NOTES ON QUALITATIVE ANALYSIS BY HORACE G. BYERS 1 PROFESSOR OF CHEMISTRY, UNIVERSITY OF WASHINGTON AND HENRY G. KNIGHT DIRECTOR OF EXPERIMENT STATION UNIVERSITY OF WYOMING NEW YORK D. VAN NOSTRAND COMPANY 23 MURRAY AND 27 WARREN STREETS 1912 \ COPYRIGHT, 1912, BY D. VAN NOSTRAND COMPANY. Nottoooti J. 8. Cushing Co. Berwick & Smith Co. Norwood, Mass., U.S.A. PREFACE THE publication of a new textbook on qualitative analysis requires an apology on the part of the authors, in view of the number and variety of those already published. We submit the following : Most of the small texts are designed, or at least serve, merely as a guide to the laboratory practice of the art of analysis. Many ignore wholly any relation between the laboratory work and the principles of general chemistry. The larger works are of two types : those which are extremely detailed labora- tory guides and those which couple the theories and laws of chemistry with directions for analytical procedure to an extent which makes them too involved for use in the ordinary labora- tory course, where a limited amount of time can be devoted to the subject. In addition there are always special conditions which confront the teachers in their various environments which render a particular arrangement desirable. In our courses we present qualitative analysis as a part of the laboratory work in general chemistry, and it therefore serves as the experimental illustration of the chemistry of the metals. We also offer an advanced course which serves to develop the art of analysis and also deals in a broader way with the rationale of the operations. We have therefore prepared these notes, which represent the kind of course we desire to give our own students, and offer them for publication chiefly for home consumption. If others have our viewpoint, they may find them useful. We have tried to keep in mind that qualitative analysis should develop laboratory technique and enable the student success- fully to carry out the identification of inorganic materials. It should also widen the knowledge and appreciation of the student of the fundamental principles of chemistry and enable him to make use of the general laws and theories as applied iii 258707 iv PREFACE to particular and individual questions. It should not be com- plete within itself, but should stimulate to wide reading and the correlation of the occurrence, preparation, and uses of the elements and their compounds with the reactions used for their separation and identification. These aims we seek to accom- plish through the reactions, the exercises, the analytical tables, and the application of general principles to particular cases, with the hope that the student will emerge from the course, not only a fair analyst, but with a broadened knowledge upon which to develop his further chemical education. In view of the very large number of our students who use the course merely as an adjunct to the chemistry of the metals, and who therefore deal chiefly with Parts II and III, we have also tried to present an elementary system which will not demand the use of platinum, and so permits a great economy in equipment for large classes. It is recognized that in the regular course of analysis of substances many complications may arise which will require modifications of treatment not provided for in these notes, but it is hoped that the matter presented will enable the student tcT recognize and deal intelligently with them when they appear. It is advised that each laboratory library or reading shelf be provided with reference books instantly available to students. These should include such books as Bottger's Qualitative Analyse, 2d Auflage, Treadwell and Hall's Analytical Chem- istry, Stieglitz' Elements of Qualitative Analysis, and Smith's (Alexander) General Chemistry. In preparing the material for these notes free use of informa- tion, wherever available, has been made, and it is impossible to properly credit the sources. No pretense to any special research -is made by the authors. Sincere thanks are due to Dr. R. E. Rose for his painstaking criticism of the manuscript and proof. H. C. B. SEATTLE, Jan., 1912 H. G. K. CONTENTS INTRODUCTION PAGE QUALITATIVE ANALYSIS i BASIS OF IDENTIFICATION i CONDITIONS PRODUCING REACTIONS . . . . . . 2 PROPERTIES USED IN IDENTIFICATION 3 (1) State of Aggregation S . . . . . 3 (2) Color x ". ........ .3 (3) Odor x 3 (4) Taste . ' X . .3 (5) Melting and Boiling Points ..... .3 (6) Spectra ^ -4 (7) Solubility^ 4 IMPORTANCE OF VARIOUS PHASES OF WORK . . 5 PART I CHEMICAL PRINCIPLES INVOLVED IN QUALITATIVE ANALYSIS . . 7 DEFINITION OF SOLUTION 7 KINDS OF SOLUTION .... 7 PHENOMENA OF SOLUTION "... 8 HYDRATES IN SOLUTION ... 9 HYDRATION OF IONS ... ...... 9 OSMOTIC PRESSURE ... 10 VAN'T HOFF'S HYPOTHESIS 13 FREEZING POINT OF SOLUTIONS 15 BOILING POINT OF SOLUTIONS 17 ACIDS, BASES, AND SALTS IN SOLUTION 18 ELECTROLYSIS 19 vi CONTENTS PAGE HYPOTHESIS OF ARRHENIUS 21 PHYSICAL EQUILIBRIUM 23 CHEMICAL EQUILIBRIUM 25 IONIC EQUILIBRIUM 27 SOLUBILITY PRODUCT 30 QUALITATIVE ILLUSTRATION 31 Group I 31 Group II 32 Subgroup A 36 Subgroup B 39 Group III . .40 Hydrolysis 42 Oxidation 48 Bead Tests 50 Group IV ........... 51 Flame Tests 52 Group V 52 Supersaturation 55 PART II METAL ANALYSIS GENERAL DIRECTIONS . . . 57 Equations ........... 57 Precipitation and Filtration 57 Washing ........... 58 Decantation 59 Evaporation 59 Amount of Sample .59 Confirmation of Tests .59 Tests of Reagents 59 Notebooks 60 THE HYDROCHLORIC ACID GROUP 60 Silver 60 Mercury 61 Lead 62 ANALYSIS OF GROUP I 63 CONTENTS Vll PAGE EXERCISES OF GROUP I .64 THE HYDROGEN SULPHIDE GROUP . 6 5 SUBGROUP A Mercury .... "5 Lead . . . - 6 7 Bismuth Copper . Cadmium 6 9 ANALYSIS OF SUBGROUP A . ^9 SUBGROUP B 7 2 Arsenic 7 2 Antimony 75 Tin ANALYSIS OF SUBGROUP B . . .... 77 EXERCISES . THE AMMONIUM SULPHIDE GROUP ... -79 Iron . . . ...... 80 Chromium Aluminium ... Manganese ^4 Zinc *5 Nickel . Cobalt 86 ANALYSIS OF GROUP III .... EXERCISES ... .... 91 AMMONIUM CARBONATE GROUP . . ... 91 Barium ... .92 Strontium 9 2 Calcium . -93 ANALYSIS OF GROUP IV .. 94 EXERCISES 95 THE SOLUBLE GROUP .... 95 Sodium .... 96 Potassium . . 9^ Ammonium 97 Magnesium 9^ ANALYSIS OF GROUP V . . 98 viii CONTENTS PART III ACID ANALYSIS PAGE INTRODUCTION . . . 101 Rules of Solubility of Salts 101 Grouping of Acids 103 GROUP I 104 Silicic Acid . . . . . . . . . . .104 Hydrosulphuric Acid . . . . . . . . .104 Thiosulphuric Acid . . . . . . . . .104 Sulphurous Acid 104 Nitrous Acid 105 Carbonic Acid 105 Hypochlorous Acid 106 ANALYSIS OF GROUP I . 106 EXERCISES ON GROUP I 107 GROUP II 107 Hydrochloric Acid 107 Hydrobromic Acid 108 Hydriodic Acid 108 Hydrocyanic Acid 109 Hydroferrocyanic Acid 109 Hydroferricyanic Acid no PREPARATION OF A SOLUTION OF THE ANIONS 115 ANALYSIS OF GROUP II no EXERCISES ON GROUP II 112 GROUP III 112 Sulphuric Acid . . . . . . . . . .112 Chromic Acid . . . . 112 Arsenic Acid 113 Arsenous Acid 113 Phosphoric Acid 114 Boric Acid 114 Hydrofluoric Acid . . . . . . . . .114 Silicic Acid . . 115 ANALYSIS OF GROUP III -. .115 EXERCISES ON GROUP III . 116 CONTENTS ix PAGE GROUP IV 117 Nitric Acid 117 Chloric Acid 118 Permanganic Acid 118 ANALYSIS OF GROUP IV 118 EXERCISES ON GROUP IV 119 GROUP V " . .119 Oxalic Acid 119 Tartaric Acid 120 Acetic Acid 120 Other Organic Acids 120 ANALYSIS OF GROUP V 121 EXERCISES ON GROUP V . . 121 PART IV SYSTEMATIC ANALYSIS PRELIMINARY EXAMINATION .122 Closed Tube Test . . . . . . . . . .123 Bead Test 124 Flame Test 125 Charcoal Test .125 PREPARATION OF THE SAMPLE .126 A Liquid ... . . . . . . . . .127 An Alloy . . . . . . . . . . .127 A Non-metallic Substance . . . - . . . .128 A. For Metal Analysis . 128 B. For Acid Analysis 129 TABLES OF ANALYSIS 130 PART V THE RARE METALS PRELIMINARY STATEMENT , .138 Flame Test ........... 139 Microcosmic Bead Test . . . 139 Grouping of Rare Elements - . 140 X CONTENTS PAGE GROUP I .140 Thallium 140 Molybdenum 141 Tungsten . . 142 Tantalum and Niobium 143 GROUP II .... -145 Gold .... . 145 Platinum 147 Palladium .... 148 Osmium 149 Iridium 149 Ruthenium and Rhodium 150 Selenium 150 Tellurium 151 Germanium 152 GROUP III . ... 152 Vanadium . . . . . . . . . . .152 Titanium 153 Zirconium 155 Uranium 156 Beryllium 157 Thorium . . . . . . . . . . .158 Cerium 159 Indium and Gallium 160 Lanthanum, Didymium, Yttrium, Scandium, Erbium, etc. . .160 GROUP IV 161 Radium . . . . . . . . . . . .161 GROUP V 161 Lithium ............ 161 Rubidium and Caesium ... . .162 APPENDIX LIST OF APPARATUS 165 REAGENTS IN SOLUTION 166 REAGENTS, SOLID 171 CONTENTS xi PAGE TABLE OF ATOMIC WEIGHTS . .172 PERIODIC SYSTEM . i?3 TABLE OF SOLUBILITIES . i?4 METRIC SYSTEM CONVERSION TABLES . i?5 DEGREE OF IONIZATION ACIDS, BASES, AND SALTS . . . i?7 INTRODUCTION THE first problem which confronts the analytical chemist is one of identification. He must determine what elements or compounds are present, whether this alone be the ultimate object of his search, or whether he be also concerned with rela- tive quantities. These two phases of analytical chemistry are usually considered separately, and are known as Qualitative Analysis and Quantitative Analysis. Yet neither is inde- pendent of the other. Even where we only desire to know what substances are present, it is impossible to ignore the relative amounts. When we wish to determine quantities exactly, it is imperative to know the constituents in order properly to select methods of separation. It follows that qualitative must always precede quantitative investigation. The complex substances with which the analyst deals may consist of mixtures or of compounds ; i.e. of substances formed by the distribution of one kind among those of other kinds, or of substances formed by chemical union of their constituents. The separation of mixtures into their component parts may at times be accomplished wholly by mechanical means, but usually, with mixtures, always with compounds, whatever degree of separation is required for identification is accomplished by chemical methods. The basis of identification is mainly empirical ; i.e. we are guided by the observation of a limited number of facts which may or may not be dependent upon scientific generalizations. It is obvious that such identification is never absolute, be- cause it would only be possible to assure ourselves of the com- plete identity between a sample of a known and one of an 2 INTRODUCTION unknown substance by comparison of all the properties of each. We tacitly assume, however, that the number of different sub- stances is finite, and that if two samples of material possess in common a certain number of distinct properties, they will be alike in all. In choosing the properties which we may use for purposes of such comparison, we must avoid those which are common to all substances or to such a large number as to be useless, also such as are incidental or easily affected by circumstances, as size, position, temperature, electrical con- dition. Such properties as may be used we call characteristic. Even in the case of characteristic properties, use must be made of several, and with increase in the number used, certainty is rapidly increased. For example, if a substance is yellow and hard, it is not safe to assume that it is sulphur ; but if, in addi- tion, we find it is also brittle and burns with a blue flame, giving rise to a distinctive odor, it becomes practically certain that the substance under consideration is wholly or partially sulphur. Ordinarily, use is made of properties of state or condition, such as color, odor, taste, and properties of reaction, or changes which may be brought about in particular forms of matter and which are peculiar to them either because of the change itself or of the properties so produced. Reactions are induced by some change in conditions under which we inspect the material investigated; the most important means by which reactions are produced are (i) change of tem- perature, (2) change of electrical condition, (3) change of chemi- cal condition by contact with other substances. Of these the last is by far the most important, and is often used in conjunction with one or both the others. Contact of substances is usually most easily secured when they are in the liquid condition, and consequently such contact is most frequently procured by bring- ing the materials into the liquid state by solution or by fusion. INTRODUCTION 3 Recognition of elements or compounds, therefore, depends upon physical properties, and usually upon such as are produced as the result of certain physical or chemical changes. Almost always the determination is based upon a quantitative differ- ence, such as degree of hardness, specific gravity, solubility, etc. The properties most frequently used in qualitative analysis are : 1 I ) State of Aggregation. The existence of a substance in the solid, liquid, or gaseous state under certain conditions is frequently characteristic, as also whether amorphous or crystalline, whether viscous or mobile. (2) Color. Often substances betray their identity by color or color changes ; e.g. gold, copper, iodine, oxides of nitrogen ; likewise, the colors produced by chemical reaction are frequently characteristic. (3) Odor. Substances may have, or may be made to produce by reaction, odors which are decidedly characteristic, and while not capable of accurate description are recognized if once familiar. Examples are the odor of chlorine, of acetic acid, of hydrogen sulphide. (4) Taste. Tastes, also, are incapable of definite measurement, but may frequently serve the analyst as means of identification, as in the case of sugar, aconite, etc. (5) Melting and Boiling Points. The transition temperature from solid to liquid or from liquid to gaseous condition is a fixed point for all pure substances, and often serves the chemist not only qualita- tively but quantitatively. 4 INTRODUCTION (6) Spectra. The vapor of each element when heated to incandescence gives rise to light waves of certain definite character. If these light waves are dispersed by a prism, as in the ordinary spectroscope, they produce bands of color which are characteristic of the element which produces them. No two elements produce the same set of color bands, and the position of the bands on the prismatic scale is an absolute means of identification. The delicacy of the spectroscope is also amazing. It is asserted that by its use the presence of one third of a millionth of a milligram of potassium may be detected with certainty. Magnesium is one of the most difficult elements to detect by this method, yet one hundred-thousandth of a milligram may be so detected. Notwithstanding this extreme delicacy, the in- strument is used to but a limited extent in elementary work. The reasons are, in part, that the method furnishes no clew to relative quantities, and indicates only the ele- ments present, not their state of combination in the material investigated. It may also be said that except for a few elements, the conditions for incandescent vapor are not easily obtained, and considerable expense is attached to the purchase and use of an instrument capable of wide application. Small spectroscopes are sometimes used to assist in the detection of alkali and alkaline earth metals. For a full discussion of the instrument and its applica- tion, the student is referred to Landauer, Craw, Lockyer and others. (7) Solubility. The solubility of substances, changes of solubility by changes in chemical combinations, and the marked physi- cal properties of certain soluble and insoluble substances produced by chemical changes are of such general appli- INTRODUCTION 5 cation that most of the operations of qualitative analysis make use of solubility changes. The most extensively used means of identification is the preparation of a solution of the substance, and then, by addition of certain known reagents, produce the separation from the solution of certain definite compounds, either as gases, liquids, or solids, which have such definite properties as leave no doubt as to the materials producing them. There are many reasons why this method is of such general application. It is simple, easy of manipulation, does not require expen- sive apparatus, does not require a high development of a scientific training, furnishes a rough idea of relative quantities, and, above all, conveys to the student a general knowledge of chemical relations which is required for quantitative manipulation. It will be seen at once that to use successfully the properties above mentioned for the purpose of identification requires thorough acquaintance with chemical substances. This acquaint- ance is best secured by a systematic classification of the sub- stances proposed for study and a thorough drill in the changes which they are likely to undergo. In an elementary course, the study is usually limited to a small number of the best known elements and such of their compounds as are most frequently encountered. The subject divides itself naturally into two parts : the detection of the metallic elements and of the acid radicals. The metals are found to fall naturally into groups, the members of which are related by similarity of behavior under specific conditions. It is obvious that such grouping is somewhat arbitrary, and the particular one used is largely a question of convenience. The grouping used in this outline is the one used most commonly, and is presented in Part II. The acid 6 INTRODUCTION radicals and non-metallic elements group themselves around common reactions much less satisfactorily than do the metals, and are less susceptible of systematic treatment. The outline suggested in Part III is more or less new, but follows the gen- eral plan found to be most satisfactory. The necessary acquaintance with the behavior of elements and radicals under definite conditions is secured by means of a series of preliminary reactions, using known materials. When familiarity with these reactions has been gained, and not till then, is it possible to 'proceed intelligently to the identi- fication of unknown substances. In analytical work the first requisite is accuracy. When an analysis is completed, the analyst must have absolute confidence in his results. As long as a doubt remains, more evidence should be secured until the doubt is resolved. An important point to be noted is that the test of skill is the detection of small quantities of substances. The second requisite is speed. Accuracy is never to be sacrificed, but the successful analyst must reach certainty without undue sacrifice of time. The study of qualitative analysis involves more than merely the development of skill in the identification of chemical indi- viduals by means of routine methods of separation. The large number of individual substances and their variations demand, in addition to the development of manipulative skill, a careful study of the rationale of the operations and of the chemical principles involved in the various steps. A sort of syllabus of these principles is presented in Part I. The acquaintanceship gained by practical experience with substances, the classification of the knowledge of the relation- ship of the substances studied, and the rationale of the processes should go hand in hand, and consequently in the following pages the attempt is made to supply at least the topics of a course of lectures and recitations which shall accompany the laboratory work and exercises furnished by Parts II and III. PART I CHEMICAL PRINCIPLES INVOLVED IN QUALITATIVE ANALYSIS SOLUTION. Since the vast majority of the operations of analysis are carried out by means of solution in water, a con- ception of the present state of knowledge of solutions is clearly important. Unfortunately there is no generally accepted defi- nition of solution, and any which may be given involves some points more or less in controversy. Perhaps as satisfactory as any is the statement, " A solution is a homogeneous mixture of two or more substances." In using the term " mixture " above, the occurrence of chemical change during the process of forma- tion of a solution is not necessarily precluded. Without enter- ing upon an academic discussion of the question of the nature of the process of solution, it may be as well to consider that chemistry recognizes three conditions under which two or more substances may exist together, viz. as mixtures, compounds, and solutions. It is clear that, by the above definition, solutions include the following mixtures : (a) Solids, liquids, and gases with gases. () Solids, liquids, and gases with liquids. (c) Solids, liquids, and gases with solids. To the analyst, b is of maximum importance, and in the fol- lowing discussion attention will be practically wholly confined to it. Also, from the viewpoint of the definition it appears that solution is to be considered as a mutual act on the part of the 7 8 QUALITATIVE ANALYSIS components. Yet we ordinarily speak of the liquid, or that component present in the larger quantity, as these/vent, and the other component as the sohite. According to the orthodox view of the subject, solution is not considered as involving chemical change ; as, for example, solution of sugar or of salt in water is not a chemical process, though chemical changes may subse- quently take place in the substance dissolved, or between the components. There are, however, certain phenomena which accompany the process which are difficult to harmonize with this view. In every case of solution, heat changes and volume changes occur. Since it happens that when the volume of a solution is less than the sum of the volume of its components, heat is usually evolved, and when the volume of the solution is greater than that of the components, a lowering of the tempera- ture is observed, it might perhaps be inferred that the energy changes are due to the volume changes. It is, however, true at times that a lowering of the temperature is coincident with con- traction of volume, and an evolution of heat with expansion of volume. The following experiment will illustrate the variability of these phenomena. EXPERIMENT I. Mix in graduated cylinders the following substances and note the change of temperature, and after the room temperature is regained, the change in volume. (a) 50 c.c. alcohol with 50 c.c. water ( + , contraction). () 50 c.c. sulphuric acid with 50 c.c. water ( + , contraction). (c) 50 c.c. acetic acid with 100 c.c. water ( , contraction). (d} 40 c.c. alcohol with 60 c.c. carbon disulphide ( , expansion). In a general way it may be said that the properties of solu- tions are not additive. A fuller discussion of this phase of the subject may be found in Ostwald's Solutions, in which also abundant references to the original literature will be found. There are many cases where the act of solution is certainly accompanied by chemical change. Examples are found in the CHEMICAL PRINCIPLES 9 solution of anhydrides to form acids, of basic oxides to form bases, and the formation of hydrates by solution of anhydrous salts. There are other cases of undoubted chemical union where separation of the substance formed has not yet been ac- complished, such as the formation of sulphurous and carbonic acids, ammonium hydroxide, etc. Rupert 1 reports the separa- tion of hydrates of ammonium hydroxide. There are also ex- amples of the formation of definite hydrates of such character that the composition varies with pressure or the temperature in a manner different from that of ordinary chemical compounds. These are represented by the so-called constant boiling mixtures of hydrochloric acid, nitric acid, hydrobromic acid, etc. with water. Jones 2 claims to have demonstrated the existence in solution of complex hydrates which cannot be isolated in a pure condi- tion, but which are to be considered as partially the cause of the failure of concentrated solutions to conform to the Ostwald dilution law (vide infra). The student will find a brief ex- position of the Hydration Theory of Jones in his work on Physi- cal Chemistry and a resume of the Solvate Theory in the American Chemical Journal, Vol. 41, p. 19. Washburn 3 has also demonstrated the existence of hydrated ions in the case of solutions of the chlorides of the alkali metals which are themselves never hydrated when crystallized from solution; and McGee 4 has shown the same to be true for the chlorides of the alkaline earths. In view of such facts, some chemists hold the view that solu- tions are chemical compounds of a less definite character than those which can be isolated in a pure condition. For a more extensive view of this phase of the subject, the student is re- 1 Jour. Am. Chem. Soc., Vol. 31, p. 866; Vol. 32, pp. 748-749; June, 1910. 2 Am. Chem.Jour., Vol. 41, p. 19. 3 Jour. Am. Chem. Soc., Vol. 31, p. 322. 4 Dissertation for Master's Degree, U. o/ IV., 1911. I0 QUALITATIVE ANALYSIS ferred to a review of the work of Werner l and to the work of Jones 2 and Kahlenberg. 3 The only general view regarding the condition of dissolved substances which has gained extensive credence is that of van't Hoff. This hypothesis has been extremely fruitful of results in systematizing the reactions with which the analyst deals and in furnishing illuminating viewpoints concerning the reasons for the results obtained. It will therefore be presented after the experimental facts upon which it rests are discussed. Osmotic Pressure If any soluble substance be put in the bottom of a cylinder and water placed upon it, gradually the solvent will take up the sub- stance, and, in the course of time, the solute will be found uni- formly distributed through the solvent. This phenomenon may be ascribed to a chemical reaction between the solvent and solute, 4 or it may be considered as due to a special kind of force called osmotic force. If there is placed between the solvent and the solute a partition of such character that the solvent can pass through and the solution formed cannot, the same action will take place. The following experiments illustrate the point : EXPERIMENT II. In a narrow cylinder place a layer of chloroform, then a thin layer of water, and on top superimpose a layer of ether. Stopper securely, mark the level of the water layer, and allow to stand. The ether will gradually pass through the water into the chloroform, increasing the volume of the lower layer. EXPERIMENT III. Three eggs of equal size may be taken and, after re- moval of the calcareous portion of the shell by means of concentrated hydro- chloric acid, treated in the following manner : One is placed in pure water, another in concentrated calcium chloride solution, and the third is kept for comparison. After twenty-four hours one will be found largely distended by the influx of pure water, the other shriveled by extraction of water. 1 Am. Ckem.Jour., Vol. 22, p. 312. 2 Am. Chem. Jour., Vol. 41, p. 19. 3 Jour. Phys. Chem., Vol. 5, p. 339. 4 Kahlenberg, Jour. Phys. Chem., Vol. 7, 1907. CHEMICAL PRINCIPLES II If the semipermeable layer is so constructed that it may be held in one position, the solvent passing through it will exert a pressure which is capable of direct measurement. This pressure is known as osmotic pressure. Such an apparatus, suitable for qualitative demonstration, is prepared as follows : EXPERIMENT IV. Bore a hole in a carrot, scrape the outer skin, fill the hole with a saturated solution of sugar, and stopper tightly with a cork carrying a glass tube four feet or more in length. Place the whole in a beaker of pure water, clamp in an upright position, and mark the level of the solution in the glass tube. In a few minutes, the level of the liquid will be seen to be rising, and the height to which it will go will depend solely upon the strength of the carrot cell walls. Such an apparatus is clearly not capable of withstanding very great pressures, and hence is not suitable for exact measurement of osmotic force. Pf effer 1 prepared osmotic cells in such a manner that he was able to measure accurately the quantitative value of this tendency of a solvent to pass into and dilute a solution. His apparatus consisted of a porous cup in the walls of which was deposited a film of copper ferrocyanide. The de- tails of the method of preparation of the apparatus are given in all works on physical chemistry and are omitted here. Using this variety of cell, Pfeffer obtained the following re- sults for sugar solutions : PER CENT SUGAR IN SOLUTION PRESSURE IN CM. OF MERCURY PRESSURE PER i PER CENT 1. 00 2.00 53-2 101.6 53-2 50.8 2.74 4.00 6.00 i5i-3 208.2 307-5 55-4 52.1 5i-3 1. 00 53-5 53-5 1 Osmotische Untersuchungen, Leipzig, 1877. 12 QUALITATIVE ANALYSIS It will be noted that the results in the third column range irregularly around a mean value and indicate that the pressure is proportional to the concentration. Pfeffer also shows the following results to be obtainable for the effect of temperature upon osmotic pressure, using a I per cent sugar solution. (Ostwald, Outlines of General Chemistry, p. 128.) TEMPERATURE C. TEMPERATURE ABSOLUTE PRESSURE PRESSURE CALCULATED 142 287.2 5 I.O 51.2 15-5 288.5 52.1 52.9 32.0 305.0 54-4 54.2 36.0 309.0 56.7 55.8 From these and similar results it appears that the temperature coefficient of osmotic pressure is the same as that for gases, viz. 0.00367. The general conclusions drawn from the results of Pfeffer are that the laws of Boyle and of Charles for gases are equally applicable to the variation of osmotic pressure with changes in concentration and temperature. Recently the work of Morse and his coworkers has developed a method of determination of osmotic pressure which leaves very little to be desired in the way of accuracy and has shown that the osmotic pressure of a sugar solution of unit concentra- tion is 24.45 atmospheres at o C. and increases with the abso- lute temperature. The following table * is instructive : 1 Am. Chem. Jour., Vol. 41, p. 274. CHEMICAL PRINCIPLES OSMOTIC PRESSURES WEIGHT SERIES SERIES SERIES SERIES SERIES SERIES NORMAL III IV V VI VIII VII CONCEN- PRESSURE /PRESSURES PRESSURES PRESSURES PRESSURES PRESSURES TRATION 4-5 10 15 20 25 O.I 2.42 2.40 2.44 2.48 2.52 2.56 0.2 4-79 475 4.82 4.91 5.02 5.10 o-3 7.11 7.07 7.19 7-33 7-45 7-57 0.4 9-35 9-43 9.58 9.78 9.96 10.12 0.5 11.75 11.82 12.00 12.29 12.49 12.73 0.6 14.12 14-43 14.54 1486 15.20 15.42 0.7 16.68 16.79 17.09 17-39 17.84 18.02 0.8 19.15 I9-3I 1975 20.09 20.60 20.73 0.9 21.89 22.15 22.28 22.94 23.31 23.66 1.3 24.45 24-53 25.06 25.42 26.12 26.33 Total pressure . . 131.71 132.68 134.75 137-49 140.52 142.24 Mean ratio of osmotic to gas pressure . . 1.074 1.065 1.O6 1 1.064 1.069 1.064 Mean loss in rotation (per cent) . . . 1-73 1.45 O.22 0.10 0.00 0.20 Since, then, osmotic pressure is proportional to the concentra- tion, it follows that Avogadro's principle applies also to this phenomenon. The Hypothesis of van't Hoff This remarkable concordance of the gas laws and those of osmotic pressure led van't Hoff 1 to the conclusion that dissolved substances exert the same pressure, in the form of osmotic pressure, as they would exert, were they gasified at the same temperature, without change of volume. It will be seen that for cane sugar the results of Morse (see table above) do not confirm this conclusion exactly, and the following modification is suggested by Morse 2 : " Cane sugar dissolved in water exerts an 1 Zeitschr.f. Physik. Chem., Vol. I, p. 481 ; 1887. 2 Am. Chem. Jour., Vol. 34, p. 93. I 4 QUALITATIVE ANALYSIS osmotic pressure equal to that which it would exert were it gasified at the same temperature and the volume reduced to that of the solvent in the pure state" The value of this view lies in the fact that, if it be correct for all substances with any kind of semipermeable membrane, we can carry over to solution all the knowledge we possess con- cerning gases, so far as these are dependent upon pressure, volume, and temperature, provided only that osmotic pressure be substituted for gaseous pressure. An instructive example is as follows : The combined gas laws may be formulated pv = RT. (See Walker's Physical Chemistry, p. 28.) If now we consider the val- ues for a gram molecule of a gas, under standard conditions, and seek the value of R, we obtain R = ^= IQ33 x 22488 = ^^ The value 22,488 is the gram molecular volume of gases in cubic centimeters as determined by Morley, and 1033 the atmos- pheric pressure in dynes. If now we substitute osmotic values as obtained by Morse 1 for a volume normal solution of cane sugar, 22.61 atmospheres at o C, R = '033x22610 = 85;554 It follows that measurements of osmotic pressure should lead to the determination of molecular weights just as in the case of gases. If, however, the attempt is made so to use it, the diffi- culty is quickly encountered that for many substances results are Obtained which are entirely at variance with facts otherwise clearly determined. For example, if we dissolve a gram mole- cular weight of potassium chloride (74.5) in 22.4 liters of water, the osmotic pressure ought to be, by the above examples, prac- tically one atmosphere. The experiment shows, however, that it is 1.88 atmospheres. Similar discrepancy is also observed in 1 Am. Chem. Jour., Vol. 34, p. 85. CHEMICAL PRINCIPLES 15 the case of practically all salts and all the ordinary- acids and bases. The explanation of this apparent failure of the facts to con- form to the van't Hoff hypothesis will be presented when some similar classes of facts have been discussed. This explanation is known as the lonization Hypothesis of Arrhenius (vide infra). Before closing the discussion of osmotic pressure, attention should be directed to the fact that it is possible to regard osmotic force only as a special manifestation of that general property of substances which for want of a better name we call chemical affinity. An article by Kahlenberg 1 presents the re- sults of experiments with various membranes and solvents and conclusions wholly at variance with those just discussed. Freezing Point of Solutions When a foreign substance is dissolved in a liquid, the freezing point of the solution is lower than that of the pure solvent, and for moderate concentrations the lowering of the pressure is proportional to the concentration. This fact was clearly demon- strated by Blagden, 2 but his discovery was neglected and for- gotten for almost a hundred years until rediscovered by Riidorff. 