A COURSE IN QUALITATIVE CHEMICAL ANALYSIS THE MACMILLAN COMPANY NEW YORK • BOSTON • CHICAGO ATLANTA • SAN FRANCISCO MACMILLAN & CO., Limited LONDON • BOMBAY • CALCUTTA MELBOURNE THE MACMILLAN CO. OF CANADA, Ltd. TORONTO A COURSE IN QUALITATIVE CHEMICAL ANALYSIS BY CHARLES BASKERVILLE, Ph.D., F.C.S. PROFESSOR IN THE DEPARTMENT OF CHEMISTRY COLLEGE OF THE CITY OF NEW YORK AND LOUIS J. CURTMAN, Ph.D. INSTRUCTOR IN THE DEPARTMENT OF CHEMISTRY COLLEGE OF THE CITY OF NEW YORK Nefo Pork THE MACMILLAN COMPANY 1910 All rights reserved Copyright, 1910, By THE MACMILLAN COMPANY. Set up and electrotyped. Published December, 1910. Nortonoh Crests J. S. Cushing Co. - Berwick & Smith Co. Norwood, Mass., U.S.A. PREFACE Experience has shown the authors that a quantitative discrimination in Qualitative Analysis for students is rarely exercised, although it is generally conceded that qualitative analysis should serve not only as a means for determining the components in an " unknown," but also, though roughly, the proportions in which the ingredients are present. This failure to gain experience in the evaluation of tests is, we believe, largely due to the fact that in the preliminary work which precedes the analysis of " unknowns," adequate provision is not made whereby the quantitative aspect as well as the qualitative meaning of the results can be simultaneously studied. In our work the quantitative feature is emphasized, first, by early acquainting the student with the fact that there is a limit in the quantity of an element in a definite volume that may be detected by a given reaction; second, by the use of known solutions which are prepared to contain definite amounts of the elements of a group in a definite volume. For example, the label on the bottle states the metallic concentration of its contents, and the student by using a specified volume knows precisely how many milligrams of each metal he is using. The advantages in the use of such a solution are: first, the size of the precipitates may be controlled by the instructor and precipi- tates of unwieldy bulk avoided; second, and this is the chief advantage, in addition to familiarizing himself with the reac- tions and separations, the student also learns the relation between the quantity of metal present and the size of the precipitate which it yields. The quantitative information thus acquired in the analysis of known solutions is subsequently applied to the V VI PREFACE analysis of unknowns. In consequence of this training, the student is able to report not only the qualitative composi- tion of his unknowns, but also approximately their quantitative composition. The value of such training cannot be over- estimated. Our students rarely find any difficulty in differ- entiating between a trace and a significant amount. In the schemes of analysis for the metals, preference has been given in a vast majority of cases to precipitation tests, because of their quantitative significance. Detailed methods for the preparation of solutions of definite strength are given for the assistance of the instructor. The value of introducing preliminary experiments in a course in Qualitative Analysis is a mooted question. We believe that they should be restricted to those which are utilized in the schemes of separation. Experience has shown us that it is a good plan to have in the laboratory several sets of bottles containing solutions of known strength of the salts of the vari- ous metals. The students are encouraged to use these to verify any of the preliminary tests in case of doubt. By comparing in special cases the results obtained with known solutions with those obtained with the unknown, a definite knowledge of that particular reaction is fixed in the student's mind. Students should be encouraged to use short cuts, and the use of the preliminary test as a means of indicating short cuts. It is assumed that the student who begins the study of Quali- tative Analysis has had a course in laboratory work in General Chemistry, and has thus become familiar with such operations as making solutions, precipitations, evaporations, ignitions, and the preparation of borax beads. He has also become familiar with the term "solubility." For these reasons such matters have received but scant attention in this book. The student should have had not a little experience in writing equations. The application of his knowledge, often meager, to the processes of oxidation and reduction, as well as to the mode of operation of reagents producing these changes, is not always clear; or his knowledge is not sufficient to cope with the cases met with in Qualitative Analysis. For this reason PREFACE VII these matters have been brought together and coordinated with some detail in the beginning of the book. The essential features of this book may be seen from the plan which is here briefly outlined. i. The chief reactions of the metals are first given with suffi- cient detail and completeness to enable the student to thor- oughly understand the basis of and the limitations to the schemes of analysis adopted. Reactions not utilized in the schemes are also given to supply information which may be turned to account in making additional confirmatory tests and in devising schemes other than those given; they also supply a number of qualitative facts upon which important quantitative methods are based. As the vast majority of students who take Qualitative Analysis subsequently pursue a course in Quantita- tive Analysis, this information supplies the foundation of fact which we believe should be given in the qualitative course. 2. An outline of the method of analysis to be employed follows. This is in the nature of a resume of the chief reac- tions, in which distinctions are emphasized with a view to their use in separations. Details in manipulation are purposely omitted in this discussion, in order that the main features and chemistry thereof may be clearly understood. 3. The scheme of analysis is then taken up. The directions are clear, especial attention being given to the amounts of reagents to be taken, as well as the most appropriate vessel to be used and its size. 4. Then follow notes. Under this head, additional informa- tion, which would obstruct the reading of the text, is supplied. This information is intended to supply the reasons for unusual details or procedures in the text and precautions that are to be taken, but it applies chiefly to matters relating to the correcting of errors and to the clearing up of doubtful results. Supplying the reasons for every step, it is believed, will go a long way toward doing away with the too frequent practice of blindly following directions. The so-called rarer elements have been omitted, and some of the commoner ones also, as their study is not essential in a VIII PREFACE first course in Qualitative Analysis. This is essentially a prac- tical book and not the place to exploit any particular hypothesis or theory. Its contents are directed toward the study of reac- tions and the operations and methods made use of in the identi- fication of unknown substances or mixtures. Discussions of the theory of electrolytic dissociation and the mass action law are now presented in courses in General Chemistry. These may be further applied in lectures, which form an integral part of the study of Qualitative Analysis, and are of especial value when accompanied by collateral reading of special works on the theoretical phases of the subject. The authors are aware that no directions, however detailed and carefully written, can replace the resourcefulness of the instructor; he should be particularly observant during the first laboratory periods when methods of work are about to be acquired, and should be quick to give personal attention to those who need it most. There are those to whom skill in manipulation comes naturally, but it should be remembered that those not so gifted may, by perseverance and constant practice, acquire an unusual degree of skill. A bright student will soon learn that he may carry on two or more operations at once, as filtration of one liquid while evaporating another, without the suggestion of an automatic teacher; but the work of the class as a whole will suffer unless the instructor is alive and richly suggestive. In the recitations, questions may be asked concerning sepa- rations and tests, other than those given in the schemes, but which are given in the descriptive portion of the book; these serve to stimulate original thinking and give opportunity for the exercise of individual ingenuity. Written quizzes have little value beyond securing figures for grading, unless the papers, after being corrected, are discussed with the stu- dents. Well-known works on the subject have been drawn from more or less. We wish especially to mention Fresenius, Pres- cott and Johnson, Treadwell, Knoevenagel, and A. A. Noyes. Conflicting statements appear in many books; in some cases PREFACE IX consultation of the original sources sufficed, in others research was necessary, to arrive at a decision. We are indebted to Mr. W. A. Hamor, who assisted in fol- lowing the proof sheets. CHARLES BASKERVILLE. LOUIS J. CURTMAN. College of the City of New York, November, 1910. TABLE OF CONTENTS PAGE Preface v Introduction (a) Analytical Chemistry Defined i (£) Terms i (c) General Directions for Laboratory Work .... 3 (d) Equations 4 (^) Oxidation and Reduction 7 (/) Classification 14 PART I The Metals. Descriptive and Analytical 17 PART II (a) The Acids. Descriptive and Analytical 104 (£) Preliminary Examination 141 PART III Complete Analysis (a) Scheme for Group III in the Presence of Organic Matter, Phos- phates, etc. .164 (3) Preparation of the Solution 168 (c) Alloys 171 (d) Treatment of Insoluble Substances 172 (>) Acid Analysis of Minerals and Metallurgical Products . .178 Appendix i. Table of Solubilities 179 2. Reagents 184 3. List of Apparatus 189 4. Preparation of Unknowns 190 Index 195 XI INTRODUCTION That branch of chemistry concerned with the problem of de- termining the composition of substances or mixtures of sub- stances and identifying them is called Analytical Chemistry. Substances are identified by their properties which, for any given substance, are unchangeable and fixed. The problem of identi- fying a substance, therefore, resolves itself into a determination of a sufficient number of the properties of the substance in ques- tion. For a mixture of substances it is frequently necessary to separate the components before the latter can be identified. The term analysis is used in Analytical Chemistry to denote the systematic examination of a substance and includes all the operations, whether analytic or synthetic, involved in the process of identifying a substance. An analysis may be either qualita- tive or quantitative. If we satisfy ourselves that a ten-cent piece consists of copper and silver, our knowledge is qualitative ; but if we ascertain how much copper and silver are contained in a given weight of this alloy, our knowledge becomes quantita- tive. For qualitative purposes, only a few of the many proper- ties possessed by a substance are utilized for its identification. In general, those which are striking and most rapidly determined are the ones selected; chief among these are color, state of aggregation, solubility, and certain chemical properties. Terms employed in Qualitative Analysis If to an aqueous solution of sodium chloride we add a water solution of silver nitrate, a white, curdy, solid substance forms and settles to the bottom of the containing vessel. The chem- ical change, as evidenced by the formation of the new sub- stance, AgCl, and brought about by mixing these solutions, is called a reaction. The AgNO3 solution employed to produce this reaction is called the reagent. The solid, insoluble AgCl is 2 QUALITATIVE CHEMICAL ANALYSIS called the precipitate. Any reaction which is accompanied by the formation of a precipitate is called a precipitation reaction and the process of forming a precipitate is known as precipita- tion. If we allow the precipitate to settle and carefully pour off the clear supernatant liquid, the process is known as decantation. If, without allowing the precipitate to settle, we pour the liquid holding the precipitate in suspension on a filter paper, supported in a funnel, or other design of filtering apparatus, the liquid will pass through the fine pores of the filter and will be thus sepa- rated from the precipitate, which will remain on the paper. This process is known as filtration and is frequently resorted to for the separation of a liquid from a solid. The precipitate on the filter is sometimes called the residue, while the liquid which passes through is called the filtrate. The equation for the reaction, omitting consideration of the water, is - NaCl + AgNO3 = AgCl + NaNO3. Assuming that an excess of silver nitrate has been used, the pre- cipitated and filtered AgCl will be wet with a solution containing NaNO3 and AgNO3. As these are soluble in water, it is evi- dent that they may be removed by treating the precipitate on the filter with water. The process of removing soluble impuri- ties from insoluble substances by treatment with water is known as washing.* All precipitates should be thoroughly washed. Much time will be saved in washing precipitates by allowing each portion of water completely to pass through the filter before adding the next. The completeness of the washing is ascertained by testing a portion of the last washings for the sub- stance it is desired to remove. In the above case, after washing several times with small amounts of water, the last portion should be tested by adding to it a little NaCl solution. If no precipitate is formed, the washing may be considered complete ; if a precipitate or cloudiness is obtained, it is an indication that all the AgNO3 has not been washed out of the precipitate. In the latter case, the washing should be continued until a negative test with NaCl is obtained. * Water is used here as the typical solvent. INTRODUCTION 3 General Directions for Laboratory Work Order and cleanliness are essential to success in qualitative analysis. It is a good plan never to put away apparatus until the latter is clean and ready for use. Beakers, evaporating dishes, and funnels, after being thoroughly washed, should finally be rinsed with distilled water, inverted, and allowed to drain and dry on a clean towel which has been spread out on the floor of the cupboard of the desk. Every piece of appara- tus in the desk should have a fixed place. Iron ware should not be kept in the same drawer or compartment with glass ware. Some attention should also be given to the arrangement of apparatus on the desk. In general, the desk space should be roughly divided into two parts - namely, one reserved for heat- ing, and the other for filtering, washing, testing, etc. Reagent bottles must never be allowed to accumulate on the desk, but should be returned to their proper places immediately after use. Accidents, which occasionally happen even to the most careful workers, must receive immediate attention; the broken glass should be collected and thrown into a special crock provided for this purpose; and the desk top should be sponged off, the apparatus cleaned, and the analysis begun anew. Vessels con- taining solutions or precipitates which are to be set aside for future examination should be properly labeled. Reagents.-The bottles on the desk are filled with solutions known as reagents. They should occasionally be wiped off with a moist rag and every effort should be made to keep their contents pure. When once an impurity is allowed to enter a reagent bottle, the reagent becomes worthless. The importance of the care of reagents will be appreciated when it is remem- bered that the value of the whole analysis is dependent upon their purity. The bottles should always be kept properly stoppered, and the stopper under no circumstances should be placed on the desk, where it is likely to take up impurities and thus eventually contaminate the reagent. The student should make it a rule always to hold the stopper between the fingers while using the bottle and to return it immediately after use; 4 QUALITATIVE CHEMICAL ANALYSIS in this way danger of contamination from the desk is avoided. Acids of two strengths will be found on the desk. Before using the reagents, the student should be sure that he is using the proper ones. If he will learn in each case the reason for adding the reagent, the error arising from the use of the wrong reagent will seldom, if ever, occur. Too much reagent is worse than adding too little. Reagents should always be added drop by drop and should never be used in great excess, unless other- wise directed; but under no circumstance is an excess of reagent to be poured back into the bottle. All operations which result in the production of fumes of any kind must be conducted under the hood, with the window of the latter almost completely closed. Equations A chemical equation is a shorthand means of expressing a reaction. The equation BaCLj + H2SO4 = BaSO4 + 2 HC1 gives us, first, the qualitative fact that when BaC^ solution is treated with H2SO4, a precipitate of BaSO4 is formed, together with hydrochloric acid; and, second, the quantitative relations between the substances reacting - namely, that 137 parts by weight of barium in solution will require 98 parts by weight of H2SO4 for its complete precipitation; or that (137 + 70) parts by weight of BaCl2, when treated with 98 parts of H2SO4, yield (137 + 96) parts by weight of BaSO4 and (2 x 36.5) parts by weight of HC1. In the case of gases, equations give us, besides qualitative and quantitative relations, those of volume also. The equation CO2 T C = 2 CO states that when one volume of CO2 is reduced by the agency of carbon, two volumes of carbon monoxide are produced. Before writing an equation, we must first know the facts. True, we may reason by analogy and foretell a reaction, and at INTRODUCTION 5 once write its equation; but such equations are to be looked upon with doubt until verified by actual experimentation. Briefly, then, to write an equation, we must know the formulas of the substances entering into the reaction, sometimes called the factors and appearing on the left side of the equality sign, and also we must know one or more of the essential products. All of the latter need not be known; they can, with a little knowledge, be " worked out." It is a good plan to designate by means of an arrow pointing downward the formulas of insoluble substances and with an arrow pointing upward the formulas of gases; eg., a glance at the equation CaCO3 + 2 HC2H3O2 = Ca(C2H3O2)2 4- H2O + CO2 shows that CaCO3 is a solid, that CO2 is a gas, and that the other substances are in solution. Any determined method, however, will answer the purpose, and as a rule it is applied only to the products of a reaction. The Law of the Conservation of Weight states that during a chemical change there is no loss in weight. From this it follows that the total weight of the factors must be equal to the total weight of the products. The Law of the Conservation of Ele- ments states the immutability of the elements during chemical change. From the combination of these two fundamental laws of chemistry, it follows that in any chemical equation, the same number of combining weights of each element must appear on both sides. This last rule, with a knowledge of the formulas of the factors and products, suffices for the writing of an equation. One need not attempt to remember the coefficients appearing in chemical equations, but should invariably work them out. The method of doing this will be explained later in connection with the writing of a number of rather complex equations given under " Oxidation and Reduction." Types of Reactions Reactions may be classed under three heads, viz., synthesis, analysis, and metathesis. A synthetic reaction is one in which 6 QUALITATIVE CHEMICAL ANALYSIS a compound is formed from its elements, or, in general, a more complex compound formed by the union of simpler ones, thus : - CaO + H2O = Ca(OH)2; Fe(CN)2 + 4 KCN = K4Fe(CN)6; H2 + C12 = 2HCL An analytical reaction is one in which some complex com- pound is broken down into its elements or into simpler com- pounds; for example:- XT r r NaCl = Na 4-Cl; CaCO3 = CaO 4- CO2. A metathetical reaction is one in which there is an exchange of radicals; these represent by far the largest number met with in qualitative analysis ; for example : - (a) AgNO3 + HC1 = | AgCl 4- HNO3; Cu(NO3)2 + H2S = ;CuS + 2HNO3; (0 BaCl2 + H2SO4 = |BaSO4 + 2HCl. Neutralization of bases by acids come under this head. NaOH + HC1 = NaCl 4- H2O; NH4OH + HC2H3O2 = NH4C2H3O2 + H2O ; Ca(OH)2 4- H2SO4 = | CaSO4 4- 2 H2O. In some cases it is necessary to know the conditions under which the reaction takes place before the equation can be written; for example, if copper nitrate is mixed with dilute H2SO4, no apparent reaction takes place, but if the mixture is boiled till SO3 fumes are given off, the following reaction occurs : - + h2SO4 = CuSO4 4- f 2 HNO3. Similarly, on boiling, the following reactions will take place: - NH4C1 + NaOH = NaCl + f NH3 4- H2O, NaCl + H2SO4 = NaHSO4 + f HC1; for, at the boiling temperature, HC1, NHg, and HNO3 volatilize as gases. In writing a complicated equation, it is convenient to consider INTRODUCTION 7 the reaction as taking place in several stages; * thus, when ammonium sulphide is added to an aqueous solution of A1C13, a precipitate consisting of the hydroxide and not the sulphide is obtained, and at the same time the evolution of H2S is notice- able. This reaction becomes easily comprehended if we con- sider it as taking place in two stages, viz.: - (i) 2 A1C13 + 3(NH4)2S = A12S3 + 6 NH4C1; but as A12S3 cannot exist in contact with water, it decomposes according to the equation (2) A12S3 + 6 HaO = 12 A1(OH)3 + f 3 Has. Combining (i) and (2), and eliminating A12S3, which appears on opposite sides of the equations, we get as the final equation 2 AlCl3+3(NH4)2S + 6 H2O= I 2 Al(OH)3 + 6 NH4Cl+f 3 H2S. Still more complicated equations will be met with which can be readily written if first resolved into two or more single equa- tions, and the latter are then combined into one equation with the elimination of common terms. Oxidation and Reduction Oxidation may be defined as a chemical change in which oxy- gen or some other electro-negative element or radical is added, or hydrogen or some other electro-positive element or radical is removed. Reduction is the reverse process. The substance which effects the oxidation is called the oxidizing agent; and that causing the reduction is called the reducing agent. A con- sideration of the following examples will make the matter clear: - Oxidation : (1)2 FeO + O = Fe2O3, (2) FeCl2 + Cl = FeCl3, (3) 2 FeSO4 + H2SO4 + Br2 = Fe2(SO4)3 + 2 HBr, (4) SnClg + 2 HgCl2 = ShCl4 + 12 HgCl; * Since we do not know what causes chemical activity, the exact mechanism of a reaction is unknown and is perhaps unknowable; but any device tending to elucidate or simplify a complicated reaction, it is believed, should be utilized, although experi- mental proof for these devices cannot always be supplied. 8 QUALITATIVE CHEMICAL ANALYSIS Reduction: (5) CuO + H2 = Cu 4- H2O, (6) 2 FeClg + SnCl2 = 2 FeCl2 + SnCl4, (7) Fe2(SO4)8 + H2 * = 2 FeSO4 + H2SO4, (8) 2 HgCl2 + SnCl2 = | 2 HgCl + SnCl4. The first four equations represent oxidation reactions, for in each case oxygen or some other electro-negative element or radical has been added. The last four are types of reduction processes, for in each case oxygen or some other electro-nega- tive element or group has been removed. If we leave out of consideration those simple cases of oxidation and reduction in which there is a direct addition of the oxidizing or reducing agent with the production of a single product, as in (i) and (2), and carefully examine the others, we find that in every case oxidation is accompanied by reduction, the oxidation of one substance always involving the reduction of another; thus, in example (4), the oxidation of SnCl2 has been accomplished at the expense of the HgCl2, which has been reduced at the same time to HgCl. In (5) the reduction of CuO has been simultane- ous with the oxidation of H to H2O. Similarly, in (6) the reduc- tion of FeCl3 has been accompanied by the oxidation of SnCl2 to SnCl4. Further examination of the above eight examples of oxidation and reduction shows that oxidation is accompanied by an increase in valence, while in reduction reactions, a lowering of valence is observed; thus, in (4) the valence of Sn has been increased from 2 to 4, while at the same time the valence of Hg has been lowered from 2 to 1. In general, we may say that any reaction in which there is an increase in valence is one of oxidation, while a reaction in which there is a lowering of valence is one of reduction. Oxidizing Agents The principal oxidizing agents are: oxygen, the halogens, HNO3, aqua regia, KC1O3, K2Cr2O7, KMnO4, Na2O2, H2O2, and PbO2. * The hydrogen is generated in contact with the Fe2(SO4)s by the action of a metal on an acid. INTRODUCTION 9 The chief reducing agents are: nascent H, SnCI2, H2S, H2SO3; and, at elevated temperatures, C, KCN, and organic matter. The halogens either add directly (i); or else may be regarded as oxidizing the acid by removing the H, setting free the acid radical, which then adds on (2); or with bases by removing the metal and thus leaving the hydroxyl radical (OH) to add on (3). Examples of such changes follow. (1) FeCI2 + Cl = FeCl3; (2) 2 FeSO4 + H2SO4 + Cl2 = Fe2(SO4)3 + 2 HC1; (3) Fe(OH)2 + KOH + Cl = Fe(OH)3 + KC1. Nitric acid acts as an oxidizing agent by virtue of its ready decomposition'according to the equation 2 HNO3 = 2 NO + H2O + 3 O. From this it is evident that 2 formular weights of HNO3 yield 3 combining weights of O. The liberated O may be regarded as acting indirectly in the same way as Cl; that is, by oxidizing the H of the acid and liberating the acid radical, which then adds on, thus : - (1) 2 HNO3 = 2 NO + H2O + 3 O ; (2) 3 H2SO4 + 30 = 3 H2O + 3 SO4. Since 2 formular weights of FeSO4 require one SO4 for com- plete oxidation, and since 3 SO4 radicals are made available by 2 HNO3, it follows that 2 formular weights of HNO3 will oxidize 6 FeSO4, to wit: - (3) 6 FeSO4 + 3 SO4 = 3 Fe2(SO4)3. Adding (1), (2), and (3), and eliminating the factors appear- ing on both sides in the equations, as 3 O and 3 SO4, the final equation becomes- 6 FeSO4+ 2 HNO3+ 3 H2SO4=3 Fe2(SO4)3+f 2 NO + 4 H2O. Aqzia regia is not only a good solvent but is also an excellent oxidizing agent. Its action is practically the same as that of Cl. It is prepared by mixing 1 part of HNO3 with 3 parts of HC1. 10 QUALITATIVE CHEMICAL ANALYSIS When heated alone, it is said to yield N0C1 and Cl, according to the equation 3 HC1 + HNO3 = NOCI + +2 H2O ; but in the presence of an oxidizable substance all of the Cl is available, so that the equation with aqua regia may be written as follows : - 3 CoS + 6 HC1 + 2 HNO3 = 3 CoCl2 + 2 NO + 4 H2O + 3 S. If the action of aqua regia on a sulphide is long continued, the liberated S will be partially or entirely oxidized to H2SO4:- S + 6 Cl + 3 H2O = SO3 + 6 HC1; SO3 + H2O = H2SO4. Potassium chlorate, KC1O3, when used in conjunction with HC1, is a powerful oxidizing agent. In effect it is similar to aqua regia and chlorine, as the following equations will indi- C3.tC I (1) 2 KC1O3 + 2HC1 = 2 HC103 + 2 KC1, (2) 2 HC1O3 = H2O + 2 C1O2 + o, (3) 2 HC1+ O = H2O + C12. Adding (1), (2), and (3), we get the following:- 2 KC1O3 +4 HC1 = 2 KC1 + 2 H2O + 2 C1O2 + Cl2. Potassium dichromate, K2Cr2O7, contains the acid anhydride CrO3; its composition may be represented by the formula 2 CrO3 + K2O. As 2 CrO3 will, on reduction, yield Cr2O3 + 3 O, it is evident that one formular weight of K2Cr2O7 possesses the same oxidizing power as two formular weights of HNO3, both yielding 3 combining weights of O. We therefore obtain the following equation for the oxidation of ferrous sulphate by K2Cr2O7: - 6 FeSO4 + K2Cr2O7 4- ^H2SO4 = 3 Fe2(SO4)3 + K2SO4 + Cr2(SO4)3 + ^H2O. The reduction of 2 CrO3 to Cr2O3 leaves the latter, as well as K2O, behind. These basic oxides readily dissolve in H2SO4 with the formation of sulphates which appear in the above equation. The amount of H2SO4 needed to balance the above reaction may be calculated from the following considerations: Cr2O3 INTRODUCTION 11 requires 3 formular weights of H2SO4 for its solution; K2O requires 1, thus making a total of 4. In addition to this quan- tity, we must consider the amount necessary to react with the 3 combining weights of O derived from the 2 CrO3; this will require 3 more H2SO4, making the total 7, which becomes the coefficient of H2SO4 in the above equations. We should also add 7 H2O to the right-hand member of the equation for a reason which must be evident to the reader. Potassium permanganate, KMnO4, behaves in acid solution as K2O + 2 MnO + 50 = 2 KMnO4; that is, 2 KMnO4 yields 5 combining weights of O. The equation for the oxidation of ferrous sulphate by KMnO4 may be written from the following considerations: 5 combin- ing weights of O will liberate 5 SO4 radicals; but as one SO4 suffices for the oxidation of 2 FeSO4, it is evident that 2 KMnO4 will oxidize 10 FeSO4, and we get the equation - 2 KMnO4 + 10 FeSO4 + 5 H2SO4 = 5 Fe2(SO4)3 "h MnO + K2O + 5 H2o. However, K2O and 2 MnO readily dissolve in H2SO4 with the formation of sulphates; three additional formular weights of H2SO4 must therefore be added to the above equation, together with 3 formular weights of water produced by the solution of the oxides in the acid. The final equation becomes - 2 KMnO4 + 10 FeSO4 + 8 H2SO4 = 5 Fe2(SO4)3 + MnSO4 + K2SO4 + 8 H2O. Sodium dioxide, Na2O2, acts as an oxidizing agent by virtue of its instability in water or when it is heated in the presence of an oxidizable substance in consequence of which it gives off oxygen:- Na2O2 + H2O = 2 NaOH + O. A solution of hydrogen dioxide, H2O2, also serves as an oxidiz- ing agent for the same reason : - H2O2 = H2O + O. 12 QUALITATIVE CHEMICAL ANALYSIS Na2O2 is, however, preferable to H2O2, for the latter can only be safely* used in diluted form (3%), while the former can be used in any concentration. NajO2 has the further ad- vantage of supplying at the same time sodium hydroxide. For oxidation in alkaline media, therefore, Na^^O-j is the better reagent. Lead dioxide, PbO2, like Na^jO^ yields O according to the equation - p^ = pbO Q It is employed to effect oxidation in acid media, as in the conversion of MnO to HMnO4; the PbO is converted by the excess of acid present into a salt. Reducing Agents Any oxidizable substance can be utilized as a reducing agent, since, in order to reduce, it must be capable of oxidation. Nascent H t acts as a reducing agent by either adding itself directly (1) or by its ability to unite with, and thus remove, the halogens contained in the compound (2): - (1) As + 3H = AsH3; (2) SbCl3 + 3 H = Sb + 3 HC1. Nascent H may be prepared by the action of an acid or alkali on certain metals: - Zn + H2SO4 = ZnSO4 -I- H2; Zn -I- 2 NaOH = Na2ZnO2 4- H2; Al + 3 NaOH = Na3AlO3 + 3 H. Nascent H may therefore be employed in both acid and alkaline media. Stannous chloride, SnCl2, acts as a reducing agent, preferably * A 30% solution has recently become very useful in the laboratory, especially in the hands of experienced chemists. We regard the above statement as particularly applicable to student work. f Nascent H is used here to designate the hydrogen which is formed when the acid and metal or alkali and metal are both in contact with the solution to be reduced. INTRODUCTION 13 in an acid solution, by virtue of the ease with which it readily oxidizes to SnCl4: - SnCl2 + 2 HgCl2 = SnCl4 4- 2 HgCl; SnCl2 4- 2 FeCl3 = 2 FeCl2 4- SnCl4; H3AsO4 4- 2 H Cl 4- SnCl2 = H3AsO3 4- H2O 4- SnCl4. Hydrogen sulphide, H2S. The reducing action of this com- pound depends upon the readiness with which it decomposes into H and S; the former then acts as nascent hydrogen, while the latter separates out in the solid state. Nearly all oxidizing agents are reduced by H2S with the separation of sulphur, thus:- 2 HNO3 + 3 H2S = 4 H2O 4-2 NO 4- 3 S. Hence, sulphides soluble in HNO3, like those of Pb, Bi, and Cu, do not liberate H2S, because the latter at once acts on the excess of HNO3 present with the liberation of S, as indicated in the equation above. Halogens, aqua regia, ferric salts, potassium permanganate, and chromates are also reduced by H2S with the separation of S, and, in some cases, with a change in color of the solution (see pp. 44-45). Sulphurous acid, H2SO3. The action of sulphites or sulphur- ous acid depends upon the ease with which they remove oxygen from compounds with the formation of sulphates and sulphuric acid, respectively, thus : - H3AsO4 4- H2SO3 = H3AsO3 4- H2SO4. While H2SO3 is an excellent reducing agent for arsenic when in the pentavalent condition, it cannot be used in a complete analysis if the alkaline earths are known to be present, because of the formation of H2SO4 as one of the products of the reduc- tion. The reduction of arsenic compounds with H2SO3 is best accomplished in a pressure bottle at ioo° C. Carbon acts as a reducing agent by virtue of its ability to oxidize to CO and CO2, thus: - CuO 4- C 4- (heat) = Cu 4- CO. 14 QUALITATIVE CHEMICAL ANALYSIS The use of potassium cyanide, KCN, as a reducing agent in the " dry way " depends upon its ability to take up O and form KCNO:- SnO2 4- 2 KCN = Sn + 2 KCNO. Classification In the analysis of a solution for metals, it has been found convenient to first separate them into groups by the use of cer- tain reagents known as group reagents. If to a solution containing all the metals in the form of salts we add a slight excess of dilute hydrochloric acid, a precipitate consisting of the chlorides of silver, mercury (-ous), and lead will form. These metals are classed together and designated as the first group. If, now, the precipitate of the chlorides of the first group is filtered off and the filtrate, which is acid from the excess of HC1 used, is treated with a stream of H2S gas, there will form a precipitate consisting of the sulphides of mercury (-ic), lead,* bismuth, copper, cadmium, arsenic, antimony, and tin. These metals are therefore classed together and col- lectively are known as the second group. If the filtrate from the second group is rendered alkaline with ammonium hydrox- ide,! and then ammonium sulphide is added, a precipitate of the hydroxides of aluminum and chromium, together with the sulphides of iron, nickel, cobalt, manganese, and zinc, will form. These constitute the third group. The filtrate from this group will contain an excess of NH4OH and some NH4Cl,t in addition to all the metals not included in the previous three groups. If to this filtrate we add ammonium carbonate in slight excess, a precipitate consisting of the carbonates of barium, strontium, and calcium will form; these constitute the fourth * Since PbCl2 is somewhat soluble in water, some of it will pass into the filtrate from Group I. and will be precipitated in the second group as sulphide. Pb, there- fore, belongs to both groups. t When the solution which is acid with HC1 is rendered alkaline with ammonium hydroxide, NH4CI forms; the presence of this salt prevents magnesium from precipi- tating along with the metals of Group III. For the same reason, magnesium is not thrown down in the fourth group. INTRODUCTION 15 group. The final filtrate will contain all the other metals not precipitated by the previous group reagents; they are mag- nesium, sodium, potassium, and ammonium, and these form the fifth group.* The division of the metals into groups is thus seen to depend upon their behavior, when in solution, towards certain reagents added in a certain order. If we were to use different reagents, the grouping would be different. It is equally important to remember that the order of the addition of the reagents is as Vital for the above' classification as the choice of the reagents; for if we were to reverse th& order,-i.e., begin with NH4OH, then add (NH4)2CO3 and then (NH4)2S, - we should get quite a different classification. It is furthermore to be noted that each reagent, taken in the order given, is capable of separating its own group from those which follow and not from those which precede it in the regular order. Below is given in tabular form the separation of the metals into groups with the formulas of the compounds which are formed. Solution containing all the metals in the form of salts. Add HC1 and filter. Precipitate: AgCl, PbCl2, HgCl. Group I. Filtrate: Groups II.-V. + excess HC1; pass in H2S and filter. Precipitate: HgS, PbS, Bi2S3, CuS, CdS, As2S3, Sb2S3, SnS. Group II. Filtrate: Groups III.-V.+ HC1 + H2S; make alkaline with NH4OH, add (NH4)2S and filter. Precipitate: A1(OH)3, Cr(OH)3, FeS, NiS, CoS, MnS, ZnS. Group III. Filtrate: Groups IV.- V.+NH4C1; add (NH4)2CO3 and filter. Precipitate: BaCO3, SrCO3, CaCO3. Group IV. Filtrate contains Mg, Na, K, NH4. Group V. * The rarer metals are not considered in the text. Special compendious treatises are necessary when they are included. 16 QUALITATIVE CHEMICAL ANALYSIS It is thus seen that - Group I. includes those metals whose chlorides are insoluble in water and in dilute acids, and are hence precipitated by HC1. Group II. includes those metals which are not precipitated by HC1, but whose sulphides are precipitated by H2S in acid solu- tions. The sulphides are, therefore, insoluble in water and in dilute acids. Group III. includes those metals which are not precipitated either by HC1 or H2S in acid solution, but are precipitated by (NH4)2S in solutions alkaline with NH4OH and in the pres- ence of NH4C1. Group IV. includes those metals unprecipitated by the reagents of Groups I., II., III., but which are precipitated by (NH4)2CO3 in the presence of NH4C1. Group V. includes metals not precipitated by the reagents of Groups I.-IV. PART I THE METALS Reactions of the Metals of Group I. The metals silver, mercury (-ous), and lead, comprising this group, are distinguished from all others by the insolubility of their chlorides in water and in dilute acids. With the exception of the nitrates and acetates, which are colorless, nearly all the salts of the metals of this group are insoluble in water. Silver i. Hydrochloric acid or a soluble chloride, when added to solu- tions of silver salts, gives a white, curdy precipitate of silver chloride (AgCl) which darkens on exposure to light. The pre- cipitate is insoluble in water, the solubility being approximately i part in 700,000 parts of water; it is insoluble in dilute acids and in dilute aqua regia, but is somewhat soluble in concentrated acids. Ammonium hydroxide readily dissolves it, with the for- mation of silver ammonia chloride [Ag(NH3)2Cl] : - AgCl + 2 NH3 = Ag(NH3)2Cl, from which AgC 1 reprecipitates on acidification with nitric acid:- Ag(NH3)2Cl + 2 HNO3 = | AgCl + 2 NH4NO3. Silver chloride also dissolves in solutions of potassium cyanide and sodium thiosulphate; when heated, it fuses without decom- position. 2. Hydrogen Sulphide and soluble sulphides precipitate black Ag2S, insoluble in cold dilute acids, alkali hydroxides, and alkali sulphides; it is soluble in hot dilute HNO3, with the formation of AgNO3 and separation of sulphur. The reaction can be con- 17 18 QUALITATIVE CHEMICAL ANALYSIS sidered as taking place in two steps, the first consisting of the solution of the sulphide with the liberation of H2S and the second of the oxidation of the H2S by the excess of HNO3 present with the formation of water, nitric oxide, and the sepa- ration of sulphur: - (i) Ag2S + 2 HNO3= 2 AgNO3 + f H2S; (2) 2 HNO3 + 3 H2S = 4 H2O + f 2 NO + 13 S. Multiplying equation (1) by 3 and adding it to (2), with the elimination of 3 H2S, which appears on opposite sides, we get as a final result: - 3 Ag2S + 8 HNO3 = 6 AgNO3 + 4 H2O + f 2 NO + 13S. Other insoluble compounds of silver are: - AgBr - yellowish white; Agl-pale yellow; AgCN - white; Ag2O - brown. Silver is readily precipitated from its solutions by the more electro-positive metals, as Cu, Zn, Hg, and Fe, as well as by various reducing agents. Mercury (-ous) 1. Hydrochloric acid and soluble chlorides give with solutions of mercurous salts a white precipitate of HgCl (calomel), in- soluble in water, the solubility being about 1 part in 300,000; it is insoluble in cold dilute acids, but dissolves in strong nitric acid and in aqua regia, the latter oxidizing it to HgCl2. Am- monium hydroxide converts calomel into a black mixture of finely divided mercury and NH2HgCl, insoluble in excess of the reagent (this is the most characteristic reaction for mercurous salts): - 2 HgCl + 2 NH3 = |(NH2HgCl + Hg) + NH4C1. The black mixture dissolves in aqua regia with the formation of mercuric chloride: - (1) NH2HgCl + 3 Cl = f N 4- 2 HC1 + HgC^; (2) Hg + 2 Cl = HgCl2. Adding (1) and (2), we get NH2HgCl + Hg + 5 Cl = 2 HgCl2 + f N + 2 HC1. THE METALS 19 2. Hydrogen Sulphide gives with solutions of mercurous salts a black precipitate consisting of a mixture of mercuric sulphide and elementary mercury; it may be assumed that the mercurous sulphide which forms first, decomposes on account of its insta- bility, thus: - 2 HgCl + H2S = | Hg2S + 2 HC1; Hg2S = |HgS + |Hg. 3. Reducing Agents, as FeSO4 or SnCl2, rapidly reduce mer- curous salts to metallic mercury : - SnC^ + 2 HgCl = SnCl4 + 2 Hg. Lead 1. Hydrochloric acid and soluble chlorides give with solutions of lead salts, which are not too dilute, a white precipitate of PbCl2, soluble in 100 parts of cold and 25 parts of boiling water; from the latter on cooling, PbCl2 separates out in the form of needles. PbCl2 is much more insoluble in dilute HC1 than in water. Ammonium hydroxide changes it to a basic chloride [Pb(OH)Cl] which is extremely insoluble in water. 2. Dilute sulphuric acid and soluble sulphates precipitate white PbSO4 which is practically insoluble in H2O (1 part in about 30,000), but much more insoluble in the presence of dilute H2SO4 or alcohol. It is soluble to some extent in HNO3, and is completely soluble in fixed alkalies and in a hot strong solu- tion of NH4C2H3O2; from these solutions PbSO4 is reprecipi- tated on adding H2SO4. 3. Potassium Chromate (K2CrO4) precipitates yellow lead chromate, readily soluble in sodium hydroxide, but insoluble in NH4OH and acetic acid: - Pb(C2H3O2)2 + K2CrO4 = | PbCrO4 + 2 KC2H3O2. 4. Hydrogen Sulphide from slightly acid solutions of lead salts precipitates black PbS. In the presence of much HC1, H2S either fails to precipitate or else produces a red insoluble com- QUALITATIVE CHEMICAL ANALYSIS 20 pound of the formula PbCl2 • 2 PbS ; the latter is converted into black PbS by treatment with (NH^S or by diluting the solu- tion and passing in more H2S. Lead sulphide is insoluble in dilute acids, alkali hydroxides, and alkali sulphides. Hot dilute HNO3 dissolves it with the formation of the nitrate and separa- tion of sulphur. Hot concentrated HNO3 oxidizes it to sul- phate. Pb(NO3)2 4- H2S = | PbS + 2 HNO3. With dilute HNO3 this reaction occurs : - 3 PbS + 8 HNO3 = 3 Pb(NO3)2 + f2NO + |3S + 4 H2O.* With concentrated HNO3, the reactions may be represented by the following equations : - (i) PbS 4- 2 HNO3 = |PbSO44- 2 NO 4- H2; (2) 2 HNO3 + 3 H2 = 2 NO 4- 4 H2O. Multiplying (i) by 3 and adding it to (2) with the elimination of 3 H2, we get - 3 PbS + 8 HNO3 = 13 PbSO4 4- f 8 NO 4- 4 H2O. In a neutral solution containing 1 part of Pb in 100,000 parts of water, H2S will distinctly reveal its presence. H2S is, there- fore, an exceedingly sensitive reagent for the detection of Pb. 5. Sodium or Potassium Hydroxide precipitates white Pb(OH)2, soluble in excess: - Pb(C2H3O2)2 4- 2 NaOH = | PbCOH^ 4- 2 NaC2H3O2; | Pb(OH)2 4- 2 NaOH = Na2PbO2 4- 2 H2O. 6. Ammonium Hydroxide precipitates a basic salt, insoluble in excess. Other difficultly soluble compounds of Pb are: - Pbl2 - yellow; PbBr2 - white ; PbSO4 - white; basic car- bonate- white, 2 PbCO3 • Pb(OH)2; PbC2O4 - white. * Resolvable into two steps, as in the case of the solution of Ag>S in HNO3. THE METALS 21 GROUP I. Outline of the Process of Analysis An examination of the foregoing reactions shows that ah the metals of this group may be precipitated by HC1. By filtering off this precipitate, we should have on the filter a mixture con- sisting of the chlorides of Ag, Hg (-ous), and Pb. In order to identify the components of this mixture, it is first neces- sary to effect their separation; this can be accomplished by taking advantage of the difference of their behavior towards hot water and ammonium hydroxide. PbCl2 is completely soluble in hot water, while the others are practically insoluble. It is thus possible, by treating the mixed chlorides with a sufficient amount of hot water, to dissolve, or extract, the PbCl2. If the quantity of Pb is large, the hot aqueous extract on cooling will deposit the characteristic needles of PbCl2, and thus the presence of Pb may be proved. If the amount is small, the water extract will require further testing to prove that it con- tains Pb. The tests with H2SO4 and K2CrO4 will prove conclu- sively whether or not lead has been extracted and therefore its presence in the original solution. Having extracted all the Pb, the residue on the filter may consist of the chlorides of Ag and Hg. These can be readily separated by,reason of the solubility of the former in ammo- nium hydroxide, so that on treating the residue on the filter with this reagent we should obtain a filtrate and a residue. The former (if Ag is present) will contain the Ag in the form of Ag(NH3)2Cl, which, on acidification with HNO3, will yield a white precipitate of AgCl. While ammonium hydroxide dissolves AgCl, it offers at the same time an indication of the presence or absence of Hg; for the latter in the form of chloride is blackened by the reagent. To confirm the presence of Hg, the black residue is taken into solution with aqua regia, whereby it is converted into HgCl2, and the latter is then tested for with SnCl2. A white precipitate of HgCl or a gray precipitate of mercury (see this reaction under Hg) proves the presence of Hg in the original solution. 22 QUALITATIVE CHEMICAL ANALYSIS SCHEME I. To the clear original solution contained in a small beaker, add dilute HC1 drop by drop with constant stirring until no further precipitation takes place (i); add 2 cc. more in excess and filter. If no precipitate forms, the absence of Ag, Hg (-ous), and large amounts of Pb (2) is indicated; in that case pass to Scheme IL The filtrate should be caught in a beaker of at least 150 cc. capacity, labeled Groups II.-V., and reserved. The precipitate is first washed with 2 cc. of dilute HC1 (3), and finally with a stream of cold water from a wash bottle. Reject the washings. The precipitate may contain PbCl2, AgCl, and HgCl. Pour through the filter containing the precipitate several small portions of hot water, using about 2 cc. at a time and allow- ing each portion to drain before adding the next. Divide the aqueous extract into two equal portions, and test it for Pb by adding to the first portion dilute H2SO4- a white precipitate is PbSO4; to the second portion add K2CrO4-a yellow precipi- tate is PbCrO4 and confirms the presence of Pb. If no precipi- tates are obtained, Pb is absent from Group I. Repeatedly wash the precipitate on the filter paper with hot water until the washings no longer react with dilute H2SO4. The residue on the filter may now consist of AgCl and HgCl. Pour a few drops of ammonium hydroxide on the filter and catch the liquid passing through in a test tube. Repeat until about 15-20 drops have been used. A blackening of the residue on the filter indicates the presence of Hg. The ammoniacal extract, if not clear (4), should be passed again through the same filter and tested for Ag by acidifying with HNO3 (5). A white precipitate or cloudiness proves the presence of silver (6). To confirm the presence of Hg, remove (7) as much as pos- sible of the black precipitate remaining on the filter to a small evaporating dish, add 1-2 cc. of aqua regia (15 drops of con- centrated HC1 to 5 drops of concentrated HNO3), and heat under the hood on wire gauze until dissolved; boil to destroy the excess of aqua regia (9), dilute with 1 cc. of water, filter (10), THE METALS 23 if necessary, and test the clear solution for Hg by adding a few drops of SnCl2. A white precipitate, which may turn gray to black, confirms the presence of Hg. NOTES 1. A white precipitate, when dilute HC1 is used, may be due to SbOCl or BiOCl; the latter, however, dissolve in excess of HC1. To insure complete precipitation, HC1 must be added in slight excess; this point is best ascertained by filtering a small portion of the mixture, and adding to the clear filtrate a drop or two of HC1, when, if the precipitation was com- plete, no further precipitate will be obtained ; if a precipitate does form, more HC1 must be added to the original solution until the test shows complete pre- cipitation. A large excess of HC1 is to be avoided on account of the appre- ciable solubility of the chlorides in an excess. 