rapidly as* to be violently explosive. The reactions with ferrous salts (2) and with oxalic acid (3) are much used in volumetric analysis. (1) 4KMnO 4 + 2H 2 SO 4 2K 2 S0 4 + 4Mn0 2 + 30 2 + 2H 2 and 2Mn0 2 -f 2H 2 S0 4 = 2MnS0 4 + 2H 2 + 2 or 4KMn0 4 + 6H 2 S0 4 = 4MnS0 4 + 2K 2 SO 4 + 50 2 + GH 2 (2) KMn0 4 + SFeCL + SBttl = MnCL + KC1 + 5FeCl 3 + 4H 2 O (3) 2KMn0 4 -f 5H 2 C 2 O 4 + 6HC1 = 2MnCl 2 + 2KC1 + SH 2 -f 10C0 2 5. Solubilities. ff. Metal. Manganese dissolves readily in dilute acids to form manganous-salts. Concentrated H 2 S0 4 dissolves it only on warming, 30 2 being 1 evolved. It combines readily with chlorine and bromine, ft. Oxides and hydroxides. All oxides and hydroxides of manganese are insoluble in water. They are soluble, upon warming, in hydrochloric acid, forming 1 man- ganous chloride; the higher oxides and hydroxides being reduced with evolu- tion of chlorine (commercial method of preparation of chlorine). Instead of hydrochloric acid, sulphuric ac'd and a chloride may be employed (HBr and HI act similarly to, and more readily than HC1). In the cold, hydrochloric acid dissolves MnO 2 to a greenish-brown solution, containing, probably, MnCl 3 or MiiCl 4 , unstable, giving chlorine when warmed and forming Mn0 2 when strongly diluted with water (Pickering, J. C., 1879, 35, 654; Nickles, A. C7i-., 1865, (4), 5, 161). Nitric and sulphuric acids dissolve manganous oxide and hydroxide to mangjanous salts. Manganese dioxide (or hydrated oxide) is insoluble in nitric acid, dilute or concentrated; concentrated sulphuric acid with heat decomposes it, evolving- oxygen and forming manganous sulphate: 2MnO 2 + 2H 2 SO 4 = 2MnSO 4 + 2H 2 O + O 2 . Manganous hydroxide is insoluble in the alkalis but soluble m solutions of ammonium salts. c. Salts. Manganous sulphide, carbonate, phosphate, oxalate, borate, and sulphite are insoluble in water, readity soluble in dilute acids. Man- ganic salts are somewhat unstable compounds, of a reddish-brown or purple-red color, becoming paler and of lighter tint on reduction to the manganous Combination. MnCL and MnS0 4 arc (1clii~/uc$cent. Man- ganic chloride, MnCl 3 , and the percliloridc , MnCl 4 , are unstable salts which 134, 6fl. MANGANESE. 179 are decomposed by water especially when hot, to MnCL and chlorine. The trichloride is greenish black, soluble in absolute alcohol and ether; while the tetrachloride is reddish-brown soluble in absolute alcohol. (W. I>. Holmes, J. Am. tioc. 29, 1277.) Manganic sulphate Mn,(S0 4 ) 3 , is soluble in dilute sulphuric, acid, but is ivdiuvd to MnS0 4 by the attempt to dissolve it in water alone; potassium manganic sulphate and other manganic alums are also decomposed by water. Alkali manganates and permanganates are soluble in water, the former rapidly changing to man- ganese dioxide and permanganate, which is much more stable in solution. In presence of reducing agents both manganates and permanganates are reduced to lower forms. K 2 MnO 4 + 8HC1 = MnCl 2 + 2KC1 + 2C1 2 + 4H 2 O 2KMnO 4 + 3MnSO 4 -f 2H 2 O = 5MnO 2 + K 2 SO 4 + 2H,SO 4 Concentrated H 2 S0 4 in the cold dissolves KMn0 4 , forming (Mn0 3 ) 2 S0. 1 (a sulphate of the heptad manganese: 2KMn0 4 + 3H 2 S0 4 = (Mn0 3 ) 2 S0 4 + 2KHS0 4 + 2H 2 (Franke, J. pr., 1887, 36, 31). If heat be applied oxygen is evolved and the manganese is reduced to the dyad (4/). 6. Reactions, a. The fixed alkali hydroxides precipitate from solu- tions of manganous salts, manganous hydroxide, Mn(OH) 2 , white, soon turning brown in the air by oxidation to manganic hydroxide, MnO(OH) . The precipitate is formed by the reaction of the negative hydroxyl ion of the alkali with the positive manganous ion : 2NH 4 OH + Mn^f 2 = Mn(OH) 2 + 2NH 4 Cl . The precipitate is insoluble in excess of the alkalis because the mangan- ous or manganic manganese does not form an acid ion. Before the manganous manganese is oxidized it is soluble in excess of ammonium salts because a complex salt is formed in which the manganese forms a part of the acid ion : + - ++ 4- - Mn(OH) 2 + 4NH 4 Cl - (NH 4 ) 2 MnCU + 2NH 4 Cl . (1). Ammonium hydroxide precipitates one-half of the manganese as the hydroxide from solutions of manganous salts, the other half being held in solution as an acid ion by the ammonium salt formed (2) (Da tu- rner, 3, 237). The presence of excess of ammonium salt prevents the precipitation of the manganese by ammonium hydroxide because the manganese is present in the acid ion (3) (separation of manganese from the metals of the third group) (Pickering, J. C. } 1879, 35, 672; Lang- bein, Z., 1887, 26, 731). Manganic hydroxide, MnO(OH), is insoluble in the alkalis or in ammonium salts. It gradually precipitates, com- pletely on exposure to the air, as a dark brown precipitate from solutions 180 MANGAXERE. 134, 65. of manganous hydroxide in ammonium salts. Alkali carbonates pre- cipitate manganous carbonate, MnC0 3 , white, oxidized in the air to the brown manganic hydroxide, and before oxidation, somewhat soluble in am- monium chloride. Strong ammonium hydroxide gradually reduces a solu- tion of potassium permanganate to manganese dioxide (10ft). (1) Mn OH) 2 + 4NH 4 C1 = (NH 4 ) 2 MnCl 4 + 2NH 4 OH (*) 2MnS0 4 + 2NH40H = (NH 4 ) 2 Mn(SO 4 ) 2 + Mn(OH) 2 (3) MnCl z + 2NH.C1 = (NH 4 ) 2 MnCl 4 &. Oxalic acid and alkaline oxalates precipitate manganous oxalate, soluble in mineral acids not too dilute. All compounds of manganese of a higher degree of oxidation are reduced to the manganous condition on warming with oxalic acid, or oxalates in presence of some mineral acid: 2KMn0 4 + 5H 2 C,0 4 + 3H 2 S0 4 = K 2 S0 4 + 2MnS0 4 + 10C0 2 + 8H 2 . Soluble cyanides, as KCN , precipitate manganous cyanide, Mn(CN) 2 , white, but darkening- in the air; soluble in excess of the precipitant by formation of double cyanides, as Mn(CN) 2 .2KCN . This solution, exposed to the air, pro- duces manga-nioyanides (analogous to ferricyanides), with oxidation of the manganese: 12(Mn(CN) 2 .2KCN) + 30 2 + 2H 2 O = 8K 8 Mn(CN) a + 4MnO(OH). Fe'" and Mn" may be separated by treating a solution of the two metals with a strong excess of KCN and then with iodine. The manganese is precipitated as Mn0 2 and the iron remains in solution (Beilstein and Jawein. B., 187'J, 12, 1528). Ferrocyanides piecipitate white inunyannux fcrroci/anide, Mn 2 Fe(CN) 6 , soluble in hydrochloric acid. Ferricyanides precipitate brown manganous frrri- cyanide, Mn s (Fe(CN) 6 ) 2 , insoluble in acids (separation, with Co and Ni , from Zn) (Tarugi, Gazzetta, 1895, 25, ii, 478). If an alkali or alkali carbonate be present, potassium ferricyanide oxidizes manganous compounds to manganese dioxide, the ferricyanide being reduced to ferrocyanide. Potassium ferro- cyanide reduces manganates and permanganates to manganous compounds. c. Nitric acid is of value in analysis of manganese compounds in that it, as a non-reducing acid, acts readily with oxidizing agents, as Pb0 2 , KC10 3 , etc., to oxidize manganous compounds to manganese dioxide or to- permanganic acid. Eeducing agents as HC1 , etc., should be absent. Sulphuric acid may be used instead of nitric acid. 2Mn(NO 3 ) 2 + 5Pb0 2 + GHN0 3 = 2HMn0 4 + 5Pb(N0 3 ) 2 + 2H 2 O 5MnS0 4 f 2KC10 3 + H 2 SO 4 + 4H 2 O = 5Mn0 2 + K 2 S0 4 + Cl a + 5H 2 SO 4 In using Pb0 2 and HN0 3 to detect manganese, the compound should first be reduced w;th hydrochloric acid, precipitated with potassium hydroxide and this precipitate dissolved in nitric acid, as Mn0 2 is not all oxidized by Pb0 2 and HN0 3 (Koninck, Z. angew., 1889, 4). d. ^Hypophosphorous acid reduces all higher forms of manganese to the manganous condition. Alkali phosphates, as Na,HPO 4 , precipitate, from neutral solutions of manganous salts, normal nia'ngamms }>1iox}>1iate, Mn s (POj r ) 2 , white, slightly soluble in water, and soluble in dilute acids. It turns brown in the air. The manganous hydrogen phosphate, MnHP0 4 , is more soluble in water, and is obtained by crystallization from a mixture of manganous sul- COLLCOfr Srf PHARMACY 134, Qg. MANGANESE. 181 phate acilulited with acetic acid and di sodium phosphate, Na2HPO4 , added till a precipitate begins to form. From the ammonium-manganese solution, freshly formed (6), phosphates precipitate nil the manganese as manyanoua ammonium phosphate. e. Hydrosulphuric acid precipitates manganous acetate but imperfectly, and not in presence of acetic acid, and does not precipitate other salts, as manganous sulphide is soluble in very dilute acids, even acetic acid. Ammonium sulphide precipitates from neutral solutions, antl forms from the recent hydroxide of mixtures made alkaline, the flesh-colored man- ganous sulphide, MnS . Acetic acid, acting on the precipitated sulphides, separates manganese from cobalt and nickel, and from the greater part of zinc. All the higher oxidized forms of manganese (in solution or freshly precipitated) are reduced to the manganous condition, with separation of sulphur (10), by hydrosulphuric acid or soluble sulphides: 4KMn0 4 + 14(NH 4 )JS + 16H 2 =: 4MnS + 4KOH + 28NH 4 OH + 5S 2 . The green manganous sulphide, MnS , crystalline, anhydrous, is formed by the action of HoS on a hot ammoniacal manganous solution not containing an excels of ammonium salts (Meineke, Z. angew., 1888, 3), also by pouring the neutral manganous solution into a hot solution of ammonium chloride and excess of colorless ammonium sulphide. The fixed alkali sulphides produce a red manganous sulphide. Soluble Sulphites precipitate from solutions of manganous salts, manganous sulphite, MnSO 3 , white, insoluble in water, soluble in acids (Gorgeu, C. r., 1883, 96, 341). Solutions of manganates or permanganates are immediately reduced to the flocculent brown-black manganese dioxide by solutions of sodium sulphite or sodium thiosulphate ; if acids be present, the reduction is complete to manganous salts. /. HC1 , HBr , and HI readily reduce the higher compounds of man- ganese to manganous salts with evolution of the corresponding halogen. When manganese dioxide is dissolved in concentrated HC1 without heat, the dark brownish colored solution contains manganese tetraohloride, MnCl 4 , and trichloride, MnCl 3 which deposits Mn0 2 on dilution with water and on warming decomposes into manganous chloride and chlorine (56) (Pickering, J. C., 1879, 35, 654, W. B. Holmes, J. Am. Soc. 29, 1277). Potassium iodide instantly reduces a solution of potassium per- manganate, forming manganese dioxide and an iodate (distinction from chloride and bromide). Potassium chlorate or bromate when boiled with concentrated nitric or sulphuric acids and manganous compounds forms manganese dioxide (c). g. Soluble arsenites precipitate manf/anous arsenite, and arsenates precipitate HHUIWIIHHIH <(rw>Hit<\ insoluble in water, soluble in acids. Arsenous acid and arsenites reduce solutions of manipulates or permanganates, forming- a brown flocculent precipitate 1 : or a colorless solution if warmed in presence of a mineral acid. //.Normal potassium chromate precipitates manganous salts, brown, soluble in acids and in ammonium hydroxide; no precipitate is formed with potassium dichroinate. . Soluble manganates and permanganates pre- 182 MANGANESE. 134, 7. cipitate manganous salts as manganese dioxide, being themselves reduced to the h ame form: 3MnSO 4 + 2KMnO 4 + 2H 2 O = 5MnO 2 + K 2 SO 4 + 2H S SO 4 . 7. Ignition with alkali and oxidizing agents, forming a bright green mass of alkaline manganate, constitutes a delicate and convenient test for man- ganese, in any combination. A small portion of precipitate or fine powder is taken. If the manganese forms but a small part of a mixture to be tested, it is better to submit the substance to the systematic course of analysis, and apply this test to the precipitate by alkali, in the fourth group. A convenient form of the test is by ignition on platinum foil with potassium or sodium nitrate and sodium carbonate (a). Ignition, by an oxidizing flame, on platinum foil, with potassium hydroxide, effects the same result, less quickly and perfectly (&). Ignition by the oxidizing flame of the blow-pipe, in a bead of sodium carbonate, on the loop of platinum wire, also gives the green color (c). (a) Mn(OH) 2 + 2KNO 3 + NjfcCO 3 = .NasMnO, + 2KNO 2 + CO 2 + H 2 O or if a small amount of KNOs is present, (fe) 3Mn(OH) 3 + 4KNO 3 + NaaCOs =2K 2 MnO 4 + Na^MnCX + 4NO + CO 2 + 3H 2 O o Mn OH) 2 + 2KOH + O 2 = K 2 MnO 4 + 2H 2 O ,1 MnvOH) 2 + NaaCOs + O 2 = NaaMnO, + H 2 O + CO 2 With beads of borax and microcosmic salt, before the outer blow-pipe flame, manganese colors the bead violet while hot, and amethyst-red when cold. The color is due to the formation of manganic oxide, the coloring- material of the amethyst and other minerals, and is slowly destroyed by application of the inner flame, which reduces the manganic to manganous oxide. 8. Detection. After the removal of the metals of the first three groups (the third group in the presence of NH 4 C1 in excess, 5fr and 6a), the Mn with Co , Ni and Zn is precipitated in the ammoniacal solution by H 2 S . By digestion in cold dilute HC1 the sulphides of Mn and Zn are dissolved, and after boiling to remove the H 2 S , Mn is precipitated as the hydroxide by excess of KOH , which dissolves the Zn . The precipitate of the man- ganese is dissolved in HN0 3 and boiled with more HN0 3 and an excess of Pb0 2 . A violet-colored solution is evidence of the presence of manganese. 9. Estimation. (1) By converting into Mn 3 4 (4c), and weighing as such. (2) By precipitating as MnNH 4 P0 4 , and after ignition weighing as Mn 2 P 2 O 7 . (3) By treating the neutral manganous salt with a solution of KMnO 4 of known strength (Qi). If some ZnS0 4 is added the action is more satisfactory (Wright and Menke, J. C., 1880, 37, 42). (J f ) By boiling the manganous com- pound with Pb0 2 and HNO 3 , and comparing the color with a permanganate solution of known strength (Peters, C. N., 1870, 33, :J5). (5) The manganous compound is oxidixed to MnO 2 by boiling with KC10 3 and HNO 3 . This is then reduced by an excess of standard H 2 O, , H,,C,O 4 or FeSO 4 , and the excess of the reagent estimated by the usual methods. (6) MnO, , obtained as in (.5), is treated with H 2 C 2 4 and the evolved CO 2 measured or weighed. (7) Mn0 2 , obtained as in (o), is boiled with HC1 and the evolved Cl estimated. 10. Oxidation. (a) Mn" is oxidized to Mn'" in alkaline mixture on exposure to the air ; to Mn lv in neutral solution by K a Mn0 4 and KMn0 4 , 135, 3. ZINC. 183 in alkaline mixture by Cl , Br , I , K,Fe(CN) , KC10 , KBrO , H.O, 1 , etc.; in acid solution by boiling with concentrated HN0 3 or H 2 S0 4 , arid KC10 3 or KBrO, . Mn VI - n is oxidized to Mn VI by fusion with an alkali and an oxidizing agent, or by fusion with KC10 3 alone (Boettger, Z., 1872, 11, 433). Mn VII ~ n is oxidized to Mn vn by warming with Pb0 2 or Pb 3 4 and HN0 3 or H 2 S0 4 . The higher oxide of lead should be in excess and reduc- ing agents should be absent as they delay the reaction; hence in analysis the manganese should be precipitated as the hydroxide or sulphide, fil- tered, washed, and then dissolved in HN0 3 or H 2 S0 4 , and boiled with the higher oxide of lead (6c). A solution of potassium manganate decomposes into potassium permanganate and manganese dioxide on standing, more rapidly on warming or dilution with water, (b) All compounds of man- ganese having a higher degree of oxidation than the dyad, (Mn" +n ) are reduced to the dyad (Mn") by H 2 C 2 4 , HH 2 P0 2 , H 2 S 4 , K 2 S , H 2 SO, , H 2 2 2 (in neutral or alkaline solution to Mn IV ), HC1 , HBr , HI , HCNS , Hg', Sn", As'", Sb'", Cu', Fe", Cr", Cr'", etc.; the reducing agents becoming respec- tively C0 2 , P v , S to S VI (depending upon the temperature, concentration, and the agent used in excess), Cl , Br , I , HCN and S VI , Hg", Sn IV , As v , Sb v , Cu", Fe'", and Cr VI . Mn IV + n is reduced to Mn IV (or Mn'") by H 3 , Asfl., 3 , SbH 3 3 , PH, 3 , Na 2 S0 3 4 , Na 2 S 2 3 4 , NH 4 OH 3 (slowly), Mn", etc. KMn0 4 is reduced to K 2 Mn0 4 on boiling with concentrated KOH : 4KMn0 4 + 4KOH = 4K,Mn0 4 + 2H 2 + 2 (Eammelsberg, ., 1875, 8, 232). 135. Zinc. Zn = 65.37. Valence two. 1. Properties. Specific gravity, 7.142 (Spring, B., 1883, 16, 2723). Melting point, 418.5 to 419.35 (Burgess, Wash. Acad. of Sci., 1-18). Boiling point, 918 (Berthelot, C. r., 1912, 134). It is a bluish-white metal, retaining its lustre in dry air, but slightly tarnished in moist air or in water. When heated to the boiling point with abundant excess of air it burns with a bluish-white flame to zinc oxide. Zinc dust mixed with sulphur is ignited by percussion (Schwarz, B., 1882, 16, 2505). At ordinary temperature it breaks with a coarse crystalline fracture. It is more malleable at 100 to 150 than at other temperatures, and at that temperature may be drawn into wire or rolled into sheets. At 205 it is so brittle that it may be easily powdered in a mortar. Zinc finds an extended use in laboratories for the generation of hydrogen. It is molded in sticks or granulated by pouring the molten metal into cold water. The pure metal is not suitable for the generation of hydrogen, as the reaction with acids proceeds too slowly (Weeren, B., 1891, 24, 1785). Com- mercial impurities render the metal readily soluble in acids, or the pure metal may be treated with a dilute solution of platinum chloride (twenty milligrams PtCl 4 per litre), or copper sulphate. Met.-illic platinum or copper is deposited upon the zinc: PtCl, + 2Zn = Pt + 2ZnCL , CuSO 4 + Zn = Cu + ZnSO 4 . 2. Occurrence. It is found as calamine (3nCO 3 ), as zinc-blende (ZnS); also associated with other metals in numerous <> es. 3. Preparation. The process usually employed consists of two operations: 'Klein, Arch. Pharm., 1880, 227, 77; Jannaesch and von Cloedt, Z. annrg., 1805, 10, 398 and 410; Carnot, C. r., 1888, 1O7, 997 and 1150. *Carnot, Bl., 1889, (3), 1, 277 ; Gorgeu, C. r., 1800, HO, 958. 3 Jones, J. C., 1878, 33, 96. 4 Hoenig and Zateck, Jf., 1863, 4, 738 ; Glaeser, Jf., 1685, 6, 329. 184 ziNC. 135, 4. (I) Roasting-: in case of the carbonate the action is: ZnCO, = ZnO -f CO,- if it is a sulphide, 2ZnS + 3O 2 = 2ZnO + 2SO 2 . (2) Reduction with distillation; after mixing- the ZnO with one-half its weight of powdered coal, it is distilled at a white heat. Its usual impurities are As, Cd , Pb , Cu , Fe and Sn . It is purified by repeated distillation, each time rejecting- the first portion, which contains the more volatile As and Cd , and the last which contains the less volatile Pb , Cu , Fe and Sn . Strictly chemically pure zinc is best prepared from the carbonate which has been purified by precipitation. 4. Oxide and Hydroxide. Zinc oxide (ZnO) is made by igniting- in the air either metallic zinc, its hydroxide, carbonate, nitrate, oxalate, or any of its org-anic oxysalts. Zinc hydroxide, Zn(OH), , is made from solutions of zinc salts by precipitation with fixed alkalis (6a). 5. Solubilities. (a) Metal. Pure zinc dissolves very slowly in acids or alkalis, unless in contact with copper, platinum or some less positive metal (Baker, J. C., 1885, 47, 349). The metallic impurities in ordinary zinc enable it to dissolve easily with acids or alkali hydroxides. In contact with iron, it is quite rapidly oxidized in water containing 1 air, but not dissolved by water unless by aid of certain salts. It dissolves in dilute hydrochloric, sulphuric * and acetic acids (Jf), and in the aqueous alkalis (2), with evolution of hydrogen; in very dilute nitric acid, without evolution of gas (3); in moderately dilute cold nitric acid, mostly with evolution of nitrous oxide (j); and, in somewhat less dilute nitric acid, chiefly with evolution of nitric oxide (5). Concentrated nitric acid dissolves zinc but slightly, the nitrate being- very sparingly soluble in nitric acid (Montemartini, Gazzstta, 1892, 22, 277). Hot concentrated sul- phuric acid dissolves it with evolution of sulphur dioxide (6). (1) Zn + H 2 S0 4 = ZnS0 4 + H 2 (2) Zn + 2KOH = K 2 Zn0 2 + H 2 (3) 4Zn + 10HN0 3 = 4Zn(N0 3 ) 2 + NH 4 NO 3 + 3H 2 O (4) 4Zn + 10HN0 3 = 4Zn(N0 8 ), + N 2 + 5H 2 (5) 3Zn + 8HN0 3 = 3Zn(N0 8 ) 2 + 2NO -f 4H 2 (6) Zn + 2H 2 S0 4 = ZnS0 4 + S0 2 + 2H 2 (&) Oxide and Hydroxide. All the agents which dissolve the metal, dissolve also its oxjde and hydroxide. (c) Salts. The chloride, bromide, iodide, chlorate, nitrate (6aq), and acetate (7aq) are deliquescent; the sulphate (7aq) is efflorescent. The chloride is readily soluble in alcohol in all proportions (Kremers, Pogg., 1862, 115, 360). The sulphide, basic carbonate, phosphate, arsenate, oxalate, and ferrocj^anide are insoluble in water; the sulphite is sparingly soluble. The ferrocyanide is insoluble in hydrochloric acid (Fahlberg, Z., 1874, 13, 380). The sulphide is almost insoluble in dilute acetic acid (sepa- ration from MnS). All zinc salts are soluble in KOH and NaOH except zinc sulphide, and all in NH 4 OH except ZnS and Zn 2 Fe(CN) 6 . 6. Reactions, a. The fixed alkali hydroxides precipitate zinc hydroxide, Zn^OH) 2 , white, soluble in excess of the precipitant forming an alkali ziticate: ZnCl 2 + 2KOH = Zn(OH) 2 + 2KC1 Zn(OH) 2 + 2 KOH = K 2 ZnO 2 + 2H 2 O Ammonia precipitates from neutral solutions free from ammonium salts, zinc hydroxide, soluble in excess of ammonia or ammonium salts forming complex zinc ammonia ions: ZnCl 2 + 2NH4OH = Zn(OH) 2 + 2NH4C1 Zn(OH) 2 + 6NH 3 = Zn(NH 3 ) 6 (OH) 2 * Muir and Jlobbs, C. A'., 1882. 45, 69. 135, 6A. ZINC. 185 The precipitate of zinc hydroxide dissolves more readily in excess of the alkalis at ordinary temperature than when heated. Unless a strong excess of the alkali be present,, boiling causes a precipitation of zinc oxide, more readily from the solution in ammonium hydroxide than in the fixed alkalis. The presence of other metals as iron or manganese makes necessary the use of much more alkali to effect solution. An alkali solu- tion as dilute as tenth Normal' does not dissolve zinc hydroxide, no matter how great an excess be added (Prescott, J. Am. Soc., 1880, 2, 29). Alkali carbonates precipitate the basic carbonate, Zn 5 (OH) 6 (C0 3 ) 2 , white, soluble in ammonium carbonate, readily in alkali hydroxides (Kraut, Z. anorg., 1896, 13, 1). Carbonates of Ba, Sr, Ca , and Mg have no action at ordinary temperatures (separation from Fe'", Al , and Cr'"), but upon boiling precipitate the whole of the zinc. ft. Alkali cyanides, as KCN , precipitate zinc cyanide, Zn(CN) 2 , white, soluble in excess of the precipitant. Alkali ferrocyanides, as K 4 Fe(CN) R , precipitate zinc ferrocyanide, Zn 2 Ee(CN) a , white (5c). Alkali ferricyanides, as K 3 Pe(CN) 6 , precipitate zinc ferricyanide, Zn 3 (Fe(CN) ) 2 , yellowish, c. See 5c. ^.Sodium phosphate, Na 2 HPO 4 , precipitates zinc phosphate, soluble in alkali hydroxides and in nearly all acids. e. Hydrosulphuric acid precipitates a part of the zinc from neutral solutions of its salts with mineral acids, and the whole from the acetate; also from other salts of zinc, by addition of alkali acetates or monochlor- acetic acid, in small excess (separation from Mn , Co , Ni , and Fe) (Berg, Z., 1886, 25, 512): ZnCl 2 + 2KC 2 H 3 2 + H 2 S = ZnS + 2KC1 + 2HC 2 H 3 2 .* That is: Zinc sulphide is not entirely soluble in dilute acids, though much more soluble in mineral acids than in acetic acid. The precipitate is white when pure. Alkali sulphides completely precipitate zinc as sulphide, both from its salts with acids and from its soluble com- binations with alkalis. Concentrated solutions of sodium sulphite precipitate solutions of zinc salts as basic zinc sulphite; or if the solutions be too dilute for immediate precipita- tion, boiling will cause the immediate formation of the bulky white precipitate of the basic sulphite (Seubert, Arch. Pharm., 1891, 229, 316). f.If a hot con- centrated zinc chloride solution be treated with ammonium hydroxide until a precipitate begins to form, a basic chloride, 2ZnCl 2 .9ZnO , will separate out upon cooling as a white precipitate (Habermann, M., 1884, 5, 432). g. Zinc salts are precipitated by solutions of alkali arsenites and arsenates, forming respectively zinc arsenite or arsenate, white, gelatinous, readily solu- ble in alkalis and acids, including arsenic acids, ft. Normal potassium chro- * In the equation for acetic acid, ab =kc, a and b, the concentrations of the H and C,II 3 O, ions respectively, are small, c is large, and k, the so-called "dissociation-constant," to which the strength of the acid is proportional, is very small. But addition of the fully-dissociated sodium acetate to the likewise completely-ionized hydrochloric acid gives a solution containing the ions in very large concentration and practically none of the non-dissociated acetic acid. To restore equilibrium the H ions of the HC1 unite with the acetic ions of the sodium acetate, leaving Na and Cl ions in the solution. The displacement of a weak acid from its salt by a strong one lies then not so much in an attraction of the strong acid by the base as in the ten- dency of the weak acid to form the non-ionized molecule. 186 ZINC. 135, 7. mate forms, with solutions of zinc salts, a yellow precipitate readily soluble in alkalis and acids, including chromic acid. No precipitate is formed with K 2 Cr 2 7 . 7. Ig-nition. With sodium carbonate, on charcoal, before the blow-pipe, com- pounds of zinc are reduced to the metallic state. The metal is vaporized, and then oxidized in the air, and deposited as a non-volatile coating-, yellow when hot and white when cold. If this coating-, or zinc oxide otherwise prepared, be moistened with solution of cobalt nitrate and again ignited, it assumes a green color (Bloxam, J. C., 1865, 18, 98). With borax or microcosmic salt, zinc compounds give a bead which, if strongly saturated, is yellowish when hot, and opaque white when cold. 8. Detection. After the removal of the first three groups, the Zn is precipitated with Co , Ni and Mn from the ammoniacal solutions by H 2 S . Digestion of the precipitated sulphides with cold dilute HC1 dissolves the Mn and Zn as chlorides. The solution is thoroughly boiled to expel the H 2 S and the zinc changed to Na 2 Zn0 2 by an excess of NaOH , which precipi- tates the manganese as the hydroxide. From the alkaline filtrate H 2 S gives a white or grayish-white precipitate evidence of the presence of Zn . 9. Estimation. (1) Zinc is weighed as an oxide, into which form it is brought by simple ignition if combined with a volatile inorganic oxyacid, otherwise it should be changed to a carbonate and then ignited. (2) It is converted into a sulphide, and after adding powdered sulphur it is ignited in a stream of hydrogen or hydrogen sulphide, and weighed as a sulphide (Kiinzel, Z., 1863, 2, 373). (3) It may be converted into ZnNH 4 P0 4 , and, after drying at 100, weighed. Ignition converts it into Zn,P,0 7 , with slight loss of zinc. (4) Volumetrically, by converting into Zn_,Fe(CN) 6 and titrating with potas- sium permanganate or by using FeCl 3 acidulated with HC 2 H 3 2 as external indicator (Voigt, Z. angew, 1889, 307). (5) By precipitation as Zn 3 (Fe(CN) ft ,) 2 , treating the precipitate with potassium iodide and titrating the liberated iodine (Mohr, DingL, 1858, 48, 115). (6') By titration in hydrochloric acid solution with K 4 Fe(CN) 6 , using a uranium salt as an indicator (Fahlberg, Z., 1874, 13, 379; Koninck and Prost, Z. angew., 1896, 568). (?) By titration in alkaline solution with Na 2 S , using a copper salt as an indicator. (8) The zinc is pre- cipitated as ZnNH 4 AsO 4 , the precipitate decomposed with HI and the liber- ated iodine titrated with standard Na 2 S 2 O 3 (Meade, J. Am. Soc., 1900, 22, 353). 10. Oxidation. Metallic zinc precipitates the free metal from solutions of Cd , Sn , Pb , Cu , Bi , Hg , Ag , Pt , An , As , Sb , Te , In , Fe *, Co , Ni, Pd, Rh, Ir, and Os (Gmelin-Kraut, Handbuch, 1875, 3, 6). Zinc with copper (zinc-copper couple, used in water analysis) reduces nitrates and nitrites to ammonia, chlorates to chlorides, iodates to iodides, ferri- cyanides to ferrocyanides, etc. (Thorpe, /. (7., 1873, 26, 541). Solutions of chromates are reduced to chromic salts, ferric salts to ferrous salts, and compounds of manganese having more than two bonds are reduced to the dyad in presence of some non-reducing acid. Zinc is precipitated as the metal from acetic solutions by Mg (Warren, C. N., 1895, 71, 92). The oxide is reduced to the metal by heating in a current of hydrogen (Deville, A. Ch., 1855 (3), 43, 477). *Davie8. J. C., 1875, 28, 3a. 136. REACTIONS OF IRON AND ZINC GROUP BASES. 187 d c fi - g PH ,0 c3 O GO -1 2 W * PH | | d c d d S W N N _ . W ,0 b C) -t c5 "z S - PH O GO I W PH rS rt O c i-H iH fc W b & ^ ^ H- 03 Is W o - i. O GO I 2 W 11 * O I o o r O o T ^ ^ PQ O O rr^ Q O CT . O ^ O ^"~ N s O3 2 ^ ii W a 5 & 6 o GO D | S PH o -*-> 1 s fl d a d !^ g d a ^ 1 | J3 O a^ss?is T3 tn .2 || 1 1 S a -9 a -oj | .ll&flitfl jHQO^Sjrt^^eU^ 188 TABLE FOR ANALYSIS OF THE ZINC G&OUP. 131 137. TABLE FOE ANALYSIS OF THE ZINC GROUP (FOURTH GROUP) (Phosphates and Oxalates being absent). Into the clear ammoniacal filtrate from the Third Group pass Hydrogen Sul- phide, and if a precipitate appears, warm until it subsides. Filter and wash with a one per cent solution of NH 4 C1 . (Test nitrate, 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 (1-4). Residue: CoS, NiS* (black). Test with the borax bead. A blue bead indicates cobalt, (132, 7). Dissolve the remainder of the sulphides in nitro-hydrochloric acid or HC1 + crystal of KC1O 3 . Evaporate to ex- S^l excess of CL_ , neutralize with H 4 OH and divide into two parts. Solution: MnCl 2 , ZnCl 2 (H 2 S,HCl). (traces of CoCL> and NiCl 2 .) Boil the solution thoroughly to remove the H 2 S , cool, and add a decided excess of potassium or sodium hydroxide and bromine water and heat (135, 6a). Filter and wash. For Cobalt: For Nickel: Precipitate : Solution : Add NaHCO 3 and H 2 O 2 ; warm Add dimethyl-gly- oxime and warm. MnO 2 (traces of Co(OH) 3 and Ni(OH) 3 . K 2 ZnO 2 gently and filter. A scarlet precipi- Dissolve in HNO 3 -f a small Test for A green color to tate shows nickel. amount of H 2 O 2 and boil. zinc by the filtrate indi- (If Co and Ni were present adding cates cobalt Or: Add excess of precipitate the manganese H 2 S . A (140). hot KOH and Br, from this solution by white pre- boil, filter, wash NH 4 OH+H 2 O 2 and boiling. cipitate If sufficient nickel (until fi 1 1 r a t e Filter, wash and redissolve (ZnS) in- be present to ob- gives no precipi- in HNO 3 +H 2 O 2 and boil.) dicates scure the blue tate with AgNOs), Boil with HNO 3 and Pb.C^ zinc. bead, add to the add solution of or PbO 2 . Manganese will solution of the hot KI and test give the reddish purple color sulphides (133, the filtrate with ofHMnO 4 (134,6c). Man- 7) an excess of ni- CS 2 . If free io- ganese may also be tested troso-/3-naphthol dine appears, for,by pouring diamine over in acetic acid so- nickel is present the Mn0 2 if Co and Ni are lution (132, 66); (133, 6/). absent, or if they are pres- filter, wash, and ent, by pouring it over the test the brick-red precipitate of Mn(OH) 3 precipitate with obtained by precipitating the borax bead. with NH 4 OH + H 2 O 2 . Manganese gives a reddish The filtrate may be purple precipitate. tested for nickel. Dark-colored original solu- tions indicating an alkali salt of manganese should be reduced by warming Study the Study 132, 6c, 135, 138, 139, Study the text at 133, 6a, 6, e and with HC1 before proceed- ing with the analysis (134, 5c and 6/). text at 135, 6a and e, 140, 141, 144, /; 132 66 and c~ 136, 138, 135 and ff. 136, 138, 139, Confirm by study of the text 139, 142, 140, 141, 144, 134, 7, 136, 138, 139, 143, 144, 145 and ff. 142, 143, 144, 145 and 145 and/. ff- * Small portions of cobalt and nickel sulphides may be dissolved by the cold dilute HC1, and will be precipitated with the Mn(OH) 2 . 140. DIRECTIONS FOR ANALYSIS WITH NOTES. 189 DIRECTIONS FOR THE ANALYSIS OF THE METALS OF THE FOURTH GROUP. 138. Manipulation. Into the warm strongly a mm onincal filtrate from the third group (128) , H^S gas is passed until complete precipitation is obtained: MnCL.2NH 4 Cl + 2NH 4 OH + H 2 S = MnS + 4NH 4 C1 + 2H 2 O (NH 4 ) 2 ZnO 2 + 2H 2 S = ZnS + (NH 4 ) 2 S + 2H 2 O 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 Hl (one part of reagent HC1 to four of water) : MnS + 2HC1 = MnCl, -f H 2 S . The black precipitate remaining undissolved contains the sulphides of Co and Ni, 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 1 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 S X , 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 HC1 , 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, 6fl), and is detected by the green color of the fused mass when testing for Cr: 3Mn(OH) 2 + 4KNO 3 -f- Na 2 CO 3 = 2K,MnO 4 + Na 2 MnO 4 + 4NO -f CO, -f 3TL..O . Too much HNO 3 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. (7) Small amounts of the Fifth Group elements are carried down with the ammonium sulphide precipitate. As much as 5 mg. of barium may be present in this precipitate (Curtman & Frankel, J. Am. Soc., 33, 724, 1911). 140. Manipulation. The black precipitate of cobalt and nickel sul- phides should first be tested with the borax bead (141, 3) 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 + 4NO + 8H a O . pivide the solution into three portions: To one portion of the solution 190 DIRECTIONS FOR ANALYSIS WITH NOTES. 141, 1. add an excess (142, #) of nitroso- /9-Naphthol, filter, and wash with hot water and then with hot HC1 (132, 65). Test the red precipitate with the borax bead for cobalt. Render the filtrate ammoniacal, filter again and test this last filtrate with H 2 S for the black precipitate of NiS (133, 66 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. (1) HN0 3 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 KC1O 3 may be used instead of HNO 3 . (2) If an insufficient amount of nitroso- /3-naphthol has been used a portion of the cobalt maj' be in the nitrate 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 C0 3 . The black precipitate so obtained will filter rapidly. (o) The test for nickel by adding KI to the mixed higher oxides of cobalt ami 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. () 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: KIO 3 + SKI + 6HC1 = 3I 2 + 6KCf + 3BVO (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 liydrosulphuric 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) 2 + 2KC1 Z&01, + 4NaOH = Na 3 Zn0 2 + gNaCl -f- 3H.O 144. ANALYSIS OF IRON AND ZINC GROUPS. 191 Filter and test the filtrate with H,S , a white or grayish-white precipitate indicates zinc (characteristic). Dissolve the well washed precipitate of Mn(OH) 2 in nitric acid and boil with an excess of lead peroxide, adding more nitric a-cid. A violet color to the nitric acid solution indicates the presence of manganese (very delicate and characteristic) : 2Mn(OH) 2 + 5Pb0 2 + 10HN0 3 = 2HMn0 4 + 5Pb(N0 3 ) 2 + GH,0 143. Notes. 1. If the H 2 S is not completely removed the Zn will be pre- cipitated as the sulphide upon adding- the NaOH , 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 O 4 + 30HN0 3 = 2HMn0 4 -f 15Pb(N0 3 ) 2 + 16H 2 6. It is very difficult to procuie PbO 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 PbO 2 has subsided. Then the unknown solution in HNO 3 may be added 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 ~M.nO., frequently fails to give permanganic acid when boiled with PbOo and HN0 3 until after reduction (134, 6c). ANALYSIS OF IRON AND ZINC GROUPS AFTER PRECIPITATION BY 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, 55 and 6a), and to insure the complete precipitation of the aluminum as the hydroxide 124, Go). 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 OH 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: 193 ANALYSIS OF IRON AND ZING GROUP. 144 o 4> ^ T3 V ^3 CO - > I 02 CO 0) fl ^ - & s 3 rz fc * S - 5 as I w w t i m % g X " ^ ! fir tJ2 * 5 ^ ; & g j&M-aB "" S W Ig Srf 13 co X CO g; ^ O o ^ w o & .*: ^ '3 fe & ^ 'S o S-8.J W I'lin w ^ d. c^- CO >-. C r^ j^.a 149. IRON AND ZINC GROUPS. 103 145. 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(OH) 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-motybdate has completely formed and settled (75, 6d). Filter and evaporate nearly to dryness to remove the nitric acid. Take up with water and a little hydrochloric acid if necessary to obtain a 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 NaCoH ;! 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 PeCl 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 194 IRON AND ZINC GROUPS. 150. to precipitate the ferric acetate as basic ferric acetate (126, 6&) 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(OH) 3 -f- SBaCL + 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 + CoC0 3 , 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 BaC0 3 . 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 151. IRON AND ZINC GROUPS. 195 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 filtrate 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). 196 CALCIUM GROUP METALS. 152. m odi ' bB^'^ o .2 ,Ci fc; >-j fl MO 03 0) _ ' ' .S 4 w ci -d 3 ~ 2 bSsI date solu- emainder ny precipi- precipitate If phos- p by drop the liquid Filter and add am- ciitate of p s r o il e e te d nt at. ra h m lize ad o t pha ide nd e he nit he ospho is ab onia, and chro d Fc epitate res me time a PO, . T m and r ph 3id m H) d a c d ip o Fe a r s w 4S . a c ^ 3 .S 8 -9 * I 43 X! -t-s S3 p. ^ s fe * o s -&. s t + H ^ c3 j-i O ^ ^2 s ^ ^^ S g -T3 r o JS 8 g fe -r M cs o fe w^. +3 1 "trw o " t rt c3 S ^ *i2 S> ^2s| !l** 2^S| iisl S^.tfJ O be feM ^ fl +3 Sl s ! ^ T3 -M iH ^ C/3 ^ '^ COS r^ ^-^ 3 .t7 N ^ T3- , ^ 9 fc 18 ^> 60 ^ cd rg C > C8^ s^ ~J Q j M * W P C J d ^ M PL, .. 2 "M ^3 . Ce 2 O 3 is white or grayish-white, soluble in acids and formed by igniting Ce 2 (CO 3 )3 , Ce_>(C 2 O 4 ) 3 or Ce6 2 in an atmosphere of hydrogen. Cerous salts are white and form color- less solutions in water. Ceric oxide, CeO 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 eerie salts, forming cerous salts. Fixed alkali hydroxides and ammonium sulphide precipitate, from solu- tions of cerous salts, the white cerous hydroxide, turning yellow by absorption of oxygen, with formation of eerie hydroxide. The precipitate is insoluble in excess of the fixed alkalis (distinction from Al and Gl). The precipitation is hindered by the presence of tartaric acid (distinction from yttrium). Ammo- nium hydroxide precipitates a basic salt; if H 2 O 2 is added before neutralizing a reddish brown precipitate is formed (delicate test for cerium). Alkali car- bonates precipitate cerous carbonate, soluble in excess of the fixed alkali car- bonates. Oxalic acid forms cerous oxalate, white, from moderately acid solutions, soluble in hot (NH4) 2 C 2 O 4 , but reprecipitated on dilution with cold water. The oxalate is less soluble in hot than in cold water. A concentrated solution of K : SO 4 forms the double sulphate, K 3 Ce(SO 4 ) 3 , white, sparingly soluble in water, insoluble in K SO, solution (distinction from Gl). NajS^O; does not precipitate cerium salts. BaCO does not precipitate cerous salts in the cold, but precipitates them completely on boiling. Ceric salts are completely precipitated by BaCO 3 in the cold. Alkali hypochlorites precipitate cerous salts as the yellow eerie hydroxide. If cerous nitrate be boiled with PbO 2 and HNO.j , eerie nitrate, a deep yellow solution, is formed (delicate test for cerium). Cerium gives no absorp- tion spectrum, but the spark spectrum shows several brilliant lines. 155. Columbium (Niobium). Cb = 93.5. 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, specific gravity, 7.06 at 15.5 Melting point, 1700 (Cir. B. of S., 1915). By ignition in the 156. DIDYMIUM. 199 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 nitrohydrochloric acid, but is readily soluble in hot con- centrated sulphuric acid, forming a colorless solution (Roscoe, C. N., 1878, 37, 25). It forms several oxides, CbO , CbO.. and Cb.O 5 . Columbic acid (anhy- dride) Cb_O 5 , 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. CbOj , black, is prepared by strongly igniting Cb>O 5 in a current of hydrogen. Cb^Os , not too strongly ignited, is soluble in acids, from which solutions NH 4 OH and (NH 4 ) 2 S pre- cipitate columbic acid containing some ammonia. By mixing CbzOs with char- coal and heating in a current of chlorine, a mixture of CbOCl :i and CbCl s is obtained. CbCl 5 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 SO 4 (distinction from tantalum). Cb^O^ not previously ignited dissolves in HF ; which solution, when mixed with KF , the HF being in excess, gives a double fluoride, 2KF.CbF 5 ; 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 titanium or tantalum compounds. Fusion of columbic acid with the alkalis gives the columbates, the potassium salt being quite soluble in water and in potassium hydroxide; the s dium 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. AgNO 3 gives a yellowish-white precipitate, CuSO 4 a green precipitate. Cb 2 O 5 in presence of HC1 or H SO,, gives a blue to brown color with Sn or Zn, due to partial reduction of the Cb (dis- tinction 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 FeSOj. ( Neodymium. Nd = 144.3. Valence three. 156. Didymmm = i _ _ n _ , T , (Praseodymium. Pr=140.9. Valence three. Specific gravity, 6.544. Melting points, Neodymium, 840?; praseodymium, 940? (Cir. B. of S., 1915). 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 praseody- mium. 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; Kruess, A., 1892, 265,. 1). Concerning the separation of didymium compounds, see Dennis and Chamot (/. 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 dif- ferent spectra, but which, united, gave the spectrum of didymium. Didymium oxide absorbs water to form the hydroxide, which absorbs CO 2 from the air, but does not react alkaline to litmus. The salts are soluble in water to a reddsih solution. The saturated sulphate solution 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 prevented by tartaric acid. Oxalic acid precipitates didymium salts completely, soluble with difficulty in HC1 . The double potas- sium 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), Schele (Z. anorg., 1898, 17, 319), Boudard (C. r., 1898, 126, 900), Demarcay (C. r., 1898, 126, 1039), and Brauner (C. N., 1898, 77, 161). 200 ERBIUM GALLIUM GLUCINUM. 157. 157. Erbium. Er = 167.7. Valence three. Erbium has been prepared in the form of a dark gray powder. Specific gravity, 4.77 at 15. (Meyer, Monatsch., 20, 793, 1899). As oxide or earth it is de- scribed by Cleve (C. r., 1880, 91, 381) as that yttrium earth the most beautiful rose colored. It forms a characteristic absorption spectrum, and a spark spec- trum with sharp lines in tfie orange and green. This earth has not been thor- oughly studied and quite probably consists of the oxides of several metals (Bois- baudran, C. r., 1886, 102, 1003; Soret, C. r., 1880, 91, 378; Crookes, C. N., 1886, 64, 13). The oxide gives upon ignition an intense green light; it is not fusible or volatile. 158. Gallium. Ga=69.9. 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 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 (C. 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, C. r., 1878, 86, 475). The oxide, Ga,O 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 4 . 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. BaC0 3 precipitates gallium salts in the cold. K 4 re(CN) 8 gives a precipitate, insoluble in ITC1 , noticeable in very dilute solutions (1-175,000). H,S does not precipitate gallium salts from solutions acid with mineral acids; from the acetate or in presence of ammonium acetate the n'lnte sulphide, Ga 2 S 8 , is precipitated; (NH 4 ) 2 S precipitates the sulphide. Gallium chloride, GaCL, , is a colorless salt, melting at 75 and volatilizing at 215 to 220. The vapor density indicates the molecule to be Ga 2 Cl 6 , which decomposes to GaCl 3 at about 400 (Friedel and Kraft, C. 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 HzSOj 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. Gluciimm (Beryllium). Gl = 9.1 . Valence two. Specific gravity, 1.85 (Humpidge, Proc. Roy. Soc., 1871, 39, 1). Melting point, 1350 ? (Cir. B. of S., 1915). It is a white malleable metal, obtainable in hexagonal crystals (Nilson and Pettersson, B., 1878, 11, 381 and 906). It was first dis- covered in 1797 by Vauquelin from beryl. The powdered metal takes fire when heated in air, burning with great brilliancy. It dissolves readily in dilute acids and also in alkalis with evolution of hydrogen. It does not decompose steam even at a red heat. It is a strongly positive element, in general properties between aluminum and the alkaline earths; as lithium is between the alka- 160. INDIUM. 201 line earths and the alkali metals. It should be classed with the alkaline earths. It is found in chrysoberyl, G1(A1O 2 ) 2 , in phenakite, GLSiO, , and in some other silicates. It is prepared by heating the chloride, G1C1 2 , with Na in a closed iron crucible (Nilson and Pettersson, /. c.); or by heating the oxide, G1O, with Mg (Winkler, B., 1890, 23, 120). The oxide, G1O , is obtained by igniting the hydroxide. It is a white infusible powder, soluble in acids and in fixed alkalis. The hydroxide is prepared by precipitating the salts with NH 4 OH , soluble in the fixed alkalis and in ammonium carbonate, concentrated; precipitated on dilution and boiling (distinction and separation from Al). The metal is soluble in acids except that when in the compact form it is scarcely attacked by HNOs . The hydroxide is soluble on continued boiling w r ith NH 4 C1 , form- ing G1C1 2 . The more common salts of glucinum are soluble in water to a solution having a sweetish taste. The carbonate and phosphate are insoluble, the oxalate and sulphate soluble, the existence of a sulphide is doubtful. Solu- tions of glucinum salts are precipitated by the alkalis, the precipitate being soluble in excess of the fixed alkalis. The alkali carbonates precipitate the carbonate, soluble in concentrated ammonium carbonate, reprecipitated on diluting, boiling and adding an excess of NH 4 OH (Joy, Am. S., 1863, (2), 36, 83). The salts are not precipitated by H 2 S , but are precipitated by (NH 4 ) 2 S as the hydroxide. BaCO 3 does not precipitate Gl salts in the cold, but precipitates them upon boiling. G1C1 2 melts at about 600 and sublimes at a white heat, forming white needles. The oxide has not been melted or sublimed. Gl usually occurs as a silicate with aluminum. The mass is fused with alkali carbonate, acidified with HC1 and the Al and Gl chlorides filtered from the SiO 2 . An excess of ammonium carbonate precipitates both metals, but redissolves the Gl . After repeating this separation several times pure glucinum hydroxide, G1(OH) 2 , is obtained upon boiling off the ammonia. The hydroxide thus obtained is ignited and weighed as the oxide. 160. Indium, In = 114.8. Valence three. Specific gravity, 7.11 to 7.28 at 20.4. Melting point, 155 (Cir. B. of S., 1915). Indium was discovered in Freiberg zinc blende by Reich and Richter (/. pr., 1863, 89, 441; 90, 175; 1864, 93, 480), by use of the spectroscope. It is found chiefly as sulphide, never native, in the Freiberg blende to the extent of about 0.1 per cent. It is found in a few other places, but in much smaller amounts (Boettger. /. pr., 1866, 98, 26). In the preparation of indium the Freiberg zinc is dissolved in HC1 or H 2 SO 4 , leaving an excess of the zinc. When no more hydrogen is evolved, the mass is digested for a day or more with the excess of Zn , whereby the indium is obtained as a precipitate with Pb , Cu , Cd , Sn , As , Fe and Zn . This precipitate is dissolved in nitric acid and evaporated with sulphuric acid; then taken up with water, separating from lead. The solution is precipitated with NH 4 OH , which precipitates the In and Fe ; this precipitate is dissolved in HC1 and boiled for some time with NaHSO;, . The indium sulphite is obtained as a fine crystalline powder, which is treated with HNO 3 and H 2 SO 4 , forming indium sulphate, from which the metal is precipitated by zinc (Bayer, A., 1871, 158, 372; Boettger, /. pr., 1869, 107, 39; Winkler, J. pr., 1867, 102, 276). Indium is a grayish-white metal, very soft, makes a good mark on paper, is ductile, easily fusible, less volatile than Zn or Cd . It is less electro-positive than Zn or Cd and hence it is precipitated from its solutions by both these elements. In the air pr in water it is rather more stable than zinc. Heated in the air it burns with a violet flame and brown smoke, forming the oxide, In 2 O 3 . Indium does not decompose water at 100. At a red heat it combines with sulphur and the halogens. By ignition with charcoal or in a current of hydrogen it is reduced to the metal from its compounds. It is soluble in HC1 and H 2 SO 4 , evolving H ; in HNO S , evolving NO . In the reactions of its salts indium deports itself quite similar to Fe"' and Al . Its most characteristic property is its spectrum; two lines, an indium a, intense blue, and an indium 0, less intense violet (Schroetter, J. pr., 1865, 95, 441). In 2 O is brown when hot, light yellow when cold, slowly soluble in cold acids, rapidly when heated. Indium salts are precipitated by the alkalis as In(OH) 3 , soluble in excess of the fixed alkalis, 202 LANTHANUM SCANDIUM. 161. reprecipitated by boiling or treating with NH 4 C1 . Tartrates prevent the precipitation by alkalis. Alkali carbonates precipitate the indium carbonate, soluble in ammonium carbonate, but reprecipitated on boiling. BaCO :i pre- cipitates the indium completely as a basi salt (separation from Co , Ni , Mn , Zn and Fe"). Phosphates form white precipitates from neutral solutions. H 2 S precipitates from neutral solutions, or solutions acid with acetic acid, yellow indium sulphide. In alkaline solutions H 2 S , or in neutral solutions (NH 4 ) 2 S , forms a white precipitate containing In 2 S 3 . Yellow In 2 S 3 boiled with (NH 4 ) 2 S X becomes white and is partly dissolved. Upon cooling the solution a bulky white precipitate separates out. K 4 Fe(CN) 6 gives a white precipitate; K 2 CrO 4 gives a yellow precipitate; K 2 Cr 2 O 7 , K 3 Fe(CN) 6 and KCNS do not form precipitates. 161. Lanthanum, La = 139.0. Valence three. Specific gravity, 6.163. Melting point, 810 ? (Cir. B. of S., 1915). In general appearance and properties very similar to Ce . It is prepared almost exclusively from cerite. By treating the mineral with an insufficient quantity of HNO 3 , a solution rich in La may be obtained. The cerium is precipitated from the solution by alkali hypochlorite. The filtrate is converted into the sulphate and separated from Ne and Pr sulphates by fractional crystallization, the latter being more soluble (Holzman, J. pr., 1858, 75, 346). Fractional precipitation with NH 4 OH is also used to separate La from Ne and Pr , the latter precipitat- ing first (Cleve, Bl, 1874, 21, 196; 1883, 39, 287). The metal is prepared from the chloride, LaCl 3 , by electrolysis or by ignition with potassium. The igni- tion point of La is higher than that of Ce ; it is also not so readily attacked by HNO 3 . In cold water La is slowly attacked, but in hot water the action is violent (Winkler, B., 1890, 23, 787). With aluminium, lanthanum forms a crystalline white alloy which is stable in air and insoluble in nitric acid (Muth- man and Beck, A., 46, 331, 1904.) The oxide, La^Os , is a white powder, readily soluble in acids; with water it forms the hydroxide, La(OH) 3 , which reacts alkaline towards litmus and absorbs CO 2 from the air. La(OH) 3 is soluble in a solution of NH 4 C1 (similar to Mg(OH) 2 ). The salts are colorless. K 2 SO 4 and H 2 C 2 O 4 form precipitates with lanthanum salts as with cerium salts. Fixed alkalis precipitate lanthanum salts as La(OH) 3 , white, insoluble in excess of the reagent and not changing color on exposure to the air (distinction from Ce). Alkali carbonates precipitate La2(CO 3 ) 3 , insoluble in excess. BaCO 3 precip- itates the salts completely in the cold. NH 4 OH precipitates basic salts. H 2 S forms no precipitate; (NH 4 ) 2 S precipitates the hydroxide. Lanthanum gives a num- ber of characteristic lines in the spark spectrum (Bettendorf, A., 1889, 256, 159). 162. Neodymium. Nd = 144.3. See Didymium (156). 163. Praseodymium. Pr = 140.9. See Didymium (156). 164. Samarium. Sa = 150.4. Valence three. Samarium was found in 1879 by Boisbaudran from didymium earths by its peculiar spectrum (C. r., 1879, 88, 323). According to Crookes,. (C. r. 1886, 102, 1464), it consists of at least two elements and is found in all yttrium earths. Its salts are light yellow, giving an absorption spectrum of six bands (Kruess, B., 1887, 20, 2144). In its chemical properties it is more similar to Nd and Pr than to Y . It is separated from Nd and Pr by the fractional precipitation of the hydroxide, basic nitrate, oxalate and sulphate; which separate before the corre- sponding Nd and Pr compounds. Melting point, 1300-1400 (Cir. B. of S., 1915). 165. Scandium. Sc = 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 167. TANTALUM TERBIUM. 203 salt remaining 1 in solution. The oxide, Sc,0 3 , is a white flocculent infusible powder, readily soluble in warm acids. The solutions of the salts show m> 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-. NaoCO;; gives a bulky white precipitate, soluble in excess of the reagent. H,S is without action, but (NH 4 ) 2 S precipitates the ]i]/1700, <1755 (Cir. B. of S., 1915); stable in air at ordi- nary temperature, but igniting when heated; attacked by vapors of Cl , Br , I and S . Sparingly soluble in dilute acids, easily soluble in concentrated acids; insoluble in the alkalis (Nilson, B., 1882, 15, 2519 and 2537; Kruess 170. TITANIUM. 205 and Nilson, B., 1887, 20, 1665). Thorium forms one oxide, ThO 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(N0 3 ) 4 , *>re deliquescent. The chloride is a white body melting at a white heat and then subliming in beautiful white needles (Kruess and Nilson, I.e.). The sulphate is soluble in five parts of cold watiT. The carbon-file, oxalate and phosphate are insoluble in water; the a.ralatc is scarcely soluble in dilute mineral acids. Alkali hydroxides or sulphides precipitate thorium hydroxide, Th(OH) 4 , 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 C0 3 readily repre- cipitates upon warming. BaC0 3 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 1 ),C,O 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,SO 4 solution, sparingly soluble in cold water, readily soluble in hot water. HF precipitates Th.F 4 , insoluble in excess, gelatinous, becoming crystalline on standing. Boiling freshly precipitated Th(OH) 4 with KF in presence of HF forms K 2 ThF 6 .4H 2 O , a heavy fine white precipitate almost insoluble in water. The distinguishing reactions of thorium are the precipitation with oxalates and with K 2 SO 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 = 48.1. Valence three and four. Titanium is found quite widely disiribut3d as lutiie, Lrookite, anatase, titanite, titaniferous iron, FeTiOs, 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', melt- ing point, 1800 (Cir. B. of S., 1915). Heated in the air it burns with an unusu- ally brilliant incandescence', sifted into the flame it burns with a blinding bril- liance. 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 com- bines 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, TiO 2 , analogous to CO 2 and SiO 2 . It occurs more or less pure in nature as rutile, brookite and anatase; it is formed by igni- tion of the hydrated titanic acid or of ammonium titanate (Woehler, J., 1849, 268). Ignition of TiO 2 in dry hydrogen gives Ti 2 O 2 , an amorphous black powder, dissolving in H 2 SO 4 to a violet-colored solution (Ebelmen, A. Ch., 1847, (3), 20, 392). TiO is formed when TiO 2 is ignited with Mg:2TiO 2 -f- Mg = TiO + MgTiO 3 (Winkler, B., 1890, 23, 2660). Other oxides have been reported. Titanic acid, TiO 2 , is a white powder, melts somewhat easier than SiO 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. TiO 2 acts as a base, forming a series of stable salts; also as an acid, forming titanates. TiCli is a colorless liquid, fuming in the air; it boils at 136.41 (Thorpe, J. C., 1880, 37, 329); it is decomposed 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 SO 4 no TiF., is volatilized (distinction from SiF 4 ). When evaporated with HF alone, TiF 4 is volatilized. The double potassium titanium fluoride, K 2 TiF 6 , formed by fusing TiO 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 alkali hydroxides, carbonates and sulphides as the hydrated titanic acid, insoluble in excess of the precipitants and in ammonium salts. BaCO 3 gives the same 206 URANIUM. 171 - precipitate. K 4 Fe(CN) gives a reddish-yellow precipitate; K,Fe(CN) 6 a yellow precipitate. NaaHPO, precipitates the titanium ahinjxt cviuplctcliL even in the presence of strong- HC1 . An acid solution of TiO, 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. ITS is without action. The solution reduces Fe"' to Fe" , CM" 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 FeSO t in the reducing flame a Mood-red bead is obtained. Titanium is very readily detected in minerals as follows. 0.1 gram of the finely powdered mineral is mixed with 0.2 gram of finely powdered sodium fluoride and 3 grams sodium pyrosulphate added without mixing. The crucible is heated until copious sulphuric acid fumes are evolved. The fused mass is rapidly cooled and heated with 2-3 c.c. dilute sulphuric acid and 10 c.c. water added. The solution is dividqd into two parts and a few drops of hydrogen peroxide added to one part. A yellow color is produced by the titanium. Chlorides, bromides and iodides interfere with this very delicate reaction (Weber, Z., 40, 799, Noyes, J. Soc. Ind., 10, 485). 171. Uranium, U = 238.2. Valence four and six. Specific gravity, 18.685 (Zimmermann, A., 1882, 213, 285). Melting point, <1850 (Cir. B. of S., 1915). Found in various minerals; its chief ore is pitch-blende, which contains from 40 to 90 per cent of UsOg . Prepared by fusing UClt 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 incan- descence to UgOg J unites directly with Cl , Br , I and S when heated; soluble in HC1 , HoSO 4 and slowly in HNO 3 . Uranous oxide, UO2 , formed by ignit- ing 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, UOs , 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 U 3 O 8 = UO 2 2UO 3 . Uranium acts as a base in two classes of salts, uranous and uranyl salts. Uranous salts are green and give green solutions, from w y hich alkalis precipitate uranous hydroxide, insoluble in excess of the alkali; alkali carbonates precipitate U(OH) 4 , soluble in (NH 4 ) 2 CO 3 ; with BaCO 3 the precipitation is complete even in the cold. H 2 S is without action; (NH 4 )J3 gives a dark-brown precipitate; K 4 Fe(CN) 6 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 HNO 3 , Cl , HC1O 3 , Br , KMnO 4 , etc. Gold, silver and platinum salts are reduced to the free metal. The hexad uranium (U^ 1 ) acts as a base, but usually forms basic salts, never normal: we have TJ0 2 (NO :J ) 2 , not TT(NO 3 ) ; UO L S0 4 , not TT($O 4 ) 3 . These basic salts w T ere formerly called uranic salts, but at present (ir0 2 )" is regarded as a basic radical and called tirunjil, and its salts are called uranyl salts, e.g., UO 2 C1 2 uranyl chloride, (TTO 2 ) S (PO 4 ) 2 uranyl orthophosphate. Solutions of uranyl salts are yellow; KOH and NaOH give a yellow precipitate, uranates, K 2 TJ 2 7 and Na.ILO, , insoluble in excess. Alkali carbonates give a yellow precipitate, soluble in excess; BaC0 3 and CaC0 3 give TJ0 3 . H 2 S does not precipitate the uranium, but slowly reduces uranyl salts to uranous salts (Formanek, A., 1890, 257, 115)1 (NH 4 )..S gives a dark-brown precipitate. K 4 Fe(CN) gives a reddish-brown precipitate. Used in the analysis and separation of uranium compounds (Fresenius and Hintz, Z. angeic., 1895, 502). Sodium phosphate gives a yellow precipitate. The hexad uranium acts as an acid toward some stronger bases. 