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 Mats. M. g. 
 
 EitrrarjJ 
 
 GIFT OF 
 
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Cornell university Library 
 QD 83.B3 
 A course in Qualitative c^^^^^^^^^ 
 
Cornell University 
 Library 
 
 The original of tiiis book is in 
 tine Cornell University Library. 
 
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A COURSE IN 
 QUALITATIVE CHEMICAL ANALYSIS 
 
THE MACMILLAN COMPANY 
 
 NXW YORK • BOSTON • CHICAGO - DALXA5 
 ATLANTA • SAN FRANCISCO 
 
 MACMILLAN & CO., Limitkd 
 
 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. 
 
 HEAD PKOFESSOR IN THE DEPARTMENT OF CHEMISTRY 
 COLLEGE OF THE CITY OF NEW YORK 
 
 AND 
 
 LOUIS J. CURTMAN, Ph.D. 
 
 ASSISTANT PROFESSOR IN THE DEPARTMENT OF CHEMISTRY 
 COLLEGE OF THE CITY OF NEW YORK 
 
 REVISED 
 
 NeiH gorfe 
 
 THE MACMILLAN COMPANY 
 
 1918 
 
 All rights reserved 
 
Copyright, 1910 and 1916, 
 By the MACMILLAN COMPANY. 
 
 Set up and electrotyped. Published December, 1910. Reprinted 
 October, 1911; July, igia; January, 1913; February, Octoberj 
 J914; January, 1915; January, 1916. 
 
 Revised edition, September, October, igi6; March, September, 1917; 
 January, 1918. 
 
 Narfnootr ^rees 
 
 J. S. Gushing Oo. — Berwick & Smith Co. 
 
 Norwood, Mass., U.S.A. 
 
PREFACE TO THE SECOND EDITION 
 
 The reception of this book by teachers of Qualitative Analy- 
 sis, calling for a number of reprintings of the first edition, has 
 been a source of gratification to the Authors. In the revision, 
 some changes and additions, as indicated by the use of the 
 book, have been made. The original plan of the book, how- 
 ever, has been maintained. Some valuable suggestions from 
 others have been utilized. The theory of Electrolytic Disso- 
 ciation and the Law of Mass Action have been outlined. New 
 methods, some proposed elsewhere and some developed by 
 original work in our qualitative laboratories, have been incor- 
 porated. All of the new methods given have been thoroughly 
 
 tried out. 
 
 THE AUTHORS. 
 
 New York City, 
 July, 1916. 
 
PREFACE TO THE FIRST EDITION 
 
 Experience has shown the authors that a quantitative dis- 
 crimination in Qualitative Analysis for students is rarely exer- 
 cised, although it is generally conceded that qualitative analysis 
 should serve not only as a means for determining the compo- 
 nents in an " unknown," but also, though roughly, the propor- 
 tions 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 famiUarizing 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 
 
vm PREFACE 
 
 analysis of unknowns. In consequence of this training, the 
 student is able to report not only the qualitative composition 
 of his unknowns, but also approximately their quantitative com- 
 position. The value of such training cannot be overestimated. 
 Our students rarely find any difiHculty in differentiating between 
 a trace and a signiiicant 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 con- 
 taining solutions of known strength of the salts of the various 
 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 
 
 IX 
 
 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. 
 
 1. 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 
 QuaUtative 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 fol- 
 lows. This is in the nature of a r^sum6 of the chief reactions, 
 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 
 
X 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 
 separations 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 oppor- 
 tunity 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 students. 
 
 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 xi 
 
 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 
 
 Prefaces v, vii 
 
 Introduction 
 
 (a) Analytical Chemistry Defined i 
 
 (^) Terms i 
 
 (c) General Directions for Laboratory Work .... 3 
 
 {d) Theory of Electrolytic Dissociation 4 
 
 (e) Law of Mass Action 12 
 
 (/) Equations 20 
 
 (g) Oxidation and Reduction 23 
 
 (k) Classification 29 
 
 PART I 
 
 The Metals. Descriptive and Analytical 33 
 
 PART II 
 
 (a) The Acids. Descriptive and Analytical 123 
 
 (d) Preliminary Examination 160 
 
 PART III 
 Complete Analysis 
 
 (a) Scheme for Group III in the Presence of Organic Matter, Phos- 
 phates, etc. . 185 
 
 (6) Preparation of the Solution 189 
 
 (c) Alloys 192 
 
 (<f) Treatment of Insoluble Substances 193 
 
 (if) Acid Analysis of Minerals and Metallurgical Products . .199 
 
 Appendix i. Table of Solubilities 201 
 
 2. Reagents 205 
 
 3. List of Apparatus 211 
 
 4. Preparation of Unknowns 212 
 
 Index 217 
 
 xiii 
 
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 AgNOg 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 Uquid which 
 passes through is called the filtrate. The equation for the 
 reaction, omitting consideration of the water, is — 
 
 NaCl + AgNOa = AgCl -1- NaNOg. 
 
 Assuming that an excess of silver nitrate has been used, the pre- 
 cipitated and filtered AgCl will be wet with a solution containing 
 NaNOj and AgNOg. 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 AgNOg 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 drbp 
 by drop and should never be used in great excess, unless otller- 
 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. 
 
 The Theory of Electrolytic Dissociation* 
 
 When a substance such as sugar is added to water and the 
 latter stirred, there results a homogeneous mixture known as a 
 solution. The dissolved substance, the sugar, is called the 
 solute ; the water in which the sugar dissolves, the solvent. In 
 a similar manner we may prepare an aqueous solution of salt. 
 An experimental study of the properties of aqueous solutions 
 discloses the fact that some conduct the electric current, while 
 others do not; e.g., a solution of sugar does not permit the pas- 
 sage of the current, while on the other hand a solution of salt is 
 found to be a gobcj conductor of electricity. It is natural to 
 suppose that the solute which is different in the two cases cited 
 is responsible for the difference in behavior. Experiment con- 
 firms this supposition. We may therefore divide substances 
 soluble in water into two classes, viz. {a) those which conduct 
 the electric current and (5) those which do not. The former 
 are called electrolytes, the latter non-electrolytes. Acids, bases, 
 and salts are electrolytes. Sugar, glycerol, urea, ethyl alcohol, 
 are non-electrolytes. 
 
 * The theoretical matter dealt with in this section is generally considered more 
 or less in preliminary courses in General Chemistry. A brief summary is here given 
 partly to refresh the student's knowledge, but chiefly to emphasize certain principles 
 which have many applications in the field of qualitative analysis. 
 
INTRODUCTION 5 
 
 A study of the freezing and boiling points of aqueous solu- 
 tions leads to a similar classification. Thus if 60 grams (the 
 molecular weight) of urea are dissolved in 1000 grams of water, 
 the resulting solution, which is said to contain one mole of urea, 
 is found to freeze at —1.86° C The same result is obtained 
 when a solution containing one mole of any other non-electro' 
 lyte is cooled to its freezing point. In general, it has been 
 found that the freezing point of a dilute solution of a non- 
 electrolyte is solely dependent upon the number of moles in 
 solution, provided that in freezing only the pure solvent sepa- 
 rates out. Thus the freezing point of a solution of 50 grams of 
 methyl alcohol in 1000 grams of water is found to be —2.90°. 
 This is what we should expect, since the number of moles in 
 solution is ff, or 1.56. And as the freezing point is lowered 
 1.86° for each mole of a non-electrolyte, the total lowering 
 should be 1,86 x 1.56= 2.90. 
 
 On the other hand, if we dissolve ^5- of a mole (5.85 grams) 
 
 of sodium chloride in 1000 grams of water and determine its 
 
 freezing point, we should find it to be — 0.350°. If sodium 
 
 chloride were a non-electrolyte, the freezing point should be 
 
 — o. 1 86. On comparing these results we see that the lowering 
 
 • -350 
 
 is — ^, or 1.88, times greater than that which we should obtain if 
 
 salt were a non-electrolyte. A comparative study of the boiling 
 points of molar solutions of electrolytes and non-electrolytes 
 reveals a similar difference, viz. that solutions of electrolytes 
 boil at a higher temperature than solutions of non-electrolytes 
 of the same molar concentration. On the basis of an experi- 
 mental study of three of the properties of solutions, viz. con- 
 ductivity, boiling point, and freezing point, we are thus led to a 
 division of soluble substances into two classes, electrolytes and 
 non-electrolytes. 
 
 Moreover, a study of the reactions of electrolytes in solution 
 shows that they behave in solution as though their constituent 
 parts were independent of each other ; thus solutions of KCl, 
 NaCl, BaClj, and FeClg each give the same precipitate with a. 
 solution of AgNOg, Ag2S04, AgCgHgOa, or AgF. In other 
 
6 QUALITATIVE CHEMICAL AH^ALYSIS 
 
 words, the chlorine in the first series of salts reacts independently 
 of the metal with which it is united to give a precipitate of AgCI 
 when any soluble salt of silver is added, the silver acting inde- 
 pendently of the acid radical with which it is united in the dry 
 state. In general, all electrolytes having a common com- 
 ponent will respond to a reaction that is characteristic of that 
 component. 
 
 All these differences in the properties of solutions of electro- 
 lytes and non-electrolytes find their explanation in the Theory 
 of Electrolytic Dissociation which was proposed by Svante 
 Arrhenius in 1887. According to this theory when an electrolyte 
 is dissolved in water, it breaks up or dissociates more or less 
 completely into electrically charged particles called tons. These 
 ions are of two kinds, viz. those carrying positive electricity called 
 cations and those bearing negative charges called anions. As 
 the resulting solution is electrically neutral, the sum of all the 
 positive charges in solution must be equal to the total number 
 of negative charges. The change which takes place according 
 to this theory when an electrolyte is dissolved in water may be 
 formulated as follows : — 
 
 HCl :;t H+ + C1-, 
 NaOH:;t:Na++OH- 
 
 NaCl:;t Na++Cl- 
 
 The above shows that dissociation is to be looked upon as a 
 reversible reaction. The extent of the dissociation varies with 
 the dilution ; and at infinite dilution, the dissociation may be re- 
 garded as complete. The -|- sign between the ions indicates 
 that they are to be regarded as independent particles with spe- 
 cific chemical and physical properties. The possession of a 
 charge of positive electricity renders the sodium ion (Na+) totally 
 different from the elementary substance sodium. The same 
 holds true for the hydrogen ion, which is distinctly different from 
 the electrically neutral hydrogen gas. Metallic sodium, as is 
 well known, reacts violently with water; sodium ions do not. 
 We usually recognize hydrogen ions by their sour taste and by 
 their property of changing the color of solutions of certain sub- 
 
INTRODUCTION 7 
 
 stances, called indicators. Hydrogen gas does not exhibit these 
 properties. 
 
 By passing a current through an aqueous solution of NaCl, 
 the electro-negative chlorine ions will travel towards the positive 
 electrode, give up their charges to the latter, and become elec- 
 trically neutral chlorine, which exhibits the well-known properties 
 of this element. The electro-positive sodium ions will at the 
 same time migrate to the negative electrode. Sodium ions 
 are not discharged at this pole owing to the fact that the 
 hydrogen ions, formed by the dissociation of the water, lose 
 their charges more readily than do the sodium ions. Conse- 
 quently on passage of the current, the hydrogen ions lose their 
 charge and electrically neutral hydrogen escapes at the negative 
 electrode. 
 
 Before taking up the applications of the ionic theory let us see 
 how it accounts for the facts mentioned in the early part of this 
 discussion. Since the depression of the freezing point is de- 
 pendent upon the number of particles of the dissolved substance 
 in a given weight of the solvent, it follows that a solution of an 
 electrolyte should depress the freezing point more than a solu- 
 tion of non-electrolyte of the same molar concentration, for the 
 reason that in the former we should have a greater number of 
 particles, each molecule of the electrolyte being capable of yield- 
 ing at least two ions. The same explanation applies to the 
 observed differences in the boiling points. The existence of 
 electrically charged ions in solutions accounts for the conduc- 
 tivity of solutions of electrolytes ; and the failure of solutions of 
 non-electrolytes to conduct the current is due to the absence of 
 ions. Finally, we have to consider how the ionic theory accounts 
 for the independent behavior of the components of electrolytes 
 in solution, which, as already mentioned, makes it possible for 
 salts having a common component to respond to a common reac- 
 tion. Since all chlorides yield chlorine ions, it is perfectly clear 
 why they should all give a precipitate with any soluble silver 
 salt, for the latter, according to the ionic theory, supplies silver 
 ions. This can be best seen by writing the equations for the 
 reaction in ionic form : — 
 
8 QUALITATIVE CHEMICAL ANALYSIS 
 
 Ag+ + NO3- + Na+ + CI- = AgCl + Na+ + NO3-. ( i ) 
 
 Ag+ + F- + NH4+ + CI- = AgCl + NH4+ + F-. (2) 
 
 If we cancel in each equation the ions which appear on oppo- 
 site sides, equations (i) and (2) become : — 
 
 Ag+ + Cl- = AgCl. 
 
 The last equation states that when silver and chlorine ions 
 are brought together, silver chloride is formed ; and as the 
 latter is a difficultly soluble substance, a precipitate of AgCl is 
 generally obtained. Hence we say that Ag ions are a test for 
 CI ions, and vice versa. From solutions, AgCl is formed only 
 by the combination of Ag and CI ions ; therefore solutions of 
 KCIO3 or CHCI3 do not yield a precipitate with silver nitrate 
 for the reason that no chlorine ions are formed by either of 
 these substances in solution. Chloroform in aqueous solution 
 does not conduct the electric current, giving no evidence of dis- 
 sociation, and hence yields no chlorine ions. Potassium chlorate 
 does ionize in solution, but the ions are- K+ and CIO3-. 
 rc Acids and Bases. In terms of the ionic theory an acid may 
 be defined as a substance which, when dissolved, in water, dis- 
 sociates with the formation of hydrogen ions. 
 
 Thus HCl:;tH+-f-Cl-. 
 
 Dibasic acids dissociate in two and tribasic acids in three 
 stages, as shown below : — 
 
 H2S04 
 
 lit H+ -1- HSO4- 
 
 HS04- 
 
 ■:^n+ + so^- 
 
 H3PO, 
 
 :i>H+ + H2P0r. 
 
 H2P04- 
 
 :;t H+ + HPO^- 
 
 HPOr 
 
 -:;tH+ + P04— . 
 
 Thus a solution of sulphuric acid may contain, besides the 
 undissociated acid, the ions H and HSO4. All the character- 
 istic properties of acids in solution — viz. (i) sour taste, (2) the 
 ability to change the colors of indicators, (3) the ability to give 
 
INTRODUCTION 9 
 
 off hydrogen gas when treated with certain metals, (4) the power 
 to neutralize bases — are due to the presence of hydrogen ions. 
 The hydrogen ions are thus the active constituents of acid solu- 
 tions. The differences in the strengths of acids are due to 
 variations in the extent to which they are dissociated. We thus 
 distinguish two classes of acids. Those which dissociate largely 
 and thus supply a high concentration of hydrogen ions, and 
 those which dissociate slightly and thus yield a low concentra- 
 tion of hydrogen ions ; the former are called strong, the latter 
 weak, acids. 
 
 The active component of solutions of bases is the hydroxyl 
 ion (OH~), and to it are attributable all tl^e characteristic prop- 
 erties of bases, viz., (i) the ability to turn^red litmus blue, (2) the 
 alkaline taste, (3) the power to neutralize acids. We may define 
 a base, therefore, as a substance which, when dissolved in water, 
 dissociates with the formation of OH ions. The following are 
 examples : — 
 
 NaOH :;tNa+ -f- OR- 
 
 ' Ca(0H)2 :^ CaOH+ -|- OH- 
 ,CaOH+ :;tCa++ + OR- 
 
 Di-acid bases dissociate in two steps as shown above. As in 
 the case of acids, we distinguish two classes of bases, depending 
 upon the degree to which they are dissociated. A strong base 
 yields a large percentage of OH ions ; a weak base dissociates 
 feebly, yielding a small concentration of OH ions. 
 
 Salts yield on dissociation negative ions of the acid and 
 positive ions of the base, as shown here : — 
 
 NaCl :;tNa+ -+- Q- 
 
 (NH4)2S04:^ NH4+ + NH4SO4-. 
 NH4SO4- :itNH4+ + S04— . 
 
 Conductivity and Dissociation. The conductivity of a solution 
 is conditioned by the presence of ions, for the latter are the 
 means by which the current is transported through the solution. 
 Therefore, the greater the extent to which an electrolyte is disso- 
 ciated in solution, the greater will be its conductivity ; and, con- 
 
lo QUALITATIVE CHEMICAL ANALYSIS 
 
 versely, the greater the conductivity of a solution containing an 
 equivalent weight of an electrolyte, the greater will be its dis- 
 sociation. In the electrical conductivity of a solution, therefore, 
 we have a ready means of determining the degree of dissocia- 
 tion. In practice the resistance of the solution is determined 
 and from the latter the conductivity is readily calculated, the 
 conductivity being the reciprocal of the resistance. 
 
 In conductivity determinations, the results are generally ex- 
 pressed on the basis of equivalent quantities. 
 
 TAe specific resistance of a solution is the resistance in ohms 
 of a cube of one centimeter edge. The specific conductivity is 
 the reciprocal of the specific resistance, and is therefore ex- 
 pressed in reciprocal ohms. The molar conductivity is the con- 
 ductivity of a solution which contains the molecular weight of 
 the substance in grams, contained between electrodes one centi- 
 meter apart. If, instead of the molecular weight, we use the 
 equivalent weight* in grams of the substance, we get the 
 equivalent conductivity. If we denote the equivalent conduc- 
 tivity by X- (lambda) and the specific conductivity by k (kappa), 
 then we have \ = k y.v, where v is the volume in cc. which 
 contains the gram equivalent of the substance. An example 
 will make this relation clear. Suppose we wish to find the con- 
 ductivity of a o. I normal solution of HCl. We first determine 
 the specific resistance, i.e., the resistance of a i cm. cube of the 
 solution when the latter is contained between two parallel elec- 
 trodes. This is found to be 28. 5 ohms. Its specific conductivity 
 
 is therefore , or 0.0351, reciprocal ohm. Since the solu- 
 
 28.5 
 
 tion is tenth normal, one equivalent will be contained in 10 liters, 
 or 10,000 cc. Therefore \ = equivalent conductivity = 10,000 x 
 0.0351 = 351 reciprocal ohms. 
 
 The equivalent conductivity is found to increase with the di- 
 lution up to a certain point and then remains constant, so far as 
 measurements show. At this point of maximum conductivity, 
 
 * In the CEise of monobasic acids and their salts, the equivalent and molar con- 
 ductivities vifill be identical. With dibasic acids and their salts, the equivalent wiU 
 be one half of the molecular weight expressed in grams. 
 
INTRODUCTION il 
 
 the electrolyte must be completely dissociated into its ions, and 
 hence no increase in conductivity results on further dilution. 
 The maximum conductivity is generally referred to as the con- 
 ductivity at infinite dilution and is usually calculated from the 
 curve representing the conductivities for known dilutions. 
 Since the conductivity is proportional to the number of ions in 
 solution, it follows that the ratio of the equivalent conductivity 
 at a given dilution, to that at infinite dilution, will give us the 
 percentage of the electrolyte that is ionized or the degree of dis- 
 sociation at the given dilution; or a = rr^-, where « denotes the 
 
 A.0O 
 
 degree of dissociation of the electrolyte at the dilution v, X, the 
 equivalent conductivity at the dilution v, and \« the equivalent 
 conductivity at infinite dilution. An example will illustrate the 
 application of the above principle. 
 
 The specific resistance of a HCl solution containing 1.825 
 grams of HCl per liter is found by experiment to be 55.55 ohms 
 
 at + 18°. K is therefore reciprocal ohms. To calculate 
 
 5S-55 
 
 the value of \„, we need only note that 1.825 g. HCl is ' , or 
 
 ■ — •, of a gram equivalent. Therefore v, the volume which will 
 
 contain one gram equivalent (36. 5 g.) of HCl, will be 20 liters, or 
 
 20,000 cc. Hence \„ = 20,000 x = 360 reciprocal ohms. 
 
 55-55 
 That is to say, the equivalent conductivity of such a solution is 
 360 reciprocal ohms. The equivalent conductivity of an HCl 
 solution at infinite dilution at 18°, obtained by extrapolation 
 from the curve of conductivities of known dilutions, is found to 
 be 384 reciprocal ohms. Therefore a = |f| = 93.75 %. That 
 is to say, a solution containing 1.825 g. HCl per liter or a 3*5 
 molar solution of HCl, 93.75 % of the HCl is dissociated into H 
 and CI ions, while but 6.25 % is in the un-ionized condition. 
 
 In the following table are given the approximate values of 
 the dissociation of the more common electrolytes met with in 
 qualitative analysis : — 
 
12 QUALITATIVE CHEMICAL ANALYSIS 
 
 ACIDS 
 
 
 
 Pekcentage Ionized in 
 
 
 
 
 0.1 Normal Solution 
 
 H+C1-; H+Br-; H+I"; 
 
 H+NOg- 
 
 - 
 
 
 90 
 
 H+ HSO4- 
 
 
 
 
 60 
 
 H+ HC2O4- 
 
 
 
 
 34 
 
 H+ HSO3- 
 
 
 
 
 20 
 
 H+ H2PO4- 
 
 
 
 
 13 
 
 H+QHgOa" 
 
 
 
 
 1.4 
 
 H+ HCO3- 
 
 
 
 
 0.12 
 
 H+HS- 
 
 
 
 
 0.05 
 
 H+CN- 
 
 
 
 
 0.0 1 
 
 BASES 
 
 
 
 
 
 K+OH-; Na+OH- 
 
 
 
 
 86 
 
 Ba+(OH-)2 
 
 
 
 
 75 
 
 NH4+OH- 
 
 
 
 
 1.4 
 
 SALTS 
 
 
 
 
 
 Type B+ A-(KN03, KCl) 
 
 
 
 
 84 
 
 Type B2+ A— or B++ A2- 
 
 (K2S04, 
 
 CaCla) 
 
 
 73 
 
 Type B3+ A or B+++ A, 
 
 rCKgFe 
 
 (CN)e, 
 
 FeCl3) 
 
 65 
 
 TypeB++A— (MgSOJ 
 
 
 
 
 40 
 
 Pure water 
 
 
 
 
 O.CXX300002 
 
 The Law of Mass Action - - ;\ .n,^^/ il ' ' /' 
 
 In the following reversible reaction, I 
 
 A + B :;> C + D, 
 
 let us denote the concentrations * of each of the substances by 
 [A], [B], [C], and [D] respectively. Let v represent the ve- 
 locity with which A and B react to form C and D and w' the 
 rate at which A and B are formed by the reaction between C 
 and D. Experiment shows that, all other things being constant, 
 V will be proportional to [A] and also to [B] and therefore to 
 their product, or 
 
 v = kxlA']x [B], 
 
 * These are expressed in moles per liter ; thus a HCl solution containing 
 
 1.825 g. HQ per liter contains — — 2 = — mole ; or i cc. contains 0.00005 mole. 
 
 36.5 20 
 
INTRODUCTION 13 
 
 where >& is a constant dependent upon the nature of the sub- 
 stances reacting as well as upon the temperature. Similarly, 
 the velocity of the reaction in the reverse direction, i.e., the 
 formation of A and B by the interaction of C and D may be 
 represented by 
 
 v' = k' X [C] X [D]. 
 
 At equilibrium, the velocities in both directions will be equal ; 
 and we have therefore 
 
 v = v' orkxlAjx [B] = y^' X [C] X [D], or 
 
 [C] x\_D'\ _k 
 [A] X [B] k' 
 
 = ^ = K. 
 
 Expressed in words the above equation states that the product 
 of the concentrations of the final substances, divided by the 
 product of the concentrations of the initial substances, is 
 a constant for any definite temperature for a reversible 
 reaction. 
 
 Where three substances are involved on each side of the equa- 
 tion, the same principle applies. Thus the equation expressing 
 the equilibrium for the reaction 
 
 A + B + CijD + E + FwouIdbe m^I|^ = K. 
 
 If two combining weights of any substance take part in the 
 reaction, then each is to be separately considered; thus the re- 
 versible reaction 
 
 A-f2B:^C-l-D may be written A-FB-|-B^C-|-D; 
 
 and the equilibrium equation becomes 
 
 [C]x[D] _,^ 
 [A]X[BP 
 
 In its most general form, the /aiv of mass action as expressed 
 in the above equation may be stated as follows: At any constant 
 temperature, equilibrium will be reached in a reversible reaction, 
 when the product of the concentrations of the final substances 
 
14 QUALITATIVE CHEMICAL ANALYSIS 
 
 divided by the product of the concentrations of the initial sub- 
 stances — each concentration raised to a power equal to the num- 
 ber of combining weights reacting — is a constant. 
 
 The Mass Action Law and Ionization. As the ionization 
 of an electrolyte is a reversible reaction, we should expect 
 that it would obey the law of mass action. Experiment 
 shows, however, that while the mass action law rigidly 
 applies to weak electrolytes, such as ammonium hydroxide 
 and acetic acid, it does not apply, for reasons which still 
 remain to be explained, to the dissociation of strongly ionized 
 acids, bases, and salts. 
 
 An example will illustrate the application of the mass action 
 law to ionization. A normal solution of ammonium hydroxide 
 is found by experiment to ionize to the extent of 0.4%, Ioniza- 
 tion takes place according to the scheme 
 
 ■ NH^OH :;t NHi+ -t- OH- (i) 
 
 At equilibrium at a given temperature, we have 
 
 [NH/] X [OH-] _ 
 
 [NH.OH] ^^ 
 
 Substituting the values in the above equation, we get 
 
 0.004 X 0.004 ^ 
 
 — - — - = 0.000016. 
 
 0.996 
 
 Influence of Dilution. Let us suppose that a normal solution 
 of NH4OH is diluted n times. Then the value of each of the 
 
 concentrations in equation (2) will be reduced to - of its original 
 
 n 
 magnitude, and we have 
 
 i[NH/]xi[OH-] 
 
 «■- '' n^ or I^lVliilOM:] < K. 
 
 i[NH,OH] «[NH,OH] 
 
 The fraction will no longer be equal to the equilibrium con- 
 stant K, because the value of the denominator has been increased 
 
INTRODUCTION 15 
 
 « times its original value. In order to restore the disturbed 
 equilibrium, reaction (i) must proceed from left to right until, 
 by an increase in the concentrations of OH" and NH4+ and a 
 consequent decrease in the undissociated NH4OH, the ratio 
 again becomes equal to K. It is evident that the net result of 
 the dilution is an increase in the dissociation. 
 
 The Common Ion Effect. The ionization of acetic acid is 
 governed by the following equilibrium equation : — 
 
 [H+] X [C^HgO,-] _ 
 [HC2H3O2] 
 
 Since a molar solution of acetic acid is ionized to the extent 
 of 0.42 <fo, we have, substituting in the above equation, 
 
 0.0042 X 0.0042 o 
 
 ^ — =^ = 0.000018. 
 
 0.9958 
 
 If to this solution we add i mole or 82 g. of NaCgHgOa per 
 liter, the salt being 53 per cent, dissociated, the concentration 
 of the C2H3O2 ions would be increased to 0.0042 + 0.53 = 0.5342. 
 If we substitute this value in the above equation, we get 
 
 — — ^ ' , which is no longer equal to the equilibrium con- 
 
 0.9958 
 
 stant 0.000018, and as a consequence the equilibrium will be 
 disturbed. The equilibrium can be restored only by an adjust- 
 ment of the values of the terms of the fraction. The denomina- 
 tor in the original equation for normal acetic acid when present 
 alone (0.9958), is so near its maximal value that any increase 
 in its magnitude (its highest possible value is i) due to a com- 
 bination of some of the H+ and CjHgOj" — to form undissociated 
 HCjHjOa — will be negligible. We therefore need only consider 
 the concentrations of the ions in restoring equilibrium. The 
 
 new value for the concentration of the CoHoO," ions is — i534i 
 
 0.0042 
 
 or 126 times as great as it was originally, and therefore to re- 
 store equilibrium, the other factor, the concentration of the 
 hydrogen ions, must be reduced to y^ of its initial value ; i.e.. 
 
i6 QUALITATIVE CHEMICAL ANALYSIS 
 
 it will be reduced to -j^g x 0.0042 = 0.000034. The net effect 
 of adding to a weakly dissociated acid a salt with an ion in 
 common is to reduce the concentration of the hydrogen ions or 
 to repress the ionization of the acid. A similar effect is pro- 
 duced when a salt with a common ion such as NH^Cl is added 
 to a weak base, such as ammonium hydroxide. The original 
 concentration of the OH ions will be reduced by virtue of an 
 increase in the concentration of the NH^ ions. 
 
 Solubility Product. In a saturated solution of AgCl equilib- 
 rium will be represented by the following equation : — 
 
 But since in a saturated solution [AgCl] is a constant, we may 
 write the above equation in the form [Ag+] x [Cl~] = K. 
 
 This equation states that for a saturated solution of a diji- 
 cultly soluble electrolyte, the product of the concentrations of its 
 ions is a constant for a given temperature. This product is 
 called the solubility product.* 
 
 With sparingly soluble electrolytes such as AgCl and BaSO^, 
 the dissociation of the dissolved substance may be regarded as 
 complete so that the concentration of the ions will be equal to 
 the solubility. An example will make this clear. The solubil- 
 ity of AgCl has been found to be 0.0000106 moles per liter. 
 The solubility product will be therefore (0.0000106).^ 
 
 The application of the principle of the solubility product 
 to precipitation may be stated as follows : Whenever the 
 product of the concentrations of the ions, which by their 
 union are capable of forming an insoluble substance, is 
 greater than the solubiUty product of that substance, pre- 
 cipitation takes place. Conversely, when this product of the 
 ionic concentrations is less than the solubility product of the 
 insoluble substance, the latter, if present, will dissolve until 
 this value is reached. 
 
 * The principle of the solubility product cannot be applied to saturated solutions 
 of soluble salts. 
 
INTRODUCTION 17 
 
 Illustrations. An aqueous solution of HgS behaves as a 
 weakly dissociated acid. The equilibrium equation is 
 
 [H+]^x[S— j _^^ 
 [H2S] 
 
 The addition of HCl to the solution will have the effect of in- 
 creasing [H+] and will correspondingly diminish [S ] (common 
 ion effect). Now if a solution of a copper salt to which some 
 HCl be added \ is treated with HjS, a precipitate is thrown 
 down. The formation of a precipitate is due to the fact that 
 the solubility product of CuS is exceedingly small and that in 
 spite of the repression of the ionization of the HjS by the free 
 HCl, the greatly diminished concentration of the S ions is still 
 sufficient to yield with the Cu ions a value greater than the 
 solubility product of CuS. On the other hand, if a zinc solu- 
 tion containing the same concentration of HCl be treated with 
 HgS, no precipitate of ZnS is obtained. The reason for this is 
 that although the solubility product of ZnS is quite small, being 
 much larger than that of CuS, the diminished concentration of 
 the S ions occasioned by the presence of the HCl is far too low 
 in value to cause the solubility product of ZnS to be exceeded, 
 however great the concentration of the Zn ions may be. If, 
 however, the solution be rendered alkaline, a precipitate of ZnS 
 is at once obtained. The addition of the base not only neutral- 
 
 • Hydrogen sulphide being a dibasic acid in solution ionizes as follows : — 
 
 H2s:1;:h+ + hs-; (0 
 
 HS-^H+-|-S — . (2) 
 
 The equilibrium equations for (i) and (2) are 
 
 [H+]x[HS-] _, [H+]x[S— ] _, 
 
 [H2S] [HS-] 
 
 Multiplying these two equations together and cancelling, we get 
 
 [H2S] 
 
 which is the equation given above. 
 
 1 2.5 CO. cone. HQ in a volume of 100 cc. of solution is a suitable concentration, 
 c 
 
i8 QUALITATIVE CHEMICAL ANALYSIS 
 
 izes the HCl present, thus causing an increase in the concentra» 
 tion of the S ions, but the HjS itself is converted into an alkali 
 sulphide whose degree of dissociation is vastly greater than that 
 of H2S. This enormous increase in the concentration of the S 
 ions accounts satisfactorily for the fact that in an alkaline solu- 
 tion the precipitation of zinc as sulphide by HjS is quantitative. 
 
 By an application of the principles involved in the solubility 
 product and the " common ion effect," the student will readily 
 comprehend why an excess of the reagent is required to effect 
 a maximum precipitation of an ion. For example, in the pre- 
 cipitation of silver ions, it is customary to add an excess of chlo- 
 rine ions. The latter will reduce the concentration of the silver 
 ions to a negligible quantity and thus afford a practically com- 
 plete precipitation. 
 
 The solubility of CaC204 in an excess of HCl may be explained 
 from the standpoint of the ionic theory and solubility product 
 principle as follows : Experiment shows that oxalic acid is a 
 weakly dissociated acid ; hydrochloric acid, on the other hand, 
 is a strongly ionized acid. In consequence of this difference, 
 the tendency, where the ions of both acids are in solution, will 
 be for the weakly dissociated acid to form. In the case under 
 consideration, some of the hydrogen ions from the HCl will 
 unite with some of the CjO^ ions derived from the CaC204 in 
 solution, to form the weakly dissociated oxalic acid. The re- 
 moval of C2O4 ions from the solution will cause the product of 
 the concentrations of the Ca and C2O4 ions in solution to fall 
 below the solubility product of CaCjO^. In consequence more 
 CaC204 will go into solution in order to restore equilibrium ; and 
 this process of the removal of CjO^ by the H ions, and the solu- 
 tion of the CaC204, will go on until all the CaC204 is dissolved. 
 The reaction may be represented as follows : — 
 
 Ca++ + CaOi— -f 2 H- = H2C2O4 -1- Ca++. 
 
 Hydrolysis. While the concentrations of OH and H ions in 
 water is exceedingly small, there being only -j-rB 7 00 ^^ ^ 
 mole for each of these ions in a liter of water, they must be 
 taken into account in dealing with the behavior of solutions of 
 
INTRODUCTION 19 
 
 salts formed by the union of a weak base with a strong acid, 
 and vice versa. Thus the student has already become familiar 
 with the fact that aqueous solutions of NajCOg and KCN possess 
 an alkaline reaction, while solutions of FeClg and CuSO^ are 
 acid to litmus. In the former case the reaction is due to the 
 presence of OH ions, while in the latter case the acidity is occa- 
 sioned by the presence of H ions in solution. 
 
 When potassium cyanide is dissolved in water, it dissociates 
 to a considerable extent into K and CN ions. We have, there- 
 fore, in solution for the moment K+, CN~ H+, and 0H~. Now 
 on encountering the CN ions, the H ions will combine with 
 them, forming the very feebly dissociated acid HCN. The 
 removal of H ions from the solution will cause more water to 
 dissociate into H and OH ions ; the former will again be taken 
 up by more CN ions, and this process will continue until the 
 equilibrium constant for HCN is reached, i.e., until 
 
 [H+]x[CN-] _ 
 [HCN] 
 
 The removal of the H ions from the solution will result in the 
 accumulation of OH ions, and these will give to the solution its 
 alkaline reaction. The reaction may therefore be represented 
 as follows : — 
 
 K+ + CN- + H+ + OH- = HCN + K++ OH" 
 
 It is evident that the OH ions will not combine with the K 
 ions for the reason that KOH, even if formed, would be largely 
 dissociated in aqueous solution. 
 
 By a similar process of reasoning it may be shown why an 
 aqueous solution of FeClg possesses an acid reaction. In this 
 case, the weakly dissociated base Fe(OH)3 is formed, which 
 removes the OH ions from the solution, thus causing an accu- 
 mulation of H ions from the progressive dissociation of water. 
 The final solution, therefore, will possess an acid reaction. We 
 may define hydrolysis as the process by which a salt of a weak 
 acid or base is decomposed into its corresponding acid and base 
 by the action of water. 
 
20 QUALITATIVE CHEMICAL ANALYSIS 
 
 Equations 
 
 A chemical equation is a shorthand means of expressing a 
 reaction. The equation 
 
 BaCLj + H2S04= BaS04+ 2 HCl 
 
 gives us, first, the qualitative fact that when BaClj solution is 
 treated with HgSO^, a precipitate of BaS04 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 BaS04 and (2 X 36.5) parts by 
 weight of HCl. 
 
 In the case of gases, equations give us, besides qualitative 
 and quantitative relations, those of volume also. The equation 
 
 CO2 + C = 2 CO 
 
 states that when one volume of COg 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 
 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 ; e.g., a glance at the equation 
 
 I CaCOg + 2 HC2H3O2 = Ca(C2H302)2 + H2O + f COj 
 
 shows that CaCOj is a solid, that CO2 is a gas, and that the 
 other substances are in solution. Any determined method. 
 
INTRODUCTION 2i 
 
 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 or gain 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 elem.ent 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 
 a compound is formed from its elements, or, in general, a more 
 complex compound formed by the union of simpler ones, thus : — 
 
 H2 + Cl2 = 2HCl; 
 CaO+H20 = Ca(OH)2; 
 Fe(CN)2 + 4 KCN = K4Fe(CN)6. 
 
 An analytical reaction is one in which some complex com- 
 pound is broken down into its elements or into simpler com- 
 pounds ; for example : — ,^ ^, ,^ „, 
 
 ^ NaCl = Na-fCl; 
 
 CaCOg = CaO -I- COj. 
 
 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 : — 
 
22 QUALITATIVE CHEMICAL ANALYSIS 
 
 («) AgNOs + HCl = I AgCl + HNO3 ; 
 
 (b) Cu(N03)2 + HaS = I CuS + 2 HNO3 ; 
 
 (c) BaCla + H2SO4 = | BaSO^ + 2 HCl. 
 
