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Maps, plates, charts, etc., may be filmed at different reduction ratios. Those too large to be entirely included in one exposure are filmed beginning in the upper left hand corner, left to right and top to bottom, as many frames as required. The following diagrams illustrate the method: Les cartes, planches, tableaux, etc.. peuvent Atre fiimAs A des taux de rAduction diff Arents. Lorsque le document est trop grand pour Atra raprodult en un ssui clichA. 11 est fiimA A partir de I'angle supArieur geuche. de gauche A droite. et de haut en bas. en prenant la nombra d'images nAcessaire. Les diagrammes suivantf illustrant la mAthode. 1 2 3 1 2 3 4 5 6 f 4 ^ US'cSy * n . Jl INTRODUCTION TO : : Qualitative 7\.tiatxsis BY W. LASH MILLER, B.A., Ph.D. Dciiioiistra/or of Che mist ly in the University of I'oionto. AND F. J. SMALL, B.A., Ph.D. Lcitinrr in Chemistry in the University of Toronto. V TORONTO : THE BRYANT PRESS 1896. PREFACE. THIS book is written prinnrily for use in tiie chemical laboratory of the University of Toronto. The laboratory course is designed to complement the work in the lecture room, by making the student familiar with the phenomena treated of in the lectures ; and, accordingly, in the present work special attention has been paid to the study of the chemical reaction, and an atiempt has been made to introduce in qualitative form the results of the recent advances in the theory of solutions. With a view of arranging the descriptive matter as concisely as possible, and, at the same time, of developing the powers of observation and comparison in the individual student, the solubilities and colors of the various salts have been presented in tabular form, and the repetition of detailed instruction avoided throughout. The tables of separation are constructed so as to necessitate continual reference to the text, and are placed with the groups of substances to which they refer; as it is the experience of the authors that when bunched in the usual way, they are made use of to the almost total neglect of the rest of the book. In conclusion, the authors desire to express their special obli- gations to the following works : — E. J. Chapman : Blowpipe Practice. A. Classen : Handbuch d. qual. chem. Analyse. W. Ostwald : Die wisse Grundlogen d. anal. Chemic. H. V. Pechmann : Anleitung z. qual. chem. Analyse. Toronto, October, 1896. TABLE OF CONTENTS. Introduction Preliminary Examination in the Dry Way Solution Determination of the Bases 31 Division into Groups The First Group I -ead Silver Mercury Separation The Reactions of the First Group. The Second Group Arsenic. Antimony 'I* n Separation II A Bismuth ....... Copper Cadmium Separation II B The Third Group Aluminium Iron Chromium Separation III A Cobalt Nickel Manganese Zinc . Separation III B I 25 29 -116 31 33 34 38 40 45 47 56 59 61 64 67 70 72 75 77 78 81 83 86 89 90 92 94 97 100 TABLE or CONTENTS. VI. The Fourth Group 104 Barium 105 Strontium 106 Calcium 107 Separation 109 The Fifth Group no Magnesium no Potassium 112 Sodium 113 Lithium 114 Ammonium 115 Separation 116 Determination of the Acids 1 17-127 Preliminary Examination 118 Division into Groups 119 Alphabetical List of the Principal Acids 120 Determination of the Alkaloids 128-134 Division into Groups . . 129 Alphabetical List of the Principal Alkaloids. . 130 Table for the Detection of a Single Acid 135 Mendelejeff's Table 136 Index 137 INTRODUCTION This book is written for the use of students who have already had a certain quantity of instruction in the subject of general Chemistry, who know what is meant by a chemi- cal formula, what qualitative and quantitative information is conveyed by a chemical equation, and who understand the methods employed in the determination of atomic and mole- cular weights. They are for the most part introduced to the study of Qualitative Analysis, not to make analysts of them (though some such course as that sketched here forms a necessary part of the education of every professional analyst), but because this course in laboratory work possesses advantages equalled by none other yet devised both in the variety of the substances and the number of the reactions it presents for study, and in the moderate nature of the demands it makes on the time and pocket of the student, on the resources of the laboratory, and on the energies of the instructor. It is not to be denied that these advantages are, to a certain extent, offset by corresponding disadvantages. A large number of bodies is studied, but all from the same point of view ; undue prominence is attached to such compounds (e.g., the insoluble salts) as are convenient for purposes of separation, while, on the other hand, as the reactions are carried out for the most part in aqueous solution, many whole classes of substances — most of the gases, and nearly all bodies decomposed by water — are necessarily excluded from the course ; and even the very reactions most met with, instead of being fair samples of the great majority of chemical reactions, are selected from a type 3 A TYPICAL ANALYSIS. particularly suited to the necessities of the method of analysis adopted. From this it follows as a necessary consequence that men whose knowledge of practical chemistry is derived merely from a laboratory course in analysis must form a very one- sided conception of the nature of a chemical reaction, which is all the harder to remove is it is founded on their own experience, and seems to fit the facts too well to need replacement by another. It is with a view of pointing out what is general in the reactions that will be met with in the laboratory, and what is exceptional and due to the fact that chemistry is there studied from the standpoint of the analyst, that this introduction is written. If a piece of silver and a piece of copper were melted together, and a chemist were asked to obtain from the alloy (a coin, for instance) the silver and the copper in the pure state, he would proceed as follows : From tables of the solubilities of the salts of silver and of copper, he knows that while the chloride of silver is very insoluble in water the corresponding cupric salt is quite soluble — the nitrates of both metals are soluble in water. His first step then would be to dissolve the coin in a mixture of nitric acid and water, thus obtaining a solution of the nitrates of the two metals, Cu+4HNO,,=Cu{NO..)„ + 2NO.,+zH„0, I y ^ to which he would add hydrochloric acid, thus "precipitat- ing " the silver, i.e., converting it into the insoluble chloride, AgNO:,-\-HCl = HNO,,+AgCl, which might be separated from the solution by mechanical means (straining through paper). From this precipitate, after freeing it, by washing, from adhering cupric nitrate solution, he might obtain the silver in the metallic state by SPECIAL NATURE OF THE REACTIONS EMPLOYED. 3 heating it mixed with a little carbonate of soda on a piece of charcoal. 2AgCl+Na^CO., + C=^2Ag+2NaCl+CO^+CO When the silver was thus disposed of, the copper might be obtained from the filtrate by precipitating with potash as the hydrate CuiNO,^)., + 2KOH = Cu{OH)^ + 2KNO.^, which on heating loses water, passing into the oxide, Cu{OH)„=CuO+H„0, from which, lastly, by heating with charcoal, the copper itself might be prepared. CuO + C = Cu+CO. The method of separation here pursued s typical of that followed wherever possible, and, in any particular case, the nearer the operations resemble those just quoted, the better from the point of view of the analyst — and the worse from that of the student of chemistry, who, it goes without saying, is the gainer by variety. Two peculiarities are common to all the reactions given above. In the first place, they are quickly finished ; in the second, they are quantitative — all the copper and silver are dissolved in the nitric acid ; all the silver is obtained as chloride, and this is all finally converted into the metallic state, so that the button of silver obtained by the analyst at the end of his work weighs, as- near as may be, the same as that originally employed in the manufacture of the coin. The first feature is common to nearly all reactions of inor- ganic salts dissolved in water ; the second, on the other hand, is the result of certain special conditions which, therefore, must be fulfilled by all reactions intended as a basis for methods of separation in analysis ; and it is, consequently, of the greatest importance to understand, as far as possible, just what these conditions are, and to learn the effect of each of them separ- ately upon the course of the chemical reaction. INFLUENCE OF INSOLUBILITY. One of the most obvious is the insolubility* of the silver chloride. If it dissolved to any noticeable extent in water, or rather in the solution of copper salts, hydrochloric acid, etc., from which it is to] be removed by filtration, then, even if the hydrochloric acid used converted all the silver nitrate into silver chloride, it is evident that only a part of the silver present would be separated in the form of solid silver chloride, the rest remaining dissolved and mixed up with the copper in the solution. There is, however, another very important respect in which this property of the silver chloride affects the reaction ; for it is very largely owing to the insolubility of this salt that the hydrochloric acid can convert the silver nitrate completely into chloride. In order to comprehend this latter important propo- sition it will be necessary to consider a few cases of chemi- cal reaction in which all the substances involved are " soluble": such, for instance, are the two bodies, hydrochloric acid and sodium nitrate, and the two others, nitric acid and sodium chloride, that would be formed from them if they reacted together in solution according to the equation : HCl+NaNO.^=HNO^+NaCl analogous to HCl + AgNO^ = HNOs+AgCl. The only difference between these two cases is that in the latter, if more than an infinitesimal quantity of the " insoluble '* silver chloride be formed, it must necessarily be " precipitated," i.e., separated in the solid form ; whereas, in the former, all four bodies are so soluble in water that if fairly dilute solutions be employed no change of the nature represented above can produce a precipitate : and, accordingly, on mixing fairly dilute solutions of the two substances mentioned (hydrochloric acid and sodium nitrate), no precipitate is formed. Is it, then, fair to say that "nothing happens" on mixing these solutions ? or that the two bodies-named " do not react " with each other ? Obviously not. That would be to make the formation of a precipitate the sole criterion of the occurrence "Short for "very slight solubility" — one]_litre of water at i8'C. dissolves about 1.7 wf. silver chloride. See page 9. REACTIONS INVOLVING SOLUBLE SALTS ONLY. of a chemical reaction. On the other hand, if anything has happened, how is one to find it out ? By evaporating the solution and seeing what salt crystallizes out ? Or by adding alcohol and (as, in general, salts are less soluble in water di- luted with alcohol than in the same quantity of pure water) thus producing a precipitate? Evidently not, for these expedi- ents would only succeed in reducing the case to one of the former type, depending on the relative solubilities of the vari- ous substances involved, while the whole object of the enquiry is to find the laws governing chemical reactions when that disturbing element is removed. An answer to the question just put may, however, be ob- tained by studying the " physical " changes which take place on mixing thetwo solutions. If there be no evolution nor absorp- tion of heat, if no contraction nor expansion, change of color, etc., etc., take place; in short, if all the measurable properties of the mixed solutions be merely the average (properly calcu- lated) of those of the two solutions mixed, it would be a safe conclusion that no reaction had taken place. If, on the other hand, the physical properties of the mixture be intermediate between those of solutions (in the same quantity of water) of hydrochloric acid and of sodium nitrate on the one hand, and of nitric acid and of sodium chloride on the other — which is the actual case — it is fair to infer that the mixture contains not one pair exclusively, but some of each of the four bodies represented by the formulae, HCl, NaNO^, HNO^, and NaCl; in other words, that a reaction does take place between the two first mentioned, but stops short before they are entirely converted into the second pair — that the reaction is " incom- plete"." A very striking instance of such a change of physical pro- perties is afforded by one of the reactions used in the analysis of iron salts, viz., the reaction between ferric nitrate and potassium sulphocyanide. The dilute solutions of each of these salts taken separately are almost colorless (that of the iron a little yellowish), while that obtained by mixing them is of a deep blood-red hue, owing to the formation of the deep red ferric sulphocyanide. That the reaction which takes 6 INCOMPLETE AND REVERSIBLE REACTJONS. place when ferric nitrate and potassium sulphocyanide are mixed in solution in the quantities represented in the equation Fe (N0-^), + 3KSCN^Fe {SCN), -\-3KNO, is not " complete," i.e., that all the ferric nitrate employed has not been converted into ferric sulphocyanide, is inferred from the observation that the color of the mixture is not nearly so deep as is that of a solution prepared by dissolving in the same amount of water the quantity of ferric sulphocyanide repre- sented on the right hand side of the cjuation above. The amount of the red salt formed varies with the temperature and other conditions ; in one case a measurement of the d' pth of color fixed the quantity of ferric sulphocyanide at tnirteen per cent, of what would be present were the reaction " complete." One of the most important results of the investigations that have been made on this matter is that the solution formed by mixing equivalent quantities of (for instance) hydrochloric acid and sodium nitrate is precisely identical in every respect with that produced from equivalent quantities of nitric acid and sodium chloride,* and the only conclusion that can be drawn from this remarkable fact (which is perfectly general, though a special case is here selected for illustration) is this, that not only can hydrochloric acid react with sodium nitrate according to the equation HCl+NaNO^=HNO^+NaCl, but also nitric acid and sodium chloride, the substances formed by this reaction, can, under suitable circumstances, react together to form again the hydrochloric acid and sodium nitrate from which they themselves were prepared : HNO^+NaCl=^Ha+NaNO^ in short, that the reaction of hydrochloric acid on sodium nitrate is " reversible." This may be indicated by replacing the sign = by the sign < *" to show that the reaction will proceed in either direction according to circumstances. *It goes without saying that the' same quantity of water must be used in each case. « CHEMICAL EQUILIBRIUM. f It is very important to obtain a clear idea of what is involved in the conception of the " reversibility " of a reaction, for it will appear from what follows that this reversibility, instead of being the exception, is the rule even among the reactions met with in the course of analysis. First, then, as to the couiiection between " reversibility " and " incom- pleteness." In the case A+B ^z^z*' C + D if the substances A and B did react completely to form C and D, the very statement that the reaction between them is' reversible implies that these two being present alone would immediately react again to form A and B ; and thus the com- position of the solution would vary from one extreme — all A and B, to the other — all C and D, the course of the reaction, resembling that of a pendulum oscillating about its point of equilibrium. As a matter of fact, however, no such phenomena are met with in chemistry ; and if the simile of the pendulum is to be retained at all, the latter must be imagined as provided with a broad vane dipping in glycerine, and prevented by friction from shooting past its point of equilibrium. Second, as to the causes determining the direction of the reaction: the following, which is a mere restatement of the case just considered, shows that the concentrations of the sub- stances involved is one very important factor. If A and B [and the same holds for C and D] be present alone, they react, and the progress of the reaction involves a diminution of their concentration in the solution and an increase of that of C and D, until a definite relation between the concentrations of the two opposite pairs is arrived at (depending in general on the chemical nature of the reacting bodies), equilibrium is reached, the reaction stops. There is no question, then, as to what will happen if in the solution at equilibrium the concentrations of both A and B be increased {.e.g., by dissolving fresh quantities of each of them in the solution) — a reaction will take place in the direction that " uses up " or diminishes the concentration of the sub- stances added, until the former relation between the various concentrations is restored ; but what if only one, say A , be ACTION OF MASS. added ? Thegresult is what might be expected, equilibrium is disturbed, and the ensuing reaction, as in the former case, serves to diminish the concentration of that substance (A) which was added to the solution. [The opposite course could onlyhavfc resulted in bringing about an explosive reaction ending in the total disappearance of C and D ; and as this would be effected by an indefinitely small addition of ^ , it was very improbable from the outset.] This influence of concentration has obtained the technical name of the " action of mass," but it should not be forgotten that it is not__the mass nor the weight of the substances involved that is of importance, but the mass per unit volume or the concentration.* The discussion of a concrete example may serve to render these conceptions more definite. In the solution formed from equivalent quantities of ferric nitrate and potassium sulpho- cyanide — or, what is the same thing, from equivalent quanti- ties of ferric sulphocyanide and potassium nitrate dissolved in the same quantity of water — and containing the four salts in equilibrium with respect to the following reaction : Fe {N0,), + 3KSCN ^:=.>' Fe iSCN), + 3KN0, if the concentration of either the ferric nitrate or of the potas- sium sulphocyanide be increased (e.g., by adding small quanti- ties of the dry salt, or a drop of the saturated solution of either) the equilibrium will be shifted and a reaction will ensue, resulting in the formation of more of the ferric sulpho- cyanide. As it is to the latter salt that the solution owes its red color, it follows that the addition of either of the first- named salts will deepen the color, and, conversely, addition of potassium nitrate will render it lighter. This forms an easily performed and very striking experiment, which simply cannot be understood from the other point of view of the chemical reaction, viz., that it depends on the " relative affinities " of the substances involved whether or not a reaction will occur ; and that when one does take place it is quantitative. From this other standpoint the formation of a red color in * In other words, it is the per cent, composition of the solution, not the scale on which the experiment is carried on, that is the determining factor. 'soluble" and "insoluble" relative terms. the first instance showed that a reaction had taken place between the two salts involved ; if, on adding more ferric nitrate, the solution became darker red, that was because the quantity previously added was insufficient to combine with the potassium sulphocyanide present. So far, so good ; but, then, what explanation for the deepening in color on addition of potassium sulphocyanide itself? It ought not to deepen ! There was too much potassium sulphocyanide there already! — The discovery that has removed all these difficulties is just this, that the "chemical affinity," whatever that very ill- defined word may really mean, is dependent, among other things, on the concentrations of the reacting substances. This explains the reversible nature of the reaction, which, again, necessitates the occurrence of an equilibrium. In the case just considered it was possible to alter arbi- trarily, within very wide limits, the concentration of any one of the four salts involved by adding the substance in question to the solution and letting it dissolve. But in this very statement of the possibilities of the case is contained also the necessary restriction. Most substances will not dissolve in all piopor- tions in water — a few liquids, e.g., nitric and acetic acids, are completely miscible with water ; but in the case of solid sub- stances and of gases there is always a limit, depending on the temperature, to the amount that will dissolve on shaking with a given quantity of solvent; in other words, an upper limit to the concentration of the solutions that can thus be prepared from them. In some instances this limiting concentration is high : water at ordinary temperature will dissolve four times its own weight of calcium chloride and three-quarters of its weight of ammonia gas ; in others it is low, one hundred parts of water dissolving seven of mercuric chloride, one-half part of the chloride of lead, and only O.00017 o^ silver chloride. As the quantity of chloride of lead that will be taken up by a test- tube full of water is small enough to easily escape notice by the eye, this and all less soluble substances are in everyday language spoken of as "insoluble." No salt, however, is strictly insoluble ; the difference between " soluble " and "insoluble" is one of degree merely, not of kind; and 10 INFLUENCE OF INSOLUBILITY. between the extremes of the chloride of calcium on the one hand, and the chloride of silver on the other, are to be found the great majority of the substances met with in the laboratory. If,then,oneofthe salts ofa chemical system be "insoluble," for instance, the silver chloride in the equilibrium, HCl + AgNO^ ^■z=z>HNO,,-\-AgCl, this insolubility wil! affect its concentration in the solution in two Vv'ays : (rt) The concentration cannot be increased above a certain very low limit ; {b) So long as there is solid silver chloride in contact with the solution, the concentration of the silver chloride dis- solved will remain constant*; and the limits thus set to alterations in the concentration of the silver chloride dissolved have a most important effect on the course of the reaction. In the case so often referred to, HCl + NaNO, :»,HNOa+Naa, if equivalent quantities of hydrochloric acid and sodium nitrate be mixed in fairly dilute solution, just about one-half will be changed into nitric acid and sodium chloride. If the nitrate of sodium be replaced by the nitrate of potassium, of ammonium, of calcium, etc., etc. — in short, by the nitrate of any metal whose chloride is soluble, the result will be found in every case to be about the same. If, however, silver nitrate be used, HCl+AgNO^^z=*'HNO.,+AgCl, long before half, or, with solutions of ordinary strength, long before the thousandth part can be converted into the chloride, the solution is already " saturated " with that salt, i.e., the concentration of the silver chloride in the solution has reached its limit, and " precipitation " has begun. From this time on, the concentration of the silver chloride in the solution * This is strictly true only when the quantities of the other substances in solu- tion are not too great. EXCESS OF THE REAGENT. II remains constant ; but as that of the nitric acid is increasing, and that of the hydrcchloric acid and of the silver nitrate as steadily decreasing — both changes which tend to assist the reverse reaction — it is obvious that at length a state of equilib- rium must be reached, in which the nitric acid formed, and the small trace of silver chloride that can dissolve, balance the hydrochloric acid and silver nitrate remaining. Evidently, the less