3 The deviations from this law in concentrated solutions are probably due to association of molecules of the solute either with each other or with molecules of the solvent. This phase of the subject is rather fully presented by Jones in the Solvate Theory resume referred to on page 9. In 1882 Raoult, 4 work- ing with solvents other than water and with water solutions of organic compounds other than salts, arrived at the generaliza- tion that equimolecular concentrations of different compounds in the same solvent effect the same depression of the freezing point 1 Jour. Phys. Chem., Vol. 10, p. 141 ; {1906. 2 Phil. Trans., Vol. 78, p. 277; 1788. 3 Pogg. Ann., Vol. 114, p. 63. 4 Compt. Rend., Vol. 94, p. 1517 ; Vol. 95, p. 188 and p. 1030. 16 QUALITATIVE ANALYSIS For example, one gram molecule per liter of most substances lowers the freezing point of acetic acid approximately 3-9C. ; of benzene 5.3C. ; of phenol 7.6 C. ; of water 1.89 C., etc. These numbers are known as the molecular depression of the freezing point. A table showing a large number of Raoult's observations is found in Ostwald's Solutions, pp. 209-212. A practical application of the principle to the determination of molecular weights was made by Beckmann, 1 and the appa- ratus for this purpose is one of the familiar objects in modern laboratory equipment. Two classes of deviations from the law of Raoult were brought out by himself and others. Some substances caused a lowering of the freezing point less than the normal, and these values are usually half as large as the molecular depression and suggest at once that molecules of the substances associate to form larger molecules. The other deviation is found most markedly in the freezing points of acids, bases, and salts when dissolved in water, though the variation is not wholly confined to this solvent, but occurs in solutions of the same substances in liquid ammonia, liquid sulphur dioxide, and to a lesser degree in some other solvents. When a substance, such as common salt, is dissolved in water, a normal solution freezes at 3.46 instead of 1.89, and this excessive lowering of the freezing point is common, to a varying degree, to all solutions of acids, bases, and salts in water. Van't Hoff introduced for salt solutions a coefficient ' /,' which represents the number by which the normal lowering must be multiplied in order to give the value actually found. Thus in the case cited, i 1.83, i.e. 1.89 x 1.83 = 3.46. By means of almost numberless experimental determinations, it gradually developed that this factor 't' varies within somewhat narrow limits for a given salt, and in that case increases with increasing dilution. The factor is never more than two for a * 1 Z.eit.f. Phys. Chem., Vol. 2, p. 638 and p. 323. CHEMICAL PRINCIPLES 17 binary salt, acid, or base, and in ternary or quaternary com- pounds may rise to nearly three or four respectively. It is obvi- ously possible that this abnormal lowering of the freezing point of water by acids, bases, and salts may be due to the formation of increased numbers of molecules by dissociation. If this be the case, then a binary compound, such as common salt, may dissociate into two portions ; a ternary into three, etc. EXPERIMENT V. In a freezing point apparatus, using a Beckman ther- mometer, determine the freezing point of pure water and solutions, in the same sample of water, of \ normal sugar and | and ^ normal salt. Boiling Points As early as 1822 Faraday 1 attempted to determine the in- fluence of dissolved substances upon the boiling point of water. Yet such were the difficulties presented by variation in the results obtained that it was not until 1889 that Beckmann 2 devised a method by which the determination could be made with sufficient accuracy for the determination of molecular weights. The most convenient as well as the most recent apparatus for the purpose is that of Menzies. 3 The influence of nonvolatile dissolved substances upon the vapor pressure of solvents, upon which the boiling point depends, has been shown to be of the same general character as that upon freezing points, viz. that equimolecular concentrations of substances produce the same increase in the boiling point of a given solvent, the molecular elevation of the boiling point being unlike for dif- ferent solvents. Thus a grain molecule of most substances dissolved in 1000 grams of the solvent raises the boiling points of certain solvents as follows : 1 Ann. Chem. Phys. 2, Vol. 20, p. 324. 2 Zeit.f. Phys. Chem. Vol. 4, p. 532. 3 Jour. Am. Chem. Soc., Vol. 32, p. 1615; 1910. !8 QUALITATIVE ANALYSIS Alcohol 1.15 Acetone 1.67 Ether 2.11 Chloroform .... 3.66 Water 0.52 Benzene 2.67 However, if acids, bases, or salts are dissolved in water, the elevation of the boiling point is found to be greater than 0.52 per gram molecule per liter, and the variation from this, which may be called the normal value, increases with increased dilution. Acids, Bases, and Salts in Solution In the discussion of Osmotic Pressure, Freezing Points, and Boiling Points, attention has been specially directed to the abnormal behavior of acids, bases, and salts when dissolved in water. The chemical behavior of these substances in water is equally unique. This peculiar behavior is well brought out by the following experiments : EXPERIMENT VI. Place a piece of zinc in each of two test tubes, and on one pour concentrated sulphuric acid (free from water) and on the other dilute sulphuric acid. EXPERIMENT VII. Grind together dry silver nitrate and dry potassium chromate, and, noticing the absence of result, add water to the mixture. EXPERIMENT VIII. Place in a test tube a mixture of dry copper nitrate and dry ammonium carbonate and heat in a Bunsen flame. Contrast this complex reaction with that which takes place between solutions of each of the same salts when mixed. Examples of a similar kind may be multiplied indefinitely. The conclusion is inevitable, that not only is the chemical behavior of such substances vastly hastened by the presence of water, but that at times the /'planes of cleavage" of the reactions are influenced by the water present. It may be imagined, of course, that this "catalytic" effect of the water is due merely to the more intimate contact of the re- CHEMICAL PRINCIPLES 19 acting substances which is secured by solution. This is not necessarily true, as may be illustrated as follows : EXPERIMENT IX. To a solution of phosphorus in carbon disulphide add a solution of iodine in the same solvent and compare the velocity of the re- action between the phosphorus and iodine with that between a small bit of dry phosphorus when dropped upon powdered iodine. EXPERIMENT X. Compare the rate and character of reaction between dry sugar and concentrated sulphuric acid with that of dilute solutions of the same substances. EXPERIMENT XI. Compare the action of dilute hydrochloric acid and a solution of dry hydrochloric acid in toluene upon marble or zinc. It is evident that in some manner water exerts a potent influ- ence upon the chemical behavior of acids, bases, and salts. This influence may be shown to be such that the portions of the com- pounds entering into reaction are related closely to those which are liberated from solutions of the same substances when sub- jected to the influence of the electric current. Electrolysis When the terminals from a source of electricity are connected by means of various materials, three classes of substances are speedily distinguished : A. Those which act as insulators and prevent the passage of electricity and are known as nonconductors. The vast majority of substances belong to this class. Examples are : nearly all pure liquids, including water ; dry substances, including acids, salts, and bases, except when highly heated ; and the nonmetals. B. Those which admit the passage of the current and yet remain essentially unchanged during the passage of the current and wholly unchanged after the source of current is disconnected. These are for the most part metals, and their mixtures, or alloys, and are known as conductors of the first class. C. Those which, while they allow the current to flow, are 20 QUALITATIVE ANALYSIS simultaneously decomposed by the current or cause decomposi- tion of the terminals and are thus changed in composition. These are known as conductors of the second class and are of the greatest interest in this connection. The substances belong- ing to this class are bases, salts, and acids highly heated or in solution ; chiefly in water solutions. It is not to the present purpose to discuss the character of the changes taking place when electric currents pass through con- ductors of the second class, but only to call attention to certain features of the process ; which is known as electrolysis. In the first place it is to be observed that the components of an electrolytic solution are separately nonconductors. EXPERIMENT XII. Prepare an electrolytic cell with platinum electrodes and a direct current source of approximately 10 volts. Into this cell introduce dry salt and pure water successively, taking care to remove the one completely before the introduction of the other. Then pour into the cell a solution of salt in water. EXPERIMENT XIII. Use sugar in place of the salt, as in Experiment XII, and note that no current passes. EXPERIMENT XIV. Use a solution of dry hydrochloric acid in dry toluene or benzene and note also that no current passes. A second point of importance to be noted is that if we make the experiment in such a way that all of a given solution is kept between the platinum terminals, electrodes, the amount of current which passes increases with increasing dilution. EXPERIMENT XV. Prepare an electrolytic cell consisting of a rectangular glass vessel (2x10x20 cm.) and place platinum terminals the whole length of the ends of the cell. Connect with source of current and ammeter, introduce concentrated hydrochloric acid to a depth of i cm., and note the current pass- ing. Then add distilled water slowly with stirring, and note increased flow. A third point to be noted is that during the process of electrol- ysis, the partition of the dissolved substances is of the same character as that which takes place in metathetical chemical CHEMICAL PRINCIPLES 21 changes. For example, it can easily be shown that when a cur- rent is passed through a solution of silver nitrate, the silver is deposited at the negative electrode and the rest of the molecule (radical NO 3 ) appears at the positive electrode, there to interact with the water, producing nitric acid and oxygen. Similarly with sodium chloride, the parts which appear at the positive and negative poles are sodium and chlorine respectively. These portions of the compounds which thus travel to the electrodes were, by Faraday, called 'ions? The Hypothesis of Arrhenius From the first observation of the distinction between the man- ner of conductivity of solutions and of metals which was made by GrothusMn 1805, various hypotheses designed to account for the process have been advanced. These need not be dis- cussed here, except to call attention to that of Clausius, 2 who assumed that some of the molecules of dissolved substances were ruptured by collision, and that the electric current seized upon these portions to effect its transport across the interven- ing space. In 1887, Arrhenius 3 published the hypothesis which is at present the prevailing one and which correlates the facts concerning acids, bases, and salts in practically the following form. Certain substances (acids, bases, and salts) when dissolved in water (and certain other solvents) are partially decomposed into parts (ions) bearing positive and negative electrical charges of equal amount. This hypothesis, primarily designed to ac- count for the process of electrolysis, is developed from the facts as follows : The so-called abnormal osmotic pressure, freezing point de- pression, and boiling point elevation are always obtained with those substances which are electrolytes and which manifest in- 1 Arrhenius's Electrochemistry, p. 21. ' 2 Pogg. Ann,, Vol. 100, p. 353. 3 Zeit. f. Phys. Chem., Vol. i, p. 631. 22 QUALITATIVE ANALYSIS creased chemical activity when placed in water. The assump- tion of dissociation in solution would account for these facts by reason of the increased number of molecules formed, thus giving rise to a greater number of particles than correspond to molecu- larly equivalent quantities. Also if the free radicals are formed by solution, chemical reaction in solution may be considered as essentially only a realignment of the ions. These phenomena are most marked in aqueous solution, but are by no means confined to it, since the same classes of sub- stances produce similar effects in liquid ammonia, liquid sulphur dioxide, and other solvents, though in less extensive fashion. The assumption of partial dissociation is necessitated by the fact that the abnormal results are ordinarily not so great as would be expected from complete dissociation. For example, the van't Hoff coefficient '/' is less than two for such a salt as NaCl in ordinary dilutions, but increases with dilution in such a way as to indicate that two is its maximum value at great dilutions. The assumption that the parts of the molecules are electrically charged is in accord with the view that a source of current maintains a constant potential of a definite magnitude and that the passage of the current consists in the discharge of the electrically charged ions against these electrodes. That the positive and negative charges must be of equal amount, though the number of ions is not necessarily equal, follows from the fact that solutions of electrolytes are as a whole electrically neutral. The hypothesis as stated seems at first to be at variance with certain fundamental principles. Taking a simple case, when a substance such as hydrochloric acid dissociates, the ions can be nothing other than hydrogen and chlorine. Hydrogen is practi- cally insoluble in water, while chlorine is colored, has a strong odor, and reacts with water, yet the solution of hydrochloric acid in water evolves no hydrogen and is perfectly colorless if pure. CHEMICAL PRINCIPLES 23 It should be noted, however, that the ions are not free elements, even when of elementary composition, since they carry charges of electricity of great amount in comparison with their size, and, moreover, are combined with water, as has been shown by the work of Jones and Washburn previously cited. In similar manner, many more or less apparent difficulties in the matter of concordance of facts with the hypothesis may be met and the general conclusion arrived at that the hypothesis furnishes a very satisfactory explanation of the phenomena presented by electrolytes. NOTE. A very clear and somewhat comprehensive discussion of the behavior of electrolytes during electrolysis and of the chemical behavior of ionic substances is presented in Chaps. XIX and XX of Smith's General Chemistry. At the same time it should be recognized that the hypothesis of Arrhenius is not the last word upon the subject. There are certain facts which are extremely difficult to harmonize with the hypothesis ; as, for example, the rapid reaction between certain compounds dissolved in pyridine, which solutions are not electro- lytes. Also the different methods of determination of the de- gree of ionization, which ought to give reasonably harmonious results, fail to do so. A discussion of this phase of the subject is presented by Kahlenberg. 1 Notwithstanding the unsatisfactory features and the fact that a not inconsiderable number of chemists deny the validity of the hypothesis, the behavior of substances during the processes of qualitative analysis is so clearly and satisfactorily expressed in terms of ions, that it is deemed advisable to use it as a tool until a more comprehensive and more satisfactory view is de- veloped. Physical Equilibrium Such changes as the distribution of solute between immis cible solvents, the solution of gases in liquids, and the deposition 1 Jour, of Phys. Chem., Vol. 5, p. 339. 24 QUALITATIVE ANALYSIS of solids from solutions are ordinarily regarded as physical changes, and these operations play important roles in qualita- tive work. A useful formulation of the conditions upon which equilibrium resulting from these changes depends may be de- duced as follows, using a simple case as a type. EXPERIMENT XIX. Make a saturated solution of iodine in dilute potassium iodide solution and to it add some chloroform. Shake vigorously and allow the chloroform to settle out. The iodine will be found distributed between the two solvents. NOTE. In using this example of physical equilibrium, it is not necessary to lose sight of the chemical nature of the union between potassium iodide and iodine. The union be- tween chloroform and iodine may be of similar though not identical nature. If we represent by 5 the concentration of iodine in potassium iodide solution, it is evident that the rate of passage into the chloroform will depend not only upon this, but also upon the attraction of chloroform for iodine which at a given temperature is constant, and may be represented by K. The velocity of this reaction is then V= KS. The tendency of iodine to go from chloroform to potassium iodide may be formulated simi- larly, V = K' S' \ and since the opposing tendencies are equal when the distribution is complete, V= V f , or S'K 1 = KS, or r- i^ r- = , and since K and K' are constant, = K. S K. S This conclusion may be verbally stated as follows : Physical equilibrium between reacting substances depends upon a con- stant ratio between the concentration of the materials. This statement is of general application, and its use leads to the laws of Dalton and of Henry with regard to the solubility of gases in liquids, and, if we keep in mind the fact that the concentration of solids is their density, to the known facts con- cerning saturated solutions. CHEMICAL PRINCIPLES 2$ Balanced Actions and Chemical Equilibrium Nearly all chemical reactions are reversible, and in analytical operations are more or less incomplete. As a rule, also, in ana- lytical operations it is desired that reactions be made as com- plete as possible. It is therefore essential that the principles of equilibrium be thoroughly understood if the operations are to be intelligently followed. A few examples of reversible reac- tions will be helpful. EXPERIMENT XVI. Into a eudiometer filled with mercury and inverted in a mercury trough introduce about 10 c.c. of dry gaseous ammonia. Pass electric sparks from a Rhumkorf coil through the gas until the volume re- mains constant ; then by means of a pipette introduce a few drops of dilute sulphuric acid, and, after noting the absorption of the undecomposed ammo- nia, continue sparking until no further contraction takes place. EXPERIMENT XVII. To a concentrated solution of magnesium chloride add ammonium hydroxide as long as a precipitate forms. To the mixture now add ammonium chloride until the precipitate is dissolved. EXPERIMENT XVIII. To 400 c.c. of water add 5 c.c. each of decinormal solutions of ammonium sulphocyanate and ferric chloride, and divide into four equal portions. To one add a few cubic centimeters of concentrated ferric chloride ; to another a few cubic centimeters of concentrated ammonium sul- phocyanate, and to the third concentrated ammonium chloride, and compare the colors of these portions with that of the fourth portion. In the cases given we are dealing with balanced actions as expressed by the following : II. MgCl 2 + 2NH 4 OH$:Mg(OH) 2 -r-2NH 4 Cl. III. FeCl 3 + 3 NH 4 CNS ^ Fe(CNS) 3 + 3 NH 4 C1. In I it will be observed that the decomposition ceases before all the ammonia is decomposed, and that when the residual ammonia is removed, the reaction proceeds in the reverse di- rection, and since the ammonia is continuously removed, the reaction proceeds to completion. In II a little investigation 26 QUALITATIVE ANALYSIS would show that precipitation ceases before the reaction reaches completion, and it is observed that the addition of ammonium chloride drives the reaction in the reverse direction with solution of the magnesium hydroxide. In III we may trace the direc- tion of change and the influence of the relative quantities of the reacting materials by the color change due to varying quantities of ferric sulphocyanate. It is evident that in each of the above cases we are dealing with reactions which can proceed in either direction, and that when under given conditions reaction ceases, i.e. equilibrium is reached, it is because the velocities in opposite directions are equal, and that, since the point where equilibrium is reached is changed by varying the relative quantities of the reacting mate- rials, the speed of a reaction is influenced by the relative amounts of the reacting substances. This principle, more or less clearly recognized by Berthollet, Williamson, and others in special cases, was first formulated as a general principle by Guldberg and Waage in I86/, 1 and is variously known as Guld- berg and Waage's Law, the Law of Mass Action, and the Law of Chemical Equilibrium. It may be stated as follows : The rate of a chemical reaction is proportional to the active mass of each of the reacting substances, and equilibrium in a react- ing system is reached when the rates of reaction in opposite senses are equal. A most useful quantitative expression of this law may be arrived at by considering a general reaction expressed by the formula A+B+^C+D, where the components form a reversible reaction and are pres- ent in a homogeneous mixture, i.e. as gases or in solution. At any given constant temperature the speed of reaction of 1 An abstract appears in Jour. Prakt. Chem. 2, Vol. 19, p. 69 (1879) ; also Ostwald's Klassiker, No. 104. CHEMICAL PRINCIPLES 27 A upon B depends upon their nature (i.e. upon the driving force of all reactions, which is ordinarily called affinity, and which is constant) and upon the concentrations. Concentration is ordinarily expressed in terms of moles per liter. If, then, we have one mole each of A and B per liter, the velocity of the reac- tion ( V) at any given instant is V= F, where F is the affinity constant. If, however, the concentration of A is c lt and B is 2 , then V c^ x c 2 x F. Similarly, if we consider C and D, the velocity of their reaction in the opposite direction is expressed by V 1 = r 3 x c x F' y where C B and c represent the concentra- tions of C and D respectively and F' the affinity constant of C and D. When equilibrium ensues, however, V must equal c x c F' V 1 ', and consequently c x c 2 x F l = c z x c x F 1 , or - -*= -~- 3 x c r C X C Since, also, F' and F are constant, we have - - = K, where ^3 X C 4: K is the affinity constant of the reversible reaction. This is the fundamental law of chemical equilibrium. It is evident that in the case of reactions such as we can write as follows : A +A -f B^C -\- C+ D, and conse- quently F= j x c l x c 2 x F and V = c s x ^ 3 x r 4 x F 1 , and the final form c \^ = K. c* x c, NOTE. Fuller discussions of the subject of balanced actions may be found in Smith's General Chemistry, Chap. XV, Walker's Physical Chemistry, Chap. XXII, and Nernst, Chap. I, Book III. Ionic Equilibrium Having presented in an elementary form the fundamental principles of equilibrium, it is now in order to apply them to ionization. Taking a simple case, acetic acid, we formulate the change taking place 28 QUALITATIVE ANALYSIS This is a reversible reaction, since by removal of the water the unionized acid is obtained, and it therefore follows that at a given dilution, there must be present not only ions of hydrogen and of the acetic radical, but also molecules of acetic acid, and these are in equilibrium according to IT It also follows that the more dilute the solution, the greater the dissociation. This is a necessary consequence if the general- ization be true; for if we have at a given dilution 100 each of hydrions and acetions, then 100 x 100 = A where by 'a' concentration of the undissociated molecules is indicated. If now the solution be diluted to 10 volumes, if no change of ionization results, we have in unit volume 10XK W. a 10 In order that equality remain must diminish in value, and 10 this can only occur by further dissociation, giving Q+ io)Q+ io) = A. a x 10 We have in conductivity measurements the most convenient means of determining the degree of ionization at different dilu- tion, and consequently the value of the ionization constant, K. An example of the derivation of the constant and a demonstra- tion that it is a constant is furnished from the conductivity measurement for acetic acid. The conductivity of acetic acid CHEMICAL PRINCIPLES at 25 at infinite dilution, that is, when it is completely dis- sociated, is yux = 364. The conductivity at other dilutions (ft t .) divided by ^ gives the percentage of ionization, and these values substituted in the formula give the values in column four of the subjoined table. V M. PER CENT IONIZATION K 8 4-34 I.I9 O.OOOOlSo 16 6.10 1.673 0.0000179 32 8.65 2. 3 8 0.0000182 64 12.09 3-33 0.0000179 128 16.99 4.68 0.0000179 256 23.82 6.56 O.OOOOlSo In making these substitutions in the formula s necessary to take into account the dilution, e.g. 0.00119 0.00119 o = 0.0000180 ; 0.99881 8 or in general for a binary electrolyte, M* K (i -M)V This is known as OstwalcTs dilution law and is found to hold rigidly for weak acids and bases. The fact that when in this manner the attempt is made to determine the value of K for strong acids and bases and for the highly dissociated salts, no constant value is obtained, has as yet received no adequate explanation. A fuller and a non- mathematical discussion of this phase of the subject is found in Walker's Physical Chemistry, pp. 224-240. 30 QUALITATIVE ANALYSIS EXERCISE The student is asked to calculate, from the data given in the appendix, the values of K for the acids, bases and, salts. This table is to be preserved and used in the subsequent discussion. The Solubility Product A most interesting deduction from the law of physical and chemical equilibrium concerns itself with the situation which ob- tains in a saturated solution of an electrolyte when in the pres- ence of the solid substance. The dissolved material consists of ions and undissociated molecules, according to the formula = K, but c lt the concen- tration of the undissociated substance, is, in the presence of the undissolved substance, the same as the value 5 in the formula o -= K, and since 5', the concentration of the solid, is a constant, o then c 1 is a constant and c 2 x c z = Kc^ or c 2 x c 3 = K. This latter value is known as the solubility product, and the conclusion is that in saturated solutions the product of the con- centrations of the ions is a constant. The importance of this conclusion is manifest when we consider that in saturated solu- tions the value c 1 is unchangeable, yet the total amount of a given substance in solution may be materially affected by addi- tion of a very soluble substance containing a common ion. Ex- tensive application of this principle is encountered in quantitative as well as in qualitative analysis. The student will have already encountered many examples of the sort, and a single experiment will be sufficient illustration. EXPERIMENT XX. Make a saturated solution of barium chloride, and after filtering, add a few cubic centimeters of concentrated hydrochloric acid. Note the precipitation of BaCl 2 and explain in terms of the solubility product. QUALITATIVE ILLUSTRATION The more important applications of the foregoing principles will be discussed in the following paragraphs, and certain addi- tional topics will be introduced as convenient illustrative mate- rial is furnished. In this connection the metal groups will be considered in order and chiefly from the standpoint of the reactions of the metal ions which are used in the group separa- tions recommended. The order in which the groups are con- sidered will be the same as that in which they are taken up in Part II. Group I The metals of this grou are silver, lead, and mercury (ous), and are considered together because their salts in solution in water form insoluble chlorides when treated with solutions con- taining chlorine ions. The soluble salts furnish the positive ions Ag + , Pb ++ , and Hg + ^j The reactions with chloride ions may then be expressed : Ag + + Cl-^AgCl + Cl->HgCl These chlorides are the most insoluble of the more usual salts of mercurous mercury and of silver, and so these metals are practically removed from the solution. Lead chloride is more soluble than lead sulphide, and consequently when the filtrate from the precipitate is treated with hydrogen sulphide, lead sul- phide is formed. Lead, therefore, belongs in both the first and second groups. The separation of the lead chloride depends 31 32 QUALITATIVE ANALYSIS upon its rapid increase in solubility with rise of temperature, while the other chlorides remain essentially insoluble. The sep- aration of silver from mercurous chloride depends upon the presence in ammonium hydroxide of ammonia molecules accord- ing to the equilibrium reaction NH 3 + H 2 O NH 4 OH, and + the formation of the complex ion Ag(NH 3 ) 2 which is soluble in water. We have then the complex reaction NH AgCl -Agei Ag + 01 + NH 8 + H 2 \\ Ag(NH 3 ) 2 Ag(NH 3 ) 2 Cl + Since the formation of the complex ions withdraws Ag ions from solution, AgCl must ionize, and consequently AgCl goes into solution. When to this solution an acid is added, the with- drawal of ammonia reverses the reaction, and silver chloride reappears. Similar complex ions with ammonia are formed with copper, cadmium, cobalt, nickel, amizinc ions. Mercurous ions react with ammonium ions to form an amido derivative NH^Hgj', which probably subsequently decomposes into the mixture + -- - - - NH 2 Hg+Hg. The following equations may be used to ex- press these changes : NH 4 OH + 2 HgCl -> NH 2 Hg 2 OH + 2 HC1 NH 2 Hg 9 OH + HC1 -> NH 2 Hg 2 Cl + H 2 O \ NH 2 HgCl + Hg Group II The special characteristics made useful for the separation of Group II from Group I and Groups III to V inclusive are that while the chlorides are soluble in the presence of a small excess of acid, the sulphides are insoluble in a solution of not too great acidity. QUALITATIVE ILLUSTRATION 33 The ions which may be formed by members of the group are as follows : POSITIVE IONS^ Hg + and Hg ++ Cd ++ Pb ++ and Pb ++++ As +++ and As +++++ Cu + and Cu ++ Sn ++ and Sn +++ + Sb +++ Complex negative ions in which the metals are present : Cu (CN) 4 Sb6 3 : Sb6 4 and SbO 3 Cd (CN) 4 As6 3 : As6 4 and AsO 3 1 1 Sn0 3 Other negative elements and radicals may take the place of the oxygen and (CN) in these ions. This is particularly true of sulphur. The precipitation of these ions is dependent upon the fact that hydrogen sulphide reduces the complex negative ions to the positive forms, and upon the exceedingly small solubility product of the sulphides., (The subject of oxidation and reduction will be discussed in connection with Group III (q.v.)). The sulphides of the group are by no means of equal insolu- bility, and the precipitation is naturally in the order of relative insolubility, if the concentration of the ions of each metal is the same, for the reason that if this be the case with a given concentra- tion of precipitant, the solubility product will be reached soon- est for the sulphide which isjgast sol n hip This_gpnsideration TiTof importance, especially when taken in connection with the influence of an excess of acid, as thft following ^vperiment will indicate. EXPERIMENT XXI. Mix together in a tall cylinder equal volumes of normal solutions of copper sulphate, cadmium sulphate, and zinc sulphate. Add dilute 34 QUALITATIVE ANALYSIS ammonium sulphide, drop by drop, allowing the precipitate to settle as fast as formed. The sulphides will form in well-defined layers in the order of the solutions named. This experiment strikingly demonstrates the order of solubil- ity of these sulphides and emphasizes the need of complete precipitation of each group, since otherwise the more soluble members might easily escape detection in the proper place, and by appearing in a subsequent group cause more or less con- fusion. The degree of acidity of the solution is a matter of impor- tance, since precipitation of the sulphides can only take place when the product of the ions exceeds the solubility product. With a given concentration of cathion, the concentration of the anion may easily be depressed in the case of a weak acid to an extent which amounts to practical elimination, as is clear from the following equation. 