IC-NRLF A,MANUAL OF CHEMICAL ANALYSIS Me GREGORY ] TT ,_ -J Aemm**msrw->nim*itii*aam*tmmiiiMiimimm m^jW<<ittoiortwnMM>o%vmMart< ,J* KM P GINN AND COMPANY 'XP Value Limited to ien Dollars iHtUST OF THIS BOOK IS- GIFT OF Professor Whitten / y, f|?3<f"liPrH ii-IP-nirilfioi; ft O o o^as&'-f'fSc-S^sL,,,* 4-B 13 I s-^ ^s:;;^^ * MANUAL OF QUALITATIVE CHEMICAL ANALYSIS BY J. F. McGREGORY \> PROFESSOR OF CHEMISTRY AND MINERALOGY IN COLGATE UNIVERSITY GINN & COMPANY BOSTON NEW YORK CHICAGO LONDON COPYRIGHT, 1903 BY J. F. McGREGORY ALL. RIGHTS RESERVED 47.9 J&rte* GINN & COMPANY- PRO- PRIETORS BOSTON U.S.A. T PREFACE An examination of most of the works on the subject of quali- tative analysis will show that they belong in one of two general classes which may be described as follows : first, the exhaustive treatise, of which such a work as that of Fresenius will serve as an example ; second, the abbreviated treatise, often very much abbreviated, in which the author attempts to cover the whole work in a few lessons. For graduate students, or for the better class of beginners who are able to devote the most of their time to the study, books of the first class may be invaluable. But the great majority of the students in our colleges do not become chemists. They study analytical chemistry at the most but a short time, and what they acquire is, and should be, to a considerable extent, of disciplinary value to them. To the great majority of our stu- dents, therefore, the exhaustive treatise, especially at the begin- ning of their course, is a means of confusion rather than an intelligent guide. The second class of text-books is also likely to fail with the average student, since it teaches him to analyze an unknown substance in such a purely mechanical way that the actual knowl- edge of the subject acquired is small, and the disciplinary value of the work becomes a minimum. This little manual was first written to meet the wants of the author's own classes, in which it has been used for several years, M44288 iv PREFACE and has now been carefully revised for this edition. In it the attempt has been made to retain the essentials of the larger works, omitting the rare metals and acids, and, at the same time, to avoid the "short cuts" so often found in smaller works. It is presupposed that the student has had a thorough course of instruction in general chemistry, at least through the non- metallic elements, before beginning this work. All laboratory work ought to be carried on under the imme- diate supervision of a competent instructor. The work should also be accompanied by a sufficient number of examinations to bring out all the essential points connected with the work. The author's own plan is to give frequent oral examinations throughout the course, especially in the earlier parts. In the introduction will be found certain definitions and general principles, which the student should know at the begin- ning of this course. The author has thought it best to omit all consideration of the dissociation theory, believing that, how- ever valuable the study of this subject may be to the chemist, its introduction as a basis of study in qualitative analysis is not to be recommended, and that its consideration should, therefore, be deferred until a later time when the student shall have a larger number of facts at his command. Special attention ought to be given to Parts I and II, which deal with simple substances only. The reactions here given are often partially or entirely omitted in a text-book ; but since they form the basis of all the more advanced portions of the work, they ought to be thoroughly mastered. In the separation of the metals in Part III only one prac- tical and well-established method is given, it being the author's PREFACE V experience where several methods are given, either that only one is used or that the student is likely to get them confused. Later in the course the student may learn other methods as it seems desirable. Part IV has been inserted in order to make the work more complete. It is not essential for all students, and, if it is found necessary to shorten the course, may be omitted. The appendix contains tables and some useful information for both students and instructor. The author desires to express his thanks to his assistant, Mr. R. B. Smith, and to all others who, either by suggestion or criticism, have so kindly assisted him in this work ; also to Pro- fessor R. W. Thomas for his careful reading and criticism of the manuscript. J. F. M. HAMILTON, N.Y., September, 1, 1903. CONTENTS PAGE INTRODUCTION . ... xi PART I REACTIONS FOR THE METALS IN SOLUTION 1 . Lead .1 Silver 3 i Mercury (Mercurous) ......... 4 ' Mercury (Mercuric) 6 Bismuth Copper ....... . 7 8 9 Arsenic 10 Antimony ... Tin (Stannous) .... Tin (Stannic) .... . Aluminum ...... Chromium .......... . 10 11 . 12 13 . 14 Iron (Ferrous) Iron (Ferric) Nickel 15 . 16 17 Cobalt . . ... . 19 Manganese Zinc 20 . 22 Magnesium ...... . Barium 23 . 24 Strontium 25 Calcium ".***&*'.'''' - Potassium . 26 27 Sodium . 28 .Ammonium . . 29 REACTIONS FOR THE ACID RADICALS IN SOLUTION . 30 Hydrochloric Acid ........ Hydrobromic Acid ........ vii 31 . 31 viii CONTENTS REACTIONS FOR THE ACID RADICALS IN SOLUTION Continued PAGE Hydriodic Acid ..... ... 32 Hydrofluoric Acid ... 32 Hydrocyanic Acid 33 Sulfocyanic or Thiocyanic Acid 34 Hydroferrocyanic Acid 34 Hydroferricyanic Acid 34 Hypochlorous Acid ......... 35 Chloric Acid .......... 35 Hydrogen Sulfid (Hydrosulfuric Acid) . . . . . 36 Thiosulfuric Acid 36 Sulfurous Acid ......... 37 Sulfuric Acid 38 Chromic Acid . . 38 Nitrous Acid ...... 39 Nitric Acid . . 40 Phosphoric Acid . 41 Arsenious Acid . . 41 Arsenic Acid . . . . - . . . . .42 Boric Acid . . 43 Carbonic Acid .......... 43 Silicic Acid .......... 44 Acetic Acid . . . . . . . . . . . 45 Oxalic Acid .......... 46 Tartaric Acid .......... 46 PART II REACTIONS FOR DRY SUBSTANCES BLOWPIPE ANALYSIS . . 48 The Effect of Heat alone 50 The Substance is heated on Charcoal ..... 55 The Substance is heated on Charcoal with Sodium Carbonate . 58 Coloration of the Flame . . . . . . 59 Coloration of the Borax or Microcosmic Bead . . . .60 The Substance is fused on Platinum Foil with Sodium Carbonate and Potassium Nitrate ....... 62 The Substance is acted upon by Sulfuric Acid . . . .62 Special Tests 66 CONTENTS ix PART III SYSTEMATIC EXAMINATION FOR METALS IN SOLUTION PAGE SIMPLE COMPOUNDS 69 Group 1 . . . 69 Group 2 ... 70 Group 3 ... 72 Group 4 . 73 Group 5 73 Group 6 .*....'. . .74 Examination for Acid Radicals ...... 75 SYSTEMATIC EXAMINATION FOR METALS IN SOLUTION MIXED COMPOUNDS 76 Preliminary Examination 77 Group 1 Lead, Silver, Mercury (Mercurous) . . . .78 Group 2 Mercury (Mercuric), Bismuth, Copper, Cadmium, Ar- senic, Antimony, Tin ....... 80 Group 2, Subdivision A 82 Group 2, Subdivision B 84 Group 3 Aluminum, Chromium, Iron 86 Phosphates, Oxalates, etc., are absent .... 88 Phosphates, Oxalates, etc., are present 90 Group 4 Nickel, Cobalt, Manganese, Zinc .... 93 GroupS Barium, Strontium, Calcium 96 Group 6 Magnesium, Potassium, Sodium, Ammonium . 98 SYSTEMATIC EXAMINATION FOR ACID RADICALS IN SOLUTION Preliminary Examination ...... .101 Preparation of the Solution . . . . . . . 102 Classification of the Acid Radicals .... 103 Acids : Group 1 104 Acids : Group 2 . . 107 Acids: Group 3 . .... 110 x CONTENTS PATTT IV PAGE SYSTEMATIC EXAMINATION OF COMPLEX SOLIDS . .111 Preliminary Examination 112 I. THE SUBSTANCE is A METAL OK AN ALLOY .... 113 A. Metals insoluble and unchanged in nitric acid . . . 113 B. Metals which form insoluble oxids by the action of nitric acid 114 C. Metals and alloys soluble in nitric acid . . . . 114 II. THE SUBSTANCE is NEITHER A METAL NOR AN ALLOY . . 115 A. The substance is partially or entirely soluble in water . 116 B. The substance is insoluble in water . . . . .116 C. The substance is insoluble in H 2 O and in HC1 . . 117 D. The substance is insoluble in H 2 O and in both HC1 and HNO 3 1 17 E. The substance is insoluble in H 2 O and in all acids . . 118 F. The substance is a silicate 120 (a) Silicates decomposed by acids 121 (b) Silicates not decomposed by acids . . . .121 G. Cyanids are present 122 APPENDIX NAMES, SYMBOLS, AND ATOMIC WEIGHTS OF THE ELEMENTS . 125 NAMES AND FORMULAS OF REAGENTS AND SOLUTIONS . . 126 PREPARATION OF REAGENTS AND SOLUTIONS .... 129 INDEX ....... 131 INTRODUCTION Analytical Chemistry treats of the composition of substances and of the methods by which we determine the same. There are two general divisions of analytical chemistry, viz. : qualita- tive and quantitative analysis. Qualitative Analysis has for its object the determination of the constituent elements of a body. It consists in the separa- tion of each of the elements, either in the free state, or, as more commonly happens, in the form of some compound which is characteristic and easily recognized. Quantitative Analysis belongs to a more advanced course, and has for its object the determination of the percentage amounts of the constituents of a body, and thus of the actual constitution of the body. Every simple inorganic substance consists of two parts. The first, which is a metal or positive radical, is chemically com- bined with the second, which is a non-metal or negative radical. The more complex substances may contain several metals, or positive radicals, and often contain more than one acid, or nega- tive radical. These may be chemical combinations or merely mechanical mixtures. By subjecting a substance to various conditions we obtain a series of phenomena which we call its reactions ; and any known substance which is employed in effecting a reaction is called a reagent. The subjecting of a substance to the action of reagents, by means of which its constituent elements are recognized, is the process employed in qualitative analysis. We may subject the substance to the action of reagents either in its original solid condition if it be a solid or in xi xii INTRODUCTION solution. These two methods of examination are known as the dry way and the wet way. The dry way may be employed for the complete analysis of simple substances, and is a valuable aid in making preliminary tests of complex substances. This comprises what is known as Blowpipe Analysis, which is fully explained in Part II of this work. The wet way is more generally used in qualitative analysis, because its reactions are, for the most part, simpler and more rapid. It can be employed with all kinds of substances. Solids and gases can generally be obtained in solution in water, or some other convenient liquid, in which they will dissolve without losing their characteristic properties. When a substance in solution is acted upon by a reagent the results are always in accordance with a law which may be stated as follows: When two substances which are in contact in solution can, under the conditions of the reaction, form a sub- stance which is insoluble or volatile, the insoluble or volatile substance will always be formed and continue to be formed until one of the factors is exhausted. An insoluble compound thus formed is called a precipitate, and precipitation is the most common form of reaction in analytical chemistry. A precipitate may be of almost any color and the color may be characteristic, or if, as is more often the case, it is not especially so, other precipitates are formed with other reagents until the combination of results is such as to determine the substance with certainty. Sometimes a precipitate which is not characteristic becomes so by its solubility or insolubility in some other reagent or in an excess of the reagent first used. Solubility in excess is confined to a few reagents. These are the alkaline hydroxids and a few alkaline salts. The treatment of a substance with a reagent sometimes results in the formation of a gas, which is recognized by its color or odor or by some other characteristic property. This INTRODUCTION xiii almost always results from the decomposition of some acid radical by means of some acid used as a reagent. In order to avoid constant mistakes, the student should thoroughly understand every reaction which he uses. In acquiring this knowledge he should, especially in the earlier and simpler portion of the course, accustom himself to the use of all the common reagents. In the more advanced portion of his course a more systematic and selective use of reagents will be necessary. The student should first be required to perform all the reactions in Part I. He should write out the chemical equation in every case in a suitable notebook, and at the time the reaction is made. He may then be given some simple solutions and, by the use of the reactions he has just been performing, find what the unknown solution contains. Having found the substance contained in a solution, he should try all the reactions given for that substance. He should also compare each reaction with all similar reactions given by other substances, noting points of difference, and in this way make each reaction as comprehensive as possible. Many students seem to feel that all that is required of them is to