 Neutralization of bases by acids comes under this head. 
 NaOH + HCl = NaCl + H2O ; 
 NH4OH + HC2H3O2 = NH4C2H3O2 + H2O ; 
 Ca(OH)2 + H2SO4 = I CaSO^ + 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 : — Cu(N03)2 + H2SO4 = CuSO^ + f 2 HNO3. 
 Similarly, on boiling, the following reactions will take place : — 
 NH4CI + NaOH = NaCl + f NH3 + H2O, 
 NaCl + H2SO4 = NaHSO^ + 1 HCl ; 
 
 for, at these temperatures, HCl, NH3, and HNO3 are evolved as 
 gases. 
 
 In writing a complicated equation, it is convenient to consider 
 the reaction as taking place in several stages ; * thus, when 
 ammonium sulphide is added to an aqueous solution of AICI3, 
 a precipitate consisting of the hydroxide and not the sulphide 
 is obtained, and at the same time the evolution of HgS takes 
 place. This reaction becomes easily comprehended if we con- 
 sider it as taking place in two stages, viz. : — 
 
 (i) 2 AICI3 + 3(NH4)2S = AI2S3 + 6 NH4CI; 
 
 but as AI2S3 does not exist in contact with water, it decomposes 
 according to the equation 
 
 (2) ALjSg + 6 H2O = 1 2 Al(OH)3 + f 3 H2S. 
 
 * Since we do not know what causes chemical activity, the exact mechanism of a 
 reaction is unknown and is perhaps unknowable ; 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. 
 
INTRODUCTION 
 
 23 
 
 Combining (i) and (2), and eliminating AljSj, which appears 
 on opposite sides of the equations, we get as the final equation 
 
 2AlCl3+3(NH^)2S+6H20 = |2Al(OH)3+6NH4Cl+t3H2S. 
 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 : 
 
 (i) 2FeO-|-0 = Fe203; 
 
 (2) FeClg + CI = FeClg ; 
 
 (3) 2 FeSO^ -I- H2SO4 -I- Bra = FejCSO J3 -1- 2 HBr ; 
 
 (4) SnCl^ + 2 HgCla = SnCI^ + \2 HgCl. 
 
 Reduction : 
 
 (5) CuO-l-H2=Cu-fH20; 
 
 (6) 2 FeCl3 -f SnCla = 2 FeClg -f- SnCl^ ; 
 
 (7) Fe2(SO,)3 4- H2 * = 2 FeSO^ -h HjSO,; 
 
 (8) 2 HgCla + SnClg = 1 2 HgCl -I- SnCl^. 
 
 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- 
 
 * The hydrogen is generated in contact with the Fe2(S04)3 by the action of a 
 metal on an acid. 
 
24 QUALITATIVE CHEMICAL ANALYSIS 
 
 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 SnClg has been accomplished at 
 the expense of the HgCl^, 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 HjO. Similarly, in (6) the reduc- 
 tion of FeClg has been accompanied by the oxidation of SnCLj 
 to SnCl^. 
 
 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 i. 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, KCIO3, ^Oc^O^, KMnO^, NajOa, HgOa. 
 and PbOg. 
 
 The chief reducing agents are : nasceijt H, SnClg, HjS, HjSOg ; 
 and, at elevated temperatures, C, KCN, and organic matter. 
 
 The halogens either add directly (i); or else maybe 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. 
 
 (i) FeClj + CI = FeClg ; 
 
 (2) 2 FeSO^ + H2SO4 + CI2 = Fe2(S04)8 + 2 HCl ; 
 
 (3) Fe(0H)2 + KOH + CI = Fe(0H)8 + KCl. 
 
INTRODUCTION 25 
 
 Nitric acid acts as an oxidizing agent by virtue of its ready 
 decomposition according to the equation 
 
 2 HNOg = 2 NO + H2O + 3 O. 
 
 From this it is evident that 2 formular weights of HNOg yield 
 3 combining weights of O. The liberated O may be regarded 
 as acting indirectly in the same way as CI ; that is, by oxidizing 
 the H of the acid and liberating the acid radical, which then 
 adds on, thus : — 
 
 (i) 2HN03 = 2NO+H20 + 3 0; 
 
 (2) 3 H2SO4 + 3 0.= 3 H2O + 3 SO4. 
 
 Since 2 formular weights of FeSO^ 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 FeS04, to wit : — 
 
 (3) 6 FeSO^ + 3 SO, = 3 Fe^CSO 4)3. 
 
 Adding (i), (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 FeSO, -I- 2 HNO3 -I- 3 H2SO4 = 3 Fe2(SOj3 + f 2 NO + 4 H2O. 
 
 Aqua regia is not only a good solvent but is also an excellent 
 oxidizing agent. Its action is practically the same as that of CI. 
 It is prepared by mixing i part of HNO3 with 3 parts of HCl. 
 When heated alone, it is said to yield NOCl and CI, according 
 to the equation 3 HCl -f- HNO3 = NOCl + Clj -f- 2 HjO ; but in 
 the presence of an oxidizable substance all of the CI is available, 
 so that the equation with aqua regia may be written as follows : — 
 
 3 CoS + 6 HCl -f- 2 HNO3 = 3 C0CI2 -I- 2 NO -I- 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 -f 6 CI -I- 3 H2O = SO3 -f- 6 HCl; 
 
 SOg + H2O = H2SO4. 
 
26 QUALITATIVE CHEMICAL ANALYSIS 
 
 Potassium chlorate, KClOj, when used in conjunction with 
 HCl, is a powerful oxidizing agent. In effect it is similar to agua 
 regia and chlorine, as the following equations will indicate : — 
 
 (i) 2KC103 + 2HC1 = 2HC103 + 2KC1; 
 
 (2) 2 HCIO3 = H2O + 2 CIO2 + O ; 
 
 (3) 2HC1+0 = H20 + Cl2. 
 
 Adding (i), (2), and (3), we get the following : — 
 
 2 KCIO3 + 4 HCl = 2 KCl + 2 H2O + 2 CIO2 + Clj. 
 
 Potassium dichromate, KjCrgO^, contains the acid anhydride 
 CrOg ; its composition may be represented by the formula 
 
 2 CrOg + KjO. As 2 CrOg will, on reduction, yield Cr203 + 3 O, 
 it is evident that one formular weight of K2Cr207 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 
 
 6 FeSO^ + K2Cr207 + xYi^^O,^ = 
 
 3 Fe2(S04)3 + KaSO^ + Cr2(S04)3 + ;irH20. 
 
 The reduction of 2 CrOg to CrgOs leaves the latter, as well as 
 KjO, 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 : Cr203 
 requires 3 formular weights of HgSO^ for its solution; KgO 
 requires i, 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 CrOg ; this will 
 require 3 more H2SO4, making the total 7, which becomes the 
 coefficient of H2SO4 in the above equations. We should also 
 have 7 HjO in the right-hand member of the equation for a 
 reason which must be evident to the reader. 
 
 Potassium permanganate, KMn04, behaves in acid solution as 
 
 KjO + 2 MnO + 5 O = 2 KMnO^; 
 that is, 2 KMn04 yields 5 combining weights of O. 
 
INTRODUCTION 27 
 
 The equation for the oxidation of ferrous sulphate by KMn04 
 may be written from the following considerations: 5 combin- 
 ing weights of O will liberate S SO4 radicals ; but as one SO^ 
 suffices for the oxidation of 2 FeS04, it is evident that 2 KMn04 
 will oxidize 10 FeSO^, and we get the equation — 
 
 2 KMn04 + 10 FeSO^ + 5 H2SO4 = 
 
 5 Fe2(S04)3 + 2 MnO + KjO + 5 H2O. 
 
 However, KgO 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 KMnOi + 10 FeS04 + 8 H2SO4 = 
 
 5 Fe2(S04)3 + 2 MnS04 + K2SO4 + 8 HjO. 
 
 Sodium dioxide, NagOg, acts as an oxidizing agent by virtue 
 of its instabihty in water or when it is heated in the presence 
 of an oxidizable substance in consequence of which it gives off 
 oxygen : — ^^q^ ^ jj^q ^ ^ NaO H + O. 
 
 A solution of hydrogen dioxide, H2O2, also serves as an oxidiz- 
 ing agent for the same reason : — 
 
 H2O2 = H2O + O. 
 
 NagOg 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. Na202 has the further ad- 
 vantage of supplying at the same time sodium hydroxide. For 
 oxidation in alkaline media, therefore, NagOg is the better reagent. 
 
 Lead dioxide, PbOg, like NajOa, yields O according to the 
 equation— Pb02 = PbO -I- O. 
 
 It is employed to effect oxidation in acid media, as in the 
 conversion of MnO to HMn04; the PbO is converted by the 
 excess of acid present into a salt. 
 
 * 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. 
 
28 QUALITATIVE CHEMICAL ANALYSIS 
 
 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 * acts as a reducing agent by either adding itself 
 directly (i) or by its ability to unite with, and thus remove, the 
 halogens contained in the compound (2): — 
 
 (i) As + 3H = AsH3; 
 
 (2) HgCla + 2 H = I Hg + 2 HCl. 
 
 Nascent H may be prepared by the action of an acid or alkali 
 on certain metals : — 
 
 Zn + H2SO4 = ZnS04 + 2 H ; 
 Zn + 2 NaOH = NagZnOa + 2 H ; 
 Al + 3 NaOH = NagAlOg + 3 H. 
 
 Nascent H may therefore be employed in both acid and alkaline 
 media. 
 
 Stannous chloride, SnClj, acts as a reducing agent, preferably 
 in an acid solution, by virtue of the ease with which it readily 
 oxidizes to SnCl^ : — 
 
 SnClg + 2 HgCla = SnCl4 + j,2 HgCl ; 
 SnClj + 2 FeClg = 2 FeClj + SnCl4; 
 • H3ASO4 + 2 HCl + SnCla = H3ASO3 + HjO + SnCl4. 
 
 Hydrogen sulphide, HjS, by virtue of the readiness with which 
 it decomposes into H and S is a reducing agent. The hydrogen 
 acts as nascent hydrogen, while the sulphur separates out in the 
 soHd state. Nearly all oxidizing agents are reduced by HgS with 
 the separation of sulphur, thus : — 
 
 2 HNO3 + 3 HjS = 4 H2O + 2 NO + 3 S. 
 
 Hence, sulphides soluble in HNO3, like those of Pb, Bi, and Cu, 
 do not liberate HjS, because the latter at once acts on the 
 
 * 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 29 
 
 excess of HNO3 present with the hberation of S, as indicated 
 in the equation above. Halogens, agua regia, ferric salts, 
 potassium permanganate, and chromates are also reduced by 
 HjS with the separation of S, and, in some cases, with a change 
 in color of the solution (see pp. 60-61). 
 
 Sulphurous acid, HjSOg, and the sulphites easily remove 
 oxygen from compounds with the formation of sulphuric acid 
 and sulphates respectively, thus : — 
 
 H3ASO4 + H2SO3 = HgAsOg + H2SO4. 
 
 While HgSOg 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 HgSOg is best 
 accomplished in a pressure bottle at ICX)° C. 
 
 Carbon acts as a reducing agent by virtue of its ability to 
 oxidize to CO and COj, thus : — 
 
 CuO + C + (heat) = Cu + CO. 
 
 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: — 
 
 SnOa + 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 HCl used, is treated with a stream of HgS gas, there 
 
30 QUALITATIVE CHEMICAL ANALYSIS 
 
 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,t 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 NH^OH and some NH^Cl.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 
 group. The final filtrate will contain all the other metals not 
 precipitated by the previous group reagents; they are magne- 
 sium, sodium, potassium, and ammonium, and these form the 
 fifth group. I 
 
 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 the order, — i.e., begin with NH^OH, 
 then add (NH4)2C03 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 viVid^a. follow and not from those which 
 
 * Since PbCla 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 HCl is rendered alkaline with ammonium 
 hydroxide, NH4Q 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. 
 
 % The rarer metals are not considered in the text. Special treatises more or less 
 elaborate are necessary when they are included. 
 
INTRODUCTION 
 
 31 
 
 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 HCl and iilter. 
 
 Precipitate: AgCl, 
 PbClj, HgCl. 
 Group I. 
 
 Filtrate : Groups II.-V. + excess HCl ; pass in HjS 
 and iilter. 
 
 Precipitate : 
 HgS, PbS, 
 BijSs, CuS, 
 CdS, AS2S3, 
 SbjSj, SnS. 
 
 Group II. 
 
 Filtrate : Groups III.-V. + HCl + HgS ; 
 make alkaline with NH^OH, add 
 (NHi)2S and filter. 
 
 
 Precipitate : 
 A1(0H)3, 
 Cr(0H)3, 
 FeS, NiS, 
 CoS, MnS, 
 ZnS. 
 
 Group III. 
 
 FUtrate: Groups IV.- 
 V. + NH4CI ; add 
 (NH4)2C03 and filter. 
 
 
 Precipitate : 
 BaCOg, 
 SrCOg, 
 CaCOg, 
 
 Group IV. 
 
 Filtrate: 
 contains 
 Mg, Na, 
 K, NH4. 
 Group V. 
 
 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 HCl. 
 
 Group II. includes those metals which are not precipitated by 
 HCl, but whose sulphides are precipitated by HjS 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 HCl or HjS in acid solution, but are precipitated by 
 (NH4)2S in solutions alkaline with NH4OH and in the presence 
 of NH4CI. 
 
 Group IV. includes those metals unprecipitated by the reagents 
 of Groups I., II., III., but which are precipitated by (NH JgCOg 
 in the presence of NH4CI. 
 
 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 concen- 
 trated acids. Ammonium hydroxide readily dissolves it, with 
 the formation of silver ammonia chloride [Ag(NH3)2Cl] : — 
 
 AgCH- 2 NH3 = Ag(NH3)2Cl, 
 
 from which AgCl reprecipitates on acidification with nitric add : — 
 
 Ag(NH3)2Cl -h 2 HNO3 = I AgCl -I- 2 NH4NO3. 
 
 Silver chloride also dissolves in solutions of potassium cya- 
 nide and sodium thiosulphate; when cautiously heated, it fuses 
 without decomposition. 
 
 2. Hydrogen Sulphide and soluble sulphides precipitate black 
 
 AggS, insoluble in cold dilute acids, alkali hydroxides, and alkali 
 
 sulphides; it is soluble in hot dilute HNO3, with the formation 
 
 of AgNOs and separation of sulphur. The reaction can be con- 
 
 B 33 
 
34 QUALITATIVE CHEMICAL ANALYSIS 
 
 sidered as taking place in two steps, the first consisting of the 
 solution of the sulphide with the liberation of HgS and the 
 second of the oxidation of the HjS by the excess of HNOg 
 present with the formation of water, nitric oxide, and the sepa- 
 ration of sulphur : — 
 
 (i) Ag2S-l-2HN03 = 2AgN03 + tHaS; 
 
 (2) 2 HNO3 + 3 H^S = 4 H2O + f 2 NO + I 3 S. 
 
 Multiplying equation (i) by 3 and adding it to (2), with the 
 elimination of 3 HgS, which appears on opposite sides, we get 
 as a final result : — 
 
 3 AgaS + 8 HNO3 = 6 AgNOg + 4 H2O + f 2 NO + 13 S. 
 
 Other insoluble compounds of silver are: — 
 
 AgBr — yellowish white; Agl — pale yellow; AgCN — 
 white; AggO — 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) 
 
 I. Hydrochloric acid and soluble chlorides give with solutions 
 of mercurous salts a white precipitate of HgCI (calomel), in- 
 soluble in water, the solubility being about i part in 300,000 ; 
 it is insoluble in cold dilute acids, but dissolves in strong nitric 
 acid and in agua regia, the latter oxidizing it to HgClj. Am- 
 monium hydroxide converts calomel into a black mixture of 
 finely divided mercury and NHjHgCl, insoluble in excess of the 
 reagent (this is the most characteristic reaction for mercurous 
 salts) : — 
 
 2 HgCl -I- 2 NH3 =|(NH2HgCl -I- Hg) + NH4CI. 
 
 The black mixture dissolves in aqua regia with the formation of 
 mercuric chloride : — 
 
 (i) NH2HgCl+3 Cl = tN + 2HCl-fHgCl2; 
 (2) Hg-F2Cl-HgCl2. 
 Adding (i) and (2), we get 
 
 NHaHgCl -f- Hg -I- 5 CI = 2 HgClj -f- f N + 2 HCl. 
 
THE METALS 
 
 35 
 
 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 = I HgjS + 2 HCl; 
 Hg,S = >|,HgS + |Hg. 
 
 3. Reducing Agents, as FeSO^ or SnClj, rapidly reduce mer- 
 curous salts to metallic mercury: — 
 
 SnCla + 2 HgCl = SnCl^ + >|' 2 Hg. 
 
 Lead 
 
 1. Hydrochloric acid and soluble chlorides give with solutions 
 of lead salts, which are not too dilute, a white precipitate of 
 PbClj, soluble in 100 parts of cold and 25 parts of boiling water; 
 from the latter on cooling, PbClg separates out in the form of 
 needles. PbClj is much more insoluble in dilute HCl 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 
 PbS04 which is practically insoluble in HgO (i 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 PbS04 is reprecipi- 
 tated on adding H2SO4. 
 
 3. Potassium Chromate (K2Cr04) precipitates yellow lead 
 chromate, readily soluble in sodium hydroxide, but insoluble in 
 NH4OH and acetic acid: — 
 
 Pb(C2H302)2 + K2Cr04 = \ PbCrO^ + 2 KC2H3O2. 
 
 4. Hydrogen Sulphide from slightly acid solutions of lead salts 
 precipitates black PbS. In the presence of much HCl, H2S 
 either fails to precipitate or else produces a red insoluble com- 
 
36 QUALITATIVE CHEMICAL ANALYSIS 
 
 pound of the formula PbClj • 2 PbS; the latter is converted into 
 black PbS by treatment with (NH4)2S or by diluting the solu- 
 tion and passing in more HgS. 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(N03)2 4- H2S =,lr PbS -I- 2 HNO3. 
 
 With dilute HNO3 this reaction occurs: — 
 
 3 PbS -h 8 HNO3 = 3 Pb(N03)2 -|-f2NO-|-|3S-|-4 ^A 
 
 * 
 
 With concentrated HNO3, the reactions may be represented 
 by the following equations : — 
 
 (i) PbS4-2HN03 = |PbS04-h2NO-|-H2; 
 (2) 2 HNO3 -h 3 H2 = 2 NO + 4 H2O. 
 
 Multiplying (i) by 3 and adding it to (2) with the elimination of 
 3 Hg, we get — * 
 
 3 PbS -I- 8 HNO3 = 1 3 PbSO^ -I- -t^ 8 NO -I- 4 HgO. 
 
 In a neutral solution containing i part of Pb in 100,000 parts 
 of water, HgS will distinctly reveal its presence. HjS is, there- 
 fore, an exceedingly sensitive reagent for the detection of Pb. 
 
 5. Sodium or Potassium Hydroxide precipitates white Pb(0H)2, 
 soluble in excess : — 
 
 Pb(C2H302)2 + 2 NaOH = | Pb(OH)2 -f 2 NaC2H302; 
 j, Pb(0H)2 -I- 2 NaOH = NajPbOa -f- 2 H2O. 
 
 6. Ammonium Hydroxide precipitates a basic salt, insoluble 
 in excess. 
 
 Other difficultly soluble compounds of Pb are : — 
 Pbl2 — yellow; PbBr2 — white; PbS04 — white; basic car- 
 bonate — white, 2 PbCOg • Pb(OH)2 ; PbCgO^ — white. 
 
 * Resolvable into two steps, as in the case of the solution of Ag2S in HNOs. 
 
THE METALS 37 
 
 GROUP I 
 Outline of the Plocess of Analysis 
 
 An examination of tlie foregoing reactions shows that all the 
 metals of this group may be precipitated by HCl. 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. PbCla 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 PbClg. If the 
 quantity of Pb is large, the hot aqueous extract on cooling 
 will deposit the characteristic needles of PbClg, 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 HgSO^ and K2Cr04 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 HgCl^, 
 and the latter is then tested for with SnClg. 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. 
 
38 QUALITATIVE CHEMICAL ANALYSIS 
 
 SCHEME I 
 
 To the clear original solution contained in a small beaker, add 
 dilute HCl 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 treat the 
 solution with HgS in accordance with directions given on page 62. 
 
 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 HCl (3), and finally with a 
 stream of cold water from a wash bottle. Reject the washings. 
 The precipitate may contain PbClg, AgCl, and HgCl. 
 
 Pour through the filter holding the precipitate several small 
 portions of hot water, using about 2 cc. at a time and allowing 
 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 HgSO^ — a white precipitate is 
 PbS04; to the second portion add K2Cr04 — a yellow precipi- 
 tate is PbCrO^ 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 HjSO^. 
 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 1 5-20 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 (8) (15 drops of con- 
 centrated HCl 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 i cc. of water, filter (10), 
 * The numbers in brackets refer to notes. 
 
THE METALS 39 
 
 if necessary, and test the clear solution for Hg by adding a few 
 drops of SnClg. A white precipitate, which may turn gray to 
 black, confirms the presence of Hg. 
 
 NOTES 
 
 1. A white precipitate, when dilute HCI is used, may be due to SbOCl 
 or BiOCl ; the latter, however, dissolves in excess of HCI. 
 
 To insure complete precipitation, HCI 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 HCI, when, if the precipitation was com- 
 plete, no further precipitate will be obtained ; if a precipitate does form, more 
 HCI must be added to the original solution until the test shows complete pre- 
 cipitation. A large excess of HCI is to be avoided on account of the appre- 
 ciable solubility of the chlorides in an excess. 
 
 2. PbCl2 is somewhat soluble in cold HjO, and while the presence of 
 HCI diminishes its solubility, a small amount always remains unprecipitated 
 by HCI and passes into the filtrate, from which it is precipitated by HjS in 
 Group n. One must therefore always look for Pb in the second group. 
 
 3. The precipitate is washed first with HCI, instead of HjO, to prevent the 
 formation of the oxychlorides of Bi and Sb. It is then washed with HjO to re- 
 move the HCI, 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 acidifyin'g 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 
 eflFected 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 (NHjHgCl i- Hg) 
 to the metallic state : — 
 
 2 AgCl -F Hg = 2 Ag -i- HgCli,. 
 
 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 after the treatment with aqua regia as described ,. 
 in note 10. 
 
40 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. (6) 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, (t) 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 CI, 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 HCl + HNOs = NOCl + 2 HjO + 2. CI. 
 
 The excess must be destroyed because its presence would oxidize the SnCl2 
 to SnCl^. As the latter does not react with HgClj, the test would be worthless. 
 10. The solution is diluted because of the possible presence of AgCl, 
 which is appreciably soluble in strong HCl. The residue, after filtration, is 
 tested for Ag by first thoroughly washing it with HjO, dissolving it in ammo- 
 nium hydroxide, and reprecipitating by acidification with HNO3. 
 
 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 41 
 
 1. 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 : — 
 
 HgClj + 2 KOH = |HgO + 2 KCl + HjO. 
 
 2. Ammonium Hydroxide produces in solutions of the mercu- 
 ric chloride a white precipitate of mercuric amido-chloride ; from 
 solutions of the nitrate, a white precipitate of mercuric amido- 
 nitrate : — 
 
 HgCl^ + 2 NH4OH = iNHgHgCl + NH4CI + 2 HjO. 
 
 3. Hydrogen Sulphide, on being passed slowly into a solution 
 of HgCl2, forms at first a white precipitate which changes on 
 further treatment with HjS to a yellow, then brown, and finally 
 to a black precipitate of HgS. These light-colored precipitates 
 are mixtures of HgCla and HgS in varying proportions ; they 
 are soluble in HNO3, and are converted by further treatment 
 with HjS or with (NH4)2S to black HgS : — 
 
 3 HgCla + 2 H^S = KHgCLj • 2 HgS) -h 4 HCl ; 
 (HgCla . 2 HgS)-l- H^S = |3 HgS -f- 2 HCl. 
 
 HgS is insoluble in dilute HCl; 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(N03)2, which is quite insoluble in dilute HNOg. 
 HgS is soluble in aqua regia with the formation of HgClj 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 
 HgClj, precipitates white HgCl, which is reduced by an excess 
 of the reagent to black metallic Hg : — 
 
 SnCl2 + 2 HgCla = | 2 HgCl + SnCli; 
 2 HgCl + SnCl2 = |2 Hg -H SnCl^. 
 
 5. Metallic Cu, Zn, or Fe, when introduced into a solution of 
 
 a mercuric salt, acidified, precipitates metallic Hg. 
 
42 QUALITATIVE CHEMICAL ANALYSIS 
 
 6. Potassium Iodide yields a red precipitate of Hgig, soluble 
 in excess of the reagent or the mercury salt. 
 
 7. If a dry mixture or a mercuric salt and NajCOg 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, arid if the dilu- 
 tion is large, the salt is decomposed with the formation of an 
 insoluble basic salt ; e.£^. — 
 
 BiClg + H2O ^ ,|, BiOCl + 2 HCl. 
 
 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 BigSg : — 
 
 3 H2S + 2 BiClg = IBigSa + 6 HCl. 
 
 The precipitate is insoluble in cold dilute acids, but dissolves in 
 boiling dilute HNO3: — 
 
 (i) BigSg + 6 HNO3 = 2 Bi(N03)3 -F 3 H^S ; 
 (2) 3 H3S + 2 HNO3 = t2 NO -1- 4 H2O + ^3 S. 
 Adding (i) and (2), and eliminating 3 HgS, we get 
 
 BigSa + 8 HNO3 = 2 Bi(N03)3 -1- '^ 2 NO -I- 4 HgO + 4, 3 S. 
 Bismuth sulphide is insoluble in (NH4)2Sa;. 
 
 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 
 iBiONOg, and the sulphate |(BiO)2S04. These are all soluble 
 in dilute inorganic acids and are changed by H2S to BigSg : — 
 
 2 BiOCl -f- 3 H2S = 1 61283 + 2 HCl + 2 HjO. 
 
THE METALS 43 
 
 4. Sodium Stannite in alkaline solution (prepared by adding 
 NaOH to SnClg till the precipitate which first forms dissolves) 
 gives with bismuth solutions a black precipitate of metallic 
 bismuth : — 
 
 (i) SnClj + 2 NaOH = | Sn(0H)2 + 2 NaCl, 
 
 (2) Sn(0H)2 + 2 NaOH = NajSnOj + 2 HjO ; 
 
 Sodium Stannite 
 
 (3) BiClg + 3 NaOH = | Bi(0H)3 + 3 NaCl, 
 
 (4) 2 Bi(0H)3 + 3 Na^SnO^ = j 2 Bi + 3 HjO + 3 NagSnOg. 
 
 5. Metallic Zn or Fe, when added to a solution of a Bi salt, 
 precipitates metallic Bi : — 
 
 2 BiClg + 3 Zn = 3 ZnClj +| 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 are insoluble in s 
 water, but are soluble in halogen acids with the formation of * 
 colorless solutions. The cupric salts, when dissolved 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(0H)2, soluble in a large excess of the strong reagent with 
 the formation of a blue liquid. The precipitate is changed on 
 boiUng 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 = I Cu(0H)2 + (NH4)2S04 ; 
 
 Cu(0H)2 + (NH J2SO4 + 2 NHg = Cu(NH8)4SOi + 2 HgO. 
 
 Deep blue solution 
 
44 QUALITATIVE CHEMICAL ANALYSIS 
 
 When CuClj is used, the blue compound formed with an excess 
 of NHg is Cu(NH3)^Cl2. The sensitiveness of the test is i 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 HjSO^ (distinction from Cd). When freshly pre- 
 cipitated, CuS is easily soluble in KCN solution ; it is also 
 soluble in hot dilute HNOg with the separation of S : — 
 
 3 CuS + 8 HNO3 = 3 Cn(NOs\ + 4 HjO H-f 2 NO + 1 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 cyanide of K and Cu : — 
 
 CuCla + 2 KCN = ,1, Cu(CN)2 + 2 KCl ; 
 2 Cu(CN)2 = ^2 CuCN + (CN)2; 
 CuCN + 3 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)4S04 potassium cyanide is added, the color will be 
 bleached, due to the formation of K3Cu(CN)4: — 
 
 (i) 2Cu(NH3)4S04 + 4KCN=,|,2Cu(CN)2-l-2K2S04+8NH3; 
 
 (2) 2 Cu(CN)2 = ^2 CuCN + (CN)2 ; 
 
 (3) CuCN + 3 KCN = K3Cu(CN)4; 
 
 (4) (CN)2 + 2 NH4OH = NH4CN + NH4CNO + HjO. 
 
 5. Potassium Ferrocyanide precipitates reddish brown cupric 
 f errocyanide : — 
 
 2 CUSO4 + K4Fe(CN)e = | Cu2Fe(CN)6 + 2 K2SO4. 
 
 * Resolvable into 2 steps (see PbS). 
 
THE METALS 45 
 
 It is insoluble in dilute acids, but soluble in NH4OH with the 
 formation of a blue solution. The reaction with K^Pe(CN)g is 
 by far the most sensitive test for Cu ; with very dilute solutions, 
 it gives a reddish brown coloration. The sensitiveness of the 
 tests is I part in 200,000. 
 
 6. Potassium Iodide yields with solutions of cupric salts a 
 yellowish white precipitate of cuprous iodide (Cujij); at the 
 same time iodine is liberated and turns the solution brown : — 
 
 2 CUCI2 + 4 KI = 4- CuJa + 4 KCl + Ij. 
 
 7. a Amino Normal Caproic Acid.* An aqueous solution of 
 this compound when added to a solution of copper not too 
 strongly acid yields a characteristic crystalline precipitate of 
 copper a normal amino caproate. This is one of the most sen- 
 sitive tests for copper and is capable of detecting 0.004 ™&- of 
 copper with certainty. 
 
 8. Certain organic substances, like glucose, reduce copper 
 solutions with the precipitation of red CugO. The test is best 
 carried out by adding to the copper solution KNaC4H^0g 
 (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 CujO is 
 obtained. The reaction consists essentially of 
 
 2 CuO + (reducing agent) = | CugO + O. 
 
 9. Many metals reduce solutions of Cu salts to the metallic 
 state; e.g., Zn, Cd, Al, Fe. If an iron nail is immersed in a 
 solution containing a Cu salt slightly acidified with HCl, a bright 
 deposit of metallic Cu is formed on the iron. As dilute a solution 
 as I 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 HCl. 
 
 Cadmium 
 The cadmium salts are for the most part colorless. The 
 nitrate, chloride, sulphate, bromide, iodide, and acetate are 
 soluble in water. 
 
 *Lyle, Curtman and Marshall, J. A. C. S., 37, (l9iS)> I47l- 
 
46 QUALITATIVE CHEMICAL ANALYSIS 
 
 1. Potassium or Sodium Hydroxide precipitates white 
 Cd(0H)2, insoluble in excess. 
 
 2. Ammonium Hydroxide precipitates white Cd(0H)2, solu- 
 ble in excess, forming complex ammonia salts: — 
 
 CdCla + 2 NH^H = | Cd(OH)2 + 2 NH^Cl; 
 Cd(0H)2 + 2 NH^Cl + 2 NHg = Cd(NH3XCla + 2 H^O. 
 
 This ammonia salt, like the corresponding Cu salt, may be 
 transposed to the double cyanide, K2Cd(CN)4, 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 = ]f Cd(CN)2 + 2 KCl; 
 Cd(CN)2 + 2 KCN = K2Cd(CN)4; 
 K2Cd(CN)4 + H2S = I 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 r61e 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 AsaOg and 
 AS2O5 are respectively the anhydrides of arsenious (HgAsOg) 
 and arsenic (HgAsO^) acids. 
 
TEE METALS 47 
 
 Reactions of the Arsenious Compounds 
 
 The arsenious compounds may be considered as derived from 
 AsgOg. The latter, on treatment with boiling water, yields 
 HgAsOg; the oxide also dissolves in alkalies with the formation 
 of soluble alkali arsenites. All other arsenites are insoluble in 
 water. 
 
 AS2O3 + 6 NaOH = 2 NagAsOg + 3 HjO. 
 
 1. NaOH, NH4OH, NajCOg, HCl, and H2SO4 do not pre- 
 cipitate As from its compounds. 
 
 2. Hydrogen Sulphide. From neutral solutions of arsenites 
 or from aqueous solutions of AsgOg, HgS does not precipitate 
 the sulphide but colors the solution yellow, which is due to the 
 formation of colloidal AsgSg : if, however, the solution is acidified 
 with HCl, a yellow precipitate of As^g immediately forms : — 
 
 2 HgAsOg + 3 HjS = ASgSg + 6 H2O. 
 
 The precipitate is insoluble in concentrated HCl (distinction 
 from Sb and Sn), hence the presence of a large quantity of free 
 acid does not interfere with the precipitation with HjS. Con- 
 centrated HNO3, aqua regia, or a mixture of concentrated HCl 
 and KCIO3 readily oxidizes AsjSg to ars^ic acid (HgAsOg — 
 soluble in water) with the separation of S. Ammonium sulphide 
 and alkali hydroxides both dissolve AsgSg (distinction from the 
 sulphides of Division A). 
 
 3 AsgSg + 10 HNOg + 4 H2O = 6 H3ASO4 -f f 10 NO -H |9 S. 
 
 The action of aqua regia on As2Sg may be represented by the 
 following equations : — 
 
 AS2S3 -t- 10 CI = 2 AsClg -f- 13 S, 
 2 AsClg + 5 H2O = AS2O5 -f- 10 HCl, 
 AsgOg + 3 H2O = 2 HgAsOi. 
 
 Adding these 3 equations, with the ehmination of substances 
 appearing on both sides, we get as the final result : — 
 
 AsgSg + 10 CI -f 8 H2O = 2 HgAsOg + 10 HCl + 1 3 S. 
 
48 QUALITATIVE CHEMICAL ANALYSIS 
 
 The action of KCIO3 + HCl may be represented thus : — 
 KCIO3 + HCl = HCIO3 + KCl ; 
 HCi + HCIO3 = HjO + CIO2 + CI. 
 It is evident that this mixture behaves somewhat like agua regia. 
 The deportment of KOH is as follows : — 
 
 AS2S3 + 6 KOH = K3ASO3 + KgAsSg + 3 HgO. 
 K3ASS3 may be regarded as derived from potassium arsenite 
 (KgAsOg), by replacement of the oxygen by S ; hence it is 
 called potassium thioarsenite. If this mixture of thioarsenite 
 and arsenite is acidified with HCl, AsjSg is reprecipitated : — 
 
 K3ASO3 + K3ASS3 + 6 HCl = 6 KCl + 3 H2O + 1 AS2S3. 
 The equation for the solution of AS2S3 in (NH^)2S is — 
 
 AsgSg + 3 (NH4)2S = 2 (NH4)3AsS3 (ammonium thioarsenite). 
 If (NH4)2S3, is used, we get ammonium thioarsenate : — 
 AS2S3 + 3 (NHJaS, = 2 (NH,)3AsS, + (3 ^ - 5)S. 
 The excess of S in the polysulphide oxidizes the thioarsenite 
 to thioarsenate. The (3;^— S)S does not separate out but dis- 
 solves in the excess of (NH4)2Sa. to form a higher polysulphide. 
 If the solution of AS2S3 in (NH4)2Sa; is acidified with HCl, 
 the As is reprecipitated as yellow AS2S5 : — 
 
 2 (NH J3ASS4 + 6 HCl = I AS2S5 + 6 NH4CI + ^3 H2S ; 
 but as this solution always contains an excess of (NH4)2S^, the 
 following reaction will also take place simultaneously : — 
 (NH J2S. + 2 HCl = 2 NH4CI + tH2S + \{x- i)S. 
 
 3. Silver Nitrate precipitates from neutral solutions of arse- 
 nites yellow silver arsenite, Ag3As03 (distinction from arse- 
 nates) : — 
 
 K3ASO3 + 3 AgNOg = \ AggAsOg + 3 KNO3. 
 The precipitate is easily soluble in acids and alkalies. 
 
 4. Magnesium mixture (solution of MgCl2 + NH4CI + NH3) 
 does not precipitate arsenites (difference from arsenates). 
 
 5. Iodine in solutions rendered alkaline with NaHC03 readily 
 oxidizes arsenites to arsenates : — 
 
 Na3 AsOg +l2 + 2NaHC08= NagAsO^ + 2 Nal + HjO + f 2 COg. 
 
TEE METALS 
 
 49 
 
 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 + HjSO^), the As compound is reduced to gaseous arsine (AsHg), which 
 escapes along with the excess of hydrogen. If the dried gases are led into a 
 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 appa.atus and the reagents employed in the test are arsenic-free. This 
 is accomplished by carrying out the test precisely as described, 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 controlled by a blank, it can- 
 not 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 AgNOj 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 : — 
 
 (1) 6 AgNOa -I- AsHg = \ AggAs • 3 AgNO, -I- 3 HNO3 ; 
 
 Yellow 
 
 (2) AggAs . 3 AgNOj -h 3 HP = HgAsOg + 3 HNO,, + 4r 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 AgNOg, HgCla 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. 
 
so QUALITATIVE CHEMICAL ANALYSIS 
 
 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 ; 
 a blackening proves the presence of arsenic. If the arsenic is in the pentad 
 state, it should first be reduced with SO3 before applying this test. 
 