1 /z 2 x s =A"=a very small number. (The small letters are v used to indicate the molecular concentration of the correspond- ing ions or molecules.) If to such a substance we add a large concentration of hydro- gen ions, then M J ""^) = ^ anc j jf x j s relatively large, y + "2 s : y becomes very nearly equal to s. Such being the case, it may happen that m xs-CuS + H 2 S. This behavior of hydrosulphides is analogous to the decomposi- tion of hydroxides, but is ordinarily more spontaneous. It is therefore somewhat simpler to formulate the reactions as if divalent sulphur ions entered directly into them. QUALITATIVE ILLUSTRATION 35 EXPERIMENT XXII. Prepare a normal solution of cadmium sulphate and divide into two portions. To one portion add a few cubic centimeters of dilute acid ; to the other some concentrated hydrochloric acid, and pass hy- drogen sulphide into both. In the one case, observe the copious precipitation, and in the other the nonformation of the sulphide. EXPERIMENT XXIII. Prepare a normal solution of ferrous sulphate and divide into two portions:. To one add a few cubic centimeters of dilute hy- drochloric acid and pass hydrogen sulphide into each. Filter the ' neutral ' solution and add a few crystals of sodium acetate and note the precipitation of more ferrous sulphide. (Explain the action of the acetate.) From these experiments, it is clear that the solution for pre- cipitation of Group II must be neither too acid nor too nearly neutral. Experience has shown that a convenient concentration is about ^ normal. A modification of this statement is neces- sary in case arsenates are present. (See discussion of sub- group B.) The separation of sub-groups A and B depends upon the am- photeric nature of arsenic, antimony, and tin. Arsenic trioxide is soluble in both acids and bases, more easily in the latter. It is therefore both acid and basic in character. The sulphides also show the same dual nature, though somewhat less markedly. As is usually the case, however, an increased valency is accom- panied by increased acidic properties, and arsenic pentoxide wholly fails to react with acids, and the pentachloride rapidly hydrolyzes in water. Similarly the pentasulphide is less stable than the trisulphide, but reacts more readily with bases. How- ever, the fact that the sulphides of arsenic fail to dissolve readily in acids is probably due rather to their very great insolubility in water than to their acid character. In solutions of alkaline sulphides the following reactions take place, with the formation of soluble salts : As 2 S 3 + 3 (NH 4 ) 2 S^2(NH,) 3 AsS 3 As 2 S 6 + 3 (NH 4 ) 2 S-*2(NH 4 ) 3 AsS 4 36 QUALITATIVE ANALYSIS These reactions are analogous to the formation of salts by the union of anhydrides and basic oxides as As 2 5 + 3K 2 0->2K 3 As0 4 SO 3 + K 2 O->K 2 SO 4 On account of the stronger acidic character of As 2 S 5 , its solu- tion is more rapid than that of As 2 S 3 . If, therefore, yellow ammonium sulphide, (NH^S*, is used with the trisulphide, the excess sulphur reacts As 2 S 3 + 2 S->As 2 S 5 and the solution is facilitated. This last consideration is of special importance, since Sb 2 S 3 and SnS 2 are almost insoluble in colorless ammonium sulphide (NH 4 ) 2 S, but by the yellow sul- phide are changed to Sb 2 S 5 and SnS 2 and then dissolve to form the soluble salts Sb 2 S 5 + 3 (NH 4 ) 2 S->2(NH 4 ) 3 SbS 4 SnS 2 + (NH 4 ) 2 S->(NH 4 ) 2 SnS 3 This difference in arsenic, antimony, and tin is in line with the general observation that in similar groups, metallic character is more strongly marked as atomic weights increase (As= 75, Sb = 120.2, Sn= 119). Sub-group A The separation of the sulphides which are insoluble in yellow ammonium sulphide offers a number of interesting considera- tions. Dilute nitric acid dissolves all of them except mercury sulphide, because by reason of its great ionization, it reverses the reaction, MX + H 2 S ^ MS + HX and is more effective than other acids because it removes the hydrogen sulphide as formed and so facilitates the reversal : H 2 S + 2 HNO, -> 2 H 9 O + S + 2 NOo QUALITATIVE ILLUSTRATION 37 In case the nitric acid used is sufficiently concentrated, sulphur is oxidized and the sulphates of lead and mercury may be formed, the latter being soluble and the former fairly insoluble in the acid solution. The sulphide of mercury is converted to the chloride by means of aqua regia which, naturally, must be completely re- moved before the characteristic test for mercuric ion is made,- since the aqua regia also converts Sn ++ to Hg+ + + Sn+ + -+ Hg + or 2 Hg ++ + Sn" + -> Sn ++++ + 2 Hg ++ The filtrate from the sulphide of mercury is treated with sul- phuric acid to convert lead nitrate to the insoluble sulphate, and the removal of nitric acid by evaporation is essential, if all the lead is to be so converted, since the reaction is reversible. Pb(N0 3 )2 + H 2 S0 4 -+ PbS0 4 + 2 The fact that no lead sulphate precipitate is formed until the concentrated acid solution is diluted is due to the primary ion- ization of the acid, giving rise to the soluble acid salt, which in turn is decomposed by dilution. H 2 S0 4 ->H + HS0 4 Pb ++ + 2 HSO 4 ^ Pb(HSO 4 )2 Pb(HS0 4 ) 2 + H 2 -> PbS0 4 + 4 H + 2 S0 4 The filtrate from the lead sulphate is rendered alkaline by means of ammonium hydroxide and the bismuth ion so con- verted into bismuth hydroxide. The mere formation of a floc- culent precipitate at this point is not necessarily indicative of bismuth because of the presence of aluminium salts and of solu- ble silicates in ordinary laboratory reagents, since both alumin- ium hydroxide and silicic acid are precipitated at this point. NOTE. The student is asked to explain why ammonium silicate, soluble in water, is not formed. ( Vide infra ' Hydrolysis ' and Smith's General Chemistry, p. 523.) 38 QUALITATIVE ANALYSIS The confirmation of the presence of bismuth is due to the hydrolysis (vide infra) of bismuth chloride and the insolubility of the bismuth oxychloride. BiCl 3 + 2 HOH Bi(OH)2Cl + 2 HC1 This reaction is easily reversible, and the presence of an excess of hydrochloric acid is prevented by evaporation or by addition of sodium acetate. NOTE. The student is asked to explain how the latter produces this result. The fol- lowing experiment beautifully illustrates the reversibility of the reaction. EXPERIMENT XXIV. In a small beaker dissolve some Bids in the small- est amount of concentrated HC1 which will effect solution. Then reverse the reaction by addition of water and then successively by HC1 and water as long as desired. The ammonium hydroxide fails to precipitate either copper or cadmium hydroxides. That this failure to precipitate is not due to the amphoteric nature of these hydroxides is shown by the fact that they are not dissolved by the alkali hydroxides. The + + + + complex ions Cu(NH 3 ) 4 and Cd(NH 3 ) 4 form hydroxides solu- ble in water, which are in equilibrium, thus : Cd ++ + 4 NH 3 + 2 OH Cd(NH 3 ) 4 (OH) 2 Cu ++ + 4 NH 3 + 2 OH ^ Cu(NH 3 ) 4 (OH) 2 . If to this system we add potassium cyanide, a curiously com- plex action takes place, which may be formulated as follows : Cd( + NH 3 ) 4 ^4NH 3 -h Cd ++ + 4 CN ^ Cd (CN) 4 +2 K + Cu"(NH 3 ) 4 ^ 4 NH 3 + Cu ++ + 4 CN ^ Cu (CN) 4 + 2 K + . It is plain that the dissociation of the complex negative ions is much smaller than that of the complex positive ions, since the point of rest is where nearly complete transformation in the di- rection (->) ensues, as is shown for the copper ions at least by the removal of the blue color. That this is true may also be QUALITATIVE ILLUSTRATION 39 shown by passing hydrogen sulphide into the blue solution, when both copper and cadmium sulphides are precipitated. If now we pass hydrogen sulphide into the colorless solution, we find a concentration of Cd ++ ions sufficient to exceed the solu- bility product for cadmium sulphide, i.e. CdxS>CdS. The cadmium sulphide, instead of copper sulphide, is precipitated in spite of its greater solubility product, as shown in Experiment XXI, p. 33. It follows, then, that Cd(CN) 4 is dissociated to a 1 1 greater extent than Cu(CN) 4 . Sub-group B The reprecipitation of the sulphides of arsenic, antimony, and tin is effected by dilute acids for the reason that the free acids H 3 AsS 4 , H 3 SbS 4 , and H 2 SnS 3 are unstable acids of the charac- ter of carbonic acid, and break down into hydrogen sulphide and As 2 S 5 , Sb 2 S 5 , and SnS 2 . The sulphides of antimony and arsenic immediately decompose, forming the trisulphides. Con- centrated hydrochloric acid dissolves the sulphides of tin and antimony, but fails to dissolve the arsenic sulphides. The arsenic sulphides, when oxidized by powerful oxidizing agents, yield the arsenic oxide which forms the strong acid H 3 AsO 4 precipitated by magnesium mixture as magnesium ammonium arsenate. The separation of antimony and tin by the precipitation with zinc and platinum is due to their marked difference of potential. NOTE. The alternative method of detection for tin and antimony, given on page 78, is to be preferred for large classes, because of the high cost of platinum, and the effort is made to supply a satisfactory method of procedure which will make the use of platinum appara- tus unnecessary. It is worthy of note that if an arsenate is present in the orig- inal mixture subjected to analysis, the negative sulphur ion will not react with the negative arsenate ion, and in order that pre^ cipitation of the sulphide may take place, other reactions must precede. These reactions may be expressed as follows : 40 Q U A LIT A TIVE ANAL YSIS Na 8 AsO 4 + 3 HCl^H 8 AsO 4 + 3 NaCl As 2 O 5 + H 2 S ^ As 2 O 3 + H 2 O + S As 2 O 3 4- 6 HC1 ^ 2 AsCl 3 + 3 H 2 O 2 AsCl 3 + 3 H 2 S ^t As 2 S 3 + 6 HC1 These reactions, with the exception of the last, being easily reversible, it is apparent that they will be facilitated by in- creased concentration of hydrochloric acid, and consequently strong acidification is desirable, if arsenates are suspected. The strong acidity of the solution is diminished by evaporation or large dilution before the precipitation is completed, for the reasons given on page 34. Group III The members of this group are characterized by the fact that when in solution in such forms that the metals furnish positive ions, they are all precipitated by ammonium sulphide and not by H 2 S in acid solution, while the members of Groups IV and V are not so precipitated. These elements form a complicated series of cations and anions not all of which are likely to be encountered in ordinary qualitative operations, but which may be tabulated roughly as follows : CATIONS ANIONS Co ++ Co +++ Co(CN) 6 Co(CN) 6 ' Ni+ + Ni+++ Ni(CN) 6 Ni(CN) 6 Zn ++ Zri6 2 Mn ++ Mn +++ Mn ++++ Mn 4 Mn 4 MnO 4 Al +++ A1O 3 A1O 2 Fe+ + Fe+ ++ FeO 3 FeO 2 FeO 4 Cr ++ Cr +++ Crb's CrO 2 CrO 4 Cr 2 O 7 Cr 2 O 8 (?) QUALITATIVE ILLUSTRATION 41 The failure of the positive ions to precipitate with hydrogen sulphide is due in part to the relatively large solubility product of the sulphides and in part to hydrolysis. For example, the following reaction takes place as indicated though the precipita- tion is not complete : Zn(C 2 H 3 2 ) 2 + H 2 S-*ZnS + 2 HC 2 H 3 O 2 I The reaction may be driven farther toward completion by the addition of sodium acetate. ' NOTE. The student is asked to compare this effect of the acetate with that noted on p. 35 (Experiment XXIII). Explain and formulate the explanation by a series of equa- tions. (Cf. p. 53, on the solubility of Mg(OH) 2 in NH 4 C1.) In the presence of hydrochloric acid the relatively large con- centration of hydrogen ions prevents any precipitation of zinc sulphide by reason of the suppression of the sulphur ions of hydrogen sulphide. If ammonium sulphide is used, we have a very much larger concentration of sulphur ions, since salts are in general highly ionized, and consequently the solubility prod- uct of zinc sulphide is exceeded, even when very minute quan- tities of zinc remain in solution, for since Zn ++ x S = K, when 1 1 the concentration of S is small, as in H 2 S solutions, then (Zn^+X^S) = K, when Zn is small, if X is very large. The same considerations apply to the other members of the group with modification for ions of chromium, aluminium, and ferric iron. Ferric ions as well as the negative oxygen complexes of the other metals do not exist in the solution for analysis of this group, if previously tested for the second group, because of the reducing properties of hydrogen sulphide; and if the iron has been oxidized by means of nitric acid, it is again reduced by ammonium sulphide and is always precipitated as the sulphide, thus: 2 FeCl 3 + (NH 4 )2S->2 FeCl 2 + 2 NH 4 C1 + S FeQ 2 + (NH^S-^FeS + 2 NH 4 C1 42 QUALITATIVE ANALYSIS Chromium and aluminium are not so reduced, yet are not pre- cipitated as trivalent sulphides, but are converted to hydroxides, thus: 2 AlClg + 3 (NH 4 ) 2 S -> A1 2 S 3 -f 6 NH 4 C1 -f 6 H 2 O - 3 H 2 S + 2 A1(OH 3 ) Hydrolysis The hydrolytic effect of water illustrated above plays a most important role in qualitative analysis and has already been en- countered in the behavior of bismuth and antimony chlorides (cf. reactions 52, p. 67 and 118, p. 75). The far-reaching effect of hydrolysis is well illustrated by the following experiment, which will also serve as a basis for the discussion of the subject. EXPERIMENT XXV. Make solutions of each of the following substances, NaCl, Na 2 SO 4 , CuSO 4 , Na 2 CO 3 , NaHCO 3 , FeCl 3 . Test each solution with red and blue litmus paper. It will be seen that the behavior of 'neutral' salts toward litmus varies widely with composition. The explanation may be had from a consideration of the law of mass action and the ionization ' constants ' of the bases and acids (cf. p. 30). The case of ferric chloride may be taken as a type. Since it is a salt, the ionization is large. Water is slightly ionized, therefore, in addition to the two equilibrium reactions, the equilibrium reactions for ferric hydroxide and hydrochloric acid must also be produced. Since the ionization constant for HC1 is very large and for Fe(OH) 3 small, then the original balance of hydrogen and hy- droxyl ions is disturbed, and an excess of hydrogen ions results. It follows that the solution is acid toward litmus. It also fol- QUALITATIVE ILLUSTRATION 43 i lows that since OH ions are withdrawn from solution to form un-ionized Fe(OH) 3 , that water must continue to ionize. Eventually, however, in spite of the withdrawal of OH ions we have '* + l ) \ l ~y) = g^ anc j t h e wa t e r ceases to ionize. // 2 This result is apparently reached before the solubility product Fe +++ x (OH) 3 = K) is exceeded, since no ferric hydroxide precipitates. NOTE. The student is asked to explain, in a similar manner, the behavior of the other solutions. In case of salts where the operation of hydrolysis produces either insoluble or slightly ionized substances from both com- ponents, the withdrawal of both hydroxyl and hydrogen ions permits the continuation of the ionization of water to a greater degree and frequently even to complete decomposition of the salt. EXPERIMENT XXVI. To a fairly concentrated solution of aluminium sulphate, add a strong solution of sodium carbonate, and note the formation of the precipitate and its change of character as the evolution of carbon dioxide proceeds. It is evident visually that the reaction is at first a simple metath- esis followed by the hydrolysis of the carbonate, this pair of reactions being analogous to those of aluminium sulphide, on p. 42. A little consideration will now make clear to the student that the reaction of salts in solution toward litmus is a function of the relative strengths (or degree of ionization) of the acids and bases from which they may be considered to have been formed. We may then distinguish four classes of salts according as they are formed from the neutralization of : 1. Strong acid and strong base neutral in solution. 2. Strong acid and weak base acid in solution. 3. Weak acid and strong base alkaline in solution. 4. Weak acid and weak base decomposed in solution. 44 QUALITATIVE ANALYSIS It will also be apparent that a more delicate indicator would serve to detect a difference of relative strength not shown by the relatively unsensitive litmus. Since aluminium, chromium, and ferric iron are all very weak base-forming elements, we find their salts with the weaker acids are all hydrolized by water. This consideration then explains the precipitation of the hydroxides of the elements by sulphides, carbonates, and acetates. With the latter salts we find less rapid hydrolysis, as would be expected, considering the strength of acetic acid as compared with carbonic acid ; and in the case of iron acetate the hydrolysis is not complete, the basic acetate being formed probably because of its great insolubility rather than because of the stability of the acetate. These considera- tions not only account for the precipitation of the group, but also for the separation of aluminium, iron, and chromium from manganese and zinc. Cobalt and Nickel Reactions. The failure of cobalt and nickel sulphides to dissolve in dilute hydrochloric acid is remark- able in view of the fact that the reaction CoCl 2 + H 2 S^CoS + 2 HC1 reaches equilibrium without precipitation of the sulphide, yet if the sulphide is formed, then a much greater concentration of acid than is sufficient to prevent its formation is required for solution. The explanation is probably to be found in the poly- merization of the sulphides after precipitation forming (CoS)^ and (NiS) x which are more insoluble than the simple molecules. The doubt in this case might be resolved, had we a method of determining the molecular weight of such solids. The remarkable similarity of properties of cobalt and nickel, which is shown, not alone in the metallic state, but also in their ionic reactions, is responsible for the somewhat involved reactions necessary for their separation. The peculiar relation of these QUALITATIVE ILLUSTRATION 45 elements, the atomic weights of which are so nearly equal, Ni 58.68, Co 58.97, is nowhere more strikingly shown than in the following experiment : EXPERIMENT XXVII. Make solutions of cobalt chloride and nickel chloride and place a portion of each to one side for later comparison. The bulk of each solution is treated with a small excess of a saturated solution of sodium bicarbonate and the precipitate allowed to settle. The supernatant liquid is poured off, and both precipitates are cooled with ice and salt and treated with a small excess of hydrogen peroxide, which is also cooled in ice and salt. The two solutions are now filtered quickly and compared in color with the original pink cobalt and green nickel solutions. The colors will be found to be interchanged, the cobalt solution is green, the nickel pink. NOTE. The nickel solution should be filtered through an ice-cold filter, but even so the pink color is very evanescent, the green color returning with simultaneous evolution of oxygen. Because of the close similarity of the ordinary salts of these elements, recourse is had to the slightly greater differences shown by them when in the trivalent form and as portions of complex negative ions. The series of reactions is identical for both elements, thus : CoCL, + 2 KCN -> Co(CN)2 + 2 KC1 Co(CN)2 + 4KCN -> K 4 Co(CN) 6 K 4 Co(CN) 6 + Br->KBr + K 3 Co(CN) 6 There is, however, a distinction between the double cobalti- cyanide and the nickel cyanide. Both are in equilibrium sys- tems as follows : Co + ^ + 3 CN ^ Co(CN) 3 + 3 KCN ^ K 3 Co(CN) 6 ^3K + + Co'(CN) 6 3 CN ^ Ni(CN) 3 + 3 KCN ^ K 3 Ni(CN) 6 46 QUALITATIVE ANALYSIS If to the systems an excess of sodium hydroxide is now added, the product of the concentration of the Ni +++ ions and the OH ions exceeds the solubility product for Ni(OH) 3 , but not for Co(OH) 3 . The inference is, of course, that the double cobalt salt is the more stable compound. A similar slight difference of the double nitrite ions is found sufficient for separation since K 3 Co(NO 2 ) 6 is insoluble in acetic acid, while the K 3 Ni(NO 2 ) 6 is soluble. NOTE. For the methods of distinguishing cobalt and nickel by means of the reaction of Tschugaeff, see Ber. d. Chern., 38 : 2520. Manganese and Zinc Separation. The separation of zinc from manganese depends upon the greater solubility product of the manganese sulphide, so that while in acetic acid solution it is easily possible to obtain Zn x S > KK, ZnS but Mn x S < K^ m MnS . Manganese is by powerful oxidation converted to the highly colored MnO 4 ion. The principles of oxidation will be dis- cussed in connection with chromium. Separation of Iron, Aluminium, and Chromium. As was ob- served on p. 35 in connection with the arsenic group, increase of valency is in general accompanied by increased tendency to form negative complex ions. This has shown itself in the for- mation of the negative complexes of trivalent cobalt and nickel and also in the negative character of the MnO 4 ion, where the manganese is heptavalent as compared with the strongly metallic character of Mn ++ . Also on p. 36 attention was directed to decreasing acidic tendencies in similar groups with increasing atomic weights (Al = 2/; 0=52; Fe=56). Both these characteristics are manifested in the group under consideration. QUALITATIVE ILLUSTRATION 47 Aluminium hydroxide is a very striking example of an am- photeric substance, i.e. one which is at the same time both acidic and basic. This is shown by the following reactions. A1(OH 8 ) 4- 3 HCl- AlClg + 3 H 2 Q A1(OH) 3 + 3 NaOH -+ Na 3 AlO 3 + 3 H 2 O Written ionically these become A1+++ + 3C1->A1C1 3 A16 3 + 3Na->Na 3 AlO 3 . Aluminium hydroxide in solution must therefore satisfy two con- stants, viz. : al(ok\ where both K and K' are extremely small, i.e. it is both a weak acid and weak base, and by consequence any tendency to form salt by autoneutralization is offset by hydrolysis. Moreover, the solubilit^ product is also small so that but minute quantities are in solution. The system may perhaps best be illustrated by the following equilibrium reaction : 3 H + 4- Al'6 3 ^ A1(OH) 3 ^A1 +++ + 3 OH It A1(OH) 3 (solid) If now to this system we add hydroxyl ions by means of a strong - base, there results the withdrawal of hydrogen ions to form water and the highly ionized aluminate, and the whole system moves in the direction indicated by (<). If, on the other hand, we add a strong acid, hydroxyl ions are removed and the system moves in the direction (>). Ammonium hy- droxide is too weakly basic to affect the system materially, and they same statement is true for the very weak acids. The fact that acetic acid (approximately ionized to the same extent as ammonium hydroxide) does dissolve the hydroxide indicates 48 QUALITATIVE ANALYSIS that more hydroxyl ions are present than hydrogen ions, i.e. A1(OH) 3 is more strongly basic than acidic. Chromium hydroxide acts in a similar manner, but is more weakly acid than aluminium hydroxide, and therefore, while it does dissolve to some extent in sodium hydroxide in the cold, its salt is completely hydrolyzed by boiling, and the hydroxide is reprecipitated. Ferric hydroxide shows so little acid char- acter as to be almost wholly unaffected by alkali solvents. Hence, the separation of iron and chromium from aluminium hydroxide and the identification of the latter. NOTE. Both iron and chromium hydroxides on fusion with bases form salts of the corresponding meta acids even as aluminium hydroxide does. This class of compounds is abundant in nature and is known as the spinels, examples of which are Fe(CrO 2 ) 2 chromite Zn(FeO 2 ) 2 franklinite Mg(AlO) 2 spinel Fe(FeO 2 ) 2 magnetite Iron and chromium both form distinct acids in a higher state of oxidation. Here also the difference of degree manifests itself clearly. Fusion with sodium carbonate and potassium nitrate converts the chromium hydroxide completely to a chro- mate, and the iron hydroxide remains unaffected except by con- version to the oxide Fe 2 O 3 by dehydration. In the presence of hydrogen peroxide, the chromium hydroxide may be converted to the chromate at ordinary temperature in solution. However, fusion with sodium peroxide converts the iron hydroxide into unstable ferrates, which are soluble in water but decompose on boiling. In slightly acid solutions chromates may be oxidized to a still greater extent, forming the free, but very unstable, acid H 2 Cr 2 O 8 , the anhydride of which, Cr 2 O 7 , is soluble, and somewhat more stable, in ether. Oxidation The r61e played by oxidation in the separation of the metals is very important, and a brief discussion of the general subject is in order. By derivation the meaning of the term is limited to QUALITATIVE ILLUSTRATION 49 such actions as involve the union of oxygen with some other substance. Examples are, of course, familiar. In the character of these changes there is, however, no essential difference from those in which union with elements other than oxygen is effected. The use of the term has thus become extended to cover those cases where an element changes its valency in a positive direc- tion. Thus iron is oxidized by chlorine to form ferrous chloride, and this in turn is further oxidized to ferric chloride, though no oxygen is involved in the operation. Considered from the ionic viewpoint any change in the electrical charge upon the ions which increases the positive charge or diminishes the negative charge is an oxidation. For example, the ferrocyanide ion is oxidized to the ferricyanide by loss of one negative charge (or electron). (See Walker's Physical Chemistry, p. 307.) While it is helpful to look upon change of ionic charge as the essential factor in oxidation and reduction of ions, and while it is possible also to interpret the formation and decomposition of complex ions in these terms, it is somewhat simpler to view these, at times somewhat complex, changes only from the standpoint of change in valency. In the fusion reaction using potassium nitrate, there is direct addition of oxygen in the case of both manganese and chro- mium. It is not possible to obtain the products of oxidation in either case in the absence of a basic material, for the reason that the free oxides are unstable at the temperatures required for their formation. We may formulate the changes as follows : 2 Cr(OH) 3 -> H 2 O+ Cr 2 O 3 + 3 O ^ 2 CrO CrO 3 + Na^Og ^ Na 2 CrO 4 + CO 2 The sodium carbonate also plays the part of a flux, or liquid solvent, rendering contact of the interacting materials more inti- 50 QUALITATIVE ANALYSIS mate and thus facilitating the reaction. All these changes may be thrown into ionic form but are not illuminated greatly by the operation. Bead Tests Some of the metal ions of the third as well as of the second group are capable of detection by the so-called bead reactions, the indications of which are tabulated on p. 124. The bead reac- tions are essentially dependent upon the conversion of the salts of meta acids into corresponding ortho salts. The simplest example is furnished by the microcosmic salt reactions, as illus- trated by the following equations : NaNH 4 HPO 4 -> NaH 2 PO 4 NaH 2 PO 4 ->NaPO 3 + I If we add to the bead, NaPO 3 , a non-volatile oxide, in small quantity, the oxide will be practically wholly dissolved, thus : NaPO 3 + M 2 O -^ NaM 2 PO 4 If the metallic oxide is bivalent or trivalent, the character of the reaction is unchanged, though the formulas of the com- pounds may be more complex. If borax is used to form the bead, there is no essential variation, and the borax glass may be considered as a mixture of the metaborate and the boric anhydride. Na 2 B 4 O 7 - 10 H 2 O -+ Na 2 B 4 O 7 + 10 H 2 cf Na2B 4 O 7 ->2 NaBO 2 4- B 2 O 3 NaBO 2 + CuO -> NaCuBO 3 All metallic oxides undergo these reactions, but not all furnish characteristically colored salts. The different valences of the metals also vary in the color of salt formed. Non-metallic oxides are of course incapable of playing the part indicated QUALITATIVE ILLUSTRATION 51 above, but with a basic bead, as Na^Og, they react in an anal- ogous manner, thus : Na 2 CO 3 + SiO 2 -> Na 2 SiO 3 4- CO 2 . Group IV The members of Group IV are characterized by being in- capable of precipitation from solution by HC1, H 2 S, or (NH^S and by formation of carbonates insoluble in ammonium salts. Only bivalent cations are encountered. The solution of the group precipitate is best effected by weak acetic acid, so that, on treatment with chromate, the reversal of the reaction is less vigorous than with a stronger acid. 2 Ba(C 2 H 3 O 2 )2 + K 2 Cr 2 O 7 4- H 2 O -> 2 BaCrO 4 + 2 KC 2 H 3 O 2 \ + 2 HC 2 H 3 2 After removal of the surplus dichromate by reprecipitation and resolution in acetic acid, the test for the presence of stron- tium ions is made on a portion of the solution by addition of a saturated solution of calcium sulphate. Since calcium sulphate is soluble only to the extent of 2 grams per liter (see Table of Solubilities, p. 174), it is evident that the precipitate of strontium sulphate can only be very slight, regardless of the amount of strontium present. This precipitate may not appear at once because of supersaturation (q.v., p. 55). The portion not used for the strontium test is treated with sodium sulphate to remove the strontium ions, and since calcium sulphate is so sparingly soluble, all but two grams per liter will precipitate with the strontium. Consequently, on treatment with ammonium oxalate, only a very slight precipitate of calcium oxalate is to be expected. The most interesting feature of this group, which also fur- nishes an interesting study of relative solubilities, is presented by the filtrate from Group IV. From the Table of Solubilities, it will be seen that the following quantities of carbonates remain 52 QUALITATIVE ANALYSIS in solution, if precipitation of each has taken place : SrCO 3 , ii mg. per liter; CaCO 3 , 13 mg. per liter; BaCO 3 , 23 mg. per liter. Small as are these quantities, a very satisfactory separa- tion and identification is possible in this filtrate. Treatment with dilute sulphuric acid should give a precipitate of barium sulphate (solubility 2.3 mg. per liter), and when this is filtered, the filtrate with ammonium hydroxide and ammonium oxalate should give a precipitate of calcium oxalate (solubility 5.6 mg. per liter). No reagent will give a precipitate of strontium of a solubility corresponding to less than 1 1 mg. per liter, but evap- oration to dryness and expulsion of the ammonium salts will furnish a residue suitable for a flame test. Flame Tests Since each of the ions of this group furnishes a characteristic flame coloration, it should be noted here that only salts which are volatile at the temperature of the Bunsen flame can be used for this purpose. It is therefore necessary to convert sulphates, chromates, etc., into the more volatile chlorides by moistening with hydrochloric acid. The volatile chlorides then dissociate in the flame, and the ions produce color waves which are the re- sultants of their spectral lines. Strontium chloride gives a somewhat evanescent flame because of its rapid volatilization. Group V The members of Group V form what is called the soluble group because relatively few of their salts are insoluble in water. With the exception of magnesium, they form univalent cations in aqueous solution. Ammonium. Solutions of ammonium hydroxide furnish a variety of reactions which depend not only upon the formation of the NH 4 ion by electrolytic dissociation, but also upon the dissociation H 2 (i) QUALITATIVE ILLUSTRATION 53 This latter dissociation manifests itself, at somewhat elevated temperatures, with ammonium salts. :2 NH 3 + H 2 S0 4 It is in consequence of the volatility of one or both of the prod- ucts of this dissociation that the most convenient test for am- monium is to heat the suspected substance with a base and when X the reaction NH 4 X + MOH -> MX + NH 4 OH -> NH 3 + H 2 O has taken place to reverse reaction (i) by absorption of the volatile ammonia in moist litmus paper. The most delicate test for the ammonium ion does not depend directly upon its ionic nature. Nessler's reagent reacts with ammonium by substitution of mercury for hydrogen. The reaction may be formulated : NH 4 OH + 2 K 2 HgI 4 + 3 KOH -> NHg 2 I + 7 KI + 4 H 2 O Sodium. Sodium ions form no very insoluble salts and there- fore dependence is usually placed upon the somewhat embar- rassingly delicate flame test. Attention is directed, however, to reactions 260 and 261 on p. 96, particularly the latter. Potassium. Potassium ions form a few insoluble salts, the one most frequently used for identification being the chlorplatinate. From the standpoint of economy, the precipitate formed by sodium cobaltinitrite is better, without material difference in efficiency. Magnesium : The presence of magnesium in this group is somewhat anomalous when one considers its place in the periodic system and the solubility of the carbonate (i g. per liter) and hydroxide (lomg. per liter). It escapes precipitation in both the third and fourth groups by reason of the presence of ammo- nium salts. The reaction between magnesium salts and ammo- nium hydroxide is easily reversible and the following experiment is very instructive. 