find out what is in the unknown solution. This is undoubtedly the goal toward which their course of study is tending ; but if, at the beginning of his course, that is all that the student desires, a much shorter method would be to ask the instructor. The value of an elementary course of study in analytical chemistry is not simply to find out what a solution or solid substance contains, but to learn how to find out what it contains. It will be observed that in Part III all tables such as are often included in a text-book and intended for aid in the sepa- ration of the metals have been omitted. This is because, in the author's judgment, the use of them in the hands of the majority of students is pernicious. When such tables are used the student almost always depends upon them rather than upon xiv INTRODUCTION the full text, and so is frequently led into error because of some- thing which has been omitted. The information given in a table is necessarily brief, and the average student can acquire such information by a few hours of study. With such informa- tion in his head he will work much faster and with more satis- faction to himself and his instructors, and when in doubt he will consult the text and not a table. Every student should be required to keep a notebook in which to record the results of all his work in the laboratory. He should also be encouraged, if not required, to make use of some large and fairly complete text-book on general chemistry for collateral reading, especially in connection with those com- pounds which he meets with in the course of his work. Such a course of action, if persisted in, will give to the student a much more comprehensive view of the subject, and will, in addition, provide him with a fund of information which will always be of value to him. The instructor should require the student to do clean and careful work. Work done in a careless way, with dirty appa- ratus and on a dirty desk, is of little or no value. Clean, intel- ligent work, accompanied by a reasonable amount of reading and study, will give to the student, even if his course is only a short one, a glimpse at least of the immense and interesting field of study and research which is always open to the chemist QUALITATIVE ANALYSIS PART I REACTIONS FOR THE METALS IN SOLUTION LEAD, Pb" Lead dissolves easily in HNO 3 with formation of Pb(NO 3 ) 2 . 3 Pb + 8 HN0 3 = 3 Pb(N0 8 ) 2 + 4 H 2 + 2 NO. It dissolves in hot concentrated H 2 SO 4 . Pb + 2 H 2 S0 4 = PbS0 4 -f 2 H 2 + S0 2 . It is not attacked by dilute H 2 SO 4 or HC1. For the reactions use lead nitrate, Pb(NO 3 ) 2 . 1. Sodium Hydroxid precipitates white Pb(OH) 2 or a white basic hydroxid, Pb 2 O(OH) 2 , according to the conditions which exist. Pb(N0 3 ) 2 + 2 NaOH = Pb(OH) 2 + 2 NaN0 3 . 2 Pb(N0 3 ) 2 + 4 NaOH = Pb 2 0(OH) 2 + 4 NaNO 3 + H 2 0. Soluble in excess of the reagent (4 vols.), easily soluble in con- centrated NaOH, forming sodium plumbite, Na 2 PbO 2 . Pb(OH) 2 + 2 NaOH = Na 2 PbO 2 + 2 H 2 0. 2. Ammonium Hydroxid precipitates a white basic salt, (PbO) 2 Pb(N0 3 ) 2 . 3 Pb(N0 8 ) 2 -1- 4 NH 4 OH = (PbO) 2 Pb(N0 3 ) 2 + 4 NH 4 N0 8 -f 2 H 2 0. Insoluble in excess of the reagent. 1 2 ... . . QUALITATIVE ANALYSIS 3. .Sodium /xr Ammonium. Carbonate precipitates white PbCO 3 . If the solution is hot a white basic carbonate, Pb 3 (OH) 2 (CO 3 ) 2 , is precipitated. This is known commercially as white lead. 3Pb(N0 3 ) 2 + 3Na 2 C0 3 + H 2 = Pb 3 (OH) 2 (C0 3 ) 2 + 6NaN0 3 + C0 2 . 4. Hydrogen or Ammonium Sulfid precipitates black PbS. Pb(N0 3 ) 2 + (NH 4 ) 2 S = PbS + 2 NH 4 N0 8 . Insoluble in cold dilute acids. Soluble in warm dilute HNO 3 , forming Pb(NO 3 ) 2 and free sulfur. 3 PbS + 8 HN0 3 = 3 Pb(N0 3 ) 2 + 4 H 2 -f 2 NO + 3 S ? Concentrated HNO 3 oxidizes PbS to PbSO 4 . 3 PbS + 8 HN0 3 = 3 PbS0 4 + 4H 2 + 8 NO. If the HNO 3 is of medium strength both reactions will go on at the same time. 5. Acid Sodium Phosphate precipitates white Pb 3 (PO 4 ) 2 . 3 Pb(N0 3 ) 2 + 2 Na 2 HP0 4 = Pb 3 (P0 4 ) 2 + 4 NaN0 3 + 2 HN0 3 . Easily soluble in HNO 3 . 6. Potassium Cyanid precipitates white Pb(CN) 2 . Pb(N0 3 ) 2 + 2 KCN = Pb(CN) 2 + 2 KN0 3 . Insoluble in excess of the reagent. 7. Potassium lodid precipitates yellow PbI 2 . Pb(N0 3 ) 2 + 2 KI = PbI 2 + 2 KN0 3 . The precipitate is soluble in boiling water (4 vols.), from which solution it crystallizes, on cooling, in golden yellow scales. 8. Potassium Chromate precipitates yellow PbCrO 4 . Pb(N0 3 ) 2 + K 2 Cr0 4 = PbCr0 4 + 2 KN0 3 . Soluble in NaOH (5 vols. of dilute or 1 vol. of concentrated) and in HNO 3 . Insoluble in acetic acid. REACTIONS FOR THE METALS IN SOLUTION 3 9. Potassium Ferrocyanid precipitates white Pb 2 Fe(CN) 6 . 2 Pb(N0 3 ) 2 + K 4 Fe(CN) 6 = Pb 2 Fe(CN) 6 + 4 KN0 3 . 10. Hydrochlorid Acid, or any soluble chlorid, precipitates white PbCl 2 . Pb(N0 3 ) 2 + 2 HC1 = PbCl 2 + 2 HN0 3 . Easily soluble in boiling water, from which, unless the solution is too dilute, it will crystallize, on cooling, in long white needles. 11. Sulfuric Acid, or any soluble sulfate, precipitates white PbSO 4 . Pb(N0 3 ) 2 4- H 2 S0 4 = PbS0 4 + 2 HN0 3 . Easily soluble in ammonium tartrate or ammonium acetate. [Add tartaric or acetic acid, and then excess of NH 4 OH.] 12. Metallic Zinc will entirely precipitate the lead in crystalline form. Pb(N0 3 ) 2 + Zn = Pb + Zn(N0 3 ) 2 . NOTE. The chemical equations have been given under Lead as a guide for the student. Such equations will generally be omitted, but the student should be required to write them for himself. SILVER, Ag' Silver dissolves easily in HNO 3 with formation of AgNO 3 . Insoluble in HC1 and in H 2 SO 4 . For the reactions use silver nitrate, AgNO 3 . 1. Sodium Hydroxid precipitates brown Ag 2 O. Insoluble in excess. Soluble in NH 4 OH, forming NH 4 AgO. 2. Ammonium Hydroxid precipitates the same. Very easily soluble in excess, forming NH 4 AgO. [For the formation of this precipitate dilute the NH 4 OH with 10 vols. of H 2 O, and use only one or two drops of the reagent.] If the silver solu- tion is very acid no precipitate will be formed. 3. Sodium Carbonate precipitates light yellow Ag 2 CO 3 . Insol- uble in excess. Soluble in NH 4 OH and in (NH 4 ) 2 CO 3 . 4 QUALITATIVE ANALYSIS 4. Hydrogen or Ammonium Sulfid precipitates black Ag 2 S. Soluble in HNO 3 . Insoluble in NH 4 OH. 5. Acid Sodium Phosphate precipitates yellow Ag 3 PO 4 . Sol- uble in NH 4 OH and in HNO 3 . 6. Potassium Cyanid precipitates white AgCN. Soluble in excess, forming AgCN(KCN). From this solution HNO 3 pre- cipitates AgCN. 7. Potassium lodid precipitates light yellow Agl. Only very slightly soluble in NH 4 OH, but easily soluble in KCN. Insoluble in HNO 3 . 8. Potassium Chromate precipitates red-brown Ag 2 CrO 4 . Sol- uble in HNO 3 and in NH 4 OH. 9. Potassium Sulfocyanate precipitates white AgSCN. Sol- uble in NH 4 OH. 10. Potassium Ferrocyanid precipitates white Ag 4 Fe(CN) 6 . Difficultly soluble in NH 4 OH. 11. Hydrochloric Acid, or any Soluble Chlorid, precipitates white AgCl. Soluble in NH 4 OH, forming (NH 3 ) 3 (AgCl) 2 , and in KCN. From these solutions HNO 3 reprecipitates the AgCl. 12. Metallic Zinc, Copper, or Mercury will precipitate the silver in crystalline form. 13. Reducing Agents, such as sulfurous acid, stannous chlorid, or ferrous sulfate, will precipitate the silver as a fine gray powder. MERCURY (Mercurous), Hg' Mercury dissolves easily in HNO 3 . If the mercury is in excess there is formed mercurous nitrate, HgNO 3 ; but if the HNO 3 is in excess, mercuric nitrate, Hg(NO 3 ) 2 , is formed. Mercury dissolves in hot concentrated H 2 SO 4 , forming mer- curic sulfate, HgSO 4 , and SO 2 . It is insoluble in HC1. For the reactions use a solution of HgNO 3 . REACTIONS FOR THE METALS IN SOLUTION 5 1. Sodium Hydroxid precipitates black Hg 2 O. Insoluble in ex- cess of the reagent. Decomposed by boiling into HgO and Hg. 2. Ammonium Hydroxid, or Ammonium Carbonate, precipitates a black mixture of amido-mercuric nitrate, HgNH 2 NO 3 , and finely divided mercury. 3. Sodium Carbonate precipitates a brownish-black basic carbonate. 4. Hydrogen or Ammonium Sulfid precipitates black HgS mixed with Hg. Soluble in aqua regia. If this precipitate is boiled with concentrated HNO 3 it is changed into a white basic compound, Hg(NO 3 ) 2 (HgS) 2 . 5. Potassium lodid precipitates yellowish-green Hgl. If an excess of the reagent is added potassium mercuric iodid, HgI 2 (KI) 2 , is formecl. This dissolves in the liquid, while metallic mercury is precipitated as a gray powder. 6. Potassium Chromate precipitates brick-red Hg 2 CrO 4 . Diffi- cultly soluble in HNO 3 . 7. Hydrochloric Acid, or a Soluble Chlorid, precipitates white HgCl (calomel). The addition of NH 4 OH changes this precipi- tate to a black mixture of amido-mercuric chlorid, HgNH 2 01, and finely divided mercury. 8. Sulfuric Acid, in not too dilute solutions, precipitates white Hg 2 S0 4 . 9. Stannous Chlorid, in very small quantity, precipitates white HgCl. In excess of the reagent the precipitate is reduced to gray metallic mercury. 10. Metallic Copper, or Zinc, in solutions slightly acidified with HC1, precipitates metallic mercury. 11. Sulfurous Acid reduces a mercurous solution to metallic mercury, which can often be collected in a globule by boiling with HCL 6 QUALITATIVE ANALYSIS MERCURY (Mercuric), Hg" Mercury and most mercurous compounds can be changed to mercuric by heating with concentrated HNO 3 . For the reactions use a solution of HgCl 2 , or Hg(NO 3 ) 2 . 1. Sodium Hydroxid precipitates yellow Hg(OH) 2 . Insoluble in excess. Soluble in warm acids. If only a few drops of the reagent are used a brown basic compound is formed. 2. Ammonium Hydroxid, or Ammonium Carbonate, precipitates white amido-mercuric chlorid, HgNH 2 Cl. 3. Sodium Carbonate precipitates a brown basic carbonate, HgC0 3 (HgO) 8 . 4. Hydrogen or Ammonium Sulfid precipitates at first a white double salt, HgCl 2 (HgS) 2 . Excess of the reagent causes this precipitate to change to yellow, orange, and finally to black HgS. Insoluble in concentrated HNO 3 . Long-continued boil- ing with HNO 3 changes it into a white basic compound. [See Mercurous 4.] 5. Potassium lodid precipitates scarlet-red HgI 2 . Easily sol- uble in excess, forming HgI 2 (KI) 2 . The precipitate is first yellow, then salmon red, and finally scarlet red. 6. Potassium Chromate precipitates, from not too dilute solu- tions, orange red HgCrO 4 . 7. Metallic Copper precipitates, from solutions acidified with HC1, gray metallic mercury. 8. Stannous Chlorid precipitates first white mercurous chlorid, HgCl, and in excess, gray metallic mercury. Other reducing agents produce a similar change. REACTIONS FOR THE METALS IN SOLUTION 7 BISMUTH, Bi'" Bismuth dissolves easily in HNO 3 , in hot concentrated H 2 SO 4 , and in aqua regia, but not in HC1. The ordinary bismuth salts are not soluble in H 2 O except in the presence of considerable free acid, usually HNO 3 or HC1. For the reactions use a solution of Bi(NO 3 ) 3 , in dilute HNO 3 . 1. Water in large quantities precipitates, if too much free acid is not present, white basic bismuth nitrate, Bi(OH) 2 NO 3 . If much free HNO 3 is present the addition of ammonium chlorid, or of HC1, will cause the precipitation of white BiOCl. These precipitates are insoluble in tartaric acid. [See Antimony 1.] 2. Sodium or Ammonium Hydroxid precipitates white Bi(OH) 3 . Insoluble in excess. Changed by boiling to yellow Bi 2 O 3 . 3. Sodium or Ammonium Carbonate precipitates white basic bis- muth carbonate (BiO) 2 CO 3 . 4. Hydrogen or Ammonium Sulfid precipitates dark brown Bi 2 S 3 . Insoluble in (NH 4 ) 2 S. Soluble in HNO 3 . 5. Acid Sodium Phosphate precipitates white BiPO 4 . Insoluble in dilute acids. 6. Potassium lodid precipitates brown BiI 3 . Soluble in excess. 7. Potassium Chromate precipitates yellow basic bismuth chro- mate 2 [(BiO) 2 CrO 4 ]Bi 2 O 3 . Insoluble in NaOH. [See Lead 8.] Soluble in HNO 3 . 8. Sodium Stannite, which is formed by adding NaOH to a solution of SnCl 2 until the precipitate first formed is dissolved, reduces the bismuth solution and forms a black precipitate, which is a mixture of Bi and Bi 2 O 3 ; or, if the reagent is added in large excess, and hot, it precipitates metallic bismuth. 8 QUALITATIVE ANALYSIS COPPER, Cu" Copper dissolves easily in HNO 3 , forming Cu(NO 3 ) 2 and NO, and in hot concentrated H 2 SO 4 , forming CuSO 4 and S^ 2 * ^ is only very slightly soluble in dilute H 2 SO 4 or HC1. For the reactions use a solution of CuSO 4 . 1. Sodium Hydroxid precipitates light blue Cu(OH) 2 . Insol- uble in excess of the reagent. Soluble in NH 4 OH. If the precipitate is boiled with an excess of the reagent it becomes black, owing to the formation of Cu(OH) 2 (CuO) 2 . 2. Ammonium Hydroxid precipitates a light blue basic salt. Easily soluble in excess of the reagent, forming CuSO 4 (NH 3 ) 4 H 2 O, which gives a deep-blue color to the solution (a very character- istic reaction). If KCN is added to this blue solution the color disappears, owing to the formation of Cu(CN) 2 (KCN) 2 . 