171 rt, 8. VANADIUM- 207 Thus we have K 2 TT2O 7 and Na^U^Oy , formed by precipitating uranyl salts with KOH and NaOH ; compare the similar salts of the hcxad chromium, K 2 Cr 2 O 7 and Na2Cr 2 O 7 . Other oxides of uranium are described, but are doubtless combinations of UO 2 and UO 3 . Zn , Cd , Sn , Pb , Co , Cu , Fe , and ferrous salts reduce uranyl salts to uranous salts. Solutions of Sn, Ft, Au, Cu, Hg and Ag are reduced to the metal by metallic uranium (Zimmcrmann, I.e.). For method of recovery of waste uranium compounds, see Laube (Z. angew., 1889, 575). 171a. Vanadium. V = 51.0. Valence two to five. 1. Properties. -Specie gravity, 5.8 (Moissan, C. r., 122); melting point 1720 (Cir. B. S., 35, 1915). A grayish non-magnetic powder; slowly oxidized in the air, rapidly on ignition with formation of V 2 O 5 . It forms with chlorine the dark brown tetrachloride. 2. Occurrence. It is often found in iron and copper ores and in some clays and rare minerals, e.g., vanadinite, (Pb 5 Cl(VO 4 )3) ; volborthite, (Cu,Ca,Ba) 3 (OH) 3 VO4 +6H 2 O) ; mottramite (a hydrous vanadate of lead, copper, and other divalent elements, of uncertain formula, (R 3 (VO 4 ) 2 .3R(OH) 2 )) . 3. Preparation. The vanadium ores are treated chiefly for the preparation of ammonium vanadate and vanadic acid. The ores are fused with KN0 3 , form- ing- potassium vanadate. This is precipitated with Pb or Ba salts and then decomposed with H.SO 4 . The vanadic acid is neutralized with NH,OH and precipitated with NH 4 C1 , in which it is insoluble. This upon ignition gives V,O pure (\Yohler, A., 1851, 78, 125). The metal is prepared from the dichlo- ride, VCL , by long-continued ignition in a current of hydrogen. 4. Oxides. Vanadium forms four oxides: VO , gray; V 2 3 , black; V0 2 , dark blue; and V,O,, , dark red to orange red. 5. Solubilities. Vanadium is not attacked by dilute HC1 or H 2 SO 4 ; concen- trated H,SO 4 gives a greenish-yellow solution; HNO 3 a blue solution. VO dis- solves in acids to a blue solution with evolution of hydrogen. V 2 O 3 dissolves in dilute HC1 to a dark greenish-black solution. Chlorine forms with V 2 O 3 , VOC1 3 and V 2 O r , . VO., dissolves in acids to a blue solution, from which solu- tions Na.CO, gives a precipitate of V 2 2 (OH) 4 + 5H 2 O , grayish-white mass, losing 4H 2 O at 100 and turning black, soluble in acids and alkalis. V 2 O 5 exists in several modifications with different solubilities in water, the red modification being- soluble in 125 parts of water at 20 (Ditte, C. r., 1880, 101, 698). Vanadic acid forms three series of salts, ortho, meta and pyro, analogous to the phosphates. Most salts are the metavanadates. The ortho compounds are quite unstable, readily changed to the meta and pyro compounds. Alkali vanaclates are soluble in water, the ammonium vanadate least soluble and not at all in NH 4 C1 . 6. Reactions. Solutions of vanadic acid produce brown precipitates with alkalis, soluble in excess to a yellowish-brown color. Potassium ferrocyanide gives a green precipitate, insoluble in acids. Tannic acid gives a blue-black solution, which is said to make a desirable ink. Ammonium sulphide precipi- tates V,S 5 , brown, soluble with some difficulty in excess of the reagent to a reddish-brown thio salt. From this solution acids reprecipitate the brown vanadic sulphide, V 2 S 5 . If to a solution of a vanadate, neutral or alkaline, solid NH 4 C1 be added, the vanadium is completely precipitated as NH 4 V0 3 , ammonium metavanadate, crystalline, colorless, insoluble in NH 4 C1 solution; upon ignition in air or oxy- gen, pure vanadic oxide, V 2 B , is obtained. 7. Ignition. Borax gives with vanadium compounds in the outer flame a colorless bead, yellow if much vanadium be present; in the inner flame a green bead, or brow r n when vanadium is present in large quantities and hot, becoming green upon cooling. All the lower oxides of vanadium ignited in air or oxygen give V 2 O 5 . 8. Detection. Vanadium will almost always be found as a vanadate (2) and is detected by the reactions used in its purification (3) ; also by the reactions with reducing agents, forming- the colored lower oxidized compounds (10). 208 YTTERBIUMYTTRIUM. 171 rt, 10. 9. Estimation. (1} It is precipitated as basic lead vanadate and dried at 100. (2} It is precipitated as ammonium vanadate, NH 4 VO 3 , in strong NH 4 C1 solution, ignited to the oxide V-Oa , and weighed. 10. Oxidation. Zn , in solutions of vanadates with dilute H 2 SOi , reduces the vanadium to the tetrad, a green to blue solution, then greenish-blue to green, the triad, and finally to lavender blue, the dyad. H 2 S reduces vanadates to the tetrad with separation of sulphur. Oxalic acid and sulphurous acid also reduce vanadates to the tetrad, the solution becoming blue. 172. Ytterbium. Yb = 173.5. 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 (/?., 1879, 12, 550; 1880, 13, 1433) describes its preparation from euxenite and its separation from Sc . 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 (J/., 1884, 5, 1). The oxide, Yb,,0 3 , is a white powder, slowly soluble in cold acids, readily upon warming. The Jiydro.riilc forms a gelatinous precipitate, insoluble in i^H 4 OH but soluble in KOH . It absorbs CO, from the air. The nitrate melts in it water of crystallization and is very soluble in water. 173. Yttrium. Y = 88.7. 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-. '., 1889, 38, 474). The metal has been prepared by electro- lysis of the chloride; also by heating the oxide, Y 2 3 , with Mg- (Winkler, B., 1890, 23, 787). Melting point, 1490 (Cir. B. of S.~ 1915). 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 elements (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 (Welsbach, M., 1883, 4, 641; Dennis and Chamot, J. Am. Soc., 1897, 19, 799). Yttrium salts are precipitated by the ai -alis and by the alkali sulphides as the hydroxide, Y(OH) ? , a white bulky pre- cipitate, insoluble in the excess of the reagents (distinction from Gl). The oxide and hydroxide are readily soluble in acids; boiling with NH,C1 causes solution of the hydroxide as the chloride. The alkali carbonates precipitate the carbonate Y 2 (CO 3 )3 , soluble in a large excess of the reagents. If the solu- tion 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 3 forms no precipitate in the cold (distinction from Al , Fe'" , Cr'" , Th , Ce , La , Nd and Pr). Hydro- fluoric 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 carbonate, decompose with KC1 , and filter from the SiO 2 . Neutralize the filtrate and precipitate the Y , La and Ce as oxalates with (NH 4 ) 2 C 2 O 1 . 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 NH,OH . Ignite and weigh as the oxide. In order to effect complete separations the operations should be repeated several times. 175. TffE CALCIUM GROUP. 09 174. Zirconium. Zr = 90. G. 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. CA, 1831, 48, 337). Melting point, 1700 ? (Cir. B. of S., 1915). The metal exists in three modifications: crystalline, graphitoidal and amorphous. The amorphous zirconium is a velvet-black powder, burning when heated in the air. Acids attack it slowly even when hot, except HF , which dissolves it in the cold. It forms but one oxide, ZrO 2 , analogous to SiO 2 and TiO> . ZrO 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 SiOa 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 K 2 SO 4 , and after resolution in acids precipitated by NH 4 OH and ignited to ZrO_> (Berlin, J. pr., 1853, 58, 145; Roerdam, C. C., 1889, 533). ZrO 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.oO ;( , Y 2 O 3 , etc., to form the mantles used in the Welsbach gas-burners (Drossbach, C. C., 1891, 772; Welsbach, \ forming hydroxides with evolution of heat. Mg oxidizes rapidly in the air 210 THE CALCIUM GROUP. 176. 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 e;icli 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. Ch., 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. Ch., 1893, 12, 125); SrC0 3 in 90,909 parts at 18 (Kohlrausch and Rose, Zeit. phys. Ch., 1893, 12, 241); CaC0 3 in 80,040 parts at 23.8 (Hollemann, /. c.); MgC0 3 in 9,434 parts (Chevalet, Z., 1869, 8, 91). The presence of NH 4 OH 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 ).,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 //.). 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, 5^), 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 186, 4. BARIUM. 211 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 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(OH) 2 in ammonium salts. 183. Solutions containing Ba , Sr , Ca , and Mg , and phosphoric, oxalic, boric, or arsenic acid, necessarily have an acid reaction, because these phosphates, oxalates, etc., are soluble only in acids; such solutions are precipitated by ammonium hydroxide or by any agent which neutralizes the solution, and, consequently, we have precipitates of this kind in the third group (145 and //.) : CaClo + H,PO 4 + 2NH 4 OH = CaHPO 4 + 2NH 4 C1 + 2H 2 O CaH 4 (PO 4 ) 2 + 2NH 4 OH = CaHPO 4 + (NH 4 ) 2 HPO 4 + 2H 2 O . If excess of the ammonium hydroxide be added the precipitate is Ca 3 (P0 4 ) 9 . Barium and strontium react like calcium. In the case of a magnesium bait 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 only at a high temperature in the case 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=l37.37. Valence two. 1. Properties. Specific gravity, 3.75 (Kern, C. N., 1875, 31, 243); melting point, 850 (Cir. B. of ., 1915). 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. c.). 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 barite, BaSO 4 , and witherite, BaCOj . 3. Preparation. (1) 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 1 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, Ba. ().n'(/r,v ami Iun1ro.ri<1c>i.- Harinm oxide is aeted upon b\ \\ater \\ith evolution of heat .iud formation of the hydroxide, \\hieh is soluble in about .'(> parts of cold \\ater and in its oxvn weight of hot water ( KVxiMislhfil and Rnehlinann. /., 1S70. :ni). Hariuiu peroxide, BaO . is very sparingly soluble in \\ater (Sehone. .1., is;?. 192. 857); soluble in aeids with formation of H,O_, . r. Salts. Most of the soluble sails of barium are permanent : the acetate is o!llorcseont. r riio elilorido, bromide. l>romnto. iodide. sul]>hide. t\MT(H'vanide, nitrato. liypo])bosphito, oblorato. neetnie, niul pbenvlsul- ]>bate, ;in freely soluble in water: the carbonate. *u!j>liatc. sulphite, chronuilr,* |u 1 ios]>hite. jihosphate, oxalate. iodate, and sitico-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. $ 1890, 143, .VJ1V, likewise the nitrate in strong hydrochloric and nitric acids. The chloride and nitrate are insolu- ble in alcohol. (). Reactions, a. The fixed alkali hydroxides precipitate only con- centrated solutions of barium salts ("iM. No precipitate is formed with ammonium hydroxide (45). The alkali carbonates precipitate barium carbonate. BaCO, , white. The precipitation is promoted by heat and by ammonium hydroxide, but is made slightly incomplete by the presence of ammonium salts (Yogel. J. /'/.. lSTn>. 7, 1,V>). Barium Carbonate BaCO, is a valuable reagent for special -pur chietly 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 tilt ra- 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 arid, and fully precipitated by sulphuric acid, the tilt rate must yield no fixed residue. This reagent removes sulphuric acid (radical) from all sulphates in solution to which it is added (c): Na.,S0 4 -f BaCO., == BaS0 4 -f Na,CO, . When salts of non-alkali metals are so decomposed, of course, they are left insoluble, as carbonates or hydroxide-, nothing remaining in solution: FeSO, + BaCO s = BaS0 4 -f FeCO s Fe,(S0 4 ) s -f ;.BaCO :i + :;H,0 ;;BaSO 4 -f 2Fe(OHV, + :;CO, The chlorides of the third group, except Fe" , are decomposed by barium carbonate; while the metals of the fourth group (/inc. manganese, cobalt, nickel), are not precipitated from their chlorides by this reagent. Tartaric *Kohlrausob and Hose. Z. j>/ij/s. ('/., 1S'.V>, 12, 241 : Schwoit/er, '/.., 18W, 29, 414. 186, 7. BARIUM. 213 810, 850 ? (C'ir. B. of S., 1915), 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. c.). 2. Occurrence. Strontium occurs chiefly in strontianite, SrCO 3 , and in celestite, SrSO, . 3. Preparation. First isolated in 1808 by Davy by electrolysis of the hydrox- ide (Tran-s. Royal $oc., 345). It is made by electrolysis of the chloride (Frey, A., 1876, 183, 367); by heating a saturated solution of SrCl, with sodium amalgam and distilling off the mercury (Franz, 7. 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, SrO 2 .8H 2 O , is made by pre- cipitating the hydroxide with H 2 O 2 ; at 100 this loses water and becomes SrO 2 , a white powder, melting at a red heat, used in bleaching works (Conroy, J. 8oc. /(?.., 1892, 11, 812). 5. Solubilities. a. Metal. Strontium decomposes water at ordinary tem- perature (Winkler, I. c.), it is soluble in acids with evolution of hydrogen, ft. 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 Zeitsctirift fur Rve&tt&udcer, 1881, 49, 257). The peroxide is scarcely soluble in water or in ammonium hydroxide, soluble in acids and in ammonium chloride. 187, 6/1. STRONTIUM. 215 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, be). The sulphate is very sparingly soluble in water (1-10,090 at 20.1) (Hollemann, Z. pliys. Ch., 1893, 12, 131); yet sullidently 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 HNO.< ; insoluble in alcohol. Strontium nuosilicate 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(OH)., , 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 hulides 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, 60). 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 216 CALCIUM. 187, 6i. chromate is about 100 times as soluble as the strontium chromate (Fre- senius and Rubbert. 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 rriuwtn. 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 , at the red end of the spectrum; the two red lines, Sr ft and Sr y, and the blue line, Sr 6 , 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 flame 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). (3) By treating the nitrates with equal volume of absolute alcohol, and ether (188, 9). For separation from barium see 186, 9. 188. Calcium. Ca = 40.07. Valence two. 1. Properties. Specific gravity, 1.6 to 1.8 (Caron, C. r., 1860, 50, 547). Melting point, 810 (Cir. B. of S., 1915). 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, forming an amalgam. 2. O3cunence. Found in the mineral kingdom *as a carbonate in marble, limestone, chalk and aragonite; 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 , 400 parts granulated zinc and 100 parts Na 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. 217 hydroxide, the carbonate (limestone), nitrate, and all organic calcium salts. The corresponding 1 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, CaO,.sH,0 , is made by adding- hydrogen peroxide or sodium peroxide to the hydroxide: Ca(OH)!, + H 2 O, == Ca0 2 + 2H,O (Conroy, J. Nor. Intl., 1892, 11, 80S).' Drying at 130 removes all Ilie watei 1 , leaving- a \vhi1c powder, CaO, , which at a red heat loses half its oxygen (Schoene, A., 3877, 192, 257). It cannot be made by heating the oxide in oxygen or with potassium chlorate (186, 4). r>. Solubilities. a. Metal. Calcium is soluble in acids with evolution of hydrogen; it decomposes water, evolving hydrogen and forming Ca(OH), . &. Oxide and hydroxide. CaO combines with dilute acids forming cor- responding salts, it absorbs C0 2 from the air becoming CaCO, .* In moist air it becomes Ca(OH) 2 , the reaction taking place rapidly and with increase of volume and generation of much heat in presence of abundance of water. The hydroxide, Ca(OH) 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, 7?., 1874, 7, 817); and in 1712 parts at 100 (Lamy, /. 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 CaCO.,), 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. S., 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 2 or SO, below 350. (Veley, J. C., 1893, 63, 82K t Goldhammer, C. C., 1888, 708; Droeze, B., 18T7, I o. 330; lioisbuudran, A. C7)., 1874. (5), 3,477 Kohlrausch and Rose, Z. phys. C?i., 1893, 12, 241 ; Raupenstrauch, M., 1885, 6, 663). 218 CALCIUM. 188, 6fl. 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, /. 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 (5&), i. e. potassium hydroxide will form a precipitate with calcium sulphate since the sulphate requires less water for its solution than the hydroxide (56 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(OH) 2 , is used as a reagent for the detection of carbon dioxide (56 and 228, 8). 6. 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 (Sidersky, 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 1 a strong 1 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 SO 4 ions must be present in solution in large excess of CO 3 ions. With strontium also an excess of SO 4 ions is required, although not so great as in the case of calcium. For barium, however, equilibrium demands that the concen- tration of CO 3 ions exceed that of SO 4 . This condition is already fulfilled when an alkali car- bonate is added to BaSO 4 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. c. See 5^. d. By the action of alkali phosphates, solutions of calcium are not distinguished from solutions of Iwriiim or strontium. e. Pure sodium sulphide, Na L .S , gives an abundant precipitate with calcium salts; even with CaSO 4 . The precipitate is Ca(OH),: CaCL + 2Na 2 S + 2H.O = Ca(OH) 2 + 2NaCl + 2NaHS . The acid sulphide, NaHS , does not precipitate calcium salts (Pelouze, A. Cli., 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. /. Calcium chloride, fused, is much used as a drying agent for solids, liquids and gases. Chlorinated lime, or bleaching powder, CaCJUO (Kingzett, J. C., 1875, 28, 404), is much used as a bleaching agent and as a disinfectant, g. Neutral or ammoniacal solutions of arsenites form a precipitate with calcium salts (distinction from barium). A solution of calcium salts including solu- tions of calcium sulphate in ammoniacal solution is precipitated by arsenic acid as CaNH 4 As0 4 (distinction from strontium after the addition of sulphuric acid) (Bloxam, C. N., 1886, 54, 16). h. 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 O T (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, CaS0 4 , plaster of Paris; which on being moistened forms the crystalline CaS0 4 .2H 2 O , 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 /?, 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 amyl MAGXE8WM. 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 = 24.32. Valence two. 1. Properties. Specific gravity, 1.75 (Deville and Caron, A. Ch., 1863, (3), 67, 346); melting point, 651 (Cir. B. of S., 1915). 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 N , P , As , S and Cl . It forms alloys with Hg and Sn, forming compounds which decompose water. 2. Occurrence. Magnesite,' MgC0 3 ; dolomite, CaMg(C0 3 ) 2 ; brucite. Mg(OH) 2 ; epsom salts, MgSO 4 .7H 2 0; and combined with other metals in a great variety of minerals. 3. Preparation. (1) By electrolysis of the chloride or sulphate (Bunsen, A., 1852, 82, 137). (2) By ignition of the chloride with sodium or potassium (Wohler, A., 1857, 101, 562). (3) Mg 2 Fe(CN) 6 is ignited with Na 2 C0 3 , 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(OH) 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 -f C0 2 + H 2 MgC0 3 + ^ (Ballo, B., 1882, 15, 3003): it is also attacked b} f he acid alkali carbonate?, as NaHC0 3 , to form MgC0 3 , Na 2 C0 3 and H (Ballo, I. c.). Soluble ID ammonium salts: Mg + 3NH 4 C1 == NH 4 MgCL + 2NH 3 + H 2 . With the halogens it acts tardily (Wanklyn and Chapman, J. C., 18G6, 19, 141). &. Oxide and hydroxide. Insoluble in water, soluble in acids. Mg(OH) 2 is soluble in 111,111 parts of water at 18 (Kohlrausch and Eose, Zeit. phys. Ch., 1893, 12, 241). In contact with water the oxide is slowly changed to the hydroxide, Mg(OH) 2 , and absorbs C0 2 from the air. Sol- uble in ammonium salts: * Mg(OH) 2 + 3NH 4 C1 = NH 4 MgCL + 2NH 4 OH . 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. Reactions, a. The fixed alkali hydroxides and the hydroxides of barium, strontium and calcium precipitate magnesium hydroxide, Mg(OH) 3; * The conditions here are the same as in the case of Mn(OH) 2 , 134, 6a, footnote. 189, 6/. MAGNESIUM. 221 white, gelatinous, from solutions of magnesium salts ; insoluble in excess of the reagent but readily soluble in ammonium salts, the magnesium pass- ++ + - +-f ing into the negative ion: Mg(OH) 2 + 4NH 4 Cl=-(NH 4 _) 2 MgCl 4 +2NH 4 OH . With ammonium hydroxide but half of the magnesium is precipitated, the remainder being held in solution in the acid ion by the ammonium ++ - + - ++ + + salt formed in the reaction: 2Mg S0 4 ~f 2NH 4 OH Mg(OH) 2 -f- (NH t )2 Mg(S0 4 ) 2 (Rheineck, DingL, 1871, 202, 268). The fixed alkali carbonates precipitate basic magnesium carbonate, Mg, (OH) 2 (C0 3 ) 3 , variable to Mg,(OH),(CO :{ ) 4 : 4MgS0 4 + INa.CO, -f H,0 = = Mg 4 - (OH),(C0 3 ) 3 + Na,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 -(- 5Na 2 C0 3 + 2H 2 == Hg 4 (OH) 2 (C0 3 ) 3 + MgH 2 (CO,) 2 + 5Na 2 S0 4 . On boiling, the acid carbonate is decomposed with escape of CO, . Ammonium carbonate docs 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). 1). 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) 3 a crystalline precipitate is slowly formed, magnesium ammonium phosphate Mg-ini 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 (5&). 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. 9 1869, 8, 125). e. Magnesium xulphide is decomposed by water, and magnesium salts are not precipitated by hydrosulphuric acid or ammonium sulphide; but Mg-0 -f- H 2 O (1-10) absorbs H 2 S , forming- in solution MgH^S. , which readily gives off HJ3 upon boiling (a very satisfactory method of preparing H.S absolutely arsenic free) (i)ivers 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: MgSO', + 2Na.,S -f 2H,O = Mg(OH), + Na,S0 4 + 2NaHS (IVIou/c. .-I. CIi.. 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 CALIFORNIA COLLEi* 222 MAGNESIUM. 189, 60. hydrochloric ac : d (7). g. Soluble arsenates precipitate magnesium salts in de- portment similar to the corresponding phosphates. 7. Ignition. Magnesium ammonium phosphate when ignited loses ammonia and water, and becomes the pyrophosphate: SMg-NHiPC^ = Mg 2 P 2 7 + H a O + 2NH 3 . The carbonate loses CO 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 -f- H,O = MgO + 2HC1; 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, MgNH 4 P0 4 , by Na 2 HP0 4 . 9. Estimation. After removal of other non-alkali metals, magnesium is pre- cipitated as MgNH 4 PO 4 , then changed by ignition to Mg 2 P 2 O 7 (magnesium pyrophosphate) and weighed as such. Separated as MgCL from KC1 and NaCl by solution in amyl alcohol, evaporated with H 2 S0 4 and weighed as MgS0 4 (Riggs, Am. $., 1892, 44, 103). It is estimated volumetrically by precipitation as MgNH 4 PO 4 , drying at about 50 until all free NH 4 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. Soc., 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, Ni, Cu, Cd , Zn , Gl , Ba , Sr , Ca , Rb , K , Na , and Li . In some cases the reaction takes place with explosive violence. From 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 , Ni , Au , Pt , and Pd (Scheibler, B., 1870, 3, 295; Villiers and Borg, C. r., 1893, 116, 1524). * Winkler, B., 1890, 23, 44, 130 and 772 ; 1891, 24, 888. t Kern, C. N , 1876, 33, 112 and 236. t Seubert and Schmidt, A., 1892, 267, 818. 190 ANALYSIS OF THE CALCIUM GROUP. 223 -M fe ^5 +-* e Pit a s ca ac P 9,1 F. c ^^te^ ^ d ^.- g a ! is Lfr-* a ^^^- s 02 00 te I 8* t-05 TJ q 02 224 btRtiCTWNS FOR AXAL78TS WITH NOTES. 191, DIRECTIONS FOR ANALYSIS OF THE METALS OF THE CALCIUM GROUP. (THE ALKALIXE EARTHS.) 191. Manipulation. To the filtrate from the fourth group in which H 2 S (192, 1) gives no precipitate (138) add NH 4 OH and ammonium carbonate as long as a precipitate is formed: BaCl 2 -f- (NH 4 ).,C0 3 = BaCO s + 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 : SrCO, + 2HC 2 H 3 2 == Sr(C 2 H,0 2 ) 2 + C0 2 + H 2 . To a small portion of the acetic acid solution add a drop of K.,Cr,0 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,0 2 ) + K 2 Cr 2 T + ELO = 2BaCr0 4 + 2X0,11,0, + 2HC 2 H 3 2 . Filter, wash the precipitate and dissolve it in HC1. Test a portion in the flame and precipitate the barium in the remainder as barium sulphate, with a drop of sulphuric acid. To the filtrate from the barium chromate add NH 4 OH 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 Sr0 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 ammonium (1-4) 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. Considerable amounts of the metals of this group, especially barium and strontium, may be precipitated with the second group on account of the formation of sulphuric acid by the oxidation of hydrogen sulphide, espec- ially by means of ferric chloride. As much as 15 mg. of barium may be pre- cipitated in this manner. A still further loss of 15 mg. of barium as well as considerable quantities of calcium and strontium as carbonates may occur during the precipitation of the iron group. Smaller quantities may be pre- cipitated as sulphate or carbonate with the ammonium sulphide group. (Curt- man and Frankel, /. Am. Soc., 33, 724 (1911).) If large quantities of the metals in the preceding groups; especially iron, are present, barium, strontium and calcium may fail to be detected for this reason, by the ammonium carbonate method of separation. The method of Curtman and Frankel should then be used. See 197. 2. Do not boil after the addition of ammonium carbonate, as this will drive off 194, . ; . DIRECTIONS FOR ANALYSIS WITH NOTES. ammonium hydroxide and carbonate, increasing the solubility of the CaCO 3 (note 3 and 178). 3. The precipitation of bivium, strontium and calcium by ammonium car- bonate in the presence of ammonium chloride is not as complete as would be desir- able in v/ry 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 carbonate, greatly re- duces the solubility of the precipitat d carbonates, yet even with th se the precipi- tation is not absolute, though more nearly so with strontium than with barium and c tlcium. i hus, in c/umtit tie analyses, if barium and calcium are precipitated as < urbonates, it must be done in the absence of ammonium chloride or sulphate, and the precipitate washed with water containing ammonium hydroxide. As much as 10 mg. of barium may remain in solution in the presence of much ammonium salts. (Curtman and Frankel, J. Am. Soc., 33, 724 (1911).) 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,Cr,O 7 , may be neglected and the filtrate from the barium, yellow, at once divided into two portions and tested for Sr and Ca . lleprecipitation 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 ) i CO 3 , an excess of ammonium hydroxide should be added to prevent the liberation of CO 2 when the ammonium car- bonate is added. 7. Strontium sulphate is so sparingly soluble in water (187, 5c) that its precipitation by CaSO 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 SrS0 4 will not precipitate solutions of calcium salts. Solutions of strontium and barium salts (except SrSO 4 ) are all precipitated by CaSO 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 (NH 4 ) 2 SO 4 ; or in other words upon a solution of CaSO, (1-287). A solution of SrSO 4 (1-10,000) may be present, but is not precipitated by i NHi) 2 C 2 O 4 . The presence of a great excess of (NH,) 2 SO 4 prevents the precipitation of traces of calcium salts by (NH 4 ) 2 C 2 O 4 . 193. Manipulation. To a portion of the filtrate from the carbonates of Ba , Sr , arid Ca add a drop or two of (NH 4 ) 2 S0 4 . A slight precipitate indicates a trace of barium. To the filtrate a few drops of (NH 4 ) 2 C 2 4 are added. A trace of calcium is indicated by a slight precipitate. Filter if a precipitate is obtained and test the filtrate for Mg with Na 2 HP0 4 . A white crystalline precipitate MgNH 4 P0 4 is evidence of the presence of magnesium. The other portion of the filtrate from the carbonates of Ba , Sr , and Ca is reserved to be tested for the alkali metals (211). 194 . No'es. 1 . B y some, magnesium is classed in the last or alkali group instead of in the alkaline earth group. It is not precipitated by the (NH^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 subdivisi n of that group than as belonging to the alkali group ( 7 > and ff. ) . 2. Traces of t?a , Sr and Ca may remain in solution after adding (NH 4 ) 2 CO 3 and warming; due to the solvent action of the ammonium salts present. To pre- 226 SEPARATION OF BARIUM, STRONTIUM, AND CALCIUM. 194, 3. vent the.se traces giving a test lor magnesium with Na2HPO 4 a drop or two of (NH 4 )2SO 4 is added to remove barium or strontium and a few drops of (NH 4 ) 2 C2O4 to remoye calcium. The precipitate (if any forms) is removed by nitration, before the Na 2 HPO 4 is added. 3. The precipitate of P^g '> V "4 does not always form rapidly, if only small amounts of Mg are present, and the solution should be allowed to stand. Rubbing the sides of the test tube with a glass stirring rod promotes the precipitation. 4. The precipitation of Mg as MgtftI 4 FO 4 is fairly delicate (1-71,492) (Kissel, Z., 1F69, 8, 173 ; but not very characteristic, as the phosphates of nearly all the metals are white and insoluble in water. Hence the reliability of this test for mag- nesium depends upon the rigid exclusion of the other metals (not alkalis) by the pre- vious processes of analysis. The precipitate should be carefully examined with a magnifying glass to ascertain if it is crystalline which is characteristic of magnesium ammonium phosphate. Small amounts of aluminum may be present at this point and a small amorphous precipitate may be aluminum phosphate. It should be filtered off and treated with a few drops of acetic acid which readily dissolves mag- nesium ammonium phosphate, but not aluminum phosphate. The filtrate should be neutralized with ammonia and a few drops of Na 2 HPO 4 added. On standing for some time small transparent crystals are deposited on the walls of the test tube if magnesium is present. 5. Lithium phosphate is not readily soluble in water or ammonium salts and may give a test for magnesium. See 210, Gd. 195. The unlike solubilities in alcohol, of the chlorides and nitrates of barium, strontium and calcium enable us to separate them very closely by absolute alcohol, and approximately by " strong alcohol," as follows: Dissolve the carbonate precipitate in HC1 , evaporate to dry ness 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 Avith alcohol (56-, 186, 187 ami 188). Residue: BaCl 2 . Dissolve in water test with CaSO 4 , SrSO 4 , K 2 Cr 2 O 7 , etc. Filtrate: SrCl 2 and CaCl 2 . Evaporate to dryness, change to nitrates by adding a few drops of HNO 3 . 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 3 ) 2 . Precipitation by CaSO 4 in water solution: flame test, etc. Filtrate: Ca(NO 3 ) 2 . Precipitation by H 2 SO 4 in alco- hol solution, by (NH 4 ) 2 C 2 O 4 , 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 ),S0 4 and a little NH 4 OH . 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 phosphate* (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 '188, 5c), 200. THE ALKALI GROUP. 227 THE ALKALI GROUP (SIXTH GROUP). Potassium. K = ;J9.10. Caesium. Cs = 132.81. Sodium. Na = 23.00. Rubidium. Rb = 85.45. Ammonium. (NH 4 )'. Lithium. Li = 6.94. 198. The metals of ilio alkalis are highly combustible, oxidizing qnickly 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, cesium 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 OH, 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). 228 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 nitrof erricyanide, 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 = 39.10. Valence one. 1. Properties Specific gravity, 0.875 at 13 (Baumhauer, B., 1873, 6, 655). Melting point, 62.3 (Cir. B. of S., 1915). 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 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.Mg-Cl, + 6H 2 O; kainite, K 2 S0 4 .MgSO 4 .MgCl2 + 6H 2 , etc. As an important constituent of many plants grape, potato, sugar-beet, tobacco, fumaria, rumex, oxalis, etc. 3. Preparation. (1) 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 3 or KOH with iron car- bide: 6KOH + 2FeC 2 = 6K + 2Fe + 2CO + 2C0 2 + 3H, (Castner, C. N., 1886, 54, 218). (//) By reduction of the carbonate or hydroxide with Fe or Mg (Winkler, B., 1890, 23, 44). 4. Oxides and Hydroxide. Potassium oande* K 2 , is prepared by carefully * The existence of tfce oxides M',Q of K, Na and Jito is disputed (Erdmann apd Koetlmer, 4* 1896, 294, 55), 205, 65. POTASSIUM. 229 heating 1 potassium with the necessary amount of oxygen (air) (Kuhnemann, C. C., 1863, 491); also by heating- K,O 4 with a mixture of K and Ag (Beketoff, C. C., 1881, 643). It is a hard, gray mass, melting above a red heat. \Yater changes it to KOH with generation of much heat. I'oldxxiuin lii/dro.ridc, KOH, is formed by treating K or K 2 O with water; by boiling a solution of K,CO :! with Ba , Sr or Ca oxides; by heating K,C0 3 with Fe.Og to a red heat and decomposing the potassium ferrate with water (Ellershausen, C. 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 supcro.t'idc, K._>O 4 , is formed when K is heated in contact with abundance of air (Harcoiirt. ,/. (7., 1862, 14, 267); also by bringing K in contact with KNO 3 heated until it begins to evolve (liolton, C. A'., 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 O and AgoO formed (Harcourt, I. c.). 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 , Fe , Ag and Pt to the oxides (Bolton, I. c.' Brodie, Proc. Roy. &oc., 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 and 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 hyHroxides. 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. &.--The potassium salts of HCN, H 4 Fe(CN) 6 , H 3 Fe(CN) 6 , and HCNS find extended application in the detection and estimation of many of the heavy metals. Tartaric acid, H 2 C 4 H 4 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 , granular-crystalline. If the solution be alkaline, acetic or tartaric acid should be added to strong acid reaction. The test must be made in absence of non-alkali bases. The precipitate is increased 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.,C 4 H 4 6 , but not by acetic acid, or at all by alcohol of fifty per cent. 230 POTASSIUM. 205, 60. Picric acid, C 6 H 2 (N0 2 ) 3 OH , or preferably its sodium salt, precipitates, from solutions not very dilute, the yellow, crystalline potassium picrate, C 6 H 2 (N0 2 ) 3 OK, soluble in 260 parts of water at 15 C. (Reichard, Z. 40, 25), insoluble in alcohol, by help of which it is formed in dilute solutions. The solution must be nearly neutral to avoid precipitation of the slightly soluble picric acid (soluble in 160 parts water). The dried precipitate 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 3 Co(NO 2 ) 6 , 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 NaN0 3 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 precipitate solutions of the metals of the first four groups as sulphides except: Hg' becomes HgS and Hg, Fe'" becomes FeS and S, and Al and Cr form hydroxides. The sulphides of arsenic, antimony and 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,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 -{- 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 mercurosum incompletely. /. 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 typo 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 * One cc. of cobaltous nitrate solution and three cc. 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. 231 reduced to Fe" salts with liberation of iodine. Arsenic acid is merely reduced to arsenous acid with liberation of iodine. Potassium chlorate is used as a source of oxygen and as an oxidizing agent in acid solutions. Sodium perchlorate, NaClO 4 , precipitates from solutions of potassium salts i>otC0 3 = BaC0 3 -f 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 232 SODIUM. 205, 8. to conceal the violet at full heat. The interposition of a blue glass, or prism filled with indigo solution, sufficiently thick, entirely cuts off the yellow light of sodium, and enables the potassium name 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 w^ater, remove the barium with NH 4 OH 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 /?, 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 (66); 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 such. (2) It is precipitated and weighed as the double chloride with platinum. (3) If present as KOH or K 2 CO 3 it is titrated with standard acid (Kippenberger, Z. angew., 1894, 495). (4) It is precipitated with H 2 SiF 6 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, KC1O 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 Rb . For oxidizing action of K 2 4 see 4. 206. Sodium. Na = 23.00. 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, 97.5 (Cir. B. of S., 1915). Boiling point, 742 (Perman, C. N., 1889, 59, 237) CALIFORNIA COLLEQi 206, 6d. SODIUM. PHARMACY 233 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 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: 2Na + 2H 2 O = 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, C. 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, NaCI; as Chili saltpeter, NaNO 3 ; 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, Na.O , is formed by burning sodium in oxygen or in air and heating again with Na to decompose the Na 2 O 2 (205, 4, footnote). Sodium hydroxide, NaOH , is formed by dissolving the metal or the oxide in water (Rosenfeld, J. pr., 1893, (2), 48, 599); by treating a solution of sodium carbonate with lime; by fusion of NaN0 3 with CaC0 3 , CaO and Na,CO 3 are formed and the mass is then exhausted with water: by igniting Na 2 C0 3 with Fe 2 O 3 , forming sodium ferrate, which is then decom- posed with hot water into NaOH and Fe(OH) 3 (Solvay, C. C., 1887, 829). It is a white, opaque, brittle crj-stalline body, melting under a red heat. The fused mass has a sp. gr. of 2.13 (Filhol, A. Ch., 1847, (3), 21, 415). It has a very powerful affinity for w r ater, gradually absorbing water from CaCl 2 (Mtiller- Erzbach, B., 1878, 11, 409). It is soluble in about 0.47 part of water according to Bineau (C. r., 1855, 41, 509). Sodium peroxide, Na,0 2 , is formed by heating sodium in CO,, free air or oxygen (Prud'homme, C. C., 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,O 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. Keactions. a. As reagents sodium hydroxide and carbonates act in all respects like the corresponding potassium compounds, which see. ft. By the greater solubility of the picrate and acid tnrlrafc of sodium, that metal is separated from potassium (205, 6ft). 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 KNO 3 , etc. d. Sodium phosphate, Na.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, 5c), soluble in acids). Solu- 234 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, 1i. 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 O at 17.5 and in 40.66 parts at 100 (Stolba, ., 1872, 11, 199); hence is not precipitated by fluosilicic acid except from very concentrated solutions (separation from K). j. Sodium chlorplatinate, Na 2 PtCl 6 , 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 hydro- chloric acid from the point of a glass rod on a slip of glass, treated with two drops of solution of chlorplatinic acid, left a short time for spontaneous evaporation and crystallization, and observed under the microscope, k. Sodium picrate soluble in 10 parts of water is used as a reagent for potassium salts (Richard, Z., 40, 377). ~k. Solution of potassium pyroantimonate, K 2 H 2 Sb 2 T , 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 IDC 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 y 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. 235 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 (Glc). 9. Estimation. (1) 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,H 2 Sb 2 O 7 and dried and weighed as Na 2 H 2 Sbo0 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 3 gas, 0.589 (Fehling, 1, 384); of the liquid, 0.6234 at (Jolly, A., 1861, 117, 181). The liquid boils at 33.7, at the liquid has a tension of 4.8 atmospheres (Bunsen, Pogg., 1839, 46, 95). Liquid 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 4 being monovalent and acting in many respects similar to K and Na . At 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 3 ; 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 -f 3KNO 3 -f 2H 2 = SKAlOo -f 3NH 3 ; by the reduction with the hydrogen of the zinc-copper couple; by boiling organic compounds containing nitrogen with KMn0 4 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 OH , 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. angeiu., 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. AMMONIUM. 07, 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 4 OH . 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, fc 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 , u , Cd , Co , Ni , and Zn , the precipitate is redissolyed by an excess of the ammonium carbonate. &. Dilute solutions of picric acid with ammonium hydroxide form in- tensely colored yellow solutions, a precipitate of ammonium picrate being 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(NO)(CN) 5 , added to a mixture of NH 4 OH 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 hydrosulphuric 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(NO)(CN) 5 . c. Ammonium nitrite, NH 4 NO 2 , is used in the preparation of nitrogen (235, 3); ammonium nitrate in the preparation of nitrous oxide, N 2 O , " laughing- gas" (237). d. Ammonium phosphate, as a reagent, acts similarly to sodium phosphate. When sodium phosphate, Na 2 HPO 4 , is used to precipitate metals in the presence of ammonium hydroxide, a double phosphate of the metal and ammonium is frequently formed as MnNH 4 PO 4 , MgNH 4 PO 4 , etc. By some chemists microcosmic salt, NaNH 4 HPO 4 , is preferred to sodium phosphate, Na,HPO 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, 6k. AMMONIUM. 237 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 OH , 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(OH) 3 is diminished by the presence of NH 4 C1 (124, 6a, and 117). ); 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. (7., 1868, 21, 29). In presence of Cu the O of the air oxidizes the nitrogen of ammonia to a nitrite (/) (Berthelot and Saint-Gilles, A. Ch., 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 cyanides (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 = NC1 3 + 4HC1 (&) 8NH 3 + 3C1 2 6NH 4 C1 + N 2 2NH 4 C1 + 301, = 8HC1 + N 2 (c) 2NH 3 + I 2 = NH 4 I + NH 2 I (d) 6NH 4 OH + 3I 2 = 5NH 4 I + NH 4 I0 3 + 3H 2 (e) 6NH 4 OH + 8HMnO 4 = 3NH 4 NO S + 8MnO(OH) a + 5H 2 O (f) 12Cu + 2NH 3 + 90 2 = 12CuO + 2HN0 2 + 2H 2 O (g) 3HN0 3 + 8A1 + 8KOH = 8KA10, + 3NH 3 + H 2 O (ft) HCN + KOH + H 2 = NH 3 + KCHO 2 (formate). 208. Caesium. Cs = 132.81. 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, C. 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(CN) 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 gra^ash-white solid, very deliquescent, absorbs CO 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 , and the acid tartrate, CsHC 4 H 4 O , being least soluble and used in preparation of the salts free from the other alkali metals. 240 RUBIDIUM LITHIUM. 208, 6 . 6. Reactions. In all its reactions similar to the other fixed alkalis. 7. Ignition. Caesium sails ceic;- ine non-luminous flame violet. The spec- trum gives two sharply defined lines, Cs a and Cs ,1, in the blue and a third faint line in the orange-red Us t , 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. (1) As the double platinum chloride; (2) as the chloride with RbCl , estimation of the amount of Cl and calculation of the relative amounts of the metals; (3) as the sulphate obtained from ignition of the acid tartrate and treatment with H,S0 4 (Bunsen, Poyg., 18G3, 119, 1). 209. Rubidium. Rb = 85.45. Valence one. 1. Properties. Specific gravitij, 1.52 (Bunsen, A., 1863, 125, 367). Melting point, 38 (Cir. B. of S., 1915); 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 Mg" or Al (Winkler, B., 1890, 23, 51; Beketoff, B., 1888, 21, c, 424). 4. Oxide and Hydroxide.- The oxide Rb 2 O 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(OH), upon Rb^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 phosphomolj'bdate. The alum is less soluble than the corresponding potassium alum. G. 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 ,,?; two less intensive in the outer red, Rb > and Rb J; a fifth Rbtin 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 Cl and calculating the proportion of the metals; (2) as the double platinum chloride. 210. Lithium. Li = 6.94. Valence one. 1 Frcpsitie;. Specific gravity, 0.5936, the lightest of all known solid bodies (Bunsen and Matthiessen, A., 1855, 94, 107). Melting point, 186 (Cir. B. of S. 1915); does not vaporize at a red heat. It is a silver-white metal with a grayish tinge; harder than K or Na, 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 in appearance and in its chemical properties to K 210, LITHIUM. 241 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 Cl , Br , I , O , S and dry CO,. . 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, triphylite, or petalite. Traces are found in a great many minerals, in mineral spring's, 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 NH 4 C1 or KC1 (Giintz, C. r., 1S9.'5, 117, 7!!2). The metal is also obtained by ignition of the carbonate with Mg 1 , but the metal is at once vaporized and oxidized. 4. Oxide and Hydroxide. It forms one oxide, Li^O , 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,C0 3 in 75 parts at 20; Li 3 PO 4 in 2539 parts pure water and 3920 parts ammoniacal water, more soluble in a solution of NH 4 C1 than in pure water (Mayer, A., i66o, i,8 ? 193). Lithium chloroplatinate and lithium picrate are very soluble in water (Richard Z., 40, 383). G. 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 N.a and Li (Rammelsberg, A. Ch., 1818, (2), 7, 157). The phosphate dissolved in HC1 is not at once precipitated by neutraliz- ing with NH 4 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-re^ 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 fed 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 /?. The color tests have an intensity intermediate between those of sodium and potassium. 8. Detection. #.// the sprctroscope. 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 Rb 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 Rb and Cs appear. Evaporate to dryness the filtrate from the precipitate of the platinum double salts, add oxalic acid and ignite, moisten with HC1, 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. 242 DIRECTIONS FOR ANALYSIS WITH NOTES. 210,9, 9. Estimation. After separation from other elements it may be weighed as a sulphate, carbonate or phosphate, Li 3 PO 4 . It may also be estimated by the comparative intensity of the lines 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 the 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 filtrate con- taining the magnesium and the alkalis to dryness, ignite gently to remove the ammonium salts. Dissolve the residue in water and add Ba(OH) 2 to precipitate the magnesium as Mg(OH) 2 (177 and 182). After filtration, the excess of barium in the filtrate 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 Reagent (207, 6fc) or by the precipitation with HgCl 2 (207, 6/). Study 207, 6, 7, 8, and 9. 214. The rare metals of the Alkali Group: lithium, rubidium, and 215. DIRECTIONS FOR ANALYSIS WITH NOTES. 243 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 nitrate 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. PART III -THE NON-METALS. 216. BALANCING EQUATIONS IN OXIDATION AND EEDUCTION. Oxidation and reduction always involves a change in valence. When the valence of an electropositive element is increased it is said to be oxidized, and conversely, when its valence is reduced, reduction has taken place. It is believed that each bond or valence is produced by the presence of a unit charge of electricity on the atom or ion. A ferrous ion may be represented as Fe ++ while a ferric ion would be Fe + + + and the oxidation or reduction of iron would therefore consist in adding to or subtracting a unit charge of electricity from the atom. Similarly the valence of negative ele- ments is proportional to the number of negative charges of electricity on their atoms. It is assumed that during oxidation and reduction unit charges of negative electricity carried by small corpuscles or electrons pass from one atom to the other. The positively charged masses of atoms are very much larger than the negative electrons. When an atom loses a nega- tive electron, its positive charge is relatively increased and it is oxidized; when it gains a negative electron it is reduced. The valence of an electro-negative element would therefore be increased when it is reduced and reduced when it is oxidized. When an element can pass from the positive to the negative condition there may be no change in valence during reduction or oxidation. 2NfH^ -f 30 = =N^" + + 3 = + 3H 2 0. The metals in salts generally act as electro-positive elements while the acid elements or radicles are electro-negative. The latter therefore may act as oxidizing agents towards the former. In general in a reaction in- volving oxidation and reduction one element is oxidized and another is reduced and the gain in valence of one element is exactly equal to the loss in valence of the other element. This is a necessary consequence of the transfer of negative electrons from one element to the other. Statement of Bonds in Plus and Minus Numbers,* according to chemical polarity, positive and negative (see 3 footnote). A bond, that is a unite 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 , * O. C. Johnson, C. N., 1880, 42, 51. See also Ostwald, Grundr. allg.Chem., 3te Aufl., 1899, S. 439. 216. BALANCING EQUATIONS IN OXIDATION AND REDUCTION. 245 that of water, (H +I ) 2 0~ n . (The plus sign is understood when no sign is written before the valence number.) 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.* The Oxidation of any element is shown by an increase, and its Eeduction 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, + H,S0 3 + 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 SO, 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). The valence of the acid element in an acid may always be found from the anhydride. In this case we have: H 2 S0 3 =S0 2 +H 2 , the valence of the sulphur in S0 2 being 4. 3Sn C' 2 + FI0 3 + 6HC1 = 3Sn C1 4 +HI + 3H 2 O Here also the three atoms of tin gain six bonds, and these are furnished by the iodine of the HIO., . 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~').] 8HMn0 4 + 5AsH 3 + 8H 2 S0 4 = 5H 3 AsO 4 + 8MnSO 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 * If there is polarity in the union of like atoms with each other in forming an elemental molecule, the sum must be zero, as in the formation of the molecules of compounds. 246 BALANCING OF EQUATIONS, 217. 25, after gaining 40. They must then have had 15 in the first place (25 40 -15). That is, the atom of arsenic in arsenous hydride has -3 bonds (As-'"H 3 ). SnCl 2 + HgCL 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. b. Oxygen has always two negative bonds. 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 3 , each atom of Bi has three positive bonds (e), and each atom of S has two negative bonds (/). 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 fol- 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 zero. Hg"(NvO-" a ) 2 .2 + 10 12 = Bi'" 2 (SviO-" 4 ) 3 .6 + 18 24 = Ba"(MnViiO-" 4 ) 2 .2 + 14 16 = Fe'"(NvO-" 3 ) 3 .3 + 15 18 = As'" 2 S-" 3 .6 6 = If the above is understood, the rule for balancing equations is easily explained. The number of bonds changed in one molecule of each shows the number 218, 4- BALANCING OF EQUATIONS. 247 of the molecules of the other which must be taken, the words each and other referring to the oxidizing and reducing agents. 218. A few equations will illustrate the application of the rule. (1) 3As 4 + 20HN0 3 + 8H,O = 12H 3 As0 4 + 20NO The arsenic in one molecule gains 20 bonds, therefore 20 molecules of HN0 3 are taken. The nitrogen loses three bonds, therefore three molecules of As 4 are taken. The valence of the nitrogen and arsenic may be found from their anhydrides: 2HNO 3 = H 2 O + N 2 O 5 and 1 H 3 AsO 4 = 3H 2 O + As 2 O 5 Equations of this kind may also be balanced by considering that the arsenic must be oxidized to the pentoxide, one molecule of arsenic requiring ten atoms of oxygen. Two molecules of nitric acid furnish three atoms of oxygen, as follows : i'HNOs = H 2 O + 2NO + 3O Three molecules of arsenic must, therefore, be taken, requiring thirty atoms of oxygen, which will be furnished by 20 molecules of nitric acid. To convert a molecule of arsenic pentoxide into arsenic acid requires 3 molecules of water. The 3 molecules of arsenic will therefore require 18 molecules of water. As the 20 molecules of nitric acid furnish 10 molecules of water, 8 more must be added. (2) 6Sb + 10HN0 3 = 3Sb 2 5 + 10NO + 5H 2 O The antimony gains five bonds, therefore five molecules of HN0 3 would be taken, and since the nitrogen loses three bonds, three of antimony would be taken, but since we cannot write Sb 2 O 5 with an odd number of atoms of antimony, we double the ratio and take six and ten. (3) 3H 2 S + 8HN0 3 = 3H 2 S0 4 + 8NO + 4H 2 The S in the first member has 2 negative bonds (a and d); in the second member it has 6 positive, gaining 8 bonds; hence 8 molecules of HNO 3 must be taken. The nitrogen in the first member has five bonds, and in the second it has two. The difference is three, therefore just three molecules of H 2 S must be taken. Further, the reaction may be explained as follows: The sulphur in the first member has two bonds (valence of two), but nega- tive because combined with hydrogen (two atoms) to form a. definite com- pound; in the second member it has six bonds (valence of six), but positive because combined with oxygen (SO 3 or JJQ H S O^' The valence of the hydrogen does not change and hence in the reaction one molecule of H 2 S gains eight bonds. The nitrogen in the first member has five bonds (valence of five), but positive because combined with oxygen (N 2 O 5 or H N~ ? ) ; in the second member it has two bonds, still positive because combined with oxygen. The valence of the hydrogen and oxygen does not change, hence in the reaction one molecule of HNO., loses three bonds. Now the number of bonds gained by the H L ,S (8) must equal the bonds lost by the HN0 3 (3). The least common multiple, twenty-four, indicates the least possible total change of valence for en oh compound; this requires that three molecules of H,S and eight of HNO 3 be taken, giving for the products three molecules of ELSO* and eight of NO with four of water to complete the equation. (4) 3Sb 2 S 8 + 28HN0 3 = 3Sb 2 5 + 9H 2 S0 4 + 28NO -f 5H 2 O In this case, both the Sb and the S in the molecule gain bonds, and must be 248 BALANCING OF EQUATIONS. 218, o. considered. It is plain (from d and e) that each atom of Sb gains 2 bonds, and the two in the molecule will gain 4. The S in Sb 2 S 3 has 2 negathe bonds, and in the second member (in H 2 SO 4 ) it has 6 positive bonds, a gain of 8. The three atoms in the molecule will gain three times eight, or 24 bonds; to this add the 4 which the Sb has gained, and we have 28 bonds gained by one molecule of Sb 2 S 3 ; hence 28 molecules of HNO 3 must be taken. We take 3 of Sb 2 S 3 for reasons explained in the first equation. Further explain as follows: In this case both the Sb and the S gain in valence (oxidized). Each atom of antimony gains two bonds, a total gain of four. Each atom of sulphur gains eight, a total gain of twenty-four; or a gain for one molecule of Sb 2 S 3 of twenty-eight bonds. As in the previous illustration, the nitrogen loses three bonds. The least common multiple, eighty-four, indicates that for the reaction each compound must undergo a change of at least eighty-four bonds. This requires for the Sb 2 S 3 three mole- cules, and for the HN0 3 twenty-eight molecules. The products are as indicated in the equation. (5) 2Ag 3 As0 4 + HZn + 11H 2 S0 4 = 2AsH 3 + GAg 1 + llZnS0 4 + 8H 2 O The silver loses three bonds, and the arsenic in changing from plus five to minus three loses eight bonds; this added to the three that the silver loses makes eleven, therefore eleven atoms of zinc are taken, and since the zinc gains two, two molecules of silver arsenate are taken. (6) 2MnO + :Pb 3 4 + COHN0 8 = 2HMn0 1 + 15Pb(N0 3 ), -}- 14H 2 The manganese gains five bonds, therefore five molecules of Pb 3 4 are taken. The three atoms of lead in one molecule of Pb 3 O 4 have in all eight bonds, but a like amount has only six in the second member, being a loss of two, there- fore two molecules of MnO are taken. (7) 2MnBr 2 + 7Pb0 2 + 14HN0 3 = 2HMn0 4 + 2Br 2 + 7Pb(N0 3 ) 2 + GH 2 The manganese gains five bonds and the bromine gains one, the two atoms gaining two, adding this to the five that the manganese gains makes a total gain of seven bonds, therefore seven of Pb0 2 are taken. The lead loses two, therefore two of MnBr, are taken. (8) MnS + 4KNO 3 + K 2 C0 3 , fusion = K 2 MnO 4 + K 2 S0 4 + 4NO + K 2 C0 3 The manganese gains four bonds and the sulphur eight, making twelve; therefore twelve of KNO 3 would be taken, and since the nitrogen loses three bonds, three of MnS would be taken, but since three is to twelve as one is to four, the latter amounts are taken. (9) 2Cr(OH) 3 + 3Mn(N0 3 ) 2 + 5K 2 CO 3 , fusion = 2K 2 Cr0 4 + 3K,MnO 4 + 6NO + 5C0 2 + 3H 2 O This is a peculiar and instructive equation. The nitrogen loses six bonds, but since the manganese in the same molecule gains four, the total loss is only two, therefore two of Cr(OH) 3 are taken. The chromium gains three, therefore three of Mn(N0 3 ) 2 are taken. (10) 3Ag + 4HN0 3 = 3Ag-N0 3 + NO + 2H 2 The rule here calls for three of silver and one of nitric acid, but three more of unreduced nitric acid are needed to combine with the silver, making four in all. (11) 2FeI 2 + 6H 2 S0 4 , cone., hot = Fe 2 (S0 4 ) 3 + 3S0 2 + 2I 2 + 6H 2 The rule here calls for two of FeI 2 and three of H 2 S0 4 , but three more of H 2 SO 4 that are not reduced are needed to combine with the iron, making six in all. (12) 3HN0 8 + 8A1 + 8KOH = 3NH 3 + 8KA10 2 + H 2 The nitrogen has five bonds in HN0 3 , and in NH 3 it has minus three, losing eight, therefore eight of aluminum are taken. The aluminum gains three, therefore three of HNO 3 are taken. 