54 QUALITATIVE ANALYSIS EXPERIMENT XXIX. Make a concentrated solution of MgCl 2 and treat with concentrated ammonium hydroxide and filter. To the precipitate, add a solution of ammonium chloride and shake. The hydroxide should dissolve. To the filtrate, add sodium phosphate ; the magnesium not precipitated by the hydroxide is shown to be considerable. To some of the original solution add ammonium chloride and then ammonium hydroxide ; no precipitation follows. This rather remarkable reversion is usually explained as fol- lows : The precipitation of Mg(OH) 2 depends upon the prod- uct of the ions being in excess of the solubility product, i.e. mg x olfi > A" Mg ( H) 2 - Ammonium hydroxide is a weak base vide infra, and being in equilibrium, ^ = K, the excess ;z^ 4 oh of ammonium ions from the highly dissociated chloride depresses the ionization to the practical elimination of hydroxyl ions. The equilibrium reaction (somewhat awkwardly expressed) Mg ++ + 2d NH 4 + OH Mg2(OH) NH 4 + C1 lit + H 5: H + It MgCl 2 2NH 4 OH Mg(OH) 2 2 NH 4 C1 A| Mg(OH) 2 is therefore driven in the direction f < t ) by the addition of am- -{) by tl monium ions and their presence before treatment with ammo- nium hydroxide prevents the movement in the direction ( >) to a sufficient extent to form Mg(OH) 2 in excess of its solubility. A similar explanation accounts for the failure to precipitate Zn(OH) 2 and Mn(OH) 2 when ammonium hydroxide and am- monium chloride are added to the third group solution in prep- aration for treatment with ammonium sulphide. (For failure to precipitate cobalt and nickel salts, see p. 32. Consult also rules of solubility, p. 101.) The fact that ammonium carbonate does not precipitate mag- nesium salts in the presence of ammonium salts might be ascribed QUALITATIVE ILLUSTRATION 55 to a similar elimination of CO 3 ions by suppression, were it not that an excess of a common ion exercises a relatively less marked effect on a highly ionized substance than upon one slightly ion- ized. It were scarcely possible to wholly suppress the ionization of ammonium carbonate by means of ammonium chloride. It is perhaps preferable to consider that magnesium ions form with the chloride the double salt NH 4 MgCl 3 (analogous to carnellite KMgClg), which is not decomposed in the presence of the CO 3 ion. The ammonium chloride also plays a role in the precipitation of magnesium by soluble phosphates in ammoniacal solution. The reaction may be expressed thus : [MgCl 2 + NH 4 OH + Na 2 HP0 4 -*MgNH 4 P0 4 + 2NaCl+ H 2 O. . * The presence of the ammonium chloride prevents precipitation of any magnesium hydroxide by weakening the basic reaction of the hydroxide. The slow formation of the phosphate precipi- tate is due to supersaturation. Supersaturation Supersaturation of solutions is frequently encountered when the substance is formed in solution, or when substances more soluble in warm than in cold solvents are dissolved to saturation at a given temperature and then brought to a lower tem- perature. The following experiment beautifully illustrates the phenomenon and also the property of independent solubility . and the effect of the presence of the solid substance. EXPERIMENT XXX. Partly fill a test tube with crystals of sodium thiosul- phate and heat until dissolved completely in the water of crystallization. Pour on top of this a layer of a solution of sodium acetate saturated at 50 C. and allow to cool. A crystal of sodium thiosulphate may now be dropped through i the acetate into the lower layer which will immediately crystallize. A crystal of sodium acetate may now be dropped into the upper layer, which will im- 56 QUALITATIVE ANALYSIS mediately crystallize. There will remain a middle portion which is not solidified by reason of the mutual dilution of the solutions by each other. The experiment illustrates the most effective means of caus- ing precipitation in supersaturated solutions. It is apparent ^ that from the equilibrium formula, = K, if S' = o, then 5 may ijj be of any magnitude, so far as the law is concerned. If, how- ever, a portion of solid substance is introduced, the material in solution must alter in quantity until the constant ratio is reached. The same result is obtained by violent agitation or by scratching the inner glass surface with a rod. The principle involved is that violent agitation initiates the formation of a minute quantity of solid, after which the operation continues according to the law of physical equilibrium. Ordinarily standing for a considerable time also produces the discharge of supersaturation. PART II METAL ANALYSIS General Directions. EQUATIONS. Every chemical reaction is capable of more or less accurate expression in the form of an equation. This consists of a representation, by means of atomic symbols, of the component parts of a reaction and of the nature and direction of the change. An equation therefore indicates the quantitative relations of the substances in so far as reaction occurs. An equation never conveys any indication of the condition necessary to produce the indicated change nor the completeness of its accomplishment. Neither does an equation indicate whether or not other reactions may be of simultaneous occurrence, nor whether the change may be immediately followed by subsequent change of the products of the reaction. Such an expression is therefore far from a complete exposition of a chemical change. Yet equations set before us in such concise form the character of the relations with which we deal that facility in their use is indispensable in a study of qualitative analysis. The student is expected, therefore, to write equations representing the chemical changes involved in the preliminary reactions. A study in equa- tion making is found in Walker's Physical Chemistry^ pp. 22-27. PRECIPITATION AND FILTRATION. The part played by precipitation in qualitative analysis has been dis- cussed in the introduction. The importance of complete pre- cipitation is mentioned on p. 34. It remains to point out the 57 58 QUALITATIVE ANALYSIS method of preparing precipitated substances for examination. The precipitant is to be added slowly, the substances thrown out of solution allowed to settle, where possible, and more precipitant added until no further change takes place. After filtering, a small amount of precipitant should always be added to the filtrate as a precautionajg^measure. Wherever possible the filtration should be from hoi solutions, both because the precipitates coagulate more readily when hot and settle out more completely, and because hot liquids pass through filters more rapidly than cold. The filtering medium is usually paper of a bibulous charac- ter, which is cut into circular form and folded to form a quadrant. It is opened between the third and fourth layers and fitted into a funnel with the help of water from a wash bottle. If the paper be fitted by its upper edge into the funnel, but not closely along the sides, the rate of filtration is materially increased. The size of the funnel should be such that the paper covers about two thirds of the side. The liquid to be filtered is poured into the paper along a glass rod held vertically, and the pouring should be at such a rate that the paper is kept continually about half full. A bit of soft rubber tubing slipped over the end of the glass rod furnishes a convenient instrument, a 'policeman,' for transferring the last portions of the precipitate into the filter. WASHING. After all the precipitate is thus transferred to the paper, it must be washed in order to free it from adhering mother liquor and to bring into the filtrate all material which prop- erly belongs in it. Wherever possible, washing should be done with hot water, for which a wash bottle with the neck wrapped in cork or asbestos is a great convenience. In all cases washing should be accomplished by means of small amounts of water added repeatedly rather than by immediate application of a large quantity. A very interesting discussion of this matter is found in Ostwald's Foundations of Analytical Chemistry, pp. 1 1-22. METAL ANALYSIS 59 DECANT ATION. Where precipitates are either gelatinous or heavy, time is frequently saved by allowing the predicate to settle and transferring the supernatant liquid to the filter without disturbing the precipitate. The latter may then be washed in the vessel in which it was formed and th^jrah water again de- canted through the filter. In this manner clogging of the filter and consequent slow filtration may be avoided, and at times it may not be necessary to transfer the precipitate to the paper. EVAPORATION OF FILTRATES. Where filtrates are to be concentrated or evaporated to dryness, it is not usually advisable to use a free flame, owing to the danger of loss by ebullition, spattering, or decomposition when dry. The most satisfactory apparatus is a laboratory water bath, but where this is not convenient, a lip beaker, preferably of porcelain, may be advantageously substituted. AMOUNT OF SAMPLE. The beginner in qualitative analysis finds it difficult to persuade himself that greater certainty is not reached by using large quantities of material. Experience teaches that small quantities are preferable, not alone by reason of greater ease in securing complete separations, but because much time is saved in manipulation. A fair general rule for un- known samples for complete analysis is to divide into four parts ; one for preliminary examination (see part IV), one for metal, and one for acid analysis, while the fourth is a reserve which in case of unsatisfactory results may be treated as the original sample. A gram of substance ought ordinarily to be sufficient for all of these analyses. CONFIRMATION OF TESTS. In the final identification, dependence is seldom to be placed upon one test alone. Each final reaction should be supplemented by additional evidence, when possible, and a sharp outlook for indications should be kept at all stages of analysis. TESTS OF REAGENTS. It is not always possible to 60 QUALITATIVE ANALYSIS secure perfectly pure reagents, and when analysis reveals only traces f a particular substance, it is frequently advisable to test the reagents used for this substance. This is particularly true for such common substances as iron, aluminium, sodium, and chlorine. This examination of the reagents is made by subject- ing the reagents used to the same series of tests as the sample which is undergoing examination. This is called "running a blank." In case the substance in question is found in the re- agents, judgment must be passed on the relative amounts found, in case purer reagents are not to be had. NOTEBOOKS. Notes are to be kept, in which are to be recorded, besides the equations already mentioned, the results of the analyses carried out and the answers to..J:he exercises which are found after each group. The student should also record any special observations made and any valuable hints, precautions, or interesting facts which are picked up in the course of study or work, and which are not to be found in the manual. Group I The Hydrochloric Acid Group Silver, Mercury (pus), and Lead GENERAL STATEMENT. The metals of this group are precipitated by hydrochloric acid lead incompletely. If lead is found in this group, it will also be found in Group IL SILVER : Solution for Reactions, Silver Nitrate. i. Hydrochloric acid or a soluble chloride precipitates white silver chloride, AgCl ; soluble in ammonia, forming Ag(NH 3 ) 2 Cl; reprecipitated by nitric acid as AgCl. Silver chloride in the light changes from white to lavender and finally to black. 1 1 The change caused by sunlight is not fully understood. Some chlorine is given off. Silver chloride precipitated in the dark has slightly different properties from that precipi- :, METAL ANALYSIS 6 1 'precipitates black silver sulphide, Ag 2 S ;. s^uble in hot cftlute nitric acid. 3. NH 4 OH precipitates white silver hydroxide, AgOH, which rapidly changes to brown silver oxide, Ag 2 O ; soluble in excess of reagent, forming Ag(NH 3 ) 2 OH and r whNH 3 ) 2 ] 2 O. 4. NaOH or KOH precipitates j upon corner oxide, Ag 2 O ; soluble in ammonia or dilute nitric aof plumbc 5. (NH 4 ) 2 S precipitates black Ag 2 S. ' See above. 6. Na2CO 3 precipitates white silver carbonate, Ag 2 CO 3 , or yellow basic carbonate. 7. K 2 CrO 4 precipitates from neutral solutions bright red silver chromate, Ag 2 CrO 4 ; soluble in nitric acid. 8. KI precipitates pale vellow silver iodide, Agl ; easily soluble in excess of reagent^fc^oluble in ammonia (!). 9. H 2 SO 4 precipitates from concentrated solutions white silver sulphate, Ag 2 SO 4 . 10. KCN precipitates white silver cyanide, AgCN ; easily soluble in excess, forming 1 1. Na 2 HPO 4 precipitates yellow silver aphosphate, Ag^PO^ ; soluble in dilute nitric acid, and ammonia. 12. Copper and some other metals precipitate metallic silver. MERCURY (ous) : Solution for Reactions, Mercurous Nitrate. 13. HC1 precipitates white mercurous chloride, HgCl; in- soluble in dilute acids ; changed by ammonia to a black mixture of Hg and HgNH 2 Cl. % 14. H 2 S precipitates a black mixture of mercury, Hg, and mercuric sulphide, HgS. K 2 S in presence of KOH dissolves the HgS, leaving Hg. HgS + Hg is insoluble in nitric acid, but soluble in aqua regia. tated in the light. Silver chloride exposed to the light is acted upon quickly by mild reduc- ing agent|j alkaline pyrogallate, hydroquinone, etc., depositing metallic silver, while that not Exposed to light is not so rapidly affected. Dry plate photography is based upon this difference. to 62 QUALITATIVE ANALYSIS 4 15. JOH or NaOH precipitates a black fnixti! Hg 2 O, and HgO ; soluble in nitric acid. 1 6. Na 2 CO 3 precipitates light yellow _^>asic carbonates, changing to gray bec^^^of decomposition into HgO, Hg, and CO 2 . *amn. 17. (NH 4 ) 2 S p used to t 5 a black mixture of Hg and HgS. See 14 above. ?oing exa- 1 8. K 2 CrO 4 preCT^fEates red mercurous chromate, Hg 2 CrO 4 ; insoluble in KOH. 19. NH 4 OH precipitates a black mixture of Hg and HgNH 2 NO 3 ; soluble in aqua regia. 20. KI precipitates yellowish green mercurous iodide, Hgl. 21. H 2 SO 4 precipitates white mercurous sulphate, Hg 2 SO 4 ; somewhat soluble in water. r|^>i 22. KCN produces^P|ray pr|^>itate of Hg and Hg(CN) 2 . * 23. SnCl 2 reduces mercurous compounds tu metallic mer- cury, Hg. % 24. Copper precipitates metallic mercury. Solution for Reactions, Lead Ac0ate or Nitrate. 25. HC1 precipitates, incompletely, lead chloride, PbCl 2 ; soluble in boiling water, crystallizing from this solution upon cooling; converted into white basic lead chloride, Pb(OH)Cl, by ammonium hydroxide. 26. H 2 S precipitates black lead sulphide PbS, even in the filtrate frffi the lead chloride precipitation ; insoluble in yellow ammonium sulphide ; soluble in warm dilute nitric acid, forming lead nitrate, Pb(NO 3 ) 2 , with the evolution of H 2 S. Part of the H 2 S is oxidized to sulphur jwith the reduction of the nitric acid to the oxides of nitrogen. Hot concentrated nitric acid oxidizes lead sulphide to white lead sulphate, PbSO 4 . 27. NH 4 OH precipitates white basic lead ciitrate, Pb 3 (OH)O 2 NO 3 , soluble mJ4NO 3 . With lead acetate solu- in H **. 63 P ordinary strength, excess of ammonium hydroxide (free from- carbonate) gives no precipitate. 28. (NH^S precipitaTO black lead sulphide, PbS. *See above. ,29. KOH or NaOH precipitate _ tehite lead hydroxide, Pb(OHA or basic salts, depending upon conditions ; soluble in excess of reagent, forming salts of plumbous acid, Na 2 PbO/ and K 2 PbO 2 ]\ soluble also in nitric acid.J 30. (NH 4 ) 2 CO 3 precipitates white basic lead carbonates; soluble in strong solution x)f potassium or sodium hydroxides or nitric afcid. 31. NagCO^precipitetes tne same. 32. K 2 CrO 4 precipitates yellow lead chromate, PbCrO 4 , " chiome yellow " ; solubldjjn sodium hydroxide ; soluble Vith difficulty in nitric acid. 33. KI precipitates yellow lead iodide, PbI 2 ; soluble in watj^, from which it crystajil^s upon cooling in shining gold yellow scales. 34. H 2 SO 4 or a soluble sulphate precipitates white lead sulphate, PbSO 4 ; soluble in alkaline ammonium tartrate, am- monium acetate, or sodium hydroxide. Lead sulphate is less soluble in water containing alcohol or dilute sulphuric acid (why ?) than in water alone. Soluble in concentrated H 2 SO 4 . 35. KCN precipitates white lead cyanide, soluble in very large excess of the reagent. Reprecipitated upon boiling. 36. SnCl 2 produces no precipitate. 37. Zn precipitates metallic lead in crystalline form. Analysis, Group I Make a mixture of the salts of the metals of this group, and separate by the following scheme of analysis : Dissolve in water with the addition of a few drops of nitric acid if necessary. .- 64 QUALITATIVE ANALYSIS The solution is tfeated with cold dilute hydrochloric IK as long as a precipitate is formed. Filter and wash twice with cold water. The precipitate contains the^members of this group. The filtrate may contain one or more metals of Groups II to V, and should be saved fiij^n the other groups are to be analyzed. (A) The precipitates washed with hot water several times. The residue may contain mercury, silver, and lead (from incom- plete washing). The filtrate contains lead chloride. (B) Divide the filtrate into several portions and apply the following tests : Treat with H 2 SO 4 ; a white precipitate indicates lead. Treat with H 2 S ; a black precipitate indicates lead. Treat with K 2 CrO 4 ; a yellow precipitate soluble in NaOH indicates lead. Wash with hot water until the filtrate no longer gives a test lead. (C) To the precipitate on the filter add ammonia ; a black residue shows the presence of mercury. me : actions I E. A white residue insoluble in NH 4 OH may be disregarded. (Why ? See re- forlead.) To the filtrate add an excess of HNO 3 ; a white precipitate shows the presence of silver. After having completed the analysis of the mixture, proceed to the analysis of unknowns Nos. I and 2. These being reported, prepare and record the following exercises : EXERCISES I. Tabulate the reactions which are common to all the mem- bers of this group. II. Devise another method for the analysis of the members of this group. III. Make a list of the important ores of lead, silver, and mercury. Define the term "ore." METAL ANALYSIS 65 What metals may be used to precipitate the ions of this group? (See Electromotive Series.) V. What are the chief uses of silver and lead as metals ? VI. What is white lead, and how is it made ? What is lunar caustic, and how is it used ? VII. What is the nature of the compound which is formed by ammonium hydroxide acting on silver salts ? Is silver am- photeric? What component of the solution of ammonium hydroxide effects the reaction? VIII. Explain the formation of basic salts and list those which are formed in -the preceding reactions. Define the term "basic salt." Group II Hydrogen Sulphide Group Mercury (ic\ Lead, Bismuth, Copper, Cadmium^ Arsenic, Antimony, and Tin GENERAL STATEMENT. The sulphides of this group are insoluble in dilute acids. The group is divided into two sub- groups : Subgroup A, those metals whose sulphides are insoluble in dilute yellow ammonium sulphide ; Subgroup B, those metals whose sulphides are soluble in dilute yellow ammonium sulphide/ Subgroup A Mercury (ic), Lead, Bismuth, Copper, and Cadmium MERCURY (ic) : Solution for Reactions, Mercuric Chloride. 38. HC1 produces no precipitate. 39u H 2 S, slowly added, forms, first a white precipitate (?); soluble in acids and in excess of the mercuric salt. By further addition of H 2 S, the precipitate becomes orange-yellow, then brown, and finally black mercuric sulphide is produced ; insolu- 66 QUALITATIVE ANALYSIS + ble in ammonium sulphide or hot nitric acid ; soluble in aqua regia. NOTE. The white precipitate is more easily secured by using mercuric nitrate. 40. NH 4 OH precipitates white mercuric amido chloride, HgNH 2 Cl; easily soluble in HC1; sparingly soluble in strong ammonium hydroxide. 41. (NH 4 ) 2 S precipitates HgS. The same changes in color as in (39) may be noted by careful addition of the reagent. 42. KOH or NaOH precipitates from cold solutions, first, reddish brown basic salts, which change to orange-yellow mer- curic oxide, HgO, when the reagent is added in excess. When heated, the yellow oxide changes to red HgO. In the presence of ammonium salts, the white "precipitate" is formed. See 40. (Explain.) 43. (NH 4 ) 2 CO 3 acts like ammonium hydroxide. 44. Na 2 CO 3 precipitates, first, a red-brown basic mercuric carbonate. With excess of reagents and heat, it is converted into yellow mercuric oxide, HgO. The basic salt formed with HgCl 2 is an oxychloride, HgCl 2 (HgO) 2 . 3 . or 4 . 45. K 2 CrO 4 precipitates an orange basic chromate, Hg 2 (OH) 2 Cr0 4 . 46. KI precipitates yellow mercuric iodide, HgI 2 ; soluble in concentrated HNO 3 and in HC1 ; soluble in solution of the iodides of the more positive metals, forming double iodides, as K 2 HgI 4 . 47. KCN precipitates from concentrated solutions white mer- curic cyanide, Hg(CN) 2 ; fairly soluble in water; easily soluble in KCN, forming a double salt, Hg(CN)2 . 2 KCN. Hg(CN) 2 is not acted upon by the hydroxides of the alkali metals. (Ex- plain. See Appendix, p. 177.) 48. SnCl 2 precipitates, in the cold, white mercurous chloride, HgCl, having a peculiar silky appearance when agitated. If METAL ANALYSIS 67 excess of the reagent is added, the mercurous chloride is reduced more or less completely to gray or black metallic mercury. 49. H 2 SO 4 produces no precipitate. 50. Fused with Na 2 CO 3 in a small tube, metallic mercury is deposited on the walls of the tube. The mercury salt should be mixed with five or six times its weight of Na^Og and heating begun slowly to drive out any moisture which may be present. This may be wiped out with a piece of filter paper. LEAD : Solution for Reactions, Lead Nitrate or Acetate. 51. The reactions for lead are given on p. 62. BISMUTH: Solution for Reactions, Bismuth Chloride. 52. Water precipitates white bismuth oxychloride, BiOCl ; soluble in hydrochloric acid, but insoluble in tartaric acid. (See Sb.) Converted by H 2 S to Bi 2 S 3 . 53. HC1 produces no precipitate. 54. H 2 S precipitates black bismuth sulphide, Bi 2 S 3 ; insoluble in dilute acids and alkalies and in ammonium sulphide, soluble in modejn?ly concentrated nitric acid. ^fc* *i J jr 58. (NH 4 ) 2 CQg Sf^Na 3 CO 3 precipitates white bismuthyl carbonate, (BiO) 2 CO 3 ; soluble in dilute acids. 59. K 2 CrO 4 precipitates yellow basic bismuth chromate, Bi 2 O(CrO 4 ) 2 ; insoluble in sodium hydroxide ; soluble in HNO 3 . 60. KI produces, in solutions not too strongly acid, dark brown bismuth iodide, BiI 3 . If the precipitate does not form 68 QUALITATIVE ANALYSIS after adding the reagent, make slightly less acid with NaOH until it appears ; soluble in excess of KI and HC1 ; not soluble in dilute nitric acid. 61. H 2 SO 4 produces no precipitate. 62. SnCl 2 produces no precipitate. 63. Zn or Fe precipitates spongy bismuth. COPPER : Solution for Reactions, Copper Sulphate. 64. HC1 produces no precipitate. 65. H 2 S precipitates black copper sulphide, CuS ; soluble in nitric acid and potassium cyanide ; insoluble in yellow ammonium sulphide. 66. NH 4 OH precipitates pale blue copper hydroxide, Cu(OH) 2 ; soluble in excess of reagent, forming a deep blue solution ; decolorized by KCN. 67. (NH 4 ) 2 S precipitates black copper sulphide. See 65 above. 68. KOH or NaOH precipitates pale blue copper hydroxide, Cu(OH)2 ; changed by boiling with excess of alkali to black copper oxide, CuO. *** 69. Na 2 CO 3 precipitates greenish blue basic copper carbonate, Cu 2 (OH) 2 CO 3 ; converted by foiling to tlack copper oxide, CuO. 70. K 2 CrO 4 precipitates reddish brown basic copper chromat) ; somewhat soluble in water ; soluble in acids. 71. KI precipitates from concentrated solutions white cuprous iodide, Cu 2 I 2 , with the liberation of iodine. 72. KCN precipitates yellowish green copper cyanide, Cu(CN) 2 ; soluble in excess of reagent, producing a double salt, K 2 Cu(CN) 4 . Not precipitated by H 2 S. (See p. 39.) 73. H 2 SO 4 produces no precipitate. 74. K 4 Fe(CN) 6 precipitates reddish brown copper ferrocyan- ide, Cu 2 Fe (CN) 6 ; insoluble in acids ; decomposed by alkalies. METAL ANALYSIS 69 75. SnC^ produces no precipitate. 76. Zn and Fe precipitate metallic copper. CADMIUM : Solution for Reactions, Cadmium Sulphate. 77. HC1 produces no precipitate. 78. H 2 S precipitates cadmium sulphide, CdS, varying in color from bright yellow to orange, according to conditions ; easily soluble in HC1, hot H 2 SO 4 , or hot HNO 3 ; insoluble in yellow ammonium sulphide and potassium cyanide. 79. NH 4 OH precipitates white cadmium hydroxide, Cd(OH)2 ; soluble in excess. If the ammoniacal solution is treated with potassium cyanide, a soluble double salt, K 2 Cd(CN) 4 , is formed from which the cadmium may be precipitated by H 2 S. See reaction 72. 80. (NH^S precipitates yellow cadmium sulphide, CdS. See 78. 8 1. KOH or NaOH precipitates white cadmium hydroxide, Cd(OH) 2 ; insoluble in excess; soluble in acids. 82. K^CrC^ precipitates, from concentrated solutions only, yellow cadmium chromate, CdCrO 4 ; soluble on addition of water. 83. KI produces no precipitate. 84. H 2 SO 4 produces no precipitate. 85. KCN precipitates white cadmium cyanide, Cd(CN) 2 ; soluble in excess, forming K^CdCCN^; reprecipitated by H 2 S as cadmium sulphide. 86. SnCl 2 produces no precipitate. Analysis, Group II, Subgroup A Make a mixture of the salts of the metals of this subgroup ; dissolve in water with the aid of hydrochloric acid, and separate and identify by the following scheme of analysis : The solution to which hydrochloric acid is added is boiled nearly to dryness to decompose any nitric acid which may be 7 QUALITATIVE ANALYSIS present. It is then diluted with water, and a small amount of hydrochloric acid is added if necessary l to dissolve any residue which may be present. NOTE. As nitric acid decomposes hydrogen sulphide, the solution should be boiled nearly to dryness with frequent additions of small quantities of hydrochloric acid to remove the nitric acid, unless it is known that no nitric acid or nitrates are present. The solution is now heated to boiling and hydrogen sulphide run inyslowly 2 until saturated. The solution is then cooled, diluted with water, and more hydrogen sulphide run in. The precipitate, which contains the metals of Group II, is filtered off and washed with hot water until a few drops of the filtrate impart only 'slight turbidity to silver nitrate solution. The filtrate contains the metals of Group III to V, and should be saved if these are to be sought. (A) The precipitate is warmed with dilute ammonium sulphide and filtered. The filtrate may contain members of Subgroup B, and should be saved if the members of this group are sought. The precipitate insoluble in ammonium sulphide contains all the members of Subgroup A. Analyzed by (B) and following : 3 (B) The precipitate is boiled with dilute nitric acid for a few minutes, filtered, and washed. The precipitate may contain mercuric sulphide (black), lead sulphate (white), or sulphur (yellow, or black, and floating). Test for mercury by adding aqua regia, boiling, evaporating to expel nitric acid, diluting with water and hydrochloric acid, and adding stannous chloride. See reaction 48. 1 The precipitation of this group is accomplished with best results if the solution contains about one part of HC1 in ten of H 2 O. 2 It is suggested that the solution be placed in an Erlenmeyer flask, fitted with a one- holed stopper, through which runs a glass tube nearly to the bottom of the flask. Attach to the H 2 S generator, loosen the stopper, and pass in H 2 S until the air is replaced, then close tightly and run in H 2 S, with frequent shaking of the flask, until action ceases. 3 If Subgroup B is known to be absent, as in the preliminary separation, (A) may be omitted. METAL ANALYSIS 71 (C) The filtrate from (B) is treated with 2 c.c. of sulphuric acid and evaporated until white fumes of sulphur trioxide are given off. Cool, add water, and filter if necessary (Filtrate D). A white precipitate indicates lead. Add concentrated ammonium acetate 1 and a little acetic acid to the precipitated lead sulphate, warm, and filter. To the filtrate add potassium chromate. A yellow precipitate soluble in sodium hydroxide shows the pres- ence of lead. (D) The filtrate from the sulphuric acid precipitate is treated with an excess of ammonia. Filter if necessary. To the precip- itate add a few drops of cone. HC1; evaporate to small volume and pour into about 500 c.c. of water. A white precipitate shows the presence of bismuth. (E) The filtrate from (D) may contain copper and cadmium. If the filtrate is blue, copper is present. If faintly blue, confirm the presence of copper, in a small portion, by 74. (F) If copper is present, decolorize with potassium cyanide and treat with hydrogen sulphide. A yellow precipitate shows the presence of cadmium. 2 After having completed the analysis of the mixture, proceed to the analysis of unknowns Nos. 3 and 4. These being reported, prepare and record the following : EXERCISES I. What reactions are common to all the metals of this sub- group ? II. Explain the presence of sulphur and lead sulphate in (B). III. If the filtrate from the mercuric sulphide in (B) were treated at once with ammonium hydroxide, what might the pre- cipitate contain ? 1 2 PbS0 4 + 2 NH 4 C 2 H 3 2 -> Pb(XH 4 ) 2 (S0 4 ) 2 + Pb(C2H 3 O 2 ) 2 . 2 If the precipitate is not yellow, it may be redissolved in dilute hydrochloric acid, boiled, made alkaline with an excess of ammonia, filtered, and reprecipitated with hydrogen sulphide. 72 QUALITATIVE ANALYSIS IV. What indications of the presence or absence of cer- tain metals are given by the color of the hydrogen sulphide precipitate ? V. What are the ions of hydrogen sulphide, which are most numerous ? Why are sulphides precipitated instead of hydro- sulphides ? Cf. Smith, General Chemistry, p. 374. VI. What are the ions present in a solution of copper hydroxide ? Does copper act as an acid-forming element ? VII. Make a table showing mineralogical names and formulas of the important ores of each metal of the group. VIII. Give the chemical reactions for at least one method of preparation of each of these metals. IX. What are the uses of calomel and corrosive sublimate, and how are they distinguished from each other ? Smith, General Chemistry, p. 654. X. What is " red precipitate " ? " White precipitate " ? XL What is Wood's metal ? Rose's metal ? What are amalgams ? XII. What are the uses of bismuth subnitrate ? of copper sulphate ? of cadmium sulphide ? XIII. What are the chief alloys of copper ? XIV. Why is copper not precipitated by hydrogen sulphide from the solution containing an excess of KCN, while cadmium is precipitated ? Why is the blue solution decolorized by KCN ? Subgroup B Metals whose Sulphides are Soluble in Dilute Yellow Ammonium Sulphide Arsenic, Antimony, and Tin ARSENIC (ous) : Solution for Reactions, Sodium Arsenite or Arse nous Chloride.^- 87. HC1 produces no precipitate. 88. H 2 S precipitates, from hydrochloric acid solution, yellow 1 Arsenous chloride is prepared by dissolving arsenous oxide in hydrochloric acid. METAL ANALYSIS 73 arsenous sulphide, As 2 S 3 ; soluble in yellow ammonium sul- phide, forming sulpho salts, such as (NH 4 ) 3 AsS 4 ; reprecipitated as As 2 S 5 and As 2 S 3 by dilute HC1 ; soluble in sodium hydroxide, ammonium carbonate, hot nitric acid, and nascent chlorine, nearly insoluble in warm concentrated hydrochloric acid. 89. NH 4 OH, KOH, or NaOH produces no precipitate. 90. (NH^S precipitates As 2 S 3 from acid solutions ; soluble in excess. See 88 above. 91. (NH 4 ) 2 CO 3 or Na^COg produces no precipitate. 92. AgNO 3 precipitates from neutral solutions silver arsenite, Ag 3 AsO 3 ; readily soluble in dilute acids, ammonium hydroxide, or ammonium salts. 93. CuSO 4 in neutral solution precipitates " Scheele's green," CuHAsO 3 ; soluble in ammonium hydroxide and dilute acids. 94. A bright strip of copper placed in a solution of arsenic (ous), made strongly acid with HC1 and boiled, gives a gray film of arsenic on the copper. " Reinsch's test." 95. Oxidizing agents, as potassium chromate and potassium permanganate, convert arsenous compounds to arsenic com- pounds. 96. Nascent hydrogen 1 reduces arsenic compounds to a colorless poisonous gas, arsine, AsH 3 , decomposed by heating 1 Marsh's Test for Arsenic. A flask containing arsenic free zinc is fitted with a funnel tube and a delivery tube of glass drawn to a small point at the end. Hydrochloric acid is added to the zinc in the flask through the funnel tube until hydrogen is freely evolved. As soon as all of the air is out of the apparatus (test) (see Instructor), light and test for arsenic or allied metals by holding a cold porcelain dish in the flame. If no spot is formed on the dish, the apparatus is ready for carrying out the test for arsenic. Add through the funnel tube a hydrochloric acid solution of an arsenic compound. Heat the delivery tube near the tip. Note the black metallic mirror of arsenic formed. Write equation. Light the issuing gas at the tip of the delivery tube and hold a cold porcelain dish in the flame. Note the spot ot black metallic arsenic. Get several spots on the dish and save for future test No. 98. Put out the flame and run the gas into a solution of silver nitrate in a test tube. A precipitate of black metallic silver is formed. Add HC1 to the solution in the test tube until a precipitate no longer forms. Filter and test the filtrate for arsenic by H 2 S. 74 QUALITATIVE ANALYSIS to metallic arsenic and hydrogen ; reduces a solution of silver nitrate to metallic silver, changing at the same time to arsenous acid, H 3 AsO 3 , which may be precipitated as arsenic trisulphide with hydrogen sulphide. 97. KCN and Na^Og fused with arsenic trioxide or tri- sulphide in a tube produce deposits of metallic arsenic on the cold parts of the tube as a shining black mirror, the KCN changing to KCNO or KCNS. The moisture should first be removed in this experiment as directed in 50. 98. Arsenic spots 1 dissolve in sodium or calcium hypochlo- rite and in hot nitric acid. ARSENIC: Solution for Reactions, Sodium Ar senate. 99. HC1 produces no precipitate. 100. H 2 S precipitates yellow arsenic sulphide, or arsenous sulphide, As 2 S 3 , and sulphur depending upon conditions. The second reaction takes place more freely when hydrogen sulphide is passed slowly through a hot solution. Both are soluble in ammonium sulphide. See 88. 101. NH 4 OH produces no precipitate. 102. KOH or NaOH produces no precipitate. 103. (NH^S precipitates yellow arsenic sulphide, As 2 S 3 , from solutions acidified with hydrochloric acid. 104. (NH 4 ) 2 CO 3 or Na 2 Co 3 produces no precipitate. 105. AgNO 3 precipitates from neutral solutions reddish brown silver arsenate, Ag 3 AsO 4 ; readily soluble in dilute acids or ammonia. 1 06. CuSO 4 precipitates greenish blue acid arsenates as CuHAsO 4 ; soluble in dilute acids and ammonia. 107. Magnesium mixture 2 precipitates white magnesium ammonium arsenate, MgNH 4 AsO 4 ; easily soluble in acids. 1 The spots produced by burning AsH 3 when a cold dish is held in the flame are called arsenic spots. 2 For preparation, see Appendix. METAL ANALYSIS 75 108. Nascent hydrogen reduces arsenic compounds to arsine, AsH 3 . See 96. 109. Reducing agents FeSO 4 , Na^Og, etc., heated with H 2 SO 4 , reduce arsenic acid to arsenous acid, H 3 AsO 3 . no. Heated with charcoal in a closed tube, a mirror of metallic arsenic is formed on the walls of the tube, in. Arsenic spots. See 98. ANTIMONY : Solution for Reactions, Antimony Chloride. 112. HC1 produces no precipitate. 113. H 2 S precipitates orange-red antimony sulphide, Sb 2 S 3 ; soluble in ammonium sulphide, forming (NH 4 ) 3 SbS 4 (see As); soluble in sodium hydroxide and in warm concentrated hydro- chloric acid. 114. NH 4 OH precipitates white antimony hydroxide, Sb(OH) 3 ; soluble in dilute hydrochloric acid or sodium hydrox- ide. 115. NaOH or KOH precipitates the same ; soluble in excess of reagent or in hydrochloric acid. 1 1 6. (NH^S precipitates from acid solutions orange-red antimony trisulphide, Sb 2 S 3 ; soluble in ammonium sulphide. See 113. 117. (NH^COg or Na 2 CO 3 precipitates white antimony hydroxide, Sb(OH) 3 ; soluble upon warming in large excess of sodium carbonate, but not in ammonium carbonate; soluble in hydrochloric acid. 1 1 8. Water precipitates white antimony oxychloride, SbOCl; soluble in acids ; converted by H 2 S into Sb 2 S 3 . 119. Nascent hydrogen (see Marsh's test) reduces antimony compounds to stibine, SbH 3 . 1 20. Antimony spots are not dissolved in hypochlorites, but are by concentrated HNO 3 . 76 QUALITATIVE ANALYSIS 1 2 1. 1 Tin 2 in presence of hydrochloric acid and platinum foil precipitates metallic antimony as a dark brown powder adher- ing to the platinum ; concentrated nitric acid causes the stain to disappear; forming meta-antimonic acid, HSbO 3 . If the liquid is touched with a glass rod dipped in an ammoniacal solution of silver nitrate, white silver meta-antimonate, AgSbO 3 , is formed. TIN (STANNOUS) : Solution for Reactions, Stannous Chloride 122. HC1 produces no precipitate. 123. H 2 S precipitates dark brown stannous sulphide, SnS ; soluble in yellow, but not in colorless, ammonium sulphide, - forming ammonium sulphostannate, (NH 4 ) 2 SnS 3 ; reprecipitated by dilute acids, forming yellow stannic sulphide, SnS 2 . 124. NH 4 OH precipitates white stannous hydroxide, Sn(OH) 2 ; insoluble in excess; soluble in dilute HC1. 125. NaOH or KOH precipitates the same; soluble in excess, forming stannites, Sn(ONa) 2 or Sn(OK) 2 . 126. Na 2 CO 3 precipitates the same as 1254 127. (NH^S precipitates from acid solutions dark brown stannous sulphide, SnS. See 123. 128. HgCl 2 is reduced by SnCL, forming white HgCl, or black Hg if SnCl 2 is in excess ; at the same time SnCl 2 is oxidized to stannic chloride, SnCl 4 . 129. Zinc in hydrochloric acid solution precipitates metallic tin as a spongy mass upon the zinc. For distinction from antimony, see 121. 