the concentration of the silver chloride, the less can be the concentration of the hydrochloric acid and of the silver nitrate at equilibrium ; and it is thus owing to the extreme insolubility of the former salt that the reaction results as it does in the almost quantitative precipitation of the silver as chloride. In his power to arbitrarily alter the concentration of the hydrochloric acid, however, the analyst has a means of making the reaction even more complete, and it has long been the practice to add " excess " of the precipitating reagent, i.e., a quantity more than sufficient to combine with all the silver present ; — the explanation of this practical recipe has, however, only comparatively recently been given. The attention of the student is particularly called to the important argument in the preceding paragraphs, by means of which a theory of " chemical affinity," invented originally to account for the occurrence of incomplete and reversible reactions, is able not only to explain the occurrence of what are, practically speaking, quantitative, non-reversible reactions, but even to give a reason for one of the conditions — excess of the reagent — long known as essential to quantitative pre- cipitation. The conception of the typical chemical reaction, as reversible and ending in an equilibrium, is of great importance for another reason. The very completeness and non-reversi- bility of the reactions employed in analytical work, which has just been shown to depend on what might almost be called the " accidental circumstance " of the insolubility of one of the salts involved, was formerly held to be the distinguishing characteristic of a "chemical" as opposed to a "physical'' la FACTORS OF CHEMICAL EQUILIBRIUM. reaction. Under this latter head were grouped such " changes of state " as the passage from water to steam, and the " phenomena of solution," e.g., of salt in water, both being reversible changes which proceed until an equilibrium is reached depending on the temperature and on the concentra- tion (mass per unit volume) of the steam, on the one hand, and of the salt in solution on the other. The discovery that chemical reactions proceed according to the same laws toward the attainment of a similar state of equilibrium modified by the same factors — among others temperature, concentration, and the special nature of the substances involved — has done much to remove this artificial distinction, which is one of those necessarily met with in the growth of every inductive science, since the vast mass of individual phenomena must first be broken up into groups if any progress is to be made, while an appreciation of the relations between these various groups can only be arrived at when the characteristics of each one of them have become properly understood. Though a very important factor of the chemical equilib- rium, the concentration of the reacting bodies is not the only one ; the temperature and the special nature of the reagents are equally important, while the storage battery affords an excel- lent example of chemical action reversed by the influence of electricity. In this Introduction, however, special attention will be given only to the concentration and to the chemical nature of the substances involved, and, unless otherwise specified, all reactions are supposed to be carried out at constant temper- ature, and without the intervention of outside electrical forces. The influence of the " chemical nature " of the various reacting substances attracted attention at a much earlier period than did that of their concentrations ; and side by side with the classification according to their composition of the substances studied under the title of Inorganic Chemistry* * The compounds containing carbon, which outnumber by far all the others together, are usually treated of apart under the title "Organic*^ Chemistry, a name Slven to denote their importance in the animal and vegetable kingdom. Carbon ioxide, however, and the carbonates are, for convenience sake, generally included among the " Inorganic " compounds. CLASSIFICATION UY CHEMICAL BEHAVIOR. »3 there has grown up a second system of classiiication based on their chemical behavior : as by far the greater num- ber of the ordinary analytical opci tions are carried out in aqueous solution, it is no doubt a consequence of this habit of working in the presence of water— which dates back to a period long before the birth of chemical analysis as such — that this second system of classiiication depends for its foundation upon the behavior of a certain number of the more commonly occurring compounds with that universally employed solvent. The reaction that serves as a starting point for the system of classiiication under consideration is that of the neutralization of acids by bases in aqueous solution. Nitric acid reacts with potash according to the equation HNO:,-\-KOH = KNO^+H.,0. Any substance that can take the place of the nitric acid in this reaction is called an Acid; while any substance capable of replacing the potash is termed a Base. Salts are sub- stances formed by the reaction of an acid with a base, according to the equation given above. Those oxides which can be formed from the acids or thj bases by the removal of water, or from which by the addition of water the latter may be obtained, have received the name of Anhydrides ; while the elements are divided into the two classes of Metals and Non-Metals, according as their hydrates are bases or acids respectively. Thus the whole classiiication, which embraces the great majority* of the inorganic compounds, turns on the distinction between Acids, Bases, and Salts. Each of these groups will be shortly considered in turn, the method of subdividing them (still with reference to the " reaction of Neutralization ") explained, and, iinaliy, their relations to the "reaction of Hydrolysis," the reverse of the reaction of neutralization, discussed. • Excluding, however, the most of those decomposed by water, see p. 23. On ^'■J'u*'" ^- '^"""^ * ^^'^'^^ account of a group of substances containing Sulphur, which exhibit the same relations in/er se as those existing between the Oxygen compounds above. '4 ACIDS. All Acids react with potash to form salts (by definition). To "explain" this as due to the "affinity" or "avidity" of the acids for the potash is as much as to say that the acids react with that substance because they want to ; similarly, to say that a given substance has " avidity " for bases is evidently merely another way of saying that it is an acid. The use of metaphorical language such as this, however, has been not withdut benefit to the science; the chemists who pictured to themselves the acids as being endowed with desires, and striving to satisfy them, naturally set themselves the question: " What will happen when two acids want the same base ? " and, what is more, were interested enough to institute an expisri- mental enquiry into the subject. In arranging for the contest they did their best to maintain a fair field, saw to it that both acids were present in equivalent quantities, and that no acci- dent, such as the formation of an insoluble salt, should intervene to aid either side ; and when they found that, even then, the major portion of the coveted base fell to one of the combatants, they ascribed his victory to his greater " strength " —and thus were the first to give definite meaning to a term which had long been in use, and which will be found of the greatest convenience in describing mjiny of the reactions in- volving the acids. To take a particular case, the distinction between a " weak" and a " strong" acid rests on the study of equilibria such as the following : . . (a) HNO ^^NaAc^=.t.HAc^ XaXO , (6) HNO.,-\-NaCl > HCl^-NaNO., in the first of which nitric and acetic acids, and in the second nitric and hydrochloric acids, are " in competition " for the base soda. If, in each of these two cases, nitric acid and the salt be mixed in equivalent quantities, the state of affairs in the two solutions at equilibrium will be found to be remark- ably different — all but a trace of the sodium acetate being converted into sodium nitrate with expulsion of the acetic acid, while very nearly half of the sodium chloride remains undecomposed. The result of these experiments might be stren(;th of acids. «5 expressed in words by saying that nitric acid has a much greater "avidity" for soda than acetic acid has, and just about the same as hydrochloric acid ; but as a great many other bases have been substituted for soda in the reactions formulated above, and as the relative " avidities " of the nitric, acetic, and hydrochloric acids for each of these other bases are found to be practically the same as for soda, it is allowable to speak of nitric acid as being " stronger " than acetic acid and " equal in strength " to hydrochloric, the mention of any particular base not being necessary where all act alike. The following table gives the relative strengths of some of the acids more frequently met with in the laboratory, com- pared with that of nitric acid arbitrarily fixed at lo units, and the figures are the numerical expression of a most important phase of the " chemical nature " of the substances to which they refer. [As recently as 1886 there were no methods avail- able for determining the relative strengths of the various acids, other than those depending on chemical reactions, such, for instance, as those just discussed. Just about that time it was found that the " strength " of an acid in aqueous solution is most intimately connected not only with its whole chemical behavior, but also with a great many of the physical proper- ties of the solution itself, and since then a quantitative theory has been elaborated to connect the physical with the chemical phenomena, by means of which, from a measurecnent of the melting point, freezing point, electrical conductivity, etc., of the solution of an acid, the numerical value of its strength, as defined above, may be calculated.] Table of the Relative Strengths of the Acids. 10. HCL HBr,HClO„HClO^, HBrO.„HNO.„ H ,S0^, HSCN 7.5 HIO, 6. H.J'0^ 5.5 H,SO„ H,PO, 2.5 H,,PO, 2. //,,/! sO, 1.5 C,H..O, 1. HF O.I C.,H^O„ 0.02. H.,S O.oi HCN The numerical values are approximate only : they hold for decinormal solutions at ordinary temperature and give the relative quantities of the acids present in equili- bria, such as (a) and (A) on page 14, |6 BASES. With respect to the Bases, much that has been said of the acids holds true ; the same distinction between " weak " and " strong " exists here as there, and the relative strength of the bases may be measured either by the study of their reactions, such as, for example, NaOH+KCl ••=► KOH+NaCl, or by measurements of the physical properties of their solutions in water, just as was the case with the acids. Potash, soda, and baryta are among the strongest bases ; ammonia in solu- tion {NHiOH ?) is much weaker, but the insolubility of the great majority of the inorganic basic hydrates renders an ac- curate determination of their relative " avidities " difficult, and the construction of a table for the commonly occurring bases analogous to that just given for the acids remains a work for the future. Before passing from this discussion of the " strength " or " avidity " of the various acids and bases to a consideration of the chemical behavior of the salts, and of the important part which they play in the reaction of neutralization, one point must be emphasized, namely, this: it is only in aqueous solution, and then only if no " insoluble " salt be concerned in the reaction, that the relative strengths of the acids (or bases) can be the sole factor determining the course of the reactions in which they may be involved. The ease with which the " weak " phosphoric snd boracic acids expel sulphuric, nitric, and hydrochloric acids from their salts in the blow-pipe reactions, and the everyday use of hydrogen sulphide in the laboratory to precipitate the sulphides of the Second Group, show con- clusively enough that the assistance of heat and of the insolu- bility or volatility of the products of the reaction may often enable certain of the weakest among the acids to drive out the strongest from their salts. A detailed consideration in this place of the various typical cases would, however, occasion too great a break in continuity, and has been relegated to a posi- tion after the chapter on the First Group ; to be carefully studied by the student after a little practice has rendered him familiar with the substances discussed. TS. «7 The next group is that of the Salts ; and naturally the first question to arise is : " Is there anything in the behavior of this class of compounds analogous to that which under- lies the classification of the acids and bases according to their strength ?" The answer to this question is contained— partly, at all events— in what has already been said on the subject of " strength " and "weakness"; and as a clear conception of the meaning of these terms is of great importance in under- standing the reactions met with in analysis, attention is par- ticularly directed to the following argument. In reactions of the type : (where HSand HIV sue two acids, and MW and MS their salts with the metal M), the acid FW, which is present at equilibrium in largest quantity, is called the " weak " acid, as it has been "expelled " from the salt MH^ by the " stronger" acid HS. Other (metaphorical) reasons might, however, with equal right have been given for the preponderance of the acid HW ; for instance, this-— that both acids " had a tendency to be formed," or better, perhaps, " wanted to be free," and that the more astute of the two succeeded in gaining his freedom and imprisoning the other with M. All metaphor aside, that acid which at equilibrium is present in the greatest quantity is the " weak " acid by definition ; and the relative " strengths " of any two acids have been found by experiment to be almost independent of the metal M of the last equation. If, however, any one of the salts MW, M'W, M"W, . . . MS, M'S, M"S, etc., experimented with, say M"W, for example, had a greater " tendency to be formed," was " weaker " or " more astute " than the others— or, to be more precise, if M"W stood in the same relation to the other sails, with respect to the reaction under consideration, that acetic acid does to nitric acid in the reaction with their sodium salts (page 14), then the quantity of the acid HW formed at equilibrium of the reaction HS+M"W ^=.>' HW+M"S v^ould necessarily be less than in the reactions where these i8 SALTS. W" m other salts were used. In other words, if the various salts differed from one another in the same manner and to anything like the same extent as do the various acids, then the simple relations which have led to the conception of the " strength " of the acids (and the same holds for the bases) would not have existed. This peculiarity of the Salt solutions is at the bottom of the most striking feature of the whole system of chemical analysis, viz.:— the existence of " tests for lead " and " tests for silver" (and, similarly, " tests for chlorides," "tests for orates," etc.), i.e., of reactions, for the most part involving precipitation, which are common to solutions of all the salts of lead and of silver, etc., respectively. This is possible only because these precipitates are almost without exception pro- duced by adding to the salt solution under examination a reagent solution of some other salt, and it would not be so generally the case if solutions of acids were used as the reagents. For instance, in the " test for lead by soluble borates," PbA+MBor = PbBor+MA, none of the salts represented having any greater " tendency to be formed " than the others, the insolubility of the lead borate {Pb Bor) determines the course of the reaction, no matter what A -and M may be. If, however, instead of a salt of boracic acid the acid itself be used, the direction of the reaction will depend not only on the insolubility of the lead PbA+H Bor ^: Pb Bor -i- HA borate, but also on the relative strengths of the two acids involved ;* and, as a matter of fact, in solutions of lead nitrate, no precipitate is produced by boracic acid. Tlus simplicity in the analytical reactions, caused by the small number of factors on which their course for the most part depends, viz., the relative solubilities of the substances involved, and the relative strength of the various acids (and * Other examples psge 50 (reactions of Group I), and page 56 (introduction to Group II). "weak" salts. 19 bases), if any be present, is somewhat marred by two circum- stances : First, the exceptional properties of a few salts, mainly deriva- tives of mercury and cadmium ; and Second, the existence of a class of compounds known as the " double salts," formed from the union of two or more such " simple salts " as have heretofore been discussed. Mercuric cyanide (soluble in water) furnishes an extreme type of a " weak" salt— if, for the want of a better, this term may be used*: solutions of cyanides convert mercuric salts quantitatively into mercuric cyanide : hydrocyanic acid, though one of the weakest of acids, completely expels nitric acid from combinations with (dyad) mercury: solutions of mercuric cyanide give no precipitates with any of the "reagents for mercury "t — in other words, equilibrium is arrived at in the reaction mentioned only when the system has passed [almost] quantitatively into the state represented on the right of the following equations : Hg{N0.,), + 2KCN = Hg{CN), + 2KN0., Hg{N0,)., + 2HCN = Hg{CN), + 2HN0,, . ' HgI„ + 2KCN = Hg(CN)., + 2KI , Among the " simple " salts, however, cases resembling even in a slight degree that just considered are extremely rare (cf. mercuric chloride, and mercuric sulphocyanide). The class of bodies known as double salts, many examples of which will be met with in the course of analytical work, are important enough to merit a short discussion here. These "double salts" are compounds formed by the union of two or more salts (hereinafter, for convenience sake, termed their " components"), familiar examples being the alums, e.g., KAl{S0^).^i2H.^0, potassium jjlumbic iodide KPbl.j, etc., etc., to which may be added for the purposes of this discus- *The new phys. chetn. theory referred to on p. 15 includes mercuric cyanide and the " weak " acids and bases in the group of bodies " slightly dissociated electro- lytically in solution. " +The only exception is the precipitation of the extremely insoluble mercuric sulphide. 90 DOUBLE SALTS, >' f m \ \' I.; sion the compounds formed by ammonia with the salts of silver, copper, etc., e.g., AgCl sNH.^ Some of these double salts, for example, alum, may be dissolved in water and crystallized from that medium ; others, again, are immediately decomposed by water into their com- ponents, and can only be dissolved in solutions of one of these; but, as might be expected, their behavior in this respect is largely influenced by the relative solubilities of the components and of the double salt itself, and the question to what extent a double salt, e.g., alum, exists as such in its solutions can be answered only by comparing the " physical " properties of the solution in question with those of solutions of each of the components separately. Measurements of this nature have been made for a great many of the double salts, and examples of every degree of decomposition in solution have been found, from the case of alum, whose solution appears to contain nothing but a mixture of the two components, to that of the double salts of potassium cyanide with the cyanides of iron (potassium ferro- cyanide and ferricyanide respectively), whose solutions show no trace of the presence of either of the simple salts from which they may be prepared. Intermediate between these two extremes — but rather nearer the ferrocyanide end of the series — come the ammoniacal silver and copper compounds above mentioned, and many of the double cyanides ; while nearer the alum end come the double salts of lead chloride with the alkali chlorides, and those formed by many of the metallic salts with the alkali citrates and tartrates. In all the instances considered there is in solution an equilibrium between the two components and the double salt and the influence of concentration or the " action of mass " is just the same here as in other similar cases. A detailed discussion of a few examples of the many complications which may be introduced into the reactions by the existence of double salts, many of them (those near the ferrocyanide end of the series) possessed of i ** great desire to be formed," will be found in the chapter en Reactions of the First Group. HYDROLYSIS. 31 A study of the reaction of Neutralization, to which the preceding five pages have been devoted, besides affording a basis for the classification of the inorganic substances accord- ing to their chemical nature, has, in connection with what has been said as to the nature of the chemical reaction itself (pp. 3-12), served to throw a good deal of light on the peculiar features of the reactions employed in chemical analysis, and on the char- acter of the system cf chemical analysis itself. A short dis- cussion of the reverse reaction, that of Hydrolysis, i.e., the action of water on some of the salts, forming from them again the acids and the bases from which they were originally prepared, will be found equally useful in defining the relations subsisting between the three classes of which so much has already been said. To begin with an example of the reaction referred to. The salt, bismuth nitrate, although obtained by dissolving the weak base bismuthous hydrate in nitric acid (and evaporating the solution), is nevertheless decomposed by pure water form- ing again nitric acid, and the (insoluble) bismuthous hydrate ; that is, the reaction 3HN0^+Bi{0H) 3<- ^3H,0-\-Bi{N0,), proceeds in either direction, according to the concentration of the nitric acid employed. If the equation just given be written in words instead of in symbols, thus : Hydrogen nitrate + Bismuth hydrate < "^ S s 8. M sd h4 '<3 »« § >■ cu (A *•* 2 1 S 9 & o lA M 1 a s 1 E s 1 s s s 4-t s" 1 >■ S ««: Ek « HJ « w "t? *^ •S * J3 8^ 2 3 O O a CO B •« I « w m en u ^ OA » -s =• o 5; ^ s ? to » = ii o vo C/l >f ^ O to to to to to oq ^ -5 ■^ a 4> ►^ r: M 2 -" > ^ C 3 THE FIRST GROUP. The metals of this group are distinguished by the great number of insoluble salts that they form. A number of these should be precipitated from pure dilute solutions of lead, silver and mercurous nitrates respectively — the reagent solutions are too concentrated, and should be diluted with their own volume of water, and about one cubic centimetre of this dilute solution used in each case. Particular attention should be paid to and a record should be kept of the colors, and state of division of these precipitates ; whether they are light, heavy, slimy, gela- tinous, granular, or crystalline ; how they behave in the light, on heating with water, with "excess" of the precipitating reagent, etc., etc. ; and any special reactions, such as the solu- tion of the " insoluble " silver salts in ammonia and of the lead salts in potash, should also be tried for the first time with material of known composition. An investigation of this nature should be carried on with pure salts of each metal as it is met with in the laboratory course ; it is only in this manner, and not from the book, that the necessary personal acquaintance with the various precipi- tates, etc., is to be gained. The following practical hints will be found useful in per- forming the various operations : Precipitation. — The reagent should be added in small quantities at a time, and between each addition the precipitate induced to subside by heating and shaking ; this should be repeated until finally a drop of the reagent produces no cloudi- ness in the clear liquid above the precipitate. Any further addition of the reagent serves only to interfere with the subse- quent course of the analysis. I^(l^(t'll(^- Filtration. — The filter paper should in every case be moist- ened with distilled water before use. As much as possible of 34 LEAD. [Group I. the liquid in the test tube should be poured on the filter paper before the main part of the precipitate is thrown on. If the liquid come through muddy, it should be poured back on the same filter, and this repeated until a clear filtrate is obtained. If the precipitate has been treated as described under " precipi- tation," no difficulty should be met with in filtration. Many precipitates, when first formed, are slimy, and stop up the pores of the filter paper ; repeated warming with water and shaking are the best remedies. Washing the precipitate. — This may be done on the filter paper by successive small portions of water from the wash bottle. The quantity of foreign matter left with the precipi- tate may be reduced almost in a geometrical ratio by each successive washing, if each fresh addition of water be made only when the last has completely drained away. The washing should be continued until an investigation of the wash-water shows that the object of the washing has been attained. The wash-waters themselves should be thrown away, and not used to dilute the first filtrate. >i»~4ri/*^(n ^hu^'/C^. LEAD. Pb. 205.36. Occurrence. — Ores : galena PbS, cerussitePiCOg, angle- site PbSO^. Metallurgy : metallic lead is obtained from galena by first roasting in air to convert part of the sulphide into oxide, and afterwards raising the temperature, when the oxide formed reacts with the undecomposed sulphide to form metallic lead, zPbO + PbS = 3P6 + SO..,. It is also ob- tained from the sulphide by reduction with charcoal or iron. In commerce : metal ; alloys, e.g., solder, and type-metal ; and litharge, PbO; red lead, P63O1; lead mttsite, Pb {NO. ^)„; sugar of lead, Pb {C.M.