218, 15. BALANCING OF EQUATIONS. 249 (13) 3BiON0 3 + 11A1 + 11KOH = 3Bi + 3NH 3 + 11KA1O 2 + H 2 The bismuth loses three bonds and the nitrogen loses eight, therefore eleven of aluminum are taken; the aluminum gains three, therefore three of the BiONO 3 are taken. ' (14) Mn(X + 4HC1 = MnCL + CL + 2H 2 The manganese loses two bonds and the chlorine gains one, but two more of unoxidized HC1 are needed to combine with the manganese, hence four are taken. (/5) 2CrI 3 + G4KOH + 27CL = 2K 2 Cr0 4 + GKI0 4 + 54KC1 + 32H 2 O The chromium gains three bonds and the iodine (in the molecule) gains twenty-four, therefore twenty-seven of C1 2 are taken and the C1 2 loses two, therefore two of CrI 3 are taken. This rule holds good in organic chemistry when all the products of the reactions are known, as the following examples will illustrate: CH 4 C- 4 H' 4 . 4+4 = CH 3 C1 C- 3 + 'H',Cl- / . 3+1 + 3-1 = CH 2 Cl a C- 2 + 2 H' 2 Cl-' a . 2+2+2-2 = CHOI, C- 1 + 8 H'Cl-' 3 . -1+3+1-3 = CCL C 4 Cl-'4. 44 = HC 2 H 8 O a H'(C a ) + 3 - 3 H' 3 O- a a . 1 + 3-3 + 3 4 = C a H 6 O (C 2 ) 1 - 5 H / 6 O- 2 . 1 5 + 6 2 = C S H 8 O 8 (Cs) - 5 + 3 H' 8 O - 2 3 5 + 3 + 8 6 = C 6 H 12 8 (C 6 )- 7 + 7 H' 12 0-V -7+7+12-12 = (1) CH 4 + 4C1 2 = CC1 4 + 4HC1 The carbon is oxidized by the chlorine from negative four to positive four, a polarity change of eight units, hence take eight molecules of chlorine; each molecule of chlorine loses two bonds, take two molecules of methane. Two is to eight as one is to four. (2) 3C 2 H 6 + 2K 2 Cr 2 T + 8H 2 S0 4 = 3HC 2 H 3 O 2 + 2K 2 S0 4 + 2Cr 2 (S0 4 ), + 11H 2 The carbon of the alcohol while possessing a valence of eight, has an oxida- tion valence of but four (minus four bonds) ; in the acetic acid the two atoms of carbon have zero bonds, that is, the combinations with negative affinity exactly equal the combinations with positive affinity; therefore take four molecules of the potassium dichromate. The two atoms of the chromium lose six bonds, take six molecules of the alcohol. Six is to four as three to two. Eight molecules of sulphuric acid are necessary to combine with the potassium and the chromium. (3) 3C,H 8 3 + 14HN0 3 = 9C0 2 + 14NO + 19H.O The three atoms of the carbon in the glycerine have minus two bonds (the negative affinity is two more than the positive affinity), and in the C0 2 a like amount has twelve bonds, a gain of fourteen. The nitrogen loses three bonds. (4) C H 12 6 + 12H 2 S0 4 = GC0 2 + 12S0 2 + 18H.O The carbon in the dextrose has zero bonds (equal positive and negative affinity combinations) and gains twenty-four bonds, while the sulphur loses two bonds. The lower ratio is one to twelve. For convenience of reference the non-metallic elements will be de- scribed in the order of their atomic weights; and the acids in the order of the degree of oxidation of the characteristic element, e. g., N before S , HC1 before HC10 , HC10 3 before HC10 4 , etc.. HYDROGEN. 219 219. Hydrogen. H= 1.008 . Valence one. 1. Properties. An odorless, tasteless gas. It is the lightest body known: One litre at 0, 760 mm. atmospheric pressure, weighs 0.08952289 gram (one crith); specific gravity, 0.06949 (Crafts, C. r., 1888, 106, 1662). It is used for filling balloons; also illuminating gas, containing about 50 per cent of hydrogen, is frequently used because it is much cheaper. It is a non-poisonous gas, but causes death by exclusion of air. It has been liquified to a colorless trans- parent liquid by cooling to 220 under great pressure and then allowing to expand rapidly (Olszewski, C. r., 1884, 99, 133; 1885, 101, 2:J8; Wroblewski, C. r., 1885, 100, 979). Critical temperature, 234.5; critical pressure, 20 atmospheres; tailing point, 243.5 (Olszewski, Phil. Mag., 1895, (5), 40, 202). It diffuses through walls of paper, porcelain, heated platinum, iron, and other metals more than any other gas (Cailletet, C. r., 1864, 58, 327 and 1057; 1865, 60, 344; 1868, 66, 847). It is absorbed by charcoal and by many metals, especially palladium; which, heated to 100 in an atmosphere of hydrogen and then cooled in that atmosphere, absorbs at- ordinary temperatures 982.14 volumes of hydrogen (Graham, J. C., 1869, 22, 419). This occluded hydrogen acts as a strong reducing agent, reducing FeCl 3 to FeCl 2 , HgCl 2 to Hg , etc. It is a better conductor of sound than air (Bender, B., 1873, 6, 665). It conducts heat seven times better than air or 480 times poorer than iron (Stefan, C. C., 1875, 529). It refracts light more powerfully than any other gas and about six times more than air. It burns with a non-luminous flame and with generation of much heat (more than an equal weight of any other substance or mixture of substances). Hydrogen forms two oxides: water, H 2 , and hydrogen peroxide, H 2 O 2 (244). 2. Occurrence. In volcanic gases (Bunsen, Pogg., 1851, 83, 197). In pockets of certain Stassfurt salt crystals (Precht, B., 1886, 19, 2326). As a product of the decay of organic material, both animal and vegetable. In combination as water and in innumerable minerals (H 2 O and OH) and in organic- compounds. 3. Formation. (a) By. the reaction of alkali metals with water, (b] By the action of superheated steam upon heated metals or glowing coals (226, 4a). (c) By dissolving aluminum or certain other metals in the fixed alkalis, (d) By the action of many metals with dilute acids (seldom HN0 3 ). By heating potassium formate or oxalate with KOH : K 2 C 2 4 + 2KOH = 2K 2 C0 3 + H 2 (Pictet, A. Ch., 1878, (5), 13, 216). 4. Preparation. (a) By the action of dilute sulphuric acid (one to eight) on commercial or platinized zinc * (135, 5a). The solution must be kept cold or traces of S0 2 and H 2 S will be evolved. (&) By the elec- trolysis of acidulated water. 5. Solubilities. Water at ordinary temperature dissolves nearly two per cent (volume) of hydrogen. Charcoal dissolves or absorbs fully ten times its volume of the gas (1). 6. Reactions. Hydrogen gas is a very indifferent body at ordinary tem- perature, combining with no other element except as it :s occluded or ab- sorbed by palladium, platinum, iron, nickel, etc.; and in the sunlight combines with chlorine and bromine. " Nascent hydrogen " (hydrogen at the moment of its generation), however, is a powerful reducing agent, and under proper * For the rapid generation of hydrogen the zinc should be granulated by pouring the molten metal into cold water. Chemically pure zinc is very slowly attacked by dilute sulphuric acid ; but the commercial zinc frequently contains sufficient impurities to insure a rapid generation of hydrogen when treated with the dilute acid. By tho addition to the granulated zinc, in a tub of water, of a few cubic centimetres of a dilute solution of platinum chloride or copper sul- phate, the zinc is made readily soluble in dilute sulphuric acid and a uniform and rapid gen- eration of hydrogen can be obtained. 219, 9. HYDROGEN. 251 conditions combines with , S , Se , Te , Cl , Br , I , N , P , As , Sb and Si with comparative readiness. The reduction of salts by nascent hydrogen in acid or alkaline solution will not be discussed here. See under the respective elements. It should be noted, however, that " nascent hydrogen " generated by different methods does not possess the same reducing properties. Sodium amalgam with acids does not give hydrogen capable of reducing silver halides; the reduction is rapid when zinc and acids are used. Neither electrolytic hydrogen nor that from sodium amalgam and acids reduces chlorates; while zinc and acids reduce rapidly to chlorides. Hydrogen generated by KOH and Al does not reduce AsV; that formed by zinc and acids gives AsH 3 . Sbv with sodium amalgam and acids gives Sb; with zinc and acids, SbH 3 (Cha- brier, C. r., 1872, 75, 484; Tommasi, BL, 1882, (2), 38, 148). Hydrogen occluded in metals as Pd , Pt , etc., is even more active than " nascent hydrogen "; often causing combination with explosive violence (Berthelot, A. Ch., 1883, (5), 30, 719; Berliner, W. A., 1888, 35, 781). Hydrogen absorbed by palladium precipitates Ag" , Au , Pt , Pd , Cu and Hg from their solutions; permanganates acidified are reduced to Mn"; Fe'" to Fe"; Crvi to Cr'"; KC10 3 to KC10; CH 3 CO 2 H to CH 3 CHO and C 2 H 5 OH; and O 6 H 3 1TC> 2 to C 6 H 5 NH 2 . The reactions are quantitative. Salts of Pb , Bi , Cd , As , Sb , W, Mo , Zn , Co , Ni , Al , Ce , U , Bb , Cs , K , Na , Ba , Sr and Ca are not reduced (Schwarzenbach and Kritschewsky, Z., 1886, 25, 374). In the presence of platinum black hydrogen reduces very much as described above; also K 3 Fe(CN) becomes K 4 Fe(CN) fi ; dilute HNO 3 becomes NH 4 NO 2 , concentrated HNO 3 be- comes HN0 2 ; Cl , Br and I combine with the hydrogen in the dark; KC10 3 and KC10 are reduced to chlorides, KC10 4 is not reduced; H 2 S0 4 , concen- trated, is reduced to H 2 SO 3 (Cooke, C. N., 1888, 58, 103). Free hydrogen very slowly acts upon a neutral solution of silver nitrate, precipitating traces of silver; and in concentrated solution with formation of Ag-NO, ; hindered by HN0 3 or KNO 3 . Solutions of Au , Pt and Cu are also acted upon (Rusself, J. C., 1874, 27, 3; Leeds, B., 18715, 9, 1456; Eeichardt, Arch. Pharm., 1883, 221, 585; Toleck and Thuemmel, B., 1883, 16, 2435; Senderens, BL, 1897, (3), 15, 991). KMnO 4 in acid, neutral, or alkaline solution slowly oxidizes hydrogen. It is not at all oxidized by nitrohj^drochloric acid, in diffused daylight, CrO 3 , at ordinary temperature, FeCl 3 , K 3 Fe(CN) 6 , HNO 3 , sp. gr. 1.42, or H 2 SO 4 , sp. gr. 1.84 (Wanklyn and Cooper, Phil. Mag., 1890, (5), 30, 431). In some cases, when hydrogen under ordinary conditions is without action, if subjected to great pressure a reducing action takes place; e.g., hydrogen at 100 atmospheres pressure precipitates Hg from HgCL (Loewen- thal, J. pr., 1860, 79, 480). 7. Ignition. Chlorine and bromine combine with hydrogen directly in the sunlight, but heat is required to effect its combination with iodine, fluorine, and oxygen. All oxides, hydroxides, nitrates, carbonates, oxalates, and organic salt ; of the following elements are reduced to the metallic or elemental state by ignition in hydrogen gas : Pb , Ag , Hg , Sn , Sb , As , Bi , Cu , Cd , Pd , Mo , Ru , Os , Rh , Ir , Te , Se , W , Fe , Cr , Co , Ni , Zn , Tl , Nb , In , V . Compounds of aluminum, manganese, and of the fifth and sixth group metals have not been reduced by hydrogen. 8. Detection. (a) Method of formation if known. (&) Its explosive union with oxygen when the mixture with air is ignited, (c) Absorption . by palladium sponge, (d) Explosive union with chlorine in the sunlight to form HC1 . (e) Separated from most other gases by its non-absorption by the chemical reagents used in gas analysis. 9. Estimation. By volume measurement, almost never by weight, except vrhen determined in its compounds by combustion to H.O . BORON BORIC ACID. 220. 220. Boron. B = 11.0 . Valence three (2). Boron does not occur free in nature. It is found chiefly as borax, Na 2 B 4 7 , and as boric acid, H 3 B0 3 , in volcanic districts. Two varieties of the element have been prepared, amorphous and crystalline. The former is changed to the latter by heating- to a white heat in presence of Al and C (Woehler and Claire- Deville, A., 1867, 141, 268). Elemental boron is prepared (a) by electrolysis; (&) by fusing B 2 3 with Al , Na or Mg-; (c) by igniting BC1 3 with hydrogen; (d) by fusing borax with red phosphorus. Specific gravity of the crystalline, 2.53 to 2.68 (Hampe, A., 1876, 183, 75); of the amorphous, 2.45. Amorphous boron is a greenish-brown, opaque powder, odorless, tasteless, insoluble in water, alcohol or ether. It is a non-conductor of electricity. Heated in air or oxygen it burns w T ith incandescence. In air it forms B.,O 3 and BN . It is oxidized by molten KOH or PbCr0 4 , with incandescence. It is dissolved by concentrated HN0 3 or H 2 S0 4 , forming boric acid. At a red heat it decom- poses steam. When heated it combines directly with S , Cl , Br , N and many metals. It forms BC1 3 with chlorine, not BC1 5 . Fused with P 2 r> it forms B 2 3 and P; with KOH , K,BO, and H; with K 2 C0 3 , K 3 B0 3 and C . Boron forms but one oxide, B,0 3 , boric anhydride. Three hydroxides are known: 2H 3 BO 3 =: B 2 O 3 .3H 2 O , orthoboric acid; 2HBO 2 = B 2 O 3 .H,0 , metaboric acid; and H 3 B 4 O T = 2B 2 O 8 ,H 2 , pyroboric acid. 221. Boric acid. H 3 B0 3 = 62.024 . H' 3 B'"0-" 3 , H B ~ ~ 1. Properties. Boron trioxide, B,0 3 , boric anhydride, is a brittle vitreous mass; sp. or. at 12, 1.8476 (Ditte, A. Ch., 1878, (5), 13, 67). Melting point, 577 (Carnelley, J. C., 1878, 33, 278). It is volatile at a very high heat (Ebelemen, A. Ch., 1848, (3), 22, 211). It has a slightly bitter taste, is hygroscopic, and shows a marked rise in temperature on solution in water (Ditte, C. r., 1877, 85, 1069). In some respects boron trioxide deports itself as a weak base. It forms a sulphide, B,S 3 , decomposed by water (Woehler and Deville, A. Ch., 1858, (3), 52, 90); a sulphate, B(HS0 4 ) 3 (D'Arcey, J. C., 1889, 55, 155); and a phosphate, BPO 4 (Meyer, B., 1889, 22, 2919). It combines with water in three proportions, forming the ortho, meta and pyroboric acids. Orthoboric acid is a weak acid, its solutions reddening litmus; at 12 it has a specific gravity of 1.5172 (Ditte, I.e.); melts at 184 to 186 (Carnelley, I. c.). Soluble in 25 parts water at 20, and in 3.4 parts at 102 (Ditte, I.e.). It is volatile in steam and in alcohol vapor. The evaporation of the water of combination of the acid carries with it from ten to fifteen per cent of the acid. 2. Occurrence. Widely distributed, but usually in very small quantities. In the rock salt deposits at Stassfurt, Germany, as boracite, Mg 7 B ]fl O so Cl 2 (62.5 per cent B 2 O S ). In the volcanic regions of Tuscany and the Liparic Islands as steam saturated with boric acid. 3. Formation. The anhydride is formed by burning the metal in air or oxygen, or by heating the acids. Orthoboric acid, H 3 B0 3 , is formed by dissolving the oxide in water; the meta acid, HB0 2 , H B = , by heating the ortho acid a little above 100 (Bloxam, /. C., 1860, 12, 177); the pyroboric acid, tetraboric acid, H 2 B 4 7 , by heating the ortho or meta acid for some time at 160 in a current of dry air (Merz, J. pr., 1866, 99, 179). 4. Preparation. (a) By evaporation of the water from the lagoons of Tuscany, which are saturated with boric acid, and recrystallization 221, 7. BORIC ACID. 253 from water, (b) The boronatrocalcite, Ca,B 6 O n .Na 2 B 4 7 + 18H 2 (45.6 per cent B 2 3 ), found in Nevada, is evaporated in lead pans with H 2 S0 4 to a stiff paste; and then treated with superheated steam in iron cylinders heated to redness. The acid passes over with the steam and is collected in lead lined chambers (Gutzkow, Z., 1874,, 13, 457). (c) Commercial borax, Na 2 B 4 7 .10H 2 , is dissolved in hot water, twelve parts, and acidi- fied with hydrochloric acid. Upon cooling, the boric acid, H 3 B0 3 , is ob- tained in small scales, which are purified by recrystallizatioii from hot water. 5. Solubilities. More soluble in hydrochloric acid solution or in alcohol than in water (1). The alcoholic solution burns with a beautiful green flame. Quite soluble in glycerine and in most alcohols and hydrocarbons, only sparingly in ether. The borates are insoluble in alcohol; those of the alkalis are soluble in water to an alkaline solution. Borates of the other metals are insoluble in water (no borate is entirely insoluble in water) ; but are usually rendered soluble by the addition of boric acid. 6. Reactions. Silver nitrate forms, in solutions of acid borates, a white precipitate of silver meiaborate, AgB0 2 , but normal borates form in part silver oxide, brown. Lead acetate gives a white precipitate of lead meta- lorate, Pb(B0 2 ) 2 ; calcium chloride, in solutions not very dilute, a white precipitate of calcium metaborate; and barium chloride, in solutions not dilute, a white precipitate of barium metaborate, Ba(B0 2 ) 2 . With aluminum salts, the precipitate is aluminum hydroxide. Borates are transposed with formation of boric acid, by all ordinary acids in some conditions even by carbonic acid. The liberated boric acid is dissolved by alcohol, and if the alcohol solu- tion be set on fire, it burns with a green flame. A solution of a borate, acidulated with hydrochloric acid to a barely perceptible acid reaction, imparts to a slip of turmeric paper half wet with it, a dark-red color, which on drying intensifies to a characteristic red color which turns dark green when moistened with a drop of alkali. 7. Ignition. Boric acid is displaced from its salts by nearly all acids including C0 2 ; but being non-volatile except at a very high heat, it dis- places most other acids upon ignition. By heating a mixture of borax, acid 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 How-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. 254 CARBON. 221, 8. 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, the 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 i n the green to Sr d in the blue. Borax, Na 2 B 4 7 , when ignited (as on a loop of platinum wire to form the borax bead) with many metallic compounds, forms a colored glass, used in the detection of certain metals (132, 7). The fused borax forms a solid brittle mass, borax glass, used in assaying and in soldering because of its power of combination with metallic oxides. 8. Detection. By conversion into the acid, if present as a salt; solution in alcohol or glycerine and burning with the formation of the green flame (very delicate, but copper salts should be removed by H 2 S and barium salts- should be removed or converted into the sulphate). Also by the red color imparted to a strip of turmeric paper. 9. Estimation. Boron compounds cannot be completely precipitated from solution by any known reagents, hence most of the methods of quantitative determination are indirect. By adding a known quantity of Na 2 C0 3 , fusing and weighing; then after determining the CO 2 subtracting its weight and that of the Na,O present (calculated from Na,C0 3 first added). The differ- ence is the weight of B 2 O 3 present. See also Will (Arch. Plianu., 1887, 225, 1101). In the presence of glycerine or mannitol, boric acid may be accurately titrated with sodium hydroxide, using phenolphthalein as an indicator: B 2 O 3 + 2NaOH = : NaBO 2 + H 2 O . Sodium carbonate must be absent or we get: 2B 2 O 3 + Na 2 CO.-i = Na^B.Oy + COo (Honig and Spitz, Z. angew., 1896, 549; Joergensen, Z. angew., 1897, 5). 222. Carbon. C = 12.0 . Usual valence four (15). 1. Properties. Carbon exists in three allotropic forms: two crystalline, diamond and graphite, and amorphous as charcoal, coke, etc. Specific gravity, diamond at 4, 3.51835 (Baumhauer, J., 1873, 237); graphite, Ceylon, 2.2o'to 2.26 (Brodie, A., 1860, 114, 6); wood charcoal, 1.57; gas coke, 1.88. Very small specimens only, of diamonds have been artificially prepared, by saturating iron with carbon at 3000. At this temperature graphite is formed and upon cool- ing under pressure the crystalline diamond form is obtained. This cooling under pressure is obtained by pouring the carbon saturated iron into a soft iron bomb, which is cooled by water (Moisson, C. r., 1893, 116, 218). Diamond is the hardest substance known. It is very strongly refractive towards light (Becquerel, A. C7., 1S77, (5), 12, 5). Fluorescence and phosphorescence of diamonds, see Kunz (C. (7., 1891, ii, 562). Ignition in an atmosphere of hydro- gen does not effect a change; in air or oxygen it burns to CO., . Graphite is a hard, gray, metal-like, opaque solid, a good conductor of electricity and a fairly good conductor of heat. It burns with difficulty. It 222, 0. CARBON. 255 is used in lead pencils, in black lead (plumbago) crucibles, as a lubricant for heavy machinery, in battery plates, lor The arc lig-ht carbon pencils, etc. Amorphous carbon is black, lighter than diamond or graphite. It is in use as coal, coke, charcoal, animal charcoal, etc.; all impure forms. Lamp-black is also amorphous carbon made from burning resin, fat, wax, coal gas, etc., with limited supply of air. It is used as a pigment in paints, in stove-black- ing, shoe-blacking, printers' ink, etc. Charcoal, preferably animal charcoal, is used for decoloring organic solutions. Charcoal absorbs many gases, hence is valuable as a disinfectant. Carbon forms two oxides: carbon monoxide, CO , and carbon dioxide, C0 2 . 2. Occurrence. Diamonds seem first to have been found in India, especially in the Golconda pits, where, as early as 1622, 30,000 laborers are said to have been employed (Walker, /., 1884, 774). Also found in other parts of Asia, in South Africa, in Brazil, etc. (Winklehner, C. C., 1888, 192; Damour, J., 1883, 774; Gorceix, J., 1881, 345; Smit, J., isso, 1400). Graphite is found in Ceylon (Wal- ther, C. C., 1890, ii, 20); in California (C. N., 1868, 17, 209); in Canada (Dawson, Am. 8., 1870, (2), 50, 130); in New York State; in New Zealand (Mac Ivor, (\ N., 1887, 55, 125); in Russia, Germany, Greenland, etc. Pure amorphous r^on occurs in nature a-; a chief product in the decomposition of organic material, air being excluded. Anthracite coal is relatively pure amorphous carbon. 3. Formation. Graphite remains as a residue when pig iron is dis- solved in acids. It forms by reducing CO with Fe 3 4 at 400. Amor- phous carbon is formed by passing CC1 4 over Na in a tube heated to red- ness (Porcher, C. A 7 ., 1881, 44, 203). 4. Preparation. Pure graphite is prepared by heating the commercial graphite on a water bath with KC10 3 and H 2 S0 4 and repeatedly washing. If it contains Si0 2 it should also be treated with NaF and H 2 S0 4 . Amor- phous carbon is prepared by heating wood, coal, or almost any organic- matter to a very high temperature in absence of air, but when so prepared it is never pure. Amorphous carbon is prepared approximately pure by heating pure cane sugar in a closed platinum crucible; then boiling in succession with HC1 , KOH , and H 2 ; then igniting to redness in an atmosphere of chlorine, cooling in the same atmosphere. A very pure form of graphite known as Acheson graphite is made in large quantities at Niagara Falls by passing a strong electric current through coke compressed into the desired form. Silicon, aluminum and other impurities are dis- tilled out at the high temperature employed. 5. Solubilities. Insoluble in water or acids Soluble in many molten metals with partial combination to form carbides. When the metal is dissolved in acids the combined carbon passes off as hydrocarbons, the excess remaining as graphite. 6. Reactions. Xot attacked by acids or alkalis. It slowly oxidizes to C0 2 when heated with concentrated H 2 S0 4 and K 2 Cr 2 T . Upon gently warming graphite with KC10 3 and HN0 3 , graphitic acid, C^H^, is said to be formed (Stingl, B., 1873, 6, 391). The important reactions of carbon require the aid of high heat and are described in the next para- graph. ACETIC ACID. 22$, ?. 7. Ignition. Unchanged by ignition in absence of air. When strongly ignited in air or oxygen it slowly burns to C0 2 . If the carbon and oxygen has been previously very thoroughly dried the action is very slow, especially with graphite. By fusion with KN0 3 or KC10 3 carbon is oxid- ized to C0 2 . With vapors of sulphur, carbon disulphide is formed ; i. e., by passing sulphur vapors over hot coals in an electrically heated furnace. In an u'mosphere of hydrogen with the electric spark, acetylene, C 2 H 2 , is formed. i-y igniting in an atmosphere of carbon dioxide. C0 2 , the whole of the carhon becomes carbon monoxide: C + C0 2 = 2CO. By simple ignition with carbon, all oxides of the elements in the follow- ing list are reduced to the elemental state (a); and if sodium carbonate is added, all of the salts of the same are likewise reduced (b). Cu , Bi , Cd , Pb, Ag, Hg, As, Sb, Sn, Pd, Mo, Ru , Os, Rh, Ir, Te, Se, w', K, Na , Rb , Cr , Fe , Mn , Co , Ni , Zn , Ti , Tl . (a) Pb 3 4 + 2C = 3Pb + 2CO 2 (6) 2PbCL + 2Na,CO 3 + C = 2Pb + 4NaCl + 3C0 2 (c) CuO + C (excess) = Cu + CO (d) C -f- 2CuO (excess) = 2Cu + CO. With excess of carbon CO is formed (c). With excess of the oxide C0 2 is formed (d). In the reduction of iron ore, the process is conducted so as to give some CO and some C0 2 . To obtain some metals in the free state (such as K and Na), special methods are adopted to exclude the air, and to produce the high temperature needed. All compounds of sulphur when ignited with carbon are reduced to a sulphide: BaS0 4 + 2C = BaS + 2CO, . 8. Detection. By its appearance; failure to react with general reagents; and by its combustion to C0 2 with oxygen (air), or with K 2 Cr 2 7 and con- centrated H 2 S0 4 (Fritsche, A., 1896, 294, 79), then by identification with Ca(OH) 2 (228, 6). 9. Estimation. By combustion to CO, and weighing 1 after absorption in KOH solution. See works on ultimate organic analysis. 223. Acetic acid. HC 2 H 3 2 = 60.032 . H I II H' 4 (C 2 )+"'-"'0-" 2 , H C C H = CH 3 C0 2 H. 1. Properties. Pure acetic acid is a colorless, crystalline, hygroscopic solid, melting at 16.5 and boiling at 118. Its specific gravity at is 1.080. It has a sharp, sour taste, an irritating burning ell'ect on the skin, and a very peiie- 223, 6. ACETIC ACID. 257 trating odor. It burns when heated nearly to the boiling point. Vinegar contains four to five per cent of acetic acid. The U. S. P. reagent contains 36 per cent of acetic acid, and has a specific gravity of 1.0481 at 15. It vaporizes from its con- centrated solutions at ordinary temperature, having the characteristic odor of vinegar. It is a monobasic acid, as the three remaining hydrogen atoms (linked to carbon) cannot be replaced by metals. 2. Occurrence. It occurs in nature in combination with alcohols in the essen- tial oils of many plants. 3. Formation. (a) During the decay of many organic compounds. (&) By gently heating sodium methylate, NaOCH 3 , in a current of carbon monoxide: NaOCH 3 + CO = CH 3 C0 2 Na (NaC 2 H 3 2 ). (c) By boiling methyl cyanide with acids or alkalis : CH 3 CN + HC1 + 2H 2 = HC 2 H 3 2 + NH 4 C1. (d) By the oxidation of alcohol: 3C 2 H G + 2K 2 Cr 2 7 + 8H 2 S0 4 == 2K 2 S0 4 + 2Cr 2 (S0 4 ) 3 + 3HC 2 H 3 0, + 11H 2 . 4. Preparation. (a) By the dry distillation of wood. (&) By the fer- nlentation of cider, beer, wine, molasses, etc. (c) Pure acetic acid is prepared by distilling anhydrous sodium acetate with concentrated sul- phuric acid. The distillate solidifies upon cooling and is termed glacial acetic acid. 5. Solubilities. Miscible in all proportions in water and alcohol. The salts of acetic acid, acetates, are all soluble in water, silver and mercurous acetates sparingly soluble. One part of silver acetate requires 115 parts of water at 10 for its solution; one part of mercurous acetate requires loo parts of water. Certain basic acetates, as Fe'" , Al , etc., are insoluble in water. Very many of the acetates are soluble in alcohol. Ammonium acetate dissolves several insoluble sulphates such as calcium and lead sulphates. 6. Reactions. The stronger mineral acids transpose the acetates, forming acetic acid. Anhydrous acetates with concentrated sulphuric acid give pure acetic acid (4), but if the sulphuric acid be in excess and heat be applied the mixture blackens with separation of carbon; and, at higher temperatures, C0 2 and S0 2 are evolved. Solution of ferric chloride forms, with solutions of acetates, a red solu- tion containing ferric acetate, Fe(C 2 H 3 2 ) 3 , which on boiling precipitates brownish-red, basic ferric acetate. The red solution is not decolored by solution of mercuric chloride (distinction from thiocyanate); but is de- colored by strong acidulation with sulphuric acid or hydrochloric acid (dis- tinction from thiocyanate and from meconate). The ferric acetate is pre- cipitated by alkali hydroxides. If acetic acid or an acetate be warmed with sulphuric acid and a little alcohol, the characteristic pungent and fragrant odor of ethyl acetate or acetic ether is obtained: HC 2 H,0, + C 2 H 5 OH = H 2 + aH 5 C 2 H 3 O 2 Acetic acid does not act as a Reducing Agent as readily as do most of 258 CITRIC ACID. 223, 7. the organic carbon compounds. Acetic acid is stable toward mild oxidiz- ing agents and is only slowly attacked by strong oxidizing agents such as chromic acid and potassium permanganate; reduces auric chloride only in alkaline solution, and does not reduce alkaline copper solution. It takes chlorine into combination slowly in ordinary light, quickly in sun- light, forming chloracetic acids. Acetic acid is a weak acid, only 0.4% being ionized in normal solu- tion and 1.4% in tenth normal solution. Glacial acetic acid dissolves sulphur which crystallizes out in needles. 7. Ignition. By ignition alone, acetates blacken, with evolution of vapor of acetone, C 3 H 6 , inflammable and of an agreeable odor. By pro- longed ignition of alkali acetates in the air, carbonates are obtained free from charcoal. By ignition with alkali hydroxides in dry mixtures, methane, marsh-gas, CH 4 , is evolved. By ignition with alkalis and arsenous anhydride, the poisonous and offensive vapor of cacodyl oxide is obtained. This test should be made under a hood with great caution and with small quantities. It is a very delicate test for acetates: 4KC 2 H 3 2 + As 2 3 = As,(CH 3 ) 4 + 2K 2 C0 3 + 2C0 2 . Butyric, propionic and valerianic acids give the same reaction. 8. Detection. (a) By its odor. (5) By the formation of the fragrant ethyl acetate upon warming with sulphuric acid and alcohol. Too much alcohol should not be used in testing for acetic acid as otherwise ethyl ether is formed, (c) By the formation of the red solution with ferric chloride (126, 6b and 152). ('d) By ignition of the dry acetate alone to acetone, CH 3 COCH 3 ; with NaOH to methane, CH 4 ; or with As,0 3 to cacodyl oxide, (e) As a delicate test for formates or acetates it is directed to warm a solution of CuCL in NaCl and add a small amount of the mate- rial under examination. Formates give a blackish-gray deposit; acetates give a bright green precipitate not changed by boiling. Both precipitates are soluble in acetic acid (Field, J. C. f 1873, 26, 575). 9. Estimation. Other volatile acids are separated by precipitation; sulphuric acid is then added and the acetic acid is distilled into water and estimated by titration with standard alkali. 224. Citric acid. H 3 C 6 H 5 7 = 192.064 . H 2 C C0 2 H H' 3 (C 6 )+ 10 - 4 H' 5 0-" 7 , H C CO ,H I H L> C C0 2 H Found in small quantities in the juices of many fruits. The chief commercial source is lemon-juice. It is a colorless, crystallizable, non-volatile solid; it ig 225, 1. TARTARIC ACID. 259 soluble in 0.75 parts of cold water, in equal parts of 90 per cent alcohol, in 1.3 parts of absolute alcohol and in 50 parts of ether. It crystallizes with one mole- cule of water in rhombic prisms which melt at 100. The citrates of the metals of the alkalis are freely soluble in water; those of iron and copper are moderately soluble; those of the alkaline earth metals insoluble. There are many soluble double citrates formed by action of alkali citrates upon precipitated citrates, or of alkali hydroxides upon metallic salts in presence of citric acid. In distinction from tartrates, the solubility of the potassium salts, non-precipitation of calcium salt in cold solution; and weaker reducing action, are to be noted. Solution of calcium hydroxide in excess (as by dropping the solution tested into the reagent) gives no precipitate with citric acid or citrates in the cold (distinction from tartaric acid), but on heating, the white calcium citrate, Ca3(C 6 H 5 O7) 2 , is precipitated (not soluble in cold potassium hydroxide solu- tion). By filtering before boiling, the tartrate and citrate may be approxi- mately separated. The chlorides of calcium and barium give no precipitate in neutral solutions (difference from tartaric acid), but if sodium hydroxide is added, calcium citrate is precipitated, insoluble in sodium hydroxide, but readily soluble in ammonium chloride. On boiling the solution in ammonium chloride, crystallized calcium tartrate is precipitated, which is now insoluble in ammo- nium chloride. Calcium citrate is soluble in acetic acid (distinction from oxalates). Solution of lead acetate precipitates white lead citrate, Pb 3 (C 6 H5O7) 2 , soluble in ammonia. Silver nitrate gives a white precipitate of silver citrate, AgsCeH^Or , which does not blacken on boiling (distinction from tartrate). For action of citric acid or citrates in hindering many of the usual analytical reactions, see Spiller, J. C., 1858, 10, 110. One part of citric acid dissolved in two parts of water, and treated with a solution of one part of potassium acetate in two parts of water, should remain clear after addition of an equal volume of strong alcohol (absence of oxalic acid and of tartaric acid and its isomers). Citric acid does not act very readily as a reducing agent; does not reduce alkaline copper solution, or silver solution; reduces permanganate very slowly. Concentrated nitric acid produces from it, acetic and oxalic acids; and diges- tion with manganese dioxide decomposes it 'with formation of acetone, acrylic and acetic acids. Concentrated sulphuric acid carbonizes citric acid with liber- ation of sulphur dioxide. Citrates carbonize on ignition, with various empy- reuniatic products, and with final formation of carbonates. On heating citric acid, it loses its water of crystallization, then fuses and decomposes with evolution of pungent fumes leaving a carbonaceous residue. By fused potassium hydroxide, short of ignition, they are decomposed with production of oxalate and acetate. 225. Tartaric acid. H 2 C 4 H 4 6 = 150.048 . H I II H C C H CH(OH)C not). Most of the tartrates are also dis- solved (and, if already dissolved, are not precipitated) by the alkali hy- droxides, owing to the formation of soluble double tartrates. The freshly precipitated oxides, hydroxides, and carbonates of the fol- lowing metals are soluble in a solution of potassium-sodium tartrate, Rochelle salt: Sb , Sn lv , Bi , Cu , Fe , Al , Cr , Co , Ni , Mn , and Zn ; Ba , Sr, Ca, and Mg to quite an extent. CdC0 3 is not dissolved (Warren, C. N., 1888, 57, 223). 6. Reactions. Solution of calcium hydroxide, added to alkaline reac- tion, precipitates from cold solution of tartaric acid, or of soluble tartrates, calcium tartrate, white, CaC 4 H 4 6 . Solution of calcium chloride with neutral tartrates gives the same precipitate. Solution of calcium sulphate forms a precipitate but slowly, or not at all (distinction from racemic acid). The precipitate of calcium tartrate is soluble in cold solution of potassium hydroxide, precipitated gelatinous on boiling, and again made soluble on cooling (distinctions from citrate), and dissolves in acetic acid (distinction from oxalate). Tartaric acid prevents the precipitation by fixed alkalis of solutions of the & alts of the following metals : Al , Bi , Co , Ni , Cr , Cu , Fe , Pb , Pt , and Zn (Grothe, /. pr., 1864, 92, 175). 225, 9. TARTARIC ACID. 261 Silver nitrate precipitates, from solutions of normal tartrates, silver tartrate, Ag 2 C 4 H 4 , white, becoming black when boiled. If the precipi- tate is filtered, washed, dissolved from the filter by dilute ammonium hydroxide into a clean test-tub**, left for a quarter of an hour on the water-bath, the silver is reduced as a mirror coating on the glass (59, lOb), distinction from citric acid. Free tartaric acid does not reduce silver salts. Permanganate is reduced quickly by alkaline solution of tartrates (distinction from citrates), precipitating manganese dioxide, brown. Free tartaric acid acts but slowly on the permanganate. Alkaline copper tar- trate, Fehling's solution (77, 6b), resists reduction in boiling solution. Chromates are reduced by tartaric acid, the solution turning green. The oxidized products, both with permanganate and chromate, are formic acid, carbonic anhydride, and water. 7. Ignition. On ignition, tartaric acid or tartrates evolve the odor of "burnt sugar, separating carbon, and becoming finally converted to carbon- ates. Strong sulphuric acid also blackens tartrates, on warming. Melted potassium hydroxide, below ignition, produces acetate and oxalate. The fixed alkali tartrates ignited in absence of air give an alkali carbonate and finely divided carbon. The mixture serves as an admirable flux for the reduction tests for arsenic (69, 7). 8. Detection. (a) By the odor of burnt sugar when ignited. (&) By the deportment of the calcium salt with cold and hot KOH (6). (c) By the formation of the silver mirror (59, 1Gb). (d) By its action as an alkali tartrate in preventing precipitation of the solutions of the heavy metals by the fixed alkalis. To test citric acid for the presence of tartaric acid, add about one cc. of ammonium molybdate solution to about one gram of the citric acid; then two or three drops of sulphuric acid and warm on the water-bath. The presence of 0.1 per cent or more of tartaric acid gives a blue color to the solution (Crismer, El., 1891, (3), 6, 23). Add to tartaric acid or a tartrate a little ferrous sulphate, then one to two drops of hydrogen peroxide, then alkali the presence of tartaric acid will be in- dicated by a deep violet color (Fenton, Chem. News, 43, 110). Heat tartaric acid with a little resorcin and concentrated sulphuric acid in a porcelain dish to 125-130 C. First, red streaks will appear, then the whole liquid turns red. Sensitiveness 0.01 mg. tartaric acid (Mohler, Bull. Soc. Chim. France [3], 4, 1890, 728). This reaction depends upon the formation of an aldehyde and its subsequent condensation with resorcin. The test can be used to distinguish between citric and tartaric acids. If c*-naphthol is used in place of resorcin, a blue liquid turning green is the result (Pifieaua, Chem. Neivs, 9i, 179). 9. Estimation. See Philipps (Z., 1890, 29, 577); Haas (C. C., 1888, 1045); Heidenhain (Z., 1888, 27, 681). tARBON MONOXIDE. 226, i. 226. Carbon monoxide. CO = 28.0 . C"0-", C = . 1. Properties. Carbon monoxide, carbonic oxide, formic anhydride, CO , is a colorless, tasteless gas. Specific gravity, 0.9678. By maintaining a pressure of 200 to 300 atmospheres at 136 and then reducing the pressure to 50 atmos- pheres the gas becomes a colorless transparent liquid (Wroblewski and Ols- zewski, A. %., 1884 (6), 1, 128). It is, when inhaled, a virulent poison, abstract- ing oxygen from the blood and combining with the haemoglobin. It burns in the air with a pale blue flame to C0 2 , but does not support combustion. Mixed with air in suitable proportions, it explodes upon ignition. It unites with chlorine in the sunlight to form phosgene, COCL . 2. Occurrence. In combination as formic acid in ants and in nettles. 3. Formation. (a) By the incomplete combustion of coal, charcoal or organic material. (b) From the reduction of metallic oxides in the blast furnace with excess of charcoal: Fe 2 3 + 3C = 2Fe -j- SCO. (c) By heating sodium sulphate with excess of charcoal (LeBlanc's soda process) : Na 2 S0 4 + 4C = Na 2 S + 4CO . See also Grimm and Eamdohr (.4., 1856, 98, 127). 4. Preparation. (a) By passing steam over charcoal at a white heat (water gas): H 2 + C = CO + H 2 (Naumann and Pistor, B., 1885, 18, 164). (b) By passing C0 2 over red hot charcoal, (c) By heating K 4 Fe(CN) 6 with concentrated H 2 S0 4 : K 4 Fe(CN) 6 + 6H,S0 4 -f 6H 2 - 2K 2 S0 4 + 3(NH 4 ) 2 S0 4 + FeS0 4 + 6CO . With dilute acid HCN is formed. (d) By heating a formate with concentrated sulphuric acid: 2KCH0 2 -f- H 2 S0 4 = K 2 S0 4 + 2CO + 2H 2 . (e) By heating an oxalate with con- centrated sulphuric acid: K 2 C 2 4 + 2H S0 4 = K S0 4 + H,S0 4 .H,0 -f CO + C0 2 . 5. Solubilities. It is not absorbed by KOH or Ca(OH) 2 (distinction from C0 2 ). It is absorbed by charcoal, cuprous chloride, and by several metals, e. g., K , Ag , and An . 6. Reactions. It is an energetic reducing agent. Combines with moist fixed alkalis to form a formate (Froelich and Geuther, A., 1880, 202, 317). In the sunlight it combines directly with chlorine or bromine. It is oxidized to C0 2 by warming with K 2 Cr.>0 7 and concentrated H 2 S0 4 ; also by palladium sponge saturated with hydrogen, and in presence of oxygen and water (Remsen and Keiser, 5., 1884, 17, 83). A solution of PdCl 2 is reduced to Pd by CO. Reduces iodine pentoxide I 2 5 at 150, I 2 5 -f~ 5CO = I 2 + 5C0 2 . 7. Ignition. When heated to redness with Na or K, carbon and an alkali carbonate are formed. Upon ignition of metallic oxides in an atmosphere of CO a reduction of the metal takes place, so far as observed the same as when the corresponding metallic forms are ignited with char- coal (Rodwell, J. C., 1863, 16, 44). 227, 4, a. OXALIC ACID. 263 8. Detection. In distinction from C0 2 by its failure to be absorbed by KOH or Ca(OH) 2 . By its combustion to CO., and detection as such. By its combination with hot concentrated KOH to form a formate. By its action on I 2 5 . It is detected in the blood by the absorption spectrum (Vogel, B., 1878, 11, 235). 9. Estimation. The measured volume of the gas is brought in contact with a solution of cuprous chloride in hydrochloric acid which absorbs the CO (Thomas, C. N., 1878, 37, 6). The gas is passed through a U-tube containing I O 5 heated to 150. The liberated iodine is absorbed in RI solution and titrated with standard Na 2 S 2 O3 solution. Jour. Am. Chem. Soc., 1900, 22, 14. 227. Oxalic acid. H 2 C 2 4 = 90.016. II II C0 2 H H' 2 (C 2 )+ 6 0-" 4 ,H C C H or | C0 2 H 1. Properties. Absolute oxalic acid, H 2 C 2 4 , is a white, amorphous solid, which may be sublimed at 150 with only partial decomposition: H 2 C 2 O 4 = C0 2 + CO + H,O . Crystallized oxalic acid, H 2 C 2 4 ,2H 2 O , effloresces very slowly in warm, dry air, and melts in its water of crystallization at 98; at which temperature the liquid soon evaporates to the absolute acid. Oxalic anhydride is not formed. 2. Occurrence. Found in many plants in a free state or as an oxalate. In sorrel it is found as KHC 2 O 4 ; in rhubarb as CaC 2 4 . As ferrous oxalate in lignite deposits; as ammonium oxalate in guano. 3. Formation. (a) By decomposition of cyanogen with water, am- monium oxalate being one of the products, (fr) By the oxidation of glycol with nitric acid, (c) By heating potassium formate above 400 (Merz and Weith, B., 1882, 15, 1507). (d) By passing C0 2 over a mixture of sodium and sand at 360 (Drechsel, BL, 1868, 10, 121). 4. Preparation. (a) By action of nitric acid sp. gr. 1.38 upon sawdust, starch, or sugar. By the continued action of concentrated nitric acid, after the sugar is all oxidized to oxalic acid, the latter is farther oxidized to C0 2 . (b) By heating sawdust with KOH or NaOH . Hydrogen is evolved, the cellulose, C 6 H 10 5 , heing converted into oxalic acid. Under certain conditions, additional products are formed. It is also formed in the oxidation of a great many organic compounds. CuH^On + 12HN0 3 = 6HAO 4 4. 12NO + 11H 2 O 3H 2 C 2 4 + 2HN0 3 = 6C0 2 + 2NO + 4H 2 O C 6 H 10 5 + GKOH + H 2 = 3K 2 C 2 4 + 9H 2 Oxalates are formed: a. By treating the oxide, hydroxide, or car- bonate with oxalic acid. In this manner may be made the oxalates of 264 OXALIC ACID. 227, 4, b. Pb, Ag, Hg', Hg", Sn". Bi , Cu", Cd , Zn , Al , Co , Ni , Mn , Fe", Fe'", Cr'", Ba , Sr , Ca , Mg , Na , K, and some others. 1). By adding oxalic acid to some soluble salt of the metal. In this manner the above oxalates may be made,, except alkali, magnesium, chromic, ferric, aluminum and stannic oxalates, which are not precipitated. Antimonous salts are precipitated, but the precipitate is basic. c. Alkali oxalates will precipitate the same solutions as oxalic acid, but many of the precipitates are soluble in excess of the alkali oxalate, and, as a rule, the salt formed is a double one, e. g., AgNH 4 C,0 4 . Ba , Ca and Sr are well-defined exceptions to this rule their precipitates, formed by this method, being normal oxalates. d. Some of the metals when finally divided are attacked by oxalic acid, hydrogen being evolved. 5. Solubilities. Oxalic acid is very soluble in water and in alcohol. Alkali oxalates are freely soluble in water, as is also chromic oxalate. Nearly all other metallic oxalates are insoluble in water or only sparingly soluble (Luckow, J. C., 1887, 52, 529). The metallic oxalates, soluble and insoluble, are transposed by dilute sulphuric, hydrochloric, and nitric acids, with formation of oxalic acid: CaC 2 4 + 2HC1 = CaCl 2 + H 2 C 2 4 . That is : the precipitated oxalates of those metals, which form soluble chlorides, dissolve in dilute hydro- chloric acid; of those metals which form soluble sulphates, in dilute sul- phuric acid; and all precipitated oxalates dissolve in dilute nitric acid Acetic acid does not dissolve precipitated oxalates, or but slightly. Certain of the oxalates dissolve, to some extent, in oxalic acid (as acid oxalates). 6. Reactions. A. With metals and their compounds. Oxalic acid and soluble oxalates precipitate solutions of many of the metallic salts. With excess of the alkali oxalates soluble double oxalates of the heavy metals are frequently formed (4). An excess of alkali oxalate transposes par- tially the alkaline earth carbonates. On the other hand, the alkali car- bonates in excess partially transpose the alkaline earth oxalates (Smith, J. C., 1877, 32, 245). See also under 6b of the respective metals. Oxalic acid is a decided reducing agent, being converted to water and carbonic anhydride (a), and the metallic oxalates to carbonates and carbonic anhydride (&), by all strong oxidizing agents. (a) 2H 2 C 2 O 4 + O 2 = 2H 2 O + 4C0 2 (6) 2K 2 C 2 4 + 2 = 2K 2 C0 3 + 2CO 2 -f Pb0 2 with oxalic acid forms lead oxalate and C0 2 . Oxalic acid has no action upon Pb 3 4 , but reduces it quickly in presence of any acici capable of changing the Pb 3 4 to PbO ? , 227, B 6. OXALIC ACID. 265 2. Oxalic arid or ammonium oxalate boiled in the sunlight with HgCl 2 gives HgCl and CO, \Gmd\ns Hand-book, 9, 118]. 3. H a As0 4 becomes H AsO, , and CO, is evolved. To prove that As v becomes As"'., add excess of potassium hydroxide, and then potassium per- manganate. The latter will be quickly decolored. 4. Bi,0- 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,0 4 is formed, and C0 2 is given off. If some non-reducing acid be present, such as H,S0 4 , it unites with the manganese, and all of the oxalic acid is converted into C0 2 . 6. Co 2 3 and Co(OH) 3 form cobaltous oxalate, and C0 2 is evolved. 7. Ni 2 3 and Ni(OH) 3 become nickelous oxalate, and CO, is evolved. 6\ 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. HCN , 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,0 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 PO, , H 3 P0 3 , and H 3 P0 4 do not act upon oxalic acid. 4. Concentrated sulphuric acid, with a gentle heat, decomposes oxalic acidy b} r 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 -f- C0 2 -f~ ^0 . 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-book, 9, 116). This reaction takes place more readily in the presence of KOH , forming KC1 and K 2 CO S . HC10 forms C0 2 and Cl . If the oxalic be in excess HC1 is formed. The action 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 Cl and HC1 , a high degree of heat and excess of oxalic acid favoring the production of HC1 (Calvert and Davieo, J. C., 1850, 2, 193). 6. Bromine decomposes oxalic acid in. alkaline mixture, forming a 266 OXALIC ACID. 227, B 7. 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 (6). 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) Agr 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(OH) 2 or in a solution of BaCL 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, (c) Warm the solution with dilute H,S0 4 and XMn0 4 . If the permanganate is not decolorized, H 2 C 2 4 is absent; if decolorized test for C0 2 with Ca(OH) 2 . 2KMnO 4 + 5H 2 C 2 O 4 + 3H 2 SO 4 = K 2 SO 4 + 2MnSO 4 + 8H 2 O + 10CO 2 . 9. Estimation. (a) It is precipitated as CaC 2 O 4 ; after washing, the Ca is determined by 188, 9, from which the oxalic acid is calculated. (6) By the amount of KMnO. which it will reduce, (c) By measuring the amount of evolved when it is oxidized in any convenient manner, usually by MnO> . (d) By the amount of gold it reduces from AuCl 3 . 228, 4. CARBON DIOXIDE. 26? 228. Carbon dioxide. C0 2 = 44.0 . (Carbonic anhydride.) Carbonic acid (hypothetical). H 2 C0 3 = 62.016 . II C IV 0~" 2 and H' 2 C IV 0-" 3 ,0 = = and H C H. 1. Properties. The xjtccific gravity of the gas CO 2 is 1.52897 (Crafts, C. r., 1888, 106, 1G2); of the liquid at 34, 1.057 (Cailletet and Mathias, C. r., 188G, 102, 1202); of the solid (hammered), slightly under 1.2 (Landolt, #., 1884, 17, 309). Critical temperature, 30.92 (Andrews, Trans. Roy. Soc., 1869, 159, 583; 1876, 166, 21). It is a heavy colorless gas; which at low temperatures, +3, and high pressure, 79 atmospheres, may be condensed to a clear mobile liquid; and upon further cooling this becomes a snow-like mass. Liquid C0 2 is more compres- sible than other liquids (Natterer, r/., 1851, 59). It diffuses through porous plates more rapidly than oxygen (Graham, C. N., 1863, 8, 79). Non-combustible and a non-supporter of combustion, except that K , Na and Mg 1 burn in the gas forming an oxide of the metal and free carbon. It is used in chemical fire engines. Non-poisonous but causes suffocation (drowning) by exclusion of air. It is taken internally without injury in soda water, etc. Liquid CO., is insoluble in water which swims on the surface. It mixes with alcohol and ether. It dissolves iodine but does not dissolve phosphorus or sulphur; it is without action upon K or Na . A spirit thermometer immersed in the liquid registers 75 (Thilorier, J. /*., 1834, 3, 109). Solid CO, at 767.3 mm. barometric pressure melts at 77.94 (Regnault, A. Ch., 1849, (3), 26, 257). When the solid is mixed with ether it gives a temperature of 98.3. 2. Occurrence. In a free state in the air, about 0.04 per cent. Found in great abundance in the form of carbonates in the earth's crust; e. g., limestone, marble, magnesite, dolomite, etc. 3. Formation. (a) By burning wood, coal, etc., in the air. (b) By burning CO . (c) By the reduction of many metallic oxides upon ignition with charcoal, (d) During fermentation or decay of organic material. (e) By the reaction between acids and carbonates. Liquid C0 2 is made by compressing the gas with pumps at a reduced temperature. Solid C0 2 is made by allowing the liquid to escape freely into woolen bags and then compressing in wooden moulds (Landolt, I. c.). 4. Preparation. CaC0 3 (chalk or marble) in small lumps is treated with hydrochloric acid in a Kipp's gas generating apparatus. The gas is passed through a solution of NaHC0 3 to remove any HC1 that may be carried over, and then dried by passing through a tube filled with fused CaCl 2 . It is also prepared on a large scale for making the liquid C0 2 , and for use in sugar factories by the ignition of limestone : CaC0 3 = CaO + C0 2 . Preparation of Carbonates. Na 2 C0 3 is made by converting NaCl into Na-,S0 4 , by treating it with H 2 S0 4 ; then by long ignition with coal and calcium carbonate, impure sodium carbonate is formed (Leblanc's process). Na.SO, + 4C + CaCO; = CaS + 4CO + Na,C0 3 268 CARBOX DIOXIDE. 228, 5. It is separated by lixiviation with water, and farther purified. The other method, known as the ammonia,, or Solvay's process, consists in pass- ing NH 3 and C0 2 into a concentrated solution of NaCl (a). The NaHC0 3 is converted into Na 2 C0 3 by heat, and the evolved CO., used over again (b). The NH 4 C1 is warmed with MgO (c), and the NH, which is given off is used over again. The MgCl 2 is strongly heated (d) and the MgO is used over a'gain, and the evolved gas sold as hydrochloric acid. This continu- ous process has nearly superseded the Leblanc process. (a) NaCl + NH 3 + H 2 + CO 2 = NaHC0 3 + NH 4 C1 (6) 2NaHCO s + heat = Na,C0 3 + C0 2 + H 2 (c) 2NH 4 C1 + MgO = MgCL + 2NH 3 + H 2 (d) MgCL + H 2 + heat == MgO + 2HC1 The other carbonates are mostly made from the sodium salt (6). 5. Solubilities. C0 2 is soluble in water, forming the hypothetical H 2 C0 3 , which reacts acid towards litmus. At 15 one volume of water absorbs 1.002 volumes of the gas (Bunsen, A., 1855, 93, 1). It is rapidly absorbed by hydroxides of the alkalis and of the alkaline earths, forming normal or acid carbonates : KOH + C0 2 == KHC0 3 or 2KOH + C0 2 - K 2 C0 3 + H 2 . The carbonates of the alkalis are soluble in water (acid alkali carbonates are less soluble than the normal carbonates), other carbonates are insoluble in water or only sparingly soluble. The presence of some other salts, especially ammonium salts, increases the solubility oi carbonates, notably magnesium carbonate (189, 5c). Many of the car- bonates are soluble in water saturated with C0 2 ; forming acid carbonates of variable composition. Boiling removes the excess of C0 2 , causing pre- cipitation of the carbonate. 6. Reactions. Dry carbon dioxide does not unite with dry calcium oxide at ordinary temperature (Birnbaum and Maher, B., 1879, 12, 1547; Scheibler, B., 1886, 19, 1973). Also at no reaction takes place between dry C0 2 and dry Na 2 , but at 400 combination takes place with incan- descence (Beketoff, El., 1880, (2), 34, 327). Carbonates of the fixed alkalis precipitate solutions of all other metallic salts: with antimony the precipitate is an oxide; with tin, aluminum, chromium, and ferricum it is an hydroxide; with silver, mercurosum, cadmium, ferrosum, manganese, barium, strontium, and calcium it is a nor- mal carbonate; with other metals a basic carbonate, except that mercuric chloride forms an oxychlori.de. Carbonic acid 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,S0 4 , and even by H 2 S , completely from carbonates of the first four groups, incompletely from those of the fifth and sixth groups (Nandin and Montholon, C. r., 1876, 33, 58). Ammonium carbonate precipitates solutions of all the non-alkali metals, 228, 6. CARBON DIOXIDE. 269 chiefly as carbonates;' excepl magnesium salts which are not at all pre- cipitated, a soluble double salt being at once formed (separation of barium, strontium, and calcium from magnesium). .With salts of silver, copper, cadmium, cobalt, nickel, and zinc the precipitate is redissolved by an excess of Hie ammonium carbonate. The decomposition of carbonates by acids is usually attended by marked effervescence of gaseous C0 2 which reddens moist litmus paper: Na 2 C0 3 + H 2 S0 4 r= Na 2 S0 4 + H 2 + C0 2 . With normal carbonates in cold solution, slight additions of acid (short of a saturation of half the base) do not cause effervescence, because acid carbonate is formed : 2Na 2 C0 3 + H 2 S0 4 = = Na 2 S0 4 + 2NaHCO n ; and when there is much free alkali present (as in testing caustic alkalis for slight admixtures of carbonate), perhaps no effervescence is obtained. By the time all the alkali is saturated with acid, there is enough water present to dissolve the little quantity of .gas set free. But if the car- bonate solution is added drop by drop to the acid, so that the latter is con- stantly in excess, even slight traces of carbonate give notable effervescence. The effervescence of carbonic acid gas, C0 2 , is distinguished from that of H 2 S or S0 2 by the gas being odorless, from that of N 2 3 by its being color- less and odorless ; from all others by the effervescence being proportionally more forcible. It should be remembered, however, that C0 2 is evolved (with CO) on adding strong sulphuric acid to oxalates or to cyanates. On passing the gas, C0 2 , into solution of calcium hydroxide (a); or of barium hydroxide (5); or into solutions of calcium or barium chloride, containing much ammonium hydroxide (c), or into ammoniacal solution of lead acetate (d), a white precipitate or turbidity of insoluble carbonate is obtained. The precipitate may be obtained by* decanting the gas (one- half heavier than air) from the test-tube in which it is liberated into a (wide) test-tube, containing the solution to be precipitated; but the opera- tion requires a little perseverance, with repeated generation of the gas, owing to the difficulty of displacing the air by pouring into so narrow a vessel. The result is controlled better by generating the gas in a large test-tube, having a stopper bearing a narrow delivery- tube, so bent as to be turned down into the solution to be precipitated, (a) C0 2 + Ca(OH) 2 = CaC0 3 + H 2 O (6) CO, + BaCOH), = BaCO 3 + H 2 O (c) CO 2 + CaCL + 2NH 4 OH = CaCO, + 2NH 4 C1 + H 2 O (d) C0 2 + Pb 2 0(C 2 H 3 2 ) 2 = PbC0 3 + Pb(C 2 H 3 2 ) 2 The solutions of calcium and barium hydroxides furnish more delicate tests for carbonic anhydride than the ammoniacal solutions of calcium and barium chlorides, but less delicate than lead basic acetate solution. The latter is so rapidly precipitated by atmospheric carbonic anhydride, that 270 fARBOy DIOXIDE. <22S. T. it cannot be preserved in bottles partly full and frequently opened, and cannot be diluted clear, unless with recently boiled water. Solutions of the acid carbonates effervesce, with escape of C0 2 , on boiling or heating, thus: iKHCO, = K 2 CO, -f H 2 + C0 2 . (Gradually, at 100.) 2NaHC0 3 = Na ; C0 3 - H,0 -f C0 2 . (Gradually, at 70; rapidly at 90 to 100.) 2NH 4 HC0 3 = (NH 4 ) 2 CO 3 + H 5 -f C0 2 . (Begins to evolve CO 2 at 36.) (NH 4 ) 4 H 2 (C0 3 ), = 2(NH 4 ) 2 CO, + H,O + CO 2 . (Begins at 49.) 7. Ignition. On ignition, the normal carbonates of the metals of the fixed alkalis are not decomposed; the carbonates of barium and strontium are dissociated slowly, at white heat,* calcium carbonate at a full red heat, forming the oxide and C0 2 . At a lower temperature, ignition changes all other carbonates to the oxide and C0 2 , except that the carbonates of silver at 250, mercury, and some of the rarer metals are reduced to the metallic state, C0 2 and oxygen being evolved. Stannous and ferrous oxides ignited in an atmosphere of C0 2 are changed to Sn0 2 and Fe . respectively, with evolution of CO (Wagner, Z., 1879, 18, 559). 8. Detection. Carbonates are detected: (a) By the sudden effervescence when treated with dilute acids, (b) By. the precipitate which this gas forms with solutions of Ca(OH) 2 , Ba(OH), , or Pb_0-C_H 0,) 2 . If but a small amount of carbonate be present, the mixture must be warmed to drive the C0 2 over into the reagent (6). A non-volatile acid as H SO H 3 P0 4 should be used, as a volatile acid might pass over with the C0 2 and prevent the formation of a precipitate, (c) Phenolphthalein detects the normal carbonate in solution of the bicarbonate (very delicate). Sodium bicarbonate fails to give a precipitate with magnesium sulphate (distinc- tion from Na 2 C0 3 ) (Patein, J. Pharm., 1892, (5), 25, 448). To detect free carbonic acid in presence of bicarbonates, a solution of 1 part of rosolic acid in 500 parts of 80 per cent alcohol may be employed, to which barium hydroxide has been added until it begins to acquire a red tinge. If 0.5 cc. of this rosolic acid solution be added to about 50 cc. of the water to be tested spring water, for instance the liquid will be colorless, or at most faintly yellowish if it contains free carbonic acid, whereas, if there be no free carbonic acid, but only double salts, it will be red (Pettenkofer, Dingl, 1875, 217, 158). Salzer (Z., 1881, 20, 227) gives a test for free carbonic acid or bicar- bonates in presence of carbonates, founded on the fact that tl ammonia reaction (207, 6) does not take place in presence of f re- borne acid or bicarbonates. This reaction is also used to detect the presence of fixed alkali hydroxides in the fixed alkali carbonates. In presence of a * Barium carbonate decomposes at 1450; strontium carbonate & 1155, and calcium Carbonate at 825, 30, 1. CYANOGEN HYDROCYANIC ACID. fixed alkali hydroxide a brown precipitate is ol>i;iinoO . The first excess of sodium hydroxide beyond the reaction gives a brown precipitate with silver nitrate (Lunge, Z. angew., 1897, 169; Bohlig, Arch. Pharm., 1888, 226, 541). 