130. Oxidizing agents, such as HNO 3 , KC1O 3 , etc., in HC1, convert stannous salts to stannic salts. 1 To perform this test a few drops of the solution are placed upon a clean platinum foil ; dip one end of a U-shaped strip of tin foil in the liquid, leaving the other end in con- tact with the platinum beyond the drop. A bright silver coin may be used instead of platinum. 2 Tin is used rather than zinc, because the latter also precipitates tin, which may be mistaken for antimony, although the deposit is much brighter. METAL ANALYSIS 77 TIN (STANNIC) : Solution for Reactions, Stannic Chloride! 131. HC1 produces no precipitate. 132. H 2 S precipitates yellow stannic sulphide; soluble in warm concentrated HC1 ; soluble in yellow ammonium sulphide, forming (NH 4 ) 2 SnS 3 . See 123. 133. NH 4 OH precipitates white metastannic acid, SnCXOH^. 134. NaOH precipitates the same; soluble in excess. 135. Na^Og precipitates the same. 136. (NH^S precipitates yellow stannic sulphide, SnS 2 ; soluble in yellow ammonium sulphide. 137. HgCl 2 produces no precipitate. 138. Zinc precipitates metallic tin. See 129. Analysis, Group II, Subgroup B Make a mixture of the salts of the metals of this subgroup, and separate by the following scheme of analysis ; 2 (A) The filtrate 3 from (A), Subgroup A, is treated with dilute HC1 to acid reaction. Filter and wash with hot water. The filtrate is discarded. The precipitate, containing all the members of this subgroup present, is treated with concentrated hydrochloric acid and warmed for a few minutes, filtered, and washed. (B) The residue may contain arsenic and sulphur; boil with concentrated HC1 and a few crystals of KC1O 3 , gradually added. Take out the floating sulphur with a glass rod ; make alkaline with ammonia, and test the solution for arsenic with magnesium mixture. (See other tests.) 1 The stannic chloride used for these preliminary reactions should be prepared from stannous chloride by treatment with HC1 and KC1O 3 and heating until chlorine is no longer evolved. Metastannic acid gives somewhat different reactions, but since it is pre- cipitated by H 2 S as stannic sulphide, it is considered sufficient to use the stannic salt given above for these preliminaries. 2 Begin the analysis for Subgroup B as given for analysis of Subgroup A. 3 If Subgroup A is known to be absent, the solution and reprecipitation with and HC1 may be omitted and the analysis begun with the H 2 S precipitate. 78 QUALITATIVE ANALYSIS (C) The filtrate from (A) may contain tin and antimony. 1 Add an excess of zinc, 2 and a piece of platinum foil, taking care that the platinum touches the zinc. (D) The black coating on the platinum is thoroughly washed with hot water and touched with a drop of nitric acid, and then a drop of ammoniacal silver nitrate. A white deposit confirms the presence of antimony. (E) The excess of zinc with the coating of tin is dissolved in HC1, boiled nearly to dryness to expel excess of acid, diluted, and treated with mercuric chloride. A white or gray precipitate confirms the presence of tin. If present in small quantities, it is best to boil ; a black precipitate of Hg after standing a few minutes confirms the presence of tin. Alternative analysis. The filtrate from (A) may be tested for antimony and tin without the use of platinum as follows: Boil to expel hydrogen sulphide and place a few drops on a bright silver surface (coin). A U-shaped strip of tin foil is placed so that one end touches the silver beyond the drop. In the presence of antimony a dark stain appears on the silver. Another portion is tested for tin by nearly neutralizing the acid, and adding a bit of granulated zinc so that some zinc re- mains after action has ceased. The residue of zinc with the adhering tin is dissolved in hydrochloric acid and tested as in (E) above. After having completed the analysis of the mixture, proceed to the analysis of unknown Nos. 5 and 6. These being reported, prepare and record the following : 1 A very satisfactory identification of antimony is carried out as follows : Take a drop of the solution containing antimony and tin and place on a platinum foil; place a U- shaped strip of tin with one end in and the other outside the drop in contact with the platinum foil. A black precipitate on the foil is conclusive evidence of the presence of antimony. 2 So that when action ceases, there will be some zinc left. METAL ANALYSIS 79 EXERCISES I. Devise another method for the separation of the members of this subgroup, if in solution by themselves. II. Of what use is KC1O 3 in (B)? What is ordinary com- mercial " arsenic " ? III. Explain why SnS is not dissolved by colorless am- monium sulphide, and why As 2 S 3 and Sb 2 S 3 are dissolved only with difficulty. IV. What are the chief ores of arsenic, antimony, and tin? V. Make a table of the antimony, arsenic, and arsenous acids, showing names and formulae. VI. How is most of the " arsenic " of commerce obtained ? VII. What is "tin plate"? How produced? Banca tin? Block tin ? VIII. What is tartar emetic ? How made ? Group III The Ammonium Sulphide Group Iron, Chromium, Aluminium, Manganese, Zinc, Nickel, and Cobalt GEN ERAL ST ATEM ENT. Iron, chromium, and aluminium are precipitated as hydroxides by ammonium hydroxide in the presence of ammonium chloride. Manganese, zinc, nickel, and cobalt are not precipitated unless phosphoric or some similar acid is present. Ammonium sulphide converts all the metals into sulphides except chromium and aluminium, which come down as hydroxides. In the presence of phosphates, borates, oxalates, silicates, fluorides, or tartrates, barium, strontium, cal- cium, and magnesium may be precipitated in this group. For test for the presence of these acids, see 'acid analysis. 80 QUALITATIVE ANALYSIS IRON (FERROUS): Solution for Reactions, Ferrous Sulphate. 139. HC1 produces no precipitate. 140. H 2 S in acid solution produces no precipitate./ 141. NH 4 OH precipitates, incompletely, ferrous hydroxide/' changing slowly to reddish brown ferric hydroxide, Fe(OH) 3 ; soluble in acids, even acetic acid ; treated with (NH 4 ) 2 S forms ferrous sulphide, FeS, black. 142. (NH 4 ) 2 S precipitates black ferrous sulphide, FeS; soluble in HC1. 143. NaOH or KOH precipitates white ferrous hydroxide/ oxidizing immediately to a dirty green, to black, and finally -to reddish brown ferric hydroxide. See 141. - 144. (NH 4 ) 2 CO 3 precipitates white ferrous carbonate, FeCO 3 , which quickly oxidizes and darkens in color, eventually changing to Fe(OH) 3 ; soluble in acids, even acetic acid. 145. NagHPC^ precipitates a mixture of acid ferrous phos- phate, FeHPO 4 , and ferrous phosphate, Fe 3 (PO 4 ) 2 , white to bluish white. By the addition of an alkali acetate, ferrous ' rift phosphate alone, Fe 3 (PO 4 ) 2 , is formed; soluble in HC1 or HNOg. 1 146. K 4 Fe(CN) 6 precipitates white potassium ferroferro- cyanide, K 2 FeFe(CN) 6 , quickly turning blue by oxidation, a small amount of ferric^^i^cyanide, Fe 4 (Fe(CN) 6 ) 3 , being formed. 147. K 3 Fe(CN) 6 precipitates dark blue ferrous ferricyanide, Fe 3 (Fe(CN) 6 )2, " Turnbull's blue"; insoluble in dilute acids. 148. KCNS produces no coloration when ferric iron is entirely absent. 149. Boiled with HNO 3 , a ferric salt is formed. 150. Chlorine water oxidizes a solution of ferrous salt to a ferric salt. 151. Barium carbonate suspended in water, when shaken 1 For this and the two following tests a ferrous salt free from ferric iron must be used. See appendix for preparation. METAL ANALYSIS 8 1 with a cold neutral or slightly acid solution of a ferrous salt, produces no precipitate containing iron. Why ? (See Table of Solubilities.) Cf. 163 and 171. IRON (FERRIC): Solution for Reactions, Ferric Chloride. 152. HC1 produces no precipitate. 153. H 2 S reduces solutions of ferric salt to the ferrous con- - dition, sulphur being set free.' 154. NH 4 OH precipitates reddish brown gelatinous ferric hydroxide, Fe(OH) 3 ; soluble in acids. 155. KOH or NaOH precipitates the same. 156. (NH^S precipitates ferrous sulphide, FeS, mixed with free sulphur. 157. (NHJ2CO 3 or Na 2 CO 3 precipitates Fe(OH) 3 . 1 58. Na 2 HPO 4 precipitates ferric phosphate, FePO 4 ; slightly soluble in acetic acid; readily soluble in HC1, HNO 3 , and H 2 SO 4 . 159. K 4 Fe(CN) 6 precipitates blue ferric ferrocyanide, Fe 4 (Fe(CN) 6 ) 3 , " Prussian blue "; insoluble in dilute inorganic acids. 1 60. K 3 Fe(CN) 6 produces no precipitate, but imparts a reddish brown color. 161. KCNS produces an intense red color of ferric sulpho- cyanate, Fe(CNS) 3 . +* 162. Most reducing agentrfas H 2 S, SO 2 , and SnCl 2 , easily reduce ferric to ferrous salts. 163. BaCO 3 suspended in water, when shaken with a cold neutral or slightly acid solution of a ferric salt, precipitates brown ferric hydroxide, Fe(OH) 3 ; soluble in HC1, and may be confirmed by 161. CHROMIUM: Solution for Reactions, Chromium Sulphate. 164. HC1 produces no precipitate. 165. H 2 S produces no precipitate in an acid solution of a chromium salt. 82 QUALITATIVE ANALYSIS 1 66. NH 4 OH precipitates bluish green chromium hydroxide, Cr(OH) 3 ; soluble with difficulty in NH 4 OH and NH 4 C1 in cold ; reprecipitated upon boiling. Soluble in acids. 167. (NH 4 ) 2 S precipitates Cr(OH) 3 . 168. KOH or NaOH precipitates Cr(OH) 3 ; soluble in ex- cess, forming green chromite salts, K 3 CrO 3 and Na 3 CrO 3 , repre- cipitated upon boiling. 169. (NH^COg or Na 2 CO 3 precipitates Cr(OH) 3 . 170. The oxides and the salts of chromium are converted into chromic acid or chromates by powerful oxidizing agents, e.g. by fusion with sodium carbonate and potassium nitrate on platinum foil, they give potassium chromate ; soluble in water. If the solution is acidified with acetic acid, and lead ace- tate added, a yellow precipitate of lead chromate is formed, PbCr, 173. H 2 S or other reducing agents, as SO 2 or alcohol, con- vert solutions of chromates, to which HC1 has been added, into green solutions of chromium salts. 174. NH 4 OH produces no precipitate. 1 The oxidation may be also accomplished by adding to a solution of the salt in ex- cess of KOH or NaOH, chlorine, bromine, or hydrogen peroxide. The insoluble chro- mium compounds may be converted to chromates also by fusion with NagO2 in iron, nickel, or silver vessels. No platinum may be used with Na2O2. 2 Reactions for chromic acid are placed here, as the chromium present as an acid is always precipitated as basic chromium in this group, and the reactions for acid chromium are very different from the reactions of basic chromium. METAL ANALYSIS 83 175. (NH^S in neutral or alkaline solutions precipitates green chromium hydroxide, Cr(OH) 3 , with oxidation of the sul- phide. In case of the polysulphide of ammonia, a thiosulphate is obtained. 176. (NH^COg or Na 2 CO 3 produce no precipitate. 177. AgNO 3 precipitates from neutral solutions red silver chromate ; soluble in nitric acid or in ammonia. 178. Lead nitrate or acetate precipitates yellow lead chro- mate, "chrome yellow," PbCrO 4 , which is soluble in NaOH, but with difficulty soluble in HNO 3 . 179. Hg(NO 3 )2 precipitates a dark red basic mercuric chro- mate, Hg 2 (OH) 2 CrO 4 ; soluble with difficulty in nitric acid. 1 80. Bismuth chloride precipitates yellowish bismuth chro- mate, Bi^CrOJgj soluble in HNO 3 . 181. Cadmium nitrate precipitates a yellow basic cadmium chromate, Cd^OH^CrC^ ; soluble in HNO 3 . 182. H 2 O 2 in acid solution produces a deep blue coloration of perchromic acid, H 2 Cr 2 O 8 . If ether is added and then H 2 O 2 and the mixture shaken, the ether takes up the perchromic acid, giving a deep blue coloration to the ether. 1 183. Bead Test? All compounds of chromium color the borax bead green. See p. 124. 184. BaCO 3 suspended in water precipitates from cold neu- tral or slightly acid solutions yellow barium chromate. ALUMINIUM : Solution for Reactions, Aluminium Sulphate. 185. HC1 produces no precipitate. 1 86. H 2 S produces no precipitate from acid solutions. 1 The more dilute the acid solution is made, the more delicate the test. 2 The platinum wire used for flame tests may be used for bead tests. Heat the wire in the flame and place in powdered borax. Heat again in the flame, and a small globule of melted borax will stick to the wire. Touch while warm to a small particle of the ma- terial to be tested and heat in the oxidizing flame until the material has dissolved and note the color. 84 QUALITATIVE ANALYSIS 187. NH 4 OH precipitates white flocculent aluminium hydrox- ide, A1(OH) 3 ; soluble in acids. 1 88. (NH 4 ) 2 S precipitates the same. 189. KOH or NaOH precipitates the same; soluble in ex- cess, but reprecipitated by boiling with NH 4 C1. Freshly pre- cipitated A1(OH) 3 is easily soluble in acids. 190. *(NH 4 ) 2 CO 3 or Na 2 CO 3 precipitates A1(OH) 3 or basic carbonates. 191. Bead Test. No colored bead is formed. 192. BaCO 3 , suspended in water, when shaken with a cold neutral or slightly acid solution, precipitates aluminium hydrox- ide, A1(OH) 3 . MANGANESE: Solution for Reactions, Manganese Sulphate. 193. HC1 produces no precipitate. 194. H 2 S produces no precipitate from acid solutions. 195. NH 4 OH in the absence of ammonium salts precipitates, incompletely, white Mn(OH) 2 ; this oxidizes quickly in the air, darkening in color. In the presence of ammonium salts this precipitate is not formed, but, upon standing, the solution soon becomes cloudy, and ultimately all the manganese is precipi- tated as dark brown MnO(OH) 2 . 196. (NH 4 ) 2 S precipitates flesh-colored manganese sulphide, MnS ; soluble in dilute mineral acids and acetic acid. It oxi- dizes and turns brown, upon standing, forming Mn 2 O 3 , MnSO 4 , and sulphur. 197. NaOH or KOH precipitates white Mn(OH) 2 ; oxidized quickly by the air ; darkening in color ; eventually forming Mn 2 O 3 . 198. (NH 4 ) 2 CO 3 or NagCOg precipitates white manganese carbonate or basic carbonates, which oxidize in the air to Mn 2 O 3 . 199. Fused with Na 2 CO 3 and KNO 3 upon platinum foil, METAL ANALYSIS 85 compounds of manganese oxidize to a bright green sodium manganate, NagMnC^ ; soluble in small amount of water ; de- composes on standing in solution to MnO 2 and KMnO 4 . 200. Bead Test. Compounds of manganese color the borax bead amethyst in the oxidizing flame ; colorless in the reducing flame. 20 1. PbO 2 when boiled with dilute H 2 SO 4 and a small quan- tity of manganese salts imparts a pink or purple color to the solution, due to the formation of permanganic acid, HMnO 4 . The color is best noted if the solution is allowed to settle for a few minutes. If chlorides are present, the dry substance should first be treated with a few drops of H 2 SO 4 , and heated until white fumes of SO 3 appear. (Why ?) 202. BaCO 3 , suspended in water, when shaken with a cold neutral or acid solution of manganese salts, produces no pre- cipitate. ZINC : Solution for Reactions, Zinc Sulphate. 203. HC1 produces no precipitate. 204. H 2 S precipitates incompletely white zinc sulphide, ZnS, from neutral solutions of zinc salts of inorganic acids, freely soluble in inorganic acids ; slightly soluble in acetic acid. 205. NH 4 OH precipitates white gelatinous zinc hydroxide, Zn(OH)2 ; soluble in excess of NH 4 OH ; reprecipitated by boil- ing. This precipitate is not formed in the presence of ammo- nium salts. 206. (NH^S precipitates white zinc sulphide ; soluble in dilute inorganic acids, but not easily in acetic acid. 207. KOH or NaOH precipitates white gelatinous zinc hydroxide, Z^OH^; soluble in excess, forming zincates, as NajjZnOa and K 2 ZnO 2 . 208. (NH 4 ) 2 CO 8 or NaaCOg precipitates white basic zinc carbonates ; soluble in large excess of reagent. 86 QUALITATIVE ANALYSIS 209. Bead Test. No colored bead is formed. 210. BaCO 3 , suspended in water, when shaken with a cold neutral or acid solution of zinc salts, produces no precipitate. NICKEL : Solution for Reactions, Nickel Nitrate. 211. HC1 produces no precipitate. 212. H 2 S produces no precipitate from neutral or acid solu- tions. 213. NH 4 OH precipitates incompletely light green nickelous hydroxide, Ni(OH) 2 ; soluble in excess, producing a blue solu- tion. Salts of ammonium prevent this precipitation. 214. (NH 4 ) 2 S precipitates black nickelous sulphide, NiS; soluble in hot dilute HC1, HNO 3 , or aqua regia. 215. KOH or NaOH precipitates light green, Ni(OH) 2 ; in- soluble in excess. It is oxidized by boiling with bromine water to Ni(OH) 3 , black. If the precipitate is filtered off and boiled with NH 4 OH, it is reduced to Ni(OH) 2 with evolution of nitro- gen. If NH 4 C1 is present, the Ni(OH) 2 dissolves. 216. (NH 4 ) 2 CO 3 or Na 2 CO 3 precipitates light green basic carbonates ; soluble in large excess of the reagent. 217. KNO 2 produces no precipitate in acetic acid solution. Cf. 227. 218. KCN precipitates yellowish green nickelous cyanide, Ni(CN) 2 ; insoluble in dilute HC1; soluble in excess of KCN. 219. Bead Test. Salts of nickel color the borax bead reddish brown in the oxidizing flame, gray in the reducing flame. 220. BaCO 3 , suspended in water, when shaken with a cold neutral or acid solution of nickel salts, produces no precipitate. COBALT : Solution for Reactions, Cobalt Nitrate. 221. HC1 produces no precipitate. 222. H 2 S produces no precipitate from neutral or acid solu- tions. 223. NH 4 OH precipitates, incompletely, blue basic cobalt METAL ANALYSIS 87 salts ; soluble in excess to a brownish red solution. Ammonium salts prevent this precipitation. 224. (NH 4 ) 2 S precipitates black cobalt sulphide, CoS ; in- soluble in cold dilute HC1; soluble in HNO 3 or aqua regia. 225. KOH or NaOH precipitates blue basic cobalt salts, which change upon warming to pink cobalt hydroxide, Co(OH)2 ; not soluble in excess of the alkali. Ammonia or ammonium salts dissolve the precipitate. 226. Na^COg precipitates red basic cobaltous carbonates which, when boiled, lose CO 2 , giving a violet color, or if the reagent is in excess, a blue color ; soluble in ammonium carbonate. 227. KNO 2 precipitates, from solutions rendered strongly acid with acetic acid, a yellow crystalline double salt, K 3 Co(NO 2 ) 6 , but usually only on long standing. 228. KCN precipitates brownish white cobaltous cyanide, Co(CN)2; soluble in excess of reagent, and reprecipitated by HC1 or H 2 SO 4 . If to the solution in excess of KCN a few drops of HC1 are added, and the solution boiled, potassium cobaltic cyanide, K 3 Co(CN) 6 , is formed, which is not reprecipi- tated by sodium hypobromite. See 218. 229. Bead Test. Compounds of cobalt color the borax bead a deep blue in both oxidizing and reducing flame. 230. BaCO 3 , suspended in water, when shaken with a cold neutral or acid solution of cobalt salts, produces no precipitate. Analysis, Group III Make up a mixture of salts : of metals of this group, dissolve, 2 and identify by the following scheme of analysis : A few drops of HNO 3 are added and the solution is boiled. If the solution turns yellow, iron is probably present. 1 The salts mentioned on p. oo are to be avoided in making this solution. In case the sample for analysis contains these acid radicals, the method of separation on p. oo must be used. 2 If difficulty is encountered in making a solution, consult pp. 'oo-oo. 88 QUALITATIVE ANALYSIS (A) Add a small quantity of NH 4 C1 and then NH 4 OH to alkaline reaction and boil. If a precipitate forms, one or more of the metals iron, chromium, and aluminium may be present. If no precipitate forms these are absent, and subsequent tests for them need not be applied. 1 (B) Whether a precipitate is formed or not in (A), add to the solution (without filtering) (NH 4 ) 2 S while still warm. Filter at once and wash the precipitate with water containing a little (NH 4 ) 2 S. The filtrate may contain metals of Groups IV and V. (If the filtrate is dark brown in color, the presence of nickel is indicated, due to some NiS remaining in colloidal solution. 2 If this is the case, the filtrate should be acidified with acetic acid and boiled for some time, filtered on a separate filter, and the precipitate tested for nickel by (D)). The filtrates are saved, if analysis of Groups IV and V is to be made. (C) Remove the precipitate from the paper and digest with cold dilute HC1; 3 filter, and wash thoroughly. The residue may contain nickel or cobalt, or both ; test by (D). If the precipitate is light colored, only sulphur is present and may be discarded. The filtrate may contain Fe, Al, Cr, Mn, and Zn, and is tested by (E). (D) The precipitate is tested for nickel and cobalt with a borax bead in the oxidizing flame. A brown bead indicates nickel, a blue bead cobalt. If the bead is violet when hot and reddish brown when cold, cobalt is present only in traces, or is absent. If cobalt is present, nickel may also be present, but the test for nickel is obscured by the deep color of the cobalt bead. 1 If desired, this precipitate may be filtered and examined as in (G), (H), and (I). The filtrate is then treated with (NH 4 ) 2 S and treated as in (C), (D), and (L). 2 If the precipitation has been performed properly, no NiS should remain in solution. 8 See Appendix for making dilute HC1. METAL ANALYSIS 89 If cobalt is present, proceed as follows : To the precipitate add a small amount of HC1 and a few drops of HNO 3 , and evaporate nearly to dryness. Add about 5 c.c. of water and filter ; discard the residue on the filter, which is sulphur. Evaporate to a small bulk, if necessary. Add NaOH, drop by drop, until a permanent precipitate is formed ; this is dissolved in a slight excess of KCN 1 and boiled. A volume of bromine water equal to the whole solution is now added, and an excess of NaOH. If a precipitate is formed, it is filtered, washed, and tested for nickel with the borax bead. Concentrate the filtrate to a small volume. Test solution with borax bead for Co. (E) The filtrate from the sulphides of nickel and cobalt is boiled with a few drops of HNO 3 to oxidize the iron and is fil- tered if necessary ; any precipitate which may be formed is dis- carded since it is sulphur. (Iron is tested for in the filtrate by taking a few drops and adding a solution of ammonium sulpho- cyanate. A deep red color indicates iron. A light red color shows that iron is present only in traces.) Add Na 2 CO 3 to the solution with stirring until a permanent precipitate just com- mences to form. Dissolve with a drop of dilute HC1. Add about two drops of acetic acid and an excess of sodium acetate, 2 and boil until, on standing, the solution above the precipitate is clear. 3 Filter hot. The precipitate contains the hydroxides of any iron, chromium, and aluminium which may be present ; to be determined by (F). The filtrate contains any manganese or zinc which may be present ; to be determined by (K). (F) The precipitate is dissolved with dilute HC1, treated with 1 A large excess of KCN tends to prevent precipitation of nickel, should it be present. 2 Five cubic centimeters may be added and, if a precipitate is formed upon boiling, more may be added so long as further addition produces a turbidity of the clear liquid. If the first addition produces no precipitate upon boiling, further addition is unnecessary. 3 The hydrolysis of the acetates of Al, Cr, and Fe is analogous to that of the carbonates and sulphides of the same metals. Cf. reactions 157, 167, 169, 175, and 176. See also p. 36. 90 QUALITATIVE ANALYSIS an excess of NaOH, boiled, and filtered. The precipitate may be a mixture of iron and chromium hydroxides. The filtrate may contain aluminium. (G) Fuse on a platinum foil with Na 2 CO 3 and KNO 3 ; a yellow color indicates chromium. Boil with water, filter, and to the filtrate add a few drops of acetic acid, to acid reaction, and lead acetate. Yellow lead chromate is formed if chromium is present. Alternative test for chromium : fuse the precipitate from (F) with sodium peroxide in a nickel, iron, or silver vessel, dissolve in water, acidify with acetic acid, and test as before. (See also reaction 170 and footnote.) (H) The residue from (G) insoluble in water is dissolved in hydrochloric acid, and tested for iron with ammonium sulpho- cyanate. (I) The nitrate from the NaOH solution (F) is acidified with HC1, an excess of NH 4 OH added, heated, and allowed to stand a few minutes. A white flocculent precipitate shows presence of aluminium. 1 (K) The filtrate (E) may contain zinc and manganese. Acidify with acetic acid and conduct H 2 S into the boiling solution. White zinc sulphide is precipitated. Filter, test the precipitate for zinc with dilute HC1. If soluble, the presence of zinc is shown. The filtrate from the ZnS precipitate is evaporated to dryness, and divided into two portions. To one add two drops of cone. H 2 SO 4 and heat until white fumes of SO 3 are given off. Cool, add dilute H 2 SO 4 , and transfer to a test tube, and fill half full with dilute H 2 SO 4 . Add PbO 2 , and heat to boiling, and allow to stand until settled. A purple or red solution indicates manganese. The other portion is mixed with NagCOg and KNO 3 and heated on a platinum foil. A green color indicates manganese. 1 A flocculent precipitate may be due to silica derived from impure sodium hydroxide. The test may be confirmed by filtration and solution of the precipitate in dilute hydro- chloric acid. METAL ANALYSIS 91 After having completed the analysis of the mixture proceed to the analysis of unknowns Nos. 7 and 8. These being reported, prepare and record the following : EXERCISES I . What reactions are com mon to all the members of this group ? II. Devise another method for the separation of nickel and cobalt. III. Give the name and composition of at least one. ore of each of the members of this group. IV. For what purposes are the following substances useful, aside from analytical work : ferric chloride, potassium ferro- cyanide, Prussian blue, potassium chromate, potassium alum, potassium permanganate, pyrolusite, zinc oxide, cobalt nitrate ? V. What is steel ? Cast iron ? Wrought iron ? Spiegel- eisen ? VI. What are the chief uses of each of the metals of this group ? VII. Why cannot you get a test for iron in K 4 Fe(CN) 6 by the methods given for this group ? VIII. Why are the carbonates, sulphides, and acetates of ferric iron, aluminium, and chromium decomposed by water ? IX. Why does ammonium hydroxide fail to precipitate man- ganese when ammonium salts are present? X. Why must organic compounds be removed before this group can be analyzed. Group IV The Ammonium Carbonate Group Barittm, Strontium, and Calcium GENERAL STATEMENT. The carbonates of this group are insoluble in water and in solutions of ammonium salts. Magnesium carbonate is insoluble in water, but soluble in a 92 QUALITATIVE ANALYSIS solution of ammonium chloride, and is placed in Group V (the soluble group). BARIUM : Solution for Reactions, Barium Chloride. 231. HC1, H 2 S, NH 4 OH, and (NH 4 ) 2 S produce no precipitate. 232. KOH or NaOH precipitates from concentrated solutions white voluminous barium hydroxide, Ba(OH) 2 ; insoluble in ex- cess of reagent, but soluble in water. 233. (NH 4 ) 2 CO 3 or Na 2 CO 3 precipitates barium carbonate, BaCO 3 ; soluble in acids, even acetic acid. 234. Na 2 HPO 4 precipitates acid barium phosphate, BaHPO 4 , easily soluble in dilute HC1 or HNO 3 . 235. K 2 CrO 4 precipitates yellow barium chromate, BaCrO 4 ; soluble in HC1; insoluble in acetic acid. 236. (NH 4 ) 2 C 2 O 4 precipitates white barium oxalate, BaC 2 O 4 ; soluble in acetic acid. 237. H 2 SiF 6 precipitates white barium fluosilicate, BaSiF 6 ; insoluble in alcohol. 238. Sulphuric acid or soluble sulphates precipitate white barium sulphate, BaSO 4 ; insoluble in acids. 239. CaSO 4 or SrSO 4 precipitates BaSO 4 . 240. Flame Test. 1 Placed upon a platinum wire and intro- duced into a Bunsen flame, a yellowish green coloration is pro- duced. STRONTIUM : Solution for Reactions, Strontium Chloride. 241. HC1, H 2 S, NH 4 OH, (NH 4 ) 2 S produce no precipitate. 242. KOH or NaOH precipitates from concentrated solutions white voluminous Sr(OH) 2 , resembling Ba(OH) 2 (q.v.\ but less soluble in water. 1 For the flame test the platinum wire should be cleaned by dipping in a solution of HC1 and heating repeatedly until no color is given to the flame. The wire is then dipped into the salt and introduced into the flame. If the color does not show immediately, dip into HC1 and try again. Some salts do not give a good flame test until after being treated with HC1 on the platinum wire. (Why ?) METAL ANALYSIS 93 243. (NH^COg or Na 2 CO 3 precipitates white strontium car- bonate, SrCOg, resembling BaCO 3 ; easily soluble in acids, even acetic. 244. Na 2 HPO 4 precipitates white acid strontium phosphate, SrHPO 4 ; easily soluble in dilute HNO 3 or HC1. 245. K 2 CrO 4 precipitates only in concentrated solutions yellow strontium chromate. The presence of acetic acid prevents the precipitation. 246. ( N H 4 )2C 2 O 4 precipitates white strontium oxalate, SrC 2 O 4 ; only slightly soluble in acetic acid. 247. H 2 SiF 6 produces no precipitate even with the addition of alcohol. 248. H 2 SO 4 precipitates white strontium sulphate, SrSO 4 ; slightly soluble in water and hence does not appear at once in dilute solutions. More soluble than barium sulphate. 249. CaSO 4 precipitates strontium sulphate. 250. Flame Test. Placed upon a platinum wire and intro- duced into a Bunsen flame, a brilliant crimson color is produced. CALCIUM : Solution for Reactions, Calcium Chloride. 251. HC1, H 2 S, NH 4 OH, (NH^S produce no precipitate. 252. KOH or NaOH precipitates from sufficiently concen- trated solutions CaCOH^; less soluble than S^OH^ (q.v.). 253. (NH^COg or Na^COg precipitates white calcium car- bonate, CaCO 3 ; slightly soluble in an excess of concentrated solution of Na^Og ; easily soluble in acids, even acetic. 254. Na 2 HPO 4 precipitates white acid calcium phosphate, CaHPO 4 ; easily soluble in dilute HC1 or HNO 3 . 255. K 2 CrO 4 produces no precipitate. 256. (NH 4 )2C 2 O 4 precipitates white calcium oxalate, CaC 2 O 4 ; insoluble in acetic acid; soluble in HC1 or HNO 3 . 257. H 2 SiF 6 produces no precipitate, even with addition of alcohol. 94 QUALITATIVE ANALYSIS 258. H 2 SO 4 precipitates from strong solutions calcium sul- phate, CaSO 4 ; more soluble in water than BaSO 4 or SrSO 4 . 259. Flame Test. Brought upon a platinum wire and intro- duced into a flame, compounds of calcium produce a brick-red flame. Compare with strontium. Analysis, Group IV Make a mixture of salts of the metals of this group. Dissolve in water with the addition of a little HC1, if necessary, and iden- tify by the following scheme for analysis : (A) Ammonium chloride, ammonium hydroxide, and ammo- nium carbonate are added and the solution warmed for about ten minutes. Filter and wash with hot water. The filtrate may con- tain members of Group V. A small portion of the precipitate should be reserved for flame tests. Dissolve the remainder in dilute acetic acid. (B) To a small portion of the acetic acid solution add K 2 Cr 2 O 7 . If no precipitate forms, proceed at once with the remainder of the solution to (E) and (F). If a precipitate forms, treat the whole solution with dichromate and filter. Wash the precipitate and test by (C). (C) Dissolve the precipitate by warming in dilute HC1 and add dilute H 2 SO 4 . A precipitate indicates barium. The pre- cipitate will appear yellow until filtered and washed. (D) The filtrate from (B) is made alkaline with NH 4 OH and reprecipitated with (NH 4 ) 2 CO 3 . If no precipitate forms, Ca and Sr are both absent. If a precipitate forms, filter and wash until white. (The precipitate may be tested for strontium and calcium by the flame test. See 250 and 259.) Dissolve the pre- cipitate in acetic acid and divide into two portions and analyze by (E) and (F). (E) To one portion of the solution from (D) add a saturated solution of CaSO 4 . Stir vigorously and allow to stand for METAL ANALYSIS 95 some time. A precipitate indicates Sr. Confirm by flame test. (F) To the other portion of the solution, add a solution of NagSC^ until precipitation is complete. Filter, and to the ni- trate add a solution of (NH 4 ^C 2 O 4 . A precipitate indicates the presence of calcium. Confirm by flame test. If no stron- tium is present, the treatment with NagSC^ may be omitted. After having completed the analysis of a known mixture, pro- ceed to the analysis of unknowns Nos. 9 and 10. These being reported, prepare and record the following : EXERCISES I. What reactions are common to all the members of this group ? II. Prepare a list of the more important minerals containing the metals of this group. III. How is lime made ? What is limewater ? Baryta water? Bleaching powder ? Superphosphate of lime ? Plaster of Paris ? IV. How are the peroxides of barium, calcium, strontium, and magnesium prepared ? What are the uses of each ? V. What is the composition of window glass ? Hard glass ? Jena glass ? VI. What are the commercial uses of heavy spar ? VII. What is the composition of plaster? VIII. What use is made of the polysulphide of calcium and how is it prepared ? Group V The Soluble Group Sodium, Potassium, Ammonium (the Alkalies}, and Magnesium GENERAL STATEMENT. As there is no reagent which will precipitate all the members of this group, it is known as the soluble group. 96 QUALITATIVE ANALYSIS SODIUM: Solution for Reactions, Sodium Chloride. 260. K 2 H 2 Sb 2 O 7 precipitates in fairly concentrated neutral or weakly alkaline solution, Na 2 H 2 Sb 2 O 7 . Rubbing with a glass rod facilitates precipitation. In acid solution, H 2 Sb 2 O 7 is pre- cipitated. 261. K 3 A1F 6 produces a flocculent precipitate, Na 3 AlF 6 . 1 The same precipitate is also produced by magnesium salts, but not by potassium and ammonium salts. 262. Flame Test. When a sodium salt is brought upon a platinum wire and introduced into a Bunsen flame, an intense yellow coloration is produced, which is not visible through a blue glass. 2 POTASSIUM : Solution for Reactions, Potassium Chloride. 263. PtCl 4 precipitates, except in very dilute solutions, yellow crystalline potassium chlorplatinate, K 2 PtCl 6 ; insoluble in alco- hol. 264. NaHC 4 H 4 O 6 or H 2 C 4 H 4 O 6 precipitates white acid potassium tartrate, KHC 4 H 4 O 6 . Agitating, or scratching the walls of the containing vessel with a glass rod, facilitates the precipitation. 