 4. Bettendorfi's Test depends upon the fact that in a solution strongly acid 
 with hydrochloric acid, stannous chloride reduces arsenic compounds to 
 metallic arsenic : — 
 
 2 AsClg + 3 SnClj = 3 SnCl^ + {r 2 As. 
 
 The test is carried out as follows : To 2 cc. of concentrated HCl in a test 
 tube, add i cc. of strong SnClj 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 HCl 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 AsjOj 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 i part of As in 200,000,000 ; 
 
 Gutzeit test will detect i part of As in 10,000,000 ; 
 
 Bettendorff test will detect i part of As in 7,000,000 ; and 
 Reinsch test will detect i 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 NagCOg, 
 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 i 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 -1- NaHCOg -I- 12 = Nal -h KI + 1 COg + HgAsOi. 
 
TEE METALS 51 
 
 Reactions of Arsenic Compounds 
 
 1. NaOH, KOH, NH^OH or NagCOj produces no precipitate. 
 
 2. Hydrogen Sulphide. From cold, moderately acid solu- 
 tions of arsenic acid, HgS precipitates, after some time, a mix- 
 ture of AS2S3 + S. The HjS first reduces the arsenic acid with 
 the separation of sulphur, and then precipitates AsjSj readily 
 from the reduced arsenic (-ic) solution : — 
 
 (i) H3ASO4+ H^S = H2O + ,|,S -f H3ASO3; 
 (2) 2 H3ASO3 -f 3 HjS = I AS2S3 + 6 H2O. 
 
 If the arsenic solution is heated, the reduction and precipita- 
 tion are hastened. If the amount of HCl is considerable, the 
 stream of HjS rapid, and the solution cold, all of the As will be 
 precipitated as ASgSg : — 
 
 2 H3ASO4 -f S HjS = 8 H2O + I AsjSg. 
 
 When, under the same conditions of acidity, the solution is 
 heated and then treated with a rapid stream of HgS, a mixture 
 of AS2S3 and AsgSg 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, HgS will rapidly precipi- 
 tate the arsenic as AS2S3 : — 
 
 H3ASO4 + H2SO3 = H2SO4 + H3ASO3 1- 
 
 As2Sg has the same solubility as AsgSg; it dissolves in 
 (N 114)23 with the formation of ammonium thioarsenate, from 
 which HCl reprecipitates AS2S5 : — 
 
 AS2S6 + 3(NH,)2S = 2(NH,)3 AsS^ ; 
 
 2(NHJ3 ASS4 + 6 HCl = IAS2S5 + 6 NH^Cl-f t3 H2S. 
 
 * 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. 
 
 t If alkaline earth metals are present, they will be precipitated by the H2SO4 
 formed. 
 
52 QUALITATIVE CHEMICAL ANALYSIS 
 
 3. Silver Nitrate precipitates from strictly neutral solutions 
 chocolate-colored AggAsO^ (distinction from arsenious and phos- 
 phoric acids) : — 
 
 3 AgNOg + NagAsO^ =1 AgsAsO^ + 3 NaNOg. 
 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 NH4MgAs04 (distinction from arsenious acid) : — 
 
 K3ASO4 + MgCla + NH4CI = INH^MgAsO^ + 3 KCI. 
 
 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-I- i2(NH4)2MoP4 + 21 HNOg = 
 
 12 HjO + 21 NH4NO3 + |(NH4)3As04 • 12 M0O3. 
 
 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^eS04, HgSOg, NagSOg, when boiled 
 with a solution of an arsenate strongly acid with HCl, reduce 
 the substance from the arsenic to the arsenious state. 
 
 Any arsenic compound, when treated with strong HCl and 
 distilled in a current of HCl gas, will yield a distillate of AsClg. 
 
 Potassium Iodide, when added to an acid solution of H3ASO4, 
 will reduce it to HgAs03 ^i*'^ ^^^ separation of iodine : — 
 H3ASO4 -f 2 HI = H3ASO3 -I- H2O -f ,1- 12. 
 
 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 AsHg takes place less 
 
TEE METALS 53 
 
 rapidly, due to the necessity of a preliminary reduction of 
 HgAs04 to HgAsOg : — 
 
 H3ASO4 + Hj = H3ASO3 + H2O. 
 With the Fleitmann test, preliminary reduction with HgSOg 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 trivalent 
 element. 
 
 Reactions of Antimonious Compounds 
 
 In carrying out the following reactions a hydrochloric acid 
 solution of SbClg may be employed. 
 
 I. Sodium Hydroxide, Ammonium Hydroxide, and Sodium Car- 
 bonate each precipitates white antimonious hydroxide, Sb(0H)3, 
 insoluble in ammonium hydroxide, but soluble in an excess of the 
 fixed caustic alkalies and in a hot solution of alkali carbonate : — 
 SbClg + 3 NH4OH = I Sb(0H)3 + 3 NH4CI; 
 Sb(OH)3 + 3 NaOH = NagSbOg + 3 H2O. 
 NagSb03 readily hydrolyzes in contact with water, yielding 
 sodium metantimonite, NaSbOj: — 
 
 Na3Sb03 + H2O = 2 NaOH + \ NaSbOg. 
 The latter is further hydrolyzed by water, yielding finally white 
 SbgOg : — 2 NaSbOa + HgO = 2 NaOH + \, Sb^Og. 
 
 SbjOg is practically insoluble in water and in nitric acid, but 
 readily dissolves in hot concentrated HCl with the formation of 
 SbClg. 
 
 Tartaric acid and the tartrates dissolve it in accordance with 
 the following equations : — 
 
 SbjOg + 2 H2C4H4O6 = 2 H(SbO)QH406 + HgO; 
 SbgOg + 2 KHC4H4O6 = 2 K(SbO)QH406 + H2O. 
 
 Tartar emetic 
 
 In both of the resulting soluble compounds of antimony, the 
 group (SbO), called antimony 1, acts as a monovalent radical 
 similar to (NO), nitrosyl, in nitrosyl sulphuric acid, H(N0)S04; 
 
54 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 SbCls containing not too much, 
 free acid a relatively large quantity of water is added, there 
 forms a white precipitate of antimony oxychloride, SbOCl: — 
 
 SbClg + HgO :^ ^ SbOCl + 2 HCl. 
 
 As indicated, the reaction is reversible, too much HCl 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)C4H406 + HCl. 
 
 The precipitate is also soluble in strong HCl and can be changed 
 directly to SbjSg by HjS : — 
 
 2 SbOCl + 3 HjS = I SbjSg + 2 H2O + 2 HCl. 
 
 3. Hydrogen Sulphide. From solutions not too strongly acid 
 with HCl, HgS precipitates red SbgSgi — 
 
 2 SbClg + 3 H2S:it|Sb2S3 + 6 HCl. 
 
 The reversibility of the reaction indicates that a high concentra- 
 tion of HCl would prevent the precipitation, also that the pre- 
 cipitate when formed would dissolve in strong HCl. It has 
 been found by experiment that HCl (i : i) readily dissolves Sb2S8 
 (distinction and method of separation from As). SbjSg is solu- 
 ble in (NH4)2S with the formation of a thio salt: — 
 
 Sb2S3 -f 3 (NH,)2S = 2 (NHJgSbSg. 
 
 AmmoDium thioantimoDite 
 
 If yellow (NH4)2S is used, the excess of sulphur in the latter 
 oxidizes the thioantimonite formed in the last reaction to thiO' 
 antimonate: — 
 
 SbaSa + 3 (NH,)2S. = 2 (NH,)3SbS, + (3^- 5)S. 
 
 The great similarity of the chemistry of the thio salts of anti- 
 mony and arsenic is thus seen. 
 
THE METALS 
 
 55 
 
 If the solution of Sh^% in (NH4)2S;, — that is, the solution 
 containing ammonium thioantimonate and an excess of (NH4)2S2, 
 — is acidified with HCl, 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)2S.:- 
 
 2 (NHJgSbS^ + 6 HCl = |Sb2S5 + 6 NH^Cl + 3 HjS ; 
 (NHJaS, + 2 HCl = 2 NH4CI +tH2S +\{x - i)S. 
 
 The sulphide of antimony is also soluble in alkalies, from solu- 
 tions of which HCl reprecipitates SbjSg. 
 
 4. Zn-Pt Couple. If a solution of antimony, acid with HCl, 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 HCl, is heated with a 
 bright iron nail, all the antimony will be deposited in the form 
 of black flocks (distinction from Sn). 
 
 Antimonic Compounds 
 
 I. Water. If a solution of SbCls, not too strongly acid with 
 HCl, is diluted with a relatively large amount of water, a white 
 precipitate of SbOjCl is formed : — 
 
 SbClg + 2 HjO^I SbOaCl + 4 HCl. 
 
 If the dilution is very great, the SbOgCl is further changed to 
 antimonic acid : — 
 
 SbOaCl + 2 H2O =|H3Sb04 + HCl. 
 
 As with SbOCl, tartaric acid prevents the precipitation of 
 SbOaCl. 
 
56 QUALITATIVE CHEMICAL ANALYSIS 
 
 2. Hydrogen Sulphide. From moderately acid solutions of 
 SbCls, HgS precipitates orange-red SbgSg : — 
 
 2 SbClg + 5 H2S =|Sb2S6 + 10 HCl. 
 The precipitate is soluble in concentrated HCl with the forma- 
 tion of SbClg and evolution of HjS : — 
 
 SbaSg + 6 HCl = 2 SbClg +t 3 HgS +|2 S. 
 
 It possesses the same solubilities as the trisulphide, dissolving 
 in (NH4)2S with the formation of ammonium thiantimonate, 
 (NH4)3SbS4; on acidifying the latter with HCl, the pentasul- 
 phide is reprecipitated : — 
 
 2 (NH4)3SbS4 + 6 HCl =|Sb2S5 + 6 NH4CI -t-f 3 H2S. 
 
 Sb2S5 also dissolves in caustic alkalies. 
 
 3. Potassium Iodide, when added to a HCl solution of SbClg, 
 reduces it with the separation of iodine [distinction from Sb 
 (-ous) compounds] : — 
 
 SbClg + 2 KI = 2 SbClg + 2 KCl +,|,l2. 
 
 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 (SbHg) 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). 
 
 SbHg is further distinguished from AsHg by the formation of black AgjSb 
 when the former is passed into a solution of silver nitrate : — 
 
 SbHg + 3 AgNOg = 4, SbAgg + 3 HNOg. 
 
 With AsHg the black deposit is due to metallic silver (see p. 49). 
 
 2. Gutzeit Test. Same as for As, the blackening being due to AggSb. 
 
 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 SbjOg. The latter is insoluble in water but is soluble in a 
 hot solution of KHC4H4O8, from which HjS precipitates red SbjSg. 
 
 Neither Fleitmann's nor Bettendorff's test are applicable to antimony 
 compounds. 
 
 Tin 
 
 The two oxides of tin, SnO and SnOg, correspond to two 
 classes of salts known respectively as the stannous and stannic 
 
THE METALS 57 
 
 compounds. In the former, tin is divalent ; in the latter tetra- 
 valent. 
 
 Reactions of Stannous Salts 
 
 The stannous salts are all colorless. Those which are solu- 
 ble in water yield solutions which have an acid reaction ; the 
 sohd salts as well as their solutions rapidly absorb oxygen from 
 the air with the formation of stannic compounds. 
 
 1. Sodium Hydroxide, Ammonium Hydroxide, or Sodium Car- 
 bonate gives a white precipitate of Sn(0H)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. 
 
 SnClj + 2 NaOH = ,|,Sn(0H)2 + 2 NaCl; 
 Sn(0H)2 + 2 NaOH = NagSnOa + 2 HjO. 
 
 The precipitate also dissolves in HCl. It possesses, like 
 A1(0H)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 2.5 per cent, of concentrated HCl) HgS 
 throws down a brown precipitate of SnS : — 
 
 SnCla + HjS = j,SnS -f 2 HCl. 
 
 SnS is soluble in strong HCl (distinction from As), nearly in- 
 soluble in colorless (NHJjS [distinction from the sulphides of 
 As, Sb, and Sn (-ic)], but is soluble in hot (NH4)2Sa; with the 
 formation of ammonium thiostannate, from which HCl precipi- 
 tates yellow stannic sulphide, SnSa : — 
 
 SnS + (NH4)2S^ = (NH4)2SnS3 + (x- 2)S ; 
 
 Brown 
 
 (NH ASnSg -F 2 HCl = 4. SnSg -f 2 NH^Cl -f \ HgS. 
 
 < Yellow 
 
 3. Mercuric Chloride, when added in excess to a solution of 
 SnClg, is reduced to white insoluble HgCl ; the SnClg is oxidized 
 at the same time to SnCl^ : — 
 
 2 HgCLj -I- SnClg = 4-2 HgCl + SnCl^. 
 
58 QUALITATIVE CHEMICAL ANALYSIS 
 
 If, however, the SnCIg is present in excess, the HgCl first 
 formed is further reduced to metallic mercury (gray or black) : — 
 
 2 HgCl + SnCla = | 2 Hg + SnCl^. 
 
 Black 
 
 As this reaction is essentially one of reduction, it is important 
 that the HgClj 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 SnClj or SnCl^, metallic 
 tin is precipitated on the zinc in the form of a gray spongy mass. 
 As the deposited tin is readily soluble in strong HCl, care must 
 be taken not to have the solution too strongly acid. 
 
 SnCl4 + 2 Zn = 2 ZnClj + | Sn ; 
 SnCl2 -I- Zn = ZnCla + >l-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 SnCl^ for the 
 following tests may be prepared by warming a rather strong 
 HCl solution of SnCl^, with KCIO3 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^ 
 
 I. Sodium Hydroxide, Ammonium Hydroxide, as well as Sodium 
 Carbonate, yield a white precipitate of stannic hydroxide, 
 Sn(OH)4, which, on drying, becomes HjSnOg. The precipitate 
 is soluble in excess of NaOH or NagCOs with the formation of 
 sodium stannate of variable composition. 
 
THE METALS 59 
 
 2. Hydrogen Sulphide. From moderately acid solutions not 
 exceeding 2.5 percent, concentrated HCl, HjS precipitates yellow 
 
 SnSg: — 
 
 SnCl^ + 2 H2S = ,|,SnS2 + 4 HCl. 
 
 SnSg is readily soluble in HCl (1:1), hence the necessity for 
 having the solution not too strongly acid. The disulphide of 
 tin readily dissolves in colorless (NH^)2S (distinction from SnS) 
 with the formation of ammonium thiostannate, from which 
 HCl precipitates SnSj (yellow): — 
 
 SnS^ + (NH,)2S = (NH,)2SnS3; 
 (NHJjSnSg + 2 HCl = 2 NH^Cl + ^HjS + ^SnSj. 
 
 On strong ignition in the air, SnSj is quantitatively converted 
 into SnOg. 
 
 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(0H)4 is 
 
 obtained: — 
 
 SnCl4 + 4 H2O 1^1 Sn(0H)4 -|- 4 HCl. 
 
 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(0H)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 : — 
 
 (i) SnCl4-|-2 H2S04=Sn(S04)2 + 4 HCl; 
 
 (2) Sn(S04)2 -F 4 H2O = 2 H2SO4 + |Sn(0H)4. 
 
 Similarly, with K2SO4 = we get 
 
 SnCl^ + 2 K2SO4 = Sn(S04)2 + 4 KCl, 
 the Sn(S04)2 being then hydrolyzed as shown above in 
 equation (2). 
 
6o 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)^, i.e., 
 HjSnOg, but differs from it in being insoluble in acids. When boiled for a 
 short time with concentrated HCl, a compound of the formula Sn505Cl2(OH)3 
 forms which, though quite insoluble in HCl, is readily soluble in water. From 
 the fact that this and similar compounds may be formed from metastannic 
 acid, the formula Sn505(OH)jj or 5 (HjSnOj) has been assigned to it. 
 Stannic hydroxide, when dried over concentrated H2SO4, has the formula 
 HjSnOj ; metastannic acid is thus seen to be a polymer of stannic hydroxide. 
 
 Reactions of Metastannic Chloride, Sn^O^Cl^OH\ 
 
 1. HCl precipitates SngOgCliCOH)^, 4H2O. 
 
 2. Prolonged boiling with water causes the precipitation of 
 all the tin as metastannic acid, insoluble in dilute acids. 
 
 3. H2SO4, K2SO4, or NagSO^ 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. HgS yields the same precipitate as with SnCl^ solutions. 
 Solutions of stannic compounds are converted into the meta- 
 stannic form by diluting and boiling : — 
 
 5 SnCl^ + 13 H^O:^ 18 HCl + Sn50sCl2(OH)8. 
 
 Conversely, metastannic compounds are converted into the 
 
 stannic form by boiling with concentrated HCl or concentrated 
 
 KOH. 
 
 Action of H2S 
 
 Besides its action ag a group reagent for precipitating the metals of Group 
 II., HjS also acts as a reducing agent. Should, therefore, an oxidizing agent 
 be present in the solution subjected to the action of HjS, it will be reduced, 
 the HjS being oxidized at the same time to elementary sulphur, which sepa- 
 
THE METALS 6i 
 
 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 FeCl,, K^CrjO,, KMnO^, 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 FeClj + HgS = 2 FeClj + 2 HCl + | S ; a 
 Y.^rj^^ solution changes by the action of H2S from reddish orange to green 
 — KjCrjO, + 3 HjS + 8 HCl = 2 CrCls + 2 KCl + 7 H^O + ,|, 3 S ; while a so- 
 lution containing KMnOi changes from a solution with purple tint to one that 
 is colorless : — 
 
 2 KMnO^ + 5 HjS + 6 HCl = 2 MnClg + 2 KCl + 8 HjO + ■Jr 5 S. 
 
 Should the concentration of nitric acid be large, it will oxidize the H^S 
 with separation of sulphur and partly with the formation of H2SO4 : — 
 
 2HN03 + 3H2S = 4H20 + |2NO + |3S; 
 3 HgS + 8 HNO3 = 3 H2SO4 H-'h 8 NO + 4 HjO. 
 
 Aqtia regia will oxidize the H^S in accordance with the equation — 
 
 Cl2+H2S=:2HC1 + >1.S. 
 
 It will be observed that in every case the presence of an oxidizing agent 
 causes the decomposition of the HjS with the separation of S ; with a large 
 amount of oxidizing agent present, the amount of S and HgSO^ 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 H^SO^ will have the effect of precipitating the alkaline 
 earths along with the metals of the second group, so that where the amount 
 of oxidizing agent is large it is advisable to eliminate the latter before pass- 
 ing in HjS. If, from the purple or orange-red color of the solution, KMnO^ 
 or KjCrjOy is suspected, reduction may be readily effected by acidifying with 
 HCl, 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 re^a, be- 
 cause of their extensive use as solvents. 
 
 An excess of HNO3 is removed by evaporating the solution to about i cc, 
 adding 3 cc. of concentrated HCl, and evaporating nearly to dryness. It can 
 then be taken up with the aid of HCl and hot water. 
 
 An excess of aqua regia is disposed of by boiling the solution down to a 
 small bulk, adding concentrated HCl and again evaporating to i cc. ; it is 
 then diluted with water and a few drops of HCl. 
 
 H2S as a Precipitating Agent. In precipitating the second group sulphides 
 with HjS, 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 HCl (notably those of Pb, Cd, 
 
62 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 HCl 
 (sp. gr. 1.2) in a volume of 100 cc. affords a satisfactOiW acidity for the 
 separation of the second and third groups by HjS. ' ■ 
 
 Precipitation with H2S 
 
 The filtrate (i) from Group I. is made j/z^^^/j/ alkaline with 
 NH4OH and then j'usi acid with dilute HCl (2); 2.5 cc. of 
 concentrated HCl are then added, the solution is heated nearly 
 to boiling and is treated with a rapid stream of HjS 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 
 HjS again until precipitation is complete.A^'^Filter, dilute the 
 filtrate with ^ its volume of water, and treat again with HjS ; 
 now filter off any precipitate formed (4). The final filtrate 
 should not give a precipitate when treated with HgS. The^'j 
 beaker containing the filtrate should be labeled Groups IH.-V., 
 at once placed on a wire gauze, and boiled until all the HjS is 
 expelled (5). 
 
 The precipitate may consist of HgS, PbS, BigSg, CuS, CdS, 
 AsgSg, SbgSg, SnS and S ; it should be washed with water con- 
 taining HgS 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 H^S in accord- 
 ance with directions ; if absent, the introduction and removal of HjS 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 mixture determined by me^ns of litmus paper and not by the quantity of 
 the reagent added. No attenlBon need be given to precipitates which .form, 
 because these are either finally dissolved or converted hy HjS 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 63 
 
 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 (NHJgSj; cannot be anything other than Cd ; on the 
 other hand, if the yellow precipitate dissolves completely in (NH4)2Sj., it must 
 be either the sulphide of As or Sn(-ic), or both. 
 
 5. As HjS 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 HjS is apparent. 
 
 Keeping a glass rod in the beaker during the boiling will facilitate the 
 removal of HoS by preventing dangerous bumping, with a consequent loss of 
 liquid. The completeness of the expulsion of the HjS may be determined by 
 holding a piece of filter paper moistened with lead acetate in the escaping vapor. 
 
 6. The precipitate is washed with HoS water to prevent the oxidation 
 of the sulphides to sulphates, which, i^ith 4he exception of PbSOj, are all 
 soluble in water. NH^NOj is added W3 present the precipitate from going 
 into the colloidal state and then passing through the filter. 
 
 The separation of Group II. into two divisions is based on 
 the difference of behavior of the sulphides towards (NH4)2Sa;. 
 
 Division II. A (the copper group) includes those sulphides 
 which are insoluble in (NH4)2S:,; these are Hg, Pb, Bi, Cu, and 
 Cd. 
 
 Division II. B (the tin group) includes those sulphides which 
 are soluble in (NH4)2S:, with the formation of thio salts. The 
 separation, however, is not sharp, which is due to the shght 
 solubility of CuS in (NH4)2S^. CuS is practically insoluble 
 in Na2S„ but the latter quite appreciably dissolves HgS. For 
 all practical purposes, the separation with (NHJgSj; is suffi- 
 ciently complete. 
 
 In the analysis of the HgS 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)2Sj„ as directed in 
 paragraph 3. In most cases, however, it is preferable to make 
 the following preliminary tests. 
 
64 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 (NH^)2Sj; diluted with an 
 equal quantity of water. Break up the precipitate with a glass 
 rod, and warm gently with constant stirring for a minute (</(? 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 (i), the copper group 
 is present. 
 
 (2) Division B (tin group). Twice pass about 10 drops of 
 (NH^)2Si, 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 HCl. In 
 another test tube,7«J/ acidify^ with dilute HCl an equal portion 
 of (NH4)2S2., diluted as abovf, 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 samfe'''appearance after acidification, the 
 absence of the Sn group is proved. 
 
 (3). 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)2Sj. diluted with 5 cc. of water ; 
 the mixture is thoroughly stirred, warmed for several minutes, 
 and filtered (5). The residue on the filter may consist of HgS, 
 PbS, BigSj, CuS, CdS and S. The filtrate may contain the thi6 
 salts of As, Sb, and Sn. A warm mixture of 2 cc. of (NH4)2Sj, 
 and an equal volume of water is poured on the residue (6), 
 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 containing 5 per cent. NH4NO3 and is then analysed 
 according to Scheme II. A. The washings are rejected. 
 
TEE METALS 
 
 65 
 
 I 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 (NHJgSx) with the separation of S. If considerable acid is 
 present, all of the (NH4)2Sj: will be decomposed, with the result that the 
 liquid which drains through the filter will be colorless and will fail to react 
 with HCl. To remedy this, add more (NHJjSx? preferably diluted with 
 NH4OH, until a colored filtrate is obtained. 
 
 3. Because of the slight solubility of CuS in (NH4)2S3., a liver-colored pre- 
 cipitate is sometimes obtained in acidifying the (NH^)2Sa. solution. 
 
 4. If the quantity of the precipitate is small, it may be treated on the filter 
 with (NHJgSx, 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 ;:he precipitate to a beaker, it must be carefully poured off before 
 adding the (NHJ„S,. 
 
 5. If the precipitate shows a tendency to pass through the filter in, the 
 colloidal condition, add several grams of NH4NO3, warm, stir, and refilter. 
 
 6. This fiirther treatment of the residue with dilute (NH^)2Sj, is given to 
 insure the complete extraction of the As, Sb, and Sn 
 
 Outline of the Separation of the Metals of Group 11. 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 HgSO^, we should expect 
 the Pb to precipitate as PbSO^; for Pb is the onl}^ 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 wat^ and the 
 PbS04 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 
 
66 
 
 QUALITATIVE CHEMICAL ANALYSIS 
 
 hydroxide is insoluble in excess. If, therefore, an excess of 
 NH4OH is added to the filtrate from the PbSO^, 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 HjS. The double cyanides are formed by adding KCN 
 to the ammoniacal solution until the blue color is discharged ; 
 the passage of HjS into this solution precipitates only the Cd as 
 yellow sulphide. 
 
 SCHEME II. A 
 
 The residue from the (NHJjSx treatment may consist of HgS, PbS, BijSg, 
 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 (i) ; boil for 
 one and a half minutes, and filter. 
 
 Residue may be 
 HgS (black) or 
 2 HgS • Hg(N03), 
 (white) + S (2). 
 Wash with water. 
 Transfer ppt. to a 
 small evap. dish ; 
 add 4CC. ag'zia re- 
 gia. pcuA boil till 
 all but S dissolves. 
 Boil a little longer 
 ■to expel CI, dilute 
 with I cc. of water, 
 and filter. To fil- 
 trate add a few 
 drops of SnClj ; a 
 white ppt., turn- 
 ing gray to black, 
 proves the pres- 
 ence of Hg. 
 
 Filtrate may contain Pb(N0g)2, Bi(NOg)g, Cu(N03)2. 
 Cd(N03)2 4- excess of HNO3. Transfer to an evap. dish, 
 add 5 cc. of cone. H2SO4, and ^aporate 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 
 PbS04(4). Wash 
 once with water, 
 and treat the ppt. 
 on the filter with 
 10 cc. of a boil- 
 ing solution of 
 
 NH.QHgO^. 
 Catch filtrate in a 
 test tube, add 2 cc. 
 acetic acid and i 
 cc. KjCrO^. A 
 yellow ppt. con- 
 firms the presence 
 ofPb. 
 
 Filtrate contains 312(804)3, CuSO^, 
 CdSO^ 4- excess of H2SO4. AddNH^OH 
 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(0H)3 (6). 
 Wash with water; 
 dissolve ppt. on the 
 filter with about 
 5 drops of cone. 
 HCl, ^nd 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). 
 
 EUltrate contains 
 Ci:(NH3)4S04, 
 Cd(NH3)4S04 -I- 
 excess of NH4OH. 
 Add KCN until the 
 solution is decolor- 
 ized, and pass in 
 H2S. A yellow ppt. 
 is CdS (8). 
 
THE METALS 67 
 
 NOTES ON SCHEME H. A 
 
 1. Stirring while heating with dilute HNO3 is important because of the 
 tendency of the S, which separates in a plastic coiidition, 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 
 PbS04 resulting from the oxidation of PbS by the use of strong HNO3 or the 
 long-continued action of the dilute acid. But, as PbSO^ is somewhat soluble 
 in hot dilute HNOg, 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 BigSj. Consequently the residue after the HNOg 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 PbSO^. As the 
 boiling point of concentrated HNO3 is 120° C. and the fuming point of HjSO^ 
 is 250° 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 ta recognize these fiimes with certainty, 
 he should show his results to his instructdr 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. PbS04 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 PbSO^ is very 
 small, it can be made distinctly visible by collecting it in the center by meaqi|^ 
 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)g; 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 wOl 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 : BiClg -f HjO^BiOCl -I- 2 HCl. 
 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. 
 
68 QUALITATIVE CHEMICAL ANALYSIS 
 
 Another confirmatory test for Bi consists in pouring on the thoroughly 
 washed white precipitate of Bi(OH)3 a solution of Na2Sn02 (prepared by add- 
 ing NaOH to I cc. of SnClj 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 H^S is passed into the solution. A little Hg may 
 find its way into the solution if the original H^S precipitate had not been 
 thoroughly washed free from HCl or chlorides ; the latter, on boiling with 
 HNO3, will yield CK^ua regm, 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 HoS, 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 (i pt. of cone, acid to 4 pts. of 
 water) on the filter. The dilute HoSO^ dissolves the CdS and leaves on the 
 filter the HgS and PbS ; the latter may be entirely changed to PbSO^. If the 
 filtrate is now largely diluted with water and treated with HjS, a yellow pre- 
 cipitate of CdS will be formed if Cd is present. 
 
 Outline of the Analysis of Group II. B (the Tin Group) 
 
 The insolubility of AsgSg in hot concentrated HCl 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 HgS. 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 AgNOg in a neutral solution as 
 Ag3As04. When the AsgSg is removed by filtration, the filtrate 
 will contain the Sb and Sn as chlorides and an excess of HjS. 
 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 HgS, 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 SnCl^ to SnClg ancj at the same time prje;. 
 cipitates the Sb in the metallic state. The filtered solution, '"»=*^ 
 containing SnClg, is then tested with HgClg solution. 
 
 The second portion is tested for Sb by the galvanic action of 
 the electric couple (Pt and Zn), which causes the antimony to 
 precipitate as a black deposit on the platinum. 
 
THE METALS 
 
 69 
 
 SCHEME II. B. (TIN GROUP) 
 
 The filtrate obtained after treating the 
 
 HjS precipitate with 
 (NH4)2Sj. will contain the thio salts of the metals of this group ; 
 viz.;(NHj3AsS4,(NH4)3SbSi,(NHj2SnS3 and excess (NHj^Sx- 
 Jtist acidify with dilute HCl (i) and filter; reject the filtrate. 
 
 The residue will consist of AsgSg, 
 
 SbgSg, SnSg, (CuS .''), and S. 
 
 Transfer the precipitate to a small beaker, add 10 cc. of con- 
 centrated HCl, and heat gently to boihng (2), with constant stir- 
 ring, for about 5 minutes. Dilute with an equal volume of water 
 and filter. 
 
 Residue is ASjS, + S (3). Wash with hot 
 water unti! the washins:-; 'i''ttT boi'ing L > expel 
 r , give only a, faint reaction with AgNO,. 
 Reject washings. Transfer ppt. to an evap. 
 dish (if water is used in transferring ppt., 
 pour it off after ppt. settles). Add 3-5 cc. 
 of cone. IINO3, stir, and heat gently until no 
 more brown fumes are given off, and until the 
 excess HNO.o is expelled. Dilute with 3 cc. 
 of water and filter through a small filter. 
 Filter wash with i cc. of water, catching 
 filtrate and washings in a test tube. Add 
 5 cc. of AgN03 (4) filter if a precipitate 
 forms. To the clear solution or filtrate add 
 one drop of phenolphthalein and then render 
 Just alkaline with NH^OH. Now add 5 % 
 acetic acid drop by drop with shaking until 
 the resulting mixture is faintly acid. A choco- 
 late-colored ppt. of AgjAsO^ confirms the 
 presence of arsenic (5). 
 
 Filtrate may contain SbClg, 
 SnCl^ + excess HCl + H.,S. 
 Boil until the H„S is com- 
 pletely expelled (6). Divide 
 the solution into 2 portions. 
 In the 1st portion, test for Sn 
 by -ivarming with an iron nail 
 for 3 minutes. Filter rapidly 
 (7) into a test tube containing 
 2-3 drops of HgCU; 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 
 (8), which is insoluble in 
 NaBrO, is Sb. 
 
 WOTES TO SCHEME H. B 
 
 1. An excess of acid is to be avoided because of the solubility of SnS„ in 
 even moderately dilute acid. 
 
 2. The mixture must not be boiled, else some AsoS^ is apt to go into 
 solution. 
 
 3. The separation of As from Sb and Sn sulphides by the use of hot con- 
 centrated HCl is not always very sharp. With mixtures consisting of a small 
 
70 QUALITATIVE CHEMICAL ANALYSIS 
 
 amount of As and a relatively large quantity of Sb, the insoluble residue may 
 contain enough Sb2Sj to give it a red color ; therefore a red residue must always 
 be examined for As. 
 
 4. If all the HCl had not been removed by washing, the addition of AgNOg 
 will yield a white precipitate of AgCl ; if this precipitate is large in amount, 
 more AgNOj should be added before filtering to insure an excess of the latter 
 in the filtrate. 
 
 5. The test for arsenic with AgNOg depends upon the formation of 
 AggAsOj, which only forms in a strictly neutral solution ; if too much acetic 
 acid is added, precipitation will fail, because of the ready solubility of Ag5As04 
 in acids as well as in alkalies. To remedy this, carefully neutralize the excess 
 of acid by the addition of dilute NH4OH. 
 
 6. Test escaping vapors with lead acetate paper. ULwv Ut.*-*.^ oLufW-O- 
 
 7. By this procedure the SnCl4 is reduced to SnClj ; 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. 
 
 8. If much Cu and little or no Sb are present, a dark red stain, easily dis- 
 tinguished from a black stain, will be promiced. 
 
 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 H^gS 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 (NH 4^)28 in solutions alkaline with ammonium hydroxide in 
 the presence of NH4CI (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. 
 
 I. Ammonium Hydroxide throws down a white gelatinous 
 precipitate of A1(0H)3, slightly soluble in excess; on boiling, 
 the dissolved hydroxide is reprecipitated. The presence of 
 
THE METALS 71 
 
 NH4CI 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, Al(OH)3 is readily soluble in 
 acids : — 
 
 AICI3+ 3 NH4OH = |A1(0H)3 + 3 NH4CI; 
 A1(0H)3 + 3 HCl = AICI3 + 3 HgO. 
 
 Aluminum hydroxide is also soluble in caustic alkalies (see 2). 
 
 2. Potassium or Sodium Hydroxide precipitates A1(0H)3, 
 soluble in excess with the formation of alkali aluminate : — 
 
 A1(0H)3 + 3 NaOH = NagAlOg + 3 VLfi. 
 
 On carefully neutralizing the alkaline solution of sodium 
 aluminate with hydrochloric acid, A1(0H)3 is reprecipitated : — 
 
 (a) Na3A103 + 3 HCl = 3 NaCl + 1 A1(0H)3. 
 
 If an excess of acid is added, the precipitate which first JEorms 
 is dissolved and we obtain : — 
 
 {b) Al(OH)3 + 3 HCl = AICI3 + 3 H2O. 
 
 It is evident that if the original sodium aluminate solution is 
 at once acidified with HCl, we shall get the net result of {a) and 
 {b). Adding {a) and {b), and eUminating A1(0H)3, which ap- 
 pears on opposite sides, we get : — 
 
 Na3A103 -f 6 HCl = AICI3 + 3 NaCl -f- 3 HgO. 
 
 If, now, we heat this solution to boiling and add ammonium 
 hydroxide in faint excess, all the aluminum will be precipitated 
 as the hydroxide : — 
 
 AICI3 -I- 3 NH4OH = I A1(0H)3 -I- 3 NH4CI. 
 
 While the addition of a large excess of solid NH^Cl will have 
 the effect of preqipitating A1(0H)3 from the aluminate, the 
 above process, viz., that of acidifying first with HCl and then 
 rendering the resulting solution faintly alkaline with ammonium 
 
72 QUALITATIVE CHEMICAL ANALYSIS 
 
 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(0H)3. 
 
 3. Ammonium Sulphide precipitates Al(OH)3 and not the 
 sulphide. AljSg 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 HjS. The action 
 of (NH4)2S on solutions of aluminum salts may be represented 
 as taking place in two steps : — 
 
 (a) 2 AICI3 -j- 3(NH4)2S = AI2S3 + 6 NH^Cl; 
 (^) AI2S3 + 6 H2O = 1 2 A1(0H)3 + 3 H^S. 
 
 Adding (a) and (i>), and eliminating AlgSg, we get as the equa- 
 tion for the final, result: — 
 
 2 AICI3 + si^n^^S + 6 H2O = |2 Al(OH)3 -F 6NH1CI + 3 HjS. 
 
 4. Sodium Carbonate also precipitates Al(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 : — 
 
 AICI3 + 3 NaCaHgOa = A1(C2H302)3 + 3 NaCl; 
 A1(C2H302)3 + H2O :;!:| Al(OHXC2H302)2 + HC2H3O2 
 
 The reagent is, in fact, hot water, which hydrolyzes ^he weak 
 salt Al( €211302)3. Cooling the solution, or the presence of an 
 excess of acetic acid, will have the effect of reversing the re- 
 action, with the result that some of the precipitate will dissolve. 
 
THE METALS 73 
 
 6. Disodium Phosphate yields with solutions of aluminum salts 
 a gelatinous precipitate of AIPO4, soluble in HCl and NaOH, 
 but insoluble in acetic acid : — 
 
 ( 1 ) NaaHPO^ + AICI3 = | AIPO4 + 2 NaCl + HCl ; 
 
 (2) Na2HP04 + HCl = NaHgPO^ + NaCl. 
 
 Adding (i) and (2), and eliminating HCl, we get : — 
 
 (3) 2 NajHPOi + AICI3 = ,], AIPO4 + NaHaPOi + 3 NaCl ; 
 {a) AIPO4 + 3 NaOH = NagAlOg + H3PO4, 
 
 {b) H3PO4 + 3 NaOH = NagPO^ + 3H2O. 
 
 Adding {a) and (b), and eliminating H3PO4, we get : — 
 AIPO4 + 6 NaOH = Na3A103 + Na3P04 + 3 H2O. 
 