229. Cyanogen, CN = 26.01. N=C C =N. A colorless, intensely poisonous gas; specific gravity, 1.8064 (Gay-Lussac, Gilb., 1816, 53, 145). The molecular weight shows the molecule to be C 2 N 2 . At ordinary atmospheric pressure it liquifies at 22 (Drion, J., 1860, 41); at 20 under four atmospheres pressure (Hofmann, B., 1870, 3, 658). The gas has an odor of bitter almonds and burns with a red color to the flame forming C0 2 and N . When cooled to about the freezing point of mercury it solidifies to a crystalline ice-like mass (Hofmann, I. c.). Critical temperature, 124 (De- war, C. 2V., 1885, 51, 27). The liquid is colorless, mobile and a non-conductor of electricity. It occurs in the gas from the coke ovens (Bunsen and Playfair, J. in:, 1847, 42, 145). It is prepared: (rt) By heating the cyanides of mercury, silver or gold: Hg(CN) 2 = Hg -f C 2 N 2 . (ft) By the dry distillation of am- monium oxalate: (NH 4 ) 2 C,O 4 = 4H 2 O + C 2 N 2 . (c) By fusing KCN with HgCL: 2KCN + Hg-Cl, = Hg -f 2KC1 + C 2 N 2 . (d) By heating a solution of CuSO 4 with KCN . Half of the CN is evolved and CuCN is formed. If the CuCN be heated with FeCl 3 or MnO, and HC 2 H 3 O 2 , the remainder of the CN is obtained. The gas is purified by absorption with aniline; oxygen, nitrogen and carbon dioxide are not absorbed (Jacquemin, A. Ch., 1886, (6), 6, 140). It combines with Cl , Br , I, S, P, and with many of the metals, reacting very much like the halogens. It dissolves in water, alcohol and ether; but gradually decomposes with formation of ammonium oxalate and carbonate (Vauquelin, A. Ch., 1823, 22, 132; Buff and Hofmann, A., 1860, " 13, 129). At 500 it combines with hydrogen to form HCN (Berthelot, BL, 1880, (2), 33, 2). With Zn it forms Zn(CN) a , rapidly at 100. With HC1 and abso- lute alcohol it forms oxalic ether, which shows cyanogen to be the nitrile of oxalic acid (Pinner and Klein, B., 1878, 11, 1481). With solution of KOH, KCN and KCNO are formed: C 2 N 2 + 2KOH = KCN + KCNO + H 2 O . Com- pare the reaction with chlorine and KOH (270). 230. Hydrocyanic acid. HCN = 27.018. H C = N. 1. Properties. Hydrocyanic acid is a clear, mobile liquid, boiling at 26. At 15 it freezes to a fibrous crystalline mass. S/>ecf/?c gravity at 19, 0.697 (Bleekrode, I' me Rnii. Nor., 1884, 37, 339). It burns with a bluish-red flame, forming H 2 O , CO 2 and N. Its index of refraction is much less than that of water (Mascart, C. r., 1878, 86, 321). It is one of the most active poisons known; of a very characteristic odor, somewhat resembling that of bitter almonds. The antidote is chlorine or ammonia by inhalation. Its water solution decomposes slowly, forming ammonium formate: scarcely at all in tlit- dark. It distils readily unchanged. The U. S. P. solution contains two per cent of HCN. It is a weak acid, scarcely reddening litmus; its salts are partially decomposed by CO* . The free acid or soluble salts when warmed 272 HYDROCYANIC ACID. 230, 2. with dilute alkalis or acids (with strong" acids in the cold) becomes formic acid and ammonia: HCN -f- 2H 2 O = HC0 2 H -f NH 3 . 2. Occurrence. The free acid dots not occur in nature, but in combination in the kernels of bitter almonds, peaches, apricots, plums, cherries and quinces; the blossoms of the peach, sloe and mountain ash; the lea\es of the peach, cherry laurel and Portugal laurel; the 3'oung branches of the peach; the stem-bark of the Portugal laurel and mountain ash; and the roots of the last-named tree, when soaked in water for a time and then distilled, yield hydrocyanic acid, together with bitter-almond oil. Potassium cyanide appears in the deposits of blast furnaces for the smelting 1 of iron ores. 3. Formation. (a) Decomposition of amygdaline by emulsine and distilla- tion. (6) By the action of the electric spark on a mixture of acetylene and nitrogen (Berthelot, (CN) 2 -f 2KCN = (KCN) 2 Hg(CN) 2 . Class I. Double cyanides which are not affected by alkali hydroxides, but are decom- posed when treated with dilute acids: ' KCN) 2 Hg(CN) 2 + 2HC1 = Hg(CN) 2 + 2KC1 -4- LHCN . These closely resemble the double iodides (potassium mercuric), and the double sulphides or thiosalts (69, 5c and 6e). The most frequently occurring of the double cyanides of this class, which dissolve in water, are given below: Potassium (or sodium) zinc cyanide, K 2 Zn(CN) 4 or (KCN) 2 Zn(CN) 2 . Potassium (or sodium) nickel cyanide, K 2 Ni(CN) 4 or (KCN) 2 Ni(CN) 2 . Potassium (or sodium) copper cyanide, K 2 Cu(CN) 3 or (KCN) 2 CuCN . Potassium cadmium cyanide, K 2 Cd(CN) 4 or (KCN)oCd(CN) 2 . Potassium (sodium or ammonium) silver cyanide, KAg(CN) 2 or KCNAgCN . Potassium (or sodium) mercuric cyanide, K 2 Hg(CN) 4 or (KCN) 2 Hg(CN)2 Potassium (or sodium) auric cyanide, KAu(CN) 4 or KCNAu(CN) 3 . Class II. Double cyanides ivhich, as precipitates, are transposed by alkali hydroxides, in dUute solution ('/), and are transposed, without decomposition, by dilute acids (6). In these double cyanides, as potassium ferrous cyanide, K 4 Fe(CN) 6 , the whole of 230, 6. HYDROCYANIC ACID. 273 the cyanogen appears to form a new compound radical with that metal whose single cyanide is insoluble in water; thus, Fe(CN)e as "ferrocyanogen," giving K4Fe(CN)e as "potassium ferrocyanide " (for the potassium ferrous cyanide). These more stable double cyanides or "ferrocyanides," etc., correspond to the platinic double chlorides or " chloroplatinates " (74, 5-), and the palladium double chlorides, or chloropalladiates (106, 5';). The most frequently occurring of the double cyanides of this class, which are soluble in water, are given below. (a) Cu 2 Fe(CN) 6 + 4KOH = 2Cu(OH) 2 + K4Fe(CN) 6 (&) K4Fe(CN) 6 + 2H 2 S0 4 = 2K,SO 4 + H 4 Fe(CN) 6 2K 3 Fe(CN) 6 + 3H 2 SO 4 = 3K 2 SO 4 + 2H 3 Fe(CN) 6 Alkali ferrocyanides, as K 4 Fe"(CN) 6 , potassium ferrocyanide. Ferricyanides, as K 3 Fe'"(CN) 6 , potassium ferricyanide. Cobalticyanid.es, as K 3 Co'"(CN)<; , potassium cobal icyani;le. Man gam cyanides, as K 3 Mn'"(CN) 6 , potas-ium manganicyanide. Chromicyanides, as K 3 fCr"')(CN) 6 , potassium chromicyanide. The easily decomposed double cyanides of Class I are, like the single cyan- ides, intensely poisonous. The difficultly decomposed double cyanides of Class II. are not poisonous. The Single Cyanides are transposed by the stronger mineral acids, more or loss readily, with liberation of hydrocyanic acid, HCN, effervescing from concentrated or hot solutions, remaining dissolved in cold and dilute solu- tions. Mercuric cyanide furnishes HCN by action of H 2 S , not by other acids. The cyanides of the alkali and alkaline earth metals are decomposed by all acids even the carbonic acid of the air and exhale the odor of hydrocyanic acid. Solution of silver nitrate precipitates, from solutions of cyanides or of hydrocyanic acid (not from mercuric cyanide) silver cyanide, AgCN , white, insoluble in dilute nitric acid, soluble in ammonium hydroxide, in hot ammonium carbonate, in potassium cyanide, and in thiosulphates uniform with silver chloride. Cold strong hydrochloric acid decomposes it with evolution and odor of hydrocyanic acid (recogni- tion from chloride); and when well washed, and then gently ignited, it does not melt, but leaves metallic silver, soluble in dilute nitric acid, and pre- cipitable as chloride (distinction and means of separation from chloride). Solution of mercurous nitrate, with cyanides or hydrocyanic acid, is resolved into metallic mercury, as a gray precipitate, and mercuric cyanide and nitrate, in solution. Salts of copper react, as stated in 77, 66; salts of lead, as stated in 57, 6&. Ferrous salts, added to saturation, precipitate from solutions of cyan- ides, not from hydrocyanic acid, ferrous cyanide, Fe(CN) 2 , white, if free from the ferric hydroxide formed "by admixture of ferric salt, and, with the same condition, soluble in excess of the cyanide, as (with potassium cyanide), (KCN) 4 Fe(CN) 2 == K 4 Fe(CN) 6 , potassium ferrocyanide (a). On acidulating this solution, it gives the blue precipitates with ferric salts (b) : (a) 2KCN + FeSO 4 = Fe(CN) 2 + K 2 SO 4 Pe(CN) 2 + 4KCN = K 4 Fe(CN) 6 (6) 3K 4 Fe(CN) 6 + 4FeCl 3 = Fe 4 (Fe(CN) fl ) + 12KC1 This production of the blue ferric ferrocyanide is made a delicate test for 274 HYDROCYANIC ACID. 230, V. hydrocyanic acid, as follows: A little potassium hydroxide and ferrous sulphate are added, the mixture digested warm for a short time; then a very little ferric chloride is added, and the whole slightly acidulated (so as to dissolve all the ferrous and ferric hydroxides), when Prus- sian blue will appear, if hydrocyanic acid was present (Link and Moeckel, Z., 1878, 17, 456. For identification of traces of hydrocyanic acid (less than 0.00002 g. in 1 c.c.) add two drops of 10% solution of sodium hydroxide, evaporate almost to dryness, cool, add one drop of a 2% solution of ferric sulphate and allow to stand in the cold for 10-15 minutes. Heat gently with two or three drops of strong hydrochloric acid and cool. The undiluted blue-green solution shows, when carefully diluted, a blue color. (G. Druce Lander und Walden, nach Leitschr. J. Unters. d. Nahrungs u. Genussm. 23 [1912] 399.) Solution of nitrophenic acid, picric acid, C 6 H 2 (N0 2 ) 3 OH , added, in a small quantity, to a neutralized solution of cyanides of alkali meta:s, on boiling( and standing), gives a blood-red color, due to picrocyanate (as KC 8 H 4 N 5 6 ). This test is very delicate, but not very distinctive, a^- var- ious reducing agents give red products with nitrophenic acid (\ogel, C. N., 1884, 50, 270). The fixed alkali hydroxides, in boiling solution, strongly alkaline, g mdu- ally decompose the cyanides with production of ammonia and formate: HCN + KOH + H 2 = KCH0 2 +NH, . Ferrocyanides and ferricya aides finally yield the same products. Dilute alkalis, not heated, transpose, as by equation a, class II above. Cyanides are strong reducing agents. The action is not so mark 3d in solution as in state of fusion (7). Permanganates are reduced by cyan- ides, and cupric hydroxide in alkaline solution forms Cu'. Solvtions of cyanides on exposure to the air take up some oxygen with formation of a cyanate: 2KCN -(- 2 = 2KCNO . Commercial potassium cyanide a ways contains some potassium cyanate. By warm digestion of a cyanide with sulphur or with yellow ammonium sulphide a thiocyanatc is formed (8). Hydrocyanic acid reduces Pb0 2 , forming Pb(CN) 2 and CN : PbO, + ;;HCN = Pb(CN), + C 2 N 2 + 2H 2 (Liebig, A., 1838, 25, 3). With HCI: and H 2 2 oxamide is formed (Altfield, J. C., 1863, 16, 94). Chlorine forms with hydrocyanic acid a cyanogen chloride (Serullas, A. Ch., 182c c , 38, 370); with iodine the reaction is not so marked, but a similar product is formed (Meyer, P., 1887, 20, III, 704). Concentrated sulphuric acid decomposes all cyanides. 7. Ignition, By fusion with fixed alkalis, cyanides and all compounds containing cyanogen yield ammonia. In state of fusion cyanides are very efficient reagents for reduction of metals from their oxides or sulphides. 231. HYDROFERROCYANIC ACID. 275 to il e metallic state (69, !). The cyiiiuttcs or thiocyanates formed in the reaction are not readily decomposed by heat alone. 8. Detection. Cyanides may he detected: (a) By the odor of the free arid upon decomposition of I he cyanide with acids. This test must be appl }d with extreme caution as the evolved HCN or CN is very poisonous. (/>) I v formation of a ferrocyanide and its reaction with ferric salts, as dese- : hed in G. (c) The production of the red ferric thiocyanate is a test for i ifdrocyanic acid, more delicate than formation of ferrocyanide. By warn, digestion this reaction occurs: 2KCN -f- S 2 = 2KCNS ; or: 2(NH 4 ) 2 S 4 + 4HCN = 4NH.CNS + 2H.S + S 2 To the material in an evaporating-dish, add one or two drops of yellow ammonium sulphide, and digest on the water-hath until the mixture is colorless, and free from sulphide. Slightly acidulate with hydrochloric acid (which should not liberate H 2 S), and add a drop of 'solution of ferric chlonde; the blood-red solution of ferric thiocyanate will appear, if hydro- cyan c acid was present (Link and Moeckel, /. c.). (d: Link and Moeckel also recommend the following test for cyanides, delic ite to 1-3,000,000. Saturate a fdter paper with a four per cent alcololic solution of guaiac; allow the alcohol to evaporate; then moisten the "paper with a one-fourth per cent solution of copper sulphate, and allow the unknown solution to trickle over this test paper. A deep blue cokn indicates the presence of a cyanide. TV detect cyanides in presence of ferri- and ferrocyanides it is directed to add tartaric acid and, in a distilling flask, pass a current of carbon dioxile, warming not above (>0. Test the distillate by the methods given above. Ferro- and ferricyanides do not yield HCN under 80 (Hilger and Tamba, Z., 1891, 30, 529; also Taylor, C. N., 1884, 50, 227). 9. Estimation. (a) The nearly neutral solution of cyanide is titrated with standard silver nitrate. No precipitate occurs as long as two molecules of alkali cyanide are present to one of silver nitrate. Soluble AgCN,KCN is form id. As soon as the alkali cyanide is all used in the formation of the doub e cyanide, the next molecule of silver nitrate decomposes a molecule of tJie double salt, forming- two molecules of insoluble silver cyanide; giving a white precipitate for the end reaction. Chlorides do not interfere (Liebig, A., 1851, 77, 102). (6) By titratioii with a standard solution of HgCl 2 , applicable in presence of cyanates and thiocyanates (Hannay, J. C., 1878, 33, 245). 231. Hydroferrocyanic acid. H 4 Fe(CN) = 215.932. H' 4 Fe"(CN)-' . Absolute hydroferrocyanic acid (230, 0, Class IT.), is a white solid, freely soluble in water and in alcohol. The solution is strongly acid to test-paper, and decomposes carbonates, with effervescence, and acetates. It is non-volatile, but absorbs oxygen from the air, more rapidly when heated, evolving hydro- cyanic acid and" depositing Prussian blue: 7H 4 Fe(CN) + O., = Fe 4 (Fe(CN) 6 ) 3 -f 2H 2 O + 24HCN . Potassium ferrocyanide is the usual starting- point in the preparation of the 276 HYDROFERROCYANIC ACID. 231. free acid or any of the salts. It is prepared by fusing together in an iron kettle nitrogenous animal matter (blood, hair, horn, hoof, etc.), commercial potash (KOH), and scrap iron. The ferrocyanide is formed when this mass is digested with water. The nitrate is evaporated to crystallization (lemon-yellow prism), soluble in four parts of water. Hydroferrocyanic acid is formed by transposition of metallic ferrocyanides in solution, with strong 1 acids (a). When the solution is heated, hydrocyanic acid is evolved; in the case of an alkali ferrocyanide, without absorption of oxygen (&). Potassium ferrocyanide and sulphuric acid are usually employed for preparation of hydrocyanic acid (c) : (a) K 4 Fe(CN) 6 + 2H 2 S0 4 = 2K 2 S0 4 + H 4 Fe(CN) 6 (6) 3H 4 Fe(CN) 6 + K 4 Fe(CN) 6 = 2K 2 FeFe(CN) G + 12HCN (c) 2K 4 Fe(CN) G + 3H 2 S0 4 = 3K 2 S0 4 + K 2 FeFe(CN) 6 + 6HCN The ferrocyanides of the alkali metals, strontium, calcium and magnesium, are freely soluble in water; of barium, sparing^ soluble; of the other metals, insoluble in water. There are double ferrocya Hides; soluble and insoluble; that of barium and potassium is soluble, but potassium calcium ferrocya&ide is in- soluble. The most of the ferrocyanides of a heavy metal and an alkali metal are insoluble. Potassium and sodium ferrocyanides are precipitated from their water solutions by alcohol (distinction from ferricyanides). The soluble ferrocyanides are yellowish in solution and in crystals, white when anhydrous. The insoluble ferrocyanides have marked and very diverse colors, as seen below. Solutions of alkali ferrocyanides, as K 4 Fe(CN) 6 , give, with soluble salts of: Aluminum, a white precipitate, A1(OH) 3 and Fe(CN) 2 (formed slowly). Antimony a white Bismuth, a white Cadmium, a white Calcium, a white Chromium, no Cobalt, a green, then gray Copper, a red-brown Gold, no Iron (Fe"), w^hite, then blue Iron (Fe'"), a deep blue Lead, a white Magnesium, a white a yellow-white Manganese, a white Mercury (Kg-'), a white Mercury (Hg"), a white Sb 4 [Fe(CN) ] 3 .25H 2 0. Bi 4 (Fe(CN) 6 ) 3 . Cd 2 Fe(CN) G (soluble in HC1). K 2 CaFe(CN) . Co 2 Fe(CN) 6 . Cu 2 Fe(CN) 6 . K 2 FeFe(CN) 6 . Fe 4 (Fe(CN) 6 ) 3 . Pb 2 Fe(CN) 6 . (NH 4 ) 2 MgFe(CN) 6 (in presence of NH 4 OH) K 2 MgFe(CN) (only in concentrated solu- tion). Mn 2 Fe(CN) 6 (soluble in HC1). Hg 4 Fe(CN) 6 (gelatinous). Hg 2 Fe(CN) , turning to Hg(CN) 2 and Fe 3 (Fe(CN) 6 ) 2 , blue. Molybdenum, a brown Nickel, a greenish-white Silver, a white Tin (Sn" and Sniv), white Uranium (uranous), brown TJranium (uranyl), red-brown Zinc, a white, gelatinous (slowly turning blue). Ni 2 Fe(CN) 6 Ag 4 Fe(ClSr) 8 (gelatinous). UFe(CN) 6 . (U0 2 ) 2 Fe(CN) 6 . Zn 2 Fe(CN) G . See Wyrouboff (A. Ch., 1876 (5), 8, 444; and 1877, (5), 10, 409). Insoluble ferrocyanides are transposed by alkalis (230, 6, Class II.) 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. HYDROFERRICYANIC ACID. 277 nevertheless reducing- agents, of moderate power: 2K 4 Fe(CN) 8 + CL = 2K 3 Fe(CN) (i + 2KC1 . PbO. with sulphuric acid forms Pb" and H 3 Fe(CN) . Ag' \vitli fixed- alkali forms an alkali ferricyanidc and metallic silver. CrVi w ith 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(CN) . Ni'" with acetic acid gives Ni" and H 3 Fe(CN) . Mn0 2 with phosphoric acid gives Mn" and H 3 Fe(CN) . Mnvn forms with potassium hydroxide Mn(X and potassium ferricyanide. \Vith sulphuric acid, manganous sulphate and hydroferricyanic acid. Ferricyanides when boiled with NH 4 OH give ferrocyanides (Playfair, J. C., 1857, 9, 128). HNO., forms first hydroferricyanic acid, then hydronitroferricyanic acid and NO. HNO 3 forms hydroferricyanic acid, and then hydronitroferricyanic acid, NO being evolved. Cl forms first hydroferricyanic and hydrochloric acids. Excess of chlorine to be avoided in preparation of ferricyanides. HC10 3 forms hydroferricyanic and hydrochloric acids. Br forms hydroferricj^anic and hj^drobromic acids. HBrO 3 forms hydroferricyanic and hydrobromic acids. I , iodine is decolored by potassium ferrocyanide, and some potassium ferri- cyanide and potassium iodide are formed. The action is slow and never complete (Cmdin's Hand-book, 7, 459). HI0 3 forms hydroferricyanic acid and free iodine. In analysis, soluble ferrocyanides are recognized by their reactions with ferrous and ferric salts and ccpper salts (see 6ft, 126 and 77). Separated from ferricyanide, by insolubility of alkali salt in alcohol. Separation of hydro- ferrocyanic acid from hydroferricyanic acid according to Ph. E. Browning and H. E. Palmer (Zeitschr. f. anorg. chem., 54, 315, nach. Zeitschr. f. anal, chem., 60 (1911) 771). Acidify 5-10 c.c. of the solution which is to be tested with acetic or hydrochloric acid and add a solution of a thorium salt. Thorium ferrocyanide will be precipitated. Shake up with finely divided asbestos, filter, wash the pre- cipitate, add sodium hydroxide to same and in this filtrate test for ferrocyanide. To the filtrate from the thorium ferrocyanide add a solution of a cadmium salt. Cadmium ferricyanide will be precipitated and is treated like thethorium precipitate. Ferrocyanides are estimated in solution with sulphuric acid by titrating with standard KMnO4 . Also by precipitation with CuSO t either for gravimetric determination or volumetrically, using a ferric salt as an external indicator. 232. Hydroferricyanic acid. H 3 Fe(CN) 6 = 214.924. H' 3 Fe"'(CN)-' 6 . Absolute hydroferricyanic acid, H 3 Fe(CN) 6 , is a non- volatile, crystallizable solid, 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 FeiCN) f; until a few drops of the liquid gives a brownish color, but no pre- cipitate with a ferric salt. The solution is evaporated to crystallization and the salt repeatedly recrystallized from water as large red prismatic crystals, very soluble in water, freely soluble in alcohol (distinction from KjFe(CN) 6 ). The free acid is made by adding to a cold saturated solution of K 3 Fe(CN) 6 three volumes of concentrated HC1 and drying the precipitate which forms, in a vacuum (Joannis, C. r., 1882, 94, 449, 541 and 725) lustrous, brown sh- green needles, very soluble in water 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- 278 H Y DROP ERHI CYAN 1C ACID. 232. cyanides are but slightly, or not at all, precipitated from their water solutions by alcohol (separation from ferrocyanides). Ferricyanides are not easily decomposed by dilute acids; but alkali hydrox- ides, either transpose them or decompose their radicals (230, 6). 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,[-Fe(CN) '0 2 forms H 4 Fe(CN) fi and H PO, . H 2 S ^orms S, then H..SO, and H,Fe(CN), ; (Wallace, I.e.). SO, forms H 2 S0 4 and H,Fe(CN) . Cl decomposes ferricyanidcs. HC1O, acts upon K 3 Fe(CN) , forming- potassium superferricyanide, K 2 Fe(CN) 8 (Skraup, A., 1877, 189, :!(ks). HI fcrms H 4 Fe(CN) (J and I . Ferricyanides in solution are detected by the reactions with ferrous and ferric salts (126, G7>). Insoluble compounds are ignited (under a hood) with a fixe 1 alkali, giving 1 an alkali cyanide, ferric oxide, and an oxide of the metal in co nbination. Detect the alkali cyanide as directed (230, 8). A ferri- cyani le is estimated by reduction to ferrocyanide with KI in presence of con- cent r.ited HC1; the liberated iodine being titrated with standard Na 2 S a O 3 . Or it is reduced to ferrocyanide by boiling with KOH and FeS0 4 , filtering, acidulating with H 2 S0 4 and titrating with KMnO 4 . 233. Cyanic acid. HCNO = 23.018. H C=N. The cyanates of the alkalis and of the fourth-group metals may be made by passii g 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 -f 2KOH = KCNO + KCN + H 2 KCN + PbO KCNO + Pb 4KCN + Pb 3 4 = 4KCNO + r,Pb The free acid may be obtained by heating cyanuric acid, H 3 C 3 N 3 O3 , to redness, better in an atmosphere of CO 2 . Cyanic acid is found in the dis- tillate H 3 C 8 N 3 8 = 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 forme d by transposing metallic cyanates with the stronger acids in the pres- ence >f water, by which it is changed into carbonic anhydride and ammonia: HCNO + H 2 O = NH 3 + CO 2 . The cyanates, therefore, when treated with hydrr-.'hloric or sulphuric acid, effervesce with the escape of carbonic anhydride (disti -iction from cyanides), the pungent odor of ci/iinic acid being perceptible: 2KCEO + 2H_,SO 4 ' + 2H 2 O = K 2 S0 4 + (NH 4 ) 2 SO 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 being insoluble or sparingly soluble. All the solutions gradrtlly decompose, with evolution of ammonia. Silver cyanate is sparingly soluble in hot water, readily soluble in ammonia; soluble, with decomposition, in dibite nitric acid (distinction from cyanide). Copper ci/anate is precipitated green i sh-y ellow. Atni.innium cyanate in solution changes gradually, or immediately when boiled, to we- 1, or carbamide, with which it is isomeric: NH 4 CNO = CO(NH,), . The latter is recognized by the characteristic crystalline laminae 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,O(HgO) 2 , not turned yellow (decomposed) by solution of sodiuri carbonate (no excess of mercuric nitrate being taken). Solution of urea, on boiling, is resolved into ammonium carbonate, which slowly vapori/cs: CH 4 N,O -f 2H 2 O = (NH.) .CO;, . 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. 280 THIOCJANIC ACID. 234. 234. Thiocyanic acid. HCNS = 59.088 . H S C = N. An aqueous solution of HCNS may be obtained by treating 1 lead thiocyanate suspended in water with H 2 S , also by treating barium thiocyanate with H 2 S0 4 in molecular proportions. The anhydrous acid is obtained by treating dry Hg(CNS) 2 with H,S . Potassium thiocyanate is formed by fusing KCN with S. Or two parts of K 4 Fe(CN) e 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 4 Fe(CN) G + 3S 2 = 4KCNS + Fe(CNS) 2 4KCN + 4K 2 S 2 3 = 4KCNS -f 3K 2 S0 4 -f K 2 S 2K 4 Fe(CN) 6 + 12K 2 S 2 S = 12KCNS + 9K 2 SO 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 1 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 products. 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(CNS) 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) 3 , decolored by solu- tion of mercuric chloride (126, 6&, 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. Lead, 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. 281 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. 1'latinic chloride, gradually added to a hot, concentrated solution of potassium thiocyanate, forms a deep-red solution of double thiocyanate of potassium and platinum (KCNS),Pt(CNS) 4 , or more properly, K,Pt(CNS) r> , potassium thfocyanoplatinate. 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,S0 4 (separation from Ag-Cl) (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 perthiQcyanogen, 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, pcrtJiiori/aiiic arid, (HCN) 2 S 3 , 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 eases the fame reaction is assumed as probable. PbO 2 and Pb 3 4 form Pb" and sulphuric acid, in acid mixture only (Hardow, ,/. (7., 1859, 11, 174). H 3 AsO 4 forms H 3 AsO 3 , hydrocyanic and sulphuric acids. Co'" forms Co" , hydrocyanic and sulphuric acids. Ni'" forms Ni" , hydrocyanic and sulphuric acids. Crvi 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 L) forms sulphuric acid and nitric oxide. HNO 3 forms sulphuric acid and nitric oxide. Cl 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 Cl . HC1O 3 forms sulphuric, hydrochloric and hydrocyanic acids. Br forms HBr and H 2 S0 4 ; but with alkalis, a bromide and sulphate. HBrO 8 forms HBr and H 2 SO 4 . HIO 3 forms H 2 S0 4 and free iodine. 235. Nitrogen. N = 14.01. Valence one to five (11). 1. Properties. Weight of molecule, N 2 , 28.08. Vapor density, 14 (Jolly, TF. 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' 3 , N 2 O , NO . N 2 O 3 , NO., and N,O, . The number of organic compounds contain- ing nitrogen is very large. The nitrogen in all compounds that are the 282 ^HYDRAZOIC ACID. 235,2. immediate products of vegetable growth bns a valence of minus three and may without change of bonds be converted into N '"H'a . 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 1 removed by red-hot copper, the C0 2 by potassium hydroxide, the ammonia and water by passing through H 2 SO 4 .* (6) Ignition of ammonium dichromate, *NH 4 ) 2 Cr 2 O 7 = N 2 + Cr 2 O 3 + 4H 2 O . (c) By heating ammonium nitrate and peroxide of manganese to about 200 (Gatehouse, C. N., 1877, 36, 118). (d) Ignition of NH 4 C1 and K 2 Cr 2 O 7 : 2NH 4 C1 + K 2 Cr 2 O 7 = ^KCl + N 2 + Cr 2 O 3 + 4H 2 U . Unless the tem- perature be carefully guarded, traces of NO are formed, which may be removed by passing the gases through FeSO 4 . (e) Action of chlorine upon NH 3 : 8NH.3 -j- OC1 2 = 6NH 4 C1 + N 2 . The NH 3 must be kept in excess to avoid the forma- tion of the dangerously explosive chloride of nitrogen, NC1 3 . (/) Removing the oxygen from the air by shaking with NH 4 OH and copper turnings.* (g) Burning phosphorus in air over water.* (h) By passing air through a mixture of FeS and sawdust; then through a pyrogallate solution, and finally through concentrated H 2 SO 4 .* (i) By shaking air with Fe(OH) 2 and Mn(OH) 2 .* (j) By passing air through an alkaline pyrogallate.* (/c) By passing air, from which UOohas been removed, mixed with hydrogen over heated platinum black, the hydro- gen having been added in just sufficient quantity to form water with all the oxygen * (Damoulin, /., 1851, 321). (Z) By warming a concentrated solution of NH 4 NO 2 or a mixture of KNO 2 and NH 4 C1: NH 4 NO 2 = N 2 + 2H 2 O . Potassium dichromate is added to oxidize to nitric acid any of the oxides of nitrogen that may be formed (Gibbs, B., 1877, 1387). (ra) By action of potassium or sodium hypobromite upon ammonium chloride: 3NaBrO + 2NH 4 C1 = N 2 +3NaBr + l^HCl -f 3H 2 O . 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 containing the so-called nitrifying bacteria growing on the roots of leguminous plants absorb nitrogen 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. (a) As free nitrogen by measuring the volume of the gas. (6) By oxidation of the organic substance with hot concentrated H 2 SO 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 CO 2 by KOH and determining the nitrogen by volume. (For Hydroxylamine, see foot-note, page 286.) 236. Hydrazoic acid (Azoimide). N 3 H = 43.038. N\ Constitution, II >NH . W Curtius, B., 1890, 23, 3023. A clear mobile liquid of 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 ; nd alcohol. An acid of marked acidity, dissolving a number of metals with * Nitrogen made from the air is not pure. It contains about one per cent of argon and smaller amounts of krypton, neon, and xenon. Because of the presence of these impurities its density is greater than that of nitrogen prepared from chemical compounds. (Ramsey and Rayleigh.) 238, <;. NITROUS OX1DK \1TRIC OXIDE. evolution of hydrogen. Its salts, the trinitrides of the metals of the alkalis and the alkaline earths, are soluble in water and crystallizable (Dennis, J. Am. /S'oc., 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. Nor.. 1S9<>, 18, 947). Hydro- nitric acid is formed by treating 1 ammonia with sodium, and the resulting 1 sodamide, NaNHo , with nitrous oxide: 2NaNH., + N 2 = NaN 3 + NaOH -j- NH 3 (Wislicenus, J3., 1892, 25, 2084). 237. Nitrous oxide. N 2 = 44.02 . !T 2 0-", N N . Nitrous oxide becomes a colorless liquid at under pressure of three atmospheres (Farady, A., 1845, 56, 157). Melts at 99 and boils at 92 (\\ 'ills, -/. C., 1S74, 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, PyrocJiemisch. Untersuch., 1885). Passed over red-hot iron N and Fe^Os 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 NO, = N 2 O + 2H,O . (ft) By passing NO through solution of S0 2 . (c) By action of HNO 3 ; sp. (jr., 1.42, diluted with an equal volume of water, upon metallic zinc, (d) A mixture of five parts of SnCl 2 , ten parts of HC1 , sp. (jr., 1.21, and nine parts of HNO 3 , sp. (jr., 1.3, is heated to boiling: 2HN0 3 + 4SnCl 2 + 8HC1 = 4SnCl 4 + N,O + 5H 2 O (Campari, J. C., 1889, 55, 569). 238. Nitric oxide. NO = 30.01 . N"0-", N = . 1. Properties. The vapor density (15) shows the molecule to be NO (Daccomo and Meyer, B., 1887, 20, 1832). Under pressure of one atmosphere it is liquified at 153.6, and under 71.2 atmospheres at 93.5, and solidifies at 167 (Olszewski, C. r., 1877, 85, 1016). Odor and taste unknown, on account of its immediate conversion into NO 2 on exposure to the air. 2. Occurrence. Not found free in nature. 3. Formation. () + 2AgCl and then Ba(NO 2 ) 2 + ZnSO, = Zn(N0 2 ) 2 + BaSO 4 . 4. Preparation. Same as above. 5. Solubilities. Silver nitrite is only sparingly soluble (120 parts of cold water). The other normal nitrites are soluble; but many basic nitrites are insoluble. Nascent hydrogen in presence of an alkali reduces nitrates to nitrites; e. g., sodium amalgam, aluminum wire in hot KOH , etc. Used in excess the nascent hydrogen reduces the nitrogen still further, forming NH 3 . G. Reactions. A. With metals and their compounds. Nitrous acid acts sometimes as an oxidizer, sometimes as a reducer; in the former case NO is iiHUttlly produced (under some conditions N 2 O , N and NH 3 are formed); in the latter case nitric acid is the usual product, but sometimes NO 2 is produced. 1. Pb0 2 becomes Pb" and nitric acid. 2. Hg' becomes Hg and nitric acid. 3. Crvi becomes Cr'" and nitric acid. 4. Co" becomes Co'" and nitric oxide. Excess of KN0 2 with acetic acid is used to separate cobalt from nickel (132, (k?). 5. Ni"' becomes Ni" and nitric acid. 6. Mn" + n becomes Mn" and nitric acid. B. With non-metals and their compounds. 1. H 4 Fe(CN) 6 becomes first H 3 Fe(CN) and then hydronitroferricyanic acid. Solution of indigo in sulphuric acid is bleached by nitrites. 2. Nitrites are decomposed by nitric acid. 3. HH,PO, becomes H 3 PO 4 and NO. J/. H,S does not displace or transpose alkali nitrites, but if acetic acid be added to liberate the nitrous acid, then S and NO are produced. H 2 S0 3 be- comes ELSO 4 and chiefly NO . With excess of H,SO 3 , N 2 O or NH 3 is formed. See Weber, For///., I860, 127, 543, and 1867, 130, 277; Fremy, C. r., 1870, 70, 61. 5. HC10 :1 becomes Cl and HN0 3 . G. HBr0 3 becomes Br and HNO 3 7. HI becomes 1 and NO . HIO 3 becomes 1 and HNO 3 . 7. Ignition. In general nitrites are changed to oxides, but with potassium and sodium nitrites a white heat is required, and with nitrites of Ag . Hg , Au and Pt the dissociation goes a step further, the free metals being produced. 8. Detection. (1) Formation of brown ring when ferrous sulphate solution and a nitrite is acidulated with acetic acid. Nitrates require a stronger acid for 241, 1. NITROGEN PEROXIDE NITRIC ACID. 285 their transposition. (2} A mixture of a nitrite and KI liberates iodine on adc'i- lion of acetic acid (nitrates requiring a stronger acid for transposition). (3) Nitrous acid with iodic acid liberates iodine, and nitric acid is produced. (4) Solution of potassium permanganate acidified with sulphuric acid is reduced by nitrites (distinction from nitrates). 9. Estimation. Acidify with acetic acid, distil and titrate the distillate with standard solution of permanganates, 240. Nitrogen peroxide (dioxide). N0 2 = 46.01 . Vapor density, 2.'! ( R-misay, ./. ('.. 1S!)0, 57, f>00). Melting- point, 10 C (Deville and Troost, C. r., 1867, 64, :.>:>?). Boils at :>l.r>4 (Thorpe, J. C., 18b(). 37, 224). Below 10 it is a white crystalline solid. Between 10 and 21.fi a liquid; nearly colcrless at !) , yellow at 0. At 21.04, orange, growing nearly black as 1h? temperature rises. The gas does not support combust io , of ordinary fuels, and is poisonous when inhaled. It dissolves in water, form- ing a greenish-blue solution containing nitrous and nitric acids. With an aqueous solution of a fixed alkali a nitrate and nitrite are formed: 2N0 2 + 2KOH = KN0 2 + KNO 3 + H 2 . 241, Nitric acid. HN0 3 = 63.018. II H'N v O-" 3 , H N = . 1. Properties. Nitric anhydride, N 2 0- , is a colorless solid, melting at C.0 with partial decomposition to N0 2 and O, and if exposed to direct sunlight decomposition begins at lower temperatures. Nitric acid, HN0 3 , has not been perfectly isolated; that containing 9D.8 per cent of HN0 3 is a colorless highly corrosive liquid (Roscoe, A., 1860, 116, 211), solidifies at 47 (Berthelot), boils at 86, but dissociation begins at a lower temperature and is complete at 255: 4HNO 3 = 4N0 2 + 2!L,O -f- 2 (Carius, B., 1871, 4, 82S). If the very dilute acid be boiled, it becomes stronger, and if a very strong acid be boiled it becomes weaker, in both cases a sp. (jr. of 1.42 and boiling point of 120 being reached; the acid then contains about 70 per cent of HNO 3 (Kolbe, A. Oh., 1867 (4), 10, 136). This is the acid usually placed on the market. The reagent usually employed has a sp. (;r. of 1.2 (Fresenius standard). The so-called fuming acid has a specific gravity of 1.50 to 1.52. The stronger acid should be kept in a cool dark place to avoid decom- position. 2. Occurrence. Found in nature as nitrates of K, Na , NH 4 , Ca , Mg , and of a few other metals, the most abundant supply coming from Chili and Bolivia as sodium nitrate, " Chili saltpeter." 3. Formation. (a) Oxidation of nitrogenous matter in presence of air, moisture and an oxide or alkali; (&) by oxidation of NO, N 2 3 or NO, by oxygen (or air) in presence of moisture; (c) from NH 3 , by passing a mixture of NH 3 and oxygen through red-hot tubes. 4. Preparation. By treating nitrates with sulphuric acid and distilling. Nitrates may be made: (a) By dissolving the metal in nitric acid, except those whose metals are not attacked by that acid, e. g., An , Pt , Al and Cr ; and also, antimony forms Sb 2 5 , arsenic, H 3 As0 4 and with excess of hot acid tin forms metastannic acid H 10 Sn 5 15 . (&) By adding HNO., to the oxides, hydroxides or carbonates. All the known nitrates can be made 286 CITRIC ACID. 241, 5. in this manner, (c) By long continued boiling, the chlorides of all ordi- nary metals are completely decomposed, no chlorine remaining, except the chlorides of Hg , Ag , Au and Pt , which are not attacked, and the chlorides of tin and antimony, which are changed to oxides. (Wurtz, Am. S., 1858, 75, 371; Johnson, Proc. Am. Ass. ScL, 1894, 163.) The anhydride is made: (a) By passing chlorine over silver nitrate: 4AgN0 3 + 2C1 2 == 4AgCl + 2N 2 5 + 2 . (&) By adding anhydrous P 2 5 to HN0 3 : 2NH0 3 + P 2 5 == 2HP0 3 + N 2 5 . 5. Solubilities. All normal nitrates are soluble. A few are decom- posed by water, e. g., Bi(N0 3 ) 3 + H 2 = = BiON0 3 + 2HNO ; , . Most nitrates are less soluble in nitric acid than in water, e. g., Cd , Pb , Ba , etc.; the barium nitrate being completely insoluble in HN0 3 , sp. gr., 1.42. Nitric acid decomposes the sulphides of all ordinary metals, except mercuric sulphide which "by long continued boiling with the concentrated acid becomes 2HgS.Hg(N0 3 ) 2 , insoluble in the acid. 6. Reactions. .4. With metals and their compounds. Nitric acid is a powerful oxidizer but unless warmed acts more slowly than chlorine. It can never be a reducer. The following products are formed: H, NH 8 , H 2 NOH *, N , N 2 , NO , HN0 2 , N0 2 . If the acid is concentrated, in excess and hot, the product is usually entirely nitric oxide, colorless, but changing to the red colored N0 2 by coming in contact with the air. Excess of the reducer, low temperatures and dilute solutions favor the production of nitrogen compounds having lower valence and of hydrogen. Xascent hydrogen usually forms NH :i , always the ultimate product if the hydrogen be produced in alkaline mixture. Nitric acid oxidizes all ordinary metals. (It does not act upon chro- mium, gold or platinum.) It forms nitrates, except in the case of tin, antimony, and arsenic, with which it forms H 10 Sn 5 15 , Sb 2 ri , and H 3 As0 4 . With the respective metals it forms Hg' or Hg", Sn" or Sn"", As'" or As v , Sb'" or Sb v , Fe" or Fe'", according to the amount of nitric acid employed. With copper it forms cupric nitrate (never cuprous); with cobalt it forms cobaltous nitrate. * Hydroxylamine, NH 2 OU, is formed by the reducing action of Sn and HC1 upon NO,N 2 O 3 , HXO 3 , etc. (Lessen, A., 1888, 252, 170); also by the action of H 2 S, SO 2 , K, ]Vn, M g, Zii, and Al upon HXO 3 . or by the action of H 2 S upon certain nitrates (Divers and Haga, C. jV., 1886, 54, 271 1. By action of sodium amalgam upon s jdium nitrite solution, XH 2 OH is produced along with nitrous oxide, free nitrogen, ammonia, sodium hyponi trite, and sodium hydroxide, the highest yield of the hydroxylamine being obtained when the nitrite solution is as dilute as one in fifty, the mix- ture being kept cold (Divers, J. C., 1899,75, 87 and 89). It is a base with an alkaline reaction and is a strong reducing agent. When pure it is a crystalline solid, odorless, melting at 33.05, boiling at 58 at 22 mm. pressure; oxidized by oxygen to HNOa (Lobry de Bruyn, B., 1892, 23; 3, 190 and 684). It is a good antiseptic and preservative. It combines with acids to form salts: NH 2 OH +HC1 =NH 2 OH.HC1 . Hydroxylamine hydrochloride is decomposed by alkalis form- ing the free base, which is decomposed by the halogens, KMnO4 , K2Cr 2 O? , BaCh and PbO? Its solution in ether reacts with sodium, forming a white precipitate ofNH?ONa, 241, B, 8. 'NITRIC ACID. 287 1. Pb0 2 is not changed. Pb 3 4 is changed thus: Pb 3 4 + 4HNO, = PbO, + 2Pb(NO ..),, + 2K.O . 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 0. , 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) G 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 = 2NO + HN0 3 -f- H 2 . S. P, PH :; , HH 2 P0 2 and HgPOg become H 3 P0 4 . That is P v ~ becomes P v . 4. S becomes H,S0 4 . H 2 S becomes first S and then H,S0 4 . H,SO. { becomes H,S0 4 ; and in general*S VI - n becomes S VI . 2HNO 3 + 3H 2 SO 3 = 2NO + 3H 2 SO 4 + H 2 O . 5. HC1, nitrohydrochloric acid: 2HN0 3 + 6HC1 = 2NO + 4H 2 + 3C1 2 (Koninck and Nihoul, Z. anorg., 1890, 477). See 269, 6#2. 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. G. 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 goncentratioii 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 acid be used. HI becomes first 1 ; then as above. 2HNO 3 + GHI , excess = 2NO + 3I 2 + H 2 O HI + HNO 3 , excess = 2NO + HIO 3 + H 2 O 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 oi-ganic bodies 288 NITRIC ACID. 241, 7. substitution products are formed, the oxides of nitrogen taking the place of the hydrogen. 7. Ignition. Nitric acid is dissociated by heat: 4HNO, = 4NO, + 2H,O + Oo , complete if at 256 (Carius, B., 1871, 4, 828). No nitrates are volatile as such. Ammonium nitrate is dissociated: NH 4 NO 3 = N 2 O + 2H 2 O. Some nitrates, e. g., those of K and Na. , are first changed to nitrites with evolution of oxygen onlv, and at an intense white heat further changed to oxides with evolution of N.,*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 , A.g , 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 Sniv , As"' to Asv f gb'" to Sbv , Fe" to Fe'" , Cr'" to Crvi , Mnvi-n to Mnvi , svi-n to Svi , etc. Heated 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 carbonaceous 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 u itric oxide, NO , which is set free from the mixture, oxidizing immediately in the air to nitrogen peroxide, N0 2 : 2KN0 3 -f 4H.SO, + 3Cu = K 2 S0 4 + 3CuSO 4 + 4H 2 + 2NO 2KN0 3 + 4H,S0 4 + GFeS0 4 = K 2 SO 4 + 3Fe a (SO 4 ), + 4H,O + 2ND The three atoms of oxygen furnished ly tiro molecules of nitrate suffice to oxidize three atoms of copper; so that 3CuO with 3HJ30 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 Wack-lrown liquid, (Fe30 4 ).,NO , decomposed by heat and otherwise un- stable: 2KN0 3 + 4H 2 S0 4 + 10FeS0 4 == K,S0 4 + 3Fe 2 (S0 4 ) s + 4H 2 + 2(FeS0 4 ) 2 NO . 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 CALIFORNIA COLlEfii 241, sd. NITRIC ACID. & PHARMAC 9 allowed to cool and a concentrated solution of FeSO^ then cautiously added to it, so that the fluids do not mix, the junction shows at first a purple, afterwards a brown color (Fresenius, QnaL Anal, IQtli ed., 387). A second method of obtaining the same brown ring is: Take sulphuric acid to a 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 ;in 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. Nitrous acid interferes with this test but the brown ring is produced when a nitrite is acidified with acetic acid while sulphuric acid is required in the case of a nitrate (239, 8). If the presence of a nitrite is suspected, the solution should first be acidi- fied with acetic acid and ferrous sulphate added. If a brown color is produced a nitrite is present. The nitrite may be removed by boiling the solution until the brown color has disappeared and does not return on adding more ferrous sulphate solution and acetic acid. After cool- ing the solution, a nitrate may be tested for by means of strong sul- phuric acid and ferrous sulphate. If strong oxidizing agents are present, excess of FeS(>4 should be added and the solution warmed. The solution should be cooled before applying the test. Metals forming insoluble sulphates should bo removed by adding Na C03 and warming. Iodides interfere by forming a brown ring of free iodine. 6. Indigo solution. In presence of HC1 heat moderately and the blue color is destroyed. Interfering substances, HC10 3 , HI0 3 , HBr0 3 , Fe'", Cr VI , Mn vn , and all that convert HC1 into C1 2 . c- Sodium salicylate is added to the solution, H 2 S0 4 is slowly added, the 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 OH . 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 -f 8KOH = 3NH S + 8KA10 2 + H 2 . Nothing interferes with this test, but action is delayed by Cl v , I v and many other oxidisers. Ammonia and ammonium salts must be removed by evaporation with before applying the test. 290 NITRIC ACID. 241, 80, e. Nitrite test. Reduce the nitrate to nitrite by warming with Al and KOH . At short intervals decant a portion of the solutioh, add a drop of KI, acidify with HC-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 v , and As v are reduced before the reduction of the HN0 3 begins: 3HN0 3 + 2AI + 5KOH = 3KN0 2 + 2KA1O 2 + 4H 2 O 2KN0 2 + 2KI + 4HC 2 H 1 2 = I 2 + 4KC 2 H 3 2 + 2H 2 O + 2ND 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 . /. Various organic compounds give characteristic color reactions with nitric acid. An excellent reagent of this kind consists of a solution of dimethyl aniline (2 drops) and p-toluidine (0.2 grams) in 50% sul- phuric acid (10 c.c.). A blood red color is produced when this reagent is brought into contact with even a dilute solution of a nitrate. The color appears as a ring test between the two liquids. Reducing agents especially ferrous salts interfere with the test. By the addition of a crystal of KC10 3 and HC1 and boiling, the interference may be overcome. KC10 3 gives a brown color but on dilution the red color of the nitrate generally appears. KC10 3 may be removed by heating with HC1. (Wood- ruff, J. Am. Soc. 19, 156. Schmidt and Lump (Ber. 43, 794) also give a color reagent which gives a red color. The reagent is a solu- tion of 0.1 gram of Di-9,10-monoxyphenanthrylamine in 1000 c.c. cone. H 2 S0 4 . The dry salt or the residue obtained by evaporation of the unknown solution is added to a few c.c. of the blue reagent which turns red in the presence of a nitrate. Oxidizing agents do not interfere. Add three drops of the solution to be tested to two drops of diphenylamine, (C 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 vl , Se 1 ^, and Fe'" interfere with this test. This test is of especial value in showing the absence of nitrates. If no color is obtained, it is certain that no nitrate is present. 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 6 H 5 OH , gives a deep red-brown color with nitric acid, by for- mation of nitrophenol (mono, di or tri), C 6 H 4 (NO 2 ;OH to C 6 H2(NO 2 ) 3 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, 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. Ths addition of potassium hydroxide d epens and brightens the c lor. According to Sprengel (J. C., 1863, 15, 396), the somewhat similar o l--r given by compounds of chlorine, bromine, iodine and by organic m ttter may be removed by adding ammonium hydroxide without diminishing the brightness of th color formed by the nitrates. 242, 3. OXYGEN. 291 i. According "to Lindo (C. N., 1888, 68, 176), resorcinal is five times more delicate a lest than phenol. T( n grammes of r< s;,rcinol are dissolved in 100 cc. of water; one dr

) 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 ;! . (tf) Neutralizing the free acid with ammonium hydroxide, and after evaporation and drying at 115, weighing as ammonium nitrate, (e) In presence of free H 2 S0 4 a ferrous solu- tion of known strength is added in excess to the nitrate and 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 n that has been formed. 242. Oxygen. = 16.000. Usual valence two. 1. Properties. A colorless, odorless gas; specific gravity, 1.10535 (Rayleigh, Proc. Roy. Soc., 1897, 204). When heated it diffuses through silver tubing quite rapidly (Troost, C. r., 1884, 98, 1427). It may be liquefied by cooling the gas under great pressure and then suddenly allowing it to expand under reduced pressure. It boils 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 118.8, and the critical pressure 50.8 atmos- pheres. Specific gravity of the liquid at - 181.4, 1.124 (Olszewski, M., 1887, 8, 73). Oxvgen is sparingly soluble in water with a slight increase in the vol- ume (Wmkler, 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, C. 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 oxygen, 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 metals 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 Chatelier, 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 88.81 per cent, it has been estimated that one-half of the crust is oxygen. Except in atmospheric air, which contains about 2;; per cent of uncombined oxygen, it is always found combined. 3. Formation. ((/) By igniting HgO . (b) By heating KC10 3 to 350, KC10 4 is produced and oxygen is evolved; at a higher temperature the KC1O 4 becomes KC1 . In the presence of MnO 2 the KC10 3 is completely changed to KC1 at 200, without forming KC10 4 , the Mn0 2 not being changed. Spongy platinum, CuO , Fe,O 3 , PbO 2 , etc.. may be substituted for Mn0 2 (Mills and Donald, J. C., 1882, 41, "l8; Baudrimont, AIII. N., 1S72, 103, 370). Spongy platinum, ruthenium, OXYGEN. 242, 3. rhodium and indium with chlorine water or with hydrogen peroxide evolve oxygen. The spongy ruthenium acts most energetically (Schoenbein, A. C/,., 1866, (4), 7, 103). (c) Action of heat on similar salts furnishes oxygen; e.g., KC10 and KC1O 2 form KC1 , KBr0 3 forms KBr , KI0 3 and KEO 4 "form KI ' and KN0 3 forms KN0 2 (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 i s changed to Cr'" , Co"' to Co" , Ni'" to Ni" , Biv to Bi'" , Fevi to Fe"' , Pbiv to Pb" , and Mn"+n to Mn"; in each case a sulphate is formed and oxygen given off: a. 2HgO (at 500) = 2Hg + O 2 &. lOKClOg (at 350) = 6KC1O 4 + 4KC1 + 3O 2 (Teed, J. C., 1887, 51, 283) 2KC10 3 (at red heat) = 2KC1 + :.O 3 2KC10 3 + nMn0 2 (at 200) = nMnO 2 + 2KC1 + 30 2 c. KC10 2 = KC1 + O 2 2KBr0 3 = 2KBr -f 3O 2 2KI0 8 = 2KI + 30 2 KIO 4 = KI + 20 2 2KN0 3 = 2KN0 2 + O 2 4KN0 2 (white heat) = 2K 2 O + 4NO + O 2 d. 2Pb 3 O 4 (white heat) = GPbO + 2 2Sb 2 O 5 (red heat) = 2Sb,0 4 + 2 Bi 2 5 (red heat) = Bi 2 O, + 2 4Cr0 3 (about 200) = 2Cr,O 3 + 30 2 4K 2 Cr 2 7 (red heat) 2Cr,,O 3 + 4K 2 Cr0 4 + 3O 2 6Fe 2 3 (white heat) = 4Fe 3 O 4 + O 2 3Mn0 2 (white heat) = Mn 3 O 4 + 2 6Co 2 O 3 (dull-red heat) = 4Co 3 O 4 -f 2 2Ni 2 3 (dull-red heat) = 4NiO -f O 2 2Ag 2 (300) = 4Ag -f 2 2BaO 2 (800) = 2BaO + O 2 e. 2K 2 Cr 2 7 + 8H 2 SO 4 = 4KCr(S0 4 ) 2 + 30 2 + 8H 2 O 4KMnO 4 + 6H 2 S0 4 = 2K 2 S0 4 + 4MnSO 4 + 50 2 + 6H 2 O 2Pb 3 O 4 -f 6H 2 S0 4 = 6PbSO 4 + 6H 2 + 2 4. Preparation. (a) By heating KC10 3 to 200 in closed retorts in the pres- ence of MnO 2 or Fe 2 3 . If KC1O 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 KC1O 3 should be "taken. (6) BaO heated in the air to 550 becomes BaO 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. (7., 1894, 66, ii, 89). (d) By passing sulphuric acid over red-hot bricks: 2H 2 S0 4 = 2S0 2 + 2H,0 -f O 2 ; the S0 2 is separated by water, and after conversion into H 2 SO 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 plaj r the same role as NO in making H 2 S0 4 (Fleitmaim, A. Ch., 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 2 O, and oxygen are formed. Then, without change of apparatus or temperature, air instead of steam is passed over the mixture of Mn 2 O 3 and NaOH . The Mn 2 0, is thus again oxidized to Na 2 Mn0 4 , and free nitrogen is liberated: 4Na 2 MnO 4 + 4H 2 (dull-red heat) = SNaOH + 2Mn 2 3 -f 30, 8NaOH + 2Mn 2 3 + air, 3(0 2 + 4N 2 ) = 4Na 2 MnO 4 + 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 liveol for three weeks, eating voraciously 243. OZONE. all the tim^, 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 1 , ultimately causing 1 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 Cl , 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; they must be formed by precipitation. 8. Detection. Uncombined oxygen is detected by its absorption by an alka- line solution of pyrogallol; by the combination \vith indigo white to form indigo blue; by its combination with colorless NO to form the brown NO 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 KC1O 3 , KMn0 4 , HgO , Au 2 O 3 , Pt0 2 , Ag.O , Sb,0 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 (CO, having been previously removed), and noting the dimunition in volume. Oxygen in combination is usually estimated by difference. 243. Ozone. :i = 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 (Poyg., 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, C. r., 1888, 107, 344; Wills, /?., 1873, 6, 769). By the oxidation of moist phosphorus at ordinary temperature (Leeds, A., 1879, 198, 30; Marig-nac, C. r., 1845, 20, 808). By electrolysis of dilute sulphuric acid, using lead electrodes Planti, C. r., 1866, 63, 181). By the action of concentrated sulphuric acid on potassium permanganate (Schoenbein, /. 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 1,000,000. In strong concentrations 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 con- cerning the physiological action, see Binz, C. C., 1873, 72. Its specific gravity 294 HYDROGEN PEROXIDE. 244. is 1.658 (Soret, A., 1866, 138, 4). It has been liquefied to a deep-blue liquid, boiling at - 106 (Olszewski, M., 1887, 8, 230). The gas is sparingly soluble in water (Carius, B., 1873, 6, 806). It decomposes somewhat into inactive oxygen at ordinary temperature, 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 PbO 2 Sn" becomes Sn0 2 Hg 1 ' becomes Hg" Bi'" becomes BL0 5 Pd" becomes Pd0 2 Cr'" becomes Crvi Fe" becomes Fe,O 3 ; in presence of KOH , K,Fe0 4 Mn" becomes MnO,; in presence of H 2 SO 4 or HNO 3 , HMnO 4 is formed. Co" becomes Co"' Ni" becomes Ni'" . With the salts of nickel and cobalt the action is slow, rapid with the moist hydroxides. K 4 Fe(CN) 6 becomes K 3 Fe(CN) 6 N 2 3 becomes HN0 3 , in absence of water N0 2 is formed SO 2 becomes H 2 SO 4 H,S becomes S and H,O , the sulphur is then oxidized to H 2 SO 4 (Pollacci, C. C., 1884, 484) P and PH 3 become H 3 P0 4 HC1 becomes Cl and H,O HBr becomes Br and ELO I becomes HIO 3 and HI0 4 (Ogier, (7. 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 HN0 2 and many other substances give the same reaction, thallium hydroxide paper is preferred by Si-hoene (/?., 1880, 13, 1508). The paper is colored brown, but the reaction is much less delicate than with potassium iodide starch paper. Chlorine, bromine, iodine and nitrous oxides do not interfere with the following test. Paper is moistened with a. 15 per cent solution of potassium iodide to which a 1 per cent alcoholic rosolic acid or phenolphthalein solution has been added until a marked opalescenco is produced. A red color is produced by exposure to ozone. The color produced by the rosolic acid is more permanent than the phenolphthalein color. It i ? estimated quantitatively by passing the gas through a neutral solution of KI and titration of the liberated iodine: O 3 + 2KI + H 2 O = O 2 + I 2 + 2KOH . Ozone may also be detected by the blue color imparted to guaiacum tincture. Hydrogen peroxide does not produce this effect unless ferrous sulphate solution is added. The most sensitive reagent is a freshly prepared 10 per cent solution of guaiacum gum in 50 per cent water solution of chloralhydrate. (Weber, Z., 43, 47). 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 $244, 6, 10. HYDROGEN PEROXIDE. 295 reacts strongly acid to litmus. The ordinary three per cent solution ean he evaporated on the water bath until it contains about HO per eent H 2 O 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 (Wolffenstein, J?., 1894, 27, iJiiOT). 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, C. 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. (or) By the electrolysis of 70 per cent H,SO t (Richarz, W. A., 1887, 31, 912). O) By the action of ozone upon ether and water (Berthelot, C. 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. BaO 2 is decomposed by dilute H 2 SO, , the BaSO t being removed by filtration. The BaO, is obtained by heating BaO in air or oxygen to low redness. At a higher heat the Ba.O., is decomposed into BaO and O (Thomsen, B., 1874, 7, 73). Sodium peroxide, Na,0, , is formed by heating sodium in air or oxygen (Harcourt, J. C., 1862, 14, 267); by adding H 2 O 2 to NaOH 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 CO, and H,SiF R with formation of H 2 O, . Na. 2 O 2 is soluble in water with generation of much heat. It is a powerful oxidizing agent. 0. 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 strong 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 . 4. Au.,0., becomes An 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). 296 HYDROGEN PEROXIDE. 244, 6, 11. 11. Cr'" becomes Cr VI in alkaline mixture (Lenssen, J.pr. y 1860, 81, 278). 12. Cr vl 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. 9 1896, 29, 1940). 13. Mn" in alkaline mixture becomes Mn0 2 . In presence of KCN a separation from Zn (Jannasch and Niederhofheim, B., 1891, 24, 3915; 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., 1855, 7, 304; Lunge, Z. angew., 1890, 6). 14. BaO , SrO , and CaO become the peroxides. 15. NaOH becomes Na 2 2 .8H 2 . 73. NH 4 OH becomes NH 4 N0 2 (Weith and Webber, B., 1874, 7, 17 and 45). 17. Ti IV is oxidized to pertitanic acid, H 2 Ti0 4 , with the production of ;i yellow color which constitutes a very delicate test for H 2 2 . B. With non-metals and their compounds. 1. K 4 Fe(CN) 6 becomes K 3 Fe(CN) f , (Weltzien, A., 1866, 138, 129); in alkaline solution the reverse action takes place: 2K 3 Fe(CN) 6 -f- 2KOH + 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). 3. H 3 P0 2 becomes H 3 P0 4 . 4. 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. Cl becomes HC1 (Schoene, /. c., page 254). It is a valuable reagent for the estimation of chloride of lime : 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 , KBr , 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; Kingzett, J. C., 1880, 37, 805). 7. Ingition. The peroxide of barium is formed by igniting- BaO to dull red- ness; strong- ig-nition causes decomposition of the BaO 2 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 or ferrous sulphate, a blue color is obtained when H 2 O2 is added. To the solution supposed to contain I^Oa add a few drops of lead acetate; then KI , starch, and a little acetic acid; with H 2 O 2 a blue color is produced (Schoen- bein, I. c.; Struve, Z., 1869, 8, 274). As confirmatory, its action on KMnO 4 and on K 2 Cr 2 O 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 O 2 (Deniges, C. r., 1890, 110, 1007; Crismer, BL, 1891, (3), 6, 22). H 2 O 2 gives some extremely delicate color tests with the analine bases (Ilosyay, B., 1895, 28, 2029; Deniges, /. Pharm., 1892, (5), 25, 591 ; Weber, Z., 43, 47). Titanium sulphate, Ti(SO 4 ) 2 , in acid solution gives an extremely delicate test for H 2 O 2 , a yellow color being produced. On the addition of caustic alkali a yellowish orange precipitate is produced which redissolves in excess of the reagent. 245, 2. FLUORINE. 