3. Sodium Carbonate precipitates a blue basic carbonate, Cu 2 (OH) 2 CO 3 . On boiling, this precipitate loses CO 2 and forms black Cu(OH) 2 (CuO) 2 . 4. Hydrogen or Ammonium Sulfid precipitates black CuS. Sol- uble in KCN and in HNO 3 . Insoluble in dilute H 2 SO 4 . [See Cadmium 4.] 5. Acid Sodium Phosphate precipitates greenish-blue Cu 3 (PO 4 ) 2 . Soluble in NH 4 OH. 6. Potassium Cyanid precipitates greenish-yellow Cu(CN) 9 . Easily soluble in excess of the reagent, forming Cu(CN) 2 (KCN) 2 . From this solution H 2 S will not precipitate the copper. [See Cadmium 6.] 7. Potassium lodid precipitates white Cu 2 T 2 and free iodin. The latter colors the precipitate brown. If H 2 SO 3 is added the precipitate appears white. REACTIONS FOR THE METALS IN SOLUTION 9 8. Potassium Sulfocyanate precipitates black Cu(SCN) 2 . If H 2 SO 3 is added in excess the copper is reduced and white Cu a (SCN) 2 is formed. 9. Po. jsium Ferrocyanid precipitates red-brown Cu 2 Fe(CN) 6 . [This is an exceedingly delicate reaction, one part of copper showing a reddish coloration in 200,000 parts of water.] 10. Metallic Iron precipitates copper from a solution. CADMIUM, Cd" Cadmium dissolves easily in HNO 3 , and slowly in H 2 SO 4 and HC1. For the reactions use a solution of Cd(NO 3 ) 2 . 1. Sodium Hydroxid precipitates white Cd(OH) 2 . Insoluble in excess of the reagent. 2. Ammonium Hydroxid precipitates the same compound. Easily soluble in excess of the reagent. 3. Sodium or Ammonium Carbonate precipitates white CdCO 3 . Insoluble in excess of the reagent. Soluble in NH 4 OH. 4. Hydrogen or Ammonium Sulfid precipitates yellow CdS. Insoluble in (NH 4 ) 2 S or in KCN. Soluble in warm dilute H 2 SO 4 or HNO 8 . [See Copper 4.] 5. Acid Sodium Phosphate precipitates white Cd 3 (PO 4 ) 2 . Sol- uble in NH 4 OH and in dilute acids. 6. Potassium Cyanid precipitates white Cd(CN) 2 . Soluble in excess of the reagent, with formation of Cd(CN) 2 (KCN) 2 . From this solution H 2 S precipitates CdS. [See Copper 6.] 7. Potassium Chromate precipitates a yellow basic chromate Cd 2 (OH) 2 CrO 4 . Since this precipitate forms slowly, use a slight excess of the reagent and allow it to stand for a few minutes. Insoluble in NaOH. [See Lead 8.] 10 QUALITATIVE ANALYSIS ARSENIC, As'" This element exists in both trivalent and pentavalent rela- tions, and has very few metallic properties. It does not dissolve in the acids to form salts, and we have already learned that it forms acids quite analogous to those of phosphorus. The principal reactions of arsenic are therefore to be found among those of the acids. Arsenious oxid, As 2 O 3 , dissolves in HC1, and this solution may be used for the reactions. 1. Hydrogen Sulfid precipitates yellow As 2 S 3 . Insoluble in HC1. Soluble in NH 4 OH or (NH 4 ) 2 CO 3 . It dissolves in (NH 4 ) 2 S, forming ammonium sulfarsenite, (NH 4 ) 3 AsS 3 , and in yellow ammonium sulfid, (NH 4 ) 2 S X , forming ammonium sulfarsenate, (NH 4 ) 3 AsS 4 . HC1 precipitates from these solutions, in the first case As 2 S 3 , in the second As 2 S 5 . The other reagents for the metals give no precipitates with arsenic solutions. If a solution containing pentavalent arsenic is treated with H 2 S it is reduced, and the arsenic precipitated as As 2 S 3 , together with sulfur. This action takes place very slowly in a cold solution, but is immediate if the solution is hot. A solution of sodium arsenate, Na 3 AsO 4 , may be used for this reaction. ANTIMONY (Stibium), Sb'" Antimony forms both trivalent and pentavalent compounds. It does not dissolve in HC1. With hot concentrated H 2 SO 4 it forms Sb 2 (SO 4 ) 3 . With dilute HNO 3 it forms Sb 2 O 3 , and with concentrated HNO 3 it forms metantimonic acid, HSbO 3 . It dissolves in aqua regia, forming SbCl 3 or SbCl 5 , according to the degree of concentration of the acids and the duration of the action. REACTIONS FOR THE METALS IN SOLUTION 11 These compounds do not dissolve in water unless free hydro- chloric or tartaric acid is present. For the reactions use a solution of SbCl 3 . 1. Water, added in excess, precipitates white antimony oxychlorid, SbOCl. Soluble in tartaric acid, so that if much of this acid is present the precipitation may not take place. 2. Sodium Hydroxid precipitates white SbOOH. Soluble in excess of the reagent, with formation of sodium metantimonite, NaSb0 2 . 3. Ammonium Hydroxid, or Sodium or Ammonium Carbonate, pre- cipitates the same. Insoluble in excess of the reagent. 4. Hydrogen Sulfid precipitates orange-red Sb 2 S 3 . Insoluble in (NH 4 ) 2 CO 3 [See Arsenic 1] and in dilute acids. Soluble in warm concentrated II Cl. Soluble also in (NH 4 ) a S, forming ammonium sulfantimonite, (NH 4 ) 3 SbS 3 , and in (NH 4 ) 2 S X , form- ing ammonium sulfantimonate, (NH 4 ) 3 SbS 4 . HC1 precipitates from these solutions Sb 2 S 3 and Sb 2 S 5 respectively. 5. Metallic Zinc, in solutions containing free HC1, precipitates the antimony as a black powder. If a piece of platinum foil is placed in the solution, in contact with the zinc, the antimony will be precipitated on the foil as a black stain. Antimony, in its pentavalent relations, is acid in its prop- erties. From such compounds the antimony may be precipi- tated by hydrogen sulfid as orange-red Sb 2 S 5 . This dissolves in warm concentrated HC1, forming SbCl 3 and precipitating sulfur. TIN (Stannous), Sn" Tin forms both stannous (Sn") and stannic (Sn"") com- pounds. It dissolves in HC1, forming SnCl 2 ; and in H 2 SO 4 , forming SnSO 4 . With very dilute HNO 3 it forms Sn(NO 3 ) 2 , some of the acid being reduced forming NH 4 NO 3 . Thus : 4 Sn + 10 HN0 3 = 4 Sn(N0 3 ) 2 + NH 4 N0 3 + 3 H 2 0. 12 QUALITATIVE ANALYSIS With concentrated HNO 3 it forms white stannic acid, H 2 SnO 3 . It dissolves in aqua regia, forming SnCl 4 . For the reactions use a solution of SnCl 2 . 1. Sodium Hydroxid precipitates white Sn(OH) 2 . Soluble in excess of the reagent, forming sodium stannite, Na 2 SnO 2 . 2. Ammonium Hydroxid, or Sodium or Ammonium Carbonate, pre- cipitates the same. Insoluble in excess of the reagent. 3. Hydrogen or Ammonium Sulfid precipitates dark brown SnS. Insoluble in (NH 4 ) 2 CO 3 [See Arsenic 1], and in (NH 4 ) 2 S if free from (NH 4 ) 2 S X . Soluble in NaOH, in HC1, and in (NH 4 ) 2 S X , forming with the latter ammonium sulfostannate, (NH 4 ) 2 SnS 3 . From this solution HC1 precipitates yellow stan- nic sulfid, SnS 2 . Stannous chlorid, and all other stannous compounds, are easily oxidized to stannic compounds. They act, therefore, as powerful reducing agents when in the presence of reducible compounds. Silver salts are reduced to metallic silver, and mer- cury salts to metallic mercury. [See Silver 13, Mercurous 9, and Mercuric 8.] Bismuth compounds are reduced to metallic bismuth. [See Bismuth 8.] Ferric compounds are changed to ferrous compounds. [See Ferric 10.] Potassium chromate, K 2 CrO 4 , and potassium permanganate, KMnO 4 , are reduced to chromium chlorid and manganese chlorid respectively. 2 K 2 Cr0 4 + 16 HC1 + 3 SnCl 2 = 4 KC1 + 2 CrCls + 3 SnCl 4 + 8 H 2 0. Many other compounds give a similar reaction. TIN (Stannic), Sn"" Stannic compounds decompose on standing and precipitate Sn(OH) 4 . If stannous chlorid be acidified with HC1, a few crystals of potassium chlorate added, and the whole boiled until the chlorous odors are driven away, the SnCl 4 thus formed may be used for the following reactions. REACTIONS FOR THE METALS IN SOLUTION 13 1. Sodium Hydroxid precipitates white stannic acid, H 2 SnO 3 . Soluble in excess of the reagent, forming sodium stannate, Na 2 SnO 3 . Soluble also in the mineral acids. 2. Ammonium Hydroxid, or Sodium or Ammonium Carbonate, pre- cipitates the same. Insoluble in excess of the reagent. 3. Hydrogen Sulfid precipitates from solutions which do not contain too large an excess of HC1, yellow SnS 2 . Soluble in concentrated HC1, and in (NH 4 ) 2 S, forming ammonium sulfostan- nate, (NH 4 ) 3 SnS 3 . Insoluble in (NH 4 ) 2 CO 3 . [See Arsenic 1.] 4. If a solution of stannic chlorid is boiled in the presence of some neutral salt, such as sodium sulfate or ammonium nitrate, metastannic acid, H 10 Sn 5 O 15 , is precipitated. This compound is a polymeric form of stannic acid, and is insoluble in HNO 3 or H 2 SO 4 . ALUMINUM, Al'" Aluminum dissolves easily in HC1, with some difficulty in H 2 SO 4 , and scarcely at all in HNO 3 . It dissolves also in NaOH and in KOH, liberating hydrogen. For the reactions use a solution of A1 2 (SO 4 ) 3 . 1. Sodium Hydroxid precipitates white A1(OH) 3 . Easily soluble in excess of the reagent, forming sodium aluminate, NaAlO 2 . From this solution it is reprecipitated by NH 4 C1. Soluble in all mineral acids and in acetic acid.* 2. Ammonium Hydroxid precipitates the same. Very slightly soluble in excess of the reagent, but reprecipitated by boiling. If NH 4 C1 is present the precipitate is not dissolved in excess of the reagent.* 3. Sodium or Ammonium Carbonate precipitates the same, liber- ating CO 2 .* 4. Ammonium Sulfid precipitates the same, liberating H 2 S.* * The presence of non-volatile organic substances, such as tartaric acid, citric acid, sugar, etc., prevents this precipitation. 14 QUALITATIVE ANALYSIS 5. Acid Sodium Phosphate precipitates white A1PO 4 . Soluble in mineral acids and in NaOH. From the solution in NaOH, NH 4 C1 precipitates the aluminum as A1(OH) 3 . 6. Sodium Acetate gives no precipitate if the solution is cold, but if added in large excess and boiled the aluminum is com- pletely precipitated as basic aluminum acetate, A1(OH) 2 (C 2 H 3 O 2 ). The solution must be neutral. If acid, neutralize with Na 2 CO 3 or NaOH. 7. Barium Carbonate precipitates white A1(OH) 3 , liberating CO 2 . CHROMIUM, Cr"' Chromium dissolves in HC1 and H 2 SO 4 , but is not soluble in HN0 3 . For the reactions use a solution of Cr 2 (SO 4 ) 3 . 1. Sodium Hydroxid precipitates gray-green Cr(OH) 3 . Soluble in excess of the reagent, giving a dark green solution and form- ing sodium chromite, NaCrO 2 . Reprecipitated by boiling or by the addition of NH 4 C1. 2. Ammonium Hydroxid precipitates the same. ^The precipitate is slightly soluble in excess of the reagent whe% cold and con- centrated, giving a reddish color to the solution. Reprecipi- tated by boiling or by the addition of NH 4 C1. 3. Sodium or Ammonium Carbonate precipitates the same, libera- ting CO 2 . The precipitate often contains some basic chromium carbonate of variable composition. 4. Ammonium Sulfid precipitates the same, liberating H 2 S. 5. Acid Sodium Phosphate precipitates gray-green CrPO 4 . Sol- uble in the mineral acids and in NaOH. 6. Sodium Acetate gives no precipitate unless iron or aluminum salts are present, in which case the chromium is partially pre- cipitated by boiling. REACTIONS FOR THE METALS IN SOLUTION 15 7. Barium Carbonate precipitates gray-green Cr(OH) 3 , liberating CO 2 . All chromium compounds, when treated with suitable oxi- dizing agents, are converted into compounds of chromic acid. A common method of oxidation is to heat the compound on a piece of platinum foil with a mixture of Na 2 CO 3 and KNO 3 , which gives the following result. 2 Cr(OH) 3 + 3 KN0 3 + 2 Na 2 C0 3 = 2 Na 2 O0 4 + 3 KN0 2 + 3 H 2 + 2 C0 2 . The reactions for chromic acid will be given with those of the other acids. IRON (Ferrous), Fe" Iron forms both ferrous (Fe") and ferric (Fe'") compounds. It dissolves easily in HC1 or in H 2 SO 4 , forming FeCl 2 and FeSO 4 respectively. It dissolves in HNO 3 , forming Fe(NO 3 ) 3 , and in aqua regia, forming FeCl 3 . For the reactions use a solution of FeCl 2 or FeSO 4 . 1. Sodium Hydroxid precipitates white Fe(OH) 2 . Insoluble in excess of the reagent. The white color of the precipitate may be seen in a freshly reduced solution, but only for a moment, since it absorbs oxygen from the air, changes first to a dirty green, and then to a red-brown color, forming Fe(OH) 3 . If ammonium salts are present the precipitation is not complete. 2. Ammonium Hydroxid partially precipitates the iron as Fe(OH) 2 . If ammonium salts are present no precipitate appears at first ; but on standing the iron is precipitated as red-brown Fe(OH) 3 . 3. Sodium or Ammonium Carbonate precipitates, under the same conditions as above, white FeCO 3 . This loses CO 2 , oxidizes very easily, and is slowly changed to Fe(OH) 3 . 4. Hydrogen Sulfid gives no precipitate in acidified solutions. In a neutral solution it gives a partial precipitation of the iron 16 QUALITATIVE ANALYSIS as black FeS. If the solution contains sodium acetate the precipitation is nearly complete. 5. Ammonium Sulfid precipitates black FeS. Easily soluble in the mineral acids. Difficultly soluble in acetic acid. The pre- cipitate oxidizes easily when exposed to the air, forming FeSO 4 and a basic ferric sulfate. 6. Potassium Cyanid precipitates light brown Fe(CN) 2 . Soluble in excess of the reagent, forming Fe(CN) 2 (KCN) 4 or K 4 Fe(CN) 6 . 7. Potassium Sulfocyanate gives no coloration unless ferric salts are present. [See Ferric 6.] 8. Potassium Ferrocyanid precipitates bluish-white potassium ferrous ferrocyanid, K 2 Fe" 3 [Fe"(CN) 6 ] 2 . This absorbs oxygen from the air and quickly becomes blue. [See Ferric 7.] 