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BY THE SAME AUTHOR 
 
 CHEMICAL LECTURE EXPERIMENTS. 
 With 224 Diagrams. Crown 8vo, $2.00. 
 
 A TEXT-BOOK OF INORGANIC CHEMIS- 
 TRY. With 146 Illustrations. Crown 8vo, $1.75. 
 
 ELEMENTARY INORGANIC CHEMISTRY. 
 
 With 108 Illustrations and 254 Experiments. 
 Crown 8vo, 90 cents. 
 
 LONGMANS, GREEN, AND CO. 
 
 NEW YORK, LONDON, AND BOMBAY 
 
A MANUAL 
 
 OF 
 
 CHEMICAL ANALYSIS 
 
 QUALITATIVE AND QUANTITATIVE 
 
 BY 
 
 G. S. NEWTH, F.I.C., F.C.S. 
 
 DEMONSTRATOR IN THE ROYAL COLLEGE OF SCIENCE, LONDON 
 ASSISTANT-EXAMINER IN CHEMISTRY, SCIENCE AND ART DEPARTMENT 
 
 NEW EDITION 
 
 LONGMANS, GREEN, AND CO, 
 
 91 AND 93 *1FTH AVENUE, NEW YORK 
 
 AND 
 
 IQOI 
 

90 
 
 PREFACE TO THE SECOND EDITION 
 
 IN bringing out this second edition of my "Manual of Chemical 
 Analysis " so soon after its publication, I have acted upon the 
 suggestion of one of my courteous reviewers, and have added to 
 the short chapter on Physico-chemical Determinations a description 
 of the " Freezing-point " and " Boiling-point " methods for determin- 
 ing molecular weights. With the exception of this addition, and 
 the correction of certain errors and misprints incidental to a first 
 edition, the book remains unaltered. I would take this opportunity 
 to thank those friends who have kindly drawn my attention to 
 such errata. 
 
 G. S. N. 
 
 June, 1899. 
 
 : i y 
 
PREFACE 
 
 A MAN once brought his son to the Royal School of Mines now 
 the Royal College of Science with the request that he might be 
 taught to " do copper." He did not want his boy to " waste his 
 time learning about oxygen and hydrogen, and ail that? but he 
 wished him simply to learn to " do copper." 
 
 Although seldom expressed with such refreshing candour, the 
 desire to do analysis without learning more than the minimum 
 amount of chemistry is still very prevalent ; and, unfortunately, 
 chemical analysis is a subject which may be, and frequently is, 
 taight and practised in such a manner as to degrade it to the level 
 of a purely mechanical and often quite unintelligible series of rule- 
 ofihumb operations. 
 
 I hope that the student whose aspirations rise no higher than to 
 lea-n to do analysis in this fashion, will not find this book suitable 
 forhim. I have done my best to make it as little of a cram-book 
 as jossible, but have endeavoured to teach analytical chemistry as 
 wel.as analysis that is, the theoretical as well as the practical side 
 of tie subject. 
 
 Vith this object in view, I have carefully avoided the use of those 
 symbolic abbreviated expressions (slang formula, they might be 
 termd), such as H 2 O (oxalic acid), H 2 T(tartaric acid), HA (acetic 
 acid) etc., which are becoming so common, and which, so far as 
 the stdent is concerned, foster those very evils of cramming which 
 we as teachers are striving to combat. If such symbols as these 
 are pemitted and recognised, why not H 2 S, HN, H 3 P, for sulphuric, 
 nitric, and phosphoric acids respectively ? And then, perhaps, 
 some sich hieroglyphic as H 2 S, HN, for sulphurous and nitrous 
 acids. 
 
 For he busy chemist to make use of such shorthand signs in 
 the privcy of his own note-books is one thing, but to print them in 
 
viii Preface. 
 
 a text-book intended for the learner, and thereby put them on the 
 same footing as such chemical formulae as H 2 SO 4 , H 2 C 2 O 4 , etc., is 
 quite another thing, and is open to the taunt of chemistry made easy. 
 
 The importance to the student of making careful notes of his 
 analytical work while in progress, cannot well be overrated, as, 
 perhaps more than anything else, this is calculated to develop in 
 him those habits of exact observation which are essential qualities 
 in a scientific man. While insisting on this point more than once, 
 I have purposely refrained from giving anything of the nature of a 
 specimen of notes, because notes which are made according to a 
 stereotyped pattern are practically of no value. Unfortunately, it 
 is the custom in most laboratories perhaps a necessary custom, 
 but none the less unfortunate to require from the student some 
 written account of his analysis by way of proof that he has con- 
 scientiously done the work. The result of this is that the notes 
 which he makes are taken primarily with a view to furnish this 
 required evidence. Points which ought to be observed and noted, 
 points which he does not at the time understand, are passed over, 
 while others which are familiar, and which are at once recognised 
 as useful evidence, are written down. Indeed, as every leader 
 has experienced, the evidence is often more or less fabricated ater 
 the completion of the analysis. 
 
 The consequence of this method is the stereotyped product so 
 painfully familiar to all who are engaged in teaching or examinhg, 
 namely, a certain number of regulation tests, arranged witl a 
 semblance of method in the everlasting three columns, healed 
 Experiment, Observation, Inference, followed by a more or less 
 slovenly copy of the time-honoured tables. 
 
 The valuelessness of such a written-out record must be p.tent 
 to all teachers, and yet we still continue to accept it in lieu of . few 
 real notes, less quickly and easily examined no doubt, becaus less 
 stereotyped, but of infinitely more value to the student himseL 
 
 In order to gain as much space as possible for purely anaytical 
 matter, and still to be able to include within the limits of i con- 
 venient volume, both qualitative and quantitative analysis,! have 
 carefully excluded all merely descriptive details which hve no 
 direct bearing upon analysis. For example, none of the prcperties, 
 either chemical or physical, of such metals as those of thealkalies 
 or alkaline earths, are made use of in ordinary analytical jr ocesses 
 for the detection or recognition of these elements, thereore any 
 
Preface. ix 
 
 descriptive account of their properties in the elemental state is 
 entirely out of place in a book which is intended only to teach 
 analytical as distinguished from general chemistry. 
 
 In the second portion of the book, devoted to quantitative 
 analysis, I have confined myself to a comparatively small number 
 of well-tried typical methods and processes, preferring to describe 
 and explain in tolerably full detail, a few quantitative determinations 
 in each of the various sections, such as shall furnish a thoroughly 
 sound course of practical study, rather than to attempt to cover 
 necessarily in more sketchy outlines a wider range of subjects. 
 Whether the choice of processes and examples I have made is as 
 good or as representative as it might be, will no doubt be a matter 
 for difference of opinion, and as the book has not been written to 
 meet any particular syllabus, this is probably a point upon which 
 hardly any two teachers will exactly agree. 
 
 The illustrations throughout have been made from original 
 photographs of the actual apparatus employed in the various 
 analytical operations described. 
 
 G. S. N. 
 
 ROYAL COLLEGE OF SCIENCE, LONDON, 
 fune, 1898. 
 
TABLE OF CONTENTS 
 
 BOOK I. 
 QUALITATIVE ANALYSIS. 
 
 PACK 
 
 CHAPTER I. Preliminary exercises Filtration; Solution; Evapo- 
 ration ; Fusion ; Precipitation ; Ignition ; Neutral- 
 isation i 
 
 ,, II. Analytical classification 13 
 
 ,, III. Reactions of the metals of Group V 19 
 
 APPENDIX TO CHAPTER III. The rare metals of Group V. . 26 
 
 CHAPTER IV. Reactions of the metals of Group IV 30 
 
 V. ,, UlA 36 
 
 VI. ., IIlB 51 
 
 ,, VII. The phosphates 63 
 
 APPENDIX TO CHAPTER VII. The rare metals of Group III. 70 
 
 CHAPTER VIII. Reactions of the metals of Group II., Division I 75 
 
 i, IX. ,, ,, ,, II., Division 2 89 
 
 APPENDIX TO CHAPTER IX. The rare metals of Group II. . 108 
 
 CHAPTER X. Reactions of the metals of Group 1 112 
 
 APPENDIX TO CHAPTER X. The rare metals of Group I. . . 1 18 
 
 CHAPTER XI. The non-metals and their acids The halogens . . 120 
 
 XII. Sulphur 137 
 
 ,, XIII. Nitrogen and phosphorus 144 
 
 ,, XIV. Carbon, silicon, boron 155 
 
 ,, XV. Systematic detection of the acids 171 
 
 ,, XVI. Preliminary tests and operations in a systematic 
 
 analysis 178 
 
 ,, XVII. The results of a qualitative analysis 186 
 
 BOOK II. 
 QUANTITATIVE ANALYSIS. 
 
 PART I. GRAVIMETRIC METHODS. 
 
 SECTION I. Preliminary manipulations 189 
 
 (i) Weighing. 
 
x " Table of Contents. 
 
 PAGE 
 
 (2) Drying and weighing a filter. 
 
 (3) Estimation of water of crystallisation. 
 
 (4) Determination of the ash of a filter. 
 
 (5) Preparation of pure salts. 
 
 SECTION II. Typical gravimetric determinations of metals . 221 
 Al, Cr, Fe, Ca, Mg, Cu, Ag, Pb, Zn, Mn, Ni, Co, K, 
 
 NH 4 , Sn, As, Sb, Cd, Hg. 
 
 SECTION III. Typical gravimetric determinations of acids . . 258 
 H 2 SO 4 , HC1, HBr, CO 2 , HNO 3 , H 3 PO 4 , SiO 2 . 
 
 S'ECTION IV. Exercises in gravimetric analysis 265 
 
 Silver coin, solder, German silver, bronze, dolomite, zinc 
 
 blende, silicate. 
 SECTION V. Electrolytic methods of analysis 292 
 
 PART II. VOLUMETRIC METHODS. 
 
 SECTION I. Preliminary manipulations 300 
 
 (1) Instruments for measuring liquids Calibration of 
 vessels. 
 
 (2) Standard solutions. 
 
 (3) Methods for ascertaining the completion of volumetric 
 reactions Indicators. 
 
 SECTION II. Volumetric methods based upon saturation 
 
 Alkalimetry, acidimetry 316 
 
 Typical analyses by means of standard acids and alkalies. 
 SECTION III. Methods based upon oxidation and reduction . 331 
 Potassium permanganate Titration with iron, ferrous salt^ 
 
 oxalic acid. 
 T)pical analyses by means of permanganate. 
 
 (1) Iron ores. 
 
 (2) Available oxygen in manganese ores. 
 
 (3) Estimation of tin. 
 
 (4) Estimation of calcium. 
 
 Potassium dichromate Titration with ferrous salts. 
 Analyses by means of dichromate. 
 
 (1) Estimation of iron in iron ores. 
 
 (2) Estimation of chromium in chrome iron ore. 
 Iodine solution Titration with thiosulphate and arsenious 
 
 oxide. 
 Estimations by means of iodine. 
 
 (1) Antimony. 
 
 (2) Arsenic. 
 (2) Tin. 
 
 Estimations bv means of iodine and thiosulnhate. 
 
Table of Contents. xiii 
 
 PAGE 
 
 (1) Sulphur dioxide in a sulphite. 
 
 (2) Available chlorine in bleaching powder. 
 
 (3) Estimation of manganese dioxide. 
 
 SECTION IV. Methods based upon precipitation .... 360 
 (I.) Precipitation with silver nitrate Chlorine, cyanogen. 
 Indirect estimation of nitric acid, calcium^ carbon 
 
 dioxide^ and cadmium. 
 
 (II.) Precipitation with sodium chloride Silver. 
 (III.) Precipitation with ammonium thiocyanate. 
 
 (1) Silver. 
 
 (2) Indirect estimation of chlorine. 
 
 (IV.) Precipitation by means of uranium acetate. 
 
 Estimation of phosphoric acid in bone ash. 
 (V.) Precipitation with sodium sulphide. 
 
 Estimation of zinc in zinc ores. 
 
 (VI.) Clark's method for estimating hardness of water. 
 APPENDIX TO SECTION IV. Estimation of copper by potassium 
 
 cyanide 374 
 
 SECTION V. Gas analysis 377 
 
 (1) Estimation by absorption and subsequent titration. 
 
 (a) Carbon dioxide in air (Pettenkofer's method). 
 
 (b) Sulphur dioxide in furnace gas. 
 
 (2) Estimation by absorption, and measurement of the 
 residual gas. 
 
 Measurement of gases. 
 
 Calibration of gas-burettes. 
 
 Correction of gaseous volumes. 
 
 Collection of samples for analysis. 
 
 Absorption in Hempel's gas-pipettes. 
 
 Analysis of mixture containing CO 2 , CO, O, N. 
 
 (3) Estimation by combustion, and determination of the 
 products. 
 
 Hydrogen by combustion with palladium. 
 
 Hydrogen by explosion. 
 
 Marsh-gas by explosion. 
 
 Analysis of mixtures containing H, CH 4 , N. 
 
 The nitrometer. 
 
 PART III. 
 
 SECTION i. Estimation of carbon, hydrogen, nitrogen, 
 
 chlorine, sulphur, and phosphorus in organic compounds . . 414 
 SECTION II. Miscellaneous physico-chemical determinations . 434 
 Specific gravity of solids and liquids. 
 Boiling-point. 
 
xiv Table of Contents. 
 
 Melting-point. 
 
 Vapour density. 
 
 Molecular weight by freezing-point. 
 