 7. Any aluminum compound, when strongly ignited in the 
 air, is converted into AI2O3; if this is moistened with a very 
 dilute solution of Co(N03)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 CrgOs and Cr03. 
 The former is basic and forms the various chromic * salts by 
 combining with acids, e.g., Cr^O^ + 6 HCl = 2 CrClg + 3 HgO ; 
 similarly by solution of CrjOg in H2SO4 and HNO3, Cr2(S04)3 
 and Cr(N03)3 are respectively formed. In all these com- 
 pounds chromium deports itself as a metal. CrOg, on the other 
 hand, is distinctly acid in character, being the anhydride of the 
 hypothetical chromic acid, HgCrO^, the salts of which are 
 known as chromates. The latter may be prepared by treating 
 CrOg with a base ; thus, sodium chromate (NagCrO^) may be 
 prepared by treating CrOg with caustic soda : Cr03+ 2 NaOH 
 = Na2Cr04 + HjO. 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 Cr™ chromic, etc. 
 
74 QUALITATIVE CHEMICAL ANALYSIS 
 
 part of the aczd radical CrO^, 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 
 CrClg, causes a precipitate of chromium hydroxide to form; 
 when added, however, to a chromate, as Na2Cr04, 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 reagents are added to sodium chromate, yellow pre- 
 cipitates are formed, due to the formation of BaCrO^ and 
 PbCrO^, 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 (Cr"") are converted into chromates (Cr^') by 
 oxidation in an alkaline solution ; conversely, chromates (Cr^^) 
 are reduced to chromic salts (Cr°^) 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 
 
 Cr203+3 = 2Cr03, 
 
 while the reduction of chromate to a chromic salt may be repre- 
 sented by 
 
 2 CrOg -I- 6 H = CrgOg -I- 3 HjO. 
 
 (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 NagCOg (which supplies the alkali) 
 and an oxidizing agent like NagOg, KClOg, or KNOg (which 
 
 * As all chromic salts may be derived from Cr203 by treatment of the latter with 
 the appropriate acid, and as the valence of Cr is the same in this oxide and its salts 
 vfe may conveniently represent all chromic salts in oxidation equations by Cr20s. 
 For a similar reason, all chromates may be represented in reduction reactions by 
 CrOs. 
 
TEE METALS 75 
 
 supplies the O). In its simplest form the equation for the 
 oxidation is — 
 
 CrjOg + 2 NagCOj + 30 = 2 Na^CrO^ + f 2 COa- 
 
 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), YLMnO^, or HjOg. The alkali first pre- 
 cipitates chromic hydroxide : — 
 
 CrClg + 3 NaOH = j Cr(0H)3 + 3 NaCl. 
 
 This then dissolves in excess, giving sodium chromite : — 
 
 Cr(0H)3 + 3 NaOH = NagCrOg + 3 HjO. 
 
 The chromite is then oxidized by the oxidizing agent to chro- 
 mate : — ■ 
 
 2 NagCrOg + 3 O + HjO = 2 NagCrO^ + 2 NaOH. 
 
 Sod, Chromite Sod. Chromate 
 
 When sodium dioxide (NagOj) 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 
 
 NagOg + H2O = 2 NaOH + O. 
 
 An excess of NagOg 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 NagCrOg + 3 NaaOj + 4 HjO = 8 NaOH + 2 NagCrO^. 
 
 The oxidation by means of sodium dioxide is to be preferred 
 to the other agents for the conversion of chromic compounds 
 to chromate. In evety 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 KjCrO^ + 2 HNOg .= KjCraO^ + 2 KNOg + HjO. 
 
76 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 : — 
 
 KaCrjOT + 2 KOH = 2 K2Cr04 + HjO. 
 
 Both chromates and dichromates, when in solution, may be 
 precipitated by solutions of Ba or Pb salts [distinction from 
 chromic (Cr™) compounds] : — 
 
 KgCrO^ + Pb(C2H302)2 = I PbCrO^ + 2 KQHgOa ! 
 
 Tello-w 
 
 KgCrjO; + 2 BaClg + Hfi = | 2 BaCrO^ + 2 KCl + 2 HCl. 
 
 yellow 
 
 (l>) Reduction of chromates to chromic compounds is effected in 
 acid solutions by any one of the many reducing agents, e.g., 
 HgS, HI, SO2, concentrated HCl, and various organic sub- 
 stances, as alcohol and oxalic acid. 
 
 With concentrated HCl the reaction is — 
 
 KaCrgO^ + 14 HCl = 2 KCl + 2 CrClg + 7 HgO + f 3 Clj. 
 
 With HjS the reduction takes place in accordance with the 
 equation 
 
 K2Cr207 + 3 HjS + 8 HCl = 2 CrClg + 2 KCl + 7 HgO + ,|, 3 S. 
 
 In this case the green solution appears turbid from the separa- 
 tion of S. 
 With sulphurous acid the equation is — 
 
 KgCraO^ + H2SO4 + 3 H2SO3 = Cr2(S04)3 -I- KjSO^ -I- 4 HgO. 
 
 In all the above cases the reduction is evidenced by a change in 
 color from orange-red to green. 
 
 Other reaction for chromates will be given in Part II., dealing 
 with the acids (see page 142). 
 
 TAe 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 
 
(f Cv, ,T THE METALS 
 
 77 
 
 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(0H)3, soluble with difficulty in 
 excess with the formation of a violet 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(0H)3, solu- 
 ble in excess in the cold to a green solution with the formation 
 of sodium chromite ; on boiling this solution, Cr(OH)g is repre- 
 cipitated (distinction from Al). The precipitate is easily solu- 
 ble in acids. 
 
 CrClg -F 3 NaOH = j, Cr(0H)3 + 3 NaCl ; 
 
 Cr(0H)3 -I- 3 NaOH = NagCrOg -I- 3 HgO ; 
 NagCrOg -|- 3 HgO (boiling) = | Cr(0H)3 -|- 3 NaOH ; 
 Cr(OH)3 -t- 3 HCl = CrClg H- 3 HjO. 
 
 3. Ammonium Sulphide precipitates Cr(0H)3, for CrjSg, like 
 AljSg, is hydrolyzed by water with the formation of Cr(OH)g 
 and the evolution of HgS : — 
 
 2 CrClg -I- 3(NH4).^S -f 6 HjO = 
 
 ]f 2 Cr(0H)3 -f- 1 3 HjS -F 6 NH^Cl. 
 
 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 i, 2, 3, 4, and 5. 
 
 5. Disodium Phosphate precipitates from solutions of chromic 
 choride green CrP04 : — 
 
 2 NagHPO^ + CrClg = 3 NaCl -I- NaH2P04 + ICrPO^. 
 
78 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. 
 
 6. 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 : — 
 
 (i) 3Na^02 + 3H20 = 6NaOH + 30; 
 
 (2) CrClg + 6 NaOH = NagCrOg + 3 NaCl + 3 HgO ; 
 
 (3) 2 NagCrOg + 3 O + H^O = 2 NagCrO^ + 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- 
 
THE METALS 79 
 
 less. Ferrous salts in solution are very unstable, for they rapidly 
 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 
 
 1. Ammonium, Sodium, or Potassium Hydroxide precipitates 
 at first white gelatinous Fe(0H)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. 
 
 FeClg + 2 NH4OH = I Fe(0H)2 + 2 NH4CI ; 
 
 White 
 
 2 Fe(0H)2 + O + H2O = \2 Fe(0H)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 alkahne solu- 
 tions, HgS 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 Law 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 
 j)roduct of the latter. 
 
8o QUALITATIVE CHEMICAL ANALYSIS 
 
 3. Ammonium Sulphide precipitates black FeS, easily soluble 
 in acids with the formation of a ferrous salt and evolution of 
 
 ^aS : — FeCla + (NH^XS = | FeS + 2 NH^Cl ; 
 FeS + 2 HCl = FeCla + 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:— FeClg + 2 KCN = ,|, Fe(CN)2-l- 2 KCl; 
 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, H^Fe(CN)g. The group 
 Fe(CN)5 is an acid radical like CrO^ 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 K2Fe''Fe(CN)5 ; 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 = j FeK2Fe(CN)6 + 2 KCl. 
 
 6. Potassium Ferricyanide produces even in very dilute solu- 
 tions of ferrous salts a dark blue precipitate, known as TurnbuU's 
 blue, which is indistinguishable in color from prussian blue : — 
 
 3 FeCl2 + 2 K3Fe(CN)g = | Fe3[Fe(CN)6Ja -t- 6 KCl. 
 
 The precipitate is insoluble in HCl, but is decomposed by caustic 
 alkalies with the formation of ferrous hydroxide and alkali ferri- 
 cyanide : — 
 
 Fe3[Fe(CN)e]2 + 6 KOH = | 3 Fe(0H)2 + 2 KgFeCCNV 
 
THE METALS 8i 
 
 The ferricyanide at once oxidizes the ferrous hydroxide, so that 
 the final products are Fe(0H)3 and K4Fe(CN)g. 
 
 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 
 HCl and KCIO3, 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 25). 
 
 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, may be 
 represented by the following equations : — . 
 
 2 FeClg -F 2 HCl -I- H2O2 = 2 FeClg -f- 2 HgO ; 
 3 FeCla + 3 HCl -f HNO3 = 3 FeClg -H f NO -h 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(S04)3 • (NH4)2S04 • 24 H2O, is violet. 
 Ferric salts in dilute aqueous solutions are readily hydrolyzed, 
 particularly on heating, with the formation of an insoluble basic 
 ferric salt which dissolves on the addition of an acid : — 
 
 Fe2(S0 J3 -1- H2O -t.\ Fe2(S04)20 -f- H2SO4. 
 
 Reactions 
 
 1. Ammonium, Sodium, or Potassium Hydroxide precipitates a 
 reddish brown gelatinous precipitate of Fe(0H)3. The precipi- 
 
82 QUALITATIVE CHEMICAL ANALYSIS 
 
 tate is unaffected by the presence of ammonium salts [distinc- 
 tion from Fe(OH)2J, 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 : — 
 
 FeClg + 3 NH4OH = |Fe(0H)3 + 3 NH^CI; 
 Fe(OH)3 + 3 HCl = FeClg + 3 HgO. 
 
 On ignition it yields FcaOg : — 
 
 2 Fe(0H)3 + (heat) = FejOg + 3 H2O. 
 
 Ignited, FegOg 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, FcjSg, is precipitated : — 
 
 2 FeCls + 3 (NHJ2S = IFcgSg + 6 NH^Cl. 
 
 The precipitate is readily soluble in hydrochloric acid with the 
 formation of ferrous chloride and the separation of sulphur : — 
 
 FeaSg + 4 HCl = 2 FeClg + f 2 U^S + |S. 
 
 3. Potassium Ferrocyanide produces with ferric salts a blue 
 precipitate known as prussian blue : — 
 
 4 FeClg + 3 K4Fe(CN)6 = |Fe4[Fe(CN)6j3 + 12 KCl. 
 
 The precipitate is insoluble in dilute HCl, 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 : — 
 
 Fei[Fe(CN)6]3 + 12 KOH = ;1,4 Fe(0H)3 + 3 K4Fe(CN)6. 
 
 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 83 
 
 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 KCl. 
 
 The reaction being reversible, its sensitiveness is increased by 
 adding an excess of the reagent. As little as i 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 Httle 
 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. Rela- 
 tively 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 color- 
 ation. 
 
 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 : — 
 
 («) FeClg + 2 Na2HPOi =|FeP04 + NaHaPO^ + 3 NaCl ; 
 {b) FeClg + NaaHPO^ + NaCaHjOa 
 
 = ^¥e?Oi + 3 NaCl + HC2H3O2. 
 
 In (a) all the iron is precipitated but not all the phosphoric 
 acid ; in (J?) both the iron and phosphoric acid are precipitated. 
 
^ 84 QUALITATIVE CHEMICAL ANALYSIS 
 
 FeP04 i^ insoluble in acetic acid, but readily dissolves in HCl 
 Caustic alkalies decompose it into Fe(0H)3 and NagPO^: — 
 
 FePO^ + 3 NaOH =| Fe(OH)3 + NagPOi- 
 
 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 : — 
 
 FeClg + 3 NaCjHgOa = Fe(C2H302)3 + 3 NaCl. 
 
 If this solution is largely diluted and boiled, all the iron will be 
 precipitated as a basic acetate : — 
 
 Fe(C2H302)3 + H20(boiling):i!:|Fe(OH)(C2H30,)2-HHC2H302. 
 
 The presence of non-volatile organic acids or sugar interferes 
 with the precipitation of Fe in reactions i, 6, 7. 
 
 8. Reduction of ferric salts to ferrous may be readily effected 
 in acid solution by reducing agents, as HgS, nascent H, SnClg, 
 H2SO3, HI, and others. The following equations illustrate 
 this : — 
 
 2FeCl3+H2S =2FeCl2-f2HCl-)-|S; 
 
 FeCl3 + H (from Zn -f- HCl) = FeCl^ + HCl ; 
 
 2 FeCl3 + SnClg = 2 FeClg 4- SnCl^ ; 
 
 2 FeClg -I- H2SO3 + H2O = 2 FeClg + 2 H CI + H2SO4 ; 
 
 FeCl3-|-HI =FeCl2 + HCl + |I. 
 
 Nickel 
 
 When in the crystalHne 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 i of cobalt. 
 
 I. Potassium or Sodium Hydroxide precipitates green gelati- 
 nous Ni(0H)2, insoluble in excess and not oxidized on exposure 
 to air : — 
 
 NiCla -I- 2 NaOH = j Ni(OH)2 + 2 NaCl. 
 
THE METALS 85 
 
 The precipitate is readily soluble in acids ; also in ammonium 
 hydroxide and ammonium salts. If the alkaline solution con- 
 taining Ni(0H)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 : — 
 
 NiCla + 2 NH^H = ,|, Ni(0H)2 + 2 NH^Cl. 
 
 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 NH^Cl + 2 NHg = Ni(NH3)^Cl2 + 2 HjO. 
 
 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 + 2 NaC2H302 + HjS = | NiS + 2 NaCl 4- 2 HC2H3O2. 
 
 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 solution. 
 
 NiCla + (NH4>,S= ,|, NiS + 2 NH^Cl. 
 
 Nickel sulphide is practically insoluble in cold HCl(sp. gr. 1.02) 
 (distinction from the sulphides of Mn, Zn, and Fe). It is also 
 
86 QUALITATIVE CHEMICAL ANALYSIS 
 
 insoluble in acetic acid, but is readily taken into solution on 
 heating with agua regia or concentrated nitric acid : — 
 
 3 NiS + 2 HNOg + 6 HC1= 3 NiCl2+ f 2 NO + 4 H^O +|3 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 CI + 4 H2O = H2SO4+ 6 HCl. 
 
 On exposure to air, moist nickel sulphide is oxidized to NiS04. 
 
 5. Potassium Cyanide gives a green precipitate of nickel 
 cyanide, readily soluble in excess with the formation of a 
 double cyanide : — 
 
 NiCla + 2 KCN = | Ni(CN)2 + 2 KCl ; 
 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): — 
 
 KaNiCCN)^ + Br + 3 KOH = \ Ni(0H)3 + 4 KCN + KBr. 
 16. Potassium Nitrite, in dilute solutions of nickel salts acid 
 with acetic acid, gives no precipitate (distinction and method of 
 separation from Co). 
 
 7. Dimethylglyoxime. From solutions of nickel salts, alkaline 
 irith ammonia, an alcoholic solution of dimethylglyoxime will 
 yield a voluminous red precipitate of the composition shown in 
 the following equation : — 
 
 rCH3-C=N0H 
 2 I _ \_ _ + NiCla + 2 NHg = 2 NH^Cl 
 
 CH3-C=N0H CH3-C = N0v 
 
 + I . I >NL 
 
 CH3-C = NOH CH3-C = N0/ 
 
 .CH3-C = NOH 
 
THE METALS 87 
 
 The presence of one part of nickel in 400,000 parts of water 
 may be detected by this reagent. Cobalt salts do not give this 
 reaction. 
 
 8. 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(B02)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 HCl are also blue. 
 
 1. Sodium Hydroxide precipitates from cold solutions a blue 
 basic salt: — 
 
 C0CI2 + NaOH = |Co(OH)Cl + NaCl. 
 
 This is converted, on warming in contact with the alkali, to pink 
 cobaltous hydroxide : — 
 
 Co(OH)Cl + NaOH = jCoCOH)^ -h 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 -1- O = |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 NH4CI -f 2 NH3 = Co(NH3)4Cl2 + 2 H2O. 
 
 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. 
 
88 QUALITATIVE CHEMICAL ANALYSIS 
 
 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 HCl (sp. gr. 1.02) and in acetic acid. It is soluble 
 in agua regia and concentrated nitric acid with the separation 
 of sulphur : — 
 
 3 CoS + 8 HNO3 = 3 Co(N03)2 + ^2 NO + 4 HgO + 13 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 : — 
 
 C0CI2 + 2 KCN = ,|,Co(CN)2 + 2 KCl; 
 Co(CN)2 + 4 KCN = KiCo(CN)6. 
 
 The latter is similar to potassium ferrocyanide, hence it is called 
 potassium cobaltocyanide. On warming the solution of the 
 double cyanide for some tinne, it changes color to a bright yel- 
 low, due to oxidation to potassium cobalticyanide [similar to 
 KgFeCCN)^]:- 
 
 2 K4Co(CN)6 + H2O + = 2 K3Co(CN)6 + 2 KOH. 
 
 The reaction takes place more rapidly if sodium hydroxide 
 and bromine or, what amounts to the same, NaBrO so- 
 lution, is added to the solution of potassium cobalto- 
 cyanide. Nickel does not form the corresponding com- 
 pound, but under these conditions it is converted into black 
 insoluble Ni(OH)3 (distinction 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 
 crystalUne precipitate of potassium nitrocobaltate, K3Co(N02)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 : — 
 
THE METALS 89 
 
 {a) C0CI2 + 2 KNO2 = Co(N02)2 + 2 KCl; 
 
 {b) 2 KNO2 + 2 HC2H3O2 = 2 HNO2 + 2 KCaHgOa; 
 
 (0 2 HNO2 = H2O + NO + NO2 ; 
 
 {d) Co(N02)2 + NO2 = Co(N02)3; 
 
 {e) Co(N02)3+ 3 KNO2 = ,|,K3Co(N02)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(N02)6, the 
 solution of cobalt should be concentrated by evaporation, the 
 mineral acid replaced by acetic, saturated with KCl, 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 
 sulphate acidified with hydrochloric acid, a voluminous red pre- 
 cipitate of cobalti-nitroso-/S-naphthol (distinction and method of 
 separation from nickel, Which, in HCl 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 are 
 nearly all colorless. 
 
 Reactions 
 
 I. Sodium or Potassium Hydroxide produces with manganous 
 salts a white precipitate of Mn(0H)2, which, on exposure to air. 
 
 rapidly oxidizes, becoming brown : — 
 
go QUALITATIVE CHEMICAL ANALYSIS 
 
 MnClg + 2 NaOH = | Mn(OH)2 + 2 NaCl ; 
 
 Mn(0H)2 + O = ,1, MnO(OH)2 (manganous acid) ; 
 MnO(OH)2 + Mn(OH)a = I MngOg + 2 H2O. 
 
 2. Ammonium Hydroxide yields with manganous solutions, in 
 the absence of ammonium salts, a partial precipitation of white 
 Mn(0H)2, oxidizing, as described in ( i ), to brown MnjOg. In the 
 presence of a sufficient amount of ammonium salts, no immediate 
 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 NH4CI and ammonium 
 hydroxide is therefore incomplete. Non-volatile organic acids 
 and sugar interfere with the precipitation of Mn(0H)2. 
 
 ""^ 3. Ammonium Sulphide precipitates light pink hydrated man- 
 ganous sulphide, which, on exposure to air, becomes dark brown, 
 due to partial oxidation to MugOg : — 
 
 MnClj + (NH J2S = I MnS -f- aq. -f- 2 NH^Cl. 
 
 The precipitate is easily soluble in dilute acids ( distinction from 
 Ni^d_Co], even in acetic acid (dist inction from Zn, as well as 
 Ni and Co). The addition of ammonium chloride assists the 
 precipitation, while the presence of oxalates and tartrates retards 
 it. On boiling with a large excess of (NH4)2S, MnS + aq. is 
 changed to a less hydrated green sulphide of the formula 
 3 MnS • H2O. 
 
 -^ 4. Lead Dioxide and Nitric Acid. If a very dilute solution of 
 manganous salt, free from HCl 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 MnSOi + 5 Pb02 + 6 HNO3 
 
 = I 2 PbS04-f 3 Pb(N03)2 + 2 HjO -I- 2 HMn04. 
 
 Purple 
 
 This reaction is sufficiently delicate to detect a trace of man- 
 ganese. 
 
THE METALS 
 
 91 
 
 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, Na2Mn04: — 
 
 Mn(0H)2 + (heat) = MnO + HjO ; 
 MnO(OH)2 + (heat) = MnOj + HgO ; 
 MnO + NagCOg + O2 = f COj + Na2Mn04 ; 
 MnOa + NaCOg + O = f CO2 + Na2Mn04. 
 
 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]: — 
 
 ZnCIa + 2 NaOH = | Zn(OH)2 -1- 2 NaCl ; 
 
 Zn(OH)2 -I- 2 NaOH = NajZnOa + 2 H2O. 
 
 Unless the solution of the zincate contains a decided excess of 
 NaOH, it will be decomposed on boiling with the reprecipitation 
 of the hydroxide : — 
 
 Na2Zn02 + 2 HgO (boiling) :5:2 NaOH + ,|,Zn(0H)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 : — 
 
 ZnCla + 2 NH4OH = |Zn(0H)2 + 2 NH^Cl ; 
 Zn(OH)2 -f 2 NH4CI + 2 NHg = 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 
 
92 QUALITATIVE CHEMICAL ANALYSIS 
 
 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, 
 HgS completely precipitates all the zinc as sulphide on boiling. 
 Warming in the presence of alkali acetate promotes the pre- 
 cipitation : — 
 
 (a) ZnClg + HjSl^lZnS + 2 HCl ; 
 
 {d) ZnClg + 2 NaCgHgOa + H2S = >l,ZnS + 2 NaCl+2 UC^H^O^. 
 
 As ZnS is soluble in HCl, the precipitation in (a) is never com- 
 plete. In equation (i>), NaCgHgOa has the effect of displacing 
 the strong HCl 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 NaCjHgOj, and 
 filtering rapidly while hot. The precipitate may then be washed 
 with hot water containing NH4C2H3O2 or NH^^NOj and HjS. 
 Zinc may also be precipitated by H2S from sodium hydroxide 
 solutions : — 
 
 Na2Zn02 + HgS = |ZnS -|- 2 NaOH. 
 
 4. Ammonium Sulphide yields in neutral and alkaline solutions 
 a white precipitate of ZnS : — 
 
 ZnCla -|-(NHJ2S = |ZnS -F 2 NH^Cl. 
 
 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 (Thdnard's 
 green). This is an excellent confirmatory test for zinc and 
 serves to distinguish zinc from aluminum. 
 
 * For an explanation of this fact according to the ionic theory and mass action 
 law see page 15. 
 
THE METALS 
 
 Outline of the Method of Analysis for Group III 
 
 93 
 
 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. We have not adopted this method 
 for the reason that under the conditions given, manganese and 
 zinc are not completely held in solution, and if present in small 
 amounts may be wholly precipitated with the trivalent metals. 
 "However, the method gives fairly satisfactory results if the 
 first precipitate of the hydroxides of Al, Cr, and Fe, containing 
 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 method adopted in this book consists in precipitating the 
 entire group with (NHjjS after rendering the solution alkaline 
 with NH^OH. Instead of (NH4)2S, the hot ammoniacal solu- 
 tion may be treated with a stream of HjS 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 HCl (i : 9), it follows 
 
94 QUALITATIVE CHEMICAL ANALYSIS 
 
 that if the third group precipitate is treated with a sufficient 
 amount of HCl (i : 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-/8-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 HCl and HjS. The greater part of the HCl and all of the 
 HgS 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 HjOj 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 NajOj are added to the main filtrate, which 
 has been freed from HgS and the greater part of the HCl 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 
 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 HCl, 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 NagOa or NaOH. When a rough estimation 
 of the amount of Mn and Fe is not desired, separation is unneces- 
 
TEE METALS 95 
 
 sary, for we can readily identify each in the presence of the 
 other. The presence of Mn may be determined by the charac- 
 teristic green bead it gives when a little of the mixture is fused 
 in a NagCOg bead in the presence of an oxidizing agent ; the 
 iron may be detected by dissolving part of the precipitate in hot 
 dilute HCl 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 HNOg, about three 
 grams of NH^Cl 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 NH^Cl, 
 while the chromium no longer acts as metal but as the acid 
 radical (CrO^), and in consequence is not precipitated by am- 
 monium hydroxide. The filtrate from the Al(OH)3 will contain 
 zinc, chromium, and a slight excess of ammonia. By rendering 
 the solution acid with acetic acid and adding BaClj, all of the 
 chromium will be precipitated as BaCr04. From the filtrate 
 the Zn may be precipitated by HgS. 
 
 Scheme of Analysis for Group UI 
 
 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' 
 HjS, contains, besides the metals of the succeeding group, an 
 excess of HCI. 
 
 Preliminary Test. To a small portion of the filtrate from 
 Group II. which has been freed from HjS, add 2-3 drops of 
 concentrated HNO3 and boil. Add about 0.5 g. of NH^Cl and 
 then ammonium hydroxide to alkaline reaction. A precipitate 
 may be Fe(0H)3 (red), A1(0H)3 (white), or Cr(0H)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, 
 
96 QUALITATIVE CHEMICAL ANALYSIS 
 
 or to the filtrate from Group II., in which ammonium hydroxide 
 produces no precipitate, (NH^)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 pre- 
 cipitate 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 exposure, Mn is present ; the latter may at once be verified 
 by the NajCOg bead + KCIO3. Failure to precipitate with 
 NH^OH and (NH^)2S in the presence of a sufficient amount of 
 NH^Cl proves the absence of Group III. In that case pass to 
 Scheme IV. 
 
 NOTES 
 
 1. As a solution containing i 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 used in the preliminary testing for Group III., as well as 
 that needed later in the preparation of the wash water for the Group III. ppt., 
 should be made as needed by treating a little dil. NH4OH with a stream of 
 HjS for several minutes. 
 
 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 NH4CI (i) rendered alkaline 
 with NH4OH and then 2 cc. of strong NH4OH in excess are added (2). 
 The mixture is heated and treated with a. stream of H„S until precipitation is 
 complete. Stir vigorously with gentle heating for a minute (3) and filter on 
 a fluted filter. [The filtrate (4) is at once made acid with acetic acid (5), 
 boiled until all the HjS 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(0H)3, Cr(0H)3, FeS, NiS, CoS, MnS, and 
 ZnS. Wash once with hot water containing NH4CI and a little (NH4)2S (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 HCl (sp. gr. 1.02), prepared by mixing 
 I part of cone. HCl with 9 parts of water, stir thoroughly (without heating)' 
 for about a minute ; allow to settle (9), and filter. 
 
THE METALS 
 
 97 
 
 Residue is NiS + CoS + S (lo). Wash on filter once with HCl (i : 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 (ii). 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. agua regia, and boil till all but a small 
 amount of black S dissolves; evaporate 7?^/ to dryness (12). Take up with 
 2 cc. of dil. HCl 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 3 cc. in excess (13).* Saturate the 
 solution with KCl (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 KNOo. AUovf the 
 ppt. to stand with frequent shaking for about \ hour and filter. Residue is 
 K3Co(NO„)i; (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 AICI3, CrClj, FeClj, MnCL, ZnCl, + HjS + excess HCl. 
 Boil down in a large evap. dish under hood to about i cc. (16) ; dilute with 
 10 cc. of water, render strongly alkaline witli clear NaOH solution, and add 
 (under hood and with caution) 2 g. of Na^Og (17). Dip finger into solution, 
 rub on thumb ; a greasy feel indicates solution is suiBciently alkaline ; if not, 
 add more NajOs. Boil with constant stirring for about i minute, add 10 cc. 
 of water, and filter. 
 
 * If a solution of dimethylglyoxime is available, proceed from this point as fol- 
 lows: Divide the solution into two equal portions. First portion test for cobalt by 
 adding an equal volume of KNO2 solution and warming. A yellow precipitate is 
 K3Co(N02)6. Confirm with borax bead. Second portion test for nickel by rendering 
 the solution alkaline with NH4OH and adding i cc. of dimethylglyoxime. A red 
 precipitate proves the presence of nickel. 
 
98 
 
 QUALITATIVE CHEMICAL ANALYSIS 
 
 Residue is Fe(0H)3 
 + MnO.,,zH20(i8). 
 Wash with hot water. 
 7esi for Mil. On a 
 Na^COo bead take 
 up a very small 
 amount of the ppt. 
 and heat ; while hot 
 dip into a little 
 powdered KCIO.; 
 contained in a small 
 watch glass. A green 
 or bluish green mass 
 of NaoMnO^ proves 
 the presence of Mn. 
 
 Test for Fe. 
 Transfer part or all 
 of the ppt. to a test 
 tube, add dil. HCl 
 and heat till solution 
 takes place. If Mn 
 has been shown to 
 be present, boil until 
 all the chlorine is 
 expelled. Cool and 
 add a few drops (19) 
 of K4Fe(CN)6; a 
 blue ppt. proves the 
 presence of Fe. 
 
 Filtrate (20) may contain NagAlO,, NagCrO^, NajZnOj 
 + excess NaOH. Render slightly acid with cautions 
 addition of cone. HNO3; add 3 grams of NH^Cl (21), 
 heat to boiling, and then add NHjOH drop by drop 
 with constant stirring until the resulting alkaline solu- 
 tion has only a faint odor of ammonia (22). (If too, 
 much NH^OH is added, boil off the excess and filter.) *■ 
 
 Residue is 
 Al(OH)3 (28). 
 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(NO,)„ (23) 
 and . igniting 
 strongly. A 
 blue mass is co- 
 balt aluminate, 
 /AL03..rCoO. 
 
 Filtrate may contain (NH4)oCr04, 
 Zn(NH3)4Cl2 -I- excess NH4OH. Acid- 
 ify with acetic acid, add a gram of 
 NaCjHsOi (25), and heat to boiling. 
 To the hot solution add BaClc, drop by 
 drop, till precipitation is complete ; allow 
 to settle and filter through a double filter. 
 
 Residue is 
 yellow BaCr04 
 (24). Confirm 
 by treating ppt. 
 on filter with 
 hot dil. HNO3. 
 Catch filtrate in 
 atest tube. Cool 
 thoroughly. 
 Add I cc. of 
 ether and i cc. 
 3% H.3O3 and 
 shake. A blue 
 color in the 
 ether layer 
 proves the pres- 
 ence of Cr. 
 
 Filtrate, which 
 should be perfectly 
 clear (26), is treated 
 with H^S. A white 
 ppt. is ZnS (27). To 
 confirm the presence 
 of Zn, filter, wash with 
 hot water, and then 
 moisten with 2 drops 
 of Co(N03)2. Wind 
 Pt wire around paper 
 apd ppt. and inciner- 
 ate. A green mass is 
 yZnO -xCoO. 
 
 NOTES 
 
 if '¥ 
 
 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 NHs to neutralize the free acid. An 
 excess of strong NH4OH is then added to form with the H;;S the (NH4)2S 
 which precipitates all the metals of this group. Treatment of the ammoniacal 
 solution with H^S is preferable to the use of (NH4)2S, for the reason that with 
 the former NiS is prevented from going into solution. If, however, (NH4)2S 
 
THE METALS 
 
 99 
 
 is preferred, the filtrate from Group II. after the addition of 2 g. of NH4CI 
 should be first rendered alkaline with (NH.i)OH, heated to boiling, and then 
 treated with an excess of colorless (NH4)2S. 
 
 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. 
 
 5. The filtrate is at once acidified with acetic acid to destroy the excess 
 of (NH4).3S, which, on standing, would oxidize to sulphate and precipitate 
 the alkaline earths. Rendering the solution acid also prevents the formation 
 of (NH4)2COg from the absorption of atmospheric COj. (NH4)2C03, if 
 formed, would also precipitate the alkaline earths. 
 
 6. The HjS 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 (when (NH4)2S has been 
 used as the precipitant) and coagulated S, and may be tested for Ni either 
 ■with a borax bead in the O. F., or by solution of the ppt. in hot dil. HNO3, 
 rendering alkaline with NH^OH and adding dimethylglyoxime. 
 
 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)„S. NH4CI is added to the wash water to prevent the precipitate 
 from passing through the filter in the colloidal condition. It is also well to 
 keep the funnel covered as much as possible during the filtering and washing 
 to minimize the oxidizing influence of the air. 
 
 9. The separation of Ni and Co from the remaining metals by the use of 
 HCl (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. The experiments of Curtman and Rothberg show that in mixtures of 
 the sulphides of Ni and Co containing as little as 3'7o CoS, a blue bead is 
 obtained. With 2.5% of CoS uncertain results are obtained, while with 2% 
 or less of CoS, the mixtures give brown beads. 
 
 12. Evaporation to i or 2 drops will suffice ; if the evaporation is carried 
 to the point of dryness, care must be taken not to ignite the residue. Should 
 the residue be accidentally ignited, redissolve in hot aqua regia and evaporate 
 again to a few drops. 
 
lOO QUALITATIVE CHEMICAL ANALYSIS 
 
 13. These conditions must be closely adhered to. KNOj 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 HCl 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 lo cc. be obtained 
 in making the test, it is recommended that the solution be evaporated to 
 10 cc. before saturating with KCl. 
 
 14. Under the conditions stated in Note 13, complete precipitation of Co 
 by KNO2 takes place after a lapse of 24 hours. By saturating the solution of 
 Co with KCl, in which K3Co(N02)6 is insoluble, and by using an excess 
 of KNO2, the precipitation of Co may be rendered complete in a half hour. 
 
 15. Not infrequently a precipitate of uncertain color is obtained with 
 NaOH. In that case the presence of Ni cannot be considered proved until a 
 characteristic Ni bead is obtained. It is however better to employ the di- 
 methylglyoxime text which is more sensitive and characteristic for Ni. 
 
 16. The solution is evaporated to i cc. to remove the excess of HCl 
 which, if present, would neutralize the NaOH next to be added. The HjS is 
 expelled at the same time. 
 
 17. NajOj added to water even in the cold decomposes, giving NaOH + O. 
 At higher temperatures the decomposition takes place violently. Na20sj 
 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 Na202 and the perchromates 
 which first form. The Na202 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^NajZnOa -I- 2 H2O. 
 
 18. Any Ni and Co dissolved by the 1.02 HCl 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)g, 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. A portion of this solution may be tested directly for Cr by 
 acidifying with HNO5, cooling thoroughly, adding i cc. each of ether and 3% 
 H2O2 and shaking. A blue color in the ether layer proves the presence of Cr. 
 
 21. The NH4CI is added to prevent a partial precipitation of Zn. 
 
 22. An excess of NH4OH is to be avoided because of the slight solubility 
 of A1(0H)5 in excess. 
 
I 
 
 THE METALS lot 
 
 23. The Co(N03)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 BaSO^ (white) to form along with the BaCrO^ 
 (light yellow), thus obscuring the test for Cr; hence the necessity of making 
 the confirmatory test. 
 
 25. The addition of NaCaH^Oa represses the ionization of the acetic acid 
 with the result that the solvent action of the latter on BaCrO., is reduced to a 
 minimum. 
 
 26. If the filtrate from the BaCr04 is not clear, filter again, through another 
 double filter, and repeat this treatment until a perfectly clear filtrate is obtained. 
 
 27. If a black ppt., due to FeS, NiS, etc., is obtained, add dil. HCl, heat, 
 and filter. Render filtrate alkaline with NaOH, add a little Na202, boil, dilute, 
 and filter. Test filtrate with H2S. A white ppt. is ZnS. 
 
 28. A slight precipitate of Al(OH)3 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)2C03 in the presence of NH4CI. 
 
 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 HgS. 
 
I02 QUALITATIVE CHEMICAL ANALYSIS 
 
 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. 
 
 1. 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 tur- 
 bidity results from the formation of barium carbonate. The 
 hydroxide is not precipitated because of its ready solubility in 
 water (i part in 20 of cold water): — 
 
 BaCla + 2 NH4OH = Ba(0H)2 + 2 NH4CI; 
 ' Ba(OH)2 + CO2 = I BaCOg + HjO ; 
 2 NH^H + CO2 = (NH4)2C03 + H2O ; 
 (NH4)2C03 + BaCla = | BaCOj -I-2 NH^Cl. 
 
 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)2C03, hence it will yield an immediate tur- 
 
 /bidity with barium salts : — 
 
 (i) (NH4)2S + BaClj = BaS + 2 NH4CI; 
 
 (2) BaS -I- 2 H2O = Ba(0H)2 + f HgS; 
 
 (3) Ba(OH)2-f-C02 = |BaC03-|-H20. 
 
 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 : — 
 
 BaClg + (NH4)2C03 = | BaCOg -I- 2 NH4CI. 
 
 * AH alkaline liquids will, on exposure to' the air, absorb CO2, with the formation 
 of carbonate proportional to the amount of CO3 absorbed; the reagent ammonium 
 hydroxide will, therefore, always contain a little (NH4)2COs, and hence will yield a 
 slight precipitate with the salts of the alkaline earths (see equations) . 
 
THE METALS 
 
 103 
 
 The precipitate is slightly soluble in NH4CI; therefore, in very 
 dilute solutions of barium salts containing much NH^Cl, ammo- 
 nium carbonate does not produce a precipitate. The precipi- 
 tate is easily soluble in acids, even in acetic and carbonic 
 acids : — 
 
 BaCOg + 2 HC^HsOa = B2L{C^li^0^\ + Hp + 1 CO^; 
 BaCOg + H2CO3 = BaH2(C03)2. 
 
 Boiling the dicarbonate decomposes it with the evolution of 
 CO2 and precipitation of the normal carbonate : — 
 
 BaH2(C03)2 + (heat) = \ BaCOg + HjO + \ CO^. 
 