^0.^).^ ; white lead, basic carbonate ; chrome yellow, Pb CrO ^. ' - Atomic Weight. — Stoech. fig. (i) 1568. 9126 grm. lead yielded 2509. 7366 grm. lead nitrate ; (ii) IO9O.6085 grm. lead yielded 1596.9459 grm. lead sulphate. Specific heat, O.0314. Volatile compounds, lead chloride, and certain organic com- pounds, e.g., lead tetra-methyl. Isomorphous relations, with Pb, Ag, Hg\ LEAD. 35 calcium, strontium, barium, magnesium, manganese, zinc, and iron in MeCO^ ; with strontium and barium in MeSO^, with calcium in Me^^Cl {P0^)^. Chemical Relations. — Lead occupies a place near the foot of the fourth column of Mendelejeff s table on the left-hand side of, but not far from, the diagonal line. The compounds in which it acts as a tetrad are, however, very few : the dioxide PbO.^^ forms salts (plumbates) analogous to the stan- nates, and more remotely to. the carbonates, while plumbic chloride PbCl^ resembles stannic chloride in being a liquid at ordinary temperatures, and in being decomposed by water, though soluble in hydrochloric acid. The plumbous salts Pb A'a, in their solubilities and isomorphous relations, resemble those of calcium, strontium, and barium more than those of the magnesium group ; and though many basic salts of lead are known, and the hydrate dissolves in caustic potash solu- tion, it is nevertheless a much stronger base than stannous hydrate, as may be seen from the stability of the plumbous salts with respect to water, and from the existence of a normal carbonate. The Compounds of Lead. NAME. Color. Acetate. ..l'b(C^H^O^)^ Bromide Pb Br^ Carbonate Pb CO^ Chloride Pb Cl^ Chromate Pb CrO^ Cyanide Pb(CN).^ FerricyanPb^ Fe ^ ( CN ) ^ P'errocyan.Pb^ Fe ( CN)^ Hydrate Pb (OH)^ Iodide :...Pb I^ Nitrate Pb (NO^)^ Oxalate PbC^O^^ Oxide PbO Phosphate ..Pb^(PO^).^ Sulphate Pb SO^ Sulphide Pb S Sulphocyan . . Pb(SCN)^ Thicsulph Pb S^O.^ white white white white yellow white brown white white yellow white white yellow white white black yellow white Solubilityin loo parts cold hot water. II (10°) OS 0.0003 0.74 insol. iiisol. s. sol. insol. R. sol. 0.08 48(10") insol. insol. insol. 0.005 insol. insol. sol. 42(102°) 30 insol. S-o insol. insol. s. sol. insol. s. sol. o-S 139(100°) insol. insol. insol. s. sol. insol. insol. dec. Remarks. Sugar of lead. S. ex. Sudacarb. ppts 2PbC0^Pb{0H)^ S. ex. Insol. chromic acid (cf. Ba). S. ex., repptd. on boiling. Insol. dil. acids. At 130°, Pb0\ soln. reacts alk. S. ex., repptd. by much water. Other oxides, Pi^a O3 , /'<^3 O4, /^i Oa S. sol cone, nitric acid. Least soluble salt of lead. Basic upon boiling. Forms PbS upon boiling. *For solubilities in salt solutions, etc., see page 50. 36 LEAD. [Group I. Lead is a bluish-white, soft, malleable metal, sp.gr. 11.38, m.p. 334''C. Tarnishes in air at ordinary temperatures from formation of plumbous oxide Pb..O. Easily soluble in dilute nitric acid, with evolution of nitric oxide NO, slightly soluble in cone, sulph. and hydrochl. acids, insol. in dil. sulph. acid. Plumbous compounds* PbX.. Colorless in solution unless the acid be colored. Plumbous sulphide. — PbS. Hydrogen sulphide from solu- tions of plumbous salts strongly acidified with hydrochl. acid precipitates brick red or yellow sulphides, eg., Pb.^SCi.^, analogous to the basic salts. Potassium plumbite. — Potash, not in excess, with soluble plumbous salts forms plumbous hydroxide, Pb{OH)„, white, easily soluble in excess of potash to form potassium plumbite, K.^PbO., : formed also by addition of potash to all insol. plumbous salts except the sulphide and ferricyanide. Potassium plumbate. — Dilute nitric acid with red lead forms dark brown amorphous plumbic oxide ; also formed by the action of an alkaline solution of bleaching powder on plumbous chloride. Plumbic oxide dissolves in hot cone, potash solution to form potassium plumbate K.^Pb 0^. Plumbous hydroxide Pb {OH).^ in a solution of potassium plumbate produces a yellow precipitate, Pb.^O^H.^0, converted into red lead on heating. Oxidation and Reduction.— Lead is converted into plumbous and higher oxides by igniting i . air. Plumbous com- pounds are oxidized to plumbic by the action of nitric acid, bleaching powder, etc. Plumbic compounds with dil. acids or reducing agents such as glycerine, sugar, and oxalic acid, yield plumbous compounds or metallic lead. Plumbous and plumbic solutions deposit metallic lead on aluminium, cad. mium, magnesium, zinc, etc. {Cf. position of lead in electro- chemical series.) Blowpipe Reactions. — B. B. on C. with carb. soda, all lead compounds, are reduced to metallic, soft malleable globule. •Most insoluble sails of lead may be produced by adding to a soluble lead salt a soluble salt of the required acid, e.^if., lead acetate witft soda carbonate precipitates lead carbonate (b^sic). Pb, Ag, Hg\ LEAD. 37 Borax and microcosmic salt beads : O.F. yellow, hot ; colorless, cold. (Platinum wire spoiled in the reducing flame !) Basic Salts, These are bodies intermediate in composition between the salts and the bases ; on treatment with acids they react, form- ing the normal salt — hence the name given to the group. Only the weaker bases form basic salts. The basic salts of lead may be prepared by the following methods (which are also for the most part applicable in the case of other metals) : (i.) By boiling lead salts with the hydrate, filtering hot and allowing to crystallize ; the basic nitrate Pb{NO.^).,. Pb {OH)., may be obtained in this manner in long needle-like crystals. (ii.) By adding to the solution of a lead salt, potash, in quantity insufficient to convert it altogether into the hydrate : — the hydrate formed reacts with the unaltered salt. (iii.) By treating many insoluble salts of lead with ammonia or with potash in quantity insufficient to dissolve the salt ; e.g., red basic chromate of lead, from the yellow chromate. The reaction ends in an equilibrium between ammonia and the normal salt on the one hand, and the basic salt and the salt of ammonia on the other. (iv.) From the salts of lead with weak acids, by hydrolysis ; lead carbonate, for instance. This mode of preparation, like (ii.), may be considered as due to a secondary reaction between the unaltered salt, and the base formed by hydrolysis. These salts are often obtained in the form of amorphous muddy-looking precipitates, whose composition varies accord- ing to the mode of their formation ; and as continued washing very often converts them entirely into the hydrates (hydrolysis), it is next to impossible to obtain them in a fit condition for quantitative analysis ; — to set up chemical formulae for such indefinite substances would be a mere waste of time. In some few cases, however, notably in that of the basic nitrates of lead, a large series of crystalline salts has been isolated. 38 SILVER. [Group I. It has been proposed to represent these basic salts by "structural formulae" analogous to those of organic chemistry —thus, instead of PbiNO^).,. Pb{OH).„ would be written NO.,-Pb-OH, intermediate between NO^-Pb-NO.^ and HO-Pb-OH; instead of Pb{NO^),. Pb{OH).,. ^PbO, the structural formula NO.^-Pb-0-Pb-O-Pb-OH ; while the formula of " white lead," instead of 2PbC0^. Pb(OH).,. would become HO-Pb-0-CO-O-Pb-O-CO-O-Pb-OH. As the molecular weights of most of these substances are unknown, the question of their formulae cannot be regarded as definitely settled. SILVER. ^^ = 107.12. f4 Occurrence. — Ores : native ; silver glance, AgS ; pyrargy- rite, Ag.,SbS.i ; in varying quantities in most specimens of galena. Metallurgy : metallic silver is obtained from its ores according to three different processes : either it is alloyed with lead and the lead removed by oxidation, or it is amalgamated with mercury and this removed by distillation, or, lastly, it is brought into solution as a salt, and the metal precipitated by means of copper. In commerce : alloys, e.g., silver coin ; lunar caustic, AgNO^. Atomic Weight. — Stoecli. fig. (i) 397.3177 grm. silver chlorate yielded on reduction 298.4230 grm. silver chloride ; 'A^/Oii^^ •WTiji^' grm. silver yielded I287.7420 grm. silver chloride. r^Specific heat, O.0570. Isomorphoiis relations, with sodium in Me^SOi^ ; with the alkalies in silver alum. Chemical Relations. — Silver is usually placed in the first column of Mendelejeff s table. Copper, silver, and gold forming a group intermediate in properties between nickel, palladium, and platinum, on the one hand, and zinc, cadmium, and mercury on the other. The similarity of silver to sodium is confined to the formulae of their salts, — both of the type MA' — to the isomorphism of a few of their compounds, and to the existence of aquadrantoxide Ag^O, andof a silver ultramarine. Pb, Ag, Hg\ SILVER. 39 The metal resembles copper in being one of the best conduc- tors of heat and electricity, and the cuprous salts in their insolubility and in their power of forming compounds with ammonia are very like the corresponding derivatives of silver. Silver occupies a place among the "noble" metals at the " negative " end of the electro-chemical series (see page 55), and is easily obtained in the metallic state from its salts. With the exception of the nitrate and acetate, all the ordinary salts of silver are insoluble in water ; many soluble double salts, however, are known. The Compounds of Silver. Solubility in 100 parts NAME. Color. cold hot Remarks. water. Acetate AgC^ff-^O., white S. sol. s. sol. Bromide Ag Br pale yel. insol. insol. Insol. in ammonia. Carbonate Ag^ CO., yel. wh. 0.003 Heated converted into oxide. Chloride Ag CI white insol. insol. sol. ammonia. Chromate Ag.^CrO.^ red brown s. sol. s. sol Cyanide 4g(CN).i white insol. insol. S. ex. Ferrocyan.AgJ'e \CN)p^ yel. wh. insol. dec. Heated forms ferricyanide. Ferric y an .A^.^Fe (CNj^ red brown insol. insol. Hydrate Ag OH ? black [vide oxide] Forms oxide on heating. Iodide As I yellow insol. insol. S. ex»; insol. in ammonia. Nitrate As^ NO.^ white 122 Ill Oxalate Ag.^C^fi^ white insol. insol. Oxide Ag.^0 gray brw'n 0.3 Other oxides Ag^O, AgO. Phosphate.... ..Ag.^PO^ yellow insol. insol. Sulphate ^gi^O^ white OS 1-5 Sulphide Ag.^S black insol. mst.l. Least sol. silver salt. Sulphocyan Ai^SC white insol. insol. S ex. Thiotulph ....Ag.^S^O^ white 50(16=) dec. S. ex., heated forms sulphide. ■It Sol. pot. cyan.: a1! insol. silver sails except sulphide. Sol. tartrrles : carbonate. Sol. ami7ionia and inammon. chlor. : all insol. silver salts except iodide and sulphide. i"Up/\,H^CrO^ +Pb{NO^)^, and the other, 2NaNO^-\-PbCr0^^z=>'Na^CrO^+Pb{NO^)^, may be seen from the following considerations: In the latter case, if solutions of the substances on the left hand of the '•^^fc. sign be mixed in equivalent quantities, just about one-half of the lead chromate will be converted into nitrate ; REACTIONS OF GROUP I. SI whereas in the former, owing to the great disparity in ** strength " between the two acids (nitric and chromic), the lead will be converted almost entirely into the nitrate ; and if the two solutions be now each separately shaken up with lead chromate, a great deal more of that salt must dissolve in the former case than in the latter, before the concentration of the lead nitrate, etc., formed, is great enough to establish equilibrium in a solution saturated with the chromate. This case, consequently, differs from the two others con- sidered in this very important particular, viz. : it is not alone the solubility of the *' insoluble " salt, but also the strength of its acid, that determines the amount dissolved at equilibrium. " If lead chromate will d'^-'olve in potassium nitrate or in sodium chloride," it was argued before, " then lead sulphate, being more soluble in water, will dissolve still more readily in those solutions." In the case just considered this conclusion will not hold ; the greater strength of the sulphuric acid more than counterbalances the greater solubility of the lead sulphate, and the sulphate of lead is very little soluble in nitnc acid, just as the phosphate, oxalate, and chromate, etc., are insoluble in the weak acetic acid. If the acid of the insoluble salt be easily volatile at the temperature of the experiment, or if it readily decompose* giving rise to a volatile decomposition product, these will be important factors in determining the course of the reaction — e.g., solution of carbonates, nitrites, sulphites, thiosulphates, etc., in acids ; similarly the greater volatility of sulphuric, nitric, and hydrochloric acids, at high temperatures, enables the non-volatile phosphoric and boracic acids to expel the former from their salts in the blowpipe reactions. Other examples under this head are : the solution of the basic hydrates in acids, and that of the insoluble lead salts in potash. The solutions of precipitates in acids or salt solutions (excluding the case where the acid of the insoluble salt is volatile) have some interesting properties. The solution forti^ed by digesting lead sulphate with nitric acid until no more will dissolve, and filtering, gives a precipitate with sul- 52 REACTIONS OF GROUP I. phuric acid on the one hand, and with a concentrated solution of nitrate of lead on the other ; in both cases the equilibrium being disturbed in such a manner as to result in the formation of more lead sulphate. A consequence of this behavior is that if a little lead nitrate solution be mixed with a compara- tively large quantity of concentrated nitric acid, and then dilute sulphuric acid added drop by drop, no precipitate will be pro- duced by the first drops, and it is not until the quantity of sulphuric acid added is greatly in excess of that able to combine with the lead dissolved that the latter is even approximately separated in the form of sulphate. The addition of water to the saturated solution of lead sulphate in nitric acid produces no precipitate, for though it decreases the concentration of the nitric acid it decreases equally that of the other three substances involved. Quite the contrary is the case when the solution contains a double salt ; in this case, for o«« substance whose concentration is diminished on one side of the equation two suffer dilution on the other, the equilibrium is disturbed, and a reaction takes place, resulting in the formation of these two. If, as is the case, in the equilibrium. NaCl+PbCl a <- NaPbCl 3» one of the components, e.g., PbCl^, be insoluble, the water added may prove insufficient to retain in solution the new quantity of this substance formed, and in such a case (which frequently occurs) addition of water will produce a precipitate. Reactions involving Two or More Insoluble Salts. (a) Two insoluble salts. — If a fairly concentrated solution of potassium chromate be poured on a small quantity of lead carbonate* the conditions for the first moment are precisely the same as in the case discussed on page 50 (where lead *The formula of lead carbonate is written PbCO^ for simplicity, the precipitate formed by sodium carbonate in lead salt solutions is really a basic carbonate. REACTIONS OF GROUP I. 53 chromate was the insoluble and sodium nitrate the dissolved salt). But long before the equilibrium K^CrO^+PbCO 3 <- ■.^ K^COa+PbCrO^ can be attained, the solution is saturated with chromate of lead, and the further progress of the reaction results only in the separation of this salt in the solid form, and the continual solu- tion of fresh quantities of lead carbonate, until all of the latter has disappeared. Then, for the first time, it is possible for the concentration of the latter to fall below the limit fixed by its solubility, and the reaction comes to a standstill when the con- centration of the lead carbonate dissolved, taken in connection with that of the remaining potassium chromate, is low enough to enable the resulting potassium carbonate and the quantity of lead chromate that can dissolve to balance the equilibrium. The result may be expressed in words by saying that the potassium chromate has " transposed " the lead carbonate — the white precipitate of the carbonate has been replaced by a yellow precipitate of the chromate of lead. A complete " transposition " of the carbonate is, however, not always the result of the reaction of potassium chromate on that salt. As long as the two *' insoluble " salts are both present in the solid form, their concentrations in the solution depend, practically speaking, only on their solubiHties and on the temperature, and is fixed by the conditions of the experi- ment ; with the chromate and carbonate of potassium, on the other hand, the case is quite different ; during the progress of the reaction the concentration of the former is continually growing less, and that of the latter greater ; so that if a large quantity of lead carbonate be present to begin with, or if the solution of potassium chromate employed be dilute, the series of reactions just considered may come to an equilibrium before all the lead carbonate has disappeared — " partial transposition "; — while the effect of largely increasing the concentration of the potassium carbonate (by addition of that salt) is to reverse the reaction and from the chromate to regain the carbonate of lead — "reverse transposition." In order that a transposition may be " reversible " the two 54 REACTIONS OF GROUP I. "insoluble" salts must not differ too much in their solubilities. The only way to compensate fot the greater insolubility of the lead chromate in the case just considered is to employ a very concentrated solution of potassium (or sodium) carbonate; and a limit is set to this by the limited solubility of the alkali car- bonates in water. Thus, though ammonium sulphide trans- poses lead carbonate instantly with formation of the black sulphide, no solution of carbonates that it is possible to prepare is concentrated enough to effect the reverse change. An important corollary to what precedes is the conclusion that, if lead nitrate be added to a mixture of ammonium sulphide and sodium carbonate, all in solution, the precipitate will consist of lead sulphide only, until all the ammonium sulphide is used up. Similarly, when silver nitrate is added to a mixture of sodium chloride and potassium chromate, the red precipitate of silver chromate is not formed until all the chlorine has been removed from the solution. Here, as elsewhere, if acids be made use of as "reagents" in place of salt solutions, the distinction between "strong" and "weak" must not be lost sight of. Although sodium chloride cannot transpose lead chromate at all, and sodium sulphate only partially, the corresponding acids (hydrochloric and sulphuric) effect the change at once. In view of the similar cases already discussed, a detailed explanation of the reactions will be omitted. (6) Four insoluble salts. — Reactions involving more than two insoluble salts are not often met with in the course of analysis. One instance, however, may be mentioned here. When the two " insoluble " salts, lead chloride and silver chromate, are mixed under water a change of color (red to yellow) indicates that chromate of lead has been formed, and it will readily be seen that equilibrium in the reaction P6C/2 ^-AgiCrO^ ^znn*" PbCrO^+2AgCl can be reached only when one or other of the salts on the left hand side of the equation has been completely dissolved. And it is equally clear that there is no means of reversing this result, unless, perhaps, at some different temperature the relative solubili- ties of the salts involved should be reversed. REACTIONS OF GROUP 1. 55 The " Electro-Chemical Series." A special case of the reaction involving two insoluble sub- stances is that in which both of these are metals : e.g., in the formation of the lead tree by tin in lead acetate solution, or of the silver tree by zinc in silver nitrate, also the deposition of mercury on a strip of copper in solutions of mercury salts, or of copper on a knife blade in cupric sulphate solutions. Zn+CnSO^ =Cu+ZnSO^ Reactions of this nature are of especial interest, as they a.T^ made use of in the voltaic cell to generate an electric current, and the great majority of them can be reversed by electric.il means. Equilibrium in these reactions depends not only on electrical forces, but also on the concentration of the salt solutions employed ; it takes a less electro-motive force to reverse the reaction formulated above {i.e., to precipitate zinc from its salts by copper) if the zinc salts be concentrated and the copper salt dilute than in the opposite case ; in general, however, it is not possible to reverse the reaction by alteration in the concentrations alone. As long as one of the ordinary [" strong "] salts be used, it does not matter which, the direction of the reaction remains the same, and it is possible to arrange the metals in a series — the electro-chemical series — such that each metal is precipi- tated from its salts by all those preceding it ; if, however, salts Mg,Al, Zn, Cd, Fe, Ni, Pb, Sn, Cu, Sb, Bi, Ag, Hg, Pd, Pt, Au such as mercuric cyanide* be employed, or if double salts be involved which have a great "tendency to be formed," the position of the metals in this series may be alteKd ; in solutions of potassium cyanide, for instance ^which form such double salts with many of the metals), the arrangement in the electro-chemical series is as follows : Zn, Cu, Cd, Sn, Ag, Ni, Sb, Pb, Hg, Pd, Bi, Fe, Pt. 'Introduction, page 19. THE SECOND GROUP. Commoner elements — {Pb), Hg", Cu, Bi, Cd, As, Sb, Sn. Rarer elements — A u, Pt, Pd, Ru, Ir, Rh, Os,J'e, Se, W, Mo. s4 Metals whose sulphides are insoluble in dilute, 2 — 3%, hydrochloric acid, and whose chlorides are soluble in water. The members of this group are precipitated as sulphides by hydrogen sulphide from the filtrate from Group I. The precipitate is often slimy, and tends to stop up the pores of the filter paper ; boiling for a few minutes with water and filtering hot is the best remedy. CULK"^ wl-rf/W^tt ^a^kt Although the solution into which the hydrogen sulphide is led must contain free hydrochloric acid, in order to bring about the precipitation of arsenic and to prevent that of members of the third group, it must not be too acid, many of the sulphides of the second group being soluble in strong hydrochloric acid. This is specially noticeable in the case of tin and antimony ; one of the methods actually employed for the preparation of pure hydrogen sulphide being to act on antimonious sulphide with concentrated hydrochloric acid : in other words, the equation 2S6C/3 +3^25 = 56,53+6^67, which represents the reaction in a slightly acid solution, must be reversed to represent that which takes place when the sulphide is treated with concentrated acid. Sb„S^ +6HCl = 2SbC/.