265. Na 3 Co(NO 2 ) 6 precipitates from neutral concentrated solutions potassio-cobaltic nitrite, K 3 Co(NO 2 ) 6 , in the form of a yellow powder. Precipitation takes place slowly from dilute solutions, but is hastened by warming gently. If the solution is alkaline, it should be made slightly acid with acetic acid before applying the test. If acid, it should be made slightly alkaline with sodium carbonate and then slightly acidified with acetic acid. 266. Flame Test. Placed upon platinum wire and heated in a Bunsen flame, compounds of potassium color the flame violet. 1 The ratio of Na to Al is i.i : i. 2 The color is not visible, but the student should note the distortion of the flame by the wire and compare this with the appearance of the flame in reaction 266. METAL ANALYSIS 97 A small amount of sodium salts obscures this test, but if the flame is observed through a cobalt glass, the sodium rays are cut off and the potassium appears reddish violet. AMMONIUM : Solution for Reactions, Ammonium Chloride. 267. PtCl 4 precipitates in concentrated solutions yellow crys- talline ammonium chlorplatinate, (NH^PtClg, insoluble in alcohol. 268. NaHC 4 H 4 O 6 or H 2 C 4 H 4 O 6 precipitates white acid am- monium tartrate, NH 4 HC 4 H 4 O 6 , from quite concentrated solu- tions upon standing. The precipitation is hastened by shaking, or scratching the walls of the containing vessel with a glass rod. 269. Na 3 Co(NO 2 ) 6 precipitates from acetic acid solution ammonium cobaltic nitrite, (NH 4 ) 3 Co(NO 2 ) 6 , in the form of a yellow powder. 270. NaOH or KOH in excess liberates ammonia, NH 3 . Recognized by (a) odor, (b) action of the gas on moist litmus, changing red litmus to blue, (c) fumes produced when a glass rod moistened with HC1 is held in escaping gas. Action is hastened by warming. 271. Ca(OH)2 liberates ammonia gas. If solid calcium oxide, CaO, and an ammonium salt are mixed in a beaker, the test for free ammonia may be made by placing over it a watch glass, upon the under side of which is placed a piece of moist red litmus. 272. Nessler's Reagent 1 is a most delicate test 2 for ammonia. It gives a distinct yellow coloration to solution containing minute quantities of ammonium compounds. If present in large quan- tities, a brown precipitate is produced. 273. All ammonium compounds are volatilized by heat. 274. Flame Test. Placed upon a platinum wire and intro- duced into a Bunsen flame, ammonium compounds give a reddish * See Appendix. 2 The test is so very delicate as to be unsuited to ordinary qualitative operations. 98 QUALITATIVE ANALYSIS violet flame when viewed through a cobalt glass. Compare with the potassium flame. MAGNESIUM : Solution for Reactions, Magnesium Sulphate. 275. PtCl 4 , Na 3 Co(NO 2 ) 6 , H 2 C 4 H 4 O 6 , NaHC 4 H 4 O 6 and Nes- sler's solution give no tests with magnesium salts. 276. NH 4 OH precipitates white magnesium hydroxide, Mg(OH) 2 ; soluble in presence of ammonium salts. 277. (NH 4 ) 2 CO 3 precipitates white magnesium carbonate, MgCO 3 ; from concentrated solution ; soluble in presence of am- monium salts. 278. KOH or NaOH precipitates white magnesium hydrox- ide, Mg(OH) 2 ; almost insoluble in water. 279. Na 2 HPO 4 in presence of NH 4 OH and NH 4 C1 pre- cipitates white crystalline magnesium ammonium phosphate, MgNH 4 PO 4 . The precipitation is slow in dilute solutions ; hastened by warming and agitation. 280. H 2 SO 4 , H 2 SiF 6 , and (NH 4 ) 2 C 2 O 4 produce no precipitate. Analysis, Group V Make a mixture of Na, K, NH 4 , and Mg salts. Dissolve a portion of the mixture and identify by the following scheme for analysis : (A) To a portion of the solution, add a small excess of NH 4 OH, and some NH 4 C1 in case it is not already known to be present, and Na 2 HPO 4 . Thoroughly agitate, or scratch with a glass rod, if a precipitate does not form at once, and allow to stand for some hours. A white crystalline precipitate indicates magnesium. 1 1 In case an unknown material is being analyzed, it may contain other cations besides those of this group. If this be true, they may be removed by treatment with ammonium carbonate, ammonium hydroxide, and ammonium chloride, and the filtrate used as directed above. METAL ANALYSIS 99 (B) A second portion of the solution is evaporated to dryness and heated to a low red heat in a porcelain dish until white fumes are no longer observed. Care must be taken to heat all portions of the dish. When cool, dissolve in as small an amount of water as possible, add a few drops of HC1, and filter if necessary. Test the filtrate for potassium with Na 3 Co(NO 2 ) 6 or PtCl 4 , and with the flame test, using cobalt glass. (C) A third portion of the solution may be treated with K 3 A1F 6 . A precipitate indicates sodium. See 261. The flame test may be used on some of the solid material obtained in (B) or on the origi- nal dry substance. A persistent yellow flame indicates sodium. 1 (D) A portion of the original material is treated with lime or a solution of KOH and the ammonia sought by (a) odor,() litmus, (c) action on a rod moistened with HC1. See 270 and 271. After completing the analysis of a known mixture, the stu- dent will proceed to the analysis of samples Nos. 1 1 and 12, after which the following exercises will be prepared and recorded : EXERCISES I. Devise another method for the analysis of this group. II. What are the more important salts of ammonium? of magnesium ? of potassium ? of sodium ? III. What are the more important compounds of these metals found in nature ? IV. Describe the Solvay process for the manufacture of soda ; the Leblanc process. V. What is Glauber's salt ? Epsom salt ? Saleratus ? Chili saltpetre ? Sal ammoniac ? Washing soda ? Gunpowder ? VI. Describe briefly the usual impurities found in common salt, and its purification. 1 The flame test for sodium is so very delicate, and the presence of sodium in small quantities is so frequent, that in reporting unknowns, considerable discrimination must be used. 100 QUALITATIVE ANALYSIS VII. What substances are present in a solution of ammonium hydroxide? See Smith's General Chemistry, p. 338 and p. 565. VIII. Write a detailed explanation of the effect of ammonium salts on the precipitation of magnesium by ammonium hydroxide. Smith's General Chemistry, p. 644. IX. What substances are present in a solution of pure sodium chloride ? PART III ACID ANALYSIS GENERAL CONSIDERATIONS. It is usual to make use of the negative ions of acids or of salts in the detection of the metalloids and non-metals. If the non-metal does not already exist in a negative ion, it is converted into such and its identity determined by the characteristic appearance or behavior. Acid analysis is, then, the detection of the non-metals. Also negative groups may at times involve elements which also reveal themselves as cations during metal analysis, e.g. arsenic, man- ganese, chromium, etc. Just as in the case of the analysis for cations, it is essential to acid analysis that the substance be brought into solution in order to facilitate identification. The appearance of this solu- tion, or the absence of certain metals, may at times render the search for certain anions unnecessary, e.g. a colorless solution is sufficient evidence of the absence of chromates or of per- manganates, while the failure to find arsenic in the course of the metal analysis renders search for arsenic and arsenous acids futile. Again the presence of lead or of barium in a soluble substance precludes the possibility of the presence of sulphuric acid. In order to assist in forming such conclusions, as well as for other purposes, the following rules of solubility will be found useful. NOTE. See also Table of Solubilities, Appendix. (i) All sodium, potassium, and ammonium salts are soluble in water, except the chlorplatinates and acid tartrates of ammo- 101 102 QUALITATIVE ANALYSIS nium and potassium, the silicofluoride of sodium and potassium, and the aluminofluoride of sodium. (2) All chlorates, nitrates, and acetates are soluble in water (basic salts excepted). (3) All carbonates, phosphates, borates, oxalates, and arse- nates, except those of the alkalies, are insoluble in water, but are soluble in dilute acids. (4) All chlorides, bromides, and iodides are soluble in water, except those of silver, lead, and mercury (ous), and mercuric iodide. (5) All sulphates are soluble except those of barium, stron- tium, and lead. (Calcium sulphate and silver sulphate are but slightly soluble.) (6) All hydroxides are insoluble except those of the alkalies and alkaline earths. It may further be remembered that the nature of the sub- stance limits the search for acids, e.g. an alloy can contain no acids and but a limited number of metals. An alloy may, however, contain non-metallic elements such as C, S, Si, P, etc. A mineral insoluble in water contains no organic acids or cyanogen compounds. It likewise contains no nitrates, chlo- rates, etc. In the qualitative analysis for acids there is no absolute sepa- ration into groups, and subsequent separation of each group, either possible or desirable. Nevertheless the acids may be advantageously grouped and group reagents suggested, so that the presence of representatives of each group may be affirmed, or denied, by single tests. Other than in this respect the iden- tification of the acids is by means of individual tests. In these notes only the acids more frequently encountered will be dis- cussed. The rare elements which form acids are briefly dis- cussed in Part IV. The following grouping is suggested : ACID ANALYSIS 103 Group I Acids the anions of which are precipitated from salts by dilute nitric acid or are decomposed by nitric acid : H 2 SiO 3 , H 2 S.HNO 2 , H 2 S 2 O 3 , H 2 SO 3 .H 2 CO 3 , HC1O. Group II Acids the anions of which are precipitated in dilute nitric acid solution by silver nitrate : HC1, HBr, HI, HCN, H 3 Fe(CN) 6 . H 4 Fe(CN) 6 . Group III Acids the anions of which are precipitated by barium chloride from neutral solutions: H 2 SO 4 . H 2 CrO 4 , H 3 PO 4 , H 3 BO 3 , H 2 C 2 4 , H 2 C 4 6 , H 3 As0 3 , H 3 AsO 4 , HF, H 4 SiO 4 . Group IV Acids the anions of which are not precipitated by silver nitrate or by barium chloride: HNO 3 , HC1O 3 , HC 2 H 3 O 2 , HMnO 4 . Group V Organic acids. Almost all of the organic acids give a " burnt " odor on heating to redness and give a black, charred residue. None of these will be considered in this connection save tartaric acid, oxalic acid, and acetic acid. These have been included in groups three and four. Oxalic acid neither chars nor gives a burnt odor on heating. Preliminary Reactions Students are advised to perform all preliminary reactions with known solutions and to write equations representing the same before proceeding to the analysis of any unknown sub- stance. For this purpose, it is best to use solutions of the alkali salts of the acids. 104 QUALITATIVE ANALYSIS Group I H 2 SiO 3 , H 2 S, HNO 2 , H 2 S 2 O 3 , H 2 SO 3 , H 2 CO 3 , HC1O. SILICIC ACID Silicic acid is precipitated from silicates soluble in water, under most circumstances, by acids as a jelly like transparent mass which always leaves an insoluble residue of silicon dioxide when evaporated to dryness and heated to no degrees centi- grade. This residue may be tested for silicon by reaction 2 or 3, p. 115. For the behavior of insoluble silicates, see Group III. HYDROSULPHURIC ACID 1. Dilute hydrochloric acid evolves hydrogen sulphide from most sulphides and will darken a bit of filter paper moistened with an alkaline solution of lead acetate. (See Appendix.) If the sulphide is insoluble, it will give the above test when the powdered solid is mixed with zinc dust and then treated with acid. 2. Silver nitrate precipitates silver sulphide from hydrogen sulphide or soluble sulphides. 3. Sodium nitroprusside, in alkaline solutions, gives a red- dish violet color with negative sulphur ions. 4. Concentrated sulphuric acid decomposes all sulphides with liberation of sulphur dioxide and sulphur. 5. Antimony salts precipitate red antimony trisulphide. SULPHUROUS AND THIOSULPHURIC ACIDS I. Dilute sulphuric acid or hydrochloric acid evolves sulphur dioxide from all sulphites in the cold. Thiosulphates when acidified also liberate sulphur dioxide as well as a milky pre- cipitate of sulphur. ACID ANALYSIS 105 2. Acid solutions of chromates and permanganates are de- composed by sulphites and sulphur dioxide ; thiosulphates also liberate sulphur. 3. Iodine solutions are decolorized by sulphurous acid. 4. Strontium chloride produces a precipitate in neutral solu- tions of sulphites, but not in thiosulphates unless very concen- trated. (Strontium thiosulphate is soluble in 3.7 parts water.) A small precipitate is nearly always produced in solutions of commercial salts, since they nearly always contain some sul- phate. 5. Ferric chloride produces a dark violet evanescent colora- tion with thiosulphates but not with sulphites. 6. On boiling with bromine water sulphites and thiosulphates are converted into sulphates, which are precipitated from acid solution by barium chloride. CARBONIC ACID 1. Dilute acids decompose all carbonates, most of them in the cold. When the gas evolved is passed into limewater or baryta water, a deposit of carbonate is obtained. 2. Concentrated acids behave in like manner, but more vigor- ously. (See also organic acids.) NITROUS ACID 1. Ferrous sulphate and concentrated sulphuric acid give nitrosyl ferrous sulphate, as in case of nitrates (q.v.). 2. Dilute acids react on nitrites with the production of brown vapors. 3. Indigo and permanganates are decolorized by nitrites in acid solutions. 4. Potassium iodide treated with nitrites in acid solution lib- erates iodine ; detected by shaking with chloroform or carbon disulphide. 106 QUALITATIVE ANALYSIS HYPOCHLOROUS ACID 1. All acids, even carbonic, liberate chlorine from hypochlo- rites. All hypochlorites are soluble in water. 2. lodo-starch paper is colored blue by hypochlorites in weakly alkaline solutions by reason of the liberation of iodine. Analysis of Group I Analysis of Nos. i and 2 of a set of unknowns is to be un- dertaken, and the outline given below should be followed. Treat some of the solid unknown with dilute sulphuric acid. An evolution of gas indicates : HCN Noted by odor of peach blossoms (caution !). Belongs in Group II. NO 2 Indicates nitrites ; noted by color, odor, solubility. H 2 S - Noted by odor and by lead acetate paper. SO 2 With separation of sulphur from thiosulphates. Noted by odor and reduction of permanganic acid. SO 2 Without separation of sulphur from sulphites. CO 2 From carbonates ; noted by turbidity produced when gas is passed into limewater. C1 2 From hypochlorites ; noted by odor and other tests. On boiling with dilute acid, if no reaction has occurred in the cold, or when action has ceased, there may be formed : CO 2 From certain carbonates ; noted as above. O 2 From peroxides ; noted by support of combustion. May be formed in the cold also. HCN From ferro- or ferricyanides. See Group II. HC 2 H 3 O 2 From acetates; noted by odor of vinegar. If the residue is colored, it may contain certain sulphides, and these may be sought by reaction I or 4, page 104. If any of the indications are observed, separate tests are ACID ANALYSIS 1 07 made for all suspected acids, according to the preliminary tests. See also exercises. A portion of the unknown is dissolved in water l and treated with hydrochloric acid, until completely decomposed, evaporated to dryness and heated gently, and redissolved as completely as possible in water with a few drops of HC1. An insoluble residue is SiO 2 and may be confirmed by 3, p. 115. EXERCISES I. If sulphides and carbonates are present together, how are they to be distinguished ? II. If on testing with lead acetate paper for sulphides a yel- low color is produced upon the paper, what is present ? III. If the gases from a sample treated with sulphuric acid are passed into acidified bichromate solution with the formation of a green color, what gases may be present ? IV. If sulphites and sulphides were both present in the same sample, what substance would be formed on acidification ? V. Make a list of the important sulphide minerals. VI. What cyanides are commercially important? Give methods of preparation of four different cyanides. Group II HC1, HBr, HI, HCN, H 3 Fe(CN) 6 , H 4 Fe(CN) 6 . Preliminary Reactions HYDROCHLORIC ACID 1. Dilute sulphuric acid produces no evidence of reaction. 2. Concentrated sulphuric acid on the dry salt gives colorless HC1, fuming in air. 1 The unknowns consist wholly of salts which are soluble in water. If samples which are not soluble are encountered, they are prepared for examination as described on p. no. The nitrate will contain sodium silicate if insoluble metasilicates were originally present. IO8 QUALITATIVE ANALYSIS 3. Silver nitrate produces a curdy white precipitate of silver chloride; soluble in ammonium hydroxide, potassium cyanide, sodium thiosulphate, and ammonium ' sesquicarbonate.' (See Appendix.) 4. The dry salt treated with concentrated sulphuric acid and potassium dichromate gives a brownish vapor of chromium oxy- chloride. If this vapor is dissolved in water, it produces chromic acid. The chromic acid may be detected by the hydrogen per- oxide reaction. See chromic acid. (Distinction from HI and HBr.) HYDROBROMIC ACID 1. Concentrated sulphuric acid on the dry salt gives rise to colored. gases (distinction from HC1) which fume in the air arid do not render water held on a glass rod turbid (distinction from HF). 2. Silver nitrate produces a curdy yellowish precipitate ; sol- uble in ammonium hydroxide, potassium cyanide, sodium thio- sulphate, but not in ammonium 'sesquicarbonate.' 3. Chlorine water liberates bromine ; soluble in chloroform or carbon disulphide, changed by excess to yellowish BrCl (distinc- tion from HI). 4. Potassium dichromate liberates no bromine in dilute sul- phuric acid solution (difference from HI). 5. Potassium dichromate and concentrated sulphuric acid on the dry salt liberate bromine. 6. Nitrites in acid solution, if dilute, free no bromine in the cold. HYDRIODIC ACID 1. Concentrated sulphuric acid reacts in the cold, liberating hydriodic acid and iodine and sulphur dioxide, hydrogen sul- phide, and sulphur. 2. Silver nitrate produces a curdy yellow precipitate ; soluble in potassium cyanide and sodium thiosulphate ; but only slightly ACID ANALYSIS 109 soluble in ammonium hydroxide and ammonium ' sesquicarbon- ate ' (difference from HC1 and HBr). 3. Lead salts precipitate yellow lead iodide. 4. Cupric salts produce a separation of cuprous iodide and iodine, made nearly white by sulphurous acid. 5. Nitrites in dilute acid solution liberate iodine. 6. Chlorine water liberates iodine which colors starch emul- sion blue or chloroform violet, in excess forms colorless solutions of iodates. 7. Potassium dichromate and both concentrated and dilute sulphuric acid liberate iodine; detected by chloroform or carbon disulphide. HYDROCYANIC ACID (PRUSSIC ACID) 1. Silver nitrate precipitates silver cyanide; soluble in excess of the alkali cyanide. 2. Ammonium sulphide boiled with alkali cyanides gives sulphocyanates. This solution, if made slightly acid, and heated until a drop gives no black precipitate with lead salts, will give a blood-red color with ferric chloride. 3. Addition of a small amount of ferrous sulphate and an excess of sodium hydroxide will, on heating, form potassium ferrocyanide. This solution, made acid with hydrochloric acid and treated with ferric chloride, leaves an insoluble precipitate of Prussian blue. 4. Sulphuric acid liberates prussic acid in the cold, which may be noted by the odor of peach blossoms (virulent poison ! caution !). HYDROFERROCYANIC ACID 1. Silver nitrate produces a white precipitate of silver ferro- cyanide. 2. Ferric salts produce a precipitate of Prussian blue. (See Treadwell and Hall, p. 100.) HO QUALITATIVE ANALYSIS 3. With exclusion of air, ferrous salts produce a white precipi- tate, which rapidly becomes blue by oxidation. HYDROFERRICYANIC ACID 1. Silver nitrate produces an orange precipitate of silver ferricyanide. 2. Ferric salts produce a brown coloration. 3. Ferrous salts produce a precipitate of Turnbull's blue. The analysis of samples 3 and 4 of set of unknowns is now undertaken, using the following scheme as a guide. Analysis of Group II For preparation of a solution of the anions, see p. 115. Acidify the solution with nitric acid, adding a slight excess, and treat with silver nitrate. If no precipitate is formed, the members of this group are not present. If a precipitate forms, a fresh portion of the solution is acidified with hydrochloric acid and a few drops of ferrous sulphate are added; an immediate blue precipitate indicates ferricyanic acid ; a white one, rapidly turning blue, indicates ferrocyanic acid. Each of these is then sought by other of the preliminary reactions. If either is present, in order to detect hydrocyanic acid some of the dry sample is placed in a porcelain dish, treated with dilute hydro- chloric acid, and covered with a watch glass, which is moistened with ammonium sulphide, and allowed to stand for some time. The cover is rinsed off with water and the rinsing acidified and treated with a drop or two of ferric chloride ; a blood-red color indicates hydrocyanic acid in the original substance. If any of the cyanogen compounds are present, 1 a larger portion of the solution is treated with an excess of nickel sulphate and 1 If ferricyanides or ferrocyanides are present, it must be observed, that on acidifying the solution with nitric acid, hydriodic acid is destroyed. In such case the presence of iodine or its salts may be detected on the dry substance by reaction I, page 108. ACID ANALYSIS III filtered ; the filtrate contains only the nickel salts of hydro- chloric, hydrobromic, and hydriodic acids, and sulphates. The nickel is removed by precipitation with sodium hydroxide and the filtrate is divided into two portions. To one portion is added a drop of chloroform and a drop of chlorine water and shaken ; a violet color indicates iodine. If present, remove by boiling with potassium dichromate and dilute sulphuric acid as long as fumes of iodine are developed, and test filtrate with chlorine water for bromide. (If very small amounts of iodides are present, they may be removed by an excess of chlorine ; and if bromine is present, the brown color will be imparted to the chloroform. If very large quantities of iodides are present, they may be removed by precipitation with copper sulphate and the test for bromides then made in the filtrate.) To the other portion silver nitrate is added and the precipitate is treated, after washing, with ammonium sesquicarbonate and boiled and filtered; the filtrate is acidified with nitric acid, and a white precipitate indicates chlorides. (The presence of chlorides may also be shown by reaction 4, p. 108.) The residue insoluble in ammonium sesquicarbonate is treated, on the filter, with dilute ammonium hydroxide and the filtrate acidified; a yellowish precipitate indicates bromides; an insoluble residue on the filter indicates iodides. A more delicate and in some respects more satisfactory method of separation of the halogens is to acidify the filtrate after the removal of the cyanides and nickel, and boil with ferric sulphate or ferric alum, an evolution of violet vapors indicating iodine. When this reaction is over, the solution is boiled with a crystal of potassium permanganate, and when the removal of the bromine is thus effected, the solution is filtered and diluted (to prevent precipitation of silver sulphate) and treated with silver nitrate, a white precipitate of silver chloride indicating the presence of chlorides. 112 QUALITATIVE ANALYSIS EXERCISES I. Make a list of the mineral chlorides which are com- mercially important, either for their own sake or for the sake of the substances prepared from them. II. In the analysis of the group by the method given, why is nickel removed, and why not use ammonium hydroxide for the purpose? See reaction No. 215. III. If the halides were present in the original substance as silver salts, devise or look up a method of separation. IV. Look up and make a synopsis of Remsen's theory of double salts of the halides. See Remsen's Advanced Course, p. 465. Group III Preliminary Reactions H 2 SO 4 , H 2 CrO 4 , H 3 AsO 4 , H 3 AsO 3 , H 3 PO 4 , H 3 BO 3 , HF, H 4 Si0 4 , H 2 Si0 3 , H 2 C 2 4 , H 2 C 4 H 4 6 SULPHURIC ACID 1. Silver nitrate precipitates silver sulphate from sufficiently concentrated solutions. (Solubility of silver sulphate, 0.58 g. per 100 c.c.) 2. Barium chloride precipitates barium sulphate ; practically insoluble in all acids. 3. Lead acetate precipitates lead sulphate ; soluble in con- centrated sulphuric acid, ammonium acetate, ammonium tar- trate. CHROMIC ACID 1. All chromates and dichromates are colored, and are reduced in acidified solution by hydrogen sulphide, sulphur dioxide, alcohol, and hydriodic acid, forming green solutions. 2. Silver nitrate precipitates silver chromate from neutral or slightly acid solutions. ACID ANALYSIS 113 3. Lead acetate precipitates lead chromate, in solutions acidified with acetic acid, in presence of ammonium acetate. 4. Hydrogen peroxide added to dilute cold solutions of chromic acid, or acidified chromates, produces a blue color, due to the formation of perchromic acid. This substance is ren- dered more stable by solution in ether. The more dilute, and the lower the temperature, the more easily is the reaction carried out. ARSENIC ACID 1 . Silver nitrate precipitates from neutral solutions chocolate- brown silver arsenate; soluble in acid and in ammonia (differ- ence in color from arsenous and phosphoric acids). 2. Hydrogen sulphide precipitates arsenic solutions only slowly. In hot, strongly acid solutions it precipitates a mixture of the tri- and pentasulphides. 3. Magnesia mixture precipitates magnesium ammonium arsenate, white and crystalline. 4. Ammonium molybdate in nitric acid solution (see ap- pendix) precipitates yellow arsenomolybdate. ARSENOUS ACID 1. Silver nitrate in neutral solutions precipitates yellow silver arsenite (difference in color from arsenic acid). 2. Hydrogen sulphide precipitates arsenous arsenic from acid solutions at once as the trisulphide. In the absence of other salts the arsenous sulphide may be colloidal. For colloidal so- lutions and their treatment see Ostwald's Scientific Founda- tions, p. 25. 3. Iodine solutions are decolorized by solutions alkaline with sodium carbonate, arsenates being formed. 4. Magnesia mixture and ammonium molybdate produce no precipitate. 114 QUALITATIVE ANALYSIS PHOSPHORIC ACID 1. Silver nitrate produces in neutral solution a yellow pre- cipitate of the orthophosphate ; soluble in nitric acid and in ammonium hydroxide. (The meta- and pyrophosphates are white. For other distinctions, see Treadwell and Hall.) 2. Hydrogen sulphide produces no visible effect on solutions of the acid or of its salts. 3. Magnesia mixture produces a precipitate of magnesium ammonium phosphate. 4. Ammonium molybdate on standing produces a quantita- tive precipitate of ammonium phosphomolybdate ; readily solu- ble in the alkalies. 5. Ferric chloride and sodium acetate produce a quantitative precipitate of ferric phosphate. (A large excess of ferric chlo- ride is to be avoided.) (Why ?) BORIC ACID (SOLUTION OF BORAX FOR REACTIONS) 1. Silver nitrate in neutral solutions precipitates silver meta- borate, changed on heating to silver oxide. Dilute solutions give the oxide directly. (Why ?) 2. Mercuric chloride produces a red precipitate. 3. Concentrated sulphuric acid and alcohol (methyl alcohol preferred), heated with borax and ignited, produce a green flame, due to the combustion of the volatile methyl orthoborate. 4. Turmeric paper, dipped in a solution of free boric acid and dried at not over 100, is turned rose-red, and the color is not removed by treatment with acids. Ammonia turns it green or blue. Borax acidified with hydrochloric acid produces the same result. (For discussion, see Treadwell and Hall, p. 311.) HYDROFLUORIC ACID (SODIUM FLUORIDE FOR REACTIONS) i. Barium chloride precipitates a flocculent, bulky precipitate from solutions of the soluble fluorides. ACID ANALYSIS 115 2. All fluorides when heated with concentrated sulphuric acid give rise to hydrofluoric acid, which acts on glass, or other sili- cates, if present, and forms silicon tetrafluoride, which in turn is decomposed by water with the formation of fluosilicic and ortho- silicic acids. This reaction is most easily carried out by placing the fluoride and acid in a test tube, and, while holding near the surface of the mixture a glass rod with a drop of water, warming gently. In the presence of a fluoride, a crust of silicic acid is formed on the water drop. SILICIC ACID (SEE AI/SO GROUP I) 1. See p. 104. 2. Silicates insoluble in water are converted into soluble so- dium silicate by fusion on platinum with sodium carbonate. 3. Silicates treated on platinum or lead with hydrofluoric acid are converted into silicon tetrafluoride, which reacts with a drop of water (using a platinum wire, with a loop to hold the water drop). See reaction 2 on HF. 4. Mineral silicates are usually, but not always, converted to soluble sodium silicate by boiling with sodium carbonate. 5. Barium salts precipitate crystalline barium silicate. OXALIC ACID (SEE GROUP V) TARTARIC ACID (SEE GROUP V) The analysis of samples 5 an d 6 of set of unknowns is now undertaken, using the following scheme as a guide. Analysis of Group III A solution is prepared by boiling the sample with a saturated solution of sodium carbonate (i gram of the sample to 15 c.c. of the solution) for fifteen minutes, replacing the water evaporated. The hot solution is diluted with an equal quantity of water and Il6 QUALITATIVE ANALYSIS filtered. 1 This solution is made slightly acid with hydrochloric acid and neutralized, drop by drop, with ammonium hydroxide, using litmus as an indicator. A portion is treated with a few drops of calcium and barium chloride solutions. (The calcium salt is added because the calcium salts of some of the group are more insoluble than the barium salts.) If no precipitate is formed, the members of this group are absent (except boric acid, the barium and calcium salts of which are fairly soluble). Test is made on the original solid material for boric acid by reaction 3, page 114, and for silicic acid by reaction 3, page 115. If a precipitate is formed, then tests for the other acids are made separately, using by preference reactions 2 and 3 for sul- phuric acid, reactions I and 4 for chromic acid (if the solution is colorless, chromic acid is absent), reactions 2 and 4 for arsenic acid, reactions 3 and 4 for arsenous acid. If either of these latter is present, it must be removed by hydrogen sulphide and the solution boiled to expel hydrogen sulphide before testing for phosphoric acid, using reactions 2, 3, and 4, p. 1 14. For oxalic and tartaric acids, see analysis of Group V. EXERCISES I. An insoluble residue nearly always remains after treatment with sodium carbonate. What substances may it contain ? II. In case the substance to be analyzed is soluble in water and the treatment with sodium carbonate is omitted, what sub- stances may be in the precipitate which are not representative of this group ? III. Give an outline of the manufacture of " superphosphate." IV. What forms of boron are found in nature and where ? 1 The residue usually consists wholly of carbonates and should dissolve completely in dilute HC1. Any portion which remains undissolved in the acid may be sulphates or silicates of difficultly decomposed minerals and may be fured with sodium carbonate in iron, or preferably platinum, and the aqueous solution of the fused mass tested for sulphates and silicates. ACID ANALYSIS Ii; V. In case of an alloy containing phosphorus or sulphur, what treatment is needed to detect them ? VI. Why is it necessary to acidify arsenic solutions in order to precipitate arsenic sulphide with hydrogen sulphide ? Group IV HN0 3 , HC10 3 , HMn0 4 , HC 2 H 3 O 2 Preliminary Reactions NITRIC ACID 1. Ferrous sulphate is oxidized by concentrated nitric acid to ferric salts, and nitric oxide is formed, which unites with an excess of ferrous sulphate to form a brown, unstable compound, riitrosyl ferrous sulphate, which decomposes on boiling. The simplest method of using the test is to add to the solution to be tested a few cubic centimeters of ferrous sulphate solution and pour down the side of the containing test tube 3-5 cubic centi- meters of concentrated sulphuric acid ; a brown ring will appear at the junction of the two liquids. If organic acids are present and are charred by the concentrated sulphuric acid, the color will not disappear on boiling. If the solution is acid, it is best to add a few cubic centimeters of the concentrated acid, cool, and then superimpose a layer of ferrous sulphate solution. Iodides, bromides, ferrocyanides, chromates, permanganates, and chlorates interfere with this test. 2. A few drops of solution of nitrates evaporated on a porce- lain dish just to dryness and gently warmed with a few drops of phenolsulphuric acid form picric acid, which, when cooled and made alkaline with ammonium hydroxide, gives an intense yellow color. 3. One cubic centimeter of brucine solution added to a mix- ture of i part of the sofution and 3 parts concentrated sulphuric acid produces a deep red color, changing to orange, then to yel- Il8 QUALITATIVE ANALYSIS low, and finally to greenish yellow. (Nitrous acid does not give this test.) 4. Iodine is not liberated from potassium iodide solutions by dilute nitric acid (see Nitrous Acid). CHLORIC ACID 1. Concentrated sulphuric acid decomposes chlorates with liberation of the greenish yellow chlorine dioxide (caution, ex- plosive) and chlorine. 2. Concentrated hydrochloric acid decomposes all chlorates with the liberation of chlorine (euchlorine, see Remsen, Ad- vanced Course, p. 1 2 1 ). 3. Reducing agents decompose chlorates in neutral, acid or alkaline solutions, e.g. nascent hydrogen. 