 4. Dilute Sulphuric Acid or any soluble sulphate produces even 
 in very dilute solutions of barium salts a heavy, white, finely 
 divided precipitate of BaSO^, practically insoluble in water (i 
 part in 400,000 parts of water) : — 
 
 BaCla + H2SO4 = I BaS04 + 2 HCl. 
 
 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 NagCOg, it undergoes partial decomposition, 
 according to the equation : — 
 
 BaSO^ + NajCOgl^t,!. BaCOg + NagSO^. 
 
 The decomposition is incomplete because of the reversibility 
 of the reaction. If, however, the mixture is filtered, and the 
 residue of BaCOg and unchanged sulphate is boiled with a fresh 
 NajCOg solution, more BaSO^ 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 NagCOg. Under these conditions, 
 the reaction proceeds to completion in one operation. On 
 cooling the melt, boiling it with water, and filtering, there will 
 
I04 QUALITATIVE CHEMICAL ANALYSIS 
 
 remain on the filter a residue of BaCOg equivalent in amount to 
 the BaSO^ 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. 
 PbS04, SrS04, and CaSO^ may be completely converted into 
 carbonate by the first method. 
 
 5. Potassium Chromate precipitates from neutral or acetic 
 acid solutions yellow barium chromate : — 
 
 K2Cr04 + BaClj = \ BaCrO^ + 2 KCl. 
 
 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 
 (KjCrgO^) only partial precipitation results : — 
 
 2 BaCIa + KaCraO^ + HgO = \2 BaCrO^ + 2 KCl + 2 HCl. 
 
 This is due to the formation of HCl, which exerts a solvent 
 action on BaCrO^; the addition of sodium acetate will render 
 the precipitation complete. 
 
 BaCrO^, like BaS04, 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 
 (i part in 2600) and completely soluble in boiling acetic acid 
 (distinction from Ca) : — 
 
 BaClg + (NH4)2C204= ,|, BaCgOi + 2 NH4CI. 
 
 7. Disodium Phosphate precipitates in neutral solutions white 
 iiocculent BaHP04 > ^^ ammoniacal solutions, sodium phos- 
 
TEE METALS 105 
 
 phate throws down the tertiary phosphate. The precipitates 
 are easily soluble in dilute acids, even in acetic acid : — 
 
 NagHPO^ + BaClj = 2 NaCH- ,J, BaHP04; 
 2 Na2HP04+3 BaCl5+2 NH3=|Ba3(P04)2+4 NaCl+2 NH^Cl. 
 
 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 HCl. 
 
 Strontium 
 
 1. Ammonium Hydroxide. Same as with Ba salts. 
 
 2. Ammonium Sulphide. Same as with Ba salts. 
 
 3. Ammonium Carbonate precipitates white SrCOg, more in- 
 soluble in water than BaCOg; in other respects, it possesses 
 about the same solubilities as BaCOg. 
 
 4. Dilute sulphuric acid or any soluble sulphate yields a white 
 precipitate of SrSO^. The precipitate is more soluble in water 
 (i part in 7000) and in acids than BaS04, and, as a consequence, 
 is precipitated from very dilute solutions only after some time ; 
 it is, however, much less soluble in water than CaS04, the latter 
 dissolving in water to the extent of i part in 500. 
 
 SrS04 is practically insoluble in a strong solution of (NH4)2S04, 
 even on boiling (distinction and method of separation from Ca). 
 
 5. Saturated CaS04 Solution yields with dilute solutions of 
 strontium salts a precipitate of SrS04, which forms only after 
 some time (distinction from Ba, which yields an immediate pre- 
 cipitate). Precipitation in this case, as well as in 4, is promoted 
 by heating, and is retarded by the addition of acids. From 
 concentrated solutions of Sr salts, an immediate precipitate 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 
 
io6 QUALITATIVE CHEMICAL ANALYSIS 
 
 Ba). From neutral concentrated solutions, however, a yellow 
 crystalline precipitate of SrCr04 forms which is soluble in acetic 
 acid. 
 
 7. Ammonium Oxalate. Same as with Ba. SrC204 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 
 
 1. Ammonium Hydroxide. Same as with Ba. 
 
 2. Ammonium Sulphide. Same as with Ba. 
 
 3. Ammonium Carbonate precipitates white amorphous CaCOg, 
 becoming crystalline on heating ; it is more insoluble in water 
 than BaCOg, but in other respects its solubilities are about the 
 same as those for BaCOg. 
 
 4. Dilute Sulphuric Acid or any alkali sulphate does not pro- 
 duce a precipitate from dilute solutions. From concentrated 
 solutions a white precipitate of CaS04 is obtained which is 
 appreciably soluble in a hot concentrated solution of (NH4)2S04 
 (distinction and method of separation from Sr). A saturated 
 solution of CaS04, 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 (i part in 170,000) and 
 in acetic acid, but is readily soluble in mineral acids : — 
 
 CaC204 + 2 HCl = CaClj + H2C2O4. 
 
 This is a most delicate test for Ca. 
 
THE METALS 107 
 
 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. 
 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 NH^Cl 
 to prevent the precipitation of Mg along with the alkaline 
 earth carbonates on adding the group reagent, (NH4)2C03. 
 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 KgCrO^. 
 After filtering the BaCr04, the filtrate will contain, besides Sr 
 and Ca, an excess of KjCrgO^. 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)2S04 and filtered, the Sr will be on the filter as 
 SrSO^, while the Ca will pass into the filtrate. From the 
 
io8 QUALITATIVE CHEMICAL ANALYSIS 
 
 latter the calcium may be precipitated as CaC204 with 
 
 The separation of Sr and Ca by the use of a boiling solution 
 of (NH4)2S04 is not complete; some CaSO^ remains undis- 
 solved, while at the same time a small amount of SrS04 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 a double filter until a perfectly 
 clear liquid is obtained ; it should then be concentrated by evaporation to 
 about 30 cc* 
 
 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 
 (NH^)2COg, and warm ; a white ppt. proves the presence of Group IV. If no 
 ppt. is obtained, the absence of more than traces (i) of the alkaline earths is 
 indicated ; in that case proceed to Scheme V. 
 
 \U 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 ; (NH^JZOg 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 BaCOg, SrCOg, and CaCOg. Wash 
 once with hot water and reject the washings. Dissolve the ppt. on the filter 
 with the 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 
 K2CrOj drop by drop till precipitation is complete. Allow the ppt.jto settle 
 and filter (using a double filter) by decantation. ' Finally, with the aid of hot 
 water, bring the ppt. on the filter. 
 
 * Any NH4Q 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 
 
 109 
 
 Residue is yel- 
 low BaCrO^. 
 Wash twice with 
 hot water and re- 
 ject washings. 
 Confirm by dip- 
 ping clean Pt 
 wire (9) mois- 
 tened with cone. 
 HCl 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 
 KjCr^O^. To remove the latter, add NH4OH to alkaline 
 reaction and then (NH J2CO3 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 SrCOg and 
 CaCOj. 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). Voxir z. very small portion 
 of this solution into a test tube and add a little CaSO^ so- 
 lution, heat to boiling, 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 (i). 
 
 (a) If Sr is present, render the remainder of the solu- 
 tion alkaline with NH^OH, add 5 cc. of (NHJ^SO^, boil 
 for a few minutes, and filter. Ppt. is SrS04. Wash with 
 hot water and confirm by flame test (8). Test filtrate 
 for Ca by adding (NHJ2C2O4. A white ppt. insol. in 
 HC2H3O2 is CaC204. Confirm by flame test. 
 
 (d) Make solution alkaline with NH4OH and add 
 (NH4)2C204. A white ppt. insol. in acetic acid is CaCjO^. 
 Confirm by flame test. 
 
 NOTES 
 
 1. Tests for traces of alkaline earths. If no precipitate is obtained with 
 (NHJjCOg, treat a small portion of the solution with dilute H2SO4, boil, and 
 allow to stand for some time. A white precipitate of BaS04 proves the pres- 
 ence of Ba^ Treat another small portion with NH4OH until alkaline, add 
 (NH4)2C204, heat to boiling, and allow to stand ; a cloudiness or white pre- 
 cipitate proves the presence of Ca. Traces of Sr are best tested for by 
 means of the spectroscope. 
 
 2. The mixture must not be boiled, because at the boiling temperature 
 (NH4)2C03 is decomposed according to the equation — 
 
 (NH4)2C03 = >|. 2 NHj + H2O + 4" 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. 
 
no QUALITATIVE CHEMICAL ANALYSIS 
 
 3. To insure completeness of precipitation, add a little (NH4)2C03 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 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 KjCrO^ should be made on a small portion of 
 this solution. If a precipitate is obtained, Ba is present and the entire solu- 
 tion should be treated with KjCrO^ as described in the scheme. If no precipi- 
 tate is obtained, Ba is absent : in that case the solution should not be treated 
 with KjCrO^ solution, but on a small portion make a preliminary test for Sr 
 with CaSO^ solution ; if present, treat the remainder of the solution accord- 
 ing to (a) ; if absent, proceed to (J)) . 
 
 6. BaCrO^, 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. 
 
 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 HCl, 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 BaSO^, but can 
 readily be distinguished from SrS04 by its failure to yield a red coloration to 
 the flame. 
 
 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 
 may be removed by one of the following methods : 
 
 (fl) Any large particles of matter adhering to the wire must first be me- 
 chanically removed. The wire is then dipped into concentrated C.P. HCl, 
 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. If 
 clean, it should give no color to the flame. 
 
 Under no circumstances must the wire, clean or otherwise, be dipped into 
 the reagent bottle of acid. The efiiciency of this method will be indicated by 
 the fact that the flame coloration becomes noticeably fainter with each treat- 
 ment. 
 
THE METALS m 
 
 1 
 
 (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 
 HCl as described in (a). 
 
 10. The Pt wire need only be dipped once into the BaCrO^ precipitate. 
 After the first heating with HCl, it should be moistened with HCl again and 
 heated. This is to be repeated several times without redipping into the 
 BaCr04 precipitate. 
 
 GROUP V 
 
 Group V. embraces the metal magnesium, the alkali metals 
 potassium and sodium, and the metallic radical ammonium 
 (NH^). 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 NH^Cl in precipitating the 
 third and fourth groups. 
 
 The test for NH^ 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 
 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 NH^, all the alkali metals give 
 characteristic flame and spectroscopic reactions. 
 
112 QUALITATIVE CHEMICAL ANALYSIS 
 
 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 NH^OH, (NH4)2S, nor (NHJjCOs 
 precipitates magnesium salts in the presence of a sufficient 
 amount of ammonium chloride, hence the classification of mag- 
 nesium with the alkali metals. 
 
 1 . Ammonium Hydroxide gives a partial precipitation of gelati- 
 nous magnesium hydroxide, Mg(OH)2, readily soluble in am- 
 monium salts : — 
 
 MgSO^ -1- 2 NH^H :^\ Mg(OH)2 + (NH4)2S04. 
 
 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 NH^ salts, or, what amounts to the same thing, 
 the non-precipitation of Mg salts by NH^OH in the presence of 
 a sufficient amount of NH^ 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(OH)2, insoluble in 
 excess and nearly insoluble in water, the solubility being i part 
 in 10,000. The washed precipitate is soluble in NH^ salts : — 
 
 MgSO^ + 2 NaOH =| Mg(OH)2 + NajSO^. 
 
 Boiling promotes precipitation. 
 
 3. Ammonium Carbonate, in the presence of NH^ salts, gives 
 no precipitate. In solutions containing no NH^ 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 MgSO^ + 4 (NH J2CO3 + H2O = 
 
 ^ Mg,(C03)3(OH)2 + t CO 2 -F 4 (NH,)2S04 
 
THE METALS 
 
 113 
 
 4. Disodium Phosphate, when added to a neutral solution of a 
 Mg salt, precipitates flocculent MgHPO^: — 
 
 NaaHPO^ + MgSO^ = | MgHPO^ + NagSO^. 
 
 If, however, NH4CI 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 : — 
 
 MgCla + NagHFO^ + NH3 = |NH4MgP04 4- 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 (i part in 13,500 at 
 23° C), but is practically insoluble in 2.5 per cent, ammonia 
 water. It is readily soluble in acetic acid. The addition of 
 NH^Cl 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 MgCjO^. The presence of NH^ 
 salts renders the precipitation incomplete. 
 
 6. HgS, (NH4)2S, and H2SO4 do not precipitate magnesium 
 salts. 
 
 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 (HgPtCle) produces in neutral or in 
 concentrated acid solutions of potassium salts a yellow crystal- 
 line precipitate of potassium chlorplatinate (KaPtClg) : — ■ 
 
 2 KCl -f- HaPtCle =|K2PtCl6 + 2 HCl. 
 
 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 appHes 
 to nearly all crystalline precipitates. The precipitate is soluble in 
 
114 QUALITATIVE CHEMICAL ANALYSIS 
 
 alkalies ; it dissolves in water to the extent of i part in lOO at 
 1 5° C. It is practically insoluble in 80 per cent, alcohol. On 
 ignition it decomposes according to the following equation : — 
 
 KaPtClg == 2 KCl +>l'Pt +f 2CI2. 
 
 Solutions of potassium iodide and potassium cyanide do not 
 give this precipitate ; they should first be changed to chloride by 
 evaporation with concentrated HCl. As NH^ salts yield a simi- 
 lar precipitate they must be removed before the test is applied. 
 
 2. Tartaric Acid (H2C4H^Og) produces in neutral solutions of 
 potassium salts, which are moderately concentrated, a white 
 crystalline precipitate of potassium acid tartrate : — 
 
 KCl + H2C4H4O6 =|KHC4H406 + HCl. 
 
 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 (NaHC^H^Og) is preferable, 
 as it does not yield any free acid as a by-product : — 
 
 KCl -I- NaHC4H406 =,1,KHC4H406 + 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. 
 
 3. Sodium Cobaltic Nitrite [Na3Co(N02)6] yields, with solu- 
 tions of potassium salts acidified with acetic acid, a yellow pre- 
 cipitate of potassium cobaltic nitrite [K3Co(N02)6J . 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 15° 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 0.00 1 mg. KCl. The presence 
 
THE METALS 115 
 
 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 NH^ 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 (HjPtClg) precipitates, under the 
 same conditions given for potassium compounds (which see), 
 a yellow crystalline precipitate of ammonium chlorplatinate, 
 (NHJaPtCle:- 
 
 2 NH4CI + HaPtClg =|(NH4)2PtClg + 2 HCl. 
 
 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 = NaaPtClg + f 2 NH3 + 2 HjO. 
 
 When strongly heated, it leaves a residue of platinum sponge 
 only (distinction from the K compound). Ammonium chlorplati- 
 nate IS somewhat less soluble in water than the corresponding 
 K compound ; it is insoluble in alcohol. 
 
 2. Sodium Cobaltic Nitrite yields with solutions of ammonium 
 salts a yellow precipitate similar to that given by potassium salts. 
 Hence before testing for potassium with this reagent, all the 
 ammonium salts must be removed. 
 
 3. Tartaric Acid (H2C4H40g) or NaHC4H40g gives from con- 
 centrated solutions of ammonium salts a crystalline precipitate 
 of NH^HC^H^Og. 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. 
 
Ii6 QUALITATIVE CHEMICAL ANALYSIS 
 
 It may be distinguished from the corresponding K compound 
 by the evolution of NHg when treated with an excess of NaOH. 
 
 4. 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 the ability of 
 the gas evolved to turn moistened red litmus paper blue : — 
 
 NH^Cl + NaOH = f NH3 + NaCl + U^O ; 
 2 NH4CI + Ca(0H)2 = f 2 NH3 + CaCla + 2 HgO. 
 
 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 HgNOg + 2 NHg = KNHjHgNOs + Hg) + NH^NOg. 
 
 5. 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(2KI.Hgl2)+NH3+3KOH = |NHg2l.H20-|-7KI + 3H20. 
 
 6. 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. 
 
 Sodium 
 
 All of the salts of sodium, with the exception of the pyroanti- 
 monate, are soluble in water. 
 
 I. Potassium Pyroantimonate Solution (KjHgSbgOy) precipi- 
 tates from neutral or slightly alkaline solutions of sodium salts 
 that are fairly concentrated, a white crystalline precipitate of 
 sodium pyroantimonate (NajHgSbjO^). Precipitation may be 
 
THE METALS 
 
 117 
 
 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 : — 
 
 KaHaSbaO^ + 2 NaCl = ], Na2H2Sb207+ 2 KCl. 
 
 The solution must not be acid, for then decomposition of the 
 reagent results with the precipitation of amorphous pyroanti- 
 monic acid : — 
 
 KaHaSbaO^ + 2 HCl = | H^SbaO^ + 2 KCl. 
 
 All other metals, with the exception of K and NH^, must be 
 removed, for they too yield precipitates with the reagent. 
 
 2. Hydrochlorplatinic Acid, and Tartaric Acid, do not precipi- 
 tate sodium salts. 
 
 3. Sodium Cobaltic Nitrite does not give a precipitate with 
 sodium salts (distinction from K and NH4). 
 
 4. Heated just below a red heat, sodium compounds are not 
 volatilized (distinction from NH^). 
 
 5. Flame Reaction. — Sodium compounds color the bunsen 
 ilame yellow even when the quantity is very small. Bunsen 
 and Kirchhoff state that as small an amount as aooooo T °f ^ 
 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 
 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 NagHPO^. 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 NajHPO^ precipitates 
 the alkaline earth metals; and as the latter are usually present 
 
ii8 QUALITATIVE CHEMICAL ANALYSIS 
 
 in small amounts in the last filtrate by reason of the slight solu- 
 bility of their carbonates in NH^Cl, they must first be removed 
 before the test can be applied. This is accomplished by adding 
 a little (NH4)2S04 and (NH4)2C204, 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 moistetied 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 NHg by heating with an 
 excess of NaOH, must, therefore, always be made on a separate 
 portion of the original substance. 
 
THE METALS 
 
 SCHEME V 
 
 119 
 
 Analysis of the Filtrate from Group IV 
 
 This will contain, besides Mg and the alkalies, traces of the 
 alkaline earths which were dissolved by NH^Cl. Divide into 
 two unequal portions. 
 
 / In ^ test for Mg. Add a few drops of (NH4)2S04 and Qilli)^Cf>^, 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 NH^OH to alkaline reaction and then 
 add Na2HP04. Shake vigorously and allow to stand several minutes. A 
 white cryst. ppt. which is soluble in acetic acid is NH4MgP04. 
 
 In \ 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 NH4CI are given off, being careful to keep the dish beloiv 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 NH^Cl (2) is given off. Cool. Transfer a small portion of the 
 residue to a watch glass, moisten with a drop of HCl, 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 i cc. of Na3Co(N02)g, 
 warm, and allow to stand for several minutes. A yellow ppt. confirms the 
 presence of K. 
 
 Test for NH^. 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 crz^wfl/ solution or 0.2 — 0.5 g. of the original solid substance 
 contained in a small beaker, add NaOH till the resulting mixture, after 
 thorough stirring, is decidedly alkaline. Heat with stirring (;). 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. 
 
120 QUALITATIVE CHEMICAL ANALYSIS 
 
 NOTES ON SCHEME V 
 
 1. Na2HP04 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 AIPO4 sometimes separates on adding Na2HP04, but the 
 latter is easily distinguished from NH4MgP04 by reason of its insolubility in 
 acetic add. In concentrating the solution to 5 cc, A1(0H)3 may separate 
 out ; in that case, it should be filtered off before adding the Na2HP04. 
 
 Should a precipitate of doubtfid form be obtained with Na2HP04 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 expeUing 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)2C08. 
 
 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 precipi- 
 tate. 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 KCl. A 
 small amount of black carbonaceous matter (due to carbonization of small 
 quantities of pyridine which is usually present in the NH4OH) is left behind 
 along with the chlorides of K and Na. This is, however, removed later 
 when the residue is treated with water and filtered before making the pre- 
 cipitation test for K. 
 
 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. 
 
 If the original material is a solid instead of a solution, it is not necessary 
 to get it into solution in order to test for NH4. To a portion of the solid 
 substance in a beaker, add NaOH in slight excess, forming a paste, and heat 
 gently with stirring. The odor of NH, evolved will prove the presence of 
 NH4 compounds. 
 
THE METALS 
 
 121 
 
 ANALYSIS OF UNKNOWN SOLUTIONS FOR ALL GROUPS 
 To 25 cc. of solution add HCl 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, 
 BijSs, CuS, 
 CdS, AS2S3, 
 SbaSj, SnS. 
 Wash and 
 treat with 
 (NHJjS^. 
 Filter. 
 Analyze 
 residue ac- 
 cording to 
 Scheme 
 11. A. 
 Analyze 
 filtrate ac- 
 cording to 
 Scheme 
 II. B. ' 
 
 Filtrate : Boil to expel H2S before proceeding 
 with the analysis of Group II. ppt. Add NH4CI, 
 make alkaline with NH4OH, and completely ppt. 
 with H2S or (NH4)2S ; filter. 
 
 Residue : 
 Al (0H)3, 
 Cr(0H)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)2C03, and filter. 
 
 Residue : 
 BaCOj, SrCOs 
 + CaCOj. 
 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 AgNOg. 
 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 BaClg from solu- 
 tions acid with HCl, viz., sulphuric acid (H2SO4) and hydrofluo- 
 silicic acid (H2SiFg); {d) Acids precipitated by BaCl2 from 
 neutral solutions only, viz., carbonic acid (H2CO3), sulphurous 
 acid (H2SO3), thiosulphuric acid (H2S2O3), phosphoric acid 
 (H3PO4), hydrofluoric acid (HF), oxalic acid (HgCjO^), boric 
 acid (H3BO3),* silicic acid (H2Si03),t tartaric acid (H2C4H4O5), 
 arsenic acid (HgAsO^), arsenious acid (HgAsOg), and chromic 
 acid (HaCrOJ. 
 
 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 AgNOg from solutions acid with 
 nitric acid. These acids follow : — 
 
 Hydrochloric acid (HCl), hydrobromic acid (HBr), hydriodic 
 acid (HI), hydrocyanic acid (HCN), hydrogen sulphide (HgS), 
 hydroferrocyanic acid [H4Fe(CN)6], hydroferricyanic acid 
 [H3Fe(CN)g], thiocyanic acid (HSCN), and nitrous acid 
 (HNO2); the last is, however, only precipitated from moderately 
 
 * Orthoboric acid (HsBOs) is here considered as representative of the acids of 
 boron. 
 
 t Metasilicic acid (HjSiOs) is taken as the type of the many silicic acids. 
 
 123 
 
124 QUALITATIVE CHEMICAL ANALYSIS 
 
 concentrated solutions and is therefore also placed in the next 
 group. 
 
 Group III. includes those acids that are not precipitated by 
 either BaClj or AgNOg, viz., nitric acid (HNOg), nitrous acid 
 (HNO2), chloric acid (HCIO3), and acetic acid (HCgHgO-t). 
 
 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 SO^ radical is united, some 
 resisting decomposition, while others are readily decomposed, 
 giving off SO3 or SOg, 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 BaS04,* insoluble in 
 water and in dilute acids even on boiling. From dilute solutions 
 a precipitate separates only on standing. Dilute HCl or HNO3 
 
 * From HCl solutions BaCla may also precipitate BaSe04 and BaSiFe; the former 
 is readily distinguished from BaS04 and BaSiFs by the fact that on boiling it with 
 concentrated HQ, chlorine is given off: — 
 
 BaSe04 + 4 HCl = BaQg -I- HaSeOs + 4. CI2 + H2O. 
 
 BaSiFe is easily recognized by its readiness to undergo decomposition when 
 heated with concentrated H2SO4 : — 
 
 BaSiFe + H2SO4 = BaS04 -f-fa HF + >^ SiFi. 
 
THE ACIDS 125 
 
 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 BaClg 
 or Ba(N03)2 ™^y ^^ obtained ; these are, however, easily dis- 
 tinguished from BaS04 by their ready solution on diluting. For 
 further properties of BaSO^, see reaction 4 under Barium. 
 
 2. Lead Acetate produces a white precipitate of PbS04, 
 soluble in a hot concentrated solution of ammonium acetate or 
 tartrate. 
 
 3. Mixed with NajCOg 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 AggS will be produced : — 
 
 NagS -I- 2 Ag 4- H2O -1- O = AggS -t- 2 NaOH. 
 
 This test is also given by other sulphur compounds. 
 
 Fluosilicates 
 
 Hydrofluosilicic acid (HgSiFg) is formed by the action of sili- 
 con tetrafluoride (SiF^) on water : — 
 
 3 SiF4 -f 4 H2O = 2 H^SiFe + 1 H^SiO^. 
 
 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 HjSiFg, decomposition sets 
 in, according to the equation — 
 
 H2SiF6 = fSiF4 + f2HF. 
 
 With the exception of the potassium and barium salts, nearly 
 all fluosilicates are soluble in water. 
 
 BaSiFg, formed by adding BaClg to a solution of a fluosilicate, 
 is a white, crystalline, insoluble substance. Its solubility at 
 17° C. is I part in 3700 parts of water. As it is sparingly solu- 
 ble in HCl, it can be precipitated by BaClj from solutions con- 
 taining this acid. For a method of distinguishing it from 
 BaS04, see footnote, page 1 24. 
 
126 QUALITATIVE CHEMICAL ANALYSIS 
 
 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 carboi^i dioxide 
 and water : — 
 
 2 NaHCOg = Na2C08+ | H2O + f COg. 
 
 1. 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 : — 
 
 CaCOg -I- 2 H CI = CaCl2 -f- HgO -|- f COg ; 
 CO2 + Ca(0H)2 = iCaCOg -I- 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 BaCOg or CaCOg. 
 The precipitate is soluble in carbonic acid as well as in all other 
 acids with the exception of HCN : — 
 
 CaCOg + H2CO3 = Ca(HC03)2. 
 
 From the solution of dicarbonate of calcium, CaCOg reprecipi- 
 tates on boiling : — 
 
 Ca(HC0g)2 = iCaCOg -I- HjO + 't>C02. 
 
 3. Silver Nitrate precipitates white silver carbonate (AggCOg), 
 which, on boiling, changes to brown silver oxide (AggO) : — 
 
 Ag2COg = j,Ag20-HtC02. 
 
THE ACIDS 127 
 
 Sulphites 
 
 The aqueous solution of sulphur dioxide, known as sulphur- 
 ous acid, is a weak dibasic acid. When boiled it decomposes, 
 giving off SOj, which ^ay 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 solutions, 
 readily oxidize on exposure to air, forming the corresponding 
 sulphates. 
 
 1. Barium Chloride precipitates from neutral solutions white 
 barium sulphite (BaSOg), readily soluble in HCl and HNO3. 
 In practice, a residue of BaS04 remains, due to the presence of 
 a small amount of sulphate originally present in the sulphite or 
 produced by tke .subsequent oxidation of the sulphite. If the 
 BaSO^ is filtered off, the clear filtrate may be shown to contain 
 sulphurofis acidi^'by adding a little bromine or concentrated 
 HNO3 and boiling, when a white precipitate of BaS04 will 
 form. The bromine or nitric acid oxidizes the sulphurous to 
 sulphuric acid, and the latter at once yields with the BaClg 
 present a precipitate of BaSO^. 
 
 Free sulphurous acid is not precipitated by BaCLj. 
 
 2. Hydrogen Sulphide, when passed into a solution of sul- 
 phurous acid, or a solution of a sulphite, acid with HCl, causes a 
 separation of sulphur with the formation at the same time of 
 pentathionic acid : — 
 
 5 H2S + 5 H2SO3 = H^SgOe + 1 5 S -f 9 H^O. 
 
 3. Dilute H2SO4 or HCl, when added to a sulphite, decomposes 
 it with the evolution of SO2 : — 
 
 NagSOg + 2 HCl = 2 NaCl -H HgO + \ SOg. 
 
 4. Silver Nitrate precipitates from neutral solutions white 
 AggSOg, which, on boihng, is decomposed, with the separation 
 of gray metallic silver : — 
 
 AgaSOg -f H2O = I 2 Ag -I- H2SO4. 
 
128 QUALITATIVE CHEMICAL ANALYSIS 
 
 5. Iodine Solutions are bleached by sulphurous acid ; this is 
 due to the reduction of iodide to hydriodic acid : — 
 
 H2SO3 + I2 + H20:;t2 HI + H2SO4. 
 
 6. Potassium Permanganate solution, acid with sulphuric acid, 
 is also decolorized by sulphurous acid : — 
 
 2 KMn04 + 5 H2SO3 = K2SO4 + 2 MnSO^^ + 2 HjSO^ + 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 KMnO^ 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 : — •^ ;-,. 
 
 K2Cr207 + 3 H2SO3 + H2SO4 = K2SO4 + Cr2(S04)3 + 4 H2O. 
 
 8. Stannous Chloride, when added to sulphurous acid or to 
 a hydrochloric acid solution of a sulphite, and the mixture 
 heated, reduces the sulphurous acid to HgS, which, after some 
 time, will precipitate the tin as stannic sulphide : — 
 
 3 SnCl2 + H2SO3 + 6 HCl = f H2S + 3 SnCl^ + 3 H2O ; 
 SnCl4 + 2 H2S = i SnS2 + 4 HCl. 
 
 9. Sulphurous acid reduces arsenic acid (H3ASO4) to arseni- 
 ous acid (HgAsOg); the action 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 HgO. The chief 
 thiosulphate is the sodium salt (Na2S203), used extensively in 
 photography because of its property of dissolving the halides of 
 silver. 
 
THE ACIDS 129 
 
 1. Dilute Hydrochloric Acid decomposes all thiosulphates with 
 the evolution of sulphur dioxide and the separation of sulphur : — 
 
 NajSaOs + 2 HCl = 2 NaCl + H2O + 1 SOj + 1 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 AgNOa + NagSaOs = | Ag^Sfi^ + 2 NaNOg ; 
 AgaSgOg + NajSaOg = 2 NaAgSaOg. 
 
 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 AggS : — 
 
 AggSaOg + H2O = I Ag^S + 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 : — 
 
 HP03+H20 = H3P04. 
 
 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 Na2HP04 = Na^PgO^ + \ HjO ; 
 2 NH^MgPOi = t HjO + \2 NH3 + MgaPaOy. 
 
I30 QUALITATIVE CHEMICAL ANALYSIS 
 
 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. 
 
 1. Barium Chloride (BaClg) precipitates from neutral solutions 
 of orthophosphates white BaHPO^, soluble in acetic, hydro- 
 chloric, and nitric acids ; from the acid solution, ammonia pre- 
 cipitates the tertiary phosphate : — 
 
 BaCla + Na2HP04 = | BaHPO^ + 2 NaCl; 
 BaHPO^-l- 2 HCl = BaCla + H3PO4; 
 H3PO,-H3NH3 = (NH,)3P04; 
 2 (NH4)3P04 + 3 BaCla = | Ba3(P04)2 -|- 6 NH^Cl. 
 
 2. Lead Acetate [Pb(C2H302)2] precipitates white lead phos- 
 phate, practically insoluble in acetic acid though soluble in 
 nitric acid : — 
 
 3 Pb(C2H302)2 + 2 NajHPO^ = 
 
 I Pb3(POj2 + 2 HC2H3O2 -I- 4 NaC2H302. 
 
 3. Silver Nitrate precipitates only from strictly neutral solu- 
 tions yellow silver phosphate (AgjPO^), soluble in mineral acids ; * 
 it is also soluble in acetic acid and in ammonium hydroxide. 
 
 2 NaaHPO^ -|- 3 AgNOg = ,|, Ag3P04 + 3 NaNOg -1- NaHaPOi- 
 
 4. Magnesium mixture (MgCl2 -I- N H4CI -f- N H^O H in slight ex- 
 cess) precipitates from solutions of orthophosphates, white crystal- 
 line ammonium magnesium phosphate (NH4MgP04-6 H2O): — 
 
 MgCl2 + NagHPO^ -I- NH3 = ,|, NH4MgP04 + 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 
 AIPO4, 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 HCl, a white precipitate of AgCl is obtained. 
 
THE ACIDS 131 
 
 separates only on standing ; on ignition, it yields magnesium 
 pyrophosphate (Mg2P207). 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 (Ag3As04). 
 
 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)3P04, 12 M0O3J. 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 more slowly and requires a higher tempera- 
 ture for its complete precipitation. 
 
 This reaction affords a means of quantitatively separating 
 phosphoric acid from the metals with which it may be combined. 
 
 6. Ferric Chloride, when added to a soluble phosphate not too 
 strongly acid, yields a buff-colored precipitate of ferric phos- 
 phate (FePO^) : — 
 
 FeCls + NagHPO^ = | FePO^ + 2 NaCl -I- HCl. 
 
 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 FePO^ is 
 insoluble, and, as a consequence, precipitation is rendered com- 
 plete : — 
 
 FeCl3-|-Na2HP04-FNaC2H302 = 3 NaCl+ jFePO^-l-HQHgOa. 
 
 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 
 
132 QUALITATIVE CHEMICAL ANALYSIS 
 
 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 z.% 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) dissolve 
 readily. 
 
 I. Concentrated HgSO^ decomposes most fluorides with the 
 liberation of hydrogen fluoride : — 
 
 (i) CaF2 + H2S04 = |CaS04 + f2HF. 
 
 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) NaaSigO^ • CaSigO^ -|- 28 HF = 
 
 Soda glass. 
 
 2NaF-|-|CaF2 + f6SiF4-l- 14H2O. 
 
 When brought in contact with water, silicon tetrafluoride reacts 
 according to the following equation : — 
 
 (3) 3 SiF, -I- 4 H2O = 2 H2SiF6 -H ,|,H4Si04. 
 
 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 HgSO^ and then hold in the escaping vapors a 
 
THE ACIDS 
 
 133 
 
 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 (SiOj) before the treatment 
 with concentrated HjSO^. 
 
 The reactions are as follows : — 
 
 2 CaFj + 2 H2SO4 = |2 CaS04+ t4 HF; 
 4HF + Si02 = tSiF4 + 2 Hfi; 
 3 SiF^ + 4 H2O = 2 HaSiFg + ^H^SiO^. 
 
 2. The Etching Test. This test, as explained above, is based 
 on the property possessed by hydrofluoric acid to dissolve SiOg 
 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 parafifine 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- 
 afifine 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 evi- 
 dent that this test is inapplicable in the presence of silicates. 
 Anhydrous 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 
 (CaFj), soluble with difficulty in HCl and HNO3, but practi- 
 cally insoluble in acetic acid. From the acid solution of calcium 
 fluoride, ammonium hydroxide does not reprecipitate the fluo- 
 ride, because of the solubility of CaFg in ammonium salts. 
 
 4. Fusion with Sodium Carbonate only partially decomposes 
 CaFj. In the presence of silica, however, the decomposition 
 may be rendered complete. 
 
134 QUALITATIVE CHEMICAL ANALYSIS 
 
 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. 
 
 1 . Barium Chloride precipitates white barium oxalate (BaCjO^), 
 soluble in oxalic and acetic acids. 
 
 2. Calcium Sulphate or Calcium Chloride precipitates white 
 crystalline CaC204, insoluble in oxalic and acetic acids, and in 
 ammonium oxalate, but soluble in hydrochloric acid. As 
 CaCgO^ is one of the most insoluble of the oxalates, CaSO^ 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 +tC02. 
 
 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 COg and water: — 
 
 2 KMnO^ + 3 H2SO4 + 5 H2C2O4 = 
 
 K2SO4 + 2 MnSO^ + 8 H2O + 1 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. 
 
THE ACIDS 135 
 
 An oxalate may be decomposed by evaporating it with a 
 mixture of HgSO^ and HNO3 until SO3 fumes are given off; 
 any organic matter present will also be destroyed at the same 
 time. 
 
 Borates 
 
 Boron forms three acids, namely, orthoboric acid (H3BO3), 
 metaboric acid (HBOj), 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. 
 
 1. Turmeric Paper Test. If a piece of turmeric paper is 
 dipped into a solution of a borate slightly acid with HCl, 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 : — 
 
 NagB^O^ + BaClg + 3 Hfi = 2 NaCl -f- 2 H3BO3 -t- 1 Ba(B02)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 (AgBOj), soluble in ammonium hydrox- 
 ide and nitric acid; warming converts it into brown silver oxide 
 (AgP). 
 
 5. Flame Tests. Free boric acid and some of its volatile 
 compounds, e.g., the fluoride (BF3), its esters, as (0113)3603 and 
 
136 QUALITATIVE CHEMICAL ANALYSIS 
 
 (€2115)3603, 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 arid the above reactions is carried out 
 by making a paste of the substance with concentrated HjSO^, 
 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 (BFg) is formed, which, when 
 brought into the bunsen flame, colors it green. 
 
 7. Concentrated HgSO^ 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 :;t 3 H2O + f (C2H5)3B03. 
 
 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 -I- Na^BiO^ = 2 NaBOa + €0(602)3. 
 
 The cobalt metaborate is blue. 
 
THE ACIDS 137 
 
 Silicates 
 
 Silica (SiOj), 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 (SiF^) : — 
 
 SiOa + 4 HF = f SiF^ + 2 H2O. 
 
 To expel silica completely with HF, the presence of concen- 
 trated HgSO^is necessary. When silica is fused with sodium 
 carbonate and the mass is extracted with water, a solution of 
 sodium silicate is obtained : — 
 
 SiOg + NajCOg = NajSiOg + f COg. 
 
 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. 
 
 I. If to a solution of sodium or potassium silicate, hydro- 
 chloric or nitric acid is added until an acid reaction results, part 
 of the siUcic acid will separate out as a gelatinous precipitate, 
 while the rest will remain in solution in the form of a hydrosol : — 
 
 NagSiOg + 2 HCl = 2 NaCl + | HgSiOg. 
 