9. Potassium Ferricyanid precipitates " Turnbull's blue," Fe" 3 [Fe'"(CN) 6 ] 2 . Insoluble in HC1. Decomposed by NaOH, forming Fe(OH) 2 , which oxidizes very rapidly, giving Fe(OH) 3 . 10. Barium Carbonate does not precipitate iron from ferrous solutions. [See Ferric 9.] IRON (Ferric), Fe"' When iron is dissolved in HNO 3 , or in aqua regia, or when ferrous salts are acted upon by oxidizing agents, such as HNO 3 , or KC1O 3 and HC1, ferric compounds are formed. 6 FeCl 2 4- KC10 3 + 6 HC1 = KC1 + 3 H 2 + 6 FeCl 3 . For the reactions use a solution of FeCl 3 . 1. Sodium or Ammonium Hydroxid precipitates red-brown Fa(OH) 3 . Insoluble in excess of the reagent. Soluble in any mineral acid. [If any non-volatile organic substance, such as tartaric acid, is present, NH 4 OH gives no precipitate.] 2. Sodium or Ammonium Carbonate precipitates the same, libera- ting CO 2 . The precipitation is only complete after boiling. REACTIONS FOR THE METALS IN SOLUTION 17 3. Hydrogen Sulfid reduces ferric salts to ferrous salts, giving a lightrcolored precipitate of sulfur. 4. Ammonium Sulfid reduces ferric salts to ferrous salts, and precipitates black FeS and sulfur. The FeS is soluble in HC1, the sulfur remaining undissolved. 5. Acid Sodium Phosphate precipitates yellowish- white FePO 4 . Soluble in HC1. Insoluble in acetic acid. 6. Potassium Sulfocyanate produces a blood-red coloration in the solution, owing to the formation of Fe(SCN) 3 . [A very deli- cate reaction.] This action does not take place in the presence of sodium acetate unless HC1 is added in excess. [See Ferrous 7.] 7. Potassium Ferrocyanid precipitates " Prussian" or " Berlin blue," Fe'" 4 [Fe"(CN) 6 ] 3 . [A very characteristic reaction.] Insol- uble in the mineral acids. Decomposed by NaOH, forming red-brown Fe(OH) 3 . 8. Sodium Acetate produces a red coloration caused by the formation of Fe(C 2 H 3 O 2 ) 3 . [If mineral acids are present they must be neutralized. This can be done best with Na 2 CO 3 .] On diluting this solution and boiling, the iron is completely precipitated as red-brown basic ferric acetate, Fe(OH) 2 (C 2 H 3 O 2 ). 9. Barium Carbonate precipitates red-brown Fe(OH) 3 , libera- ting CO 2 . [See Ferrous 10.] 10. Stannous Chlorid, or any other reducing agent, reduces ferric salts to ferrous salts. NICKEL, Ni" Nickel forms nickelous (Ni") and a few nickelic (Ni'") com- pounds. It dissolves slowly in HC1 and in H 2 SO 4 , forming NiCl 2 and NiSO 4 respectively, and readily in HNO 3 , forming Ni(N0 3 ) 2 . For the reactions use a solution of Ni(NO 3 ) 2 . 18 QUALITATIVE ANALYSIS 1. Sodium Hydroxid precipitates apple-green Ni(OH) 2 . Insol- uble in excess of the reagent. Soluble in NH 4 C1. If sodium hypochlorite, NaOCl, or bromin water with excess of NaOH, is added to this precipitate, it is oxidized to black Ni(OH) 3 . 2. Ammonium Hydroxid precipitates the same from a neutral solution. Easily soluble in excess of the reagent to a light blue solution. If ammonium salts are present, or some free acid, by neutralizing which ammonium salts would be formed, no precipitate appears. 3. Sodium or Ammonium Carbonate precipitates an apple-green basic carbonate of variable composition. Soluble in (NH 4 ) 2 CO 3 to a blue solution. 4. Hydrogen Sulfid gives no precipitate if the solution contains a free mineral acid. If the solution is neutral it gives a partial precipitation of black NiS. Sodium acetate added in excess to the nickel solution forms nickel acetate, from which solution H 2 S precipitates all the nickel as black NiS. 5. Ammonium Sulfid precipitates the same. Slightly soluble in excess of the reagent to a dark brown solution, from which the NiS can be reprecipitated by boiling or by the addition of acetic acid. Insoluble in dilute HC1 or acetic acid. Soluble in warm HNO 3 or in aqua regia. 6. Acid Sodium Phosphate precipitates apple-green Ni 3 (PO 4 ) 2 . Easily soluble in dilute acids. 7. Potassium Cyanid precipitates yellow-green Ni(CN) 2 . Sol- uble in excess of the reagent, forming Ni(CN) 2 (KCN) 2 , and reprecipitated from this solution by dilute HC1. If NaOH is added to the latter solution, and then bromin water in excess, black Ni(OH) 3 is precipitated, liberating cyanogen bromid, CNBr. (Poison! Work under a hood.) [See Cobalt 7.] 8. Potassium Ferrocyanid precipitates green Ni 2 Fe(CN) 6 . Insol- uble in dilute acids. REACTIONS FOR THE METALS IN SOLUTION 19 9. Potassium Ferricyanid precipitates yellow-brown Ni 3 [Fe(CN) 6 ] 2 . Insoluble in dilute acids. 10. Potassium Nitrite gives no precipitate in a nickel solution. [See Cobalt 10.] COBALT, Co" Cobalt in its chemical relations very closely resembles nickel. It dissolves in the mineral acids, forming the corresponding salts. These in solution, or when they contain water of crystallization, are red, but on losing water become blue. For the reactions use a solution of Co(NO 3 ) 2 . 1. Sodium Hydroxid precipitates a blue basic salt. If excess of the reagent is added, and the whole boiled, the precipitate is changed to red Co(OH) 2 . On standing, this slowly oxidizes to brown Co(OH) 3 . Sodium hypochlorite and brpmin water give reactions similar to those with nickel. 2. Ammonium Hydroxid precipitates from a neutral solution a blue basic salt. Soluble in excess of the reagent to a red-brown solution. If ammonium salts are present no precipitate appears, but the solution becomes red brown. 3. Sodium or Ammonium Carbonate precipitates a red-lilac basic carbonate of variable composition. Soluble in excess of the (NH 4 ) 2 CO 3 to a red solution, which slowly becomes brown by oxidation. If acid sodium carbonate is used as the reagent it precipitates normal cobalt carbonate, CoCO 3 . 4. Hydrogen Sulfid gives no precipitate if the solution contains a free mineral acid. In a neutral or alkaline solution, or in one containing sodium acetate, it precipitates black CoS. 5. Ammonium Sulfid precipitates black CoS. Insoluble in excess of the reagent, in dilute HC1, and in acetic acid. Soluble in warm HNO 3 and in aqua regia. 6. Acid Sodium Phosphate precipitates blue Co 3 (PO 4 ) 2 . Sol- uble in the mineral acids and in NH 4 OH. 20 QUALITATIVE ANALYSIS 7. Potassium Cyanid precipitates red-brown Co(CN) 2 . Soluble in excess of the reagent, forming Co(CN) 2 (KCN) 4 , and reprecipi- tated from this solution by dilute HC1. If the solution in KCN is boiled for some time potassium cobalticyanid, K 3 Co(CN) 6 , is formed. [Analogous to potassium ferricyanid.] HC1 gives no precipitate in this solution. If NaOH and bromin water are added to the solution in KCN the same compound, K 3 Co(CN) 6 , is formed, and cobalt is not precipitated. [See Nickel 7.] (Since the commercial cobalt salts often contain traces of nickel, a very slight precipi- tate will generally be formed.) 8. Potassium Ferrocyanid precipitates bluish-green Co 2 Fe(CN) 6 . Soluble in concentrated HC1 to a blue-green solution. 9. Potassium Ferricyanid precipitates brown Co 3 [Fe(CN) 6 ] 2 . Insoluble in HC1. 10. Potassium Nitrite, added in excess to a cobalt solution which has been previously acidified with acetic acid, precipi- tates yellow cobaltic-potassium nitrite, Co(NO 2 ) 3 (KNO 2 ) 3 . The reaction is represented by the following equation : Co(N0 3 ) 2 + 7 KN0 2 + 2 H(C 2 H 3 2 ) = Co(N0 2 ) 3 (KN0 2 ) 3 4- 2 KN0 3 + 2 K(C 2 H 3 2 ) + H 2 O + NO. The precipitate forms slowly in dilute solutions, and so should be allowed to stand some time. The precipitate is somewhat soluble in pure water, but insoluble in the presence of KNO 2 . [See Nickel 10.] MANGANESE, Mn" Manganese forms four classes of compounds, two in which it is basic, and two in which it is acid. These are the manganous (Mn") and manganic salts (Mn'"), manganates (Mn vi ) and per- manganates (Mn vii ). It dissolves easily in most acids, forming manganous salts, which are the common ones. For the reactions use a solution of MnSO 4 . REACTIONS FOR THE METALS IN SOLUTION 21 1. Sodium Hydroxid precipitates white Mn(OH) 2 . Insoluble in excess of the reagent. Soluble in NH 4 C1. The precipitate oxidizes slowly in the air, forming brown Mn(OH) 3 . 2. Ammonium Hydroxid precipitates the same in a neutral solu- tion. In a solution containing ammonium salts, or a free acid, no precipitate is formed at first ; but on standing, the solution soon oxidizes, and all the manganese is finally precipitated as brown Mn(OH) 3 . 3. Sodium or Ammonium Carbonate precipitates white MnCO 3 . Boiling makes the precipitation complete. 4. Hydrogen Sulfid gives no precipitate in either neutral or acid solutions. 5. Ammonium Sulfid precipitates flesh-colored MnS. Soluble in dilute mineral acids and in acetic acid. Insoluble in NH 4 C1, in the presence of which the precipitation is complete. 6. Acid Sodium Phosphate precipitates white Mn 3 (PO 4 ) 2 . Sol- uble in the mineral acids and in acetic acid. If the precipitate is dissolved in HC1, an excess of NH 4 OH added, and the whole boiled, a light rose-colored crystalline precipitate of MnNH 4 PO 4 is formed. 7. Potassium Ferrocyanid precipitates white Mn 2 Fe(CN) 6 . Easily soluble in H 2 SO 4 and in HNO 3 , and with difficulty in HC1. 8. Potassium Ferricyanid precipitates brown Mn 3 [Fe(CN) 6 ] 2 . 9. If a manganese compound is heated on a piece of platinum foil with a mixture of Na 2 CO 3 and KNO 3 it will be oxidized, forming green sodium manganate, Na 2 MnO 4 . Thus : MnS0 4 + 2 Na 2 C0 3 + 2 KN0 3 = Na 2 MnO 4 + Na 2 S0 4 + 2 KN0 2 + 2 CO 2 . 22 QUALITATIVE ANALYSIS If the green mass is dissolved in water with the addition of a few drops of acetic acid, the color of the solution will change to red, owing to the formation of sodium permanganate, NaMnO 4 , and dark brown MnO 2 will be precipitated. 10. If a small quantity of red lead, Pb 3 O 4 , is placed in a test-tube with 2 cc. of concentrated HNO 3 , a few drops of the manganese solution added, and the whole carefully warmed, the solution becomes red from the formation of permanganic acid. 2 MnSO 4 + 5 Pb 3 4 + 26 HN0 3 = 2 HMn0 4 + 2 PbS0 4 + 13 Pb(N0 3 ) 2 + 12 H 2 0. 11. Manganic and permanganic acids and their salts are easily reduced to manganous salts in the presence of reducing agents, such as H 2 SO 3 , H 2 S, or nascent hydrogen. ZINC, Zn" Zinc dissolves easily in most acids, forming the correspond- ing salt. It also dissolves in NaOH, forming sodium zincate, Na 2 ZnO 2 . For the reactions use a solution of ZnSO 4 . 1. Sodium Hydroxid precipitates white Zn(OH) 2 . Easily sol- uble in excess of the re~agent, forming sodium zincate, Na 2 ZnO 2 . 2. Ammonium Hydroxid precipitates the same. Easily soluble in excess of the reagent, forming ZnSO 4 (NH 3 ) 4 . Ammonium salts prevent the precipitation. 3. Sodium Carbonate precipitates a white basic carbonate of variable composition, but usually Zn 2 (OH) 2 CO 3 . Boiling makes the precipitation complete. If acid sodium carbonate is used as the reagent it precipitates normal zinc carbonate, ZnCO 3 . 4. Ammonium Carbonate precipitates the same. Soluble in excess of the reagent. Ammonium salts prevent the pre- cipitation. REACTIONS FOR THE METALS IN SOLUTION 23 5. Hydrogen Sulfid gives no precipitate in solutions containing a free mineral acid. In a neutral or alkaline solution, or one acidified with acetic acid, it precipitates white ZnS. 6. Ammonium Sulfid precipitates the same from any solution. If NH 4 C1 is present the precipitation is complete. 7. Acid Sodium Phosphate precipitates white Zn 3 (PO 4 ) 2 . Sol- uble in dilute acids and in NH 4 OH. 8. Potassium Cyanid precipitates white Zn(CN) 2 . Soluble in excess of the reagent, forming Zn(CN) 2 (KCN) 2 . From this solution (NH 4 ) 2 S precipitates white ZnS. 9. Potassium Ferrocyanid precipitates white Zn 2 Fe(CN) 6 . In- soluble in dilute acids and in NH 4 OH. 10. Potassium Ferricyanid precipitates brownish-yellow Zn 8 [Fe(CN) 6 ] 2 . Soluble in HC1 and in NH 4 OH. MAGNESIUM, Mg" Magnesium dissolves easily in all acids, forming the corre- sponding salts. If heated in the air it takes fire quite easily, and burns with an intensely white light, forming MgO. For the reactions use a solution of MgSO 4 . 1. Sodium Hydroxid precipitates white Mg(OH) 2 . Soluble in NH 4 C1. The presence of ammonium salts prevents the precipitation. 2. Ammonium Hydroxid gives a partial precipitation of the same in a neutral solution. If an ammonium salt or a free acid is present no precipitate appears. 3. Sodium Carbonate precipitates a white basic carbonate of variable composition. If the precipitate is boiled it has the composition Mg 3 (OH) 2 (CO 3 ) 2 . If ammonium salts are present no precipitate is formed. 24 QUALITATIVE ANALYSIS 4. Ammonium Carbonate gives no precipitate if ammonium Halts are present. 5. Acid Sodium Phosphate precipitates white MgHPO 4 , which by boiling changes to Mg 3 (PO 4 ) 2 . If NH 4 C1 is added to the solution, and then NH 4 OH in excess, the reagent precipitates white crystalline MgNH 4 PO 4 . This precipitate forms slowly in a dilute solution and is complete only after standing some hours. 