2. PbCl2 is somewhat soluble in cold H2O, and while the presence of HC1 diminishes its solubility, a small amount always remains unprecipitated by HC1 and passes into the filtrate, from which it is precipitated by H2S in Group II. One must therefore always look for Pb in the second group. 3. The precipitate is washed first with HC1, instead of H2O, to prevent the formation of the oxychlorides of Bi and Sb. It is then washed with H2O to remove the HC1, which would interfere with the solution of PbCl2 in hot water. 4. If all the Pb has not been extracted, it will be changed by ammonia to the insoluble basic compound Pb(OH)Cl, which frequently passes through the filter. As this precipitate is soluble in HNO3, it does not interfere with the test for Ag. 5. In acidifying a solution, it is imperative thoroughly to mix the solu- tion after the addition of the acid, and then to test it with litmus. A solution in a test tube can be mixed by placing the thumb over the mouth of the tube and shaking. If the solution is contained in a beaker, thorough mixing is effected by stirring with a glass rod. 6. A cloudiness or turbidity is as conclusive a reaction as the formation of a large precipitate, provided the precaution is taken to see that both the solution to be tested and the reagent are perfectly clear. It is important to remember that when the amount of Hg is large and that of silver relatively small, ammonium hydroxide may fail to extract any AgCl, owing to the fact that the latter is reduced by the mercury of the black mixture (NH2HgCl + Hg) to the metallic state : - 2 AgCl + Hg = 2 Ag + HgCl2. When, therefore, a large black residue is obtained with ammonium hydroxide, and the test for Ag is negative, it becomes necessary to recover any Ag the black mixture may contain as described in note io. 24 QUALITATIVE CHEMICAL ANALYSIS 7. When the precipitate is large, a small amount may be removed with a glass or horn spatula. If too small to be handled in this way, recourse may be had to one of the following methods : (a) The funnel containing the filter is held horizontally with its rim resting against the edge of an evaporating dish or beaker, and the precipitate is washed out by directing a forceful stream of water from a wash bottle against the filter, at the same time giving the funnel a rotary motion, ^b) By carefully perforating the apex of the filter with a platinum wire and gradually enlarging the hole (in this way clog- ging of the stem with filter paper is avoided); the precipitate can then be readily washed out with a forceful stream of water from a wash bottle. In this case, as in (a), the water should be carefully decanted after the precipitate has settled and then the precipitate should be treated with the solvent directed to be used, (c) Where the amount is very small and firmly adheres to the filter, it can be gotten into solution by the following procedure: Remove the filter from the funnel, close it up, and dry it by pressing it between the folds of several thicknesses of filter paper; then unfold it, tear away portions to which no precipitate adheres, and spread the rest of the paper with the pre- cipitate uppermost on the bottom of an evaporating dish. Pour the solvent on the filter paper, heat, and stir with a glass rod till solution takes place; dilute with a little water and filter off the paper. 8. Aqua regia should always be prepared in small amounts immediately before use, and the cold mixture brought in contact with the substance to be dissolved and then heated. In contact with a substance it is capable of dis- solving, it acts like Cl, and as the latter is an oxidizing agent, the chloride formed will be that of the highest valence of the metal capable of existing under the conditions. 9. Prolonged boiling has the effect of destroying aqua regia in accord- ance with the equation 3 HC1 + HNO3 = N0C1 + 2 H2O + 2 Cl. The excess must be destroyed because its presence would oxidize the SnClj- to SnCl4. As the latter does not react with HgCl2, the test would be worthless. 10. The solution is diluted because of the possible presence of AgCI, which is appreciably soluble in strong HC1. The residue, after filtration, is tested for Ag by first thoroughly washing it with H2O, dissolving it in ammo- nium hydroxide, and reprecipitating by acidification with HN03. Reactions of Metals of Group II. Division A Mercury in Mercuric Salts Most of the salts of mercury are colorless and poisonous. The aqueous solutions of the normal salts have an acid reaction due to hydrolysis; they all volatilize on ignition. THE METALS 25 i. Potassium or Sodium Hydroxide gives a precipitate which at first is brownish but rapidly changes on further addition of the reagent to yellow HgO, insoluble in excess: - HgCl2 + 2 KOH = | HgO + 2 KC1 + H2O. 2. Ammonium Hydroxide produces in solutions of the mercuric chloride a white precipitate of mercuric amido-chloride; from solutions of the nitrate, a white precipitate of mercuric amido- nitrate : - HgCl2 + 2 NH4OH = | NH2HgCl 4- NH4C1 + 2 H2O. 3. Hydrogen Sulphide, on being passed slowly into a solution of HgCl2, forms at first a white precipitate which changes on further treatment with H2S to a yellow, then brown, and finally to a black precipitate of HgS. These light-colored precipitates are mixtures of HgCl2 and HgS in varying proportions; they are soluble in HNO3, and are converted by further treatment with H2S or with (NH4)2S to black HgS : - 3 HgCl2 + 2 H2S = (HgCl2 • 2 HgS) + 4 HC1; (HgCl2 • 2 HgS) + H2S = 13 HgS + 2 HC1. HgS is insoluble in dilute HC1; also in hot dilute HNO3 (differ- ence from the sulphides of Pb, Bi, Cu, and Cd). Prolonged boiling with strong HNO3 converts it into the white compound 2 HgS • Hg(NO3)2 which is quite insoluble in dilute HNO3. HgS is soluble in aqua regia with the formation of HgCl2 and the separation of sulphur; it is practically insoluble in (NH4)2S, but completely soluble in Na2S in the presence of sodium hydroxide. 4. Stannous Chloride, added in small quantity to a solution of HgCl2, precipitates white HgCl, which is reduced by an excess of the reagent to black metallic Hg : - SnCl2 + 2 HgCl2 = 12 HgCl + SnCl4; 2 HgCl + SnCl2 = 12 Hg + SnCl4. 5. Metallic Cu, Zn, or Fe, when introduced into a solution of a mercuric salt, acidified, precipitates metallic Hg. 26 QUALITATIVE CHEMICAL ANALYSIS 6. Potassium Iodide yields a red precipitate of Hgl2, soluble in excess of the reagent or the mercury salt. 7. If a dry mixture of a mercuric salt and Na2CO3 is heated in an ignition tube, a sublimate of metallic Hg will be formed in the upper portion of the tube. Bismuth Nearly all the bismuth salts are white or colorless. The aqueous solutions always have an acid reaction, and if the dilu- tion is large, the salt is decomposed with the formation of an insoluble basic salt; e.g.: - BiCl3 4- H2O^| BiOCl 4- 2 HC1. This reaction is very characteristic of bismuth salts and is inter- fered with by the presence of much acid. Because of their tendency to hydrolyze with water, aqueous solutions of bismuth salts can only be prepared with the aid of an acid. 1. Hydrogen Sulphide precipitates black Bi2S3: - 3 H2S 4- 2 BiCl3 = | Bi2S3 4- 6 HC1. The precipitate is insoluble in cold dilute acids, but dissolves in boiling dilute HNO3 : - (1) Bi2S8 4- 6 HNO3 = 2 Bi(NO'3)3 4- 3 H2S; (2) 3 H2S + 2 HNO3 = f 2 NO + 4 H2O 4- 13 S. Adding (1) and (2), and eliminating 3 H2S, we get BijSg 4- 8 HNO3 = 2 Bi(NO3)3 2 NO + 4 H2O +13 S. Bismuth sulphide is insoluble in (NH^S^. 2. Potassium, Sodium, or Ammonium Hydroxide precipitates white Bi(OH)3, insoluble in excess. Its insolubility in excess of NH4OH distinguishes it from Cu and Cd; the precipitate is soluble in dilute acids, however. 3. Water, when added in large amount to Bi salts, precipi- tates white basic salts; the chloride gives | BiOCl, the nitrate | BiONO3, and the sulphate |(BiO)2SO4. These are all soluble in dilute inorganic acids and are changed by H2S to Bi2S3: - 2 BiOCl 4- 3 H2S = | Bi2S3 4- 2 HC1 4- 2 H2O. THE METALS 27 4. Sodium Stannite in alkaline solution (prepared by adding NaOH to SnCl2 till the precipitate which first forms dissolves) gives with bismuth solutions a black precipitate of metallic bismuth : - (i) SnCl2 + 2 NaOH = | Sn (OH^ + 2 NaCl, (2) Sn(OH)2 + 2 NaOH = Na2SnO2 + 2 H2O ; Sodium Stannite (3) BiCl3 + 3 NaOH = | Bi(OH)3 + 3 NaCl, (4) 2 Bi(OH)3 + 3 Na2SnO2 = 12 Bi + 3 H2O + 3 Na^nOg. 5. Metallic Zn or Fe, when added to a solution of a Bi salt, precipitates metallic Bi: - 2 BiCl3 + 3 Zn = 3 ZnCl2 + | 2 Bi. Copper Copper forms two classes of salts ; viz., the cuprous com- pounds, in which Cu is monovalent; and the cupric compounds, in which Cu is divalent. The former are very unstable, being readily oxidized to the cupric compounds; they aredjjsoluble in water, but are soluble in halogen acids with the formation of colorless solutions. The cupric salts, when dissolvect'in water, yield blue or green solutions which have an acid reaction. Reactions of the Cupric Salts 1. Sodium or Potassium Hydroxide precipitates light blue Cu(OH)2, soluble in a large excess of the strong reagent with the formation of a blue liquid. The precipitate is changed on boiling to black CuO. In the presence of sufficient tartaric, citric, or arsenic acid, NaOH fails to precipitate Cu salts. 2. Ammonium Hydroxide, largely diluted and added cautiously to solutions of copper salts, precipitates a light blue basic salt, readily soluble in excess, producing a deep blue solution, due to the formation of a cupric ammonia salt: - CuSO4 + 2 NH4OH = |Cu(OH)2 + (NH4)2SO4; Cu(OH)2 + (NH4)2SO4 + 2 NH3 = Cu(NH3)4SO4 4- 2 H2O. Deep blue solution 28 QUALITATIVE CHEMICAL ANALYSIS When CuCl2 is used, the blue compound formed with an excess of NH3 is Cu(NH3)4C12. The sensitiveness of the test is 1 part in 25,000; it therefore is an exceedingly good test except for traces of the metal. 3. Hydrogen Sulphide produces in Cu solutions a black pre- cipitate of CuS which is insoluble in dilute acids and alkalies, insoluble in sodium sulphide, but somewhat soluble in ammonium sulphide, especially if the latter is yellow and hot; it is insoluble in hot dilute H2SO4 (distinction from Cd). When freshly pre- cipitated, CuS is easily soluble in KCN solution; it is also soluble in hot dilute HNO3 with the separation of S: - 3 CuS + 8 HNO3 = 3 Cu(NO3)2 + 4 H2O + f 2 NO + I 3 S.* When exposed to the air in moist condition, CuS oxidizes to CuSO4. 4. Potassium Cyanide produces a yellow precipitate of Cu(CN)2, which immediately decomposes into white cuprous cyanide (CuCN) and cyanogen; on adding an excess of the reagent, the precipitate dissolves, with the formation of a com- plex cym^^of K and Cu : - ^^CuCl2 + 2 KCN = | Cu(CN)2 + 2 KC1; 2 Cu(CN)2 = 12 CuCN +(CN)2; CuCN KCN = K3Cu(CN)4 (potassium cuprous cyanide). From solutions of K3Cu(CN)4, H2S does not precipitate CuS (distinction from Cd). If to the deep blue solution of Cu(NH3)4SO4 potassium cyanide is added, the color will be bleached, due to the formation of K3Cu(CN)4: - (1) 2 Cu(NH3)4SO4 + 4 KCN = |2Cu(CN)2+2 K2SO4+8 NH3; (2) 2 Cu(CN)2 = 12 CuCN + (CN^; (3) CuCN + 3 KCN = K3Cu(CN)4; (4) (CN)2 + 2 NH4OH = NH4CN + NH4CNO + H2O. 5. Potassium Ferrocyanide precipitates reddish brown cupric ferrocyanide: - 2 CuSO4 + K4Fe(CN)6 = | Cu2Fe(CN)6 + 2 K2SO4. * Resolvable into 2 steps (see PbS). THE METALS 29 It is insoluble in dilute acids, but soluble in NH4OH with the formation of a blue solution. The reaction with K4Fe(CN)6 is by far the most sensitive test for Cu; with very dilute solutions, it gives a reddish brown coloration. The sensitiveness of the test is i part in 200,000. 6. Potassium Iodide yields with solutions of cupric salts a yellowish white precipitate of cuprous iodide (Cu2I2); at the same time iodine is liberated and turns the solution brown : - 2 CuCl2 d- 4 k-I = Cu2!2 ~b 4 KC1 + I2' 7. Certain organic substances, like glucose, reduce copper solutions with the precipitation of red Cu2O. The test is best carried out by adding to the copper solution KNaC4H4O6 (Rochelle salt) and NaOH, the latter being added until the resulting solution assumes a deep blue color. On boiling and adding a small quantity of glucose, a red precipitate of Cu2O is obtained. The reaction consists essentially of 2 CuO + (reducing agent) = Cu2O + O. 8. Many metals reduce solutions of Cu salts to the metallic state; eg., Zn, Cd, Al, Fe. If an iron nail is immersed in a solution containing a Cu salt slightly acidified with HC1, a bright deposit of metallic Cu is formed on the iron. As dilute a solution as 1 part in 120,000 of water will give this test. Cu salts, when ignited in the bunsen flame, impart to it a green color which is intensified if the Cu solution contains HC1. Cadmium The cadmium salts are for the most part colorless. The nitrate, chloride, sulphate, bromide, iodide, and acetate are soluble in water. 1. Potassium or Sodium Hydroxide precipitates white Cd(OH)2, insoluble in excess. 2. Ammonium Hydroxide precipitates white Cd(OH)2, solu- ble in excess, forming complex ammonia salts: - CdCl2 + 2 NH4OH = | Cd(OH)2 + 2 NH4C1; Cd(OH)2 + 2 NH4C1 + 2 NH3 = Cd(NH3)4C12 + 2 H2O. QUALITATIVE CHEMICAL ANALYSIS 30 This ammonia salt, like the corresponding Cu salt, may be transposed to the double cyanide, KjCd^N)^ by a KCN solution. 3. Hydrogen Sulphide precipitates, from solutions not too strongly acid, yellow CdS, insoluble in cold dilute acids, alkali hydroxides, and (NH4)2S; and insoluble in KCN (distinction from Cu). It is soluble in hot dilute HNO3 with the separation of S (reaction similar to Pb, which see), and soluble in hot dilute H2SO4 (difference from Cu). From hot slightly acidulated solutions, H2S precipitates red CdS. 4. Potassium Cyanide yields a white precipitate, Cd(CN)2, soluble in excess, with the formation of K2Cd(CN)4, from which H2S precipitates CdS (distinction from Cu): - CdCl2 + 2 KCN = | Cd(CN)2 + 2 KC1; Cd(CN)2 + 2 KCN = K2Cd(CN)4; K2Cd(CN)4 + H2S = | CdS + 2 KCN + 2 HCN. Reactions of the Metals of Group II. Division B Arsenic Arsenic forms two series of compounds. Those in which it plays the role of a trivalent metal are known as the arsenious compounds; and those in which the metal is pentavalent are known as the arsenic compounds. The two oxides As2O3 and As2O6 are respectively the anhydrides of arsenious (H3AsO3) and arsenic (H3AsO4) acids. Reactions of the Arsenious Compounds The arsenious compounds may be considered as derived from As2O3. The latter, on treatment with boiling water, yields H3AsO3 ; the oxide also dissolves in alkalies with the formation of soluble alkali arsenites. All other arsenites are insoluble in water. As2O3 + 6 NaOH = 2 Na3AsO3 + 3 H2O. THE METALS 31 i. NaOH, NH4OH, NagCOg, HC1, and H2SO4 do not precipi- tate As from its compounds. 2. Hydrogen Sulphide. From neutral solutions of arsenites or from aqueous solutions of As2O3, H2S does not precipitate the sulphide but colors the solution yellow, which is due to the formation of colloidal As2S3; if, however, the solution is acidified with HC1, a yellow precipitate of As2S3 immediately forms: - 2 H3AsO3 + 3 H2S = As2S3 + 6 H2O. The precipitate is insoluble in concentrated HC1 (distinction from Sb and Sn), hence the presence of a large quantity of free acid does not interfere with the precipitation with H2S. Concen- trated HNO3, aqzta regia, or a mixture of concentrated HC1 and KC103 readily oxidizes As2S3 to arsenic acid (H3AsO4- soluble in water) with the separation of S. Ammonium sulphide and alkali hydroxides both dissolve As2S3 (distinction from the sul- phides of Division A). 3 As2S3 + io HNO3 + 4 H2O = 6 H3AsO4 + f io NO + 19 S. The action of aqua regia on As2S3 may be represented by the following equations: - As2S3 + io Cl = 2 AsC16 + | 3 S, 2 AsCL + 5 H2O = As2O5 + io HC1, As2O5 + 3 H2O=2 H3AsO4. Adding these 3 equations, with the elimination of substances appearing on both sides, we get as the final result: - As2S3 + 10 Cl+ 8 H2O = 2 H3AsO4 + 10 HC1 +13 S. The action of KC1O3 + HC1 may be represented thus : - KC1O3 + HC1 = HC103 + KC1; 2 HCIO3 = H2O + C1O2 + Cl. It is evident that this mixture behaves like aqua regia. The deportment of KOH is as follows: - As2S3 -I- 6 KOH = K3AsO3 + K3AsS3 + 3 H2O. 32 QUALITATIVE CHEMICAL ANALYSIS 3AsS3 may be regarded as derived from potassium arsenite (K3AsO3), by replacement of the oxygen by S; hence it is called potassium thioarsenite. If this mixture of thioarsenite and arsenite is acidified with HC1, As2S3 is reprecipitated : - K3AsO3 + K3AsS3 + 6 HC1 = 6 KC1 + 3 H2O + | As2S3. The equation for the solution of As2S3 in (NH4)2S is- As2S3 4- 3 (NH4)2S = 2 (NH4)3AsS3 (ammonium thioarsenite). If (NH^Sj is used, we get ammonium thioarsenate: - As2S8 + 3(NH4)2SiP = 2(NH4)8AsS4 + (3x- 5)S. The excess of S in the polysulphide oxidizes the thioarsenite to thioarsenate. The (3^ - 5)S does not separate out but dis- solves in the excess of (NH4)2S^ to form a higher polysulphide. If the solution of As2S3 in (NH4)2SJ. is acidified with HC1, the As is reprecipitated as yellow As2S6: - 2(NH4)3AsS4 4- 6 HC1 = | As2S6 +6NH4Cl + f3 H2S; but as this solution always contains an excess of (NH4)2S;r, the following reaction will also take place simultaneously: - (NH^S, + 2 HC1 = 2 NH4C1 + f H2S + |(x - i)S. 3. Silver Nitrate precipitates from neutral solutions of arse- nites yellow silver arsenite, Ag3AsO3 (distinction from arse- nates) : - K3AsO3 + 3 AgNO3 = | Ag3AsO3 4- 3 KNO3. The precipitate is easily soluble in acids and alkalies. 4. Magnesium mixture (solution of MgCl2 4- NH4C1 4- NH3) does not precipitate arsenites (difference from arsenates). 5. Iodine in solutions rendered alkaline with NaHCO3 readily oxidizes arsenites to arsenates : - Na3AsO34-I2 + 2 NaHCO3 = Na3AsO44-2 Nai4-H2O4-f 2 CO2. Special Tests for the Detection of Minute Amounts of Arsenic 1. The so-called Marsh Test for arsenic and its modifications is based on the fact that when arsenic compounds are introduced into a hydrogen generator (Zn + H2SO4), the As compound is reduced to gaseous arsine (AsH3), which escapes along with the excess of hydrogen. If the dried gases are led into a THE METALS 33 hard glass tube heated to redness, the arsine will be decomposed and a de- posit of metallic arsenic will form in the tube just beyond the part heated. If the gases are ignited, they burn with a bluish white flame; if a piece of cold porcelain is held in this flame, it will receive a black coating which is readily soluble in a solution of sodium hypobromite (NaBrO) (distinction from Sb). In making this test, it is necessary to run what is called a blank to make sure that the apparatus and the reagents employed in the test are arsenic-free. This is accomplished by carrying out the test precisely as de- scribed, except that no arsenic compound is added; and if the materials are arsenic-free, no mirror will be formed. Unless the results of this test are con- trolled by a blank, it cannot be considered trustworthy, for ordinary pure zinc and sulphuric acid, as well as glass tubing, usually contain sufficient arsenic to yield a positive result. For the detection of arsenic in wall paper as well as in reagents, medicinal preparations, and foods, the following tests are employed. The first two are modifications of the Marsh test; but all have the advantage over the Marsh test in that they are more rapid and require no special apparatus. 2. The Gutzeit Test for arsenic depends upon the fact that arsine colors a solution of silver nitrate (i : i) first yellow and finally black. To carry out this test, put in a test tube a few pieces of arsenic-free zinc and cover with about 3 cc. of dilute sulphuric acid. Place near the top of the tube a plug of cotton, stopper the tube with a loosely fitting cork which has been covered with 2 folds of filter paper moistened with AgNO3 solution (i : i), and allow to stand several minutes. If no darkening of the paper is produced, the blank is satisfactory. Now remove the plug and stopper, and introduce a very small amount of the solution or substance to be tested for arsenic. Replace the plug and stopper, and allow to stand again for several minutes. A yellow stain which quickly becomes black proves the presence of arsenic. Fre- quently, especially with solutions of silver nitrate less concentrated than i : i, the yellow stage is not seen. The reactions involved are the following: - (i)6 AgNO3 + AsH3 = | Ag3As • 3 AgNO3 + 3 HNO3; Yellow (2) Ag3As • 3 AgNO3 4- 3 H2O = H3AsO3 + 3 HNO3 4- f 6 Ag. Black Phosphine, stibine, and hydrogen sulphide interfere with this reaction ; the last can be guarded against by moistening the cotton plug with lead acetate solution. Instead of AgNO3, HgCl2 may be used; this yields a yellow stain. 3. Fleitmann Test. As the Gutzeit test does not distinguish between arsenic and antimony, Fleitmann devised a method by which the arsenic alone can be detected. This consists in generating the arsine in an alkaline solution. Method. Into a test tube place a small piece of arsenic-free Zn or Al; cover with a few cubic centimeters of NaOH and heat to boiling to start the reaction, then remove from flame, add the arsenic solution, cap with stopper covered with paper moistened with silver nitrate solution, and allow to stand; 34 QUALITATIVE CHEMICAL ANALYSIS a blackening proves the presence of arsenic. If the arsenic is in the pentad state, it should first be reduced with SO2 before applying this test. 4. Bettendorff's Test depends upon the fact that in a solution strongly acid with hydrochloric acid, stannous chloride reduces arsenic compounds to metallic arsenic: - 2 AsC13 + 3 SnCl2 = 3 SnCl4 + f 2 As. The test is carried out as follows : To 2 cc. of concentrated HC1 in a test tube, add 1 cc. of strong SnCl2 solution; then add a few drops of the solution to be tested for As and heat gently. A brown color or precipitate indicates the presence of arsenic. Antimony is not reduced under these conditions. The addition of a small piece of tin foil will have the effect of hastening the reaction, but must not be used if bismuth or antimony is known to be present. 5. Reinsch Test. If a solution containing arsenic, to which } of its bulk of concentrated HC1 has been added, is boiled with a strip of bright copper foil, the latter becomes coated with a gray deposit of copper arsenide (Cu5As2). If the foil is removed, washed, and dried between the folds of filter paper, and then slowly and carefully heated in a dry test tube, a white crystalline subli- mate of AsjOs will form; the latter can readily be recognized by examining with a lens or confirmed by dissolving in boiling water and applying the Fleitmann test. Sensitiveness of the Special Arsenic Tests Marsh test will detect 1 part of As in 200,000,000; Gutzeit test will detect 1 part of As in 10,000,000; Bettendorff test will detect 1 part of As in 7,000,000 ; and Reinsch test will detect 1 part of As in 40,000. Any solid substance containing arsenic, when mixed intimately with four times its weight of a mixture of KCN and Na2CO3, and heated in an ignition tube, will yield a black mirror of metallic arsenic on the cooler part of the tube. The sensitive- ness of this test is 1 part in 8000. Oxidation of arsenious to arsenic compounds can be accom- plished by the addition of iodine; the reaction proceeds best in a solution alkaline with sodium dicarbonate: - KH2AsO3 + NaHCO3 + I2 = Nai + KI + f CO2 4- H3AsO4. Reactions of Arsenic Compounds 1. NaOH, KOH, NH4OH or Na2CO3 produces no precipitate. 2. Hydrogen Sulphide. From cold, moderately acid solu- THE METALS 35 tions of arsenic acid, H2S precipitates, after some time, a mix- ture of As2S3 + S. The H2S first reduces the arsenic acid with the separation of sulphur, and then precipitates As2S3 readily from the reduced arsenic (-ic) solution: - (1) H3AsO4 + H2S = H2O + | S + H3AsO3 ; (2) 2 H3AsO3 + 3 H2S = J, As2S3 + 6 H2O. If the arsenic solution is heated, the reduction and precipita- tion are hastened. If the amount of HC1 is considerable, the stream of H2S rapid, and the solution cold, all of the As will be precipitated as As2S6 : - 2 H3AsO4 4- 5 H2S = 8 H2O + | As2S5. When, under the same conditions of acidity, the solution is heated and then treated with a rapid stream of H2S, a mixture of As2S3 and As2S5 is obtained. To rapidly precipitate the arsenic existing in the pentad state, it should be first reduced by adding sulphurous acid (H2SO3) to the cold solution and boiling till the excess of SO2 is expelled.* From the resulting reduced solution, H2S will rapidly precipi- tate the arsenic as As2S3: - H3AsO4 + H2SO3 = H2SO4 + H3AsO3f. As2S5 has the same solubility as As2S3; it dissolves in (NH4)2S with the formation of ammonium thioarsenate, from which HC1 reprecipitates As2S6 : - As2S5 + 3(NH4)2S = 2(NH4)3 AsS4; 2 (NH4)3 AsS4 + 6 HC1 = I As2S5 + 6 NH4C1 + f 3 H2S. 3. Silver Nitrate precipitates from strictly neutral solutions chocolate-colored Ag3AsO4 (distinction from arsenious and phos- phoric acids): - 3 AgNO3 + Na3AsO4 = | Ag3AsO4 + 3 NaNO3. * The reduction is best accomplished by adding the H2SO3 to the cold solution contained in a pressure bottle, stoppering, and heating in a boiling water bath for one hour. The bottle should be thoroughly cooled before opening. f If alkaline earth metals are present, they will be precipitated by the H2SO4 formed. 36 QUALITATIVE CHEMICAL ANALYSIS The precipitate is easily soluble in acids and in ammonium hydroxide. 4. Magnesia mixture yields (better when the solutions are cold) with neutral or ammoniacal solutions a white crystalline precipitate of NH4MgAsO4 (distinction from arsenious acid): - K3AsO4 4- MgCLj + NH4C1 = | NH4MgAsO4 + 3 KC1. The precipitate is soluble in acids, but insoluble in 2.5 per cent ammonia water. 5. Ammonium Molybdate, when added in great excess to a hot nitric acid solution of arsenic acid, yields a yellow precipi- tate of ammonium arsenomolybdate of variable composition (distinction from arsenious acid): - H3AsO4 + I2(NH4)2MoO4 + 21 HNO3 = 12 H2O + 21 NH4NO3 + |(NH4)3 AsO4- 12 MoO3. The presence of NH4NO3 favors this reaction. The precipitate is soluble in ammonium hydroxide, also in an excess of H3AsO4 ; hence the necessity of having an excess of the reagent. Phos- phates, or phosphoric acid, give a precipitate of similar appear- ance, hence they should be absent in making the test. 6. Reducing Agents, like FeSO4, H2SO3, Na^Og, when boiled with a solution of an arsenate strongly acid with HC1, reduce the substance from the arsenic to the arsenious state. Any arsenic compound, when treated with strong HC1 and distilled in a current of HC1 gas, will yield a distillate of AsC13. Potassium Iodide, when added to an acid solution of H3AsO4, will reduce it to H3AsO3 with the separation of iodine: - H3AsO4 + 2 HI = H3AsO3 + H2O +112. This test serves to detect arsenic acid in the presence of arse- nious acid. Special Tests for the Detection of Small Amounts. In the Marsh and Gutzeit tests, the reduction to AsH3 takes place less rapidly, due to the necessity of a preliminary reduction of H3AsO4 to H3AsO3: - H3AsO4 + H2 = H3AsO3 + H2O. THE METALS 37 With the Fleitmann test, preliminary reduction with H2SO3 is necessary. Antimony Like arsenic, antimony forms two series of compounds, viz., antimonic salts, in which antimony acts as pentavalent; and antimonious compounds, in which antimony behaves as a triva- lent element. Reactions of Antimonioits Compounds In carrying out the following reactions a hydrochloric acid solution of SbCl3 may be employed. i. Sodium Hydroxide, Ammonium Hydroxide, and Sodium Car- bonate each precipitates white antimonious hydroxide, Sb(OH)3, insoluble in ammonium hydroxide, but soluble in an excess of the fixed caustic alkalies and in a hot solution of alkali carbonate : - SbCl3 + 3 NH4OH = | Sb(OH)3 + 3 NH4C1; Sb(OH)3 + 3 NaOH = Na3SbO3 4- 3 H2O. Na3SbO3 readily hydrolyzes in contact with water, yielding sodium metantimonite, NaSbO2:- Na3SbO3 + H2O = 2 NaOH NaSbO2. The latter is further hydrolyzed by water, yielding finally white Sb2O3: - 2 NaSbO2 + H2O = 2 NaOH + Sb2O3. a a 1 A Q Sb2O3 is practically insoluble in water and in nitric acid, but readily dissolves in hot concentrated HC1 with the formation of SbCl3. Tartaric acid and the tartrates dissolve it in accordance with the following equations : - Sb2O3 + 2 H2C4H4O6 = 2 H(SbO)C4H4O6 + H2O ; Sb2O3 + 2 KHC4H4O6 = 2 K(SbO)C4H4O6 + H2O. Tartar emetic. In both of the resulting soluble compounds of antimony, the group (SbO), called antimonyl, acts as a monovalent radical similar to (NO), nitrosyl, in nitrosyl sulphuric acid, H(NO)SO4; 38 QUALITATIVE CHEMICAL ANALYSIS tartar emetic may, therefore, be called potassium antimonyl tar- trate. The solubility of antimony compounds in tartaric acid is of great analytical importance. 2. Water. If to a solution of SbCl3 containing not too much free acid a relatively large quantity of water is added, there forms a white precipitate of antimony oxychloride, SbOCl: - SbCl3 + H2OZJ 4 SbOCl + 2 HC1. As indicated, the reaction is reversible, too much HC1 having the effect of reversing the reaction; precipitation may be has- tened by heating. The precipitate is easily distinguished from the corresponding bismuth compound by its solubility in tartaric acid: - SbOCl + H2C4H4O6 = H(SbO)C4H4O6 4- HC1. The precipitate is also soluble in strong HC1 and can be changed directly to Sb2S3 by H2S : - 2 SbOCl + 3 H2S = 4 Sb2S3 + 2 H2O + 2 HC1. 3. Hydrogen Sulphide. From solutions not too strongly acid with HC1, H2S precipitates red Sb2S3: - 2 SbCL + 3 H2S±:} Sb2S3 + 6 HC1. The reversibility of the reaction indicates that a high concentra- tion of HC1 would prevent the precipitation, also that the pre- cipitate when formed would dissolve in strong HC1. It has been found by experiment that HC1 (i : i) readily dissolves Sb2S3 (distinction and method of separation from As). Sb2S3 is solu- ble in (NH4)2S with the formation of a thio salt: - Sb2S3 + 3(NH4)2S = 2 (NH4)3SbS3. Ammonium thioantimonite. If yellow (NH4)2S is used, the excess of sulphur in the latter oxidizes the thioantimonite formed in the last reaction to thio- antimonate : - Sb2Ss + 3(NH4\S4 = 2(NH4)gSbS4 + (3x- 5)S. The great similarity of the chemistry of the thio salts of anti- mony and arsenic is thus seen. THE METALS 39 If the solution of Sb2S3 in (NH4)2Sj.- that is, the solution containing ammonium thioantimonate and an excess of (NH4)2Sa. - is acidified with HC1, we get (as with the As compound) the higher sulphide precipitated and at the same time a separation of sulphur resulting from the decomposition of the excess of (NH4^:- 2(NH4)3SbS4 + 6 HC1 = | Sb2S5 + 6 NH4C1 + 3 H2S; (NH^ + 2 HC1 = 2 NH4C1 + f H2S +1(^ - i)S. The sulphide of antimony is also soluble in alkalies, from solu- tions of which HC1 reprecipitates Sb2S3. 4. Zn-Pt Couple. If a solution of antimony, acid with HC1, is poured into a dish containing a piece of platinum foil and a piece of zinc is added so that it touches the platinum, there will form on the platinum foil a black deposit or stain of metallic antimony. Even in dilute solutions this characteristic test can be applied. If arsenic is known to be present, the test should be carried out under a hood because of the possible formation of arsine. 5. If a solution of antimony, acid with HC1, is heated with a bright iron nail, all the antimony will be deposited in the form of black flocks (distinction from Sn). Antimonic Compounds 1. Water. If a solution of SbCl5, not too strongly acid with HC1, is diluted with a relatively large amount of water, a white precipitate of SbO2Cl is formed: - SbCl5 + 2 H2O | SbO2Cl + 4 HC1. If the dilution is very great, the SbO2Cl is further changed to antimonic acid: - SbO2Cl + 2 H2O = | H3SbO4 + HC1. As with SbOCl, tartaric acid prevents the precipitation of SbO2Cl. 40 QUALITATIVE CHEMICAL ANALYSIS 2. Hydrogen Sulphide. From moderately acid solutions of SbCl6, H2S precipitates orange-red Sb2S5: - 2 SbCl5 + 5 H2S =| Sb2S6 + io HC1. The precipitate is soluble in concentrated HC1 with the forma- tion of SbCl3 and evolution of H2S : - Sb2S5 4- 6 HC1 = 2 SbCl3 + f 3 H2S + 12 S. It possesses the same solubilities as the trisulphide, dissolving in (NH4)2S with the formation of ammonium thioantimonate, (NH4)3SbS4; on acidifying the latter with HC1, the pentasul- phide is reprecipitated: - 2(NH4)3SbS4 + 6 HC1 = | Sb2S5 + 6 NH4C1 + f 3 H2S. Sb2S5 also dissolves in caustic alkalies. 3. Potassium Iodide, when added to a HC1 solution of SbCl5, reduces it with the separation of iodine [distinction from Sb (-ous) compounds] : - SbCl5 + 2 KI = 2 SbCl3 + 2 KC1 + 112. Special Tests for Small Amounts of Antimony 1. The Marsh Test. This is carried out in the same manner as directed for arsenic. The stibine (SbH3) which forms is decomposed in the hot tube with the separation of metallic antimony in the form of a black mirror which is insoluble in a solution of NaBrO (distinction from As). SbH3 is further distinguished from AsH3 by the formation of black Ag3Sb when the former is passed into a solution of silver nitrate : - SbH3 + 3 AgNO3 = SbAgs + 3 HNO3. With AsH3 the black deposit is due to metallic silver (see p. 33). 2. Gutzeit Test. Same as for As, the blackening being due to AgsSb. 3. The Reinsch Test, when applied to antimony compounds, yields a black coating on the copper foil, which, when dried and ignited, gives a non-crystal- line sublimate of Sb2O3. The latter is insoluble in water but is soluble in a hot solution of KHC4H4O6, from which H2S precipitates red Sb2S3. Neither Fleitmann's nor Bettendorff's tests are applicable to antimony compounds. Tin The two oxides of tin, SnO and SnO2, correspond to two classes of salts known respectively as the stannous and stannic THE METALS 41 compounds. In the former, tin is divalent; in the latter, tetravalent. Reactions of Stannozis Salts The stannous salts are all colorless. Those which are solu- ble in water yield solutions which have an acid reaction; the solid salts as well as their solutions rapidly absorb oxygen from the air with the formation of stannic compounds. i. Sodium Hydroxide, Ammonium Hydroxide, or Sodium Car- bonate gives a white precipitate of Sn(OH)2, which is readily soluble in excess of NaOH with the formation of sodium stan- nite ; it is insoluble in excess of the other precipitants. SnCl2 + 2 NaOH = | Sn(OH)2 + 2 NaCl; Sn(OH)2 + 2 NaOH = Na2SnO2 + 2 H2O. The precipitate also dissolves in HC1. It possesses, like A1(OH)3, both acid and basic properties; substances of this character are called amphoteric substances. 2. Hydrogen Sulphide. From moderately acid solutions (con- taining not more than 3 per cent, of concentrated HC1) H2S throws down a brown precipitate of SnS : - SnCl2 + H2S = | SnS + 2 HC1. SnS is soluble in strong HC1 (distinction from As), nearly in- soluble in colorless (NH4)2S [distinction from the sulphides of As, Sb, and Sn (-ic)], but is soluble in hot (NH^S^. with the formation of ammonium thiostannate, from which HC1 precipi- tates yellow stannic sulphide, SnS2 : - SnS + (NH4)2S, = (NH4)2SnS3 + (x - 2)S; Brown. (NH4)2SnS3 + 2 HC1 = | SnS2 + 2 NH4C1 + f H2S. Yellow. 3. Mercuric Chloride, when added in excess to a solution of SnCl2, is reduced to white insoluble HgCl; the SnCl2 is oxidized at the same time to SnCl4 : - 2 HgCI2 + SnCl2 = 12 HgCl + SnCl4. 42 QUALITATIVE CHEMICAL ANALYSIS If, however, the SnCl2 is present in excess, the HgCl first formed is further reduced to metallic mercury (gray or black): - 2 HgCl + SnC^ = | 2 Hg + SnCl4. Black. As this reaction is essentially one of reduction, it is important that the HgCl2 solution contain no oxidizing agent. Stannic compounds do not give this reaction; it thus serves to distin- guish stannous from stannic compounds. 4. Bismuth salts are reduced by an alkaline solution of stan- nous salts with the precipitation of black metallic bismuth (see under Bismuth, reaction 4). 5. Metallic Zinc. When metallic zinc is introduced into a hydrochloric acid solution of either SnC^ or SnCl4, metallic tin is precipitated on the zinc in the form of a gray spongy mass. As the deposited tin is readily soluble in strong HC1, care must be taken not to have the solution too strongly acid. SnCl4 4- 2 Zn = 2 ZnCl2 4-1 Sn; SnClg 4- Zn = ZnCl2 4-1 Sn. Stannic Compounds With the exception of the sulphide all the stannic compounds are either colorless or white. They are generally obtained from stannous salts by oxidation; thus, a solution of SnCl4 for the following tests may be prepared by warming a rather strong HC1 solution of SnCl2, with KC1O3 added in small portions, until the solution becomes yellow, and then boiling off the excess of chlorine; after a slight dilution with water, the solution is ready for use. Reactions of a Solution of SnCl± 1. Sodium Hydroxide. Ammonium Hydroxide, as well as Sodium Carbonate, yield a white precipitate of stannic hydroxide, Sn(OH)4, which, on drying, becomes H2SnO3. The precipitate is soluble in excess of NaOH or Na2CO3 with the formation of sodium stannate of variable composition. THE METALS 43 2. Hydrogen Sulphide. From moderately acid solutions not exceeding 3 per cent, concentrated HC1, H2S precipitates yellow SnS2 : - SnCl4 + 2 H2S = | SnS2 + 4 HC1. SnS2 is readily soluble in HC1 (1:1), hence the necessity for having the solution not too strongly acid. The disulphide of tin readily dissolves in colorless (NH4)2S (distinction from SnS) with the formation of ammonium thiostannate, from which HC1 precipitates SnS2 (yellow): - SnS2 + (NH^S = (NH4)2SnS3; (NH4)2SnS3 + 2 HC1 = 2 NH4C1 + f H2S + | SnS2. On strong ignition in the air, SnS2 is quantitatively converted into SnO2. 3. Mercuric Chloride gives no precipitate with stannic salts (distinction from stannous). 4. Hydrochloric or sulphuric acid does not precipitate stan- nic salts from solutions that are moderately concentrated (dis- tinction from metastannic compounds); when, however, the solutions are diluted and boiled, a precipitate of Sn(OH)4 is obtained: - SnCl4 + 4 H2O^ | Sn(OH)4 + 4 HC1. 5. Potassium or Sodium Sulphate. From cold solutions no precipitate is obtained with these reagents (distinction from metastannic compounds); but, on boiling, a precipitate of Sn(OH)4 is obtained. Reactions 4 and 5 can be explained on the assumption that the oxy-salts of stannic tin first form, but being unstable in dilute solutions are decomposed into stannic hydroxide, thus : - (1) SnCl4 + 2 H2SO4 = Sn(SO4)2 + 4 HC1; (2) Sn (SO4)2 + 4 H2O = 2 H2SO4 + | Sn (OH)4. Similarly, with K2SO4 we get SnCl4 + 2 1^504 = Sn(SO4)2 + 4 KC1, the Sn(SO4)2 being then hydrolyzed as shown above in equa- tion (2). 44 QUALITATIVE CHEMICAL ANALYSIS Metastannic Compounds There are two forms of stannic compounds; viz., the normal and the metastannic forms. The normal may be considered as derived from stannic hydroxide, formed by the action of NaOH on SnCl4. It is readily soluble in acids. The metastannic compounds are derived from metastannic acid, a white substance obtained by the action of hot dilute HNO3 on metallic tin; it has the same empirical formula as the partially dehydrated Sn(OH)4, i.e., H2SnO3, but differs from it in being insoluble in acids. When boiled for a short time with concentrated HC1, a compound of the formula SnsO5Cl2(OH)8 forms which, though quite insoluble in HC1, is readily soluble in water. From the fact that this and similar compounds may be formed from metastannic acid, the formula Sn5O5(OH)10 or 5 (H2SnO3) has been assigned to it. Stannic hydroxide, when dried over concentrated H2SO4, has the formula H2SnO3; metastannic acid is thus seen to be a polymer of stannic hydroxide. Reactions of Metastannic Chloride, SnhOfJlfOH\ 1. HC1 precipitates Sn6O6Cl4(OH)6, 4 H2O. 2. Prolonged boiling with water causes the precipitation of all the tin as metastannic acid, insoluble in dilute acids. 3. H2SO4, K2SO4, or Na2SO4, precipitates a white substance which changes, on washing with water, to metastannic acid (distinction from stannic chloride). 4. KOH precipitates metastannic acid, which is converted by an excess of the concentrated reagent to a potassium salt; the latter is soluble in water and in dilute KOH solution. 5. NH4OH precipitates metastannic acid. 6. H2S yields the same precipitate as with SnCl4 solutions. Solutions of stannic compounds are converted into the meta- stannic form by diluting and boiling: - 5 SnCl4+ 13 H2Oz£ 18 HC1 + Sn5O5Cl2(OH)8. Conversely, metastannic compounds are converted into the stannic form by boiling with concentrated HC1 or concentrated KOH. Action of H2S Besides its action as a group reagent for precipitating the metals of Group IL, H2S also acts as a reducing agent. Should, therefore, an oxidizing agent be present in the solution subjected to the action of H2S, it will be reduced, the HoS being oxidized at the same time to elementary sulphur, which sepa- THE METALS 45 rates in a finely divided state, and, in some cases, partly to sulphuric acid. Among the oxidizing agents likely to be present in the filtrate from Group I. are FeCl3, K2Cr2O7, KMnO4, HNO3, and aqua regia. The reduction of the first three substances is accompanied by a change in color of the solution, thus frequently affording an indication of their presence. For example, a solution containing ferric salts undergoes a change from a yellow or reddish yellow to a colorless condition - 2 FeCl3 + H2S = 2 FeCl2 + 2 HC1 + S; a K2Cr2O7 solution changes by the action of H2S from reddish orange to green - K2Cr2O7 + 3 H2S + 8 HC1 = 2 CrCl3 + 2 KC1 + 7 H2O + 3 S ; while a so- lution containing KMnO4 changes from a solution with purple tint to one that is colorless: - 2 KMnO4 + 5 H2S + 6 HC1 = 2 MnCl2 + 2 KC1 + 8 H2O + f 5 S. Should the concentration of nitric acid be large, it will oxidize the H2S with separation of sulphur and partly with the formation of H2SO4: - 2 HNO3 + 3 H2S = 4 H2O +^2NO+ f 3 S ; 3 H2S + 8 HNO3 = 3 H2SO4 + 18 NO + 4 H2O. Aqua regia will oxidize the H2S in accordance with the equation - Cl2 + H2S = 2 HC1 + S. It will be observed that in every case the presence of an oxidizing agent causes the decomposition of the H2S with the separation of S ; with a large amount of oxidizing agent present, the amount of S and H2SO4 will be formed in quantity sufficient to seriously interfere with the analysis. A large quantity of sulphur is undesirable because it complicates and masks the results ; and the presence of H2SO4 will have the effect of precipitating the alkaline earths along with the metals of the second group, so that where the amount of oxidixing agent is large it is advisable to eliminate the latter before pass- ing in H2S. If, from the purple or orange-red color of the solution, KMnO4 or K2Cr2O7 is suspected, reduction may be readily effected by acidifying with HC1, adding alcohol, and boiling. For most purposes, it will only be neces- sary to consider the presence of a large excess of HNO3 or aqua regia, be- cause of their extensive use as solvents. An excess of HNO3 is removed by evaporating the solution to about 1 cc., adding 3 cc. of concentrated HC1, and evaporating nearly to dryness. It can then be taken up with the aid of HC1 and hot water. An excess of aqua regia is disposed of by boiling the solution down to a small bulk, adding concentrated HC1 and again evaporating to 1 cc.; it is then diluted with water and a few drops of HC1. H2S as a Precipitating Agent. In precipitating the second group sulphides with HoS, it is exceedingly important that the solution have a certain approxi- mately definite acidity. If too great a concentration of acid is present, com- plete precipitation will be impossible, owing to the appreciable solubility of some of the sulphides in moderately strong HC1 (notably those of Pb, Cd, 46 QUALITATIVE CHEMICAL ANALYSIS and Sn). On the other hand, if the acid concentration is too small, certain metals of the third group will also precipitate, as Zn, Ni, and Co. By experiment it has been found that a concentration of 2.5 cc. of HC1 (sp. gr. 1.2) in a volume of 100 cc. affords a satisfactory separation of the second and third groups by H2S. Precipitation with H2S The filtrate (1) from Group I. is made slightly alkaline with NH4OH and then just acid with dilute HC1 (2); 2.5 cc. of concentrated HC1 are then added, the solution is heated nearly to boiling and is treated with a rapid stream of H2S for a few minutes. Without filtering (3), add enough cold water to make the total volume 100 cc., cool to room temperature, and pass in H2S again until precipitation is complete. Filter, dilute the filtrate with | its volume of water, and treat again with H2S; now filter off any precipitate formed (4). The final filtrate should not give a precipitate when treated with H2S. The beaker containing the filtrate should be labeled Groups III.-V., at once placed on a wire gauze and boiled until all the H2S is expelled (5). The precipitate may consist of HgS, PbS, Bi2S3, CuS, CdS, As2S3, Sb2S3, SnS and S; it should be washed with water con- taining H2S and about 5 per cent. NH4NO3 (6) until the wash- ings are only faintly acid. Reject the washings. NOTES 1. A preliminary test for the second group should first be made on a small portion of the filtrate, in order to determine whether or not Group II. is present. If present, the entire filtrate should be treated with H2S in accord- ance with directions ; if absent, the introduction and removal of H2S will thus be avoided. In that case pass to Scheme III. 2. The solution must be thoroughly stirred with a glass rod during the addition of the ammonium hydroxide and acid, and the acidity or alkalinity of the mixture determined by means of litmus paper and not by the quantity of the reagent added. No attention need be given to precipitates which form, because these are either finally dissolved or converted by H2S into sulphides. If a large excess of acid is known to be present, it should be removed by evaporation and the solution then brought to the proper condition of acidity as directed in the procedure above. THE METALS 47 3. The solution is not diluted at once to 100 cc. because of the possible presence of arsenic, which comes down best in a hot strongly acid solution. It is a good plan to mark with a label the level at which the beaker will hold 100 cc.; the dilution can then be made without resorting to a measuring cylinder. 4. The color of the H2S precipitate sometimes affords an indication of the metals present. If black, it may be due to Pb, Cu, Hg, or to all of them ; if yellow, to Cd, Sn(-ic), or As ; if orange, to Sb. A yellow precipitate which is insoluble in (NH4)2Sj, cannot be anything other than Cd ; on the other hand, if the yellow precipitate dissolves completely in (NH4)2S^, it must be either the sulphide of As or Sn(-ic), or both. 5. As H2S in solution readily oxidizes in contact with air to S and H2SO4, and as the latter will precipitate alkaline earths, the necessity for immediately expelling the H2S is apparent. Keeping a glass rod in the beaker during the boiling will facilitate the removal of H2S by preventing dangerous bumping, with a consequent loss of liquid. The completeness of the expulsion of the H2S may be determined by holding a piece of filter paper moistened with lead acetate in the escaping vapor. 6. The precipitate is washed with H2S water to prevent the oxidation of the sulphides to sulphates, which, with the exception of PbSO4, are all soluble in water. NH4NO3 is added to prevent the precipitate from going into the colloidal state and thus passing through the filter. The separation of Group II. into two divisions is based on the difference of behavior of the sulphides towards (NH4)2Sj.. Division II. A (the copper group) includes those sulphides which are insoluble in (NH^S*; these are Hg, Pb, Bi, Cu, and Cd. Division II. B (the tin group) includes those sulphides which are soluble in (NH^S^ with the formation of thio salts. The separation, however, is not sharp, which is due to the slight solubility of CuS in (NH4)3Sr CuS is practically insoluble in Na2Sa., but the latter quite appreciably dissolves HgS. For all practical purposes, the separation with (NH4)2Sa. is suffi- ciently complete. In the analysis of the H2S precipitate it is sometimes just as well to assume the presence of both divisions and to treat the well-washed precipitate at once with (NH4)2S.