 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 
 
138 QUALITATIVE CHEMICAL ANALYSIS 
 
 of great analytical importance, since it enables the analyst, 
 early in 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- 
 sihcic acid : — 
 
 NaaSiOg+a NH4CI+2 HgO = |H2Si03+2 NaCl + 2 NH4OH ; 
 Na2Si03+(NH4)2C03 + 2H20 = |H2Si03+Na2C03+2NH40H. 
 
 3. Barium Chloride precipitates white BaSiOg, soluble in 
 acids. 
 
 4. Silver Nitrate precipitates yellow Ag2Si03, 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 NH^Cl, a canary-yellow 
 precipitate is produced. 
 
 6. Sodium Metaphosphate Bead Test. When a silicate is 
 fused in a metaphosphate bead prepared from microcosmic salt 
 (NH^NaHPO^), 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 : — 
 
 CaSiOg + NaPOg = CaNaPO^ + Si02. 
 
 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. 
 
TEE ACIDS 139 
 
 concentrated HCl * 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 120° C. until no more acid fumes 
 are given off. The dried residue is first treated with 10 cc. of 
 concentrated HCl 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 HjSO^ 
 and about 5 cc. of hydrofluoric acid. It is then gently heated 
 to dryness and finally ignited. By this treatjnent all the silicic 
 acid will be driven off as SiF^. Any residue is tre&,ted accord- 
 ing to page 197 for insoluble substances. f 
 
 {b) 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 ten 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, FeaOs, or BaS04, or of all three substances. 
 
I40 QUALITATIVE CHEMICAL ANALYSIS 
 
 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 120° 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 167). 
 
 (c) 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 2 cc. of concentrated HjSO^, and evaporate 
 on a hot plate until SO3 fumes are evolved. 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 
 SO,. 
 
 2. Silver Nitrate does not react with free tartaric acid, but 
 from solutions of normal tartrates it precipitates white curdy 
 silver tartrate (AgjC^H^Og), 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. 
 
THE ACIDS 
 
 141 
 
 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 (CaC^H^Og), soluble in acids, including acetic acid; the 
 precipitate is soluble in cold caustic soda free from carbonate, 
 from which, on boiling, CaC4H406 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 CaClg 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 74). 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 
 
142 QUALITATIVE CHEMICAL ANALYSIS 
 
 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 ^6 and 74. 
 
 1. Barium Chloride precipitates yellowish white barium chro- 
 mate (BaCr04), 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 (PbCrO^) : — 
 
 Pb(C2H302)2 + NajCrO^ = \ PbCrO^-f 2 NaCaHgOa. 
 
 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 AgNOg + NagCrO^ = | AgaCrO^ + 2 NaNOg. 
 
 Silver chromate is readily soluble in nitric acid and ammo- 
 nium hydroxide. From slightly acid concentrated solutions a 
 reddish brown crystalline precipitate of AgjCrgO^ is formed 
 which possesses about the same solubilities as AggCrO^: — 
 
 2 AgNOg + KjCraOy = ],Kz^(Zx^O^ -V 2 KNO3. 
 
 The chromates of silver are readily converted into the chloride 
 by treatment with HCl. 
 
 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 CSg, 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 (CgHgOH), when added to a chromate solu- 
 tion acidified with HCl or H2SO4, and the mixture is boiled. 
 
TEE ACIDS 
 
 143 
 
 causes a reduction of the chr ornate 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 : — 
 
 KgCraO^ + 3 CgHgOH + 8 HCl = 
 
 2 CrClg + 7 H2O + 2 KCl + f 3 CH3CHO. 
 
 This method is frequently used to reduce chromates to chromic 
 salts before treating a solution containing a chromate with HgS. 
 For the equation for the reduction of chromates by concentrated 
 HCl, see page 76. 
 
 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 CrOg • HjOj, which, however, is very unstable. 
 One part of K2Cr04 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. BaClg 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 
 
144 QUALITATIVE CHEMICAL ANALYSIS 
 
 chlorides dissolve in aqua regia with the exception of silver 
 chloride. Fusion with sodium carbonate transposes * all insolu- 
 ble chlorides : — 
 
 2 AgCl + Na2C03 = ,|, 2 Ag + 2 NaCl + f COj + f O. 
 
 The chlorides of barium, sodium, and potassium are practi- 
 cally insoluble in concentrated hydrochloric acid. 
 
 1. Silver Nitrate precipitates from nitric acid solutions, whiter 
 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 ZnClg : — 
 
 2 AgCl -}- Zn = ZnClg -I- 1 2 Ag. 
 
 2. Lead Acetate precipitates white lead chloride (PbClj). For 
 its properties, see reaction i under Lead. 
 
 3. Concentrated HgSO^ and Mn02, 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 : — 
 
 Mn02 4-2 NaCl +2 H2SO4 = MnSO^ + NagSO^ + 2 HgO + fClj. 
 
 4. Potassium Dichromate and Concentrated Sulphuric Acid. 
 If a dry mixture of a chloride and powdered KgCrjO^ is treated 
 with concentrated H2SO4 and heated gently, chromyl chloride 
 (CrOgClg), a reddish brown gas, will be evolved : — 
 
 KjCraOy +4 NaCl -H 6 H2SO4 = 
 
 f 2 CrOgCla + 2 KHSO4 -f 4^NaHS04 + 3 HgO. 
 
 * 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. 
 
THE ACIDS 
 
 145 
 
 If the gas IS absorbed by ammonium hydroxide, the latter 
 will be colored yellow owing to the formation of ammonium 
 chromate : — 
 
 CrOaClg + 4 NH^OH = (NH4)2Cr04 + 2 NH4CI + 2 H^O. 
 
 The presence of chromic acid in the ammonium hydroxide 
 solution may be shown by rendering it acid with acetic acid and 
 adding Pb(C2H302)2, when a yellow precipitate of PbCrO^ 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 reac- 
 tion with the bromide is as follows : — 
 
 6 KBr + KaCrjO^ + 1 1 HjSO^ = 
 
 t 3 Br2+ 8 KHSO4 + Cr2(S04)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 NaCOg, 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 
 Na2S203. 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 
 
146 QUALITATIVE CHEMICAL ANALYSIS 
 
 a few drops of carbon disulphide.(CS2) or chloroform (CHCI3), 
 in which bromine is soluble, are added, and the mixture is 
 shaken, the CSg or CHCI3 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 + CI = KCl + 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 HgSO^, when added to a bromide, and the 
 mixture is heated, causes the liberation of bromine with hydro- 
 bromic acid. 
 
 5. Concentrated HgSO^ and MnOj, when added to a bromide, 
 and the mixture is heated, causes 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 abiUty 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 -1- 8 HNO3 = f 3 Bra -I- 6 KNO3 -|- f 2 NO -f 4 Hfi. 
 
 Iodides 
 Solubilities. With the exception of Hglj, Hgl, Agl, Cul, and 
 the sparingly soluble Pblj, all iodides are soluble in water. 
 
THE ACIDS 
 
 147 
 
 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. 
 
 1. 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 HgSO^, 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 CI 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 Uberation of iodine : — 
 
 KI + CI = KCl + I. 
 
 The free iodine may be recognized by shaking the mixture 
 with a few drops of CSg or CHCI3; 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 -f 6 H2O + 5 CI2 = 10 HCl -I- 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 CSg violet or starch paste blue. 
 
148 QUALITATIVE CHEMICAL ANALYSIS 
 
 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, yield 
 a dirty brown precipitate of cuprous iodide (Cul) ; at the same 
 time iodine is set free : — 
 
 2 CuSO^ + 4 KI = 2 K2SO4 + 1 2 Cul + 1 2 I. 
 
 7. Ferric Salts also liberate iodine from iodides, being at the 
 same time reduced to the ferrous state : — 
 
 Fe2(S04)3 + 2 KI = 1 12 + 2 FeSO^ + K2SO4. 
 
 8. Potassium Dichromate, when added to an iodide solution 
 acid with sulphuric acid, causes iodine to be set free : — 
 
 In all of the above cases the liberated iodine may be detected 
 by shaking the mixture with one cc. of CHCI3 or CSj, which 
 acquires a violet color. 
 
 9. Mercuric Chloride, when added to an iodide solution, pre- 
 cipitates scarlet mercuric iodide (Hglj), soluble in an excess of 
 alkali iodide : — 
 
 HgCla -f 2 KI = I Hgl2 + 2 KCl ; 
 Hgl2-f2KI = K2HgI,. 
 
 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 4- O = NaCNO. 
 
THE ACIDS 
 
 149 
 
 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 + t(CN)2; 
 Pb(CN)2=Pb + 2C + tN2. 
 
 Caution : All tests for cyanides which involve the evolution of 
 a gas should he 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 HjS with the precipitation 
 of mercuric sulphide and the production of HCN. 
 
 1. 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 + 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 : — 
 
 2AgCN=|2Ag-|-t(CN)2. 
 
 2. Sulphuric Acid, dilute, when added to a solution of a cya- 
 nide [except Hg(CN)2J, 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 -I- 2 H2SO4 -I- 2 H2O = C0SO4 + (NH4)aS04 -I- f 2 CO ; 
 K4Fe(CN)6 -}- 6 HjSO^ + 6 HgO = 
 
 FeSO^ -I- 2 K2SO4 H- 3 (NH4)2S04 + 1 6 CO. 
 
ISO QUALITATIVE CHEMICAL ANALYSIS 
 
 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 : — 
 
 (i) FeSO^ + 2 NaOH = ^ Fe(0H)2 + NajSO^ ; 
 
 (2) Fe(0H)2 + 2 KCN = ,|, Fe(CN)2 + 2 KOH ; 
 
 (3) Fe(CN)2 + 4 KCN = K,Fe(CN)e ; 
 
 (4) 3 K4Fe(CN)6 + 4 FeClg = | Fe4[Fe(CN)6]3 + 12 KCl. 
 
 4. Formation of Thiocyanate. To a solution of an alkali cya- 
 nide add a little (NH^)2S3. and evaporate the solution on a 
 water bath to dryness. The residue, which will now consist of 
 ammonium thiocyanate (NH^SCN), is treated with one or two 
 drops of dilute HCl, 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)3j. 
 
 If it is desired to detect hydrocyanic acid evolved from an 
 insoluble cyanide on treatment with hot dilute HgSO^, cover the 
 test tube containing the mixture with a crucible cover, on the 
 under side of which has been placed a drop of (NH4)2Sj., 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)g] is a colorless crystalline 
 solid, easily soluble in water; on exposure to air the solution 
 
THE ACIDS 151 
 
 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. 
 
 1. Silver Nitrate precipitates white silver ferrocyanide, insol- 
 uble in dilute nitric acid and in ammonium hydroxide, but soluble 
 in potassium cyanide solution : — 
 
 4 AgNOg + KiFe(CN)e = | Ag^FeCCN)^ -f- 4 KNOg. 
 
 On ignition, the precipitate is ' decomposed with the separa- 
 tion of silver and evolution of cyanogen gas : — 
 
 Ag,Fe(CN)6 = 1 4 Ag + 1 2(CN)2 + I FeQ + f Nj. 
 
 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 : — 
 
 2K,Fe(CN)s + 3H2S04 = 
 
 I K2Fe[Fe(CN)6J + 3 KjSO^ -f- f 6 HCN. 
 
 Concentrated sulphuric acid, when heated, completely decom- 
 poses ferrocyanides with the evolution of carbon monoxide : — 
 
 K4Fe(CN)6 + 6 H2SO4 + 6 Hfi = 
 
 FeSO^ -f 2 K2SO4 -I- 3(NH,)2S04 -f- 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)g]3 + 12 NaOH = 4 Fe(0H)3 -|- 3 Na4Fe(CN)e. 
 
 The residue is dissolved in acid and the metal is tested for in 
 the solution obtained. The filtrate is acidified with HCl and 
 tested for the ferrocyanogen radical by adding ferric chloride. 
 
152 QUALITATIVE CHEMICAL ANALYSIS 
 
 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 HgS. 
 
 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 HjO = 
 
 Fe2(S04)3 -I- 3 K2SO4 -1- 6 (NH4)2S04 + f 12 CO. 
 
 1. Silver Nitrate gives with ferricyanides a reddish brown pre- 
 cipitate of silver ferricyanide, 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 TurnbuU's 
 blue, Fe8[Fe(CN)gj2, insoluble in acids. 
 
 3. Reducing Agents, such as HjS, SOj, and HI, readily reduce 
 ferricyanides to ferrocyanides in alkaline solutions : — 
 
 2 K3Fe(CN)e -1- KaS = 2 K4Fe(CN)g + ^S. 
 
 Thiocyanates 
 
 Alkali thiocyanates are readily prepared by heating the cy- 
 anide with sulphur : — 
 
 KCN-1-S=KCNS; 
 
TEE ACIDS 
 
 153 
 
 or by heating an alkali cyanide or hydrocyanic acid with ammo- 
 nium polysulphide : — 
 
 KCN + (NH4)2S, = KCNS + {^B.^\S^^. 
 
 Solubilities. Nearly all the thiocyanates are soluble in water 
 with the exception of those of silver, mercury, copper, and gold. 
 
 1. Silver Nitrate precipitates white, curdy silver thiocyanate, 
 insoluble in dilute nitric acid, but soluble with difficulty 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 + = H20 + |S. 
 
 HjS reacts with bases forming hydrosulphides and sulphides, 
 which, if HgS is looked upon as a dibasic acid, may be regarded 
 as acid and normal sulphides : — 
 
 NaOH + HgS = NaSH + H2O ; 
 2 NaOH + HaS = NagS + 2 HjO. 
 
 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. 
 
154 QUALITATIVE CHEMICAL ANALYSIS 
 
 1. Hydrochloric Acid, when moderately strong (i : i), decom- 
 poses nearly all sulphides with the evolution of HgS ; the latter 
 may be recognized by its odor and by its property of turning 
 lead acetate paper black : — 
 
 ZnS + 2 HCl = ZnCla + tHjS. 
 
 The few sulphides undecomposed by HCl are attacked by a 
 mixture of zinc and HCl with the liberation of HjS. 
 
 2. Silver Nitrate precipitates from solutions of sulphides or 
 hydrogen sulphide, black silver sulphide (AgjS), insoluble in 
 cold but soluble in hot dilute nitric acid; it is insoluble in 
 ammonium hydroxide. 
 
 3. Oxidizing Solvents, such as concentrated HNOg, agua regia, 
 HCl + KCIO3, 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 CI2 + 4 H2O = HaSOi + 6 HCl ; 
 3 PbS + 8 HN03(hot and concentrated) = 
 
 1 3 PbS04'+ f 8 NO + 4 H2O. 
 
 4. Sodium Nitroprusside [Na2(N0)Fe(CN)g] 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 dichromatCj potassium permanganate, ferric salts, etc., on 
 HgS, see page 61. 
 
 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. 
 
THE ACIDS 155 
 
 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. 
 
 1. 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 (NOj) * : — 
 
 (i) NaN02 + H2S04=NaHS04H-HN02; 
 
 (2) 3 HNO2 = HNO3 -f f 2 NO -1- H2O ; 
 
 (3) N0 + = N02. 
 
 Nitrites are also decomposed by acetic acid with gentle heat- 
 ing ; the NO2 given off may be recognized by its turning starch- 
 iodide paper blue. 
 
 2. Silver Nitrate precipitates, from solutions of nitrites which 
 are not too dilute, white silver nitrite (AgN02), 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 potassium nitrite, a yellow pre- 
 cipitate of potassium cobaltic nitrite [K3Co(N02)8j. 
 
 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 CSj or CHCI3 violet. 
 
 5 . Potassium Permanganate solution, when warmed, is bleached 
 by a solution of a nitrite slightly acid with dilute sulphuric 
 acid : — 
 
 2 KMnO^ 4- 5 HNO2 + 3 HjSO^ = 
 
 K2SO4 -1- 2 MnSO^ -I- 5 HNO3 + 3 H2O. 
 
 * The older nomenclatuie is applied to the oxides of nitrogen as being more 
 distinctive. 
 
156 QUALITATIVE CHEMICAL ANALYSIS 
 
 GROUP III 
 The acids of this group are not precipitated by either AgNO, 
 or BaClg. 
 
 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 NaCjHgOa = NajCOg + 1 {CUsXCO. 
 
 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. 
 
 1. 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 acohol, 
 amyl acetate, having an odor resembling pear essence, will be 
 produced : — 
 
 CH3 • COO[H + HOjCoHg :^ CH3 • COOCaHg -|- HjO. 
 
 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 (AgCgHgOa), sparingly soluble in cold water (1.04 
 
THE ACIDS 157 
 
 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 : — 
 
 ¥&{C^Ufi^\ + H.p:;>|Fe(OH)(C2H302)a + HQHgOa. 
 
 Nitrates 
 
 Solubilities. , The nitrates, with the exception of a few basic 
 nitrates, such as BiONOg, 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 : — 
 
 KN03 = KN02 + tO. 
 
 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 (NOj) : — 
 
 (i) NaN03-fH2S04 = NaHS04-l-HN03; 
 (2) 2 HNO3 = \2 NO2 + H2O -I- fO. 
 
 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 : ist, 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 
 
158 QUALITATIVE CHEMICAL ANALYdfS 
 
 in the excess of unoxidized FeS04 with the formation of a 
 brown unstable compound : — 
 
 (i) H2S04 + NaN03 = HN03 + NaHS04; 
 
 (2) 2 HNO3 + 6 FeS04 + 3 H2SO4 = 
 
 3Fe2(SOj3 + 4H20 + 2NO; 
 
 (3) 2 FeS04+ NO=(FeS04)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 incUned 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 colora- 
 tion will be produced in 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 HCl that has been recently 
 boiled, a few drops of a solution of indigo in sulphuric acid are 
 added, and the mixture is then boiled, the blue coloration will 
 persist, if the HCl 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 HCl like 
 KCIO3, 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 : — 
 
 (i) NaN03 + H2SO4 = NaHSO^ + HNO3 ; 
 
 (2) 3 Cu + 8 HN03= 3 Cu(N03)2 + t2 NO +4 HgO; 
 
 (3) N0-F-0 = N02. 
 
 5. Reduction to Ammonia. If a solution of a nitrate is made 
 strongly alkaline with NaOH, and a few pieces of aluminum or 
 
TEE ACIDS 
 
 IS9 
 
 zinc and iron filings are added, and the mixture is heated, am- 
 monia 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 : — 
 
 KC103 = KCl+t3 0. 
 AgC103 = AgCl+t3 0. 
 
 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 (ClOj), 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 ClOj : — 
 
 3 KCIO3 + 2 H2SO4 = KCIO4 + 2 KHSO4 + H2O + f 2 CIO2. 
 
 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 KCIO3. One should never look down 
 into a test tube, especially when performing this test. Equal 
 consideration is due one's neighbor. 
 
 2. Concentrated HCl, when added to a chlorate, is oxidized 
 with the production of chlorine and chlorine peroxide : — 
 
 KCIO3 -F 2 HCl = KCl + f CI + f CIO2 + HjO. 
 
i6o QUALITATIVE CHEMICAL ANALYSIS 
 
 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 + HCl. 
 
 ACID ANALYSIS 
 
 Preliminary 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, e.g., oxiding agents like KjCrgO^, HgAsO^, FeCl^ 
 cannot exist with reducing agents like sulphites and iodides. 
 It is also desirable that the first three of the following prelimi- 
 nary tests be carried out before commencing the systematic 
 search for acids. 
 
THE ACIDS 
 
 i6i 
 
 PRELIMINARY TESTING OF SOLIDS 
 
 I. Heat a small quantity of the substance in a tube closed at 
 one end. 
 
 Observation 
 
 Indication 
 
 Decrepitation. 
 
 j^>. f- — 
 
 NaCl, Pb(N03)2, K2SO4, zinc blende, 
 and other substances. 
 
 Carbonization, 
 
 accompanied by burnt 
 
 Organic matter, tartrates, and other 
 
 odor and formation of water. 
 
 organic acids and salts. 
 
 Water given off. 
 
 Water mechanically inclosed, water 
 
 
 
 of hydration, and hydroxide. 
 
 Gases given ofE 
 
 : — 
 
 
 
 O — kindles a spark ; 
 
 Chlorates, peroxides, certain oxides, 
 
 Colorless 
 
 
 nitrates, etc. 
 
 and 
 
 CO — burns with blue 
 
 Oxalates. 
 
 odorless. 
 
 ilame; 
 
 
 
 CO2 — turns limewater 
 
 Carbonates, oxalates, and organic 
 
 
 turbid. 
 
 matter. 
 
 
 N H3 — turns red litmus 
 
 NH4 salts and organic compounds 
 
 
 blue; 
 
 containing N. 
 
 
 SO2 — recognized by 
 
 Sulphur, sulphides, sulphites, and cer- 
 
 
 odor; 
 
 tain sulphates. 
 
 Colorless 
 
 (CN)2 — recognized 
 
 Cyanides. 
 
 with 
 
 by odor,* and burns 
 
 
 odor. 
 
 with reddish flame ; 
 
 
 
 H2S — recognized by 
 
 Moist sulphides. 
 
 
 odor; 
 
 
 
 Acetone — recognized 
 
 Acetates. 
 
 
 by odor. 
 
 
 
 NOa — reddish brown ; 
 
 Nitrates of heavy metals. 
 
 Colored 
 
 turns starch-KI 
 
 
 gases. 
 
 paper blue. 
 
 
 
 CI, Br, I — recognized 
 
 Chlorides, bromides, and iodides in 
 
 
 . by color and odor. 
 
 the presence of oxidizing agents. 
 
 A sublimate fa 
 
 nns: — 
 
 
 White. 
 
 
 NH4 salts, HgCl, HgClj, AS2O3, SbiQs, 
 and certain organic compounds. 
 
 Yellow. 
 
 
 AS2S3, HgO (accompanied by globules 
 ofHg). 
 
 • Smell cautiously. 
 
l62 
 
 QUALITATIVE CHEMICAL ANALYSIS 
 
 Observation 
 
 Indication 
 
 Yellow, becoming red when 
 
 Hgl,. 
 
 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). 
 
 PbCrO^. 
 
 Yellow (hot), white (cold). 
 
 ZnO. 
 
 2. Put a small quantity of the substance in a test tube, add a 
 little dilute HCl, and heat gently. 
 
 Observation 
 
 Indication 
 
 CO, 
 
 effervescence, turns limewater turbid. 
 
 Carbonates. 
 
 SO2 
 
 recognized by odor. 
 
 Sulphites. 
 
 SO2 
 
 accompanied by a separation of S. 
 
 Thiosulphates. 
 
 H^S 
 
 : recognized by odor and by lead acetate paper. 
 
 Sulphides. 
 
 NO, 
 
 : 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 fames are evolved which redden 
 litmus. 
 
 Br, I, mixed with HBr, SO2, and per- 
 haps HjS. 
 
 Halogen acids from their salts. 
 Iodides and bromides. 
 
 * Smell cautiously by fanning vapor with hand towards nose. 
 
THE ACIDS 
 
 163 
 
 Observation 
 
 Indicatioii 
 
 Chlorine — bleaching litmus. 
 
 Chloride and oxidizing agent together. 
 
 CrOjClj— (reddish brown). 
 
 Chromate and chloride together. 
 
 ClOj — yellow color of gas and 
 
 Chlorate. 
 
 H2SO4; gas explosive. 
 
 
 HF — etches glass ; yields SiF^, which 
 
 Fluoride. 
 
 turns " water bead " turbid. 
 
 
 HC2H3O2 — recognized by odor. 
 
 Acetate. 
 
 NO2 — recognized by odor and by 
 
 Nitrates, nitrites. 
 
 turning starch-KI paper blue. 
 
 
 SOj — recognized by odor. 
 
 Sulphite, thiosulphate, or reducing 
 
 
 agent acting on HjSO^. 
 
 SO2 — accompanied by blackening. 
 
 Organic matter or tartrate. 
 
 CO — (without blackening) recog- 
 
 Oxalate, cyanides, ferro- and ferri- 
 
 nized by burning with blue flame. 
 
 cyanides. 
 
 COj — turns limewater turbid. 
 
 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. 
 
 (b) 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 moistened with two drops 
 
 
 of very dilute Co(N03)2 solution 
 
 
 and heated again gives a 
 
 
 blue mass. 
 
 AI. 
 
 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), 
 
 
 near residue, and not volatile. 
 
 Sn. 
 
 Reddish yellow (hot), yellow (cold). 
 
 Pb. 
 
 Orange (hot), light yellow (cold). 
 
 Bi. 
 
 Reddish brown, cold and hot. 
 
 Cd. 
 
164 
 
 QUALITATIVE CHEMICAL ANALYSIS 
 
 5. Mix with anhydrous sodium carbonate and heat on char- 
 coal. 
 
 Observation 
 
 Indication 
 
 (a) Metallic globule forms without incrustation : — 
 
 — t 
 
 yellow, 
 
 Au. 
 
 red, 
 
 Cu. 
 
 white and malleable. 
 
 Ag. 
 
 (6) 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 Flams 
 
 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 HCl, 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. 
 
 Na. 
 
 Deep red. 
 
 Li or Sr. , 
 
 Reddish yellow. 
 
 Ca. 
 
 Green or greenish yellow. 
 
 Ba, Cu, or HjBOs. 
 
 Pale blue. 
 
 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 HQ, and introduce into the colorless bunsen flame. 
 
THE ACIDS 165 
 
 Method of Acid Analysis 
 
 The method employed for the detection of the acids is differ- 
 ent from that used in the systematic examination for metallic radi- 
 cals. 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 num- 
 ber of acids which must be looked for is very much reduced ; 
 first, from a knowledge of the solubilities and metallic content 
 of the substance ; and, second, by the results furnished by the 
 preliminary experiments just given. The reagents BaClj and 
 AgNOg, when properly applied, are valuable in that they give 
 indications of the presence or absence of whole groups ; e.£^., if to 
 a moderately concentrated and neutral solution of the substance, 
 BaCl^ or CaClj 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 HCl will dis- 
 close the presence or absence of the following acids : HjCOg, 
 H2SO3, H2S2O3, HjS, HCN, and HNO2. In the course of the 
 analysis for metals, HgAsOg, H3ASO4, and H2Cr04 will be de- 
 tected, ^i'or 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 
 solution," may be obtained by boiling the finely powdered sub- 
 stance with an excess of NagCOg solution, with constant stirring, 
 for several minutes (i). If ammonia is given off, boil with the 
 addition of more Na2C03 until no more of this gas is evolved, 
 and then filter. 
 
 * The solution should also contain no considerable quantity of NH4 salts, else, 
 borates, fluorides, and tartrates may not precipitate. 
 
i66 
 
 QUALITATIVE CHEMICAL ANALYSIS 
 
 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 
 JHjPOi and SiOj. 
 
 Acidify with HNO3, 
 evaporate to dry- 
 ness, extract with 
 hot dil. HNO3, and 
 filter. Test residue 
 with NaPOs bead or 
 with HF in a plati- 
 num crucible. Test 
 filtrate for PO4 with 
 
 (NH4)2M004. 
 
 2nd part. Test for 
 HF. 
 
 Filtrate (prepared solution) will conjain the acids in 
 form of sodium salts -|- an excess of NajCOg and is to 
 be used for the acid tests unless otherwise directed. 
 
 Preliminary Tests for the Acid Groups 
 /usi 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 BaClj 
 and CaCla solutions (3). A white precipitate shows the 
 presence of the Group I. (4) ; acidify with HCl. If 
 the ppt. does not dissolve, H2SO4 is present. If the 
 group is present, test separate portions of the prepared 
 solutions for H8PO4, H3BO3, HF, H2C2O4, H2C4H4O6, 
 H2Si03. 
 
 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. 
 HBr and HCNS may also 
 possibly be absent. If 
 Group II. is shown to 
 be present, test separate 
 portions of the prepared 
 solutions for HI, HBr, 
 HCl, HCN, H4Fe(CN)8, 
 H3Fe(CN)6, and HCNS. 
 
 To filtrate or solution in 
 which AgNOs produces no 
 ppt. in HNO3 solution, 
 add more AgN03 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 HClOj, HNOt, 
 and HC2H3O2. 
 
TEE ACIDS 167 
 
 NOTES 
 
 1. If the substance is soluble in water and contains no heavy metals, the 
 treatment with Na^COg may be dispensed with. If acid (preferably nitric 
 acid) has been used in getting the substance into solution, and the latter is 
 boiled, HjCOs, H^SOa H^SgOg, HjS, 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 NajCOg 
 solution, may leave a residue consisting of the phosphates and fluorides of 
 certain metals which are not readily transposed by boiling with NagCOg. 
 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 cru- 
 cible with NajCOgj 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 HjS into the acidified 
 solution, filtering, and boiling out the HjS from the filtrate. The latter may 
 then be treated with NajCOj in the manner already described. In the absence 
 of nitrates and chlorates and in the presence of only metals which are pre- 
 cipitated by H2S, a solution for the acid tests may be prepared by saturating 
 with HjS a suspension of about i g. of the substance in about 50 cc. of water. 
 Heat to boiling and filter. The filtrate after boiling to expel the HgS is used 
 in small portions for the acid tests. Nitrates and chlorates, if present, would 
 oxidize the H2S, forming H2SO4 ; while chlorates would be reduced to chlo- 
 rides, thus making the tests for these acids of no value. 
 
 2 Boiling with NajCOg 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 NajCOg. On acidifying this solution, a precipitate may be 
 obtained. It should be filtered olF, rejected, and the filtrate again boiled to 
 drive out any H2S that may be liberated before neutralizing the solution with 
 ammonium hydroxide. All the COj must be expelled after acidifying, other- 
 wise BaCOg and Ag2C05 will precipitate when the group reagents are added. 
 Too much HNOg should be avoided, as this will form with the ammonia next 
 to be added an unnecessarily large amount of NH^NOg, 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 oi 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 presence 
 
1 68 QUALITATIVE CHEMICAL ANALYSIS 
 
 of much ammonium salts interferes more or less with the precipitation of all 
 the acids of this group with the exception of HjjSO^ and H2C2O4. 
 
 5. If only a cloudiness is obtained with AgNOg, it indicates a trace of 
 chlorides which should not be reported. AgaCgO^ is difficultly soluble in 
 HNO3 ; an excess of this acid should be added before adding the AgNOj 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; AgjS, black; AggFeCCN)^, reddish brown; AgCl, 
 AgCN, AgSCN, and Ag^FeCCN)^ are white; and AgBr is yellowish white. 
 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 HjSO^ ; the latter is, however, diffi- 
 cultly soluble in water. The color of the precipitate forming at the neutral 
 junction of the two liquids wiU often indicate which of the acids of Group I. 
 are present; if yellow, it may be Ag3P04 or AgjAsOj; if brownish red, 
 AgjAsOj.; if purplish red, Ag2Cr04; 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 HCl and warm. A carbonate, if present, 
 will evolve COg, which may be recognized by its property of 
 rendering turbid a drop of limewater supported in a glass 
 tube. 
 
TEE ACIDS 169 
 
 NOTES 
 
 1. Sulphites, if present, would evolve SOj, which would also render lime- 
 water turbid. Sulphides and nitrites also liberate gases on treatment with 
 dilute HCl. To avoid the interference of these substances, use a strong 
 solution of KaCrjO^, instead of acid, and warm the mixture. COj alone will 
 be evolved. KjCrjO^ 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. 
 
 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 HCl. 
 In the presence of a sulphite, SOj will be evolved; this gas 
 may be readily recognized by its odor and by its property of 
 bleaching a very dilute solution of KMnO^ (see reaction 6 
 under Sulphites). 
 
 Thiosulphates 
 
 These are detected in the preliminary tests with dilute HCL 
 In the presence of a thiosulphate, SOg 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 HCl to acid reaction, boil to expel COj, 
 filter if necessary, and to the filtrate add BaClg. 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 NagCOg, and heat on charcoal before 
 
lyo QUALITATIVE CHEMICAL ANALYSIS 
 
 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 
 HCl, boil out the COg, 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 
 test tube, add concentrated HjSO^, 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 H^SiO^. (See footnote, page 
 124.) 
 
 Chromates 
 
 (a) Solutions of a dichromate or a chromate possess an 
 orange or yellow color which is very characteristic. Acidifica- 
 tion with dilute HCl, followed by treatment with HgS, 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 presence of a chromate may be detected 
 by acidifying the prepared solution with HNO3, thoroughly 
 cooling, and then adding i cc. of ether and i cc. of 3 % HjOg 
 and shaking. A blue color in the ether layer proves the pres- 
 ence of a chromate. Reaction 6, page 143. 
 
 {d) 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, e.^., strong HCl and a little alcohol (see 
 reaction 5 under Chromates), KI, and NagSOg ; the last is oxi- 
 dized at the same time to sulphate. If the prepared solution is 
 colorless, chromates cannot be present. 
 
 (c) Precipitation Test. In the absence of sulphates, phos- 
 phates, oxalates, and tartrates, acidify the " prepared solution " 
 with acetic acid, boil to expel COj, filter if necessary, and to the 
 clear solution add i g. of NaCgHgOg and a little Pb(C2H302)2 
 
THE ACIDS 171 
 
 solution. A yellow precipitate of PbCrO^ proves the presence 
 of a chromate. 
 
 If the above-mentioned acids are present, acidify the prepared 
 solution with HNO3, boil to expel COg, filter if necessary, and 
 render the resulting filtrate just alkaline with ammonium hy- 
 droxide ; add CaClg, 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 BaClj, 
 and filter off the BaSO^ on two folds of filter paper. To the 
 filtrate add several grams of NaCgHgOg to completely replace 
 the nitric acid by acetic acid, and heat, when a yellow precipi- 
 tate of BaCrO^ will be formed. 
 
 Arsenites 
 
 Arsenites are detected in the analysis for metals. In HCl so- 
 lutions, HjS yields an immediate precipitate of AsgSg. Arsen- 
 ites are not precipitated by either magnesia mixture or ammonium 
 molybdate. In a strictly neutral solution, AgNOg produces a 
 yellow precipitate of AggAsOg (phosphates respond to the same 
 test). 
 
 Arsenates 
 
 From acid solutions, HjS slowly yields a yellow precipitate. 
 From strictly neutral solutions, AgNOg precipitates reddish 
 brown AggAsO^. 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 
 distinguished from arsenites, even in the presence of phos- 
 phates, by the ability of arsenates to liberate iodine from 
 KI in a solution acid with HCl. For the detection of Iodine, 
 see p. 147. 
 
172 QUALITATIVE CHEMICAL ANALYSIS 
 
 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 10 cc. of ammonium molybdate solution, and 
 warmed ; in the presence of a phosphate, a yellow precipitate 
 of (NH4)3P04 • 12 M0O3 will be formed. 
 
 (d) Or the prepared solution is acidified with HCl, 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 NH^MgPO^ forms 
 in the presence of a phosphate. 
 
 HOTES 
 
 1. The solution is evaporated with concentrated HNO3 to oxidize any 
 reducing agent that may be present and which would interfere with the 
 (NH4)2Mo04 test; it also converts at the same time any meta- or pyro-phos- 
 phate to the ortho form, which alone is precipitated by (NH4)2Mo04. 
 
 2. If arsenic has been found, it should be removed by rendering the pre- 
 pared solution strongly acid with HCl, heating to boiling, and passing in a 
 stream of HgS for 20 minutes ; then filter, boil out the H^S from the filtrate, 
 add HNO3, and evaporate nearly to dryness; extract the residue with boiling 
 dilute HNO3 and add to the somewhat cooled solution an excess of am- 
 monium molybdate. A yellow precipitate shows the presence of PO4. The 
 solution to be tested for PO4 should not be above 70° C, as there is danger 
 of decomposing the reagent, with the resulting precipitation of white M0O3. 
 
 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 COg, and filter if necessary ; warm the filtrate and 
 add an equal volume of a saturated CaSO^ solution. A white 
 crystalline precipitate indicates the presence of an oxalate. 
 
THE ACIDS 
 
 173 
 
 Confirm by filtering off the precipitate, washing it with water, 
 dissolving it in hot dilute H2SO4, and adding a drop of dilute 
 KMn04 solution. In the presence of an oxalate, the KMnO^ 
 will be bleached. (See reaction 4 under Oxalates.) 
 
 BOTES 
 
 The solution is rendered acid with acetic acid to prevent the precipitation 
 of carbonates and phosphates. CaFj may be precipitated, but may usually be 
 distinguished from CaC204 by the fact that the former is gelatinous while the 
 latter is crystalline. It is, however, better to make the confirmatory test. 
 
 Fluorides 
 
 1. The etching test (see page 133) is not applicable in the 
 presence of silicates or silica. 
 
 2. The test depending upon the formation of SiF^ and the 
 detection of the latter by its property of rendering a water 
 "bead," held on the loop of a platinum wire, turbid (see 
 page 1 32), is applicable to fluorides in the presence of SiOj or 
 silicates. To insure a positive test, a little sand or dry sodium 
 silicate should be added. 
 
 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 NagSiOg. 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 HCl and dry it; this may be conveniently accompUshed 
 by placing it on the outside of a test tube containing water 
 which has just been heated to boiling. In the presence of a 
 
174 QUALITATIVE CHEMICAL ANALYSIS 
 
 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 HgS 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 
 in an evaporating dish with solid Na^jSOg, adding dilute HCl, 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 NajCOj ; dilute, 
 and filter. Iodides, if present, may be removed by precipitation with AgNOj. 
 after rendering the solution acid with HNO3. FeClg, if present, will color the 
 turmeric paper brown on concentration, but will give a brown insteaf' 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 NaPOa ^^^^ test may be applied to the original sub- 
 stance. 
 