4. Ferrous salts are oxidized by chlorates to ferric salts on boiling with dilute sulphuric acid. PERMANGANIC ACID 1. All permanganates are soluble in dilute acids and are re- duced to manganese dioxide in alkaline solution and to manga- nous salts in acid solution by reducing agents, SnCl 2 , H 2 SO 3 , etc. 2. Hydrogen peroxide in acid solution decolorizes permanga- nates with a lively evolution of oxygen. ACETIC ACID (See GROUP V) Analyses of samples 7 and 8 are now undertaken according to the following scheme. Analysis of Group IV If the aqueous solution is colored red by permanganic acid, it is decolorized by hydrogen peroxide and a portion evaporated to dryness on the water bath and tested for chlorates by re- action i. If chlorates are present, a second portion is treated ACID ANALYSIS 119 with sulphur dioxide while being boiled. After cooling, it is tested for nitric acid by reactions I, 2, or 3 for nitric acid. A third portion is acidified and treated with hydrogen sulphide until the oxidizing acids are removed; filtered and evaporated nearly to dryness and treated for acetic acid as in Group V. EXERCISES I. Devise a method for removing the acids which interfere with the ferrous sulphate test for nitric acid. II. If dry permanganates are treated with concentrated sul- phuric acid, what is formed, and why is the operation dangerous ? III. How is potassium permanganate made? IV. How are chlorates made ? V. Look up the history of euchlorine and state what is its bearing on the demonstration of the elementary nature of chlo- rine. See Alembic Club reprints, No. 13. Group V H 2 C 2 O 4 , H 2 C 4 H 4 O 6 , HC 2 H 3 O 2 , etc. Preliminary Reactions OXALIC ACID 1. Concentrated sulphuric acid heated with an oxalate evolves carbon monoxide and carbon dioxide. The former is detected by passing the mixed gases through baryta water or sodium hy- droxide and igniting the insoluble monoxide. 2. Calcium salts, even calcium sulphate, precipitate oxalates ; soluble in hydrochloric, but not in acetic acid. 3. Silver nitrate precipitates white curdy silver oxalate; soluble in ammonium hydroxide and dilute acids. 4. Heat decomposes the oxalates with but slight carbonization, some of them leaving carbonates of the metals and others the metals themselves. 120 QUALITATIVE ANALYSIS TARTARIC ACID 1. Ignition of the tartrates leaves a residue of carbon, gives an odor of burnt sugar, and leaves either a carbonate of the metals or the metals themselves. (See Treadwell and Hall, P- 3170 2. Concentrated sulphuric acid, when heated with tartrates, gives carbon monoxide, carbon dioxide, and sulphur dioxide. The gases may be passed through acidified bichromate, then through alkaline solution, and the issuing gas burned. 3. Barium chloride precipitates barium tartrate. ACETIC ACID 1. Concentrated sulphuric acid sets acetic acid free from its salts, and, if alcohol is added at the same time, ethyl acetate is produced, which may be noted by its " fruity " odor. 2. Ignition decomposes all acetates, leaving behind oxides, carbonates, or metals, with the evolution of combustible gases, but not always with the formation of free carbon. 3. Ferric chloride in neutral solutions produces a blood-red ferric acetate, decomposed by boiling into a reddish precipitate of basic ferric acetate. OTHER ORGANIC ACIDS Nearly all other organic acids char on ignition and leave carbonates of the alkali metals when such are present. These acids are not usually present in inorganic mixtures, and their discussion will not form a part of this treatment. No unknown samples are given for Group V, since the acids which are here presented are found also in Groups III and IV. When organic matter is present, the search for these may be conducted as given below. ACID ANALYSIS 121 Analysis of Group V A portion of the solution to be tested for members of this group may be treated with a solution of calcium sulphate, and a precipitate soluble in hydrochloric acid without evolution of carbon dioxide indicates the presence of oxalic acid. The re- mainder is evaporated to near dryness on the water bath, and a portion may be tested for tartaric acid by reaction 2. Another portion may be treated for acetic acid as in reaction I, or it may be made exactly neutral and treated with ferric chloride as in reaction 3. EXERCISES I. How is sugar of lead manufactured ? For what purpose is it used ? II. What is " cream of tartar," and how is it purified ? III. How is oxalic acid made commercially? IV. What is pyroligneous acid, and what substances may be obtained from it ? When the student is made somewhat familiar by a study of the preceding reactions and the twenty (or more, if desired) samples containing members of single groups are analyzed, he may be considered ready to undertake the complete analysis of unknown substances. This involves all the steps already given, and some modifications which are given in Part IV. The directions for group separations are given in abbreviated and tabular form for convenience in use. PART IV SYSTEMATIC ANALYSIS The complete analysis of a substance may be considered as involving three steps : 1. Preliminary examination. 2. Examination for metals (cations). 3. Examination for non-metals (anions). A preliminary examination is needed in order to proceed in- telligently to get the substance into solution. It also serves to give information which renders subsequent examination more accurate and interesting, and at times renders such analysis brief or even unnecessary. It is also of value where the pres- ence of some one element is to be determined and recourse to a complete analysis is not desirable. It is not essential that all the steps in the following outline be rigidly followed in every case when a complete examination is contemplated, but they are to be followed so far as to enable the worker to proceed in- telligently in the later work. The closed tube test is, however, never to be omitted, since it serves to detectahe presence of organic matter which must always be removed before a mineral analysis is undertaken. The sample is first powdered, and then small portions of the homogeneous "sample are taken for the separate tests. 1 As a general rule, it is wise to reserve at least one half of the sample. Of the other half, one third is reserved for the metal analysis, one third for the acid analysis, and the remaining portion used for the preliminary examination. 1 If the substance is a liquid or alloy, see p. 127. 122 SYSTEMATIC ANALYSIS 123 Closed Tube Test A hard glass test tube is to be used or a small tube sealed at one end. A small amount of the substance is placed in the clean and dry tube, and heated gently at first and then to the highest heat of the Bunsen flame. Careful observation is made of any visual phenomena, and the odor and combustibility of the issuing gases are tested from time to time ; the latter is best done by means of a match or glowing splinter. The phenomena which may occur and their interpretation are given below : PHENOMENA INDICATION 1 . Moisture without carbonization. 2. Carbonization with odor and com- bustible gases. 3. Sublimate. (a) Black, with garlic odor. (b) Black or gray, forming globules when rubbed, or forming a mirror. (c) Black, turning red when rubbed. (d) Black, with violet vapors. (e) Reddish brown vapors, cool- ing to a yellow solid. (/) Yellow to red. (g) White, easily volatile. (h) White, volatile with diffi- culty. 4. Gases. (a) Kindles a glowing splinter. () Brownish red, with charac- teristic odor. (c) Colorless, with character- istic odor. d Characteristic odor. Water. Organic matter. Arsenic. Mercury. Mercury sulphide. Iodine. , Sulphur or persulphides. Sulphur or arsenic sulphide. Arsenious oxide, calomel, corrosive sublimate, or ammonium salts. Antimony oxide. Oxygen from peroxides, chlorates, or nitrates. Nitrogen peroxide from nitrates or nitrites. Sulphur dioxide from sulphates, sul- phides, or sulphites. Ammonia from ammonium salts or a cyanide. 124 QUALITATIVE ANALYSIS PHENOMENA INDICATION (e) Characteristic odor, black- ening lead acetate paper. (/) Burning with a blue flame. (g) Causing turbidity in drop of limewater. ^ Colored characteristic. Hydrogen sulphide from moist sul- phides. Carbon monoxide from an oxalate or tartrate, or methane from an acetate. Carbon dioxide from carbonates or organic matter. Chlorine, bromine, iodine. 5. Non-volatile residue may change color ; if so, the indications are as follows : ORIGINAL COLOR COLOR HOT COLOR COLD INDICATION White Yellow White Zinc oxide White or yellow White or yellow White or yellow Yellow or brown Yellow Brownish red Brownish red Black Pale yellow Deep yellow Pale yellow Brownish red Tin oxide Lead oxide Bismuth oxide Iron oxide Yellow or red Green Green Chromium Pink, green, or blue Black Black Nickel, copper, cobalt 6. The substance melts and remains liquid with or without expulsion of vapors alkali salts. The Bead Test A bead of borax is made on a loop at the end of a straight platinum wire, and after cooling, is moistened with the tip of the tongue, and a very little of the solid substance to be tested is caused to adhere to the bead, and the whole introduced into the oxidizing flame of the Bunsen burner. After the color of the transparent bead is observed, it is brought into the reducing flame, and after some minutes the color is again observed. The results and their indications follow : OXIDIZING FLAME REDUCING FLAME INDICATION Blue Blue Cobalt Green Green Chromium Greenish blue Amethyst Brownish red Brownish yellow Red Colorless Gray Bottle-green Copper Manganese Nickel Iron SYSTEMATIC ANALYSIS 125 If the substance is heated in the oxidizing flame with the microcosmic bead and floats on the bead as an undissolved skeleton, it indicates silicon dioxide or silicates. Several of the rarer elements also give colored beads. (See Part V, p. 139-) The Flame Test A little of the substance is moistened with concentrated sul- phuric acid and introduced into the lower edge of a Bunsen flame by means of a platinum wire. If the flame is colored yellow, observe it through a cobalt glass. When the coloration of the flame has ceased, moisten the wire with concentrated hydro- chloric acid and heat again. This treatment results in the vola- tilization of the sulphates of the alkalies, while the less volatile salts of the alkaline earths are converted into the more volatile chlorides. The characteristic flame colorations are as follows (see also Part V, p. 139): COLOR INDICATION Yellow ; invisible through blue glass. Carmine ; violet through blue glass. Scarlet, masked by barium flame ; voilet through blue glass. Yellowish red ; greenish gray through blue glass. Yellowish green. Emerald green. Azure blue. Light blue. Violet ; purple through blue glass. Sodium. Lithium. Strontium. Calcium. Barium or borates. Copper. Copper chloride. Arsenic compounds. Potassium. The Charcoal Test a. A small portion of the substance is heated on a piece of charcoal before the blowpipe. If deflagration takes place it indi- cates a nitrate, nitrite, or chlorate, or some other substance rich in oxygen. If the substance melts and runs into the charcoal, it indi- cates an alkali salt. An evolution of a garlic odor indicates arsenic. 126 QUALITATIVE ANALYSIS b. The substance is mixed with twice its own amount of so- dium carbonate, and, after moistening slightly, is heated on char- coal before the blowpipe in the reducing flame. There may be obtained : PHENOMENA INDICATION A. Metal without incrustation : Malleable button. Gray particles. B. Metal with an incrustation : Brittle button. Malleable button. C. Incrustation without metal : White yellow when hot. White garlic odor. Brown. D. White infusible mass : Heated with cobalt nitrate blue. Heated with cobalt nitrate pink. No color change. E. Sulphur compounds are reduced to sulphides. The melted mass is scraped off and placed on a silver coin, with a drop of water. A brown stain. Gold, silver, tin, copper, lead. Iron, nickel, cobalt. Antimony with white sublimate. Bismuth with yellow sublimate. Lead with yellow sublimate. Zinc. Arsenic. Cadmium. Aluminium. Magnesium. Calcium, strontium, barium. Sulphur compounds. (For a very full discussion of the use of the blowpipe on charcoal, see Moses and Parsons's Mineralogy, p. 97 et seq.} These preliminary reactions will give enough information concerning the character of the sample to enable the analyst to proceed with the preparation of the solution for the wet analysis, if such is necessary. Preparation of the Sample We distinguish three cases : i. The sample is a liquid, i.e. a solution. SYSTEMATIC ANALYSIS 127 Evaporate a portion to dryness on a water bath. If a residue is obtained, test the solution for nitric acid and for organic matter. If found, evaporate to dryness and ignite and redis- solve in water, with addition of acid if necessary. Analyze ac- cording to the general method given on page 130. If no residue is found, test need only be made for acids, ammonia, and hydro- gen peroxide. 2. The sample is an alloy or metal. Break the sample into small pieces or roll it into thin sheets. Take a one-gram portion and treat with nitric acid (one part con- centrated HNO 3 to one part water). A. No apparent action takes place. Treat the sample with aqua regia and examine for the noble metals and aluminium. B. Complete solution takes place. Evaporate to near dryness and dilute with water and analyze as in the group separations for metals, p. 1 30 et seq. The only acids which need be looked for are silicic, phosphoric, and sulphuric. C. Complete solution does not take place. Evaporate to dry- ness as in B ; dilute with water, filter, and separate as in the fol- lowing table : RESIDUE FILTRATE SiO 2 , SnO 2 , Sb 2 O 5 , P 2 O 5 , Bi 2 O 3 , traces of Pb, Cu, Fe, etc. Wash with water, treat with am- monium sulphide, boil, and filter. Analyze as in B above. RESIDUE FILTRATE SiO 2 , Bi 2 S 3 , PbS CuS Dissolve in HC1 and filter. Add filtrate to ammonium sulphide solu- tion and examine residue for silicic acid. (NH 4 ) 3 SbS 4 , (NH 4 ) 3 P0 4 , (NH 4 ) 2 SnS 3 . Treat with HC1 and filter. Exam- ine filtrate for phosphoric acid. Ex- amine residue for Group II (metals). 128 QUALITATIVE ANALYSIS 3. The substance is nonmetallic. A. Solution for metal analysis. Treat very small portions of the finely powdered and thor- oughly mixed substance in a test tube, successively with water, dilute hydrochloric acid, concentrated hydrochloric acid, dilute nitric acid, concentrated nitric acid, and aqua regia until solu- tion is obtained. If complete solution is obtained, a larger quan- tity about one fourth gram of the substance is dissolved in the solvent found necessary, and, if nitric acid or aqua regia was used, is evaporated to almost complete dryness ; dilute with water and analyze as directed in the general analysis of the groups on page 1 30 et seq. If the treatment above outlined is not sufficient to effect solu- tion, the residue must be fused. 1 In the absence of the sub- stances described in a, b, c, etc., below, fusion may be performed as follows : Take about one half gram of the residue insoluble in acids and four to five times the quantity of sodium carbonate and fuse on platinum until a quiet fusion is obtained. Then boil with water and filter. The aqueous solution may be acidi- fied and added to the solution obtained by acids if no precipitate is formed; otherwise it must be analyzed separately. The residue insoluble in water is washed and treated with strong hydrochloric acid and this solution added to the one obtained by acids if no precipitate is -formed ; otherwise it is analyzed separately. a. Ignition with free access of air will dispose of sulphur and carbon. A test may be made by heating in a test tube with CuO and passing the gas evolved through very dilute acidified permanganate and into limewater. iThe ordinary substances insoluble in water and acids which may be found in the residue for fusion are C, S, BaSO 4 , SrSO 4 , CaSO 4 , PbSO 4 , AgCl, CaF 2 , SiO 2 , natural and ignited oxides of Al, Cr, Fe, and Sn, many silicates, silicon, and silicides. The pre- liminary examination will have indicated which of these may be expected to be present, and the treatment of the insoluble residue must be varied accordingly. SYSTEMATIC ANALYSIS 129 b. Silver chloride will be dissolved by a solution of potas- sium cyanide and reprecipitated by ammonium sulphide. c. Chromic iron and chromic oxide are but little affected by fusion with sodium carbonate, and may be brought into solution by fusion with an oxidizing flux, such as sodium carbonate and potassium nitrate, or sodium peroxide. In the latter case a platinum vessel must not be used, but the fusion is to be carried out in a nickel, iron, copper, or silver vessel. d. In case of alumina and difficultly soluble iron oxides, the best flux is acid potassium sulphate. e. In case of calcium fluoride, it is best to treat with sul- phuric acid and then deal with the sulphate. /. In case of stannic oxide or of silicon or silicides, it is best to fuse in a silver vessel with potassium hydroxide and then dissolve in water. B. Solution for acid analysis. It is impossible to prepare a solution suitable for all the acid tests. Group I tests are all made upon separate small portions of the solid substance. A solution of the remaining acids may be prepared by boiling the sample for fifteen minutes with a saturated solution of sodium carbonate and filtering as described in Part III, p. 115. The filtrate will contain the acid radicals of all the common acids, except where the original substance is barium sulphate and certain silicates. The group reagent tests are then made upon this solution, and each group found represented is analyzed as per directions in Part III. GENERAL ANALYSIS OF THE METALS Group I To the neutral or acid solution add dilute hydrochloric acid, with constant stirring, until a precipitate is no longer formed. Filter and wash twice with cold water. The nitrate may con- tain members of Groups II to V. The precipitate is analyzed for metals of Group I as indicated below : The precipitate is washed with boiling water until no more dissolves. Precipitate: AgCl, HgCl Filtrate: PbCL NH 4 OH. Divide into several portions. (#) Treat with H 2 SC>4 ; a white Residue: Hg and Filtrate : precipitate indicates lead. HgNH.Cl Ag(NH^) 2 Cl () Pass in H 2 S ; a blaclc pre- A black " residue Acidify with cipitate indicates lead. shows the pres- HNO r A white (c} Treat with K 2 CrO 4 ; a ence of mercury precipitate shows yellow precipitate indicates lead. (ous). the presence of Ag. Group II The nitrate from Group I is evaporated nearly to dryness, with frequent additions of small quantities of HC1 to expel any HNO 3 which may be present. Dilute with water and add HC1 to about one part in ten, heat to boiling, and pass in H 2 S slowly as long as a precipitate forms. * Filter. Cool, dilute, and pass 1 If arsenates are present, acidify strongly and pass H 2 S ; filter and evaporate filtrate nearly to dryness, dilute, and again pass H 2 S until the precipitation is complete. 130 S GENERAL ANALYSIS OF THE METALS 133 o-a III |ii ^It Is g {I g-2 o S -2 *0 ^-l*l a ll I M = 3 gf"S c? -3 Erf ' *0 *l|f 5 f a8 - !l4i!sll l gl 'n:^oE-3^=S OX:T) ' IsK.--^ els lfes !8ili^ ^^<:S i^:^ color shows the f manganese. _ v o u-s.s H-'IsJ 13 a.2 - S"S fi ii! U 'S O C - '* 'i J3 W 1) 3 ^ a"5 ^.s sj -o - -y < g E I . ^fi 2 d fill be Residue iS (bla oS (blac C S Tes a bora Bro ad shows t presence of nickel. Blue bead shows th presence of cobalt. N icke may be present. wit n b ^ dp itate : Ni ( OH Test with a borax be Brown bead shows t esence of nickel. JT s Co(CN} id with HCI o a small bul h a borax bea icates Co. te wi n Filtrate Make vaporat nd test Blue i 134 QUALITATIVE ANALYSIS Whether a precipitate forms or not, add to the solution while still warm (without filtering) (NH 4 ) 2 S. Filter at once and wash the precipitate with water containing a little (NH 4 ) 2 S. The filtrate l may contain metals of Group IV and V. The pre- cipitate contains all the metals of Group III which are present. (If phosphoric, silicic, boric, hydrofluoric, or tartaric acids are known to be present, use the following scheme for analysis.) A. If tartaric or oxalic acid is present, evaporate the solu- tion or filtrate from Group II to dryness and fuse with Na 2 CO 3 and KNO 3 , dissolve the fused mass in water, acidify with HC1, pass in H 2 S to reduce any chromic acid if present to a chromium salt, and treat as given below, B. B. If tartaric and oxalic acid are absent, boil the solution or filtrate from Group II to expel any H 2 S present. Add a few drops of HNO 3 and boil. If the solution becomes yellow, iron is probably present. Add a small quantity of NH 4 C1 and NH 4 OH to alkaline reaction and boil. If a precipitate forms, one or more of the metals Fe, Cr, or Al are present. If no precipitate forms, these are absent, and subsequent tests need not be applied. In either case add to the solution while still warm (without filtering) (NH 4 ) 2 S. Filter and wash with water containing a little (NH 4 ) 2 S. The precipitate is to be analyzed for Group III, p. I35. 2 The filtrate may contain metals of Groups IV and V. The precipitate may contain hydroxides of Al, Cr, and the sulphides of Fe, Co, Ni, Mn, and Zn, and phosphates and similar compounds of Ba, Ca, Sr, &$.d Mg. 1 If the filtrate is dark brown in color, due to some NiS remaining in solution, the presence of Ni is indicated. If this is the case, the filtrate should be acidified with acetic acid and boiled for some time filtered on a separate filter and tested directly for Ni with a borax bead. The filtrate contains members of Groups IV and V. See Ostwald, Scientific Foundations, p. 24. 2 If phosphoric, silicic, boric, or hydrofluoric acids are known to be absent, analyze by the short method on page 133. GENERAL ANALYSIS OF THE METALS 135 $2 1*81.5 ~ c/2 2 -5 - !l s -;!iiftl If -Sll- 8 BsCJ^ii TO trt -* ii ^^-^r^ .rtti ' 1=5 ^ x-^""^ " fl^fii" _ <= 411 Li* so a% Os, Rh, Ru, Ir, Se, Te, Ge, and (Mo) The elements listed in this group are all precipitated by H 2 S in acid solution. The first two are sometimes dis- cussed with the common metals because of their frequent use in the laboratory and in the arts. They are placed by us among the rare elements partly because of the infrequent use of the ordinary qualitative methods for their detection, fire assay methods being usually employed, and also because of the association in nature of platinum with several other ele- ments of the group. Selenium and tellurium properly are nonmetals, but their place here is secured by the formation of sulphides. Detailed methods of se'paration will not be given, but the reader is referred to Bottger's Qualitative Analyse, 2d Auflage, p. 473, for the separation of the platinum metals, or to Mylius and Dietz, Ber. d. deutsch. Chem. Ges., Vol. 31, p. 3187, and to Bottger, p. 480, for the remainder of the group, or Noyes and Bray, Jour. Am. Chem. Soc., Vol. 29, pp. 137-205. Gold. Gold is usually found native, though in small amounts, disseminated through quartz, sand, or other minerals. It is also 146 QUALITATIVE ANALYSIS found associated with silver, as a telluride. All the compounds are dissociated by heat, giving the free element. It is liberated from its salts through displacement by nearly all other metals. It forms alloys readily with many elements which are extremely useful in the arts, especially in coinage and in jewelry. It forms an amalgam with mercury, which is used as a means of collecting the gold from the minerals with which it is asso- ciated. It also unites readily with chlorine, and forms a double salt with potassium cyanide, both of which properties are made use of in metallurgical operations. The detection of its presence in minerals, as well as the estimation of its amount, is ordinarily performed by the fire assay method, for which consult works on assaying. It forms two series of salts, behaving as a univalent and as a trivalent metal. It is not attacked by acids, but if treated with aqua regia it dissolves, forming auric chloride, AuCl 3 , soluble in water. The solution may be used for the following reactions : 1. KOH or NaOH precipitates, if added slowly, reddish brown Au(OH) 3 , soluble in excess, forming an aurate. 2. NH 4 OH precipitates fulminating gold, A^NH^OH^ or AuN 2 H 3 3H 2 O. 3. H 2 S precipitates, in the cold, black gold sulphide, Au 2 S 2 , soluble in (NH 4 ) 2 S 4 with some difficulty; readily soluble in K 2 S, insoluble in acids ; soluble in aqua regia. In boiling solutions metallic gold is precipi- tated. 4. FeSO 4 precipitates metallic gold as a brown powder. 5. H 2 C 2 O 4 precipitates metallic gold, reaction hastened by warming. 6. SnCl 2 precipitates, from not too acid solutions, metallic gold as a reddish purple substance Purple of Cassius). THE RARE METALS 147 Platinum. Platinum occurs widely distributed, but in small amounts, chiefly in the native state, but usually alloyed with smaller quantities of palladium, osmium, iridium, rhodium, and ruthenium. It is not largely produced, and by reason of its varied application, is continually increasing in value. On account of its small chemical activity, it is used very extensively in labora- tories as wire, foil, dishes, stills, etc. It occludes gases readily, and in the finely divided condition is an excellent catalytic agent. As a laboratory reagent its salts are frequently used, and would be used more extensively were they less expensive. It finds application in the arts as the leading wires for electric light bulbs, since its coefficient of expansion and that of glass are nearly equal ; in the manufacture of standard electrical instruments and in photographic processes. It forms two series of salts in which it acts as a bivalent and as a tetravalent metal. Both as a biva- lent and tetravalent element, it forms a part of complex negative ions. The bivalent forms, especially, are unstable, and it is mainly with the tetravalent compounds that we have to deal. All the compounds, like those of gold, are decomposed by heat, leaving the metal. The element is rarely sought analytically in the wet way, being obtained from its ores by the fire assay method when its detection and estimation are desired. It is insoluble in acids, but may be obtained from associated gangue by digestion with aqua regia as chlorplatinic acid, H 2 PtCl 6 . The solution may be used for the following characteristic reactions : * 1. Potassium or ammonium salts precipitate the difficultly soluble chlorplatinates. See Potassium reactions, p. 96. (Used to separate platinum from gold.) 2. H 2 S precipitates slowly in cold but rapidly in hot solu- tions, brown PtS ; insoluble in acids ; soluble in aqua regia ; difficultly soluble in the polysulphides of the 1 For the preparation of chlorplatinic acid in pure condition, see Treadwell and Hall, P- 234. 148 QUALITATIVE ANALYSIS alkali metals and ammonium, from which it is re- precipitated on acidification. 3. FeSO 4 H 2 C 2 O 4 , and SnCl 2 do not precipitate metallic platinum. The platinum metals, Pd, Os, Rh, Ru, and Ir, occur practically always as constituents of the crude platinum found in nature and are separated from it in its purification. Their separation from each other is a difficult analytical operation, and the student has already been referred elsewhere for details. See p. 145. The present discussion will therefore be very brief and confine itself to the mode of obtaining a solution of each and to a very few of the most striking characteristics. Palladium. Palladium is sometimes alloyed with gold and silver as an ore, but is usually associated with platinum. It is sometimes used to make graduated scales for astronomical in- struments, to plate silverware, and in dentistry. It is used also to manufacture automatic gas lighters, because of its power of occlusion, which property also makes it useful for gas analysis. While palladium forms two series of compounds in which it is bivalent and tetravalent, only the former are sufficiently stable for analytical purposes. It is the only metal of the group which dissolves in nitric acid, and the aqueous solution of palladous nitrate presents the following reactions : 1. H 2 S precipitates black PdS from acid and neutral solu- tions, insoluble in ammonium sulphide, but soluble in boiling concentrated HC1. 2. KI produces a black precipitate, PdI 2 , soluble in excess of precipitate and also in ammonium hydroxide, a characteristic reaction. 3. Hg(CN) 2 precipitates Pd(CN) 2 , insoluble in HC1, but soluble in KCN and NH 4 OH. THE RARE METALS 149 Osmium. Osmium finds some technical application as a lamp filament, and its tetroxide is used in bacteriological operations. The element manifests a valency of 2, 3, 4, 6, and 8, the last being the best known in the form of the oxide OsO 4 , which is known as osmic acid, though it has scarcely any acid character. When platinum alloys are treated with aqua regia, osmium and iridium remain undissolved, and the osmium may be separated by volatilization in a current of oxygen. If in a finely divided condition, it is soluble in aqua regia and can be driven out of the solution by distillation. It is characterized by an offensive chlorinelike odor and irritating reaction with the mucous mem- brane of eyes and throat. The solution of osmic acid in KOH may be used for the following reactions : 1. Treated with nitric acid, the osmic acid is liberated, and its odor may be noted. 2. H 2 S precipitates in HC1 solution the sulphide, insoluble in (NH^S. 3. Indigo is decolorized by osmic acid. 4. SnCl 2 precipitates a brown precipitate soluble in HC1 to a brown solution. Iridium. Osmium iridium alloy is used because of its hard- ness, infusibility, and indifference to reagents, as pen points, watch and compass bearings, and iridium itself may be used for knife edges, for balances, and similar purposes. Iridium remains nearly insoluble when the platinum alloys are digested in aqua regia. The residue not volatile in a current of oxygen can be iridium and ruthenium, and may be fused in a silver crucible with sodium hydroxide, and the fused mass, after removal from the crucible, may be dissolved in aqua regia, giving a solution of lg which will give the following reactions : i. H 2 S first decolorizes the solution and finally precipitates Ir 2 S 3 , readily soluble in (NH 4 ) 2 S. 150 QUALITATIVE ANALYSIS 2. NaOH changes the color of the solution from dark red to green ; on warming it changes to a sky-blue color. 3. Reducing agents usually change the color to an olive- green. Ruthenium. Neither the element nor its compounds find any extensive use in the arts. It is therefore only likely to be en- countered in the investigation of platinum ores. It is insoluble in aqua regia. Rhodium. This metal likewise finds no special commercial application and is therefore likely to be of interest only in the analysis of platinum ores for which the student is referred to the paper of Mylius and Dietz. See p. 145. Selenium. This element resembles sulphur closely in many of its properties and is therefore found replacing it in many of its compounds. It is of quite frequent occurrence, but almost invariably in very small quantities. It is also sometimes found in deposits of native sulphur as a sulphide or selenide. On roasting or burning it is converted to the solid oxide SeO 2 which does not readily oxidize to the higher form and consequently re- mains as flue dust or solid deposit in the acid chambers in these processes. As an element, it finds application because of the effect of light upon its electrical conductivity, in the manufac- ture of electrical devices. It exists in two allotropic modifica- tions and two series of compounds analogous to those of sulphur. Selenium-bearing flue dust may be fused with potassium hy- droxide in a silver crucible and extracted with water, and the solution of K 2 SeO 3 and K 2 Se used for the following reactions : i. H 2 S produces in dilute acid solution a lemon-yellow precipitate of selenium and sulphur, soluble in am- monium sulphide. H 2 SeO 3 + 2 H 2 S -> 3 H 2 O + 2 S + Se. The acidifying of the original solution liberates THE RARE METALS 151 H 2 Se, detected by its characteristic odor of decayed cabbage or horse-radish. i. Reducing agents precipitate the selenium in the red variety, changed by continued boiling to black. 3. Chlorine or aqua regia converts the selenious acid to selenic acid, which forms with BaCl 2 insoluble BaSeO 4 . The latter on boiling with HC1 is re- duced, with evolution of chlorine, to BaSeO 3 , solu- ble in acid. 4. Flame test : Heated in the reducing flame, selenium and its compounds color the flame cornflower-blue (a shred of asbestos serves in lieu of a wire). 5. Codeine added to concentrated H 2 SO 4 containing traces of selenium will give a green color. Tellurium. Tellurium is of less frequent occurrence than selenium, but it is more important because of its occurrence with gold in certain ores. Since the gold volatilizes.with the tellurium on heating, these ores require special treatment. Considerable quantities of tellurium are obtained at the Baltimore and Omaha works of the American Smelting and Refining Co. as a by- product. 1 No commercial uses have been found for it. Tellu- rium compounds when introduced into the human system impart a peculiarly offensive odor to the breath. Tellurium forms two series of compounds analogous to those of sulphur and selenium. The following reactions are characteristic : 1. A fragment of tellurium heated gently with concentrated sulphuric acid produces a beautiful carmine color. 2. Flame test : Tellurium and its compounds, heated in the reducing flame, give a blue flame. 3. HNO 3 oxidizes tellurium to tellurous acid, which is but little soluble in water, but is dissolved readily by 1 This product probably contains two unidentified elements. See Flint, Am. Jour, of Set., Vol. 30, p, 1209; 1910. 152 QUALITATIVE ANALYSIS bases. It is also soluble^in strong nitric acid, form- ing tellurium nitrate, which is hydrolyzed by water. 4. Tellurium compounds heated with sodium carbonate, out of contact with air, produce Na 2 Te ; soluble in water to a red solution, from which tellurium may be precipitated by a current of air. 5. H 2 S and other reducing agents precipitate tellurium, or TeS 2 , which is soluble in ammonium sulphide. Germanium. This extremely rare element belongs to the carbon family and in properties lies between tin and carbon. Its chief interest lies in the prophecy of its properties by Men- delejeff in 1871 and its discovery fourteen years later. 1 It is precipitated from its soluble salts by hydrogen sulphide in strongly acid solution, soluble in ammonium sulphide. The sulphide is white. Group III V, Ti, Zr, Be (67), Th, Ce, La, Di (Nd t Pr\ Yt, Yb, Sc, Er y and other still rarer elements Of these elements, several are so very rare as to render their discussion in these notes of but little value, and the student is referred elsewhere when their detection is rendered neces- sary. They are all precipitated by ammonium hydroxide and ammonium sulphide along with the other elements of the third group, although part of them are acid-forming elements and are precipitated because of the hydrolysis of their ammonium salts. Vanadium. This element is found quite widely distributed, but in conjunction with the compounds of other elements and in small amounts. The commonest minerals are Vanadinite, (Pb 6 (VO 4 ) 3 Cl), Carnotite, (K 2 O 2 U 2 O 3 . V 2 O 5 ), Descloizite, ((PbZn) 2 (OH)V0 4 ). 