6. Ammonium Oxalate gives no precipitate in dilute solutions. In concentrated solutions it gives a white precipitate of MgC 2 O 4 . Soluble in NH 4 C1. [See Calcium 6.] BARIUM, Ba" The metal barium has little practical value and is very difficult to obtain in the metallic state. It oxidizes easily in the air and decomposes water at the ordinary temperature. Its salts are easy to form and many of them are soluble in water. For the reactions use a solution of BaCl 2 . 1. Sodium Hydroxid precipitates, if the solution is not too dilute, white Ba(OH) 2 . Somewhat soluble in cold water, much more so in hot. 2. Ammonium Hydroxid gives no precipitate in barium solutions. 3. Sodium or Ammonium Carbonate precipitates white BaCO 3 . Somewhat soluble in NH 4 C1. Soluble in water containing CO 2 , forming the acid carbonate, BaH 2 (CO 3 ) 2 . The precipitation can be made complete by adding NH 4 OH in slight excess and boiling. 4. Acid Sodium Phosphate precipitates a white acid phosphate, BaHPO 4 . If NH 4 OH is present it forms BaNH 4 PO 4 . Sol- uble in dilute acids and reprecipitated by NH 4 OH. REACTIONS FOR THE METALS IN SOLUTION 25 5. Potassium Chromate precipitates yellow BaCrO 4 . Soluble in HC1 and HNO 3 . Insoluble in NaOH [See Lead 8], and in acetic acid [See Strontium 5], 6. Sulfuric Acid, or any soluble sulfate, precipitates white BaSO 4 . This precipitation takes place even in extremely dilute solutions. Insoluble in all acids and alkalies. [See Strontium 6.] 7. Ammonium Oxalate precipitates, if the solution is not too dilute, white BaC 2 O 4 . Dilute solutions give no precipitate. [See Calcium 6.] Soluble in HC1 and HNO 3 . 8. Hydrofluosilicic Acid precipitates white BaSiF 6 . Somewhat soluble in water. Insoluble in alcohol and in dilute acids. [See Strontium 8.] 9. If a barium compound is heated on a platinum wire in an oxidizing flame it imparts a pale green color to the flame. If the compound is a chlorid, or is moistened with HC1, the color is more distinct. STRONTIUM, Sr" Strontium very closely resembles barium in its chemical properties. For the reactions use a solution of SrCl 2 . 1. Sodium Hydroxid precipitates white Sr(OH) 2 . Somewhat soluble in water, but less so than Ba(OH) 2 . 2. Ammonium Hydroxid gives no precipitate. 3. Sodium or Ammonium Carbonate precipitates white SrCO 3 . Its properties are like those of BaCO 3 . [See Barium 3.] 4. Acid Sodium Phosphate precipitates white SrHPO 4 . Like BaHPO 4 . [See Barium 4.] 5. Potassium Chromate gives no precipitate at first, but after a time, if the solution is neutral and not too dilute, yellow SrCrO 4 26 QUALITATIVE ANALYSIS is precipitated. Insoluble in alcohol even when dilute, so that if alcohol is added to the solution the precipitate appears at once. Soluble in acetic acid. [See Barium 5.] 6. Sulfuric Acid, or any soluble sulfate, precipitates white SrSO 4 . Slightly soluble in water, so that if the solution is very dilute the precipitate does not appear immediately. [See Barium 6 and Calcium 5.] A concentrated solution of Na 2 CO 3 or (NH 4 ) 2 CO 3 converts it into SrCO 3 . 7. Ammonium Oxalate precipitates white SrC 2 O 4 . Somewhat soluble in water, but less so than BaC 2 O 4 . 8. Hydrofluosilicic Acid gives no precipitate in moderately dilute solutions. [See Barium 8.] 9. If a strontium compound is heated on a platinum wire in an oxidizing flame it imparts a crimson color to the flame. If the compound is a chlorid, or is moistened with HC1, the color is more distinct. CALCIUM, Ca" Calcium very closely resembles strontium and barium in its chemical properties. For the reactions use a solution of CaCl 2 . 1. Sodium Hydroxid precipitates white Ca(OH) 2 . Slightly sol- uble in water, but much less so than Sr(OH) 2 or Ba(OH) 2 . 2. Ammonium Hydroxid gives no precipitate. 3. Sodium or Ammonium Carbonate precipitates white CaCO 3 . Its properties are like those of BaCO 3 . [See Barium 3.] 4. Acid Sodium Phosphate precipitates white CaHPO 4 . If NH 4 OH is present the normal salt, Ca 3 (PO 4 ) 2 , is precipitated. Soluble in dilute acids and reprecipitated by NH 4 OH. 5. Sulfuric Acid, or any soluble sulfate, precipitates white CaSO 4 . Somewhat soluble in water, so that if the solution REACTIONS FOR THE METALS IN SOLUTION 27 is very dilute no precipitate appears. A concentrated solution of Na 2 CO 3 or (NH 4 ) 2 CO 3 converts it into CaCO 3 . 6. Ammonium Oxalate precipitates, even from very dilute solu- tions, white CaC 2 O 4 . Soluble in HC1 or HNO 3 . Insoluble in water and acetic acid. [See Barium 7.] The presence of NH 4 OH hastens the precipitation, which, if the solution is cold and dilute, is complete only after long standing. 7. Hydrofluosilicic Acid gives no precipitate even if alcohol is added. [See Barium 8.] 8. If a calcium compound is heated on a platinum wire in an oxidizing flame it imparts a yellowish-red color to the flame. If the compound is a chlorid, or is moistened with HC1, the color is more distinct. THE ALKALI METALS Potassium and sodium are the common elements belonging to the alkali group of metals. They are very strong bases, and form salts with every acid known. The salts are all soluble in water to some extent, and so form no precipitates with the com- mon reagents, most of which are compounds of these metals. There are a few compounds which are difficultly soluble in water, and these are precipitated when they are produced in sufficiently concentrated solutions. Most of the salts of these metals are either insoluble, or difficultly soluble, in alcohol, so that the addition of this reagent often helps the formation of a pre- cipitate. The compound radical ammonium, NH 4 , forms a series of compounds analogous to those of potassium and sodium. POTASSIUM (Kalium), K' For the reactions use a solution of KC1. 1. Acid Sodium Tartrate, HNa(C 4 H 4 O 6 ), precipitates, if the solution is not too dilute, white crystalline HK(C 4 H 4 O 6 ). The 28 QUALITATIVE ANALYSIS precipitate forms slowly, but may be hastened by shaking. [See Ammonium 2.] 2. Hydrofluosilicic Acid precipitates, if the solution is not too dilute, white K 2 SiF 6 . Insoluble in dilute acids and in alcohol. 3. If to 2 cc. of a solution of sodium nitrite there are added 1 cc. of acetic acid and 5 drops of a solution of cobalt nitrate a deep orange-yellow liquid is formed. This precipitates from the potassium solution yellow cobaltic-potassium nitrite, Co(NO 2 ) 8 (KNO a ) 8 . [See Cobalt 10.] 4. Platinum Chlorid * precipitates from neutral or slightly acid solutions yellow K 2 PtCl 6 . Soluble in 100 parts of water. Insoluble in alcohol. All potassium salts, when heated on a platinum wire in an oxidizing flame, impart a reddish-violet color to the flame. This color appears red when seen through a blue glass. SODIUM (Natrium), Na' For the reactions use a solution of NaCl. 1. Hydrofluosilicic Acid precipitates, after long standing if the solution is dilute, or upon addition of alcohol, white Na 2 SiF 6 . 2. Acid Potassium Pyroantimonate, K 2 H 2 Sb 2 O 7 , precipitates, in neutral solutions which do not contain other metals, white Na 2 H 2 Sb 2 O 7 . The precipitation is slow, but may be hastened by shaking. 3. All sodium salts, when heated on a platinum wire in an oxidizing flame, impart a bright yellow color to the flame. This color is not seen through blue glass. A crystal of potassium bichromate appears colorless in this yellow light. * The reactions with platinum chlorid may be omitted. REACTIONS FOR THE METALS IN SOLUTION 29 AMMONIUM, (NH 4 )' For the reactions use a solution of NH 4 C1. 1. Sodium Hydroxid, or any soluble base, when heated with an ammonium compound, decomposes it, liberating NH 3 . This may be recognized by its characteristic odor, or by the white clouds of NH 4 C1 which are formed if a rod moistened with HC1 is held in the escaping gas. 2. Acid Sodium Tartrate precipitates, if the solution is not too dilute, HNH 4 (C 4 H 4 O 6 ). The precipitation may be hastened by shaking. [See Potassium 1.] 3. Platinum Chlorid precipitates, from neutral or slightly acid solutions, yellow (NH 4 ) 2 PtCl 6 . Soluble in 170 parts of water. Insoluble in alcohol. 4. Nessler's Reagent, a solution of HgI 2 (KT) 2 with an excess of KOH, precipitates, even from extremely dilute solutions, brown Hg 2 I(NH 2 )O. This reaction is best shown by filling a test-tube nearly full of water, adding two or three drops of the ammonium solution, and then the reagent. REACTIONS FOR THE ACID RADICALS IN SOLUTION For the reactions for the acid radicals it is better to use solu- tions of the salts derived from the acids rather than the acids themselves, although the latter may sometimes be used. The salts are usually neutral in their action on litmus paper, while the acids and most acid salts turn the blue litmus paper red. The free acids may be further distinguished by leaving no residue when a few drops are evaporated to dryness on a piece of platinum foil. The reactions are similar to those for the metals except that the solutions which were then used for the reactions now become the reagents, and the reagents then used are now the solutions for the reactions. In testing for the acid radicals, therefore, only a few of the more characteristic reactions will be given. For other reactions the student is referred to those given under the different metals. The student must always consider the nature not only of the reagent used but also of the substance in the solution. If either is reducing in its action this will manifest itself in the precipitation. [See Silver 13, Mercurous 9, Bismuth 8, etc.] The relation which the metal in each bears to the acid radical in the other must also be considered. If the metal in the unknown solution forms a precipitate with the acid radical of the reagent no information regarding the acid radical of the unknown substance can be obtained by this particular reaction. This does not often occur, but when it does the other reactions must be relied upon for proving the constitution of the acid radical. 30 REACTIONS FOR THE ACID RADICALS IN SOLUTION 31 HYDROCHLORIC ACID, HC1 For the reactions use a solution of NaCl. 1. Lead Acetate precipitates white PbCl 2 . Soluble in boiling water or in a large quantity of cold water, so that if the solu- tion is very dilute the precipitate may fail to appear. 2. Silver Nitrate precipitates white AgCl. Easily soluble in NH 4 OH and in KCN, and reprecipitated from these solutions by HN0 3 . 3. Mercurous Nitrate precipitates white HgCl. By the addition of NH 4 OH this precipitate becomes black, owing to the forma- tion of amido-mercuric chlorid, HgNH 2 Cl, mixed with finely divided mercury. The changing of this precipitate from white to black by the addition of NH 4 OH is characteristic of mercurous compounds rather than of chlorids. HYDROBROMIC ACID, HBr For the reactions use a solution of KBr. 1. Lead Acetate precipitates white PbBr 2 . Somewhat soluble in water, but not as easily soluble as PbCl 2 . 2. Silver Nitrate precipitates yellowish-white AgBr. Soluble with some difficulty in NH 4 OH, but easily soluble in KCN. Insoluble in dilute acids. 3. Mercurous Nitrate precipitates yellowish-white HgBr. The precipitate becomes black on adding NH 4 OH. 4. Chlorin Water liberates bromin from many of its compounds, coloring the solution red brown. If a little carbon disulfid, CS 2 , is added to the solution, and the whole well shaken, the bromin dissolves in the CS 2 and colors it red brown. 32 QUALITATIVE ANALYSIS HYDRIODIC ACID, HI For the reactions use a solution of KI. 1. Lead Acetate precipitates yellow PbI 2 . Soluble in boiling water, from which solution it crystallizes, on cooling, in golden yellow scales. 2. Silver Nitrate precipitates light yellow Agl. Only very slightly soluble in NH 4 OH, but easily soluble in KCN. Insoluble in HNO 3 . 3. Mercurous Nitrate precipitates yellowish-green Hgl. [See Mercurous 5.] 4. Mercuric Chlorid precipitates scarlet-red HgI 2 . Soluble in KI. [See Mercuric 5.] 5. Bismuth Nitrate precipitates brown BiI 3 . 6. Copper Sulfate precipitates white Cu 2 I 2 together with free iodin, which colors the precipitate brown. If H 2 SO 3 is added the precipitate appears white. 7. Chlorin or Bromin Water liberates iodin from most of its compounds. If a little CS 2 is added to the solution and the whole well shaken, the iodin dissolves in the CS 2 and colors it violet. If starch paste is added to the solution, the iodin colors it deep blue. HYDROFLUORIC ACID, HF For the reactions use a solution of KF. 1. Lead Acetate precipitates white PbF 2 . Soluble in HNO 3 . 2. Silver Nitrate gives no precipitate. (Distinction between fluorids and the other halogen salts.) 3. Barium Chlorid precipitates white BaF 2 . Soluble in HC1 or in HNO Q . REACTIONS FOR THE ACID RADICALS IN SOLUTION 33 4. Calcium Chlorid precipitates white CaF 2 . Scarcely soluble in any dilute acid. 5. All fluorids are decomposed by concentrated H 2 SO 4 , lib- erating HF. This acid unites with the silicon in glass, forming SiF 4 . Hence if a fluorid, together with some concentrated H 2 SO 4 , is heated for a moment in a clean test-tube, and the tube then emptied and cleaned, it will be found to have been etched. To show this reaction the solution must be fairly concentrated. HYDROCYANIC ACID, HCN For the reactions use a solution of KCN. 1. Lead Acetate precipitates white Pb(CN) 2 . Insoluble in KCN. 2. Silver Nitrate precipitates white AgCN. Soluble in KCN and in NH 4 OH, and reprecipitated from these solutions by HNO 3 . 