 Molecular weight by boiling-point. 
 
 APPENDIX 
 
 Table of elements and atomic weights 453 
 
 Reagents System of standard strengths 454 
 
 Table of hardness 459 
 
 Expansion of water between o and 25 460 
 
 Tension of aqueous vapour from 5 to 25 461 
 
 Table of factors for the correction of gaseous volumes . . 462 
 
CHEMICAL ANALYSIS 
 
 BOOK I. 
 
 QUALITATIVE ANALYSIS. 
 
 CHAPTER I. 
 PRELIMINARY EXERCISES. 
 
 THE first step that the student must take in approaching the 
 subject of analytical chemistry, is that of making himself practically 
 familiar with certain simple operations or manipulations which he 
 will constantly be required to carry out in the course of his work, 
 and upon the dexterous and cleanly performance of which much 
 of his success as aa analyst will depend. If he has not had previous 
 experience in practical chemistry, therefore, he should carefully go 
 through the following exercises. 
 
 i. Filtration. The method by which a liquid is separated 
 from any solid substance with which it is mechanically mixed, is 
 most usually that of filtering the mixture through porous paper, 
 known 'AS filter-paper. 
 
 EXERCISE i. Fold a circular filter-paper into half, and then at 
 right angles into half again. Open this into a cone having one 
 thickness of paper on one side and three on the other. This cone 
 is then placed in a glass funnel of such a size that the glass will 
 project slightly above the paper. The paper is then moistened 
 with distilled water, which should not be poured out of the funnel 
 again, but allowed to run through. After being cautiously pressed 
 into the glass funnel, the paper should fit close to the glass all 
 round, leaving no air-spaces. If this is not the case, either another 
 funnel of the right angle (60 degrees) should be selected, or another 
 filter-paper folded so that the cone shall be of the same angle as 
 the funnel. This can be done by making the second fold of the 
 
 B 
 
2 Qualitative Analysis. 
 
 paper not quite at right angles to the first. In this way a cone 
 will be formed having either a more acute or more obtuse angle 
 than 60 degrees, as the paper is opened out one side or the other. 
 The funnel is supported by a metal or wooden stand. 
 
 Now place some diluted hydrochloric acid in a small beaker, 
 and stir into it, by means of a glass rod, a quantity of finely 
 powdered charcoal. When thoroughly mixed, pour upon the filter. 
 When slowly pouring from a wide vessel like a beaker, there is 
 risk of some of the liquid being spilt by running down the outside 
 of the vessel, as shown in Fig. i. If it be poured quickly, it is 
 
 FIG. i. 
 
 FIG. 2. 
 
 likely to splash over the funnel. To prevent both of these acci- 
 dents, the liquid should be poured down against a glass rod held 
 lightly against the edge of the beaker, and in such a position that 
 the liquid does not strike at once against the apex of the paper 
 cone (Fig. 2). 
 
 The filtrate (i.e. the liquid which passes through the filter) may 
 be received in another beaker, which should be placed close against 
 the stem of the funnel, so that the liquid shall run down against 
 the glass. In this way splashing is prevented. The filtrate should 
 be perfectly clear, the whole of the solid being retained on the 
 filter. When all the liquid has passed through, the charcoal and 
 the filter-paper are both still soaked with the hydrochloric acid. 
 In order to remove this, and so to make the separation of the solid 
 from the liquid complete, the filter and its contents must be washed 
 
Preliminary Exercises. 
 
 with distilled water.* This is done by directing a fine stream of 
 water from a wash-bottle into the funnel, working downwards from 
 the upper edges of the paper, and so washing the charcoal down 
 into the apex of the filter (Fig. 3). Each washing must be allowed 
 to drain right through before more water is used. This must be 
 continued until the filtrate is entirely free from acid, which may 
 be ascertained by allowing one or two drops of it to fall upon a 
 piece of blue litmus paper. 
 
 In practice, the size of the filter should bear a rational relation 
 to the quantity of solid matter to be separated from a liquid. This 
 is more especially impor- 
 tant when the material re- 
 tained upon the filter has 
 to be washed. If the 
 amount of solid is small, 
 the filter used should be 
 proportionately small, and / 
 the washing operation will 
 be more quickly and effec- 
 tually accomplished than if 
 an unduly large filter is 
 employed. 
 
 2. Solution. This 
 term is applied both to the 
 act of dissolving and to the 
 product obtained by dis- 
 solving. 
 
 EXERCISE 2. Place a 
 little powdered potassium 
 carbonate in a test-tube, 
 and add a small quantity 
 of water. In a few moments 
 the salt will have entirely 
 dissolved. The salt has 
 undergone solution in water. The product is a solution of potas- 
 sium carbonate. The water is called the solvent. The process of 
 solution is accelerated by heating the liquid, and it takes place more 
 quickly the more finely the solid is powdered. 
 
 Put a similar quantity of potassium carbonate into another 
 test-tube, and add a little dilute nitric acid. The salt again under- 
 goes solution, the acid here being the solvent. But in this case 
 there is a radical difference. First, a -visible difference, in that the 
 
 * In the following exercises, and in all analytical operations, distilled 
 water must always be employed ; and when beakers, test-tubes, etc., are 
 washed up after use, they must be finally rinsed with distilled water. 
 
 FIG. 3 . 
 
Qualitative A nalysis. 
 
 act of solution is accompanied by an effervescence, or rapid evolu- 
 tion of gas ; and second, an invisible difference ; for the resulting 
 liquid is not a solution of potassium carbonate, but of potassittm 
 nitrate. In the first case, the process is not accompanied by any 
 chemical change ; the operation is therefore called simple solution : 
 the original substance is present in the liquid, and can be obtained 
 in its former state by evaporating the water. In the second case, 
 the process is distinguished as chemical solution, because chemical 
 action took place between the substance dissolved and the solvent, 
 and the original substance cannot be got back by evaporating the 
 solvent. 
 
 3. Evaporation. The process of changing from the liquid 
 to the gaseous or vaporous state is known as evaporation. This 
 operation is greatly accelerated by the application of heat. When 
 it takes place without the aid of external heat, the process is spoken 
 of as spontaneous evaporation. 
 
 EXERCISE 3. Pour the two solutions obtained in Exercise 2 
 into separate porcelain evaporating-dishes, and heat them gently 
 by means of a Bunsen with a "rose" burner (as shown in Fig. 4). 
 Continue the operation until all the liquid 
 has evaporated away and a dry residue 
 is left. This is called evaporating to dry- 
 ness. As the condition of dryness is ap- 
 proached, the flame must be turned down 
 more and more, to prevent the substance 
 from " sputtering." Try to conduct the 
 operation so that as little as possible of 
 the residue is lost in this way. 
 
 The two residues may now be exa- 
 mined by one simple test, which will 
 prove that the one from the watery solu- 
 tion is the same as it was before being 
 dissolved, and that the other is quite dif- 
 ferent. Add to each a few drops of dilute 
 F i Gi 4 . nitric acid : the first dissolves with effer- 
 
 vescence, as did the original potassium 
 carbonate ; the other is unacted on by the acid. 
 
 Sometimes it is necessary to carry on the operation of evapora- 
 tion more carefully than can be done by heating the dish in the 
 manner described. In this case the process is conducted upon 
 a steam-bath. Water is boiled in a metal vessel (resembling a 
 saucepan), and the evaporating-dish, supported by a metal ring 
 which forms the cover, is heated by the steam. The following 
 exercise is a case in point. 
 
 EXERCISE 4. Dissolve some crystals of ammonium nitrate in 
 a little water ; place half the solution in a dish, and evaporate it 
 
Preliminary Exercises. 5 
 
 over a rose burner. Evaporate the other portion in a dish upon ;i 
 steam-bath. Note the difference in the results in the two cases. 
 
 4. Fusion is the term used to denote the process of changing 
 a substance from the solid to the liquid state by the action of heat. 
 Thus, when lead is heated it enters into a state of fusion, or, shortly, 
 it fuses or melts. Fusion must not be confounded with solution. 
 Chemical action often takes place when one of the reacting sub- 
 stances is in a condition oi fusion, which is incapable of taking 
 place when they are only in solution. For example 
 
 EXERCISE 5. Dissolve a small piece of potassium hydroxide 
 (caustic Potash} in water, and add to the colourless solution a minute 
 quantity of powdered manganese dioxide. No chemical action 
 takes place. 
 
 Place a similar piece of potassium hydroxide in a dry test-tube, 
 and heat it : the solid fuses to a colourless liquid. Drop into the 
 fused mass a few particles of the manganese dioxide. Chemical 
 action at once takes place, resulting in the formation of the deep 
 green-coloured compound, potassium manganate. (This reaction 
 is used as a test for manganese compounds.) 
 
 EXERCISE 6. Place a small quantity of powdered barium 
 sulphate in a test-tube, add water, and boil for a minute or two. 
 If the amount of barium sulphate is quite small, it will be easy to 
 see that practically none of it dissolves. Allow it to settle, and pour 
 a few drops of the liquid upon a watch-glass, and set it to evapo- 
 rate to dryness on a steam-bath. 
 
 Treat another similar quantity of the barium sulphate with 
 dilute hydrochloric acid, and evaporate a few drops of the liquid 
 in the same way. The result of these two operations will prove 
 that barium sulphate is insoluble in either water or hydrochloric add. 
 
 Next dissolve a little sodium carbonate in water, and add to the 
 clear solution a few particles of barium sulphate ; boil the liquid, 
 and observe that no change takes place. 
 
 Now carefully mix a small quantity of barium sulphate with 
 about five times as much sodium carbonate ; place the powder in 
 a platinum crucible, supported on a pipe-clay triangle in the manner 
 shown in Fig. 5, and heat strongly with a blowpipe. When the 
 mass has been in complete fusion for a few minutes, allow the 
 crucible to cool. Then place it on its side in a small beaker with 
 a little water, and warm gently. The mass in the crucible will soon 
 become disintegrated, some of it dissolving, while a part remains 
 undissolved. Filter the liquid as in Exercise I, washing the residue 
 upon the funnel until the filtrate no longer restores the blue colour 
 to reddened litmus paper. Now pour a few drops of dilute hydro- 
 chloric acid upon the residue on the filter, receiving the liquid 
 which passes through in a fresh beaker or test-tube. Observe that 
 effervescence at once takes place. But this residue cannot be 
 sodium carbonate, because that salt, being soluble in water, has 
 
6 Qualitative Analysis. 
 
 been all removed ; neither can it be barium sulphate, for that com- 
 pound has been shown to be insoluble in dilute hydrochloric acid. 
 By the process of fusion, the sodium carbonate and barium sulphate 
 underwent a chemical reaction, resulting in the formation of sodium 
 sulphate and barium carbonate. The former salt, being soluble in 
 water, was dissolved in that liquid along with the excess of sodium 
 carbonate. The barium carbonate is insoluble in water, but dis- 
 solves in dilute hydrochloric acid, forming barium chloride (soluble) 
 
 and carbon dioxide, which 
 escapes as gas. Therefore, 
 by fusion the insoluble 
 barium sulphate is converted 
 into soluble barium carbo- 
 nate. 
 
 5. Precipitation. 
 When chemical action 
 takes place between sub- 
 stances in solution, and one 
 of the products of the action 
 is insoluble, the latter sub- 
 stance is thrown out of solu- 
 tion, or precipitated. The 
 substance so thrown down is 
 termed a precipitate. 
 
 EXERCISE 7. Dissolve 
 a minute quantity of sodium 
 chloride (common salt) in 
 water in a test-tube. In 
 another tube dissolve a small 
 crystal of silver nitrate, and 
 mix the two solutions to- 
 gether. The two compounds 
 react upon each other, form- 
 ic. 5. " ing sodium nitrate (soluble 
 in water) and silver chloride 
 (insoluble in water). The insoluble white precipitate is therefore 
 the silver chloride. 
 
 If, in the above example, the two substances are mixed in a 
 particular proportion, there may have been exactly the amount of 
 sodium chloride necessary to supply chlorine enough to unite with 
 the whole of the silver in the silver nitrate used. In this case there 
 would be nothing left in solution but sodium nitrate, i.e. no excess 
 of either silver nitrate or of sodium chloride. Ascertain if this 
 happened to be the case in Exercise 7, by the following experi- 
 ment : 
 
Preliminary Exercises. 7 
 
 EXERCISE 8. Filter the mixture obtained above, and divide 
 the filtrate into two portions. To one add a single drop of a 
 solution of sodium chloride, (i) If a precipitate is formed, it 
 proves that there is some silver nitrate present, and that therefore 
 an excess of this compound was used in Exercise 7. Continue 
 adding the sodium chloride solution one drop at a time,* shaking 
 or stirring the liquid after every addition, so long as it produces 
 further precipitation. 
 
 (2) If no precipitate is thrown down by adding sodium chloride, 
 add to the second portion of the filtrate a single drop of silver nitrate 
 solution. If this gives a precipitate, it proves that sodium chloride 
 is present, and that therefore an excess of this substance had been 
 employed in Exercise 7. Continue adding the silver solution drop 
 by drop, with constant stirring, 30 as to hit off as nearly as possible 
 the exact point when it just ceases to produce any further pre- 
 cipitate. 
 