„ as directed in paragraph 3. In most cases, however, it is desirable to make the following preliminary tests. 48 QUALITATIVE CHEMICAL ANALYSIS Preliminary examination of the H^S precipitate to determine the presence of- (i) Division A (copper group). By means of a glass spatula, put a very small amount of the precipitate into a small evapo- rating dish; add about 15 drops of (NH4)2SZ diluted with an equal quantity of water. Break up the precipitate with a glass rod, and warm gently with constant stirring for a minute {do not boil}. If the precipitate completely dissolves, the Cu group is absent and the main precipitate is analyzed for the Sn group only (Scheme II. B). If a residue remains (1), the copper group is present. (2) Division B (tin group). Twice pass about 10 drops of (NH4)2SZ, diluted with an equal volume of water, through the precipitate, which has been thoroughly washed and from which most of the water has drained; catch the liquid which passes through in a test tube (2) and just acidify with dilute HC1. In another test tube,/?^ acidify with dilute HC1 an equal portion of (NH4)2SZ, diluted as above, and compare the results. A colored precipitate in the first tube which is different from that of the second proves the presence of the Sn group (3). If the two tubes present the same appearance after acidification, the absence of the Sn group is proved. Separation of the Divisions of Group II. If both divisions are shown to be present by the above tests, the entire precipitate is treated in a small beaker (50 cc. capacity) (4) with 10 cc. of (NH4)2SZ diluted with 5 cc. of water; the mixture is thoroughly stirred, warmed for several minutes, and filtered. The residue on the filter may consist of HgS, PbS, Bi2S3, CuS, CdS and S. The filtrate may contain the thio salts of As, Sb, and Sn. A warm mixture of 2 cc. of (NH^S* and an equal volume of water is poured on the residue (5), and the liquid which passes through is united with the filtrate containing the thio salts of Division B, which is then analyzed according to Scheme II. B. The residue is washed twice with hot water and is analyzed according to Scheme II. A. The washings are rejected. THE METALS 49 NOTES 1. A residue of S is not to be taken as indicating the presence of the Cu group. 2. If the precipitate has not been washed free of acid, the latter will de- compose the (NH4)2SX, with the separation of S. If considerable acid is present, all of the (NH4)2SX will be decomposed, with the result that the liquid which drains through the filter will be colorless and will fail to react with HC1. To remedy this, add more (NH4)2SX, preferably diluted with NH4OH, until a colored filtrate is obtained. 3. Because of the slight solubility of CuS in (NH4)2SX, a liver-colored pre- cipitate is sometimes obtained in acidifying the (NH4)2SX solution. 4. If the quantity of the precipitate is small, it may be treated on the filter with (NH4)2SX, using small portions at a time and allowing each portion to drain through before adding a fresh portion; in this way a maximum of ex- traction with the minimum amount of solvent is secured. If water is used in transferring the precipitate to a beaker, it must be carefully poured off before adding the (NH4)2SX. 5. This further treatment of the residue with dilute (NH4)2SX is given to insure the complete extraction of the As, Sb, and Sn. Outline of the Separation of the Metals of Group II. A Of the sulphides of Group II. A, only HgS is insoluble in hot dilute HNO3; the addition of this reagent to the mixed sul- phides will therefore result in the solution of all the sulphides as nitrates, with the exception of HgS, which will be left behind as an insoluble residue. If, now, to the filtrate containing the nitrates of Pb, Bi, Cu, and Cd we add H2SO4, we should expect the Pb to precipitate as PbSO4; for Pb is the only metal of this group whose sulphate is insoluble in water. It is, however, appreciably soluble in nitric acid, and as the filtrate contains an excess of this acid, its removal is necessary if complete precipi- tation of the Pb as sulphate is desired. This is accomplished by adding concentrated H2SO4 and evaporating the solution until SO3 fumes appear ; it is then diluted with water and the PbSO4 is filtered off. The filtrate will now contain Bi, Cu, and Cd as sulphates, and an excess of H2SO4. A glance at the reactions of the salts of these metals shows that NH4OH pre- cipitates all three of them, but that Bi is the only one whose hydroxide is insoluble in excess. If, therefore, an excess of 50 QUALITATIVE CHEMICAL ANALYSIS NH4OH is added to the filtrate from the PbSO4, Bi(OH)3 will be precipitated and may be separated from the remaining two metals by filtration. The ammoniacal filtrate will be colored blue if Cu is present. Cu and Cd may be separated by taking advantage of the difference of behavior of their double cyanides towards H2S. The double cyanides are formed by adding KCN to the ammoniacal solution until the blue color is discharged; the passage of H2S into this solution precipitates only the Cd as yellow sulphide. SCHEME II. A The residue from the (NH4)2S^ treatment may consist of HgS, PbS, Bi2S3, CuS, and CdS. Wash the precipitate into a small beaker, pour off the water, add 15-30 cc. of dilute HNO3, and heat with constant stirring (1) ; boil for one and a half minutes, and filter. Residue may be HgS (black) or 2 HgS • Hg(NO3)2 (white) + S (2). Wash with water. Transfer ppt. to a small evap. dish; add 4 cc. dil. aqua regia, and boil till all but S dissolves. Boil a little longer to expel Cl, dilute with 1 cc. of water, and filter. To fil- trate add a few drops of SnCl2; a white ppt., turning gray to black, proves the pres- ence of Hg. Filtrate may contain Pb(NO3)2, Bi(NO3)3, Cu(NO3)2, Cd(NO3)2 + excess of HNO3. Transfer to an evap. dish, add 5 cc. of cone. H2SO4, and evaporate under a hood until SO3 fumes are given off (3) ; cool and cautiously pour the contents of dish into a beaker containing 25 cc. of water. Rinse what remains in the evap. dish with a little water into the same beaker; stir, allow to settle, and filter. Residue is PbSO4 (4). Wash once with water, and treat the ppt. on the filter with a boiling solution of NH4C2H3O2. Catch filtrate in a test tube, acidify with acetic acid, and add K2CrO4. A yellow ppt. con- firms the presence of Pb. Filtrate contains Bi2(SO4)3, CuSO4, CdSO4 + excess of H2SO4. AddNH4OH to alkaline reaction and then in slight excess. A deep blue coloration proves the presence of Cu (5). Allow any ppt. which forms at the same time to settle, and filter. Residue is Bi(OH)3 (6). Wash with water; dissolve ppt. on the filter with about 5 drops of cone. HC1, and catch drops in a clean beaker. Add 50 cc. of water, warm, and allow to stand for a few minutes; a white ppt. or cloudi- ness is BiOCl (7). Filtrate contains Cu(NH3)4SO4, Cd(NH3)4SO4 + excess of NH4OH. Add KCN until the solution is decolor- ized, and pass in H2S. A yellow ppt. is CdS (8). THE METALS 51 NOTES ON SCHEME II. A 1. Stirring while heating with dilute HNO3 is important because of th6 tendency of the S, which separates in a plastic condition, to inclose portions of the sulphides, with the result that the latter are protected from the solvent action of the acid. 2. Besides the substances mentioned, the residue may consist of a little PbSO4 resulting from the oxidation of PbS by the use of strong HNO3 or the long-continued action of the dilute acid. But, as PbSO4 is somewhat soluble in hot dilute HNO3, enough Pb passes into solution for its detection in the next operation. A black residue is not to be taken as proof of the presence of Hg; it may be sulphur mechanically inclosing small quantities of the black sulphides, as PbS, CuS, and Bi2S3. Consequently the residue after the HNO3 treatment, whether it is black or white, must be tested for Hg. 3. The object of evaporating the solution until SO3 fumes result is to com- pletely remove the HNO3, which has a solvent action on the PbSO4. As the boiling point of concentrated HNO3 is 1200 C. and the fuming point of H2SO4 is 2500 C., it is evident that all the HNO3 will have been removed when the solution is boiled until dense white fumes of SO3 are given off. The student should not look for SO3 fumes until the bulk of his solution has been reduced to about 3 cc. If he is unable to recognize these fumes with certainty, he should show his results to his instructor before proceeding with the next step. Unless this operation is properly conducted, it will not be possible to completely separate the Pb, with the result that the tests for Bi and Cd will be interfered with. 4. PbSO4 is a heavy white powder; a white coarsely crystalline precipi- tate (a basic sulphate of Bi) sometimes separates, hence the necessity of making a confirmatory test for Pb. Where the amount of PbSO4 is very small, it can be made distinctly visible by collecting it in the center by means of a rotary motion imparted to the beaker. 5. The deep blue solution produced by an excess of ammonium hydroxide is sufficiently sensitive for the detection of Cu. A much more delicate test consists in acidifying a portion of the solution with acetic acid and adding K4Fe(CN)6; a brown precipitate or coloration proves the presence of Cu. 6. The formation of a white precipitate at this point is not proof of the presence of Bi; for if all the Pb had not been removed in the previous opera- tion it will precipitate here. A confirmatory test for Bi must therefore always be made. 7. It frequently happens that a precipitate or cloudiness is not obtained' even when Bi is present. This is due to the presence of an excess of acid, which reverses the direction of the reaction: BiCls+H2O ^2 BiOCl + 2 HC1. QUALITATIVE CHEMICAL ANALYSIS 52 This may be remedied by adding NH4OH drop by drop to neutralize the excess of acid, care being taken, however, to keep the solution acid. Another confirmatory test for Bi consists in pouring on the thoroughly washed white precipitate of Bi(OH)s a solution of Na2SnO2 (prepared by add- ing NaOH to I cc. of SnCl2 till the precipitate which first forms dissolves), when if Bi is present it will be blackened due to the precipitation of metallic Bi. 8. The presence of Hg or Pb in the solution to be tested for Cd will yield a black precipitate when H2S is passed into the solution. A little Hg may find its way into the solution if the original H2S precipitate had not been thoroughly washed free from HC1 or chlorides; the latter, on boiling with HNO3, will yield aqua regia, which dissolves some of the HgS. The presence of Pb is due to its incomplete removal in the previous operation. A black precipitate obtained with H2S, in testing for Cd, may be examined for this metal by filtering the precipitate, thoroughly washing it with hot water, and finally treating it with hot sulphuric acid (1:6) on the filter; if the fil- trate is now largely diluted with water and treated with H2S, a yellow precipi- tate of CdS will be formed if Cd is present. Outline of the Analysis of Group II. B (the Tin Group) The insolubility of As2S3 in hot concentrated HC1 is made use of to separate it from the sulphides of Sb and Sn; the latter two dissolve in this acid, forming the corresponding chlorides with an evolution of H2S. The residue, after filtra- tion, is taken into solution with concentrated HNO3, and the presence of arsenic, now in the form of arsenic acid, is con- firmed by its precipitation with AgNO3 in a neutral solution as Ag3AsO4, and also by the formation of yellow ammonium arsenomolybdate with ammonium molybdate solution. When the As2S3 is removed by filtration, the filtrate will contain the Sb and Sn as chlorides and an excess of H2S. As the tests for Sb and Sn can be made in the presence of each other, it is needless to effect their separation. The filtrate, after boiling to expel the H2S, is divided into two portions. An iron nail is placed in the first portion and the liquid warmed. The iron acts on the acid present, liberating hydro- gen, which reduces SnCl4 to SnCl2 and at the same time pre- cipitates the Sb in the metallic state. The decanted solution, containing SnCl2, is then tested with HgCl2 solution. The second portion is tested for Sb by the galvanic action of THE METALS 53 the electric couple (Pt and Zn), which causes the antimony to precipitate as a black deposit on the platinum. The filtrate obtained after treating the H2S precipitate with (NH^ will contain the thio salts of the metals of this group; viz., (NH4)3AsS4, (NH4)3SbS4, (NH4)2SnS3 and excess (NH4)2Sr Just acidify with dilute HC1 (i) and filter ; reject the filtrate. The residue will consist of As2S5, Sb2S5, SnS2, (CuS?), and S. Transfer the precipitate to a small beaker, add io cc. of con- centrated HC1, and heat gently to boiling (2), with constant stir- ring, for about 5 minutes. Dilute with an equal volume of water and filter. SCHEME II. B (TIN GROUP) Residue is As2Ss + S (3). Wash with hot water until the washings give only a faint re- action with AgNO3. Reject washings. Trans- fer ppt. to an evap. dish (if water is used in transferring ppt., pour it off after ppt. settles). Add 3-5 cc. of cone. HNO3, stir, and heat gently until no more brown fumes are given off. Dilute with 3 cc. of water and filter. Divide the filtrate into two unequal portions. To 4 contained in a test tube add 4 cc. of AgNO3 (4) ; filter if a ppt. forms. To the filtrate in a test tube, very carefully add NH4OH one drop at a time and gently agitate the upper portion of the liquid (5). At the neutral zone a chocolate-colored ppt. of Ag3AsO4 forms. Add the (0 portion to 10 cc. of ammonium molybdate solution, heat gently, but do not boil (6). A yellow ppt., appearing after some time, confirms the presence of arsenic. Filtrate may contain SbCl3, SnCl4 + excess HC1 + H2S. Boil until the H2S is com- pletely expelled (7). Divide the solution into 2 portions. In the 1st portion, test for Sn by warming with an iron nail for 3 minutes. Filter rapidly (8) into a test tube containing 2-3 drops of HgCl2; a white ppt., which may turn gray or black, proves the presence of Sn. In the 2nd portion, test for Sb. Pour it into a small evap. dish containing Pt foil in contact with a piece of Sn or Zn. A black stain on the Pt (9), which is insoluble in NaBrO, is Sb. NOTES TO SCHEME H. B 1. An excess of acid is to be avoided because of the solubility of SnS2 in even moderately dilute acid. 2. The mixture must not be boiled, else some As2S5 is apt to go into solution. 3. The separation of As from Sb and Sn sulphides by the use of hot con- centrated HC1 is not always very sharp. With mixtures consisting of a small 54 QUALITATIVE CHEMICAL ANALYSIS amount of As and a relatively large quantity of Sb, the insoluble residue may contain enough Sb2S5 to give it a red color; therefore a red residue must always be examined for As. 4. If all the HC1 had not been removed by washing, the addition of AgNOs will yield a white precipitate of AgCl; if this precipitate is large in amount, more AgNO3 should be added before filtering to insure an excess of the latter in the filtrate. 5. The test for arsenic with AgNOs depends upon the formation of Ag3AsO4, which only forms in a strictly neutral solution; if too much NH4OH is added, precipitation will fail, because of the ready solubility of Ag3AsO4 in alkalies as well as in acids. To remedy this, transfer the solution to a small beaker, add a piece of litmus paper and carefully add acetic acid drop by drop with constant stirring till the solution becomes neutral. 6. The ammonium molybdate solution must not be boiled, else white MoO3 may separate; the latter must not be mistaken for the yellow precipitate. The mixture, after heating, should be placed in a beaker of hot water and allowed to remain in the latter for 5-10 minutes. 7. Test escaping vapors with lead acetate paper. 8. By this procedure the SnCl4 is reduced to SnCl2; as the latter rapidly oxidizes on exposure to air, particularly in a hot solution, the necessity of rapidly filtering into a test tube containing the reagent is apparent. 9. If much Cu and little or no Sb are present, a dark red stain, easily dis- tinguished from a black stain, will be produced. The cleaning of the Pt foil is easily accomplished by first washing it with water and then pouring on it concentrated HNO3. Reactions of Metals of Group III The metals of this group are distinguished from those of the first and second groups by the fact that they are not precipitated by H2S from solutions containing 2.5 per cent, of hydrochloric acid, sp. gr. 1.2. They are associated together in one group be- cause of their common property of being completely precipitated by (NH4)2S in solutions alkaline with ammonium hydroxide in the presence of NH4C1 (distinction from Groups IV. and V.). Aluminum The aluminum salts are nearly all colorless; the salts of the halogen acids, and the nitrate, sulphate, and acetate, are soluble in water. THE METALS 55 i. Ammonium Hydroxide throws down a white gelatinous precipitate of A1(OH)3, slightly soluble in excess; on boiling, the dissolved hydroxide is reprecipitated. The presence of NH4C1 diminishes the solvent action of ammonia on the hydrox- ide ; hence, to completely precipitate aluminum by ammonia, the latter should be added only in slight excess and the mixture boiled until the liquid has but a faint odor of the reagent. When freshly precipitated, A1(OH)3 is readily soluble in A1C13 + 3 NH4OH A1(OH)3 + 3 NH4C1; A1(OH)3 + 3 HC1 = A1C13 + 3 H2O. Aluminum hydroxide is also soluble in caustic alkalies (see 2). 2. Potassium or Sodium Hydroxide precipitates A1(OH)3, soluble in excess with the formation of alkali aluminate: - Al(OH)g + 3 NaOH = Na3AlO3 4- 3 H2O. On carefully neutralizing the alkaline solution of sodium aluminate with hydrochloric acid, A1(OH3) is reprecipitated: - (a) Na3A103+ 3 HC1=3 NaCl + | A1(OH)3. If an excess of acid is added, the precipitate which first forms is dissolved and we obtain: - (£) A1(OH)3 + 3 HC1 = A1C13 + 3 H2O. It is evident that if the original sodium aluminate solution is at once acidified with HC1, we shall get the net result of (a) and (d). Adding (a) and (£), and eliminating A1(OH)3, which ap- pears on opposite sides, we get: - Na3AlO3 + 6 HC1 = A1CL + 3 NaCl + 3 H„O. If, now, we heat this solution to boiling and add ammonium hydroxide in faint excess, all the aluminum will be precipitated as the hydroxide : - A1C13 + 3 NH4OH = I Al(OH)g + 3 NH4CL While the addition of a large excess of solid NH4C1 will have the effect of precipitating A1(OH)3 from the aluminate, the above process, viz., that of acidifying first with HC1 and then 56 QUALITATIVE CHEMICAL ANALYSIS rendering the resulting solution faintly alkaline with ammonium hydroxide, is the more common procedure and the one to be generally recommended. It is important to remember that the direct addition of am- monium hydroxide to a sodium aluminate solution will not pre- cipitate A1(OH)3. 3. Ammonium Sulphide precipitates A1(OH)3 and not the sulphide. A12S3 may be prepared in the dry way, but, on bring- ing it in contact with water, it at once hydrolyzes with the for- mation of the hydroxide and the evolution of H2S. The action of (NH4)23 on solutions of aluminum salts may be represented as taking place in two steps : - (a) 2 A1C13 + 3(NH4)2S = A12S3 + 6 NH4C1; (0 A12S3 + 6 H2O = | 2 A1(OH)3 + 3 Adding (a) and (£), and eliminating A12S3, we get as the equa- tion for the final result: - 2 AlCl3 + 3(NH4)2S + 6 H2O = 12 Al(OH)3 + 6 NH4Cl + 3 H2O. 4. Sodium Carbonate also precipitates A1(OH)3. In the presence of non-volatile organic acids, as tartaric, citric, and malic acids, as well as certain organic matter containing (OH) groups, as sugars and starch, ammonium hydroxide, sodium carbonate, and ammonium sulphide fail to precipitate aluminum salts. 5. Alkali Acetate. If an excess of alkali acetate is added to a slightly acid or neutral solution of an aluminum salt, and the mixture is largely diluted with water and boiled, a bulky pre- cipitate of basic aluminum acetate will be thrown down : - A1C13 + 3 NaC2H3O2 = A1(C2H3O2)3 + 3 NaCl; A1(C2H3O2)3 + H2O I A1(OH)(C2H3O2)j + HC2H3O2. The reagent is, in fact, hot water, which hydrolyzes the weak salt A1(C2H3O2)3. Cooling the solution, or the presence of an excess of acetic acid, will have the effect of reversing the reaction, with the result that some of the precipitate will dissolve. THE METALS 57 6. Disodium Phosphate yields with solutions of aluminum salts a gelatinous precipitate of A1PO4, soluble in HC1 and NaOH, but insoluble in acetic acid : - (i) Na2HPO4 + A1C13 = | A1PO4 + 2 NaCl + HC1; (2) Na2HPO4 + HC1 = NaH2PO4 + NaCl. Adding (1) and (2), and eliminating HC1, we get: - (3) 2 Na2HPO4 + AICI3 = IA1PO4 + NaH2PO4 + 3 NaCl; («) A1PO4 + 3 NaOH = Na3AlO3 + H3PO4, (b) H3PO4 + 3 NaOH = Na3PO4 + 3 H2O. Adding («) and (b), and eliminating H3PO4, we get: - A1PO4 + 6 NaOH = Na3A103 + Na3PO4 + 3 H2O. 7. Any aluminum compound, when strongly ignited in the air, is converted into A12O3; if this is moistened with a very dilute solution of Co(NO3)2 and again strongly heated, a blue mass is obtained, due to the formation of cobalt aluminate. This reaction serves as an excellent confirmatory test for aluminum. Chromium The two principal oxides of chromium are Cr2O3 and CrO3. The former is basic and forms the various chromic * salts by combining with acids, e.g., Cr2O3 + 6 HC1 = 2 CrCl3 + 3 H2O ; similarly by solution of Cr2O3 in H2SO4 and HNO3, Cr2(SO4)3 and Cr(NO3)3 are respectively formed. In all these com- pounds chromium deports itself as a metal. CrO3, on the pther hand, is distinctly acid in character, being the anhydride of the hypothetical chromic acid, H2CrO4, the salts of which are known as chromates. The latter may be prepared by treating CrO3 with a base; thus, sodium chromate (Na2CrO4) may be prepared by treating CrO3 with caustic soda: CrO3 + 2 NaOH = Na2CrO4 + H2O. In the chromates, chromium plays the * In this case the ending -ic refers to the element when acting as a base, i.e., its electro-positive properties dominate. This may perhaps best be shown by the valence as Cr111, chromic, etc. 58 QUALITATIVE CHEMICAL ANALYSIS part of the acid radical CrO4, the reactions of which are differ- ent from those given by Cr when a constituent of a chromic salt. An example will serve to illustrate this difference. Ammonium hydroxide, when added to a chromic salt, like CrCl3, causes a precipitate of chromium hydroxide to form; when added, however, to a chromate, as Na^CrO^ no precipi- tate results. Further, if to chromic chloride we add a solution of barium chloride or lead acetate, no precipitate results; while if the same reagants are added to sodium chromate, yellow pre- cipitates are formed, due to the formation of BaCrO4 and PbCrO4, respectively. The distinction is further noted when we compare the aque- ous solutions of chromic salts and chromates; the former pos- sess a green or violet color, while the latter are nearly always yellow. Chromic salts (Crrn) are converted into chromates (CrVI) by oxidation in an alkaline solution ; conversely, chromates (CrVI) are reduced to chromic salts (Crni) by reduction in an acid medium. The essential change may best be seen by considering only the oxides as taking part in the reactions ;* thus, the oxidation of a chromic salt to a chromate is given by the equation Cr2O3 + 30 = 2 CrO3, while the reduction of chromate to a chromic salt may be repre- sented by 2 CrO3 + 6 H = 2 Cr2O3 + 3 H2O. (a) Oxidation of Chromic Salts to Chromates. The oxidation is always carried out in an alkaline medium. In the dry way the oxidation may be accomplished by fusing a chromic com- pound with a mixture of Na2CO3 (which supplies the alkali) and an oxidizing agent like Na2O2, KC103, or KNO3 (which * As all chromic salts may be derived from Cr2Og by treatment of the latter with the appropriate acid, and as the valence of Cr is the same in this oxide and its salts, we may conveniently represent all chromic salts in oxidation equations by CroOg. For a similar reason, all chromates may be represented in reduction reactions by CrOg. THE METALS 59 supplies the O). In its simplest form the equation for the oxida- tion is - Cr2O3 + 2 Na2CO3 + 30 = 2 Na2CrO4 4- f 2 CO2. The oxidation may be carried out in an alkaline solution by using any one of the many oxidizing agents, such, for instance, as Br (or any halogen), KMnO4, or H2O2. The alkali first pre- cipitates chromic hydroxide: - CrCl3 4- 3 NaOH = | Cr(OH)3 + 3 NaCl. This then dissolves in excess giving sodium chromite: - Cr(OH)3 + 3 NaOH = Na3CrO3 + 3 H2O. The chromite is then oxidized by the oxidizing agent to chromate: - 2 Na3CrO3 + 3 O + H2O = 2 Na2CrO4 4- 2 NaOH. Sod. Chromite. Sod. Chromate. When sodium dioxide (Na2O2) is used, it is needless to first make the solution alkaline, because the sodium compound in contact with water is decomposed, yielding NaOH and O, according to the equation Na2O2 4- H2O = 2 NaOH 4- O. An excess of Na2O2 will therefore yield the excess of NaOH necessary to convert the chromic salt to sodium chromite, while the oxygen liberated at the same time will oxidize the chromite to chromate: - 2 Na3CrO3 + 3 Na2O2 4- 4 H2O = 8 NaOH 4- 2 Na2CrO4. The oxidation by means of sodium dioxide is to be preferred to the other agents for the conversion of chromic compounds to chromate. In every case the oxidation is accompanied by a change in color from green to yellow. If a solution of a chromate is acidified, the color changes from a yellow to an orange-red, due to the formation of a dichro- mate : - 2 K2CrO4 4- 2 HNO,= K2Cr2O7 4- 2 KNO3 4- H2O. a o u a { q ' a 60 QUALITATIVE CHEMICAL ANALYSIS Conversely, if a base is added to a dichromate solution, the color changes from orange-red to yellow, due to the formation of a chromate: - K2Cr2O7 + 2 KOH = 2 K2CrO4 4- H2O. Both chromates and dichromates, when in solution, may be precipitated by solutions of Ba or Pb salts [distinction from chromic (Cr111) compounds] : - K2CrO4 + Pb(C2H3O2)2 = | PbCrO4 4- 2 KC2H3O2; Yellow K2Cr2O7 + 2 BaC^ 4- H2O = 12 BaCrO4 4- 2 KC1 + 2 HC1. Yellow (b) Reduction of chromates to chromic compounds is effected in acid solutions by any one of the many reducing agents, e.g., H2S, HI, SO2, concentrated HC1, and various organic sub- stances, as alcohol and oxalic acid. With concentrated HC1 the reaction is - K2Cr2O7 + 14 HC1 - 2 KC1 + 2 CrCl3 4- 7 H2O 4- f 3 CLj. With H2S the reduction takes place in accordance with the equation K2Cr2O7 4- 3 H2S 4* 8 HC1 = 2 CrCl3 4~ 2 KC14- 7 H2O 4- 3 S. In this case the green solution appears turbid from the separa- tion of S. With sulphurous acid the equation is - K2Cr2O7 4- H2SO4 4- 3 H2SO3 = Cr2(SO4)3 4- K2SO4 4- 4 H2O. In all the above cases the reduction is evidenced by a change in color from orange-red to green. Other reactions for chromates will be given in Part II., deal- ing with the acids (see page 122). The Chromic Salts Of the common salts of chromium, the sulphate and the chloride exist in two forms : one is very readily soluble in water; while the other, which has been ignited, is neither soluble in THE METALS 61 water nor acids. The nitrate exists in one form only and is easily soluble in water. All aqueous solutions of chromium salts have either a green or violet color, which varies with the concentration and other conditions. A solution containing as little as i part of chromium in 10,000 parts of water will have a distinct bluish green color. Reactions of the Chromic Salts 1. Ammonium Hydroxide produces a greyish green or blue gelatinous precipitate of Cr(OH)3, soluble with difficulty in excess with the formation of a red solution, from which, on boiling, Cr(OH)3 is reprecipitated. The precipitate is easily soluble in acids and in sodium hydroxide (see 2). 2. Sodium or Potassium Hydroxide precipitates Cr(OH)3, solu- ble in excess in the cold to a green solution with the formation of sodium chromite; on boiling this solution, Cr(OH)3 is repre- cipitated (distinction from Al). The precipitate is easily solu- ble in acids. CrCl3 + 3 NaOH = | Cr(OH)3 + 3 NaCl; Cr(OH)3 + 3 NaOH = Na3CrO3 + 3 H2O ; Na3CrO3 + 3 H2O (boiling) = | Cr(OH)3 + 3 NaOH ; Cr(OH)3 + 3 HC1 = CrCl3 + 3 H2O. 3. Ammonium Sulphide precipitates Cr(OH)3, for Cr2S3, like A12S3, is hydrolyzed by water with the formation of Cr(OH)3 and the evolution of H2S : - 2 CrCl3 + 3(NH4)2S -f- 6 H2O = |2Cr(OH)3+ f 3 H2S + 6 NH4C1. 4. Sodium Carbonate also precipitates the hydroxide. The presence of non-volatile organic acids, like tartaric and citric acids, as well as organic matter containing (OH) groups, as sugar and starch, interferes with reactions 1, 2, 3, 4, and 5. 5. Disodium Phosphate precipitates from solutions of chromic choride green CrPO4: - 2 Na2HPO4 + CrCl3 = 3 NaCl + NaH2PO4 + | CrPO4. 62 QUALITATIVE CHEMICAL ANALYSIS The precipitate is easily soluble in inorganic acids, but is prac- tically insoluble in cold dilute acetic acid, although it is soluble in a large excess of 50 per cent, acetic acid. Sodium Dioxide. If a solution of a chromic salt is treated with a sufficient amount of sodium dioxide and boiled, all of the chromium will be converted into sodium chromate. The reac- tion may be represented by the following equations: - (1) 3 Na2O2 4- 3 H2O = 6 NaOH + 3 O ; (2) CrCl3 4- 6 NaOH = Na3CrO3 4- 3 NaCl 4- 3 H2O ; (3) 2 Na3CrO3 4- 3 O 4- H2O = 2 NagCrO* 4- 2 NaOH. 7. If Sodium Acetate is added to a solution of a chromium salt, no precipitate is produced even on boiling. If, however, the solution contains relatively large amounts of iron (ferric) and aluminum, the chromium will be almost completely precipitated as a basic acetate on boiling (compare the corresponding reac- tion for aluminum). Should the iron and aluminum be present in small and the chromium in relatively large amounts, the pre- cipitation will be incomplete and in the filtrate will be found some of the Al, Cr, and Fe. The important deduction from these facts is, that in the pres- ence of a large amount of chromium it is necessary, in order to completely precipitate aluminum and iron as basic acetates, that one of the latter metals be present in large excess. Iron Iron, as an electro-positive element, forms two distinct classes of salts; viz., the ferrous compounds, in which iron is divalent, and the ferric salts, in which iron is trivalent. As the two classes exhibit a difference in behavior when treated with the same reagents, we shall consider them separately. The Ferrous Compounds When they contain " water of crystallization " the ferrous salts are green, and when anhydrous, they are white. The aqueous solutions, except when concentrated, are almost color- less. Ferrous salts in solution are very unstable, for they rapidly THE METALS 63 absorb oxygen from the air and are converted into basic ferric salts, difficultly soluble in water. Oxidizing agents readily change ferrous salts to ferric compounds. Reactions i. Ammonium, Sodium, or Potassium Hydroxide precipitates at first white gelatinous Fe(OH)2, which, on exposure to the air, is rapidly oxidized, becoming first dirty green, then black, and finally a reddish brown; the last is ferric hydroxide, and the other colors are doubtless due to varying mixtures of ferrous and ferric hydroxides. FeCl2 + 2 NH4OH = | Fe(OH)2 + 2 NH4C1; White. 2 Fe(OH)2 + O + H2O = 12 Fe(OH)3. Reddish brown. In the presence of much ammonium chloride, ammonium hy- droxide fails to yield an immediate precipitate; but on exposure of the ammoniacal solution to the air, ferric hydroxide is finally thrown down. If the air is excluded, ammonium hydroxide does not precipitate ferrous salts in the presence of a sufficient quantity of ammonium salts (distinction from ferric salts). The property of not being precipitated by ammonium hydrox- ide in the presence of a sufficient amount of ammonium salts is not peculiar to ferrous salts alone, but is shared alike by the salts of nickel, cobalt, manganese, zinc, and magnesium.* 2. Hydrogen Sulphide, in acid solution, gives no precipitate. From neutral solutions a slight precipitate of FeS results; if, however, considerable sodium acetate is present, a larger, though still incomplete, precipitation is obtained. From alkaline solu- tions, H2S completely precipitates the iron as black ferrous sulphide. * An explanation of this fact may be found in the Theory of Electrolytic Dissoci- ation and the Laiv of Mass Action. The presence of NH4CI diminishes the concen- tration of the (OH) ions derived from the ammonium hydroxide to such an extent as to yield with the ferrous iron present an amount of Fe(0H)2 less than the solubility product of the latter; in other words, the amount of Fe(0H)2 which can form under these conditions is less than that required to saturate the solution, and as a conse- quence no precipitate forms. 64 QUALITATIVE CHEMICAL ANALYSIS 3. Ammonium Sulphide precipitates black FeS, easily soluble in acids with the formation of a ferrous salt and evolution of h2S:- FeCl2 + (NH4)2S = | FeS + 2 NH4C1/, FeS + 2 HC1 = FeCl2+ f H2S. ' When moist, it readily oxidizes in the air, becoming first ferrous sulphate and finally brown basic ferric sulphate. To prevent this oxidation, the precipitate should be washed with water con- taining ammonium sulphide. The presence of ammonium chlo- ride assists the precipitation. 4. Potassium Cyanide precipitates brown ferrous cyanide, soluble in excess with the formation of potassium ferrocya- nide : - FeCj2 + 2 KCN = । Fe(CN)2 + 3 KC1; Fe(CN)2 + 4 KCN = K4Fe(CN)6. The solution of potassium ferrocyanide does not give any of the reactions of ferrous salts; it is therefore not a ferrous salt, but the potassium salt of ferrocyanic acid, H4Fe(CN)6. The group Fe(CN)6 is an acid radical like CrO4 in chromates and differs distinctly in its behavior from iron, existing as the simple metallic component of salts. 5. Potassium Ferrocyanide precipitates, in the complete ab- sence of air, white K2FeHFe(CN)6; under ordinary atmospheric conditions, however, a light blue precipitate is obtained, due to partial oxidation; on prolonged exposure, it is completely con- verted into a dark blue precipitate of prussian blue: - FeCl2 + K4Fe(CN)6 = | FeK2Fe(CN)6 + 2 KC1. 6. Potassium Ferricyanide produces even in very dilute solu- tions of ferrous salts a dark blue precipitate, known as Turn- bull's blue, which is indistinguishable in color from prussian blue: - 3 FeCl2 + 2 K3Fe(CN)6 = |Fe3[Fe(CN)6]2 + 6 KC1. The precipitate is insoluble in HC1, but is decomposed by caustic alkalies with the formation of ferrous hydroxide and alkali ferri- cyanide : - Fe3[Fe(CN)6]2 + 6 KOH = |3 FeCOH), + 2 K3Fe(CN)3. THE METALS 65 The ferricyanide at once oxidizes the ferrous hydroxide, so that the final products are Fe(OH)3 and K4Fe(CN)6. 7. Potassium thiocyanate gives no reaction with ferrous salts {distinction from ferric salts). Oxidation of Ferrous to Ferric Salts. It has been already stated that solutions of ferrous salts are very unstable, oxidizing gradually on exposure to air to ferric compounds. The oxida- tion can be more rapidly accomplished by the use of oxidizing agents in acid solution, as the halogens, aqua regia, a mixture of HC1 and KC1O3, nitric acid, potassium permanganate, potassium dichromate, and hydrogen dioxide. The equations for the oxi- dation of ferrous salts by nearly all of these oxidizing agents have been given under Oxidation and Reduction (see page 9). In oxidizing with nitric acid, the strong acid should be added drop by drop to the boiling acid solution of ferrous salt until no further darkening of the solution is evident. The oxidizing action of hydrogen dioxide and aqua regia, respectively, are given by the following equations: - 2 FeCL + 2 HC1 + H2O2 = 2 FeCl, + 2 H,O ; zj a U O a J 3 FeCl2 + 3 HC1 + HNO3 = 3 FeCl3 + f NO + 2 H2O. Ferric Salts Most of the ferric salts, as the chloride, nitrate, and sulphate, yield solutions with a yellowish brown color, which varies in intensity with the concentration and temperature of the solution, as well as with the quantity of free acid present. The ferric ammonium alum, Fe2(SO4)3 • (NH4)2SO4 • 24 H2O, is violet. Fer ric salts in dilute aqueous solutions are readily hydrolyzed, par- ticularly on heating, with the formation of an insoluble basic ferric salt which dissolves on the addition of an acid: - Fe2(SO4)3 + H2O Fe2(SO4)2O + H2SO4. Reactions 1. Ammonium, Sodium or Potassium Hydroxide precipitates a reddish brown gelatinous precipitate of Fe(OH)3. The precipi- 66 QUALITATIVE CHEMICAL ANALYSIS tate is unaffected by the presence of ammonium salts [distinc- tion from Fe(OH)2], and is soluble in acids, but is insoluble in an excess of sodium hydroxide (distinction from Al and Cr); it is also insoluble in an excess of ammonium hydroxide : - FeCl8 + 3 NH4OH = | Fe(OH)3 + 3 NH4C1; Fe(OH)3 + 3 HC1 = FeCl3 + 3 H2O. On ignition it yields Fe2O3 : - 2 Fe(OH)3 + (heat) = Fe2O3 + 3 H2O. Ignited, Fe2O3 is difficultly soluble in dilute acids, but dissolves on prolonged treatment with hot concentrated hydrochloric acid. 2. Ammonium Sulphide gives with acid solutions a precipitate consisting of FeS + S. From ammoniacal solutions, black ferric sulphide, Fe2S3, is precipitated: - 2 FeCl3 + 3 (NH4)2S = | Fe2S3 + 6 NH4CL The precipitate is readily soluble in hydrochloric acid with the formation of ferrous chloride and the separation of sulphur: - Fe2S3 + 4 HC1 = 2 FeCl2 + f 2 H2S + S. 3. Potassium Ferrocyanide produces with ferric salts a blue precipitate known as prussian blue : - - 4 FeClg + 3 K4Fe(CN)6 = | Fe4[Fe(CN)6]3 + 12 KC1. The precipitate is,insoluble in dilute HC1, but dissolves in oxalic acid, as well as in a great excess of the precipitant, with the formation of a blue solution. Prussian blue is decomposed by caustic potash, the products being ferric hydroxide and potassium ferrocyanide: - Fe4[Fe(CN)6]3 + 12 KOH = |4 Fe(OH)3 + 3 K4Fe(CN)e. In making this test for iron, it is important that the solutions contain only a small amount of strong acid, as the latter would partially decompose the reagent with the formation of a small quantity of iron salt, which, reacting with the unchanged portion THE METALS 67 of the reagent, would yield a blue coloration. Neutral solutions containing I part of iron in 500,000 parts of water will give this reaction. When only small amounts of iron are present, a blue or green coloration, instead of a blue precipitate, is obtained. 4. Potassium Ferricyanide does not precipitate ferric salts, but produces a brown coloration (distinction from ferrous salts). 5. Potassium Thiocyanate gives with solutions of ferric salts a deep red coloration, due to the formation of ferric thiocyanate, which is soluble in water: - FeCl3 + 3 KCNS Fe(CNS)3 + 3 KC1. The reaction being reversible, its sensitiveness is increased by adding an excess of the reagent. As little as 1 part of iron in 1,600,000 parts of water can be detected by this reagent. The delicacy of the test may be further increased by adding a little pure ether and shaking; the ether extracts, and thus concen- trates, the colored body. Nitric and chloric acids also give with the reagent a red coloration, but the latter, when due to these substances, is destroyed by adding alcohol and heating. The presence of relatively large amounts of alkali acetate, organic acids, like tartaric, acetic, and oxalic, as well as phosphoric, arsenic, and boric acids, interfere with the reaction in neutral, though not in strongly acid, solutions. The addition of acid in making the test is therefore advisable. Mercuric chloride bleaches the red coloration. 6. Disodium Hydrogen Phosphate in neutral or slightly acid solutions of ferric salts containing a relatively large amount of sodium acetate, produces a buff-colored precipitate of ferric phosphate: - (a) FeCl3 + 2 Na2HPO4 = |FePO4 + NaH2PO4 +3 NaCl; (£) FeCl3 + Na2HPO4+ NaC2H3O2 = 1 FePO4 + 3 NaCl+ HC2H3O2. In (a) all the iron is precipitated but not all the phosphoric acid; in (£) both the iron and phosphoric acid are precipitated. 68 QUALITATIVE CHEMICAL ANALYSIS FePO4 is insoluble in acetic acid, but readily dissolves in HC1. Caustic alkalies decompose it into Fe(OH)3 and Na3PO4: FePO4 + 3 NaOH = | Fe(OH)3 + Na3PO4. Treatment with ammonium hydroxide or hot water effects a par- tial hydrolysis into the hydroxide. 7. Sodium or Ammonium Acetate, when added in excess to a slightly acid solution of a ferric salt, causes the solution to take on a reddish brown color, due to the formation of ferric ace- tate : - FeCl3 4- 3 NaC2H3O2 = Fe(C2H3O2)3 + 3 NaCl. If this solution is largely diluted and boiled, all the iron will be precipitated as a basic acetate: - Fe(C2H3O2)3+ H2O(boiling) Fe(OH yC2H3O2>2 + H C2H3O2. The presence of non-volatile organic acids or sugar interferes with the precipitation of Fe in reactions 1, 6, and 7. 8. Reduction of ferric salts to ferrous may be readily effected in acid solution by reducing agents, as H2S, nascent H, SnClg, H2SO3, HI, and others. The following are the equations : - 2 FeCl3 + H2S = 2 FeCl2 + 2 HC1 + |S; FeCl3 + H (from Zn + HC1) = FeCl2 + HC1; 2 FeCl3 + SnCl2 = 2 FeCl2 + SnCl4; 2 FeCl3 + H2SO3 + H2O = 2 FeCLj + 2 HC1 + H2SO4; FeCl3 + HI = FeCl2 + HC1 + |l. Nickel When in the crystalline condition or in aqueous solutions, the nickel salts are green ; when anhydrous, they are yellow. The green solutions can be rendered colorless by admixture with cobalt compounds in the proportion of 3 of nickel to 1 of cobalt. 1. Potassium or Sodium Hydroxide precipitates green gelati- nous Ni(OH)2, insoluble in excess and not oxidized on exposure to air:- NiCl2 + 2 NaOH = | N^OH^ + 2 NaCl. THE METALS 69 The precipitate is readily soluble in acids, also in ammonium hydroxide and ammonium salts. If the alkaline solution con- taining Ni(OH)2 in suspension is treated with bromine or chlorine and the mixture boiled, black nickel (-ic) hydroxide is formed: - Ni(OH)2 + NaOH + Br = | Ni(OH)3 + NaBr. 2. Ammonium Hydroxide, when considerably diluted and added in small quantity, causes a green turbidity, due to the formation either of a basic salt or the hydroxide: - NiCl2 + 2 NH4OH = | Ni(OH)2 + 2 NH4C1. The precipitate is readily soluble in excess or in the presence of ammonium salts, with the formation of a blue solution contain- ing a nickel ammonia salt: - Ni(OH)2 + 2 NH4C1 + 2 NH3 = Ni(NH3)4Cl2 + 2 H2O. Therefore, in the presence of sufficient ammonium salts, Ni is not precipitated by ammonium hydroxide. 3. Hydrogen Sulphide yields no precipitate in solutions of nickel containing mineral acids or much acetic acid. If, how- ever, the acetic acid solution contains a relatively large amount of sodium acetate, or if the solutions are rendered ammoniacal, hydrogen sulphide will completely precipitate the nickel as black nickel sulphide: - NiCl2 -I- 2 NaCoHoOo + H2S = | NiS 4- 2 NaCl + 2 HC2H3O2. & Lt U ii ii I A O A 4. Ammonium Sulphide gives with neutral or alkaline solu- tions of nickel salts a black precipitate of NiS, somewhat soluble in excess, especially in the presence of free ammonia, with the formation of a dark brown solution (distinction from Co). If this brown solution is acidified with acetic acid and boiled, NiS is reprecipitated. The presence of large quantities of ammo- nium salts prevents the solution of NiS in (NH4)2S. NiCl2 + (NH^S = | NiS + 2 NH4C1. Nickel sulphide is practically insoluble in cold HC1 (1.02) (distinction from the sulphides of Mn, Zn, and Fe). It is also 70 QUALITATIVE CHEMICAL ANALYSIS insoluble in acetic acid, but is readily taken into solution on heating with aqua regia or concentrated nitric acid: - 3 NiS + 2 HNO3 + 6 HC1 = 3 NiCl2 + f 2 NO + 4 H2O + 13 S. The sulphur, which separates in a plastic condition, often appears black because of the presence of some NiS inclosed in it. If the treatment with aqua regia is continued for some time, all the sulphide will be dissolved and the sulphur will be con- verted into sulphuric acid : - S + 6 Cl + 4 H2O = H2SO4 + 6 HC1. On exposure to air, moist nickel sulphide is oxidized to NiSO4. 5. Potassium Cyanide gives a green precipitate of nickel cyanide, readily soluble in excess with the formation of a double cyanide: - NiCl2 + 2 KCN = | Ni(CN)2 + 2 KC1; Ni(CN)2 + 2 KCN = K2Ni(CN)4. If the solution of the double cyanide is made strongly alkaline with NaOH, and then treated with bromine or chlorine and gently heated, decomposition of the double cyanide results with the precipitation of black nickel(-ic) hydroxide (distinction and method of separation from Co): - K2Ni(CN)4 + Br + 3 KOH = | Ni(OH)3 + 4 KCN + KBr. 6. Potassium Nitrite, in dilute solutions of nickel salts acid with acetic acid, gives no precipitate (distinction and method of separation from Co). 7. Borax Bead Test. A borax bead, when fused with a nickel compound in the oxidizing flame, is colored reddish brown, due to the formation of Na2Ni(BO2)4. In the reducing flame, the nickel is reduced to the metallic state, imparting a gray color to the bead. Cobalt The cobalt salts, when in the crystallized condition or in aque- ous solution, are reddish pink; in the anhydrous form, they are usually blue. The concentrated aqueous solutions in the pres- ence of HC1 are also blue. THE METALS 71 i. Sodium Hydroxide precipitates from cold solutions a blue basic salt: - CoCl2 + NaOH = |Co(OH)Cl + NaCl. This is converted, on warming in contact with the alkali, to pink cobaltous hydroxide: - Co(OH)Cl + NaOH = | Co(OH)2 + NaCl. The precipitate is insoluble in excess, but readily soluble in ammonium salts; hence, the presence of ammonium salts in sufficient quantity interferes with the precipitation. On expo- sure to the air, the pink hydroxide oxidizes to black Co(OH)3 [resemblance to Fe(-ous) and Mn, and distinction from Ni] : - 2 Co(OH)2 + H2O + 0 = ^2 Co(OH)3. 2. Ammonium Hydroxide, in the absence of ammonium salts, produces the same precipitate as in (i), but the latter readily dissolves in excess of the reagent to a brownish solution, which, on exposure to air or on boiling, changes to a red solution, due to the formation of a complex ammonia compound: - Co(OH)2 + 2 NH4C1 + 2 NH3 = Co(NH3)4C12 + 2 H20. As in the case of Al, Cr, and Fe(-ous), the precipitation of Co as hydroxide is interfered with by the presence of non- volatile organic acids or sugar. 3. Hydrogen Sulphide. Same as with Ni. 4. Ammonium Sulphide precipitates in neutral or alkaline solutions black CoS, insoluble in excess (distinction from Ni), insoluble in HC1 (sp. gr. 1.02) and in acetic acid. It is soluble in aqua regia and concentrated nitric acid with the separation of sulphur: - 3 CoS + 8 HNO3 = 3 Co(NO3)2 + f 2 NO + 4 H2O + | 3 S. 5. Potassium Cyanide gives in neutral solutions a light brown precipitate of cobaltous cyanide, easily soluble in excess to a brown solution with the formation of a double cyanide: - CoCl2 + 2 KCN = | Co(CN)2 + 2 KC1; Co(CN)2 + 4 KCN = K4Co(CN)6. QUALITATIVE CHEMICAL ANALYSIS 72 The latter is similar to potassium ferrocyanide, hence it is called potassium cobaltocyanide. On warming the solution of the double cyanide for some time, it changes color to a bright yel- low, due to oxidation to potassium cobalticyanide [similar to K3Fe(CN)6] : - 2 K4Co(CN)6 + H2O + 0 = 2 K3Co(CN)6 + 2 KOH. The reaction takes place more rapidly if sodium hydroxide and bromine or, what amounts to the same, NaBrO solution, are added to the solution of potassium cobaltocyanide. Nickel does not form the corresponding compound, but under these conditions it is converted into black insoluble Ni(OH)3 (distinc- tion and method of separation from Co). 6. Potassium Nitrite produces, when added in excess to a not too diluted solution of cobalt acidified with acetic acid, a yellow crystalline precipitate of potassium nitrocobaltate, K3Co(NO2)6. With dilute solutions of cobalt, the mixture should be warmed and allowed to stand for at least twelve hours in order to get complete precipitation. The reaction may be represented as taking place in several stages : - (a) CoCl2 + 2 KNO2 = | Co(NO2)2 + 2 KC1; (0) 2 KNO2 + 2 HC2H3O2 = 2 HNO2 + 2 KC2H3O,; (r) 2 HNO2 = H2O + NO + NO2; \d) Co(NO2)2 + NO2 = Co(NO2)3; (*) Co(NO2)3 + 3 KNO2 = | K3Co(NO2)6. The precipitate is somewhat soluble in water, but is prac- tically insoluble in a solution saturated with a potassium salt. It is insoluble in alcohol and in an excess of KNO2 solution. Hence, for a rapid precipitation of cobalt as K3Co(NO2)6, the solution of cobalt should be concentrated by evaporation, the mineral acid replaced by acetic, saturated with KC1, and then treated with an excess of KNO2 solution. If the mixture is now warmed and vigorously shaken, complete precipitation may be secured in a half hour. 7. Nitroso-P-naphthol, dissolved in 50 per cent, acetic acid, yields with a hot solution of cobalt, preferably the chloride or THE METALS 73 sulphate acidified with hydrochloric acid, a voluminous red pre- cipitate of cobalti-nitroso-/3-naphthol (distinction and method of separation from nickel, which, in HC1 solution, does not give a precipitate). 