3. Copper Sulfate precipitates greenish-yellow Cu(CN) 2 . Sol- uble in KCN, from which solution H 2 S will not precipitate the copper. 4. Cadmium Nitrate precipitates white Cd(CN) 2 . Soluble in KCN, from which solution H 2 S precipitates yellow CdS. 5. If a few drops of (NH 4 ) 2 S X are added to a solution of KCN, and the solution boiled for a moment, potassium sulfocy- anate, KSCN, is formed. If HC1 is now added in excess, ferric chlorid will produce a blood-red coloration, owing to the formation of Fe(SCN) 3 . 6. If a small quantity of NaOH is added to a solution of KCN, then three or four drops each of FeSO 4 and FeCl 3 , and finally HC1 in excess, Prussian blue is formed. [See Ferric 7.] 34 QUALITATIVE ANALYSIS SULFOCYANIC OR THIOCYANIC ACID, HSCN For the reactions use a solution of KSCN. 1. Silver Nitrate precipitates white AgSCN. Soluble in NH 4 OH. 2. Mercurous Nitrate produces a gray precipitate, which is a mixture of Hg and Hg(SCN) 2 , with perhaps some HgSCN. 3. Mercuric Nitrate precipitates white Hg(SCN) 2 . 4. Copper Sulfate precipitates, from a concentrated solution, black Cu(SCN) 2 . If the solution is dilute, an emerald-green coloration is produced. If H 2 SO 3 in excess is added to this solution, and the whole boiled, the copper is reduced and white Cu 2 (SCN) 2 is precipitated. [See Copper 8.] 5. Ferric Chlorid produces a blood-red coloration in the solu- tion, owing to the formation of Fe(SCN) 3 . This is a very characteristic reaction. HYDROFERROCYANIC ACID, H 4 Fe(CN) 6 For the reactions use a solution of K 4 Fe(CN) 6 . 1. Lead Acetate precipitates white Pb 2 Fe(CN) 6 . 2. Silver Nitrate precipitates white Ag 4 Fe(CN) 6 . Insoluble in dilute NH 4 OH. 3. Copper Sulfate precipitates red-brown Cu 2 Fe(CN) 6 . This reaction can be shown in a very dilute solution. 4. Ferric Chlorid precipitates Prussian blue, Fe 4 [Fe(CN) 6 ] 3 . Decomposed by NaOH, forming red-brown Fe(OH) 3 . HYDROFERRICYANIC ACID, H 3 Fe(CN) 6 For the reactions use a solution of K 3 Fe(CN) 6 . 1. Silver Nitrate precipitates red-brown Ag 3 Fe(CN) 6 . Soluble in NH 4 OH. REACTIONS FOR THE ACID RADICALS IN SOLUTION 35 2. Ferrous Sulfate precipitates Turnbull's blue, Fe 8 [Fe(CN) 6 ] 2 . Insoluble in HCL 3. Ferric Chlorid gives a red-brown solution but no precipitate. 4. Zinc Sulfate precipitates brownish-yellow Zn 3 [Fe(CN) 6 ] 2 . Soluble in HC1 and in NH 4 OH. HYPOCHLOROUS ACID, HC10 All hypochlorites are soluble in water, and so the acid radical cannot be precipitated. If a concentrated solution of a hypo- chlorite is boiled, oxygen is liberated. If a dilute acid is added to the solution, chlorin is liberated. For the reactions use a solution of NaClO. 1. Lead Acetate, to which NaOH has been added until the precipitate first formed is dissolved, precipitates brown PbO 2 . Pb(C 2 H 3 2 ) 2 + 2 NaOH + NaClO = Pb0 2 + 2 NaC 2 H 3 2 + NaCl + H 2 0. The precipitation is hastened by boiling. 2. Silver Nitrate gives a white precipitate of AgCl, silver chlorate being formed at the same time. 3 AgN0 3 + 3 NaClO = 2 AgCl + AgC10 3 + 3 NaN0 3 . 3. If a piece of litmus paper is moistened with a few drops of the solution and then exposed to acid fumes, the color will be bleached. If the moistened paper is breathed upon, the CO 2 in the breath will effect the same change. CHLORIC ACID, HC10 3 The chlorates are all soluble in water and so form no pre- cipitates. For the reactions use a solution of KC1O 3 . 1. Silver Nitrate gives no precipitate with a chlorate, but if H 2 SO 3 is added to the solution, the chlorate is reduced to a chlorid, and AgNO 3 then gives a white precipitate of AgCl. 36 QUALITATIVE ANALYSIS 2. Hydrochloric Acid decomposes the chlorates, giving chlorin peroxid and chlorin. Thus : 2 HC1 + KC1O 3 = KC1 + H 2 + C10 2 + Cl. The chlorin peroxid and chlorin dissolve in the solution, coloring it yellow; but if the solution is boiled these gases will pass off, giving what is called a " chlorous odor." This mixture possesses great oxidizing power. HYDROGEN SULFID (Hydrosulfuric Acid), H 2 S For the reactions use a solution of H 2 S or (NH 4 ) 2 S. 1. Lead Acetate precipitates black PbS. Soluble in warm dilute HN0 3 . 2. Silver Nitrate precipitates black Ag 2 S. Soluble in warm dilute HNO 3 . 3. Antimony Chlorid precipitates orange-red Sb 2 S 3 . Soluble in (NH 4 ) 2 S and reprecipitated by HC1. 4. All soluble sulfids, and most insoluble ones, are decom- posed by warm H 2 SO 4 , liberating H 2 S, which may be detected by its odor ; also by the brown or black stain on a piece of paper moistened with lead acetate and held in the escaping gas. THIOSULFURIC ACID, H 2 S 2 3 The salts of this acid, which are called thiosulfates, were formerly called hyposulfites. The free acid does not exist. For the reactions use a solution of Na 2 S 2 O 3 . 1. Lead Acetate precipitates white PbS 2 O 3 . Soluble in an excess of Na^SgOg. Decomposed by boiling, forming. PbS. [See Sulfurous Acid 1.] 2. Silver Nitrate precipitates white Ag 2 S 2 O 3 . Easily soluble in an excess of Na 2 S 2 O 3 , forming the double salt, AgNaS 2 O 3 . REACTIONS FOR THE ACID RADICALS IN SOLUTION 37 The precipitate quickly beeomes black, especially if warm, being reduced to Ag 2 S. Ag a S A + H 2 = Ag 2 S + H 2 S0 4 . 3. Barium Chlorid precipitates, from a concentrated solution, white BaS 2 O 3 . Soluble in a large quantity of water. Decom- posed by HC1, liberating sulfur dioxid and sulfur. 4. Ferric Chlorid produces a violet color in the solution. The color is not permanent and the solution soon becomes cloudy, owing to the reduction to FeCl 2 and the liberation of sulfur. 5. Hydrochloric Acid decomposes the thiosulfates, liberating SO 2 and giving a precipitate of free sulfur. [See Sulfurous Acid 4.] SULFUROUS ACID, H 2 S0 3 This is a weak acid and exists only in a dilute solution. It easily decomposes when heated, forming SO 2 and H 2 O. Its salts are much more 'stable, but are all decomposed by dilute acids. A solution of a sulfite, on standing, becomes partially oxidized, forming a sulfate. For the reactions use a solution of Na 2 SO 3 . 1. Lead Acetate precipitates white PbSO 3 , which is not decom- posed by boiling. [See Thiosulf uric Acid 1.] Soluble in dilute HN0 3 . 2. Silver Nitrate precipitates white Ag 2 SO 3 , which is decom- posed by boiling, forming black metallic silver. 3. Barium Chlorid precipitates white BaSO 3 . Easily soluble in dilute HC1. The solution in HC1 is often incomplete, owing to the presence of sulfates, which precipitate insoluble BaSO 4 . 4. Hydrochloric Acid decomposes the sulfites, liberating SO 2 . If a little potassium permanganate, KMnO 4 , is now added, it is at once decolorized, owing to reduction. Thus: 2 KMn0 4 + 5 S0 2 + 2 H 2 O = 2 MnS0 4 + K 2 S0 4 + 2 H 2 S0 4 . Most other mineral acids give a similar reaction. 38 QUALITATIVE ANALYSIS 5. If a solution of H 2 SO 3 is boiled with a little stannous chlorid and HC1 it is first reduced to H 2 S, the SnCl 2 being oxidized to SnCl 4 . The H 2 S then precipitates yellow SnS 2 . SULFURIC ACID, H 2 S0 4 For the reactions use a solution of Na 2 SO 4 . 1. Lead Acetate precipitates white PbSO 4 . Easily soluble in ammonium tartrate or ammonium acetate. [See Lead 11.] 2. Barium Chlorid precipitates white BaSO 4 . Insoluble in all dilute acids. 3. Calcium Chlorid precipitates, in not too dilute solutions, white CaSO 4 . A concentrated solution of Na 2 CO 3 or (NH 4 ) 2 CO 8 converts the precipitate into CaCO 3 , which dissolves in dilute HC1, liberating CO 2 . CHROMIC ACID, H 2 Cr0 4 Chromium ordinarily acts like the metals, forming compounds with the acid radicals. It may be oxidized by fusion with Na 2 CO 3 and KNO 3 , forming a compound with the metal, in which chromium is found in the acid radical. [See page 15.] For the reactions use a solution of K 2 CrO 4 . 1. Lead Acetate precipitates yellow PbCrO 4 . Soluble in NaOH. Insoluble in acetic acid. 2. Silver Nitrate precipitates red-brown Ag 2 CrO 4 . Soluble in HN0 3 and in NH 4 OH. 3. Mercurous Nitrate precipitates brick-red Hg 2 CiO 4 . Soluble in HNO 3 . 4. Barium Chlorid precipitates yellow BaCrO 4 . Insoluble in NaOH and in acetic acid. REACTIONS FOR THE ACID RADICALS IN SOLUTION 39 5. Take a dilute solution of hydrogen dioxid, H 2 O 2 , acidify with HC1, add a little ether (about half an inch deep in the test- tube), then two or three drops of the chromate solution, and shake. A portion of the chromate will be oxidized by the H 2 O 2 , forming an unstable blue compound, which is supposed to be perchromic acid, HCrO 4 . This dissolves in the ether, which rises to the surface, giving it a rich blue color. 6. Nitric Acid converts the yellow K 2 CrO 4 into red potassium dichromate, K 2 Cr 2 O 7 . This salt, which may be regarded as an acid chromate, gives in most cases the same reactions as the normal chromate. 7. The chromates, and especially the dichromates, when treated with H 2 SO 4 , form sulfates and liberate oxygen. They are therefore powerful oxidizing agents, and are used as such, especially in organic chemistry. k (\ v o - 1 - '))! SOi, -* K^Ou ^^i r^V). *vAfcG*-3 ^J ' 2 ^ -j ' ' '> ^ * X> NITROUS ACID, HN0 2 This acid does not exist in the free state. Even when liber- ated in a dilute solution it is easily decomposed, giving nitric acid, nitric oxid, and water. Its salts, the nitrites, are quite stable, but they are all decomposed by dilute acids, forming HNO 3 and liberating NO. For the reactions use a solution of KNO 2 . 1. Silver Nitrate precipitates, in a concentrated solution, white AgN0 2 . 2. Cobalt Nitrate, to which has been added acetic acid, precipi- tates in an excess of the solution, yellow Co(NO 2 ) 8 (KNO 2 ) 3 . [See Cobalt 10.] Sodium nitrite does not give this reaction. 3. If KI is added to a solution of a nitrite, together with a little starch paste and a few drops of dilute H 2 SO 4 , iodin is liberated, which colors the starch paste blue. 40 QUALITATIVE ANALYSIS 4. If a little FeSO 4 is added to a solution of a nitrite, and then a few drops of dilute acetic acid, the whole becomes brown, from the NO which is liberated dissolving in the FeSO 4 . [See Nitric Acid 1.] 5. Nitrous acid is capable of oxidation to nitric acid in the presence of oxidizing agents. If KMnO 4 is added to a solution of a nitrite acidified with H 2 SO 4 , it is decolorized, owing to its reduction. Thus : 4 KMn0 4 + 10 KNO 2 + 11 H 2 S0 4 = 7 K 2 S0 4 + 4 MnSO, + 10 HN0 3 + 6 H 2 0. NITRIC ACID, HN0 3 All nitrates are soluble in water, and so form no precipitates with the metals. They are all decomposed by H 2 SO 4 , liberating HNO 3 . Nitric acid is an oxidizing agent, and the tests which indicate its presence are connected with an oxidizing action. For the reactions use a solution of KNO 3 . 1. Mix some of the nitrate solution with an equal volume of FeSO 4 . Incline the tube a little and carefully pour down the side some concentrated H 2 SO 4 , and where the mixture meets the surface of the acid a brown ring of color will appear. The brown compound, which is due to a solution of NO in FeSO 4 , is decomposed by heat, liberating the NO. [See Nitrous Acid 4.] The action is threefold: (1) the liberation of HNO 3 by the action of H 2 SO 4 on the nitrate ; (2) the oxidation of FeSO 4 by the HNO 3 , liberating NO ; and (3) the absorption of NO by FeSO 4 , forming the unstable brown compound. 2. If concentrated H 2 SO 4 is mixed with the nitrate solution, a few fragments of copper added, and the whole boiled, NO is liberated, which, combining with the oxygen in the air, forms red-brown fumes of NO 2 . These will be more easily seen by looking down through the mouth of the tube. The action is analogous to that in the first test. REACTIONS FOR THE ACID RADICALS IN SOLUTION 41 PHOSPHORIC ACID, H 3 P0 4 For the reactions use a solution of Na 2 HPO 4 . 1. Lead Acetate precipitates white Pb 3 (PO 4 ) 2 . Easily soluble in HNO 3 . 2. Silver Nitrate precipitates yellow Ag 3 PO 4 . Soluble in N1I 4 OH and in HNO 3 . 3. Barium Chlorid precipitates white BaHPO 4 . Soluble in dilute HC1 and reprecipitated by NH 4 OH. 4. Magnesium Sulfate, to which has been added NH 4 C1 and then NH 4 OH in excess, precipitates white crystalline MgNH 4 PO 4 . This precipitate forms slowly, and in a dilute solution is com- plete only after standing some hours. 5. Ammonium Molybdate, (NH 4 ) 2 MoO 4 , with an 'excess of HNO 3 precipitates yellow ammonium phospho-molybdate, (NH 4 ) 3 PO 4 (MoO 3 ) 12 . Soluble in NH 4 OH and in excess of Na 2 HPO 4 . The precipitation is hastened by warming. 2 4.0 = ARSENIOUS ACID, H 3 As0 3 We have already learned that arsenic may be precipitated as a sulfid. [See Arsenic 1.] It forms no other salts in which it acts as a metal, but acts like an acid, forming salts with the metals. The alkaline salts only are soluble in water. These may be formed by dissolving the oxid, As 2 O 3 , in a solution of an alkaline hydroxid. For the reactions use a solution of K 3 AsO 3 or Na 3 AsO 3 . 