297 9. Estimation. (a) By measuring the amount of oxygen liberated with MnO (Hanriott, BL, 1885, (2), 43, 468). (ft) "By the amount of standard KMnO 4 reduced, or by measuring the volume of oxygen set free: 2KMnO 4 -f 3H 2 SO 4 + 5H 2 O 2 = K 2 SO 4 + 2MnSO 4 + 8H 2 O -f- 5O 2 . (c) By decomposition of KI in presence of an excess of dilute H 2 SO 4 ; and titration of the liberated iodine with standard Na 2 S 2 O 3 . (t/) Dissolve a weighed sample of BaO 2 in dilute HC1 , add K 3 Fe(CN) 6 ; transfer to an azotometer and add KOH . The volume of oxygen is a measure of the amount of H 2 O 2 (Baumann, I. c.). 245. Fluorine. F = 19.00. Valence one. Since Davy's experiments in 1813, many others have attempted the isolation of fluorine. In his zeal the unfortunate Lonyet fell a victim to the poisonous fumes which he inhaled. Faraday, (lore, 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, anct "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, C. r., 103, 202 and 256; J. C., 50, 1886, 849 and 976; A. Ch., 1891, (6), 24, 224). 1. Properties. A gas of light greenish-yellow color and strong pungent odor; Specific gravity, 1.313 (Moissan, C. R., 109). When cooled to a temperature of about - 187 it condenses to a mobile yellow liquid. Specific gravity of this liquid is 1.14. At 223 C., fluorine solidifies to a pale yellow solid. The solid loses its yellow color and becomes perfectly white at 252 (Moissan and Dewar, C. R., 1903, 136). 2. Occurrence. Fluorine does not occur free in nature, but is found in con- siderable quantities in combination with calcium in the mineral fluorspar, CaF 2 . Its other fairly common compounds are cryolite, Na 3 AlF 6 and apatite, Ca 5 (PO 4 ) s F . 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, forming SiF 4 , silicon fluoride, a gaseous compound, generally prepared by the_ action of concentrated sulphuric acid on calcium fluoride and silicic anhydride. (ft). 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 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 6 (d) (Offer- mann, Z. angeiv., 1890, 617). This acid is used as a reagent; forming metallic fluosilicates (silicofluorides), soluble and insoluble (246). 298 HYDROFLUORIC ACIDFLUOSILICIC ACID. 246. a. CaF 2 + H 2 SO 4 = CaSO 4 + 2HF 6. 2CaF 2 + SiO 2 + 2H 2 SO 4 = 2CaSO 4 + 2H 2 O + SiF 4 c. SiF 4 + 4H 2 O = Si(OH) 4 + 4HF (not remaining free) d. 2HF + SiF 4 H 2 SiF 6 c and d. 3SiF 4 + 4H 2 O = Si(OH) 4 + 2H 2 SiF 6 246. Hydrofluoric acid. HF = 20.008. 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, , 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. . 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, &); a gas which does not attack glass, but when passed into water produces fluosilicic acid, H 2 SiF 6 (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., + :2KHS0 4 = CaS0 4 + K 2 SO 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 yellowish-green color is imparted to the flame. 247. Fluosilicic acid. H 2 SiF 6 = 144.316. Fluosilicic acid * (hydrofluosilicic acid), (HF) 2 SiF 4 , or H 2 SiF 6 , is soluble in water and forms metallic fluosilicates (silico fluorides] , 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. * 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 sand 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. 249, 4. SILICON SILICON DIOXIDE. 299 Potassium fluosilicate is precipitated translucent and gelatinous. Ammonium hydroxide precipitates silicic acid with formation of ammonium fluoride. With roncentra ed sulphuric acid, they disengage hydrofluoric acid, HF. By heat, they are resolved into fluorides and silicon fluoride: BaSiF 6 = BaF 2 + SiF 4 . 248. Silicon. Si = 28.3. Valence four (15). There are three modifications of silicon: () Amorphous. A dark brown powder; specific :;:;). In the fused bead of microcosm ic salt particles of silica swim iiiHliHxnlri'd. If a silicate be taken, its base will, in most cases, be dissolved out, leaving a " skclclini <>f silica" un- dissolved in the liquid bead. But with a head 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 1 the latter in fine powder with three parts of precipitated calcium carbonate, and one-half part of ammonium chloride, and healing in a platinum crucible to redness for half an hour, avoiding 1 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 SiO., 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 -f- 3H = H 2 SiO, + 2H 2 SiF (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 = 3i.04. 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 44.2 1.814 (Damien, 1881). At ordinary temperatures it is brittle and easily pulverized. At 44.1 (Burgess, Wash. Acad. of Sci., 1-18) it melts, but may be cooled in some instances (under an alkaline liquid) as low as +4 with- out 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 mole- cule to P 2 (Meyer and Biltz, B., 1889, 22, 725). Specific gravity of the red amor- phous modification at 0, 2.18 (Jolibois, C. r., 149, 287-289). 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 1 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, G. N., 1879, 39, 144). Heated in absence of air above the boiling point it sublimes as a. colorless gas, depositing lustrous transparent crystals (Blondlot, C. r., 1866, 63, 397). At low temperatures phosphorus oxidizes slowly in the air with a characteristic odor, probably due to the formation of ozone and phosphorous oxide, P 2 O 3 (Thorpe and Tutton, ,7. C., 1890, 57, 573). It ignites spontaneously in the air at 44.5, burning with a bright yellowish white light producing much heat. From the finely divided state, as from the evaporation 302 PHOSPHORUS. 250, 2. 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 HoS , SO, , CS 2 , Br , Cl , 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 ( ly -talline form, action beginning at 260. If ordinary red phosphorus is heated to 4oJ for sixty hours, in absence of air, it becomes red pyromorphic phosphorus wlio.se density is 2.37. Heated to 725 this phosphorus melts, and if suddenly cooled becomes a violet-colored mass (Jolibois, C. r., 149, 287-289). Heated in the air from 250 to 260 it takes fire (Schroetter, I. c.). A black crystalline metallic variety of phosphorus is described by Hittorf (/. c.); also Remsen and Kaiser (Am., 1882, <% 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 (Rossel and Frank, B.. 1894, 27, 52). Red 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 ELSO 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, SiO 2 , with the charcoal is preferred, in which case the whole of the phosphorus is reduced (45). The vapor density of the gaseous oxide indicates the molecule to be P 4 6 . tiitccific yntvUy 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 O 4 and P (T. and T., J. C., 1891, 59, 1019). It reacts slowly with cold \vater, forming H 3 PO 3 ; with hot water the reaction is violent and PH 3 is evolved. Upon exposure to the air it oxidizes to P^O G . The acid, H 3 P0 3 , is a crystalline solid, very deliquescent, melting at 74 (Hurtzig and Geuther, A., 1859, 111, 171). It is a dibasic arid, forming no tribasic salts (Amat, (\ 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 (Railton, J. C., 1855, 7, 216; Michaelis, J. 0., 1875, 28, 1160). Neither meta nor pyro- phosphorous acids are known, but a number of pyrophosphites have been pre- pared (Amat, C. 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 O 3 is formed together with P 2 5 when phosphorus is ignited in the air. H 3 PO 3 is formed together with H 3 P0 4 when phosphorus is oxidized with HNO 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: 3CuS0 4 + P 4 + 6ELO = Cu 3 P 2 + 2H 3 PO 3 + 3H 2 SO 4 (Schiff, A., 1860, 114, 200). 4. Preparation To prepare phosphorous anhydride, P,O, , phosphorus is burned in a tube with an insufficient supply of air (Thorpe and Tutton, I.e.). The acid, H 3 P0 3 , is prepared by dissolving the anhydride in cold water; by decomposing' PC1 3 with water (Hurtzig and Geuther, I. c.}. r>. Solubilities. The acid is miscible in water in all proportions. Alkali phosphites are soluble in water, most others are insoluble (distinction from hypophosphites) . 6. Reactions. 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 becomes Hg-Cl and then Hg , CuCL becomes CuCl then Cu (Rammelsberg, J. C., 187;t, 255, 1. H7POPHOSPHORIC ACID PHOSPHORIC ACID. 307 26, 13). Concentrated H 2 S0 4 with heat forms H,P0 4 and S0 2 (Adie, J. C., 1891, 59, 230). H,S0 3 forms H 2 S and H 3 P0 4 (Woehler, A., 1841, 39, 252). Nascent hydrogen (Zn and H,S0 4 ) produce PH 3 (Dusart, C. r., 1856, 43, 11 :>(>). 7. Ignition. The acid is decomposed by ignition, forming HPO ;i and P or PH, (Yigier, Bl, 1869, (2), 11, 125; Hurtzig and Geuther, 7. r.). Phosphites Jin- decomposed by heat, leaving a pyrophosphate and a phosphide and evolving PH 3 or H (Rammelsberg 1 , /?., 1876,9, 1577; and Kraut, A., 1875, 177, 274). 8. Detection. By oxidation to H 3 P0 4 and detection as such. Tt is distin- guished from hypophosphorons acid by reducing CuSO t to Cu, while the latter forms Cu 2 H,.; 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 3 PO 4 and estimation as such. 254. Hypophosphoric acid. H 4 P 2 6 = 162.112. II II H' 4 P IV 2 0-" 6 ,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, 1 1 1 and 271); also by oxidizing phosphorus with dilute HN0 3 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 3 PO 3 and HPO 3 , and at 120 gives H 4 P 2 7 and PH 3 (Joly, C. r., 1886, 102, 110 and 760). It is oxidized to H 3 PO 4 by warm HNO, , slowly by KMn0 4 in the cold, rapidly when heated. It is not oxidized by H 2 2 , chlorine water or H 2 CrO 4 ; HgCl, becomes HgCl (Amat, C. r., 1890, 111, 676). It is not reduced by Zn and H 2 SO, (distinction from H 3 P0 2 and H 3 P0 3 ). With a solution of silver nitrate it gives a white precipitate which does not blacken in the light (distinction from H 3 P0 2 and H 3 P0 3 ). 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 = 98.064 . II H' 3 P v O-" 4 , H P H. I I H 1. Properties. Phosphoric anhydride, P 2 O 5 *, is a white, flakey, very delique- scent solid, fusible, subliming undecomposed at a red heat. It is very soluble in water, forming three varieties of phosphoric acid: ortlto, H 3 P0 4 ; meta, HPO^; * According to Tildeii and Barnett (J". C., 1896, 69, 154) the molecule is P 4 O 10 not P 2 O 5 ; P 4 O a not P U O 3 1 Thorpe and Tutton, J. C., 1891, 59, 1032) ; ana P 4 S 6 not P a S 3 (Isambert, C.r., 1886, 1O, 1386). 308 PHOSPHORIC ACID. 255, 2. and pyro, H 4 P 2 O T . Orthophosphoric acid is a translucent, feebly crystallizable and very deliquescent soft solid. Specific t/niriti/, 1.88 (Schiff, A., 1860, 113, 183); melting point, 41.75 (Berthelot, BL, 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 PO 4 , primary, monobasic or mono- metallic phosphates: M' 2 HPO 4 , secondary, dibasic or dimetallic phosphates; and M' 3 P0 4 , tertiary, tribasic, trimetallic or normal phosphates. The first two are acid salts, but Na 2 H?0 4 is alkaline to test paper. Metaphosphoric acid, HPO 3 , H O P O , is a white waxy solid, volatile at a red heat II 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. pliys. Ch., 1890, 6, 122). O II II Pyrophosphoric acid, H 4 P 2 O 7 , H O P P O H,isa glass-like I l O 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 3 PO 4 . Heated to redness HPO 3 is formed. It forms two classes of salts: M'oH->PoO 7 and M' 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 a nil yd rule, P.O.-, , is formed by burning phosphorus in great excess of air; also by burning phosphorus in NO , N0 2 , or C10 2 . Orthophosphoric acid, H 3 P0 4 , is formed by long exposure of phosphorus to moist air, or by oxidation with HN0 3 ; by oxidation of H S PO 2 or H 3 PO S with the halogens, HN0 3 , HC10 3 , etc.; by treating P 2 r> , HP0 3 , or H 4 P 2 O T with boiling water; by combustion of PH 3 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 w r hen 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 () can be rerifiecl, that is, pure phosphoric acid can be separated; but the changes shown in equations (ft) and (c) do not comprise the whole of the material taken. In the operation (&) 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,0 4 + 2H 3 P0 4 6. Na 2 HP0 4 + 2HNO 3 = 2NaNO 8 + H 3 PO 4 and Na 2 HPO 4 -f HNO 3 = NaNO, + NaH 2 PO 4 c. 2CaHP0 4 + 2HC1 = CaCL + CaH 4 (P0 4 ) 2 Metaphosphoric odd is formed by treating P 2 0., with cold water; by decom- position of lead metaphosphate with H 2 S or of the barium salt with H 2 SO 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,P 2 O 7 , with H 2 S or of the' corresponding barium salt with H 2 S0 4 ; or by heating H 3 PO 4 to a little above 200 until no yellow silver phosphate, Ag 1 : ,PO l , is obtained on dissolving- in water and Ireatment with silver nitrate after neutralization with NH 4 OH . 255, 5. i'HOMi>noitir ACID. 309 4. Preparation. To prepare P 2 5 , phosphorus is burned in a slow cur- rent of dry oiygen healing to about 300, then in a more rapid current of the gas, and iinally the P,0- is distilled in an atmosphere of oxygen (Shenstone, \Vnll*' Die., 1894, IV, 141). H 3 P0 4 is prepared by warming phosphorus, one part, with nitric acid, */;. 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 HNO.j 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 NaPO, 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 T ; 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,H,P 2 7 may be prepared by titrating a sat- urated solution of Na 4 P 2 T 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. 9 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 oxychloride. All phosphates are dissolved or transposed by HNO., , HC1 , or H 2 S0 4 , and all are dissolved by HC 2 H,0 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, :n<> PHOSPHORIC ACID. 255, 6^4. except silver, are soluble in solutions of alkali pyrophosphates, as double pyro phosphates soluble in water (distinction from orthophosphates). Ferric iron as a double pyre-phosphate loses the characteristic properties of that metal (Persoz, J. C., 1849, 1, 183). Phosphates are insoluble in alcohol. 0. Reactions. A. With metals and their compounds. Phosphoric acid dis- solves some rnetals, 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- CO 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 b} r 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 1 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, and ferric chloride; no precipitates with barium or calcium chloride, or w r ith 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 in neutral solution form silver orthophosphate, Ag 3 P0 4 , yellow ; with metaphosphates, silver m eta- phosphate, AgP0 3 , white ; and with pyrophosphates, silver pyrophosphate, Ag 4 P 2 7 , white, all soluble in ammonium hydroxide. Silver metaphos- phate 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 (6) means of distinguishing the acid phosphates from the normal. (a) Na 2 HP0 4 + 3AgN0 3 = Ag 3 PO 4 + 2NaNO 3 + HNO 3 (6) Na 3 PO 4 + 3AgNO 3 = Ag 3 PO 4 + 3NaNO 3 Free orthophosphoric acid forms no precipitate with reagent silver nitrate, because silver phosphate is soluble in dilute HNO 3 . With lead acetate or nitrate, Na 2 HP0 4 forms Pb a P0 4 , white, insoluble in acetic acid, as are also the phosphates of aluminum and ferricum. With * A solution containing 5 p. c. ferric chloride, mixed with one-fourth its volume of a 10 p. c. soluti'-n of orthophosphoric acid, requires that near half of the latter be 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. c. solution of meta- phosphoric acid, form a precipitate, to dissolve which, <20 cc. of the same m; taphosphoric acid solution c r 5 cc. of a 24 p. c. solution of hydrochloric acid are required. Four cc. of a 5 p. c. solution of silvern itrate 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. K r atschmer and Sztankovansky, Z M 1883, 21, 520.] 255, 6.4. PHOSPHORIC ACID. 311 PbCL, the precipitate always contains a chloride. Free phosphoric arid. H,P0 4 , 1'onns an acid phosphate, PbHP0 4 (Heintz, Pogg., 1848, 73, lli)). 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 HNO, . 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. 5 (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. K 3 P0 4 + 12(NH 4 ) 2 MoO 4 + 21HNO 3 = (NH 4 ) 3 PO 4 .12MoO3 + 21NH,NO 3 + 12H 2 O . 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. 312 PHOSPHORIC ACID. 255, 6Z?. 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 9 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 2 Mg 3 , recognized by odor of PH 3 , 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 O 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. Bimetallic orthophosphates, M' 2 HP0 4 , are changed to normal pyrophosphates upon ignition; also tribasic orthophosphates when one-third of the base is volatile, e. #., MgNH ( PO 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,O 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 , Fe , Cu , Co , Ni , Mn , Grl and IT 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 1 , 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,0 7 and AgPO, 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 T (("listing- 256, 3. SULPHUR. 313 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 HNO. { in H,P0 4 . Sodium metaphosphate does not give a precipitate with ZnS0 4 cold and in excess; with Na 4 P 2 0- and Na 2 H,P,0 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. (I)) By precipitation and weighing as lead phosphate, Pb 3 (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 = 32.06 . Usual valence two, four and six (14). 1. Properties. Sulphur is a solid, in yellow, brittle, friable masses (from melt inn); 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, ./. pr., 1852, 55, 101). The si>cH/i<- gravity of native sulphur is 2.0748 (Pisati, B., 1874, 7, 361). Melting point rhombic 114.5. Boiling point, 444.53 (Callendar and Griffiths, C. N., 1891, 6 , 2). Vapor density at 1160 is 34, indicating that the molecule is S 2 (Bineau, C. r., 1859, 49, 799); but at lower temperatures the molecule seems to vary from 2 to ; 8 . Sulphur is polymorphous, existing in various crystalline forms, rhombic, mono- clinic and triclinic systems, and also in amorphous conditions. It is also classified by the relative solubilities of the various forms in carbon disulphide. In chemical activity, volatility and other properties it stands as the second member of the Oxygen Series: O, 16,000; S, 32.07; Se, 79.2; and Te, 127.5. On being heated it melts at 114.5 to a pale yellow liquid; as the temperature rises it grows darker and thicker, until at about 180 it is nearly solid, so that the dish may be inverted without spilling. At 260 it again becomes a liquid as at first; and at 444.53 it boils and is converted into a brownish-red vapor. If it is slowly cooled, exactly the same physical changes take place in the reverse order, becoming thick at 180 and thin again at 114.5, and at lower temperatures solid. If, at a temperature near its boiling point, it is poured into cold water, it forms a soft, ductile, elastic string, resembling india-rubber. In a few hours this ductile sulphur changes back to the ordinary form, the change evolving heat. But if poured into water from the ether liquid form that is, at 114.5 it forms only ordinary, brittle sulphur. In contact with air sulphur ignites at 248 (Hill, C. N., 1890, 61, 125); burning in air or oxygen with a pale blue flame and penetrating odor to SO 2 . The isolated oxides of sulphur are S0 2 , S0 3 , S 2 O 8 and S 2 O 7 . Sulphur and oxygen combine directly to form SO 2 and SO 3 ; the former by burning sulphur in oxygen, the latter by the action of ozone upon SO,; also by burning sulphur with oxygen under several atmospheres pressure. S 2 3 is made by dissolving sulphur^in sulphur dioxide; S 2 O 7 by the action of the electric discharge upon a mixture of SO 3 and O . 2. Occurrence. (a) Found in a free state, and as SO, in volcanic districts. (&) As H 2 S in some mineral springs, (c) As a sulphide: iron pyrites, FeS_, ; copper pyrites, CuFeS 2 ; orpiment, As 2 S 3 ; realgar, As.,S 2 ; zinc blende, ZnS; cinnabar, HgS; galena, PbS. (d) As a sulphate: gypsum, CaSO 4 .2ELO; heavy spar, BaSO 4 ; kieserite, MgS0 4 ,H 2 O; bitter spar (Epsom salts), MgSO 4 ,7H 2 O; Glauber salt, Ta 2 S0 4 ,10H 2 O , etc. 3. Formation. (a) By decomposing polysulphides with HC1 (Schmidt, Phar- tnnwutiscJic Chortle, 1898, 175). (&) By adding an acid to a solution of a thio- sulphate. (e) By the reaction between SO 2 and H 2 S: 2S0 2 + 4H 2 S = 3S 2 -f 4H 2 O . (d) By the decomposition of metallic sulphides with nitric acid; 2Bi 3 S, + 16HN0 3 = 4Bi(N0 3 ), + 3S 2 + 4NO + 8H 2 O . 314 suLPHrR. 256, 4. 4. Preparation. (a) The native sulphur is separated from the clay and rock in which it is embedded, partly by melting- and partly by distillation. (&) From FeS 2 by heating- in close cylinders 3FeS 2 = Fe 3 S 4 + S 2 ; or at a higher temperature: 2FeS 2 = 2FeS + S 2 . Much of the sulphur contained in pyrite* is converted into and utilized as sulphuric acid. 5. Solubilities. Ordinary (not precipitated) sulphur is soluble in carbon di- sulphide; the ductile variety is insoluble. There are several allotropic forms of sulphur. Samples of commercial sulphur are almost never found which are entirely soluble or insoluble in carbon disulphide. Forms of sulphur insoluble in CS 2 are changed to soluble forms upon heating to the melting point; also amorphous sulphur insoluble in CS 2 (formed by adding acids to thiosul- phates or SO 2 to H 2 S) is changed to the soluble form by mixing with a solution of H 2 S in water. It dissolves readily in hot solutions of the hydroxides of potassium, sodium, calcium or barium, forming polysulphides and thiosul- phates: 3Ca(OH) 2 + 6S 2 = 2CaS 5 + CaS 2 O 3 + 3H 2 O . These can be separated by alcohol, in which the sulphides dissolve. These products are also readily decomposed by acids with separation of sulphur (method of preparation of precipitated sulphur). Precipitated sulphur (in analysis, HC1 upon (NH 4 ) 2 S X ) is soluble in benzol or low boiling petroleum ether; of value in analysis for the removal of the sulphur to detect the presence of traces of As or Sb sulphides (Fresenius, Z., 1894, 33, 573). 6. Reactions. A. With metals and their compounds.. Sulphur does not combine with metals without the aid of heat (see 7), except that under very great pressure (6500 atmospheres) it combines with Pb , Sn , Sb , Bi , Cu , Cd , Fe , Zn , and Mg (Spring, B., 1883, 16, 999). Flowers of sulphur boiled with SnCL gives SnS and SnCl 4 ; with HgN0 3 almost exactly one-half of the mercury is precipitated as HgS . No action with sulphates of Cd , Fe", Mn", Ni and Zn ; with acid solutions of SbCl 3 and BiCl 3 ; or with solutions of As v and As'" (Vortmann and Padberg, B., 1889, 22, 2642). Sulphur boiled with hydroxides of K , Na , NH 4 , Ba , Ca, Sr, Mg, Co, Ni, Mn, Hg", Bi, Cu', Cu", Cd , Pb , Ag , and Hg' forms sulphides and thiosulphates; also some sulphates are formed. No action with hydroxides of Fe, Zn and Sn (Senderens, BL, 1891, (3), 6, 800). B. With non-metals and their compounds. 1. HCN warmed with sulphur or a polysulphide becomes a thiocyanate: 2KCN + S 2 = 2KCNS or 4HCN + 2(NH 4 ) 2 S 4 = 4NH 4 CNS + 2H 2 S + S 2 . 2. HN0 3 becomes NO and H 2 S0 4 . Strong acid and long continued boiling are necessary to the complete oxidation of the sulphur. The crystallized variety is attacked with much greater difficulty than the amorphous or flowers (Saint-Gilles, A. Ch., 1858, (3), 54, 49). 3. Red phosphorus combines readily at ordinary temperature, forming P 2 S 3 or P 2 S 5 , depending upon the relative amounts of the elements used. Ordinary phosphorus combines explosively. See 252, 6. Tribasic sodium or potassium phosphate when boiled with sulphur forms alkali polysul- phide and thiosulphate, changing the phosphate to dibasic phosphate (Filhol and Senderenj, (7, r,, 1883, 96, 1051), 257, 1. HYDROSL'Li'in me A<:I/>. 315 4. H 2 S0 4 , concentrated and hot, becomes S0 2 from both the S and the H 2 S0 4 : 4H 2 S0 4 + S 2 == 6S0 2 + 4H 2 . 803" when added to S at 12 forms the blue hyposulpburous anhydride, S 2 3 (not the anhydride of thiosulphuric acid, S 2 2 ). S0 2 reacts with S even at ordinary tempera- tures, forming thiosulphuric acid and tri or tetrathionic acid (Colefax, J. C., 1892, 61, 199). 5. Cl in presence of water forms HC1 and H 2 S0 4 . HC10 3 becomes HC1 and H 2 S0 4 . 6. Br in presence of water becomes HBr and H 2 S0 4 . HBr0 3 becomes HBr and H 2 S0 4 . 7. Sulphur does not appear to have any action upon iodine or upon iodine compounds. 7. Ignition. In the air, at ordinary temperatures, finely divided sulphur is very slightly oxidized, by ozone, to sulphuric acid; at 248 it begins to oxidize rapidly to sulphurous anhydride, burning with a blue flame. Sulphur, when fused with the following elements, combines with them to form sulphides: Pb , Ag , Hg , Sn , As, Sb , Bi , Cu , Cd , Zn , Co, Ni , Fe , Sr, Ca, Mg, K, Na , In, Tl , Pt , Pd , Rh , Ir , Li, Ce , La, Ne , Pr . Svi n becomes Svi when fused with alkaline carbonate and nitrate or chlorate. That is, free sulphur, S , or any compound containing sulphur with valence less than six, is oxidized to a sulphate if fused with an alkaline nitrate or chlorate, nitric oxide or a chloride being formed and carbon dioxide escaping. 8. Detection. (a) By burning in the air to a gas having the odor of burning matches. (&) By its solubility in CS 2 . (c) By formation of H 2 S0 4 with oxidizing agents, (d) By the formation of sulphides upon fusion with metals, (e) By the blackening of silver coin after boiling with alkali hydroxide, (f) Formation of reddish-purple with sodium nitroferricyanlde after boiling with alkali hydroxide, (g) In organic compounds by heating with Na and testing the Na 2 S with sodium nitro- ferricyanide (Vohl, B.', 1876, 9, 875). 9. Estimation. Sulphur is usually estimated by oxidation to a sul- phate and weighing as BaS0 4 . 257. Hydrosulphuric acid. H 2 S = 34.076. H' 2 S-", H S H . 1. Properties. Molecular weight, 34.070. Vapor density, 17. Boil in g point, 61.8. Freezing point, 85.56. Under a pressure of 14.G atmospheres it be- comes a liquid at 11.11 (Faraday, A., 1845, 56, 156). It is a colorless poisonous gas. It burns readily, forming sulphur dioxide and water: 2H 2 S + SO,, = 2SO 2 + 2H 2 . The aqueous solution slowly decomposes upon exposure to the air with separation of sulphur. The gas is readily expelled from its aqueous solution by boiling; slowly when exposed at ordinary temperature. Both the gas and the water solutions have a feebly acid reaction towards moist litmus paper. They also possess a strong characteristic odor, resembling that of rotten eggs. In acid or in alkaline solutions it is a strong reducing agent. See 6. HYDROSULPHURIC ACID. 257, 2. 2. Occurrence. Found free in volcanic gases and frequently in mineral springs. While the inhaled gas is poisonous, the mineral waters containing it are reputed to be a healthful beverage. 3. Formation of Hydrosulphuric Acid. (a) By direct union of the elements when passed over pumice stone heated to 400 (Corenwinder, A. Cli^ 1852, (3), 33:, 77). (6) Heating paraffin or tallow with sulphur (Fletcher, C. X., 1879, 40, 154); and by passing illuminating gas through boiling sulphur (Taylor, ('. A., 1883, 47, 145). (c) The sulphur in coal becomes H 2 S in the process of gas- making, (d) From steam and sulphur at 440. (e) Often occurs in nature from reduction of gypsum by decaj'ing organic matter (Myers, J. pr., 1869, 108, 123). (f) Transposition of sulphides by hydracids or by dilute phosphoric or dilute sulphuric acid, (y) Decomposition of organic compounds containing sulphur. Formation of Sulphides. (1) By fusion of the metals with sulphur, see 256, 7. (2) By action of H 2 S upon the free metals, hydrogen being evolved. With Hg- and Ag this occurs at ordinary temperature, but with most metals a higher temperature is needed. (3) Action of BLS on metallic oxides or hydroxides. Those sulphides which are decomposed bj- water (e. (/., A1,S 3 , Cr,S 3 ) are not formed in its presence, but by action of H 2 S upon the oxide at a red heat. (4) By action of soluble sulphides upon metallic solutions. The ordinary sulphides of the first four groups are formed thus, except ferric salts, which are precipitated as FeS , and aluminum and chromic salts as hydroxides. (5) By action of CS 2 upon oxides at a red heat. (6) By action of free sulphur upon oxides at a red heat. (7) By the action of charcoal upon the oxyacids of sulphur at a red heat in presence of an alkaline carbonate. To prepare a sulphide absolutely arsenic free, take BaSO 4 , 100 grams; coal, pulverized, 25 grams; and NaCl, 20 grams, mix, ram into a clay crucible and ignite to a white heat for several hours (Winkler, Z., 1888, 27, 26). (8) By the action of zinc amalgam on sulphuric acid (Walz, C. A'., 1871 23, 245). (9) As a reagent for the formation of metallic sulphides in analysis it is recommended by Schiff and Tarugi (#., 1894, 27, 3437), Schiff (., 1895, 28, 1204), and Tarugi (Gazzetta, 1895, 25, i, 269), to use ammonium thioacetate, CH 3 COSNH 4 ; prepared by distilling a mixture of phosphorus pentasulphide and glacial acetic aid (300 grams each) with 150 grains of cracked glass. A large distilling flask is used and the distillate is collected to 103. It is then dissolved in a slight excess of ammonium hydroxide, diluting to three volumes from one volume of the acid. Salts of the metals of the first two groups in acid solution are readily precipitated as sulphides upon warming with this reagent. 1. 2Fe + S 2 = 2FeS 2. 2Ag + H 2 S = Ag 2 S + H 2 5. Pb(OH) 2 + H 2 S = PbS + 2H C O 4Fe(OH) 3 + 6H 2 S = 4FeS + S 2 + 12H 2 4. 4FeCl 3 + 6(NH 4 ) 2 S = 4FeS + S 2 + 12NH 4 C1 5. 2CaO + CS 2 2CaS + CO 2 6. 4CaO + 3S 2 4CaS + 2S0 2 7. K 2 S0 4 + 20 = K 2 S + 2C0 2 4. Preparation. For laboratory purposes it is nearly always made by adding H 2 S0 4 or HC1 to FeS . The ferrous sulphide is prepared either by fusion of the iron with the sulphur, or by bringing red hot iron rods in contact with sticks of sulphur,, and is made to drop into tubs of cold water. Dilute H,S0 4 should be used:* FeS + H,S0 4 == FeS0 4 + HJS . Concentrated H 2 S0 4 has no action on FeS , unless heated and then S0 2 is evolved: 2FeS + 10H 2 S0 4 = Fe 2 (S0 4 ) 3 + 9S0 2 + 10H 2 ; and frequently free sulphur is formed by the action of the H 2 S upon the S0 2 first formed. * If the acid is diluted with eleven volumes of water ferrous sulphate crystals will not be deposited. 257, 5. HYDROSULPHURIC ACID. 317 The colorless ammonium sulphide, (NH 4 ) 2 S, is prepared by saturating ammonium hydroxide with H 2 S until a sample will no longer give a pre- cipitate with a solution of magnesium sulphate; showing that ammonium hydroxide is no longer present. Upon standing the solution gradually becomes yellow with formation of the poly sulphides or yellow ammonium sulphide, (NH 4 ) 2 S X This may be hastened by the addition of sulphur (Bloxam, J. C., 1895, 67, 277). Sodium sulphide, Na 2 S , is prepared by neutralizing an alcoholic solution of NaOH with H 2 S and then adding an equal amount of NaOH and allowing to crystallize ; air being excluded. The various polysulphides, Na 2 S 2 to Na 2 S 5 , arc prepared by boiling the normal sulphide with the calculated amounts of sulphur (Boettger, A., 1884, 223, 335; Geuther, A., 1884, 224, 201). 5. Solubilities. At 15 water dissolves 2.66 volumes of the gas H 2 S . Sulphides which dissolve in dilute H 2 S0 4 evolve H 2 S , e. g., CdS , FeS , MnS , ZnS , etc. But if a sulphide requires concentrated H 2 S0 4 for its solution; S and S0 2 are formed or S0 2 alone; e. g., Bi 2 S 3 , CuS, HgS . If concentrated H 2 S0 4 be used upon a sulphide that might have been dis- solved in the dilute acid, then no H 2 S is evolved: ZnS + 4H 2 S0 4 = ZnS0 4 + 4S0 2 + 4H,0 . Or with a small amount of water present: 2ZnS + 4H 2 S0 4 = 2ZnS0 4 + S 2 + 2S0 2 + 4H 2 . The sulphur of the zinc sul- phide is oxidized to free sulphur and that of the sulphuric acid is reduced to sulphur dioxide. HgS is almost insoluble in HN0 3 , dilute or concen- trated, readily soluble in chlorine, nitrohydrochloric acid, or chloric acid if kot. Most other sulphides are soluble in hot HN0 3 (74, 6e). Long continued boiling with water more or less completely decomposes the sul- phides of Ag , As , Sb , Sn , Fe , Co , Ni , and Mn ; no effect with sulphides of Hg, Au, Pt, Mo, Cu, Cd, and Zn (Clermont and Frommel, A. Ch. y 1879, (5), 18, 203). As a reagent, hydrosulphuric acid, gas or solution in water finds ex- tended application in the analytical laboratory. The grouping of the bases for analysis depends very largely upon the relative solubilities of the sulphides. Hydrosulphuric acid in alkaline solution, alkali sulphide or polysulphide, is a scarcely less important reagent., being especially valuable in the subdivision of the metals of the second group. The sulphides of the first four groups are insoluble. Hydrosulphuric acid transposes salts of the first two groups in acid, neutral, and alkaline mixtures, except arsenic, which is generally imperfectly precipitated un- less some free acid or salt that is not alkaline to litmus be present. The result is a sulphide, but mercurosum forms mercuric sulphide and mer- cury, and arsenic acid may form 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 318 HYDROSULPHURIC ACID. 257, 6. solutions of their normal salts traces of cobalt, nickel, manganese, and zinc (135, 60) are precipitated. 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 sul- phide and sulphur. The precipitates have strongly marked colors that of zinc being white; manganese, flesh colored; those of iron, copper, and lead, black; arsenic stannic and cadmium, yellow; antimony, orange-red; stannous, brown; mer- cury, successively white, yellow, orange, and black. 6. Reactions. A. With metals and their compounds, Some metals are converted into sulphides on being treated with hydrosulphuric acid; e. g., Ag , Cu , Hg , etc. The alkali polysulphides slowly attack many metals with formation of sulphides: Sn becomes M' 2 SnS 3 ; Ag becomes Ag 2 S, no action with colorlesc (NH 4 ) 2 S ; Ni forms NiS ; Fe , FeS; Cu, CuS and then Cu 2 S (with colorless ammonium sulphide, (NH 4 ) 2 S , Cu 2 S is formed with evolution of hydrogen) (Priwozink, A., 1872, 164, 46). The hydroxides or non-ignited oxides of Pb", Ag , Hg", Sb , Sn , Bi'", Cu, Cd, Fe", Co", Ni", Mn", Zn, Ba, Sr , Ca, Mg, K, Na, and NH 4 unite with moist H 2 S at ordinary temperature to form sulphides without change of the valence of the metal. In other cases the valence of the metal is changed, usually with liberation of sulphur. 1. Pb"+ n becomes PbS and S . 2. As v in acid solution forms some As 2 S 3 and S . See 69, 6e. 3. Hg' becomes HgS and Hg . 4. Cr vl becomes Cr'" and S , if the H 2 S be in excess : 2K,Cr 2 7 + 8H 2 S = 4Cr(OH) 3 -f 3S 2 + 2K 2 S + 2H,0 . 5. Fe"' becomes Fe" and S : 4FeCl 3 + 2H 2 S =: 4FeCl 2 + 4HC1 + S 2 . If the solution be alkaline FeS is precipitated : 4FeCl 3 -f 6K..S = 4FeS + 12KC1 + S 2 . 6. Co"+ n becomes Co" and S . 7. Ni"+ n becomes Ni" and S . 8. Mn" +n becomes Mn" and S. In alkaline solution with excess of KMn0 4 , an alkali sulphate is formed and Mn0 2 : 8KMn0 4 -f 3K 2 S = 3K 2 S0 4 + 4K 2 + 8Mn0 2 (Schlagdenhafen, Bl., 1874, (2), 22, 16). In the above reactions, if an alkaline sulphide be used instead of hydro- sulphuric acid, the metal will be precipitated as a sulphide with the * The normal fixed alkali sulphides (Na 2 S, K 2 S), precipitate solutions of calcium and mag- nesium salts as the hydroxides : Ca(C 2 H 3 O 2 ^ + 2Na 2 S + 2BT 2 O = Ca(OH) 3 -f 2XaC 2 H 3 O 3 + 2NaHS. No reaction with the acid fixed alkali sulphides (NaHS, KHS) or with ammonium sulphides (Pelouze, A. C/i M 1866, (4), 7, 172). 257, 7. HYDROSULPHURIC ACID. 319 formation of an alkali hydroxide; except that the arsenic will remain in solution (69, 5c) and the chromium will be precipitated as the hydroxide. Dry H 2 S has no action on the dry salts of Pb , Ag , Hg , As , Sb , Sn , Bi , Cu , Cd , or Co ; nor does it redden dry blue litmus (Hughes, Phil. Mag., 1892, (5), 33, 471). Many insoluble sulphides, freshly precipitated, transpose the solutions of other metallic salts. In some cases the action is quite rapid at ordinary tem- perature, in others long-continued heating 1 (several hours) at 100 is necessary. PdS is formed by action of PdCL with sulphides of all the metals following in the series below named, but PdS is not transposed by solutions of the metals following-. Silver salts form Ag 2 S with sulphides of the metals following in the series but not with sulphides of Pd and Hg , etc.; Pd , Hg , Ag , Cu , Bi , Cd , Sb , Sn , Pb , Zn , Ni , Co , Fe , As , Tl and Mn (Schiirmann, A., 1888, 249, 326). B. With non-metals and their compounds. 1. H 3 Fe(CN) 6 becomes H 4 Fe(CN) 6 and S. Proof: Boil to expel the excess of hydrosulphuric acid, then add ferric chloride (126,' 66). 2. HN0 3 becomes NO and S . If the HN0 3 be hot and concentrated the sulphur is oxidized to sulphuric acid. 3. H 2 S has no reducing action on the acids of phosphorus. 4- H 2 S0 3 becomes pentathionic acid, H 2 S 5 6 , and sulphur: 10H 2 S0 3 -j- 10H,S == 2H 2 S 5 6 + 5S 2 + 18H 2 . With excess of H 2 S the product is entirely free sulphur from both compounds: 2H 2 S0 3 + 4H 2 S = 3S 2 + 6H 2 (Debus, J. (7., 1888, 53, 282). H 2 S0 4 , dilute no action; concentrated and hot, S and S0 2 are formed: 2H 2 S0 4 + 2H 2 S = S 2 + 2S0 2 + 4H 2 ($256, 6J54). 5. Cl with H 2 S in excess forms HC1 and S ; with Cl in excess forms HC1 and H,S0 4 . HC10 3 with H 2 S in excess forms HC1 and S ; with HC10 3 in excess HC1 and H 2 S0 4 . 6. Br with H..S in excess forms HBr and S ; with Br in excess HBr and H 2 S0 4 . HBrO, with H 2 S in excess forms HBr and S ; with HBr0 3 in excess HBr and H,S0 4 . 7. I becomes HI and S (Filhol and Mellies, A. Ch., 1871, (4), 22, 58). HIO ;! becomes HI and S . 7. Ignition. Dry hydrosulphuric acid gas is not decomposed when heated to 350 to 360. At this temperature AsH 3 in presence of potassium polysulphide, K 2 S 3 , liver of sulphur, is decomposed: 2AsH 3 + 3K 2 S 3 = 2K,AsS 3 + 3H 2 S; thus furnishing a ready means of purifying- H 2 S for lexicological work (69, 6'6) (Pfordten, B., 1884, 17, 2897). If air be excluded some sulphides may be sublimed unchanged; c. f/., HgS , As.,S 8 , As 2 S 6 , Sb,S 8 , etc. In some cases part of the sulphur is separated, leaving a sulphide of a lower metallic valence: 2FeS 2 = 2FeS + S 2 . Some sulphides remain unchanged upon ignition in absence of air; e. g., FeS , MnS , CdS , etc. All sulphides suffer some change on being ignited in the air; some slowly, others rapidly; Sb 2 S 8 , CuS , A1 2 S 3 , Cr,S 3 , etc., evolve SO, and leave 20 HYDROSULPHURIC AC'lh. 257, 8. the oxide of the metal; HgS , Ag-,8 , etc., evolve S0 2 and leave the free metal. All sulphides, as well as all other compounds of sulphur, when fused with KN0 3 or KC10 3 in presence of an alkali carbonate are oxidized to an alkali sulphate; forming- NO or KC1 and evolving CO, . The metal is changed to the carbonate, oxide or the free metal (228, 7). When ignited on charcoal with sodium carbonate or (distinction from sulphates) if ignited in a porcelain crucible with sodium carbonate soluble sodium sulphides are obtained. The production of the sodium sulphide is proved by the black stain of Ag 2 S , formed on metallic silver by a moistened portion of the fused mass. (Compounds of selenium and tellurium, 112 and 113.) 8. Detection. (a) The odor of the gas constitutes a delicate and char- acteristic test when not mixed with other gases having a strong odor. (&) The gas blackens filter paper moistened with a solution of lead ace- tate, delicate and characteristic. In the detection of traces of the gas, a slip of bibulous paper, so moistened, may be inserted into a slit in the smaller end of a cork, which is fitted to the test-tube, wherein the material to be tested is treated with sulphuric acid; the tube being set aside in a warm place for several hours. If any oxidizing agents are present as chromates, ferric salts, manganic salts, chlorates, etc. hydrosulphuric acid is not generated, but instead sulphur is separated, or sulphates are formed (6). (c) The gas blackens silver nitrate solution, delicate but PH.{ , AsH 3 , and SbH 3 also blacken silver nitrate solution, (d) By its reducing action upon nearly all oxidizing agents with separation of sul- phur, which is detected according to 256, 8. KMn0 4 is perhaps the most delicate test but the least characteristic, (e) Its oxidation to a sulphate is characteristic in absence of other sulphur compounds. This method is usually employed with sulphides not transposed by dilute H 2 S0 4 ; chlorine, nitrohydrochloric acid or bromine being the usual oxidizing agents. Also, these sulphides and certain supersulphides, attacked with difficulty by acids, as iron pyrites and copper pyrites, are reduced and dissolved, with evolution of Jiydrosulphuric add, by dilute sulphuric acid with zinc. The gas, with its excess of hydrogen, may be tested by method (/). (/) Sodium nitroferricyanide Na 2 [Fe(CN) 5 (NO)].2H 2 also known as sodium nitroprusside gives a very delicate and characteristic test for H 2 S as an alkali sulphide. The gas is passed into ammonium hydroxide; and to this mixture a 20 per cent solution of the reagent is added, pro- ducing a transient reddish-purple color. Free H 2 S , dilute, remains colorless; a concentrated solution gives a blue color, due to the reducing action of the H 2 S on the ferricyanide. Caustic alkali hinders the reac- tion. 0.000018 grain of H 2 S , as gas or alkali sulphide, can be detected by this reagent (Reichard, Z., 43, 222). (g) The most delicate test for hy- drogen sulphide involves the formation of methylene blue (E. Fisher, Ber., 16, 2234). By this test 0.02 mg. of hydrogen sulphide in a liter will give a blue color after standing half an hour. The test is carried out as follows : to the solution to be tested is added one-tenth of its volume 258, 5. THIOSULPHURIC ACID. 321 of concentrated HC1, then a small amount of dimethylparaphenylendiamine sulphate (NH 2 .C H 4 .N(CH 3 ) 2 .H 2 S0 4 ) and after it' has dissolved a drop or two of dilute ferric chloride solution. For method of separation of the various sulphur compounds from each other consult Kynaston (J. C., 1859, 11, 166), Bloxam (C. N., 1895, 72, 63), Votocek (Ber., 1907, 40, 414) and Autenrieth and Windaus (Z., 1898, 295). 9. Estimation. Sulphides are usually oxidized to H 2 S0 4 (by chlorine, bromine, or nitrohydrochloric acid, or by fusion with KN0 3 and Na 2 C0 3 ) precipitated with BaCl 2 and weighed as BaS0 4 . 258. Thiosulphuric acid. H 2 S 2 3 = 114.136. Difhionous acid. II H' 2 (S 2 ) IV 0-" 3 , H S S H.* 1. Properties. Thiosulphuric acid, H 2 S 2 O 3 (formerly called hyposulphurous acid), has not been isolated; but it almost certainly exists in dilute solutions, when a dilute weak acid is added to a solution of sodium thiosulphate, Na 2 S 2 3 , soon beginning to decompose into H 2 S0 3 and S (Landolt, B., 1883, 16, 2985). The thiosulphates are not particularly stable compounds, some decomposing almost immediately upon forming; e. g., mercury thiosulphates. Alkali thio- sulphates decompose upon heating into sulphate arid polysulphide: 4Na 2 S 2 O 3 == 3Na,S0 4 + NsLjSs . Other salts give also S and H 2 S . Boiling solution of a thiosulphate gives a sulphate and H 2 S or a sulphide of the metal. 2. Occurrence. Not found in nature. 3. Formation. Thiosulphates are formed by the oxidation of alkali or alkaline earth polysulphides by exposure to the air or by SO 2 or K 2 Cr,,0 7 : 2CaS 5 + 30 2 = 2CaS,0 3 + 3S 2 ; 4Na 2 S 5 + 6S0 2 = 4Na,S 2 O 3 + 9S 2 ; 2K,S 5 + 4K 2 Cr 2 7 + 13H 2 O = 5K 2 S 2 O 3 + 8Cr(OH) 3 + 2KOH (Doepping, A., 1843, 46, 172; Gueront, C. r., 1872, 75, 1276). Also by heating ammonium sulphate with phosphorus pentasulphide (Spring, /?., 1874, 7, 1157). 4. Preparation. Thiosulphates are prepared by boiling sulphur in a solu- tion of normal alkali sulphite: 2Na 2 SO 3 + S 2 = 2Na,S 2 O 3 . Fixed alkali or alkaline earth hydroxides with sulphur also form thiosulphates: 3Ca(OH)o + 6S 2 = 2CaS 5 + 'CaS 2 O 3 + 8H,O (Filhol and Senderens, C. r., 1883, 96, 839; Senderens, C. ?\, 1887, 104, 58). Commercial sodium thiosulphate is prepared by passing SO, into " soda waste " suspended in water, calcium thiosulphate being formed. This is treated with sodium sulphate, filtered and evaporated to crystallization. 5. Solubilities. The larger number of the thiosulphates are soluble in water; those of barium, lead and silver being only very sparingly soluble. The thio- sulphates are insoluble in alcohol. They are decomposed, but not fully dis- solved, \>y acids, the decomposition leaving a residue of sulphur. * Bunte, B., 1874, 7, 646. T'HIOSVLPH'URIC ACID. 258, (5. Alkali thiosulphate solutions dissolve the thiosulphates of lead and silver; also the chloride, bromide and iodide of silver, and mercurous chloride; th.6 iodide and sulphate of lead; the sulphate of calcium, and some other precipi* tates by formation of soluble double thiosulpliatcs: Ag 2 S 2 3 + Na 2 S 2 3 = 2NaAgS 2 0. AgCl + Na 2 S 2 3 = NaAgS 2 3 + NaCl PbS0 4 + 3Na 2 S 2 3 =Na 4 Pb(S 2 3 ) 3 + Na 2 S0 4 6. Reactions. A. "With metals and their compounds. With soluble thio- sulphates, solutions of lead and silver salts are precipitated as thiosulphates, white, soluble in excess of alkali thiosulphate. These precipitates decompose upon standing, rapidly on warming, into sulphides and sulphuric acid: Ag.S.Oj -f- H,O = AgoS -f H.,S0 4 . Soluble mercury salts with sodium thiosulphate fornTa white precipitate, almost instantly turning black with decomposition to mercuric sulphide. Na 2 S,O 3 blackens HgCl , a portion of the mercury going into solution, colorless, reprecipitated black upon warming. Acid solutions of arsenic and antimony are precipitated by hot solution of Na 2 S 2 O 3 as sulphides, As 2 S 3 and Sb 2 S 3 (a separation from tin,* which is not precipitated) (6e, 69, 70 and 71). Solutions of copper salts with thiosul- phates, on long standing, precipitate cuprous salt, changed by boiling to cuprous sulphide and sulphuric acid (separation from cadmium, 78, Ge). Solutions of ferric salts are reduced to ferrous salts with formation of sodium tetrathionate: 2FeCl 3 + 2Na 2 S 2 O 3 = SFeCL + 2NaCl + Na 2 S 4 O 8 ; used as a quantitative method of estimation, with a few drops of potassium thiocyanate as an indicator. Chromic acid (chromates in acid solution) are reduced to chromic salts with oxidation of the thiosulphate. Permanganates in neutral solution become manganese dioxide, in acid solu- tion the reduction is complete to manganous salt, a sulphate and dithionate being formed (Luckow, Z., 1893, 32, 53). Barium chloride forms a white precipitate of barium thiosulphate, BaS 2 3 , nearly insoluble in water; 100 parts of water dissolve 0.2675 part of BaS 2 O 3 H 2 O at 17.5. Calcium chloride forms no precipitate (distinction from a sulphite). B. With non-metals and their compounds. When thiosulphates are decom- posed by acids, the constituents of thiosulphuric acid are dissociated as sul- phurous acid and sulphur. Nearly all acids in this way decompose thiosul- phates: 2Na 2 S 2 O 3 + 4HC1 = 4NaCl + 2H 2 S0 3 -f S, . Thiosulphates are reducing agents even stronger and more active than the sulphites to which they are so easily converted. This reduction is illustrated by the action on arsenic compounds, on ferric salts and on chromates and permanganates as given above. Also the halogens are reduced to the halide salts forming a tetrathionate: 2Na 2 S 2 3 + I 2 = 2NaI + Na 2 S 4 O 6 . If chlorine or bromine be in excess the tetrathionate is further oxidized to a sulphate: Na 2 S 2 3 + 4CL -f 5H 2 = Na 2 S0 4 + H 2 S0 4 + 8HC1 . Chloric, bromic and iodic acids are first reduced to the corresponding halogens and then with an excess of the thiosulphate to the halides, always accompanied with the separa- tion of sulphur. Nitric acid is reduced to nitric oxide w T ith the separation of sulphur. 7. Ignition. On ignition, or by heat short of ignition, all thiosulphates are decomposed. Those of the alkali metals leave sulphates and poly sulphides (a), others yield sulphurous acid with sulphides, or sulphates, "or both. The capacity of thiosulphates for rapid oxidation, renders their mixture with chlorates, nitrates, etc., explosive, in the dry way. Chlorates with thiosulphates explode violently in the mortar. Cyanides and ferricyanides, fused with thin- sulphates, form thiocyanates, which may be dissolved by alcohol from other products. By fusion on charcoal with Na,CO 3 , thiosulphates form sulphides (&) and (c); and by fusion with an alkali carbonate and nitrate or chlorate, * According to Vortmann (M., 1886, 7, 418) sodium thiosulphate may be used instead of hydro- sulphuric acid in the second group of bases. An excess of the reagent is to be avoided and nitric acid should be absent. 259. HYPOSULPHUROUS ACID. 323 a sulphate is formed (d). By ignition of a metallic salt with Na^Os in a dry test-tube the characteristic colored sulphide of the metal is obtained (Landauer, B., 1872, 5, 406). (a) 4Na 2 S 2 3 = Na 2 S 5 + 3Na 2 S0 4 (&) Na 2 S 2 3 + ETa 2 C0 3 + 2C = 2Na 2 S + 3C0 2 (c) 2PbS 2 3 + 4Na 2 C0 3 + 5C = 4Na 2 S + 2Pb + 9C0 2 (d) 3Na 2 S 2 3 + 3Na,C0 3 + 4KC10 3 = 6Na 2 S0 4 -f 4KC1 + 3C0 2 8. Detection. In analysis, thiosulphates are distinguished by giving a pre- cipitate of sulphur with evolution of sulphurous anhydride when their solu- tions are treated with hydrochloric acid; by their intense reducing power, shown in the blackening of the silver precipitate; and by non-precipitation of calcium salts. The precipitation of sulphur with evolution of sulphurous anhydride, by addition of dilute acids as hydrochloric or acetic is characteristic of thiosulphates. It will be understood, however, that in presence of oxidizing agents, which can be brought into action by the acid, sulphides w r ill likewise give a precipitate of sulphur. In the presence of a sulphate and sulphite the thiosulphate is detected as follows: Add BaCl 2 and NH 4 C1 in excess, then HC1 to solution of all but the BaSO.i . Filter and treat the filtrate with iodine, forming BaSO 4 of the sulphite and BaSiO,, of the thiosulphate. Filter and add bromine to the filtrate, which then forms BaS0 4 (Smith, C. N., 1895, 72, 39). Sulphides, sulphites and thiosulphates may be separated as follows: Add to the neutral solution cadmium carbonate, shake and filter off the cadmium sulphide and excess of cadmium carbonate. Test the filtrate for sulphide with sodium nitro- prusside and again add cadmium carbonate until the sulphide is entirely removed. Add strontium nitrate to the filtrate, allow to stand overnight and filter off the strontium sulphite. Test the precipitate for sulphurous acid. Test the filtrate for thiosulphuric acid by acidifying with HC1 and warming. (Autenrieth and Windaus, Z., 1898, 295.) 9. Estimation. By titration with a standard solution of iodine, or by titrating the iodine liberated by a standard solution of potassium dichromate (125, 10, and 279, 657). 259. Hyposulphurous acid, H 2 S0 2 = 66.076. (Hydro sulphurous or diihionous acid.) H' 2 S"-0-" 2 , H S H. II Obtained by Schiitzenberger (G. r., 1869, 69, 196) by the action of zinc on sulphurous acid: Zn + 2SO 2 + H 2 O ZnSO + H 2 SO 2 . The sodium salt is formed by treating a concentrated solution of sodium acid sulphite with zinc filings: Zn + 3NaHSO 3 = ZnSO 3 + Na 2 SO 3 + NaHSO 2 + H 2 O . In the forma- tion of the free acid or of the sodium salt no hydrogen is evolved. It is a very unstable compound, a strong reducing agent, rapidly absorbs oxygen from the air, becoming sulphurous acid or a sulphite. According to Bernthsen (B., 1881, 14, 438) the sodium salt does not contain hvdrogen. He gives the formula as Na 2 S 2 4 : Zn + 4NaHSO 3 = ZnS0 3 + Na 2 SO 3 + Na 2 S 2 O 4 + 2H 2 O . It is used in the preparing of indigo white for the printing of cotton fabrics. See also Dupre, J. C., 1867, 20, 291, 324 DITH IONIC ACIDTRITHIONIC ACID. 260. 260. Dithionic acid. H 2 S,0 6 = 162.136. II II H' 2 (S 2 ) x O-" 6 , H S S H . II II Known only in the form of its salts and as a solution of the acid in water. The free acid or the anhydride has not been prepared. The manganous salt is prepared by the action of a solution of sulphurous acid upon manganese dioxide at a low temperature: Mn0 2 + 2H 2 S0 3 = MnS,O 6 + 2H 2 . Similar results are obtained with nickelic or ferric oxides (Spring and Bourgeois, BL, 1886, 46, 151). The acid is obtained by treating the manganous salt with Ba(OH) 2 and the nitrate from this with the calculated amount of H 2 S0 4 . It is a colorless solution and may be evaporated in a vacuum until it has a specific gravity of 1.347. It decomposes upon further heating: H 2 S 2 O 6 = H 2 S0 4 -f SO, . All other thionic compounds decompose upon heating witli separation of sulphur . By exposure to the air dithionic acid is oxidized to sulphuric acid. All dithionates are soluble in water and may be purified by evaporation and crystallization (Gelis, A. Ch., 1862, (3), 65, 230). Dithionic acid is also prepared by carefully adding a potassium iodide solu- tion of iodine to sodium acid sulphite (Hoist and Otto, Arch. Phann., 1801, 229, 171); Spring and Bourgeois (Arch. Pharm., 1891, 229, 707) contradict the above statement. 261. Trithionic acid. H 2 S 3 6 = 194. 196 . II II H' 2 (S 3 ) 10 0- 2 6 , H S S S H. II II The free acid and anhydride are not known. The potassium salt is prepared by boiling potassium acid-sulphite with sulphur (a) ; by treating potassium thiosulphate with sulphurous acid (&) (no action with sodium thiosulphate) (Baker, C. N., 1877, 36, 203; Villiers, C. r., 1889, 108, 402); by the action of iodine on a mixture of sodium sulphite and thiosulphate (c) (Spring, B. t 1874, 7, 1157): (a) 12KHS0 3 + S 2 = 4K 2 S 3 6 + 2K 2 S0 3 + 6H 2 O (6) 4K 2 S 2 3 + 6S0 2 = 4K,S 3 6 + S 2 (c) Na 2 S0 3 + Na 2 S 2 3 + I 2 = Na 2 S 3 6 + 2NaI The acid is prepared by adding perchloric or fluosilicic acid to the potassium salt. The acid is quite unstable; at low temperature in a vacuum it decom- poses into S0 2 , S and H 2 SO 4 . The salts are quite stable; they are not oxidized by chloric or iodic acids, while the free acid is rapidly oxidized by these acids. Fixed alkalis or sodium amalgam change the trithionate to sulphite and thio- sulphate (Spring, I.e.). 263. TETRATHIONIC ACIDPENTATHIONIC ACID. 325 262. Tetrathionic acid. H 2 S,0 6 = 226.256. II II H' 2 (S 4 ) 10 0- 2 6 , H S S S S H. II II The salts are soluble in water and are comparatively stable. They are best obtained in crystalline form by adding" alcohol to their solutions in water. The acid has not been isolated but it is much more stable than the tri or pentathionic acids. In dilute solution it can be boiled without decomposition. The concentrated solution decomposes into H 2 S0 4 , S0 2 and S . Tetrathionates are prepared by adding iodine to the thiosulphates: 2BaS 2 3 + I 2 = BaS 4 Q + BaI 2 (Maumene, C. r., 1879, 89, 422). The lead salt is obtained by the oxidation of lead thiosulphate by lead peroxide in presence of sulphuric acid: 2PbS 2 O 3 + PbO 2 + 2H 2 S0 4 = PbS 4 6 + 2PbSO 4 + 2H 2 (Chancel and Diacon, <7. pr., 1863, 90, 55). To obtain the acid the lead should be removed by the necessary amount of sulphuric acid, and not by hydrosulphuric acid, which causes the formation of some pentathionic acid. A number of other oxidizing- agents may be used to form the tetrathionate from the thiosulphate (Fordos and Gelis, C. r., 1842, 15, 920). Sodium amalgam reconverts the tetra- thionate into the thiosulphate: Na 2 S 4 Os + 2Na = 2Na 2 S 2 3 (Lewes, /. (7., 1880, 39, 68; 1881, 41, 300). Tetrathionic acid is also formed with pentathionic acid in the reactions between solutions of H 2 S and SO 2 (Wackenroder's solution, A., 1846, 60, 189). See also Curtius and Henkel (J. #r., 1888, (2), 37, 137). The acid gives no precipitate of sulphur when treated with potassium hydroxide (distinction from pentathionic acid). 263. Pentathionic acid. H 2 S 5 6 = 258.316 . II II H' 2 (S 5 ) 10 0- 2 6 , H S S S S S H. II II Only known in the salts and in the solution of the acid in water. It is formed by the action of H 2 S upon S0 2 in the presence of water (a) ; by the action of water on sulphur chloride (6); by the decomposition of lead thiosulphate with H 2 S (Persoz, Pogg., 1865, 124, 257) : a. 10H 2 S0 3 + 10H 2 S = 2H 2 S 5 6 + 5S 2 + 18H 2 O 6. 10S 2 C1 2 + 12H 2 = 2H 2 S 5 6 + 5S 2 + 20HC1 The filtrate from the decomposition of SO, by H 2 S is known as Wackenroder's solution (Arch. P/mrw., 1826, 48, 140). It has been shown to contain the tri and tetrathionic acids in addition to the pentathionic acid (Debus, C. N., 1888, 57, 87). Pentathionic acid may be concentrated in a vacuum until it has a specific gravity of 1.6; farther concentration or boiling heat alone decomposes it into H 2 SO 4 , SO 2 and S . The solution of the acid does not bleach indigo. When treated with a fixed alkali hydroxide an immediate precipitate of sulphur is obtained (distinction from H.,S 4 O ): 4H-,S,O 6 + 20NaOH = <>Na 2 SO 3 + 4NaoS.,O 3 + 3S, + 14H,O (Takamatsu and Smith, J. f'., 1880, 37, 592); or if the NaOH be added short of neutralization: 10H.,S : .O 6 + 20NaOH = 10Na 2 S 4 O. + 5S 2 + 20H 2 . Neutralization of pentathionic acid with barium carbonate gives barium tetrathionate and sulphur (Takamatsu and Smith, J. C., 1882, 41, 162; Lewes, J, C., 1881, 39, 68), See also Spring, A., 1879, 199, 97. TABLE OF THIONIC ACIDS. 264. *~ ^> ^s o o 1 4 cc "11 a* 1 SS "^ 5 s| e :- f il 1 w I ^ a 'S -B ^ -S ,0 '-p 1 Pentathionic acid, H a i:| || tg| |3 | 1^1 13$ 53 g *| fH iif |1| if i a !t| Kl l-l il. i l l i| N| %! * !i:l ! I I 111 HI - a P. c o> o -P -S as OTJoJ 45 c ? -P 5 fc <1 t* <1 < gradually turning b: heating, with evolu HCN. On warming, whitish precipitate. Whito iirpcirtitato of sul Decolori/ed. without i of dilute of H^ssCV 1 |l i i s 2 > s- | O "3 53 M T ."S ? <.& .& = i as B" 60 if! l i |. I | 3 d" _g" o?J3 o C-S 2 ; 1 -P oii fe S ^ -Prrt'S. O ^ ^ O "t- ga 1 3 ^ * etrathioni< I 8 tf 1 5L ! S 'I ^1 ; ^1 lg| | ^ o"S o ^2 S *- ^ -1! i 1 35 ' s |1 ^ 1 '23 y> y> II EH O* 'S'O "3 S 53 ^ O -t- 3^ - ^ S "o fc fc KH ^'" fc < ^ p * ^i S J be . C *-* 2^ 4 1 Us | |1 o s* n C x *, T3 -2 ,Q ** ^5 B i ^S s 1 -'o '3 M <3? '^ * o |3 | 2^ li I" 5 OB ^ 'S S P^ >1 o ^ 4 .2 I * ** | 1 ^S Trithioc ! fi 1! P Hi g. +3 S O C Jjo ^2 5 f-t a 3ts Bf S a "o >-P fl-Q OQD'-M fe S h 5? fe ^5 32 0^ 0' 1 O : IB ffl * |J "O i - - i- '. .- ,-4 ; ' i . -S 3 1 5:51 1 1 o B S -H -^ . ' 8 '3 '3 ^ H .2 S i S 2 & 2 & J3 P, 03 P, P< ft o 2 S O O C ^ o '3 4 P is '?, 'fl > , 3 ! 1 z. 1 si bo ^3 S ^ > ^J $/ "t^ S s d " "p a a >> .So 2 -a g ' e . -C i 1 _3 i ft M 5 E 3 fl-P ~ 3 -22 --i 1 I i T | 3 s tS - 265, 4. SULPHUROUS ACID. 327 265. Sulphurous anhydride. S0 2 = 64.06. Sulphurous acid. H 2 S0 3 = 82.076. II S IV 0-" 2 and H' 2 S IV 0-" S ,0 S = and H S H. 1. Properties. Sulphurous anhydride, S0 2 , sulphur dioxide, is a colorless gas of a strong suffocating' odor of burning sulphur. Specific gravity of the liquid at 0, 1.4338 (Cailletet and Matthias, C. r., 1887, 104, 1563) ; of the gas at and 760 mm. pressure, 2.2369 (Leduc, C. r., 1893, 117, 219). It is liquefied at atmos- white wooly solid. Cooled to 76.1 it becomes a snow-white solid (Faraday, C. r., 1861, 53, 846). The dry gas is not combustible in the air, does not react acid to litmus, but in presence of water it has a marked acid reaction. The gas and the free acid, not the salts, are quite poisonous, due to the absorption of the S0 2 by the blood and oxidation to H 2 S0 4 . The gas is soluble in water, form- ing probably sulphurous acid, H 2 SO 3 . The pure acid has not been isolated, but forms salts mono and dibasic as if derived from such an acid (Michaelis and Wagner, B., 1874, 7, 1073). It has a strong odor from vaporization of sulphurous anhydride, which is soon completely expelled upon boiling. The acid oxidizes slowly in the air, forming H 2 SO 4 , hence sulphurous acid usually gives reactions for sulphuric acid. Light seems to play an important part in this oxidation (Loew, Am. 8., 1870, 99, 368). The moist gas or a solution of the acid is a strong bleaching agent, however not acting alike in all cases. Wool, silk, feathers, sponge, etc., are permanently bleached: also many vegetable sub- stances, straw, wood, etc.; yellow colors and chlorophyll are not bleached; red roses are temporarily bleached, immersion in dilute H 2 S0 4 restoring the color. 