 The exact point at which precipitation is complete is not equally 
 easy to determine in all cases. Some precipitates are heavy, 
 granular, or crystalline, and settle quickly ; others again are light 
 or flocculent, and only subside slowly and imperfectly, so that it is 
 difficult to see whether the addition of more of one of the solutions 
 does or does not produce any additional precipitate. In such cases 
 the liquid should be filtered, and the filtrate tested by adding a few 
 drops more of the precipitant. 
 
 Very often several substances present together in one and the 
 same liquid, form insoluble compounds with another which is 
 added. These will not be all precipitated simultaneously, but in a 
 certain order one after the other, the precipitation of one being 
 more or less complete before that of the next begins. The substance 
 being added, first selects the compound present for which it has 
 the greatest chemical affinity, and afterwards that with which it 
 unites less eagerly. This being so, it will be evident that unless 
 care be taken to ensure complete precipitation, it might easily 
 happen that the whole of one of the substances present in the 
 solution escaped precipitation. It is of the utmost importance, 
 therefore, in analysis, to be quite sure that precipitation is as com- 
 plete as possible. On the other hand, the reckless addition of pre- 
 cipitants is a fault which must be as carefully guarded against, as 
 it is almost as fruitful a source of trouble as the other. 
 
 In most instances, also, it is essential to wash the precipitate 
 until it is quite free from any of the soluble substances present in 
 
 * When solutions are to be added drop by drop, it is best to use a pipette ; 
 that is, a piece of ordinary glass tube drawn to a point at one end, and about 
 6 or 8 inches long. 
 
8 Qualitative Analysis. 
 
 the liquid, as explained in Exercise i. A precipitate may be re- 
 moved from the filter either by means of a spatula (preferably 
 platinum, but, failing this, either glass or porcelain ; iron should 
 never be used), or by pushing a glass rod through the apex of the 
 filter, and then washing the precipitate through by means of the 
 wash-bottle, or by dissolving it off by pouring into the funnel the 
 liquid to be used for its solution. 
 
 EXERCISE 9. Add a solution of sodium carbonate to a solution 
 of barium chloride, until precipitation is just complete. Barium 
 carbonate is precipitated, and sodium chloride remains in solution. 
 Pour the mixture upon three separate filters, and wash the pre- 
 cipitate on each until quite free from sodium chloride (see Exercises 
 7 and 8), getting the precipitate well down into the apex. 
 
 Take the first filter, and remove a portion of the precipitate with 
 a spatula. If the quantity in the funnel is small, then carefully draw 
 the paper cone out of the funnel, spread it open upon a flat sheet of 
 glass, and scrape off as much of the precipitate as possible with the 
 spatula, and transfer it to a test-tube. Dissolve it by adding a few 
 drops of hydrochloric acid. 
 
 Through the apex of the second filter push a glass rod, and wash 
 the precipitate through into a test-tube, using a fine jet of water, 
 and as little of it as possible. Dissolve this also by adding a few 
 drops of the same acid. 
 
 Upon the third filter pour a small quantity of hydrochloric acid, 
 collecting the filtrate in a test-tube. Pour the filtrate back over 
 the precipitate once or twice, until the whole has dissolved. 
 
 6. Ignition. Strictly speaking, this word carries with it the 
 idea of combustion. In common speech it signifies the act of 
 "setting fire" to an inflammable substance ; and in more scientific 
 language we speak of the ignition temperature, or the igniting- 
 point of a body, meaning thereby the temperature, to which it is 
 necessary to raise it in order that combustion may be initiated. 
 Unfortunately, in analytical phraseology the term ignition is used 
 in a somewhat slipshod way to denote a variety of operations 
 where substances are simply strongly heated, and where the idea 
 of combustion is altogether excluded. In this book the words heat 
 or strongly heat will be used instead of ignite to signify these 
 operations. 
 
 EXERCISE 10. Strongly heating in an open dish. Place a 
 little solution of ammonium chloride in a small evaporating-dish, 
 and evaporate to dryness. Then strongly heat the dish with the 
 dry residue until no more white fumes (consisting of the volatilising 
 ammonium chloride) are evolved. If the dish has been heated all 
 over there should then be nothing left in it. The complete 
 vaporisation of the salt is more quickly and certainly accomplished 
 
Preliminary Exercises, 9 
 
 by using a small platinum capsule or crucible in which to heat the 
 residue obtained by evaporating the solution to dryness. 
 
 EXERCISE n. Strongly heating in a tube closed at one end. 
 Place a minute quantity of mercuric oxide in a small test-tube 
 (4 inches x f %), and apply heat to the compound. Note the change 
 of colour; also that it gradually disappears, and that a sublimate 
 collects on the cool part of the tube, having a white metallic 
 appearance. Test the evolved oxygen by means of a glowing.splint 
 of wood. By means of a paper " spill " rub the metallic sublimate, 
 and (if necessary, with a pocket lens) see the globules of mercury. 
 
 EXERCISE 12. Heating in the blowpipe flame. Select a piece 
 of small tubing of lead glass, and heat it in a blowpipe flame, 
 holding the glass in the extreme tip of the flame until it is red hot. 
 Then gradually bring it further into the flame, and observe that 
 when the glass reaches the inner cone of the flame a film begins to 
 appear upon the red-hot portion. On withdrawing the glass to the 
 tip of the flame again, this film gradually disappears. Bring the 
 glass once more into the inner cone of flame, and when the film 
 has again made its appearance, remove the glass and allow it to 
 cool. It will then be seen that what appeared like a film when it 
 was hot, is a black shining metallic-looking deposit in the glass. 
 This deposit is metallic lead. The lead compound in the glass, 
 when heated in the inner cone of flame, is reduced to the metallic 
 state ; and when, after being so reduced, it is heated in the tip of 
 the flame (i.e. in the outer cone or sheath of the flame), the metal is 
 again oxidised. The inner flame is therefore called the reducing 
 flame, and the outer cone is distinguished as the oxidising flame.* 
 
 EXERCISE 13. Heating on charcoal in the blowpipe flame. 
 Select a close-graine<J piece of charcoal, as free as possible from 
 cracks, and file a flat surface upon it with a broad, flat file.t On 
 the flat part scoop a small hollow, and place in it a little red-lead 
 mixed with about an equal quantity of sodium carbonate. Heat 
 this mixture in the inner blowpipe flame, holding the blowpipe 
 and the charcoal in the manner shown in Fig. 6, so that the flame 
 shall play along the surface of the charcoal. Very quickly the lead 
 oxide will be reduced to the metallic state, and appear in the form 
 of brilliant silvery globules. When the charcoal is removed, it 
 will be seen that surrounding the cavity there is a yellowish deposit, 
 or incrustation. This consists of lead oxide. If the outer tip of 
 the flame be directed upon this incrustation it will quickly disappear, 
 and will impart a bluish colour to the end of the flame. 
 
 Pick out one or two of the globules of metal, and gently strike 
 one with a small hammer, or with a pestle, upon some hard surface. 
 
 * The memory of the beginner may be aided by the alliteration, Outer, 
 Oxidising. The inner flame is a reducing agent by reason of the fact that 
 within the cone there is an excess of strongly heated coal-gas ; whereas in the 
 outer flame there is an excess of heated atmospheric oxygen. 
 
 f- Specially prepared rectangular blocks of charcoal (6 inches long and i 
 square inch section) ;ire sold for the purpose. One such block can be used 
 many times. 
 
10 
 
 Qualitative A nalysis. 
 
 Note whether the metal is hard and brittle, or soft and malleable. 
 Also further identify the metal as lead by rubbing it upon a piece 
 of paper, which will be marked by it much as by an ordinary 
 pencil. 
 
 EXERCISE 14. Heat on another piece of charcoal a crystal or 
 two of zinc sulphate with a little sodium carbonate in the inner 
 blowpipe flame. No metallic globules are formed in this case, 
 because zinc is too easily oxidised ; but an incrustation appears on 
 the charcoal, which is canary-yellow while hot, but turns white on 
 cooling. Touch the incrustation with a single drop of a solution 
 of cobalt nitrate, and again heat it, using the outer flame. The 
 
 FIG. 6. 
 
 incrustation then becomes green. Notice that the incrustation is 
 not driven off by being thus heated, because zinc * oxide is not 
 volatile. 
 
 7. Fusion with Borax. When borax is strongly heated, it 
 melts to a clear vitreous mass. In this condition it is capable at a 
 high temperature of dissolving many metallic compounds, giving in 
 some cases characteristically coloured glasses. 
 
Preliminary Exercises. II 
 
 EXERCISE 15. Twist the end of a piece of platinum wire into 
 a small round loop or eye,* and pick up a little borax upon it by first 
 heating the wire and then dipping it while hot into the powdered 
 salt. On heating the borax upon the wire in a blowpipe flame, it 
 first swells up, and finally fuses, forming a transparent colourless 
 bead of borax glass. Allow the bead to cool, and touch it with a 
 glass rod which has been dipped into a solution of any cobalt salt, 
 so as to bring only a minute quantity of the cobalt compound upon 
 the bead. Heat the bead once more, and notice that as it melts 
 the borax loses its transparent appearance. When again allowed 
 to cool, the bead will appear of an azure blue colour. 
 
 If too much of the cobalt salt was employed, the bead may 
 appear almost black ; in this case a part of it may be shaken off 
 when it is fluid, and more borax picked up and melted with what 
 remains of the original bead upon the wire. If too little cobalt is 
 present, the colour will be correspondingly pale. The colour of 
 the bead is best examined by holding it against a white object 
 (such as the bottle of borax itself) in a good light. 
 
 Fuse the bead again, holding it first in the outer flame, and 
 afterwards in the inner flame, and see that in each case when 
 cold the blue colour remains the same. 
 
 EXERCISE 16. Make another borax bead, and touch it with 
 a small quantity of a solution of manganous sulphate. Heat this 
 in the outer blowpipe flame. After cooling, examine the colour 
 carefully. Pale violet, lilac, purple, or amethyst. Heat the bead 
 again, holding it in the inner flame. Notice that it gradually 
 loses its opacity ; that as it is heated, something in the fused mass 
 which seems to give it an appearance of muddiness clears away, 
 and the molten globule looks clear. When it is in this condition 
 remove it, and when cold it will be found to have lost its colour 
 entirely. Manganese compounds there- 
 fore give a purplish bead in the outer 
 flame, which becomes colourless upon 
 being heated in the reducing flame. 
 
 8. Neutralisation. \Vhen an acid 
 is carefully mixed with an alkali (the 
 substances being in solution), a point is 
 
 * For greater convenience, as well as economy, 
 a short piece of wire (about 2 inches) should be 
 fixed into a glass tube, about the same length, 
 to serve as a handle. The glass tube is first drawn 
 out to a point, and the wire inserted into the fine 
 end. On bringing this into a blowpipe flame, 
 the glass fuses round the wire and holds it. 
 Two or three of these should be made, and a 
 convenient plan is to fit the glass tube into a 
 cork, so that when not actually in use the wires 
 
 can be kept in small test-tubes containing dilute hydrochloric acid, as in 
 Fig- 7- 
 
12 Qualitative Analysis. 
 
 reached when the mixture no longer possesses the properties of 
 either the acid or the alkali. The solution is then said to be 
 neutral. The point of neutrality is ascertained by the use of 
 certain sensitive colouring matters which have their colour changed 
 by acids and alkalies. The commonest of these is litmus, the 
 solution of which in water has a purple colour, capable of being 
 turned red by acids, and blue by alkalies. The yellow colour of 
 turmeric is changed to brown by alkalies, but is not altered by 
 acids, therefore this can only be used to indicate alkalinity, and 
 will not discriminate between a neutral and an acid liquid. 
 
 EXERCISE 17. Add a few drops of litmus solution to a little 
 dilute hydrochloric acid in a beaker standing upon a piece of white 
 paper, or a white tile. Add to the red liquid some solution of 
 sodium hydroxide, adding it cautiously in small quantities, with 
 constant stirring, until the colour of the litmus is just turned blue. 
 The liquid is now alkaline. By means of a glass rod moistened 
 with the dilute acid, introduce a minute additional quantity of the 
 acid, so as to cause the colour of the litmus to become of a purple 
 tint. The solution is then neutral, and the least trace of either 
 acid or alkali will at once turn it red or blue, as the case may be. 
 (Instead of adding litmus solution, papers tinted with litmus may 
 be dipped into the liquid.) 
 
 All substances which redden litmus are not acids, although all 
 acids will redden litmus. That is to say, there are many things 
 which have an acid or sour taste which do not belong to that class 
 of compounds which chemists call adds. An acid may be denned 
 as a compound containing hydrogen -which can be displaced by a 
 metal, when the latter is presented to it in combination as a 
 hydroxide. The hydrogen which is thus displaced 'is not liberated 
 as free hydrogen, but unites with the hydrogen and oxygen of the 
 metallic hydroxide to form water, while the metal takes the place 
 of the hydrogen thus displaced from the acid. 
 
CHAPTER II. 
 ANALYTICAL CLASSIFICATION. 
 
 THE word analysis, in its strict meaning, signifies the breaking 
 up or separation of a compound substance into its constituent 
 parts. It is the true antithesis of the word synthesis, which 
 means the building up of a compound from its constituents. For 
 example, when an electric current is passed through a solution of 
 common salt, the compound is split up or separated into its two 
 component elements, sodium and chlorine. The operation is there- 
 fore " analytical." When sodium and chlorine are brought together 
 under suitable conditions, the two elements unite, and reproduce 
 the compound sodium chloride, the process in this case being a 
 synthetical one. 
 