^ +3H^S. Consequently, in order to ensure the complete conversion into their sulphides of the metals of the second group, hydrogen THE SECOND GROUP. 57 sulphide must be passed in until a portion of the filtrate after dilution with water gives no further precipitate with the gas. Mixed with the sulphide in the second group precipitate is often found a quantity of free sulphur, arising from the oxida- tion of a portion of the hydrogen sulphide. This indicates the presence of oxidizing agents — nitric, chromic, or arsenic acids, ferric chloride, etc., etc. — in the solution, and may in great part be prevented by boiling the filtrate from Group I. with a few drops of alcohol ; if free nitric acid is known to be present, the solution should be evaporated nearly to dryness and water added before leading in the hydrogen sulphide. Group II. A. As.Sh.Sn. The members of this group resemble one another in that their sulphides, with solutions of the sulphides of the alkali metals '«//oie' ammonium sulphide contains ammonium ortho-sulphantimoniate. If now to a solution of an alkaline sulphantimonite or sulphantimoniate there be added an acid, e.g., sulphuric acid, a salt (sulphate) of the alkali will be formed and one of the sulphides of antimony precipitated, with the liberation of hydrogen sulphide. 2(NH,)^SbS^ + 3H,SO,=3{NH,),SO,+Sb,S^+3H,S 2{NJ\),SbS^+3f^tSO,=i{NH,)^SO,+Sb,S,+3H,S reactions analogous to the decomposition of the alkaline antimonites and antimoniates by acids : — 2{NH^)^SbO^+3H,SO,^-.3iNH,)„SO.-irSb,0,-^3H,0 2{NH^)^SbO^ + 3H„SO,=3{NH^hSO^+Sb„0,+3H^O In the case of arsenious sulpnide the reactions are : Soln. in yd. am. sulph : As^S^ + ^iJVffiU 5a = 2(AW4)3^j5a + 3S, Pptn. by dil, sulph. acid : 2(Nat}^AsS3 + 3ff,SOi = 3(NIfi)^SOi + AsiSa + 3H^S for the stannous sulphide : Solution : SnS+2NfftSff+Sit={N/fi)2SnS3 + //^S Precipitation : (NffJ^SftS^ + H^S04,=^{NHi)iS0i + SnS^ + H^S The sulphides of this group are also soluble in solution of potassium hydrate (and some of them in ammonia, sodium carbonate, ammonium carbonate, etc), forming a mixture of sulpho and oxy-salts, from whicii, by the addition of acids, there is icprecipitated the same sulphide that was dissolved. For example : As, S6,Sh.] ARSENIC. 59 ARSENIC. As = 74.52. OccuKRENCE. — Orcs : realgar, As^S^; orpiment, ^s^Sg ; and mispickel, or arsenical pyrites, FeAsS. Metallurgy: the sulphide is roasted with iron, whereby the arsenic is set free and passes off in the form of vapor. In commerce : white arsenic, As.,0^; fly powder, suboxide ; Paris green and Scheele's green (arsenites of copper). Atomic Weight.— S/o^cA. fig.: 22.173 grtn* arsenic tri- chloride converted 39.597 grm. metallic silver into silver chloride. Specific heat, O.0830. Volatile compounds : the metal ; arsine; arsenic trioxide, ^45^ O3 and /I S40a; arsenic trichloride, bromide, etc.; and certain organic compounds. Isomorphous relations : with antimony and bismuth in the free state, and in Mcc^S^ ; with antimony in Me^Og, ^Ag^S.Me^Sa and Cu^MeS^ ; with phosphorus in K^MeO^ and in the apatite group. Chemical Relations. — In its molecular weight {As^\ in the formulae and crystalline forms of its compounds, and in the existence of a gaseous hydride, arsenic closely resembles' phosphorus and antimony, its next neighbors in the fifth column of Mendelejeff's table. Like other members of the same column — bismuth excepted — the hydrates of its two oxides, As^Og and ^45405, have acid properties ; and the chloride AsCl^ (.AsCl^ unknown), though soluble in a small quantity of water, or in the presence of hydrochloric acid, is decomposed by much water with formation of arsenious and hydrochloric acids. The sulpho -acids of arsenic are referred to on page 58. The Compounds of Arsenic. Arsenic is a steel-grey brittle metal, sp. gr. 4.71, volatile at 450° C. ; not oxidized in dry air, but gradually coated with arsenious anhydride ^4X303 in moist air at ordinary temperatures ; heated, it burns to form arsenious oxide. Arsenic dissolves readily in nitric or nitro-hydrochloric €o ARSENIC. [Group II. A. acid, slowly in cone, sulph. acid (forming arsenic acid), not in hydrochloric acid ; soluble in solution of potash. NAME. ARSE^riOUS. Bromide As Br^ Chloride ■ .As Cl^ Hydride 4s H^ Iodide As I^ Oxide As^ O3 Sulphide As^S^ ARSENIC Hydrate H^AsO^\H^O Oxide As^O^ Sulphide As:^S^ Color. Solubility in 100 parts cold hot water. white dec. white sol. gas Svols(o°i) brick red s. sol. white 4 yellow insol. white 16.7 white ISO yellow insol. dec. dec. insol. dec. 9.'S insol. 50 very sol. insol. Remarks. m.p. 20' C, b.p. 220° C. tn.p. \%''C.,h.p. I34''C. See page 68. Sublimes. Sol. dil. hydrochl. acid. Sol. amtn. sulph. and in [alk. carb. Other sulph: realgar As^S,i Arsenious compounds AsX^, colorless in solution. Arsenious sulphide, AsoS.^ ; hydrogen sulphide, (or sodium thiosulphate on boiling) in solutions of arsenious acid, acidu- lated with dilute hydrochloric acid, precipitates arsenious sulphide. In the absence of hydrochloric acid no precipitate is formed, but the solution is colored yellow. Aisenic compounds AsX^, colorless in solution. Arsenic sulphide, As^S^ ; precipitated from boiling hydro- 'Chloric acid solution of arsenic acid by hydrogen sulphide. At ordinary temperatures, sulphur, arising from the reduction of arsenic to arsenious acid, is at first precipitated ; when tliis reduction is complete, arsenious sulphide is precipitated. Arsenites (derivatives of ihf^^sOa, etc.), obtained by dissolv- ing metallic arsenic, arsenious anhydride, or other arsenious compounds, in potash. A solution of an arsenite, e.g., arsenious acid neutralized with a few drops of ammonia, gives the following characteristic reactions : (i.) Ferric hydrate precipitates all the arsenic, as basic ferric arsenite — (antidote for arsenic !). (ii.) Silver nitrate produces a yellow precipitate of silver arsenite, Ag-^AsO^ ; soluble in Ammonia and in nitric acid. (iii.) Cupric sulphate yields no precipitate; on careful addi- tion of potash, however, a green precipitate, CuHAsO^ As, Sd, Sn.] ARSENIC. 6l (Scheele's green), is formed ; addition of ammonia instead of potash produces Paris green NH^CuAsO^. (iv.) Stannous chloride, in hydrochloric acid solution, pre- cipitates amorphous arsenic, mixed with tin (brown black). Arseniates (derived from //3/lsO^, etc.), obtained by the oxidation of arsenious acid and arsenites. A solution of arsenic acid, neutralized with a few drops of ammonia, gives the following reactions : (i.) Silver nitrate produces a red-brown precipitate of silver arseniate, Ag^AsO^, soluble in nitric acid and in ammonia. (ii.) Ammonium molybdate and nitric acid, on warming, {cf. phosphoric acid) yield a yellow precipitate of arsenio- molybdate of ammonia, {NH j.^AsO ^ + ^L2 M ^^0.^; soluble in ammonia, reprecipitated by nitric acid. (iii.) Ammonium chloride, ammonia, and then magnesium sulphate precipitate white crystalline ammonium-magnesium- arseniate (MgNH ^)AsOi. Oxidation and Reduction. — Arsenious acid is oxidized to arsenic acid by bromine water, chromates, iodine, cone, nitric acid, permanganates, etc. H^AsO^+Cl^-\-H^O = H ^AsO^-{■ 2HCI. Arsenic cpds. are reduced to arsenious by potassium iodide in hydrochl. acid solution, by sulphurous acid, hydrogen sulphide, sodium thiosulphate, ferrous chloride, stannous chloride, etc. Arsenic and arsenious acids are reduced to metallic arsenic by stannous chloride in hydrochl. acid solution, and by metallic copper, cadmium, zinc, etc. Blowpipe Reactions. — B. B. on C. with carb. soda or potass, cyan.; all compounds of arsenic volatilize; the presence of arsenic is recognized by an odor of garlic. In a bulb tube, with potassium cyanide, free arsenic condenses and collects at the mouth of the tube, forming a steel-colored mirror. ANTIMONY. 56 = 119.00. ' Occurrence. — Ores : stibnite, S62S3; allemontite, As^Sbf^ ; and Valentin "te, Sb^itO^. Metallurgy : the metal is obtained 6a ANTIMONY. [Group II. A. from the sulphide by heating with iron filings ; in the preparation of tartar emetic the sulphide is first roasted to convert it into the oxide, and then dissolved in cream of tartar. In commerce : metal ; alloys, e.g., type metal, pewter, Babbit metal, etc.; tartar emetic C^H^K{ShO)Oa ; and Naples yellow (basic plumbous antimoniate). Atomic Weight. — Stoech. fig.: I4.4802 grm. antimonious sulphide reduced with hydrogen yielded IO.3476 grm. metallic antimony. Specific heat, O.052. Volatile compounds : antimonious chloride, bromide, and iodide, antimony trioxide, and certain organic compounds. Isomorphous relations : with arsenic in Me„0^, in Me^S^, in ^Agc,SMe^S,^,\n Cu^MeS^, with bismuth in Me^S;^. Chemical Relations. — Antimony occupies a place in the fifth column of MendelejefTs table, and in forming a gaseous hydride ShH^, acids derived from the two oxides Sb-^O^ and Sbi^O^, and in the isomorphous relations of many of its compounds it closely resembles nitrogen phosphorus and arsenic. It stands below them in the column, however, and in forming a normal sulphate, and nitrate, and many basic salts, shows analogy to the more metallic bismuth. Like arsenic, germanium, and tin, it forms soluble sulpho-salts with the alkalies (see page 57). The Compounds of Antimony. Solubility in loo part> • • name. Color. cold 'lot Remarks. water ANTIMONTOUS. Bromide Sb Br^ white dec dec. ni.p. 90° C; b.p. 280° No ShBr^. Chloride Sb Cl^ white sol. dec. m.p., j2''C;b.p. 230; also.WC/5 Hydrate Sb(OH)^ white insol. insol. Sol. in alkali hydrates. Iodide Sbl^ dec. dec. Trimorphous:red,green,and yel. Nitrate Sb\,NO,,\ white dec. dec. Basic salts with water. Oxide Sb^O^ white insol. insol. Other oxides SbO^, Sb^ 0^ . Oxychloride Sh OCi white insol. insol. Sol. f "l. hydrochl. acid. Sulphate Sb,(SO^)a white dec. dec. Basic salts with water. Sulphide Sb^S^ orange insol. insol. Also Sb^ S^. Hydride SbH^ gas 4vol(lo"') insol. See page 68. u4s, Sir, Sn.] ANTIMONY. 63 Antimony is a bluish-white, lustrous metal, sp. gr. 6.97, m.p. 432^C. At ordinary temperatures it oxidizes slowly in moist air ; it burns with formation of antimonious oxide, Sb^O^. It dissolves slowly in hot hydrochl. or sulphuric acids ; it is acted on by nitric acid with formation of oxides, and by nitro-hydrochloric with formation of oxides and chlorides. Antimonious compounds SbX^, colorless in solution; mostly insol. in water owing to formation of basic salts, soluble in acidulated water. Antimonious sulphide Sb.yS^, precipitated from not too strongly acidulated solutions of antimonious compounds by hydrogen sulphide, sod'um sulphide, and sodium thio- sulphate on boiling. In neutral solutions a coloration only is produced. Aniimonious acid and antimonites. — Potash, ammonia, or ammonium carbonate, precipitates from solutions of anti- monious salts, in the absence of tartrates and citrates, — Antimonious oxide Sb^O^, soluble in excess of the reagent to form potassium, (or ammonium), antimonite KSbO„. Upon evaporation, or upon long standing, this solution decomposes with separation of antimc ious oxide. Potassium perman- ganate, and potassium dichromate, in acid solution, are reduced (decolorized) by antimonious acid. Metantimonic acid; HSbO^; white, insol., obtained by treating metallic antimony with an excess of cone, nitric acid, and evaporating ; soluble in potash. Pyroantimonic acid, H^Sb^O^. — Metantimonic acid dis- solves in hydrochl. acid, forming a solution of antimony penta- chloride, Sb^Cl^, from which addition of much water precipi- tates pyroantimonic acid. Oxidation and Reduction. — Antimonious compounds are oxidized to antimonic by action of bromine water, nitric acid, chrcmates, permanganates, etc. Antimonic compounds are reduced to antimonious by the action of hydrogen iodide or of stannous chloride. Antimonic and antimonious com- pounds are reduced to metallic antimony by metallic zinc. 64 TIN. [Group II. A. As, Blowpipe Reactions. — B. B. on C. with carb. soda or potassium cyanide, all antimony compounds are reduced to a white brittle globule of metallic antimony. This rapidly oxidizes, producing white fumes, and leaving a white crystalline deposit on the charcoal. TIN S«. = 116.78. Occurrence. — Ores : the only ore of tin is cassiterite, or tin stone, SnO^. Metallurgy : the ore is first roasted to drive off sulphur arsenic and other impurities, then reduced with charcoal, and finally purified by distillation. In commerce : metal ; alloys, e.g., solder, bronze, Britannia metal, etc. ; various compounds used in dyeing, e.g., tin salt SnCl^2H20, preparing salts, Na^SnO^ + sH^O, and pink salt, SnCl^ + zNH ^Cl. Atomic Weight. — Stoech. fig.: 45.8323 grm. of tin oxi- dized by nitric acid yielded 58.2519 grm. tin dioxide. Specific heat, 0.055. Volatile compounds : stannous chloride ; stannic chloride, bromide, and iodide. Isomorphous relations, with titanium, silicium, and zirconium in MeOz ; with silicium, etc., in K^MeFg. Chemical Relations. — In Mendelejeffs table tin stands between germanium and lead, and, like them, forms two series of compounds, stannous SnX^, and stannic SnX^. Stannous hydrate is soluble in potash, and also in strong acids, in this resembling the hydrate of lead ; it is, however, a much weaker base than the latter, all its salts being decomposed by water, with formation of basic compounds. Tin dioxide behaves as the anhydride of an acid, " stannic acid," H-^^SnOs, corre- sponding to plumbic acid, and more remotely to silicic and carbonic acids. The tetrachloride resembles those of the other members of the same column, in being a liquid with a low boiling point. Both it and the dichloride form crystalline compounds with water, but are decomposed by much water ; they are soluble in hydrochloric acid. Like germanium and the other members of Group II. A, it forms a series of sulpho-salts (see page 58). Js, S6, Sn.] )r a y e TIN. The Compounds of Tin. 65 NAME. Color. STANNOUS. Bromide Sn Br^ 2H^ O Carbonate (basic) Chloride Sn Cl^zH^O Ferrocyan. . . . Sn^Fe (CN) Ferricyan. .Sn^{Fe(CN)f^)^ Hydrate Sn(OH) Iodide Sn I^2H^0 Nitrate. .Sn(NOn)^.2oH^O Solubility in loo parts cold hot water. Oxalate Sn C^O^ Oxide Sn Phosphate (basic) Sulphate Sn SO^ Sulphide SnS STANNIC. Chloride Sn Cl^ Hydrate H^Sn 0^ Nitrate Sn(NO^)^ Phosphate (basic) Sulphate. ..Sn(S0^)^2H^0 white white white w. (gela.) white white red. yel, white white black white white br'wn bl liquid white white white white sol. dec. 271 insol. insol. insol. dec. deliq. insol. insol. insol. sol. insol. sol. s. sol. dec. insol. dec. dec. dec. dec. insol. insol. insol. dec. deliq. insol. insol. insol. sol. insol. dec. s. sol. dec. insol. dec. Remarks. Dec. by much water. Carb. soda ppts. hydrate. Dec. by much water. Sol. in alkalies. Sol. in alkali iodides. Dec. by much water. Other oxide, SnO^. Sol't'ns deposit basic salts Other sulphide, SnS^ yel. 5«C/4.S^a0cryst. Sol. in alkalies. Sol. in tartrates : all insol. stannous compounds except the sulphide. Tin is a lustrous,white metal, sp.gr. 7.30, m.p. 2^2^C.; tar- nishes slowly in air ; burns at white heat to form stannic oxide SnOz' Dissolves slowly in dil. hydrochl. and sulph. acids, rapidly if these acids be hot and concentrated, readily in dil. nitric acid, or in nitro-hydrochl. acid. Stannous compounds, SnX.^, colorless in solution. Stannous hydrate SniPH)^. — The fixed alkalies, ammonia, alkali carbonates, barium carbonate, and potassium cyanide precipitate stannous hydrate ; soluble in excess of fixed alkali (jtot in ammonia or carbonates, distinction from antimony), forming stannites, e.g., K^SnO^. Stannic compounds, SnX^, colorless in solution. Stannic hydrate H ^SnO ^ (stannic acid). — The fixed alkalies, ammonia, alkali carbonates, sodium sulphate, and potassium iodide precipitate stannic hydroxide insol. in ammonia and ammonium carbonate, soluble in alkali carbonates, forming stannates. I 66 TIN. [Group II A. Metastannic acid, HioSn^Oi ^{7).*— The action of nitric acid on metallic tin is important. Very dilute acid con- verts it into stannous nitrate, with simultaneous formation of ammonium nitrate ; cone, nitric acid, on the other hand, con- verts it into metastannic acid, which has the same composition as stannic acid, but differs from it, however, in the following respects : (i.) it is completely insol. in nitric acid ; (ii.) boiled with cone, hydrochl. acid it is converted into metastannic chloride insol. in the cone, acid, but soluble in water, from which solu- tion it is precipitated upon addition of sulph. acid ; (iii.) boiled with solution of caustic soda it forms sodium metastannate, insol. in excess, but soluble in water; (iiii.) by fusion with alkalies it is converted into salts of ordinary stannic acid ; on the other hand, stannous chloride boiled with hydrochl. acid, or even on standing, forms metastannic acid. Oxidation and Reduction. — Stannous compounds are oxidized to stannic by nitric acid, sulphuric acid (hot), chlorine, bromine, iodine, mercuric chloride, etc. (Stannous chloride is a powerful reducing agent, reducing even nitrous acid, sulphurous acid, and potassium ferrocyanide.) Stannic compounds are reduced to stannous by metallic copper and tin. Stannous and stannic compounds deposit metallic tin on aluminium, cadmium, magnesium, zinc, etc. Blowpipe Reactions.— B. B. on C. with carb. soda (more readily upon addition of potassium cyanide), all com- pounds of tin are reduced to a malleable metallic globule. At the same time, a slight white incrustation, SnO.^, is formed on the charcoal, which, moistened with cobalt nitrate and ignited, becomt;s blue-green. 'Isomeric with stannic acid, molecular weight unknown. SEPARATION IN GROUP II A. In order to separate the members of Group II A from the Group II. precipitate, the latter, after being well washed, is treated with a little (2 — ^cc.) of the yellow ammonium sulphide solution, warm, but not boiling, according to either of the two methods described for the solution of lead chloride on page 45. 1^ port'"" ^ the solution so obtained is filtered off and acidified with dilute sulphuric acid. If the precipitate consist merely of light yellow, milky, finely divided sulphur, burning without residue, nothing has dissolved in the ammonium sulphide, Group II A is absent ; if, however, a curdy yellow to brown* precipitate be formed, the whole of the filtrate must be treated with excess of dilute sulphuric acid, and the precipitate analyzed for members of Group II A, while the residue from the ammonium sulphide solution must be thoroughly washed with successive small quantities of warm ammonium sulphide before testing for Group II B. These washings should be thrown away, and not used to dilute the Group II A solution. Arsenic may be separated from Antimony and Tin either (a) by warming the sulphides As^S^,SbiS^,SnS^ with ammonium carbonate solution, which dissolves the sulphide of arsenic, leaving the other two; or (b) by warming them with concentrated hydrochloric aoid, which leaves the arsenic and dissolves the others. Neith-*r of these methods is very satisfactory, however ; in the former a little antimony, and in the latter a little arsenic, goes into solution. Tin may be separated from Arsenic and Antimony by dissolving the sulphides in warm dilute hydrochloric acid, with addition of a crystal of potassium chlorate, boiling the solution to expel free chlorine, adding sodium thiosulphate and boiling again ; whereupon the arsenic and antimony are precipitated <*' *The dark color is due to the presence of cupric sulphide, which is slightly soluble in yellow ammonium sulphide. Solutions of the polysulphides of potassium and sodium do not dissolve the sulphide of copper, but dissolve that of mercury. 68 SEPARATION. [Group II A. as the sulphides As^S.^ and 56.. Sa, while the tin remains in solution. To bring about complete precipitation the sodium thiosulphate must be present in large excess, and the boiling must be continued some time. From the precipitate the arsenic may be removed by either of the methods given above. The only really accurate separation, however, is that known as " Marsh's Test." Nascent hydrogen reduces com- pounds of tin to the metal, while those of arsenic and antimony are converted into the gaseous hydrides Arsine AsH.^, and Stibine SbH ^, respectively. The two gases may be distinguished by their action on silver nitrate solution, or by burning them and examining the products of their combustion. The methods of separation and identification based on these facts go by the name of " Marsh's Test," and may be carried out as follows : (a) Silver nitrate test. — A flask of 200 cc. capacity, the "generator," is fitted by means of a perforated cork to a glass tube twice bent at right angles and dipping into a dilute solution of silver nitrate in a test-tube ; (a second tube, the " funnel tube," widened at the upper end and passing through the cork almost to the bottom of the flask, is very convenient for adding successive small portions of the liquid to be analyzed). Into the generator is put a little pure granulated zinc ; and the substance supposed to contain arsenic, antimony, or tin, dissolved in excess of dilute hydrochloric or sulphuric acid, is added. If arsenic or antimony be present, the gas evolved will produce a black precipitate of metallic silver or silver antimonide, respectively, AsH^ + 6AgNO^+3H^ 0= SAg-^H^AsO^ + eHNO^ SbH^ + iAgNO, = Ag^Sb+sHNO^ in the silver nitrate solution.* As, however, the precipitate might consist of silver sulphide (due to hydrogen sulphide from sulphur compounds in the generator), the presence of the two elements first mentioned must be proved by a further examination as follows : the black precipitate is re- * In concentrated (50%) solutions of silver nitrate, arsenic produces a yellow pre- cipitate of Ag^As^ which is decomposed by water into arsenious acid and metallic silver. As^ S6, Sft.] SEPARATION. 69 moved from the solution by filtration, washed, warmed with dilute hydrochloric acid Ag,Sb+6HCl = SbCl^ + ^AgCl+3H„ and the solution tested for antimony ; while the filtrate, after being freed from silver by the addition of dilute hydrochloric acid and filtering, is tested for arsenic. Any tin present will remain in the generator, the contents of which should be dissolved in dilute hydrochloric acid, and the solution tested for stannous salts by mercuric chloride, potassium perman- ganate, hydrogen sulphide, etc. (6) Combustion test. — For this purpose the bent tube of the generator should be replaced by a short upright tube, drawn out at the upper end to a fine opening — the gas jet. Before igniting the escaping gas, a portion should be collected in a test-tube, inverted over the jet and conveyed with the orifice downward, and closed with the thumb, to the neigh- borhood of a flame ; if the gas take fire with a slight pufT, and then burn quietly, the apparatus is free from air; if, how- ever, the contents of the test-tube ignite with a whistling noise, the air has not been completely driven out of the generator, and the application of a light to the jet would, in all probability, cause an explosion. When the apparatus is finally freed from air, the gas should be ignited at the jet (if the generator be provided with a funnel tube, this is the time to add the solution), and a porcelain evaporating dish held in the flame. If arsine or stibine be present, the gas will burn with a bluish-white flame (^45^03, Sbi^Os), and black spots of me- tallic arsenic or antimony will be formed on the porcelain. The arsenic spots are iJrown-black, and shining, soluble in alkaline solution of sodium hypochlorite, and turned yellow by warming with ammonium sulphide. In dilute nitric acid they dissolve, and if silver nitrate, and then a little ammonia, be added to the drop of nitric acid solution, a yellow precipitate Ag^AsO^ is formed. The antimony spots, on the other hand, are dull black, insol- uble in the hypochlorite, and turned orange by ammonium sulphide. Dilute nitric acid turns them white, Sb^Og, and, if to this white spot silver nitrate and then ammonia be added, it is blackened with formation of Ag^O, more quickly on warming. to BISMUTH*. [Group II B. BISMUTH. Bt=207.34. Occurrence. — Ores: native; bismuth ochre, Bi^O^) and bismuth glance, Btj S3. Metallurgy: metallic bismuth is obtained from its ores by roasting, and subsequent reduction with iron or charcoal. In commerce : metal ; alloys, e.g., Wood's metal, Rose's metal, etc.; and " magisteriumbismuthii," Bi(NO^)„ sH.O. Atomic Weight. — Stoech. fig. : (i.) 28.5866 grm. bismuth oxidized by nitric acid yielded 26.310 gr bismuth oxide ; (ii.) I6.6450 grm. bismuth oxide yield<.J 25.2551 grm. bismuth sulphate. Specific heat, O.0288. Tsomorphous relations : with arsenic and antimony, in the metal and in Me^S^. Chemical Relations. — Bismuth, standing at the foot of the fifth column in Mendelejeffs classification, forms no hydride, but resembles the other members of the column in its derivatives of the types BiX^ and BiX^ . Bismuthous hydrate is a weak base (c/. the hydrates of Group III A) ; its salts with the stronger acids are all decomposed by water (with formation of basic salts), while no salts with the weaker acids, e.g., thiosulphuric, hydrocyanic, etc., are known ; it does not, how- ever, act as an acid toward potash. The hydrate of the pent- oxide HBiO^ {cf. HNO^) is a very weak acid; its salts with the alkalies have been prepared by the action of bromine or chlorine on bismuthous hydrate suspended in potash or soda solution ; halogen compounds of the type BiX^ have not been prepared. The Compounds of Bismuth. Bismuth is a reddish-white, brittle metal, often showing iridescence from superficial oxidation, s^.