1. Silver Nitrate precipitates, in a neutral solution, yellow Ag 3 AsO 3 . Soluble in NH 4 OH, in HNO 3 , and in NH 4 NO 3 . If the ammoniacal solution is boiled for some time, metallic silver is precipitated, a portion of the arsenite being oxidized to an arsenate. 42 QUALITATIVE ANALYSIS 2. Copper Sulfate precipitates Scheele's green, CuHAsO 3 . Soluble in NH 4 OH and in acids. If an excess of NaOH is added to a solution of an arsenite, then a few drops of CuSO 4 , and the whole boiled, red Cu 2 O is precipitated, a portion of the arsenite being oxidized to an arsenate. (Distinction from arsenates.) [See Arsenic Acid 2.] 3. Reinsch's Test. If a piece of metallic copper is placed in an arsenite solution acidified with HC1 and warmed, a gray film of copper arsenid, Cu 5 As 2 , is formed on the surface of the copper. 4. Stannous Chlorid, to which has been added at least two volumes of concentrated HC1, and then warmed, precipitates black metallic arsenic. 5. Hydrogen or Ammonium Sulfid gives no precipitate in a neutral or alkaline solution of an arsenite. If the solution is acid, or is rendered acid, yellow As 2 S 3 is precipitated. [See Arsenic 1.] ARSENIC ACID, H 3 As0 4 Salts of this acid are derived from As 2 O 5 or by oxidation of the arsenites. The reactions for arsenic acid are very closely analogous to those for phosphoric acid. For the reactions use a solution of Na 3 AsO 4 . 1. Silver Nitrate precipitates, in a neutral solution, red-brown Ag 3 AsO 4 . Soluble in NH 4 OH and in HNO 3 . 2. Copper Sulfate precipitates light blue CuHAsO 4 . Soluble in NH 4 OH. Not decomposed by boiling in an alkaline solution. [See Arsenious Acid 2.] 3. Magnesium Sulfate, to which has been added NH 4 C1 and then NH 4 OH in excess, precipitates white crystalline MgNH 4 AsO 4 . The precipitate forms slowly, especially in a dilute solution, but may be hastened by shaking. Soluble in HC1 and reprecipitated REACTIONS FOR THE ACID RADICALS IN SOLUTION 43 by NH 4 OH. If (NH 4 ) 2 S or H 2 S is added to the solution in HC1, the arsenate is reduced and yellow As 2 S 3 mixed with free sulfur is precipitated. The action is hastened by warming. (Distinction from phosphates.) 4. Ammonium Molybdate gives no precipitate in the cold. If the solution is warmed, yellow ammonium arseno-molybdate, (NH 4 ) 3 AsO 4 (MoO 3 ) 12 , is formed. Insoluble in NH 4 OH. 5. Hydrogen or Ammonium Sulfid in a solution acidified with HC1 reduces an arsenate to an arsenite, and then precipitates yellow As 2 S 3 mixed with free sulfur. [See Arsenic 1.] BORIC ACID, H 3 B0 3 This is a weak acid. Only a few salts are known, and most of these are from the derivatives, metaboric acid, HBO 2 , and pyroboric acid, H 2 B 4 O 7 . For the reactions use a solution of borax, Na 2 B 4 O 7 . 1. Lead Acetate precipitates, in not too dilute solutions, white. Pb(BO 2 ) 2 . Soluble in excess of the reagent. 2. Silver Nitrate precipitates white AgBO 2 . If the solution is too dilute, the precipitate is yellow or brown, from the presence of Ag 2 0. 3. Barium Chlorid precipitates, in not too dilute solutions, white Ba(BO 2 ) 2 . Soluble in excess of the reagent and in NH 4 C1. 4. If a little alcohol is added to a solution of a borate in a small porcelain dish, then a little concentrated H 2 SO 4 , and the whole warmed, ethyl borate (boric ether) is formed. This is inflammable, and so the mixture may be ignited, when it will be seen to burn with a green-bordered flame. CARBONIC ACID, H 2 C0 3 This is a weak acid, existing only in a dilute solution. Its salts are common and quite stable. For the reactions use a solution of Na, 2 CO 3 . 44 QUALITATIVE ANALYSIS 1. Lead Acetate precipitates white PbCO 3 or, if the solution is hot, a white basic carbonate. 2. Silver Nitrate precipitates light yellow Ag 2 CO 3 . Soluble in NH 4 OH and in (NH 4 ) 2 CO 3 . 3. Barium Chlorid precipitates white BaCO 3 . Soluble in water containing CO 2 , forming BaH 2 (CO 3 ) 2 , and reprecipitated by boiling. To show this solubility fill a large test-tube containing the precipitate with water, shake the mixture, pour out about one half, and lead CO 2 through the liquid until the precipitate dissolves. 4. All carbonates, whether soluble or insoluble in water, are decomposed with effervescence by dilute HC1 and other acids, liberating CO 2 . The presence of the CO 2 may be confirmed by decanting the heavy gas into a test-tube containing a little lime-water, Ca(OH) 2 . On shaking with the gas, the lime-water becomes milky, owing to the precipitation of CaCO 3 . Baryta- water, Ba(OH) 2 , may be used in place of the lime-water. SILICIC ACID, H 4 Si0 4 This acid exists only in a dilute solution. On attempting to concentrate the solution, it loses a molecule of water and forms metasilicic acid, H 2 SiO 3 . This in turn decomposes on heating to 130, forming H 2 O and SiO 2 . The silicates occurring in nature are generally salts of the polysilicic acids. Only the alkaline silicates are soluble in water. For the reactions use a solution of Na 4 SiO 4 . 1. Hydrochloric Acid (concentrated) precipitates, in not too dilute solutions, white gelatinous H 4 SiO 4 . In dilute solutions this precipitate will appear only after long standing. It is somewhat soluble in water and in HC1, so that a portion of it remains in the solution. REACTIONS FOR THE ACID RADICALS IN SOLUTION 45 2. Ammonium Carbonate or Ammonium Chlorid precipitates the same. 3. Barium Chlorid precipitates white Ba 2 SiO 4 . 4. Calcium Chlorid precipitates white Ca 2 SiO 4 ACETIC ACID, H(C 2 H 3 2 ) This and the following acids belong to the division known as Organic Chemistry. They are, however, frequently used in the process of analysis, and it is therefore desirable to know how to recognize them. Acetic acid is a liquid, boiling, when pure, at 119. It has a characteristic pungent odor like that of vinegar. It is a monobasic acid, and its salts are all soluble in water, although one or two of them may be partly precipitated if the solutions are sufficiently concentrated. For the reactions use a solution of Na(C 2 H 3 O 2 ). 1. Silver Nitrate precipitates, from concentrated solutions, white Ag(C 2 H 3 2 ). Soluble in NH 4 OH. 2. Ferric Chlorid gives, in a neutral solution, a deep red colora- tion, but no precipitate. If this solution is now boiled, a pre- cipitate of red-brown basic ferric acetate, Fe(OH) 2 (C 2 H 3 O 2 ), appears and the solution becomes colorless. 3. Sulfuric Acid gives no precipitate, but liberates acetic acid, which may be recognized by its odor. 4. If a few drops of alcohol are added to a solution of an acetate, then a little concentrated sulfuric acid, and the whole warmed, ethyl acetate (acetic ether) is formed. This is a vola- tile, ethereal liquid, having an agreeable and characteristic odor somewhat like that of apples. 46 QUALITATIVE ANALYSIS OXALIC ACID, H 2 C 2 04 Oxalic acid is a white crystalline solid, soluble in water. It is a dibasic acid. The oxalates of the alkalies are soluble in water, while most of the others are insoluble. For the reactions use a solution of (NH 4 ) 2 C 2 O 4 . 1. Lead Acetate precipitates white PbC 2 O 4 . Soluble in HNO 3 . 2. Silver Nitrate precipitates white Ag 2 C 2 O 4 . Soluble in HN0 3 and in NH 4 OH. 3. Barium Chlorid precipitates, in not too dilute solutions, white BaC 2 O 4 . Somewhat soluble in water. Soluble in HC1 and in acetic acid. 4. Calcium Chlorid precipitates white CaC 2 O 4 . Insoluble in water, in acetic acid, and in NH 4 OH. Soluble in HC1. TARTARIC ACID, H 2 (C 4 H 4 6 ) Tartaric acid is a white crystalline solid, soluble in water. It is a dibasic acid. The normal tartrates of the alkalies are easily soluble in water, while the acid salts dissolve with difficulty. Most of the other tartrates are insoluble in water, but many of them dissolve in excess of the alkaline tartrates, forming double salts. For the reactions use a solution of sodium potassium tartrate (Rochelle salt), %aK(C 4 H 4 O 6 ). 1. Lead Acetate precipitates white Pb(C 4 H 4 O 6 ). Soluble in HNO 3 and in NH 4 OH. 2. Silver Nitrate precipitates white^ Ag 2 (C 4 H 4 O 6 ). Soluble in HNp 3 am\inNH 4 OH. 3. Barium Chlorid, when added in excess, precipitates white Ba(C 4 H 4 O 6 ). Soluble in acetic acid. REACTIONS FOR THE ACID RADICALS IN SOLUTION 47 4. Calcium Chlorid, when added in excess, precipitates white Ca(C 4 H 4 O 6 ). Soluble in all acids and in NH 4 C1. Insoluble in NH 4 OH. 5. If a few drops of silver nitrate are added to a solution of a tartrate in a carefully cleaned test-tube, then NH 4 OH added drop by drop until the precipitate first formed is nearly dis- solved, and the whole gently warmed, the silver tartrate will be reduced and metallic silver deposited as a brilliant mirror upon the glass. (A very characteristic reaction.) PART II REACTIONS FOR DRY SUBSTANCES BLOWPIPE ANALYSIS Matter, at the ordinary temperature, nearly always exists in the solid state. For purposes of analysis it is much more con- venient to have it in the liquid condition. This may be obtained by solution or fusion. Since the latter often requires a very high temperature the former is almost universally employed. In obtaining a solution of a given substance, it sometimes happens that a change in composition occurs, since solution may be attended by chemical change. In order to know the original composition of a substance we may have to analyze it in its original condition. The methods employed for this pur- pose belong to that part of analytical chemistry known as Blowpipe Analysis, and, while it is not desirable to present an exhaustive treatise at this time, a knowledge of the simpler operations belonging to this part of the subject is indispensable to the chemist. By these methods all simple inorganic substances and many organic compounds may be completely analyzed, and many important facts about the more complex compounds may be learned. Some of the phenomena, while perfectly evident, are so slight as to be easily overlooked, unless carefully observed. They are, however, simple and easy to follow, and if carefully observed the results are accurate and conclusive. This part of the work, therefore, is of great value in developing the powers of observation. 48 REACTIONS FOR DRY SUBSTANCES 49 Nearly all of the results are produced by heat, either alone or with reagents. The operations should be carried on system- atically and the student taught to make the proper deductions from each operation before going on to the next. For the heat a Bunsen lamp is employed. This is supple- mented by an instrument called a blowpipe. The lamp should be furnished with an inner tube to be inserted for use with the blowpipe. The latter is used for producing both the oxidizing and reducing flame. The proper manner of using the blow- pipe, and the ability to produce both the oxidizing and reducing flame, should be thoroughly mastered. The operation, which is not very difficult, but is apt to puzzle the student at first, is as follows. Insert the blowpipe tube in the Bunsen lamp and turn down the flame so that it will be from four to live centimeters in length. Bring the tip of the blowpipe into the flame about a third part of the width of the flame and near the end of the inner tube. If a fairly strong current of air is now sent through the blowpipe, a long, pointed, blue, oxidizing flame is produced. This flame is used for fusion and oxidation. The hottest part of the flame is about midway between the point of the inner blue cone and the extreme tip of the flame. The point of maximum oxidation is at the extreme tip of the flame, or even just beyond this if the temperature is found to be sufficiently high. If the tip of the blowpipe is held just outside the gas flame, and a gentle stream of air is sent through the blowpipe, the inner blue cone will be surrounded by a luminous mantle, form- ing the reducing flame. The point of maximum reduction is just within the point of the luminous mantle. The reducing flame is not nearly so hot as the oxidizing flame. The proper use of the blowpipe can best be learned by prac- tice under the guidance of a competent instructor. The other apparatus, as well as the reagents to be used, will be described as they are employed. 