 But the word analysts has come to bear a wider meaning, 
 and to include all the various processes and operations which 
 chemists make use of in order to find out what any compound is 
 composed of, or to enable them to identify the substance, quite 
 irrespective of whether or not the process involves the breaking up 
 of the body into its component parts. Thus, a chemist will often 
 recognise a substance by its particular crystalline form, or from 
 some other characteristic appearance it may present when examined 
 under a microscope (microscopic analysis}. Or sometimes he can 
 detect the presence of certain elements in an unknown substance, 
 by examining the light which is emitted when the compound is 
 strongly heated (spectrum analysis). 
 
 Reactions. Most analytical operations, however, involve 
 some chemical change. These changes are called reactions. 
 When the change is effected by strongly heating the substance, 
 it is described as a dry reaction, or a reaction in the dry way. 
 This is to distinguish this class of reactions from those which take 
 place between substances that are dissolved, either in water or 
 some other liquid, and which are sometimes spoken of as wet 
 reactions, or reactions in the wet way. 
 
14 Qualitative Analysis. 
 
 Most analytical reactions are "double decompositions," in which 
 one of the products of the chemical action is either markedly 
 different from the others and from the reacting compounds, in its 
 solubility, or its colour ; or where it is evolved as a gas having 
 properties by which it may be readily identified. For instance, the 
 two compounds barium chloride and sodium sulphate are soluble 
 in water, forming colourless solutions ; if these are mixed together, 
 " double decomposition " takes place, resulting in the formation of 
 sodium chloride and barium sulphate, thus 
 
 BaCI 2 + Na 2 SO 4 = 2NaCl + BaSO 4 
 
 The barium sulphate is practically insoluble in water, and con- 
 sequently is precipitated. Now, if we know some property belong- 
 ing to this precipitate of barium sulphate which is so characteristic 
 of the compound that we could thereby identify it and distinguish 
 it from all other white precipitates, then this reaction between 
 barium chloride and sodium sulphate can obviously be used as a 
 means of testing for the presence of either a soluble barium salt 
 or a soluble sulphate. For if, on adding a solution of sodium 
 sulphate to an unknown solution, barium sulphate were precipitated, 
 the unknown liquid must have contained a soluble barium salt ; or, 
 on the other hand, if we add barium chloride to an unknown solu- 
 tion and obtain barium sulphate again, then this unknown solution 
 must have contained a sulphate.* 
 
 Reagents. The materials that are used to bring about 
 analytical reactions are termed reagents. Thus, in the illustration 
 given above, the sodium sulphate is the reagent when it is added 
 to the unknown solution in order to test for barium ; while the 
 barium chloride is the reagent when it is used to test for a sulphate. 
 Some reagents are capable of causing reactions of a similar 
 character with a number of substances ; such are often known as 
 general reagents. Others, again, are employed because they pro- 
 duce a characteristic reaction with some one substance in 
 particular ; these are distinguished as special reagents. 
 
 Reagents are the tools with which the analyst works, and upon 
 the intelligent and skilful use of them everything depends. In most 
 laboratories the student finds himself supplied with all the necessary 
 reagents ready prepared ; but for the help of those who may require 
 to make them up, brief directions for doing so are given in the 
 Appendix. 
 
 Analytical Classification. Substances are usually divided^ 
 
 * Sulphuric acid being included, as hydrogen sulphate. 
 
Analytical Classification. 15 
 
 into two classes, namely, (i) Metals, and (2) Acid-radicals. 
 These are also sometimes called positive radicals and negative 
 radicals respectively. When analysing such a compound as sodium 
 chloride, NaCl, the sodium and the chlorine are each separately 
 detected : the sodium is the metal (or positive radical), and the 
 chlorine is the acid (or negative) radical. But in such a case as 
 sodium nitrate, NaNO 3 , we do not separately detect the sodium, 
 the nitrogen, and the oxygen, but the sodium and the negative or 
 acid radical represented by the formula NO 3 . Or, again, when 
 such a compound as ammonium sulphate, (NH 4 ) 2 SO 4 , is submitted 
 to analysis, we do not separately test for the elements nitrogen, 
 hydrogen, sulphur, and oxygen, but for the positive radical NH 4 , 
 and the acid-radical SO 4 . 
 
 Sometimes the radicals, whether metals or acid radicals, may 
 be detected by being actually isolated, in which case they are 
 recognised by their known properties in the free state. For 
 example, from the compound lead chloride, PbCl 2 , it is easy to 
 isolate both the lead and the chlorine. The metal lead so obtained 
 is readily identified by its familiar physical properties, while the 
 gas chlorine is equally easily distinguished by its own well-known 
 characteristics. 
 
 In some cases, where a radical is incapable of isolated existence, 
 it may be detected by the separation of some product of its 
 decomposition. Thus, in such a compound as ammonium 
 carbonate, (NH 4 ) 2 CO 3 , neither the positive radical NH 4 , nor the 
 acid-radical CO 3 can exist in the free or uncombined state. But 
 we detect the presence of the former by the evolution of ammonia, 
 NH 3 , and the latter by the expulsion of carbon dioxide, CO 2 , from 
 the compound. 
 
 In the large majority of cases, however, whether the various 
 radicals are capable of isolated existence or not, they are detected 
 by causing them to pass into fresh combinations with certain 
 reagents, whereby new compounds are formed which are readily 
 recognised by their known properties. Thus, in the case of sodium 
 chloride above quoted, instead of isolating the chlorine, we can 
 employ the reagent silver nitrate, AgNO 3 . When this is added to 
 a solution of sodium chloride, double decomposition takes place, 
 and silver chloride, AgCl, is formed, which, being insoluble in water, 
 is precipitated. Silver chloride has properties by which it is easily 
 recognised, hence by the formation of this compound we can detect 
 the presence of the chlorine in sodium chloride. 
 
 In analytical classification the term " metal " includes, besides 
 
i6 
 
 Qualitative A nalysis. 
 
 the metals proper, certain metalloidal elements, such as arsenic, 
 selenium, and others (which, strictly speaking, are not true metals, 
 but which lie on the borderland between the metals and the non- 
 metals), and also the compound positive radical ammonium, NH 4 . 
 These "metals," then, are divided into a number of groups, 
 based on the behaviour of their compounds towards certain 
 reagents. 
 
 As, however, more than one scheme or plan of analysis is 
 possible, due either to a different choice of reagents or to their use 
 in a different order, so the analytical classification of the metals 
 which is followed by some chemists varies somewhat from that 
 used by others. 
 
 In this book the following arrangement will be adopted : 
 
 Group I. 
 
 Group II. 
 
 Group III. 
 
 Group IV. 
 
 Group V. 
 
 Silver 
 
 Mercury 
 
 Aluminium 
 
 Barium Ammonium 
 
 Mercury 
 
 Lead 
 
 Chromium Strontium Sodium 
 
 Lead* 
 
 Bismuth 
 
 Iron Calcium Potassium 
 
 
 
 Copper 
 
 Nickel Magnesium 
 
 Thallium f 
 
 Cadmium 
 
 Cobalt 
 
 
 
 Tungsten 
 
 Antimony 
 
 Manganese 
 
 
 Lithium 
 
 
 Arsenic 
 
 Zinc 
 
 
 Rubidium 
 
 
 Tin 
 
 
 
 Caesium 
 
 
 Gold 
 
 Beryllium 
 
 
 
 
 j* lat inum 
 
 Zirconium 
 
 
 
 
 
 Thorium 
 
 
 
 
 Ruthenium 
 
 i Cerium 
 
 
 
 
 Rhodium 
 
 I Scandium 
 
 
 
 
 Palladium 
 
 Yttrium 
 
 
 
 
 Osmium 
 
 Lanthanum 
 
 
 
 
 Iridium 
 
 Ytterbium 
 
 
 
 
 Tellurium 
 
 Titanium 
 
 
 
 
 Selenium 
 
 Tantalum 
 
 
 
 
 Molybdenum 
 
 Niobium 
 Uranium 
 
 
 
 
 
 , Indium 
 
 
 
 
 
 1 Thallium 
 I Vanadium 
 
 
 
 In the regular course of analysis, the groups are separated in 
 the order in which they are here numbered. The " silver, lead, 
 mercury" group is separated first, and the "ammonium, sodium," 
 etc., group last. In studying the reactions of the metals, however, 
 
 * The reason why certain metals are placed in more than one division will 
 appear later. 
 
 f The substances printed in small type are usually called rare elements. 
 
Analytical Classification. 17 
 
 it is more usual to begin with the fifth group, for the reason that 
 the compounds of these metals are less complex, and the student 
 is therefore led on gradually from what is comparatively simple 
 to that which is more difficult. Because of this, some chemists 
 prefer to number the groups in the reverse order, that is, in the 
 order in which the preliminary study of them is made, instead of 
 the order in which they are actually disposed of in the course of 
 analysis. 
 
 The reagents by means of which the elements are separated 
 into these groups, and which are known as group reagents, must 
 be used in regular order. Each is only capable of separating its 
 own family of metals from those coming after it in the series, and 
 not those going before. For example, the group-reagent for 
 Group II. is only capable of separating the metals of this family 
 from those of III., IV., and V., but not from those of Group I. 
 If, therefore, the metals of Group I. are not first separated, they may 
 be precipitated along with the members of the second family by 
 the group-reagent for that family. 
 
 The various group-reagents (or general reagents) and the 
 particular compounds of the metals which are precipitated by them 
 are indicated in the following table : 
 
 GROUP I. General reagent, Hydrochloric acid, HC1, 
 
 precipitates 
 
 AgCl, Hg'sCl,, PbCl 2 (partially soluble). 
 1, H 2 WO 4 .)* 
 
 GROUP II. General reagent, Sulphuretted hydrogen, H 2 S, 
 
 precipitates in acid solution 
 
 PbS, Hg S, Bi S , CuS, CdS \ Insoluble in ammo- 
 (Ru 2 S 3( Rh 2 S 3 , PdS, OsS.) / nium sulphide. 
 
 Sb 2 S 3 , As 2 S 3 , Sn n S and Sn iv S 2 , AuS, PtS 2 ] Soluble in 
 
 TeS 2( Se, MoS,) e 
 
 GROUP III. General reagent, Ammonium sulphide, (NH ( ) S, 
 
 precipitates in presence of ammonium chloride and ammonia 
 (a) Hydrated ( AlaCHO),,, Cr 2 (HO) 6 . 
 
 compounds ( Be < H )" Zr ( HO )4- T h(HO) 4> Ce(HO) 3 , Sc(HO) 3 , Y(HO),, 
 LS I La(HO) 3 , Yb(HO) 3 , H 2 TiO 3 , H 3 TaO 4 , H 3 NbO 4 .) 
 
 * The compounds represented by the formulas in small type are those of 
 the so-called rare elements. They are included in the table in order to give, in 
 a bird's-eye view as it were, not only their position in the scheme of classifica- 
 tion, but also the composition of their compounds, which are precipitated by 
 the group-reagents. 
 
 C 
 
1 8 Qualitative Analysis. 
 
 \ FeS, NiS, CoS, MnS, ZnS. 
 
 () Sulphides j (UO a S, InS, T1 2 S) (Vanadium 
 
 converted into soluble ammonium thiovanadate*). 
 
 GROUP IV. General reagent, Ammonium carbonate, 
 (NHj)^CO,, precipitates in presence of ammonium chloride 
 and ammonia 
 
 BaCO . SrCO , CaCO . 
 
 GROUP V. No general reagent. The group consists of 
 (NH 4 ), Na, X, Mg. (Li, Rb, Cs.) 
 
 * Vanadium belongs to the "arsenic and antimony" family in the natural 
 classification of the elements. The sulphide is, however, not precipitated by 
 H 2 S, V 2 O 5 being thereby reduced to V 2 O 4 , which gives a blue colour to the 
 liquid. " NH 4 C1, in presence of ammonia, precipitates white ammonium meta- 
 vanadate, NH 4 VO 3 , but ammonium sulphide converts this into the soluble 
 ammonium thiovanadate, which gives a brown colour to the solution. The 
 true group-reagent, therefore, does not actually precipitate this metal. 
 
CHAPTER III. 
 REACTIONS OF THE METALS OF GROUP V. 
 
 THIS group contains the alkali metals (ammonium being regarded 
 as a metal), and also the element magnesium, which is more nearly 
 allied to the metals of the alkaline earths. The members of this 
 family are not precipitated by any group-reagent, but they are 
 ^with the exception of ammonium) separately tested for in the 
 solution which is obtained after the metals of Groups I. to IV. have 
 been removed. By referring to the table on p. 17, it will be seen 
 that, in the course of separating the various groups, certain am- 
 monium compounds are employed, therefore it will be obvious that 
 it is necessary to test for this "metal " in the substance under ex- 
 amination before adding any ammoniacal compounds. 
 
 Ammonium, NH . 
 
 DRY REACTIONS. When heated alone in a glass tube, ammo- 
 nium salts undergo change. 
 
 (a) If the acid is readily volatile, the salt dissociates, but the 
 ammonia and the volatile acid, as they together pass away from 
 the heated area, immediately reunite, reproducing the original 
 compound, which then settles or condenses on the cool part of the 
 tube, forming a sublimate. 
 