^r 9.9, m.p. 26y°C.,b.p. I300°C. At ordinary temperature not oxidized in air ; burns to form bismuth oxide, Bi^O^; dissolves in cold nitric acid without evolution of hydrogen, in hot cone, sulph. acid with evolution of sulphur dioxide ; insol. in dil. hydrochl. and sulph- uric acids. Hg, Bi, Cu, C./.] BISMUTH . 7» Solubility i n too parts NAME, Color. cold hot Remarks. water. BISMUTHOUS. Acetate ..Bi(C^H^O^)^ white sol. sol. Bromide Bi Br^ yellow deliq. dec. Carbonate (basic) white inso . insol. Dec. above ioo°C. Chloride BiCl^ white dec. dec. Oxychlor. sol. dil. hydrochl. Chromate (basic) yellow insol. insol. Sol. dil. acids. [acid. FerrocyanBiJ^Fe{CN)t).i wh. yel. insol. insol. Insol. in dil. acids. Ferricyan . Bi Fe( CN )a br. ye!. insol. insol. Insol. in dil. acids. Hydrate Bi (OH)^ white- insol. dec. T orms oxide on heating. Iodide Bi I^ NitrateBi(NO^)^Srf^O dark brown sol. dec. S. ex. white dec. dec. Basic salt sol. dil. nitric. (iC, Oxalate.... Bi^{C„Oi)s Oxide Bi^O^ white insol. dec. yellow insol. insol. Other oxides BiO, Bi.^O^, Phosphate BiPO^ white insol. insol. [/?,-, o„ Sulphate .... Bi^{SO^^ white dec. dec. Basic salt sol. sulph. acid. Sulphide Bi^S^ black insol. insol. S. tartrates : chromate, ferricyanVierrocyan, and hydrate. Bismuthous compounds. — BiX^, colorless in solution. Bismuthous hydrate, Bi {OH)^, precipitated from solutions of bismuthous salts by alkali hydrates and by potassium cyanide. Bismuth monoxide, BiO, black, precipitated from solutions of bismuthous compounds by addition of potassium stannite* or by addition of solution of dextrose and heating. Bismuthic acid, HBiOn, red. Obtained by addition of bromine water to a cone, solution of potash in which bismuth- ous oxide is suspended ; it decomposes readily on heating. Oxidation and Reduction. — Bismuthous compounds to bismuthic by heating in the air, or by the action of chlorine or bromine in cone, potash solution. Bismuthous salts are reduced to bismuth monoxide by potassium stannite, and by dextrose ; bismuthous compounds deposit metallic bismuth upon copper, cadmium, iron, lead, tin, zinc, etc. Blowpipe Reactions. — B. B. on C. with carb. soda, all compounds of bismuth are reduced to a brittle easily fusible globule, producing at the same time an incrustation, orange yellow, hot ; lemon yellow, cold. Borax and microcosmic salt beads : in both flames, faintly yellow, hot ; colorless, cold. 'Prepared by the addition of potash to stannous chloride in sufficient excess to dissolve the precipitate, Sn(OH)^, at first formed. 7* COPPER. [Group II B. COPPER. C» = 63.i2. OccuRRENCii. — Ores : native ; red copper ore Cn„0 ; mala- chite, Cn(OH).,CuCO.y, copper glance, CuS; and copper pyrites CuFeSs- Metallurgy: the metal is obtained from the oxide by reduction with charcoal ; the sulphide is first roasted to convert part of the sulphide into oxide, then upon raising the temperature the oxide formed reacts with the undecom- posed sulphide to set free copper, according to the equation 2 C«0-f C«5 = 3C«-f-SOy. In commerce: metal; alloys, e.g., brass, bronze, German silver, etc. ; blue vitriol, CuSO^^H^O ; and Paris green, CuHAsO^. AtomicWeight. — Stoech.fig.: (i.)20.327O grm.cupric oxide yielded 16. 2279 grm. metallic copper, (ii.) Electrolytically, one grm. of copper is equivalent to 8.406 grm. of silver. Specific heat, 0.093. Isomorphous relations : with iron, nickel, cobalt, manganese, zinc, and magnesium, in MeSO^jH^^O ; with these together with mercury and cadmium, mMeSO . jK" 2 SO 4 Sfi^O.* Chemical Relations. — Copper forms two series of salts : the cuprous, CuX, white, mostly insoluble in water, are com- parable to the salts of silver, to the mercurous, aurois, and to the thallious salts ; while the cupric, CuX„, are more like those of the magnesium group ; numerous insoluble basic cupric salts are known, the solutions of the normal salts react acid to litmus, and the hydrate loses water and is converted into the oxide below 100° C. Of the third series of derivatives of copper, corresponding to the oxide Cu„0^, little is known. Copper salts combine with ammonia. The metal itself is one of the best conductors of heat and electricity. The Compounds of Copper. Copper is a soft, malleable, reddish metal ; sp. gr. 8.94, m. p. 1054° C. In dry air at ordinary temperatures unchanged ; in moist air it becomes gradually coated with basic carbonate. It dissolves in acids without evolution of hydrogen : slowly in *Ammonia may replace potassium in these double salts. ^c. yiv. Cu, G/.] COPPER. 73 dilute hydrochloric and sulphuric acids, readily in dilute nitric acid, and with evolution of sulphur dioxide in hot concentrated sulphuric acid. NAMK. CUPKIC Acetate Cii{ C^H^ 0^)^lH^O Bromide Cu Br^zH.fi Carbonate (basic) Chloride Cu Ci.,2H^0 Chromate (basic) Cyanide Cu\CN) Ferrocyan . . . Cu j Fe{ C„JV)„ Ferricyan . . . C'« „ ( Fe CJIV„ ) Hydrate ' .Cu (0H)\ Nitrate. . . Cu (NO.^).fiH^O Oxalate . ., Cu t\ t)^ Oxide Cu O Phosphate Cu^{PO^) Sulphate ...Cu SO^sH^O Sulphide Cu S ■Sulphocyan Cu {SCN)^ Thiosulph Cu S^O^ Color. Solubility in loo parts cold hot water blue black blue or grn. green br. yellow br. yellow red br. yel. grey pale blue blue blue wh. black blue grn. blue black black white 7 (20") deliq. s. sol. 60 s. sol. insol. insol. insol. insol. very sol, insol. insol. insol. 40 0.0001 insol. insol. 20(100 ) deliq. s. sol. very sol. s. sol. insol. insol. insol. (^ec. very sol. insol. insol. insol. 203 insol. insol. dec. Remarks. Green basic bromides. Forms oxide on heating. Cone, solutions green. S. ex. ; dec. on standing. Insol. in dil. acids. Insol. in dil. HCl, Forms oxide on heating. Forms oxide on heating. Heated evolves oxygen. S. sol. amm. sulphide. Dec. on warming. Dec. on warming. .Sol. in ammonia : all insol. cupric salts except sulphide. Sol. in pot. cyan. : all insol. cupric salts. Sol. in tartrates : all insol. cupric salts except the sulphide and ferrocyanide. Sol. in ammon. chl. : cupric cyanide and carbonate. Cuprous compounds. — CuX, colorless in solution (except "where the acid is colored), insoluble in water ; easily soluble in concentrated ammonia, concentrated hydrochloric, hydro- bromic, etc., acids; reprecipitated upon addition of water ; no sulphate or nitrate known. Cuprous cyanide CuCN, white, insoluble. In solutions of cupric salts, potassium cyanide produces a red precipitate of cupric cyanide ; converted on boiling into cuprous cyanide with evolution of cyanogen gas, CnN„. Cuprous hydrate CuOH, yellow, insoluble ; from alkaline solutions of cupric salts by solution of dextrose ; heated, red cuprous oxide {formed. Cuprous iodide Cul, white, insoluble; from solutions of .cupric salts by potassium iodide ; (precipitate colored dark brown from separation of free iodine). This precipitation is 74 COPPER. [Group II B> Hg, complete only in the presence of such reducing agents as sodium thiosulphate, ferrous sulphate, etc. Cuprous sulphocyanate, CuSCN, white, insoluble. Potassium sulphocyanate in solutions of cupric salts precipitates cupric sulphocyanate (black), converted into cuprous sulphocyanate upon warming, or in presence of sulphur dioxide. Cuprous thiosulphate, Cu^S^O^, white, insoluble. Sodium thiosulphate in concentrated solutions of cupric salts precipi- tates cupric thiosulphate, yellow, which upon standing changes slowly into cuprous thiosulphate, white, and finally into cuprous sulphide, black. Cupric compounds CuX^, anhydrous, white; dilute solu- tions, blue; concentrated solution of cupric chloride, green. The dilute blue solution of this salt becomes green upon addition of concentrated hydrochloric acid. Salts of cupric acid (derived from Cu„0^ ?). Cupric nitrate added to water holding bleaching powder in suspen- sion produces a precipitate at first green, afterwards violet,, which, upon standing or washing, becomes converted into blue cupric hydrate. Oxidation and Reduction. — Cuprous compounds are oxidized to cupric by the action of heat and by acids. Cupric compounds ate converted into cuprates by the action of bleaching powder. Cupric compounds are reduced to cuprous by the action of stannous chloride, arsenious acid', or dextrose in alkaline solution, also by potassium iodide, potassium cyan., and potassium sulphocyan. Cupric and cuprous solutions deposit metallic copper upon aluminium, cadmium, iron, lead, tin, zinc, etc. See p. 55. Blowpipe Reactions.— 5. B. on C, with carb. soda, all copper compounds are reduced to metallic copper. Borax bead : O.F., green, hot ; blue, cold ; R.F., colorless, hot ; reddish brown, cold. Microcosmic salt bead : O.F., green, hot ; blue, cold ; R.F., green, hot ; reddish brown, cold. All copper compounds, moistened with dil. hydrochloric acid, impart a green or greenish blue color to the flame. Hg, Bi, Cu, Cd.1 CADMIUM. 75 CADMIUM. Crf=lll.i8. Occurrence. — Ores : greenockite, CdS ; to 'he extent of two or three per cent, in most zinc ores. Metalluyyy : obtained as by-product in the manufacture of zinc. In commerce : metal ; alloys, e.g., Wood's metal, cadmium amalgam ; cad- mium iodide, Cdl„ (photography) ; and cadmium yellow, CdS, Atomic Weight. — Stoech. fig.: 64.2051 grm. cadmium sulphate yielded 44.4491 grm. cadmium sulphide. Specific heat, 0.0567. Volatile compounds : metallic cadmium, cadmium chloride, bromide, and iodide. Isomorphous relations : with copper, magnesium, zinc, mercury, iron, nickel, cobalt, and manganese in MeSO ^,K ^SO ^,6H „0. Chemical Relations. — In Mendelejeff's table cadmium stands in the second column, between zinc and mercury ; it resembles the former in its position in the electro-chemical series, in the appearance of the metal, and in the isomorphous relations of some of its salts. It forms but one series of salts,^ whose solutions react acid to litmus ; the chloride is not decomposed by water, and no insoluble basic salts are known. Unlike zinc, it does not dissolve in potash. Cadmium iodide is a rather "weak " salt (see Introduction, page 19). The Compounds of Cadmium. Cadmium is a white malleable metal, sp. gr. 8.67, m.p. 320. Not oxidized in air at ordinary temperatures; burns in air to foim cadmium oxide, CdO (brown). Dissolves slowly in concentrated hydrochloric and sulphuric acids with evolution of hydrogen, readily in concentrated nitric acid with evolution of oxides of nitrogen. Ail compounds of cadmium are colorless in solution ; except those whose acids are colored. Oxidation and Reduction. — All compounds of cadmium deposit metallic cadmium upon ^inc, magnesium, and alu- minium. 76 CADMIUM. [Group II h. Solubility in loo parts NAME. Color. cold hot water. Remarks. ^ce/afe Cd{CiH^O^),^.bH„0 white very sol. very sol. Bromide CdBr^zH^O white deliq. deliq. Carbonate Cd CO^ white insol. insol. No basic carbonate. Chloride Cd Cl^iH^O white 140 150 Chr ornate CdCrO^ yellow very sol. very sol. Cyanide Cd[CN)^ white insol. insol. S. ex. Ferrocyan .... Cd.^Fe (CiV)^ white insol. insol. Sol. dil. hydrochl. acid. Ferricyan ...CdJFel, ON) e ) ^ yellow insol. insol. Sol. dil. hydrochl. acid. Hydrate Cd (OH)^ white insol. insol. S. alkali hydrate (c/. Zn) Iodide Cdl^.iH^O white 93(20°) 133(100°) Nitrate . ..Cd{NO.i),,. a,H^ white deliq. deliq. Oxalate CdC^O^ white 0.007 O.OI sol. ammonium oxalate. Oxide CdO brown insol. insol. Phosphate Cdgf^PO^^ white insol. insol. Sol dil. acids. Sulphate ..ACd SO^.p/^ white 60 very sol. Nocomp'd CdSO^,.^a,J. Sulphide CdS yellow insol. insol. Sol. is%hotsulph. acid. Sol. in ammonia and in pot. cyan.: all insol. cadmium salts except sulphide. Sol. in tartrates : all insol. cadmium salts except ferrocyan., phosphate, and sulphide. Sol. amnion, chlor.: all insol. cadmium salts except ferrocyan., hydrate, and sulphide. Blowpipe Reactions. — B. B. on C. with carb. soda, all compounds of cadmium are reduced. The metal vaporizes, forming an incrustation brown in the oxidizing flame, volatile in the reducing flame. Borax and microcosmic salt beads : in both flames, yellow, hot ; colorless, cold. irLfr CjlScJU^^^^Mi «:. ddt^^^^^^c^i'' ^*^/^ 7^ SEPARATION IN GROUP II B. To the precipitate of the sulphides of Group II B, washed, and, if necessary, freed from members of Group II A by the treatment described on page 67, add dilute nitric acid and warm. Residue : Mercuric sulphide, black, soluble in aqua regia. Sulphur :-iHNO^+H„S = ?). Solution : Nitrates of lead, bismuth, copper, and cadmium. To a part of the solution in nitric acid add a little dilute sulphuric acid ; if no precipitite be formed, even after addition of alcohol shaking and standing for a minute. Lead is absent ; if a precipitate be formed, the reagent must be added to the whole of the nitric acid solution, and the precipitated lead sulphate removed by filtration. Then add ammonia in excess White precipitate : Bismuthous hydrate. Blue sol'ition : Copper (soluble cupric ammonium salts). To detect the presence of Cadmium : If copper be absent, add to the ammoniacal solution (freed from the hydrate of bismuth by filtration) ammonium sulphide. A yellow precipitate indicates cadmium, CdS. If copper be present, one of two methods may be adopted. (a) To the blue solution from which any precipitate of bismuthous hydrate has been removed by filtration, add potas- sium cyanide until decolorized, then ammonium sulphide ; a yellow precipitate indicates cadmium (the black cupric sulphide is soluble in potassium cyanide solution). (6) Precipitate the sulphides of copper and cadmium from the blue solution (using hydrogen sulphide or ammonium sulphide), filter and wash the precipitate and warm it with dilute sulphuric acid, which dissolves the sulphide of cadmium, leav- ing that of copper undissolved ; filter, and test the filtrate for cadmium. For Confirmatory Tests, see the description of each metal. THE THIRD GROUP. Commoner elements.— Fe, Cr, A I, Ni, Co, Zn, Mn. Rarer elements. — Ur, Be, Th, Zr, Ce, Ti, Y, V, etc. vl al The third group consists of those metals which are not precipitated by hydrogen sulphide in dilute hydrochloric acid solution, but which are precipitated — some as sulphides* others as hydrates — by ammonium sulphide from a solution containing ammonium chloride. The more commonly occurring metals of this group fall naturally into two classes : Group III A , those precipitated as hydrates, viz., iron, chromium, aluminium ; and Group III B, those precipitated as sulphides, viz., nickel, cobalt, manganese, zinc. To separate these, the Group III precipitate may be dissolved in hydrochloric acid, and the members of the first subdivision precipitated either by treatment with barium car- bonate in the cold, or by boiling with sodium acetate. It is, however, much more convenient to precipitate the hydrates of class A directly by adding ammonium chloride and ammonia to the filtrate from the second group (any iron present being previously oxidized by nitric acid), thus avoiding altogether the use of ammonium sulphide solution in Group III A. The use of ammonium chloride prevents the precipitation of magnesia and of members of Groups IV and ""/ b, ihe ammonia. A certain amount of the chloride of ammonium is formed in any case, when ammonia is added to tiii^ filtrate from the second group (containing hydrochloric ac; '.) ; lit this, while preventing complete precipitation of magnesna, manganese, etc., might be insufficient in quantity to keep them altogether in solution ; and the addition of a further portion of ammonium chloride guards against the incon- THE THIRD GROUP. 79 venience of obtaining these metals partly in Group III A, and partly in Groups IV and V, respectively. Group III A. Iron, chromium, and aluminium are precipitated as hydrates of the formula M{OH)^. These three, however, are not the only metals to form hydrates of this type, others, for example, being cobalt, nickel, bismuth, and antimony. All of these hydrates are alike in being very weak bases — none of them can drive ammonia from solutions of its salts, and some of them even act as acids toward potash — a similarity that appears to depend rather on the formula of the hydrate than on the position of the metal in Mendelejeff's table, as in general, where an element forms more than one hydrate, that derived from the highest oxide is always the more acid or less basic* As has been pointed out in the Introduction, this weakness of the bases entails striking consequences in the chemical behavior of the solutions of their salts; the chlorides are all decomposed by water, though soluble in hydrochloric acid, the sulphides and most of the soluble salts are decom- posed by water, especially on boiling, forming either the hydrate or some basic compound, while no carbonate3,t nitrites, sulphites, thiosulphates, etc., are known ; the car- bonates, thiosulphates, etc., of the alkalies bringing about a precipitation of the hydrates. The triacid bases fall into two natural classes : (i.) Iron, chromium, aluminium, cobalt, and manganese ; forming alums and spinelles. (ii.) Bismuth and antimony; connected by the isomerism of their oxides, sulphides and salts, with arsenic and phosphorus. The latter, by reason of the insolubility of their sulphides in dilute hydrochloric acid, fall in Group II ; the reason that the others, iron, chromium, and aluminium, come down in Group III A, while cobalt and manganese pass into the next sub-group, lies in the difference between * Compare, for example, the hydrates corresponding to the monoxide, sesquioxide, and trioxide of chromium, and those of the various oxides of iron, manganese, bismuth, ^gen. Potassium ferrocyanide, K^Fe{CN)^, obtained by dissolving ferrous cyanide Fe{CN)^ in excess of potassium cyanide, or Al, Fe, Cr.] IRON. H I by the action of potassium cyanide on metallic iron or its oxides. A solution of ferrocyanide of potash upon oxidation yields potassium ferricyanide K a Fe{CN)^. A solution of either of these salts gives none of the characteristic reactions of iron ; heated with concentrated sulphuric acid, however, the double salts are decomposed and the corresponding sulphates are formed. Solubility in loo parts NAME. Color. cold hut Remarks. water. FERROUS. Acetate ... .Fe(C^HgO^)<^ Bromide. . . . FeBr^lH^O white very sol. very sol. green sol. very sol. Carbonate Fe CO., Chloride Fe Cl^./^H^O grey white insol. insol. Dec. in air to hydrate white 130 very sol. Cyanide FeiCN).^ yel. red insol. insol. S. ex. forms A\Fe(CN)n Ferrocyan Fe^Fe(CN)f^ light blue insol. insol. Dark blue on standing. Ferricyan ..Fe^(Fe(CN)f^)^ dark blue insol. insol. Turnbull'sbluen.s. acids Hydrate Fe (OH)^ grn. white insol. insol. Iodide Fel^.sn^O dark grn. very sol. dec. Nitrate. .Fe (N03)^.6H^O white sol. dec. Oxalate Fe CjO^ white 0.02 0.03 Sol. in dil. acids. Oxide Fe O black insol. insol. Phosphate Fe^( F0^)^.8ff^ blue wh. insol. insol. Sulphate Fe SO^.J/f^O green 60 333 Green vitriol. Sulphide FeS black insol. s. sol. Other sulphides : FeS... Thiositlph ..FeS^Oa-SHiO green very sol. very sol. - 1 A mm. Sulphate. ...(6U^0) pale green 17 very sol. FERRIC. Acetate Fe(C^H^0^)3 dull red very sol. dec. Bromide Fe Br^ red s. sol. sol. Chloride Fe Cl^ garnet red deliq. deliq. Low temp. Fe^Cl^, Chromate (basic) brown insol. dec. Ferrocyan. . Fe^(Fe{CN)a)i dark blue insol. insol. Prussian blue, insol. dil. Ferricyan Fe Fe( CN)^ brown sol. sol. [acids. Hydrate Fe {OH)^ red brown insol. insol. Insol. alkali hydrates. Nitrate... Fe(N0.^)^.\%H.t0 white very sol. very sol. Oxalate Fci(C^Ot)^ white insol. insol. Sol. oxalic acid. Oxide />sO, grey insol. insol. Other oxide Fs^O^. Phosphate ...Fe PO^.^H^ yellow wh. insol. insol. Sol. in dil. acids. Sulphate..Fe^(SO^)a.^H^O red brown very sol. dec. SulphocyanFe(SCN) ^iH^O red deliq. deliq. ^ Amm. Sulphate. ..(2^H^0) white 14(20°) 400(100°) Ferric ammon. alum. Sul. in amm. chlor. : ferrous carbonate and hydrate. Sol. in pot. cyan, (hot): all insol. ferrous salts; ferric cyanide, ferrocyan., and phosphate. Sol. in tartrates : ferrous carbonate, hydrate, phosphate (diflf. sol.) ; ferric cyanide, ferrocyan. , hydrate, and phosphate. ,%, ^. IMAGE EVALUATION TEST TARGET (MT-3) 1.0 1.1 l££|2| |2.5 2.0 1^ IU4 1 l.25|,.4_,,.o ^ 6" ► VQ Photographic Sciences Corporation 23 WfST MAIN STRKT WEBSTER, N.Y. 14580 (716)872-4303 m v iV <^ ;^ ^ 86 CHROMIUM. [Group III A. Ox'DATioN AND REDUCTION. — Fcrrous compounds are oxidized to ferric, by the oxygen of the air (hence reagent bottle of ferrous sulphate invariably contains ferric sulphate), by chromic acid, potassium permanganate, silver nitrate, chlorine, bromine, nitric acid, etc. loFeSOi, -{-sKMnOi ^-SH^SO^ = sFe^iSOi)^ +K^S0i+2MnS0i-V8H^0 Ferric compounds are reduced to ferrous, by the action of iron, zinc, or tin, by stannous chloride, hydrogen sulphide, ammonium sulphiije, dextrose, etc. 2Fea^-\-{NHJ^S = 2FeCl^+2NH^Cl+S^ Blowpipe Reactions. — B. B. on C. with carb. soda, all compounds of iron are reduced to a black metallic powder attracted by the magnet. Borax bead: O.F., red, hot; yellow, cold ; R.F., colorless. Microcosmic salt bead : bottle green, hot ; almost colorless, cold. CHROMIUM. Cr=51.77. Occurrence.— Or«: chrome iron ore, FeCr^O^ ; croco- isite, P6C/'04. Metallurgy: chrome iron ore, mixed with quick- lime, is heated in a reverberatory furnace with free access of air, whereby calcium chromate is formed ; dilute sulphuric acid converts t^his into soluble dichromate, which with potassium carbonate yields red chromate of potash (potass, dichromate), from which all chromium compounds areobtained. In commerce : potassium (and sodium) dichromate, K„Cr20j ; chrome alum,. K^SO^.Cr^{SOjQ.24H^O; chrome yellow, PbCrO^. Atomic Weight. — Stoech. fig. : 6.6595 S^"^' 5 ammonium bichromate, upon heating, yielded 4.0187 grm. cttiomic oxide. Specific heat, O.122. Volatile compounds : chromic trichloride, CrC/g ; chronious chloride, CrCl^ ; chromium oxychloride, CrO,,Cl^. Isomorphou^ relations: with sulphur and iron, in sulphates and ferrates ; with iron and aluminium, in alums, spinelles, and oxides (c/. iron). Chemical Relations. — Chromium forms three series of compounds: chromous CrX^, chromic Cr^a, and chromates* Al, Fe, Cr.] CHKOMIUM. ^ The chromous compounds are very easily oxidized, e.g., by the air — and are not met with in analysis. They resemble in chemi- cal behavior, formulae, and crystalline form, the corresponding ferrous compounds. The chromic compounds resemble tho^je of the sesquioxides of aluminium, iron, manganese, and cobalt, there being a series of chromic alums and of members of the spinelles group. In their basicity the chromic hydrates stand between alumina and ferric hydrate, some forms being soluble and others insoluble in potash. The chromates correspond to the sulphates, the dichromates to the pyrosulphates {not to the bisulphates, or acid sulphates), chromium trioxide to sulphur trioxide, chromyl chloride to sulphuryl chloride, etc. The more metallic nature of chromium is evidenced by the existence of alkali chlorochromates undecomposed by cold water. The Compounds of Chromium. Solubility in loo parts NAME. Color. cold hot Remarks. water. CHROMIC. Acetate Ct\C^H^O^).