8. A borax bead, when fused with cobalt compounds either in the oxidizing or reducing flames, is colored blue. This test is not masked by the presence of moderate amounts of nickel. Manganese The manganese salts, which may be formed by the solution of the oxide MnO in acids, are colored pink in the crystallized condition as well as in concentrated aqueous solutions. In the anhydrous state, with the exception of the sulphide, they ate nearly all colorless. Reactions i. Sodium or Potassium Hydroxide produces with manganous salts a white precipitate of Mn(OH)2, which, on exposure to air, rapidly oxidizes, becoming brown: - MnCl2 + 2 NaOH = |Mn(OH)2 + 2 NaCl; Mn(OH)2 + O = |MnO(OH)2 (manganous acid); MnO(OH)2 + |Mn(OH)2=4Mn2O3+ 2 H2O. 2. Ammonium Hydroxide yields with manganous solutions, in the absence of ammonium salts, a partial precipitation of white Mn(OH)2, oxidizing, as described in (i), to brown Mn2O3. In the presence of a sufficient amount of ammonium salts, no immedi- ate precipitate forms; but, on exposure to air, MnO(OH)2 is thrown down. The separation of manganese from any or all of the trivalent metals of this group by means of NH4C1 and ammonium hydroxide is therefore incomplete. Non-volatile organic acids and sugar interfere with the precipitation of Mn(OH)2. 3. Ammonium Sulphide precipitates light pink manganous sulphide, which, on exposure to air, becomes dark brown, due to partial oxidation to Mn2O3 : - MnCl2 + (NH4)2S = | MnS + 2 NH4C1. 74 QUALITATIVE CHEMICAL ANALYSIS The precipitate is easily soluble in dilute acids (distinction from Ni and Co), even in acetic acid (distinction from Zn, as well as Ni and Co). The addition of ammonium chloride assists the precipitation, while the presence of oxalates and tartrates retard it. 4. Lead Dioxide and Nitric Acid. If a very dilute solution of manganous salt, free from HC1 or chlorides, is boiled with a gram of lead dioxide and a few cubic centimeters of con- centrated nitric acid, and allowed to settle, the clear super- natant liquid will be colored purple, due to the formation of permanganic acid: - 2 MnSO4 + 5 PbO2 + 6 HNO3 = 12 PbSO4 + 3 Pb(NO3)2 4- 2 H2O 4- 2 HMnO4 Purple. This reaction is sufficiently delicate to detect a trace of man- ganese. 5. If a sodium carbonate bead is fused with a very small amount of a manganese compound in the oxidizing flame, or if the fused mass, while hot, is quickly dipped into a little pow- dered potassium chlorate, a bluish green or green mass will be formed, due to the formation of sodium manganate, Na2MnO4 - Mn(OH)2 + (heat) = MnO + H2O; MnO(OH)2 + (heat) = MnO2 + H2O ; MnO 4- Na2CO3 4- O2 = f CO2 4- Na2MnO4; MnO2 4- Na^Og + O = f CO2 4- NagMnO^ Zinc Most of the zinc salts are colorless; some are soluble in water, and the others are dissolved by acids. 1. Sodium or Potassium Hydroxide precipitates white gelati- nous zinc hydroxide, readily soluble in excess with the formation of sodium zincate [similar to Al, distinction from Fe(-ic) and Mn] : - ZnCl2 4- 2 NaOH = | Zn(OH)2 + 2 NaCl; Zn(OH^ 4- 2 NaOH = NagZnOg 4- 2 H2O. THE METALS 75 Unless the solution of the zincate contains a decided excess of NaOH, it will be decomposed on boiling with the reprecipitation of the hydroxide: - Na2ZnO2 + 2 H2O (boiling)^ 2 NaOH + |Zn(OH)2. 2. Ammonium Hydroxide yields with solutions of zinc salts, in the absence of ammonium salts, a partial precipitation of zinc hydroxide, readily soluble in excess in the presence of ammo- nium salts with the formation of a complex ammonia salt: - ZnCl2 + 2 NH4OH = | Zn(OH)2 4- 2 NH4C1; Zn(OH)2 + 2 NH4C1 + 2 NH3 = Zn(NH3)4Cl2 + 2 H2O. 3. Hydrogen Sulphide, when passed into neutral solutions of zinc salts of inorganic acids, incompletely precipitates white zinc sulphide (ZnS). A partial precipitation is also obtained from solutions containing a small amount of free mineral acid. From solutions of zinc acetate, or from neutral solutions of salts of strong acids containing a moderate amount of sodium acetate, H2S completely precipitates all the zinc as sulphide on boiling. Warming in the presence of alkali acetate promotes the pre- cipitation : - (a) ZnCl2 + H2S | ZnS 4- 2 HC1; (£) ZnCl2 + 2 NaC2H3O24- H2S = |ZnS + 2 NaC14-2 HC2H3O2. As ZnS is soluble in HC1, the precipitation in (#) is never com- plete. In equation (£), NaC2H3O2 has the effect of displacing the strong HC1 by the weak acetic acid, in which ZnS is prac- tically insoluble. The tendency of ZnS to pass through the filter may be over- come by precipitating the sulphide in a nearly boiling solution of acetic acid containing a moderate excess of NaC2H3O2, and filtering rapidly while hot. The precipitate may then be washed with hot water containing NH4C2H3O2 or NH4NO3 and H2S. 4. Ammonium Sulphide yields in neutral and alkaline solutions a white precipitate of ZnS : - ZnCl2 + (NH4)2S = I ZnS + 2 NH4C1. 76 QUALITATIVE CHEMICAL ANALYSIS ZnS is readily soluble in dilute mineral acids, but is insoluble in acetic acid and in caustic alkalies. 5. Any dried zinc compound, when moistened with dilute cobalt nitrate solution and ignited, will yield a green mass, due to the formation of a double oxide of Co and Zn (Th6nard's green). This is an excellent confirmatory test for zinc and serves to distinguish zinc from aluminum. Outline of the Method of Analysis for Group III. From an examination of the foregoing reactions, it becomes evident that if ammonia is added to a solution containing all the metals of this group, in the presence of a sufficient amount of ammonium chloride, all the trivalent metals (assuming the iron to be in the ferric state), viz., Fe, Al, and Cr, will be precipitated as hydroxides, while the remaining metals will be left in solution. This method would seem a desirable one for the separation of the third group into two divisions, and such a plan is, in fact, adopted by some chemists; but the fact that under these condi- tions manganese and zinc are not completely held in solution, and if present in small amounts may be wholly precipitated with the trivalent metals, makes this method of separation in- complete; therefore we have not adopted it. It should, however, be remembered that the method gives fairly satisfactory results if the first precipitate of the hydroxides of Al, Cr, and Fe, con- taining some Mn and Zn, is dissolved and reprecipitated, and the second filtrate is united with the first. Another reaction which may be utilized in separating the third group into two divisions is the basic acetate precipitation. This, it will be remembered, is based on the fact that in a nearly neutral solution containing a large excess of sodium acetate, a large amount of boiling water precipitates the basic acetates of ferric iron, aluminum, and chromium, while the remaining diva- lent metals are left in solution. This method of separation, one of the oldest in analytical chemistry, is exceedingly valuable in some cases, but it is not to be employed as a general method, because, as already pointed out (see under chromium, reaction 7), of its uncertainty in the presence of chromium. THE METALS 77 The method adopted in this book consists in precipitating the entire group with (NH^S after rendering the solution alkaline with NH4OH. Instead of (NH4)2S, the hot ammoniacal solu- tion may be treated with a stream of H2S until the precipita- tion is complete. In either case, the precipitate will consist of the hydroxides of aluminum and chromium, and the sulphides of iron, nickel, cobalt, manganese, and zinc. Since only the sulphides of Ni and Co are insoluble in HC1 (i : 9), it follows that if the third group precipitate is treated with a sufficient amount of HC1 (1:9) and filtered, there will remain on the filter the sulphides of Ni and Co, while in the filtrate will be found the chlorides of Al, Cr, Fe, Mn, and Zn. The residue, consisting of NiS and CoS, is next examined with a borax bead in the oxidizing flame, when, if not too great an amount of nickel is present, a blue bead will be obtained, indicating the presence of cobalt. To separate nickel and cobalt existing as sulphides, we must first get them into solution; this is accomplished by heating with aqua regia, which converts the sulphides into soluble chlorides. From the solution of the chlorides, the cobalt may be separated from the nickel by precipitation with either potas- sium nitrite or nitroso-/3-naphthol. The nickel in the filtrate may be precipitated with NaOH and the resulting green hy- droxide verified with a borax bead test. The main filtrate contains, besides the chlorides of Al, Cr, Fe, Mn, and Zn, an excess of HC1 and H2S. The greater part of the HC1 and all of the H2S are expelled by boiling down to a few cc. It will be recalled that in the cold the hydroxides of Al, Cr, and Zn are soluble in excess of sodium hydroxide, forming, respectively, an alumi- nate, chromite, and zincate; while the hydroxides of Fe(-ic) and Mn are insoluble. If an oxidizing agent like Br or H2O2 is present, the chromite is converted into the yellow chromate, which is not precipitated on boiling. So that if a decided excess of NaOH and a little Na2O2 are added to the main filtrate, which has been freed from H2S and the greater part of the HC1 by evaporation to a few cc., and the mixture is boiled, diluted, and filtered, there will remain on the filter the hydroxides of Mn and Fe(-ic), while in the filtrate we should have sodium zincate, sodium 78 QUALITATIVE CHEMICAL ANALYSIS aluminate, and sodium chromate; the last will be evidenced by the yellow color which it imparts to the alkaline solution. In the residue the separation of Fe and Mn may be accomplished by dissolving the precipitate in HC1, nearly neutralizing the solution, and precipitating the Fe as basic acetate. From the filtrate, the Mn may be precipitated by adding Br and boiling, or by the addi- tion of an excess of Na^ or NaOH. When a rough estimation of the amount of Mn and Fe is not desired, separation is unneces- sary, for we can readily identify each in the presence of the other. The manganese may be verified by the characteristic green bead it gives when a little of the mixture is fused in a Na2CO3 bead in the presence of an oxidizing agent, while the iron may be detected by dissolving part of the precipitate in hot dilute HC1 and adding a few drops of potassium ferrocyanide, when a blue precipitate of prussian blue will be obtained. The filtrate from the Mn and Fe precipitate will contain sodium chro- mate, sodium aluminate, and sodium zincate, as well as an excess of NaOH. If this solution is acidified with HNO3, about three grams of NH4C1 are added, and then it is rendered slightly alka- line with ammonium hydroxide, only the Al will be precipitated. The zinc does not precipitate because of the presence of NH4C1, while the chromium no longer acts as metal but as the acid radical (CrO4), and in consequence is not precipitated by am- monium hydroxide. The filtrate from the Al(OH)3will contain zinc, chromium, and a slight excess of ammonia. By rendering the solution acid with acetic acid and adding BaCl2, all of the chromium will be precipitated as BaCrO4. From the filtrate the Zn may be precipitated by H2S. Scheme of Analysis for Group III. This scheme is applicable only in the absence of non-volatile organic matter and interfering acids such as phosphoric acid. The filtrate from Group II., having been boiled to remove the H2S, contains, besides the metals of the succeeding group, an excess of HC1. Preliminary Test. To a small portion of the filtrate from Group II. which has been freed from H2S, add 2-3 drops of THE METALS 79 concentrated HNO3 and boil. Add about 0.5 g. of NH4C1 and then ammonium hydroxide to alkaline reaction. A precipitate may be Fe(OH)3 (red), A1(OH)3 (white), or Cr(OH)3 (greenish blue). If the amount of Mn in the solution is large, a small precipitate of MnO(OH)2 (brown) may also be obtained. If no precipitate forms, the absence of Al, Cr, and Fe is proved. If a precipitate is obtained, it is rapidly filtered and to this filtrate, or to the filtrate from Group IL, in which ammonium hydroxide produces no precipitate, (NH4)2S is added; a precipitate proves the presence of one or more of the remaining members of Group III., viz., Ni, Co, Mn, and Zn. The color of this precipitate sometimes affords an indication of the metals present. If it is black, Ni or Co, or both, are present; if white, Zn is present, and Ni and Co are absent. If pink, becoming brown on expo- sure, Mn is present; the latter may at once be verified by the Na2CO3 bead + KC103. Failure to precipitate with NH4OH and (NH4)2S in the presence of a sufficient amount of NH4C1 proves the absence of Group III. In that case pass to Scheme IV. NOTES 1. As a solution containing 1 part of Cr in 10,000 parts of water shows a distinct bluish green coloration, a colorless solution need not be tested for Cr. 2. The (NH4)2S solution, before using, should be tested with a little NiCL solution in a separate test tube; it should give an immediate black precipitate. 3. Solutions of Ni and Co may be mixed in such proportions as to yield an almost colorless solution; an almost colorless filtrate from the ammonium hydroxide precipitate does not, therefore, prove the absence of Ni or Co. SCHEME III. If the preliminary tests have shown the presence of Group III., the entire filtrate from Group II. is treated with 2 grams of NH4C1 (i) and is heated to boiling. Remove the flame and add ammonium hydroxide with constant stirring until the solu- tion is alkaline (2), then add (NH4)2S until precipitation (as shown by test) is complete. Stir vigorously with gentle heating for a minute (3) and filter on a fluted filter. [The filtrate (4) is 80 QUALITATIVE CHEMICAL ANALYSIS at once made acid with acetic acid (5), boiled until all the H2S is expelled (6), and filtered (7). The clear filtrate is received in a beaker, labelled "Groups IV. and V.," covered and reserved.] The main ppt. may consist of A1(OH)3, Cr(OH)3, FeS, NiS, CoS, MnS, and ZnS. Wash once with hot water containing NH4C1 and a little (NH^S (8), and discard washings. With the aid of a spatula, transfer ppt. to a beaker. Carry to a hood and add 60-80 cc. of HC1 (sp. gr. 1.02), prepared by mixing 1 part of cone. HC1 with 9 parts of water, stir thoroughly (without heating) for about a minute; allow to settle (9), and filter. Residue is NiS + CoS + S (10). Wash on filter once with HC1 (1:9) and reject washings. Test ppt. with borax bead in the O. F. A blue bead proves the presence of Co. Nickel may, however, be also present (11). A reddish brown bead proves the presence of Ni and the absence of a relatively large amount of cobalt. In either case, transfer ppt. to a small evaporating dish, carry to a hood, add 5-10 cc. of dil. aqua regia, and boil till all but a small amount of black S dissolves ; evaporate just to dryness (12). Take up with 2 cc. of dil. HC1 and an equal vol. of hot water; heat if necessary to effect solution. Filter through a very small filter into a test tube. Add NaOH drop by drop to the filtrate till a slight but permanent ppt. forms. Dissolve the ppt. in acetic acid and add about 0.5 cc. in excess (13). Saturate the solution with KC1 (14) by adding the salt in small amounts and shaking after each addition until no more dissolves. Decant the clear solution into another test tube and to the latter add an equal vol. of KNO2. Allow the ppt. to stand with frequent shaking for about | hour and filter. Residue is K3Co(NO2)6 (yellow). To filtrate add NaOH to alkaline reaction; a green ppt. (15) which yields a brown borax bead in the O. F. proves the presence of Ni. Filtrate contains A1C13, CrCl3, FeCl2, MnCl2, ZnCl2 + H2S + excess HC1. Boil down in a large evap. dish under hood to about i cc. (16) ; dilute with io cc. of water, render strongly alkaline with clear NaOH solution, and add (under hood and with caution) 2 g. of Na2O2 (17). Dip finger into solution, rub on thumb; a greasy feel indicates solution is sufficiently alkaline; if not, add more Na2O2. Boil with constant stirring for about 1 minute, add 10 cc. of water, and filter. THE METALS 81 Residue is Fe(OH)3 + MnO2,arH2O(i8). Wash with hot water. Test for Mn. On a Na2CO3 bead take up a very small amount of the ppt. and heat; while hot dip into a little powdered KC103 contained in a small watch glass. A green or bluish green mass of Na2MnO4 proves the presence of Mn. Test for Fe. Transfer part or all of the ppt. to a test tube, add dil. HC1 and heat till solution takes place, cool and add a few drops (19) of K4Fe(CN)6; a blue ppt. proves the presence of Fe. Filtrate (20) may contain Na3A103, Na2CrO4, Na2ZnO2 + excess NaOH. Render slightly acid with cautious addition of cone. HNO3; add 3 grams of NH4C1 (21), heat to boiling, and then add NH4OH drop by drop with constant stirring until the resulting alkaline solu- tion has only a faint odor of ammonia (22). (If two much NH4OH is added, boil off the excess and filter.) Residue is A1(OH)3 (27). Wash with hot water several times and con- firm by wind- ing Pt wire around a por- tion of the paper and ppt., moistening with several drops of Co(NO3)2(23) and igniting strongly. A blue mass is co- balt aluminate, 7AI2O3 • arCoO. Filtrate may contain (NH4)2CrO4, Zn(NH3)4Cl2 + excess NH4OH. Acid- ify with acetic acid, add a gram of NaC2H3O2 (25), and heat to boiling. To the hot solution add BaCl2, drop by drop, till precipitation is complete ; allow to settle and filter through a double filter. Residue is yellow BaCrO4 (24). Confirm by dissolving on filter with dil. HNO3. A yellow solution confirms the presence of Cr. Filtrate, which should be perfectly clear (26), is treated with H2S. A white ppt. is ZnS. To con- firm the presence of Zn, filter, wash with hot water, and then moisten with 2 drops of Co(NO3)2. Wind Pt wire around paper and ppt. and inciner- ate. A green mass is yZnO • xCoO. NOTES 1. NH4CI is added, first, to prevent Mg from precipitating with the Third Group metals; second, because it aids in the precipitation and filtration of the sulphides by preventing them from going into the colloidal condition. 2. The solution is made alkaline with NH3 to neutralize the free acid, which would otherwise decompose the (NH4)2S next to be added. An ex- cess is to be avoided, as it increases the solubility of Al (OH) 3, Cr(OH)3, and NiS. If H2S is to be used instead of (NH4)2S, add 2 cc. cone. NH4OH in excess, and then pass in H2S until the precipitation is complete. 3. Vigorous stirring and heating of the precipitate will have the effect of rendering it more compact and easier to filter. 4. If the filtrate has a dark brown or black color, Ni is probably present. 82 QUALITATIVE CHEMICAL ANALYSIS 5. The filtrate is at once acidified with acetic acid to destroy the excess of (NH4)2S, which, on standing, would oxidize to sulphate and precipitate the alkaline earths. Rendering the solution acid also prevents the formation of (NH4)2CO3 from the absorption of atmospheric CO2. (NH4)2CO3, if formed, would also precipitate the alkaline earths. 6. The H2S is expelled because, like (NH4)2S, it is capable of being oxidized partially to H2SO4 on exposure to air. 7. The residue obtained may consist of NiS and coagulated S, and may be tested for Ni with a borax bead in the O. F. If a satisfactory test for Ni is obtained at this point, it is needless to go through the separation of Ni and Co. 8. The precipitate is rapidly filtered with the aid of a fluted filter in order to prevent the atmospheric oxidation of the sulphides to sulphates; for the same reason, it is recommended that the wash water contain a little (NH4)2S. NH4C1 is added to the wash water to prevent the precipitate from passing through the filter in the colloidal condition. 9. The separation of Ni and Co from the remaining metals by the use of HC1 (1:9) is not complete; small amounts of Ni and Co may pass into solu- tion, while portions of FeS and other acid-soluble sulphides may be mechani- cally inclosed by S and thus escape solution by the acid. 10. A black residue does not prove the presence of Ni or Co for the rea- son stated in note 9; it may be FeS inclosed by S. It is also well to remem- ber that metals of the Second Group, that have not been completely precipi- tated by H2S, will appear at this point. 11. A small amount of Co can, by the bead test, be readily detected in the presence of a relatively large amount of Ni. The test for Ni is, however, effective in the presence of a relatively small amount of Co. 12. Evaporation to 1 or 2 drops will suffice; if the evaporation is carried to the point of dryness, care must be taken not to ignite the residue. 13. These conditions must be closely adhered to. KN02 precipitates Co best in a concentrated solution acid with acetic acid; no mineral acid is per- missible, owing to the solubility of the precipitate therein. The free HC1 is neutralized by the addition of NaOH in slight excess and the latter in turn is neutralized by acetic acid. Should a larger volume than 10 cc. be obtained in making the test, it is recommended that the solution be evaporated to 10 cc. before saturating with KC1. 14. Under the conditions stated in Note 13, complete precipitation of Co by KN02 takes place after a lapse of 24 hours. By saturating the solution of Co with KC1, in which K3Co(NO2)6 is insoluble, and by using an excess of KN02, the precipitation of Co may be rendered complete in a half hour. THE METALS 83 15. Not infrequently a precipitate of uncertain color is obtained with NaOH. It is important to remember that the presence of Ni cannot be con- sidered proved until a characteristic Ni bead is obtained. 16. The solution is evaporated to i cc. to remove the excess of HCI, which, if present, would neutralize the NaOH next to be added. The H2S is expelled at the same time. 17. Na2O2 added to water even in the cold decomposes, giving NaOH + O. At higher temperatures the decomposition takes place violently. Na2O2 should, therefore, be added in small portions to the cold solution with con- stant stirring; the final mixture, which should be strongly alkaline, must be boiled for a minute to decompose the excess of Na2O2 and the perchromates which first form. The Na2O2 furnishes the oxygen necessary for the oxidation of the chromite to chromate. It is important to remember that unless the solution is strongly alkaline, some Zn will be precipitated on boiling and diluting the mixture, due to the reversibility of the reaction, thus : - Zn(0H)2+2 NaOH ^Na2ZnO2 + 2 H2O. 18. Any Ni and Co dissolved by the 1.02 HC1 will appear at this point; their presence, however, does not interfere with the tests for Fe and Mn. 19. A decidedly blue precipitate should be obtained if Fe is present. A blue coloration or a brown or white precipitate is not to be taken as proof of the presence of Fe. In making this test, care must be taken not to add more than a few drops of K4Fe(CN)6, as the precipitate is soluble in an excess. 20. If the filtrate is yellow, Cr is present; if colorless, the test for Cr need not be made. 21. The NH4C1 is added to prevent a partial precipitation of Zn. 22. An excess of NH4OH is to be avoided because of the slight solubility of Al(0H)s in excess. 23. The Co(NO3)2 solution must be very dilute; if strong, it will, on ignition, yield black CoO, which will obscure the blue color of cobalt alumi- nate. The same applies to the confirmatory test for Zn. 24. It not infrequently happens that the solution contains sufficient sul- phates to cause a precipitate of BaSCh (white) to form along with the BaCrO4 (light yellow), thus obscuring the test for Cr; hence the necessity of making the confirmatory test. 25. The addition of NaC2H3O2 weakens the acetic acid with the result that the solvent action of the latter on BaCrO4 is reduced to a minimum. 26. If the filtrate from the BaCrO4 is not clear, filter again, through another double filter, and repeat this treatment until a perfectly clear filtrate is obtained. 84 QUALITATIVE CHEMICAL ANALYSIS 27. A slight precipitate of Al(OH)s is nearly always obtained, being de- rived from the reagents as well as from the action of NaOH on the glass. Judgment must therefore be exercised in reporting the presence of Al in the substance analyzed. GROUP IV. THE ALKALINE EARTHS The alkaline earth metals, barium, strontium, and calcium, are distinguished from the metals of the preceding group by the fact that their salts are neither precipitated by H2S nor by (NH4)2S ; they are grouped together and are distinguished from Group V. by reason of their common property of being precipi- tated by (NH4)2CO3 in the presence of NH4C1. As the "analytical" grouping happens to be identical with their classification according to the Periodic Law, the order of variation in properties becomes an easy matter to remember; for, in most cases, the solubilities of the compounds of strontium are intermediate between those of barium and calcium. Unless the acid radical imparts a color, the salts of the alkaline earths are white or colorless, and, for the most part, insoluble in water. The sulphides, like those of aluminum and chromium, can only exist in the dry state; when treated with water, they are at once hydrolyzed with the formation of the hydroxide and the evolution of H2S. Reactions of the Salts of Barium Many of the salts of barium are insoluble in water; excepting the sulphate and fluosilicate, all are, however, soluble in dilute hydrochloric acid. i. Ammonium Hydroxide (free from carbonate), when added to a solution of barium salts, does not yield a precipitate. If, however, the clear alkaline mixture is exposed to the air, or if ammonium hydroxide from the reagent bottle* is used, a turbidity results from the formation of barium carbonate. The hydrox- * All alkaline liquids will, on exposure to the air, absorb COj, with the formation of carbonate proportional to the amount of CO2 absorbed; the reagent ammonium hydroxide will, therefore, always contain a little (NH^COs, and hence will yield a slight precipitate with the salts of the alkaline earths (see"equations). THE METALS 85 ide is not precipitated because of its ready solubility in water (i part in 20 of cold water) : - BaCl2 + 2 NH4OH = Ba(OH)2 + 2 NH4C1; ^^(OH^CO^ |BaCO3 + H2O; 2 NH4OH + CO2 = (NH4)2CO3 + H2O; (NH4)2CO3 + BaCl2 = | BaCO3 + 2 NH4C1. 2. Ammonium Sulphide (free from carbonate) does not pre- cipitate barium salts; on standing in the air, or with (NH4)2S from the reagent bottle, a slight turbidity results from the forma- tion of barium carbonate. Reagent (NH4)2S, being an alkaline liquid, will, in consequence of absorption of atmospheric CO2, contain a little (NH4)2CO3, hence it will yield an immediate tur- bidity with barium salts: - (1) (NH^S + BaCl2 = BaS + 2 NH4C1; (2) BaS + 2 H2O = Ba(OH)2 + f H2S ; (3) Ba(OH)2 + CO2 = | BaCO3 + H2O. 3. Ammonium or Sodium Carbonate produces in neutral or alkaline solutions of barium salts a white amorphous precipitate of BaCOg, which, on standing or heating, becomes crystalline : - BaCl2 + (NH4)2CO3 = | BaCO3 + 2 NH4CL The precipitate is slightly soluble in NH4C1; therefore, in very dilute solutions of barium salts containing much NH4C1, ammo- nium carbonate does not produce a precipitate. The precipi- tate is easily soluble in acids, even in acetic and carbonic acids: - BaCO3 + HC2H3O2 = Ba(C2H3O2)2 + H2O + f CO2; BaCO3 + H2CO3 = BaH2(CO3)2. Boiling the dicarbonate decomposes it with the evolution of CO2 and precipitation of the normal carbonate : - BaH2(CO3)2 + (heat) = | BaCO3 4- H2O + f CO2. 4. Dilute Sulphuric Acid or any soluble sulphate produces even in very dilute solutions of barium salts a heavy, white, finely 86 QUALITATIVE CHEMICAL ANALYSIS divided precipitate of BaSO4, practically insoluble in water (i part in 800,000 parts of water): - BaCLj 4- H2SO4 = | BaSO4 + 2 HC1. The precipitate is insoluble in alkalies, and is nearly insoluble in dilute but is somewhat soluble in strong acids. Boiled with a strong solution of Na2CO3, it undergoes partial decomposition, according to the equation: - BaSO4 + Na2CO3 ^|BaCO3 4- Na2SO4. The decomposition is incomplete because of the reversibility of the reaction. If, however, the mixture is filtered, and the residue of BaCO3 and unchanged sulphate is boiled with a fresh Na^COg solution, more BaSO4 will be converted to carbonate. By repeating this process a sufficient number of times, one can transform all the sulphate to carbonate. As the carbonate, after thorough washing, is easily soluble in acids, it will be seen that this procedure offers a means of getting an insoluble sulphate into solution. A better and more expeditious method of rendering the re- action complete consists in fusing the sulphate of barium with several times its weight of Na2CO3. Under these conditions, the reaction proceeds to completion in one operation. On cooling the melt, boiling it with water, and filtering, there will remain on the filter a residue of BaCO3 equivalent in amount to the BaSO4 taken; the carbonate is then taken into solution with dilute hydrochloric acid. The method of fusion with alkali carbonate just outlined is of general application and is employed where it is desired to take into solution substances which are insoluble in water and in acids. The sulphates of strontium and calcium, though not as insoluble as that of barium, are sufficiently insoluble to be classed with insoluble substances and may be got into solution by the fusion method. PbSO4, SrSO4, and CaSO4 may be completely converted into carbonate by the first method. 5. Potassium Chromate precipitates from neutral or acetic acid solutions yellow barium chromate : - K2CrO4 + BaCl2 = | BaCrO4 4- 2 KC1. THE METALS 87 The precipitate is practically insoluble in water (i part in 250,000) and in acetic acid (distinction from Sr and Ca), but is soluble in mineral acids. With potassium dichromate zK2Cr2O7) only partial precipitation results : - 2 BaCl2 + K2Cr2O7 + H2O = 12 BaCrO4 + 2 KC1 + 2 HC1. This is due to the formation of HC1, which exerts a solvent action on BaCrO4; the addition of sodium acetate will render the precipitation complete. BaCrO4, like BaSO4, is best precipitated in a boiling solution ; for under these conditions the precipitate is obtained in a form which can be readily filtered and washed, without passing through the pores of the filter. 6. Ammonium Oxalate precipitates from moderately dilute solutions white barium oxalate, somewhat soluble in water (1 part in 2600) and completely soluble in boiling acetic acid (distinction from Ca) : - BaCl2 + (NH4)2C2O4 = | BaC2O4 + 2 NH4C1. 7. Disodium Phosphate precipitates in neutral solutions white flocculent BaHPO4; in ammoniacal solutions, sodium phos- phate throws down the tertiary phosphate. The precipitates are easily soluble in dilute acids, even in acetic acid: - Na2HPO4 4- BaCl2 = 2 NaCl + | BaHPO4; 2 Na2HPO4+3BaCl2 + 2NH3=| Ba3(PO4)2 + 4NaC14-2NH4Cl. 8. Flame Reaction. Barium salts, preferably the chloride, when heated on a platinum wire in the bunsen flame, impart to it an apple-green color; frequently it is yellowish green, due to sodium as an impurity, The reaction becomes more delicate if the wire is first moistened with concentrated HC1. Strontium 1. Ammonium Hydroxide. Same as with Ba salts. 2. Ammonium Sulphide. Same as with Ba salts. 3. Ammonium Carbonate precipitates white SrCO3, more in- soluble in water than BaCO3; in other respects, it possesses about the same solubilities as BaCO3. 88 QUALITATIVE CHEMICAL ANALYSIS 4. Dilute sulphuric acid or any soluble sulphate yields a white precipitate of SrSO4. The precipitate is more soluble in water (1 part in 7000) and in acids than BaSO4, and, as a consequence, is precipitated from very dilute solutions only after some time; it is, however, much less soluble in water than CaSO4, the latter dissolving in water to the extent of 1 part in 500. SrSO4 is practically insoluble in a strong solution of (NH4)2SO4, even on boiling (distinction and method of separation from Ca). 5. Saturated CaSO4 Solution yields with dilute solutions of strontium salts a precipitate of SrSO4, which forms only after some time (distinction from Ba, which yields an immediate precipitate). Precipitation in this case, as well as in 4, is pro- moted by heating, and is retarded by the addition of acids. From concentrated solutions of Sr salts, an immediate precipi- tate is obtained. 6. Potassium Chromate does not yield a precipitate with dilute solutions of strontium salts or with concentrated solutions acid with acetic acid (distinction from and method of separation from Ba). From neutral concentrated solutions, however, a yellow crystalline precipitate of SrCrO4 forms which is soluble in acetic acid. 7. Ammonium Oxalate. Same as with Ba. SrC2O4 is only sparingly soluble in acetic acid. 8. Disodium Phosphate. Same as with Ba. 9. Flame Reaction. Strontium salts, preferably the chloride, when heated on a platinum wire in the bunsen flame, impart to it a deep red color. Calcium i. Ammonium Hydroxide. Same as with Ba. 2. Ammonium Sulphide. Same as with Ba. 3. Ammonium Carbonate precipitates white amorphous CaCO3,"" becoming crystalline on heating; it is more insoluble in water than BaCO3, but in other respects its solubilities are about the same as those for BaCO3. THE METALS 89 4. Dilute Sulphuric Acid or any alkali sulphate does not pro- duce a precipitate from dilute solutions. From concentrated solutions a white precipitate of CaSO4 is obtained which is appreciably soluble in a hot concentrated solution of (NH4)^SO4 (distinction and method of separation from Sr.) A saturated solution of CaSO4, of course, does not precipitate Ca salts (dis- tinction from Sr and Ba). 5. Potassium Chromate does not yield a precipitate from dilute neutral solutions or from concentrated solutions acid with acetic acid. 6. Ammonium Oxalate produces a white crystalline precipi- tate of calcium oxalate immediately from strong solutions and slowly from dilute solutions of calcium salts. The presence of free ammonia, or heating, facilitates the precipitation. The pre- cipitate is practically insoluble in water (1 part in 170,000) and in acetic acid, but is readily soluble in mineral acids: - CaC2O4 + 2 HC1 = CaCl2 + H2C2O4. This is a most delicate test for Ca. 7. Disodium Phosphate gives the same reaction as with Ba. 8. Flame Test. Calcium salts, preferably the chloride, when heated on a platinum wire in the bunsen flame, impart to it a yellowish red color. Outline of the Method of Analysis for Group IV With certain mixtures it is possible, with a little practice, to detect all the metals of this group when occurring together by the simple flame reactions, as the characteristic colors do not all appear at the same time; the latter fact is due to the differ- ence in the volatility of the chlorides. By an analysis of the flame colorations with the spectroscope it is not difficult to de- tect all the alkaline earths, even when they are all present to- gether in the same solution. But as the spectroscopic and flame tests do not distinguish between significant amounts and mere traces due to accidental impurity, they cannot be relied on to determine the composition of an unknown substance. QUALITATIVE CHEMICAL ANALYSIS 90 They are, however, exceedingly valuable as confirmatory tests and for the detection of traces. If the filtrate from Group III., concentrated to a few cc., fails to yield a flame coloration, the absence of Group IV. would be proved, although the reverse would not hold. If the solution to be analyzed for Group IV. is the filtrate from Group III., it will contain a sufficient amount of NH4C1 to prevent the precipitation of Mg along with the alkaline earth carbonates on adding the group reagent, (NH^COg. The precipitated carbonates are dissolved in acetic acid and from this diluted solution the barium is separated from the re- maining metals of this group by precipitation with K2CrO4. After filtering the BaCrO4, the filtrate will contain, besides Sr and Ca, an excess of K2Cr2O7. By reprecipitating the Sr and Ca as carbonates, and filtering, they can be separated from the excess of chromate. If the carbonates are now dissolved in acetic acid, and the resulting solution is boiled with a solution of (NH4)2SO4 and filtered, the Sr will be on the filter as SrSO4, while the Ca will pass into the filtrate. From the latter the calcium may be precipitated as CaC2O4 with (NH4)2C2O4. The separation of Sr and Ca by the use of a boiling solution of (NH4)2SO4 is not complete; some CaSO4 remains undis- solved, while at the same time a small amount of SrSO4 goes into solution. The necessity for making confirmatory flame tests is therefore apparent. SCHEME IV. The filtrate from Group III. which has been acidified with acetic acid, boiled, and filtered from the coagulated sulphur and NiS, as described under Scheme III., should be perfectly clear; if cloudy, it should be boiled again for a few minutes and repeatedly filtered through the same filter until a perfectly clear liquid is obtained; it should then be concentrated by evaporation to about 40 cc.* * Any NH4CI which separates out during the evaporation should be filtered off and rejected. The removal of an unnecessarily large excess of ammonium salts at this point is a decided advantage, because it reduces the amount of material that must be removed by volatilization in the subsequent examination for Group V. THE METALS 91 Preliminary Test for Group IV. To a small portion of the clear filtrate in a test tube add NH4OH to alkaline reaction and then a slight excess of (NH4)2CO3, and warm ; a white ppt. proves the presence of Group IV. If no ppt. is obtained, the absence of more than traces (1) of the alkaline earths is indicated; in that case proceed to Scheme V. If the preliminary test shows the presence of Group IV., the entire filtrate contained in a beaker is made alkaline with NH4OH and is heated nearly to boiling; (NH4)2CO3 is then added in slight excess and the mixture is warmed (but not boiled) (2). The ppt. is allowed to settle and then filtered. The filtrate (3) should be received in a small beaker labelled Group V., covered, and reserved. The ppt. may consist of BaCO3, SrCO3, and CaCO3. Wash once with hot water and reject the washings. Dissolve the ppt. on the filter with die least amount of hot dilute acetic acid (4). Make the volume up to 40-50 cc. by dilution with water (5), heat to boiling, and, while boiling, add K2CrO4 drop by drop till precipitation is complete. Allow the ppt. to settle and filter (using a double filter) by decantation. Finally, with the aid of hot water, bring the ppt. on the filter. Residue is yel- low BaCrO4. Wash twice with hot water and re- ject washings. Confirm by dip- ping clean Pt wire (9) mois- tened with cone. HC1 into ppt., and hold in flame. Do this repeat- edly; a green coloration ap- pearing after some time con- firms the pres- ence of Ba (10). Filtrate (6) may contain Sr and Ca as acetates and K2Cr2O7. To remove the latter, add NH4OH to alkaline reaction and then (NH4)2CO3 till precipitation is complete; heat and filter. If no ppt. forms, the absence of more than traces of Sr and Ca is indicated. Residue is SrCO3 and CaCO3. Reject filtrate. Wash ppt. with hot water until the washings are no longer yellow; reject washings. Dis- solve ppt. on filter in the least amount of hot dil. acetic acid, and dilute the resulting solution with an equal volume of water. Preliminary Test for Sr (7). Pour a very small portion of this solution into a test tube and add a little CaSO4 so- lution, shake vigorously, and allow to stand for a few min- utes. (a) A slowly forming ppt. or cloudiness indicates the presence of Sr; proceed according to (a). (b) No ppt. or cloudiness proves the absence of Sr; proceed ac- cording to (£). (a) If Sr is present, render the remainder of the solu- tion alkaline with NH4OH, add 5 cc. of (NH4)2SO4, boil for a few minutes, and filter. Ppt. is SrSO4. Wash with hot water and confirm by flame test (8). Test filtrate for Ca by adding (NH4)2C2O4. A white ppt. insol. in HC2H3O2 is CaC2O4. Confirm by flame test. (b) Make solution alkaline with NH4OH and add (NH4)2C2O4. A white ppt. insol. in acetic acid is CaC2O4. Confirm bv flame test. 92 QUALITATIVE CHEMICAL ANALYSIS NOTES 1. Tests for traces of alkaline earths. If no precipitate is obtained with (NH4)2CO3,"treat a small portion of the solution with dilute H2SO4, boil, and allow to stand for some time. A white precipitate of BaSO4 proves*the pres- ence of Ba. Treat another small portion with NH4OH until alkaline, add (NH4)2C2O4, heat to boiling, and allow to stand; a cloudiness or white pre- cipitate pfoves the presence of Ca. Traces of Sr are best tested for by mea^s of the spectroscope. 2. The mixture must not be boiled because at the boiling temperature (NH4)2CO3 is decomposed according to the equation - (NH4)2CO3 = f 2 NH3 + H2O + f CO2. The carbonates, which are first thrown down as an amorphous precipitate, are converted by heating and stirring into the crystalline form which can be readily filtered and washed. 3. To insure completeness of precipitation, add a little (NH^aCQg to the filtrate ; if a precipitate forms, add it to the first; if no precipitate forms, pre- cipitation was complete. 4. This is accomplished by placing a test tube under the stem of the fun- nel and pouring on the precipitate a hot mixture of say 5-10 cc. of dilute acetic acid and an equal volume of water, allowing the acid to pass through the filter and pouring the same acid repeatedly through the filter till all the precipitate is dissolved. 5. A preliminary test with K2CrO4 should be made on a small portion of this solution. If a precipitate is obtained, Bajs present and the entire solu- tion should be treated with KoCrO^ as described in the scheme. If no precipi- tate is obtained, Ba is absent: in that case the solution should not be treated with K2CrO4 solution,"Eut on a small portion make a preliminary test for Sr with CaSO4~soIution; if present, treat the remainder of the solution accord- ing to (a) ; if absent, proceed to (£). 6. BaCrO4, even when precipitated in a boiling solution, may pass through ( the filter, yielding a cloudy filtrate; when this is the case, the filtrate must be , boiled again and refiltered. -J 7. It is important to use only a portion of the liquid for the preliminary test. If by mistake the entire filtrate is used, the test for Ca obviously cannot be made. 8. The confirmatory test for Sr is made by moistening the precipitate with concentrated HC1, dipping the wire into it and holding in the flame. A deep red coloration confirms the presence of Sr. If all the Ba had not been com- pletely precipitated as chromate, it will appear here as white BaSO4, but can readily be distinguished from SrSO4 by its failure to yield a red coloration to the flame. THE METALS 93 9. A perfectly clean platinum wire must impart no color to the colorless bunsen flame. When this is not the case, an impurity is indicated, and this mar be removed by one of the following methods: \a) Any large particles of matter adhering to the wire must first be me- chanically removed. The wire is then dipped into concentrated C.P. HC1, contained in a weighing or small specimen tube, and then held in the flame for several seconds. The acid dissolves and thus removes some of the ad- hering material and partly converts some of the still remaining impurity on the wire into chlorides which are volatilized in the flame. Repeat this opera- tion several times; finally dip wire into fresh acid, and hold in flame. It should, if clean, give no color to the flame. Under no circumstances must the wire, clean or otherwise, be dipped into the reagent bottle of acid. The efficiency of this method will be indicated by the fact that the flame coloration becomes noticeably fainter with each treat- ment. (b) Should the above treatment, however, fail to cleanse the wire, it must be dipped while red hot into borax, and heated until a bead forms. By prop- erly manipulating the wire in the flame, the bead can be made to travel back and forth several times over the entire length of the wire. It is then shaken off. Should any solid material then adhere to the wire, it can now be readily removed by scouring with sand. The wire is then treated with concentrated HC1 as described in (a). 10. The Pt wire need only be dipped once into the BaCrO4 precipitate. After the first heating with HC1, it should be moistened with HC1 again and heated. This is to be repeated several times without redipping into the BaCrO4 precipitate. GROUP V. Group V. embraces the metal magnesium, the alkali metals potassium and sodium, and the metallic radical ammonium (NH4). Magnesium is closely allied from an analytical stand- point to the alkaline earths, for its hydroxide, carbonate, and phos- phate are insoluble in water. It has been placed in Group V. for the reason that in the course of complete analysis it will be found in the last filtrate along with the alkali metals. This, of course, is due to the presence of NH4C1 in precipitating the third and fourth groups. The test for NH4 is not made on the final filtrate, because the latter will always contain ammonium salts added in the form of reagents in the regular course of analysis. The test, there- fore, must always be made on a portion of the original 94 QUALITATIVE CHEMICAL ANALYSIS substance. The other alkali metals, lithium, caesium, and rubidium, belong to this group, but have not been included because of their rare occurrence. Omitting consideration of magnesium, which serves as a bridge between groups IV. and V., it may be stated that the chief characteristic of the alkali metals is the fact that nearly all their salts are soluble in water; thus, the chloride, sulphate, sulphide, nitrate, phosphate, oxalate, carbonate, and hydroxide are soluble in water; indeed, their aqueous solutions have been used as reagents. Excluding NH4, all the alkali metals give characteristic flame and spectroscopic reactions. Magnesium Magnesium salts are colorless. With the exception of the hydroxide, carbonate, phosphate, arsenate, and arsenite, all the salts of magnesium are soluble in water. They do not color the bunsen flame. Neither NH4OH, (NH^S, nor (NH^COg precipitates magnesium salts in the presence of a sufficient amount of ammonium chloride, hence the classification of mag- nesium with the alkali metals. i. Ammonium Hydroxide gives a partial precipitation of gelati- nous magnesium hydroxide, Mg(OH)2, readily soluble in am- monium salts : - MgSO4 + 2 NH4OH^:| MgCOH^ + (NH^SO^ The precipitation is only partial in consequence of the forma- tion of an ammonium salt as a by-product of the reaction. The precipitate is readily soluble in acids. The solubility of the precipitate in NH4 salts, or, what amounts to the same thing, the non-precipitation of Mg salts by NH4OH in the presence of a sufficient amount of NH4 salts, is a phenomenon of the same order as that already met with in the cases of ferrous, man- ganese, and zinc compounds. 2. Sodium, Potassium, or Calcium Hydroxide completely pre- cipitates, in the absence of NH4 salts, Mg(0H)2, insoluble in THE METALS 95 excess and nearly insoluble in water, the solubility being i part in 10,000. The washed precipitate is soluble in NH4 salts: - MgSO4 + 2 NaOH = |Mg(OH)^ + Na2SO4. Boiling promotes precipitation. 3 Ammonium Carbonate, in the presence of NH4 salts, gives no precipitate. In solutions containing no NH4 salts, a white basic salt precipitates on standing or on boiling. The composi- tion of the precipitate is variable, depending upon the condi- tions of concentration and temperature: - 4 MgSO4 + 4(NH4)2CO3 + H2O = } Mg^COsWOHX + f CO2 + 4(NH4)sSO4. 