50 QUALITATIVE ANALYSIS I. THE EFFECT OF HEAT ALONE The substance is heated in a piece of hard glass tubing closed at one end. It should be heated in the Bunsen flame, at first gently, then strongly. A. Water is given off. This is recognized by its condensation in small drops in the upper part of the tube, and indicates the following about the substance. (a) It is a deliquescent salt. A deliquescent substance is one which absorbs moisture from the atmosphere. The water is driven off at a comparatively low temperature, and usually in small quantities. (b) It contains enclosed water. The substance often decrepi- tates, or crackles, when heated. This is caused by the bursting of the particles. The amount of water is usually small. This reaction is particularly characteristic of NaCl and certain other halogen salts ; also of other salts which do not contain water of crystallization. (c) It contains chemically combined water. The substance may be an hydroxid, an acid salt (generally of a volatile acid or one easily decomposed), an ammonium salt, or an organic com- pound. The steam, as it passes out of the tube, should be tested with a piece of moistened litmus paper. A neutral reac- tion usually indicates an hydroxid; an acid reaction, an acid salt; an alkaline reaction, an ammonium salt. The compound often shows a permanent change of color. If it blackens and gives off empyreumatic odors, it indicates an organic substance. This action may require a fairly high temperature. The amount of water varies, but is usually not very great. (d) It contains water of crystallization. A portion of the water comes off at or below 100, but the last portion may require a much higher temperature. The amount of water is REACTIONS FOR DRY SUBSTANCES 51 usually relatively large. Some alums, borates, and phosphates swell up considerably while giving off their water. B. A gas is given off. This may consist of one or more of the following gases. (a) Oxygen. This is recognized by the igniting of a glow- ing splinter when introduced into the tube. It indicates that the substance was a nitrate, a peroxid, or some highly oxidized salt, such as a chlorate, bromate, iodate, dichromate, or permanganate. (b) Ammonia. This is easily recognized by its odor and the white fumes of NH 4 C1 which are formed when a glass rod moistened with HC1 is held in the escaping gas. It indicates an ammonium salt. (c) Carbon dioxid. This is recognized by the turbidity which is caused when lime-water [Ca(OH) 2 ] is exposed to the gas. It indicates a carbonate or an organic compound. The latter usually blackens by heating. All carbonates give this except normal carbonates of the alkali metals. These, especially the commercial carbonates, sometimes contain, as impurity, small quantities of the acid carbonates, which give this reaction. A convenient instrument for showing this is made from a piece of glass tubing drawn out at one end to a capillary and bent like a siphon. This is attached to the closed tube by a piece of rubber tubing, the capillary end being placed in the lime-water and the closed tube heated. (d) Carbon monoxid. This gas burns with a bright blue flame, by which it may be recognized. This indicates an oxalate or a formate. The latter blackens when heated. The flame of carbon monoxid does not always appear blue, because of the presence of some impurity (such as sodium). It may also fail to appear because of the presence of water in the form of steam. 52 QUALITATIVE ANALYSIS (e) Sulfur dioxid. This may be recognized by its odor, that of burning sulfur. It indicates a sulfate or a sulfite. (f) Hydrogen sulfid. This is recognized by its odor, and by blackening a piece of paper moistened with lead acetate. It indicates a sulfid containing water. (g) Nitrous oxid. This supports combustion nearly as well as oxygen and may be recognized by the same test. It indicates ammonium nitrate. (h) Nitrogen trioxid or tetroxid. These gases may be recog- nized by their brownish-red color and a peculiar odor which is like that of nitric acid. They usually indicate a nitrate or nitrite of the heavy metals. The alkaline salts do not give this reaction. (i) Chlorin. This is recognized by its yellow color and its odor; also by its bleaching action. It indicates certain chlorids and hypochlorites. (/) Bromin. This is recognized by its red-brown color and its odor, which is much like that of chlorin. It indicates certain bromids, and other bromin compounds. (k) lodin. This is recognized by its deep violet-colored vapor. It indicates iodin, an iodid, or some other iodin compound. (I) Cyanogen. This is recognized by its odor (Poison !), which is like that of KCN, and by the crimson color of its flame. It indicates a cyanid of one of the less basic metals. (m) Organic gases. These may usually be recognized by their inflammability, the flame being more or less luminous. They indicate an organic substance. C. A sublimate is formed. Some substances when heated pass directly from the solid to the gaseous state. When this gas comes in contact with the colder surface of the upper part of the tube it condenses again to a solid, forming a sublimate. This REACTIONS FOR DRY SUBSTANCES 53 usually begins to form at a distance of one or more centimeters from the substance, and the line of formation is usually a sharp one, the sublimate shading off gradually above it. Sometimes \\hen a substance fuses, a film of the melted material may extend up the tube for some distance from the substance. This must not be mistaken for the sublimate. The sublimates vary much in color. 1. A white sublimate is formed by the following substances. (a) Ammonium salts. If two or three drops of a solution of NaOH are placed in the tube with the substance and warmed (the original substance may be treated this way in a test-tube), ammonia is given off, which may be recognized by its odor. (b) Mercurous chlorid. The sublimate is yellow while hot, but becomes white on cooling. (c) Mercuric chlorid. This is much like the mercurous chlorid, but fuses before it sublimes. (d) Arsenic trioxid. This gives a sublimate of octahedral crystals. If a bit of charcoal is placed in the tube with the substance and heated, a black mirror of arsenic is produced. 2. A colored sublimate is formed by the following substances. (a) Arsenic. This gives a black shining mirror, and may be formed by the element itself, or by its compounds in the pres- ence of a reducing agent, like carbon. The vapor has a peculiar garlic-like odor. (b) Antimony sulfid. This sublimes only at a very high temperature, the sublimate being black when hot, and reddish brown when cold. (c) Mercuric sulfid. This forms a black sublimate which shows red when rubbed with a glass rod. (d) lodin. This gives a black sublimate and a deep violet- colored vapor. 54 QUALITATIVE ANALYSIS (e) Arsenic sulfid. This is dark reddish brown while hot, and yellowish red when cold. (/) Sulfur. This may come from free sulphur or from cer- tain sulfids. The sublimate is brownish yellow while hot, and sulfur yellow when cool. It burns easily, giving off sulfur dioxid. (g) Mercuric iodid. This forms a yellow sublimate, which soon changes to red, especially if rubbed with a glass rod. D. The substance changes color. Many substances, upon heating, do not change in composition, but change in appearance. Upon cooling, the original color appears. Many salts decompose on heating, leaving an oxid of the metal which may have a bright color. This may exhibit some of the phenomena given below, and thus give a clew to the original substance. The following are characteristic. (a) The substance is white, becomes yellow when hot, and white again on cooling. This indicates zinc oxid, ZnO. (b) The substance is white, becomes yellowish brown when hot, and is a dirty pale yellow on cooling. At a high temperature it is infusible and luminous. It indicates stannic oxid, SnO 2 . (c) The substance is orange or light yellow, becomes brown red when hot, and yellow when cold. It is fusible at a high temperature. It indicates lead oxid, PbO. (d) The substance is yellow, becomes orange yellow or red brown when hot, and yellow when cold. It is fusible at a high temperature. It indicates bismuth oxid, Bi 2 O 3 . (e) The substance is red or red brown, becomes a very dark red brown, almost black, when hot, and red when cold. It is infusible. It indicates ferric oxid, Fe 2 O 3 . (/) The substance is yellowish red, becomes dark brown or black when hot, and red when cold. It decomposes when REACTIONS FOR DRY SUBSTANCES 55 strongly heated, forming a black or gray sublimate, and giving off oxygen. It indicates mercuric oxid, HgO. E. The substance fuses without decomposition. This generally indicates that the substance is an alkaline salt, though a very few other salts do this. F. The substance carbonizes. Water is usually given off to- gether with gases having a characteristic odor. This indicates an organic compound. Only a few organic compounds can be completely determined by these actions. They are the more common acids, and metallic salts of organic acids. (a) Acetates. These give off aceton, which has a characteris- tic odor somewhat suggestive of vinegar. (b) Formates. These give off CO, which burns with a bright blue flame. (c) Tartrates. These give an odor like that of burnt sugar. These salts of metallic bases and organic acids in decompos- ing by heat always leave a carbonate of the metal, which by further heating will form an oxid of the metal, unless the latter belongs to the alkali group. The carbonate may be recognized by the effervescence when a drop of HC1 is added. H. THE SUBSTANCE IS HEATED ON CHARCOAL This gives the effect of heat in the presence of a strong reducing agent, the hot charcoal. A shallow cavity is made in a piece of soft-wood charcoal, the substance is placed in this, and heated with the blowpipe flame, at first gently, afterwards strongly. It is best to hold the charcoal somewhat inclined toward the flame, so that in case an incrustation should be formed it may . be observed more readily. If the substance is very light and dry, so that it is liable to be blown away, it may be moistened with water, or, in some cases, a small piece of borax may be fused with the substance. 56 QUALITATIVE ANALYSIS A. The substance fuses easily without decomposition, and sinks into the charcoal. This indicates a salt of the alkali metals or some of the salts of the alkaline earths. B. The substance yields a metallic bead without any incrustation. (a) The bead is white. This indicates tin, aluminum, or silver. Tin is very easily fusible (232), aluminum requires quite a high temperature (655), and silver fuses only with great difficulty (960). (b) The bead is red. This indicates copper and requires a very high temperature (1080) and often long-continued heating. The bead is malleable, which distinguishes it from CuO, which is red, but brittle. (c) The bead is yellow. This indicates gold and requires a very high temperature (1061). C. The substance yields a metallic bead with an incrustation. (a) The bead is white, soft, and malleable. The incrustation is yellow and volatile. This indicates lead. (b) The bead is white, soft, and malleable. The incrustation is red brown and volatile. This indicates cadmium. (c) The bead is white, rather hard, and malleable. It is pretty well covered with the incrustation, which is yellow when hot, and white when cold. This indicates zinc. (d) The bead is white, hard, and brittle. The incrustation is white and volatile. This indicates antimony. (e) The bead is white, hard, and brittle. The incrustation is yellow and volatile. This indicates bismuth. D. The substance is infusible, dark brown or black in color, gives no incrustation, and is more or less easily attracted by a magnet. This indicates iron, chromium, nickel, cobalt, or manganese. These may be distinguished by the borax bead. [See Coloration of the Borax or Microcosmic Bead.] REACTIONS FOR DRY SUBSTANCES 57 Molybdenum, tungsten, and some of the platinum metals are infusible, and would be found here. They are not magnetic. E. The substance deflagrates, or burns up quickly. This indi- cates a nitrate, or some highly oxidized salt, such as a chlorate, bromate, or iodate. Only the alkaline salts show this action in a marked degree ; other salts are scarcely to be recognized by this test. F. The substance decrepitates. This indicates a crystalline salt, which may contain enclosed water but does not usually contain water of crystallization. Sodium chlorid and other halogen salts show this action best. G. The substance volatilizes. (a) It forms a white, very volatile incrustation, and has a strong garlic odor. This indicates arsenic or some of its coin- pounds. (Poison !) (b) Those substances which form a sublimate in the closed tube are volatile on charcoal, and some of them give an incrus- tation which is volatile and similar in color to the sublimate. H. The substance burns. (a) The substance is a metal which burns with a brilliant white light, leaving a white infusible oxid. This indicates magnesium. (b) The substance is a metal which burns with bright scintil- lations, leaving a dark brown or black oxid, which is magnetic. This indicates iron. (c) The substance is a metal which burns with a bright white light, leaving an oxid which is yellow while hot, and white when cold. This indicates zinc. (d) The substance burns with a blue flame, giving off SO 2 , which may be recognized by its odor. This indicates sulfur. 58 QUALITATIVE ANALYSIS I. The substance is infusible, white, and highly luminous when strongly heated. It should be allowed to cool somewhat, and the residue then moistened with a drop or two of a solution of Co(NO 3 ) 2 , and again ignited strongly. The mass, on cooling, should then show a characteristic color as follows. (a) Slue. This indicates aluminum oxid and alkaline phos- phates or borates. Silicon dioxid and certain silicates when heated very strongly show this same reaction. * (b) Flesh color. This indicates magnesium oxid. (