 2. Evolution of SIF^. In a platinum crucible, treat a mixture 
 of equal parts of the dry substance and CaFg (free from SiOg) 
 with a little concentrated HgSO^, 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 CaF,. 
 
THE ACIDS 
 
 Tartrates 
 
 175 
 
 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 HjSO^ 
 (see reaction i under Tartrates). 
 
 1. Concentrate the prepared solution to about 0.5 cc, just 
 acidify with acetic acid, and add 2 cc. of KCjHgOg ; shake vigor- 
 ously and allow the mixture to stand. A white crystalline pre- 
 cipitate may be KHC4H4O5. Confirm by dissolving in a few 
 drops of dilute KOH solution, and precipitate the tartrate with 
 a little AgNOg; 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 HCl. Remove 
 the metals of Groups I. and II., if present, by means of HjS ; 
 and those of Group III. with NH^OH and (NH4)2S (Al and 
 Cr will, of course, not be precipitated). The clear filtrate is 
 acidified with HCl, boiled to expel HgS, and is finally rendered 
 alkaline with NH4OH. An excess of CaCla 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 CaC4H40e, CaC204, and Ca3(P04)2, is treated with a cold, 
 strong NaOH solution to dissolve out the CaC^H^Oe, 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 i drop of NH4OH and a little AgNOg, 
 and warm. In the presence of a tartrate, a black precipitate or 
 a silver mirror will be formed. 
 
176 QUALITATIVE CHEMICAL ANALYSIS 
 
 GROUP II 
 
 Iodides 
 
 1. Iodides, if present, will be detected in the preliminary ex- 
 amination of the substance with concentrated HjSO^. 
 
 2. To a portion of the original or prepared solution acidified 
 with HCl, add a little potassium nitrite solution or chlorine 
 water and then about i cc. of CS2 or CHCI3, and shake vigor- 
 ously. An iodide, if present, will color the CSj or CHCI3 violet. 
 
 3. AgNOj in HNO3 solution precipitates yellow Agl, prac- 
 tically insoluble in NH^OH. 
 
 BOTES 
 
 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 Substantes^^ page 195. 
 
 Bromides 
 
 1. Bromides, if present, will be detected in the preliminary 
 examination of the substance with concentrated HjSO^. 
 
 2. To a portion of the original or prepared solution, acid with 
 HCl, add I cc. of CSg or CHCI3. Cautiously treat with small 
 amounts of Cl water and shake vigorously after each addition. 
 In the presence of bromides, the CHCI3 or the CSg 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 
 CHCI3 or CSj. If the amount of iodide present is large, chlorine water should 
 be added until an intense violet color is produqed in the CHCI3 or CSj ; 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 CS^ 
 and Cl water. If the CSj is still colored a deep violet, the operation is re- 
 peated until only a faint pink color is imparted to the CSj ; now, on adding a 
 little more Cl water and shaking, a brown or reddish color will be produced if 
 
THE ACIDS 
 
 177 
 
 a bromide is present. If the amount of iodide in the original solution is 
 small, as is shown by the feint purple color of the CSj, 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 CSj 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 "/nsoluiie Substances,''^ page 195. 
 
 Thiocyanates, cyanides and ferrocyanides interfere with the detection of 
 bromides by the chlorine water test for the reason that these acids are readily 
 oxidized by the reagent added to liberate the Br. This difficulty may be 
 overcome by the method devised by Curtman and Wickoif which is based on 
 the fact that in a solution slightly acid with H2SO4, cuprous sulphate (prepared 
 by adding HjSOg to CuSO^) precipitates cyanides, ferrocyanides, ferricyanides, 
 and thiocyanates (also iodides), leaving in solution bromides and chlorides. 
 ■ Method. The solution to be tested, which should be neutral or slightly acid 
 with H2SO4, is treated with 15 cc. saturated solution of SOj. Heat to boiling 
 and while hot add slowly 2 N — CuSO^ until an excess is added. The solu- 
 tion should be blue ; a green color shows insufficient CuSO^. Filter while hot 
 and wash the ppt. twice with hot water, adding the washings to the filtrate. 
 Boil down the filtrate to j-io cc. to concentrate the solution and to expel the 
 excess SO2. (A slight white ppt. which may separate should be discarded.) 
 Transfer solution to a test tube and cool. Now add i cc. 3 N — HgSO^ and i 
 cc. i^fc KMnO^ and shake. Add 0.5 cc. CSj and shake again. A yellow 
 color in the CSj layer proves the presence of Br. The acid and KMn04 are 
 added and shaken first in order that the CSj may not be in contact with the 
 KMnO^ longer than is necessary, since they react with the formation of a little 
 MnOg which with vigorous shaking may dissolve in the CS2, yielding a color 
 indistinguishable from that given by small amounts of Br. 
 
 Chlorides 
 
 In the absence of bromides, iodides, cyanides, ferrocyanides, 
 and thiocyanates, a white precipitate, obtained with AgNOg 
 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 AgNOg 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 
 
178 QUALITATIVE CHEMICAL ANALYSIS 
 
 the formation of a white curdy precipitate shows the presence 
 of a chloride. 
 
 Chlorides in the presence of Bromides and Iodides may be 
 detected by one of three following methods : — 
 
 1. 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 KjCrgO^ ; 
 add 5 cc. of concentrated HjSO^, and heat. Absorb any fumes 
 that may be evolved in dilute NH^OH. The latter will be 
 colored yellow if a chloride was originally present (see reaction 
 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 
 KMnO^ and the liberated Br is removed by boiling ; any resid- 
 ual substance which will give with AgNOs 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 HjSO^, 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, KMn04 solution is 
 added in a quantity sufficient to give the solution a purple color which does 
 not disappear on boiling. The KMnO^ 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 KMnO^, a few drops of alcohol are added, the mixture 
 is boiled with stirring, and the brown hydrated MnOj is filtered off. The fil- 
 trate, which should be colorless, is treated with a few drops of AgNOg. A 
 ■white precipitate, insoluble in HNOg and soluble in ammonium hydroxide, 
 proves the presence of a chloride. 
 
 3. Vortmann's Method consists in acidifying the prepared 
 solution with acetic acid, adding PbOj, and boiling until no more 
 
THE ACIDS 179 
 
 bromine and iodine are given off (as shown by tests) and the 
 solution on settling is colorless. 
 
 All the hydrobromic acid and part of the hydriodic acid are 
 oxidized by the PbOg; the remainder of the iodine, combined in 
 the form of lead iodide, settles on the bottom of the beaker 
 along with the excess of PbOg added. Filter and wash the 
 precipitate with hot water, and test the filtrate for chlorides 
 with AgNOg. 
 
 NOTES 
 
 Cyanides, ferrocyanides, ferricyanides, and thiocyanates 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 AgNOj, 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 
 NajCOg, extracting the melt with water, and filtering. The filtrate will con- 
 tain NaCl, NaBr, NaT, and an excess of NajCOg. 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 
 HCl and a little FeClg is then added. In the presence of a 
 ferrocyanide, a blue precipitate of prussian blue is obtained. 
 
 Ferricyanides 
 
 1. To a small portion of the prepared solution acidified with 
 HCl, add a freshly prepared solution of FeSO^ ; the formation 
 of a dark blue precipitate of Turnbull's blue proves the presence 
 of a ferricyanide. 
 
 2. From a nitric acid solution, AgNOg precipitates reddish 
 brown Ag3Fe(CN)e. 
 
 Thiocyanates 
 
 I. Acidify a portion of the prepared solution with HCl and 
 add FeClgj a deep red coloration, due to the formation of 
 
i8o QUALITATIVE CHEMICAL ANALYSIS 
 
 ferric thiocyanate, proves the presence of a thiocyanate. The 
 solution is acidified with HCl to prevent the interference of 
 (i) acetic acid, which, with FeClg, 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 
 FeCIg. 
 
 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 HCl a little SnClj sufficient in 
 amount to reduce any I or Br present to their corresponding halogen acids, 
 boiling, and then absorbing the HCNS, which distills over, in a test tube con- 
 taining FeClj, 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 HCl and concentrated H2SO4. 
 
 2. To a portion of the prepared solution, add 2 cc. of NaOH 
 solution, and treat with a little FeSO^ and a few drops of 
 FeClg; heat gently for a short time and then acidify with HCl. 
 A blue precipitate of Fe4[Fe(CN)g]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, proceed as follows : Put into a small distill- 
 ing flask about ij cc. of water that has been saturated with CO2, add an 
 excess of solid NaHCOj, and finally some of the powdered original substance. 
 Quickly stopper the flask and distill under a hood, catching the distillate in 
 ahttle 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 relatively 
 stronger carbonic acid. 
 
 The method of Barnebey may also be employed to advantage. This test 
 
 depends upon the fact that alkali cyanide solutions have a solvent action on 
 
 N 
 CuS. Render 10 cc. — CuSO. alkaline with NH.OH and treat with a few 
 ID * * 
 
 bubbles of HjS. Divide the suspension of CuS thus formed into 2 portions ; 
 
 and to one add a little of the prepared solution and shake. Compare the 
 
TEE ACIDS i8i 
 
 tubes. A bleaching of the color in the treated tube shows the presence of a 
 cyanide. 
 
 Sulphides 
 
 1. Most sulphides will have been detected in the preliminary 
 treatment with dilute HCl by the evolution of HjS, 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 HCl 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 I are usually decomposed by this treatment, yielding HjS, 
 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 AggS 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 HjSO^, and cool 
 thoroughly in a stream of running water. Incline the tube and 
 carefully add 2 to 3 cc. of a strong freshly prepared FeSO^ solu- 
 tion, and allow the mixture to stand. In the presence of a 
 nitrate, a brown coloration will form at the junction of the two 
 liquids. 
 
i82 QUALITATIVE CHEMICAL ANALYSIS 
 
 NOTKS 
 
 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 AgaSOi (free from nitrates). Chromates will be reduced by FeS04, 
 yielding green Cr3(S04)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 Na2S03 and dilute H2SO4, boil- 
 ing until the solution is green, and then precipitating the Cr and Mn salts 
 with an excess of NaaCOs- Filter, acidify the filtrate with dilute H2SO4, and 
 make the test on the resulting solution. Chlorates interfere on account of 
 CIO2, which will form on adding concentrated H2SO4 ; these will also be 
 reduced by the above treatment. Ferro- and ferri-cyanides yield with FeS04 
 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 H0SO4, 
 heating the mixture to boiling, and adding BaCU to precipitate the H2SO4. 
 The H2SO4 and BaCl2 are added to form heavy BaS04, 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 coloration 
 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 HCl ; the NO2 fumes given off may be readily 
 detected by their property of turning starch-iodide paper blue. 
 
 2. Brown Coloration Test. The same as with nitrates except 
 that in this case acetic acid or dilute H2SO4 may be used instead 
 of concentrated HgSO^. 
 
THE ACIDS 183 
 
 Acetates 
 
 1. 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 heat gently. The char- 
 acteristic odor of amyl acetate ("pear essence") indicates the 
 presence of an acetate. If ethyl alcohol is used, the odor of 
 ethyl acetate will be made evident on warmingj the mixture (see 
 reaction 2 under Acetates). 
 
 Chlorates 
 
 1. Chlorates are recognized by their behavior when treated 
 with concentrated HjSO^. 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 ClOj will be evolved ; the evolution is accom- 
 panied by a slight explosion if the gas is sufRciently heated. 
 
 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 
 AgNOg and filtering. To the filtrate, the volume of which 
 should be 50 cc, add 5 cc. cone. HNO3 and S cc. saturated SOj 
 solution. Heat. A white precipitate of AgCl proves the pres- 
 ence of a chlorate. (See reaction 4 under Chlorates.) 
 
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 
 
 i8s 
 
1 86 QUALITATIVE CHEMICAL ANALYSIS 
 
 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 HgS has been expelled, add an ex- 
 cess of NagCOg, boil vigorously for a moment, and filter. Render 
 the filtrate slightly acid with acetic acid, boil off the COg, and 
 then add an equal volume of a saturated CaS04 solution ; a white 
 crystalline precipitate proves the presence of an oxalate. 
 
 Organic matter and oxalates may both be removed by the 
 following procedure : To the residue obtained by evaporating 
 the filtrate from Group II. to dryness, add S cc. cone. Hj^SO^ and 
 heat gently until the mass has completely charred. Cool. Add 
 5 cc. cone. HNO3 and heat (gently at first) until fumes of SO3 
 are given off. Cool. Add 5 cc. more cone. HNO3 and evapo- 
 rate again to SO3 fumes. Repeat this treatment with HNO3 
 and evaporating to SO3 fumes until the solution is either color- 
 less or possesses only a faint straw color. (Three treatments 
 are generally sufficient.) Cool. Cautiously dilute with 25 cc. 
 water, boil to expel gases and filter usine a double^ filter. 
 Analyze the filtrate for Groups III; and V. and for Ca, if the 
 latter is not found in the residue. The residue may consist of 
 BaS04, SrSO^, CaS04 and anhydrous Cr2(S04)3. Boil the 
 residue for 10 minutes with 20 cc. of a saturated NajCOg solu- 
 tion and filter. Wash the residue until the washings after 
 being acidified with HCl, fail to give a test for sulphates. 
 Reject filtrate and washings. Heat residue with 15 cc. dil. 
 HNO3 and filter. Analyze filtrate for Cr, Ba, Sr, and Ca. 
 
 Test for Phosphates. To about 2 cc. of the filtrate from 
 Group II., from which the HgS 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)3P04 • 12 M0O3 
 proves the presence of a phosphate. 
 
 Silicates, if present, should have been detected in the prelim- 
 inary examination of the solid substance with a NaPOg bead ; and 
 should have been removed, preferably before proceeding with 
 
PHOSPHATE SEPARATION 187 
 
 the metallic analysis, by evaporating the solution acid with HCl 
 or HNO3* to dryness, desiccating at 120°, extracting the residue 
 with a fewcc. of concentrated HCl,* diluting, heating, and filter- 
 ing off the dehydrated and insoluble silica. The latter may then 
 be verified by the NaPOg bead test, or by treatment in a plati- 
 num crucible with CaFg and concentrated HjSO^, heating, and 
 testing the escaping vapor with a drop of water (see reaction i, 
 page 132). 
 
 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 
 NH^Cl 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 HCl 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 ^ft Jhe phosphate has been precipitated ; for 
 
 * If metals of the first group are present, use HNO3 instead of HCl. 
 
i88 QUALITATIVE CHEMICAL ANALYSIS 
 
 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 NaCjHgOj, 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 83 and 84). 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 HjS 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)g. 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 cautiously, with 
 constant stirring, dilute HCl, drop by drop, until a clear solu- 
 tion is obtained; then add 8 g. of NH4C2Hg02 and 8 cc. of 
 50 per cent, acetic acid. If the solution is not red, add FeClg 
 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)g, showing that 
 an excess of Fe (-ic) is present ; if a light-colored precipitate is 
 obtained with NH^OH, more FeClj 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 
 
PREPARATION OF THE SOLUTION 189 
 
 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. NajOj are added, the mixture is boiled, 
 with stirring, and finally filtered. The residue, consisting of 
 Fe(0H)3, is rejected. The filtrate may contain NagAlOg, 
 NagCrO^, Na3P04, 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 ^^^ then 
 rendering alkaline with ammonium hydroxide; the white gelati- 
 nous precipitate of AIPO4 or A1(0H3) is filtered off, washed 
 with hot water several times, and the presence of aluminum is 
 confirmed by igniting with a few drops of Co(N03)2 in the 
 usual way. The filtrate will contain the chromium as Na2Cr04. 
 
 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 
 
I90 QUALITATIVE CHEMICAL ANALYSIS 
 
 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 extracts 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 HCl and 
 hot concentrated HCl. If these fail to effect solution, try the 
 action of dilute and concentrated HNO3 on separate small por- 
 tions; if these also fail, add HCl to the mixture containing 
 HNO3, thus forming aqua regia, and heat. If still insoluble, 
 
PREPARATION OF THE SOLUTION 191 
 
 examine it by the method given for "Insoluble Substances" 
 page 197. 
 
 NOTES 
 
 1. During the treatment with HCl, indications of the presence of certain 
 acids will be given (see '■^Preliminary Examination^'' page 162). 
 
 2. If Hg and As are present, boiling with HCl 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 HCl 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 HCl 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 HCl treatment causes gelatinous silicic acid to separate, evapo- 
 rate the mixture on the water bath to dryness, dehydrate by heating to 120° C. 
 for half an hour, extract with 3 cc. of hot concentrated HNO3 or HCl, 
 dilute, heat, and filter off the SiOz. The filtrate is then examined for the 
 metals. 
 
 6. If HNOg has been used as a solvent, boil the liquid down to about 
 I cc, dilute with 20 cc. of water, and if the solution clouds on dilution, clear 
 with a few drops of HNOg 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 i cc. to destroy 
 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 HCl and 
 again evaporated down to i 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 boihng HCl or HNO3 : 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 HCl ; the latter, on the other 
 hand, dissolves the oxides of Sn and Sb, as well as MnOj, all of which are 
 not dissolved by HNO3. 
 
192 
 
 QUALITATIVE CHEMICAL ANALYSIS 
 
 METALS AND ALLOYS 
 
 From 0.5 to I gram of the metal or alloy, in the form of shav- 
 ings, foil, filings, or turnings, is treated with 20 cc. 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. 
 
 (d) If a white residue is left which is insoluble on dilution 
 and heating, it may consist of hydrated SnOj or SbgOg, or both, 
 admixed with arsenic in the form of tin arsenate, phosphorus in 
 the form of tin phosphate, bismuth as BijOg, and traces of Cu 
 and Pb. Filter. 
 
 Filtrate. Evap- 
 orate to drive oflF 
 excess of HNO3. 
 Add HCl 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. 
 Treat residue with 
 hot dil. HNO3 and 
 unite filtrate with 
 solution of residue 
 2. 
 
 Dry and fuse residue in a porcelain crucible with 4 
 times its weight of a mixture of equal parts of NajCOa and 
 S ; cool, extract melt with hot water, and filter. 
 
 Residue 2 is CuS, ihSa, 
 PbS. Dissolve in hot dil. 
 HNO3 and combine with 
 the corresponding solution 
 obtained from the first fil- 
 trate and proceed as di- 
 rected in the analysis of the 
 main filtrate in Scheme II. 
 A. 
 
 Filtrate will contain the 
 As, Sb, and Sn as thio-salts 
 -I- an excess of Na2S. Just 
 acidify with HCl, filter, and 
 reject filtrate. Residue may 
 consist of AS2S5, SbjSs, and 
 SnSa + S ; analyze accord- 
 ing to Scheme II. B. 
 
 * Minute amounts of Sb dissolve completely in HNOg; silver alloys containing a 
 very small amount of Pt are completely dissolved by HNO3. 
 t A black residue of carbon or graphite is sometimes left. 
 
INSOLUBLE SUBSTANCES 
 
 NOTES 
 
 193 
 
 1. If solution of the alloy does not readily take place, and Pt and Au are 
 absent, treat with HCl ; the latter is more satisfectory than HNO3 as a solvent 
 for Al, since the latter is only difficultly soluble in HNOs. HNO3 is used in- 
 stead of HCl, first, because it is the better solvent for metals and alloys ; second, 
 because treatment with HCl would convert any P, S, and As usually present as 
 phosphide, sulphide, and arsenide, respectively, into PH3, H2S and AsHs, which 
 would be lost by volatilization. HNO3 oxidizes these elements to their corre- 
 sponding acids, viz. . H3PO4, H2SO4 and H3ASO4. Only these acids together 
 with HjSiOs need to be tested for in the analysis of alloys. 
 
 2. A small white residue may be boiled with concentrated HCl and the 
 resulting 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 HgCla. 
 
 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 
 ^nto solution by the action of the acids taken singly or together. 
 
 The most common insoluble substances are the following : — ■ 
 
 C, S, Ag3Fe(CN)6, Ag^FeCCN)^, AgCN, AgCl, AgBr, Agl, 
 BaSO^, SrSO^, CaSO^, PbSO^, PbCla, fused PbCr04, ignited or 
 anhydrous chromic salts, ignited and native oxides, as AlgOg 
 (corundum), FegOg, SnOj (cassiterite), CrjOg, CrgOg • FeO 
 (chrome-iron ore), CaFj, Sb204, Fe4[Fe(CN)5]3, SiOg, and cer- 
 tain silicates. 
 
 Carbon is generally recognized by its black color, insolubility 
 in agua regia, and combustibility when heated strongly on plati- 
 num foil. When heated with KNO3, deflagration ensues with 
 the formation of KjCOg ; this method is not applicable to graph- 
 ite, the presence of 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 SOg 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. 
 
194 QUALITATIVE CHEMICAL ANALYSIS 
 
 Treatment with ag'ua 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.* 
 
 PbSO^ 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 
 HjSO^ or KjCrO^ ; in the second, test for CI by diluting, addi- 
 fying with HNO3, and adding AgNOg; and in the third, test 
 for SO4 by acidifying with HCl, filtering if necessary, and add- 
 ing BaClg. 
 
 Sulphates of the Alkaline Earth Metals are best fused in a 
 platinum crucible with five times their weight of anhydrous 
 NagCOg ; 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 NajSO^ and an excess of 
 NagCOg. 
 
 SrS04, CaS04, and PbSO^ may be quantitatively converted 
 into the corresponding carbonates by prolonged boiling with a 
 concentrated NajCOg solution. If the residue after filtering 
 is thoroughly washed free from alkali, it may then readily be 
 dissolved by acid. 
 
 BaS04 requires several treatments for its complete transfor- 
 mation by NajCOg solution. 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 151 under Ferrocyanides, reaction 3. 
 
INSOLUBLE SUBSTANCES 
 
 195 
 
 Fused PbCrO^, CrgOg, Chrome Iron Ore and Ignited Chromic 
 Salts are best fused with NagOg in a nickel crucible or with a 
 mixture of NagCOg and NaNOg. By this treatment, soluble 
 Na2Cr04 is formed, which after treatment with water is sep- 
 arated by filtering and tested for in the filtrate. 
 
 SnOj and SbgO^ are got into solution by fusing in a porcelain 
 crucible with three times their weight of NagCOg 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 accord- 
 ing to Scheme II. B. 
 
 AljOg and FegOg are fused in platinum with KHSO4 ^^ KgSgO^, 
 whereby they are converted into soluble sulphates. Fusion with 
 NajCOg, followed by acid treatment, also takes these oxides in 
 solution. 
 
 Silver Halides may be treated by one of the following two 
 methods : — 
 
 1. Fuse with NagCOg 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 HjSO^ 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. 
 
 SiOj and Silicates are recognized by the " skeleton " NaPOg 
 bead. 
 
 Silicates are usually decomposed by fusing with five times 
 their weight of a mixture of equal parts of anhydrous NajCOg 
 and KjCOg, to which about o.i 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 SOg fumes are given off. The resi- 
 due will be CaS04 ; extract it with water for some time and filter. 
 Test the filtrate for Ca with (N 114)20204. 
 
196 QUALITATIVE CHEMICAL ANALYSIS 
 
 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 : — 
 
 1. 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 NaPOj which is unaffected by the reduc- 
 ing flame ; at the same time, indications of SiOg 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 NagCOg 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 pres- 
 ence of Sn. Divide the HNO3 solution into two portions. To 
 the first add HCl; a white precipitate soluble in NH^OH shows 
 the presence of Ag. To the second portion add dilute HgSO^ ; 
 a white precipitate is PbS04. If no globule is obtained and the 
 white substance is not blackened by (N 114)23, 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 HCl 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 SiOg and Cr, 
 the following may be used : — 
 
 SiOg. Mix the finely powdered substance in a platinum cru- 
 cible or lead tube with a small quantity of CaFj (SiOg free), add 
 
SCHEME FOR INSOLUBLE SUBSTANCES 197 
 
 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, SiOg is present. 
 
 Chromium. Prepare a NagCOg bead. Take up a little of the 
 substance mixed with KCIO3 and heat. Place bead in about 
 I cc. of water and heat. A yellow solution indicates Cr. 
 
 Scheme for the Treatment of Insoluble Substances 
 
 If the substance contains Pb salts, these may readily be 
 removed by repeatedly digesting with hot (NH4)2C4H40g or 
 NH4C2H3O2 solution, filtering, and testing the filtrate for Pb, 
 SO4, and CI. 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 given only when Ag salts are shown 
 to be present by preliminary test with Zn + H2SO4. The KCN 
 extract is tested for Ag by adding (NH4)2S, filtering off the 
 AggS, washing, and dissolving in hot dilute HNO3. To the clear 
 solution, add HCl ; 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 KjCOj and Na2C03 + 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 SiOj, 
 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 
 boiUng water. Break up the mass with pestle during extrac- 
 tion. Allow finally to settle, and filter. 
 
198 
 
 QUALITATIVE CHEMICAL ANALYSIS 
 
 Filtrate may contain NajSiOj, Na2Cr04, 
 NaF, NajPOi, NasAlOs, NajSOi, Na^SnOj, 
 NajAsO^, (NaSbOj), NajMnO^, NajBOj 
 and Na2C03, as well as K salts of these 
 acids. Divide into two equal portions. 
 
 ist portion. Acidify with HCl 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 SiOa by 
 heating at 120° C. till all HCl is driven off. 
 Extract the residue with 25 cc. of cone. HCl, 
 stir thoroughly, add 25 cc. of water, and 
 boil with stirring ; filter. 
 
 Residue : 
 
 SiOn. 
 
 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. 
 
 Residue may consist of BaCOg, 
 SrCOg, CaCOg, possibly AlgOg, 
 MgO, FCgOg, SnOj, andunattacked 
 SiOj, and in some cases nearly 
 any metal or its oxide ; wash sev- 
 eral times with hot water. Treat 
 with hot dil. HNO, and filter. 
 
 Filtrate is to 
 be united with 
 HCl 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 SiOj, 
 SnOj, and 
 AI2O3. Fuse in 
 a Ni crucible 
 with NaOH. 
 Extract with 
 water and filter. 
 Filtrate con- 
 tains Na2Sn03, 
 NasAlOs + 
 NaaSiOs. Test 
 for Al and Sn, 
 if not already 
 found. 
 
 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. NH^Cl, and the resulting mixture is then 
 thoroughly mixed with 8 grams of alkali-free CaCOg 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 CaClj, which decomposes the siUcate 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- 
 
ACID ANALYSIS OF MINERALS, ETC. 199 
 
 tion may be hastened by crushing any lumps with a pestle. 
 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)2C03 to complete precipita- 
 tion, and finally a little (NH^)2C204. 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, siHcates, phosphates, borates, 
 sulphates, fluorides, and chlorides. Carbonates and sulphides 
 may be detected by treatment with HCl, and silica or sili- 
 cates by the NaPOg 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 
 19s). In the absence of sulphides, the test for sulphates may 
 also be made by fusing the original substance with NagCOg, ex- 
 tracting the melt with boiling water, and filtering ; the filtrate, 
 after acidifying with HCl and boiling to drive out the COg, is 
 then treated with BaClg. If sulphides are present, boil the 
 finely powdered substance, with constant stirring, with a satu- 
 rated solution of Na2COg, then filter, acidify the filtrate with 
 HCl, and add BaCl2 ; or if the original substance dissolves com- 
 pletely in HCl, the resulting solution may be treated with BaCl2. 
 For the detection of fluorides in the presence of silicates, see 
 test 3 under Fluorides, page 173. For the detection of borates 
 in silicates undecomposed by concentrated H2SO4, see reaction 
 6 under Borates, page 1 36. 
 
 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 NajCOg, extracting the mass 
 with boiling water, and filtering. The filtrate will then contain 
 
200 QUALITATIVE CHEMICAL ANALYSIS 
 
 NagSiOj, NaF, NaBOg + the excess of NagCOg. A portion of 
 this solution, after slightly acidifying with HCl, may be tested 
 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 173). 
 
TABLE OF SOLUBILITIES 201 
 
 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. 
 
 Prelifninaty 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 border lines 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 i6th German edition of Fresenius' Quali- 
 tative Analysis. 
 
202 
 
 QUALITATIVE CHEMICAL ANALYSIS 
 
 SOLUBILITY 
 
 Oxide 
 
 Chromate . . . . 
 Sulphate. .... 
 Phosphate . . 
 
 Borate 
 
 Oxalate 
 
 Fluoride , 
 
 Carhonate . . . 
 
 Silicate 
 
 Chloride .... 
 Bromide . . . . , 
 
 Iodide 
 
 Cyanide 
 
 Ferrocyanide 
 Ferricyanide , 
 Thiocyanate . 
 
 Sulphide 
 
 Nitrate 
 
 Chlorate 
 
 Tartrate 
 
 Citrate 
 
 Malate , 
 
 Succinate . . . 
 Benzoate . . . . 
 Salicylate.... 
 
 Acetate 
 
 Formate . . . . . 
 Arsenite . . . . , 
 Arsenate..... 
 
 Pi 
 
 W 
 
 Wi 
 
 W13.IB 
 
 W 
 
 W3 
 
 w 
 
 w 
 
 W 
 W 
 
 w 
 
 w 
 w 
 w 
 w 
 w 
 w 
 
 W5.6.Y.22.4 
 w 
 
 w 
 
 w 
 w 
 
 w 
 
 s 
 s 
 < 
 
 w 
 w 
 w 
 
 W2i.„ 
 
 w 
 
 w 
 
 w 
 w 
 w 
 
 w 
 w 
 
 w-a 
 A 
 
 W 
 
 w-a 
 w-a 
 
 w&a 
 w-a 
 
 w-a 
 W 
 
 w-a 
 A 
 
 W 
 
 W-A 
 w-a 
 W-I 
 
 Au 
 a 
 A 
 A-I 
 A 
 a 
 W 
 
 W-A. 
 
 A 
 w— a 
 w-aj, 
 w-a 
 
 w 
 w-a 
 
 W 
 
 W 
 
 A 
 
 W 
 
 u 
 
 A&I 
 
 a 
 
 W&I, 
 
 W&I 
 w&i 
 
 a-i 
 
 W 
 
 w-a 
 A 
 
 A 
 A-I 
 
 W 
 
 W 
 
 w-a 
 
 A 
 
 W 
 
 w-a 
 A 
 
 W 
 
 ai9 
 W 
 
 Notes to Table of Solubilities 
 
 1. 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 
 
 203 
 
 TABLE 
 
 1 
 
 
 3 
 
 1 
 
 1 
 
 
 1 
 U 
 
 
 1 
 
 n 
 
 .3 
 
 s 
 
 
 
 
 
 S 
 g 
 
 1 
 
 1 
 1 
 
 u 
 
 1 
 
 i 
 
 
 
 
 a 
 
 A 
 
 a 
 
 Am 
 
 A 
 
 A 
 
 A 
 
 a 
 
 a 
 
 
 a 
 
 a 
 
 a&i 
 
 A42 
 
 Oxide 
 
 
 w 
 
 a 
 
 A-I 
 
 a 
 
 w-a 
 
 w 
 
 a 
 
 a 
 
 
 
 a 
 
 
 a 
 
 Chromate 
 
 w„ 
 
 W 
 
 W-A 
 
 A-I 
 
 w-a 
 
 W27 
 
 Wso 
 
 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 
 
 Wa 
 
 I 
 
 W-I 
 
 A-I 
 
 Wjs 
 
 W 
 
 W-A33 
 
 W 
 
 W36 
 
 W3,.^ 
 
 W 
 
 w«, 
 
 W-A,3 
 
 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 
 
 ^23 
 
 A 
 
 A 
 
 Ab9 
 
 agi 
 
 a 
 
 A 
 
 aae 
 
 ago 
 
 341 
 
 341 
 
 A44.4B 
 
 Sulphide 
 
 w 
 
 w 
 
 w 
 
 W 
 
 WSQ 
 
 W 
 
 W 
 
 W34 
 
 w 
 
 
 w 
 
 
 
 
 Nitrate 
 
 w 
 
 w 
 
 w 
 
 w 
 
 w 
 
 w 
 
 w 
 
 w 
 
 w 
 
 
 
 W 
 
 
 
 Chlorate 
 
 w-a 
 
 w. 
 
 a 
 
 a 
 
 w-a 
 
 a 
 
 w 
 
 a 
 
 w-a 
 
 
 
 a 
 
 
 346 
 
 Tartrate 
 
 w 
 
 w 
 
 a 
 
 a 
 
 a 
 
 w-a 
 
 w 
 
 
 a 
 
 
 
 
 
 
 Citrate 
 
 
 w 
 
 w-a 
 
 w-a 
 
 a 
 
 w-a 
 
 w 
 
 
 
 
 
 w 
 
 W 
 
 
 Malate 
 
 w-a 
 
 a 
 
 a 
 
 a 
 
 a 
 
 w-a 
 
 w 
 
 
 w 
 
 
 
 
 a 
 
 
 Succinate 
 
 w 
 
 a 
 
 w-a 
 w-a 
 
 a 
 w-a 
 
 a 
 
 w-a 
 
 a 
 w 
 
 
 w 
 
 
 
 
 
 
 Benzoate 
 
 Salicylate 
 
 w 
 
 W 
 
 w 
 
 W25 
 
 w-a 
 
 w 
 
 Wsj 
 
 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 
 
 ^■-, ■*w^.,-''V"^ 
 
 Hydrogen sodium carbonate, W. 
 Tricalcium phosphate, A. 
 Ammonium magnesium phosphate, A. 
 Potassium aluminum sulphate, W. 
 Ammonium aluminum sulphate, W. 
 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. 
 
 10, 
 II, 
 12. 
 
 13 
 14 
 
 IS 
 
204 QUALITATIVE CHEMICAL ANALYSIS 
 
 1 8. 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. 
 
REAGENTS 
 
 20S 
 
 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. 
 
 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 
 AcMs 
 
 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. HNO^, 3 N, sp. gr. i.ii, 19 % HNO3 by weight. 
 
 Cone. H^SOi, sp. gr. 1.84, 98 % H2SO4. 
 
 Dil. H^SO^, 3 N, sp. gr. 1.09, 13 % H2SO4. 
 
 Cone. HF, 40 %. < 
 
 Aeetie Acid, S N, sp. gr. 1.04, 30 % by weight. Dilute 285 cc. 
 glacial acetic acid to a liter. 
 
 Sulphurous Acid, H^SO^, a solution saturated at 15" contains 
 approximately 16.5 % HgSOg. 
 
 Tartaric Acid, 2 N, 150 g. in i liter. 
 
2o6 QUALITATIVE CHEMICAL ANALYSIS 
 
 Aqua Regia, i part of cone. HNOgto 3 parts cone. HCl; to 
 be prepared only when needed. 
 
 H^S gas is prepared by the action of HCl (i :i) 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 4 N solution will be 
 4 X 40 X 1^ = 177.7 g- in I liter. 
 
 Potassium hydroxide, KOH, 4 N. 
 
 The grade used for analytical purposes contains about 20 % 
 water ; hence the quantity needed for a 4 N solution will be 
 4 X 56 X f = 280 g. in I liter. 
 
 Barium, hydroxide, Ba(0H)2, saturated solution. 
 
 Calcium hydroxide, Ca(0H)2, saturated solution. 
 
 Salts 
 
 Ammonium acetate, NH^CjHgOj. 3 N. 250 g. in a liter. 
 
 Ammonium carbonate, (NH4)2C03 SO^ free. Dissolve, without 
 heating, 192 g. of the powdered salt in a mixture of 80 cc. of 
 NH40H(sp. gr, 0.90) and 500 cc. of water. When solution is 
 complete, dilute to i liter. The strength is approximately 4 N. 
 
 Ammonium chloride, NH^Cl, 4 N. 214 g. in i 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. M0O3 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)2C204 • HjO. 35.54 g. in i liter. 
 
 Ammonium sulphide (colorless), (NH4)2S. Saturate 3 parts 
 of NH4OH with HgS, add 2 parts of ammonium hydroxide, 
 and dilute with an equal volume of water. 
 
REAGENTS 207 
 
 Ammonium sulphide (yellow), (NH4)2S:,. Digest the color- 
 less undiluted (NH4)2S with flowers of sulphur in the proportion 
 of I g. to the liter and then dilute with an equal volume of water. 
 
 Ammonium sulphate, {'^Yi^^O^, N. 100 g. in i liter. 
 
 Barium chloride, BaClj • 2 HgO, W. 122.17 g- in i liter. 
 
 Bromine water, saturated solution. 
 
 Calcium chloride, CaClj, anhydrous, N. 55.6 g. in i liter. 
 
 Calcium sulphate, CaSO^ • 2 HgO, saturated solution. 
 
 Chlorine water, saturated solution. 
 
 Cobalt nitrate, Co(N03)2 • 6 HgO, for confirmatory tests for Al 
 and Zn. 0.5 g. in i liter. 
 
 Copper sulphate, CuSO^ ■ 5 HgO. 2 N. 249.6 g. in i liter. 
 
 Dimethylglyoxime. Dissolve 5 g. in 500 cc. hot 95 % alcohol. 
 
 Ferric alum, Y&^O^^ • (NH4)2S04 • 24 HgO, saturated solu- 
 tion. 
 
 Ferric chloride, FeClg ■ 6 HgO,* 2 N. 180 g. in i liter. 
 
 Ferrous sulphate, FeSO^ • 7 HjO. To be prepared in small 
 amounts as needed. 
 
 Hydrochlorplatinic acid, HgPtCle • 6 HjO. 10 % solution. 
 
 Hydrogen dioxide, 3 %. 
 
 Lead acetate, Pb(C2H302)2 • 3 HjO,! N. 189.5 g. in i liter. 
 
 Magnesia mixture. Dissolve no g. of MgCl2 • 6 HjO and 
 280 g. of NH4CI 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. 
 