^ very sol. dec. Basic acetate deposited. Bromide Cr Br^ dark gr. very sol. very sol. Dissolves slowly. Carbonate . . . ^ (basic) red br. insol. insol. Chloride CrCl., dark red very sol. very sol. Dissolves slowly. Cliroinate Cr 0^ br. yel. insol. dec. Cyanide Cr(CN)^ bl. grn. insol. insol. S. ex. on warming. Hydrate Cr(OH)^ pale bl. insol. insol. Sol. cold potash. Nitrate. . . Cr {JV0:i)„9H^ purple very sol. very sol. Oxide Cr^ 0^ green insol. insol. Other oxides Cr O3, Cr 0. Phosphate Cr^{PO^)^ green insol. inscl. Sulphate 0-.,(.VC> J,,. iSiV^O violet sol. sol. Sulphide Cr„S^ grey grn. insol. insol. Fuse hydrate and sulphur. Chromyl chloride.. Cr O^Cl^ red dec ■ dec. Pot. chrom K^CrO^ \ellow 5° 6o Pot. dichrom K^Cr.^0,, red lO ICX) Chrome alum purple 14 Sol'n grn. on heating. Sol. in pot. cyan.: all insol. chromic salts. Sol. in tartrates : cyanide and hydrate. Chromium is a light green, crystalline, difficultly fusible powder ; sp. gr. 7.3, m.p. above that of platinum ; not oxidized in air at ordinary temperatures; heated in air, or in the presence of potash, it forms the oxide ; it dissolves readily in 88 CHROMIUM. [Group III A. dilute hydrochloric and sulphuric acids, but is insoluble in nitric acid. Chromic compounds CrXa, exist in two modifications in solution : green and violet ; only the latter crysta'Mze. Chromic chloride, CrCl a', obtained by reducing potassium dichromate or chromic acid with hydrogen sulphide or alcohol in the presence of dilute hydrochloric acid. In the latter case the solution should be heated until the color changes to a pure green, after which the excess of acid and alcohol must be boiled off. 2K.,Cr^O,+6H.,S+i6HCl=4CrCl.:,-i-4KCl+iiH^O+3S^ Chromic hydrate, Cr(PH)^ ; precipitated from chromic salts {e.g., chloride) by fixed alkalies, ammonia, alkali carbon- ates, barium carbonate, ammonium sulphide, or potassium cyanide. The precipitate, is soluble in excess of cold alkali hydrate, almost insoluble in ammonia. The purple color sometimes observed in the filtrate from Group III A arises from the presence of small quantities of ammonium chromium double salts. Chromyl chloride CrOc,Cl.;.. — If equal quantities of finely divided potassium dichromate and sodium chloride be mixed, and heated wi*h concentrated sulphuric acid in a test-tube provided with a cork and gas delivery tube, the latter dipping into caustic potash solution, brown vapors of chromyl chloride will distill over ; these react with the potash to form potassium chromate and potassium chloride. Potassium chromite KCrO^, in solution emerald green ; from chromic chloride by adding potash in sufficient excess (in the cold) to dissolve the precipitate at first formed. Upon boiling this solution, or by treating with acids, ammonium salts, or hydrogen sulphide, a chromic hydrate* is precipitated (c/. aluminium). With alkaline solution of zinc hydrate, zinc chromite ZnCr„0^, yellow, is precipitated. Potassium chromate K^CrO^, from chiomic hydrate or •The hydrate precipitated on boiling a dilute solution of p tassium chromite con- tains less water than the " normal " hydrate, and is much less soluble. - M, Fe, Cr.] SEPARATION. 89 oxide by fusing with a mixture of sodium nitrate and carbonate. With acids a red color is produced, arising from the formation of dichromates. The original color is restored by the addition of alkalies. Chromium chromate Cr^^Ofi.CrO^, brown yellow; from chromic chloride with potassium chromate : consequently formed if potassium dichromate be partially reduced by hydro- chloric acid and alcohol, and alkali be added. Oxidation and Reduction.— Acid solutions of chromates or dichromates are reduced to chromic compounds (green) by almost all reducing agents, e.^., sulphurous acids, sodium thiosulphate, hydrogen sulphide, alcohol, oxalic acid, ferrous and stannous salts, etc. Chromic compounds in alkaline solution, i.e., as alkali chromitcs, can be oxidized to chro- mates by chlorine or bromine water, bleaching powder, lead dioxide, etc. Blowpipe Reactions.— B. B. on C. with carb. soda ; chromium compounds are all reduced to chromic anhydride, Cr^O.^, green. Borax and microcosmic salt beads: both flames, yellow green, hot; emerald green, cold. SEPARATION IN GROUP III A. Wash the precipitate of ferric, chromic, and aluminiuni hydrates, and boil for several minutes in a test-tube with a :Solution of potassium hydrate ; filter. Solution : potassium aluminate KAIO^ ; from this a white, gelatinous precipitate of alumina may be obtained by barely acidifying with hydrochloric acid and adding ammonium carbonate. Residue: Fe{OH)^, CrOH^. Iron may be detected by dissolv- ing a part in dilute hydrochloric acid and adding potassium sulphocyanate or ferrocyanide ; and chro- mium in another portion by oxidizing in alkaline solution, acidifying with acetic acid and adding lead acetate. 90 COBALT. [Group III B. To oxidize the chromic hydrate it may be suspended in solution of potash and warmed with either bromine water, lead dioxide, or bleaching powder ; or the hydrate may be fused with nitre and sodium carbonate, and the fused mass dissolved in water, and filtered (ferric hydrate separated). COBALT. Co = 58.4. Occurrence.— Or« ; cobalt glance, {CoFejAsS^ ; speiss cobalt, {CoFeNt)As^ ; and almost always with nickel in the ores of the latter. Metallurgy : consult text-book. In com- merce : smalt (silicate); Th^nard's blue (aluminate) ; and cobalt nitrate, Co(iV03)„ . 6//., 0. Atomic Weight, — Stoech. Jig. : upon reduction with hydrogen, 88.5908 grm. cobaltous oxide yielded 26.40 grm. metallic cobalt. Specific heat, O.106. Isomorphous relations : with magnesium, zinc, iron, nickel, and manganese, in MeSO ^yH ^O / with these, together with copper, cadmium, and mercury, in MeSO^.K^SO^.yH^O; with iron in A/e3(P0 J3.8//2O, and with iron, aluminium, etc., in alums. Chemical Relations. — Cobalt forms two series of salts : the«cobaltous salts CoX.^, which resemble those of nickel to an extraordinary degree, correspond to the ferrous salts ; the hydrate, however, is a weaker base than the corresponding hydrate of iron, and from solutions of its salts potash in small quantity precipitates basic cobaltous salts, which, by con- tinued action of water, may be converted into the hydrate. The cobaltic compounds CoX^, with the exception of certain double salts (e.g., cobaltic potassium nitrite and potassium cobalticyanide), are teadily decomposed by water, and only a few of them have been obtained in the crystalline state. A very large number of ammonio-cobalt compounds are known.. The metal itself is slightly magnetic. Co,Ni,Mn,Ztt.] COBALT. 9' The Compounds of Cobalt. NAME. COBALTOUS. Acetate Co {CiH^Ot)^ Broniide.. Co Br^.bH^O Carbonate Co CO^ Chloride ....Co Cl^.dH^O Chromate (basic) Cyanide Co{CN)i Ferrocyan . ...Co^Fe{Cl^)^ Ferricyan...Co.^(Fe(CN)a)i Hydrate Cr (OH).^ Iodide Co I^.dH^O Nitrate . . Co (NO^)i.6HiO Oxalate Co CjO^ Oxide Co O Phosphate Cos(PO^)».Sff■■ ,, NICKEL. iVt = 58.3. OccukVENCE.— Om ; nickel blende, NiS ; millerite, NiS^ ; nickeline, iVtVls; and almost always with cobalt in its ores. Metallurgy : consult text-book. In commerce : nickel matt ; alloys, e.g., German silver, armor-plates, etc. Atomic Weight.— Stoech. fig.: 28.5943 grm. nickel oxide yielded 22.4730 grm. metallic nickel. Specific heat, O.109. Isomorphous relations : with magnesium, zinc, cobalt, iron, and Co, Ni, iMh, Ztto.] NICKEL. 93 magnesium, in MeSOi.jH^O ; with these together with copper, cadmium,and mercury in MsSO^.KoSO^.GHijO; with magnesium in Me^S^O^. ChemicalRelations.— Nickel formsbut one series of salts, of the type NiX^, which strongly resemble the cobaltous salts ; from their solutions in water, however, potash does not precipi- tate basic salts. Ammonia-nickel compounds exist, but are formed much less readily than are the corresponding derivatives of cobalt. After iron, nickel has the greatest magnetic perme- ability. The Compounds of Nickel. NAME. Color. Ace/a/e Ni(C^H,0.i)^ Bromide Ni Br^ . zHiO Carbonate (basic) Chloruie NiCl^.Z //g O Chromate (basic) Cyanide Ni(CN)^ Ferroiyan Ni^Fe (CN)n Ferricyan Ni.AFe(CN)n)i Hydrate Ni(OI/)^ Iodide. Nil^.dH^O Nitrate Ni^NO^)^ . dH^O Oxalate Ni C^O Oxide Ni O Phosphate. . .Nic^PO^)^ . tH,^0 Sulphate '. . NiSo^^ .tH^O Sttlphide NiS Sulphocyan ..Ni(SCN\^ . ^^^0 Thiosulphate . Ni^S^O^ . oH^O green green green green red br. grn. wh. g:n. wh. grn. yel. pale grn. green green green black green green black br. yel. green Solubility in loo parts cold hot I6 deliq. insol. sol. insol. insol. insol. insol. insol. deliq. .5° insol. insol. insol. 37 insol. sol. sol. very sol. deliq. insol. sol. insol. insol. insol. insol. insol. deliq. very sol. insol. insol. insol. 62 (70°) insol. sol. dec. Remarks. Sol. amm. carb. Anhydrous, yellow. S. ex. Insol. dil. acids. Insol. alkali hydrates. Sol. in amm. oxal. Other oxides Ni^O^. Sol. dil. acids. Anhydrous, yellow. S. sol. ammon. sulph. Forms sulphide. SjI. in ammonia, or in pot. cyan.: all insol. nickel salts except sulphide. Sol. in ammon. chlor.: carbonate, hydrate, oxalate, and phosphate. Sol. in tartrates : all insol. nickel salts except ferrocyan. and sulphide. Nickel is a silver-white, lustrous, malleable metal ; sp. gr. 8.82, m.p. 1500" C. In air at ordinary temperatures unchanged ; heated, forms nickel oxide ; dissolves readily in dilute nitric acid, slowly in dilute hydrochloric and sulphuric acids ; insol- uble in concentrated nitric acid (cf. iron). Nickelous compounds, NiX 2 : dilute solutions and hydrated crystals, green ; anhydrous crystals, yellow. 94 MANGANESE. [Group in B. Nickel cyanide, Ni{CN)s : from solutions of nickel salts with potassium cyanide ; precipitate soluble in excess {KCN)^ Ni(CN)„, reprecipitated upon addition of a few drops of dilute hydrochloric acid (c/. cobalt). Nickel hydrate, NiiOH)^: from solutions of nickel salts with alkali hydrates or ammonia; soluble in excess of ammonia, forming compounds which, upon exposure to the air, show a change of color similar to that of the cobaltamines. From these solutions the nickel can be slowly reprecipitated by potassium or sodium hydrate (distinction from cobalt). Oxidation and Reduction. — In alkaline solution nickel- ous compounds are oxidized to nickelic oxide by chlorine water, bromine water, sodium hypochlorite, and sodium hypo- bromite. Nickelic oxide is reduced to nickelous by oxalic, nitrous, phosphorous, sulphurous, hydrochloric, hydrobromic, hydroferrocyanic, etc., acids. Nickel compounds deposit metallic nickel upon zinc, cadmium, tin, etc. Blowpipe Reactions.— B. B. on C. with carb. soda, all compounds of nickel are reduced to a metallic black powder attracted by the magnet. Borax bead: 0. F., purple- red or violet, hot ; yellow brown, cold. R. F., grey or color- less. Microcosmic salt bead : both flames, reddish brown, hot ; reddish yellow, cold. MANGANESE. M» = 54.53. Occurrence.— Orw : pyrolusite, MnO^ ; braunite, A/WgOa; manganite or wad, {MnO)OH ; hausmanite, Mn^O^ ; man- ganese spar, MnCOj ; in almost all ores of iron, and in many silicates, coloring them. Metallurgy : the metal is obtained with difficulty from the oxides by reduction with charcoal. In comnfrce : black oxide of manganese, MnOt^ ; potassium permanganate, KMnO^ ; and in alloys, e.g., manganese bronze. Atomic Weight.— Stoech. fig.: (i.) 80.566igrm. silver per- manganate required 42.2506 grm. potassium bromide to convert it into silver bromide ; (ii.) IO.6730 grm. manganese oxide yielded 22.6875 grm. manganese sulphate. Specific heat. C(i,Ni, Mn, Zti.] MANGANKSE. 95 0.109. Isotnorphous relations : with magnesium, zinc, iron, nickel, and cobalt, in MeSOi.yH^O ; with these, together with copper, cadmium, zinc, and mercury in MeSO ^.K .:tSO ^.bH M ; with iron and chromium in the alums and in Me^O.^ ; with iron and aluminium, in MeO{OH) ; with iron, chromium, and sulphur in K^MeO^ ; and with chlorine in KMeO^. Chemical Relations. — Manganese forms six series of salt. Manganous, MnX„, resembling ferrous salts; manganic, MnXn, like cobaltic salts, easily decomposed by heat and on boiling with water (only a few, e.g., the sulphate and phosphate, have been obtained in the crystalline form) ; manganese tetrachloride, bromide, etc., MnX^, obtained by the action of the corresponding acids on manganese dioxide, and known only in solution (a double salt, MnF^ . 2KF, has been obtained in the solid form) ; the manganites, e.g., CaMn.O^^, prepared by the action of bleaching powder on manganous nitrate; the manganates, e.g., KMn^O^, cotVGSipondxng to the chromates and sulphates, from manganese dioxide by fusing with potash and nitre ; these latter are soluble in water con- taining potash, but decomposed by pure water with precipita- tion of manganic hydrate and formation of the permanganates, e.g., KMnO^, corresponding to the perchlorates. ' The Compounds of Manganese. Manganese is a greyish-white, brittle, difficultly fusible metal, sp.gr. 7.2, m.p. igoo'' C; oxidizes readily in the air at ordinary temperatures and decomposes water without heating. It dissolves readily in dilute acids. Manganous compounds. — MnX„ ; colorless or slightly pink in solution. Manganous carbonate, MnCO a', precipitated from solutions of manganous salts by alkali carbonates ; in air oxidizes to form blown manganic oxyhydrate MnO{OH). Manganous cyanide, Mn{CN)„ : precipitated from mangan- ous solutions by potassium cyanide ; soluble in excess of alkali cyanide to form the double cyanide Mn{CN)^ . zKCN; 96 MANGANESE. [Group III B. exposed to the air potassium manganicyanide is formed (c/. iron and cobalt). Mandamus hydrate, MniPH)^'. precipitated from manganous salts by fixed alkali hydrates; soluble in ammonium salts with formation of ammonio-manganous compounds, from which, exposed to the air, manganic oxyhydrate MnO.OH, brown, gradually separates. Manganous sulphide, MnS.H^O : precipitated from mangan- ous and manganic solutions by ammonium sulphide. With excess of ammonia and absence of ammonia salts a green pre- cipitate 3 A/«5 . 2H;. .9 is formed. ' m NAME. MANGANOUS. Acetate . . . Mn {.C<^H^O^)^ Bromide . .Mn 8r,.^H^0 Carbonate Mn CO3 Chloride Mn Cl^.^H^O Chromate Mn Cr 0^ Cyanide Mn {CN), Ferrocyan . . . Mn Fe { CN)^ Ferricyan .Mn^^Fe[^CN)^)^ Hydrate Mn (OH)^ Iodide Mn 1^.^/1^ O Nitrate.. Mn (NOa)^.6H,tO Oxalate Mn C^O^ Oxide MnO Phosphate Mn^^PO^)^ Sulphate ....Mn SO^ . TH^ O Sulphide Mn S Thiosulph Mn^S,0 MANGANIC. Hydrate Mn 4 0, ( OH) j Phosphate ...Mn PO^.H^O Sulphate Mn^(SO^^ Pot. Permang. ... K Mn O^ Color. pink pink pink pink brown brown white brown white pink white pink green white pink green white brown pink green violet Solubility In 100 parts cold hot water. 30(15°) deliq. o.oi sol. insol. insol. insol. insol. insol. deliq. deliq. 0.04 insol. s. sol. very sol. insol. deliq. sol. s. sol. dec. 6.S Remarks. very sol. deliq. s. sol. 650 insol. insol. insol. insol. insol. deliq. deliq. 0.1 insol. s. sol. very sol. insol. dec. insol. sol. dec. very sol Sol. amm. carb. S. ex. Insol. dil. hydrochl. acid Brown on standing. Decomposes in air. S. sol. ex. Olhttox. Mn^OAtMnO^,. Sol. dil. acids. [Mn^O.,. Also MnSO^.SH^O. Hydrated sulphide, flesb Forms sulphide. [color> Sol. in ammon. chlor. : hydrate, cyanide, chromate, and carbonate. Sol. in pot. cyan. : cyanide and carbonate. S. tartrates : hydrate, cyanide, chromate, and carbonate. Manganic compounds, MnX^ : purple-red in solution. Manganates, green, soluble in water. In alkaline solution Co, Ni, Afn, Zn.] ZINC. 07 manganates remain unchanged ; treated with dilute nitric or sulphuric acid they rapidly undergo transformation into per- manganates (violet). iAfn(OH)^ + 4AW0j + Na, CO, = 2/C^MnO^ + Na^MnO^ + ^NO + CO, + iH^ 2A'2 AfnOi + aHNO^ - 2KMnOi + MnOs+ ^KNO^ + 2H^0. Oxidation and Reduction. — Manganous compounds may be oxidized to manganic upon heating, or upon simple expo- sure to the air, e.g., hydrate, carbonate, cyanide, etc. Man- ganous compounds are oxidized to manganates and perman- ganates by fusion with carb. soda and potassium chlorate or nitrate ; by heating in alkaline solution with bromine, or in acid solution with red lead.^ Manganous compounds are oxidized to manganites by solution of bleaching powder. Manganic compounds are reduced to manganous by warming with acids. Permanganates and manganates are reduced to manganous compounds by ferrous sulphate, o/alic acid, sodium thiosulphate, alcohol, etc., in acid solution ; permanganates to manganates and finally to manganese dioxide by alkaline solutions of alcohol, stannous chloride, etc. Blowpipe Reactions. — All manganese compounds, fused with carb. soda on a platinum wire, yield a green-colored bead (turquoise enamel). Borax and microcosmic salt beads ; 0. F., violet, hot ; amethyst-red, cold ; R. F., slowly decolorized. ZINC. Zn = 64.77. Occurrence. — Ores : zinc blende, ZnS ; smithsonite, ZnCO.^; and siliceous calamine, Z«St04 .HjO. Metallurgy: the metal is obtained from zinc blende by roasting (to convert the sulphide into oxide), and subsequent reduction with char- coal. In commerce: metal; alloys, e.g., brass, bronze, etc.; white vitriol, ZnSO^ . 7H.^0 ; zinc chloride, ZnCl^ ; and zinc white, ZnO, .-•'-, * Mix a littlfe red lead with 5-10 cc. dil. nitric acid in a test-tube, heat almost to boiling and add a drop of dilute solution of manganous sulphate. Very delicate test for manganese if properly carried out. Essential to the success of the experiment is the absence of all reducing agents {e.g:, hydrochloric acid, excess of manganous salt). ii 98 ZINC. r [Group III B. • Atomic Weight.— Stoech. fig.: (i.) 29.6754 grm. zinc dissolved in dilute sulphuric acid evolved IO.0934 litres hydrogen at normal temperature and pressure ; (ii.) 16. 0316 grm. zinc yielded 2O.2608 grm. zinc oxide ; (iii.)electrolytically, one grm. of zinc is equivalent to 8.30 grm. silver. Specific heat, O.og^. Volatile compounds: metallic zinc, zinc chloride, and zinc organic compounds. homorpiious relations : with cadmium in MeS ; with beryllium in MeSiO^ ; with magnesium, calcium, iron, and manganese, in MeCO.^ ; with magnesium, iron, nickel, cobalt, and manganese, in MeSO^ .jH.O; with copper, magnesium, cadmium, mercury, iron, nickel, cobalt, manganese, and chromium, in MeSO ^ . K„SO ^ .dH .,0 ; and with iron and magnesium in the spinelles. Chemical Relations. — With acids, zinc forms salts of the type Z«A'o, many of which are 'somorphot s with the corre- sponding salts of magnesium, chromium, manganese, iron, cobalt, and nickel ; the chloride is partially decomposed on boiling with water, and the hydrate dissolves in solutions of potash or soda forming alkali zincates. Though strongly resembling magnesium, zinc is one of the heavy metals (sp. gr. 7.5), and correspondingly its hydrate and sulphide are insoluble in water, and its oxide is easily reduced by carbon. The Compounds of Zinc. Zinc is a bluish white metal, sp. gr. I.14, ni.p. 45o''C. ; not oxidized in air at ordinary temperatures ; heated in air forms zinc oxide ; if impure, it dissolves readily in dilute acids, very slowly in concentrated nitric acid. (The solution of pure zinc in dilute acids may be hastened by touching it with a strip of copper or platinum standing in the acid, or by adding copper sulphate to the solution. Solution of the pure metal in fixed alkalies is promoted by presence of iron filings.) Zinc compounds, MX./, colorless in solution, except the -acid be colored. Potassium zincate, K^ZnO^, white: from solutions of zinc salts by potash in sufficient excess to dissolve the precipitate*. Co, Ni, Mn, Ztt.] ZINC. 99 Zn{OH)^, first formed. Dilute solutions of zincates on boil- ing precipitate zinc oxide ; concentrated solutions are not decomposed by heat. NAME. Color. Solubility in loo parts cold hot Acffate Zii (C„H,iO„)., Bromide Zii A'/o Carbonate (basic) Cltloride Zn Cl„ Chroinate [basic) Cyanide Zn (CN)„ Ferrocyan. . . . Zn„Fe (CW)„ Ferricyan ... Zn.^( Fe( CJV) ,,).•, Hydrate Zn (0H)\\ Iodide ... Zn I^\ Nitrate ....Zn {NO.^ ) ., . 6//., O Oxalate Zn C^'O^ Oxide Zn O Phosphate Zn^iFO^)^ Sulphate Zn SO^^.TlF^O Sulphide Zn S Sulphocyan Zn {SCN)„ Thiosulph ....Zn S„O^.H,^d white white white white yellow white white yellow white white white white white white white white white white Remarks. S. sol. amin. carb. B.P. 400" C. S. ex. S. sol. dil. hydrochlacid. Sol. dil. hydrochlor. acid. Sol. alkali hydrates. Sol. dil. acids. " White vitriol." Insol. dil. acetic acid. I Forms sulphide. Sol. in ammonia : all insol. zinc salts except sulphide and ferrocyanide. Sol. in pot. cyan. : all insol. zinc salts except sulphide, ferrocyan. and hydrate. Sol. in tartrates : all insol. zinc salts except sulphide, ferrocyan. and phosphate. Sol. in ammon. chlor. : all insol. zinc salts except sulphide, ferocyan. .and phosphate Blowpipe Reactions. — B. B. on C with carb. soda, compounds of zinc are reduced to metallic state. The metal oxidizes to form an incrustation on the charcoal, yellow when hot. white when cold ; moistened with cobalt nitrate and re- ignited, the incrustation assumes a green color. SEPARATION IN GROUP III B. Treat the washed precipitate of the sulphides of nickel, cobalt, manganese, and zinc with cold dilute* (three per cent.) hydrochloric acid. Filter. Black residue (soluble in aqua regia) : NiS, CoS ; detected by blowpipe. Solution, MnCla, ZnCl^. — Add potash in excess aud digest without warming. Filter. Precipitate MniPH)^. — Oxidize to permanganic acid by red lead and nitric acid. Solution KoZnO^. — Ammonium sulphide gives white precipitate of zinc sulphide To separate nickel and cobalt several methods are employed, all depenamg on the fact that cobalt forms salts derived from the sesquioxide, while nickel does not. Of these the most con- venient is that with potassium cyanide, involving the formation of potassium cobalticyanide, as follows : — To the solution containing salts of nickel and cobalt (e.g., NiCl^ and CoCl^), add potassium cyanide in quantity slightly more than sufficient to dissolve the precipitate first formed : — CoCl^ + 2KCN = 2KCl + CoiCN)„. (Nickel : the same.) CoiCN)n+2KCN = Co{CN)^.2KCN. {K\cke\ : the same.) Adda few drops of dilute hydrochloric acid and boil a minute : HCl-\-KCN = HCN + KCl Co(CN).,.2KCN + HCl+2KCN = K.