 [Generally, however, a certain small amount of the dissociated 
 portions of the compound escapes recombination : e.g. heat a small 
 quantity of ammonium chloride in a dry test-tube ; notice that 
 white fumes are produced, which sublime up the tube. Now hold 
 a moistened red litmus paper in the mouth of the tube, and it will 
 be turned blue, showing that a portion of the ammonia escapes 
 from the tube before it meets the hydrochloric acid from which 
 it has been dissociated. For a moment discontinue heating, and 
 presently the blued paper will be reddened, for the molecules of 
 hydrochloric acid which have lost their partners (the escaped 
 ammonia) now make their way up the tube and act on the litmus 
 paper.] 
 
2O Qualitative Analysis. 
 
 (V) If the acid is non-volatile, or volatile only at a high tempera- 
 ture, then the ammonium salt is decomposed, ammonia being 
 evolved, while the acid remains. 
 
 [E.g. Heat a little ammonium sulphate or phosphate in a test- 
 tube ; ammonia is rapidly evolved, and may be detected by its 
 characteristic smell.] 
 
 (<r) The ammonium salts of certain oxyacids which readily part 
 with oxygen (such as ammonium nitrate, nitrite, chromate) are also 
 decomposed by heat, the ammonia being oxidised to nitrogen or 
 oxides of nitrogen. 
 
 [E.g. Heat a few crystals of ammonium nitrate in a test-tube. 
 Examine the gas with a taper and a glowing splint of wood.] 
 
 NH 4 N0 3 = 2H 2 + N 2 
 NH 4 NO 2 = 2H 2 O + N 2 
 (NH 4 ) 2 Cr 2 7 = Cr 2 3 + 4H 2 O + N 2 
 
 Ammonium is separated from the other members of the group 
 by evaporating the solution to dryness, and strongly heating the 
 residue until the ammonia is completely expelled, which may generally 
 be regarded as accomplished when fumes are no longer given off. 
 
 WET REACTIONS. Ammonium salts are all soluble in water, 
 therefore it is only in concentrated solutions that any precipitations 
 with reagents can be formed. Use ammonium chloride. 
 
 Caustic alkalies (NaHO or KHO) and oxides or hydroxides 
 of metals of the alkaline earths (e.g. CaO, Ba(HO) 2 ), when heated 
 with an ammonium salt, cause the evolution of ammonia gas, NH 3 . 
 
 (NH 4 ) 2 SO 4 + 2NaHO = Na 2 SO 4 + 2H 2 O + 2NH 3 
 2NH 4 C1 + CaO = CaCl 2 + H 2 O + 2NH 3 
 
 In practice, sodium hydroxide solution is added either to the solid 
 salt or to its solution in water, and the mixture gently warmed. 
 The evolved ammonia may be recognized (i) by its characteristic 
 odour if present in sufficient quantities ; (2) by its power of restoring 
 the blue colour to moist reddened litmus paper, or of turning tur- 
 meric paper brown ; (3) by the formation of white fumes when a 
 glass rod moistened with strong hydrochloric acid is held in the 
 mouth of the test-tube. 
 
 [In special cases, as in the examination of natural waters, minute 
 traces of ammonia are detected by the use of Nessler's solution (a 
 solution of potassium mercuric iodide in potash), which gives either 
 a brown precipitate or a coloration, according to the amount of 
 ammonia present, 2(HgI 2 ,2KI) -I- sKHO + NH 3 = NHg 2 "I,H 2 O 
 + 7KI +2H 2 O.] 
 
Group V. 2 1 
 
 Platinum chloride, PtQ 4 , precipitates from concentrated 
 solutions a yellow crystalline compound, ammonium chloro- 
 'platinate, (NH 4 ) 2 PtCl fi (sometimes called ammonium platinic 
 chloride, 2NH 4 C1, PtCl 4 ), soluble in 170 parts of water at 10 ; 
 insoluble in alcohol and ether. This compound is distinguished 
 from the similar potassium salt in that, when strongly heated, it 
 leaves a residue of spongy platinum only. 
 
 Tartaric acid, H 2 (C 4 H 4 O, ; ), or hydrogen sodium tartrate, 
 HNa(C 4 H 4 O 6 ), produces in concentrated solutions a white pre- 
 cipitate of hydrogen ammonium tartrate, H(NH 4 )(C 4 H 4 O 6 ). Soluble 
 in water, in mineral acids, and in alkalies; insoluble in alcohol. 
 This compound is distinguished from the corresponding potassium 
 salt by the fact that when strongly heated, the carbonaceous residue 
 is without any alkaline reaction. 
 
 [When tartaric acid is used, the acid previously in combination 
 with the ammonia is liberated by the double decomposition, thus 
 
 NH 4 C1 + H 2 (C 4 H 4 6 ) = H(NH 4 )(C 4 H 4 6 ) + HC1 
 
 And as the precipitate is soluble in mineral acids, the delicacy of the 
 
 reaction is increased by employing hydrogen sodium tartrate as the 
 
 reagent, in which case no free acid is formed in the reaction, thus 
 
 NH 4 C1 + HNa(C 4 H 4 G ) = H(NH 4 )(C 4 H 4 O 6 ) + NaCl] 
 
 Sodium, Na. 
 
 DRY REACTION. Sodium compounds, when heated upon a 
 platinum wire in a Bunsen flame, undergo volatilisation, and 
 impart to the flame a brilliant golden yellow colour. This flame- 
 reaction is the most characteristic and delicate test for this metal.* 
 WET REACTIONS. All sodium salts are soluble ; sodium platino- 
 chloride is soluble in water, in alcohol, and in ether. Hydrogen 
 sodium tartrate also is freely soluble in water. Sodium pyroanti- 
 monate,t however, is less soluble in water than the corresponding 
 
 * When the light emitted by heating a sodium salt in the Bunsen flame is 
 examined by the spectroscope (see p. 27), it is found to be monochromatic, i.e. 
 to consist of one colour only, namely, pure yellow light. Any coloured object 
 whose colour contains no yellow in its composition, when illuminated only by 
 the light emitted by the sodium flame, appears black. The red colouring matter 
 in blood, for example, contains no yellow ; hence blood, when seen only in the 
 light of the sodium flame, looks black: for this reason, the hands and face of a 
 person illuminated only by the sodium flame appear black and dirty in propor- 
 tion as the individual has a high colour or not. Indigo, and many other 
 common coloured materials, similarly absorb the yellow light emitted by sodium, 
 hence if the sodium flame be viewed through a thin stratum of such a coloured 
 solution, the yellow light is entirely intercepted. For the use that is made of 
 
 this property, see potassium (p. 22). 
 f Formerly misnamed sodiu 
 
 iium metantimonate. For a comparison of phos- 
 phates, arsenates, and antimonates, see Newth's " Inorganic Chemistry," p. 458. 
 

 
 22 
 
 Qualitative Analysis. 
 
 potassium salt, and is therefore precipitated by the addition of a 
 strong solution of potassium pyroantimonate to a strong solution of 
 a sodium salt, such as sodium chloride, thus 
 
 H 2 K 2 Sb 2 O 7 + 2NaCl = H 2 Na 2 Sb 2 O, + 2KC1 
 
 Potassium, K. 
 
 DRY REACTION. When potassium compounds are heated upon 
 a platinum wire in a Bunsen flame, they impart to the flame a pale 
 violet or lilac colour. This delicate tint, however, is completely 
 masked by the intense yellow colour which the presence of even 
 minute quantities of sodium compounds impart to the flame. 
 
 Introduce a fragment of potassium nitrate into the Bunsen 
 flame upon a loop of clean platinum wire ; * notice the lilac colour 
 imparted to the flame. Now look at the flame through a potassio- 
 scopefi ,and observe that it appears a brilliant crimson- red colour. 
 Upon another wire introduce a particle of sodium chloride into the 
 flame, and notice that when this is examined through the potassio- 
 scope, the intense golden yellow light is absolutely cut off, and is 
 invisible. Now touch the wire containing the nitre with a fragment 
 of sodium chloride, and again bring it into the flame. The yellow 
 of the sodium completely overpowers and masks the violet of the 
 potassium when viewed direct, but if looked at through the potassio- 
 scope, the red colour due to the potassium shines up as brilliantly as 
 before, while the yellow is completely intercepted \ (see also p. 33). 
 
 * By merely touching the wire with the fingers, it contracts sufficient sodium 
 compounds to give the yellow flame. To clean it, it should be dipped in hydro- 
 chloric acid and heated until it ceases to 
 impart any colour to the flame. 
 
 f The potassioscope is a little instru- 
 ment consisting of a flat circular glass cell 
 filled with a solution of one of the rose- 
 aniline blue colouring matters. It possesses 
 the advantage over the older indigo prism, 
 in that it not only makes it possible to de- 
 tect potassium in the presence of sodium 
 (which a solution of indigo does equally 
 well), but, owing to the fact that it entirely 
 intercepts the red light emitted by lithium, 
 barium, strontium, and calcium, it ren- 
 ders it possible to detect potassium with 
 certainty, even when any of these other 
 elements are present. 
 
 When studying flame reactions, it is 
 often of the greatest convenience to use a 
 stand on which to support the platinum 
 wire, so that the hands may be free ; a 
 simple stand is readily constructed as 
 shown in Fig. 8. A piece of glass tube 
 or glass rod is inserted in a large cork 
 (rubber, being heavier, makes a steadier 
 loot), and a piece of galvanized iron wire is twisted two or three times round 
 
 FIG. 8. 
 
Group V. 23 
 
 WET REACTIONS. Most potassium salts are soluble in water. 
 Use a solution of potassium chloride. 
 
 Platinum chloride, PtCl 4 , produces, with concentrated solu- 
 tions of potassium salts, a yellow crystalline precipitate of potassium 
 chloro-platinate (or potassium platinic chloride), K 2 PtCl 6 , soluble 
 in 1 10 parts of water at 10 (therefore more soluble than the 
 corresponding ammonium compound). Soluble in alkalies (there- 
 fore the solutions used should be acid) ; nearly insoluble in alcohol ; 
 quite insoluble in a mixture of alcohol and ether (therefore the . 
 precipitation of this compound is promoted by the addition of 
 alcohol). 
 
 Hydrogen sodium tartrate, HNa(C 4 H 4 O 6 ), gives, with solu- 
 tions of potassium salts, a white crystalline precipitate of hydrogen 
 potassium tartrate, HK(C 4 H 4 O 6 ) ; soluble in much water, and also 
 in acids and alkalies (therefore the solution should be both concen- 
 trated and neutral). The precipitate is insoluble in alcohol. 
 
 Hydrofluosilicic acid (or silico-flnoric acid), H 2 SiF 6 , 
 throws down a white precipitate of gelatinous appearance, consisting 
 of potassium silicofluoride, ICjSiFg, sparingly soluble in water. 
 
 Magnesium, Mg. 
 
 In the " natural classification " of the elements, magnesium is 
 associated with the metals of the alkaline earths (Be, Ba, Sr, Ca) 
 on the one hand, and with zinc and cadmium on the other. Its 
 position along with the alkalies in Group V. of the analytical 
 classification, is simply because it differs from the other members of 
 its own natural family in that the presence of ammonium chloride 
 prevents the precipitation of magnesium hydroxide by ammonia 
 in Group III., and also of magnesium carbonate by the group 
 reagent of Group IV. 
 
 DRY REACTION. When magnesium salts are strongly heated 
 in the outer blowpipe, a white infusible residue of the oxide is left. 
 If, after cooling, the residue be moistened with a drop or two of 
 cobalt nitrate solution and again strongly heated in the outer blow- 
 pipe flame, the mass acquires a pink colour. This reaction is 
 reliable only in the absence of other metallic oxides. 
 
 WET REACTIONS. Of the common salts of magnesium, the 
 sulphate, chromate, nitrate, and halogen salts are soluble in water. 
 One prominent characteristic of magnesium compounds is the 
 
 the rod with one end projecting at right angles to the upright. The little 
 glass tube into which the platinum wire is fused, is then slipped over the project- 
 ing iron wire. This arrangement admits of the wire being raised or lowered as 
 desired, while at the same time it readily remains in any position. 
 
24 Qualitative Analysis. 
 
 readiness with which they form " double " salts, many of which are 
 soluble in water. Use magnesium sulphate. 
 
 Alkaline hydroxides (NH 4 HO, KHO, NaHO, Ca(HO) 2 , or 
 Ba(HO) 2 ) precipitate from solutions of magnesium sulphate or 
 chloride, white magnesium hydroxide, Mg(HO) 2 . Almost insoluble 
 in water ; soluble in acids, soluble in ammonium chloride 
 
 Mg(HO) 2 + 4NH 4 C1 - 2NH 4 HO + MgCl 2 ,2NH 4 Cl (soluble 
 double salt) 
 
 .Owing to the solubility of magnesium hydroxide in ammonium 
 chloride, only -half the magnesium is precipitated from magnesium 
 chloride by means of ammonia, thus 
 
 2MgCl 2 + 2NH 4 HO = Mg(HO) 2 + MgCl 2 ,2NH 4 Cl 
 If ammonium chloride is previously present in sufficient quantity to 
 combine with the whole of the magnesium hydroxide, the alkaline 
 hydroxides give no precipitate. 
 
 Alkaline carbonates (K 2 CO 3 , Na 2 CO 3 , (NH 4 ) 2 CO 3 ) pro- 
 duce in solutions of magnesium salts, in the absence of ammonium 
 salts, precipitates of basic carbonates of magnesium, the composition 
 of which varies with conditions of temperature and concentration. 
 The precipitate with (NK 4 ) 2 CO 3 only separates out after a short 
 time. In the presence of ammonium chloride tJiese reagents give no 
 precipitate. 
 