 Phenolphthalein, i % solution in 50 % ethyl alcohol. 
 
 Potassinm. acetate, KC2Hg02. Saturated solution. 
 
 Potassium chromate, K2Cr04, N. 97. 3 g. in i liter. 
 
 Potassium cyanide, KCN, N. 65.2 g. in i liter. 
 
 Potassium dichromate, K2Cr207, N. 73.8 g. in i liter. 
 
 Potassium ferrocyanide, K4Fe(CN)6, N. 105.7 g- in i liter, 
 
 N ' 
 
 Potassium iodide, KI, — • 83.1 g. in i liter. 
 2 
 
 Potassium nitrite, KNOg- 500 g. in i liter. 
 
 Potassium permanganate, KMnO^, N. 79.1 g. in i liter. 
 
 * Should contain a IMe free HCl. t The solution should contain some free acetic acid 
 
2o8 QUALITATIVE CHEMICAL ANALYSIS 
 
 Potassium thiocyanate, KCNS, N. 97.2 g. in i liter. 
 
 N 
 Silver nitrate, AgNOj, — • 42.5 g. in i liter. 
 
 4 
 
 Silver sulphate, AggSO^. Saturated solution. 
 
 Sodium acetate, NaCjHgOg, 4 N. 328 g. in i liter. 
 
 Sodium carbonate, NagCOs (dry). Saturated solution. 
 
 Sodium cobaltic nitrite, Na3Co(N02)6. Dissolve 100 g. NaNOg 
 in 300 cc. distilled water, slightly acidify with acetic acid, and 
 then add 10 g. of Co(N03)2 • 6 HgO. 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, Na2FeNO(CN)5 • 2 HgO. ID % solution. 
 
 Sodium phosphate, NajHPO^ • 12 HjO, N. 119 g. in i liter. 
 
 Sodium stannite, prepared as needed by adding to a little SnClg 
 solution sufficient NaOH solution to redissolve the precipitate 
 which first forms. 
 
 Sodium thiosulphate, NagSjOg - 5 HjO, N. 124 g. in i liter. 
 
 N 
 
 Stannic chloride, SnCl4, — • 32.7 g. in i liter. 
 
 N 
 Stannous chloride,* SnClg • 2 HgO, — • 56.5 g. in i liter. 
 
 Stannous chloride (for Bettendorff Test). Dissolve 113 g. of 
 SnClg • 2 HgO in 75 cc. of cone. HCl, and add a few pieces 
 of C.P. tin foil and keep in a glass stoppered bottle. 
 
 Starch paste. Prepared as needed by mixing about i 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 CSg is added. 
 
 * The solution should be strongly acid with HCl; the addition of a little C.P. tin 
 foil prevents the oxidation of the reagent. 
 
REAGENTS 209 
 
 Solvents 
 
 Alcohol, amyl, (CgHnOH), C.P. Used in small amounts in 
 the test for acetate. 
 
 Alcohol, ethyl, (CjHgOH). 95 %, sp. gr. 0.815. 
 Benzol, CgHg, useful for dissolving sulphur. 
 Chloroform, CHCI3, used for dissolving iodine. 
 Carbon disulphide, CSg, used for dissolving iodine. 
 Ether, ethyl, (C2H5)20, solvent for fats and oils. 
 
 Dry Reagents 
 
 Aluminum turnings, pure. 
 
 Ammonium chloride, NH^Cl, C.P. 
 
 Ammonium nitrate, NH4NO3, C.P. 
 
 Borax, Na2B^07 • 10 H2O, C.P. and powdered. 
 
 Calcium carbonate, CaCOg, alkali free. 
 
 Calcium fluoride, CaFj, Si02 free. 
 
 Copper, strips. 
 
 Ferrous sulphate, FeS04 • 7 HjO, C.P. 
 
 Fusion mixture (Na2C03 + KgCOg, dry and C.P.). 
 
 Iron filings. 
 
 Iron nails. 
 
 Lead dioxide, Pb02, free from Mn. 
 
 Litmus paper, blue and red ; to be kept in stoppered bottles. 
 
 Manganese dioxide, MnOj, C.P. and powdered. 
 
 Microcosmic salt, NaNH4HP04 ■ 4 H2O. 
 
 Paraffine, m.-p. 124°. 
 
 Potassium acid sulphate, KHSO4, fused, C.P. in small lumps. 
 
 Potassium chlorate, KCIO3, C.P. powdered. 
 
 Potassium chloride, KCl, C.P. 
 
 Potassium cyanide, KCN, pure. 
 
 Potassium dichromate, K2Cr207, C.P. powdered. 
 
 Potassium ferricyanide, K3Fe(CN)g, C.P. 
 
 Potassium nitrate, KNO3, C.P. fine crystals. 
 
 Sand, sea. 
 
2IO QUALITATIVE CHEMICAL ANALYSIS 
 
 Silica, SiOj, purified. 
 
 Sodium acetate, NaCgHgOa, C.P. 
 
 Sodium acid carbonate, NaHCOg, C.P. 
 
 Sodium carbonate, NagCOg, anhydrous, C.P. powdered. 
 
 Sodium dioxide, NagOg, C.P. 
 
 Sodium sulphite, NagSOg • 7 HgO, pure anhydrous. 
 
 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 
 
 211 
 
 LIST OF APPARATUS 
 
 2 nests of beakers, Griffin, 1-4. 
 I graduated cylinder, 10 cc. 
 I graduated cylinder, 50 cc. 
 I 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. 
 I watch glass, 12.5 cm. 
 
 1 florence flask, f . b. 50 cc. 
 
 2 evaporating dishes, 10 cm. 
 2 evaporating dishes, 6.5 cm. 
 I porcelain crucible. 
 
 I horn spatula. 
 
 I rubber stopper, one hole, 
 
 No. I. 
 I funnel cleaner. 
 I sponge. 
 
 I test tube cleaner. 
 
 ^ box gummed labels, #217. 
 
 I doz. sheets filter paper, 18.5 
 
 cm., S. & S. 595. 
 I pkg. filter paper, 12.5 cm., 
 
 S. & S. 595. 
 I doz. fluted filters, 12.5 cm., 
 
 S. & S. 588. 
 I test tube rack. 
 I test tube holder. 
 I filtering stand. 
 I box matches, safety. 
 I pair forceps (small). 
 I pipestem triangle. 
 
 1 retort stand (2 rings). 
 
 2 bunsen burners, with hose. 
 I blowpipe. 
 
 1 stick charcoal. 
 
 2 pieces wire gauze, 10 cm. 
 
 square. 
 I triangular file. 
 I platinum wire. 
 I platinum foiL 
 I towel 
 
212 QUALITATIVE CHEMICAL ANALYSIS 
 
 PREPARATION OF UNKNOWNS 
 
 In the making up of unknowns, stock solutions of the concen- 
 tration I cc. = lOO 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 
 I cc. NaCl solution, 2 cc. Ca(N03)2, and i 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. NH^. 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 i gram. The minimum will depend upon the scheme 
 of analysis employed. It may be exceedingly small if the most 
 sensitive tests are used, e.£^., 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 
 
 213 
 
 TABLE EMPLOYED IN THE PREPARATION OF STANDARD 
 STOCK SOLUTIONS* 
 
 Group Substancb 
 
 fokmulak 
 Weight 
 
 solobility of 
 Salt in 100 Pts. 
 OF Cold Water 
 
 Per cent. 
 Metal 
 
 Quantity of 
 Salt to be Dis- 
 solved IN I 
 Liter to give 
 Strength i cc. 
 
 = 100 MG. OF 
 
 Metal 
 
 IAgN03 . . . . 
 
 170 
 
 V. s.i 
 
 63.S 
 
 157 
 
 HgNOa-HjO . . 
 
 280 
 
 sol. in pres. of 
 HNOs 
 
 71 -s 
 
 140 
 
 Pb(N03)2 . . . 
 
 331 
 
 48 
 
 62.S 
 
 160 
 
 Pb(QH302)2-3H20 
 
 379 
 
 46 
 
 54.6 
 
 183 
 
 IIHg(N03)2 4(H20) 
 
 333 
 
 sol. in pres. of 
 HNOs 
 
 60 
 
 167 
 
 HgClj 
 
 271 
 
 7-4 
 
 74 
 
 I3S' 
 
 Bi(N03)3-sH20 . 
 
 484 
 
 sol. in pres. of 
 HNOs 
 
 43 
 
 233 
 
 Cu(N03)2-6H20 . 
 
 29s 
 
 V. S. 
 
 21.S 
 
 465 
 
 CuCl2-2H20. . . 
 
 170 
 
 120 
 
 37 
 
 270 
 
 CUS04SH20 . . 
 
 249 
 
 40 
 
 25 
 
 400 
 
 Cd(N03)2-4 H2O . 
 
 308 
 
 V. s. 
 
 36 
 
 278 
 
 CdCl2-2H20 . . 
 
 219 
 
 , 140 
 
 SI 
 
 196 
 
 3CdS04-8H20. . 
 
 769 
 
 V. s. 
 
 43-5 
 
 230 
 
 AS2O3 
 
 198 
 
 4 
 
 75-S 
 
 ( y 
 
 NajHAsOa . . . 
 
 170 
 
 V. s. 
 
 44 
 
 227 
 
 Na2HAs04i2H20 
 
 402 
 
 28 
 
 18.7 
 
 ( y 
 
 AS2O5 
 
 230 
 
 ISO 
 
 65 
 
 153 
 
 SbCl3 
 
 226^ 
 
 sol. in pres. of 
 HCl 
 
 53 
 
 188 
 
 SiiCl2-2H20 . . 
 
 225 
 
 V. S. 
 
 S3 
 
 189 
 
 SnCIi-s 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. 
 ' 33 S- i" I l'*er HCl (i : i) gives strength i cc. = 25 mgs. As. 
 
 * 267g. in I liter will give strength i cc.=so mgs. As. 
 
 * Taken from an article published in School Sci, and Math., Vol. X., No. 6, by one of us 
 (L.J.C.). 
 
214 
 
 QUALITATIVE CHEMICAL ANALYSIS 
 
 TABLE EMPLOYED IN THE PREPARATION OF STANDARD 
 STOCK SOLUTIONS — Continued 
 
 
 
 
 
 
 
 
 
 
 Salt to be Dis- 
 
 Gkoot Substanck 
 
 fosmular 
 Weight 
 
 Solubility of 
 Salt IN 100 Pts. 
 OF Cold Water 
 
 Per cent. 
 Metal 
 
 solved IN I 
 Liter to give 
 Strength i cc. 
 
 = 100 MG OF 
 Metal 
 
 III Al2(S04)s-l8H20 . 
 
 666 
 
 107 
 
 8.1 
 
 ( y 
 
 AiCla-eHjO . . . 
 
 242 
 
 74 
 
 II. I 
 
 ( Y 
 
 A1(N03)3-8H20 . 
 
 261 
 
 V. s. 
 
 10.3 
 
 970 
 
 Cr2(S04)3-i8H20 . 
 
 716 
 
 V. s. 
 
 14.6 
 
 690 
 
 K2Cr2(S04)4-24H20 
 
 1000 
 
 20 
 
 S-2 
 
 ( )» 
 
 Cr(N03)3-9H20 . 
 
 400 
 
 V. s. 
 
 13 
 
 770 
 
 CrCls-eHjO. . . 
 
 266.5 
 
 V. s. 
 
 19.6 
 
 Sio 
 
 FeS04-7H20 . . 
 
 278 
 
 60 
 
 20 
 
 Soo 
 
 Fe(N03)3-9H20 . 
 
 404 
 
 v.s. 
 
 14 
 
 71S 
 
 FeCl3-6H20. . . 
 
 270 
 
 V. s. 
 
 20.7 
 
 482 
 
 Ni(N03)2-6H20 . 
 
 291 
 
 so 
 
 20 
 
 500 
 
 NiClz-eHjO. . . 
 
 238 
 
 V. s. 
 
 25 
 
 400 
 
 NiS047H20 . . 
 
 280 
 
 106 
 
 21 
 
 47S 
 
 Co(N03)2-6H20 . 
 
 291 
 
 V. s. 
 
 20 
 
 500 
 
 CoClj-eHjO. . . 
 
 238 
 
 v.s. 
 
 24-S 
 
 407 
 
 CoS04-7H20 . . 
 
 281 
 
 so 
 
 31 
 
 475 
 
 MnS04-4H20 . . 
 
 223 
 
 123 
 
 2S 
 
 400 
 
 MnCl2-4H20 . . 
 
 198 
 
 ISO 
 
 28 
 
 360 
 
 Mn(N03)2-6H20 . 
 
 287 
 
 v.s. 
 
 19 
 
 527 
 
 ZnS04-7H20 . . 
 
 288 
 
 13s 
 
 22.5 
 
 445 
 
 Zn(N03)2-6H20 . 
 
 298 
 
 v.s. 
 
 22 
 
 455 
 
 ZnClz 
 
 136 
 
 V. s. 
 
 48 
 
 208 
 
 1 620g. in I liter will give strength i cc.=so mg. Al. 
 2450 g. in I liter will give strength i cc.=5o mg. Al. 
 < 192 g. in I liter will give strength i cc.=io mg, of Cr. 
 
PREPARATION OF UNKNOWNS 
 
 2IS 
 
 TABLE EMPLOYED IN THE PREPARATION OF STANDARD 
 STOCK SOLUTIONS — Continued 
 
 
 
 
 
 Quantity of 
 
 
 
 
 
 Salt to be Dis- 
 
 Group Substance 
 
 formular 
 Weight 
 
 Solubility of 
 Salt m 100 Pts. 
 OF Cold Water 
 
 Per cent. 
 Metal 
 
 solved IN I 
 
 Liter to give 
 
 Strength i cc.= 
 
 100 MG. OF 
 
 Metal 
 
 IVBaCl2-2H20. . . 
 
 244 
 
 41 
 
 56 
 
 179 
 
 Ba(C2H302)2-H20 . 
 
 273 
 
 63 
 
 50 
 
 200 
 
 Sr(N03)r4H20 . 
 
 284 
 
 40 
 
 31 
 
 324 
 
 Sr(N03)2 . . . 
 
 212 
 
 39 
 
 41.3 
 
 242 
 
 Sra2-6H20 . . . 
 
 266 
 
 106 
 
 33 
 
 304 
 
 • CaCl2 
 
 III 
 
 V. s. 
 
 36 
 
 278 
 
 Ca(N03)2-4H20 . 
 
 236 
 
 V. s. 
 
 17 
 
 590 
 
 VMgS047H20 . . 
 
 246 
 
 77 
 
 9-7 
 
 ( y 
 
 Mg(N03)2-6H20 . 
 
 256.5 
 
 200 
 
 9-4 
 
 1060 
 
 MgCl2-6H20 . . 
 
 203.5. 
 
 365 
 
 11.9 
 
 837 
 
 NaCl 
 
 58 
 
 35 
 
 40 
 
 250 
 
 Na2HP04i2H20 . 
 
 358 
 
 9.3 
 
 13 
 
 ( y 
 
 NaNOa .... 
 
 8S 
 
 80 
 
 27 
 
 371 
 
 KCI 
 
 IS 
 
 32 
 
 52 
 
 192 
 
 KHSO4 .... 
 
 136 
 
 V. s. 
 
 28.5 
 
 350 
 
 KNO3 
 
 lOI 
 
 31 
 
 39 
 
 257 
 
 NH4CI .... 
 
 S3 
 
 33 
 
 34 
 
 294 
 
 (NH4)2S04 . . . 
 
 132 
 
 76 
 
 27.5 
 
 365 
 
 NH4NO3 .... 
 
 80 
 
 200 
 
 22.5 
 
 445 
 
 (NH4)2HP04 . . 
 
 132 
 
 36.5 
 
 27.5 
 
 36s 
 
 LiCl 
 
 42 
 
 80 
 
 16.7 
 
 600 
 
 LiNOs .... 
 
 69 
 
 48 
 
 10 
 
 ( y 
 
 1 SIS g- in 
 
 I liter will give 
 
 strength i oc.= 
 
 50 mg. Mg. 
 
 
 2 77 g. in I 
 
 Q ^-. 
 
 liter will give s 
 
 "rength i cc.=i 
 
 D mg. Na. 
 
 
 * S°°S- ill I liter will give strength i cc, =50 mg. Lj. 
 
INDEX 
 
 Acetate, basic, of Al, 72. 
 Cr, 78. 
 Fe (-ic), 84. 
 separation, 93. 
 Acetates, action of heat on, 156. 
 detection, 183. 
 reactions, 156. 
 solubilities, 156. 
 Acid analysis, general examination, 165. 
 defined, 8. 
 
 preliminary examination, 160. 
 Adds, detection in minerals, 199. 
 dissociation, 12. 
 division into groups, 123. 
 first group, analytical, 168. 
 descriptive, 124. 
 preliminary examination for, 160. 
 second group, analytical, 176. 
 descriptive, 143. 
 special tests for, 168. 
 third group, analytical, 181. 
 descriptive, 156. 
 Agate mortar, 138. 
 Alcohol, amyl, rs6, 183. 
 ethyl, 156. 
 
 reduction of chromates by means of, 
 61, 142. 
 Alkalies, characteristics of, in. 
 detection in Scheme V., 119. 
 detection in silicates, 198. 
 salts of, reactions, 113. 
 solubilities, in. 
 Alkaline earths, characteristics of, loi. 
 detection in Scheme IV., 108. 
 salts of, reactions, loi. 
 solubilities, loi. 
 Alloys, analysis of, 192. 
 Aluminate of sodium, decomposition with 
 add, 71. 
 decomposition with NH4CI, 71. 
 Aluminvun, detection in scheme of analy- 
 sis, 98. 
 salts, characteristics, 70. I 
 
 Aluminum salts, reactions, 70. 
 
 solubilities, 70. 
 Ammonia, complex salts of, 43, 46, 85, 
 87, 91. 
 
 in drinking water, 116. 
 Ammonium arseno-molybdate, 52. 
 
 carbonate, reagent, 206. 
 
 chlorplatinate, 115. 
 
 -magnesium arsenate, 52. 
 
 -magnesium phosphate, 112. 
 
 phospho-molybdate, 131. 
 
 salts, detection, ng. 
 reactions, 115. 
 solubilities, 115. 
 
 sulphide, cplorless, 206. 
 yellow, 207. 
 Amyl alcohol, 156, 183. 
 Analysis, add, general, 163. 
 
 add, of minerals, 199. 
 
 for metals of all groups, 121. 
 
 of alloys, 192. 
 
 of Group in. (metals) in the presence 
 of phosphates, etc., 185. 
 
 of insoluble substances, 193. 
 
 qualitative, i. 
 
 quantitative, i. 
 Anions, 6. 
 
 Antimonic compounds, reactions, 55. 
 Antimonious compounds, reactions, 53. 
 Antimony, detection in scheme of analy- 
 sis, 69. 
 
 by Gutzeit Test, 56. 
 
 by Marsh Test, 56. 
 
 by Reinsch Test, 56. 
 Apparatus, Ust of, 211. 
 Aqua Regia, action of H2S on, 61. 
 
 as an oxidizing agent, 25, 40. 
 
 as a solvent, 25, 40, 190. 
 
 preparation of, 25, 40, 206. 
 
 removal of excess, 61, 191. 
 Arsenates, detection, 52, 69, 171. 
 
 reactions of, 51. 
 
 reduction with KI, 52. 
 
 217 
 
2l8 
 
 INDEX 
 
 Arsenates, reduction with SOj, 29, 51, 52. 
 
 special tests for, 52, 171. 
 Arsenic, adds of, 46. 
 oxides, 46. 
 
 sensitiveness of special tests for, 50. 
 Arsenic tests, special : 
 BettendorfE, 50. 
 Fleitmann, 49. 
 Gutzeit, 49. 
 Marsh, 49. 
 Reinsch, 50. 
 tests for, in reagents, 49. 
 in wall paper, 49. 
 Arsenites, detection, 49, 69, 171. 
 oxidation of, 50. 
 reactions, 47. 
 solubilities, 47. 
 special tests for, 49, 171. 
 Arsine, 49. 
 
 Barium hydroxide, reagent, 206. 
 
 salts, detection in scheme of analysis, 
 108. 
 reaction of, 102. 
 solubilities, 102. 
 
 sulphate, decomposition of, 103. 
 Base defined, 9. 
 Bases, list of, 206. 
 
 dissociation, 12. 
 Basic acetate of aluminum, 72. 
 chromium, 78. 
 iron (-ic), 84. 
 Basic acetate separation, 93. 
 Bead tests with borax, 164. 
 
 tests with NaPOs, 138, 174, 195, 196. 
 BettendorflE test for As, 50. 
 Bismuth salts, basic, 42. 
 
 reactions, 42. 
 
 solubilities, 42. 
 
 test in scheme of analysis, 66. 
 Blowpipe tests on charcoal, 163, 164. 
 Borates, detection in salts, 173. 
 
 detection in sihcates, 136. 
 
 reactions, 135. 
 
 solubilities, 135. 
 Borax bead tests, 164. 
 Boric acid, detection, 173. 
 Bromides, detection, 176. 
 
 insoluble, 145, 195. 
 
 reactions, 145. 
 
 solubilities, 145. 
 
 Cadmium salts, detection in scheme of 
 analsrsis, 66. 
 
 reactions, 46. 
 
 solubiUties, 45. 
 Calcium salts, detection in scheme of 
 analysis, 108. 
 
 reactions, 106. 
 
 sulphate, analysis of, 103. 
 Carbon, detection of, 193. 
 
 removal of, 193, 197. 
 Carbonaceous residue, 141, 161, 185. 
 Carbonates, detection, 168. 
 
 reactions, 126. 
 
 solubilities, 126. 
 Carbonic add, detection, 126, 168. 
 Cations, 6. 
 Chlorates, as oxidizing agents, 26, 91. 
 
 behavior on ignition, 159. 
 
 detection, 183. 
 
 reactions, 159. 
 
 solubilities, 159. 
 Chlorides, detection in absence of bro- 
 nudes, 177. 
 
 detection in presence of bromides and 
 iodides, 178. 
 
 insoluble, treatment of, 144, 195. 
 
 reactions, 144. 
 
 solubilities, 143. 
 Chlorine water, reagent, 207. 
 Chlorplatinate, ammonium, 115. 
 
 potassium, 113. 
 Chromates, as oxidizing agents, 26. 
 
 conversion to dichromates, 75. 
 
 detection, 170. 
 
 reactions, 142. 
 
 reduction of, 76. 
 
 solubilities, 141. 
 Chrome-iron ore, treatment of, 195. 
 Chromium, insoluble compounds, treat- 
 ment of, 143. 
 
 oxidation to diromates, 74. 
 
 reactions, 77. 
 
 salts, detection in scheme of analysis, 
 
 97- 
 solubiUties, 76. 
 Chromyl chloride test for chlorides, 
 
 178. 
 Classification of metals, 29. 
 Closed tube tests, 161. 
 Cobalt nitj-ate test for Al, 73. 
 nitrate test for Zn, 92. 
 
INDEX 
 
 219 
 
 Cobalt salts, detection in scheme of 
 analysis, 96. 
 
 oxidation to cobaltic compounds, 
 87, 88. 
 
 reactions, 87. 
 Colloidal PisSi, 47. 
 Colloidal state, 65. 
 Common ion effect, 15. 
 Conductivity, 9. 
 
 at infinite dilution, 11. 
 
 eqmvalent, 10. 
 
 molar, 10. 
 
 specific, 10. 
 Copper salts, detection in scheme of 
 analysis, 66. 
 
 reactions, 43. 
 
 solubilities, 43. 
 Csranides, action of heat on, 148. 
 
 as reducing agents, 29. 
 
 detection, 180. 
 
 reactions of simple, 149. 
 
 solubilities, 148. 
 
 Decantation, 2. 
 
 Decompositionof alkaline earth sulphates, 
 103, 194. 
 
 antimony tetroxide, 195. 
 
 cassiterite, 195. 
 
 chromium compounds (insol.), 195. 
 
 fluorspar, 195. 
 
 ignited oxides, 195. 
 
 Prussian blue, 151, 194. 
 
 silicates, 138, 195. 
 
 silver halides, 195. 
 Dehydration of silicic add, 139. 
 Dichromates, conversion to chromates, 
 76. 
 
 reduction with H2S, 61, 76. 
 
 reduction with alcohol, 61, 76. 
 Directions for laboratory work, 3. 
 
 Electrolytes, 4. 
 Electrolytic dissociation, 4. 
 Equations, compUcated, writing of, 22. 
 
 defined, 20. 
 
 factors of, 20. 
 
 methods of writing, 20. 
 
 products of, 20. 
 Ethyl alcohol, 61, 142, 183. 
 Examination, general, for adds, 165. 
 
 preliminary (general), 160. 
 
 Ferric salts, characteristics, 81. 
 
 detection in scheme of analysis, 96. 
 
 reactions, 81. 
 
 reduction to ferrous, 84. 
 Ferricyanides, action on heating, 152. 
 
 detection, 179. 
 
 reactions, 152. 
 
 solubiUties, 152. 
 Ferrocyanides, detection, 179. 
 
 reactions, 151. 
 
 solubilities, 151. 
 Ferrous salts, characteristics, 78. 
 
 detection in scheme of analysis, 96. 
 Ferrous salts, oxidation to ferric, 81. 
 
 reactions, 79. 
 Filtrate defined, 2. 
 Flame tests, 164, 196. 
 Fleitmann test for As, 49. 
 Fluorides, characteristics, 132. 
 
 detection in the absence of SiOj, 173. 
 
 detection in the presence of SiOj, 
 
 173- 
 reactions, 132. 
 solubiUties, 132. 
 Fluosihcates, detection, 170. 
 reactions of, 125. 
 
 Glass, etching of, 133. 
 
 Group reagents, 30. 
 
 Group I., adds, analytical, 168. 
 
 descriptive, 123, 124. 
 Group n., adds, analytical, 176. 
 
 descriptive, 123, 143. 
 Group m., acids, analytical, 181. 
 
 descriptive, 124, 156. 
 Group I., metals, analytical, 38. 
 descriptive, 33. 
 Group n., A, metals, analytical, 66. 
 descriptive, 40. 
 Group H., B, metals, analytical, 69. 
 descriptive, 46. 
 Group HI., metals, analytical, 96. 
 descriptive, 70. 
 Group IV., metals, analytical, 108. 
 descriptive, loi. 
 Group v., metals, analytical, 119. 
 descriptive, iix. 
 Groups, division of adds into, 123. 
 
 metals into, 29, 30,31. 
 Gutzeit test for As, 49. 
 Sb, s6. 
 
220 
 
 INDEX 
 
 Hart's method, 178. 
 Heating in dosed tube, 161. 
 on charcoal, 164. 
 with cone. H2SO4, 162. 
 Hydrogen dioxide, oxidation with, 27. 
 
 nascent, as a reducing agent, 28. 
 Hydrogen sulphide as Group H. reagent, 
 6r. 
 reducing agent, 61. 
 detection, 181. 
 Hydrolysis, salts of antimony, 54. 
 bismuth, 42. 
 iron (-ic), 81. 
 
 Insoluble substances, treatment of, 193. 
 
 scheme of analysis for, 197. 
 Introduction, i. 
 Iodides, characteristics of, 147. 
 
 detection, 176. 
 
 insoluble, treatment of, 196. 
 
 reactions, 147. 
 
 removal of, in preparation of solution, 
 190. 
 
 solubilities, 146. 
 Ions, 6. 
 Iron, see/erroMx odA ferric. 
 
 Laboratory work, directions for, 3. 
 Law of the Conservation of Elements, 2r. 
 Weight, 21. 
 Mass Action, 13. 
 Lead acetate, reagent, 207. 
 
 peroxide as an oxidizing agent, 27. 
 salts, detection in small amounts, 36. 
 scheme of analysis, 
 38, 66. 
 reactions, 35. 
 
 removal of, in the analysis of insoluble 
 substances, r97. 
 
 Magnesia mixture, reagent, 207. 
 Magnesium salts, detection in scheme of 
 analysis, irg. 
 reactions, ir2. 
 removal of, 120. 
 solubilities, 112. 
 Manganate of soditmi, 91. 
 Manganese salts, detection of small 
 amounts, 90. 
 in scheme of analysis, 96. 
 oxidation to permanganic add, 90. 
 reactions, 89. 
 
 Marsh test for As, 49. 
 Sb, s6. 
 Mass action law, 12. 
 
 and ionization, 14. 
 influence of dilution, 14. 
 Medidnal preparations, arsenic in, 49. 
 Mercmic salts, detection in scheme of 
 analysis, 66. 
 reactions, 40. 
 Mercurous salts, detection in scheme of 
 analysis, 38. 
 reactions, 34. 
 Metals, division of, into groups, 31. 
 Group I., reactions of, 33. 
 
 scheme of analysis for, 38. 
 
 Group n.. A, reactions of, 40. 
 
 scheme of analysis for, 66. 
 
 Group II., B, reactions of, 46. 
 
 scheme of analysis for, 69. 
 
 Group in., reactions of, 70. 
 
 sdieme of analysis for, 96. 
 Group IV., reactions of, loi. 
 
 scheme of analysis for, 108. 
 Group v., reactions of, in. 
 scheme of analysis for, 119'. 
 Metals and alloys, analysis of, 192. 
 Metaphosphate bead test for SiOj, 138, 
 
 174, 19s, 196- 
 Metastannic compounds, reactions of, 60. 
 
 conversion to staimic compounds, 60. 
 Metathesis, 21. 
 Minerals, analysis for adds, 199. 
 
 powdering of, 138. 
 
 Nessler's reagent, 116. 
 
 Nickel salts, characteristics of, 84. 
 
 detection in scheme of analysis, 96. 
 
 reactions, 84. 
 Nitrates, behavior on ignition, 157. 
 
 detection, 181. 
 
 distinction from chlorates, 160. 
 
 reactions, 157. 
 
 solubiUties, 157. 
 Nitric acid, as an oxidizing agent, 25. 
 
 action on ferrous salts, 25. 
 
 action on H2S, 61. 
 
 action on iodides, 190. 
 
 detection of free, 159. 
 Nitrites, behavior on ignition, 155. 
 
 detection, 182. 
 
 reactions, 15s. 
 
INDEX 
 
 221 
 
 Nitrites, solubilities, 155. 
 Non-electrolytes, 4. 
 
 Organic matter, detection, 161, 185. 
 
 interference of, 185. 
 
 removal of, 186. 
 Orthophosphoric add, 129. 
 Oxalates, behavior on ignition, 134. 
 
 detection, 172, 186. 
 
 interference of, 185. 
 
 reactions, 134. 
 
 removal of, 186. 
 
 solubilities, 134. 
 Oxidation defined, 23. 
 Oxidizing agents : 
 
 aqua regia, 25. 
 
 halogens, 24. 
 
 hydrogen dioxide, 27. 
 
 lead peroxide, 27. 
 
 list of, 24. 
 
 nitric acid, 25. 
 
 potassium chlorate, 26. 
 
 potassitmi dichromate, 26. 
 
 potassium permanganate, 27. 
 
 removal of, before precipitating with 
 H2S, 61. 
 
 sodium dioxide, 27. 
 
 Permanganates as oxidizing agents, 27. 
 Phosphate, meta-, test for SiOa, 138, r74, 
 19s, ig6. 
 separation, 188. 
 Phosphates, characteristics of, 129. 
 detection, 172, 186. 
 interference of, 185. 
 reactions, 129. 
 removal of, 188. 
 scheme of analysis in the presence of, 
 
 188. 
 solubilities, 130. 
 Phosphoric acids, 129. 
 Platinum wire, cleaning of, no. 
 Potassium chlorate, 26. 
 Potassium dichromate, 26. 
 Potassium, permanganate, 26. 
 
 salts, detection in scheme of analysis, 
 III. 
 reactions, 113. 
 solubihties, 113. 
 Powdering minerals, 138. 
 Precipitates, washing of, 2. 
 
 Precipitation, defined, 2. 
 
 completeness of, 39. 
 
 with H2S, 62. 
 Preliminary tests, closed tube, 161. 
 
 with cone. H2SO4, 162. 
 
 with dil. HCl, 162. 
 Preparation of solution for adds, 165. 
 of solids, 189. 
 
 Qualitative analysis, i. 
 Quantitative analysis, i. 
 
 Reactions of metals of Group I., 33. 
 Group n.. A, 40. 
 B,46. 
 Group in., 70. 
 Group IV., loi. 
 Group v.. III. 
 Reactions, reversible, 42, 54, 60, 72, 81, 
 83, 84, 100, 103, 156. 
 types of, 21. 
 Reagent defined, i. 
 Reagents, detection of arsenic in, 49. 
 dry, Ust of, 209. 
 group, 29. 
 Ust of, 205. 
 use and care of, 3. 
 Redudng agents, list of, 24. 
 Reduction defined, 23. 
 with carbon, 29. 
 
 hydrogen sulphide, 28. 
 nascent hydrogen, 28. 
 potassium cyanide, 29. 
 stannous chloride, 28. 
 sxdphurous add, 29, 51. 
 Reinsch test for As, 50. 
 Sb, s6, 
 Removal of bromides, 190. 
 iodides, 190. 
 organic matter, 186. 
 oxalates, 186. 
 phosphates, 188. 
 silicates, 137, 186. 
 Residue defimed, 2. 
 
 Salts, dissociation of, 12. 
 
 Scheme I., metals of the silver group, 
 
 38. 
 n., A, metals of the copper group, 
 
 66. 
 n., B, metals of the tin group, 69. 
 
222 
 
 INDEX 
 
 Scheme HI., metals of the iron group, 
 96. 
 rv., the alkaline eaiths group, 108. 
 v., the alkalies, 119. 
 Separation of 2A from 2B, 63, 64. 
 Silica, tests for, 138, 174, 186, 196. 
 SiUcates, alkalies in, 198. 
 
 decomposition by acids, 138. 
 
 by fusion with alkali carbonate, 
 
 139- 
 by HF, 140. 
 detection, 174, 186. 
 reactions, 137. 
 removal of, 137, 186. 
 Silver halides, removal of, 197. 
 salts, characteristics, 33. 
 
 detection in scheme of analysis, 
 
 38. 
 reactions, 33. 
 treatment of, 195. 
 Smith, J. L., method for alkahes, 198. 
 Sodiimi aluminate, decomposition by 
 adds, 71. 
 cobaltic nitrite, reagent, 208. 
 dioxide as an oxidizing agent, 27. 
 manganate, 91. 
 peroxide, 27. 
 phosphate bead test, 138, 174, 19s, 
 
 196. 
 plumbite test for sulphides, 154. 
 salts, detection in scheme of analysis, 
 119. 
 reactions, u6. 
 solubUities, 116. 
 stannite, reagent, 208. 
 zincate, 91. 
 Solids, preUminary testing of, i6r. 
 Solubilities, table of, 201. 
 Solubility product, 16. 
 Solution of insoluble substances, 193. 
 preparation of, for analysis, 189. 
 preparation of, for add analysis, 
 165. 
 Solvents, Ust of, 209. 
 Stannic, meta, compounds, 60. 
 salts, characteristics, 58. 
 
 detection in scheme of analjrsis, 
 
 69. 
 reactions, 58. 
 solubilities, 58. 
 Stannite of sodium, reagent, 208. 
 
 Stannous salts, characteristics of, 57. 
 
 detection in scheme of analysis, 69. 
 
 reactions, 57. 
 
 reduction by means of, 28. 
 
 solubilities, 57. 
 Starch iodide reaction, 148. 
 
 paste, preparation of, 208. 
 Stibine, 56. 
 
 Strontium salts, detection in scheme of 
 analysis, 108. 
 reactions, 105. 
 
 sulphate, decomposition of, 103. 
 Substances, insoluble, 193. 
 Sulphates, characteristics of, 124. 
 
 decomposition of alkaline earth, 103, 
 194. 
 
 detection, 169. 
 
 reactions, 124. 
 
 solubilities, 124. 
 Sulphides, behavior on ignition, 153. 
 
 detection, 154, 181. 
 
 insoluble, 154. 
 
 reactions, 154. 
 
 sodium plumbite test for, 154. 
 
 solubilities, 153. 
 Sulphites, characteristics of, 127. 
 
 detection, i6g. 
 
 reactions, 127. 
 
 solubihties, 127. 
 Sulphur, removal of, 193. 
 Sulphmic acid, detection of free, 124. 
 
 preliminary testing with, 162. 
 Synthesis, 21. 
 
 Systematic treatment of insoluble sub- 
 stances, 196. 
 
 Table for the preparation of unknowns, 
 213. 
 
 of solubilities, 202. 
 Tartrates, behavior on ignition, 141. 
 
 characteristics of, 140. 
 
 detection, 175. 
 
 reactions, 140. 
 
 solubilities, 140. 
 Thenard's green, 92. 
 Theory of Electrolytic Dissodation, 4. 
 Thio-antimonate, 56. 
 
 -antimonite, 54. 
 
 -arsenate, 48. 
 
 -arsenite, 48. 
 
 -cyanates, detection, 179. 
 
INDEX 
 
 223 
 
 Thio-cyanates, reactions, 153. 
 solubilities, 153. 
 -stannate, 57. 
 -sulphate, detection, 169. 
 reactions, 129. 
 Tin, see stannic, stannous, and meta- 
 
 starmic. 
 Turmeric paper, 210. 
 
 Unknowns, preparation of, 212. ^ 
 
 Vortmann's method, 178. 
 
 Wall paper, test for arsenic in, 49. 
 Washing precipitates, 2. 
 
 Zinc salts, characteristics of, 91. 
 
 detection of, in scheme of analysis, 96. 
 
 reactions, 91. 
 
 solubilities, 91. 
 Zinc-platinum couple, action of, on SbCU, 
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