^Co{CN)^+KCl+H^ (Nickel: no reaction). *One part of the reagent solution " dilute hydrochloric acid" and two parts of SEPARATION. lOI To the solution add dilute hydrochloric acid drop by drop : the potassium cobalticyanide forms soluble cobalticyanic acid ; while the nickel is precipitated as cyanide, soluble in excess of hydro- chloric acid Ni{CN) ,.2KCN + 2HCI = Ni{CN) „ + 2HCN + 2KCI Ni(CN)^+2HCl = NiCl^+2HCN Or, to the solution of potassium cobalticyanide and nickel potassium cyanide, add a little potash and then bromine water : nickel is precipitated as the black nickelic hydrate Ni{OH)^, while the cobalticyanide is not decomposed. Nickel may be detected in the presence of cobalt by the following reactions, which, however, do not lead to a separation of the two elements. To the solution of the chlorides, add ammonium chloride, ammonia and potassium ferricyanide : a blood-red color (due to the formation of a cobaltamine compound) indicates cobalt ; if the mixture be boiled, the presence of nickel will be shown by the formation of a copper-red precipitate. ANALYSIS OF GROUP III IN PRESENCE OF PHOSPHATES, ETC. The composition of the precipitate formed by the third group reagents, is modified by the presence of certain sub- stances : tartaric acid, citric acid, and many other non-volatile organic substances (especially those whose structural formulae contain the hydroxyl group, for instance sugar), hinder, and, if present in sufficient quantity, completely prevent the precipi- tation of alumina and chromic and ferric hydrates ; while phosphoric and oxalic acid, as soon as the solution is made alkaline by ammonia, bring about a precipitation of calcium, strontium, barium, and magnesium as phosphates or oxalates respectively. For this reason it is necessary, first of all, to I02 SEPARATION. [Group III. make a preliminary investigation as to the presence or absence of these substances, according to the scheme given in the chapter on the detection of the acids. If they be present, they may be removed as follows before proceeding with the analysis : To remove organic substances: the filtrate from Group II is evaporated to dryness and then gently ignited in a porcelain crucible, and the residue dissolved in hydrochloric acid; or, if phosphorus be present, in nitric acid, filtered and analyzed according to the usual method. To remove phosphoric acid : the nitric acid solution is warmed in an evaporating dish with granulated tin, or better tin-foil, and excess of nitric acid, until a portion tested with ammonium molybdate shows the absence of phosphoric acid. [As the nitric acid evaporates it must be replaced ; if this be not attended to, aluminium, chromium, and iron may go into the precipitate.] The whole is then heated for a moment to boiling, the solution poured off from the precipitate, diluted with water, lead (from the tin-foil) precipitated by hydrogen sulphide, filtered, and the filtrate examined for Group III as usual. To analyze the precipitate produced by the third group reagents in a solution containing phosphoric acid : Wash separately the precipitates of Group III A and Group III B, mix them, and warm them with ammon um sulphide. Filter and wash. The precipitate may contain the sulphides of iron, manganese, cobalt, nickel, and zinc, the hydrates of aluminium and chromium, and the phosphates of aluminium, chromium, calcium, strontium, barium, and mag- nesium. Dissolve this precipitate in hot dilute hydrochloric acid with addition of a crj'stal of potassium chlorate, remove chlorine by boiling and any sulphur by filtration ; add sodium carbonate until a precipitate is produced, then just enough dilute hydrochloric acid to dissolve it again, and, lastly, excess of sodium acetate solution slightly acidified with acetic acid. Warm for some time and filter hot. Precipitate FePO^, AlPO^, CrPO,. The iron and chro. With phosphates.'] SEPARATION. 103 mium may be detected as described on page 89 ; the aluminium by boiling the precipitate with potassium hydrate for some minutes, filtering, acidifying the filtrate with hydrochloric acid and adding excess of ammonia, which causes a white precipitate of aluminium phosphate insoluble in acetic acid. Filtrate, CoCL_, NiCL_, MnCL_, ZnCl.,, CaCl., SrCl„, BaCl., . M^CL,, and either Fed. ^, CrCL^ and AlCl.^ orH^PO^, Add ferric chloride drop by drop so long as a precipi- tate results and until the liquid turns red, and digest for a time at a gentle heat. Remove the precipitate FePOi by filtration. To the filtrate, or solution not precipitated by ferric chloride, add amonium chloride etc., proceedmg as directed on page 89. THE FOURTH GROUP. Barium, strontium, anc* calcium, the metals of this group, resemble one another so closely that it is more satisfactory to treat of the chemical relations of all three together in this place than to consider each one under a separate head. In Mendelejeff's table they are the next-door neighbors of the alkali metals, and resemble the latter in their power of decom- posing cold water, in being readily oxidized in the air, and in the strongly basic properties of their hydrates ; they differ in forming insoluble carbonates, sulphates, and phosphates. Of the three hydrates, that of barium is the strongest base ; it may be melted at a red heat without losing water, and the carbonate of barium gives off but little carbon dioxide at its melting temperature — contrast the corresponding compounds of calcium. Strontium in this, as in other respects,(for instance, in the solubilities of its salts) stands midway between the other two. Each of these metals forms only one series of salts ; but as none of the latter are volatile enough to admit of a vapor density measurement, and as, until recently, no specific heat determination of any of the metals themselves had been made, the atomic weights of the elements could be deduced only from a study of their isomorphous relations. In order to show their analogy, from a crystallographic point of view, with the members of the magnesium group, the salts of barium, strontium, and calcium have been given the general formula MX 3. The atomic weights thus decided on have been confirmed by a subsequent determination of the specific heat of calcium. Barium, strontium, calcium, and magnesium are termed the "alkaline earths," standing intermediate between the alkalies on the one hand, and the " true earths," e.g., alumina. Ba, Sr, (7a. ] BARIUM. »05 on the other. Magnesium, however, by reason of the solubility of its carbonate iji ammonia salts, is separated from the others in the scheme of analysis adopted, and has a place in the fifth analytical group. BARIUM. Ba = 136.40. Occurrence.— Oy^s : heavy spar, BaSO^ ; and witherite, BaCOs. Metallurgy : the metal is obtained by electrolysis of the chloride, or by reducing the latter with sodium or potas- sium. In commerce : permanent white (paint), /iaSO^ ; barium nitrate, Ba{NO^)^ ; barium dioxide, BaO.^ (manufacture of hydrogen peroxide). Atomic Weight. — Stoech. fig. : I24.1929 grm. barium chloride convert I28.8934 grm. silver into silver chloride. Iso- morphous relations : with calcium, strontium, and lead, in MeCOa (arragonite series), and in M^SO^. The Compounds of Barium. NAME. Bromide BaBr^ . 2H^ O Carbonate BaCO^ Chloride BaCl^-zH^O Chromate BaCrO^ Cyanide Ba {CN)^ Ferrocyan . . . Ba^Fe ( CN)^ Ferricyan ■ . .,Bag(Fe{CJV)^)i Hydrate ....Ba (OH)^ .Sff^O Iodide Bal,^ . iH^O Nitrate Ba{NO\)\. Chloride . . . AfgC/.,.fi //y white 130 367 Chromale M'^CrO j white sol. very .sol. Cyanide M,ir(CN).i white .10]. sol. Dec. on eva(Joration. Ferroeyan . . .A/jf.^Fe (t'A'),, white 25 sol. Ferrieyan . M^r^ ( Fe ( CW ) „ ) , white sol. very sol. Hydrate Mg (01/)., white s. sol. s. sol. Sol'n shows alkaline reaction. Iodide ''/fj'i white deliq. deliq. Nitrate. . M/fi NO.,)., .6/1^0 while 2CO very sol. Oxalate M.K'<-\0^ white 0.07(15') 0.08 Insol. acetic acid. . Oxide J/^O while 0.00 1 insol. Phosphate M^i^/IPO, white s. sol. s. sol. Sulphate ^/f:SO^ TH^O Sulphide M^S white 25.8 (0°) 71.4 Epsom sails. l)rown s. sol. dec. Sii/phoeyan Mi;^SCN)., while deliq. very sol. Thiosulph . MgS^O.^.en^O white 125 (l8«) very sol. .Inim. phosphate white 0.005 insol. 1 Mg{NHi)rO^. 6H.0. • Sol. in tartrates : carhonate, hydrate, oxalate. Sol. in ammon. chl.: carb.onate, hydrate, oxalate, phosphate. Magnesium is a malleable silver- white metal, sp. gr. I.75, m.p. 75o°C. ; it oxidizes slowly in the air at ordinary temper- tures ; heated, it burns to form the oxide MgO. It decom- poses water slowly at loo^'C. Magnesium compounds, MgX^; colorless in solution, except when the acid is colored. Magnesium phosphates, white. " Ordinary " sodium phos- phate Na^HPOi precipitates from concentrated solutions of magnesium salts, MgHPO ^'yixom dilute solutions, Mg^{PO^)^ only on heating; from solutions containing ammonia and 112 POTASSIUM. [Group V. ammonium salts, MgNH^PO^ (from very dilute solutions only on long standing or on rubbing the walls of the test-tube with a glass rod). Blowpipe Reactions. — If the magnesium compound be first ignited, then moistened with cobalt nitrate, and ignited again, the mass assumes a pale rose-red color. POTASSIUM. /^ = 38.84. Occurrence. — Sources of potassium : sylvine, KCl ; carnallite, KCLMgCL.6H^0 ; wood ashes (carbonate). Metallurgy : reduction of the carbonate by ignition with carbon ; also by electrolysis. In commerce : caustic potash, KOH ; the bromide, iodide, chlorate, nitrate (saltpetre), cyanide, ferrocyanide, chromate, etc., etc. NAME. Acetate A'C^H^O^ Bromide . . . KBr Carbonate K^C0^.2H^0 Carbonate KHCO^ Chloride KCl Chromate A'j CrO^ Cyanide A'CA'" Ferrocyan . K\Fe(CN)^ . 3//^ O Ferricyan Kc^Fe[ CN)^ Hydrate '....KOH Iodide KI Nitrate KNO.;, Oxalate A'^C^O^ Oxide A'a <5 Phosphate A'aiOO^ Phosphate A'^HPO^ Sulphate K.iSO^ Sulphate KHSO^ Sulphi.: K.Ji Sulphocyan KSCN Thiosulph K.iS^O.,.]A,H^O Tartrate KHC^H^O Platinichloride Pot. Antimonyl. Tartrate.. .. Color. white white white white white yellow white yellow red white white white white white white white white white white white white white yellow white Solubility in 100 parts cold ' hot water. deliq 25 25 I. 32 50 very sol. 28 40 200 120 25 50 very sol. sol. very sol. I2.S very sol. very sol. very sol. very sol. 1.2(25=) 0.9(10°) 7- deliq. 100 100 sol. 57 60 dec. 100 80 very sol. 2CO 200 67 very sol. sol. very sol. 25 very sol. very sol. very sol. very sol. 6.6(100°) 5.2(100°) 53 Remarks. Bicarbonate. Normal tart'te very sol. K^PtCl^ Tartar emetic. Atomic Weight.— S^occ/f. fig.\ 2IO.85508 grm. silver required I45.70775 grm. potassium chloride to convert it into Mg,K,Na,Li,(NH,).-] SODIUM. "3 silver chloride ; 487.66o5 grm. potassium chlorate on reduction yielded 3I3.8175 grm. potassium chloride. Specific heat, O.166. Isomorphous relations : with the other alkali metals, silver and thallium, in the alums; with sodium and silver in MeCl ; with ammonia in many salts and double salts. SODIUM. Na = 22.87 Occurrence. — Sources of sodium: rock salt, NaCl; sea water. Metallurgy : see potassium. In commerce : metal ; caustic soda, NaOH ; washing soda., N a ^CO^ioHi,0 ; baking soda NaHCO.^; common salt, NaCl ; Chili saltpetre, NaiYO 3; Glauber's salts, NaoSO^.ioH„0 ; " hypo," (sod. thiosulphate), Na^SnOg^HM; hypophosphite, iVa/Z-iPOo -^2 0; borax, Na^B^O.; water-glass, Na2Si^0^); bichromate, Na^Cr^O.,. 2H„0; etc. I -A NAME. SODIUM. Acetate ..Na C^ff^0.^iII^O Borate . . ..ATa^B^^Oj. 121/^0 Bromide Na Br.^H^O Carbonate . . . Na^ Co^ . loH^ ( Carbonate Na HCO Chloride Na CI Chromate Na^ CrO^... \oH^ O BichromateNa^ Cr^ O^ . 2J/^ O Cyamde Na CN Ferrocyan Na^Fe ( CN ) Ferricyan Na.j^Fe^CN)^ Hydrate .Na OH Iodide Na I./^H^O Nitrate. NaNO^ Oxalate Na^C^O\ Oxide Na^U Phosphate . .normal (i2H^0) Phosphate sec. ( 1 2H2 Sulphate Na^ SO^ Sulphate NaH SO^ Sulphide Na^^^ Sulphocyan Na SCN Thiosulph. ..Na^S^O^.sH^ Platinichloride Pyroantimoniate Color. white white white white white white yellow red white yellow red white white white white white white white white white white white white red white Solubility in 100 parts cold hot water. 33 200 6 2CX) 75 112 21 10 420 dec. 35 sol. 40 sol. sol. sol. sol. sol. 20 sol. 20 80 60 210 80 300 2CX3 4 very sol. 20 4 sol. 3 96 very sol, sol. 42 very sol. sol. deliq. 102 deliq. very sol. very sol. insol. very sol. insol. Remarks. Borax. Na^PlCl^ Na^H^Sb^O.;.6H.i.O 114 LITHIUM. [Group V, Atomic Weight.— 86.03122 grm. sodium chloride con- verted 156.86206 grm. silver into silver chloride. Specific heat, 0. 293. Isomorphous relations : with the other alkali metals, silver and thallium, in the alums; with potassium and silver, in Me CI', with silver, in Me^SO^. ' LITHIUM = 7.o. Occurrence.— Ores : lepidolite, or lithia mica, a complex silicate ; triphylline, a phosphate containing iron, etc.; found also in many mineral waters. Metallurgy : by electrolysis of the chloride. In commerce : carbonate, Li^COg. Atomic Weight.— Stoech. fig.: I5.5533 grm. lithium car- bonate, with dil. sulphuric acid gave 9.2414 grm. carbon dioxide. Specific heat, O.941. Isomorphous relations: with the other alkali metals, silver and thallium, in the alums. NAME. Acetate Li C^H^O Bromide ... Li Br Carbonate Li^ CO3 Chloride LiCl.2H^0 Chromate Li^ Cr O^ Hydrate Li OH Nitrate ..Li NO^ Oxalate Z«, Cj 0\ Phosphate Li^PO^ Sulphate Li^SO^.H^O Color. Solubility 1 cold n 100 parts hot water. white white white 35° 140 0.1 300 white yellow white 65 very sol. s. sol. 125 very sol. s. sol. white white 48 .7 very sol. sol. white white 0.04 35 s. sol. 28 Remarks. AMMONIUM. Occurrence. — Sources of ammonia : at the present day ammonia is obtained exclusively as a by-product in the manufacture of coal gas. In commerce : " ammonia," solution of NH^ in water; sal-ammoniac, NH^Cl; microcosmic salt, //iV«(iV//JP04 . 4H2O ; carbonates ; sulphate. Mg, K, Na, Li, (NH^).-] AMMONIUM. "5 Nessler's test for ammonia: add potassium iodide (35 grm.) to a solution of mercuric chloride (13 grm.), until the precipi- tate at first formed is almost completely redissolved ; then add potash (100 grm.), let settle, and pour off the clear supernatant liquid (dilute to one litre). The solution so prepared is known as " Nessler's reagent," and with ammonia, or its salts, pro- duces a brown precipitate, NHgJ; most delicate test for ammonia. NAME. Ammonia NH^ Acetate NH^C^H^O^ Bromide NH^ Br Carbonate {NH^)^ CO^ Chloride NH^ CI Chromate {NH^)iCrO^ Cyanide NH^ CN Ferrocyan\NH^)^ Fe(CN)^ Ferricyan. {NH^)^ Fe(CN)^ Iodide NHJ Nitrate Nlf^ NO^ Oxalate {NU^^ C^O^ Phosphate . . . (NH^^^ H PO^ Sulphate (NH^)^ SO^ Sulphide {NH^)^S Sulphocyan NH^ SCN Thiosulph (A^^Ji ^aOg Platinichloride Mag. Molybdate Color. Solubility in loo parts cold hot water. gas 87.5(0') white deliq. white very sol. white sol. white 33 yellow very sol. white sol. yellow sol. red very sol. white deliq. white 200 white 5 white 20 white 77 colorless sol. white very sol. white 125 yellow 0.67 white 0.007 19(56") deliq. very sol. sol. 77 dec. very sol. sol. very scl. deliq. very sol. 41 sol. 98 sol. very sol. very sol. 1.25 Remarks. Solubilities for 760 mm. In air forms polysulphides [yel. {NH^)<, Pt CI, 6- {NH^).;, PO^. loMO, [3^2^ SEPARATION IN THE FIFTH GROUP. Detection of Magnesium {in the absence of lithitim). To the filtrate from Group IV add a little ammonia, enough ammo- nium chloride to redissolve any precipitate formed, and, lastly sodium phosphate. Precipitate : magnesium ammonium phosphate ; white. Separation of Magnesium from Lithium : Prepare a neutral solution of chlorides, free from ammonia, as follows : — evapor- ate to dryness the filtrate from Group IV, ignite gently, dissolve the residue in dilute hydrochloric acid, evaporatr* to dryness again, and dissolve in a little water. To the solution so obtained add ammonium oxalate, boil, and add acetic acid in excess. Precipitate : magnesium oxalate, white. Solution : lithium salts (precipitated by sodium phos- phate). Ammonia: salts of ammonium on boiling with potash evolve ammonia gas, recognized by its odor, by forming white fumes with hydrochloric acid, and by its action on litmus and on turmeric paper. Sodium, potassium, and lithium are most easily detected by the colorations they impart to the non-luminous flame. See p. J^5. DETERMINATION OF THE ACIDS. The method of grouping and analyzing the acids is iden- tical with that already employed in the determination of the bases. As in the former case, the acids are divided into classes by their behavior with certain group reagents. This division, however, is much less complete than is that of the bases, and at best serves but to group the acids into a few large classes, each member of which must be tested for in a separate portion of the solution. Before beginning with the wet tests involved in grouping the acids, it is advisable to subject a little of the solid sub- stance to a preliminary examination, by heating it on platinum foil, and by warming it in a test-tube with strong sulphuric acid ; whereby, in a majority of cases, valuable infor- mation as to the nature of the acids present may be obtained. A knowledge of the bases present is also of great service, as by consulting tables of the solubilities of the salts of the bases in question, it is often possible to greatly restrict the number of acids to be tested for. For example, if the salt submitted for analysis be soluble in water, all of those acids which form insoluble salts with the metals present are necessarily excluded. If any bases other than the alkalis be present, they must first be removed, because of difficulties which their presence might introduce into the separation. This may conveniently be effected by precipitating with sodium carbonate, and filtering ; after which the filtrate must be carefully neutralized before testing for the acids. The substance to be analyzed may be brought into solu- tion by the use of water alone, or of potash (for obvious i I ii8 THE ACIDS. reasons acids are rarely employed). In case it prove insoluble in these solvents, it should be fused with sodium carbonate (thus forming the soluble sodium salt of the acid), the fused mass digested with hot water, the insoluble carbonate filtered off, and the excess of sodium carbonate in the filtrate carefully neutralized with dilute hydrochloric acid. PRELIMINARY INVESTIGATION. Behavior on heating alone: A small portion of the substance is cautiously heated on platinum foil in the Bunsen flame : blackening indicates organic substances. Behavior on treating with conc. sulphuric acid : (a) Evolution of a colored gas : Reddish brown, bromates, bromides, nitrites ; pale yellow, chlorates, hypochlorites ; violet, iodides. (6) Evolution of a colorless gas possessed of a distinct odor : Odor of vinegar, acetates ; hydrochl. acid, chlorides ; peach- blossom odor, cyanides, ferricyanides, ferrocyanides, sulphocy- anides ; hydrofluoric acid, fluorides ; nitrogen pentoxide, nitrates » hydrogen sulphide, sulphides ; sulphur dioxide, sulphites, thiosul- phates ; odor of burnt sugar, tartrates. (c) Evolution of a colorless, odorless gas : Carbonates, citrates, oxalates. (d) No evolution of gas : " Arseniates, arsenites, benzoates, borates, chromates, phosphates, salicylates, silicates. DIVISION INTO GROUPS. I. The Calcium Group. — To the neutral solution, from which all bases but the alkalies have been removed, add calcium nitrate, if a precipitate appear, boil, cool thoroughly and filter. Precipitate : calcium salts of the acids of the first group ; those in brackets slightly soluble. To the precipitate add acetic acid : — Solution, with evolution of gas : carbonate, sulphite. Solution : arseniate, arsenite, borate, citrate, phosphate, (tartrate). Residue, soluble in dil. hydrochl. acid : fluoride, (iodate), oxalate. Residue, insoluble in dil. hydrochl. acid : silicate, (sulphate). II. The Silver Group.— To the filtrate from Group I add silver nitrate, filter. Precipitate: silver salts of the acids of the second group. To the precipitate add nitric acid, and boil : — Solution : acetate, hypophosphite, sulphide, thiosulphate. Residue, insol. in ammonia : bromide (slightly sol.), iodide. Residue, soluble in ammonia: benzoate, bromate (bromide), chloride, chromate, cyanide, ferricyanide, ferrocyan- ide, hypochlorite, iodate, salicylate, sulphocyanate. i'< III. Acids whose Calcium and Silver Salts are SOLUBLE. — Chloric, nitric, nitrous, perchloric. REACTIONS OF THE ACIDS. Acetic Acid, C^H^O^. — All acetates are soluble in water (silver and mercurous acetates sparingly soluble). Sulphuric acid, hot, cone: characteristic odor of vinegar ; upon addition of a little alcohol, odor of ethyl acetate. Ferric chloride : deep red color from formation of ferric acetate ; color not destroyed by mercuric chloride (c/. ferric sulphocyanate). On boiling with excess of acetate, basic ferric acetate (reddish brown) is precipitated. Arsenic Acid, H^AsO^. — All arseniates, except those of the alkali metals, insol. in water. For reactions, see page 6i. Arsenious Acid, H^AsO^. — All arsenites, except those of the alkali metals, insol. in water (calcium and barium arsenites sparingly soluble). For reactions, see page 60. Benzoic Acid, C^H^O.^. — Colorless, glistening thin lamellae or needles. Most of the benzoates soluble in water and in alcohol. Sulphuric acid, cone: brown colored solution, from which benzoic acid is precipitated by water. Ferric chloride : in neutral solutions, flesh-colored precipi- tate — basic iron benzoate. BoRACic Acid (metaboracic acid H^BO^). — All borates, except those of the alkali metals insol. in water. Heated to i6o°C., metaboracic acid is converted into pyroboracic acid Flame test : a fragment of a borate, moistened with a drop of cone, sulph. acid and a little alcohol, and ignited, green flame ; this may also be obtained by moistening first with cone, sulph. acid, then with glycerine, and igniting upon a platinum wire. THE ACIDS. 121 Bromic Acid, HBrO^. — All bromates soluble in water (silver, lead, and mercurous bromates sparingly soluble) ; on heating they liberate oxygen with formation of bromides. Sulphuric acid, hot cone: liberation of bro'-.ine. Silver nitrate : white precipitate AgBrO^, decomposed by hydrochl. acid with evolution of bromine (distinction from bromides). Potassium iodide in acid solution : liberation of iodine (blue color with starch). Carbonic Acid [//gCOg?].— Carbonates and bicarbonates of the alkali metals soluble in water, all others insoluble ; those of the alkali earth metals soluble in water saturated with carbon dioxide. AU acids decompose carbonates with effer- vescence ; the carbon dioxide liberated may be recognized by the turbidity it produces in a drop of lime water on a glass rod. Chloric Acid HCIO^. — All chlorates soluble in water. Heated they liberate oxygen with formation of chlorides (dis- tinction from nitrates). Sulphuric acid, cone : Evolution of green-yellow chlorous oxide, C/0 2. [Caution: explosive — only small quantities to be experimented with.] Aniline sulphate, in acid solution, is immediately colored an intense blue by a fragment of a chlorate. Chromic Acid, H^CrO^. — Chromates of the alkali metals, of magnesium, calcium, and zinc, soluble in water ; strontium and mercuric chromates sparingly soluble ; the others insoluble. For reactions, see page 89. Citric Acid, Ceffg^f— Citrates of the alkali metals soluble in water ; those of iron, copper, and tin sparingly soluble ; the others insoluble, but for the most part soluble in alkali citrates, with formation of double salts. Sulphuric acid, cone. : Decomposition with liberation of carbon dioxide and carbon monoxide (inflammable). No charring, except on prolonged heating (distinction from tartrates). . 1^ laa THE ACIDS. Calcium hydrate : white precipitate of calcium citrate formed upon prolonged heating (distinction from tartrates). Ferricyanic Acid, H3F<;(CN)8-— Ferricyanides of the alkalies, of the alkaline earths, and of magnesium, soluble in water (barium ferricyanide sparingly soluble) ; other ferricyan- ides insol. Sulphuric acid, hot, cone. : evolution of hydrocyanic acid iq.v.). The solution should be tested for iron. Ferrous sulphate : dark blue precipitate (Turnbull's blue). Ferric chloride ' no precipitate, brown coloration. Ferrocyanic Acid, //4F^(CW)8.