4. Disodium Phosphate, when added to a neutral solution of a Mg salt, precipitates flocculent MgHPO4: - Na2HPO4 + MgSO4 = | MgHPO4 + Na2SO4. If, however, NH4C1 and ammonia are added to the solution of a Mg salt before adding the sodium phosphate, a character- istic white crystalline precipitate of ammonium magnesium phosphate forms: - MgCl2 + Na2HPO4 + NH3 = | NH4MgPO4 + 2 NaCl. From dilute solutions the precipitate forms slowly, but it may be hastened by cooling and vigorously stirring the mixture. The precipitate is slightly soluble in water (1 part in 13,500 at 230 C.), but is practically insoluble in 2.5 per cent, ammonia water. It is readily soluble in acetic acid. The addition of NH4C1 prevents the formation of the hydroxide when ammo- nium hydroxide is added. This is a most delicate test for Mg. 5. Ammonium Oxalate gives with dilute solutions of magne- sium salts no precipitate; from concentrated solutions, however, it yields a white precipitate of MgC2O4. The presence of NH4 salts renders the precipitation incomplete. 6. H2S, (NH^S, and H2SO4 do not precipitate magnesium salts. 96 QUALITATIVE CHEMICAL ANALYSIS Potassium With the exception of the acid tartrate, cobaltic nitrite, chlor- platinate, and perchlorate, nearly all the salts of potassium are soluble. i. Hydrochlorplatinic Acid (H2PtCl6) produces in neutral or in concentrated acid solutions of potassium salts a yellow crystal- line precipitate of potassium chlorplatinate (K^PtClg):- 2 KC1 + H2PtCl6 = | K2PtCl6 + 2 HC1. From moderately dilute solutions the precipitate separates out only after standing, but may be hastened by cooling, stirring, or shaking the mixture vigorously in a test tube. This applies to nearly all crystalline precipitates. The precipitate is soluble in alkalies; it dissolves in water to the extent of i part in 100 at 15° C. It is practically insoluble in 8o per cent, alcohol. On ignition it decomposes according to the following equation: - K2PtCl6 = 2 KC1 + | Pt + f 2 Cl2. Solutions of potassium iodide and potassium cyanide do not give this precipitate; they should first be changed to chloride by evaporation with concentrated HC1. As NH4 salts yield a simi- lar precipitate, they must be removed before the test is applied. 2. Tartaric Acid (H2C4H4O6) produces in neutral solutions of potassium salts, which are moderately concentrated, a white crystalline precipitate of potassium acid tartrate: - KC1 + H2C4H4O6 = | KHC4H4O6 + HC1. Precipitation may be hastened by vigorously shaking the mix- ture. The precipitate is soluble in alkalies and in mineral acids. A solution of sodium acid tartrate (NaHC4H4O6) is preferable, as it does not yield any free acid as a by-product: - KC1 + NaHC4H4O6 = | KHC4H4O6 + NaCl. The solubility of the precipitate in water at 15° C. is i part in 222. The reaction is a little more than twice as sensitive as i. Ammonium salts must be absent as they yield a similar pre- cipitate with the reagent. THE METALS 97 3. Sodium Cobaltic Nitrite [Na3Co(NO2)6] yields, with solu- tions of potassium salts acidified with acetic acid, a yellow pre- cipitate of potassium cobaltic nitrite [K3Co(NO2]6. From dilute solutions the precipitation may be hastened by warming the mixture. The precipitate is soluble in water to the extent of i part in 11,000 at 150 C. and is therefore the most sensitive of the reactions mentioned. The test cannot be applied in the presence of NH4 salts for the reason that the latter yield a sim- ilar precipitate. 4. Flame Reaction. Potassium salts, preferably the chloride and nitrate, when heated on a platinum wire in the bunsen flame, impart to it a violet color. The sensitiveness of this reac- tion was given by Bunsen to be 0.001 mg. KC1. The presence of even small amounts of sodium compounds interferes with this reaction by masking the color. If, however, the flame is viewed through several thicknesses of cobalt glass, the latter will absorb the yellow rays and thus permit the violet color to be seen. 5. When heated just below a red heat, potassium chloride is not volatilized (distinction from NH4 salts). Ammonium The ammonium salts very closely resemble the potassium compounds. They have the same crystalline form and, in gen- eral, about the same solubility. 1. Hydrochlorplatinic Acid (H2PtCl6) precipitates, under the same conditions given for potassium compounds (which see), a yellow crystalline precipitate of ammonium chlorplatinate, (NH4)2PtCl6: - 2 NH4C1 + H2PtCl6 = |(NH ^PtClg + 2 HC1. The precipitate may easily be distinguished from the correspond- ing potassium compound by the fact that it is decomposed by an excess of NaOH with the evolution of NH3: - (NH4)2PtCl6 + 2 NaOH = Na2PtCl6 + j 2 NH3 + 2 H2O. When strongly heated, it leaves a residue of platinum sponge only (distinction from the K compound). Ammonium chlorplati- 98 QUALITATIVE CHEMICAL ANALYSIS nate is somewhat less soluble in water than the corresponding K compound; it is insoluble in alcohol. 2. Tartaric Acid (H2C4H4O6) or NaHC4H4O6 gives from con- centrated solutions of ammonium salts a crystalline precipitate of NH4HC4H4O6. The precipitate is soluble in acids and in alkalies, and is very much more soluble in water than the cor- responding K compound; for this reason it is not a good test. It may be distinguished from the corresponding K compound by the evolution of NH3 when treated with an excess of NaOH. 3. Sodium Hydroxide. All ammonium compounds, when heated with an excess of caustic soda, potash, or lime, undergo decomposition with the evolution of ammonia gas; the latter may be detected by its characteristic odor or by its ability to turn moistened red litmus paper blue: - NH4C1 + NaOH = f NH3 + NaCl + H2O ; 2 NH4C1+ Ca(OH)2 = 12 NH3 + CaCl2 + 2 H2O. The evolved NH3 may be further recognized by holding in the escaping vapor a piece of filter paper moistened with mer- curous nitrate solution. Ammonia, if present, will blacken the paper in accordance with the following reaction : - 2 HgNO3 + 2 NH3 = |(NH2HgNO3 + Hg) + NH4NO3. 4. For the detection of minute amounts of ammonia, such, for instance, as are present in drinking water, an alkaline solu- tion of mercuric potassium iodide, known as Nessler's reagent, is used. With this solution, a yellow coloration is obtained which deepens in color, becoming brown with relatively greater amounts. With still greater amounts a brown precipitate is obtained. 2 (2 KI. Hgl2)+ NH3+ 3 KOH = |NHg2I • H2O + 7 KI + 3 H2O. 5. All ammonium salts are volatilized at a temperature just below a red heat (distinction and method of separation from Na and K salts), some undergoing decomposition at the same time. THE METALS 99 Sodium All of the salts of sodium, with the exception of the pyroanti- monate, are soluble in water. i. Potassium Pyroantimonate Solution (K2H2Sb2O7) precipi- tates from neutral or slightly alkaline solutions of sodium salts that are fairly concentrated, a white crystalline precipitate of sodium pyroantimonate (Na2H2Sb2O7). Precipitation may be hastened by shaking the mixture vigorously in a test tube. Sodium pyroantimonate is soluble in boiling water to the extent of i part in 300: - K2H2Sb2O7 + 2 NaCl = 1 Na2H2Sb2O7 + 2 KCL A A A i T A A A I The solution must not be acid, for then decomposition of the reagent results with the precipitation of amorphous pyroanti- monic acid : - K2H2Sb2O7 + 2 HC1 = | H4Sb2O7 + 2 KC1. All other metals, with the exception of K and NH4, must be removed, for they too yield precipitates with the reagent. 2. Hydrochlorplatinic Acid, Tartaric Acid and Sodium Cobaltic Nitrite do not precipitate sodium salts. 3. Heated just below a red heat, sodium compounds are not volatilized (distinction from NH4). 4. Flame Reaction. - Sodium compounds color the bunsen flame yellow even when the quantity is very small. Bunsen and Kirchhoff state that as small an amount as ^00 WW °f a milligram of sodium will give a flame test. To distinguish between a trace and a significant amount, attention must be given to the intensity and duration of the coloration. Outline of the Method of Analysis for Group V. As the special tests for most of the metals of Group V. are not interfered with by the presence of the others, it is needless in most cases to effect their separation before applying the test; thus, the precipitation test for potassium may be made in the 100 QUALITATIVE CHEMICAL ANALYSIS presence of sodium, and the test for Mg may be made in the presence of all the alkali metals, for the latter are not precipitated by Na2HPO4. The test for Mg is therefore carried out on a small portion of about one-third of the filtrate from Group IV. It must, however, be remembered that Na^PC^ precipitates the alkaline earth metals; and as the latter are usually present in small amounts in the last filtrate by reason of the slight solu- bility of their carbonates in NH4C1, they must first be removed before the test can be applied. This is accomplished by adding a little (NH4)2SO4 and (NH4)2C2O4, boiling, and filtering off the precipitated alkaline earths in the form of sulphates and oxalates. The filtrate, after concentration, may then be tested for magnesium. The remaining two-thirds of the original filtrate is used for the detection of Na and K. As this filtrate contains a large amount of ammonium salts accumulated in the course of the analysis, and as the latter interfere with the precipitation tests for K by yielding similar precipitates, it is necessary to remove them before the test for K is made. This is accomplished by taking advantage of the fact that at a temperature just below a red heat, all NH4 salts are volatilized. The NH4 salts removed, the residue is moistened with a little water, and the flame tests are applied. If an intense yellow coloration is obtained which persists for some time, the presence of Na is proved; the flame may then be further examined for potassium by viewing it through several thicknesses of cobalt glass. If a violet-colored flame is obtained in the first place, Na is absent. In either case, a confirmatory test for potassium should be made by any one of the precipitation tests. Ammonium is not tested for in this group for the previously mentioned reason that NH4 compounds in the form of reagents have been added to the solution in the course of the analysis. The test, which consists in liberating NH3 by heating with an excess of NaOH, must, therefore, always be made on a separate portion of the original substance. THE METALS 101 SCHEME V. Analysis of the Filtrate from Group IV. This will contain, besides Mg and the alkalies, traces of the alkaline earths which were dissolved by NH4C1. Divide into two unequal portions. In | test for Mg. Add a few drops of (NH4)2SO4 and (NH4)2C2O4, boil and filter. Reject any ppt. which may form (i). Concentrate the filtrate by evap. to 5 cc. (should any ammonium salts crystallize out, filter and reject). To filtrate contained in a test tube add NH4OH to alkaline reaction and then add Na2HPO4. Shake vigorously and allow to stand several minutes. A white cryst. ppt. which is soluble in acetic acid is NH4MgPO4. In j test for Na and K. Evap. in a large evap. dish, if the vol. of the solution is large, to about 15 cc. Transfer to a small dish and continue to evap. over a wire gauze until sputtering occurs; then cover dish with a watch glass and gently heat until the mass is perfectly dry. By means of a glass rod, transfer to dish particles of salt adhering to the watch glass, place uncovered dish on a pipestem triangle, and ignite (under hood) until no more fumes of NH4C1 are given off, being careful to keep the dish below a red heat and to heat the sides and rim as well as the bottom of the dish. With the aid of a glass rod scrape any salt adhering to the sides of the dish into the center, and stir up salt at the bottom as much as possible; ignite again until no more NH4C1 (2) is given off. Cool. Transfer a small portion of the residue to a watch glass, moisten with a drop of HC1, dip clean Pt wire into it, and hold in flame. A violet coloration proves the presence of K and absence of Na. An intense yellow coloration (3) which persists for some time proves the presence of Na; in that case, examine flame for K either with several thicknesses of cobalt glass or better with a spectroscope (4). Test for K. The remainder of the residue in the dish is dissolved in the least amt. of hot water (3 cc.) just acidified with acetic acid, and is filtered through a small filter. To filtrate in a test tube add a few drops of Na3Co(NO2)6, warm, and allow to stand for several minutes. A yellow ppt. confirms the presence of K. Test for NH4. This test is always made on a small portion of the original substance and never on the filtrate from Groups III. and IV. To about 5 cc. of the original solution contained in a small beaker, add NaOH till the resulting mixture, after thorough stirring, is decidedly alkaline. Heat with stirring (5). Ammonia, if present, will be made evident by its characteristic odor and by its ability to turn blue a piece of red litmus paper held above the beaker. 102 QUALITATIVE CHEMICAL ANALYSIS NOTES ON SCHEME V. 1. Na2HPO4 precipitates all the metals except the alkalies and As, so that the test for Mg can only be made in the absence of these metals. A floccu- lent precipitate of A1PO4 sometimes separates on adding Na2HPO4, but the latter is easily distinguished from NH4MgPO4 by reason of its insolubility in acetic acid. In concentrating the solution to 5 cc., A1(OH)3 may separate out; in that case, it should be filtered off before adding the Na2HPO4. Should a precipitate of doubtful form be obtained with Na2HPO4 and it is desired to confirm the presence of Mg, the precipitate on the filter should be treated with a little acetic acid, the resulting clear filtrate made alkaline with NH4OH, vigorously shaken, and allowed to stand for several minutes. If Mg is present, a white crystalline precipitate will form. It is not customary to remove the Mg before testing for the alkalies. Should it be desired, however, to estimate the amount of K and Na, the Mg may be removed by first expelling NH4 salts by evaporation of the solution to dryness and thoroughly igniting the residue. The latter is then dissolved in water and the Mg is precipitated with Ba(OH)2 solution, filtered, and the Ba in the filtrate removed by precipitation with sulphuric acid, or with NH4OH and (NH4)2CO3. 2. Unless the last traces of NH4 salts are removed, the precipitation test for K will be worthless for the reason that NH4 salts yield a similar precipitate. During the ignition, the bottom of the dish must not be allowed to reach a red heat else there will be danger of volatilizing NaCl and KC1. 3. A fleeting yellow coloration is not to be taken as evidence of the pres- ence of Na. The coloration should persist for at least 6 seconds. 4. The test for K should be confirmed always by a precipitation test. 5. Stirring the mixture while heating is important because of the tendency of the mixture to bump and the consequent danger of having the strongly alkaline mixture spurted into the eyes. Care should be taken in heating and stirring to prevent any of the alkaline liquid from coming in contact with the litmus paper. Failure to observe this precaution will vitiate the test. THE METALS 103 ANALYSIS OF UNKNOWN SOLUTIONS FOR ALL GROUPS To 25 cc. of solution add HC1 and filter. Residue: AgCl, PbCl2, HgCl. Analyze ac- cording to Scheme I. Filtrate contains Groups II.-V. Bring to proper conditions of acidity and pass in H2S. Filter. Residue: HgS, PbS, Bi2S3, CuS, CdS, As2S3, Sb2S3, SnS. Wash and treat with (NH4)2S*. Filter. Analyze residue ac- cording to Scheme II. A. Analyze filtrate ac- cording to Scheme II. B. Filtrate: Boil to expel H2S before proceeding with the analysis of Group II. ppt. Add NH4C1, make alkaline with NH4OH, and completely ppt. with H2S or (NH4)2S ; filter. Residue: A1(OH)3, Cr(OH)3, FeS, NiS, CoS, MnS, ZnS. Analyze ac- cording to Scheme III. Filtrate is at once acidified with acetic acid and boiled to expel H2S, and then filtered. Make clear filtrate alkaline with NH4OH, add (NH4)2CO3, and filter. Residue: BaCO3, SrCO3 -f- CaCO3. Analyze ac- cording to Scheme IV. Filtrate will con- tain Mg and the " al- kalies." Divide into two portions : | test for Mg according to Scheme V; | test for Na and K accord- ing to Scheme V. Test for NH4 is made on a separate portion according to Scheme V PART II THE ACIDS It is customary to arrange the acids into groups according to their deportment with two reagents, viz., BaCl2 and AgNO3. Three groups are thus distinguished. Group I. includes those acids whose barium or calcium salts are insoluble in water; they are therefore precipitated from neutral solutions by BaCl2. This group is divided into two parts, namely: (a) Acids precipitated by BaCl2 from solu- tions acid with HC1, viz., sulphuric acid (H2SO4) and hydrofluo- silicic acid (H2SiF6); (£) Acids precipitated by BaC^ from neutral solutions only, viz., carbonic acid (H2CO3), sulphurous acid (H2SO3), thiosulphuric acid (H2S2O3), phosphoric acid (H3PO4), hydrofluoric acid (HF), oxalic acid (H2C2O4), boric acid (H3BO3),* silicic acid (H2SiO3),f tartaric acid (H2C4H4O6), arsenic acid (H3AsO4), arsenious acid (H3AsO3), and chromic acid (H2CrO4). Group II. includes those acids whose barium salts are soluble in water, but whose silver salts are insoluble in nitric acid; they are therefore precipitated by AgNO3 from solutions acid with nitric acid. These acids follow : - Hydrochloric acid (HC1), hydrobromic acid (HBr), hydriodic acid (HI), hydrocyanic acid (HCN), hydrogen sulphide (H2S), hydroferrocyanic acid [H4Fe(CN)6], hydroferricyanic acid [H3Fe(CN)6], thiocyanic acid (HSCN), and nitrous acid (HNO2); the last is, however, only precipitated from moderately * Orthoboric acid (H3BO3) is here considered as representative of the acids of boron. j Metasilicic acid (HjSiOs) is taken as the type of the many silicic acids. 104 THE ACIDS 105 concentrated solutions and is therefore also placed in the next group. Group III. includes those acids that are not precipitated by either BaCl2 or AgNO3, viz., nitric acid (HNO3), nitrous acid (HNO2), chloric acid (HC1O3), and acetic acid (HC2H3O2). Sulphates With the exception of the sulphates of lead, barium, strontium, and calcium, all normal sulphates are soluble in water. Silver and mercury (-ous) sulphates are, however, difficultly soluble. Nearly all basic sulphates are insoluble, but readily dissolve in hydrochloric or nitric acids. Alkali and alkaline earth sulphates are not decomposed when ignited gently in a closed tube, but at higher temperatures more or less decomposition takes place. The behavior of other sulphates on being heated varies with the nature of the metal with which the SO4 radical is united, some resisting decomposition, while others are readily decomposed, giving off SO3 or SO2, or both, and leaving the oxide of the metal. Free sulphuric acid is recognized even in the presence of a sulphate by its property, when concentrated, of removing the elements of water from organic substances and leaving a charred residue; thus, if to a solution containing free H2SO4 a little cane sugar is added and the mixture is evaporated just to dry- ness, preferably on a steam bath, a black residue will be ob- tained. i. Barium Chloride precipitates white BaSO4,* insoluble in water and in dilute acids even on boiling. From dilute solutions a precipitate separates only on standing. Dilute HC1 or HNO8 * From HC1 solutions BaCL may also precipitate BaSeO4 and BaSiF6 ; the former is readily distinguished from BaSO4 and BaSiF6 by the fact that on boiling it with concentrated HC1, chlorine is given off : - BaSeO4 4- 4 HC1 = BaC12 4- H2SeO3 4- CI2 4- H2O. BaSiF6 is easily recognized by its readiness to undergo decomposition when heated with concentrated H2SO4 : - BaSiF6 4- H2SO4 = BaSO4 4- f 4HF 4-1 SiF4. 106 QUALITATIVE CHEMICAL ANALYSIS should be added before the reagent in order to prevent the pre- cipitation of chromates, sulphites, and carbonates. Strong acid must not be used, for otherwise a crystalline precipitate of BaCl2 or Ba(NO3)2 may be obtained; these are, however, easily dis- tinguished from BaSO4 by their ready solution on diluting. For further properties of BaSO4, see reaction 4 under Barium. 2. Lead Acetate produces a white precipitate of PbSO4, soluble in a hot concentrated solution of ammonium acetate or tartrate. 3. Mixed with Na2CO3 free from sulphur compounds, and heated on charcoal with the reducing flame of a blowpipe, all sulphates are reduced to sulphides. If the fused mass is placed on a silver coin and then moistened with a drop of water, a black stain of Ag2S will be produced : - Na^ 4- 2 Ag + H2O + O = Ag2S + 2 NaOH. This test is also given by other sulphur compounds. Fluosilicates Hydrofluosilicic acid (H2SiF6) is formed by the action of sili- con tetrafluoride (SiF4) on water: - 3 SiF4 + 4 H2O = 2 H2SiF6 + | H4SiO4. If the silicic acid, which is formed at the same time, is filtered off, the filtrate will contain an aqueous solution of hydrofluo- silicic acid. Both the acid and its salts are decomposed on heat- ing. On evaporating a solution of H2SiF6, decomposition sets in, according to the equation - H2SiF6 = f SiF4 4-f 2 HF. With the exception of the potassium and barium salts, nearly all fluosilicates are soluble in water. BaSiF6, formed by adding BaCl2 to a solution of a fluosilicate, is a white, crystalline, insoluble substance. Its solubility at 170 C. is 1 part in 3700 parts of water. As it is sparingly solu- ble in HC1, it can be precipitated by BaC^ from solutions con- taining this acid. For a method of distinguishing it from BaSO4, see footnote, page 105. THE ACIDS 107 Carbonates Carbonic acid (H2CO3) is a weak dibasic acid which is only- known in solution. With bases it yields an important class of stable salts known as carbonates. On ignition, the carbonates of calcium and strontium are decomposed, while the normal alkali carbonates are but slightly affected ; ammonium carbon- ate volatilizes on heating. Nearly all of the carbonates are white, and, with the exception of the carbonates of the alkali metals, all the normal salts are insoluble in water. The aqueous solutions of the carbonates and dicarbonates of the alkalies possess an alkaline reaction. Sodium dicarbonate, on ignition, is changed to the normal salt with evolution of carbon dioxide and water: - 2 NaHCOg = Na2CO3 4- f H2O + f CO2. i. All the acids, excepting HCN, decompose carbonates with effervescence, due to the evolution of CO2; the latter may be recognized by its property of rendering turbid a drop of lime- water held in the escaping gas: - CaCO3 + 2 HC1 = CaCl2 4- H2O + f CO2; CO2 4- Ca(OH)2 = | CaCOg + H2O. This constitutes the chief reaction for carbonates from an analytical standpoint. 2. Barium or Calcium Chloride, when added to a solution of a normal carbonate, gives a white precipitate of BaCO3 or CaCOg. The precipitate is soluble in carbonic acid as well as in all other acids with the exception of HCN : - CaCO3 + H2CO3 = Ca(HCO3)2. From the solution of dicarbonate of calcium, CaCO3 reprecipi- tates on boiling: - Ca(HCO3)2 = | CaCO3 4- H2O 4- f CO2. 3. Silver Nitrate precipitates white silver carbonate (Ag2CO3), which, on boiling, changes to brown silver oxide (Ag2O): - Ag2CO3 = | Ag2O 4- f CO2. 108 QUALITATIVE CHEMICAL ANALYSIS Sulphites The aqueous solution of sulphur dioxide, known as sulphur- ous acid, is a weak dibasic acid. When boiled it decomposes, giving off SO2, which may be easily recognized by its odor. Neutralized by bases, it forms sulphites, all of which are insolu- ble or nearly so in water with the exception of those of the alkali metals. The solid salts, as well as their aqueous solu- tions, readily oxidize on exposure to air, forming the correspond- ing sulphates. i. Barium Chloride precipitates from neutral solutions white barium sulphite (BaSO3), readily soluble in HC1 and HNO3. In practice, a residue of BaSO4 remains, due to the presence of a small amount of sulphate originally present in the sulphite or produced by the subsequent oxidation of the sulphite. If the BaSO4 is filtered off, the clear filtrate may be shown to contain sulphurous acid by adding a little bromine or concentrated HNO3 and boiling, when a white precipitate of BaSO4 will form. The bromine or nitric acid oxidizes the sulphurous to sulphuric acid, and the latter at once yields with the BaC^ present a precipitate of BaSO4. Free sulphurous acid is not precipitated by BaCl2. 2. Hydrogen Sulphide, when passed into a solution of sul- phurous acid, or a solution of a sulphite, acid with HC1, causes a separation of sulphur with the formation at the same time of pentathionic acid : - 5 H2S -I- 5 H2SO3 = H2S6O6 + | 5 S + 9 H2O. 3. Dilute R2SO4 or HC1, when added to a sulphite, decomposes it with the evolution of SO2: - Na2SO3 4- 2 HC1 = 2 NaCl + H2O + f SO2. 4. Silver Nitrate precipitates from neutral solutions white Ag2SO3, which, on boiling, is decomposed with the separation of gray metallic silver : - Ag2SO3 + H2O = 12 Ag + H2SO4. THE ACIDS 109 5. Iodine Solutions are bleached by sulphurous acid; this is due to the reduction of iodine to hydriodic acid: - H2SO3 + I2 + H2O 2 HI + H2SO4. 6. Potassium Permanganate solution, acid with sulphuric acid, is also decolorized by sulphurous acid : - 2 KMnO4 + 5 H2SO3 = K2SO4 + 2 MnSO4 + 2 H2SO4 + 3 H2O. Since sulphur dioxide, which may be liberated from a sulphite (see reaction 3), is much heavier than air, it may be decanted into another test tube containing a small amount of exceedingly dilute KMnO4 solution. Now, on thoroughly mixing the gas and permanganate solution, the latter will be bleached. 7. Potassium Dichromate, when added to sulphurous acid, is reduced to a chromic salt; the reduction is accompanied by a change in color to green : - K2Cr2O7 + 3 H2SO3 + H2SO4 = K2SO4 + Cr2(SO4)3 + 4 H2O. 8. Stannous Chloride, when added to sulphurous acid or to a hydrochloric acid solution of a sulphite, and the mixture is heated, reduces the sulphurous acid to H2S, which, after some time, will precipitate the tin as stannic sulphide: - 3 SnCl2 + H2SO3 + 6 HC1 = f H2S + 3 SnCl4 + 3 H2O ; SnCl4 + 2 H2S = | SnS2 + 4 HCL 9. Sulphurous acid reduces arsenic acid (H3AsO4) to arseni- ous acid (H3AsO3); the reaction is preferably conducted in a closed pressure bottle heated in a water bath. It also reduces ferric to ferrous salts. 10. Sulphites, when heated with sodium carbonate on char- coal, behave in the same way as sulphates (see reaction 3). Thiosulphates Thiosulphuric acid is unknown, for, when liberated from its salts, it at once breaks down into SO2, S, and H2O. The chief thiosulphate is the sodium salt (Na2S2O3), used extensively in photography because of its property of dissolving the halides of silver. 110 QUALITATIVE CHEMICAL ANALYSIS i. Dilute Hydrochloric Acid decomposes all thiosulphates with the evolution of sulphur dioxide and the separation of sulphur: - Na^Og + 2 HC1 = 2 NaCl + H2O + f SO2 + | S. This reaction serves to distinguish this class of salts from sul- phites, which do not give a separation of sulphur when so treated. 2. Silver Nitrate precipitates white silver thiosulphate, easily soluble in an excess of sodium thiosulphate with the formation of a complex salt: - 2 AgNO3 + Na2S2O3 = | Ag2S2O3 + 2 NaNOg; Ag2S2O3 + Na2S2O3 = 2 NaAgS2O3. Boiling decomposes the double salt with the separation of silver sulphide and sulphur. Silver thiosulphate, almost as soon as formed, owing to its instability, becomes yellow, then brown, and finally black, with the formation of Ag2S : - Ag2S2O3 + H2O = | Ag2S + H2SO4. Phosphates Three phosphoric acids are known, viz., orthophosphoric acid (H3PO4), metaphosphoric acid (HPO3), and pyrophosphoric acid (H4P2O7). The most stable in solution, as well as the most important of these, is the ortho- acid. The others are converted into this form by boiling with water, for example: - HPO3+H2O = H3PO4. Metaphosphoric acid, or the acetic acid solution of a meta- phosphate, is distinguished from the other two by its charac- teristic property of coagulating albumen. Pyrophosphoric acid and its salts are formed by heating the ortho- acid or its mono-hydrogen or mono-ammonium salts; thus:- 2 H8PO4 = f H2O + H4P2O,; 2 Na2HPO4 = Na4P,O, + f H,O; 2 NH4MgPO4 = f H2O + f 2 NH3 + Mg2P2O7. THE ACIDS 111 With the exception of the alkali salts, all normal pyro- and ortho-phosphates are insoluble in water. All of the phosphates of the metals of Group IV. are soluble in acetic, hydrochloric, and nitric acids. Those of the trivalent metals of Group III. are insoluble in acetic but soluble in hydrochloric acid. i. Barium Chloride (BaCl2) precipitates from neutral solutions of orthophosphates white BaHPO4, soluble in acetic, hydro- chloric, and nitric acids; from the acid solution, ammonia pre- cipitates the tertiary phosphate : - BaCl2 + Na2HPO4 = | BaHPO4 + 2 NaCl; BaHPO4 + 2 HC1 = BaCl2 + H3PO4 ; H3PO4 + 3 NH3 = (NH4)3IjO4; 2(NH4)3PO4 + 3 BaCl2 = | Ba3(PO4)2 + 6 NH4C1. 2. Lead Acetate [Pb(C2H3O2)2] precipitates white lead phos- phate, practically insoluble in acetic acid though soluble in nitric acid: - 3 Pb(C2H3O2)2 + 2 Na2HPO4 = + 2 HC2H3O2+ 4 NaC2H3O2. 3. Silver Nitrate precipitates only from strictly neutral solu- tions yellow silver phosphate (Ag3PO4), soluble in mineral acids ;* it is also soluble in acetic acid and in ammonium hydroxide. 2 Na2HPO4 + 3 AgNO3 = | Ag3PO4 + 3 NaNO3 + NaH2PO4. 4. Magnesium mixture(MgCl2 + NH4Cl + NH4OH in slight ex- cess) precipitates from solutions of orthophosphates, white crystal- line ammonium magnesium phosphate (NH4MgPO4-6 H2O): - MgCl2 4- Na2HPO4 + NH3 = | NH4MgPO4 + 2 NaCl. The precipitate is soluble in acids, including acetic acid; from solutions of the latter ammonium hydroxide reprecipitates the double phosphate (method of purification and separation from A1PO4, which is insoluble in acetic acid). The precipitate is slightly soluble in water, but is insoluble in 2.5 per cent, ammonia water. From very dilute solutions, the precipitate * With HC1, a white precipitate of AgCl is obtained. 112 QUALITATIVE CHEMICAL ANALYSIS separates only on standing; on ignition, it yields magnesium pyrophosphate (Mg2P2O7). For further details concerning the precipitate, see reaction 4 under Magnesium. Arsenates yield a similar precipitate, but the latter is rendered reddish brown on treatment with silver nitrate, due to its conversion into silver arsenate (Ag3AsO4). 5. Ammonium Molybdate, when added to a warm nitric acid solution of a phosphate, yields a canary-yellow precipi- tate of ammonium phosphomolybdate of variable composition [(NH4)3PO4, • 12 MoO8]. Precipitation may be hastened by heating and by having an excess of ammonium nitrate present. The precipitate is soluble in excess of phosphoric acid or of alkali acid phosphate; hence, to secure complete precipitation, a large excess of the reagent is necessary. It is also soluble in alkalies, including ammonium hydroxide. Arsenic acid yields a similar compound with this reagent, but the precipitate forms slowly and requires a higher temperature for its complete precipitation. This reaction affords a means of quantitatively separating phosphoric acid from the metals with which it is combined. 6. Ferric Chloride, when added to a soluble phosphate not too strongly acid, yields a buff-colored precipitate of ferric phos- phate (FePO4): - FeCl3 + Na2HPO4 = | FePO4 + 2 NaCl + HC1. As ferric phosphate is soluble in hydrochloric acid, the pre- cipitation by ferric chloride in the above reaction is never com- plete ; by adding an excess of sodium acetate, however, the hydrochloric acid is replaced by acetic acid, in which FePO4 is insoluble, and, as a consequence, precipitation is rendered complete: - FeCl3 + Na2HPO4 + NaC2H3O2 = 3 NaCl +1 FePO4 + HC2H3O2. This reaction is utilized in removing phosphoric acid from solutions. 7. Metallic Tin, when added to a nitric acid solution of a phosphate, precipitates the phosphate as stannic phosphate, the THE ACIDS 113 excess of the tin separating at the same time as insoluble meta- stannic acid. This reaction, like the one above, may be em- ployed for the separation of the PO4 radical from the metals. Fluorides Hydrogen fluoride, the water solution of which is hydrofluoric acid, at 19.5 ° C. is a colorless, highly corrosive liquid. Its aque- ous solutions attack the skin, producing painful sores; it must therefore be handled with care. As its chief property is its ability to etch glass, it must be kept in ceresine, hard rubber, or platinum vessels. With bases, it yields salts, known as fluorides. The fluorides of the alkali metals, with the exception of that of lithium, are soluble in water. Those of the alkaline earth group are either insoluble or sparingly soluble in water. The fluorides of Cu, Pb, Zn, and many other heavy metals, are only slightly soluble, while those of Al, Ni, Co, Ag, Sb, and Sn(-ous) dis- solve readily. 1. Concentrated H2SO4 decomposes most fluorides with the liberation of hydrogen fluoride: - (1) CaF2 + H2SO4 = |CaSO4 + f 2 HF. The reaction proceeds more rapidly if the mixture is heated. The HF may be recognized by its ability to etch glass (see re- action 2). If the reaction is carried out in a test tube, the hydrofluoric acid which is set free attacks the glass, with the evolution of silicon tetrafluoride: - (2) Na2Si3O7 • CaSi3O7 4- 28 HF = soda glass. ' 2 NaF + | CaF2 + f 6 SiF4 + 14 H2O. When brought in contact with water, silicon tetrafluoride reacts according to the following equation : - (3) 3 SiF4 + 4 H2O = 2 H2SiF6 + | H4SiO4. The above three reactions may be utilized in testing for a fluo- ride. One need only heat the substance in a test tube with concentrated H2SO4 and then hold in the escaping vapors a 114 QUALITATIVE CHEMICAL ANALYSIS drop of water held on the loop of a platinum wire. In the pres- ence of a fluoride, the " water bead " will become turbid owing to the formation of gelatinous silicic acid. The test may also be carried out in lead or platinum vessels if the fluoride is first intimately mixed with ignited silica (SiO2) before the treatment with concentrated H2SO4. The reactions are as follows : - 2 CaF2 + 2 H2SO4 = 12 CaSO4 + f 4 HF; 4 HF + SiO2 = f SiF4 + 2 H2O ; 3 SiF4 + 4 H2O = 2 H2SiF6 + H4SiO4. 2. The Etching Test. This test, as explained above, is based on the property possessed by hydrofluoric acid to dissolve SiO2 or glass. In a platinum crucible or lead dish, mix, with the aid of a piece of wood, some of the powdered substance with concen- trated sulphuric acid. Cover with a watch glass that has been coated on the convex side with paraffine and through which some characters have been scratched. Put a little water on the upper concave side of the watch glass to prevent the par- affine from melting during the heating, and gently heat the crucible or dish, preferably on a water bath. After some time, remove the watch glass, warm it, and wipe the paraffine off. If a fluoride is present in the substance under examination, the glass will be corroded or etched in those places where the glass has been exposed to the liberated hydrofluoric acid. It is evident that this test is inapplicable in the presence of silicates. Anhy- drous HF does not etch glass. 3. Calcium Chloride, added to an aqueous solution of a fluo- ride, gives a white gelatinous precipitate of calcium fluoride (CaF2), soluble with difficulty in HC1 and HNO3, but practi- cally insoluble in acetic acid. From the acid solution of calcium fluoride, ammonium hydroxide does not reprecipitate the fluoride, because of the solubility of CaF2 in ammonium salts. 4. Fusion with Sodium Carbonate only partially decomposes CaF2. In the presence of silica, however, the decomposition may be rendered complete. THE ACIDS 115 Oxalates With the exception of those of the alkali metals, magnesium and chromium, nearly all the oxalates are insoluble or sparingly soluble in water; they all are soluble in mineral acids and in many cases in an excess of alkali oxalate with the formation of double salts. i. Barium Chloride precipitates white barium oxalate (BaC2O4), soluble in oxalic and acetic acids. 2. Calcium Sulphate or Calcium Chloride precipitates white crystalline CaC2O4, insoluble in oxalic and acetic acids, and in ammonium oxalate, but soluble in hydrochloric acid. As CaC2O4 is one of the most insoluble of the oxalates, CaSO4 is an excellent reagent for the detection of this acid. 3. Concentrated Sulphuric Acid, when added to an oxalate in the solid state, decomposes it with the evolution of CO and CO2: - H2C2O4 + H2SO4 = H2O + H2SO4 + f CO +1CO2. If the mixed gases are passed through limewater or sodium hydroxide, the CO2 will be absorbed and the escaping CO may be recognized by the characteristic blue flame with which it burns. 4. Potassium Permanganate Solution, when added to a hot sulphuric acid solution of an oxalate, is bleached because of its reduction to a manganous salt, the oxalic acid being oxidized at the same time to CO2 and water : - 2 KMnO4 4- 3 H2SO4 + 5 H2C2O4 = K2SO4 + 2 MnSO4 + 8 H2O + f 10 CO2. 5. Behavior of Oxalates on Ignition. At a red heat, all oxa- lates are decomposed with the evolution of CO and CO2. The oxalates of the alkalies and alkaline earths are converted by ignition into carbonates with little or no carbonization. Mag- nesium oxalate yields MgO when heated. All other oxalates leave either a residue of metal or an oxide, depending upon the ease with which the oxide is reduced. 116 QUALITATIVE CHEMICAL ANALYSIS Rapid decomposition of an oxalate may be accomplished by fusing it in a porcelain crucible with NH4NO3; any organic matter present will also be destroyed at the same time. Borates Boron forms three acids, namely, orthoboric acid (H3BO3), metaboric acid (HBO2), and tetraboric acid (H2B4O7); the last two may be obtained from the first by careful heating. Salts of the ortho-acid hydrolyze in a water solution, so that the salts of boric acid we are concerned with are either of the meta- or pyro- type. Borates of the alkali metals alone are soluble in water, yielding solutions which have an alkaline reaction owing to par- tial hydrolysis; all other borates are either insoluble or spar- ingly soluble in water, but are readily soluble in mineral acids and in ammonium salts. i. Turmeric Paper Test. If a piece of turmeric paper is dipped into a solution of a borate slightly acid with HC1, and the paper is then dried by placing it on a watch glass and heat- ing the latter on a water or steam bath, the paper assumes a reddish brown color. If the paper is now moistened with a drop of caustic soda solution, the color changes to a greenish black. 2. Barium Chloride precipitates from concentrated solutions flocculent barium metaborate, soluble in excess of barium chlo- ride, in ammonium chloride, and in acids : - Na2B4O7 + BaC^ + 3 H2O = 2 NaCl + 2 H3BO3 + | Ba(BO2)2. 3. Calcium Chloride gives with borates reactions precisely similar to those produced by BaCl2. 4. Silver Nitrate precipitates from cold concentrated solutions white silver metaborate (AgBO2), soluble in ammonium hydrox- ide and nitric acid; warming converts it into brown silver oxide (Ag2O). 5. Flame Tests. Free boric acid and some of its volatile compounds, e.g., the fluoride (BF3), its esters, as (CH3)3BO3 and THE ACIDS 117 (C2H5)3BO3, when brought into the bunsen flame, impart to it a characteristic green color. 6. Concentrated Sulphuric Acid, when added to a borate, de- composes it with the liberation of free boric acid (H3BO3). The test depending upon this and the above reactions is carried out by making a paste of the substance with concentrated H2SO4, taking up some of the mixture on the loop of a platinum wire, and holding it in the flame, when, if a borate is present in the substance being examined, the characteristic green color will be observed. This test does satisfactorily for most borates. For silicates containing boron, which are not decomposed by con- centrated H2SO4, it is necessary to mix the mineral with a little calcium fluoride before adding the sulphuric acid; under these conditions, volatile boron fluoride (BF3) is formed, which, when brought into the bunsen flame, colors it green. 7. Concentrated H2SO4 and Alcohol. If concentrated sulphuric acid is added to a borate and then a little methyl or ethyl alco- hol, and the mixture is stirred and lighted, the resulting flame will be found to be green at its borders, due to the formation of volatile methyl or ethyl borate : - H3BO3 + 3 C2H5OH 3 H2O + 4 (C2H5)3BO3. The concentrated H2SO4 performs the double function of liberating the boric acid and absorbing the water formed in the above reversible reaction, thus causing the latter to proceed from left to right. 8. Behavior on Ignition. Borates of the alkali metals, when heated, swell up (escaping of water), and finally fuse with the formation of a colorless glass; the latter, possessing as it does an excess of acid oxide, readily unites with metallic oxides on heating, with the formation of metaborates (borax beads) having characteristic colors. Thus, with cobalt oxide or any cobalt compound we obtain CoO + Na2B4O7 = 2 NaBO2 + Co(BO2)2. The cobalt metaborate is blue. 118 QUALITATIVE CHEMICAL ANALYSIS Silicates Silica (SiO2), the anhydride of silicic acid, occurs abundantly in a more or less pure state in nature as quartz, rock crystal, flint, agate, sand, etc. It is insoluble in water and in all acids with the exception of hydrofluoric acid; the latter dissolves it with the formation of gaseous silicon tetrafluoride (SiF4): - SiO2 4- 4 HF = j SiF4 + 2 H2O. To expel silica completely with HF, the presence of concen- trated H2SO4 is necessary. When silica is fused with sodium carbonate and the mass is extracted with water, a solution of sodium silicate is obtained : - SiO2 + Na2CO3 = Na^SiOg + f CO2. All silicates are insoluble in water with the exception of those of sodium and potassium, which are soluble in water in the presence of free alkali. 1. If to a solution of sodium or potassium silicate, hydro- chloric or nitric acid is added until an acid reaction results, part of the silicic acid will separate out as a gelatinous precipitate, while the rest will remain in solution in the form of a hydrosol: - Na2SiO3 + 2 HC1 = 2 NaCl + X H2SiO3. <00 J 4 O If the sodium silicate solution is very dilute, the silicic acid set free may remain entirely in solution. Precipitated silicic acid is somewhat soluble in water and in acids, and is readily soluble in alkali hydroxides and carbonates. If an acid solution of an alkali silicate, which may contain more or less of precipitated silicic acid in suspension, is evapo- rated to dryness on a boiling water bath, the precipitated, as well as the dissolved, silicic acid loses water and becomes insolu- ble in acids; therefore, on extracting the dried residue with hy- drochloric acid and filtering, practically all the silicic acid (about 99 per cent.) will be left on the filter in a more or less dehy- drated state. The more complete the dehydration, the more insoluble does the resulting silicic acid become. This property of silicic acid of becoming insoluble in acids on dehydration is THE ACIDS 119 of great analytical importance, since it enables the analyst, at the outset of the analysis, to completely separate the silicic acid from the metals with which it was originally combined. 2. Ammonium chloride or ammonium carbonate, when added to a solution of alkali silicate, causes a precipitation of meta- silicic acid: - Na2SiO3+2 NH4Cl + 2 H2O = | H2SiO3+2 NaCl +2 NH4OH ; Na2SiO3+(NH4)2CO3+2 H2O = |H2SiO3+Na2CO3 + 2 NH4OH. 3. Barium Chloride precipitates white BaSiO3, soluble in acids. 4. Silver Nitrate precipitates yellow Ag2SiO3, soluble in acids and in ammonium hydroxide. 5. Ammonium Molybdate solution containing an excess of nitric acid yields, with solutions of silicates, a yellow solution. On heating in the presence of much NH4C1, a canary-yellow precipitate is produced. 6. Sodium Metaphosphate Bead Test. When a silicate is fused in a metaphosphate bead prepared from microcosmic salt (NH4NaHPO4), the bases are dissolved to a transparent bead, while the silica in the form of a " skeleton " of the original mass remains undissolved as an opaque body : - CaSiOo + NaPOo = CaNaPO4 + SiO2. Treatment of Insoluble Silicates (a) Silicates decomposed by Acids (not including hydrofluoric acid). The finely ground silicate * is heated in a casserole with * Powdering Minerals. Since substances are more readily soluble in a state of fine powder than in the form of lumps, the process of powdering is always resorted to before the analysis of minerals, slags, and ores is begun. This is accomplished by first wrapping up in a clean towel a number of selected specimens of the substance, placing the latter on a plate of hard steel and breaking them up with several sharp blows from a hard steel hammer. The resulting mixture of powder and coarse parti- cles is then introduced into a diamond steel mortar, in small portions at a time, and crushed to a coarse powder; the latter is then thoroughly mixed, and about two grams or more are reduced to an extremely fine state of division by grinding in an agate mortar until the entire quantity passes through a 100-mesh sieve. 120 QUALITATIVE CHEMICAL ANALYSIS concentrated HC1* and is boiled until decomposition is com- plete. As a result of the action of the acid, part of the silicic acid separates in the gelatinous form. The mixture is then evaporated to dryness, preferably on a steam bath, and the silica is completely dehydrated by heating the nearly dry mass in an air oven maintained at 1200 C. until no more acid fumes are given off. The dried residue is first treated with 10 cc. of concentrated HC1 and heated to dissolve the bases, some of which may have been rendered difficultly soluble by the dehy- dration process; water is then added, the mixture is heated again with stirring, and finally is filtered. The filtrate will con- tain all of the metals in the form of chlorides. The residue will contain practically all of the silicic acid and may be con- taminated with small amounts of the bases, chiefly iron and aluminum as oxides. To test the purity of the silicic acid, the precipitate and a portion of the filter retaining the precipitate are placed in a platinum crucible, moistened with a little con- centrated ammonium nitrate solution to facilitate the combus- tion of the filter, and ignited until all of the paper is consumed. The crucible is then carried to the hood, placed on a pipestem triangle, and is treated with a few drops of concentrated H2SO4 and about 5 cc. of hydrofluoric acid. It is then gently heated to dryness and finally ignited. By this treatment all the silicic acid will be driven off as SiF4. Any residue is treated accord- ing to page 172 for insoluble substances, f (£) Silicates which are slightly or not attacked appreciably by acids in the above treatment are decomposed by fusing them in a platinum crucible, unless reducible metals are shown to be present, with six times their weight of a mixture of sodium and potassium carbonates until the mass is in a state of quiet fusion. By this treatment, the silica will be converted into sodium silicate, while the bases will be variously converted into car- bonate, oxide, or metal, depending upon their nature. After cooling, the crucible with its contents is placed in a casserole or * When metals of the first group are known to be present, it is preferable to use nitric acid. t The residue may consist of AI2O3, Fe2O3, or BaSO4, or of dll three substances. THE ACIDS 121 evaporating dish, and is treated with an excess of hydrochloric acid.* The latter takes the bases into solution and at the same time decomposes the alkali silicate with a partial precipitation of silicic acid. The mixture is then evaporated to dryness, dehydrated at 1200 C., extracted with hot concentrated acid, diluted with water, and the silicic acid is filtered off and purified as already outlined. The filtrate is examined for bases and acids except those removed by the above treatment (see page 148). (4) Decomposition of Silicates with Hydrofluoric Acid. Under a hood, treat about one gram of the finely powdered mineral contained in a platinum crucible or dish with about 10 cc. of hydrofluoric acid and a few drops of concentrated H2SO4, and evaporate on a steam bath to dryness. By this procedure, the silicate will be decomposed and the metals are left as sulphates. Treatment for the Detection of Alkalies in Silicates. If undecomposed by acids, except hydrofluoric, apply J. Lawrence Smith's method, which is given under the head of " Insoluble Substances." The hydrofluoric acid method is also suitable for the detection of the alkalies in insoluble silicates. Tartrates Solubilities. The normal tartrates of the alkali metals, as well as those of aluminum and ferric iron, are soluble in water; nearly all others are insoluble in water, but are soluble in hy- drochloric and nitric acids, and for the most part are soluble in an excess of alkali tartrate with the formation of double salts. 