 Hydrogen disodium phosphate, HNa 2 PO 4 , precipitates 
 hydrogen magnesium phosphate, HMgPO 4 , and tri-magnesium 
 phosphate, M*g 3 (PO 4 ) 2 . In the presence of ammonium chloride, 
 however, the double ammonium magnesium phosphate is thrown 
 down as a white crystalline precipitate, NH 4 MgPO 4 . It is 
 appreciably soluble in water, but insoluble in ammonia; hence 
 ammonia must be previously added. 
 
 In very dilute solutions the precipitation only takes place on 
 long standing. It is accelerated by stirring with a glass rod, the 
 deposition first appearing where the rod has rubbed the glass 
 vessel. The precipitate is soluble in acids, even acetic acid, but 
 reprecipitated by ammonia. 
 
 SYSTEMATIC PLAN OF ANALYSIS FOR THE METALS OF 
 
 GROUP V. 
 
 After having carefully gone through the various reactions for 
 the metals of Group V., the student should proceed to the examina- 
 tion of a few solutions containing mixtures of two or more salts of 
 these metals. 
 
Detection of the Metals of Group V. 25 
 
 The various special tests by which the individual members of 
 this group are recognised are not, for the most part, interfered with 
 by the presence of the rest of the group. Therefore a complete 
 separation of all the metals is not necessary. Thus, the test for 
 ammonium (evolution of ammonia by heating with sodium 
 hydroxide) can be made in the presence of Mg, K, and Na com- 
 pounds ; and obviously must be made in a separate portion of the 
 solution under examination from that in which sodium is to be 
 tested for, as it involves the addition of a sodium compound. 
 
 The test for magnesium, likewise (precipitation of NH 4 MgPO 4 ), 
 may be made in the presence of all the other members ; and 
 clearly must be made also in a separate portion of the solution, as 
 it involves the addition of both ammonium and sodium compounds. 
 Similarly, the flame tests for potassium and sodium are not inter- 
 fered with by the presence of ammonium or magnesium. The 
 flame reaction for potassium, however, unless examined by the aid 
 of the potassioscope, must always be corroborated by the forma- 
 tion of potassium chloroplatinate. But before this test can be 
 applied, ammonium salts must first be removed. 
 
 Solutions may be examined for the metals of Group V., 
 NH 4 , Na, K, Mg, by the following system : 
 
 DETECTION OF THE METALS OF GROUP V. 
 
 Operation i. To a portion of the solution add NaHO, and heat 
 in a test-tube. The evolution of ammonia (detected by its odour, 
 and its action on test-papers) proves the presence of NH 4 . 
 
 Operation 2. To a second portion add NH 4 C1, NH 4 HO, and 
 HNa 2 PO 4 . A white crystalline precipitate of NH 4 MgPO 4 proves 
 the presence of Mg. 
 
 Operation 3. Evaporate another (and larger) portion to dryness 
 in a porcelain dish. If ammonium salts are present * (already 
 ascertained in Operation i), they must be removed. For this pur- 
 pose, scrape the residue out of the dish and strongly heat it on the 
 lid of a platinum crucible (or a piece of platinum foil), until, on 
 momentarily withdrawing it from the flame, fumes are no longer 
 visible. 
 
 Dissolve the residue in a small quantity of water, and add one 
 drop of HC1. Dip a clean platinum loop into the solution, and heat 
 
 * In a complete analysis, ammonium salts are always present here, as they will 
 have been introduced in the process of separating the other groups. Under 
 these circumstances, therefore, the operation of removing ammonium com- 
 pounds is always necessary. The substance under analysis is tested for 
 ammonium before the ammoniacal reagents are introduced. 
 
26 Qualitative Analysis. 
 
 it in a Bunsen flame. An intense * yellow coloration proves the 
 presence of Na. A lilac colour indicates the presence of K. 
 
 In either case examine the flame through the potassioscope ; a 
 crimson flame indicates K. 
 
 Add to the solution of the residue a few drops of PtCl 4 , and stir 
 with a glass rod. A yellow precipitate of K 2 PtCl 6 confirms K. 
 
 APPENDIX TO CHAPTER III. 
 
 LITHIUM, RUBIDIUM, AND CAESIUM. 
 
 These three elements are usually placed in the category of rare 
 metals. It must be remembered, however, that there are degrees 01 
 rarity ; and while the compounds of rubidium and caesium are 
 certainly among the very rare substances with which the chemist 
 comes into contact, those of lithium, on the other hand, are very 
 widely distributed and are much more frequently met with.f 
 
 Lithium, Iii. 
 
 DRY REACTION. Lithium salts impart to the flame a brilliant 
 carmine-red colour. 
 
 WET REACTIONS. All the common salts are readily soluble in 
 water, except the carbonate, phosphate, and oxide, which are 
 soluble with difficulty. The chloride and nitrate are soluble in a 
 mixture of alcohol and ether (distinction from Na and K, the 
 chlorides and nitrates of ivhich are not soluble}. 
 
 Na 2 CO 3 and K 2 CO 3 precipitate Li 2 CO 3 from cold moderately 
 concentrated solutions (i part dissolves in 100 parts of water). 
 
 HNa 2 PO 4 gives a white precipitate, on boiling, of Li 3 PO 4 . The 
 precipitation is complete in the presence of NaHO. 
 
 PtCl 4 gives no precipitate (distinction from NH 4 , K, Rb, Cs). 
 
 HNa(C 4 H 4 O 6 ), hydrogen sodium tartrate, gives no precipitate. 
 
 Rubidium, Kb, and Caesium, Cs. 
 
 The compounds of these metals present the very closest 
 resemblance to those of potassium, and there are scarcely any 
 
 * More or less of a yellow flame is usually obtained, owing to the presence 
 of traces of sodium compounds as impurities in the reagents previously used in 
 separating the groups in the course of a complete analysis. 
 
 f Perhaps a rough idea of the relative rarity of the compounds of these 
 metals might be gained by a comparison of their cost. Lithium salts can be 
 obtained for about 12^. per pound, while rubidium and caesium salts cost 
 about 5-y. per drachm, i.e. at the rate of ^64 per pound. 
 
The Rare Metals of Group V. 27 
 
 chemical reactions by which they can be distinguished. The 
 separation of these metals is based on the different degrees of 
 solubility of their chloro-platinates. In this respect, as well as in 
 their other properties,* rubidium stands intermediate between 
 potassium and caesium. 
 
 When heated in a Bunsen flame, rubidium and caesium salts 
 impart to it a lilac colour, which to the unaided eye is absolutely 
 indistinguishable from that produced by potassium compounds ; 
 and when compounds of these metals, as well as those of lithium, 
 are mixed with comparatively minute quantities of sodium salts, the 
 colours they give to the flame are completely overpowered and 
 masked by the yellow of the sodium. By means of the spectroscope, 
 however, not only are the apparently identical colours given by 
 potassium, rubidium, and caesium proved to consist of light of 
 different quality or composition, but the presence of any or all of 
 them is easily and certainly detected even when admixed with 
 sodium salts. 
 
 The spectroscope is an instrument by means of which the light 
 emitted by strongly heated substances can be examined after it has 
 been made to pass through a glass prism. Its use depends upon 
 the fact that different coloured lights possess different degrees of 
 refrangibility ; that is to say, different coloured rays of light are 
 bent out of their straight course, by passage through a prism, at 
 different angles. Ordinary white light is composed of rays of all 
 degrees of refrangibility, hence, when such light passes through a 
 prism, the various coloured rays are separated, and spread out in the 
 order of their refrangibility, the least refrangible red at one 
 extreme, to the deep violet at the other. This familiar " rainbow " 
 coloured band of light is called the continuous spectrum. 
 
 In the spectroscope the light is passed through a narrow slit at 
 one end of a small telescope, and an image of the slit is received 
 upon a glass prism. This bends the light out of its straight course, 
 and spreads it out into the various colours of which it is composed. 
 If white light be admitted, then the continuous spectrum is seen, 
 which is an infinite number of images of the slit arranged side by 
 side ; if such a monochromatic light as that given by heating sodium 
 salts in the flame be used, then one image of the slit is seen in that 
 part of the spectrum which corresponds to the particular degree of 
 refrangibility of the light. In this case we say that the spectrum 
 of sodium consists of one line (that is, one image of the slit) in the 
 yellow, or one yellow iine.\ The light which passes out of the prism 
 is usually examined by a second telescope, which can be revolved 
 round the prism so that it can sweep the whole length of the 
 spectrum. 
 
 When the red colour imparted to the Bunsen flame by a 
 
 * Atomic weights, melting-points, etc. Also in the optical properties of 
 their crystallized salts (Tutton, J. C. S., May, 1896). 
 
 t When this sodium line is examined by a higher dispersive power, it is 
 found to consist of a group of lines. 
 
28 
 
 Qualitative A nalysis. 
 
 lithium salt is examined by the prism, it is seen to consist of light 
 of two degrees of refrangibility, therefore two images of the slit are 
 seen, one in the bright orange (a little on the red side of the sodium 
 line), and the other an extremely brilliant image in the red. The 
 spectrum of lithium, therefore, is one bright orange line and one 
 brilliant red line (see Fig. 9). 
 
 Similarly, when the lilac flame of potassium is examined in this 
 way, the light is found to be composed of three colours ; therefore 
 
 FIG. 9. 
 
 three lines are seen one extremely bright line far down in the red ; 
 a second, less brilliant, about halfway between the first and the 
 yellow sodium (or D) line ; while the third is far away in the deep 
 violet, and not easy to see.* 
 
 If a mixture of salts of sodium, potassium, and lithium be heated 
 in a flame, and the light examined by the spectroscope, f then all 
 
 * The sodium D line will be also visible, not only on account of the presence 
 of traces of sodium compounds in the, other salts, but owing also to the 
 presence of particles of salt which are always floating about in the air, and 
 which tinge the Bunsen flame. 
 
 t Where the equipment of the laboratory will allow, every student should 
 have the opportunity of examining the spectra of the metals of the alkalies and 
 alkaline earths. 
 
The Rare Metals of Group V. 29 
 
 six lines are seen ; sodium, one yellow ; potassium, two red and one 
 violet ; lithium, one red and one orange. 
 
 In such a simple mixture as this there is no difficulty in 
 identifying the various lines of the different elements, but in more 
 complex spectra it is only possible to do so by recording the exact 
 position they each occupy on some fixed scale. 
 
 When the spectra of rubidium and caesium are in this way com- 
 pared with that of potassium, a striking difference is at once 
 observed. In the first place, they are obviously much more complex. 
 Rubidium shows two strongly marked red lines (hence the name 
 of the element) very close together, and near to the red line of 
 potassium. Then a number of lines in the orange-red and in the 
 green, and lastly two characteristic and brilliant violet lines. 
 
 Caesium gives red lines (not so far down as those of rubidium) 
 orange and green lines, and two most characteristic brilliant 
 lines in the bright blue part of the spectrum (ccesium signifies 
 sky -blue). 
 
 In the chart (Fig. 9), the most prominent lines* in the spectra 
 of the five metals of the alkalies are compared together. The 
 letters at the top refer to lines in the solar spectrum mapped by 
 P>aunhofer, which correspond to those given by known metals. 
 Thus, the D line in the solar spectrum corresponds to the yellow 
 line of sodium. 
 
 * Speaking generally, it may be said that at higher temperatures than can be 
 obtained in the Bunsen flame, lines will appear in the spectrum of a particular 
 metal which are not seen at the lower temperature. 
 
CHAPTER IV. 
 REACTIONS OF THE METALS OF GROUP IV. 
 
 THIS group consists of the three metals, barium, strontium, 
 and calcium, known as the metals of the alkaline earths. 
 . The group-reagent, ammonium carbonate, precipitates the 
 carbonates of the metals in an ammoniacal solution, even in the 
 presence of ammonium chloride. Hence the group-reagent separates 
 these three metals from magnesium. 
 
 Barium, Ba. 
 
 DRY REACTION. Barium compounds, heated on platinum wire 
 in the Bunsen flame, impart a pale apple-green colour to the flame, 
 which becomes more distinct if the substance on the wire is 
 moistened with strong hydrochloric acid. The test is not very 
 reliable. 
 
 Barium sulphate, BaS0 4 (also SrSO 4 and CaSOJ, when heated 
 on charcoal or with carbon, is reduced to the sulphide. 
 
 WET REACTIONS. Of the common salts of barium, the chloride, 
 bromide, iodide, nitrate, chlorate, acetate, and sulphide are soluble 
 in water. 
 
 Ammonium carbonate, (NH 4 ) 2 CO 3 , group-reagent (also 
 Na 2 CO 3 and K 2 CO 3 ) precipitates barium carbonate, BaCO 3 , as a 
 white amorphous powder. Insoluble in water ; readily dissolved, 
 with evolution of carbon dioxide, by dilute acids ; slighjly soluble in 
 NH 4 C1. 
 
 It is also dissolved by a solution of carbon dioxide in water, 
 forming hydrogen barium carbonate, H 2 Ba(CO 3 ) 2 .* 
 
 * On this account complete precipitation of barium carbonate cannot be 
 accomplished by hydrogen ammonium carbonate, HNH 4 CO 3 , for carbon 
 dioxide is evolved, which dissolves a portion of the precipitate 
 
 2 HNH 4 CO 3 + BaCl 2 = BaCO 3 '+ 2 NH 4 C1 + H 2 O f- CO 2 
 
 And since the solution of ammonium carbonate used as the group-reagent 
 consists chiefly of HNH 4 CO 3 (because the normal salt undergoes decomposition 
 into the hydrogen salt and free ammonia when in aqueous solution), it is always 
 necessary to first add NH 4 HO before applying the group-reagent. 
 