2. Occurrence. Found free in volcanic gases (Ricciardi, B., 1887, 20, 464). 3. Formation. (a) By burning sulphur in air. (ft) By heating sulphur with various metallic oxides, (c) By decomposition of thiosulphates with HC1. (rf) By burning H 2 S or CS 2 in air. (e) By the action of hot concentrated sulphuric acid on metals, carbon, sulphur, etc. (f) By heating sulphur with sulphates. (fir) By decomposition of sulphites with acids: (a) S 2 + 20 2 = 2S0 2 (&) Mn0 2 + S 2 = MnS + SO 2 2Pb 3 4 + 5S 2 6PbS + 4S0 2 (c) 2Na 2 S 2 3 -f 4HC1 = 4NaCl + 2S0 2 + S 2 + 2H 2 (d) 2H 2 S + 30 2 = 2S0 2 + 2H 2 CS 2 + 3O 2 = 2SO, + CO, (e) Cu + 2H 2 S0 4 = CuS0 4 + S0 2 -f 2H 2 O S 2 + 4H 2 S0 4 = 6S0 2 + 4H 2 C + 2H 2 S0 4 = 2SO, + C0 2 + 2H 2 O (f ) FeS0 4 + S 2 = FeS + 2S0 2 (g) Na 2 S0 3 + 2H 2 SO 4 = 2NaHSO 4 + SO 2 -f H 2 O 4 Preparation. (a) By heating moderately concentrated sulphuric acid with copper turnings: Cu + 2H 2 SO 4 = CuSO 4 -f SO 2 + 2H 2 O . The gas is dried by passing through concentrated sulphuric acid. (&) By heating a mixture of sulphur and cupric oxide in a hard glass tube, (c) In a Kipp's generator by decomposing cubes composed of three parts calcium sulphite and one part of calcium sulphate, with dilute sulphuric acid (Neumann, B., 1887, 20, 1584). Preparation of sulphites. The sulphites of the ordinary metals are usually made by action of sulphurous acid upon the oxides or hydroxides of the metals. are normal, except mercurous, which is acid, ancl chromium, aluminum 328 SULPHUROUS ACID. 265, 5. and copper, which are basic. Sulphurous acid precipitates solutions of metals of the first and second groups, except copper and cadmium. The sulphites of the alkalis precipitate solutions of the other metals except chromium salts; and some normal sulphites may be made in this manner. The sulphites of silver, mercury, copper and ferricum (known only in solution) are unstable, the sulphurous acid becoming sulphuric at the expense of the base, which is reduced to a form having a less number of bonds. With the unstable stannous sulphite the action is the reverse. (See 6 A.) All sulphites by ex- posure to the air slowly absorb oxygen, and are partially converted into sulphates. 5. Solubilities. One volume of water at dissolves 68.861 volumes of sul- phurous anhydride; at 20, 36.206 volumes (Carius, A., 1855, 94, 148); or at 20 , 0.104 part by weight (Sims J. C., 1862, 14, 1). One volume of alcohol dissolves at 15, 116 vol. SO2 . Charcoal absorbs 165 volumes, camphor 308 volumes, glacial acetic acid 318 volumes of the gas. Liquid sulphurous anhydride dis- solves P , S , I , Br and many gases. The sulphites of the metals of the alkalis are freely soluble in water; the normal sulphites of all other metals are insoluble, or but very slightly soluble in water. The sulphites of the metals of the alkaline earths, and some others, are soluble in solution of sulphurous acid, the solution being precipitated on boiling. The alkali bases form acid sulphites (bisulphites), which can be obtained in the solid state, but evolve sulphurous anhydride. The sulphites are insoluble in alcohol. They are decomposed by all acids except carbonic and boric, and in some instances, hydrosulphuric. 6. Reactions. A With metals and their compounds. Sulphurous acid reacts with Zn , Fe , Sn , and Cu to form hyposulphurous acid, H 2 SO, (Schiitzenberger, C. r., 1869, 69, 196). With Zn in the presence of HC1 it is reduced to hydrosulphuric acid: 3Zn + 6HC1 + H 2 S0 3 = 3ZnCl 2 + H 2 S -f- 3H 2 . Free sulphurous acid precipitates solutions of first and second group metals except those of copper and cadmium; solutions of other metallic salts are not precipitated owing to the solubility of the sulphites in acids. Alkali sulphites precipitate solutions of all other metallic salts. The precipitates, mostly white, are soluble in acetic acid. The precipitates of Pb , Hg , Ba , Sr , and Ca are usually accompanied by sulphates, due to the fact that soluble sulphites nearly always contain sulphates (4). Solution of lead acetate precipitates, from solutions of sulphites, lead sulphite, PbS0 3 , white, easily soluble in dilute nitric acid ; and not blacken- ing when boiled (distinction from thiosulphate). Solution of silver nitrate gives a white precipitate of silver sulphite, Ag 2 S0 3 , easily soluble in very dilute nitric acid or in excess of alkaline sulphite, and turning dark- brown when boiled, by formation of metallic silver and sulphuric acid. Solution of mercurous nitrate with sodium sulphite gives a gray precipi- tate of metallic mercury. Solution of mercuric chloride produces no change in the cold; but on boiling, the white mercurous chloride is precipi- tated, with formation of sulphuric acid. Still further digestion, with sufficient sulphite, reduces the white mercurous chloride to gray metallic mercury (58, 6e). Solution of ferric chloride gives a red solution of ferric sulphite, Fej(SO g ) a ; or, in more concentrated solutions, a yellowish precipitate of 265, 6//, 2. SULPHUROUS ACID. 329 basic ferric sulphite, also formed by addition of alcohol to the red solu- tion. The red solution is decolored on boiling; the acid radical reducing the basic radical, and forming ferrous sulphate. Solution of barium chloride gives a white precipitate of barium sul- BaSO ;! , easily soluble in dilute hydrochloric acid distinction from , which is imdis^olved, and should be filtered out. Now, on adding to the filtrate nitrohydrochloric acid, a precipitate of barium sulphate is obtained evidence that sulphite has been dissolved by the hydrochloric acid: BaS0 3 + 2HC1 = BaCL + H 2 SO ;! BaCl 2 + H 2 S0 3 + C1 2 + H 2 O = BaS0 4 + 4HC1 One part of barium sulphite is dissolved by 46,000 parts of water at 18. Calcium chloride reacts similar to barium chloride, the precipitate of calcium sulphite being less soluble in water than the corresponding sul- phate. One part of calcium sulphite is dissolved by 800 parts of water at 18 while one part of the strontium salt is dissolved by 30,000 parts water at 18. Sulphurous acid and sulphites are active reducing agents by virtue of their capacity for oxidation to sulphuric acid and sulphates. The reactions with silver, mercury and ferricum given above illustrate the reducing action, and the following should also be noted: Pb0 2 becomes lead sulphate. As v forms arsenous and sulphuric acids. Sb v forms Sb'". Cu" becomes cuprous sulphate. Cr VI forms chromic sulphate. Co'" forms cobaltous sulphate. Ni'" forms nickel sulphate. Mn"+ n forms manganous sulphate. With Mn0 2 in the cold, manganous dithionate, MnS 2 6 , is formed (Gmelin's Hand-look., 2, 174). With stannous chloride sulphurous acid acts as an oxidizing agent, form- ing stannic sulphide and stannic chloride or stannic chloride and hydro- sulphuric acid, according to the amount of hydrochloric acid present (71, l). Sulphites are decomposed by heat into oxides and sulphurous anhydride: CaSO 3 CaO + SO,; or into sulphates and sulphides: 4NaoS0 3 = 3Na.,S0 4 + Na 2 S. 8. Detection. Free sulphurous acid is detected by its odor and by its decolorizing action upon a solution of KMn0 4 or I (Hilger, J. (7., 1876, 29, 443). The reaction with iodic acid is also employed as a test for sulphurous acid (as well as for iodic). A mixture of iodic acid and starch is turned violet to blue by traces of sulphurous acid or sulphites in vapor or in solution, the color being destroyed by excess of the sulphurous acid or the sulphite. Sulphites are distinguished from sulphates by failure to precipitate with BaCl 2 in presence of HC1 . After removal of the BaS0 4 by filtration the sulphite is oxidized to sulphate by chlorine water and precipitated by the excess of BaCl 2 present. For separation from sulphides and thiosulphates see 258, 8. Normal potassium sulphite, K 2 S0 3 , is alkaline to litmus but when treated with BaCl 2 gives a neutral solution. The acid sulphite, KHS0 3 , is neutral to litmus but with BaCl 2 gives an acid solution: 2KHSO n + BaCL = BaS0 3 + 2KC1 + S0 2 '+ H 2 6 (Villiers, C. r., 1887, 104, 1177). 9. Estimation. (rt) After converting into H 2 S0 4 by HN0 3 or Cl it is precipi- tated by BaCL and weighed as BaSO 4 . (6) The oxidation is effected by fusing with Na 2 C0 3 and KNO 3 (equal parts), (c) A standard solution of iodine is added, and the excess of iodine determined by a standard solution of N"a 3 S 2 3 . 266, 4. SULPHURIC ACID. 266. Sulphuric acid. H 6 S0 4 = 98.076. H',S VI 0~" 4 , H - S H . II 1. Properties. Absolute sulphuric acid, H 2 S0 4 , is a colorless oilv liquid (oil of vitriol): spec-Hie gravity, l.S.'iTl at 15 (Meiidelejeff, #., 1884, 17, 2541). According 1 to Marignac (.1. 1'Ji.. isr>:{. (:;), 39, 184), it begins to boil at about 290, ascending' to :-i:!S with partial decomposition. At temperatures much below the boiling- point (160) it vaporizes from open vessels, giving off heavy, white, suffocating vapors, exciting- coughing- without giving premonition by odor. At ordinary temperature it is non-volatile and inodorous. At low tem- peratures it solidifies to a crystalline mass. The freezing point is greatly influenced by the amount of water present. When the acid contains one mole- cule of water, H 2 SO 4 .H 2 O , the melting point is highest, -f-7.5 (Pierre and Puchot, A. Ch., 1874, (5), 164). H 2 S0 4 is a very strong acid and, because of its high boiling point, displaces all the volatile inorganic acids; on the other hand it is displaced, when heated above its boiling point, by phosphoric, boric, and silicic acids. It is a dibasic acid, forming two series of salts, M'HS0 4 and M' 2 S0 4 . It is miscible with water in all proportions with production of heat; it abstracts water from the air (use in desiccators), and quickly abstracts the elements of water from many organic compounds, and leaves their carbon, a char- acteristic charring effect. It dissolves in alcohol, without decomposing it but if in sufficient proportion producing ethylsulphuric acid, HC 2 H 5 S0 4 . Sulphuric anhydride, S0 3 , is a colorless, fibrous or waxy solid, melting- at 14.8 (Rebs, A., 1888, 246, 379), boiling- at 46 (Schulz-Sellak, B., 1870, 3, 215), and vaporizing with heavy white fumes in the air at ordinary temperatures. It is very deliquescent, and on contact with water combines rapidly, forming sulphuric acid with generation of much heat. 2. Occurrence. Found free in the spring water of volcanic districts. Found combined in g-ypsum, CaS0 4 + 2HoO; in heavy spar, BaS0 4 ; in celestite, SrSO 4 ; in Epsom salts, MgS0 4 + 7H 2 0; in Glauber salt, Na 2 S0 4 + 10H 2 O , etc. 3. Formation. (a) By electrolyzing- H,O , using Pt electrodes with pieces of S attached (Becquerel, C. r., 1863, 56, 237)". (6) By oxidizing S or SO 2 in presence of water by Cl , Br , HNO 3 , etc. (c) By heating S and H 2 to 200. (d) By adding H 2 6 to S0 3 . (e) By passing- a mixture of S0 2 and over platinum spong-e and then adding- water. 4. Preparation. Industrially, sulphuric acid is made by utilizing the S0 2 evolved as a by-product in roasting various sulphides e. g., iron and copper pyrites, blende, etc. (a) and (&); or by burning sulphur in the air to form the S0 2 . The S0 2 is oxidized and converted into sulphuric acid by two distinct processes known as the contact and the chamber process. In the contact process the S0 2 , after careful purification, arsenic especially being removed, is passed together with oxygen through a contact mass 33$ SULPHURIC ACID. 266, 4. containing finely divided platinum or other catalytic reagent maintained at the proper temperature. The S0 2 unites with the oxygen to form S0 3 , which -is absorbed in dilute sulphuric acid. This process is espe- cially advantageous for making concentrated or fuming sulphuric acid. In the chamber process the S0 2 is passed into a large leaden chamber and brought into contact with HN0 3 , steam, and air. The HN0 3 first oxidizes a portion of the S0 2 (c) ; the steam then reacts upon the NO. , forming HN0 3 and NO (d). This NO is at once oxidized again by the air to N0 2 , so that theoretically no nitric acid is lost, but all is used over again. Practically, traces of it are constantly escaping with the nitrogen intro- duced as air, so that a fresh supply of nitric acid is needed to make up for this loss. The dilute acid known as chamber acid is concentrated first in lead pans and then in platinum or silica pans. Commercial sulphuric acid known as oil of vitriol has a sp. gr. of 1.83 and contains 93% K,S0 4 ; when heated to 338 a 98% acid distills over. The absolute H 2 S0 4 cannot be made by evaporation or distillation; it still contains about two per cent of water. It may be made by adding to water, or to the H 2 S0 4 containing the two per cent of water, a little more S0 3 or H 2 S 2 7 than would be needed to make H 2 S0 4 ; then passing perfectly dry air through it until the excess of S0 3 is removed, leaving absolute H 2 S0 4 . Fuming pyrosulphuric, or Nordhausen sulphuric acid, H 2 S 2 7 , is made by solution of sulphuric anhydride in sulphuric acid (e) ; by drying FeS0 4 + 7H,0 until it becomes FeS0 4 + H 2 , and then distilling (/). Sulphuric anhydride is made by the action of heat on sodium pyrosulphate, Na 2 S 2 7 (g), prepared by heating NaHS0 4 to dull redness ; by distilling pyrosulphuric acid, the anhydride is collected in an ice-cooled receiver; by heating H 2 S0 4 with P 2 5 (7i): (a) 2ZnS + 30 2 = 2ZnO + 2S0 2 (&) 4FeS 2 + 11O 3 = 2Fe 2 3 -f SSO 3 (c) S0 2 + 2HNO S = H 2 S0 4 + 2N0 2 (d) 3N0 2 + H 2 = 2HN0 3 + NO (e) H 2 S0 4 + S0 3 = H 2 S 2 7 (0 4FeS0 4 + H 2 = 2Fe 2 3 + H 2 S 2 7 + 2SO, (g) Na 2 S 2 7 = Na 2 S0 4 + S0 3 (ft) H 2 S0 4 + P 2 5 = 2HP0 3 + S0 3 Sulphates are made: (a) by dissolving the metals in sulphuric acid; (#) by dissolving the oxides or hydroxides; (c) by displacement. All salts containing volatile acids are displaced by sulphuric acid and a sulphate formed (except the chlorides of mercury). The excess of acid may generally be expelled by evaporation, or the cry stills washed with cold water or alcohol. The insoluble sulphates are best made by precipita- tion. 266, 6A. SULPHURIC ACID. ' 333 5. Solubilities. Sulphuric acid is miscible with water in all proportions; the concentrated acid with generation of much heat. Sulphuric acid transposes the salts of nearly all other acids, forming sulphates, and either acids (as hydrochloric acid, $269, 4) or the products of their decomposi- tion ( as with chloric acid, 273, 6). Chl<>i < of silver, tin, and antimony are with difficulty transposed by sulphuric acid, and chlorides of mercury not at all. Also, at temperatures above about 300 phosphoric and silicic acids (and other acids not volatile at this temperature) transpose sulphates, with vaporization of sulphuric acid. The sulphates of Pb , Hg', Ba, Sr, and Ca are insoluble, those of Hg' and Ca sparingly soluble. Sulphuric acid and soluble sulphates precipi- tate solutions of the salts of Pb , Hg', Ba , Sr , and Ca ; Hg' and Ca salts incompletely. The metallic sulphates are insoluble in alcohol which pre- cipitates them from their moderately concentrated aqueous solutions. Alcohol added to solutions of the acid sulphates precipitates the normal sulphates, sulphuric acid remaining in solution: 2KHS0 4 = K 2 S0 4 + H 2 S0 4 . PbS0 4 is soluble in a saturated solution of NaCl in the cold, depositing after some time crystals of PbCl 2 , complete transposition being effected. A solution of PbCl 2 in NaCl is not precipitated on addition of H 2 S0 4 (Field, J. C., 1872, 25, 575). 6. Reactions. A. With metals and their compounds. Sulphuric acid, dilute, has no action on Pb , Hg , Ag , Cu *, and Bi . Au , Pt , Ir , and Kh are not attacked by the acid, dilute or concentrated; other metals are attacked by the hot concentrated acid with evolution of S0 2 . The fol- lowing metals : Sn , Th , Cd , Al , Fe , Co , Ni , Mn , Zn , Mg , K , and Ka are attacked by the acid of all degrees of concentration ; the dilute rapidly and the cold concentrated slowly, with evolution of hydrogen; the hot concentrated with evolution of S0 2 . The degree of concentration and the temperature may be regulated so that the two gases may be evolved in almost any desired proportions. A secondary reaction frequently takes place, the metal decomposing the S0 2 forming H 2 S or a sulphide; and the H 2 S decomposing the S0 2 with separation of sulphur (Ditte, A. Ch., 1890, (6), 19, 68; Muir and Adie, J. C., 1888, 53, 47). Sulphuric acid or soluble sulphates react with soluble barium salts to give barium sulphate, white, insoluble in hydrochloric or nitric acids. This insolubility is a distinction from all other acids except selenic and fluo- silicic. The precipitate formed in the cold is very fine and difficult to separate by nitration; if formed in hot acid solution and then boiled it is retained by a good filter. In dilute solution for complete precipitation the mixture should stand for some time. Solutions of lead salts give a * Andrews, J. Am. Soc,, 1896, 18, 251. 334 SULPHURIC ACID. 266, 6B. white precipitate of lead sulphate not transposed by acids except H 2 S (5), soluble in the fixed alkalies. The presence of alcohol makes the precipi- tation quantitative (57, 9). Solution of calcium salts not too dilute form a white precipitate of calcium sulphate (188, oc). Dilute sulphuric acid does not oxidize any of the lower metallic oxides. Concentrated sulphuric acid is an oxidizing agent. \Yhen hot it liberates one atom of oxygen and is reduced to sulphurous acid, which is decom- posed with the evolution of sulphur dioxide and water. The concentrated acid with the aid of heat effects the following changes: Hg 2 forms mercuric sulphate, and sulphurous anhydride is evolved. SnCl 2 forms, first, sulphurous anhydride, then hydrosulphuric acid, stannic chloride at the same time being produced. Fe" is changed to Fe 2 (S0 4 ) a by hot concentrated sulphuric acid. Mn"+ n forms MnS0 4 and . That is, all compounds of manganese having a degree of oxidation above the dyad are reduced to the dyad with evolution of oxygen. Potassium permanganate dissolves in cold concentrated sulphuric acid with formation of a green solution of a sulphate of the heptad manganese, (Mn0 3 ) 2 S0 4 (134, 5c). Similarly the hot concentrated acid also reduces Pb 1 ^ to Pb", Co'" to Co", Ni'" to Ni", Fe VI to Fe'", and Cr VI to Cr'", oxygen being liberated (oxidized) and the metal reduced while the bonds of the S0 4 radical are not changed; a sulphate of the metal being produced. B. With non-metals and their compounds. When dilute sulphuric acid transposes the salts of other acids, no other change occurs if the acid set free be stable under the conditions of its liberation. In ordinary reactions sulphuric acid never acts as a reducing agent. 1. Many organic acids and other organic compounds are decomposed by the hot concentrated acid, the elements of water being abstracted and carbon set free. Continued heating of the carbon with the hot concen- trated acid oxidizes it to C0 2 with liberation of S0 2 . H 2 C 2 4 becomes C0 2 , CO , and H 2 . The bonds of the H 2 S0 4 remain unchanged. K 4 Fe(CN) 6 with dilute H 2 S0 4 forms HCN : 2K 4 Fe(CN) G + 3H S0 4 = 6HCN + K 2 FeFe(CN), + 3K 2 S0 4 . Cyanates are decomposed into C0 2 and NH 3 : 2KCNO + 2H 2 S0 4 + 2H 2 = K 2 S0 4 + (NH 4 ) 2 S0 4 + 2C0 2 . Thiocyanates are also decomposed by concentrated sulphuric acid. 2. Nitrites are decomposed with formation of nitric acid and NO : 6KN0 2 + 3H 2 S0 4 = 3K 2 S0 4 + 2HNO :] + 4NO + 2H 2 . $266, 8. SULPHURIC ACID. 335 3. H 3 P0 2 or hypophosphites are oxidized to phosphoric acid with re- duction of the sulphuric acid to sulphurous acid and then to sulphur. 4. Sulphur is slowly changed by hot concentrated sulphuric acid to sulphurous acid with reduction of the sulphuric acid to the same com- pound. Hydrosulphuric acid with hot concentrated sulphuric acid is oxidized to sulphur with reduction of the sulphuric acid to sulphurous acid. Further oxidation may take place as indicated ahove. 5. Chlorates are transposed and then decomposed when treated with concentrated sulphuric acid : 3KC10 3 + 2H 2 S0 4 == 2KHS0 4 + KC10 4 + 2C10 2 + H L> . 6. HBr forms Br and S0 2 . No action except in concentrated solution. 7. HI forms I and S0 2 . 7. Ignition. All sulphates fused with a fixed alkali carbonate are transposed to carbonates (oxide or metal if the carbonate is decomposed by the heat used, 228, 7) with formation of a fixed alkali sulphate (method of analysis of insoluble sulphates). If the sulphate, or any other compound containing sulphur, is fused in the presence of carbon, as fusion with a fixed alkali carbonate on a piece of charcoal, the resulting mass contains an alkali sulphide, which, when moistened, blackens metallic silver. The sulphates of Cu , Sb , Fe , Hg , Ni and Sn are completely decomposed at a red heat: 2FeSO 4 = Fe 2 O 3 + S0 3 + S0 2 ; 2CuS0 4 = 2CuO + 2S0 2 -f O 2 . A white heat decomposes the sulphates of Al , Cd , Ag , Pb , Mn and Zn . An ordinary white heat has no action on the sulphates of the alkalis and alkaline earths; but at the most intense heat procurable the sulphates of Ba , Ca and Sr are changed to oxides; and at the same temperature K 2 S0 4 and Na 2 SO 4 are completely volatilized, preceded by partial decomposition. Lead sulphate heated in a current of hydrogen is reduced according to the following equation: 2PbS0 4 + 6H 2 = Pb + PbS + SO 2 + 6H,0 . After a distinct interval the remainder of the sulphur is removed as H 2 S: PbS -f- H 2 = Pb + H : S (Rodwell, J. C., 1863, 16, 42). Potassium sulphate heated in a current of hydrogen is reduced to potassium acid-sulphide: K 2 S0 4 -f- 4H- = KOH + KHS + 3H 2 O (Berthelot, A. Cft,, 1890, (6), 21, 400). Potassium acid- sulphate, KHSO 4 , heated to 200 evolves H 2 SO 4 . The sodium acid-sulphate decomposes more readily. 8. Detection. Free sulphuric acid or the soluble sulphates are detected by precipitation in hot hydrochloric acid solution with barium chloride, forming the white, granular, insoluble barium sulphate. The sulphates insoluble in water are decomposed for analysis (1st) by long boiling with solution of alkali carbonate; and more readily (2d) by fusion with an alkali carbonate. In both cases there are produced alkali mil jrtid lex soluble in water, and carbonates soluble by hydrochloric or nitric acid, after removing the sulphate (a). If the fusion be done on charcoal, more or less dcoxidation will occur, reducing a part or the whole of the 336 PERSULPHURIC ACID. 266, 9. sulphate to sulphide (7), and the carbonate to metal (as with lead, 57, 7), or leaving the metal as a carbonate or oxide (7, 222 and 228). a. BaS0 4 + Na 2 C0 3 = Na 2 S0 4 (soluble in water) + BaC0 3 (soluble in acid). A mixture of H,S0 4 and a sulphate may be separated by strong alcohol, which precipitates the latter. A test for free sulphuric acid, in distinction from sulphates, may be made by the use of cane sug-ar, as follows: A little of the liquid to be tested is concentrated on the water-bath; then from two to four drops of it are taken on a piece of porcelain, with a small fragment of white sugar, and evaporated to drj'ness by the water-bath. A greenish-black residue indicates sulphuric acid. (With the same treatment, hydrochloric acid gives a brownish-black, and nitric acid a yellow-brown residue.) A strip of white glazed paper, wet with the liquid tested, by immersing it several times at short intervals, then dried in the oven at 100, will be colored black, brown or reddish, if the liquid contains as much as 0.2 per cent of sulphuric acid. 9. Estimation __ (a) By precipitation as barium sulphate and weighing as such. The solution should be hot and acidified with hydrochloric acid, and the mixture should be boiled a few minutes after the addition of the barium chloride. (6) By precipitation as barium sulphate with an excess of an hydro- chloric acid solution of barium chromate (three per cent hydrochloric acid). Add NH 4 OH , fill to a definite volume, and filter through a dry filter-paper. Transfer an aliquot portion to an azotometer with H 2 2 , and after acidifying, determine the oxygen evolved (Baumann, Z. anc/ew., 1891, 140) (244, 6A, 12) (e) When present in small amounts in drinking water by a photometric method (Hinds, C. N., 1896, 73, 285 and 299). 267. Persulphuric acid. HS0 4 = 97.068. 1. The anhydride. The anhydride, S 2 O 7 , was discovered by Berthelot (C. r., 1878, 88, 20 and 71). It is obtained by the action of the silent electric discharge upon a mixture of equal volumes of dry SO 2 and O . At 0, it consists of flexible cyrstalline needles, remaining stable for several days. When heated it decomposes into SO 3 and O . With SO 2 it combines to form SO 3 : S 2 O 7 +SO2 = 3SO 3 . Although in its reactions it acts as a strong oxidizing agent, it is weaker than chlorine or ozone; oxalic acid and chromium salts are not oxidized (Traube, B., 1889, 22, 1518, 1528; 1892, 25, 95). 2. The Acid. The acid was first prepared by Marshall, who electrolyzed cold fairly dilute sulphuric acid (J . C., 9, 771). Hydrogen is liberated at the cathode while the HSO 4 anions discharged at the anode unite to form persulphuric acid, the following reaction taking place. The acid may also be formed by the action of H 2 O 2 on concentrated H 2 SO 4 . Water solutions of the acid decompose very rapidly. Solutions of the acid, in concen- trated sulphuric acid, are more stable. 3. Salts. The potassium salt, K 2 S 2 O S , is prepared by the electrolysis of a sat- urated solution of KHSO 4 with a current of 3 to 3.5 amperes. It is a white crys- talline powder, which may be recrystallized from hot water with almost no decom- position. Continued heating of the solution effects decomposition. One hundred parts of water dissolve 0.564 part of the salt at and 4.08 parts at 40. The ammonium salt is prepared by the electrolysis of a saturated solution of ammonium sulphate. One hundred parts of water dissolve 58.2 parts of the salt at 0. It can be recrystallized from water if the solution not heated above 60. It forms monoclinic crystals. The dry salt is stable at 100. It is used in the cyanide process for the recovery of gold (Elbs, Z. annew., 1897, 195). The potas- sium is the least soluble of the persulphate salts. A solution of KoCOs gives an abundant crystalline precipitate of K 2 S 2 Og from a solution of the ammonium salts. The barium salt, BaS 2 O s .4H 2 p , is fairly soluble and may be prepared by rubbing the ammonium salt with barium hydroxide. , 1. CHLORINE. 33? 3. Reactions. All persulphates when dissolved in water are decomposed slowly in the cold and more rapidly on heating, oxygen, free sulphuric acid and a sulphate being formed. 2K 2 S 2 O S + 2H 2 O = 2K 2 SO 4 + 2H 2 SO 4 + O 2 . A large proportion of the oxygen escapes as ozone, which may be identified by its odor arid action on starch iodide paper. Ammonium persulphate in water solution decomposes slowly at the ordinary temperature without the evolution of oxygen. 8(NH 4 ) 2 S 2 O S + 6H 2 O = 14NH,HSO 4 + 2H 2 SO 4 + 2HNO ;! . Persulphates act as strong oxidizing agents. Salts of Ag' , Mn" , Co" , Ni" and Pb" ore oxidized in the presence of alkalies to the peroxides of these metals. If ammonia and a little silver nitrate are added to a strong solution of ammonium persulphate, nitrogen is rapidly evolved and the solution becomes heated to boiling. The silver peroxide first formed oxidizes the ammonia with liberation of nitrogen. (Z. rtujs. Ch., 87, 255, 1901.) Fe" and Ce'" are oxidized to Fe'" and Ce"" salts. KI is rapidly oxidized; K 4 Fe(CN) 6 becomes K 3 Fe(CN) 6 ; Alcohol is slowly oxidized to aldehyde, rapidly on warming; organic dyes are slowly bleached. 4. Caio's Acid. By adding a solid persulphate to concentrated sulphuric acid at 0, a solution is obtained possessing strong oxidizing properties. It may also be obtained by adding 30 per cent hydrogen peroxide (perhydrate) to concentrated sulphuric acid: H 2 SO 4 + H 2 O 2 = H 2 O + H 2 SO 5 . The acid HgSO 5 is known as monopersulphuric acid [Z. angew., 1898, 845: Ber., 34, 853 (1901); 41, 1839 (1909)]. 5. Detection. Persulphates are tested for by their oxidizing properties and formation of the peroxides of some metals. They are distinguished from hydro- gen peroxide by the fact that persulphates do not decolorize potassium perman- ganates and do not produce a yellow color with titanium sulphate. 268. Chlorine. Cl = 35.46. Valence one, three, four, five, and seven. 1. Properties. Molecular weight, 70.92. Vapor density, 35.8. The molecule contains two atoms, C1 2 . Under ordinary air pressure it liquefies at 33.6 and solidifies at 102 (Olszewski, M., 1884, 5, 127). Under pressure of six atmos- pheres it liquefies at 0. It is a greenish-yellow, suffocating 1 gas, not com- bustible in oxygen, burns in hydrogen (in sunlight combines explosively), forming HC1 . On cooling an aqueous solution of the gas to 0, crystals of C1 2 .10H 2 O separate out (Faraday, Quart. Jour, of Sci., 1823, 15, 71). Chlorine when passed into a solution of KOH produces, if cold, KC1 and KC1O , if hot, KC1 and KC10 3 : 2KOH + CL KC1 + KC10 + H 2 0; GKOH + 3C1 2 = 5KC1 + KC1O., + 3H 2 O . Passed into an excess of NH 4 OH , NH 4 C1 and N are formed: 8NH 4 OH + 3C1 2 = 6NH 4 C1 + N_ + 8H,O; if chlorine be in excess chloride of nitrogen is formed: NH 4 OH -f 3C1 2 = NC1 3 + 3HC1 + H 2 O . The NC1 3 is one of the most dangerous explosives known; hence chlorine should never be passed into NH 4 OH or into a solution of ammonium salts without extreme caution. Chlorine bleaches litmus, indigo and most other organic coloring matter. The three elements, chlorine, bromine and iodine, resemble each other in almost all their properties, reactions and combinations, differing (as do their atomic weights, 35.45, 79.95, 126.85) with a regular progressive variation; so that their compounds present themselves to us as members of progressive series. In several particulars fluorine (atomic weight, 19.05) corresponds to the first member of this series (13). Two oxides of chlorine have been isolated: CLO , hypochlorous anhydride (270), and C10 2 , chlorine dioxide. The latter is made by the addition of H 2 SO 4 to KC1O 3 at 0. It is a yellowish-green gas, condensing at to a red- brown liquid. At 59 it becomes a crystalline solid, resembling K,Cr 2 7 . It may be preserved in the dark, but becomes explosive in the sunlight. 338 CHLORINE. 208, 2. The most important acids containing chlorine are discussed under the sections following. They are: Hydrochloric acid, HC1 . Hypochlorous acid, HC10 . Chlorous acid, HC10 2 . Chloric acid, HOICK . Perchloric acid, HC10 4 . 2. Occurrence. It does not occur free in nature, but its salts are numerous, the most abundant being NaCl . 3. Formation. (a) By the action of HC1 upon higher oxides as indi- cated in 269, 6A. The usual class-room or laboratory method is illus- trated by the following equations: Mn0 2 + 4HC1 = MnCL + C1 2 -f 2H,O Mn0 2 + 2NaCl + 3H,SO 4 = MnSO 4 + 2NaBSO 4 + 01, + 2H 2 (b) By fusing together NH 4 N0 3 and NH 4 C1 : 4NH 4 N0 3 + 2NH 4 C1 = 5N, + CL + 12H 2 O . (c) By ignition of dry MgCL in the air: 2MgCl + 2 = : MgO + 2C1 2 (Dewar, J. 8oc. Ind., 1887, 6, 775). (d) Some chlorides" are disso< iated by heat alone: 2AuCl 3 = 2Au + 8C1 2 . 4. Preparation. (a) Weldon's process: Mn0 2 is treated with HC1 , an-i the MnCL, formed is precipitated as Mn(OH), by adding Ca(OH) 2 . The Mn OH) 2 is warmed by steam, and air is blown into it, oxidizing it again to MnO, , and by repeating this process the same manganese is used over again. See ] unge and Prett (Z. angew., 1893, 99) for modification of this method, using E NO, . (&) Deacon's process: HC1 , mixed with air, is passed over fire-bricks moistened with CuCl 2 and heated to about 440. The heat first changes the CuOL to CuCl , evolving chlorine; then the oxygen of the air, aided by the HC1 , oxi- dizes the CuCl to CuCl 2 . It is not certain that the explanation is correct. It is only known that the hydrochloric acid which is passed into the appj ratus comes out as free chlorine, and that the copper chloride (small in amount) does not need renewing, (c) Electrolytic process: Chlorine is very largely pro luced as a by-product in the manufacture of caustic soda by the electrolysis of common salt. A number of processes have been developed in some of which the fuse I salt is electrolyzed while in others the electric current is passed through strong brine. 5. Solubilities. The maximum solubility of chlorine in water is at 10. At one volume of water dissolves 1.5 volumes of chlorine; at 10 three volumes; at 30 1.8 volumes (Eiegel and Walz, J., 1846, 72). B-iling completely removes the chlorine from water. The chlorine acts upo:i the water to a small extent, forming hydrochloric and hypochlorous acid : C1 2 H- H 2 O s=> H Cl+H CIO The reaction is reversible, but if an alkali is present the hydrogen ions are removed and all of the chlorine reacts with the water: C1 2 + H 2 O + 2OH = Cl + CIO + 2H 2 O Only chlorides and hypochlorites remain in solution. On acidifying .the solution the chlorine is again liberated. 268, 651. CHLORINE. 339 6. Reactions. A. With metals and their compounds. Chlorine is one of the most powerful oxidizing agents known, becoming always a chloride or hydrochloric acid. All metals are at lacked by moist chlorine, forming chlo ides, many of them combining with vivid incandescence. With per- fectly dry chlorine many of the metals are not at all attacked. Sn , Sb , :nd As are rapidly attacked, forming liquid chlorides (Cowper, J. (7., 188c, 43, 153; Veley, J. C., 1894, 65, 1). In the presence of acids the oxid ition of the metal takes place to the same degree as when that metallic compound is acted upon by HC1 (269, 6A); a chloride is formed having the same metallic valence that would have resulted from treating the oxide or hydroxide with hydrochloric acid, e. g., adding HC1 to Co 2 3 makes CoCl , not CoCl 3 , hence adding chlorine to metallic cobalt makes CoCL and not GoClo, . In alkaline mixture usually the highest degree of oxidation possible is attained, as indicated by the following: 1. Pb" becomes Pb0 2 and a chloride in alkaline mixture. With PbCl 2 , it is cli.imed that the unstable PbCl 4 is formed (Sobrero and Selmi, A. Ch., 1850, (3), 29, 162; Ditte, A. Ch., 1881, (5), 22, 566). 2. Hg' becomes Hg" in acid and in alkaline mixture; also HC1 or a chloride. 3. As"' becomes As v in acid and in alkaline mixture. Some water must be present or the reverse action takes place, forming AsCl. 5 (269, 6A2). 4- Sb"' becomes Sb v and a chloride with acids and alkalis. 5. Sn" becomes Sn IV and a chloride with acids and alkalis. 6. Mo VI ~ n becomes Mo VI and a chloride with acids and alkalis. 7. Bi"' becomes Bi v and a chloride with alkalis only. 8. Cn' becomes Cu" and a chloride with alkalis and with acids. 9. Cr'" becomes Cr VI and a chloride in alkaline mixture only. 10. Fe" becomes Fe'" and a chloride with acids and alkalis, but with alkalis it is also further oxidized to a ferrate. 11. Co" becomes Co(OH) 3 and a chloride with alkalis only. 12. Ni" becomes Ni(OH) 3 and a chloride with alkalis only. IS. Mn" becomes Mn0 2 and a chloride with alkalis only. See Ditte, /. c. s for formation of MnCl 4 . B. With non-metals and their compounds. 1. H 2 C 2 4 in acid mixture: H 2 C 2 4 + C1 2 == 2C0 2 + 2HC1 , the H 2 C 2 4 must be in excess and hot (Guya-rd, Bl., 1879, (2), 31, 299); in alkaline mixture: K 2 C 2 4 + 4KOH + C1 2 = 2K 2 C0 3 + 2KC1 + 2H 2 . HCN becomes CNC1 and HC1 (Bischoff, B., 1872, 5, 80). HCNS forms NH 3 , H 2 S0 4 , C0 2 , and other variable products, and HC1 (Liebig, A., 1844, 50, 337). H 4 Fc(CN) 6 becomes H 3 Fe(CN) and HC1 ; an excess of Cl finally decom- poses the H 3 Fe(CN) 6 . 340 CHLORINE. 268, 62. 2. Chlorine does not appear to have any oxidizing action upon the oxides or acids of nitrogen. 3. Phosphorus and all lower oxidized forms become ILP0 4 with forma- tion of HC1 . 4- Sulphur and all its lower oxidized forms are oxidized to H 2 S0 4 with formation of HC1 . In an alkaline solution a sulphate and a chloride are formed. With H 2 S , S is first deposited, which an excess of Cl oxidizes to H S0 4 . A sulphide in an alkaline mixture is at once oxidized to a sul- phate without apparent intermediate liberation of sulphur. 5. In alkaline mixture chlorine oxidizes chlorites, and hypochlorites to chlorates with formation of a chloride : KC10 2 + 2KOH + CL == KC10 3 + 2KC1 + H 2 . With NaOH a hypochlorite is formed if cold, if hot a chlorate : 2NaOH + C1 2 = NaCIO + NaCl + H 2 GNaOH + 3C1 2 = NaC10 3 + SNaCl + 3H 2 O 6. Chlorine does not oxidize bromine in acid mixture, in alkaline mix- ture a bromate and a chloride are formed. HBr in acid solution becomes free bromine, in alkaline mixture a bromute ; hydrochloric acid or a chloride being formed. 7. Iodine is oxidized to HI0 3 in acid mixture, forming HC1 ; in an alkaline mixture a periodate and a chloride are formed. From hydriodic acid or iodides, iodine is first liberated, followed by further oxidation as indicated above: 2HI + C1 2 = 2HC1 + I 2 ; I, + 5C1 2 + 6H 2 = 2HI0 3 + 10HC1 ; KI -f 8KOH + 4C1 2 = KI0 4 + 8KC1 + 4H 2 . By comparing the oxidizing action of Cl with that of Br and I, the following facts will be observed, and should be carefully considered. The elements chlorine, bromine, and iodine have an oxidizing power in reverse order of their atomic weights, chlorine being the strongest. That is, if all three have the same oxidizing effect, the chlorine acts with the greatest rapidity; and in some cases, as with cuprous salts, the chlorine oxidizes while the iodine does not. Their hydracids are reducing agents graded in the reverse order. If any increase of bonds takes place in presence of an acid, by chlorine, bromine or iodine, the same increase always occurs in presence of a fixed alkali. But the oxidation frequently goes further in presence of a fixed alkali. Thus, with chlorine and potassium hydroxide we form Pb0 2 , Ni(OH) 3 , Bi 2 5 , Co(OH) 3 , K 2 Fe0 4 , and Mn0 2 , which cannot be formed in presence of an acid. It is very important to remember that those oxides which are formed ~by chlorine, in presence of a fixed alkali, but not in presence of an acid, are the only ones which can be reduced by hydrochloric acid. And further, that this reduction proceeds not always to the original form, never proceeding beyond that number of bonds capable of being formed in presence of an acid. Thus, 269, 30. HYDROCHLORIC ACID. 341 any lead salt, with potassium hydroxide and chlorine, forms PbO, , and this treated with hydrochloric acid again forms the lead salt, PbCl, . And ferrous chloride with potassium hydroxide and chlorine forms K 2 Fe0 4 , in which iron is a true hexad, and K 2 Fe0 4 with hydrochloric acid forms, not the ferrous chloride with which we began, but ferric chloride, for it could only be oxidized to that point in presence of an acid. The above is true for bromine and iodine, as well as for chlorine. 7. Ignition. See 1. 8. Detection. Free chlorine is recognized by its odor, by its liberation of iodine from potassium iodide, by its bleaching action upon litmus, indigo, etc., and by its action as a powerful oxidizing agent (see above). Chlorine acts on metallic mercury in the cold, producing the insoluble mercurous chloride: 2Hg + C1 2 = 2HgCl . As hydrochloric acid does not act upon metallic mercury it may be separated from chlorine by this reaction, the mixture being shaken with mercury until the chlorine is removed. Silver nitrate precipitates one- sixth of the chlorine as chloride : 3C1 2 + 6AgNO 3 + 3H 2 O = 5AgCl+AgClO 3 + 6HNO 3 . 9. Estimation. (a) It is added to a solution of potassium iodide and the lib- erated iodine determined by standard sodium thiosulphate. (6) It is converted into a chloride by reducing agents, and estimated by the usual methods (269, 8). 269. Hydrochloric Acid. HC1 = 36.468. H'Cl-', H 01. 1. Properties. Vapor density, 18.22. At ordinary pressure it liquefies at -82.9 and solidifies at -112.5 (Weber, Z. Anorg., 74, 297 ; 1912). At 10 under pressure of 40 atmospheres it condenses to a colorless liquid (Faraday, TV., 1845, 155). Critical temperature, 52.3; critical pressure, 86 atmospheres (Dewar, C. N., 1885, 61, 27). Dissociated into H and Cl at about 1500, but combines again upon cooling (Deville, C. r., 1865, 60, 317). It is a colorless gas, having an acrid, irritating odor. Readily absorbed by water. The chemically pure concentrated acid has usually a specific gravity of 1.20, and contains 39.11 per cent HC1 (Lunge and Marchlewski, Z. angew., 1891, 4, 133). The U. S. P. acid has a specific gravity of 1.163 at 15 and contains 31.9 per cent HC1 . A concen- trated solution of HC1 gives off gaseous HC1 faster than H 2 O ; a dilute solution gives off H 2 O faster than HC1 , as a final result in both cases an acid sp. (jr. 1.1 distils unchanged at 110 and contains 20.18 per cent HC1 (Bineau, A. Cli., 1843, (3), 7, 257). 2. Occurrence. Found native only in the vicinity of volcanoes. Found as a chloride in many minerals, sodium chloride being the most abundant. 3. Formation. (a) All chlorides except those of mercury are trans- posed by H 2 S0 4 ; silver chloride must be heated nearly to the boiling point of the H 2 S0 4 before the action begins. Lead, antimony and tin chlorides are slowly transposed. 342 HYDROCHLORIC ACID. 2f9, 3&. (ft) By the action of sunlight on a mixture of H and Cl , or by heatmg the mixture to 150. (c) Platinum black, palladium, charcoal, and some other sub- stances which rapidly absorb gases will cause the union of the hydrogc a and the chlorine, (d) When hydrogen is passed over the heated chlorides of the most of the metals of the first four groups, the metals are set free and "tydro- chloric acid is formed, (e) Slowly formed by the action of chlorine upon water in the sunlight; rapidly by its action upon reducing acids srch as H 2 C 2 4 , HH 2 P0 2 , H 2 S, H 2 S0 3 , etc.: HH 2 PO 2 + 2C1 2 + 2H 2 O = H 3 P0 4 + 4HC1 . Chlorides may be made: (a) By direct union of the elements, riostly without heat. Whether an ous or ic salt is formed depends upca the amount of chlorine used, (b) By the action of hydrochloric acid upon the corresponding oxides, hydroxides, carbonates, or sulphites. The solutions formed may be evaporated to expel excess of acid. If the chloride? thus formed contain water of crystallization it cannot be removed by heat alone, for part of the acid is by this means driven off, and a basic salt rei mins. If the anhydrous chloride is desired, it may always be made by (a 1 , and when thus formed may be sublimed without decomposition, (c) Chi >rides of the first group are best made by precipitation, (d) Metals solu )le in hydrochloric acid evolve hydrogen and form chlorides. In these cases ous, and not ic, salts are formed, (e) Many chlorides may be formed by bringing HgCl 2 in contact with the hot metal. 4. Preparation. For commercial purposes, made by treating NaC; with H 2 S0 4 and distilling. 5. Solubilities. Hydrochloric acid (gas) is very soluble in waier as stated in (1); forming in its solutions of various strengths the hydro- chloric acid of commerce. Its combinations with metals, forming chlor- ides, are for the most part soluble in water. AgCl and HgCl are insoluble in water. PbCL is only slightly soluble in cold water (57, 5c). These three chlorides constitute the first or silver group of metals, and ar^ pre- cipitated from their solutions by hydrochloric acid or soluble chlorides (61). The following chlorides not commonly met with are inso'uble: cuprous chloride CuCl, aurous chloride AuCl, thallous chloride TlCi and platinous chloride, PtCL . The following oxychlorides are insoluble : BiOCl, SbOCl and Hg,Cl 2 . Solutions of lead salts are not precipitated by mercuric chloride; green chromic chloride is incompletely precipitated and a sulphuric acid solution of molybdenum oxychloride not at dl by silver nitrate. The chlorides of Sb'" , Sn" , and Bi require the presence of some free acid to keep them in solution. AsCl 3 , PC1 3 , SbCl . and SnCl 4 are liquids at ordinary temperature. The first two are decorr posed by water liberating HC1 : AsCl 3 + 3H 2 = H 3 As0 3 + 3HC1. A sat -rated solution of bismuth nitrate is precipitated by HC1 as the oxychloride ( 7 3, 6/) which is readily soluble in excess of HC1 . Hydrochloric acid increases the solubility of the chlorides of Pb , Hg , Ag , Sb , Au , Pt , Bi and Cu' ; it decreases the solubility of Cd , Cu" , Co, Ni , Mn, Th , Ba, Sr, 269, A2. HYDROCHLORIC ACID. 343 Ca , Mg , Au , K and NH 4 . Chlorides of Th , Ba , Na , K and NH 4 are nearly insoluble in strong HC1 ( Ditto, C. r., 1SS1, 92, 242; A. Ch. f 1881, (5), 22, 551; Berthelot, A. Ch. f 1881, (5), 23, 8(5). Chlorides of Li, Ca and Sr are soluble in absolute alcohol or amyl alcohol. Silver chloride is readily soluble in ammonium hydroxide (separation from lead and mercurous chlorides) (59, 6a); lead chloride is soluble in fixed alkali hydroxides (57, 6a). H( 1 dissolves or transposes all insoluble oxalates, carbonates, hypophos- phitcs, phosphates, and sulphites. Sulphides of Fe", Mn , and Zn are dissolved readily; those of Pb , Ag, Sb , Sn , Bi , Cu , Cd , Co, and Ni if the rcid be concentrated; As 2 S 3 and As a S r , are insoluble in the cold con- centrated acid, very slowly soluble in the hot concentrated acid; HgS , red, is insoluble; black, very slowly soluble in the hot concentrated acid. HgS0 4 is only partially transposed by HC1 (58, 6/), BaS0 4 not at all. The ; nsoluble sulphates of Pb , Hg', Sr , and Ca are slowly but completely dissolved by the hot concentrated acid. Many of the metallic chlorides are soluble in alcohol, a few are soluble in ether. 6. Reactions. A. With metals and their compounds. Hydrochloric acid acts upon the following metals, forming chlorides with evolution of hydrogen : Pb (slowly but completely), Sn , Cu (very slowly), Cd , Fe , Cr , Al . C!o , Ni , Mn , Zn , and the metals of the fifth and sixth groups : Ag , Hg , As , Sb , Au , Pt , and Bi are insoluble in HC1 (Ditte and Metzner, A. Ch., 1893, (6), 29, 389). The following metallic oxides and hydroxides are acted upon by hydro- chloric acid, forming chlorides of the metal without reduction, water be- ing the only by-product : Pb" , Ag , Hg , As'" (only with very concentrated acid), Sb, Sn, Au'", Pt , Mo VI , Bi'", Cu, Cd, Fe , Al , Cr"', Co", Ni", Mn". Zn , Ba , Sr , Ca , Mg , K , and Na . The ignited oxides unite with HC1 more slowly than when freshly precipitated or when dried at 100. Ignited Cr,0 3 is insoluble in HC1 : other ignited oxides, as Fe 2 3 , A1 2 3 , etc., require very long continued boiling with the HC1 to effect solution. The following metallic compounds are attacked by hydrochloric acid with reduction of the metal and evolution of chlorine: 1. Pb" +n becomes PbCL ; no action with a chloride in presence of a three per cent solution of acetic acid, while bromine is completely set free from a bromide by Pb0 2 in presence of three per cent of acetic acid (detection of a chloride in presence of a bromide) (Vortmann, M ., 1882, 3, 510; B., 1887, 15, 1106). 2. As v becomes AsCl 3 . (The presence of very concentrated HC1 is required; Fresenius, Z., 1862, 1, 448; Smith, J. Am. Soc., 1895, 17, 682 and 735.) 344 HYDROCHLORIC ACID. 265, 6 AS. 3. Bi v becomes BiCl, . 4- Cr vl becomes CrCl 3 . With K 2 Cr 2 7 , bromine is completely liberated from a bromide in presence of 4 cc. of H 2 S0 4 to 100 cc. of water. The chlorine of a chloride is not liberated, and the bromine may be removed by boiling. Test the solution for a chloride (Dechan, J. C., 1886, 49, 682). Dry HC1 does not reduce Cr VI but combines with it to form the volatile Cr0 2 Cl 2 , chlorochromic anhydride (method of detecting a chloride in the presence of a bromide). 5- With the exception of ferrates the salts of iron are not reduced by hydrochloric acid. 6. Co"+ n becomes CoCl 2 . 7. Ni"+ n becomes NiCl 2 . 8. Mn" +n becomes MnCL . Mn0 2 with small amounts of dilute H 2 S0 4 (1-10) may be used to detect a chloride in presence of an iodide or bromide. Boiling the mixture removes the iodine first, then the bromine; while the chlorine is not set free until considerable H 2 S0 4 has been added (Jones, C. N., 1883, 48, 296). A mixture of KHS0 4 and KMn0 4 completely liber- ates the bromine from a bromide in the cold. A chloride remains unde- composed until warmed. Aspirate off the bromine, warm and collect the chlorine (Berglund, Z., 1885, 24, 184). B. With non-metals and their compounds. 1. No reducing action with H 2 C 2 4 , H 2 C0 3 , HCN, HCNS, H 4 Fe(CN) 6 , and H,Fe(CN) e . 2. HN0 2 forms chiefly NO and Cl . HN0 3 forms N0 2 C1 and Cl , or NOC1 and Cl , or merely N0 2 and Cl . In case excess of HC1 is used the reaction is: 2HN0 3 + 6HC1 = 2NO + 3C1 2 + 4H 2 (Koninck and Nihoul, Z. anorg., 1890, 477). Dry HC1 gas, passed into a cold mixture of con- centrated H 2 S0 4 and HN0 3 , reacts according to the following equations: 2HC1 + 2HN0 3 = 2H 2 + 2N0 2 + C1 2 (Lunge, Z. angew., 1895, 4, 8, and 11). 3. No reducing action with H 2 S , H 2 S0 3 , or H 2 S0 4 . With thiosulphates the unstable H 2 S 2 3 is liberated which decomposes as follows : 2Na 2 S 2 3 -(- 4HC1 = 4NaCl + S 2 .+ 2S0 2 + 2H 2 . Sulphates of Ag and Kg' are completely transposed by HC1 , those of Ba , Sr , and Ca not at all, all others partially (Prescott, C. N., 1877, 36, 179). 4. With an excess of HC1, hypophosphites, phosphites, and phosphates are dissolved or transposed without reduction. 5. Hypochlorous acid forms chlorine and water : HC10 + HC1 = H 2 -j- C1 2 . Chloric acid forms C10 2 , C1 2 , and Cl in varying proportions, but with HC1 in excess the following reaction takes place : KC10 3 + 6HC1 ;= KC1 + 3C1 2 + 3H 2 (Koninck and Nihoul, Z, anorg., 1890, 481). 269, 8c. HYDROCHLORIC ACID. 345 6. KBr0 3 is decomposed by boiling with HC1 , the bromine being set free: 2KBrO, + 12HC1 = 2KC1 + Br, + 5C1 2 + 6H,0 (Kaemmerer, J. pr., 1862, 85, 452). 7. With HI0 3 , ICl-j and Cl are formed, no action in dilute solutions: HI0 3 + 5HC1 = IC1 3 + Cl, + 3H 2 (Ditte, A., 1870, 156, 336). According to Bugarsky (Z. anorg., 1895, 10, 387) KHI 2 G with dilute H 2 S0 4 does not liberate chlorine from a chloride even on boiling (separation from a bromide). 7. Ignition. The chlorides of metals are, generally, more volatile than the other compounds of the same metals: example, ferric chloride. Insoluble chlorides are readily transposed by fusion with sodium carbonate: PbCL + Na,C0 3 = PbO + 2NaCl + CO 2 . If the carbonate be mixed with charcoal, or if the fusion is done on a piece of charcoal, the metal is also reduced: 2PbCL + 2Na 2 CO, + C = 2Pb + 4NaCl + 3CO 2 . Heated in a bead of microcosmic salt, previously saturated with copper oxide in the inner blow-pipe flame, chlorides impart a blue color to the outer flame, due to copper chloride. Dry sodium sulphate at 150 is transposed by dry HC1 (Colson, C. r., 1897, 124, 81). Gaseous HC1 transposes potassium and sodium sulphates completely at a dull-red heat. With the sulphates of the alkaline earths the transposition is nearly complete (Hensgen, B., 1876, 9, 1671). The silver halides heated with bismuth sulphide on charcoal before the blow-pipe give distinguishing colored incrustations: Agl , bright red; Ag-Br , deep yellow; AgCl , white (Goldschmidt, C. C., 1876, 297). 8. Detection. (a) In its soluble compounds, when not in mixtures with bromides and iodides, hydrochloric acid is readily detected by pre- cipitation with solution of silver nitrate, as a white curdy precipitate, opalescence if only a trace be present, turning gray on exposure to the light. The properties of the precipitate of silver chloride are given in 59, 5c and 6/. It is of analytical interest in that it is freely soluble in ammonium hydroxide (considerably more freely than the bromide, and far more freely than the iodide of silver); soluble in hot, concentrated solution of am- monium carbonate (which dissolves traces of bromide, and no iodide of silver); insoluble in^iitric acid, temporarily soluble in strong hydrochloric acid, precipitating again on dilution. It should be observed, that it is appreciably soluble in solutions of chlorides. (b) A test for traces of free hydrochloric acid, in distinction from metallic chlorides, is made by heating the solution with MnO, , without adding an acid, and distilling into a solution of potassium iodide and starch. Larger proportions of HC1 are more frequently separated by distilling it intact. (c) Gaseous hydrochloric acid (formed by adding sulphuric acid to dry chlorides, 3a) is readily detected by the white fumes formed when brought in contact with ammonia vapor. Also by bringing a stirring rod moist- 'd with silver nitrate in contact with the hydrochloric acid gas, Con- 346 HYDROCHLORIC ACID. 269, 8d. firm by proving the solubility of the white precipitate in ammonium hydroxide. (d) The reaction with chromic anhydride is in use as a test for hydro- chloric acid, more especially in presence of bromides: (a) 2HC1 + Cr0 3 = Cr0 2 Cl 2 (chlorochromic anhydride) + H 2 O (ft) 4NaCl + K 2 Cr 2 7 + 3H 2 S0 4 = 2Cr0 2 Cl a + 2Na 2 S0 4 + K 2 S0 4 + 3H 2 To obtain a rapid production of the gas, so that it may be recognized by its color, the operation may be made as follows: Boil a mixture of solid potassium dichromate and sulphuric acid, in an evaporating-dish until bright red, and then add the substance * to be tested, in powder- obtained, if necessary, by evaporation of the solution. If chlorides are present, the chromium dioxydichloride rises instantly as a bright brownish- red gas. The distinction from bromine requires, however, that the mate- rial, which should be dry, should be distilled, by means of a tubulated flask or small retort, the vapors being condensed in a receiver, and neutral- ized with an alkali (c and -d). The chromate formed makes a yellow solu- tion (bromine, a colorless solution). As conclusive evidence of chlorine, the chromate (acidified with acetic acid), with lead acetate, forms a yellow precipitate (bromide, a white precipitate, if any) : (c) Cr0 2 Cl 2 + 2H 2 =: H 2 Cr0 4 + 2HC1 (d) CrO 2 Cl 2 + 4(NH 4 )OH = (NH 4 ) 2 Cr0 4 + 2NH 4 C1 + 2H 2 (e) To detect a chloride in the presence of a cyanide or thiocyanate, add an excess of silver nitrate, filter and wash. To the moist precipitate add a few drops of silver nitrate (318, 24) and then several cubic centi- meters of concentrated sulphuric acid and boil for two or three minutes. The silver cyanide and thiocyanate are completely dissolved with decom- position, while the silver chloride is not changed except on long continued boiling. The student should confirm by tests on known material. According to Borchers (C. N., 1883, 47, 218), to detect a chloride in the presence of a cyanide or a thiocyanate add silver^iitrate, filter, wash, and boil the precipitate with concentrated nitric acid to complete oxida- tion of the cyanogen compound. See Mann (Z., 1889, 28, 668) for detec- tion of a chloride in presence of an alkali thiocyanate by use of CuS0 4 and H 2 S . (/) If a solution containing iodides, bromides, and chlorides be boiled with Fe 2 (S0 4 ) 3 , all the iodine is liberated and may be collected in a solution of EJ and estimated with standard Na 2 S 2 3 . The solution should * With the chlorides of mercury no brown fumes are obtained as these chlorides are not transposed by the sulphuric acid; and the chlorides of lead, silver, antimony, and tin are so slowly transposed that the formation of the chromium dioxydichloride may escape observation. Before relying upon this test the absence of the above named metals should be assured. 269, 8k. HYDROCHLORIC ACID. 34 1 } be cooled to about 60 and a slight excess of KMn0 4 added. The bromine is all liberated and may be collected in NH 4 OH and estimated as a bromide after reduction with S0 2 . The chloride may now be detected in the filtrate and may be estimated by one of the usual methods. Aspiration aids the removal of the iodine and bromine (Weiss, C. C., 1885, 634 and 712; Hart, C. N., 1884, 50, 268). (g) Villiers and Fayotte (C. r., 1894, 118, 1152, 1204 and 1413) detect a chloride in presence of an iodide and bromide by passing the liberated halogens into a solution of aniline in acetic acid (400 cc. of a saturated water solution of aniline to 100 cc. of glacial acetic acid) use 3 to 5 cc. of this solution for each test. Iodine gives no precipitate; bromine gives a white precipitate; and chlorine a black precipitate. If the bromide be present in large excess, add silver nitrate, digest the precipitate with ammonium hydroxide, add hydrogen sulphide and test the filtrate as the original solution. Liberate the halogen with KMn0 4 and H Q S0 4 . (70 Deniges (Bl, 1890, (3), 4, 481; 1891, (3), 5, 66) uses H 2 S0 4 and Fe"' to liberate the iodine, and K 2 Cr0 4 to liberate the bromine; then after boiling off the I and Br he adds KMn0 4 to liberate the chlorine. The iodine he detects with starch paper, the bromine fumes are absorbed on a rod moistened with KOH , which then gives an orange-yellow color with aniline. The chlorine he collects as the bromine and obtains a violet color with aniline. (i) Dechan (J. 0., 1886, 50, 682; 1887, 51, 690) removes iodine of iodides by distilling with a concentrated solution of K 2 Cr 2 7 ; then the bromine of bromides by adding dilute H 2 S0 4 and again distilling. The chloride is precipitated by AgN0 3 after dilution and addition of HN0 3 . (/) Vortman (If., 1882, 3, 510; Z., 1886, 25, 172) detects chlorine in presence of bromine and iodine as follows: The solution containing the halogens combined with the alkali or alkaline earth metals is heated with acetic acid and peroxide of lead until the supernatant liquid is colorless and has no longer the slightest odor of iodine or bromine; in this way the whole of the bromine and part of the iodine are driven off, the remainder of the latter remaining as ioclate of lead along with the excess of lead peroxide. This is filtered off, the precipitate washed with boiling water, and the chlorine precipitated from the filtrate by addition of silver nitrate. (k) The halogens may also be very readily separated by means of potassium persulphate, K 2 S 2 8 (318, 15). In dilute acetic acid iodine is liberated while bromides and chlorides are not oxidized. On acidifying with H..S04 bromine is liberated, while there is no action on chlorides if the strength of the sulphuric acid does not exceed 2N. If the free iodine and bromine have been removed by CS 2 or boiling, the chlorine may be 348 HYPOCHLOROUS ACID. 269, 9. precipitated by means of AgN0 3 . As in the presence of chlorates the iodine is oxidized to I0 3 , C10 3 must be absent. If present, the halogens must be precipitated as silver salts and reduced with metallic zinc. The following reactions take place: 2K1 + K 2 S 2 O 8 = 2K 2 SO 4 + I 2 . 