1. Concentrated Sulphuric Acid, when added to a tartrate and the mixture is heated, causes a charring with the evolution of SO2. 2. Silver Nitrate does not react with free tartaric acid, but from solutions of normal tartrates it precipitates white curdy silver tartrate (Ag2C4H4O6), readily soluble in nitric acid and ammonium hydroxide. If the tube containing the ammoniacal solution of silver tartrate is heated gently, preferably in a boil- * If metals of Group I. are known to be present, it is preferable to use nitric acid. 122 QUALITATIVE CHEMICAL ANALYSIS ing water bath, a deposit of silver forms on the sides of the tube in the form of a mirror. 3. Calcium Chloride, added in excess to a concentrated solution of a neutral tartrate, precipitates white crystalline calcium tar- trate (CaC4H4O6), soluble in acids including acetic acid; the precipitate is soluble in cold caustic soda free from carbonate, from which, on boiling, CaC4H4O6 reprecipitates in a gelatinous form which redissolves on cooling. The precipitation of cal- cium tartrate is interfered with by the presence of ammonium salts. If CaC^ is not added in excess, a white amorphous precipitate forms; this dissolves in excess of the normal tartrate with the formation of a double salt. 4. Potassium Salts, when added to a neutral solution of a tar- trate, give no precipitate; if, however, the resulting solution is rendered acid with acetic or citric acid, a crystalline precipitate of "cream of tartar" forms at once or after some time, depend- ing upon the concentration of the tartrate solution. For prop- erties of this salt, see reaction 2 under Potassium. Free tartaric acid or sodium acid tartrate solutions of moderate strengths give an immediate precipitate with potassium salts. 5. Behavior on Ignition. Tartaric acid and tartrates, when heated, decompose with the evolution of inflammable vapors possessing the odor of burnt sugar; a carbonaceous residue is left at the same time, consisting of carbon in the case of tartaric acid or of a mixture of carbon and alkali carbonate in the case of alkali tartrates. The heavy metal tartrates on heating may leave a residue consisting of the oxide of the metal or of the metal itself. Chromates Chromic acid and chromates have already been mentioned in connection with the metal chromium (see page 58). They are all red or yellow. Solubilities. The chromates of the alkalies, magnesium and calcium, are soluble in water. Nearly all the others are insoluble; most of these dissolve in nitric acid. The acid solutions are THE ACIDS 123 always red, owing to the presence of a dichromate. The color of chromates, even in very dilute solutions, is easily visible, and hence, in the absence of other colored substances, furnishes a delicate test for their presence. For a discussion of the reduc- tion of chromates to chromic salts and of the oxidation of the latter to chromates, see under Chromium, pages 58 and 60. 1. Barium Chloride precipitates yellowish white barium chro- mate (BaCrO4), difficultly soluble in water but soluble in hydro- chloric and nitric acids. For other properties, see reaction 5 under Barium. 2. Lead Acetate precipitates from neutral or acetic acid solu- tions yellow lead chromate (PbCrO4): - Pb(C2H3O2)2 + Na2CrO4 = | PbCrO4 + 2 NaC2H3O2. This is practically insoluble in water, acetic acid, and ammonium hydroxide, but is soluble in caustic soda solution, from which acetic acid reprecipitates the chromate; it is also soluble in nitric acid. 3. Silver Nitrate precipitates from strictly neutral solutions purplish red silver chromate : - 2 AgNO3 + Na2CrO4 = Ag2CrO4 + 2 NaNO3. Silver chromate is readily soluble in nitric acid and ammo- nium hydroxide. From slightly acid concentrated solutions a reddish brown crystalline precipitate of Ag2Cr2O7 is formed which possesses about the same solubilities as Ag2CrO4: - 2 AgNO3 + K2Cr2O7 = | Ag2Cr2O7 + 2 KNO3. The chromates of silver are readily converted into the chloride by treatment with HC1. 4. Potassium Iodide, when added to a dichromate or a nitric acid solution of a chromate, is oxidized with the liberation of iodine; a drop of CS2, when added and shaken with the mix- ture, acquires a violet color, due to the extraction of iodine. The chromic acid is reduced at the same time. 5. Ethyl Alcohol (C2H5OH), when added to a chromate solu- tion acidified with HC1 or H2SO4, and the mixture is boiled, 124 QUALITATIVE CHEMICAL ANALYSIS causes a reduction of the chromate to a chromic salt, as indi- cated by a change in color of the solution from orange-red to green. The alcohol is oxidized at the same time to acetalde- hyde: - K2Cr2O7 + 3 C2H6OH + 8 HC1 = 2 CrCl3 + 7 H2O + 2 KC1 + f 3 CH3CHO. This method is frequently used to reduce chromates to chromic salts before treating a solution containing a chromate with H2S. For the equation for the reduction of chromates by concentrated HC1, see page 6o. 6. Hydrogen Dioxide Test. If to a mixture consisting of 5 cc. of dilute hydrogen dioxide solution, slightly acid with dilute sul- phuric acid, and 2 cc. of ether, a little chromate solution is added, and the mixture is shaken, the ether layer will acquire an intense blue color, due to the presence of some such additive compound as CrO3 • H2O2, which, however, is very unstable. One part of K2CrO4 in 40,000 parts of water is said to be the sensitiveness of this reaction. 7. Insoluble Chromium Compounds, when fused with sodium carbonate to which a little potassium chlorate has been added, and the mass is extracted with water, will yield an aqueous solution containing the chromium as chromate, while the resi- due will contain the other metals. (Note, however, the conduct of Manganese under similar conditions.) GROUP II. This group comprises those acids which yield with silver nitrate precipitates insoluble in nitric acid. BaCl2 does not precipitate them. Chlorides Solubilities. All chlorides are soluble in water with the exception of those of silver, copper(-ous), and mercury (-ous). The oxychlorides of mercury (-ic), bismuth, and antimony are also insoluble. Lead chloride is sparingly soluble in cold water. The normal chlorides of bismuth and antimony require the presence of free acid to keep them in solution. All insoluble THE ACIDS 125 chlorides dissolve in aqzta regia with the exception of silver chloride. Fusion with sodium carbonate transposes* all insolu- ble chlorides: - 2 AgCl + Na2CO3 = 12 Ag + 2 NaCl + f CO2 + f O. The chlorides of barium, sodium, and potassium are practi- cally insoluble in concentrated hydrochloric acid. i. Silver Nitrate precipitates from nitric acid solutions white, curdy silver chloride (AgCl), which darkens on exposure to light; the precipitate is soluble in ammonium hydroxide and carbonate, sodium thiosulphate, and potassium cyanide. Besides the fusion method with sodium carbonate above men- tioned, silver chloride may be tested for chlorine by treating it with zinc and sulphuric acid, and allowing the action to con- tinue for several minutes. If, now, the mixture is filtered, the filtrate will contain the chlorine in the form of ZnCl2 : - 2 AgCl + Zn = ZnCl2 + 12 Ag. 2. Lead Acetate precipitates white lead chloride (PbCl2). For its properties, see reaction I under Lead. 3. Concentrated H2SO4 and MnO2, when added to a chloride, and the mixture is heated, oxidize the chloride, with the evolution of chlorine; the latter is recognized by its color and odor as well as by its ability to bleach moist litmus or indigo paper: - MnO2 + 2 NaCl + 2 H2SO4 = MnSO4 + Na^C^ + 2 H2O + f Cl2. 4. Potassium Dichromate and Concentrated Sulphuric Acid. If a dry mixture of a chloride and powdered K2Cr2O7 is treated with concentrated H2SO4 and heated gently, chromyl chloride (CrO2Cl2), a reddish brown gas, will be evolved : - K2Cr2O7 + 4 NaCl + 6 H2SO4 = 12 CrO2Cl2 + 2 KHSO4 + 4 NaHSO4 + 3 H2O. * When an insoluble salt is treated with sodium carbonate, whereby the acid radical is converted into a soluble sodium salt, the compound is said to be transposed. 126 QUALITATIVE CHEMICAL ANALYSIS If the gas is absorbed by ammonium hydroxide, the latter will be colored yellow owing to the formation of ammonium chromate: - CrO2Cl2 + 4 NH4OH = (NH4)2CrO4 + 2 NH4C1 + 2 H2O. The presence of chromic acid in the ammonium hydroxide solution may be shown by rendering it acid with acetic acid and adding Pb(C2H3O2)2, when a yellow precipitate of PbCrO4 will be obtained. This test for a chloride is of special value, for by its means chlorides may be detected in the presence of bromides. Iodides, if present in not too large amounts, do not interfere. The re- action with the bromide is as follows: - 6 KBr + K2Cr2O7 + 11 H2SO4 = t 3 Br2 + 8 KHSO4 + Cr2(SO4)3 + 7 H2O. The liberated bromine does not color the ammonium hydroxide, and hence does not interfere with the test. Iodides behave in the same way as bromides. Bromides Solubilities. All bromides are soluble in water with the exception of those of silver, mercury (-ous), copper (-ous), and lead, the last being only sparingly soluble in cold water. Solu- ble bromides of the heavy metals are easily transposed by boil- ing with sodium carbonate solution. The insoluble bromides are tested for the halogen by fusion with Na2CO3, extracting the melt with water, and filtering. The filtrate will then contain the bromide as NaBr. 1. Silver Nitrate precipitates yellowish white silver bromide (AgBr), which darkens on exposure to light; it is insoluble in nitric acid and in cold ammonium carbonate; it dissolves with difficulty in cold ammonium hydroxide, but easily in KCN and Na2S2O3. Silver bromide may be decomposed by Zn and sul- phuric acid in the same way as indicated for AgCl. 2. Chlorine Water, when added in small amounts to a solution of a bromide, decomposes it with the liberation of bromine. If THE ACIDS 127 a few drops of carbon disulphide (CS2) or chloroform (CHC13), in which bromine is soluble, are added, and the mixture is shaken, the CS2 or CHC13 will acquire a yellow or reddish color, depending upon the amount of bromine present. The test is exceedingly sensitive; I part of bromine in 1000 parts of water suffices to give a distinctly visible result. KBr + Cl = KC1 + Br. Care must be exercised in performing the test to add the chlorine water one drop at a time, and to shake after each ad- dition ; for, if an excess is added, colorless bromine chloride (BrCl) may form. 3. Potassium Dichromate, in the presence of cold dilute sul- phuric acid, does not liberate bromine from bromides (distinc- tion from iodides). 4. Concentrated H2SO4, when added to a bromide, and the mix- ture is heated, causes the liberation of bromine with hydro- bromic acid. 5. Concentrated H2SO4 and MnO2, when added to a bromide, and the mixture is heated, cause the liberation of bromine; the latter is recognized by its color, odor, and by its property of turning starch yellow and starch-iodide paper blue. 6. Potassium Nitrite, when added to a bromide acidified with dilute sulphuric acid, does not liberate bromine (distinction from iodides). 7. Potassium Permanganate, when added to a bromide solu- tion acid with sulphuric acid, and the mixture is boiled, causes the liberation of bromine, recognizable by its odor, color, and by its ability to turn starch-iodide paper blue. 8. Nitric Acid decomposes bromides, with the exception of AgBr, on heating, with the liberation of bromine: - 6 KBr + 8 HNO3 = f 3 Br2 + 6 KNO3 + f 2 NO + 4 H2O. Iodides Solubilities. With the exception of Hgl2, Hgl, Agl, Cui, and the sparingly soluble Pbl2, all iodides are soluble in water. 128 QUALITATIVE CHEMICAL ANALYSIS Some of the insoluble ones dissolve in an excess of alkali iodide with the formation of double salts. Soluble iodides of the heavy metals are best tested for the halogen by first transposing by boiling with sodium carbonate solution. The insoluble iodides -like the insoluble chlorides and bromides - are best trans- posed by fusion with sodium carbonate. i. Silver Nitrate, when added to a solution of an iodide, pre- cipitates yellow amorphous silver iodide (Agl), insoluble in nitric acid, and only very sparingly soluble in ammonium hy- droxide and cold ammonium carbonate (distinction from chlo- rides). It is soluble in potassium cyanide and sodium thiosul- phate solutions. Silver iodide may be decomposed by Zn and H2SO4 in the manner indicated for AgCl and AgBr. 2. Concentrated H2SO4, acting alone on an iodide, yields hy- driodic acid and iodine, mixed with various reduction products of sulphuric acid, depending upon the temperature and the rela- tive proportions of acid and iodide. 3. Concentrated H2SO4 and Manganese Dioxide, when added to an iodide, and the mixture is heated, liberate iodine. The reac- tion is similar to those in which Br and Cl are set free by the same reagents. 4. Chlorine Water, when added drop by drop to a solution of an iodide, decomposes it with the liberation of iodine: - KI + Cl = KC1 + I. The free iodine may be recognized by shaking the mixture with a few drops of CS2 or CHC13; the latter solvents will ex- tract the iodine and acquire a reddish violet color. Free iodine may also be identified by the blue colored compound it yields when treated with starch paste. In liberating iodine with chlorine water, care must be taken to avoid adding an excess, otherwise the liberated iodine will be oxidized to colorless iodic acid: - I2 + 6 H2O + 5 Cl2 = 10 HC1 + 2 HIO3. 5. Potassium Nitrite, when added to a solution of an iodide acid with sulphuric acid, causes the separation of iodine, recog- nized by coloring CS2 violet or starch paste blue. THE ACIDS 129 In carrying out the starch-iodide reaction, it is important to have the iodine solution very dilute; if it is at all concentrated, a nearly black precipitate, instead of a fine blue coloration, will be obtained. Potassium nitrite has an advantage over chlorine as an iodine liberator, as an excess does not hinder the reaction. 6. Cupric Salts, when added to a solution of an iodide, yields a dirty brown precipitate of cuprous iodide (Cui) ; at the same time iodine is set free : - 2 CuSO4 + 4 KI = 2 K2SO4 + 12 Cui + 12 I. 7. Ferric Salts also liberate iodine from iodides, being at the same time reduced to the ferrous state: - Fe2(SO4)3 + 2 KI = 112 + 2 FeSO4 + K2SO4. 8. Potassium Dichromate, when added to an iodide solution acid with sulphuric acid, causes iodine to be set free : - K2Cr2O7+6KI + 7 H2SO4=4K2SO4 + Cr2(SO4)3 + |3l2 + 7 H2O. In all of the above cases the liberated iodine may be detected by shaking the mixture with a few drops of CHC13 or CS2, which acquire a violet color. 9. Mercuric Chloride, when added to an iodide solution, pre- cipitates scarlet mercuric iodide (Hgl2), soluble in an excess of alkali iodide: - HgC^ + 2 KI = | Hgl2 + 2 KC1; Hgl2 + 2 KI = K2HgI4. Cyanides Solubilities. The alkali and alkaline earth cyanides and mer- curic cyanide are soluble in water; nearly all others are insol- uble. The cyanides of the heavy metals dissolve in an excess of alkali cyanide with the formation of complex double salts. Heated with exclusion of air, the cyanides of the alkalies and alkaline earths fuse without decomposition. In contact with air, they oxidize with the formation of cyanate: - NaCN + O = NaCNO. 130 QUALITATIVE CHEMICAL ANALYSIS It is in consequence of the readiness with which they oxidize that the reducing power of cyanides is due. The cyanides of the heavy metals when heated in a closed tube undergo decom- position, the products varying with the nature of the metal: - Hg(CN)2=Hg + f(CN)2; Pb(CN>2 = Pb + 2 C + f N2. Caution : All tests for cyanides which involve the evolution of a gas should be performed under the hood. The cyanides of the noble metals on heating break up into the metal and cyanogen gas. By this means, the cyanides of silver and mercury allow themselves to be readily detected. Mercuric cyanide differs in many respects from the other water- soluble simple cyanides. It does not yield a precipitate with silver nitrate and is not decomposed by cold dilute sulphuric acid; it is, however, decomposed by H2S with the precipitation of mercuric sulphide and the production of HCN. i. Silver Nitrate, when added to a simple cyanide, excepting Hg(CN)2, yields a white precipitate of silver cyanide (AgCN), easily soluble in excess of alkali cyanide with the formation of a double cyanide: - AgCN 4- KCN = KAg(CN)2. AgCN is also soluble in ammonium hydroxide and in sodium thiosulphate, but is insoluble in nitric acid. On ignition, it is decomposed with the evolution of cyanogen gas and the separa- tion of silver: - 2 AgCN = 12 Ag 4-f (CN^. 2. Sulphuric Acid, dilute, when added to a solution of a cya- nide [except Hg(CN)2], decomposes it with the liberation of HCN, readily recognized by its odor. If the dilute acid is heated, it is capable of decomposing the insoluble cyanides. If the acid is concentrated and hot, it will attack all cyanides whether simple or complex : - Co(CN)2 4- 2 H2SO4 + 2 H2O = CoSO4 +(NH4)2SO4 + f 2 CO; K4Fe(CN)6 4- 6 H2SO4 4- 6 H2O = FeSO4 4- 2 K2SO4 4- 3 (NH^SO* 4- f 6 CO. THE ACIDS 131 The above equations, which are typical, show that the metals are converted into sulphates, that carbon monoxide is produced, and that all the nitrogen is converted into ammonium sulphate. 3. Formation of Ferrocyanide. If a solution containing an alkali cyanide is made strongly alkaline with sodium hydroxide and a little ferrous sulphate and ferric chloride are added, and the mixture is gently heated and finally made acid with hydro- chloric acid, a precipitate of prussian blue will be formed. The reactions which take place are the following: - (1) FeSO4+2 NaOH = | Fe(OH)2 + Na2SO4; (2) Fe(OH)2 + 2 KCN = | Fe(CN)2 + 2 KOH ; (3) Fe(CN)2 + 4 KCN = K4Fe(CN)6; (4) 3 K4Fe(CN)6 + 4 FeClg = | Fe4[Fe(CN)6]3 4- 12 KC1. 4. Formation of Thiocyanate. To a solution of an alkali cya- nide add a little (NH4)2Sj. and evaporate the solution on a water bath to dryness. The residue, which will now consist of ammonium thiocyanate (NH4SCN), is treated with one or two drops of dilute HC1, partly to insure the destruction of any undecomposed sulphide and partly because the presence of hydrochloric acid assists the final reaction [see reaction 5 under Iron (-ic)] ; if a drop of ferric chloride is now added, a blood-red coloration will be produced because of the formation of ferric thiocyanate [Fe(SCN)3]. If it is desired to detect hydrocyanic acid evolved from an insoluble cyanide on treatment with hot dilute H2SO4, cover the test tube containing the mixture with a crucible cover, on the under side of which has been placed a drop of (NH4)2S4., and allow the action to continue for several minutes; then invert the cover and dry on the water bath, and proceed as directed above for the formation of ferric thiocyanate. Ferrocyanides Hydroferrocyanic acid [H4Fe(CN)6] is a colorless crystalline solid, easily soluble in water; on exposure to air the solution 132 QUALITATIVE CHEMICAL ANALYSIS becomes blue from decomposition. The salts of ferrocyanic acid are much more stable than the acid. Solubilities. The ferrocyanides of the alkalies and alkaline earth metals are soluble in water; nearly all the others are in- soluble or sparingly soluble in water and in cold dilute acids. i. Silver Nitrate precipitates white silver ferrocyanide, insol- uble in dilute nitric acid and in ammonium hydroxide, but soluble in potassium cyanide solution : - 4 AgNO3 4- K4Fe(CN)6 = | Ag4Fe(CN)6 4- 4 KNO3. On ignition, the precipitate is decomposed with the separa- tion of silver and evolution of cyanogen gas: - Ag4Fe(CN)6 = 14 Ag 4- f 2 (CN)2 + | FeC2 + f N2. 2. Sulphuric Acid, when cold and dilute, does not decompose ferrocyanides; on heating to boiling, however, partial decom- position sets in with the liberation of part of the cyanogen as hydrocyanic acid: - 2 K4Fe(CN)6 + 3 H2SO4 = | K2Fe [Fe (CN)6] + 3 K2SO4 + f 6HCN. Concentrated sulphuric acid, when heated, completely decom- poses ferrocyanides with the evolution of carbon monoxide: - K4Fe(CN)6 4- 6 H2SO4 + 6 H2O = FeSO4 4- 2 K2SO4 4- 3 (W SO4 + f 6 CO. 3. Ferric Salts, added to a slightly acid solution of a ferro- cyanide, yield a precipitate of prussian blue. 4. The Solution of Insoluble Ferrocyanides is accomplished by boiling the compound with sodium hydroxide and filtering, when the metal will be left on the filter in the form of hydroxide, while the filtrate will contain the acid radical in the form of sodium ferrocyanide: - Fe4[Fe(CN)6]3 4- 12 NaOH = 4 Fe(OH)3 4- 3 Na4Fe(CN)6. The residue is dissolved in acid and the metal is tested for in the solution obtained. The filtrate is acidified with HC1 and tested for the ferrocyanogen radical by adding ferric chloride. THE ACIDS 133 If the metal is one whose hydroxide is soluble in excess of sodium hydroxide, as Zn, it may be separated from the ferro- cyanide by passing carbon dioxide into the alkaline solution, boiling, and then filtering off the basic carbonate of zinc. In the case of lead ferrocyanide, the lead may be precipitated from the alkaline solution by a stream of H2S. Fusion with sodium carbonate decomposes ferrocyanides. Ferricyanides Solubilities. All ferricyanides are insoluble in water and in cold dilute acids with the exception of those of the alkalies and alkaline earths. Heated to redness, ferricyanides decompose, the products being iron carbide, cyanide, nitrogen, and cyano- gen ; the last burns with a characteristic purplish flame. Sul- phuric acid, when dilute and warm, causes partial decomposition with the evolution of HCN. When concentrated and hot, it effects a complete decomposition with the liberation of CO. The equation as given by Treadwell is as follows: - 2 K3Fe(CN)6 + 12 H2SO4 + 12 H2O = Fe2(SO4)3 + 3 K2SO4 + 6(NH4)2SO4 + f 12 CO. 1. Silver Nitrate gives with ferricyanides a reddish brown pre- cipitate of silver ferrocyanide, insoluble in nitric acid, but solu- ble in ammonium hydroxide and in potassium cyanide. 2. Iron Salts. Ferric salts give no precipitate, but a dark coloration; ferrous salts give a blue precipitate of Turnbull's blue, Fe3[Fe(CN)6]2, insoluble in acids. 3. Reducing Agents, such as H2S, SO2, and HI, readily reduce ferricyanides to ferrocyanides in alkaline solutions : - 2 K3Fe(CN)6 + K2S = 2 K4Fe(CN)6 + | S. Thiocyanates Alkali thiocyanates are readily prepared by heating the cya- nide with sulphur: - KCN +S = KCNS; QUALITATIVE CHEMICAL ANALYSIS 134 or by heating an alkali cyanide or hydrocyanic acid with ammo- nium polysulphide: - KCN + (NH4)2S. = KCNS + (NH^. Solubilities. Nearly all the thiocyanates are soluble in water with the exception of those of silver, mercury, copper, and gold. i. Silver Nitrate precipitates white, curdy silver thiocyanate, insoluble in dilute nitric acid and in ammonium hydroxide. 2. Ferric Salts give with alkali thiocyanate solutions a blood- red coloration, due to the formation of ferric thiocyanate. The color is destroyed by mercuric chloride. Sulphides Hydrogen sulphide is a colorless gas having a characteristic and unmistakable odor. Its solution in water possesses a feeble acid reaction, but it is very unstable, oxidizing readily in contact with air with the separation of sulphur : - H2S + O = H2O +1S. H2S reacts with bases forming hydrosulphides and sulphides, which, if H2S is looked upon as a dibasic acid, may be regarded as acid and normal sulphides : - NaOH + H2S = NaHS + H2O ; 2 NaOH + H2S = Na^ + 2 H2O. Behavior on Ignition. Out of contact with the air, most sulphides remain unchanged. The sulphides of arsenic and mercury when heated sublime unchanged. Tin disulphide and iron disulphide give off part of their sulphur. All sulphides on being heated in contact with air (roasted) are oxidized either to oxides or to sulphates. Solubilities. The sulphides of the alkalies are soluble in water; those of the alkaline earths, aluminum and chromium, are hydrolyzed by water with the formation of hydroxides; all other normal sulphides are insoluble in water. 1. Hydrochloric Acid, when moderately strong (1:1), decom- poses nearly all sulphides with the evolution of H2S; the latter THE ACIDS 135 may be recognized by its odor and by its property of turning lead acetate paper black : - ZnS + HC1 = ZnCl2 + f H2S. The few sulphides undecomposed by HC1 are attacked by a mixture of zinc and HC1 with the liberation of H2S. 2. Silver Nitrate precipitates from solutions of sulphides or hydrogen sulphide black silver sulphide (Ag2S), insoluble in cold but soluble in hot dilute nitric acid; it is insoluble in ammonium hydroxide. 3. Oxidizing Solvents, such as concentrated HNO3, aqua regia, HC1+ KC1O3, when used to dissolve a sulphide, do not liberate hydrogen sulphide, but cause a partial separation of sulphur and a partial oxidation of the sulphide to sulphuric acid: - HgS + Cl2= HgCl2 + | S; S + 3 Cl2 + 4 H2O = H2SO4 + 6 HC1; 3 PbS + 8 HNO3 (hot and concentrated) = ^3PbSO4 + f 8NO + 4H2O 4. Sodium Nitroprusside [Na2(NO)Fe(CN)5] imparts to nor- mal alkali sulphide solutions a reddish purple color. An aqueous solution of H2S does not give this reaction. 5. Sodium Plumbite (prepared by adding caustic soda solu- tion in excess to a lead salt) will detect a sulphide even in the presence of free alkali or carbonate, by producing a brown or black precipitate. This test is exceedingly sensitive. 6. For the oxidizing effect of the halogens, nitric acid, potas- sium dichromate, potassium permanganate, ferric salts, etc., on H2S, see page 45. Insoluble Sulphides. Sulphides insoluble in acids, when fused in a small nickel crucible with sodium hydroxide, are decom- posed with the formation of sodium sulphide. If the mass is placed on a silver coin and then moistened, a black stain of silver sulphide (Ag2S) will be produced. 136 QUALITATIVE CHEMICAL ANALYSIS Nitrites Solubilities. Most nitrites are soluble in water; silver nitrite, however, is only sparingly soluble in water. On ignition, nitrites undergo decomposition with the production in general of nitro- gen oxides and the oxide of the metal. i. Sulphuric Acid, when dilute, decomposes all nitrites (dis- tinction from nitrates) with the evolution of nitric oxide (NO); the latter immediately oxidizes in contact with the air to brown nitrogen peroxide (NO2)* : - (i) NaNO2+H2SO4=NaHSO4 + HNO2; (2) 3 HNO2 = HNO3 + f 2 NO + H2O ; (3) NO + O = NO2. Nitrites are also decomposed by acetic acid with gentle heat- ing; the NO2 given off maybe recognized by its turning starch- iodide paper blue. 2. Silver Nitrate precipitates, from solutions of nitrites which are not too dilute, white silver nitrite (AgNO2), difficultly soluble in cold but more easily soluble in hot water. 3. Cobalt Salts solutions acidified with acetic acid give with moderately strong solutions of nitrites a yellow precipitate of potassium cobaltic nitrite [K3Co(NO2)6]. 4. Potassium Iodide, when added to a solution of a nitrite, and the mixture is acidified with acetic acid, produces a separation of iodine; the latter may be recognized by turning starch paste blue, or by coloring a drop of CS2 or CHC13 violet. 5. Potassium Permanganate solution, when warmed, is bleached by a solution of a nitrite slightly acid with dilute sulphuric acid: - 2 KMnO4 + 5 HNO2 + 3 H2SO4 = K2SO4 + 2 MnSO4 + 5 HNO3 + 3 H2O. * The older nomenclature is applied to the oxides of nitrogen as being more distinctive. THE ACIDS 137 GROUP III. The acids of this group are not precipitated by either AgNO3 or BaCl2. Acetates Solubilities. All normal acetates are readily soluble in water. Some basic acetates, such as those of iron(-ic), aluminum, and chromium, are practically insoluble, while the normal silver and mercurous salts are only sparingly soluble. On ignition, acetates decompose with little or no charring and with the production of a combustible gas. The alkali acetates, on ignition, are converted into carbonate and acetone: - 2 NaC2H3O2 = Na2CO3 + f (CH3)2CO. All other acetates behave similarly. If the carbonate is un- stable at the ignition temperature, the oxide is produced; if the latter is unstable, then the metal alone is left as a residue. i. Sulphuric Acid, whether dilute or concentrated, liberates acetic acid from its salts; the acid, being volatile, is easily rec- ognized by its characteristic odor. 2. Alcohol and Concentrated Sulphuric Acid. If to a cooled mixture of an acetate and concentrated sulphuric acid, a little ethyl alcohol is added and then the mixture is gently heated, ethyl acetate will be formed; the latter is easily recognized by its fruity odor. If amyl alcohol is used instead of ethyl alcohol, amyl acetate, having an odor resembling pear essence, will be produced: - CH3- COO[H + HO]C2H5^CH3- COOC2H5 + H2O. Acetic acid. Ethyl alcohol. Ethyl acetate. The reversible reaction is made to proceed almost entirely from left to right by the presence of the concentrated sulphuric acid, which removes the water as soon as it is formed. 3. Silver Nitrate gives with moderately concentrated solutions of an acetate or acetic acid a white crystalline precipitate of silver acetate (AgC2H3O2), sparingly soluble in cold water (1.04 138 QUALITATIVE CHEMICAL ANALYSIS parts of the salt dissolve in 100 parts of water at 20° C.), more readily soluble in hot water, and easily soluble in ammonium hydroxide. 4. Ferric Chloride, when added to an alkali acetate solution, produces a reddish brown solution, due to the formation of ferric acetate. If this solution is largely diluted and boiled, all of the iron will be precipitated as a basic ferric acetate : - Fe(C2H3O2)3+ H2O^|Fe(OH)(C2H3O2)2 + HC2H3O2. Nitrates Solubilities. The nitrates, with the exception of a few basic nitrates, such as BiONO3, are all soluble in water. Barium nitrate, however, is only sparingly soluble in water and is almost insoluble in moderately strong nitric acid. All nitrates on igni- tion undergo decomposition, the alkali and alkaline earth ni- trates yielding nitrites with the evolution of oxygen : - KNO3= KNO2 + fO. At higher temperatures, the nitrites are decomposed with the production of the oxides of nitrogen and a residue consisting of the oxide or peroxide of the metal. The nitrates of the heavy metals yield at a red heat nitrogen peroxide and oxygen. Heated on charcoal, all normal nitrates deflagrate. 1. Concentrated Sulphuric Acid, when added to a solid nitrate, causes the evolution of nitric acid. If the mixture is heated, the nitric acid is decomposed with the production of brown fumes of nitrogen peroxide (NO2):- (1) NaNO3 + H2SO4 = NaHSO4 4- HNO3; (2) 2 HNO3 = f 2 NO2 + H2O + f O. 2. Ferrous Sulphate, when added to a cool mixture of a nitrate solution and concentrated sulphuric acid, produces a deep brown color. The reaction may be considered as taking place in three stages: 1st, the liberation of nitric acid by the action of con- centrated sulphuric acid on the nitrate; 2nd, the reduction of the free nitric acid by the ferrous sulphate, resulting in the pro- duction of nitric oxide ; and, 3rd, the solution of the nitric oxide THE ACIDS 139 in the excess of unoxidized FeSO4 with the formation of a brown unstable compound : - (i) H2SO4 + NaNO3 = HNO3 + NaHSO4; (2) 2 HNO3 + 6 FeSO4 + 3 H2SO4 = 3 Fe2(SO4)3 + 4 H2O + 2 NO; (3) 2 FeSO4 4- NO = (FeSO4)2NO (brown compound). This test, which is exceedingly delicate, is carried out as fol- lows : To about 3 cc. of the nitrate solution contained in a test tube, add an equal volume of concentrated sulphuric acid, mix, and thoroughly cool under running water. Hold the tube in an inclined position and cautiously add a few cc. of a strong, freshly prepared ferrous sulphate solution (free from nitrates), and allow the tube to stand. If a nitrate is present, a brown ring will form at the zone of contact of the two liquids. Nitrites give the same reaction, but may be carried out with dilute instead of concentrated sulphuric acid. The colored com- pound is destroyed on warming. 3. Indigo Solution. If to a little HC1 that has been recently boiled, a few drops of a solution of indigo in sulphuric acid are added, and the mixture is again boiled, the blue coloration will persist, if the HC1 contains no free chlorine. If to the blue solu- tion a nitrate in the form of a solid or in solution is added, and the liquid is boiled, the indigo will be bleached. As the bleaching of the indigo is caused by the chlorine which is liberated, any other oxidizing agent, which will yield chlorine with HC1 like KClOg, will produce the same result. 4. Copper Filings, when added to a nitrate, and the mixture is heated with concentrated sulphuric acid, cause the production of brown nitrogen peroxide fumes : - (1) NaNO3 + H2SO4 = NaHSO4 + HNO3; (2) 3 Cu + 8 HNO3 = 3 Cu(NO3)2 + f 2 NO + 4 H2O; (3) NO + O = NO2. 5. Reduction to Ammonia. If a solution of a nitrate is made strongly alkaline with NaOH, and a few pieces of aluminum or 140 QUALITATIVE CHEMICAL ANALYSIS zinc and iron filings are added, and the mixture is heated, ammo- nia gas will be evolved. Nitrites give the same reaction. 6. Free Nitric Acid may be recognized even in the presence of nitrates by adding to the solution a few quill cuttings and evaporating to dryness on the water bath. The quills will be found to have a yellow color, due to the formation of xantho- proteic acid; the same compound is formed when concentrated nitric acid is brought in contact with the skin. Chlorates The chlorates are all soluble in water. On prolonged igni- tion, they are decomposed, giving off oxygen and leaving a residue of the chloride of the metal: - KC1O3 = KC1 + f 3 O ; AgC103 = AgCl + f 3 O. > In consequence of the readiness with which they decompose with the liberation of oxygen, chlorates are valuable oxidizing agents. When they are mixed with organic matter and heated, deflagration results. 1. Concentrated Sulphuric Acid decomposes chlorates with the production of chlorine peroxide (C1O2), a very unstable, green- ish yellow gas, which, on being warmed, violently explodes. The sulphuric acid acquires at the same time a deep yellow color, due to dissolved C1O2: - 3 KCIO3 + 2 H2SO4 = KC1O4 + 2 KHSO4 4- H2O + f 2 C1O2. In carrying out the test, it is preferable first to warm a little concentrated sulphuric acid in a test tube and then to drop in a very small crystal of KC1O3. One should never look down into a test tube, especially when performing this test. Equal consideration is due one's neighbor. 2. Concentrated HC1, when added to a chlorate, is oxidized with the production of chlorine and chlorine peroxide: - KCIO3 + 2 HC1 = KC1 + j Cl + f C1O2 + H2O. THE ACIDS 141 3. Aniline Sulphate. If to a solution of a chlorate in con- centrated sulphuric acid a drop of aniline sulphate solution is added, a deep blue coloration will be developed; the color may- be intensified by the addition of a few drops of water. This reaction is exceedingly delicate and may be used to distinguish a chlorate from a nitrate. 4. Reducing Agents, as sulphurous acid or the alkali sulphites in acid solutions, change the chlorates to chlorides : - HCIO3 + 3 H2SO3 = 3 H2SO4 + HC1. ACID ANALYSIS P reliminary Examination Before proceeding with the analysis for the acids, the student should first complete his examination for metals, the results of which will, by a proper use of the table of solubilities, restrict the number of acids to be looked for. An example will make this clear. If the substance under examination is soluble in water and lead has been found, none of the acids which form insoluble salts with lead need be looked for, viz., carbonic, sulphuric, hydrogen sulphide, chromic, oxalic, etc. Again, if the original substance is insoluble in water, but soluble in hydrochloric acid, and barium has been found, one need not look for sulphuric acid. Further, if silver has been detected in the metallic analysis of a substance soluble in water, it is evident that all the acids of Group II. need not be looked for. It is also well to remember that certain acids cannot coexist in solution, eg., oxidizing agents like K2Cr2O7, H3AsO4, FeCl3 cannot exist with reducing agents like sulphites and iodides. It is also desirable that the first three of the following pre- liminary tests be carried out before commencing the systematic search for acids. 142 QUALITATIVE CHEMICAL ANALYSIS PRELIMINARY TESTING OF SOLIDS 1. Heat a small quantity of the substance in a tube closed at one end. Observation Indication Decrepitation. Carbonization, accompanied by burnt odor and formation of water. Water given off. Gases given off: - O - kindles a spark; Colorless and CO - burns with blue odorless. | flame; CO2-turns limewater turbid. ' N H3-turns red litmus blue; SO2 - recognized by odor; Colorless (CN)2 - recognized with by odor,* and burns odor. with reddish flame; H2S - recognized by odor; Acetone - recognized by odor. 'NO2 -reddish brown; Colored turns starch-KI gases. ' paper blue. Cl, Br, I - recognized - by color and odor. A sublimate forms: - White. Yellow. NaCl, Pb(NO3)2, K2SO4, zinc blende, and other substances. Organic matter, tartrates, and other organic acids and salts. Water mechanically inclosed, water of hydration, and hydroxide. Chlorates, peroxides, certain oxides, nitrates, etc. Oxalates. Carbonates, oxalates, and organic matter. NH4 salts and organic compounds containing N. Sulphur, sulphides, sulphites, and cer- tain sulphates. Cyanides. Moist sulphides. Acetates. Nitrates of heavy metals. Chlorides, bromides, and iodides in the presence of oxidizing agents. NH4 salts, HgCl, HgCl2, As2O3, Sb2O3, and certain organic compounds. As2S3, HgO (accompanied by globules of Hg). * Smell cautiously, THE ACIDS 143 Observation Indication Yellow, becoming red when Hgl2. rubbed. Reddish brown drops, yellow S. when cold. Black accompanied by garlic odor. As. Black accompanied by violet I. vapor. Metallic globules or mirror. Hg- Substance changes color: - Becomes black. Salts of Cu, Ni, Co, Mn. Becomes dark red. Salts of Fe. Black (hot) and red (cold) ac- Hg salts. companied by metallic globules. Dark red (hot), yellow (cold). PbCrO4. Yellow (hot), white (cold). ZnO. 2. Put a small quantity of the substance in a test tube, add a little dilute HC1, and heat gently. Observation Indication CO2: effervescence, turns limewater turbid. Carbonates. SO2: recognized by odor. Sulphites. SO2: accompanied by a separation of S. Thiosulphates. H2S : recognized by odor and by lead acetate paper. Sulphides. NO2: brown, turns starch-KI paper blue. Nitrites. HCN : recognized by odor.* Cyanides. 3. Heat a small portion in a test tube with concentrated sulphuric acid. Observation Indication Acid fumes are evolved which redden litmus. Br, I, mixed with HBr, SO2, and per- haps H2S. Halogen acids from their salts, Iodides and bromides. * Smell cautiously by fanning vapor with hand towards nose, QUALITATIVE CHEMICAL ANALYSIS 144 Observation Indication Chlorine - bleaching litmus. CrO2Cl2 - (reddish brown). C1O2 - yellow color of gas and H2SO4. Gas explosive. HF - etches glass ; yields SiF4, which turns " water bead " turbid. HC2H3O2 - recognized by odor. NO2 - recognized by odor and by turning starch-KI paper blue. SO2 - recognized by odor. SO2 - accompanied by blackening. CO - (without blackening) recog- nized by burning with blue flame. CO2 - turns limewater turbid. Chloride and oxidizing agent together. Chromate and chloride together. Chlorate. Fluoride. Acetate. Nitrates, nitrites. Sulphite, thiosulphate, or reducing agent acting on H2SO4. Organic matter or tartrate. Oxalate, cyanides, ferro- and ferri- cyanides. Carbonates and oxalates. 4. Heat alone on charcoal with blowpipe. Observation Indication (a) Substance fuses and runs into charcoal. Salts of Na, K, and Li. (£) Substance decrepitates. NaCl and other compounds. (c) Substance deflagrates. Chlorates, nitrates. (d) Substance is infusible; residue mois- tened with water reacts alkaline. Ba, Sr, Ca, Mg. Residue moisted with two drops of very dilute Co(NO3)2 solution and heated again gives a blue mass, Al. green mass, Zn. pink mass. Mg. (^) Produces an incrustation that is white accompanied by garlic odor. As. Yellow (hot), white (cold). Zn. Yellowish brown (hot), white (cold), hear residue, and not volatile. Sn. Reddish yellow (hot), yellow (cold). Pb. Orange (hot), light yellow (cold). Bi. Reddish brown, cold and hot. Cd. THE ACIDS 5. Mix with anhydrous sodium carbonate and heat on char- coal. Observation Indication (a) Metallic globule forms without incrustation: - yellow, Au. red, Cu. white and malleable. Ag- (£) Metallic globule with incrustation: - malleable, Pb, Sn. brittle. Sb, Bi. (c} Dark and brittle magnetic mass. Fe, Co, Ni. 6. Make borax bead test; introduce first in oxidizing and then in reducing flame. Oxidizing Flame Reducing Flame Indication Blue. blue. Co. Greenish blue. red - opaque. Cu. Green. green. Cr. Yellow. green. Fe. Brown. gray - opaque. Ni. Violet. colorless. Mn. 7. Moisten substance with concentrated HC1, take up a small portion on the loop of .a platinum wire, and hold in the flame.* Observation Indication Intense yellow which persists for several seconds. Deep red. Reddish yellow. Green or greenish yellow. Pale blue. Na. Li or Sr. Ca. Ba, Cu, or H3BO3. As. * Certain substances, like the sulphates of the alkaline earths, are not volatilized in the flame; in such cases, it is well to hold them first in the reducing flame, then moisten with HC1, and introduce into the colorless bunsen flame. 146 QUALITATIVE CHEMICAL ANALYSIS Method of Acid Analysis The method employed for the detection of the acids is different from that used in the systematic examination for metallic radicals. We cannot, as was done with the metals, divide the acids into groups by certain reagents and then separate the various group precipitates into their component acids. For the most part, the procedure consists in independently and separately testing for each of them. From the list of acids this would seem a long and tiresome task, but in actual practice the number of acids which must be looked for is very much reduced; first, from a know- ledge of the solubilities and metallic content of the substance; and, second, by the results furnished by the preliminary experi- ments just given. The reagents BaCl2 and AgNO3, when properly applied, are valuable in that they give indications of the presence or absence of whole groups; e.g., if to a moderately concentrated and neutral solution of the substance, BaCl2 or CaCl2 is added and no precipitate results, the absence of all the members of Group I. may be inferred.* However, these reagents cannot be used to separate the acids in the manner in which group reagents are employed to precipitate metallic groups. GENERAL EXAMINATION FOR ACIDS Preliminary treatment of the sample with dilute HC1 will dis- close the presence or absence of the following acids: H2CO3, H2SO3, H2S2O3, H2S, HCN, and HNO2. In the course of the analysis for metals, H3AsO3, H3AsO4, and H2CrO4 will be de- tected. For the examination for acids, it is desirable in most cases to have a solution which shall contain the acids in the form of sodium salts. Such a solution, known as the " prepared solu- tion," may be obtained by boiling the finely powdered substance with an excess of Na2CO3 solution, with constant stirring, for several minutes (i). If ammonia is given off, boil with the ad- dition of more Na2CO3 until no more of this gas is evolved, and then filter. * The solution should also contain no considerable quantity of NH< salts, else borates, fluorides, and tartrates may not precipitate. THE ACIDS 147 Residue will con- tain the hydroxides, carbonates, and basic carbonates of the metals (except the alkalies, As and Sb) ; it may also contain phosphates, fluorides, and silicates. Re- serve this residue, and if these acids are not found in the fil- trate, divide it into two parts. 1st part. Test for PO4 and SiO2. Acidify with HNO3, evaporate to dry- ness, extract with hot dil. HNO3, and filter. Test residue with NaPO3 bead or with HF in a plati- num crucible. Test filtrate for PO4 with (NH4)2MoO4. 2nd part. Test for HF. Filtrate (prepared solution) will contain the acids in form of sodium salts + an excess of Na2CO3 and is to be used for the acid tests unless otherwise directed. Preliminary Tests for the Acid Groups Just acidify a small portion of the prepared solution by the careful addition of dil. HNO3; filter if necessary (2) and boil the filtrate until all of the CO2 is expelled. Render faintly alkaline with ammonia, and boil off any excess of the latter that may have been added; filter again if necessary. Divide this solution into 2 parts. 1st part. Test for Group I. by adding a little BaCl2 and CaCl2 solutions (3). A white precipitate shows the presence of the Group I. (4) ; acidify with HC1. If the ppt. does not dissolve, H2SO4 is present. If the group is present, test separate portions of the prepared solutions for H3PO4, H3BO3, HF, H2C2O4, H2C4H4O6, H2SiO3. 2nd part. Test for Group II. Render the solution acid with HNO3 and add an excess of silver nitrate. A ppt. proves the presence of Group II. Note the color of the ppt. (5) and filter. Residue. Wash several times on filter with water. Transfer ppt. to a test tube and shake vigorously with an excess of dilute NH4OH. If complete solution takes place, the absence of HI and H4Fe(CN)6 is shown. HBrand HCNS may also possibly be absent. If Group II. is shown to be present, test separate portions of the prepared solutions for HI, HBr, HC1, HCN, H4Fe(CN)6, H3Fe(CN)6, and HCNS. To filtrate or solution in which AgNO3 produces no ppt. in HNO3 solution, add more AgNO3 to insure complete precipitation; fil- ter if necessary, and to the filtrate contained in a test tube carefully add 3 to 5 drops of ammonia. Agi- tate the upper portion of the liquid and note the color of any ppt. which may form at the neutral zone (6). Group III. Test separate portions for HC103, HNO3, and HC2H3O2. 148 QUALITATIVE CHEMICAL ANALYSIS NOTES 1. If the substance is soluble in water and contains no heavy metals, the treatment with Na2CO3 may be dispensed with. If acid (preferably nitric acid) has been used in getting the substance into solution, and the latter is boiled, H2CO3, H2SO3, H2S2O3, H2S, HCN, and HC2H3O2 will be driven off or decomposed, and therefore should not be looked for; the presence of all of these acids will have been revealed, however, in the preliminary examination. Substances insoluble in water but soluble in acids, when boiled with Na2CO3 solution, may leave a residue consisting of the phosphates and fluorides of certain metals which are not readily transposed by boiling with Na2CO3. Substances insoluble in acids should be fused in a platinum crucible if no reducible metals are present, otherwise in either a nickel or a porcelain crucible with Na2CO3; the melt is then extracted with boiling water and the solution is filtered. The filtrate will correspond to the " prepared solution " and should be used for the acid tests. The residue is tested for phosphates and fluorides. In certain cases where As and Sb are known to be present, it may be necessary to remove these metals by passing H2S into the acidified solution, filtering, and boiling out the H2S from the filtrate. The latter may then be treated with Na2CO3 in the manner already described. 2. Boiling with Na2CO3 will leave all of the metals in the residue, with the exception of the alkalies, As, Sb, and small amounts of metals slightly soluble in excess of Na2CO3. On acidifying this solution, a precipitate may be obtained. It should be filtered off, rejected, and the filtrate again boiled to drive out any H2S that may be liberated before neutralizing the solution with ammonium hydroxide. Too much HNO3 should be avoided, as this will form with the ammonia next to be added an unnecessarily large amount of NH4NO3, in which the Ca or Ba salts of nearly all the acids of Group I., especially the borate, fluoride, and tartrate, are soluble. 3. CaCl2 is also added because the fluoride, tartrate, and oxalate of calcium are much more insoluble than the corresponding salts of Ba. 4. If no precipitate is obtained, the absence of all the acids of Group I. is proved with the exception of boric acid, which is precipitated only from rather concentrated solutions. It should, however, be remembered that the pres- ence of much ammonium salts interferes more or less with the precipitation of all the acids of this group with the exception of H2SO4 and H2C2O4. 5. If only a cloudiness is obtained with AgNO3, it indicates a trace of chlorides which should not be reported. Ag2C2O4 is difficultly soluble in HNO3; an excess of this acid should be added before adding the AgNO3 to prevent its precipitation. The color of the silver precipitate, together with its solubility in ammonium hydroxide, affords important indications of the acid present: Agl is yellow; Ag2S, black; Ag3Fe(CN)6, reddish brown; AgCl, AgCN, AgSCN,and Ag4Fe(CN)6 are white ; and AgBr is yellowish white. THE ACIDS 149 Of these, only the sulphide, iodide, and ferrocyanide are insoluble in ammo- nium hydroxide; the bromide and thiocyanate are difficultly soluble in this reagent. 