— Ferrocyanides of the alkali metals, and of magnesium, calcium, and strontium; soluble in water ; all others insol. Sulphuric acid, hot, cone. : evolution of hydrocyanic acid iq.v.), forming a solution containing iron. Ferrous sidphate : light blue precipitate, becoming darker on standing. Ferric chloride : deep blue precipitate (Prussian blue). Cupric sulphate : red-brown precipitate. Hydriodic Acid, HI. — Silver, lead, mercurous and mercuric iodides, insol. in water ; all other iodides soluble, (except bismuth iodide, decomposed); the insol. iodides dissolve in excess of potassium iodide with formation of double salts. Sulphuric acid, hot, cone. : liberation of iodine (violet vapors). Oxidizing agents, e.g., potassium dichromate in acid solu- tion, liberation of iodine (blue color with starch, violet with carbon disulphide or coal oil). - Silver nitrate : light yellow precipitate, insol. in ammonia (distinction from chlorides and bromides). Cupric sulphate : brownish white precipitate, (Cu I). Solu- tion to be tested for free iodine. Hydrobromic Acid, HBr. — Bromides of silver, lead, and mercury (ous), insol. in water ; mercuric bromide slightly soluble ; other bromides easily soluble. THE ACIDS. "3 Sulphuric acid, hot, cone. : liberation of bromine (brown fumes). Oxidizing agents, e.g. tTpoX2iS&\wm dichromate in acid solution, liberation of bromine (yellow color with starch, or with carbon disulphide or coal oil). Silver nitrate, yellow white precipitate, sparingly soluble in ammonia, (distinction from chlorides). Hydrochloric Acid HCl, chlorides of silver and mercury ifitts), insoluble in water ; lead chloride sparingly soluble, the others easily soluble. Sulphuric acid, hot, cone. : liberation of hydrochloric acid gas (colorless acrid fumes ; white cloud with ammonia.) Potassium dichromate and concentrated sulphuric acid : form- ation of chromyl chloride ; see page 88. Hydrocyanic Acid, HCN, cyanides of alkalies, alkaline earths, magnesium, and mercury soluble in water ; the others insoluble, often soluble, however, in excess of alkali cyanide with formation of double salts. Sulphuric acid, hot, cone. : evolution of hydrocyanic acid gas (poisonous, colorless gas; with pungent, so-called " peach blossom " odor). Prussian blue test : to a solution of a cyanide add a few drops each of potash, of ferrous sulpliate, and of ferric chloride ; boil, and acidulate with dilute hydrochloric acid. Prussian blue is formed. Sulphocyanate test : to a solution of a cyanide add a few drops of ammonium sulphide, heat gently in an evaporating dish until the solution becomes colorless, acidulate with dilute hydrochloric acid, and add ferric chloride— blood-red coloration. Hydrofluoric Acid, HF. — Fluorides of the alkali metals, of silver, tin and mercury {ic), soluble in water; of coppen bismuth, cadmium, iron, and zinc, sparingly soluble ; those of the alkaline earth metals and of magnesium, insoluble. Sulphuric acid, hot, cone. : evolution of hydrofluoric acid, recognized by its pungent odor and by its etching action upon glass. Caution ! 'I ^li "4 THE ACIDS. Hypochlorous Acid, f/C/0.— All hypochlorites (except that of silver) are soluble in water. They are decomposed by dilute acids with liberation of free chlorine, which may be recognized by its odor, color, and bleaching action upon moist litmus paper. Silver nitrate : white precipitate, silver hypochlorite ; quickly decomposing into silver chloride and silver chlorate. Aniline sulphate, in dilute sulphuric acid solution, upon addition of a fragment of a hypochlorite, blue coloration. HVPOPHOSPHOROUS AciD, HgPOo.— All hypophosphites are soluble in water. Heated with nitric acid they form ortho- phosphates. Ctipric sulphate : heated to 7oC°, black precipitate of copper hydride ; decomposed upon heating, or upon treatment with dil. hydrochl. acid. Potassium permanganate, chromic acid, and a solution of potassium iodide containing iodine, are all reduced (decolorized) by hypophosphites. iV/v^y ;M^m^^ : black precipitate of reduced silver. Iodic Acid, H/0 ,,.—Iodates of the alkali metals soluble in water; other iodates insol., or very sparingly soluble. Sulphuric acid, hot, cone. : no evolution of iodine. • Ferrous sulphate and cone, sulphuric acid : liberation of iodine (see hydriodic acid). Nitric Acid, HNO^.—PAX nitrates are soluble in water. Brown ring test : one centimetre each of the solution of a nitrate and of ferrous sulphate solution are mixed in a test-tube and cone, sulphuric acid added, allowing it to fall along the wall of the tube : at the boundary between the two liquids a brown ring or coloration will be observed, arising from the solution of nitric oxide in ferrous sulphate. Phenyl sulphate (i part phenol, 4 parts cone, sulph. acid, and 2 parts water) : brown-red coloration, becoming yellow or green upon addition of ammonia. Nitrous Acid, HN0^.—\\\ nitrites are soluble in water —silver nitrite sparingly soluble. THE ACIUS. l«f Brown ring test : obtained by the action of ferrous sulphate on a nitrite, without addition of sulphuric acid ^distinction from nitrates, q.v.). Potassium iodide uitii dil. sulph, acid: liberation of free iodine (see hydriodic acid). Oxalic Acid, CM^O.^. — Oxalates of the alkali metals and of magnesium and iron (ferric), soluble in water ; chromic and stannic oxalates, sparingly soluble ; the others, insoluble. Sulphuric acid, hot, cone. : evolution of carbon dioxide and carbon monoxide (inflammable). Orthophosphoric Acid, HaPO^. — Di and tri-metallic phosphates, with the exception of those of the alkali metals, insoluble in water ; mono-metallic phosphates, usually some- what soluble. Ammonium molybdate: mix in a test-tube one cubic centimetre of ammonium molybdate solution, with one drop of the solution of the phosphate, and add dilute nitric acid until the precipitate at first formed (MoO^) is redissolved. The liquid assumes a yellow color, and on warming (not boiling !) a canary-yellow precipitate of ammonium phospho- molybdate (N H ^) ^PO ^.izMoO ^ is formed. "Magnesia mixture " (prepared by the addition of ammonia to a solution of magnesium sulphate containing ammonium chloride in sufficient quantity to prevent the formation of a precipitate) : white crystalline precipitate of magnesium ammonium phosphate. Silver nitrate, yellow precipitate, Ag^^PO^', (distinction from pyrophosphoric acid). Pyrophosphoric Acid, H^P.fi^ .—Pyrophosphates of the alkali metals soluble in water ; all others insol. Upon heat- ing in acid solution, pyrophosphates are converted into ortho- phosphates. Ammonium molybdate : no precipitate (c/. orthophosphoric). *' Magnesia mixture'' : no precipitate (c/. orthophosphoric). s" Silver nitrate : white precipitate. - 126 THE ACIDS. Salicylic Acid, C^H^O^ ; colorless four-sided prisms, or long needle-like crystals, odorless, but possessed of a sour- sweet taste. Salicylic acid and its salts are sparingly soluble in cold water, easily soluble in hot water and in alcohol. Ferric chloride : intense violet coloration. Alcohol and concentrated sulphuric acid on warming: odor of wintergreen. Heated with lime : odor of phenol, (carbolic acid). Silicic kcin [HSiOg,!']. — Silicates of the alkali metals are soluble in water, all others insol. Hydrochloric acid : decomposition with separation of jelly- like mass, insol. when once dried. Silica skeleton : If a small fragment of i. silicate be dissolved in the microcosmic salt bead, the silicic acid is liberated and floats undissolved in the bead, which forms an opalescent glass upon cooling. Sulphuric Acid, Yi^SO^. — Sulphates of barium, lead, and strontium, insol. in water , of silver, calcium, and mercury {ous), sparingly soluble; the others, soluble. r Barium chloride : white precipitate {BaSOJ insoluble in all acids. . ,, /,. --/jc,;-.^,. ..• ',-;.^ ■:;., ,i.:,.,,. ■ •■.,:• All Acids containing Sulphur. — B. B. on C. with carb. soda form sodium sulphide. The fused mass, moistened l^ 7d placed on a silver coin, produces a black stani. " Silver coin test." Sulphuretted Hydrogen, HgS.— Sulphides of the alkalies, of the alkaline earths, and of magnesium, soluble in water ; aluminium and ferric sulphides decomposed ; all other sulphides insol. Sulphuric acid, hot, cone: evolution of sulphuretted hydrogen ; (odor ; lead acetate paper blackened). Potassium nitroprusside* and potash solution : violet red coloration, which disappears on standing. "Prepared by heating potassium ferrocyanide with nitric ac'!d, and neutralizing with potash. THE ACIDS. i«7 Sulphurous Acid, [f/ 2 S.^^. 3?]— Alkali sulphites, soluble; all others insoluble in water. All acids : evolution of sulphur dioxide (odor). Barium chloride : white precipitate, soluble in dilute hydrochloric acid. If to this solution a little dilute nitric acid be added, insoluble barium sulphate is formed upon heating. Silver coin test : see page 126. SuLPHOCYANic AciD, {HCNS) — Sulphocyanates of the alkalies, alkaline earths, magnesium, iron, manganese, zinc, cobalt, copper, and mercury, (tc) soluble in water, all others insoluble. Ferric chloride: blood red coloration, decolorized by mer- curic chloride (distinction from acetic acid), by arsenious acid, by oxalic acid, and by sodium acetate (color restored by hydro- chloric acid). Tartaric Acid, C^H^Oq. — Tartrates of alkali metals, manganous, ferric, cobalt, stannous, and antimonious tartrates, soluble in water ; acid tartrates of potassium, ammonium, and calcium, sparingly soluble ; all other tartrates insoluble in water, but often dissolved in excess of alkali tartrates with formation of double tartrates. Sulphuric acid, hot, cone. : charring, and odor of burnt sugar. Silver mirror : Dissolve the tartrate in ammonia, add silver nitrate, and allow to stand for a few minutes in a warm place. The inside of the test tube becomes coated with metallic silver. Thiosulphuric Acid, H^S^O^. — Thiosulphates of silver, lead, and barium, sparingly soluble in water ; all others easily soluble. The insoluble thiosulphates are, in most cases, easily dissolved by excess of alkali thiosulphates. All acids : decomposition with evolution of sulphur dioxide, and precipitation of sulphur (distinction from sulphites). Silver nitrate : white precipitate, silver thiosulphate, grad- ually becoming black from formation of silver sulphide. Potassium permanganate \ iodine solution: both decolorized. Silver coin test : see page 126. THE ALKALOIDS. The term "alkaloids" is restricted in chemistry, to a somewhat extensive class of organic bases, which are usually found, uncombined, in nature, in certain plants. They contain carbon, hydrogen, oxygen, and nitrogen in varying propor- tions, and are related chemically to the pyridine and quinoline bases. Often grouped with the alkaloids, on account of their physiological action, are substances which, chemically, are widely different, e.g., glucosides, albuminoids, etc. ; a few of the more common of these are included in a supplementary list at the end of the alkaloids proper. The alkaloids are colorless compounds possessed of bitter taste and characteristic physiological action. They are usually insoluble in water, readily soluble in acids, with which they form well defined salts. From their solutions most alkaloids may be precipitated by potassium mercuric iodide, by phos- phomolybdic acid, and in some cases by dilute potash. The first two precipitates vary in color from white to yellow and are insoluble in dilute acids. No satisfactory and complete method for grouping the alkaloids has been devised. The following scheme, proposed by Fresenius, will be found of assistance in roughly subdividing them, after which each member of the sub-groups must be tested for separately. DIVISION INTO GROUPS. ! VoLATiLK : coniine, nicotine, (recognized by their distinctive odors). Non-Volatile : brucine, morphine, strychnine, etc. (a) To a portion of the solution of the non-volatile alkaloid, add potash to slight alkaline reaction ; allow to stand five minutes, and, if a precipitate appear, add potash to ' strongly alkaline reaction. The precipitate dissolves: (atropme*), cocaine, morphme. The pyecipiiate is insoluble : aconitine, brucine, cinchonine, narcotine, quinine, strychnine, veratrine. (b) To another portion of the solution add a few drops of dilute sulphuric acid, neutralize with concentrated solution of bicarbonate of soda, and allow to stand fifteen minutes. Precipitate : cinchonine, narcotine, quinine. No precipitate : brucine, strychnine, veratrine., •Precipitated by polash from concentrated solutions only. REACTIONS OF THE ALKALOIDS.* AcoNiTiNE, C3^//^3.Y0i„,- Colorless powder, difficultly soluble in water, easily soluble in acids and alcohol. Sulp/iicric acid, cone. : solution with yellowish brown color- ation, which becomes bright yellow upon addition of dilute nitric acid. Atropine, C^ .//ogiVOa.— Lance-shaped colorless crystals ; difficultly soluble in cold water, easily soluble in alcohol and in acids. Sulphuric acid, hot, cone. : odor of orange blossoms. Brucine, C„3f/„,A^„0^.4/f„0.— Colorless, prismatic cry- stals ; difficultly soluble in water, easily soluble in alcohol and in dilute acids. Sulphuric acid, cone, containing nitric acid :t intense red coloration. Nitric acid, cone, {sp.gr. I.4) : blood-red coloration, quickly <:hanging to yellowish red. Chlorine or bromine water : pink coloration, yellowish brown on addition of ammonia. CiNCHONiNE,Ci3//2.,iY,,Oi.— Transparent,lustrousprisms, or white powder; insoluble in water, soluble in alcohol and in acids. Potassium ferrocyanide flocky-white precipitate, soluble in excess (distinction from quinine). An^'i''?]"''^* "^ sometimes arranged as follows : (a) Volatile : nicotine, coniine ; (*) Alkaloids of opium : cocaine, codeine, morphine, narcotine, papaverine; (<) Strych- nine and related compounds : brucine, strychnine ; (d) Quinine and related com- pounds : aconitine, atropine, cinchonine, colchicine, quinine, veralrine. tThis solution (recommended by Erdmann) is prepared as follows : to 100 c.c. water add 5 dropsnitric acid, sp. gr. 1.25 ; to 20 c.c. cone. -.ulph. acid add 10 drops .of this dilute solution of nitric acid. THE ALKALOIDS. Chlorine or bromine ivater, then ammonia: white precipi tate. Cocaine, C\, Ho ^iVO^.— Colorless prisms; difficultly Sf^luble in water, soluble in alcohol and in acids. Ammonia: white precipitate, soluble in excess. CoDEiNK,C, „H., ixVO^.H.O.— Colorless, rhombic crystals; difficultly soluble in water, easily soluble in alcohol and in acids. Sulphuric acid cone, containing nitric acid : blue colora- tion. Nitric acid ccic. : yellow solution. Colchicine, C.„//„„.VO„ —Yellowish-white resinous ma- terial, soluble in water and in alcohol. Nitric acid cone. : violet coloration ; upon addition of water, yellow. Sulphuric acid cone. : cold, yellow coloration ; upon heating, red. If the sulphuric acid contain a trace of nitric acid, the color of the solution changes gradually from green to yellow. ... Conine, Cg/Zi-iV.— Transparent, oily liquid, possessed of a nauseating odor and taste ; difficultly soluble in water, easily soluble in alcohol. Potassium mercuric iodide : resinous precipitate, becoming crystalline on standing. Chlorine water : turbidity (distinction from nicotine). Morphine, C^,H^,,NO^.H.,0.—Co]or\ess, clinorhombic prisms; difficultly soluble in water, easily soluble in alcohol and in acids. Nitric acid, cone. : orange yellow coloration. Ferric chloride : deep blue coloration. NARCEfNE, Ca 3^2 uArO„.2H,0.-Long glistening needles, difficultly soluble in cold, easily in hot water or alcohol. 132 THE ALKALOIDS. Sulphuric acid couc. solution with brown color ; with excess of acid, bright yellow. Iodine, dissolved in potassium iodide solution : the solid alkaloid colored blue (distinction from all other opium alkaloids). Nakcotine, CooH.y.YO-. — Colorless, lustrous prisms; insoluble in water and in alcohol, slightly soluble in acids. Sulphuric acid cone, containing nitric acid : orange yellow coloration, changing to red and finally to violet. Bromine water : yellowish green coloration; upon addition of ammonia, yellowish red. NicoTiNK, C,„//,4A'._,. — Colorless oil ; soluble in water, in alcohol and in acids. yiitric acid cone: red coloration. Potassium mercuric iodide • resinous precipitate, becoming crystalline on standing. Pafaverine, C.,f,H .-^XO^. — Colorless, needle-shaped crys- tals, insoluble in water and in cold alcohol, soluble in warm alcohol and in acids. Sulphuric acid cone. : cold, no coloration, heated, deep violet blue. PiPEKiNE, C, ;//,„xV0.3. — Glassy four-sided prisms, insol. in water, fairly soluble in alcohol. • Sulphuric acid cone. : solution with yellow color, becoming dark brown and finally greenish-brown on standing. QuiNiDiNE, C.^^,H., Ar„0„.— Glistening four-sided prisms, very difficultly soluble in water, easily in alcohol ; reactions as under quinine, with one exception, viz : — Potassium iodide : in neutral solutions, white powdery pre- cipitate. Quinine, CaoZ/g^A^O;..— White amorphous powder, diffi- cultly soluble in water, easily soluble in acids. The solutions of quinine and its salts are phosphorescent. THE ALKALOIDS. 133 Bromine water and then ammonia : green, flocky, precipi- tate, soluble in excess of ammonia to fcrm emerald green solu- tion ; neutralized with dilute hydrochloric acid, the solution becomes blue, and with excess of acid, violet or red. ' Potassium ferrocyanidc : dark red coloration. Strychnine, C...{H .....X J)... — White glistening rhombic prisms, very difficultly soluble in water and in alcohol, easily- soluble in dilute acids. Sulphuric acid cone, dissolves strychnine to form a clear solution. If to a small portion of this solution a fragment of potassium dichromate be added, a violet coloration, sur- rounding the dichromate, may be observed. This coloration is produced also (but less distinctly) by potassium permangan- ate, and by red lead. Potassium fcrricyanide : yellowish green coloration. Thebaine,Ci,,//._,,.,,.VO,.,.— White quadratic scales,insol.in water, easily in alcohol and in acids. Sulpliuric acid cone: solution with blood red color, gradually becoming yellowish red. ferric chloride : no color reaction (distinction from morphine). Nitric acid, sp.gr. 1.4 : solution with yellow color. Veratkine, C.^„H^,,N0„. — Colorless prisms; insol. in water, soluble in alcohol and in acids. Sulphuric acid cone: yellow coloration, becoming bright red on standing. Upon addition of bromine water this solu- tion be'comes purple. Hydrochloric acid cone: colorless solution cold, colored deep red upon boiling. Caffeine or Theine, CyHi„N^O„H„0.— White, glisten- ing, silky needles ; insol. in alcohol, soluble in water and in acids. 134 THE ALKALOIDS. Chlorine or bromine water evaporated to dryness with a caffeine solution, and the reddish-brown mass moistened with ammonia : violet coloration. DiGiTALiNE.— White, amorphous powder ; insol. in water, easily soluble in alcohol. Sulphuric acid cone. : brown solution ; on standing, dark red ; with bromine water this solution gives a beautiful violet- red coloration. PiCROTOXiNE, CisffiaOoffoO.— White, glistening, four- sided prisms, or needles ; easily soluble in water, and in alcohol ; insol. in dilute acids. Sulphuric acid cone: golden yellow coloration; upon addition of potassium bichromate, violet, and then green. Salicine, C^^H-^^O.. — White, glistening needles or lam- ellae ; easily soluble in alcohol and in water. Sulphuric acid, cone. : colciS salicine blood red without dissolving it. Nitric acid cone. : colorless solution, cold ; yellow upon heating. TABLE FOR THE DETECTION OF A SINGLE ACID. I. To A PORTION OF THE NeUTKAL SOLUTION ADD BaRIUM Chloride. White precipitate, insol. in hydrochloric acid : silicate, sulphate. White precipitate, soluble in hydrochloric acid : borate, carbonate (effervescence), phosphate, sulphite (odor of sul- phur dioxide), tartrate, thiosulphate, {separation of sul- phur). Yellow precipitate, insol. in acetic acid : chromate. White precipitate, insol. in acetic acid: oxalate. II. To another portion of the solution add Silver Nitrate. Precipitate, soluble in dil. nitric acid and in ammonia : white, borate, carbonate, oxalate, silicate ; yellow, arsen- ite, phosphate ; red, arseniate ; dark red, chromate. Precipitate, insol. in dilute nitric acid, but soluble in ammonia : white, chloride, cyanide, ferrocyanide {diffi- ctiltly soluble), hypochlorite, sidphocyanate ; y^Wowi, brom- ide {diffictdtly soluble) ; orange red, ferrocyanide. Precipitate, insol. in dilute nitric acid and in ammonia : white, (ferrocyanide) ; yellow, {bromide), iodide ; black, sulphide, (soluble in concentrated nitric acid). III. Acids not precipitated by Barium Chloride or Silver Nitrate. Brown ring test (page 124) : nitrate, nitrite. Reaction with sulphuric acid (page 121) : chlorates. *For the method of bringing the substance to be analyzed into solution, end for preliminary examination, see page Ii8. < H PJ I— I 1 N t w MM in t>. X CO 'H t^ r-i J.-i o -.c o o ■r. ci • tH rH I— 1 1-* OS <; b« UO ^^-T^ l-H ^ 1— t o > ^ O ^ C>00 I.-: x d lo O O O O 1-t l-( c: CO GO Oi rH rH ■ cr; vO O 1- c^ M in ro CTn ►— ( o ^ cc J.-^ 1^ ci >o HH ;i r-t cc o t^ n^ > 1 C< ^ r-r O j CO 01 00 00 00 in po -^ o o • n 51 o 1—1 T-i ad ui -t co' t^ 1— t o tH « o t- C^ (M X CO > ^ :s; T-l O c tc rH CO »0 t~ C5 rH tH X o rH Cq Q^ ^ ^ a. ^ -^ < C/O ^ CQ H ON N GO c^ q C^ ■^ «>. • 7\ T^ 00 t^ r-i d t^ W5 ^ > O r-^ (M '^ t- C. 1-1 o CO I-. c< ^ •^ 1-1 CM ^ ^ ^ Ct^ h- q 00 j>. 00 q N f 00 . C5 CO Oi -t t^ rH CO 00 hH O (M CO O OC tH CO OS ^.H 0:; tH T^ T-\ ^ « S 8^ -TS O N) CO U 03 ^ t^ -^ q 00 00 00 q O iH t-^ n^ 00 TjH ci HH n: ::^ in CO 00 CO ■ INDEX Acetic acid, 120. Acids : determination, 117. Acids: reactions, 120-127. Aconitine, 130. Alkaloids, 128. Allemontite, 61. Alum, 81. Aluminium, 8t. Ammonium, 114. Amm. sulph., yel., 57. Anglesite, 34. Antimonyt 61. Apatite, 107. Arragonite, 107. Arsenic, 59. Arsenic acid, 61, ° 120. Arsenious acid, 60, 120. Asbestos, no. Atomic weights (talile), 136. Atropine, 130. Babbitt metal, 62. Barium, 105. Bases grouping, 32. Basic salts, 37. Benzoic acid, 120. Bismuth, 70. Bleaching powder, 107. Blende, 97. Blowpipe, Use of, 27. Blue vitriol, 72. Boracic acid, 120. Borax, 113. Borax beads (colors), 26. Brass, 72, 97. Br.nunile, 94. Britannia metal, 64. Bromic acid, 121. Bronze, 64, 97. Brown ring test, 124. . Brucine, 130 Bunsen's Method of Reduction, 27. Cadmium, 75. Caffeine, 1 33. Calamine, 97. Calcite, 107. Calcium, 107. Calomel, 40. Carbonic acid, 121. Carnallite, 112. Cassiterite, 64. Cerussite, 34. Chili saltpetre, 113. Chloric acid, 121. Chromic acid, 121, 87. Chromium, 86. Cinchonine, 130. Citric acid, 121. Clay, 81. > Cobalt, 90. Cocaine, 131. Coelestine, io6. Colchicine, 131. Confirmatory tests, 45. Codeine, 131. Coniine, 131, Copper, 72. Corrosive sublimate, 43. Corundum, 81. Crocoisite, 86. Cryolite, 81. Cuprous and cupric, 73. Delicacy of reactions, 47. Digitaline, 134. Dolomite, no. Electrochemical series, 55. Epsom salts, no. Feldspar, 81. Ferric) anic acid, 122. Ferrocyanic acid, 122. Ferrous and ferric, 83. Filtration, 33. Flame tests, 25. Fluorspar, 107. F'ly powder, 59. (lalena, 34. Geiman silver, 72, 92. Glauber's salts, 1 13. (ireenockite, 75. (Jreen vitriol, 83. Gypsum, 107. Haematite, 83. liausmannite, 94. Heavy spar, 105. Hydrogen sulphide, 126. Hydriodic acid, 122. Hydrobromic acid, 122. Hydrochloric acid, 123. Hydrocyanic acid, 123. Hydrofluoric acid, 123. "Hypo," 113. •38 INUKX. Hypuchloric acid, 124. lIyp()|)hosphorouft ncid, 124. Iodic acid, 124. Iron, 83. Lead, 34. Lepidolite, 114. Limestone, 107. Litharge, 34. Lithium, 114. Lunar caustic, 38. Magisterium l)isniuthii, 70. Magnesia mixture, 125. Magnesitc, no. Magnesium, iio. Malachite, 72. Manganese, 94. Marble, 107. Marsh's test, 68. Mendelejeffs table 136. Mercury, 40. Metastannic acid, 66. Mica, 81. Microcosmic salt, 114. Microcosniic salt beads, 26. Millerite, 92. Mispickel, 59. , Morphine 131. Naples yellow, 62. Narceine 131. Narcotine 132. Nessler's test, 1 1 S- Nickel, 92. Nicotine 132. Nitric acid, 124. Nitrous acid, 124. Orpiment, 59. Orthophosphoric acid, 125. Oxalic acid, 125. Oxidizing agents, 44. Papaverine 132. Paris green, 72. I'etmanent white, 105. Pewter, 62. Phosphates in Group III., lOi. Phosphor beads (cwlors), 26. Phosphoric acid, 125. Picrotoxine, 134. Pink salt, 64. Piperine 132. Plaster of Paris, I07. Potassium, 112. Precipitation, 33. Preparing salts, 64. Prussian blue, 83. Prussian blue test, 123. Pyrargyrite, 38. Pyrites, 83. Pyrolusite, 94. Pyrophosphoric acid, 125. Quinidine 132. <^)uinine 132. Realgar, 59. Reducing agents, 44. Rock salt, 113. Rose's nietal, 70. Sal ammoniac, 1 14. Salicine 134. Salicylic acid, 126. Sanphire, 81. Scheele's green, 59. Silicic acid, 126. Silver, 38, Silver coin test, 126. Smalt, 90. Smithsonito, 97. Soda, 113. Sodium, 113. Solder, 34, 64. Solution, 28. Spinelle, 81, no. Stannic and Stannous, 65, Slibnite, 61. ^ Strontianite, 106. Strontium, 106. Strychnine 133. Sulphides, 81, 126. Sulphocyanic acid, 127. Sulpho-salts, 57. Sulphuretted hydrogen, 126. Sulphuric acid, 126. Sulphurous acid, 127. Superphosphates, 107. Syivine, 112. Tartar emetic, 62. Tartaric acid, 127. Thebaine, 133. Theine, 133. Thenard's blue, 90. Thiosulphuric acid, 127. Tin, 64. Transposition, 53. Triphylline, 114. Turnbull's blue, 85. Type metal, 34, 62. Ultramarine, 81. Valenlinite, 61. Veratrine 133. Vermillion, 40. Vitriol blue, 72. '• green, 83. " white, 97. Wad, 94. Washing precipitate, 34. Water-glass, n3 Weak bases, 79. White vitriol, 97. Witherite, 105. Wood's metal, 70, 75. Yellow amnion, sulph., 57. Zinc, 97.