1- Jour, f.prakt. Chem., Vol. 34, p. 177. THE RARE METALS 153 The element is used as an alloy in the manufacture of va- nadium steel, which has very high tensile strength. The com- pounds are also used in porcelain manufacture for coloring the enamel green, in calico printing and silk dyeing, and in the preparation of an indelible ink. The use of vanadium in steel has greatly increased the value of vanadium ores. The element forms a series of five oxides, of which the pen- tavalent form is the most important. A solution of vanadium suitable for examination may be prepared as follows : Fuse the vanadium mineral with KOH and extract with water. The solution is filtered and the second group metals removed by H 2 S. The filtrate is boiled with a few drops of nitric acid to oxidize the vanadium to vanadic acid. The solution may be used as follows : 1. H 2 S causes no precipitate, except sulphur, but changes the color to blue. Zinc produces the same reaction. 2. Ammonium hydroxide and ammonium chloride slowly precipitate ammonium metavanadate, NH 4 VO 3 . 3. H 2 O 2 added to the nearly colorless alkali vanadates in faintly acid solution produces a brownish red color. 4. Bead test : Vanadium compounds color the bead brown- ish yellow in the oxidizing flame, green in the reducing flame. Titanium. Titanium belongs in the carbon family and is fairly widely distributed, both in its own mineral forms as Rutile, Ti(TiO 4 ), and Titanic iron, FeTiO 3 , Titanite, CaSiTiO 3 , etc., and as an impurity in many iron ores. It is therefore likely to be encountered in the analysis of iron ores or of rocks. The metal itself finds but little use except for demonstration of its proper- ties, and its presence in other ores is considered undesirable because of its refractory character, though a titanium steel may be made which possesses many desirable characteristics. Its 154 QUALITATIVE ANALYSIS ores are sometimes used in lining puddling furnaces, to color porcelain yellow, in coloring artificial teeth, etc. The titanite is used as a gem. Its occurrence and uses being so relatively frequent and since it is also found in some sands, in certain varieties of coal, in meteorites, occasionally in ashes, bones, etc., it is sometimes classed among the common elements and its study and analytical reactions provided for in the regular groups. It forms four oxides, of which the dioxide is the most impor- tant and from which its natural compounds are derived. Its ores are not easily rendered soluble by fusion in sodium car- bonate, and it is best to fuse in potassium acid sulphate and dis- solve the fused mass in cold water. The titanium dissolves as Ti(SO 4 ) 2 , and the solution may be acidified with hydrochloric acid and used for the following characteristic reactions : NOTE. A preparation of pure titanium sulphate is most easily made as follows : The fused mass is dissolved in cold water and filtered and then heated to boiling. The pre- cipitated TiO(OH) 2 is filtered hot and redissolved in concentrated sulphuric acid. 1. KOH, NaOH, NH 4 OH, (NH 4 ) 2 S, and carbonates, in- cluding BaCO 3 , precipitate in the cold orthotitanic acid, readily soluble in acids; hot, metatitanic acid is precipitated ; difficultly soluble in acids. 2. NaC 2 H 3 O 2 forms Ti(C 2 H 3 O 2 ) 4 , completely hydro lyzed by boiling. (This hydrolyzing of salts occurs on boiling the salts of ordinary minerals acids, hence the necessity of using cold water in preparing the solution above.) 3. H 2 O 2 added to not too strongly acid solutions produces a yellow to orange-red color, due to TiO 3 . (See also Vanadium.) Ferric chloride, of course, inter- feres with this reaction, as do also the chromates. 4. Tin or zinc in acid solutions reduce the titanium to the trivalent form, giving a violet color. Titanium salts are not reduced by H 2 S. THE RARE METALS 155 5. Bead reaction: Titanium beads are colorless in the oxi- dizing flame, but violet in the reducing flame. By addition of tin the bead is reduced more quickly. Zirconium. Zirconium silicate ZrSiO 4 (Zircon) is the com- monest natural form of this element, though it occurs in a number of other minerals. The oxide of the metal is unaltered by the oxyhydrogen flame and glows brilliantly when heated. It is con- sequently used instead of lime in the Drummond light in light- houses and is the chief constituent of the glowing filament in the Nernst lamp. These uses are of diminishing importance because of the introduction of other methods of lighting. The mineral zircon is sometimes cut and used as a gem under the name hya- cinth. Less valuable gems of the same material are known as jargon. The element forms two oxides and corresponding salts, of which those corresponding to ZrO 2 are more important. A convenient method of obtaining a solution for examination is as follows: Fuse the pulverized mineral substance with sodium carbonate at as high a heat as possible. After cooling, digest the mass with cold water and boil. The sodium zirconate is hydrolyzed and remains undissolved. The residue may be dis- solved in sulphuric acid and after removal of second group metals, if present, the solution is precipitated with NH 4 OH and filtered. The precipitate may be redissolved in sulphuric acid and used for the following tests: l i. NH 4 OH, NaOH, KOH, and (NH^S precipitate white Zr(OH) 4 ; insoluble in excess ; soluble in acids. If precipitated hot, it is difficultly soluble in dilute but readily soluble in concentrated acid. l If iron and alumina are also present, they may be removed by precipitation with ex- cess of NH 4 OH in cold solution and filtering. The precipitate dissolved in sulphuric acid is treated with an excess of ammonium carbonate solution and again filtered. The filtrate is boiled, and zirconium hydroxide is precipitated. 156 QUALITATIVE ANALYSIS 2. H 2 O 2 precipitates white Zr 2 O 5 from solutions not too acid. The precipitate evolves chlorine when treated with HC1. No color is produced. (See Ti.) 3. K 2 SO 4 precipitates from a concentrated not too acid solution white basic zirconium sulphate ; insoluble in dilute HC1. (Th and Ce give similar reactions.) Uranium. This element occurs in a number of rare minerals, the most widely distributed of which is pitchblende, or uraninite, a mineral which is 75 to 8$ per cent uranium oxide (or uranyl uranate), but contains also nitrogen, helium, thorium, zirconium, calcium, iron, copper, lead, etc. It is most famous as the source from which the radioactive elements radium and polonium were obtained by Professor and Madame Curie. It is found most abundantly at Joachimstal, Bohemia, but also occurs in many places in the United States. The metal is used to some extent in the manufacture of gun steel. The oxide is used to produce a glass of a peculiar greenish yellow fluorescence, and this furnishes its most extensive application. The salts are used to produce orange and black tints for enamels and china paint- ing. Some uranium also is used in photography and in the preparation of the Welsbach mantles. (See Cerium and Thorium.) The element forms a complex series of oxides and corresponding compounds, of which the ones which chiefly concern the analyst are uranous salts, where the element is tetravalent and the uranyl, or basic, salts, where it is hexavalent. A solution suitable for examination for uranium reactions may be prepared from its ores by digesting the finely pulverized material with aqua regia, and evaporation to dryness on the water bath. The residue is extracted with water and dilute hydrochloric acid and treated with H 2 S to remove second- group metals. The filtrate is boiled with nitric acid and filtered. The filtrate is treated with an excess of ammonium hydroxide THE RARE METALS 157 and again filtered. The precipitated ammonium uranite will dissolve in an excess of ammonium carbonate. This solution may be filtered and acidified with HC1 and the yellow solution used for the following reactions : 1. NH 4 OH precipitates yellow ammonium uranate (NH^UgOy, soluble in alkali carbonates. 2. K 4 Fe(CN) 6 causes a brown precipitate or, in dilute solutions, a brown coloration, insoluble in dilute acid. 3. H 2 O 2 produces a yellowish white precipitate in con- centrated solutions, if the solution is kept neutral by drop-by-drop addition of ammonia. 4. Zinc produces a green coloration. 5. Bead test: The uranium bead is yellow green in both oxidizing and reducing flames and is fluorescent in the former. Beryllium. This element occurs rather widely distributed, but usually in small quantities, and the specially fine crystals of its compounds are cut and used as gems. The most im- portant minerals are beryl (emerald, aquamarine), Be 3 Al 2 (SiO 3 ) 6 , and chrysoberyl (Alexandrite, cat's eye), BeAl 2 O 3 ; also gadol- inite, Be 2 FeYSi 2 O 10 . It is bivalent in all its compounds, and its salts are sweetish in taste ; hence the name Glucinum. It may be obtained in so- lution by fusion of beryl with sodium carbonate and treatment of the fused mass with hydrochloric acid. The solution is filtered from the silicic acid and evaporated to dryness and dis- solved in dilute acid. This solution is treated with an excess of ammonium carbonate and filtered cold. The solution is fil- tered and boiled and the precipitate of basic carbonate dissolved in HC1. The solution will not be quite free from aluminium, but will serve for the following reactions : 158 QUALITATIVE ANALYSIS 1. NH 4 OH and(NH 4 )2S precipitate Be(OH) 2 not soluble in excess, but readily soluble in acids or alkali bases. 2. H 2 C 2 O 4 and (NH 4 ) 2 C 2 O 4 produce no precipitate (which distinguishes it from Th, Ce, Zr, Y, La, Di, etc.). 3. (NH 4 )2CO 4 precipitates the carbonate, readily soluble in excess, but not so readily soluble in alkali carbon- ates. Thorium. The purest form of thorium found in nature is thorite, ThSiO 4 ; but it is also found associated with other rare elements in gadolinite (see Be) and in monazite sand (see also Ce), from which nearly all commercial thorium compounds are derived. Many other rare minerals also contain it. Thorium compounds are of increasing interest, because of the use of the oxide in the formation of gas mantles for the Welsbach light. The mantles are prepared by saturating a cotton form with nitrates or bromides of the rare earths. The cotton web is burned away, and the somewhat fragile shell of oxides which re- mains constitutes the mantle and is chiefly thorium oxide, ThO 2 . The illuminating power of the mantle is said to be at its maxi- mum when it consists of 99 per cent thorium oxide and I per cent cerium oxide, though the mixture usually contains oxides of several other of the rare metals. Thorium salts are also of great interest because of their radioactivity and their decomposition into thorium " emanations." For detailed information, consult works on Radioactivity. Thorium minerals are usually decomposed by sulphuric acid, and a solution suitable for the characteristic tests may be obtained by digesting thorite, or monazite sand, or discarded Welsbach mantles, in concentrated H 2 SO 4 , heating the mass to redness. Cool and extract with cold water. The filtrate may be treated with oxalic acid and the precipitated oxalates dissolved in hot THE RARE METALS 159 ammonium oxalate. The filtrate on acidification with dilute HC1 will give a precipitate of the oxalate, and this on ignition is converted into almost pure thorium oxide. This oxide dis- solved in concentrated sulphuric acid reacts as follows : (This solution requires long digestion.) 1. NH 4 OH and other bases precipitate Th(OH) 4 ; insol- uble in excess ; readily soluble in dilute acid. 2. K 2 CO 3 precipitates the white carbonate ; soluble in ex- cess ; rendered turbid by heating, and redissolving when cooled. 3. Oxalic acid and ammonium oxalate react as indicated in the preparation of the solution (see above). 4. H 2 O 2 gives a white precipitate, soluble in sulphuric acid, the solution giving the perchromate test with chromic acid. Zirconium gives the same reaction. 5. No bead coloration is formed with thorium compounds. For fuller information concerning the analysis of monazite sand, see Zeit. f. Angew. Chem., Vol. 14, p. 655 ; 1901 ; also Ber., Vol. 35, p. 2826; also Jour. Am. Chem, Soc., Vol. 24, p. 901 ; 1902. For separa- tion of cerium and thorium and zirconium, see Zeit. f. Anal. Chem., Vol. 36, p. 676, and Vol. 37, p. 94, . Prakt. Chem. (2), Vol. 66, p. 59. Cerium. Cerium occurs in monazite sand, associated with lanthanum, neo- and praesodymium, thorium, and other rare elements, as the phosphate, and as the silicate in cerite with lanthanum and didymium, and many other rare minerals. The chief interest of cerium lies in its relation to thorium in the Welsbach mantles and in the use of the oxalates in medicine as a nerve sedative in hysteria and pregnancy and as a remedy for sea sickness. The element may be extracted from its ores by digestion with 160 QUALITATIVE ANALYSIS hot sulphuric acid (see Thorium) and a solution prepared as fol- lows: Cool the hot mass from which the excess H 2 SO 4 has been ex- pelled, digest with cold water, and filter. The filtrate is treated with H 2 S to remove second group metals, and after filtration is treated with oxalic acid. The precipitate is digested with ammonium oxalate. (See Thorium.) The residual oxalates are ignited and consist of cerium, lanthanum, and didymium oxides, and are dis- solved in concentrated HC1. The solution is treated with KOH in excess and chlorine passed into the mixture until the precipi- tate is deep yellow in color. The mixture is now filtered. The filtrate contains the lanthanum and didymium. See below. The eerie oxide is dissolved in HC1 (i-i) with evolution of chlorine, and the clear solution of CeCl 3 used for the following reactions : 1. NH 4 OH and other bases precipitate Ce(OH) 3 ; insoluble in excess. 2. H 2 O 2 in neutral solution precipitates an orange-yellow CeO 3 , made lighter in color by boiling. 3. Bead test : Cerium salts color the oxidized bead of both borax and microcosmic salt orange-yellow while hot, pale yellow when cold ; the reduced bead is colorless. Indium and Gallium. These elements are of extremely rare occurrence, mainly in zinc ores, and were discovered by means of the spectroscope. The student is referred for fuller informa- tion regarding them, to Browning's Introduction to the Rarer Elements. Lanthanum, Didymium {Neo and Prceso), Yttrium, Scandium, Erbium, etc. These elements are extremely rare, and together with several others, some of which are of doubtful existence, are extracted from the cerium and thorium and beryllium min- erals. In case one of these elements is found in a mineral and it is desired to investigate for the rarer elements, the student is THE RARE METALS l6l referred to the system of analysis of Noyes, Tech. Quart., Vol. 16, pp. 93-131, and Jour. Am. Chem. Soc., Vol. 29, p. 137; 1907. Also a suggested course of procedure is given in Bottger's Quali- tative Analyse, 2d Auflage, p. 208 et seq. Group IV Ammonium Carbonate Group Rd Radium occurs in extremely minute quantities in the two uranium ores, pitchblende and carnotite, from which it was ex- tracted with almost infinite pains by Madame Curie. It natu- rally has no place in these notes, except as it deserves mention because of the close resemblance of its chemical behavior to that of barium. Group V The Soluble Group Li, Rb, Cs Lithium. This element is found very widely distributed in very small amounts in tourmaline, muscovite, epidote, and ortho- clase, which are common silicate minerals. It occurs also in larger quantities in the somewhat rare lepidolite, or lithia mica, triphylite, petalite, etc. Because of this wide distribution it finds its way into many surface waters where its detection is important. Traces of it are also found in many plant ashes, although large amounts in soil seem to be detrimental to plant growth. It is present particularly in seaweed, tobacco, cocoa, coffee, sugar cane, and in milk, blood, and muscular tissue. Its salts find use in medical preparations, and mineral waters con- taining it in weighable quantities are considered valuable be- cause of the solubility of lithium urate. They are therefore supposed to assist in the removal of uric acid from the tissues of the body. 162 QUALITATIVE ANALYSIS Most of the famous mineral spring waters, as Durkheim, Kissingen, Ems, Karlsbad, Buffalo Lithia, and numerous others, owe their curative value, real or supposed, to the presence of these salts. Its salts 'have the general behavior of those of sodium and potassium, but it somewhat resembles the alkaline earth metals in that it forms a somewhat difficultly soluble car- bonate and phosphate. In the regular course of analysis, it is found along with sodium and potassium in the fifth group filtrate, and if the solu- tion is evaporated to dryness, the dry chloride of lithium will dissolve in alcohol or in amyl alcohol. The alcoholic solution may be evaporated and the residue dissolved in water. If lithium is present, it will give the following reactions : 1. Flame test : Lithium salts color the flame carmine, violet through blue glass and invisible through green glass. They also give a very characteristic spectrum, the red line of which is nearer the yellow than that of potassium. 2. H 2 PtCl 6 forms no precipitate. 3. Na 2 HPO 4 precipitates Li 3 PO 4 from moderately concen- trated solutions, and if the solution is made slightly alkaline with NaOH and evaporated to dryness, the phosphate remains almost quantitatively insol- uble in water containing a little ammonia. 4. (NH 4 ) 2 CO 3 precipitates from the concentrated solution Li 2 CO 3 (solubility 13 g. per liter at 18). Rubidium and Cczsium. These elements occur widely dis- tributed, but in minute amounts in the minerals and mineral waters which contain potassium and lithium. Their chief claim to interest lies in their discovery by means of the spectroscope in the hands of its inventors, Kirchhoff and Bunsen. They form even more insoluble chlorplatinates than does THE RARE METALS 163 potassium, and in the regular course of analysis are sought in the potassium chlorplatinate precipitate. After ignition, the precipitate is extracted by water and the solution evaporated to near dryness. The presence of rubidium may be recognized by two dark red lines near the potassium line, and two lines in the violet. Caesium may be recognized by two sky-blue lines. Rubidium colors the flame red and caesium sky-blue. APPENDIX LIST OF APPARATUS The student will need the following set x>f apparatus to carry on the work outlined in this manual : Asbestos board, 4x4 inches. Beakers, Nos. i, 2, and 3. Blue glass. Bottle (rubber stopper) for NaOH. Bottles, reagent, glass-stoppered. Bunsen burners, with two feet of rub- ber tubing. Crucible tongs. Evaporating dishes. Filter papers. Flasks, Erlenmeyer : 60 c.c. 100 c.c. 300 c.c., for H 2 S precipitation, with stopper. Flasks, Florence : 100 c.c., for hydrogen generator (with rubber stopper, funnel tube, and rubber connections). 750 c.c., for wash bottle, with stopper. Besides the above apparatus, the student should supply him- self with the following : Forceps. Funnel rack. Funnels. Glass tubing. Platinum foil. Platinum wire. Rubber tubing. Sand bath. Spatula, bone. Stand with two rings. Test-tube brush. Test-tube holder. Test-tube rack. Test tubes, large. Test tubes, small. Triangles, plain. Watch glasses. Wire gauzes. Apron, rubber or denim. Half sleeve, pair. Matches. Notebook. Towel. 165 1 66 QUALITATIVE ANALYSIS REAGENTS IN SOLUTION (Needed in analysis) ACIDS r Acetic, Sp. Gr. 1.044: one part glacial acid and two and one half parts of water. Aqua regia : one part of nitric and three parts of hydrochloric. Hydrochloric, concentrated, Sp.Gr. 1.20: 36.5 per cent HC1. Hydrochloric, dilute : one part concentrated acid and four parts of water. Hydrochloric, dilute, for separation of Co and Ni from Mn and Zn : one part concentrated acid to twenty of water. Hydrogen sulphide ; is used either as a gas or saturated solution. Nitric, concentrated, Sp. Gr. 1.42 : 68 per cent HNO 3 . Nitric, dilute : one part concentrated acid and four parts of water. Sulphuric, concentrated Sp. Gr. 1.84; 98 per cent H 2 SO 4 . Sulphuric, dilute : one part concentrated acid and four parts of water. (Add the acid to the water with con- stant stirring.) Sulphurous acid : used either as a gas or saturated solution. ALKALIES : Ammonium hydroxide, concentrated, Sp. Gr. 0.90 : 28 per cent NH 3 . Ammonium hydroxide, dilute : one part concentrated am- monia and three parts of water. Potassium hydroxide: 100 grams of the solid to one liter of water. Sodium hydroxide : 100 grams of the solid to one liter of water. SALTS AND OTHER REAGENTS IN LIQUID FORM : Ammonium acetate : 1000 c.c. of concentrated ammo- nium hydroxide, 1250 c.c. of glacial acetic acid, neu- APPENDIX 167 tralize with ammonia or acetic acid; or a saturated solution of the salt in water. Ammonium sesquicarbonate : 500 grams of the crystals of ammonium carbonate in 600 c.c. of water and 80 c.c. of concentrated ammonia; after solution is complete, make up to a liter with water. If the salt is dry, use 200 grams instead of 500. Ammonium carbonate : 500 grams of the crystals in 600 c.c. of water ; after solution is complete, make up to a liter with water. If the salt is dry, use 200 grams instead of 500. Ammonium chloride : 100 grams of the salt in a liter of water. Ammonium molybdate : dilute 100 c.c. of ammonia (Sp. Gr. 0.90) with 150 c.c. of water and dissolve in this 50 grams of molybdic acid ; dilute 250 c.c. of concen- trated nitric acid with 500 c.c. of water, and pour slowly into the first solution with constant stirring ; allow to stand in a warm place for forty-eight hours and decant the clear supernatant liquid for use. Ammonium oxalate : a saturated solution in water. Ammonium sulphate : 250 grams of the crystallized salt in a liter of water. Ammonium sulphide, colorless : saturate a liter and a half of concentrated ammonium hydroxide with hydrogen sulphide; add a liter of concentrated ammonium hy- droxide and two liters of water. Ammonium sulphide, yellow : add to the above solution about 75 grams of powdered roll sulphur. Barium carbonate : 60 grams suspended in a liter of water. Barium chloride : 60 grams of the crystallized salt in a liter of water. Barium hydroxide : 50 grams of the crystallized salt in a liter of water. 168 QUALITATIVE ANALYSIS Bromine water : 50 grams of potassium bromide in 500 c.c. of water and shake with 10 grams of bromine until dissolved. Brucine solution : 0.2 grams dissolved in 100 c.c. concen- trated sulphuric acid. Calcium hydroxide : a saturated solution in water. Calcium sulphate : a saturated solution in water. Chlorine water : water saturated with chlorine gas. Keep in a dark place or in a dark bottle. Ether ; commercial. Ethyl alcohol : 95 per cent. Ferric chloride : 100 grams to a liter of water. Ferrous sulphate : 200 grams per liter ; add 5 c.c. H 2 SO 4 and some pieces of metallic iron (tacks) ; keep stop- pered. Hydrogen peroxide : 3 per cent solution. Indigo : place five parts of fuming sulphuric acid in a beaker which is immersed in water, and then add slowly with constant stirring one part of powdered indigo ; cover the beaker and allow to stand forty-eight hours, and then pour into twenty times its volume of water ; stir thoroughly and filter if necessary. Lead acetate : 100 grams of the crystals in a liter of water. Magnesia mixture : 90 grams of magnesium chloride, 240 grams ammonium chloride, 50 c.c. concentrated am- monium hydroxide ; dilute to a liter. Methyl alcohol ; commercial. Mercuric chloride : 50 grams to a liter of water. Nessler's reagent : 7 grams of potassium iodide are dis- solved in 20 c.c. of water, and to this solution is added gradually a solution of 3.2 grams of mercuric chloride in 60 c.c. of water until a permanent precipitate is formed; 120 c.c. of concentrated potassium hydroxide APPENDIX 169 solution is added and allowed to stand until settled; decant from the precipitate as needed. Phenol sulphonic acid : dissolve 24 grams of phenol crystals in a mixture of 148 c.c. of concentrated sulphuric acid and 12 c.c. of water; keep in the dark. Potassium chromate : 100 grams to a liter of water. Potassium cyanide : 30 grams to a liter of water. Potassium ferricyanide : 10 grams to a liter of water. Potassium ferrocyanide : 1 5 grams to a liter of water. Potassium fluoaluminate (reagent for sodium) : treat an ex- cess of freshly precipitated A1(OH) 3 with concentrated HF in a lead or platinum dish. Allow to stand at room temperature for two days. Add an equal volume of a saturated solution of potassium acetate ; boil and filter. Add an amount of 50 per cent alcohol equal to the volume of nitrate. Filter again, if necessary, and preserve in a glass container. Potassium iodide : 20 grams to a liter of water. Potassium nitrite : 50 grams to a liter of water. Potassium sulphate : 85 grams to a liter of water. Potassium sulphocyanate : 50 grams to a liter of water. Silver nitrate : 25 grams to a liter of water. Silver sulphate : a saturated solution. Sodium acetate : a saturated solution. Sodium ammonium phosphate : 70 grams to a liter of water. Sodium carbonate : 1 50 grams of the anhydrous salt or 250 grams of the crystallized salt to the liter. Sodium cobaltic nitrite : dissolve 100 grams of sodium nitrite in 300 c.c. of water, add acetic acid to slight acid reaction, and then add 10 grams of cobalt nitrate. Allow to stand several hours and filter if not clear. The solution decomposes slowly, therefore it is ad- visable to prepare only in small quantities. 170 QUALITATIVE ANALYSIS Stannous chloride : heat an excess of tin with concen- trated hydrochloric acid ; dilute with four volumes of water and keep in stoppered bottles with an excess of tin. Tartaric acid : 100 grams to the liter. Zinc sulphate: 140 grams of the crystallized salt to the liter. OTHER SOLUTIONS USED IN PRELIMINARY EXPERIMENTS Aluminium sulphate : 30 grams to the liter. Antimony chloride : 25 grams of the salt, 250 c.c. of concentrated hydrochloric acid, and sufficient water to make up to a liter. Arsenous chloride: 35 grams of arsenous oxide, 50 c.c. concentrated hydrochloric acid, and sufficient water to make up to a liter. Bismuth chloride : 30 grams to a liter of water and enough hydrochloric acid to dissolve the precipitate formed. Boric acid : a saturated solution. Cadmium chloride : 25 grams to a liter of water. Calcium chloride : 25 grams to a liter of water. Chromium sulphate : 30 grams to a liter of water. Cobalt nitrate : 30 grams to a liter of water. Copper sulphate : 30 grams to a liter of water. Lead acetate (alkaline): two parts lead acetate, one part of ammonium acetate in water ; make alkaline with ammonium hydroxide. Magnesium sulphate : 100 grams to a liter of water. Manganese sulphate : 40 grams to a liter of water. Mercurous nitrate : 60 grams to a liter of water. Potassium bichromate : 50 grams to a liter of water. APPENDIX 171 Potassium bromide : 30 grams to a liter of water. Potassium chloride : 50 grams to a liter of water. Potassium chromate : 50 grams to a liter of water. Potassium iodide : 30 grams to a liter of water. Sodium arsenate : 30 grams to a liter of water. Sodium arsenite : 30 grams to a liter of water. Sodium phosphate : 120 grams to a liter of water. Stannic chloride : 1 5 grams of stannous chloride are treated in an evaporating dish with 50 c.c. of hydro- chloric acid and 5 grams of potassium chlorate, and heated until chlorine is no longer evolved ; the solu- tion is then diluted to 500 c.c. Strontium chloride : 30 grams to a liter of water. REAGENTS IN SOLID FORM Aluminium foil. Aluminium wire. Ammonium chloride. Ammonium nitrate. Ammonium sulphate. Ammonium sulphite. Barium chloride. Barium hydroxide. Borax. Calcium carbonate (marble). Copper foil. Copper wire. Cotton, absorbent. Ferrous sulphate. Ferrous sulphide. Iron filings. Iron wire. Lead acetate. Lead dioxide. Oxalic acid. Potassium carbonate. Potassium chlorate. Potassium cyanide. Potassium nitrate. Silver nitrate. Sodium ammonium phosphate. Sodium carbonate. Sodium hydroxide. Sodium sulphate. Sodium sulphide. Sodium sulphite. Sulphur (powdered or flow- ers). Tartaric acid. Zinc, granulated. Zinc, sheet. QUALITATIVE ANALYSIS TABLE OF ATOMIC WEIGHTS (1912) ELEMENT SYMBOL ATOMIC WEIGHT o = 16 ELEMENT SYMBOL ATOMIC WEIGHT 0=16 Aluminium . . . Antimony .... Ar^on . IP A ' 27.1 I2O.2 30 88 Molybdenum . . Neodymium . . . Neon .... Mo Nd Ne 96.0 144-3 2O 2 Arsenic As jy.uu 74.. QO Nickel . Ni c;8.68 Barium Ba I 37 37 Nitrogen N - 14 01 Bismuth .... Boron F l j/ "Jl 208.0 I I.O Osmium .... Oxygen .... Os o 190.9 16.0 Bromine .... Cadmium .... Caesium .... Calcium .... Carbon Br & Ca c 79.92 II2-4 I32.8I 40.07 12 O Palladium . . . Phosphorus . . . Platinum .... Polonium .... Potassium Pd P Pt Po , K; 106.7 31.04 195.2 (?) 2Q IO Cerium .... Ce I -dO 25 Prassodymium . Pr jy.iu 1 40 6 Chlorine .... Chromium .... Cobalt Cl Cr Co; 3546 52.0 rg Q7 Radium .... Rhodium .... Rubidium Ra Rh Rb 226.4 102.9 gt AC Columbium . . . Copper Dysprosium . . . Erbium Cb Cu Dy Er }j.y/ 93.5 63-57 162.5 l67 7 Ruthenium . . . Samarium . . . Scandium . . . Selenium . Ru Sm Sc Se *O'*r3 IOI.7 150.4 44.1 7Q.2 Europium .... Eu 1 12. Silicon Si - 28.3 Fluorine .... Gadolinium . . . Gallium .... Germanium . Glucinum .... Gold . . . F - Gd Ga Ge Gl Au 19.0 157-3 6 9 . 9 72.5 9.1 IQ7 2 Silver Sodium .... Strontium . . . Sulphur .... Tantalum. . . . Tellurium Na Sr S Ta Te 107.88 23.00 87-63 32.07 I8I.5 127.6* Helium He *y/ * 3. QQ Terbium .... Tb I tJQ.2 Hydrogen .... Indium . H In j.yy 1. 008 114 8 Thallium .... Thorium Tl Th 204.O 232 Iodine I 126 92 Thullium . Tm i68.c Iridium Ir - IQ3.I Tin Sn 1 19.0 Iron Fe A yj'* CC.8C Titanium .... Ti 48.1 Krypton .... Lanthanum . . . Lead Kr La Pb 82.92 139.0 207. 1 Tungsten .... Uranium .... Vanadium W U V 184.0 238.5 KI.O Lithium Li 6 QA. Xenon Xe I 3O 2 Lutecium .... Magnesium . . . Manganese Lu Mg Mh w -y*F 174.0 24.32 r/i Q-3 Ytterbium . . . Yttrium .... Zinc Yb Yt Zn 172.0 89.0 6c 37 Mercury .... Hg j^f.yj 200.6 Zirconium . . . Zr u io/ 90.6 * Flint, Am. four. ScL, Vol. 30, p. 1209, 1910, gives atomic weight as 124.3. APPENDIX 173 M j^ ^ i^ .-, LT^ CO >-r> 2 ?1 . 2 s o^ 4 a. u IT N aj " o u "ii ii , 11 "ll *\\ O fa g U 3 J= T3 0^ PH OH &~& g ON - 1 8 vd o. CO II tX. u O fa J! u a S II \ * * m * Ok NO o 2 ON q 2 vO CO Tj- OO 00 I O 1? c/3 u u II OJ II 1 II TJ n II 0. q q 2 ^ S i 6 00 1 o II 1! u ^r II & Ti II H iri H (2 O OH ^ < U IH OH H s rj- ir> o . . ON *? a. 2 OO S\ ON . . 1 ft II U II c/3 II H u vO oo to ' M vd d d o" CO O CO CO Q O O d d 3 1-1 to 2 M tN M Tf CO Cx H " vO vO CO coco W ^ fl CO IO COCJ co to oo r o' o* oq N COCO ON to oo"* d d tN CJ 33 8? d o" d o' 10 ^J- 0=0 d d If Si cc < to SI N ON CJ M S8 d d vo' CJ vO w 1? to (N d d 00' H \Q oo? 9l II d d O.OOII 0.047 CD 'g |lj , COO 00 IT) \O oo COW ol rt PQ s* vq ON CO CJ o r M O d d ^" CO oo' d f " CJ oq q 0^8 o o IN q CO O d d CO H d d 00 CO O co O O d d CJ M O co O O d d < '6 CO to tx ON ^VO o || H M CO O d d \ d d q ci co ON CO 00 O co d cog d d ?f p ^-ON tx O 4 O d d H to o PQ J2 O I 4? vO J 3* d d M VO oTo? d d to co ft cooo ci d to d o* r! d o q q o o \j~) O d d c? 10 d d to CO b M N J 3 a tx CO oo o" VO M to to M o' o' CO rf CO co 10 CO o 10 CJ oo S) M vO 00 A" vq to M VO* M M Svo* (NO - H _j CD rt vO CJ 1>. ON ONH ^"8 O^Tf VO vS* COTf ^. CO to OO M CJ CO coct CO 00 J^ to to 00 vO IN CO -j- ^_J cots. txvO vO d oo O \O H vO H M CO CO O ONH " M M VO M to ON ON vO oo vq to q vO Tj- vO to vo d 00 00 CO CO vo" d CJ to vq co cJoo w co ci oTvq O ON oo' to H H vO o u 03 d u 03 o" o" C/2 9 u f 8" APPENDIX 175 TABLE OF METRIC SYSTEM Length Weight Volume Notation Kilometer .... Kilogram .... Kiloliter 1000. Hectometer . . . Hectogram .... Hectoliter . . . . 100. Decameter .... Decagram .... Decaliter 10. Meter Gram Liter I. Decimeter .... Decigram .... Deciliter .1 Centimeter .... Centigram .... Centiliter .01 Millimeter .... Milligram .... Milliliter .oof METRIC MEASURES WITH ENGLISH EQUIVALENTS MEASURES OF LENGTH I Millimeter, mm. = 0.03937 inches. i Centimeter, cm. = 10 mm. = 0.3937 inches, i Decimeter, dm. = 10 cm. = 3.9371 inches. i Meter, m. = 10 dm. = 39.3708 inches. I Kilometer, km. = 1000 m. = 0.6214 miles. MEASURES OF VOLUME i Cubic centimeter, c.c. = 0.06103 cubic inches. ( 61.027 cubic inches, or, i Cubic decimeter (liter), 1. = ioooc.c.= ] . < 1.057 U. S. quarts. i Cubic meter = 1000 1. = 35.3166 cubic feet. MEASURES OF WEIGHT I Milligram, mg. = 0.0154 grains, i Gram, g. = 1000 mg. = 15.432 grains. I Kilogram (kilo) = 1000 g. = 2.2046 pounds (av.). ENGLISH MEASURES WITH METRIC EQUIVALENTS I Inch = 25.399 millimeters. i Foot = 0.3048 meters. i Mile = 1.609 kilometers. i Cubic inch = 16.386 cubic centi- meters. i Cubic foot =28.315 liters. i Quart = 0.9463 liters. i Grain = 0.0648 grams. i Ounce (av.) = 28.3496 grains, i Pound (av.) = 0.4536 kilograms. 176 QUALITATIVE ANALYSIS APOTHECARIES' WEIGHTS AND MEASURES WITH METRIC EQUIVALENTS i Grain = 0.0648 grams. i Scruple = 1.296 grams. i Drachm = 3.888 grams. i Ounce troy = 31.1035 grams. I Pound troy =373.2418 grams. i Minim = 0.0616 cubic centimeters. i Fluid drachm = 3.6965 cubic centimeters. i Fluid ounce = 29.572 cubic centimeters. i Pint =473.11 cubic centimeters. TO CONVERT A TO B MULTIPLY BY Inches Centimeters 2.54 Feet Meters 0.305 Miles Kilometers 1 .609 Meters . Inches 39-37 Kilometers Miles 0.621 Square inches Square centimeters 6.452 Square yards Square meters 0.836 Square centimeters . . . Square inches 'I55 Square meters Square yards 1.196 Cubic inches Cubic centimeters 16.386 Cubic yards Cubic meters 0.765 Cubic centimeters .... Cubic inches 0.061 Cubic meters Cubic yards 1.308 Fluid ounces Cubic centimeters 2 9-57 Quarts Liters 0.946 Cubic Centimeters . . . Fluid ounces 0.034 Liters. . Quarts 1-057 Grains Milligrams 64.799 Ounces (av.) Grams 28.35 Pounds (av.) Kilograms 0.373 Grams Grains 15-432 Kilograms Pounds 2.205 APPENDIX 177 DEGREE OF IONIZATION OF ACIDS, BASES, AND SALTS Unless otherwise specified, the figures give the per cent ionized in a normal solution at 18, calculated from the electrical conductivities : ACIDS PER CENT SALTS PER CENT HNO 3 82.0 KC1 75.0 HN0 3 (cone., 62%) .... 9.6 NH 4 C1 74.0 HC1 78-4 NaCl 67.6 HC1 (cone., 35%) 13.6 HgCl 2 UNIVERSITY OF CALIFORNIA LIBRARY