2KBr + K 2 S 2 O 8 + H 2 SO 4 = 2K 2 SO 4 + Br 2 + H 2 SO 4 . KC1 + K 2 S 2 O 8 + H 2 SO 4 (1.5 2N) = No action. 2Agl + 2AgBr + 2AgCl + 3Zn = 6Ag + Znl 2 + ZnBr 2 + ZnCl 2 . 9. Estimation. (a) It is precipitated by AgNO 3 , washed, and after igni- tion, weighed as AgCl . (6) By a standard solution of AgNO 3 . A little Na 2 HPO 4 , or, better, K 2 Cr 2 O 7 , is added to the chloride to show the end of the reaction. When enough AgNO 3 has been added to combine with the chlorine the next addi- tion gives a yellow precipitate with the phosphate, or a red with the chromate. 270. Hypochlorous acid. HC10 = 52.468 . H'Cl'O-", H Cl . 1. Properties. Hypochlorous anhydride, CLO , is a reddish-yellow gas, con- densing- at about 20 to a blood-red liquid, which boils at about 17 (Pelouze, A. Ch., 1843, (3), 7, 176). Rise of temperature causes decomposition, explo- sively, into chlorine and oxygen (Balard, A. Ch., 1834, 57, 225). Molecular iccnjltt, 86.9. Vapor density, 43.5 at 10. The acid, HC1O , has not been isolated. Its aqueous solution smells like C1 2 O , decomposing rapidly, especially in the sun- light, into Cl and HC10 3 . 2. Occurrence. Not found in nature. 3. Formation. (a) By adding chlorine to HgO in the presence of water: 2HgO + 2C1 2 -f- H 2 = Hg 2 OCl 2 + 2HC1O (Carius, A., 1863, 126, 196). (b) By adding five per cent nitric acid to calcium hypochlorite and distilling at a low temperature (Koffer, A., 1875, 177, 314). (c) By passing chlorine into the sulphates of Mg , Zn , Al , Cu , Ca or Na: Na 2 SO 4 + C1 2 -f- H 2 O = NaHS0 4 + NaCl + HC1O . (d) By heating a mixture of KC10 3 and HoCoO 4 to 70 (Calvert and Davies, A. Ch., 1859, (3), 55, 485). 4. Preparation. For commercial purposes, as a bleaching agent and as a disinfectant; used as calcium hypochlorite with calcium chloride, chlorinated lime, made by bringing chlorine in contact with calcium hydroxide, without heating. Lunge and Schoch (B., 1887, 20, 1474) give the formula Ca"^ 1 to chlorinated lime. See also Kraut (A., 1882, 214, 244). Also as sodium hypochlorite, made by treating 1 sodium hydroxide with chlorine short of satu- ration in the cold: 2NaOH + C1 2 = NaCIO + NaCl + H 2 O . The sodium hypochlorite-and-chloride mixed as formed by chlorine in solution of sodium hydroxide or sodium carbonate, or by double decomposition between solution of the calcium hypochlorite-and-chloride and solution of sodium carbonate is pharmacopceial, under the name of solution of chlorinated soda (NaCl. NaCIO). 5. Solubilities. Hypochlorites are all soluble in water and are decomposed by heating. 6. Reactions. The hypochlorites are all unstable. They are decomposed by nearly all acids, including CO 2 : 2Ca(ClO) 2 + 2CO 2 = 2CaC0 3 + 2C1 2 + O 2 ; 271, 9. CHLOROUS ACID. 349 4NaClO + 4HC1 = 4NaCl + 2H 2 O + 2C1 2 -f O ? . They are very powerful oxidiz- ing agents, acting in acid solution as free chlorine, as the above equations indicate. Hypochlorites act as chlorine in alkaline mixture ($268,6) (Fresenius, Z. angew., 1895, 501). On warming, all hypochlorites when in solution are converted into chlorides and chlorates: 3NaC10 = 2NaCl + NaClO 3 . In the presence of 40 per cent or more of caustic potash, potassium hypochlorite decomposes into chloride with evolution of oxygen (Winteler, Z. Gngew., 33, (1902), 778). When shaken with mercury, hypochlorites or free hypochlorous acid produce a reddish basic mercuric chloride, insoluble in water, soluble in HC1 (distinction from free chlorine, which produces white mercurous chloride insoluble in HC1). 7. Ignition.- A 1 1 hypochlorites are decomposed by heat: 2KC10 = 2KC1 + O 2 . 8. Detection.- Although silver hypochlorite is soluble in water, it decom- poses very quickly, so 1h;i1 on adding- silver nitrate to sodium hypochlorite the final reaction is as follows: SNaCIO + :*AgNO 3 = 2AgCl + AgC10 3 -f- 3NaN0 3 . When KC1O is shaken with Hg , yellowish-red Hg.OCL is formed; the other potassium salts of chlorine, i.e., KC1 , KC1O, , KC10 3 and KC10 t , have no action upon Hg . An indigo solution is decolored by hypochlorites, while KMnO 4 is not decolored. If arsenons acid be present, the indigo solution is not decolored until the arsenons acid is all oxidized to arsenic acid. 9. Estimation. It is estimated as AgCl after reduction with Zn and H.,S0 4 . Rosenbaum (Z. anyew., 1893, 80) gives a method for estimating the various chlorine compounds in chlorinated lime. 271. Chlorous acid. HC10 2 = 68.468. H'Cl'"0-" 2 , H Cl = . 1. Properties. The anhydride, C1 2 3 , has not been isolated and the free acid is known only in solution, and this generally contains some HC1O 3 . It has an intense yellow color and is very unstable. 2. Occurrence. Neither the acid nor its salts are found in nature. 3. Formation An impure chlorous acid is said to be formed when KC10 3 is treated with HN0 3 and As,0 3 , C^H^O^ or C 6 H 8 (Millon, A. Ch., 1843, (3), 7, 298; Schiel. A., 1859, 109, 318; Carius, A., 1866, 140, 317). Chlorites of a number of metals have been made by adding the bases to a water solution of the acid; also from KC10, by transposition. 4. Preparation. KC1O 2 is prepared by adding an aqueous solution of C10 2 of known strength to the proper quantity of KOH , and evaporating in a vacuum The crystals of KC1O 3 which are formed in the reaction are removed and the mother liquor is crystallized from alcohol. 5. Solubilities. All chlorites which have been prepared are soluble in water, lead and silver chlorites sparingly soluble. 6. Reactions. Chlorouc acid or potassium chlorite in dilute acid solution is a powerful oxidizing agent, acting similar to chlorine. 7. Ignition. Chlorites when heated evolve oxygen and leave a chloride, or first a chloride and a chlorate (Brandau, A., 1869, 151, 340). 8. Detection. A concentrated solution of a chlorite gives a white precipitate with silver nitrate, fairly readily soluble in more water. KMnO, is decolored, a brown precipitate being formed. A solution of indigo is decolored even in presence of arsenous acid (distinction from hypochlorous acid). Chlorites when slightly acidulated give a transient amethyst tint to a solution of ferrous sulphate. 9. Estimation. By reduction to chloride and estimation as such. By meas- 350 CHLORINE PEROXIDE CHLORIC ACID. 272. uring the amount of ferrous iron oxidized to the ferric condition: 4FeSO 4 + HC1O 2 L- 2H 2 S0 4 = 2Fe 2 (SO 4 ) 3 + HC1 + 2H 2 O . 272, Chlorine Peroxide, C10 2 = 67.46. ci iv o-" 2 , o== cl ~~~ cl:=0 or = C1=:0 *- Chlorine peroxide, C1O 2 , at ordinary temperature, is a dark greenish-yellow gas. In concentrated solution it has very much the odor of nitrous acid. Cooled in a mixture of ice and salt it condenses to a bromine-red liquid; and in a mixture of solid C0 2 and ether it forms a mass of orange-yellow, brittle crystals. When warmed to about 60 it explodes with violence. In direct sunlight at ordinary temperature it decomposes slowlj- into chlorine and oxygen, while in the dark it is quite stable. In contact with many substances, as phosphorus, sulphur, sugar, ether, turpentine, etc., it explodes at ordinary temperature. In moist condition it bleaches blue litmus-paper without pre- viously reddening it. One volume of water absorbs about 20 volumes of the gas at 4 (Millon, A. 07*., 1843, (3), 7, 298). The solution in \vater contains HC10 2 and HC10 3 . It is prepared by carefully adding KC10 3 to cold concentrated H 2 S0 4 ; the mixture is then carefully warmed to 20, later somewhat higher. The gas is con- densed in a tube cooled by a mixture of ice and salt: 3KC1O 3 + 2H 2 S0 4 = 2KHSO 4 + KC1O 4 + H 2 O + 2C1O 2 (Millon, 1. c.). It is also made by warming a mixture of oxalic acid and potassium chlorate. When prepared in this man- ner it is mixed with C0 2 : 2KC1O, + 2H 2 C 2 O 4 = K 2 C 2 O 4 + 2H 2 O + 2C1O 2 + 2CO 2 (Calvert and Davies, A., 1859, 110, 344). It is also formed, mixed with chlorine, when KC1O 3 is warmed with HC1 . HI is oxidized to I; SO 2 to H 2 S0 4 . Indigo is bleached even in presence of As 2 O 3 . 273. Chloric acid. HC10 3 = 84.468. H'Cl v O-" 3 , H Cl 21 1. Properties. A solution of chloric acid may be evaporated in a vacuum until its specific gravity is 1.282 at 14. The composition is then HC10 3 .7H 2 O , containing 40.1 per cent HC1O 3 (Kaemmerer, Pogg., 1869, 138, 390). Farther attempts at concentration on heating to 40 result in evolution of chlorine and oxygen, forming HC1O 4 : 8HC1O 3 = 4HC1O 4 + 2H 2 O+3O 2 + 2C1 2 (Serullas, A. Ch., 1830, 45, 270). Its solution in the cold is odorless and colorless; first reddening and then bleaching litmus. It is a strong oxidizing agent, paper soaked with the acid takes fire on drying. The anhydride, C1 2 O 6 , has not been isolated. 2. Occurrence. Does not occur in nature. 3. Formation. The free acid may be formed by adding an excess of H 2 SiF 6 to a hot solution of KC1O 3 ; the filtrate is evaporated in vacuo, the excess of H 2 SiF 6 volatilizes, leaving the HC1O 3 . Many chlorates are formed by treating the metallic hydroxides with the free acid. Also by the action of Ba(Clp 3 ) 2 upon the sulphate of the metal whose chlorate is required; or by the action of the chloride of the chlorate needed, upon a solution of AgClO 3 . * Pebal, A., 1875, 177, 1. 273, GBL CHLORIC ACID. 351 4. Preparation. ijy adum!* 1 H,SO.i in molecular proportions 1o a solution of Ba(C10 3 ) 2 . Chlorates of the fifth and sixth group metals are prepared by passing 1 chlorine into the respective hydroxides dissolved or suspended in water. By repeated crystallization the chlorate is separated from the chloride which is also formed: 6KOH + 3CL = 5KC1 -f KC10 3 + 3ELO . 5. Solubilities. All chlorates are soluble in water, the chlorates of Hg , Sn , and Bi require a little free acid. Mercurous and ferrous chlorates are very unstable. Potassium chlorate is the least soluble of the stable metallic chlorates; soluble in about 21 parts water at 10 (Blarez, C. r.> 1891, 112, 1213). (]. Reactions. A. With metals and their compounds. Chloric acid attacks Mg evolving hydrogen and forming a chlorate only. With Zn , Fe, Sn, and Cu some chloride is also formed. With Zn and H 2 S0 4 the reduction to chloride is complete, and with sodium amalgam no reduction whatever (Thorpe, J. C., 1873, 26, 541). With the zinc-copper couple * the reduction to a chloride is rapid and complete. The hot concentrated acid attacks all metals. With oxides or hydroxides the acid forms chlor- ates provided a chlorate of that metal can by any means be formed. Free chloric acid is a strong oxidizing agent, and if an excess of the reducing agent is used, it is converted into hydrochloric acid, or a chloride. With the aid of heat the chloric acid splits up, forming some chlorine and oxides of chlorine. Hg' forms Hg". As"' forms As v . Sb'" forms Sb v . Sn" forms Sn iv . Cu' forms Cu". Cr"' forms Cr VI , chromic salts are readily oxidized to chromic acid on boiling with KC10 3 and HN0 3 . Fe" forms Fe'" (a distinction from perchloric acid) (Carnot, C. r., 1896, 122, 452). Mn" forms Mn IV , manganous salts are rapidly oxidized to Mn0 2 on warm- ing with KC10 3 and HNO, . Salts of lead, cobalt, and nickel do not appear to be oxidized on boiling with KC10 3 and HN0 3 . B. With non-metals and their compounds. 1. H 2 C 2 4 forms C0 2 and varying proportions of Cl and HC1 . Heat and excess of oxalic acid favors the production of HC1 (Guyard, Bl, 1879, * Gladstone and Tribe's copper-zinc couple is prepared by treating thin zinc foil with a 1 per cent solution of copper sulphate until the zinc is covered with a black deposit of reduced cop- per. When washed and dried it is ready for use. 352 CHLORIC ACID. 273, (2), 31, 299). All oxalates are decomposed, C0 2 and a chlorate or chloride of the metal being formed. Carbonates are all transposed. HCNS forms H 2 S0 4 , HCN , and HC1 . H 4 Fe(CN) 6 first forms H ;i Fe(CN) 6 and HC1 ; a great excess of HC10 3 decomposes the H,Fe(CN) 6 . 2. HN0 2 forms HN0 3 and Cl . Nitrites are transposed and oxidized, forming chlorates or nitrates of the metal. 3. PH 3 , HH 2 P0 2 , and H,P0 3 form H,P0 4 and HC1 . Hypophosphites and phosphites are transposed and then oxidized, H 3 P0 4 and a chlorate or a chloride of the metal being produced. 4. S VI ~ n forms S VI and HC1 ; that is, the sulphur of all compounds becomes H 2 S0 4 with formation of HC1 . All sulphides, sulphites, thio- sulphates, etc., are transposed, forming a chlorate, chloride, or sulphate of the metal. 5. HC1 in excess forms only Cl and H 2 (269, 6B5). NaCl warmed with HC10 3 evolves 01 , leaving only NaC10 3 . 6. HBr forms Br and HC1 . KBr warmed with HC10 3 evolves Br , leav- ing only KC10 3 . 7. I and HI form HI0 3 and HC1 . Soluble iodides form iodic acid or an iodate. 7. Ignition. All chlorates are resolved by heat into chlorides and oxygen: 2KC10 3 = 2KC1 + 30 2 . Some perchlorate is usually formed as an intermediate product: 2KC10 3 = KC10 4 + KC1 + 2 (Serullas, A. Ch., 1830, (2), 45, 270). In presence of various metallic oxides, etc., the oxygen is separated more easily, the metallic oxides remaining unchanged. With manganese dioxide, the oxygen of potassium chlorate is obtained at about 200; ferric oxide, platinum black, copper oxide, and lead dioxide may be used (242, 3). If chlorates are rapidly ignited some chlorine is given off (Spring and Prost, PL, 1889, (3), 1, 340). When triturated or heated with combustible substances, charcoal, organic substances, sulphu". sulphites, cyanides, thiosulphates, hypophosphites, reduced iron, etc. chlorates violently explode, owing to their sudden decomposition, and the simultaneous oxidation of the combustible material. This explosion is more violent than with corresponding mixtures of nitrates. Alkali chlorates when fused with an alkali, or an alkali carbonate, and a free metal or a lower oxide, or salt of the metal, generally oxidizes it to a higher oxide, or to a salt having an increased number of bonds; and the chlorate is reduced to a chloride e. g., Mn vl ~ n becomes Mn VI . That is, any compound of manganese having less than six bonds is oxidized to the hexad (a). Cr'" becomes Cr VI (6). As v ~ n becomes As v (c). Pb lv ~ n CALIFORNIA COLLEGE of PHARMACY 274, 3. PERCHLORIC ACID. 353 becomes Pb IV (d). Co'"- n becomes Co"' (e). C IV ~ n becomes C 1V (/). P v ~ n becomes P v (g). I v ~ n becomes I v (h). S VI ~ n becomes S VI (t). (a) 3Mn s 4 + 18KOH + 5KC10, = 9K,Mn0 4 -f 5KC1 + 9H.O (6) 2CrCl 3 + lONaOH + NaC10 3 = 2Na 2 CrO 4 + TNaCl + 5H 2 O (c) 3As 4 + 36KOH -f 10KC1O 3 = 12K,As0 4 + 10XC1 + 18H 2 O (d) 3Pb 3 O i + Na.COs + 2NaC10 3 = 9Pb0 2 + 2NaCl + Na 2 CO 3 (e) GCoCL + 12KOH + KC1O 3 = 3Co 2 3 + 13KC1 + 6H 2 (0 3K,C 4 H 4 O e + 5KC10 3 = 5KC1 + 3K 2 CO S + 9C0 2 + 6H 2 O (g) 3Pb(H 2 P0 2 ) 2 + 18KOH + 5KC10 3 = 3Pb0 2 + K 3 F0 4 + 5KC1 + 15H 2 O (ft) ZnI 2 + K,C0 3 + 2KC1O 3 = ZnO +, 2KIO 3 + 2KC1 + C0 2 (i) 3K,S 5 G + 12K 2 C0 3 + 10KC1O 3 = 15K 2 SO 4 + 10X01 + 12CO 2 8. Detection, ^u) Dry chlorates when warmed with concentrated sul- phuric acid, detonate evolving yellow fumes : 3KC10 3 + 2H 2 S0 4 2KHS0 4 + KC10 4 + 2C10 2 + H 2 . This action is modified by reducing agents; some acting rapidly, increase the detonation; others acting slowly, lessen it. (6) HC10 3 , like HN0 3 , decolors indigo solution and gives colors with brucine. diphenylamine, paratoluidine, and phenol similar to those formed by HN0 3 . (c) By ignition a chloride is left: 2KC10 3 = 2KC1 + 30 2 . (d) It is changed to a chloride by nascent hydrogen: 2KC10 3 + ^^n -|- 7H 2 S0 4 == 6ZnS0 4 + K 2 S0 4 + 2HC1 + 6H 2 0; or by reducing acids or bases: 2KC10 3 + H 2 S0 4 + 6HJ30, = K 2 S0 4 + 6H 2 S0 4 + 2HC1 . The resulting HC1 is then identified in the usual manner. Chlorides, if origin- ally present, should first be removed by silver nitrate. 9. Estimation. (a) Reduction to a chloride and estimation as such. (&) Addi- tion of HC1 and XI and estimation of the liberated iodine with standard 274. Perchloric acid. HC10 4 = 100.468. = H'Cl VII 0-" 4 . H Cl = = 1. Properties. Specific gravity, 1.782 at 15. The anhydrous HC10 4 is a color- less oily liquid, volatile but cannot be distilled without partial decomposition, often with explosive violence. Only its solution in water can be safely handled. Paper, charcoal, ether, phosphorus, and many other substances when brought in contact with the anhydrous acid take fire. The dilute acid is very stable, not being- easily reduced (Berthelot, A. Ch., 1882, (5), 27, 214). It does not bleach, but merely reddens blue litmus paper. 2. Occurrence. Not found in nature. 3. Formation. (a) By electrolysis of a solution of Cl or HC1 in water (Riche, C. r., 1858, 46, 348). (6) KC1O 4 is formed by electrolysis of KC1O 3 , using platinum electrodes (Lidoff and Tichomiroff, J. (7., 1883, 44, 149). (c) KC10 3 is heated with an excess of H,SiF 6 , after cooling and filtering, the filtrate is carefully distilled (Roscoe, J. C., 1863, 16, 82; A., 1862, 121, 346) (d) By treating the sulphate of the metal, the perchlorate of which is desired, 354 BROMINE. 274, 4, with Ba(ClO 4 ) 2 in molecular proportions, (e) By treating the chloride of the metal, the perchlorate of which is desired, with AgC10 4 in molecular propor- tions. 4. Preparation. KC10 4 is made by carefully heating KC10 3 until no more oxygen is evolved: 2KC10 3 = KC1 -f KC1O 4 + O 2 (7). The residue is dissolved in water and upon cooling 1 crystals of KC1O 4 separate. The free acid, nearly- pure, is obtained by cautiously distilling KC10 4 with concentrated H 2 S0 4 . 5. Solubilities. All of the perchlorates of the ordinary metals are soluble in water, and all are deliquescent except NH t ClO 4 KC1O 4 , Pb(C10 4 ) 2 and HgClO 4 (Serullas, A. Oh., 1831, 46, 362). Potassium perchlorate is soluble in 142.9 parts of water at 0, in 52.5 parts at 25, and in 5 parts at 100 (Muir, C. N., 1876, 33, 15). KC1O 4 is insoluble in alcohol (distinction from NaC10 4 ) (Schloessing, A. Ch., 1877, (5), 11, 561). 6. Reactions. Iron and zinc evolve hydrogen when treated with perchloric acid. The acid reacts with the hydroxides of many metals to form per- chlorates. It is not reduced by HCl , HNO 3 , H 2 S or SOo . Iodine is oxidized to HIO 4 with liberation of chlorine: I 2 + 2HC1O 4 = 2HI0 4 + C1 2 . A solution of indigo is not decolored by HC10 4 even after the addition of HCl (distinction from all other oxyacids of chlorine). It is not reduced by the zinc-copper couple (distinction from chlorate). Sodium perchlorate, NaClO 4 , is used as a reagent to precipitate potassiuir salts. 7. Ignition. Perchlorates strongly ignited evolve oxygen and leave a chloride (242, 3). 8. Detection. In presence of a hypochlorite, chlorite, chlorate and chloride boil thoroughly with HCl; the first three are decomposed, leaving chloride and perchlorate. Remove the chloride with AgN0 3 and fuse the evaporated filtrate with Na 2 C0 3 . Dissolve the fused mass in water and test for a chloride; its presence indicates the previous presence of a perchlorate. Perchlorates may also be separated from the other chlorine acids by passing SO2 gas, which reduces all the chlorine acids excepting perchloric acid. Blattner & Brasseur, Ch. Z., 24, 793. 9. Estimation. (a) After being changed to a chloride a,s indicated above, it is estimated in the usual manner. (6) It is fused with zinc chloride and the amount of chlorine liberated measured by the amount of iodine set free from a solution of potassium iodide (separation from chlorate, chlorides 'and nitrates). (c) KC1O 4 is heated to 200 with HPO 3 and KI ; the iodine liberated showing the amount of perchlorate present (Gooch and Kreider, Am. S., 1894, 48, 33; and 1895, 49 , 287). 275. Bromine. Br = 79.92. Valence one and five. 1. Properties. Molecular weight, 159.8; vapor density, 80; specific gravity, 3.18828 at 0; boiling point, 58.7. At 7.3 it becomes a brown solid (Burgess, Wash. Acad. of Sc., 1-18). At ordinary temperatures bromine is a brown-red, intensely caustic liquid, freely evolving brown vapors, corrosive vapors of a suf- focating chlorine-like odor. As a solid it is still darker in color. It reacts with KOH in all respects similar to chlorine (268, 1). Indigo, litmus and most other organic coloring matters are bleached. A solution of starch is colored slightly yellow. Bromine decomposes hydrosulphuric acid with separation of sulphur, and sub- sequent production of sulphuric acid; changes ferrous to ferric salts, and (in presence of water) acts as a strong oxidizing agent. It displaces iodine from iodides, and is displaced from bromides by chlorine; its character being intermediate between that of chlorine and that of iodine. No oxides of bromine have, with certainty, been isolated. The well-estab- lished acids are: Hydrobromic, HBr; hypobromous, HBrO; bromic, HBrO 3 , 275, 6^4, 11. BROMINE. 355 2. Occurrence. Not found free in nature. As a bromide in sea water, mother liquor from salt wells, mineral springs, and in a few minerals. 3. Formation (a) Hydrobromic acid or any soluble bromide is wanned with MnO 2 and H,S0 4 . (b) Any soluble bromide is treated with chlorine water and the solution warmtd. 4. Preparation. The bromine of commerce is obtained chiefly from the mother liquor of the salt works: (a) By treating- with MnO., and H',S0 4 : MgBr,, + Mn0 2 '+ 2H 2 S0 4 = MgSp 4 + MnS0 4 + Br 3 + 2H 2 . (b) By leading- a current of steam and chlorine into the bottom of a vessel filled with coke, into which a stream of the mother liquor flows from above: MgBr 2 + CL = MgCl 2 -f Br 2 . (c) By adding- to the mother liquor a mixture of Mg(OH) 2 , suspended in water and saturated with chlorine, rendering acid and distilling- in a current of steam: Mg(C10 3 ) 2 -f GMgBr, + 12HC1 = 7MgCl 2 + 6H 2 + 6Br 2 . ((/) By electrolysis of the mother liquor at a low temperature and then distilling- in a current of steam. Commercial bromine is freed from chlorine by adding- KBr and distilling-. If iodine be present it is first removed as Cul . 5. Solubilities. Bromine dissolves in 30 parts of water at 15, forming an orange-yellow solution (Dancer, J. C., 1862, 15, 477). Its water solution is not permanent, but slowly decomposes: 2Br 2 + 2H 2 O = 4HBr -f O 2 . Much more soluble in HC1 , HBr ', Kbr , BaCl 2 , SrCl 2 , and in many other salts than in water. Soluble in carbon disulphide, chloroform, ether and alcohol. Readily removed from its solution in water by shaking with carbon disulphide or chloro- form, imparting a brown color to the solvent. 6. Reactions. A. With metals and their compounds. Bromine unites directly with gold, platinum, and all ordinary metals to form bromides. It combines with metallic mercury forming the insoluble mercurous bromide. Silver salts are precipitated, yellow-white, as bromide and bromate : 6AgN0 3 + 3Br 2 + 3H 2 == 5AgBr + AgBr0 3 + 6HN0 3 . In the following metallic compounds the valence of the metal is changed; the bromine being reduced to HBr or, if in alkaline mixture, to a bromide. The reaction is less violent than with chlorine. 1. Pb" becomes Pb0 2 in alkaline mixture only. 2. Hg' becomes Hg" in acid and in alkaline mixture. 3. As'" becomes As v in acid and in alkaline mixture. With AsH 3 and a solution of bromine in water H 3 As0 3 is first formed, and if the bromine be in excess the final products are H 3 As0 4 and HBr . 4. Sb'" becomes Sb v in acid and in alkaline mixture. 5. Sn" becomes Sn IV in acid and in alkaline mixture. 6. Bi"' becomes Bi 2 5 in alkaline mixture only. 7. Cu' becomes Cu" in acid and alkaline mixture. S. Cr'" becomes Cr VI in alkaline mixture only. 9. Fe" becomes Fe'" in acid mixture; in alkaline mixture the iron is further oxidized to a ferrate, HBr or a bromide being formed. 10. Co" becomes Co"' in alkaline mixture only. 11. Ni" becomes Ni'" in alkaline mixture only (Kilpius, J. C., 1876, 29, 742). 356 BROMINE. 275, 6A, 12. 12. Mn IV ~ n becomes Mn IV in alkaline mixture only. B. With non-metals and their compounds. 1. H 2 C 2 4 becomes a carbonate and a bromide in alkaline mixture. An excess of hot saturated oxalic solution changes Br to HBr . HCNS forms, among other products, H 2 S0 4 and a bromide in acid mix- ture, and a sulphate and a bromide in alkaline mixture. H 4 Fe(CN) 6 in acid mixture forms H 3 Fe(CN) 6 and HBr , in alkaline mix- ture a ferricyanide and a bromide (Wagner, J. C., 1876, 29, 741). 2. HN0 2 becomes HN0 3 and HBr if dilute and cold. 3. PH 3 , HH 2 P0 2 and H 3 P0 3 become H 3 P0 4 and HBr with acids, and a phosphate and a bromide in alkaline mixture. P and Br unite to form PBr 3 or PBr 5 , depending upon relative amounts of the elements present. The phosphorus bromides are decomposed by water, forming HBr and the corresponding acids of phosphorus. 4. S, H 2 S , H 2 S0 3 , H 2 S 2 3 , S VI ~ n becomes H 2 S0 4 and HBr with acids, a sulphate and a bromide in alkaline mixture. 5. Br does not act as an oxidizing agent upon the compounds of chlorine, but may, at low temperatures, combine with chlorine to form a chlorine bromide, BrCl (Bornemann, A., 1877, 189, 183). 6. In alkaline mixture hypobromites by boiling are oxidized to bromates with formation of a bromide. 7. Iodine becomes an iodate and a bromide in alkaline mixture; the elements may combine to form the unstable bromiodide, IBr (Bornemann, /. c.). HI and iodides form I and HBr , but in alkaline mixture an iodate and a bromide are produced. 7. Ignition. Warming 1 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. Ether 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,O 3 . (ft) It is estimated by the amount of As 2 O 3 which it oxidizes in alkaline solution, (f) It is converted into HBr by H Z S or H 2 S0 3 , and then precipitated by and weighed as AgBr . 276, QA. BYDROBROMIC ACID. 357 276. Hydrobromic acid. HBr = 80.928. H'Br-', H Br. 1. Properties. Molecular weight, 149.9. Vapor density, 39.1. A colorless gas, condenses to a liquid at 09 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 is 1.78; containing 82.02 per cent HBr, 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. (jr. 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 O 4H 3 P0 4 -f- 20HBr . (b) By action of H 3 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 SO 4 added in molecular proportions, (d) By passing a mixture of Br and H over platinum sponge. (e) By action of Br on H 3 PO 2 . (f) By adding Br to Na 2 S0 3 . Metallic bromides are formed: (1) 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 , Na , Ba , Sr and Ca are made by the action of bromine on their hydrox- ides and subsequent fusion: 6KOH -f 3Br 3 = KBr0 3 + 5KBr + 3H 2 2KBr0 3 (ignited) = 2KBr + 30 a 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 (liecoura, C. r., 1890, 110, 784). (&) H 2 S0 4 is added to a concen- trated solution of KBr; after twenty-four hours the greater portion of the KHSO 4 has crystallized out. The remaining liquor is then distilled. The product usually contains traces of H 2 SO 4 . (c) By passing bromine into hot paraffme (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 OH . 6. Keactions. 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 358 HYDROBROMIC ACID. 276,6.1,7. calcium and the alkali groups. It unites with salt forming oxides and hydroxides to produce bromides without change of valence: PbO + 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 3 As0 4 4- 2HBr = H 3 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 not Fe" , and Br . 6. Cr VI 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) (Dechan, J. 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 Rivals, C. r., 1897, 124, 859 and 954). Silver nitrate solution precipitates, from solutions of bromides, silver bromide, 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 decom- posed by chlorine in the cold, rapidly when heated in a stream of chlorine. Solution of mercurous nitrate precipitates mercurous bromide, HgBr, yellowish-white, soluble in excess of alkali bromides. Solutions of lead salts precipitate, from solutions not very dilute, lead bromide, PbBr 2 , white. B. With non-metals and their compounds. 1. H,Fe(CN) ( . 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: 2K 4 Fe(CN) 6 + Br 2 == 2K 3 Fc(CN) a + 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. HYDItOBROMIC ACID. 359 3. Phosphorus compounds arc not reduced. 4. 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 -j- 2H 2 = H.,SO. t -f- 8, 46, 348). (b) By the action of chlorine on iodine in the presence of much water. The HC1 formed cannot be expelled by boiling without decomposing the HIO 3 . It must be removed by the careful addition of Ag 2 . (c) By adding water to IC1 3 and w r ashing with alcohol: 2IC1 3 -4- 3H 2 O = HIO 3 + 5HC1 + IC1 . ((/) KIO 3 is made by treating iodine with KOH: 3l a + 6KOH = KIOj + r>KI + 3H 2 O . And then washing with alcohol to remove the KE . (e) By heating potassium chlorate and iodine: 10KC10 3 + 6I 2 + 6H 2 O :== 6KHI 2 O 6 + 4KC1 + 6HC1 (Bassett, J. C., 1890, 57, 760). (f) By boiling iodine with barium hydroxide until neutral, filtering and decomposing with sulphuric acid (Steven- son, C. N., 1877, 36, 201). (g) By the action of I upon AgNO 3 : 5Ag-N0 3 + 31, + 3H,0 = 5 Agl + 5HN0 3 + HIO a . lodates of the alkalis and alkaline earths are easily made by the action of iodine on the hydroxides, and separation by alcohol or by crystallization from the iodides which are formed in the reaction. All iodates may be made by action of the acid on the hydroxides or carbonates. 4. Preparation. (a) Iodine is oxidized by boiling with nitric acid sp. gr. 1.52, and removing the excess of the nitric acid by evaporation. (6) By adding a slight excess of H 2 S0 4 to Ba(IO 3 ) 2 and removal of the excess of H 2 SO 4 by the careful addition of Ba(IOi) 2 (. HBr forms Br and I . 7. HI forms I from both acids. The addition of tartaric acid to a mix- ture of KI and KI0 3 is sufficient to give an immediate test for free iodine with CS 2 . It must be remembered that an iodide alone rendered acid will give a test for free iodine after a short time. 8. Morphine reduces iodic acid with separation of iodine. 7. Ignition. Potassium and sodium iodates on ignition form iodides and evolve oxygen (Cook, J. C., 1894, 65, 802). Many other iodates evolve oxygen but the iodide formed is further decomposed as stated in 275, 7. Iodates in dry mixture with combustible bodies are reduced, on heating or concussion, with detonation, but much less violently than chlorates or nitrates. 8. Detection. It is usually detected, after acidulation, by treatment with some reducing agent for the formation of free iodine. H 2 S0 3 is often employed because it acts rapidly and in the cold; but traces of HI0 3 frequently escape detection for the least excess of H 2 S0 3 at once reduces the iodine to colorless hydriodic acid. A desirable reagent for this reduc- tion is one that will act rapidly in the cold, and in no case cause the further reduction to hydriodic acid. The following reducing agents have been used : K 4 Fe(CN) B acidulated with dilute H 2 S0 4 , H 3 As0 3 , Cud , FeS0 4 , morphine sulphate and uric acid. -To detect KI0 3 in KI it is recom- mended by Schering (J. (7., 1873, 26, 191) to add a crystal of tartaric acid to the solution. The formation of a yellow zone is indicative of an iodate. Hydrochloric acid may be used, but if it contains a trace of chlorine it will give the test for an iodate. Iodine frequently occurs in nitric acid as iodic acid. Hilzer (J. (7., 1876, 29, 442) directs io add equal volumes of water, carbon disulphide, and then coarse zinc filings. It may be necessary to warm the solution slightly. Biltz (C. C., 1877, Sfi) dilutes the HNO r> with water, adds starch solution and then H 2 S solution drop by drop. A blue zone is formed if HI0 3 be present. 0. Estimation. (a) B.v precipitation with AgN0 3 , and after drying at 100 weighing- as Ag-10,. (ft) By reducing- to an iodide and estimating- as snch. (r) By treating- with KI acidulated with H 2 SO, , and titrating the iodine lib- erated with standard Na 2 S 2 O, . 372 PERIODIC ACID. 282. 282. Periodic acid. HI0 4 = 191.928. H H H \ I / 000 II XI / H'I VII 0-" 4 or H' 5 I VII 0-" 6 , H 1 = orH I H. The anhydride, I 2 7 , has not been isolated, and but one acid is known in the free condition, HIO 4 ,2H 2 O or H 5 IO 6 . 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 1 named acids: HI0 4 , H 3 I0 5 , H 3 I0 6 , H 4 I 2 9 , HJ.O^ , H 12 I 2 13 , H 10 I 4 O 19 , H 10 I 6 26 (Rammelsberg, Pogg., 1865, 134, 368, 499). The free periodic acid, H 5 IO C , is prepared: (a) By oxidizing iodine with per- chloric acid: 2HC1O 4 + I 2 + 4H,0 = 2H r> I0 6 + C1 2 (Kaemmerer, Pogg., 1869, 138, 406). (I)) By heating iodine or barium iodide with a mixture of barium oxide and barium peroxide, digesting with water, and transposing the Ba 5 (I0 6 ) 2 thus obtained with the calculated amount of sulphuric acid (Ram- melsberg, Pogg., 1869, 137, 305). (c) By conducting chlorine into sodium iodate in presence of sodium hydroxide: NalO, + 3NaOH -f C1 2 = Na 2 H 8 IO e + 2NaCl . This acid periodate dissolved in water with a little nitric acid and then precipitated with silver nitrate, forms the silver salt, AgoH-IOe . This precipitate is dissolved in nitric acid and evaporated on the water-bath, when orange-colored crystals of silver met a periodate are formed according to the following: 2Ag 2 H 3 'lO 6 + 2HN0 3 = 2AgI0 4 + 2AgN0 3 + 4H 2 O . Water decom- poses this precipitate: 2AgIO 4 + 4H 2 O = H 3 IO 6 + Ag 2 H 3 I0 6 . Or the silver periodate, AgI0 4 , is decomposed by Cl or Br (Kaemmerer, ?. c., p. 390). The silver salts vary in color: AgI0 4 is orange: Ag 2 HIO 5 , dark brown; Ag 4 I 2 9 , 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 CO 2 and I . H 3 PO 2 becomes H 3 P0 4 and HI . H 2 S becomes S and HI . H 2 S0 3 becomes H 2 SO 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,SO 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 -f I 2 + 8H,SO + H 2 O 2HI0 4 + 7H 2 S0 3 = I 3 + 7H,S0 4 + H 2 O HI0 4 + 4H 2 S0 3 = HI + 4H 2 S0 4 HC1 becomes Cl and IC1 8 HI forms I from both acids. According to Lautsch (J. pr., 1867, 100, 86), its behavior with mercurons nitrate is characteristic. The pentasodic periodate, Na 5 IO a , gives a light- yellow precipitate, Hg 5 IO, , 283. COMPARATIVE REACTIONS OF HALOGEN COMPOUNDS. 373 CO - 000 I 3 s I f i ooc OQ d {4 S i 3 1 o M -2 -2 02 2 'S, 'a '3, 000 rrt w 1 '3 2 '3 8- |-f sf ft ft ft (V ^ *> o o *C^3 * e 5 B f: 5Z G Q B B : 2 , CB eS -2 $ _e rf eS ^ 73 o* 'p, '3 2 Dissolv* to B 'ft 1 PH ft O H* z V S 2 i | ^ IN B 6 S 1 1 6 1 5 m H +.J 4 s ]p, '3 I d d * S o . d S ft 1 ? si w I d p pa ;, B oa s 3 S w^. BE' d ^ 5 i xA Q 1 A 3 * 1 B c 'a 'a ft "O 3 g M d "0 i G ft o 2 ft o ft O O W S ^ 03 2 5 :l 5 a fc B S a B B 1 o -d ^ 3 2 . Q) at V g "3 P _0 1 d 1 o A S H T3 5 ^ 1 1 K bo v S s 5 hn ^ M d ft? 1 "o M* to ft M ttt H 9 00 S 1 fc B M O B B M -' MM* unities. 1 3 ^ 1 . cipitate. cipitate. 1 ,2 P 3 i ^ 03 <5 3 | g I* "S, ^ o 2 T bo PQ PQ w 1 i .2 ft o ft PQ H S T PL, P & B B B BQ DQ ^ rQ a 3 -s 0) 1 s ; 9) | Id ^ 1 "$ 2 S : M* 2 'ft > ^j bo ; g o 1 i Q 8 PH q B B 4, i 3 1 p 1 ^ 1 A.gNO 3 , in excess. AfiNOj, excess of the tion tested. NH 4 OH, to the silve cipitate. HgCl 2 , in excess. HgClj, excess of the tion tested. HgNO 3 , excess. IJ.Clo. CuS0 4 , with H a SO s . I lid Br. HI. Chromate, with H 2 S< KMnO 4l dilute neut lfa,SO 3 , with H 2 SO 4 . w" PART IV -SYSTEMATIC EXAMINATIONS. REMOVAL OF ORGANIC SUBSTANCES. L Before the liquid reagents can be applied, solids must be dissolved. The methods of inorganic analysis do not provide against interference by organic compounds; and, in general, it is impossible to conduct inorganic analysis in material containing organic bodies. It is, therefore, first necessary to ascertain if the unknown substance contains organic matter. This has been ascertained in the preliminary examination by the blackening or charging of the substance when heated in the closed tube. A burnt odor is also excellent evidence of the presence of organic matter. If organic matter is found to be present, it may be removed as follows: 1st, by combustion at a red or white heat, with or without oxidizing reagents; 2d (in part), by oxidation with potassium chlorate and hydrochloric 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. The residue is treated according to 301. 286. Oxidation with potassium chlorate and hydrochloric acid on the icater-bath 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. 287. The action of sulphuric irith 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 292. PRELIMINARY EXAMINATION OF SOLIDS. 375 is now placed in the tubulure, and nitric acid added in small portions, gntdu- ally raising the heat, for about half an hour so as to expel the chlorine, and n t vaporize chlorides. The material is now transferred to a platinum dish and heated until the sulphuric acid begins to vaporize. 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 nearly all of the sulphuric acid, using the lowest possible heat at the close. Cool and dilute with a few c.c. of water. If a residue remains undissolved boil for a few minutes, cool, filter and wash with dilute sulphuric acid. The residue will contain (a) undecomposed mineral matter such as silicates, (6) all the lead, strontium and barium as sul- phates, (c) some of the calcium, bismuth, antimony and tin and (d) practically all of the chromium as the insoluble anhydrous sulphate. 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. H\i mall/six, 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. Thins 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, 172). PKELIMINAEY EXAMINATION OF 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 1 flame. 292. 77* r flame produced &// ihc Wow-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 * Jf the unknown is a solution the solid may be obtained by evaporation on the water bath. 376 PRELIMINARY EXAMINATION 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 1 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 1 1n- 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 Tuhe 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 the side of the coal; a small fragment 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 inclined at an angle of about twenty-five degrees and the flame made to play horizontally upon the assay. The substance is first heated in the oxidizing flame and then in the reducing flame. Any metallic beads which form are allowed to cool and carefully examined for malleability, fusibility and color. 300 PRELIMINARY EXAMINATION OF SOLIDS. 377 Any escaping gases are to be tested for their odor; the changes 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. Some 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 1 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 \vhich 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 exces's, 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 378 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 1 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 otjier reductions effected, as directed in 77, 7, and elsewhere. CONVERSION OF SOLIDS INTO LIQUIDS. 301. Before the fluid reagents can be applied, solids must be dissolved. To obtain a complete solution, the following steps must be observed: First. The solid remaining after removal of organic matter, or in the absence of organic, the entire solid, is treated as follows unless it is a metal or alloy (see 303). 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 dissolved. If a solu- tion 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 w r ashed, 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, HNO 3 should be added first, for HC1 would form an insoluble chloride. If the substance contain tin (especially an alloy of tin) HNO, would form insoluble metastannic acid, H 10 Sn 5 15 , in which case HC1 should be used first. Fourth. Should a residue remain it is to be digested with nitrohydro chloric 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 with 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- r-ited by filtration, should be preserved separately, as above directed, and analyzed separately. The separate results, in many cases, indicate the original combination of each metal. 303. TREATMENT OF A METAL OR AN ALLOY. 379 TREATMENT OF A METAL OR AN ALLOY. 303. As most metals are converted into soluble compounds by nitric acid, this is the best acid to use for the solution of metals and alloys. Of the common metals, all are oxidized and all, except tin, arsenic and antimony are converted into nitrates which are soluble in water. Arsenic is converted by strong nitric acid into arsenic acid which is readily soluble in water. Antimony is converted into the pentoxide which is insoluble in water or nitric acid and tin is converted into metastannic acid which is also insoluble in water and nitric acid. When an alloy containing antimony and, especially, tin is treated with nitric acid, considerable portions of the other metals remain in the insoluble residue with the tin and antimony. The action of nitric acid on these metals may be represented as follows : 6Sb + 10HN0 3 = 3Sb 2 O 5 + 10NO + 5H 2 O . 15Sn + 2OHNO 3 + 5H 2 O = 3Hi Sn 5 Oi 5 + 20NO . 3Pb + 8HN0 3 = 3Pb(N0 3 ) 2 + 2ND + 4H 2 O . Bi + 4HNO 3 = Bi(NO 3 ) 3 + NO + 2H 2 O . Gold and platinum are not attacked by nitric acid. Of the rare elements, titan- ium and tungsten are converted into insoluble hydroxides like tin, tungsten being converted into the yellow insoluble tungstic acid (H 2 WO.i) . Alloys may also con- tain such insoluble constituents as ferro-chrome or ferrosilicon which separate as a black residue. Carbon, as well as gold and platinum, also remains as a black insoluble residue. Method of Procedure. About one gram of the metal or alloy in the form of sawings or drillings is placed in a beaker. Add 10 cc. concentrated nitric acid and 5 cc. water. Cover the beaker with a watch crystal and warm on the water bath until action ceases. If any of the alloy remains unacted upon add a little more nitric acid and water and continue the heating. When the action is complete evaporate to dryness on the waterbath. Add 5 cc. concentrated HNO 3 and 20 cc. water, digest on the water bath, filter, and wash the residue. The solution is tested for the metals of all the groups. The residue is digested in strong hydrochloric acid and if necessary in aqua regia. If a residue still remains it may be silicic, titanic or tungstic acid. Transfer to a latinum crucible and add a few cc. of hydrofluoric acid and warm under a hood, ilicic acid is volatilized. Add a few drops of H 2 SO 4 and warm until the hydro- fluoric acid is volatilized. Titanium and tungsten dissolve on diluting with water. If a considerable residue still remains, it is probably metastannic acid. Filter this and wash and transfer the paper with the precipitate to a porcelain crucible, sup- ported on a pipestem triangle, and burn the paper. Add six times the weight of the precipitate of a mixture of equal parts of sulphur and dry sodium carbonate. Place the lid on the crucible and heat with a small flame of the Bunsen burner until the sulphur melts. Continue the heating until the blue flame of the burning sulphur which escapes around the lid of the crucible is almost gone. Allow the crucible to cool with the lid on. Place the crucible with its contents in a beaker and dissolve the fused mass in warm water. The tin will be in the solution as sodium thiostannate. Any insoluble sulphides of copper, lead, etc., are filtered off and the solution acidified with hydrochloric acid. The precipitated sulphides may be examined for tin and antimony or added to the remainder of the sulphides of the sub-group. If considerable tin is present in the alloy, it may be more advantageous to dis- solve it by the following procedure. One gram of the alloy which has been cut into small shavings with a clean knife or otherwise powdered is placed in a beaker and 15 cc. concentrated hydrochloric acid added. The beaker is warmed on the water bath, one drop of concentrated nitric acid is added and the heating con- tinued. When the action again ceases, another drop of nitric acid is added. This is repeated several times until the alloy is completely dissolved. By this process all metals in the alloy are converted into chlorides. Some of these chlorides, espe- cially lead chloride, are insoluble in hydrochloric acid and remain as a white crys- talline deposit. The tin is converted into stannic chloride and remains in solution 380 TREATMENT OF A METAL OR AN ALLOY. 303, B. unless too much nitric acid has been added, which converts the tin into the insoluble metastannic acid. The solution is diluted with an equal volume of water and the insoluble chlorides of the first group filtered off and analyzed in the usual manner. The remaining metals are tested for in the filtrate, as usual. This solution may be turbid, espe- cially when cold, due to the action of water on SbCl 3 . H 2 S is passed without filter- ing as the precipitate is readily converted into the sulphide. 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-brow r n or black powder or mass, insoluble in the hydrochloric acid. If such a residue exists, decant tcJiile Jwt, 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 5 ; but without filtering proceed with the separation and detection of the tin and antimony by the usual process. The Dark-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 bro\vnish-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. * Traces may sometimes be dissolved. 309. SEPARATION OF ACIDS FROM BASES. 381 SEPARATION OF THE ACIDS FROM THE BASES. 304. The preliminary examination of the solid material in the dry way will give indications drawing 1 attention to certain acids. Solutions can be evapo- rated to obtain a residue for this examination. Thus, detonation (not the decrepitation caused by water in crystals) indicates chlorates, nitrates, bro- mates, iodates. Explosion or deflagration will occur if these, or other oxygen- fiiniishing salts as permanganates, chromates are in mixture with easily combustible matter (273, 7). Hypophosphites, heated alone, deflagrate in- tensely. A brownish-yellow capo-r indicates nitrates or nitrites (241, 7); a grcc ~ Q *s B E PC ^ 03 ,0 i 1 N 3 S ^ S .2 .s e . ^ a g OOO r-i 1 H * S !! 02 > KT (U ^ LO. PRELIMINARY EXAMINATION OE SOLIDS | 4s? 1 if 1 OD | 1 I* '1 ie substance suffers no change: Absence of volatile bodies (including combined w r ater), those which change color on heating. ie substance changes color: Organic compounds blacken from separation of carbon, Cu and Co salts blacken at high heat. sS ^ s ^e S ^ to ^ ^ ^ r^ o d - 7^ iilij tiij* *S to ** cc - ^ *! .= 2 ^ ^ 11 to "O r2 ^ S Sls| ^ 2 r< "S 2 ^S ^ 1 1 3 1 1 1 11 &tsl y isii if. , r-J .^ d H J- ^ 'S G -* O "o r }. ^ rt ? ^^ a |?s 1| M W ^ g ro ^ O ,Q *> g ft -g 5 | 'g 'g 3 rt ^ ^ o o o o o d - a & ~y v & f< N (Xi pq FH o OQ e substance fuses: Most alkali salts and numerous other salts. Many s 'cc "o 02 br; p "g O 1' fl 1 fi .2 1 15 02 te substance sublimes, partially or wholly: H,O of crystallization, combination, or absorption. Sublimate condensing in cold part of the tube. Hg (58, 7), flrray, easily rubbed to globules. HgCL first melts, then forms white crystalline sublimat HgCl , without melting, forms a sublimate, yellow while HgS , a black sublimate, turning red on trituration. As , steel-gray sublimate; garlic odor. AsoO 3 sublimes in white octahedral crystals, does not i As 2 S 3 , sublimate nearly black while hot, reddish-yelloir irl Sb 2 S 3 fuses yellow; forms white, amorphous sublimates. CO ? gUj FS g g^ GOO SXi *H d CO * o > >> ^ te; 1 ^ cij a) d W . |.sf I "E rrt ,D CO? ass 'd 03 M S > M +j ^j cJ a 0) fit tb o H H 310. PRELIMINARY EXAMINATION OF SOLID8. 383 How when solidified. )r peroxide. A small cad-paper. Recognizec w p c3 O ^ 3 1 nd alkaline reaction on poisonous. vapor. modified by oxidation litmus-paper. o 3 e-radish (113). i 'H 9 *v <-~ fl Ctf ^S cu o 'O M "o "S NH 4 salts, those not decomposing, white sublimate (207, 7). FeCl 3 slowly sublimes as a reddish-yellow stain (126, 7). S , free or by reduction of sulphide, gives reddish-brown drops, H 2 C 2 4 , a heavy white vapor and crystalline sublimate. I , a violet vapor and blue-black sublimate. 3 substance evolves d gas or vapor: indicates the presence of a nitrate, chlorate, bromate, iodai piece of coal placed upon the assay glows upon being heated. H 2 S , from hydrated sulphides and some sulphites, blackens by its odor. S0 2 , from sulphites, thiosulphates, certain sulphates, etc. Ee bleaching effect. NH 3 , from its compounds which decompose, characteristic odor, litmus. CN , recognized by characteristic odor and violet flame. Intenst Oxides of Nitrogen, from nitrates or nitrites, reddish-brown, acr Jj o Bj QQ s C3 O> 1 g i % o in of the changes stated above as occurring in operation I. ar Oxides are obtained from metals, except from Ag , Au and J S and sulphides yield S0 2 . Eecognized by its odor and action As yields a sublimate of As,O 3 . Garlic odor. Sb yields a sublimate (white), of Sb 2 3 and Sb 2 5 . Bi , a sublimate, dark-brown while hot, lemon-yellow when col Te, gray sublimate of tellurous anhydride (Te0 2 ). Se and selenides evolve SeO 2 , odor resembling that of rotten he Hg , sublimate of metallic mercury. e substance decrepitates: Crystals as NaCl. (If finely pulverized, the decrepitation is a e substance deflagrates: Nitrates, Chlorates, lodates, Hypophosphites, Permanganates, g i -. 03 H 5* g V wj U i* CJ -9 i i $ ,3 5 tr* pq o M 3* ** & o +j ^ ^ O s /I is "^ 5* -C; ti | *& H O ncandescent: ue is alkaline white hot), n C ^ 15 S & r 1 | g 1 o rs: ^ 2" 'c q ^ ^ S " es fi =C _d > "w ^o ?3 1^ P^ aj ^ .2 |S f 3 ^ J obtained: S V CO . OT CS w fi 1^ *> ^ "s -i e * ? ^> o ^j S ^ -9 S % GO *^ ^^ ^> ^3 ^ ^ ^ 1 w .o" stance forms an incrusta lemon-yelloiD while hot jlatile with bluish flam r5 i_ S ^0 1. | 5 L. O M ^ rt rz j a) C r^ O O ^ v P ^ ^ ill , faint yellow while hot, \ llic bead is formed: Pb , . ss or incrustation is coloi ) , yellowish-green. ) , bluish-green. 5 , (Zi>% dark-green. O 3 , Si0 2 , phosphates, , fle fill-color or pink. ) , brick-red. ' , CaO , firr2/. c grains are obtained: Sb , brittle. , Cu , Sn , Au , malleable isiblc magnetic powder is i s ao g i | dq" g" o" " 8310. PRELIMINARY EXAMINATION OF SOLIDS. 385 A 'c reddish-yellow. s, P o ^H ,fl OJ ^ OJ r; V f^ o * CO O-i fl o fit o fl V 'o rt o ith glycerine and r a 'fl 3 'o fl .- ^ rt .ti ^ V 2 s d a to +-> rfl j2 02 "*i M a; fl S -M o3 ^ *^ ^ 60S >i o o b fl ,fl 02 d P rt ^ flU r^ P " fl v S c$ "> ^ ^ ^a ^ a s , o" .2 +> ^^ g p 4-* w S [o * M O CO -* > * B Cb *r^ o ed o "^ ^ S ^J *r! TJ nu K; o 13 . a Is fi o "o ^> fi ? 1 o p ^3 ^ S CO e CO ceo -^S 02 (Bor Wi iu phate. H a zi 413 * a ^ N fit c S d? 85 a rC d3 to : w g gs to ^ to in i bB , ^ ed o M O R. s GROUPING OF THE METALS. 387 388 TABLE FOR THE SEPARATION OF THE METAL*. Pb Hg' Ag As'" As* Sb* Sb'" Sniv Sn" AiT" p t iv Hg" | Cu' 1 Cu' 2 Cd | A ' CO r*T -3 Fe' | Fe/ fe Co Ni Mn" Zn Ba Sr Ca Mg K Na NH 4 313. TABLE FOR REVIEW OF PbCI 2 HgCI AgCI HgCI AgCI f H 2 S0 4 = PbSO* White. J H 2 S = PbS Black. I K 2 Cr0 4 = PbCr0 4 Yellow. I Kl = Pbl 2 Yellow. "I o fNH 2 HgCI + Hg Black. j S [(NH 3 ) 3 (AgCI) 2 \ Add HN0 3 -|AgCI White. ' As 2 S s ^ (NH 4 ) 4 As 2 S 5 2 As 2 S s 'o 2 f H 3 As0 4 i As S (NHJ,AsS A cd As^S K eg CO "trt HOgOg V 1*1147 3 MO*J 4 OL O2V5 S Sb 2 S 6 (NH 4 ) 3 SbS 4 1 1 Sb 2 S s 5 SbCI 5 "^ ft Sb 2 S 3 c5 3 5 a. SnS 2 2 A (NH 4 ) 2 SnS 3 8 i= SnS 2 5 SnCI 4 c^J SI ^ j ^i rf x ^p SnS ^ O^* a. Au 2 S 3 S Solution. 1 1 Au 2 S 3 i AuCI 3 2! * f PtS 2 'S Solution. S 5 PtS 2 PtCI 4 IB^. s 2 So p 1 MoS 3 1 (NH 4 ) 2 MoS 4 <^ MoS 3 fig MoCl e Hi = / \ / / ^N \ ^ HgS "3 HgS HgS Dissolve in nitrohydrochlori I ~ PbS >, PbS |j Pb(N0 3 ) 2 ^ j PbS0 4 Confirm b. Bi 2 S 3 < Bi 2 s 3 a* Bi(N0 3 ) 3 c|jf Bi(N0 3 ) 3 S 1 Cu 2 S | Cu 2 S 1 Cu(NO,) 2 ?2x] Cu(N0 3 ) 2 *"| xS / ^ 1 d ^ CuS CdS r CuS CdS Cd(N0 3 ) 2 Sg2 ^^'-S o o o 3 Blood red. I Test original solution (acid) with KCNS for Fe'" and with K 3 FefCN) e for Fe"-j Fe 3 [Fe(CN) 6 ] 2 B]ue. fa. l_l> r a. 1 b Test with borax bead. Blue bead. Add NaHC0 3 and H 2 2 , Green solution. Test with borax bead. Brown bead. Heat with J M! ,ni^ ' add Kl. Br and NaOH i N|(OH )3 f Free | i n C3 2 } Or add nitroso-/3 naphthol. ' Co Red precipitate. Ni - Solution. \ Test with 1 AddNH 4 OH ' filter and borax add H bead. 9 S -j NiS. Black l) 2 j Boil with Pb0 2 and HN0 3 [HMn0 4 Purple. ) a -< Add H 2 S -ZnS White. Dissolve in HCI and add H 2 S0 4 ]BaS0 4 White. V a n \ g i .. 1. Add CaS0 4 , set aside ten minutes -} SrS0 4 , White. |0a)a SrC0 3 CO Sr(C 2 H 3 2 i 2 " o,fl Moisten SrS0 4 with HCI and apply flame test. ** |5 ^0.2 U i-* 3 2. Add K 2 S0 4 , boil, set aside ten minutes. I 3 2 U 5 ^ CaC0 3 Ca(C 2 H 3 2 ) 2 j S^ Filter and add CaC white, soluble in HCI. ( Nn 4 ) 2 u 2 U 4 H,PO.,, White. -Apply flame test using cobalt glass. Violet. i After removal of Mg apply flame test, yellow. . FIRST TABLE. }314. fe H H 's t t w_ CO PREL M M w ! H CO 000 a a w oa 00 O) 5 tf-a .2 g "S 13 .S . S S 3J w >, K i ^J TH ^ -. ^ d cs 11: =1^ S ^ M . 2 ^ "3 '6 co ~S8 314, AOTDS. TABLE. 301 tj _P _0 02 co rJ cd CO T ' ' c: pj PU P) 03 C 1 r\ nd 03 r\ C od S Q) -*j ^ O S p o g 'be 'C "o o 03 T? J cj rt "c~ /ti E/) QV UJ a "-i O * FH 03 -u ^ tM C c -( BP ^2 2 Q p p, j PH ;-3 JS H S co Si ooo cyanates PH I ^ S s I norides. 314. ACIDS. FIRST TABLE. 393 'o nj fl a; o3 03 g .-i O r^i O >-s ns O) C3 gg S s 3' .. bo 8 a -S a JH ^j o; cd ^ ^ ,2. * 5 8 8 o >PH - -v ^ GO O c3 ^ a s ^ & M 1. 1 1 1 1 o ^'S ^ > So"- P rd fl a o H 111 a rd 03 (1) CO O fl 5 g .a 1 -4 ' cC f-, fee 3 'o C3 c "G sd ^r X fl '-!-! fcL H ^0 i>T fl r- 1 ~f-> <-/' - ^^ ^ s cc s vx > ^O K- 'S ^ 1 g H N r^ 'o c3 P, o cd g 1 03 o 'i-i 03 w 1 s cfl a> 't-t o s ^ fc ^ ^Pn 3 /' EH p- ^ 2 ^ 3 nd >* r^ " O 02 "d 3 03 a> H ^j 'S fl 3. ci ^ fn t +j be fl ^-> fl IH rr o '? . '> d3 O a fl rG PM OQ rP g G "5 4J P 8 - ,. : .g .. - oa o PH aj C 3 "fl CO CO a> CO . O PH i 0) rO ce I bi) a> 3 O CO 8. r 3 'o _0 t 1 o fl 03 0> ^ 1 i 9 i I 394 ACIDS. FIRST TABLE. >3l4 cq 02 CD 03 S g a ^ r 1 J "3 SCO O 000 JE3 r3 fc CO _ g M o "S ' CO O OQ 4S J5 I H Q CO 2 5.23 CD O "opw CO rn "" 3 H i 8 (~j 33 QJ ^tj 15 S 03 CD OOO ^ . -e CO O 03 ) ^ Q^ CD O3 ^ ^"3 r-H CO g bX) ug oo * r.s. ^ O CO o M o ^ "73 O5 . 1 g s ^'S ^ fc ^ CD _, s| -(J r^ 1 OQ ^ *2 ^ W .S o 1-2 r25 fl o o I! a <- P-i .-s s q 3 *o ^ 03 O ." 3 a r5 r^ B 5 "" 1 5| ,H^^ 1 ^-g > .S O^aj gp^cO^ MH O HH O ^ ^ X M fl W PI C* o r2* ' 7>> s S CD M p < W p p P 1 1 < M rn i S* V ^ s S *^3 1 1 1 ^ O 1 I 1/2 H & g W S 314. ACIDS. FIRST TABLE. CD M rc _r- O QJ CO ^ ^ ^ QO "* CD c3 CQ ^ * bo s g M ^ CD iZ n H CO CD P-t 'Jj Pi -r; j 4 _ tO 8 J " fllll CJ _, O * Sfill^ .S I pq l 3 3 s.a s CO O 2 * o3 O a 2 P r^l P-, I s *H ~tj 'S d co 2 ^ CD S rO g ^ ,0 PH CO S'S^'S'S 5 o o o be I fc* rS p, r> 3 9 ntrated o f-H a *'~i rS ^ fl S . - I-H O B o> S 03 r^ r-P ^ 22 g > ^ j? IT^ CD 2 3f I '2 I 1 o- O H rCn O PH OO <3 O ^ CD r-H CQ o ^ ^ g c^o ' J S * * *." "ll ill! . 1 CD (.S3 3 ^g r^j PL, t^ O C/3 CD o I . . V O rr-5 co rt CD rt ^ ^3 > CD i- co CO S 3i co bJO co -P .S '^ H-J ej, 2 S3 o 03 S3 ' PH a -M O 3 g '-S B " ^0^1 * s a^ j s l^-a fill- I I I 396 ACIDS. FIRST TABLE. 314. bJD O S o C3 C3 P-i -2 g s .-s ^^ - s ^31< 1 .- I a * .B"J ?r P o ca o g S o >H .r 1 03 O ^ p i o . r& & ^ PH . vS ^11 '3 "^ *y ^ c 03 O PH 2 ^^ 'C 'S 2 ^ .-a 2 ^ -PH* g "p nd -g "" 'S * OT rj - ansposed by concentrated liberated. Acidify and bJD bC 3 coo o a ' , CO 8 o coo o s * C03^ of c +H 03 1 1 P fl P rH OJ O ;? tJ PH -g> rf rn H ; g be ^I'S ^ & o |1i ra in y S be ^^ - ^ ic ^ J .S co =+H O be o Is 42 111 CD J* 'S 1 2 -3 be 3 o P=3 ^ r^ PH S^ QJ ^ ' bO p fl .2 !. o s r^J f>^ 02 5f| SB! ^ a, - -g ^8 111 a -q 1 i **& . cS ^* cc < g -S "3 P o h S ~ ^ s r^ ^o f< -I- 3 r* ^ -3^0 .^^J S 13 fl p_i 02 cd O -M I"! a ^ ^ ^ ^ ,CJ OS |>-|- O O .S ' ^3 O ^^ d ^ -M O) S o-S 31 "S m o 02 rj P fl S H3 O g ^ ^ PH p 2 >> 13 ^ : 0) ~M > Cfi i i -M 'S '^ rC 'O ^ S *!< OQ 314. 6 ACID*. FIRST TABLE, 39' 1 09 O rt S 3 sjs > q_, U 03 S o y fe bo a| Q g ^ * .2 , ^ 3 QQ j (-^H *"""* t'Z* (M 03 r- O g^ 1 M al ^-32 O r5 ' P< c5 ^ o o o3 12 ^ o Onj eg ci 03 o S oj'1 fi w^i 5 s -^ P^ 00 O V^ "3 nce o 6 and a lo* 33 I g 3 - ft ^ te* '3 3 1 2 E^"" Sgsl 05 O ^ d 03 ^ o S W no "S ^ o -a ^ w d S a ^ | |1 8 fi "3 3 g O S d Is s a | i 1 rrl O ^ O -U g s a^^ .2 ^ a d ^j T3 O 02 5 s "S S S I 6 8 5 * -2 a 2 ft'S Is S a s-'Sl 3 -o ^"2 P 5 a -a -s o fslll si!?* ._-_. 316. ACIDS. THIRD TABLE. 399 CO bfl bfl Cu u S 'S 03 -*-> be - -a- 1 -Mils II . I l fl " r tift ~-2 iiiill 02 Q> ^ > c >> S r2 "S, . a 8 "3 P. g & - -^ >> s S >, s I- 3 s d III T3 0) ra 1rf2 .2 g S g ii' ^ ft > o '3 'S S 3 P, A 1-; 3 -0 2 Sg 8 |I| lu ' E ^11 c P ? o .Z c 55 O o .25 400 ACIDS. FOURTH TABLE. 317, 1. $317. TABLE FOR IDENTIFICATION AND SEPARATION OF THE COMMONLY OCCURRING ANIONS ( ACIDS).* CO. S0 2 N 2 H 2 S HCN C 2 H 4 0, 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 tests 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 filtration. The filtrate contains all the anions (acids) and the excess of C0 3 " . Acidulation with HN0 3 sets free C0 2 , and Si0 2 is precipitated; identified in the microcosmic salt bead. The filtrate is made ammoniacal. 3. Ca(N0 3 ) 2 solution precipitates: insoluble; H 2 S0 4 liberates HF . soluble, reappearing with NH 3 ; decolors KMn0 4 solution, h Fe" + Fe"" + OH' gives 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 MgNH 4 P0 4 . 4. In the filtrate from 3. Ba(N0 3 ) 2 precipitates: CrO/'(Cr 2 7 ") as BaCr0 4 , yellow, soluble in HC1; the yellow color of the solution becoming green on boiling with alcohol. F as CaF 2 "J insoluble in V in acetic dilute C 2 4 " as CaC 2 4 j acid; HC1 CN' as Ca(CN) 2 v /heated Pruss C 4 H 4 6 " as CaC 4 H 4 6 with K' *3 tassiu 8 In th HAs0 3 " as CaHAs0 3 ~(-j H 2 S pre< HAs0 4 " as CaHAs0 4 HP0 4 " as CaHP0 4 * From Chem. Prakt. Abegg and Herz (1900), Breslau, Page 113 ; reviewed by Fresenius, Z. t 1900, 39, 566. by OH' / coloration, gives with \ Prussian 318, 2. NOTES ON THE DETECTION OF ACIDS. 401 S0 4 " as BaS0 4 \ f unchanged, remains insoluble in HC1. SiF 6 " as BaSiF 6 \ insoluble on 1 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 nitrate from 4. is exactly neutralized with HNO ;! *; Zn(N0 3 ) 2 then precipitates : Fe(CN)/" as Zn 3 [Fe(CN) 6 ] 2 brownish-yellow dissolved ( brown Fe(CN)/'" as Zn 2 Fe(CN) 6 white Fe'" and H* ( blue. G. A few drops of the nitrate from 5. are treated with as little Fe'" as possible : Eed f 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 ) , ni ,. ... v much Cl-water. Br brown coloration, does not disappear with j The second portion is evaporated to dryness with K,Cr 2 7 , fused, and the mass after cooling distilled with concentrated H,S0 4 ; appearance of oily brown drops of Cr0 2 Cl 2 , forming Cr0 4 " with water: Cl' . 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 NO/. 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"; SO/ and NO/; NO/ and CN'; NO/ and S"; NO/ and I'; NO/ and HAs0 3 "; .S" and HAs0 4 " , 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 cooling (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 HCl at this point. 402 NOTES (XV 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 w r ater. 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 CO 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 CaCO s . 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 KC1O 3 , K,Cr 2 O 7 , KMn0 4 , 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 H, 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, SiO 2 . 10. Meta- or pyrophospliates 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 bij 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, 15. NOTES ON THE DETECTION Of ACIDS. 