6. Silver nitrate also precipitates from neutral solutions all the acids of Group I. with the exception of HF and H2SO4; the latter is, however, diffi- cultly soluble in water. The color of the precipitate forming at the neutral junction of the two liquids will often indicate which of the acids of Group I. are present; if yellow, it may be Ag3PO4 or Ag3AsO3; if brownish red, Ag3AsO4; if purplish red, Ag2CrO4; if white, the oxalate, silicate, or borate. SPECIAL TESTS FOR THE ACIDS The metallic analysis and preliminary tests completed, the student should, with the aid of the table of solubilities, thought- fully prepare a list of acids which are likely to be present and hence to be looked for. No acid should be excluded which is compatible with the solubility and metallic content of the sub- stance. Minerals, as a rule, need not be tested for organic and cyanogen acids, and, if insoluble, for nitrates, chlorates, bro- mides, and iodides. Alloys contain no acids as such; they may, however, contain acid-forming elements such as S, P, and Si, which, by treatment with suitable oxidizing agents, will yield the corresponding acids. Carbonates Treat a small portion of the finely ground substance in a test tube with dilute HC1 and warm. A carbonate, if present, will evolve CO2, which may be recognized by its property of rendering turbid a drop of limewater supported in a glass tube. NOTES 1. Sulphites, if present, would evolve SO2, which would also render lime- water turbid. Sulphides and nitrites also liberate gases on treatment with dilute HC1. To avoid the interference of these substances, use a strong solution of K2Cr2O7, instead of acid, and warm the mixture. CO2 alone will be evolved. K2Cr2O7 oxidizes sulphides and sulphites, and is without action on nitrites. The same end may be attained by treatment with acid and then passing the evolved mixed gases through bromine water. 2. Certain carbonates are not readily decomposed by cold dilute acid, e.g., magnesite, dolomite, and the carbonates of the heavy metals. They are all decomposed, however, on warming the acid. 150 QUALITATIVE CHEMICAL ANALYSIS 3. In making the test, care must be exercised to prevent any acid, which may be mechanically carried up the tube in the form of spray, from coming in contact with the drop of limewater. 4. The drop of limewater should be examined shortly after exposure to insure the non-formation of soluble calcium dicarbonate. 5. Where the amount of CO2 liberated is small, it is necessary to heat the acid to drive out the CO2, which otherwise would remain wholly in solution and thus escape detection. Sulphites Treat a small quantity of the solid substance with dilute HC1. In the presence of a sulphite, SO2 will be evolved; this gas may be readily recognized by its odor and by its property of bleaching a very dilute solution of KMnO4 (see reaction 6 under Sulphites'). Thiosulphates These are detected in the preliminary tests with dilute HC1. In the presence of a thiosulphate, SO2 is evolved, accompanied by a separation of S. Sulphates Sulphates will have been indicated in the preliminary testing for the groups of acids. To a small portion of the "prepared solution," add dilute HC1 to acid reaction, boil to expel CO2, filter if necessary, and to the filtrate add BaCl2. A white precipitate indicates the presence of sulphates or fluosilicates. To confirm the presence of sulphates, dry the precipitate, mix it with a little anhydrous Na2CO3, and heat on charcoal before the blowpipe. Remove the residue from the charcoal, place it on a bright silver coin, and add a drop of water. A black stain confirms the presence of a sulphate. Fluosilicates Acidify a small portion of the prepared solution with dilute HC1, boil out the CO2, filter if necessary, and to the clear fil- trate add BaCl2 to complete precipitation. Allow to stand for several minutes and filter. Wash and completely dry the pre- cipitate at a low temperature. Transfer the precipitate to a THE ACIDS 151 test tube, add concentrated H2SO4, heat, and hold in the escap- ing gases a drop of water held on the loop of a platinum wire. In the presence of a fluosilicate, the drop of water will become turbid owing to the formation of H4SiO4. (See footnote, page 105.) (a) Solutions of a dichromate or a chromate possess an orange or yellow color which is very characteristic. Acidifica- tion with dilute HC1, followed by treatment with H2S, will cause a change in color to green, accompanied by a separation of S; hence this acid will be detected in the precipitation of the second group of metals. (£) The change in color, which is an evidence of reduction, may be brought about by a variety of reducing agents in the presence of free acid, ^., strong HC1 and a little alcohol (see reaction 5 under Chromates'), KI, and NagSOgj the last is oxi- dized at the same time to sulphate. If the prepared solution is colorless, chromates cannot be present. (f) Precipitation Test. In the absence of sulphates, phos- phates, oxalates, and tartrates, acidify the " prepared solution " with acetic acid, boil to expel CO2, filter if necessary, and to the clear solution add 1 g. of NaC2H3O2 and a little Pb(C2H3O2)2 solution. A yellow precipitate of PbCrO4 proves the presence of a chromate. If the above-mentioned acids are present, acidify the prepared solution with HNO3, boil to expel CO2, filter if necessary, and render the resulting filtrate just alkaline with ammonium hy- droxide ; add CaCl2, warm, shake vigorously, and allow to stand half an hour and filter. The precipitate may consist of the tartrate, oxalate, and phosphate of calcium. The filtrate may contain sulphates besides the chromate. To remove the former, acidify with HNO3, heat to boiling, add a slight excess of BaCl2, and filter off the BaSO4 on two folds of filter paper. To the filtrate add several grams of NaC2H3O2 to completely replace the nitric acid by acetic acid, and heat, when a yellow precipi- tate of BaCrO4 will be formed. Chromates 152 QUALITATIVE CHEMICAL ANALYSIS Arsenites Arsenites are detected in the analysis for metals. In HC1 so- lutions, H2S yields an immediate precipitate of As2S3. Arsen- ites are not precipitated by either magnesia mixture or ammonium molybdate. In a strictly neutral solution, AgNO3 produces a yellow precipitate of Ag3AsO3 (phosphates respond to the same test). Arsenates From acid solutions, H2S slowly yields a yellow precipitate. From strictly neutral solutions, AgNO3 precipitates reddish brown Ag3AsO4. Precipitates are obtained with both magnesia mixture and ammonium molybdate. The last two tests apply only in the absence of phosphates. For distinctions between phosphates and arsenates, see reactions 4 and 5 under Phos- phates. If arsenic is found in the metallic analysis, it is usually present as arsenite or arsenate. In the absence of oxidizing agents, as chromic and nitrous acids, arsenates are further dis- tinguished from arsenites, even in the presence of phosphates, by the ability of arsenates to liberate iodine from KI in a solution acid with HC1. For the detection of Iodine, see p. 128. Phosphates Phosphates will have been detected in a complete analysis be- fore precipitating the third group of metals. (a) In the absence of arsenates, a small portion of the pre- pared solution is strongly acidified with concentrated HNO3, then it is evaporated nearly to dryness, diluted with water, treated with io cc. of ammonium molybdate solution, and warmed; in the presence of a phosphate, a yellow precipitate of (NH4)3PO4- 12 MoO3 will be formed. (^) Or the prepared solution is acidified with HC1, boiled to expel CO2, filtered if necessary, made alkaline with ammonium hydroxide, filtered again if a precipitate forms, and the clear, cooled filtrate is treated with magnesia mixture and thoroughly shaken. A white crystalline precipitate of NH4MgPO4 forms in the presence of a phosphate. THE ACIDS NOTES 153 1. The solution is evaporated with concentrated HNO3 to oxidize any reducing agent that may be present and which would interfere with the (NH4)2MoO4 test; it also converts at the same time any meta- or pyro-phos- phate to the ortho form, which alone is precipitated by (NH4)2MoO4. 2. If arsenic has been found, it should be removed by rendering the pre- pared solution strongly acid with HC1, heating to boiling, and passing in a stream of H2S for 20 minutes; then filter, boil out the H2S from the filtrate, add HNO3, and evaporate to dryness; extract the residue with boiling dilute HNO3 and add to the somewhat cooled solution an excess of ammonium molybdate. A yellow precipitate shows the presence of PO4. The solution to be tested for PO4 should not be above 700 C., as there is danger of decom- posing the reagent, with the resulting precipitation of white MoO3. Oxalates Oxalates should be detected before proceeding with the pre- cipitation of Group III. Slightly acidify some of the prepared solution with acetic acid, boil out the CO2, and filter if necessary; warm the filtrate and add an equal volume of a saturated CaSO4 solution. A white crystalline precipitate indicates the presence of an oxalate. Confirm by filtering off the precipitate, washing it with water, dissolving it in hot dilute H2SO4, and adding a drop of dilute KMnO4 solution. In the presence of an oxalate, the KMnO4 will be bleached. (See reaction 4 under Oxalates.') NOTES The solution is rendered acid with acetic acid to prevent the precipitation of carbonates and phosphates. CaF2 may be precipitated, but may usually be distinguished from CaC2O4 by the fact that the former is gelatinous while the latter is crystalline. It is, however, better to make the confirmatory test. Fluorides i. The etching test (see page 114) is not applicable in the presence of silicates or silica. 2. The test depending upon the formation of SiF4 and the detection of the latter by its property of rendering a water "bead," held on the loop of a platinum wire, turbid (see 154 QUALITATIVE CHEMICAL ANALYSIS page 113), is applicable to fluorides in the presence of SiO2 or silicates. 3. Silicates which are not decomposed by concentrated H2SO4 may be tested for fluorides by fusing with 6 to 8 times their weight of a mixture of equal parts of sodium carbonate and po- tassium carbonate, extracting the melt with water, and filtering. The filtrate will contain all of the F as NaF, as well as the silica in the form of Na2SiO3. Acidify with acetic acid and filter off any precipitate which forms. To the filtrate, add CaCl2 and allow the mixture to stand for some time. Collect the precipi- tate on a filter, dry, and apply the tests for a fluoride. Borates i. Turmeric Paper Test. Dip a piece of turmeric paper into a small portion of the original or prepared solution acidified with HC1 and dry it; this may be conveniently accomplished by placing it on the outside of a test tube containing water which has just been heated to boiling. In the presence of a borate, the turmeric paper assumes a permanent reddish brown color, which, on treatment with a drop of caustic alkali, is changed to a greenish black color. 2. The Flame Test (see reactions 6 and 7 under Borates} may be conducted on a portion of the original substance, pro- vided Ba and Cu are both absent. If these metals are present, the test may be applied by first removing the Cu with H2S and then the Ba with sulphuric acid. If the test is carried out in a test tube provided with a stopper, through which passes a glass tube drawn out near the end, and the mixture is heated, only the vapors of boron ester will escape. If the issuing gas is lighted, it will burn with a green flame. The advantage of this form of apparatus is that, as neither Ba nor Cu form volatile compounds under these conditions, they do not interfere. NOTES ON THE TURMERIC TEST Oxidizing agents like chlorates, chromates, and iodides interfere with this test by destroying the turmeric. HNO3 is an exception. Chlorates and chromates may both be reduced by treating the original substance contained THE ACIDS 155 in an evaporating dish with solid Na2SO3, adding dilute HC1, and warming after the reaction has proceeded for some time, to drive out the excess of SO2. Filter, if necessary, and boil the filtrate with a slight excess of Na2CO3; dilute, and filter. Iodides, if present, may be removed by precipitation with AgNO3 after rendering the solution acid with HNO3. FeCl3, if present, will color the turmeric paper brown on concentration, but will give a brown instead of a greenish black color when the dried paper is treated with caustic soda. It is evident that if the prepared solution is used, Fe cannot be present. Silicates 1. The NaPO3 bead test may be applied to the original sub- stance. 2. Evolution of SiF4. In a platinum crucible, treat a mixture of equal parts of the dry substance and CaF2 (free from SiO2) with a little concentrated H2SO4, and heat under a hood. A drop of water, held on the loop of a platinum wire, when brought near the mouth of the crucible, will be rendered turbid by the escaping SiF4. About 2 cc. of aqueous HF may be used instead of the CaF2. Tartrates The presence of tartrates will be indicated when the substance under examination is heated in a closed tube, as well as by the characteristic behavior when heated with concentrated H2SO4 (see reaction I under Tartrates}. i. Concentrate the prepared solution to about 0.5 cc., just acidify with acetic acid, and add 2 cc. of KC2H3O2; shake vigor- ously and allow the mixture to stand. A white crystalline pre- cipitate may be KHC4H4O6. Confirm by dissolving in a few drops of dilute KOH solution, and precipitate the tartrate with a little AgNO3; dissolve the precipitate in a slight excess of ammonium hydroxide and carry out the silver mirror test as described in reaction 2 under Tartrates. 2. If no heavy metals are present and the substance is soluble in water, the silver mirror test may at once be applied. 3. If heavy metals are present, dissolve the original substance in water or in the least possible amount of dilute HC1. Remove 156 QUALITATIVE CHEMICAL ANALYSIS the metals of Groups I. and II., if present, by means of H2S; and those of Group III. with NH4OH and (NH4)2S (Al and Cr will, of course, not be precipitated). The clear filtrate is acidified with HC1, boiled to expel H2S, and is finally rendered alkaline with NH4OH. An excess of CaCl2 is then added and the mixture is shaken vigorously, allowed to stand for a short time, and finally filtered. The precipitate, which may consist of CaC4H4O6, CaC2O4, and Ca3(PO4)2, is treated with a cold, strong NaOH solution to dissolve out the CaC4H4O6, then it is stirred thoroughly, diluted, and filtered. If a precipitate forms on heating the clear filtrate to boiling, a tartrate is indicated. Confirm by filtering while hot, wash the precipitate, and transfer it to a test tube. Add 1 drop of NH4OH and a little AgNO3, and warm. In the presence of a tartrate, a black precipitate or a silver mirror will be formed. GROUP II. Iodides i. Iodides, if present, will be detected in the preliminary ex- amination of the substance with concentrated H2SO4. 2. To a portion of the original or prepared solution acidified with HC1, add a little potassium nitrite solution or chlorine water and then about I cc. of CS2 or CHC13, and shake vigor- ously. An iodide, if present, will color the CS2 or CHC13 violet. 3. AgNO3 in HNO3 solution precipitates yellow Agl, prac- tically insoluble in NH4OH. NOTES Nitrous acid - i.e., a nitrite in acid solution - is preferable to chlorine as an iodine liberator for the reason that an excess of the former does not hinder the reaction (see reaction 4 under Iodides'). The liberated iodine may also be recognized by the blue compound it forms with starch paste. Insoluble iodides are tested for the halogen by one of the methods given under " Insolu- ble Substances," page 172. Bromides i. Bromides, if present, will be detected in the preliminary examination of the substance with concentrated H2SO4. THE ACIDS 157 2. To a portion of the original or prepared solution acid with HC1, add 1 cc. of CS2 or CHC13. Cautiously treat, with small amounts of Cl water and shake vigorously after each addition. In the presence of bromides, the CHC13 or the CS2 will acquire a reddish or yellow color, depending upon the amount of bro- mide present (see reaction 2 under Bromides'). NOTES Iodides, if present, interfere with the test by imparting a violet color to the CHC13 or CS2. If the amount of iodide present is large, chlorine water should be added until an intense violet color is produced in the CHC13 or CS2; the liquid is then carefully decanted or the aqueous portion is removed to another test tube by means of a pipette, and there treated with fresh portions of CS2 and Cl water. If the CS2 is still colored a deep violet, the operation is re- peated until only a faint pink color is imparted to the CS2; now, on adding a little more Cl water and shaking, a brown or reddish color will be produced if a bromide is present. If the amount of iodide in the original solution is small, as is shown by the faint purple color of the CS2, more chlorine water should be added, without decanting the liquid, and the mixture should be shaken after each small addition; in the presence of a bromide, a character- istic brown color will finally be observed in the CS2 layer. Chlorine water exercises a selective action, liberating practically all the iodine first; an excess will oxidize the latter to colorless iodic acid; and on further addition of chlorine water, bromine will be liberated. Insoluble bromides are treated as directed under " Insoluble Substances" page 172. Chlorides In the absence of bromides, iodides, cyanides, ferrocyanides, and sulphocyanides, a white precipitate, obtained with AgNO3 in a solution acid with HNO3, is proof of the presence of chlorides. Chlorides in the presence of iodides and absence of bromides, cyanides, and ferricyanides are tested as follows : To the HNO3 solution, add AgNO3 to complete precipitation, filter, and wash; digest the precipitate for several minutes with cold ammonium hydroxide and filter ; finally acidify the filtrate with HNO3, when the formation of a white curdy precipitate shows the presence of a chloride. 158 QUALITATIVE CHEMICAL ANALYSIS Chlorides in the Presence of Bromides and Iodides may be detected by one of three following methods: - i. The Chromyl Chloride Method. In a small, dry distilling flask, place a mixture of some of the powdered original sub- stance, or the residue obtained by the evaporation to dryness of a portion of the prepared solution, and powdered K2Cr2O7; add 5 cc. of concentrated H2SO4, and heat. Absorb any fumes that may be evolved in dilute NH4OH. The latter will be colored yellow if a chloride was originally present (see reac- tion 4 under Chlorides). 2. Hart's Method. Principle : HI is oxidized by a ferric salt and the I set free is boiled off; HBr is then oxidized with KMnO4 and the liberated Br is removed by boiling; any resid- ual substance which will give with AgNO3 a white precipitate insoluble in HNO3 and soluble in ammonium hydroxide, must be a chloride. The method is not reliable for the detection of very small amounts of chlorides in the presence of relatively large amounts of the others. Method. The solution contained in an evaporating dish is rendered slightly acid with dilute H2SO4, then treated with a concentrated solution of ferric alum, and the mixture boiled until no more iodine is given off. This point may be determined by holding in the escaping vapors a piece of paper moistened with starch paste, which, in the presence of iodine, will be colored blue. When the expulsion of the iodine is complete, KMnO4 solution is added in a quantity sufficient to give the solution a purple color which does not disappear on boiling. The KMnO4 oxidizes the HBr, setting bromine free, and this halogen escapes with the steam. The boiling is continued until a piece of moistened starch-iodide paper, held in the vapor, is no longer turned blue, showing the absence of bromine. If the solution is now purple, showing an excess of KMnO4, a few drops of alcohol are added, the mixture is boiled with stirring, and the brown hydrated MnO2 is filtered off. The fil- trate, which should be colorless, is treated with a few drops of AgNO3. A white precipitate, insoluble in HNO3 and soluble in ammonium hydroxide, proves the presence of a chloride. 3. Vortmann's Method consists in acidifying the prepared solu- tion with acetic acid, adding PbO2, and boiling until no more bromine and iodine are given off (as shown by tests) and the solution on settling is colorless. THE ACIDS 159 All the hydrobromic acid and part of the hydriodic acid are oxidized by the PbO2; the remainder of the iodine, combined in the form of lead iodide, settles on the bottom of the beaker along with the excess of PbO2 added. Filter and wash the precipitate with hot water, and test the filtrate for chlorides with AgNO3. NOTES Cyanides, ferrocyanides, and ferricyanides interfere with the tests for the halides ; they must therefore be removed before the tests are applied. This is accomplished by completely precipitating both cyanides and halides with AgNO3, then filtering, drying, separating the precipitate from the filter, and igniting in a dish or crucible. By this procedure, the cyanogen compounds are decomposed with the separation of Ag, while the silver halides remain unchanged. The latter are best got into solution by fusing them with Na2CO3, extracting the melt with water, and filtering. The filtrate will contain NaCl, NaBr, Nai, and an excess of Na2CO3. The solution is acidified with HNO3 and the tests for the halogen acids are made as given above. Or the residual silver halides may be treated with Zn and dilute sulphuric acid, and the action allowed to continue for half an hour. The halogens go into solution as Zn salts, and, after filtering, the filtrate is tested for the halogen acids as given above. Ferrocyanides A small portion of the prepared solution is acidified with HC1 and a little FeCl3 is then added. In the presence of a ferro- cyanide, a blue precipitate of prussian blue is obtained. Ferricyanides i. To a small portion of the prepared solution acidified with HC1, add a freshly prepared solution of FeSO4; the formation of a dark blue precipitate of Turnbull's blue proves the presence of a ferricyanide. 2. From a nitric acid solution, AgNO3 precipitates reddish brown Ag3Fe(CN)6. Thiocyanates i. Acidify a portion of the prepared solution with HC1 and add FeCl3; a deep red coloration, due to the formation of ferric thiocyanate, proves the presence of a thiocyanate. The 160 QUALITATIVE CHEMICAL ANALYSIS solution is acidified with HC1 to prevent the interference of (i) acetic acid, which, with FeCl3, would give a red coloration owing to the formation of ferric acetate; and (2) to prevent tartaric acid and other hydroxy- acids from combining with FeCl8. NOTES Ferri- and ferro-cyanides interfere by yielding precipitates or blue solutions which may completely mask the red color due to HCNS. Iodine, set free by oxidizing agents which may be present, also interferes with this test by the color it imparts to the solution. All these interfering substances may be removed by distilling the HCNS. This is accomplished by adding to a por- tion of the prepared solution acidified with HC1 a little SnCl2 sufficient in amount to reduce any I or Br present to their corresponding halogen acids, boiling, and then absorbing the HCNS, which distils over, in a test tube con- taining FeCl3, when characteristic red Fe(CNS)3 will be formed. Cyanides 1. Cyanides will have been detected by the odor of HCN in the preliminary examination with HC1 and concentrated H2SO4. 2. To a portion of the prepared solution, add 2 cc. of NaOH solution, and treat with a little FeSO4 and a few drops of FeCl3; heat gently for a short time and then acidify with HC1. A blue precipitate of Fe4[Fe(CN)6]3 proves the presence of a cyanide (see reaction 3 under Cyanides'). NOTES The presence of a ferricyanide, ferrocyanide, or thiocyanate interferes with test 2. When these are present, treat some of the solid substance in a small distilling flask with dilute H2SO4, and warm gently under a hood, catching the HCN that is given off in a test tube containing 1 cc. of (NH4)2SX. Pour the contents of the tube into a small evaporating dish and evaporate on the water bath to dryness. To the residue add 1 drop of dilute HC1 and a few drops of FeCl3; a red coloration proves the presence of a cyanide. The HCN which distils over may also be caught in a test tube containing NaOH solution and test 2 applied. A more difficult and perhaps better method is to put into a small distilling flask about 15 cc. of water that has been saturated with CO2, adding an excess of solid NaHCO3, and finally some of the powdered original substance. Quickly stopper the flask and distil under a hood, catching the distillate in THE ACIDS 161 a little NaOH solution, and apply test 2. The HCN .contained in the dis- tillate is derived only from the simple cyanide by the action of the rela- tively stronger carbonic acid. Sulphides 1. Most sulphides will have been detected in the preliminary treatment with dilute HC1 by the evolution of H2S, which may be recognized by its odor and by its property of turning lead acetate paper black. 2. If aqua regia or strong HNO3 is used to get a sulphide into solution, the latter will be oxidized to sulphate with more or less separation of sulphur. 3. If treatment with HC1 does not effect the decomposition of a sulphide, add Zn and dilute H2SO4 to the substance con- tained in a test tube, loosely stoppered with a cork covered with filter paper moistened with lead acetate, and allow the mixture to stand for some time. Sulphides which do not respond to test 1 are usually decomposed by this treatment, yielding H2S, which will blacken the lead acetate paper. 4. Sulphides unattacked by acids should be fused with a little NaOH on a porcelain crucible cover; if the melt is placed on a bright silver coin and moistened with a drop of water, a black stain due to Ag2S will form. It must be remembered, however, that sulphates in the pres- ence of organic matter may be reduced to sulphides when fused with NaOH, and thus give the final test. GROUP III. Nitrates i. Acidify a portion of the prepared solution or the concen- trated water extract of the original substance with dilute H2SO4, then add an equal volume of concentrated H2SO4, and cool thoroughly in a stream of running water. Incline the tube and carefully add 2 to 3 cc. of a strong freshly prepared FeSO4 solu- 162 QUALITATIVE CHEMICAL ANALYSIS tion, and allow the mixture to stand. In .the presence of a nitrate, a brown ring will form at the junction of the two liquids. NOTES Chromates, iodides, bromides, chlorates, ferricyanides, ferrocyanides, and permanganates interfere with the test. Nitrites give the same reaction (see reaction 2). Iodides and bromides in contact with concentrated H2SO4 are partially oxidized with the liberation of free I and Br; these, by coloring the solution, interfere with the test. The halides may be removed by precipita- tion with Ag2SO4 (free from nitrates). Chromates will be reduced by FeSO4, yielding green Cr2(SO4)3, which will obscure the brown color. Permanganates by their strong purple color will mask the reaction. Both chromates and per- manganates may be removed by adding solid Na2SO3 and dilute H2SO4, boil- ing until the solution is green, and then precipitating the Cr and Mn salts with an excess of Na2CO3. Filter, acidify the filtrate with dilute H2SO4, and make the test on the resulting solution. Chlorates interfere on account of C1O2, which will form on adding concentrated H2SO4; these will also be reduced by the above treatment. Ferro- and ferri-cyanides yield with FeSO4 blue precipitates, and hence interfere with the reaction. These may be re- moved by the addition of ferrous and ferric salts and a little dilute H2SO4, heating the mixture to boiling, and adding BaCl2 to precipitate the H2SO4. The H2SO4 and BaCI2 are added to form heavy BaSO4, which will have the effect of carrying down the blue precipitates, which are difficult to filter when alone. Although provision is made for the removal of interfering elements, these are of rare occurrence in mixtures ordinarily met with. The ring test is therefore the one most frequently employed for the detection of the nitrates. 2. Reduction to Ammonia. Render either the aqueous extract of the original substance or the prepared solution strongly alka- line with NaOH, and boil with stirring until no more ammonia is given off. Add some aluminum turnings, or a mixture of granulated zinc and iron filings, and heat again. In the presence of a nitrate or nitrite, the odor of ammonia will be evident (see reaction 5 under Nitrates}. Nitrites 1. Nitrites will have been detected in the preliminary exami- nation with dilute HC1; the NO2 fumes given off may be readily detected by their property of turning starch-iodide paper blue. 2. Brown Ring Test. The same as with nitrates except that THE ACIDS 163 in this case acetic acid or dilute H2SO4 may be used instead of concentrated H2SO4. Acetates i. Acetates are detected in the preliminary treatment of the original substance with concentrated H2SO4 by the odor of vinegar. 2. Treat some of the original solid substance contained in a small evaporating dish or beaker with i to 2 cc. of amyl alcohol and 5 cc. of concentrated H2SO4, and gently. The charac- teristic odor of amyl acetate ("pear essence") indicates the pres- ence of an acetate. If ethyl alcohol is used, the odor of ethyl acetate will be made evident on warming the mixture (see reac- tion 2 under Acetates). Chlorates i. Chlorates are recognized by their behavior when treated with concentrated H2SO4. Heat about i cc. of concentrated H2SO4 in a test tube, remove from the flame, and under a hood {pointing mouth of tube toward back of hood) add a very minute amount of the original substance. In the presence of a chlorate, greenish yellow C1O2 will be evolved; the evolution is accom- panied by a slight explosion. 2. In the absence of halogen acids, a small portion of the solid substance is ignited in a small porcelain dish at a tempera- ture just below a red heat; it is then cooled, extracted with water, transferred to a test tube, and finally treated with a few drops of silver nitrate. A white curdy precipitate of AgCl proves the presence of a chlorate. 3. If halogen acids are present, they must be removed by adding to the boiling solution acidified with HNO3 an excess of AgNO3 and filtering. To the filtrate add a crystal of Na2SO3. A precipitate of AgCl indicates the presence of a chlorate. (See reaction 4 under Chlorates I) PART III ANALYSIS OF GROUP III. (METALS) IN THE PRESENCE OF ORGANIC MATTER, PHOSPHATES, OXALATES, ETC. The phosphates, fluorides, oxalates, borates, and silicates of the metals of Groups III. and IV., including Mg, are soluble in mineral acids, but are precipitated when the free acid which holds them in solution is neutralized by ammonium hydroxide. Should any of these acids be present in the original solution, they will offer no difficulties in the analysis of Groups I. and II., for in these the solution is kept acid. On proceeding, however, to precipitate Group III. the solution is first rendered alkaline, and, as a consequence, there will be precipitated along with the metals of Group III. part or all of the metals of Group IV. as phosphates, oxalates, etc., depending upon the quantity of these acid radicals present. It is evident, therefore, that a different procedure from that given must be followed for the analysis of Group III. if these acids are present. It will be recalled that certain non-volatile organic acids and compounds, as tartaric acid, citric acid, sugar, and starch, hinder the precipitation of the trivalent elements Al, Cr, and Fe(-ic) as hydroxides and basic acetates. For this reason, before proceed- ing with the Third Group analysis, it is necessary to test for non- volatile organic matter, and, if found, to remove it. The presence of organic matter will have been indicated on heating a small portion of the original substance in a tube closed at one end. Blackening of the residue, accompanied by a burnt odor, indicates the presence of organic matter. If the substance under examination is a solution, evaporate a small portion to 164 THE PHOSPHATE SEPARATION 165 dryness, heat to a dull red heat, and look for indications of organic matter. Test for an Oxalate. To a small portion of the filtrate from Group II., from which the H2S has been expelled, add an ex- cess of Na2CO3, boil vigorously for a minute, and filter. Render the filtrate slightly acid with acetic acid, boil off the CO2, and then add an equal volume of a saturated CaSO4 solution; a white crystalline precipitate proves the presence of an oxalate. Organic matter and oxalates may both be removed by the following procedure: Mix the residue obtained by evaporating the filtrate from Group II. to dryness with four times its weight of NH4NO3, and add this in small portions to about 2 g. of NH4NO3 that have been previously heated to fusion in a porce- lain crucible. Heat until the mass is in a state of quiet fusion, then raise the temperature slightly to decompose the excess of NH4NO3, and cool. Dissolve the melt in water, using a little HC1, if necessary, to effect solution. If the filtrate is orange or yellow, a chromate is indicated; the latter will have been formed in consequence of the oxidizing action of NH4NO3 on any chromic salt present. To reduce the chromate to a chromic salt, concentrate the solution by evaporation to about § cc., add an equal volume of concentrated HC1 and a little alcohol, and heat to boiling until the solution is green (see reaction under Chromates, page 123). Evaporate to a small bulk to remove the excess of HC1, and if phosphates are absent analyze for Groups III., IV., and V. in the usual way. Test for Phosphates. To about 2 cc. of the filtrate from Group II., from which the H2S has been removed, add a few drops of concentrated HNO3 and evaporate nearly to dryness; take up with a little dilute HNO3, transfer to a test tube, add an equal volume of ammonium molybdate solution, and heat gently {do not boil). A yellow precipitate of (NH4)3PO4 • 12 MoO3 proves the presence of a phosphate. Silicates, if present, should have been detected in the prelimi- nary examination of the solid substance with a NaPO3 bead ; and should have been removed, preferably before proceeding with the metallic analysis, by evaporating the solution acid with HC1 166 QUALITATIVE CHEMICAL ANALYSIS or HNO3* to dryness, desiccating at 1200, extracting the residue with a few cc. of concentrated HC1,* diluting, heating, and filter- ing off the dehydrated and insoluble silica. The latter may then be verified by the NaPO3 bead test, or by treatment in a plati- num crucible with CaF2 and concentrated H2SO4, heating, and testing the escaping vapor with a drop of water (see reaction 1, page 113). If, however, the analysis has been begun without regard to the presence of silicates, the filtrate from Group II. should be tested for this acid radical, and, if found, removed by the pro- cedure just outlined, before proceeding to precipitate Group III. Borates and fluorides are usually held in solution by the NH4C1 present, and hence in their presence no modification of the usual scheme need be made. Outline of Method to be Used in the Presence of Phosphates Oxalates, silicates, fluorides, borates, and non-volatile organic matter having been disposed of, it only remains to provide a method of analysis for the Third Group metals which shall include the presence of phosphates. The scheme which follows is based upon the fact that of the phosphates of Groups III. and IV., only those of Al, Cr, and Fe(-ic)are insoluble in acetic acid; if, therefore, the iron is oxidized and the free HC1 is replaced by acetic acid, part or all of the trivalent metals pres- ent will be precipitated as phosphates, depending upon the quantity of phosphoric acid present. If the amount of PO4 is less than that required to combine with all of the Fe(-ic), Al, and Cr, the precipitate which forms will contain all the PO4. If, on the other hand, the quantity of PO4 present exceeds that required to unite with the trivalent metals, it will be evident that more trivalent metals will have to be added to completely pre- cipitate the PO4. The metallic radical used for this purpose is Fe(-ic), partly because its phosphate is the least soluble in acetic acid, but chiefly because it is possible, when a salt of ferric iron is used, to tell when all the phosphate has been precipitated ; for * If metals of the first group are present, use HNO3 instead of HC1. THE PHOSPHATE SEPARATION 167 when this condition is reached, any excess of ferric salt added will yield with the acetate radical present red ferric acetate, which can readily be recognized by its color. Now, on adding an excess of NaC2H3O2, largely diluting, boiling the solution, and rapidly filtering, the separation of all of the trivalent metals, including the excess of iron added, together with all the phos- phoric acid, is accomplished (see reactions 6 and 7 under Iron, pages 67 and 68). The filtrate, now free from PO4, is concen- trated by evaporation, and then treated for the remaining metals of Groups III. and IV. in the usual way. The Phosphate Separation If phosphates are shown to be present, the entire filtrate from Group II. is boiled until all of the H2S is expelled, a few drops of concentrated HNO3 are then added, and the solution boiled for several minutes to insure the complete oxidation of the iron present. Test a separate small portion, about 2 cc., for iron by adding a few drops of K4Fe(CN)6. A blue pre- cipitate proves the presence of iron. The remainder of the solution is transferred to a beaker of 500 cc. capacity and is treated with ammonium hydroxide, added drop by drop with vigorous stirring, until a slight precipitate is produced which persists after stirring for 3 minutes. Now add cazttiously, with constant stirring, dilute HC1, drop by drop, until a clear solu- tion is obtained; then add 8 g. of NH4C2H3O2 and 8 cc. of 50 per cent, acetic acid. If the solution is not red, add FeCl3 solution drop by drop, with stirring, until the solution assumes a deep red color, avoiding an excess. In the presence of a pre- cipitate, the color of the solution may be seen by filtering a small portion; the filtrate should give, when made alkaline with NH4OH, a reddish brown precipitate of Fe(OH)3, showing that an excess of Fe(-ic) is present; if a light-colored precipitate is obtained with NH4OH, more FeCl3 should be added. Now dilute the solution with hot water to 400 cc., heat rapidly to boiling, and boil for 3 minutes only. Allow the precipitate to settle, filter on a large fluted filter contained in a 10 cm. funnel, and wash 168 QUALITATIVE CHEMICAL ANALYSIS with hot water. The residue may consist of the phosphates and basic acetates of Al, Cr, and Fe(-ic), and it may also con- tain small amounts of Ni, Mn, and Zn. The filtrate, which should not have a yellow color, is at once evaporated in a large evaporating dish to 50 cc., and any precipitate which separates out is filtered off and rejected. The filtrate, now concentrated and free from PO4, is analyzed as usual for Groups III., IV., and V.; the tests for Al, Cr, and Fe should be omitted, as these metals will be in the residue from the basic acetate separation. The precipitate, consisting of the phosphates and basic acetates of Al, Cr, and Fe (-ic), is transferred to a beaker with the aid of 20 cc. of water, 2 g. Na2O2 are added, the mixture is boiled, with stirring, and finally filtered. The residue, consisting of Fe(OH)3, is rejected. The filtrate may contain Na3AlO3, Na2CrO4, Na3PO4, and an excess of NaOH. If the solution is yellow, chromium is present; if colorless, Cr is absent. Alumi- num is detected by acidifying the solution with HNO3 and then rendering alkaline with ammonium hydroxide ; the white gelati- nous precipitate of A1PO4 or A1(OH)3 is filtered off, washed with hot water several times, and the presence of aluminum is confirmed by igniting with a few drops of Co(NO3)2 in the usual way. The filtrate will contain the chromium as Na2CrO4. PREPARATION OF THE SOLUTION The preliminary tests completed, the next step in the system- atic examination is to get the substance into solution. This is accomplished by the use of the solvents, water, nitric acid, hydrochloric acid, and aqua, regia. In determining the solvent, it is advisable to experiment with small portions of the original substance at first, finally treating, after the proper solvent has been found, about one gram of the original material for the analysis. With mixtures, more than one solvent may be re- quired. In such a case, it is a good plan to keep, and analyze separately, portions dissolved by different solvents; the addi- tional labor involved will be compensated by the information PREPARATION OF THE SOLUTION 169 which this procedure will supply concerning the manner in which the metallic and acid radicals are united. Treatment with Water Treat a small quantity of the finely powdered substance with about 25 cc. of water and heat to boiling. If solution takes place, treat one gram of the sample in the same way and analyze the resulting solution for the metals and acids. Test the aque- ous solution with litmus; if alkaline, the presence of a car- bonate, hydroxide, sulphide, phosphate, borate, or cyanide is indicated; if acid, it may point to an acid salt, free acid, or the salt of a heavy metal. If no solution appears to have taken place, filter, and evaporate some of the clear filtrate on a watch glass to dryness; if only a slight residue remains, the substance may be considered insoluble in water; if a moderate amount of residue is left, it indicates that the mixture contains a water- soluble component. In that case, treat a gram sample with boiling water and filter. Analyze the aqueous extract for acids and bases. Treat the residue with acids as given below. NOTES If iodides or bromides, particularly the former, have been indicated in the preliminary test with cone. H2SO4, the original substance, whether it dissolves wholly or in part in water, must be treated for the removal of these halides before the analysis for the metals is begun. This is accomplished by treating the substance with HNO3 and heating until no more I is evolved. If iodides are not removed, there will be danger of forming explosive brownish black nitrogen iodide, in making the preliminary test for Group III., because of the action of iodine set free by the treatment with HNO3, on the ammonia which is next added. Treatment with Acids If the substance is insoluble in water, treat it or the residue from the water treatment successively with hot dilute HC1 and hot concentrated HC1. If these fail to effect solution, try the action of dilute and concentrated HNO3 on separate small por- tions; if these also fail, add HC1 to the mixture containing HNO3, thus forming aqua regia, and heat. If still insoluble, 170 QUALITATIVE CHEMICAL ANALYSIS examine it by the method given for "Insoluble Substances" page 172. NOTES 1. During the treatment with HC1, indications of the presence of certain acids will be given (see "Preliminary Examination" page 143). 2. If Hg and As are present, boiling with HC1 may cause these elements to be lost by volatilization; the remedy is to be sought in the use of HNO3, which oxidizes them into compounds which are not readily volatile. 3. If complete solution with HC1 is obtained, the absence of Ag, Hg (-ous), and large amounts of Pb is indicated. Evaporate the solution nearly to dry- ness to expel most of the acid, dilute, and analyze the resulting solution, beginning with Scheme II. 4. Treatment with concentrated HC1 may cause the precipitation of Pb in the form of crystalline needles of PbCl2; when this is the case, filter them off, dissolve in boiling water, and test for Pb. 5. If the HC1 treatment causes gelatinous silicic acid to separate, evapo- rate the mixture on the water bath to dryness, dehydrate by heating to 1200 C. for half an hour, extract with 3 cc. of hot concentrated HNO3 or HC1, dilute, heat, and filter off the SiO2. The filtrate is then examined for the metals. 6. If HNO3 has been used as a solvent, boil down the liquid to about 1 cc., dilute with 20 cc. of water; and if the solution clouds on dilution, clear with a few drops of HNO3 and analyze the solution for all groups. 7. When aqua regia is employed, the smallest possible amount should be used; the solution should then be evaporated down to about 1 cc. to de- stroy the excess, diluted with 15 cc. of water, and the chlorides of Group I. filtered off and analyzed. The filtrate is treated with 5 cc. of concentrated HC1 and again evaporated down to 1 cc., diluted somewhat, and analyzed for the metals, beginning with Group II. 8. With few exceptions, the following substances, while insoluble in water, are dissolved by boiling HC1 or HNOs: all phosphates, arsenates, arsenites, borates, carbonates, oxalates, and tartrates (the alkali salts are soluble in water) ; also the oxides, hydroxides, sulphides of the heavy metals, alumina, magnesia, and a number of metallic iodides and cyanides. Oxides of Al, Fe, and Cr which have heated intensely do not dissolve readily in these acids. 9. Because of its oxidizing action, HNO3 dissolves sulphides and most metals and alloys which are not attacked by HC1; the latter, on the other hand, dissolves the oxides of Sn and Sb, as well as MnO2, all of which are not dissolved by HNO3. METALS AND ALLOYS 171 METALS AND ALLOYS From 0.5 to 1 gram of the metal or alloy, in the form of shav- ings, foil, filings, or turnings, is treated with 20cc. of HNO3 (sp. gr. 1.2) and is heated gently (under a hood) until no more red fumes of NO2 are given off; it is then diluted with an equal volume of water, heated again for a few minutes, and filtered if necessary. If complete solution takes place, the absence of Au, Pt, Sb, and Sn is shown; * in that case, expel the excess of HNO3 by evaporation, dilute with water, and analyze the solu- tion for all groups except IV. and V. Mg, however, must be included. (a) If a metallic residue is left, it is probably Pt or Au f, or both. (^) If a white residue is left which is insoluble on dilution and heating, it may consist of hydrated SnO2 or Sb2O5, or both, admixed with arsenic in the form of tin arsenate, phosphorus in the form of tin phosphate, bismuth as Bi2O3, and traces of Cu and Pb. Filter. Filtrate. Evap- orate to drive off excess of HNO3. Add HC1 to ppt. 1st group and fil- ter. Analyze resi- due for Group I. Treat filtrate with H2S and filter. Analyze filtrate for Group III. and Mg. Dissolve residue in hot dil. HNO3 and unite with corre- sponding solution from first residue. Dry and fuse residue in a porcelain crucible with 4 times its weight of a mixture of equal parts of Na2CO3 and S ; cool, extract melt with hot water, and filter. Residue is CuS, Bi2S3, PbS. Dissolve in hot dil. HNO3 and combine with the cor- responding solution ob- tained from the first filtrate and proceed as directed in the analysis of the main fil- trate in Scheme II. A. Filtrate will contain the As, Sb, and Sn as thio-salts + an excess of Na2S. Just acidify with HC1, filter, and reject filtrate. Residue may consist of As2S5, Sb2S5, and SnS2 + S ; analyze accord- ing to Scheme II. B. * Minute amounts of Sb dissolve completely in HNO3; silver alloys containing a very small amount of Pt are completely dissolved by HNO3. f A black residue of carbon or graphite is sometimes left. QUALITATIVE CHEMICAL ANALYSIS 172 NOTES 1. If solution of the alloy does not readily take place, and Pt and Au are absent, treat with HC1; the latter is more satisfactory than HNO3 as a solvent for Al, since the latter is only difficultly soluble in HNO3. HNO3 is used in- stead of HC1, first, because it is the better solvent for metals and alloys ; second, because treatment with HC1 would convert any P, S, and As usually present as phosphide, sulphide, and arsenide, respectively, into PH3, H2S and AsH3, which would be lost by volatilization. HNO3 oxidizes these elements to their corre- sponding acids, viz.: H3PO4, H2SO4 and H3AsO4. These acids only need to be tested for in the analysis of alloys. 