Group IV. 31 
 
 H 2 SO 4 , or any soluble sulphate, produces a white granular 
 precipitate of BaSO 4 , practically insoluble in water, insoluble also 
 in acids and alkalies. (Boiling concentrated H 2 SO 4 slowly dis- 
 solves it, forming hydrogen barium sulphate, H 2 Ba(SO 4 ) 2 .) In- 
 soluble in solutions of (NH 4 ) 2 SO 4 . BaSO 4 , being practically insoluble 
 in water, is precipitated by a saturated solution of SrSO 4 , although 
 such a solution contains only i part of salt in 7000 parts of water. 
 
 Potassium chromate, K 2 CrO 4 , produces a primrose-yellow 
 precipitate of barium chromate, BaCrO 4 , practically insoluble in 
 water (distinction from SrCrO 4 , which is MODERATELY soluble, ana 
 CaCrO 4 , which is VERY soluble in water]. It is insoluble in acetic 
 acid (distinction from SrCrO 4 ) ; soluble in HNO 3 and in HC1. 
 By boiling with K 2 CO 3 , barium chromate is converted into the 
 carbonate 
 
 BaCrO 4 + K,CO 3 = BaCO 3 4- K 2 CrO 4 
 
 Hydrofluo silicic acid, H 2 SiF 6 , gives a white crystalline pre- 
 cipitate of barium silicofluoride, BaSiF 6 , slightly soluble in water, 
 but insoluble on the addition of alcohol (distinction from SrSiF 6 
 and CaSiF , which are readily soluble in water). 
 
 Strontium, Sr. 
 
 DRY REACTION. When heated in the Bunsen flame, volatile 
 strontium salts, such as SrCl 2 , Sr(NO 3 ) 2 , impart a rich crimson 
 colour to the flame ; other salts require to be moistened upon the 
 wire with strong HCI. Strontium sulphate must be converted into 
 the sulphide, which is then dissolved in HCI and the chloride 
 tested in the flame.* 
 
 WET REACTIONS. The same common salts of strontium as of 
 barium are soluble in water. The chromate and sulphate are 
 somewhat soluble. 
 
 (NH 4 ) 2 CO 3 (Na 2 CO 3 and K 2 CO 3 ) precipitates SrCO 3 , exactly 
 similar to the- barium compound in its reactions. 
 
 H 2 SO 4 or soluble sulphates precipitate SrSO 4 . The precipitate 
 is slightly soluble in water (i : 7000), confterably soluble in HCI 
 or HNO 3 , but almost insoluble in a solution of (NH 4 ) 2 SO 4 (dis- 
 tinction from CaSO 4 ). SrSO 4 is precipitated by a solution of 
 
 * Small quantities of SrSO 4 upon a filter may be converted into SrS by first 
 drying the filter, then folding it into a small roll, and twisting a platinum wire 
 round it so as to hold it in a flame. The paper is burnt to an ash, and the 
 latter is then heated upon the wire in an ordinary smoky gas-flame, when the 
 reduction of the sulphate takes place. The ash is then touched with a glass 
 rod dipped in strong HCI, and on bringing it into a Bunsen flame the crimson 
 flame is seen. 
 
\ 
 
 32 Qualitative Analysis. 
 
 CaSO 4 * (a saturated solution of which only contains i part in 
 430 parts of water) ; the precipitation does not take place at once 
 in cold solutions, but appears quickly on heating. 
 
 K ,CrO 4 and H 2 8iF 6 (see barium reactions with these reagents) 
 
 Calcium, Ca. 
 
 DRY REACTIONS. Calcium compounds, when heated in a Bunsen 
 flame, impart to it a reddish colour, especially if previously moistened 
 with hydrochloric acid. The presence of strontium masks the red 
 colour given by calcium compounds. When calcium carbonate is 
 strongly heated it loses carbon dioxide, and is converted into 
 calcium oxide 
 
 CaCO 3 = CO 2 + CaO 
 
 This process is earned on in the lime-kilns during the operation of 
 " burning lime." Limestone (calcium carbonate] is thus converted 
 into lime (calcium oxide}. SrCO 3 and BaCO 3 undergo similar 
 decomposition ; SrCO 3 less readily than CaCO 3 , while BaCO 3 
 requires prolonged heating to a white heat to effect the change. 
 
 WET REACTIONS. The same common salts of calcium are 
 soluble in water as of strontium and barium ; and, generally 
 speaking, calcium salts are more readily soluble. Thus, the chloride 
 and nitrate are extremely deliquescent ; the sulphate dissolves to 
 the extent of i : 430 (compare Sr and Ba). The solubility of the 
 oxalates, however, is in the reverse order, calcium oxalate being 
 the most insoluble, and barium oxalate the most soluble. A 
 peculiarity of calcium salts is that many of them appear to be less 
 soluble in hot than in cold water. 
 
 (NH 4 ) 2 CO 3 (also Na 2 CO 3 and K 2 CO 3 ) precipitates CaCO 3 , 
 similar to the barium and strontium compounds in its reactions. 
 
 H SO, or soluble sulphates,t added to a strong solution of a 
 calcium salt,J give an immediate precipitate of calcium sulphate. 
 
 * SrSO 4 is itself soluble in water to the extent of i : 7000, but obviously the 
 addition of a solution of CaSO 4 or H 2 SO 4 to this solution does not precipitate 
 the SrSO 4 . Neither is the strontium precipitated from this solution by any 
 reagent which forms with it a compound which is more soluble in water than 
 is the sulphate, e.g. oxalic acid or oxalates. 
 
 f The soluble sulphate must be more soluble than calcium sulphate ; SrSO 4 
 therefore, being less soluble, would not give a precipitate with a solution of a 
 calcium salt, although, being more soluble than BaSO 4 , it throws down this 
 compound from solutions of barium salts. For the same reason, a solution of 
 CaSO 4 will throw down SrSO 4 and BaSO 4 from solutions of strontium and 
 barium salts. 
 
 % Only in the case of calcium salts which are more soluble than calcium 
 sulphate. 
 
Group IV. 33 
 
 From more dilute solutions the precipitate only separates after 
 some time, or, if still more dilute, not at all. The precipitate is 
 insoluble in alcohol, therefore the addition of this liquid in 
 considerable bulk favours the precipitation. 
 
 Calcium sulphate is readily soluble in a concentrated solution 
 of ammonium sulphate, especially when hot (distinction from 
 SrSO 4 and BaSO 4 ). Boiling with potassium carbonate easily 
 converts it into calcium carbonate. 
 
 Oxalic acid, H 2 C 2 O 4 , or soluble oxalates, gives a white 
 crystalline precipitate of calcium oxalate. The precipitate is 
 soluble in mineral acids ; hence, when oxalic acid is used as the 
 precipitant, precipitation is not complete, owing to the liberation 
 of the mineral acid of the calcium salts, thus 
 
 CaCl, + H 2 C 2 4 = CaC 2 4 + 2HC1 
 
 Calcium oxalate is insoluble in NH 4 HO, therefore, if the solution 
 be first rendered alkaline, the whole of the oxalate is thrown down. 
 In practice it is usual to employ ammonium oxalate, (NH 4 ) 2 C 2 O 4 . 
 In cold dilute solutions precipitation is not immediate. The pre- 
 cipitate is insoluble in water and in acetic acid. Strontium and 
 barium oxalates are soluble, to a small extent, both in water and 
 in acetic acid. 
 
 SPECTRA 'OF BARIUM, STRONTIUM, AND CALCIUM. 
 
 The spectra given by these metals, when their salts are heated 
 in the Bunsen flame, are extremely characteristic. For calcium, 
 the chloride may be used, with occasional moistening with hydro- 
 chloric acid. For barium and strontium, the nitrates, or, better 
 still, the chlorates give the best result, although in the absence of 
 these salts the chlorides can be employed. 
 
 From the charts of their spectra given in Fig. 10, it will be seen 
 that they each give lines in the red portion of the spectrum. Owing 
 to this fact, if salts of these metals are heated in a Bunsen flame, 
 and the flame examined through an indigo prism, the flame will 
 appear red\\\ each case, because indigo does not intercept or absorb 
 the red portion of the spectrum. The potassioscope, on the other 
 hand, has the property of cutting off the red rays emitted by these 
 metals. It absorbs all the red part of the spectrum down very 
 nearly to the prominent red line (K.o) of potassium (compare the 
 chart on p. 28), beyond the lithium and caesium red lines. Hence, 
 if flames which are coloured by the alkaline earths are examined 
 through this instrument, no red \\'\\\ be visible. 
 
 | 
 
34 Qualitative Analysis. 
 
 The importance of this lies in the following fact : When Groups 
 IV. and V. are separated by precipitating the carbonates of barium, 
 strontium, and calcium, minute traces of these metals pass through 
 into Group V., as their carbonates are not absolutely insoluble. 
 Hence, when the final residue is obtained (in which the tests for 
 
 i 
 
 FIG. io. 
 
 potassium are made), it will contain these traces of these alkaline 
 earths. If this residue be tested in the flame, and the flame 
 examined by the potassioscope, the presence of these members of 
 Group IV. will not vitiate the reaction ; but with the indigo prism 
 the red light they emit would be mistaken for, and returned as that 
 given by potassium.* 
 
 SEPARATION OF GROUPS IV. AND V. 
 
 The solution is first rendered alkaline by the addition of NH 4 HO. 
 NH 4 C1 is then added (to prevent the precipitation of magnesium 
 carbonate), after which ammonium carbonate is added until the 
 carbonates of the metals of Group IV. are completely thrown down. 
 The mixture may be gently warmed. [It must not be boiled, or 
 the precipitated carbonates will react with the NH 4 C1, forming 
 soluble chlorides, while NH 3 and GQ 2 will escape with the steam ; 
 thus, BaCOs + 2NH 4 C1 = Bad, + CO 2 + 2NH 3 + H 2 O.] The 
 mixture is filtered. The filtrate is examined for the metals of 
 Group V. by the method given on p. 25, while the precipitate is 
 treated in the following way : 
 
 * The only element which could be mistaken for potassium, when using the 
 potassioscope, is the exceedingly rare metal rubidium, whose spectrum contains 
 red lines still lower down than those of potassium. 
 
Group IV. 35 
 
 SEPARATION OF THE METALS OF GROUP IV. 
 
 The precipitate, consisting of PaCO 3 , SrCO 3 , and CaCO 3 , is washed, 
 and then dissolved in a small quantity of warm dilute acetic acid, 
 H(C 2 H 3 O 2 ). To the solution thus obtained potassium chromate, 
 K 2 CrO 4 , is added ; the mixture gently warmed and filtered.* 
 
 The precipitate con- ' The filtrate f contains strontium and calcium 
 sists of yellow barium acetates. Add a strong solution of 
 chromate, BaCrO 4 . (NH 4 ) 2 SO 4 , and boil for a short time to 
 
 ensure the solution of CaSO 4 . Filter. 
 
 Boil the precipitate with 
 K 2 CO 3 , which converts 
 
 BaCrO 4 into BaCO 3 . -^he precipitate con- The solution contains 
 Filter and wash the j t f s SQ calcium sulphate. 
 
 P K e Tr a n fr n- 'T Add ammonium oxa- 
 
 S*riLfaJlS: "'** 'he precipitate! late (NH ) 3 C 3 O 4 (ov 
 and confirm bv SrSO, thoroughly and con- , NH 4 )HO and oxa- 
 or H.,SiF 6 hrm . b y the flame i he acid, H 2 C 2 O 4 ). 
 
 reaction. A white precipitate 
 
 of CaC 2 O 4 confirms 
 calcium. 
 
 Other methods of separating the metals of Group IV., based on 
 the insolubility of certain of their common salts in alcohol, may- 
 be employed. Thus, the nitrates of barium and of strontium are 
 insoluble in alcohol, while that of calcium is soluble in this liquid 
 or mixtures of alcohol and ether). 
 
 The precipitated carbonates are dissolved in dilute HNO 3 , and 
 the solution evaporated to dryness. The residue, consisting of the 
 mixed nitrates, is then treated with a mixture of alcohol and ether, 
 in about equal volumes, which dissolves the Ca(NO 3 ).,, leaving 
 Ba(N0 3 ) 2 and Sr(NO 3 ).. 
 
 Again, barium chloride is insoluble in the same mixture, whereas 
 the chlorides of strontium and calcium are both soluble. Therefore, 
 if the mixed carbonates are dissolved in dilute HC1, and the solu- 
 tion evaporated to dryness, the SrCl 2 and CaCl 2 may be dissolved 
 out, leaving a residue of BaCl 2 . 
 