403 the sulphur with the reagent decomposing" the sulphide will cause the oxida- tion of a portion of the sulphur 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. 1,}. \ 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. I'l- 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, Sd. 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 \vater, 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 five 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. Separation by the persulphate method. To ten cc. of the original solution, add slight excess of Na 2 CO 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 O 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 CF 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 solu- tion 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 F 2 SO4 , previously diluted with water, a little more K 2 S 2 O 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 completely removed. If bromine is present, add one-half gram of K 2 S 2 O 8 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 8 and boil five min- utes 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, * In consequence of the relative commercial 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 b found in the iodides of commerce, and occurring 1 from the difficulty of exact s."pa-uti< D in ; !r- ir.anu hictiirc of k>;ii!^ f com Ice! p. 404 NOTES ON THE DETECTION OF ACIDS. 318, 16. cool and add a few drops of silver nitrate; a white, curdy precipitate of silver chloride indicates the presence of Cl. If too much silver nitrate is added, a white crystalline precipitate may be formed, but will dissolve upon dilution and warming. If ClO's is present, the above procedure cannot be followed, for the I' would be oxidized in IO' 3 . In this case it is necessary to precipitate the CF , 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 AgCl , 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 ZnCl 2 , ZnBr 2 , ZnI 2 , according to the above method, starting at the beginning. Even if no heavy metals are present, Na^COs 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 SO 4 , Cl , HNO 2 , HNO 3 , or other oxidizer). In presence of a great excess of Br' , CuSO 4 , KNO 2 , or KgCl is an excellent test tor I 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 1 agent (Note 11). The filtrate is then to be treated by either method above given. If). The separation by ammonium hydroxide, as a solvent of the silver pre- cipitates AgCl , Ag'Br , AgT 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. 8r/). 17. The direct removal of iodides In/ 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). 78. 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), bv remaining in solution during the precipitation by silver nitrate (273, 5). //). Chloric acid is separated from nitric acid after finding that silver nitrate (/Ires 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. 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, 6?>; 152; 223, 6). The solution tested must contain 110 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 AgGl , 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. 8$. M. Dechan's method (269, 8i) consists (/) in boiling the mixture with a solution of 10 grammes of K 2 Cr 2 O 7 , dissolved in 100 cc. of water, which lib- 319, 1. NOTES ON THE DETECTION OF ACIDS. 405 e rates and expels all oi the iodine without disturbing the bromine and chlo- rine 5K 2 Cr 2 T + OKI = Cr 2 3 + 8K 2 Cr0 4 + 3I 2 (2) 8 cc. of a dilute solution of sulphuric acid (consisting- of equal volumes of H^SO^ ;>. (jr. 1.84, and water) are added to 100 cc. of the dichromate solution, and 011 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. 24. In the detection of chlorides in presence of cyanides and thiocyanates by the decomposition of the silver salts with concentrated sulphuric acid (269, 8c), 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. 2~). 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 KMn0 4 and MnO, . In the presence of a fixed alkali the manganate solution is green and does not rapidly change to permanganate. The mang-anates 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. PEINCIPLES. 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). 406 PRINCIPLES. 319, 'Z. 2. All nitrates, chlorates and acetates are soluble, but salts of cuprosum, 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 SnO(OH) 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. 9 Fe 4 O s (OH) 6 , and Cu 3 2 (OH) 2 ]. This precipitate re- dissolves in eight cases, forming, if potassium hydroxide be used . . . K 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 SnO(OH) 2 , with Sb v , SbO(OH),, 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. Com- plex ammonium compounds are formed, such as (NH 3 ) 2 Ag f OH and (NH,) 4 Zn++ (OH) 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. Cuprous chloride is also insoluble. (For action on higher oxides, etc., see 269, QA.) 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, 6.4.) 7. The iodides of lead, silver, mercury and cuprosum are insoluble. Hydriodic acid and soluble iodides precipitate solutions of lead, silver, ir.rrcury and cuprosum. Cupric salts are precipitated as cuprous iodide witli liberation of iodine. Ferric salts are reduced to ferrous, arsenic arid to arsonous acid, Sb v to Sb'" with liberation of iodine. (Bismuth, stannoiis 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. BisnuiUi iodide is black; stunnous, antimonous and arsenous iodides are yellowish reel..* 319, 1.1. PRINCIPLES. 407 8. The sulphates of lead, mercurosum, barium, strontium and calcium arc 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 othr 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 257,6.4). (b) 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. 6. A normal hydroxide with Sn", Al , Cr'" and Fe'"; with Sn IV , SnO(OH), ; with Sb v , SbO(OH), . 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 , HIO a , H 2 C 2 4 , KC 2 H 3 2 , HN0 3 , H 3 P0 4 , H 2 S0 4 , and even by H 2 S , com- pletely from carbonates of first four groups, incompletely from those of the fifth and sixth groups (Nandin and Montholon, C. r. f 1876, 83, 58). 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. 408 PRINCIPLES. 319, 12. 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 jat 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 , Ft , Sn0 2 , Sb 2 3 , As 2 3 , Bi,0 3 , CuO , CdO , Fe 3 4 , Cr 2 3 , A1 2 3 , CoO, 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. &. Alkali hydroxides ignited in air at a white heat are not changed. Other hydroxides evolve H 2 and leave as in (a). c. Alkali carbonates ignited in air at a white heat are not changed. Other carbonates evolve C0 2 and leave as m (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, ft. 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 the chlorides of these metals. 320. EQUATIONS. 409 320. 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 illustrated 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 Chemie. 1. Sb -f HN0 3 2. As, + HN0 3 3. As,0 3 + HN0 3 4. Mn(OH), -f- Pb0 2 + HN0 3 5. MnS0 4 + Pb 3 4 + H 2 S0 4 , dilute 6. MnOo + KN0 3 + K 2 CO 3 , fusion 7. S a + KN0 3 + K 2 C0 3 , fusion 8. MnS + KN0 3 + K,C0 3 , fusion 9. Mn^ + Pb 3 4 -{- HNO, 10. Cr(OH) 3 + KN0 3 + K 2 C0 3 , fusion. 11. Pb 3 (As0 4 ) 2 + 2n 4- H 2 S0 4 , dilute 12. Cu 2 As 2 O T -f Zn + H 2 S0 4 , dilute 13. Pb(NO 3 ) 2 + Al + KOH 14. Cu(NO 3 ) 2 + Al + KOH 15. Bi(N0 3 ) 3 + Al + KOH 16. Hg 10 2 (N0 3 ) 8 + Al + KOH 17. MnS + Mn(N0 3 ) 2 + K,C0 3 , fus. 18. Mn 3 5 + Pb 3 4 + HN0 3 19. Fe + HoS0 4 , con., hot. 20. KI + KI0 3 -f H,S0 4 , dilute 21. MnS0 4 + KMn0 4 + H 2 S0 4 , dilute 22. (NaCl + K 2 CrO 4 -f H 2 S0 4 ), dry, hot 23. KN0 3 + FeS0 4 + H 2 S0 4 , con., cold 24. K 2 Cr 2 0(Cr0 4 ) 3 + HC1 , hot 25. Hg- 8 0(N0 3 ) 6 + 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 s O 4 , white heat 29. NaH,PO 2 , ignition 30. Fe 8 9 (AsO 3 ), + FeS -f HC1 31. FeBr 2 + HN0 3 32. Sn + HN0 3 , hot 33. KOH + Br 2 , hot 34. FeI 2 + H 2 S0 4 , sp. gr. 1.84, hot 35. KBr + KBrO 3 + H 2 S0 4 , dilute 36. FeSO 4 + KMnO 4 + H 2 S0 4 , dilute 37. K 2 Cr 2 0(CrOJ 3 + KOH + Br 2 38. 4Hg 2 0,(N 2 5 ) 3 + Al -f KOH 39. Ag- 3 AsO 3 -f- SnCl 2 -f HC1 , sp. gr. 1.18 40. Co 2 3 , ignition, white heat 41. H 2 S -f HNO S , sp. gr. 1.42, hot 42. Hg- 3 (As0 4 ) 2 + FeS + HC1 43. Fe 8 O u (As0 3 ) 2 + KOH + Cl a 44. FeI 2 + HN0 3 , sp. gr. 1.48, hot 45. Cr 2 (S0 4 ) 3 -f Cr(N0 3 ) 3 + K 2 C0 3 , fusion 46. Pb,(As0 4 ) 2 + Zn + H 2 S0 4 , dilute 47. KOH + C1 2 , cold 48. KBr + KI0 3 + H 2 SO 4 , dilute 49. (Cr 2 OHCl 5 + K 2 Cr0 4 + H 2 SO 4 ), dry, hot 50. Zn 4 3 (N0 3 ) 2 + FeS0 4 + H 2 S0 4 , concentrated, cold 51. Hg 3 (As0 4 ) 2 + SnCL + HC1, sp.gr. 1.18 52. Mn 3 O 6 , ignition 53. Fe 2 2 S0 8 + Zn + H 2 SO 4 , dilute 54. Bi 2 S 3 + HN0 3 , dilute, hot 55. Hg 3 As0 4 + FeS + HC1 56. Cr 2 (OH) 4 S0 4 + KOH + C1 2 57. Fe(H 2 P0 2 ) 2 + HNO 3 58. Cr 2 3 + KC10 3 + K 2 CO 3 , fusion 59. Cu 5 2 (As0 4 ) 2 + Zn + H 2 S0 4 , dil. 60. KOH -f C1 2 , hot 61. MnaOu + KC10 3 + K 2 C0 8 , fusion 62. HIO, + SnCL + HC1 63. Bi 12 13 (N0 3 ) 10 + FeS0 4 + H 2 SO 4 , con., cold 64. CrO 3 , ignition 65. KMnO 4 + H 2 C 2 4 + H 2 S0 4 , dilute 66. FeAs0 4 + SnCL + HC1 , sp. gr. 1.18 67. Fe 3 Cl 8 + FeS + HC1 68. 5CuO.As 2 5 + Fe + HC1 69. HIO 3 -f H 2 C 2 4 , hot 70. (Cr 2 (OH) 5 Cl + K 2 Cr 2 7 + H,S0 4 ), dry, hot 71. Fe(NO 3 ) 2 + FeS0 4 + H 2 SO 4 , con., cold 72. Ag 2 SO 4 + Zn 73. H 2 S0 3 + HN0 3 , sp. gr. 1.42 74. FeAsO 4 + FeS -f HC1 75. Pb(AsO 2 ) a + KOH + CL 76. Fe(NO 3 ), + HNO 3 77. Mn s O 5 + Mn(NO,) 2 + K 2 CO g , fusion 78. Fe 8 O 9 (As0 3 ) 2 -f KOH + Br 2 79. Pb 10 8 (OH) 8 (NO s ) 6 + Al + KOH 410 PROBLEMS IN SYNTHESIS. 321 321. PROBLEMS IN SYNTHESIS. For the sake of a more thorough drill in the principles of oxidation antf 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. 322. TAPLE OF XOLrniLlTJES. 411 322. TABLE OF SOLUBILITIES.* Showing the classes to which the compounds of the commonly occurriiu] elements belong in respect to their solubility in water, hydrochloric acid, nitric acid, or aqua regia. PRELIMINARY REMARKS. 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 borntp 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. * Tho following: table of solubilities, is taken from Fresenius Qualitative Analysis, Well's translation of 16th German edition. 412 TABLE OF SOLUBILITIES. SOLUBILITY I Potassium. Sodium. Ammonium. 1 3 3 1 Strontium. Calcium. Magnesium. Aluminum. Chromium. d d N Manganese. Nickel. Cobalt. Oxide W W W W w W-A A A A&I A a, 7 A A Chromate. W, w w a w-a w-a w a w w a a Sulphate.. W 18 .,, W ^14-2'SO I I W-I W ^13M4 W&I 16 W W W W Phosphate W W 8 w 8 -n a a A u a, 3 a a a a a a Borate .... Wj w w a a a w-a a a a a a a Oxalate... W, W w a a 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 10 W A A A A A A A A Silicate... W W a a a a a-i a a a a a Chloride.. w, 7 W 86 W,,.,. W W W W w W&I W W W W Bromide . . W w W w w w w w w&i w w AV w Iodide W w W W w w w w w w AV AV w Cyanide. . . W w w w-a w w w a A a a-i a-i FerrocyVle 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 W W W w W-A 45 a a a-i A,. A a ]8 a, 9 Nitrate... W W W W W w w w W w w AV W Chlorate . . W w w w W w w w w w w W w Tartrate . . W -'T'1S48 W 7 w. a a. 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 4T w w w w Succinate. w w w w-a w-a w-a w w-a w-a w AV w- Benzoate.. w w W w w w w Salicylate. ir 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 a a a Arsenate.. W W w a a a a a a a a a a 322. TAFLE. TABLE OF SOLUBILITIES. 413 d 1 Ferrous. | I I Mercuroue Mercuric. c a Bismuth. Cadmium. 2 Platinum. Stannous. Stannic. Antimonoi a a A 34 A A A a a a a a&i A 42 Qfcide w a A-I a w-a w a a a a Chromate W 20 V 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 w a w a Oxalate w-a v w a w-a a w w-a w w w Fluoride A a A a a A a a Carbonate a r a a a Silicate W " i I W-I A-I W 28 W W-A 33 W W 85 W 37 . 38 W W 40 W-A 43 Chloride w v. i w-i a-i w w w-a W w w w-a Bromide W i i W-A A A w a W a i w w w-a Iodide a-i I a W a a W w Cyanide i i i a i 1 i Ferrocy'de I - i w-a i Ferricy'de w w i a A w a w-a a w Sulphocy'de A a a 23 A A A 2 a,i a A a 38 a 39 at, a 4 , ^44-45 Sulphide w w W W w a W W W* w w Nitrate w w w w w w w w w w Chlorate w-a W.j. a a w-a a w a w-a a *M Tartrate w vr 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 R w-a a a w-a a w Benzoate w-a w-a w Saiicylate w V w W 26 w-a w W 32 w w w w Acetate w V w w-a w w w w w w Formate a il a a a a A a Arsenlte a B a a a a a a a a Arsenate TABLE OF SOLUBILITIES. 322. (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 ?s 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. 323. REAGENTS. 415 323. 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 grains 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 O 2 (300, 5N), sp. gr. 1.04, 30 per cent acid. Arsenic, H 3 AsO 4 .y H.O (15, y a H 3 As0 4 ~ 5). Flue-silicic, H a SiF. , 247. 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. Hydrosulphuric, H 2 S , saturated aqueous solution, 257, 4. lodic, HIO 3 (15, i/ 2 , HI0 3 -7- 6). Nitric, HN0 3 (315, 5N, sp. gr. 1.165), sp. gr. 1.2, 32 p. c. acid. Nitrohydrochloric, about one part of concentrated HN0 8 to three parts HC1 . Nitrophenic, C 8 H 2 (NO,) a OH (picric acid). Oxalic, HoC 2 O 4 .2HoO , crystals dissolved in 10 parts water. Phosphoric, H 3 PO 4 (16, N/2). Sulphuric, H 2 SO 4 , concentrated, sp. gr. 1.84. Sulphuric, dilute (245, 5N, sp. gr. 1.153), one part acid to five parts water. Sulphurous, H 2 SO 3 , saturated aqueous solution. Tartaric, H 2 C 2 H 4 8 , crystals dissolved in three parts water. Alcohol, C 2 H 6 , sp. gr. 0.815, about 95 p. c. Aluminum Chloride, A1C1 S (22, N/2). Nitrate, Al(NO 3 ) 3 .7y 2 H,O (58, N/2). Sulphate, A1 2 (SO 4 ) 3 .18H 2 O (55, N/2). Ammonium Carbonate, (NH 4 ) 2 CO 8 (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 eig-ht parts water. Hydroxide, NH 4 OH (85NH 8 , 5N, sp. gr. 0.964), sp. gr. 0.96, 10 p. c. NH 3 . Ammonium Molybdate, (NH 4 ) 2 MoO 4 (36MoO, , N/2, 75, 6d), 150 g. salt in one litre of NH 4 OH , pour this into one litre of HN0 3 , sp. gr. 1.2. Ammonium Oxalate, (NH 4 ) 2 C 2 O 4 .i H,.2H,0 (43, N/2). Nitrate, Cu(NO 3 ) 2 .6H 2 O (74, N/2). Sulphate, CuSO 4 .5H 2 O (62, N/2), in 10 parts water. Cuprous Chloride, CuCl (50, N/2, use HC1). Ferric Chloride, FeCl 3 (27, N/2), 20 parts water to one part metal. Nitrate, Fe(NO 3 ) 3 .9H 2 O (67, N/2). Ferrous Sulphate, FeS0 4 .7H 2 O (80, N/2"), use a few drops of H 2 SO 4 . Gold Chloride, HAuCl 4 .3H 2 O , solution in 10 parts water. Hydrogen Peroxide, 3 p. c. solution. Indigo Solution, 6 parts fuming- H 2 SO 4 to one part indigo, pulverize, st'r and cool, allow to stand 48 hours and pour into 20 parts water. Lead Acetate, Pb(C 2 H 3 2 ) 2 .3H 2 O (95, N/2), dissolve in 10 parts of water. Chloride, PbCl, , saturated solution, N/7. Nitrate, Pb(NO 3 ) 2 (83, N/2). Magnesia Mixture: MgSO 4 , 100 g-.; NH 4 C1 , 200 g-.; NH 4 OH , 400 cc.; H,O , 800 cc. One cc. = 0.01 g-. P. Magnesium Chloride, MgCL.6H,O (51, N/2). Nitrate, Mg(N0 3 ),.6H 2 O (64, N/2). Sulphate, MgSO 4 .7H 2 O (62, N/2), in 10 parts of water. Manganous Chloride, MnCl,.4Ho6 (50, N/2). Nitrate, Mn(N6 3 ),.6H,O (72, N/2). Sulphate, MnSO 4 .7H 2 O (69, N/2). Mercuric Chloride, HgCl, (68, N/2), in 16 parts of water. Nitrate, Hg(N0 3 ), (81, N/2). Sulphate, HgSO 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.6H 2 O (59, N/2). Nitrate, Ni(NO 3 ) 2 .6H,O (73, N/2). Sulphate, NiS0 4 .6H 2 O (66, N/2). Palladous Sodium Chloride, Na,>PdCl 4 , in 12 parts water. Potassium Arsenate, K s As0 4 (26, % K 3 As0 4 -4- 5). Arsenite, KAsO 2 (24, %~KAsO a + 3). Bromate, KBr0 3 (14, i//KBr0 3 -4- 6). Bromide, KBr (60, N/2). Carbonate, KoCO 3 (207, 3N). Chlorate, KC10, , the dry salt. Chloride, KC1 (37, N/2). 323. REAGENTS. 417 Potassium Chromate, K 2 CrO 4 (49, N/2), in 10 parts water. Cyanide, KCN (33, N/2), in four parts water. Bichromate, K 1 ,Cr 2 O 7 (38, i/ 2 , K 2 Cr,O 7 ^-4), i n 10 parts water. Ferrocyanid'e, k 4 Fe(CN) c .3H 2 O (53, N/2), 12 parts water. Ferricyanide, K 3 Fe(CN) 6 (55, N/2), in 10 parts water. Hydroxide, KOH (312 [90 p. c. KOH], 5N). lodate, KIO 3 (18, % KIO 3 4- 6). Iodide, KI (83, N/2), dissolve in 20 parts water. Mercuric Iodide, K 2 HgI 4 , Nessler's solution, 207, 6fc. Nitrate, KNO., (50, N/2), the crystallized salt. Nitrite, KNO, , the dry salt. Pyroantimonate, K 2 H 2 Sb 2 O 7 .6H 2 , see 70, 4c. Permanganate, KMnO 4 (16, i/ 2 KMn0 4 -h 5). Thiocyanate, KCNS (49, N/2), in 10 parts water. Hydrogen Sulphate, KHSO 4 , fused salt. Sulphate, K,S0 4 (44, N/2), in 12 parts of water. Platinic Chloride, H,PtCl .GH 2 O , in 10 parts of water. Silve. Nitrate, AgNO 3 (43, N/4), in 20 parts of water. Sulphate, Ag,S0 4 , saturated solution, N/13. Sodiv.ni Acetate, NaC,H 3 O,,.3H 2 O , in 10 parts of water. Carbonate, Na 2 CO ? (159, 3N), one part anhydrous salt or 2.7 parts of the crystals, fra^COg.lOHoO , in 5 parts of water. Chloride, NaCl (29, N/2). Tetraborate, Na 2 B 4 O 7 .10H 2 O, lorax, the crystallized salt. Hydroxide, NaOH (220 [90 p. c. NaOH], 5N), dissolve in 7 parts of water. Hypochlorite NaCIO, 270, 4. Nitrate, NaN0 3 (43, N/2). Phosphate, Na 2 HPO 4 .12H 2 O (GO, N/2), dissolve in 10 parts of water. Phosphomolybdate, 75, Qd. Sulphate, (35, N/2). Sulphide, Na 2 S , one part NaOH saturated with H a S to one part of NaOH , unchanged. Acid Sulphite, the dry salt. Sulphite, Na 2 SO 3 .7H 2 O (63, N/2), in 5 parts of water. Acid Tartrate, NaHC 4 H 4 O , in 10 parts of water Thiosulphate, Na 2 S.,O.,.5H,O , in 40 parts of water. Ststnric Chloride, SnCl 4 (33, N/2). Stanr.ous Chloride, SnCL.2HoO (56, N/2), in 5 parts water strongly acid with HC1. Strontium Chloride, SrCL.GHoO (67, N/2). Nitrate, Sr(NO 3 ) 2 "(53, N/2). Sulphate, SrS0 4 , a saturated aqueous solution. Zinc Chloride, ZnCl 2 (34, N/2). Nitrate, Zn(N0 3 ) 2 .6H,O (74, N/2). Sulphate, ZnS0 4 .7H 2 O (72, N/2). INDEX. Acetates, detection of 258 ignition of 267 with ferric salts 157 Acetic acid 256-258 estimation of 258 formation of 257 glacial 257 occurrence of 257 preparation of . 257 properties of 256 reactions of 257 solubilities of 257 Acids, detection of, notes on 401 displacement of weak by strong ... 185 effect of concentrated sulphuric upon 390 list of 13 precipitated by barium and cal- cium chlorides 398 preparation of 408 separation from bases 381 table of, precipitated by silver nitrate 399 table of separation of 400 Alabandite 177 Alabaster 216 Alkali carbonates, with third and fourth group salts 143 group 227 hydroxides, action on double cyanides 272 hydroxides, detection of in pres- ence of carbonates 270 hydroxides, reactions with 227 Alkalis, on third and fourth group metals 141 Alkali metals 5 Alkaline earth metals 5 earth metals in presence of phos- phates 226 earths, relative solubilities of 210 Alkali sulphides, as reagents. . . .317, 318 action of, on stannic salts 86 action of, on stannous salts 85 Alloys, analysis of 379 with copper , , , , 104 Alluvial sand 91 Alpha iron 154 Aluminum 144-148 acetate 146 compounds, ignition of 147 detection of 147, 166 distinction from chromium 150 estimation of 147 hydroxide, formation and prop- erties 145 hydroxide, solubility in ammo- nium chloride 164 occurrence of 144 oxidation of 148 oxide and hydroxides 145 phosphate, separation of 146, 147 preparation of 144 properties of 144 reduction of 147, 148 salts, reactions of 145 salts, with hydrosulphuric acid. . . . 146 salts, with phenylhydrazin 146 separation of, from Cr and 4th group by basic acetates 145 separation of, from iron by Na 2 S 2 O 3 and Na 2 SO 3 146 separation of, from glucinum 201 solubilities 144 Alums 147 Ammonia, occurrence 235 formation of, from nitric acid 286 preparation of 235 properties of 235 Ammonium 235-239 arsenomolybdate 62, 98 benzoate, in separation of Cu from Cd 107 carbonate, as a reagent 236 carbonate, in separation of As, Sb and Sn 120 chloride, as a reagent 237 chloride, in the third group 164 chloride, with PtCl 4 95 compounds, solubilities of 235 cyanate in formation of urea 279 detection of ,,,,,,,,,,,,, , 238 419 420 INDEX. Ammonium, directions for detection 242 estimation of 238 hydroxide, as a reagent 236 hydroxide, as a distinguishing reagent for the first group 53 hydroxide, detection by mercuric chloride 238 hydroxide, preparation and prop- erties of 235 molybdate, preparation of 98 molybdate, test for phosphates.. . . 311 molybdate, with arsenic acid 67 oxidation of 239 phosphomolybdate 98 picrate, formation of 236 polysulphide, formation of 237 salts, detection by Nessler's re- agent 237 salts, ignition of 238 solution to be tested for 242 sulphate, in separation of stron- tium and calcium 226 sulphide, as a reagent 237 sulphide, formation of 236 sulphide, preparation of 316 sulphide, on iron and zinc groups 191 sulphide, yellow, formation of 115 sulphide, yellow, in separation of cobalt and nickel 190 sulphide, yellow, in cupric salts. . . 118 test for nitric acid 288-289 thioacetate as a substitute for hydrosulphuric acid 316 Analysis of alkali group 242 proximate 14 operations of 13, 20 ultimate 14 Anatase 205 Anglesite 29 Anions, table of separations of 400 Antimonic acid 76 distinction from antimonous 123 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 Antimonous sulphide 74 sulphide, precipitation of 77 Antimony 72-82 acids of 72 compounds, reduction with char- coal 80 detection of, in alloys 379 detection of . 80 detection of traces of 123 distinction from arsenic 78 estimation of 81 in the test for aluminum 106-167 metal with hydrosulphuric acid ... 66 mirror 65 notes on analysis of 123 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 121 separation from arsenic 65 separation from tin by sodiu.n thiosulphate 78 separation from tin 81 solubility of 73 spots 66 sulphide, separation from arse i- ous sulphide 123 sulphide, separation from sta i- nous sulphide 123 with iodine 66 Apatite 216, 297 Aragonite 216 Argentite 46 Argol, purification of 260 Argyrodite 137 Arrhenius 20 Arsenates, distinction from arsei- ites 70, 71 separation from phosphates 402 Arsenic 56-72 acid, precipitation by hydrosul- phuric acid 114 acid, reduction by hydrosi 1- phuric acid and hydriodic aci i. 61 acid, reduction with sulphuro is acid 60 acid, with ammonium molybdato. 67 acid, with molybdates 62 INDEX. 421 Arsenic acid, with nitric acid 66 acid with silver nitrate 67 antidote for 62 compounds, ignition of 69 compounds, with concentrated hydrochloric acid 61 compounds, with magnesium salts 61 compounds, with stannous chlo- ride 89 detection of 70 detection of, in poisoning 68 distinction from antimony 78 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 122 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 65 separation from antimony by peroxide of hydrogen 121 separation from Sb and Sn by use of thiosulphates 60 spots, formation of 64 spots, properties of 66 sulphide, separation of, from Sb 2 S 3 123 sulphides with ammonium car- bonate 120 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 123 Arsenopyrite 57 Arsenous hydride 64 oxide, crystals, identification of . . . 67 sulphide, solubilities of 58 sulphide, with HC1 gas 67 Arsine 64 from alkaline mixtures 64 Arsine reactions with KOH 123 separation from stibine 65 with hydrosulphuric acid 60, 65 Asbestos 299 Asbolite 167 Atomic weights, table of 1 Avogadro's Hypothesis 21 Azoimide (hydronitric acid) 282 Azomide 8, 282 Barite 211 Barium 211-214 carbonate, action on ferric salts. 156-157 carbonate, as a reagent 212 carbonate, as a reagent for third and fourth groups 143 carbonate, as a reagent to precip- itate chromium 150 carbonate, and ferric salts 156-7 carbonate, to separate phos- phates from third, fourth and fifth groups 194 chloride, separation of, from SrCl 2 and CaCl 2 by HC1 212 detection of 214 estimation of 214 hydroxide, formation of 211 iodide, properties 370 occurrence of 211 oxide, preparation of . . 211 peroxide, ignition of 296 peroxide, preparation 211 preparation of 211 properties of 211 salts, separation of sulphites from sulphates 213 salts, spectrum of 213 separation of, from Sr, Ca and Mg by sulphates 213 solubilities of 212 strontium and calcium, separa- tion of by alcohol 226 sulphate, separation 215 Bases, alkali 12 alkaline earth 12 copper, group of 11 definition of 3 fifth group of 12 first group of 11 fourth group of 12 iron group of 12 need for separation from acids . 380, 381 second group of 11 silver group of 11 sixth group of 12 422 INDEX. Bases, third group of 12 tin group of 11 zinc group of 12 Bauxite 144 Beryl 200 Beryllium 200 Beta iron 154 Bismite 100 Bismuth 100-104 blowpipe, reactions of 103 chloride, sublimation of. 103 detection as iodide 103 detection by alkaline stanmte 103 detection by cinchonine 102 detection in alloys 379 detection of 103 detection of traces of 102 dichromate 103 estimation of 103 hydroxide, solubility in glycerol ... 101 iodide, stability toward water 103 nitrate, precipitation with HC1. . . . 101 nitrate, reactions 101 notes on analysis of 130 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, comparisons with CuandCd 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 Bismuthinite 100 Bismutite 100 Bitter spar 313 Black Band 155 Blowpipe, examination of solids. . . . 386 Blue vitriol 105 Bonds, plus and minus 244-245 Borates, green flame by ignition of 253 in analysis 54 reactions of 253 Borax 252 bead, formation of 254 Borax, bead, test for Mn 189 bead, use of 377 Boric acid 252-254 estimation of 254 formation of 252 occurrence of 252 preparation of 252 properties of 252 solubility of 253 Boron 252 Braunite 177 Bromates, detection of 361 estimation of 362 ignition of 361 preparation of 360 solubilities of 361 Bromic acid 360-362 properties of 360 reactions of 361 Bromides, detection of 359 detection in presence of iodides. 403, 404 * estimation of 360 formation of 357 ignition of 359 solubilities of 357 with first group metals 358 Bromine 354-356 detection of 356 estimation of 356 formation of 355 occurrence of 355 preparation of 355 properties of 354 reactions with 355 solubilities of 355 Brookite 205 Brown ring, test for nitric acid 288 Brucine, reactions with nitric acid . . 290 Brucite 220 Cacodyl oxide, test for acetates .... 258 Cadmium 110-112 detection of 112 estimation of 112 hydroxide 110 notes on analysis of 131 occurrence of 110 oxide 110 preparation of 110 properties of 110 reactions of, comparison with Bi and Cu 112 salts, absorption by porous sub- stances, separation from Cu. ... 112 INDEX, 423 Cadnuam salts, fused with K 2 S 112 salts, with alkaline tartrates, sepa- ration from Cu Ill salts, with alkalis Ill salts, with ammonia Ill salts, with barium carbonate Ill salts, with pyrophosphates, sepa- ration from Cu Ill salts, reactions with Na_>S 2 O ; , separation from Cu 112 salts, reduction of by metals 112 salts, reduction of by ignition 112 separation from Cu by H 2 S in presence of KCN 107 separation from Cu by KCNS Ill separation from Cu by glycerol ... 105 separation from Cu by Na^S^O ; andNa 2 SO 3 112 solubilities of 110 Caesium 239-240 Calamine 183 Calaverite 138 Calcium 216-220 carbonate in spring water 217 carbonate, solubility of 224 detection of 219 detection of by spectrum 219 estimation of 219 group. . 209 group, directions for analysis of ... 224 hydroxide, formation and prop- erties 217 hydroxide, formation by Na2S. ... 219 hydroxide, to detect CO 2 218 oxide, formation and properties ... 216 occurrence of 216 peroxide 217 preparation of 216 properties of 216 salts, separation of oxalic from phosphoric acid by 218 salts with Na 2 S 219 separation from Ba and Sr by amyl alcohol 217 separation from Ba and Sr by (NH 4 ) 2 SO 4 217 solubilities of 217 sulphate, separation from stron- tium sulphate 215 sulphate, solubility in ammo- nium sulphate 226 sulphate, to detect strontium 219 Calomel 37 Carbon 254-256 amorphous 255 Carbon, detection of 256 dioxide 267-271 dioxide, absorption by Ca(OH) 2 . . . 269 dioxide, detection in sodium car- bonate 270 dioxide, detection by calcium hy- droxide 218 dioxide, distinction from H 2 S, SO 2 , N 2 O 3 , etc 269 dioxide, formation of 267 dioxide, occurrence of 267 dioxide, properties of 267 monoxide 262-263 preparation of 255 properties of 254 reactions of 255 reduction by ignition with 256 relations of 10 solubilities of 255 Carbonates, acid, decomposition of 236 decomposition of, by acids 270 detection of 270 detection of traces 402 estimation of 271 ignition of 270 occurrence of 267 preparation of 267 reactions with 268 Carbonic oxide, formic anhydride. . 262 Carnallite 228 Cassiterite 82 Cassius' purple 93 Castor & Pollux 239 Celestite 214, 331 Cement 299 Cementite 154 Cerargyrite 46 Cerite 198, 199, 202 Cerium 198 Cerussite 29 Chalcocite 104 Chalcopyrite 45, 104 Chalk 216 Chamber process for sulphuric acid manufacture 331 Chili saltpeter 233, 363, 369 occurrence of 285 Chloric acid 350-353 formation of 350 preparation of 351 properties of 350 separation of, from nitric acid .... 403 Chlorates, detection of 353 distinction from nitrates 404 estimation of . 353 424 INDEX. Chlorates, formation from chlorine. 340 ignition of 352 oxidation by ignition of 352 preparation of 351 reactions with 351 solubilities of 351 Chlorides, detection of 151 detection of, in presence of bro- mides 346, 347, 403 detection of, in presence of cy- anides or thiocyanates 346, 405 formation of 341-342 ignition of 345 Chloride of lime, formation of 348 estimation of, by H 2 O 2 296 Chlorine 337-341 action on metals 339 as an oxidizer 339 detection of 341 estimation of 341 formation of 338 occurrence of 338 peroxide, formation and proper- ties 350 properties of 337 solubilities of 338 Chlorochromic test for chlorides . . . 346 anhydride 151 Chlorous acid, formation and de- tection 349 properties of 349 Chromates 151, 152 in test for HC1 151 reduction of, by hydrochloric acid 151 reduction of by H 2 S 150 use in separation of barium 213 with antimonous salts 78 with As'" 151 with ferrous salts 161 Chrome-ironstone 148 Chromic acid, detection of 152 formation of 151 identification of 405 Chromite 148 Chromium 148-153 distinction from aluminum 150 estimation of 152 hydroxide, solubility in ammonium hydroxide 165 and manganese in third group separation 166 metal, solubility of 149 occurrence of 148 oxidation of . 152 Chromium oxides and hydroxides . . 148 oxide, solubilities of 149 preparation of 148 properties of 148 reduction of 152 salts, solubilities of 149 salts, reaction of 149 separation from Al and Fe by H 2 O 2 152 separation from fourth group 150 separation from Fe by Na-jSgOs and Na 2 SO 3 146-147 Chromous salts 149 Cinchonine as a test for bismuth. . . 102 Citric acid 258-259 detection of oxalic acid in 259 distinction from tartaric 259 properties and reactions 259 Cinnabar 37, 313 Clay iron stone 155 Coal, anthracite 255 Cobalt 167-172 bead test 172 cobalticyanide separation from nickel 169 detection of 172 detection of by means of am- monium thiocyanate 170 detection of in presence of Ni by H 2 O 2 190 estimation of 172 hydroxide 165 metal, solubilities of 168 nitrate, effect of ignition with 377 occurrence of 167 oxidation of 173 oxides and hydroxides 167 phosphate, a distinction from Ni 171 preparation of 167 properties of 167 reduction of 173 salts, solubilities of 168 salts, with alkalis 168 salts, with barium carbonate 169 separation from nickel by ether ... 168 separation from nickel by KNO 2 170 separation from nickel by KMnO 4 . 172 separation from nickel by NH 4 CNS 172 separation from nickel by ni- troso-/3-naphthol 170, 190 Cobaltite 57, 167 Colloidal sulphides of fourth group 189 Color, flame tests 377 Columbite 198,203 INDEX. 425 Columbium 198-199 distinction from Ti 207 properties and reactions of . . . .198-199 separation from tantalum 203 Contact process for sulphuric acid manufacture 331 Copper 104-110 acetoarsenite 109 analysis of, notes 128 arsonite 109 compounds with cyanogen 107 detection of 109 detection of, in alloys 379 detection of traces of, with H 2 S . . . 108 detection of, with HBr 108 electrical conductors 104 estimation of 109 ferrocyanide, formation of 107 group, metals of 56, 100 hydroxide of 104 occurrence of 104 oxides of 104 precipitation of, by iron wire 110 preparation of 104 properties of 104 pyrites 204, 313 reactions of, comparison with Bi andCd 112 reduction by ignition 112 reduction of, by KCNS 107 salts, detection by potassium xanthate 107 salts, reaction with zinc-plati- num couple 110 salts, reduction of, with H 3 PO 2 .... 107 salts, separation of, from Cd by Na 4 P 2 O 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 Na 2 S 2 O 3 andNa^COs 112 separation of, from Cd by H 2 S in presence of KCN 107 separation from Cd by nitroso- /3-naphthol 107 separation from Cd by ammo- nium benzoate 107 separation from Pd 106 traces, loss of 118 traces of, with K 4 Fe(CN) 6 107 Corundum 144 Cream of tartar, formation of 260 FAGE Crocoite 29, 148 Crookesite 204 Cryolite 144, 297 Crythrite 57 Cuprammonium salts 106 Cupric hydroxide in NH 4 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 Na^Oa 108 salts, reduced by SO 2 109 sulphide, colloidal 108 sulphide, formation of 108 sulphide, separation from Cd by H 2 SO 4 108 sulphide, solubility in (NH 4 ) 2 S X .... 108 sulphide, solubility in KCN 108 sulphide, with K 2 S 118 sulphide, with (NH 4 ) 2 Sz 115 Cuprite 104 Cuprous iodide 108 oxide, formation of, by glucose. . . . 105 salts, oxidation of, by As 2 O 3 110 salts, separation, from Cd by S. . . 107 salts, with metallic sulphides 108 sulphide, formation by NaoS 2 O 3 . . . 108 thiocyanate, formation of 107 Cyanates, detection of, in presence of cyanides 279 Cyanic acid 279 Cyanide of silver, distinction from chloride 273 Cyanides, detection as thiocyanate. . 275 double, dissociated by acids 272 double, not dissociated by acids. . . 272 estimation of 275 guaiacum test 275 ignition of 274 preparation of 272 reactions with 272 simple, with mineral acids 273 solubility of 272 transposition by acids 275 Cyanogen properties and reac- tions . 271 Danger and Flandin, detection of arsenic 69 Daubreelite 148 Decomposition of organic mate- rial 374-375 426 INDEX. Dialysis, separation from organic material by 375 Diamond 254 Diaspore 144 Didymium 199 Didymium Earths 202 Dimethylaniline, test for nitric acid 290 Dimethylglyoxime, test for nickel 176 Diphenylamine, test for nitric acid 290 Dissociation, electrolytic 20 Dithionic acid, formation and properties 324 Dolomite 220, 267 Dragendorff's reagent 102 Electrolytic dissociation 20 Enargite 57 Epsom salts 220, 313, 331 Equations illustrating oxidation and reduction 409 rule for balancing 246 Erbium 200 Ethyl acetate, odor of 257 Euxenite 137, 198, 202, 207 Everett's salt 157 fatty material, removal of 375 Feldspar 144 Ferric acetate, formation of 257 acetate, separation of from chro- mium 157 basic nitrate, separation from aluminum 161 and ferrous compounds, distinc- tion 165 hydroxide, antidote for arsenic .... 62 phosphate, formation of 159 salts, detection of traces 158 salts, with acetates 157 salts, with BaCOa 156 salts, with HI and iodides 161 salts, with H 3 PO 2 159 salts, with H 2 S 150 salts, with KCNS 158 salts, with K 3 Fe(CN) 6 158 salts, with K 4 Fe(CN) 6 157-158 salts, with stannous chloride 89 salts, separation from ferrous sul- phate 156 Ferric thiocyanate, distinction from ferric acetate 157 hindrance to reactions of 158 Ferricyanides, in distinction be- tween Co and Ni. . . .170 Ferricyanides, reactions of 278 Ferrite 154 Ferrocyanides, detection of 277 detection and estimation 279 reactions of 276-277 Ferrotellurite 138 Ferrous iron, detection of, in ferric salts 158 in the third group 164 in the third group with phos- phates 194 salts, traces in ferric salts 158 salts, with chromates 161 salts, with HNO 3 159 salts, with KCN 157 Ferrous salts, with K 3 Fe(CN) 6 158 salts, with K 4 Fe(CN) 6 157 sulphate, with gold salts 93 First group metals, table of 52 Fixed alkalies 227 alkali hydroxides on stibine 79 alkalis with salts of tin 84 Flame, blowpipe, production of 376 or color tests 385 oxidizing and reducing 375-376 reactions with copper salts . . 109 Flint 299 Fluor-spar 216, 297 Fluorides, solubilities of 298 Fluorine 297-298 Fluosilicates, formation of 298-299 Fluosilicic acid 298-299 in detection of potassium 231 in separation of Ba, Sr and Ca . . . 213 Formates, formation from cyan- ides 274 Fourth group, directions for anal- ysis 189 reagents 142 sulphides colloidal 189 table of 183 Fresenius and Babo, detection of arsenic 68 Froehde's reagent 99 Fulminating gold 92 Gadolinite 202, 207 Galena 29, 45, 313 Gallium (eka-aluminium) 200 Gamma iron 154 Garnierite 173 Gas-laws 21 Gases, absorption of by palladium ... 133 Germanium, properties and reac- tions. . . 137 INDEX. 427 Germanium sulphide 118 Glass, etching by hydrofluoric acid. 313 Glauber's salts 313, 331 Glucinum (Beryllium) 200-201 distinction from yttrium 207 separation from aluminum 201 separation from cerium 198 Glucose, in formation of cuprous oxide 105 Gold 91-93 detection in alloys 379, 380 detection of 93 distinction from Pd 133 estimation of 93 fulminating 92 notes on analysis 125 occurrence, properties, etc 91 reduction by ferrous sulphate 93 reduction with oxalic acid 92 salts with alkalis 92 salts with stannous chloride 89 separation from Ir 135 Graphite 254 Greenockite 101, 110 Gypsum 216, 219, 313, 331 Haematite 155 Halogens 9 as oxidizers 340 compounds, comparative table of 373 hydracids as reducers 340 separation of by persulphate method 347, 403 Hausmannite 177 Heat, upon substances in closed tubes 376,382 upon substances in open tubes. 376, 383 Heavy spar 313, 331 Hydriodic acid 365-368 action on antimonic salts 78 action on arsenic salts 61 action on ferric salts 161 as a reducer 366, 367 formation of 365 Hydrobromic acid 356-360 detection of Cu with 108 formation of 357 occurrence of 357 preparation of 357 properties of 357 reactions of 357 Hydrochloric acid 341-348 action on Sb-jSn 77 action on bismuth nitrate 102 Hydrochloric arid, effect of excess in second group 114 formation of 341 formation from MgCl 2 222 gas on arsenic sulphide 67 occurrence of 341 preparation of 342 properties of 341 reactions with 343 solubilities of 342 Hydrocyanic acid 271-275 formation of 272 occurrence of 272 on PbO 2 274 preparation of 272 properties of 271 solubilities of 272 Hydrof erricyanic acid 277-279 Hydroferrocyanic acid 275-277 separation from hydroferri- cyanic acid 277 Hydrofluoric acid 298 Hydrofluosilicic acid (fluosilicic acid) 298 Hydrogen 250, 251 absorption by Pd sponge 13& detection of 251 estimation of 251 formation of 250 nascent 251 occluded 251 occurrence of 250 preparation of 250 properties of 250 reactions with 250 reducing action of, with ignition. . . 251 solubilities of 250 peroxide, detection of 296 peroxide, estimation of 297 peroxide, estimation of bismuth with 104 peroxide, formation of 295 peroxide, occurrence of 295 peroxide, on sulphides of arsenic and antimony 121 peroxide, preparation of 295 peroxide, properties of 294 peroxide, reactions with 295 peroxide, reagent to separate Co from Ni 190 peroxide, separation from ozone. . . 243 peroxide, separation of Al, Fe and Cr with 152 peroxide, solubilities of 295 peroxide, with arsenic 71 428 INDEX. Hydrosulphuric acid 315-320 action on copper salts 108 action on ferric salts 160 dissociation of 114 formation of 316 gas as a reagent 113 gas on antimony 67 gas on arsenic 67 occurrence of 316 on aluminum salts 146 on stannic salts 86 on stannous salts 85 on third and fourth group salts . 142, 164 preparation of 316 properties of 315 uses as a reagent 317 with arsenic acid 114 with oxidizing agents 114 Hydrosulphurous acid 323, 324 Hydroxylamine, formation and properties 286 Hydrozoic acid 8, 282 Hypobromous acid, formation and properties 360 Hypochlorites, detection of 404 formation of 348 formation from chlorine 344 on arsenic 66 Hypochlorous acid 348, 349 Hypoiodous acid, existence of 363 Hyposulphites, detection of 305 ignition of 305 Hypophosphites in formation of PH S 305 Hypophosphoric acid 307 Hypophosphorous acid 304-306 estimation of 306 formation of 304 preparation of 305 properties of 304 reactions of 305 solubilities of 305 with bismuth salts 102 Hyposulphurous acid 323-324 Imperial green 109 Indigo test for nitric acid 289 Indium 201 Ink, common 157 sympathetic 168 lodates detection of 371 estimation of 371 formation of 369 ignition of 371 lodates, reactions of 370 lodicacid 369-371 formation of 369 preparation of 369 properties of 369 reactions of 370 Iodide of nitrogen 363 Iodides, decomposition by HNO 3 . . . 290 detection as PdI 2 133 detection of 368 estimation of 369 formation of 365 ignition of 368 occurrence of 365 reactions of 366 separation of, from bromides and chlorides by KMnO 4 181 solubilities of 365 Iodine 362-364 detection of 364 estimation of 364 formation of 363 liberation by copper salts 108 occurrence of 363 on antimonous salts 78 o"n antimony 66 on arsenic 66 preparation of 364 properties of 362 reactions of 363 separation from Br by Pd 134 solubilities of 363 Ions 21 lonization and solution 20-24 Iridium 134, 135 Iron 157-162 alpha 154 and zinc groups 141 beta 154 detection of 165, 166 detection of traces in copper 157 detection of traces 157, 158 estimation of 162 gamma 154 group 144 group, separation from Co, Ni, and Mn by ZnO 161 hydroxides 155 in relation to metals 6 occurrence of 154 oxidation of 162 oxides, 154, 155 preparation of 154 properties of 153 pyrites 313 INDEX. 429 Iron, reduction 162 salts, ignition of 161 salts, solubilities of 156 salts, with alkalis 156 salts, with nitroso-/3-naphthol 157 salts, separation from Al as basic nitrate 161 separation from Al and Cr by nitroso-/3-naphthol 157 separation from Cr and Al 157 separation from Ni by xanthate ... 175 solubilities of 155 Kainite 228 Kieserite.. . 313 Lanthanum 202 Lead 29-36 acetate, properties of 32 chloride 34 chloride, precipitation of 53 chromate, formation of 35 compounds, ignition of 35 detection in alloys 379 detection of 36 estimation of 36 in the test f or Al 167 iodide, formation and proper- ties 35 notes on analysis of 129 occurrence of 29 oxidation and reduction 36 oxides of 29 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 for 33 tests for 54 Leblanc-soda process 259 Lepidolite 241 Light, action on silver salts 50 Lime, slacked 217 Limestone (CaCO 3 ) 216, 219, 267 Lithium 234-236 Limonite 155 Linnaeite 167 Lollingite 57 Magnesia mixture. . . 146 Magnesite 220, 267 Magnesium 220-222 as a reducing agent 222 detection of 222 estimation of 222 hydroxide, formation 220 occurrence of 220 oxalata, separation of, from K andNa 221 oxide, formation of 220 preparation of 220 properties of 220 removal for detection of sodium. . . 242 salts, with ammonium salts 221 salts, with arsenic acid 61 salts, with Na 2 S 221 salts, solubilities of 220 Magnetite 155 Malachite 104 Manganates, identification 405 Manganese 177-182 detection of,. . . . '. 182, 191 estimation of 182 hydroxides of 177 hydroxides, solubilities of 178 ignition of 182 in third group 164, 166-167, 189 occurrence of 177 oxidation of 182 oxidation to permanganic acid .... 180 oxides 177 oxides, solubilities of 178 preparation and properties 177 reduction of 182, 183 reduction by sulphites 181 salts, reactions with oxalic acid . . . 180 salts, solubilities of 178 salts, with alkalis 179 salts, with sulphides 181 separation from zinc with acetic acid 189 solubilities of 178 with KI 181 Manganic acid 177 Magnesite 177 Marble 216,267 Marcosite 155 Marsh's test 62 Mass act ion, law of 23, 38 Mayer's reagent 43, 238 Melanconite .104 430 INDEX. Mercurammonium compounds ... 39 Mercuric chloride with stannous chloride 88 sulphide, formation and proper- ties 41 sulphide, with K 2 S 118 Mercury 37-45 chlorides 42 compounds, ignition of 43 detection and estimation of 44 iodides 42 metallic, analysis of 379 occurrence of 37 oxidation of 45 oxides 37 preparation and properties of 37 salts, reactions 39, 43 salts, solubilities of 38 solubilities of 37 sulphide, analysis of 128 Metals, classification 10 grouping 387 table of separation 388 Metaphosphoric acid 308 Metastannic acid 83 Mica 299 Microcosmic salt 236 use in ignition 377 Milk of lime 217 Millerite 173 Molybdates in analysis 54 with phosphates 98 Molybdenite 97 Molybdenum 97-99 deportment in second group 99 detection of 99, 124 estimation of 99 ignition tests 99 notes on analysis of 124 occurrence of 97 oxides and hydroxides 97 preparation and properties 97 reduction tests 99 solubilities of 97 Molybdic acid 97 Molybdite 97 Mottramite 206 Monazite 204 Monazite sand . . .199 Nascent hydrogen on nitric acid. . . . 286 Neodymium 199, 203 Nessler's reagent 43, 237 Niccolite 57, 173 Nickel 173-176 detection of 176 detection of, in presence of Co by KI 190 distinction from cobalt 175 estimation of 176 hydroxides 173 hydroxides with KI 176 ignition of 175 occurrence of 173 oxidation of 177 oxides 173 properties and preparation 173 reduction 177 salts with alkalis 174 separation from Co, cyanide method 169-170 separation from Co, by nitroso- j8-naphthol 173 separation from Co, by KNO2 170-171 separation from Co, by sulphide 175 separation from Co, by xan- thate 175 separation from cobalt by NH 4 CNS 172 solubilities of 173 solubility of NiS in ammonium sulphide 175 test for by means of dimethylglyox- ime 176 xanthate, separation from Fe 175 Niobium (Columbium) 198-199 Nitrates, decomposition by igni- tion 288 distinction from chlorates 404 occurrence of 285 preparation of 285 proof of absence 402 solubilities of 286 Nitric acid 285-291 as an oxidizer 286 Brown ring test 288 decomposition of, by HC1 287 detection of 288 detection by diphenylamine 289 detection by reduction to NH 3 ..286, 289 detection by reduction to nitrite . . 289 dissociation by heat 288 estimation of 291 formation of 285 indigo test 289 in separation of Sn, Sb and As . . . 121 sodium salicylate test 289 test for by dimethylaniline and diphenylamine 290 INDEX. 431 Nitric acid, with phenol 290 with pyrogallol 291 with brucine 290 occurrence of 285 on antimony 66 on arsenic 66 preparation of 285 products of reduction 286 properties of 285 Nitric anhydride, formation of 286 oxide 104, 221, 283 Nitrites, decomposition by ignition. 284 detection of 284 test for nitric acid 288-290 Nitrof erricyanides 278 Nitrogen 281, 282 chloride 62, 120,337 combination with elements 282 detection and estimation 282 family 7 formation, occurrence 282 peroxide 285 properties 281 Nitroso-/3-naphthol, separation of CoandNi 170, 190 separation of Cu from Cd 107 with iron salts 157 Nitroprussides 278 Nitrous acid 284-285 as an oxidizer. . . 284 as a reducer 284 formation of 284 occurrence of 284 properties of 284 reactions with 284 solubilities of 284 Noble metals, enumeration 7 Nordhausen sulphuric acid 332 Notes on detection of acids 401 on analysis of calcium group. . .224226 on analysis of third group 164 Opal 299 Orangite 204 Order of laboratory study 25 Organic substances, removal of . 374, 375 Orpiment 57, 313 Orthoclase 144 Osmium 133 Osmotic pressure 20, 21 Oxalates, decomposition by ignition of 402 decomposition by oxidation 402 detection of 266 distinction from tartrates 260, 389 Oxalates, estimation of 266 ignition of 266 in 3d, 4th and 5th groups 194 reactions of 264 solubilities of 264 Oxalic acid 263-266 as a reducer 264 decomposition of by H.SO 4 265 formation of 263 in separation of gold 92 occurrence of 263 preparation and properties of ^263 solubility of 264 Oxidation, balancing equations in 244, 245 Oxidizing flame 375 Oxygen 291-293 as a poison 292, 293 combinations with ignition 293 detection of 293 estimation of 293 foriuLiuan of 291 occurrence of 291 preparation of 292 reactions with 292 Ozone 293,294 separation from H 2 Q 2 295 Palladium 133, 134 distinction from gold and plati- num 134 separation from copper 106 sponge 133 Palladous iodide in analysis 133 Paris green 62, 109 Pearlite 154 Pentathionic acid, formation and properties ' 325 Pentlandite 173 Perchlorates, preparation and properties 353, 354 Perchromic acid 153 Periodic acid 372 system, table of 2 Permanganates identification 405 action on antimonous salts 78 Permanganic acid 178 Persulphate method of separating the halogens 347, 403 Persulphuric acid 336 Petalite 241 Phenol reaction for nitric acid 290 Phenylhydrazine, on aluminum salts .145 432 INDEX. Phosgene, formation 262 Phosphates, changes by ignition. . . 312 detection 165, 312, 402 distinction between primary, secondary and tertiary 310 estimation of 313 in presence of third and fourth group metals.. 143, 193, 194, 196, 197 occurrence of 308 reaction with ammonium molyb- date 193,311 separation as ferric phosphate 193 solubilities of 309 Phosphides, formation of 312 Phosphine 304 Phosphoric acid 307-313 preparation of 309 properties of 307 Phosphoric anhydride, formation of 308 Phosphorite 216 Phosphorous acid 306, 307 detection of 307 preparation and properties of 306 Phosphorus 301-304 detection and estimation of 304 in combination with the halogens 304 occurrence and preparation of .... 303 properties of 301, 303 use in match-making 302 Phosphotungstates 137 Picric acid, in detection of potas- sium 230 Pitch-blende 206 Plaster of Paris (calcium sulphate). . 219 Platinized asbestos 93 Platinum 93-97 apparatus, care of 95 black 93 chloride, as a reagent 94 detection of 96, 124, 379-380 distinction from palladium 134 estimation of 96 iridium alloys, properties 134 notes on the analysis of 125 occurrence of 94 preparation and properties 93, 94 reduction of 95, 96 sponge 93 Polarity 3 Potassium 228-232 as a reducing agent 232 bichromate, in test for stron- tium and calcium 225 carbonate, as a reagent 229 PAGE Potassium chlorate, in preparation of oxygen 292 chloride with platinum chloride ... 95 cyanide with copper salts 107 cyanide with ferrous salts 157 detection of 229, 232 estimation of 232 ferricyanide, formation of 277 ferrocyanide, formation of 273, 276 hydroxide, as a reagent 229 iodate, in separation of alkaline earths 213 iodide, as a reagent 240 iodide, in separation of AgCl from SbCl 3 122 iodide, in the test for nickel 190 iodide, on nickelic hydroxide 175 iodide, on permanganates 181 nitrite in separation of cobalt from nickel 170, 171 occurrence, preparation and properties of 228 picrate 230 pyroantimonate 73, 224 salts, flame test 231 thiocyanate with copper salts 107 thiocyanate with iron salts 158 xanthate, for detection of copper 107 Powder of algaroth 75 Praseodymium 199, 202 Precipitates, formation and re- moval of 17, 18 Principles 405 Problems in molecular propor- tions 19 in synthesis 410 Proustite 46, 57 Prussian-blue, formation of . . . 158, 274 Purple of Cassius 89, 93 Pyrargyrite 46 Pyrite 45, 155 Pyroantimonic acid 73 Pyrochlor 198 Pyrogallol, as a test for nitric acid 291 Pyrolusite 177 Pyromorphite 29 Pyrophosphoric acid, formation. . . . 308 Pyrosulphuric acid, formation 332 Pyrrhotite 155, 173 Quartz 299 Reagents, care in the addition of . . . 17 list of... . 415 INDEX. 433 Realgar 57, 313 Reducing flame, description of 375 Reduction, balancing equation" in 244, 245 with charcoal 376, 377, 381 Reinsch's test for arsenic 67 Rhodium, distinction from iridium 135 properties and reactions 132 Rhodocrosite 177 Rochelle salts, composition of 260 Rosolic acid as a test for carbon dioxide 270 Rubidium, properties and reac- tions 240 Ruby 144 Rule for balancing equations 246 Ruthenium, properties and reac- tions 129 Rutile 205 Salt 233 Saltpeter, occurrence 285 Samarium, properties and reactions 202 Sand 299 Sapphire 144 Scandium, properties and reac- tions 202-203 Scheele's green and Schweinfurt's green 109 Scheelite 136 Selenic acid, separation from sul- phuric acid 140 Selenite 216 Selenium, properties and reac- tions 139, 140 Siderite 155 Silica (silicon dioxide) 300 detection and estimation of 301 in the microcosmic bead 301 in the third group 167 removal of 402 solubilities of 300 Silicates, decomposition by igni- tion 300 in analysis 54 Silicic acid 299-301 Silicon 299 distinction from tantalum 203 Silico-fluoride (fluosilicate) 298 Silicon fluoride, formation 297, 298 preparation and properties 298 separation from thorium 205 Silver 45-50 arsenate and arsenite, formation . . 62 bromate, properties of 361 Silver chloride, formation and prop- erties 48 cyanate in distinction from cyanides 279 detection of 50, 380 estimation of 50 in presence of mercury salts 55 iodate, properties of 370 mirror, formation by tartrates .... 261 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 281 Sipylite 207 Skutterudite 167 Smaltite 57, 167 Soapstone 299 Soda lime on stibine 79 process, Le Blanc's 267 process, Solvay's 268 Sodium 232^235 amalgam, action with arsenic 64 as a reducing agent 235 detection of 73, 234 estimation of 235 flame test 234 hydroxide, formation of 233 nitroferricyanide as reagent. . . 236, 320 occurrence of 233 phosphate as reagent 233 Sodium phosphomolybdate as re- agent.. 98, 238 preparation and properties of. .232, 233 pyroantimonate 73, 80 ' pyrophosphate with copper and cadmium 107 salicylate test for nitric acid 289 sulphide, preparation of 317 thiosulphate on cupric salts 108 thiosulphate with antimony salts. . 78 thiosulphate with third group metals 146, 147 Solids, conversion into liquids 328 decomposition upon ignition. . . 382, 383 effect on ignition with cobalt nitrate 384 preliminary examination of 375 separation of 17 table for preliminary examina- tion 382 Solubility, degrees of 15, 16 434 INDEX. Solubility-product 24 Solution and ionization 20-24 Solvay soda process 268 Sonnenschein's reagent 98 Sperrylite 94 Stannic salts, solubilities 84 sulphide, formation and proper- ties of 86 Stannite 82 alkali, as a test for bismuth 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 125 salts, solubilities 84 salts with silver nitrate 87 salts with sulphurous acid 86 sulphide, formation and proper- ties 85 Stephanite 46 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 Stibnite 72 Strontianite 214 Strontium 214-216 detection of 216, 219 estimation of 216 hydroxide, formation 214 occurrence of 214 preparation and properties of 214 sulphate, distinction from CaSO 4 215 sulphate, separation from BaSO 4 215 Sulphates, detection of 335 estimation of 336 ignition of 335 preparation of 332 reduction by ignition with car- bon 256 solubilities of 333 Sulphites, detection of 330 distinction from sulphates , . 330 estimation of 330 ignition of 330 interference in test for oxalates . . . 402 preparation of 327 separation from sulphates by Ba salts 213 solubilities of . 328 Sulphides, detection 320 estimation of 321 formation of 316 ignition of 319 reactions of 318, 319 solubilities of 28, 317 Sulphur 313, 315 combinations on ignition of 315 detection and estimation of 315 formation of 313 in the tin group 118 occurrence of 313 oxidation by reagents 314, 315 oxides 313 precipitation of 54, 114, 115 preparation and properties of . .313, 314 reactions in forming sulphides. . . . 314 relations of 9 separating copper from cadmium 107 solubilities of 314 Sulphuric acid 331-336 detection in presence of sulphates 335 formation and occurrence of 331 manufacture of 331 properties of 331 reactions with 333-335 separation from Se 140 separation from Te 139 anhydride, preparation of 331 Sulphurous acid 327-330 on arsenic acid 60 and sulphites as reducers 329 occurrence of 327 preparation and properties of 327 formation of 327 reduction of cupric salts 108 solubilities of 328 on stannous salts 86 Sylvanite 138 Synthesis, problems in 410 Table for acids as precipitated by barium and calcium chlorides . . . 398 for acids precipitated by silver nitrate 399 for acids, preliminary 390 for analysis in presence of phos- phates by the use of alkali ace- tates and ferric chloride 196 for analysis in presence of phos- phates by use of ferric chloride and barium carbonate 107 for analysis of the Silver Group (first) 52 INDEX. 435 Table for analysis of the Copper Group (second) 126 for analysis of the Tin Group (second) 116 for analysis of the Iron Group (third) 163 for analysis of the Zinc Group (fourth) 188 for analysis of the Calcium Group (fifth) 223 of grouping of the metals 387 of separations of the metals 388 of separation of the ammonium sulphide precipitates of the Iron and Zinc Groups 192 of solubilities 411 Tannic acid with iron salts 157 Tantalite 198, 203 Tantalum, distinction from silica . . 203 distinction from titanium 203 properties and reactions of 203 separation from columbium 203 Tartar emetic, composition of 260 Tartaric acid 259-261 in detection of potassium 229 distinction from citric acid 259 formation and properties 259-260 Tartrate calcium, deportment with water 260 detection of 260 distinction from citrates 260 distinction from oxalates 260, 401 estimation of 261 Tartrates, ignition 261 reactions 260 solubilities 260 Tellurite 138 Tellurium 138-139 distinction from selenium .... 139, 140 properties and reactions of .... 138, 139 separation from sulphuric acid. . . . 139 Tenorite 104 Terbia 203 Terbium 203-204 Tetradynite 138 Tetrathionic acid, formation and properties 325 Thallious iodide 204 Thallium, properties and reactions . 204 Thioacetate in formation of sul- phides 316 Thiocyanates, reactions with 280 test for cobalt by means of 170 Thiocyanic acid as a reducer 281 properties of 280 Thionic acids, table of compari- sons 326 Thiosulphates, detection of 323 distinction from sulphates and sulphites 323 estimation of 323 ignition of 323 formation and properties of 321 Thiosulphuric acid 321-323 Third group reagents 142 Thorite 204 Thorium 204-205 Tin 82-89 creaking of 82 detection of 88, 122, 379 estimation of 88 Group, metals of 56 Group, separation from Copper Group 115 Group, sulphides with (NH 4 ) 2 S.r. . . 115 occurrence of 82 oxidation of 88 oxides and hydroxides 82 preparation and properties of 82 notes on the analysis of 124, 125 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 123 . separation from arsenic 118 solubilities of 83 sulphides, colloidal 115 with antimony and with arsenic . . 87 Tinstone 82 Titanif erous iron 205 Titanite 205 Titanium 205-206 distinction from columbium 206 distinction from tantalum 203 properties and reactions of 205 separation from thorium 205 Triphylite 241 Trithionic acid, formation and properties 324 Tungsten, properties and reactions 136, 137 Turnbuii's blue. . . 155 Unit of quantity 23 Uranium, properties and reactions. 206 Urea, from ammonium cyanate 279 436 INDEX. Valence, negative 3 Valentinite 72 Vanadinite 206 Vanadium 206, 207 Volatile alkali (ammonium hy- droxide) 221 Volborthite. . .... 206 Wad 167 Water, action on bismuth salts 101 action on antimonous salts 75 Welsbach burners 209 Witherite 211 Wolframium (tungsten) 136 Wulfenite . . 97 Ytterbium properties and reac- tions 207 Yttrium.. . 207 PAGE Zincates, formation of 184 Zinc 183-186 blende: 183, 313 blende (Freiburg) 201 detection and estimation of 186 Family 5 granulated 63, 183 Group, table for analysis 188 Group, comparative reactions 187 hydroxide and oxide 184 ignition of 186 occurrence of . . 183 oxidation of 186 platinized 183, 250 preparation and properties 183 reduction of 186 salts, solubilities and reactions of 184 sulphide, formation in presence of acetic acid 184 Zircon 209 Zirconium . , .209 LITERATURE OF THE CHEMICAL INDUSTRIES On our shelves is the most complete stock of technical, industrial, engineering and scientific books in the United States. 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