2. A small white residue may be boiled with concentrated HC1 and the re- sulting liquid divided into two portions. One portion is tested with Pt and Zn couple for Sb. The other is heated with an iron nail for some time and the clear decanted solution tested for Sn by the addition of HgCl2. 3. A portion of the HNO3 filtrate may be tested for H3PO4 and H2SO4, and, if found, P and S reported. INSOLUBLE SUBSTANCES By an insoluble substance we mean one which cannot be got into solution by the action of the acids taken singly or together. The most common insoluble substances are the following: - C, S, Ag3Fe(CN)6, Ag4Fe(CN)6, AgCN, AgCl, AgBr, Agl, BaSO4, SrSO4, CaSO4, PbSO4, PbCl2, fused PbCrO4, ignited or anhydrous chromic salts, ignited and native oxides, as A12O3 (corundum), Fe2O3, SnO2 (cassiterite), Cr2O3, Cr2O3 • FeO (chrome-iron ore), CaF2, Sb2O4, Fe4[Fe(CN)6]3, SiO2, and cer- tain silicates. Carbon is generally recognized by its black color, insolubility in aqua regia, and combustibility when heated strongly on plati- num foil. When heated with KNO3, deflagration ensues with the formation of K2CO3; this method is not applicable to graphite, which may be determined by its physical properties. Sulphur is recognized (in the preliminary testing) by the for- mation of a yellow sublimate and evolution of SO2 when heated in a glass tube. When S and C are present, it is desirable to remove them by roasting in an open porcelain crucible. INSOLUBLE SUBSTANCES 173 Treatment with aqua regia will have converted the simple and complex cyanides, as well as all the halides of Ag, into AgCl; the latter dissolves to a large extent in the strong acids, but separates out again when the latter are diluted. Long treatment with aqua regia will dissolve prussian blue, but the following method is preferable for complex insoluble cyanides in general: Boil the substance with a strong solution of NaOH, dilute and filter; the residue will contain the heavy metal as hydroxide, while the filtrate will contain the acid radical in the form of the Na salt and may be examined by the methods already given.* PbSO4 and PbCl2 may be dissolved by treatment with hot, strong NH4C2H3O2 solution. The extract is divided into three portions: in the first, test for Pb by the addition of a little H2SO4 or K2CrO4; in the second, test for Cl by diluting, acidi- fying with HNO3, and adding AgNO3; and in the third, test for SO4 by acidifying with HC1, filtering if necessary, and add- ing BaCl2. Sulphates of the Alkaline Earth Metals are best fused in a platinum crucible with five times their weight of anhydrous Na2CO3; the melt is then completely extracted with hot water, filtered, and the residue is thoroughly washed with water. The residue will consist of the carbonates of alkaline earths, which may readily be got into solution with hydrochloric acid, and the resulting solution tested in the usual way. The water ex- tract will contain the acid radical as Na2SO4 and an excess of Na2CO3. SrSO4, CaSO4, and PbSO4 may be quantitatively converted into the corresponding carbonates by prolonged boiling with a concentrated Na2CO3 solution. If the residue after filtering is thoroughly washed free from alkali, it may then readily be dissolved by acid. BaSO4 requires several treatments for its complete transfor- mation by Na2CO3. One treatment changes about 80 per cent, of this sulphate into carbonate. * For method of analysis of insoluble double cyanides not precipitated by excess of NaOH, see page 133 under Ferrocyanides, reaction 3. 174 QUALITATIVE CHEMICAL ANALYSIS Fused PbCrO4, Cr2O3, Chrome Iron Ore, and Ignited Chromic Salts are best fused with Na2O2 in a nickel crucible or with a mixture of Na2CO3 and NaNO3. By this treatment, soluble Na2CrO4 is formed, which after treatment with water is separated by filtering and tested for in the filtrate. SnO2 and Sb2O4 are got into solution by fusing in a porcelain crucible with three times their weight of Na2CO3 mixed with an equal quantity of S. The melt is extracted with hot water and filtered. The filtrate will contain the Sb and Sn in the form of thio-salts; it is just acidified and the precipitate treated according to Scheme II. B. A12O3 and Fe2O3 are fused in platinum with KHSO4 or K2S2O7, whereby they are converted into soluble sulphates. 'Fusion with Na2CO3, followed by acid treatment, also takes these oxides in solution. Silver Halides may be treated by one of the following two methods: - i. Fuse with Ma^Og in a porcelain crucible. The product will consist of metallic silver and the sodium salts of the halides ; extract with water and filter. Test the residue for Ag and the filtrate for halogens. 2. Treat with Zn and dilute H2SO4 in a crucible or small dish; allow the action to continue for 20 minutes, and then filter. Test the residue for Ag and the filtrate for halogens. SiO2 and Silicates are recognized by the "skeleton" NaPO3 bead. Silicates are usually decomposed by fusing with five times their weight of a mixture of equal parts of anhydrous Na^Og and K2CO3, to which about o. 1 g. of KNO3 is added. This will be taken up more fully in the systematic treatment. CaF2 is decomposed by heating the finely ground material with concentrated H2SO4 in a platinum dish or crucible, and evaporating until no more SO3 fumes are given off. The resi- due will be CaSO4; extract it with water for some time and filter. Test the filtrate for Ca with (NH^CjO^ INSOLUBLE SUBSTANCES 175 Systematic Treatment of Insoluble Substances Before proceeding with the systematic treatment of a residue insoluble in acids, it is desirable to make the following prelimi- nary tests: - i. Examine the residue carefully with a lens and determine whether or not the substance is homogeneous. 2. Determine whether free C and S are present; if present, remove by roasting. 3. Chromic oxide is green and will be made evident by yield- ing a green bead with NaPO3 which is unaffected by the reduc- ing flame; at the same time, indications of SiO2 or of a silicate will also be obtained. 4. If the substance is white, treat it with a little (NH4)2S. If it blackens, the presence of Ag or Pb salts is indicated; con- firm as directed in (5). 5. If black or colored, mix a small amount with Na2CO3 and heat on charcoal with a reducing flame; a lustrous malleable globule shows the presence of either Pb, Ag, or Sn. Flatten the globule in a mortar and heat with dilute HNO3. A clear solution indicates the absence of Sn ; a white residue, the presence of Sn. Divide the HNO3 solution into two portions. To the first add HC1; a white precipitate soluble in NH4OH shows the pres- ence of Ag. To the second portion add dilute H2SO4; a white precipitate is PbSO4. If no globule is obtained and the white substance is not blackened by (NH4)2S, the absence of Pb and Ag is shown. 6. Flame Test. Take up some of the material on a mois- tened Pt wire and hold in the reducing flame for some time. The reducing flame will change the sulphates of the alkaline earths to sulphides. Moisten the wire with a drop of HC1 and hold in the colorless bunsen flame. Alkaline earths impart their characteristic colorations to the flame. 7. If test 3 above is found unsatisfactory for SiO2 and Cr, the following may be used : - SiO2. Mix the finely powdered substance in a platinum cruci- ble or lead tube with a small quantity of CaF2 (SiO2 free), add 176 QUALITATIVE CHEMICAL ANALYSIS concentrated H2SO4, and warm. Hold in the escaping vapors a drop of water on the loop of a Pt wire; if the drop becomes turbid, SiO2 is present. Chromium. Prepare a Na2CO3 bead. Take up a little of the substance mixed with KC1O3 and heat. Place bead in about i cc. of water and heat. A yellow solution indicates Cr. If the substance contains Pb salts, these may readily be removed by repeatedly digesting with hot (NH4)2C4H4O6 or NH4C2H3O2 solution, filtering, and testing the filtrate for Pb, SO4, and Cl. The residue, which should be thoroughly washed and tested until free from Pb salts [shown by wash water no longer reacting with (NH4)2S], is then treated with KCN solution to dissolve AgCl, AgBr, Agl, and AgCN, filtered, and washed. The treatment with KCN is made only when Ag salts are shown to be present by preliminary test with Zn + H2SO4. The KCN extract is tested for Ag by adding (NH^S, filtering off the Ag2S, washing, and dissolving in hot dilute HNO3. To the clear solution, add HC1; a white precipitate is AgCl. If S and C are present, heat in an open porcelain crucible till all C and S are oxidized. Mix the substance, now free from Pb and Ag salts, in a platinum crucible with six times its weight of a mixture of equal parts of anhydrous powdered K2CO3 and Na2CO3 + o.i g. of KNO3. (If reducible metals have not been removed, treat in a Ni crucible. Porcelain cannot be used, as it gives up SiO2, Ca, and Al to the melt). Heat over a blast lamp till the mass is in a state of quiet fusion. Remove flame. When cool, transfer crucible to a casserole or evaporating dish, and extract with boil- ing water. Break up the mass with pestle during extraction. Allow finally to settle and filter. Scheme for the Treatment of Insoluble Substances INSOLUBLE SUBSTANCES 177 Filtrate may contain Na2SiO3, Na2CrO4, NaF, Na3PO4, Na3A103, Na2SO4, Na2SnO3, Na3AsO4, (NaSbO3), Na2MnO4, and Na2CO3, as well as K salts of these acids. Divide into two equal portions. ist portion. Acidify with HC1 and add a small portion to the HNO3 solution of the residue; if no ppt. forms, unite the two fil- trates, evap. to dryness, dehydrate SiO2 by heating at 120° C. till all HC1 is driven off. Extract the residue with 3 cc. of cone. HC1, stir thoroughly, add 25 cc. of water, and boil with stirring; filter. Residue may consist of BaCO3, SrCO3, CaCO3, and possibly A12O3, SnO2, and unattacked SiO2; wash several times with hot water. Treat with hot dil. HNO3 and filter. Filtrate is to be united with HC1 solution of aqueous extract of melt unless a ppt. forms; in that case, keep the solutions separate. The ppts. produced by the same group reagents may be united and examined together. Residue may consist of SiO2, SnO2, and A12O3. Fuse in a Ni crucible with NaOH. Extract with water and filter. Filtrate con- tains Na2SnO3, Na3A103 + Na2SiO3. Test for Al and Sn, if not already found. Residue SiO2. Heat filtrate to boiling and treat with H2S. Without fil- tering, dilute with cold water to 100 cc. and saturate again with H2S. Filter. Analyze residue for II. A and II. B. Analyze filtrate for all other groups. 2d portion. Treat for acids. Detection of Alkalies in Insoluble Silicates The J. Lawrence Smith Method. One gram of the finely ground mineral is first pulverized in an agate mortar with its own weight of C.P. NH4C1, and the resulting mixture is then thoroughly mixed with 8 grams of alkali-free CaCO3 and heated in a covered platinum crucible, gently at first and finally to a dull red heat, for 40 minutes. The crucible should be placed in a hole made in a piece of thick asbestos board in such a way that only two-thirds of the crucible can be directly heated by the burner. The mass does not fuse but sinters. The active agent is fused CaCl2, which decomposes the silicate with the formation of chlorides of the alkali metals. After cooling, the crucible with its contents is transferred to a casserole, boiled with water, and the CaO is allowed to slake. The last opera- tion may be hastened by crushing any lumps with a pestle. 178 QUALITATIVE CHEMICAL ANALYSIS After standing for one hour, the mixture is filtered, and the filtrate is freed of lime by rendering it alkaline with ammonium hydroxide, heating, adding (NH4)2CO3 to complete precipita- tion, and finally a little (NH^C^O^ Filter, evaporate the filtrate to dryness, and ignite the residue to drive off NH4 salts. The residue is then treated in the usual way for K and Na. Acid Analysis of Minerals and Metallurgical Products With a few exceptions, minerals and slags need only be tested for sulphides, carbonates, silicates, phosphates, borates, sulphates, fluorides, and chlorides. Carbonates and sulphides may be detected by treatment with HC1, and silica or sili- cates by the NaPO3 bead test. In the HNO3 solution of the finely powdered substance, the tests for phosphates, chlorides, and sulphates (in the absence of sulphides) may be made. If the chloride is present in an insoluble form, as AgCl, it should be treated by one of the methods already mentioned (see page 174). In the absence of sulphides, the test for sulphates may also be made by fusing the original substance with Na^COg, ex- tracting the melt with boiling water, and filtering; the filtrate, after acidifying with HC1 and boiling to drive out the CO2, is then treated with BaClg. If sulphides are present, boil the finely powdered substance, with constant stirring, with a satu- rated solution of Na2CO3, then filter, acidify the filtrate with HC1, and add BaCl2; or if the original substance dissolves com- pletely in HC1, the resulting solution may be treated with BaCl2. For the detection of fluorides in the presence of silicates, see test 3 under Fluorides, page 154. For the detection of borates in silicates undecomposed by concentrated H2SO4, see reaction 6 under Borates, page 117. The tests for borates and fluorides, when these occur together in combination with silicates, may be carried out in one sample by fusing about one gram with Na^Og, extracting the mass with boiling water, and filtering. The filtrate will then contain Na2SiO3, NaF, NaBO2 + the excess of Na2CO3. A portion of this solution, after slightly acidifying with HC1, may be tested TABLE OF SOLUBILITIES 179 for boric acid with turmeric paper. The remainder of the aque- ous extract is then tested for a fluoride, as described in test 3 under Fluorides (see page 154). TABLE OF SOLUBILITIES* Showing the classes to which the compounds of the commonly occurring elements belong in respect to their solubility in water, hydrochloric acid, nitric acid, or aqua regia. P reliminary Remarks For the sake of brevity, the classes to which the compounds belong are expressed by letters. These have the following signification: W or w, soluble in water. A or a, insoluble in water, but soluble in hydrochloric acid, nitric acid, or in aqua regia. I or i, insoluble in water, hydrochloric acid, or nitric acid. Further, substances standing on the borderlines are indicated as follows: W-A or w-a, difficultly soluble in water, but soluble in hydrochloric acid or nitric acid. W-I or w-i, difficultly soluble in water, the solubility not being greatly increased by the addition of acids. A-I or a-i, insoluble in water, difficultly soluble in acids. If the behavior of a compound to hydrochloric and nitric acids is essentially different, this is stated in the notes. Capital letters indicate common substances used in the arts and in medicine, while the small letters are used for those less commonly occurring. The salts are generally considered as normal, but basic and acid salts, as well as double salts, in case they are important in medicine or in the arts, are referred to in the notes. The small numbers in the table refer to notes on the following pages. * Taken from Wells' translation of the 16th German edition of Fresenius' Quali- tative Analysis. 180 QUALITATIVE CHEMICAL ANALYSIS Potassium | Sodium Ammonium | Barium | Strontium Calcium । Magnesium | Aluminum Chromium | Zinc | Manganese | | Nickel | Cobalt Oxide W w W W w W-A A A A&I A A A Chromate W. w-a w-a a w a a Sulphate w,3.is w WH*20'30 I I W-I W Wia-ii W&Ii5 W w w w Phosphate w W8 w8-12 a a All au a a a a a a Borate w a a a a a a a a a Oxalate w3 w w a A w-a w-a a Fluoride w w-a w-a A-I a-i w w w-a a w-a wa Carbonate w4 Ww w A A A A A A A A Silicate w w a a a a a a Chloride W3J wn.M W w W w w W&I w w w W Bromide w W w w w w w w&i w w w w Iodide w w w w w w w w w Cyanide w w a a a-i a-i Ferrocyanide w w w w-a w w w A-I a i i Ferricyanide w w w w w a i i i Thiocyanate w w w w w Sulphide w w w w w-a4B a a a-i A16 A a18 a19 Nitrate w w w w w w w Chlorate w w w w Tartrate w W7 A Citrate Malate Succinate Benzoate Salicylate w w w w-a w-a w-a w Acetate w w w w w w w W Formate Arsenite w Arsenate w w w a a a a a a a a a a SOLUBILITY Notes to Table of Solubilities i. Potassium dichromate, W. 2. Potassium borotartrate, W. 3. Hydrogen potassium oxalate, W. 4. Hydrogen potassium carbonate, W. 5. Hydrogen potassium tartrate, W. 6. Ammonium potassium tartrate, W. 7. Sodium potassium tartrate, W. 8. Ammonium sodium phosphate, W. 9. Acid sodium borate, W. TABLE OF SOLUBILITIES 181 TABLE Ferrous Ferric Silver Lead Mercurous Meruric | Cupric Bismuth | Cadmium | Gold Platinum | Stannous | Stannic Antimonious a A a A«4 A A A a a a a a&i a42 Oxide w a A-I a w-a w a a a a Chromate W2O w W-A A-I w-a W27 WM w w w w a Sulphate a A a a a a a a a a a w-a Phosphate a a a a a a w-a a Borate a a a a a a a a a w a w a Oxalate w-a w w a w-a a w w-a w w w Fluoride A a A a a A a a Carbonate a a a a a Silicate w I W-I A-I w W-Ag3 w ^37.38 w Ww w-a43 Chloride w w i w-i a-i w w w-a w w w w-a Bromide w w i W-A A A w a w a i w w w-a Iodide a-i I a W a a w w Cyanide i i i a i i i Ferrocyanide I w i w-a i Ferricyanide w w i a A w a w-a a w Thiocyanate A a a23 A A ^29 a31 a A a36 a39 a41 a4i A^.^ Sulphide w w W w w W WM w w Nitrate w w w w w w w w w w Chlorate w-a W22 a a w-a a w a w-a a a46 Tartrate w w a a a w-a w a Citrate w w-a w-a a w-a w w w Malate w-a a a a a w-a w w a Succinate w a w-a a a w-a a w Benzoate w-a w-a w Salicylate w w w w-a w Ws, w w w w Acetate w w w w-a w w w w w w Formate a a a a a a A a Arsenite a a a a a a a a a a Arsenate io. Hydrogen sodium carbonate, W. ii. Tricalcium phosphate, A. 12. Ammonium magnesium phosphate, A. 13. Potassium aluminum sulphate, W. 14. Ammonium aluminum sulphate, W. 15. Potassium chromium sulphate, W. 16. Zinc sulphide, as a sphalerite, soluble in nitric acid with separation of sulphur; in hydrochloric acid only upon heating. 17. Manganese dioxide, easily soluble in hydrochloric acid; insoluble in nitric acid. 182 QUALITATIVE CHEMICAL ANALYSIS 18. Nickel sulphide is rather easily decomposed by nitric acid; very difficultly by hydrochloric acid. 19. Cobalt sulphide, like nickel sulphide. 20. Ammonium ferrous sulphate, W. 21. Ammonium ferric chloride, W. 22. Potassium ferric tartrate, W. 23. Silver sulphide, only soluble in nitric acid. 24. Minium is converted by hydrochloric acid into lead chloride; by nitric acid into soluble lead nitrate and brown lead peroxide which is insoluble in nitric acid. 25. Tribasic lead acetate, W. 26. Mercurius solubilis Hahnemanni, A. 27. Basic mercuric sulphate, A. 28. Mercuric amido-chloride, A. 29. Mercuric sulphide, not soluble in hydrochloric acid, nor in nitric acid, but soluble in aqua regia upon heating. 30. Ammonium cupric sulphate, W. 31. Copper sulphide is decomposed with difficulty by hydro- chloric acid, but easily by nitric acid. 32. Basic cupric acetate, partially soluble in water, and com- pletely in acids. 33. Basic bismuth chloride, A. 34. Basic bismuth nitrate, A. 35. Sodium auric chloride, W. 36. Gold sulphide is not dissolved by hydrochloric acid, nor by nitric acid, but it is dissolved by hot aqua regia. 37. Potassium chlorplatinate, W-I. 38. Ammonium chlorplatinate, W-I. 39. Platinum sulphide is not attacked by hydrochloric acid, is but slightly attacked by boiling nitric acid (if it has been precipitated hot), but is dissolved by hot aqua regia. 40. Ammonium stannic chloride, W. 41. Stannous sulphide and stannic sulphide are decomposed and dissolved by hot hydrochloric acid, and are con- verted by nitric acid into oxide, which is insoluble in an excess of nitric acid. Sublimed stannic sulphide is dissolved only by hot aqua regia. TABLE OF SOLUBILITIES 183 42. Antimonious oxide, soluble in hydrochloric acid, not in nitric acid. 43. Basic antimonious chloride, A. 44. Antimony sulphide is completely dissolved by hydro- chloric acid, especially upon heating; it is decomposed by nitric acid, but dissolved only to a slight degree. 45. Calcium antimony sulphide, W-A. 46. Potassium antimony tartrate, W. 47. Hydrogen calcium malate, W. 184 QUALITATIVE CHEMICAL ANALYSIS REAGENTS With a few exceptions, all reagents should be of the highest purity obtainable and each sample lot tested before use. The fact that the bottle bears the label C. P. is no guarantee of its purity. It is especially important that the reagent be tested for the presence of the acid or basic radical it is employed to detect; e.g., " arsenic free zinc " should be tested for arsenic by the Gutzeit or Fleitmann test before being used. Sodium carbonate, employed in relatively large amounts for fusion purposes, should be of a high degree of purity, and should be tolerably free from foreign bases and acids. Solutions Acids Cone. HCl, sp. gr. 1.2, 39% HCl by weight. Dil. HCl, 3 N, sp. gr. 1.05, 10% HCl by weight. Cone. HNO& sp. gr. 1.42, 70% HNO3 by weight. Dil. HNOZ, 3 N, sp. gr. 1.10, 10% HNO3 by weight. Cone. H^SO^, sp. gr, 1.84, 98% H2SO4. Dil. H2SO±, 3 N, sp. gr. 1.09, 13 % H2SO4. Cone. HF, 40 %. Acetic Acid, 2 N, 120 g. of glacial acetic acid in 1 liter. Tartaric Acid, 2 N, 150 g. in 1 liter. Aqua Regia, 1 part of cone. HNO3 to 3 parts cone. HCl; to be prepared only when needed. H^S gas is prepared by the action of HCl ( 1: 1) on FeS; the gas should be washed by passing it through water before using. Bases Cone. Ammonia, sp. gr. 0.90, 28 % NH3. Dil. Ammonia, sp. gr. 0.96, 10% NH3. Sodium hydroxide, NaOH, 4 N. As the material used for qualitative purposes contains about 10% of water, the amount needed for a 4N solution will be 4X4OX -1gQ-= 177.7 g- in 1 liter. Potassium hydroxide, KOH, 4 N. REAGENTS 185 The grade used for analytical purposes contains about 20% water; hence the quantity needed for a 4 N solution will be 4X56x| = 280 g. in 1 liter. Barium hydroxide, Ba(OH)2, saturated solution. Calcium hydroxide, Ca(OH)2, saturated solution. Salts Ammonium acetate, NH4C2H3O2. Add 1000 cc. of ammonium hydroxide (sp. gr. 0.90), slowly and with constant stirring, to 1250 cc. of glacial acetic acid. Ammonium carbonate, (NH4)2CO3SO4 free. Dissolve, without heating, 192 g. of the powdered salt in a mixture of 80 cc. of NH4OH (sp. gr. 0.90) and 500 cc. of water. When solution is complete, dilute to 1 liter. The strength is approximately 4 N. Ammonium chloride, NH4C1, 4 N. 214 g. in 1 liter. Ammonium molybdate solution. To a mixture of 271 cc. of cold distilled water and 144 cc. of NH4OH (sp. gr. 0.90), add 100 g. MoO3 and stir till solution is complete; slowly add this solution with constant stirring to a mixture of 489 cc. HNO3 (sp. gr. 1.42) and 1148 cc. of water. Allow the mixture to stand for 24 hours and then decant the clear liquid into a bottle. Ammonium oxalate, (NH4)2C2O4 • H2O. 35-54 g. in 1 liter. Ammonium sulphide (colorless), (NH4)2S. Saturate 3 parts of NH4OH with H2S, add 2 parts of ammonium hydroxide, and dilute with an equal volume of water. Ammonium sulphide (yellow), (NH^S^.. Digest the color- less undiluted (NH4)2S with flowers of sulphur in the proportion of 1 g. to the liter and then dilute with an equal volume of water. Ammonium sulphate, (NH4)2SO4, N. 100 g. in 1 liter. Barium chloride, BaCl2 • 2 H2O, N. 122.17 S' *n 1 liter. Bromine water, saturated solution. Calcium chloride, CaCl2, anhydrous, N. 55.6 g. in 1 liter. Calcium sulphate, CaSO4 • 2 H2O, saturated solution. Chlorine water, saturated solution. Cobalt nitrate, Co(NO3)2 • 6 H2O, for confirmatory tests for Al and Zn. 0.5 g. in 1 liter. 186 QUALITATIVE CHEMICAL ANALYSIS Ferric alum, Fe2(SO4)3 • (NH^SC^ • 24 H2O, saturated solu- tion. Ferric chloride, FeCl3 • 6 H2O, * 2 N. 180 g. in 1 liter. Ferrous sulphate, FeSO4 • 7 H2O. To be prepared in small amounts as needed. Hydrochlorplatinic acid, H2PtCl6 • 6 H2O. 10% solution. Hydrogen dioxide, 3 %. Lead acetate, Pb(C2H3O2)2 • 3 H2O,f N. 189.5 g. in 1 liter. Magnesia mixture. Dissolve no g. of MgCl2 • 6 H2O and 280 g. of NH4C1 in a liter of distilled water; when solution is complete, add 261 cc. of ammonium hydroxide (sp. gr. 0.90), then add enough water to make the volume 2 liters. Mercuric chloride, HgCl2. Saturated solution. Potassium acetate, KC2H3O2. Saturated solution. Potassium chromate, K2CrO4, N. 97.3 g. in 1 liter. Potassium cyanide, KCN, N. 65.2 g. in 1 liter. Potassium dichromate, K2Cr2O7, N. 73.8 g. in 1 liter. Potassium ferrocyanide, K4Fe(CN)6, N. 105.7 g. in 1 liter. N Potassium iodide, KI, -. 83.1 g. in 1 liter. 2 Potassium nitrite, KNO2. 500 g. in 1 liter. Potassium permanganate, KMnO4, N. 79.1 g. in 1 liter. Potassium thiocyanate, KCNS, N. 97.2 g. in 1 liter. N Silver nitrate, AgNO3, -. 42.5 g. in 1 liter. 4 Silver sulphate, Ag2SO4. Saturated solution. Sodium acetate, NaC2H3O2, 4 N. 328 g. in 1 liter. Sodium carbonate, Na2CO3 (dry). Saturated solution. Sodium cobaltic nitrite, Na3Co (NO2)6. Dissolve 100 g. NaNO2 in 300 cc. distilled water, slightly acidify with acetic acid, and then add 10 g. of Co(NO3)2 • 6 H2O. Allow the solution to stand for 24 hours and filter if necessary. As the solution does not keep very well, only small amounts should be prepared at a time. Sodium nitroprusside, Na^eNC) (CN)5 • 2 H2O. 10 % solution. Sodium phosphate, Na^PC^ • 12 H2O, N. 119 g. in 1 liter. * Should contain a little free HC1. t The solution should contain some free acetic acid. REAGENTS 187 Sodium stannite, prepared as needed by adding to a little SnCl2 solution sufficient NaOH solution to redissolve the precipitate which first forms. Sodium thiosulphate, Na2S2O3- 5 H2O, N. 124 g. in 1 liter. N Stannic chloride, SnCl4, -. 32.7 g. in 1 liter. 2 N Stannous chloride* SnCl2 • 2 H2O, -. 56.5 g. in 1 liter. Stannous chloride (for Bettendorff Test). Dissolve 113 g. of SnCl2 • 2 H2O in 75 cc. of cone. HC1, and add a few pieces of C. P. tin foil and keep in glass stoppered bottle. Starch paste. Prepared as needed by mixing about 1 g. of powdered starch with a little cold water to form a thin paste and then adding it to 200 cc. of boiling water; boil for a minute, cool, and use. The solution does not keep, owing to the growth of molds. It may be kept for some time, however, if a pre- servative such as CS2 is added. Solvents Alcohol, amyl, (C5HnOH), C. P. Used in small amounts in the test for acetate. Alcohol, ethyl, (C2H6OH). 95%, sp. gr. 0.815. Benzol, C6H6, useful for dissolving sulphur. Chloroform, CHC13, used for dissolving iodine. Carbon disulphide, CS2, used for dissolving iodine. Ether, ethyl, (C2H5)2O, solvent for fats and oils. Dry Reagents Aluminum turnings, pure. Ammonium chloride, NH4C1, C. P. Ammonium nitrate, NH4NO3, C. P. Borax, Na2B4O7 • io H2O, C. P. and powdered. Calcium carbonate, CaCO3, alkali free. Calcium fluoride, CaF2, SiO2 free. Copper, strips. * The solution should be strongly acid with HC1; the addition of a little C. P. tin foil prevents the oxidation of the reagent. 188 QUALITATIVE CHEMICAL ANALYSIS Ferrous sulphate, FeSO4 • 7 H2O, C. P. Fusion mixture (Na2CO3 + K2CO3, dry and C. P.). Iron filings. Iron nails. Lead dioxide, PbO2, free from Mn. Litmus paper, blue and red; to be kept in stoppered bottles. Manganese dioxide, MnO2, C. P. and powdered. Microcosmic salt, NaNH4HPO4 • 4 H2O. Paraffine, m.-p. 1240. Potassium acid sulphate, KHSO4, fused, C. P. in small lumps. Potassium chlorate, KC1O3, C. P. powdered. Potassium chloride, KC1, C. P. Potassium cyanide, KCN, pure. Potassium dichromate, K2Cr2O7, C. P. powdered. Potassium ferricyanide, K3Fe(CN)6, C. P. Potassium nitrate, KNO3, C. P. fine crystals. Sand, sea. Silica, SiO2, purified. Sodium acetate, NaC2H3O2, C. P. Sodium acid carbonate, NaHCO3, C. P. Sodium carbonate, NagCO^ anhydrous, C. P. powdered. Sodium dioxide, Na2O2, C. P. Sodium sulphite, Na^SOg • 7 H2O, pure. Starch, potato. Sulphur, flowers. Tin foil, C. P. Turmeric paper; to be kept in glass-stoppered bottles. Zinc, granulated, C. P. Zinc, granulated, arsenic-free. APPARATUS 189 LIST OF APPARATUS 2 nests of beakers, Griffin, 1-4. 1 graduated cylinder, 10 cc. 1 graduated cylinder, 50 cc. 1 wash bottle with fittings, 750 cc. 4 funnels, 6.5 cm. 1 funnel, 10 cm. 2 pieces cobalt glass. 1 doz. test tubes (15 cm.) 3 ft. glass rod. 3 ft. glass tubing. 2 specimen bottles, 50 cc. 2 watch glasses, 10 cm. 2 watch glasses, 5 cm. 1 watch glass, 12.5 cm. 1 florence flask, f. b. 50 cc. 2 evaporating dishes, 10 cm. 2 evaporating dishes, 6.5 cm. 1 porcelain crucible. 1 horn spatula. 1 rubber stopper, one hole, No. 1. 1 funnel cleaner. 1 sponge. I test tube cleaner. | box gummed labels, #217. 1 doz. sheets filter paper, 18.5 cm., S. & S. 595. 1 pkg. filter paper, 12.5 cm., S. &S. 595. 1 doz. fluted filters, 12.5 cm., S. & S. 588. 1 test tube rack. 1 test tube holder. 1 filtering stand. 1 box matches, safety. 1 pair forceps (small). 1 pipestem triangle. 1 retort stand (2 rings). 2 bunsen burners, with hose. 1 blowpipe. 1 stick charcoal. 2 pieces wire gauze, 10 cm. square. 1 triangular file. 1 platinum wire. 1 platinum foil. 1 towel. 190 QUALITATIVE CHEMICAL ANALYSIS PREPARATION OF UNKNOWNS In the making up of unknowns, stock solutions of the concen- tration i cc. = 100 mg. of metal are prepared. The quantity of salt necessary to dissolve in a liter to yield this strength is given in column 5 of the table below. By means of burettes or pipettes definite quantities of these standard solutions are measured out into student "unknown" bottles, homeopathic vials of 50 cc. ca- pacity. For the analysis the student uses 25 cc. of his solution, the other half being reserved in case the analysis is to be re- peated. The amounts of standard solutions pipetted out should be such as to yield a suitable concentration when the volume is diluted to 50 cc., i.e., when the unknown bottle is filled. An example will make this clear. Pipette out into unknown bottle 1 cc. NaCl solution, 2 cc. Ca(NO3)2, and 1 cc. of NH4NO3, and then fill the bottle with distilled water. Since the student uses only 25 cc. of this solution, this quantity will contain 50 mg. Na, 100 mg. Ca, and 50 mg. NH4. Qualitative unknowns may be prepared of such a strength that the total weight of metal in 25 cc. never exceeds 1.5 grams, though it should usually be kept within 1 gram. The minimum will depend upon the scheme of analysis employed. It may be exceedingly small if the most sensitive tests are used, e.g., the spectroscopic tests for the alkali and alkaline earth metals, the KCNS test for Fe, and the Marsh and Gutzeit tests for As and Sb. But if it is desired that the student report roughly the relative proportions of the ingredients present, precipitation methods will be largely used, which, by the size of the precipitates they yield, give indications of the approximate quantities of the metals present. In the latter case, the minimum quantity of metal present in 50 cc. will have to be much larger than it is in the first case. PREPARATION" OF UNKNOWNS 191 TABLE EMPLOYED IN THE PREPARATION OF STANDARD STOCK SOLUTIONS* Group Substance Formulae Weight Solubility of Salt in ioo Pts. of Cold Water Per cent. Metal Quantity of Salt to be Dis- solved in 1 Liter to give Strength i cc. = IOO MG. OF Metal I AgNO3 . . . . 170 V. S.1 63-5 I57 HgNO3H2O . . 280 sol. in pres, of hno3 7i-5 140 Pb(NO3)2 . . . 331 48 62.5 160 Pb(C2H3O2)2-3H2O 379 46 54-6 183 II Hg(NO3)24(H2O) 333 sol. in pres, of HNOS 60 167 HgCl2 271 7-4 74 I352 Bi(NO3)3-5 H2O . 484 sol. in pres, of hno3 43 233 Cu(NO3)2-6H2O . 295 V. S. 21-5 465 CuC12-2H2O. . . 170 120 37 270 CuSO4-5H2O . . 249 40 25 400 Cd(NO3)2-4 H2O . 308 V. s. 36 278 CdCl2-2H2O . . 219 140 5i 196 3CdSO4-8H2O. . 769 V. s. 43-5 230 As2O3 198 4 75-5 ( )3 Na2HAsO3 . . . 170 V. s. 44 227 Na2HAsO4 i2 H2O 402 28 18.7 ( y As2O5 230 150 65 153 SbCl3 226 sol. in pres, of HC1 53 188 SnCl2-2 H2O . . 225 V. S. 53 189 SnCl4-5 H2O . . 350 V. S. 34 294 SnCl4 260 V. S. 46 218 1 Very soluble. 2 This amount readily dissolves in i liter of water containing 50 g. of NaCl. 8 332- in 1 liter HC1 (1:1) gives strength 1 cc. = 2£ mgs. As. 4 267 g. in 1 liter will give strength 1 cc.=so mgs. As. * Taken from an article published in School Set. and Math., Vol. X., No. 6, by one of us (L.J.C.). 192 QUALITATIVE CHEMICAL ANALYSIS TABLE EMPLOYED IN THE PREPARATION OF STANDARD STOCK SOLUTIONS - Continued Group Substance Formulae Weight Solubility of Salt in ioo Pts. of Cold Water Per cent. Metal Quantity of Salt to be Dis SOLVED IN 1 Liter to give Strength i cc. = IOO MG. of Metal III A12(SO4)3-i8H2O . 666 107 8.1 ( r A1C1/.6 H2O . . . 242 74 11.1 ( )2 Al(NO3)3-8 H2O . 261 v. s. 10.3 970 Cr2(SO4)3i8H2O . 716 V. s. 14.6 690 K2Cr2(SO4)4-24H2O 1000 20 5.2 ( )8 Cr(NO3)3-9 H2O . 400 V. s. <3 770 CrCl3-6 H2O . . . 266.5 V. s. 19.6 570 FeSO4-7 H2O . . 278 60 20 500 Fe(NO3)3-9 H2O . 4°4 V. s. 14 715 FeCl3-6H2O. . . 270 V. s. 20.7 482 Ni(NO3)2-6H2O . 291 50 20 500 NiCl2-6H2O . . . 238 V. s. 25 400 NiSO4-7H2O . . 280 106 21 475 Co(NO3)2-6H2O . 291 V. s. 20 500 CoC12-6H2O. . . 238 V. s. 24-5 407 CoSO4-7H2O . . 281 50 21 475 MnSO4-4H2O . . 223 123 25 400 MnCl2-4 H2O . . 198 150 28 360 Mn(NO3)2-6H2O . 287 V. s. 19 527 ZnSO4-7H2O . . 288 135 22.5 445 Zn(NO3)26H2O . 298 V. s. 22 455 ZnCl2 136 V. s. 48 208 1 620g. in i liter will give strength i cc.=5o mg. Al. 2 450 g. in 1 liter will give strength 1 cc.=$o mg. Al. 8 192 g. in 1 liter will give strength 1 cc. = io mg. of Cr. PREPARATION' OF UNKNOWNS 193 TABLE EMPLOYED IN THE PREPARATION OF STANDARD STOCK SOLUTIONS - Continued Group Substance Formulae Weight Solubility of Salt in ioo Pts. of Cold Water Per cent. Metal Quantity of Salt to be Dis- solved in 1 Liter to give Strength i cc.= IOO MG. OF Metal IV BaCl2-2 H2O . . . 244 41 56 179 Ba(C2H3O2)2H2O 273 63 5° 200 Sr(NO3)2-4H2O . 284 40 31 324 Sr(NO3)2 . . . 212 39 41.3 242 SrCl2-6H2O . . . 266 106 33 304 CaCl2 III v. s. 36 278 Ca(NO3)2-4H2O . 236 V. s. 17 590 VMgSO4-7H2O . . 246 77 9-7 ( V Mg(NO3)2-6H2O . 256.5 200 9-4 1060 MgCl2-6H2O . . 203-5 365 11.9 837 NaCl 58 35 40 250 Na2HPO4i2 H2O . 358 9-3 13 ( )2 NaNO3 .... 85 80 27 371 KC1 75 32 52 192 KHSO4 .... 136 v. s. 28.5 350 KNO3 IOI 3i 39 257 NH4C1 .... 53 33 34 294 (NH4)2SO4 . . . 132 76 27-5 365 NH4NO3 .... 80 200 22.5 445 (NH4)2HPO4 . . 132 36.5 27-5 365 LiCl 42 80 16.7 600 LiNO3 .... 69 48 IO ( )8 1 5I5 g- in 1 liter will give strength i cc.=5o mg. Mg. 2 77 g- in i liter will give strength i cc. = iomg. Na. 3 500 g. in 1 liter will give strength 1 cc. =50 mg. Li. INDEX Acetate, basic, of Al, 56. Cr, 62. Fe (-ic), 68. separation, 76. Acetates, action of heat on, 137. detection, 163. reactions, 137. solubilities, 137. Acid analysis, general examination, 146. preliminary examination, 141. Acids, detection in minerals, 178. division into groups, 104. first group, analytical, 149. descriptive, 105. preliminary examination for, 141. second group, analytical, 156. descriptive, 124. special tests for, 149. third group, analytical, 161. descriptive, 137, Agate mortar, 119. Alcohol, amyl, 137, 163, 187. ethyl, 137, 187. reduction of chromates by means of, 45> I23- Alkalies, characteristics of, 93. detection in Scheme V., 101. detection in silicates, 177. salts of, reactions, 96. solubilities, 94. Alkaline earths, characteristics of, 84. detection in Scheme IV., 91. salts of, reactions, 84. solubilities, 84. Alloys, analysis of, 171. Aluminate of sodium, decomposition with acid, 55. decomposition with NH4CI, 55. Aluminum, detection in scheme of analy- sis, 81. salts, characteristics, 54. reactions, 54. solubilities, 54. Ammonia, complex salts of, 27, 29, 69, 7 b 75- in drinking water, 98. reagent, 184. Ammonium arseno-molybdate, 36. carbonate, reagent, 185. chlorplatinate, 97. -magnesium arsenate, 36. -magnesium phosphate, 95. phospho-molybdate, 112. salts, detection, 101. reactions, 97. solubilities, 97. sulphide, colorless, 185. yellow, 185. Amyl alcohol, 137, 163, 187. Analysis, acid, general, 147. acid, of minerals, 178. for metals of all groups, 103. of alloys, 171. of Group III. (metals) in the presence of phosphates, etc., 167. of insoluble substances, 172. qualitative, I. quantitative, 1. Antimonic compounds, reactions, 39. Antimonious compounds, reactions, 37. Antimony, detection in scheme of analy- sis, 53- by Gutzeit Test, 40. by Marsh Test, 40. by Reinsch Test, 40. Apparatus, list of, 189. Aqua Regia, action of H2S on, 45. as an oxidizing agent, 9, 24. as a solvent, 9, 24, 169. preparation of, 9, 24, 184. removal of excess, 45, 170. Arsenates, detection, 36, 53, 152. reactions of, 34. reduction with KI, 36. reduction with SO2, 13, 35, 36. special tests for, 36, 152. 195 196 INDEX Arsenic, acids of, 30. oxides, 30. sensitiveness of special tests for, 34. Arsenic tests, special: Bettendorff, 34. Fleitmann, 33. Gutzeit, 33. Marsh, 32. Reinsch, 34. tests for, in reagents, 33. in wall paper, 33. Arsenites, detection, 32, 53, 152. oxidation of, 34. reactions, 30. solubilities, 30. special tests for, 32, 152. Arsine, 32. Barium hydroxide, reagent, 185. salts, detection in scheme of analysis, 91. reaction of, 84. solubilities, 84. sulphate, decomposition of, 86. Bases, list of, 184. Basic acetate of aluminum, 56. chromium, 62. iron (-ic), 68. separation, 76. Bead tests with borax, 145. tests with NaPOa, 119, 155, 174, 175. Bettendorff test for As, 34. Bismuth salts, basic, 26. reactions, 26. solubilities, 26. test in scheme of analysis, 50. Blowpipe tests on charcoal, 144, 145. Borates, detection in salts, 154. detection in silicates, 117. reactions, 116. solubilities, xi6. Borax bead tests, 145. Boric acid, detection, X54. Bromides, detection, 156. insoluble, 127, 174. reactions, 126. solubilities, 127. Cadmium salts, detection in scheme of analysis, 50. reactions, 29. solubilities, 29. Calcium salts, detection in scheme of analysis, 91. reactions, 88. sulphate, analysis of, 86. Carbon, detection of, 172. disulphide, 187. removal of, 175, 176. Carbonaceous residue, 122, 142, 164. Carbonates, detection, 149. reactions, 107. solubilities, X07. Carbonic acid, detection, 107, 149. Chlorates, as oxidizing agents, 10, 74. behavior on ignition, 140. detection, 163. reaction, 140. solubilities, 140. Chlorides, detection in presence of bro- mides, 157. detection in presence of bromides and iodides, 158. insoluble, treatment of, 125, 174. reactions, 124. solubilities, 124. Chlorine water, reagent, 185. Chlorplatinate, ammonium, 97. potassium, 96. Chromates, as oxidizing agents, IO. conversion to dichromates, 59. detection, 151. reactions, 122. reduction of, 60. solubilities, 122. Chrome-iron ore, treatment of, 174. Chromium, insoluble compounds, treat- ment of, 124. oxidation to chromates, 58. reactions, 61. salts, detection in scheme of analysis, 81. solubilities, 60. Chromyl chloride test for chlorides, 158. Classification of metals, 14. Closed tube tests, 142. Cobalt nitrate test for Al, 57. nitrate test for Zn, 76. salts, detection in scheme of analysis, 80. oxidation to cobaltic compounds, 71, 72. reactions, 70. INDEX 197 Colloidal AS2S3, 31. Colloidal state, 47. Copper salts, detection in scheme of analysis, 50. reactions, 27. solubilities, 27. Cyanides, action of heat on, 129. as reducing agents, 14. detection, 160. reactions of simple, 129. solubilities, 129. Decantation, 2. Decomposition of alkaline earth sulphates, 86, 173. antimony tetroxide, 174. cassiterite, 174. chromium compounds (insol.), 174. fluorspar, 174. ignited oxides, 174. Prussian blue, 132, 173. silicates, 119, 174. silver halides, 174. Dehydration of silicic acid, 118. Dichromates, conversion to chromates, 60. reduction with H2S, 45, 60. reduction with alcohol, 45, 60. Directions for laboratory work, 3. Equations, complicated, writing of, 7. defined, 4. factors of, 5. methods of writing, 4. products of, 5. Ethyl alcohol, 45, 123, 137, 187. Examination, general, for acids, 146. preliminary (general), 141. Ferric salts, characteristics, 65. detection in scheme of analysis, 80. reactions, 65. reduction to ferrous, 68. Ferricyanides, action on heating, 133. detection, 159. reactions, 133. solubilities, 133. Ferrocyanides, detection, 159. reactions, 131. solubilities, 132. Ferrous salts, characteristics, 62. detection in scheme of analysis, 80. Ferrous salts, oxidation to ferric, 65. reactions, 63. Filtrate defined, 2. Flame tests, 145. Fleitmann test for As, 33. Fluorides, characteristics, 113. detection in the absence of SiO2, 153. detection in the presence of SiC>2, 153. reactions, 113. solubilities, 113. Fluosilicates, detection, 150. reactions of, 106. Glass, etching of, 114. Group reagents, 14. Group I., acids, analytical, 149. descriptive, 104. Group II., acids, analytical, 156. descriptive, 124. Group III., acids, analytical, 161. descriptive, 137. Group I., metals, analytical, 22. descriptive, 17. Group II., A, metals, analytical, 50. descriptive, 24. Group II., B, metals, analytical, 53. descriptive, 30. Group III., metals, analytical, 78. descriptive, 54. Group IV., metals, analytical, 91. descriptive, 84. Group V., metals, analytical, 101. descriptive, 93. Groups, division of acids into, 104. metals into, 14,15, 16. Gutzeit test for As, 33. Sb, 40. Hart's method, 158. Heating in closed tube, 142. on charcoal, 144. with cone. H2SO4, 143. Hydrogen dioxide, oxidation with, 11, nascent, as a reducing agent, 12. Hydrogen sulphide as Group II. reagent, 45- reducing agent, 44. detection, 161. Hydrolysis, salts of antimony, 38. bismuth, 26. iron (-ic), 65. 198 INDEX Insoluble substances, treatment of, 172. scheme of analysis for, 176. Introduction, I. Iodides, characteristics, of, 127. detection, 156. insoluble, treatment of, 174. reactions, 127. removal of, in preparation of solution, 169. solubilities, 127. Iron, see ferrous and ferric. Laboratory work, directions for, 3. Law of the Conservation of Elements, 5. Weight, 5. Lead acetate, reagent, 186. peroxide as an oxidizing agent, 12. salts, detection in small amounts, 20. scheme of analysis, 22, 50. reactions, 19. removal of, in the analysis of insoluble substances, 176. Magnesia mixture, reagent, 186. Magnesium salts, detection in scheme of analysis, 101. reactions, 94. removal of, 102. solubilities, 94. Manganate of sodium, 74. Manganese salts, detection of small amounts, 74. in scheme of analysis, 80. oxidation to permanganic acid, 74. reactions, 73. Marsh test for As, 32. Sb, 40. Mass action law, 63. Medicinal preparations, arsenic in, 33. Mercuric salts, detection in scheme of analysis, 50. reactions, 24. solubilities, 24. Mercurous salts, detection in scheme of analysis, 22. reactions, 18. Metals, division of, into groups, 14. Group I., reactions of, 17. scheme of analysis for, 22. Group II., A, reactions of, 24. scheme of analysis for, 50. Metals, Group II., B, reactions of, 30. scheme of analysis for, 53. Group III., reactions of, 54. scheme of analysis for, 79. Group IV., reactions of, 84. scheme of analysis for, 91. Metals of Group V., reactions of, 93. scheme of analysis, 101. Metals and Alloys, analysis of, 171. Metaphosphate bead test for SiO2, 119, I55» 165, 174. Metastannic compounds, reactions of, 44. conversion to stannic compounds, 44. Metathesis, 6. Minerals, analysis for acids, 178. powdering of, 119. Nessler's reagent, 98. Nickel salts, characteristics of, 68. detection in scheme of analysis, 80. reaction, 68. Nitrates, behavior on ignition, 138. detection, 161. distinction from chlorates, 141. reactions, 138. solubilities, 138. Nitric acid, as an oxidizing agent, 9. action on ferrous salts, 9. action on HoS, 45. action on iodides, 169. detection of free, 140. Nitrites, behavior on ignition, 136. detection, 162. reactions, 136. solubilities, 136. Organic matter, detection, 142, 165. interference of, 165. removal of, 165. Orthophosphoric acid, 110. Oxalates, behavior on ignition, 115. detection, 153, 165. interference of, 164. reactions, 115. removal of, 165. solubilities, 115. Oxidation defined, 7. Oxidizing agents: aqua regia, 9. halogens, 9. hydrogen dioxide, 11. INDEX 199 Oxidizing agents: lead peroxide, 12. list of, 8. nitric acid, 9. potassium chlorate, 10. potassium dichromate, 10. potassium permanganate, II. removal of, before precipitating with H2S, 45. sodium dioxide, 11. Permanganates as oxidizing agents, 11. Phosphate, meta-, test for SiO2, U9> I55» 165, 174. separation, 165. Phosphates, characteristics of, 110. detection, 152, 165. interference of, 164. reactions, no. removal of, 167. scheme of analysis in the presence of, 167. solubilities, in. Phosphoric acids, no. Platinum wire, cleaning of, 93. Potassium chlorate, 10. Potassium dichromate, IO. Potassium, permanganate, 11. salts, detection in scheme of analysis, 101. reactions, 96. solubilities, 96. Powdering minerals, 119. Precipitates, washing of, 2. Precipitation, defined, 2. completeness of, 23. with H2S, 46. Preliminary tests, closed tube, 142. with cone. H2SO4, 143. with dil. HC1, 143. Preparation of solution for acids, 146. of solids, 168. Qualitative analysis, I. Quantitative analysis, I. Reactions of metals of Group I., 17. Group II., A, 24. B, 30. Group III., 54. Group IV., 84. Group V., 93. Reactions, reversible, 26, 38, 44, 56, 65, 67, 68, 83, 86, 137. types of, 5. Reagent defined, I. Reagents, detection of arsenic in, 33. dry, list of, 187. group, 14. list of, 184. use and care of, 3. Reducing agents, list of, 9. Reduction defined, 7. with carbon, 13. hydrogen sulphide, 13. nascent hydrogen, 12. potassium cyanide, 14. stannous chloride, 12. sulphurous acid, 13, 35. Reinsch test for As, 34. Sb, 40. Removal of bromides, 169. iodides, 169. organic matter, 165. oxalates, 165. phosphates, 167. silicates, 118, 165. Residue defined, 2. Scheme I., metals of the silver group, 22. II., A., metals of the copper group, 50. II., B., metals of the tin group, 53. III., metals of the iron group, 79. IV., the alkaline earths group, 91. V., the alkalies, 101. Separation of 2A from 2B, 47, 48. Silica, tests for, 119, 155, 165, 175. Silicates, alkalies in, 177. decomposition by acids, 119. by fusion with alkali carbonate, 120. by HF, 121. detection, 155, 165. reactions, 118. removal of, 118, 165. Silver halides, removal of, 176. salts, characteristics, 17. detection in scheme of analysis, 22. reactions, 17. treatment of, 174. Smith, J. L., method for alkalies, 177. INDEX 200 Sodium aluminate, decomposition by acids, 55. decomposition by NH4CI, 55. dioxide as an oxidizing agent, 11. manganate, 74. peroxide, 12. phosphate bead test, 119, 155, 174, 175- plumbite test for sulphides, 135. salts, detection in scheme of analysis, 101. reactions, 99. solubilities, 99. stannite, reagent, 186. zincate, 74. Solids, preliminary testing of, 142. Solubilities, table of, 179. Solution of insoluble substances, 172. preparation of, for analysis, 168. preparation of, for acid analysis, 146. Solvents, list of, 187. Stannic, meta, compounds, 44. salts, characteristics, 42. detection in scheme of analysis, 53- . reactions, 42. solubilities, 42. Stannite of sodium, reagent, 186. Stannous salts, characteristics of, 41. detection in scheme of analysis, 53. reactions, 41. reduction by means of, 12. solubilities, 41. Starch iodide reaction, 129. paste, preparation of, 187. Stibine, 40. Strontium salts, detection in scheme of analysis, 91. reactions, 87. sulphate, decomposition of, 86. Substances, insoluble, 172. Sulphates, characteristics of, 105. decomposition of alkaline earth, 86, 173- detection, 150. reactions, 105. solubilities, 105. Sulphides, behavior on ignition, 134. detection, 135, 161. insoluble, X35. Sulphides, reactions, 134. sodium plumbite test for, 135. solubilities, 134. Sulphites, characteristics of, 108. detection, 150. reactions, 108. solubilities, 108. Sulphur, removal of, 175. Sulphuric acid, detection of free, 105. preliminary testing with, 143. Synthesis, 5. Systematic treatment of insoluble sub- stances, 175. Table for the preparation of unknowns, 190. of solubilities, 179. Tartrates, behavior on ignition, 122. characteristics of, 121. detection, 155. reactions, 121. solubilities, 121. Thenard's green, 76. Thio-antimonate, 40. -antimonite, 38. -arsenate, 32. -arsenite, 32.. -cyanates, detection, 159. reactions, X33. solubilities, 134. -stannate, 41. -sulphate, detection, 150. reactions, 109. Tin, see stannic, stannous, and meta- stannic, Turmeric paper, 188. Unknowns, preparation of, 190. Vortmann's method, 158. Wall paper, test for arsenic in, 33. Washing precipitates, 2. Zinc salts, characteristics of, 74. detection of in scheme of analysis, 80. reactions, 74. solubilities, 74. Zinc-platinum couple, action of on SbCU, 39- r| 'HE following pages contain advertisements of books by the same author or on kindred subjects. STANDARD BOOKS ON CHEMISTRY PUBLISHED BY THE MACMILLAN COMPANY ABEGG and HERZ. 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