 * Instead of adding the chromate at once to the whole of the solution, a 
 small separate portion may be used, and the presence or absence of the Ba 
 thus ascertained. If present, then the whole of the solution is treated as above, 
 and the precipitation of barium chromate may then be taken as confirmatory. 
 If absent, then a second small portion of the so'lution is taken, and tested for Sr 
 by adding CaSO 4 . If Sr is present, then the separation of Sr and Ca is carried 
 out as shown above; but if CaSO 4 gives no precipitate even on warming, 
 thus proving the absence of Sr, the remainder of the solution is at once tested 
 for Ca by the oxalate reaction. 
 
 f If barium has been separated as chromate, the filtrate will be coloured 
 yellow by the excess of K 2 CrO 4 used, and a white precipitate may appear 
 yellow. The process is more thorough if the carbonates of Sr and Ca are 
 again thrown down by adding ammonium carbonate. The precipitate is then 
 washed free from chromate, dissolved in a few drops of HO, and the sulphates 
 separated as directed above. 
 
CHAPTER V. 
 THE METALS OF GROUP III. 
 
 THIS group comprises the metals aluminium, chromium, iron, 
 manganese, zinc, nickel, and cobalt, besides a number of 
 the rare metals (see Classification, p. 16). 
 
 Although all these metals are precipitated by the group-reagent, 
 namely, ammonium sulphide, in the presence of ammonia, the 
 compounds that are thrown down are not all of similar composition, 
 as was the case in Group IV. 
 
 The addition- of ammonia to a solution containing the metals 
 of this group, results in the precipitation of the hydroxides (or 
 hydrated oxides) of Al, Cr, and Fe, corresponding to the oxides 
 R 2 O 3 .* Thus, taking aluminium in potash alum as an example 
 
 K 2 S0 4 ,A1 2 (S0 4 ) 3 + 6NH 4 HO = K 2 SO 4 + 3(NH 4 ) 2 SO 4 + A1 2 (HO), 
 
 In the presence of ammonium chloride, the hydroxides of the 
 metals Mn,t Zn, Ni, and Co (which have the general formula 
 R(HO) 2 ) are not precipitated by ammonia, owing to the formation 
 of double chlorides. 
 
 If ammonium sulphide be added to the mixture after the 
 ammonia, the already precipitated hydroxides of aluminium and 
 chromium are unchanged, but the ferric hydroxide is converted 
 into ferrous sulphide, with the elimination of sulphur ; thus 
 
 Fe 2 (HO) 6 + 3H 2 S = 2FeS + 6H 2 O + S 
 and the remaining metals are also thrown down as sulphides. 
 
 * The separation of the metals of Group III. is complicated by the fact that 
 if Mn, Zn, Ni, and Co are present in the form of phosphates, the addition ol 
 ammonia causes the partial precipitation of these phosphates. And, further, 
 if the members of Group IV. and Mg are present as phosphates, they also are 
 thrown down by ammonia, and therefore appear in Group III. The special 
 treatment of phosphates will be discussed later. 
 
 f In the case of manganese, the separation is not complete under all 
 conditions (see reactions of manganese). 
 
Group III. Division A. 37 
 
 The result, therefore, of adding NH 4 C1, NH 4 HO, and (NH 4 ) 2 S is 
 as follows : 
 
 Precipitated by] 
 
 he) 
 
 ' r e^riuje convertea into f eb 
 
 Precipitated on the 
 addition of (NH 4 ) a S 
 
 Fe 2 (HO) 6 converted into FeS 
 
 MnS 
 
 ZnS 
 NiS 
 
 CoS 
 
 fhe metals of Group III., therefore, may be subdivided into 
 two families, based upon their behaviour towards ammonia : Divi- 
 sion A, consisting of the three metals, aluminium, chromium, and 
 iron ; and Division B, of manganese, zinc, nickel, and cobalt. 
 
 REACTIONS OF THE METALS OF GROUP III. DIVISION A. 
 Aluminium, Al. 
 
 DRY REACTION. When aluminium compounds are strongly 
 heated on charcoal in the outer flame, aluminium oxide is formed, 
 and if this be moistened with a solution of cobalt nitrate, and again 
 strongly heated, either upon the charcoal or upon a loop of plati- 
 num wire, the mass assumes a rich blue colour, due to the formation 
 of cobalt aluminate. 
 
 This test is, however, greatly masked if other metallic oxides 
 which are coloured are present at the same time. It may be 
 employed as a confirmatory test when aluminium is separated from 
 iron and chromium in the course of analysis. 
 
 WET REACTIONS. Of the common salts of aluminium, the 
 chloride, A1 2 C1 6 , and sulphate, A1 2 (SO 4 ) 3 , are soluble in water. 
 The important salts, however, are the double sulphates of 
 aluminium with ammonium or potassium, known as ammo- 
 nium alum, (NH 4 ) 2 SO 4 ,A1 2 (SO 4 ) 3 ,24H 2 O, and potassium alum, 
 K 2 SO 4 ,A1 2 (SO 4 ) 3 ,24H 2 O, respectively.! A solution of either of 
 these alums may be used for the following reactions. 
 
 * In reality the compounds precipitated are the hydrated sulphides, ex- 
 pressed by the general formula R(HS)(HO), or RS,H 2 O. To avoid unnecessary 
 complication in reactions, the molecule of H 2 O maybe left out of consideration. 
 
 f The alums constitute a large class of double sulphates, having the general 
 formula M 2 SO4,Ro(SO4)3,24H 2 O, where M is a monovalent element or group, 
 such as potassium or ammonium, and R is either aluminium, iron, chromium, 
 or manganese. The commonest of all the salts is potassium aluminium alum, 
 K 2 SO4,A1 2 (SO 4 ) 3 ,24H 2 O ; this is the salt, therefore, that is distinguished by 
 the single word ALUM. From their formulae it may be seen at a glance that 
 
38 Qualitative Analysis. 
 
 NH ( HO throws down a white translucent precipitate of the 
 hydrated oxide, or hydroxide,* A1 2 O 3 ,3H 2 O, orA! 2 (HO) 6 . Soluble 
 in a large excess of the reagent, but on gently boiling, the hydroxide 
 is entirely precipitated. [Prolonged boiling, however, causes partial 
 dissociation of the ammonium salt in solution, ammonia escapes, 
 and free acid is liberated, which then begins to dissolve the preci- 
 pitate.] In the presence of ammonium chloride, the precipitation 
 of A1 2 (HO) 6 by ammonia is complete. The precipitate is readily 
 soluble in mineral acids, and in acetic acid. 
 
 KHO or NaHO produces the same precipitate, which readily 
 combines with an excess of the reagent, forming a soluble aluminate 
 of potassium or sodium (Al 2 O 3 ,3Na 2 O or Na 6 Al 2 O 6 ). These 
 aluminates are decomposed by acids, even by such feeble acids 
 as carbonic acid or hydrosulphuric acid (sulphuretted hydrogen), 
 with reprecipitation of the aluminium hydroxide, as this compound 
 is not soluble in sodium carbonate or sodium sulphide. Thus, 
 with carbonic acid 
 
 Al 2 O 3 ,3Na 2 O -f 3CO 2 + 3H 2 O = 3Na 2 CO s + A1 2 O 3? 3H 2 O [or A1 2 (H O) 6 ] 
 
 In the case of stronger acids, such as HC1, the same action 
 takes place, but any excess of the acid beyond that required to 
 combine with the sodium of the aluminate at once re-dissolves the 
 A1 2 (HO) 6 . When, therefore, this acid is used, a slight excess is 
 added, and the aluminium hydroxide is re-precipitated by means 
 of ammonia. 
 
 Sodium and potassium aluminates are also decomposed by am- 
 monium chloride, with the precipitation of aluminium hydroxide ; 
 the precipitation is complete on boiling. The compound thrown 
 down under these circumstances consists mainly of the di-hydrated 
 oxide, A1 2 O S ,2H 2 O ; thus 
 
 Al,O 3 ,3Na 2 O + 6NH 4 C1 = 6NaCl + 6NH 3 + H 2 O + A1 2 O 3 .2H 2 O 
 
 these salts are composed of a molecule of each of the two sulphates, together 
 with twenty-four molecules of water. With compounds of this description 
 there is, unfortunately, a tendency in certain quarters to add the formulae of 
 the two salts together, and then to divide the numerals by their greatest com- 
 mon measure ; thus, KsSO^Al^SO^s^HoO = KaAlofSO^^HaO, which, 
 divided by 2 - KA^SO^.iaHjO. Presumably this plan is adopted with 
 a view to simplification, but as it obscures the origin and the nature of the 
 compounds, and as there is not the smallest evidence that such formulae are 
 more exact representations of the molecular constitution of the compounds, 
 their use is greatly to be deprecated. 
 
 * Three hydrated oxides of aluminium are known, obtainable by precipi- 
 tation under different circumstances, namely, AljOs.sHsO, AliOj.aHjO, 
 and A1 2 O 3 ,H 2 O; these may also be formulated A1 2 (HO), A1 2 O(HO) 4 , and 
 A1 2 O 2 (HO) 2 respectively. 
 
Group III. Division A, 39 
 
 BaCO suspended in water, precipitates A1 2 (HO) 6 , carbon di- 
 oxide being evolved. The precipitation is complete even in the 
 cold.* If alum or aluminium sulphate is used, the precipitate is 
 mixed with insoluble barium sulphate 
 
 A1 2 (S0 4 ) 3 + 3BaC0 3 + 3H 2 O = A1 2 (HO) 6 + 3 BaSO 4 + 3 CO 2 t 
 
 K,CO ; and Na.CO, precipitate an uncertain mixture of the 
 hydroxide and basic carbonates. 
 
 (NH 4 ) 2 S precipitates aluminium hydroxide, with evolution of 
 sulphuretted hydrogen (compare Fe). 
 
 A1 2 (S0 4 ) 3 + 3(NH 4 ) 2 S + 6H 2 = A1 2 (HO) 6 4-3(NH 4 ) 2 SO 4 + 3 H 2 S t 
 
 [Aluminium forms no sulphide in the wet way. A1 2 S 3 (obtained 
 by the union of Al and S) is decomposed instantly by water, forming 
 the trioxide, and evolving H 2 S.] 
 
 Chromium, Cr. 
 
 DRV REACTIONS. Chromium compounds impart to a borax 
 bead a grass-green colour, when heated either in the outer or inner 
 blowpipe flame. 
 
 When fused in a platinum capsule with five or six times their 
 weight of a mixture consisting of i part of KNO 3 and 2 parts of 
 dry Na 2 CO 3 or K 2 CO 3 (or i part of KC1O 3 with 6 parts of Na 2 CO 3 ), 
 chromium compounds are converted into alkaline chromates, which 
 appear as a yellow mass, soluble in water to a yellow solution. 
 In the case of chromic oxide, for instance, Cr 2 O 3 , the reaction is the 
 following : 
 
 Cr a O s + 2K. 2 CO 3 + KC1O 3 = 2K 2 CrO 4 4- KC1 -f 2CO 2 
 The chief natural source of chromium is the mineral chrome iron 
 ore, Cr 2 O 3 ,FeO. When this is fused with either of the above 
 mixtures, the same reaction takes place as regards the chromium, 
 while the iron is changed to Fe 2 O 3 ; thus 
 
 6Cr 2 O 3 ,FeO + i2K 2 CO 3 + 7KC1O 3 = i2K 2 CrO 4 + 3Fe 2 O 3 + 
 
 ;KC1 + I2CO 2 
 
 * In the presence of certain organic acids, as oxalic, tartaric, or citric acids, 
 aluminium hydroxide is only more or less imperfectly precipitated by the 
 above-mentioned reagents, owing to the formation of soluble double salts 
 of the organic acid with aluminium and the alkali metal ; such, for example, 
 us the double tartrate of aluminium and sodium, Na 2 (C 4 H 4 O 6 ),Al 2 (C 4 H 4 O 6 )3. 
 This applies also in the case of the corresponding chromium and iron com- 
 pounds. 
 
 f In these equations simple aluminium sulphate is given instead of alum, 
 in order not to unnecessarily load the equation with materials taking no part 
 in the reaction. 
 
40 Qualitative Analysis. 
 
 [Sodium peroxide, Na a O 2 , may be substituted as the oxidizing 
 material, in which case the fusion should be carried out in a silver 
 capsule.] 
 
 WET REACTIONS. The two best-known classes of chromium 
 salts are derived from the two oxides, namely 
 
 Chromium sesquioxide (or chromic oxide), Cr 2 O 3 
 Chromium trioxide (chromic anhydride), CrO 3 
 
 Chromic oxide, Cr 2 O 3 , is basic, uniting with acids to form the 
 chromic salts, such as chromic hydroxide, Cr 2 (HO) 6 or Cr 2 O 3 ,3H 2 O ; 
 chromic chloride, CrCl 3 ; chromic sulphate, Cr 2 (SO 4 ) 3 ; double potas- 
 sium and chromium sulphate (chrome alum}, K 2 SO 4 ,Cr 2 (SO 4 ) 3 , 
 24H 2 O. Of these salts, the hydroxide alone is insoluble in water. 
 
 Chromium trioxide, CrO 3 , is the anhydride of the hypo- 
 thetical chromic acid,* H 2 CrO 4 , which gives rise to salts known as 
 chromates, analogous in constitution to the sulphates. 
 
 Chromates of the metals of Groups IV. and V. are all soluble in 
 water, except BaCrO 4 . The other chromates are insoluble. 
 
 a. Chromic Salts. These salts are mostly of a purplish 
 or violet-grey colour when solid, giving either a purple or green 
 solution when dissolved, the colour depending upon the conditions 
 of solution. Thus chrome alum dissolved in cold water gives a 
 purple solution, which on boiling turns green, t and on long standing 
 again becomes purple. 
 
 NH,HO produces a bluish or greenish-grey precipitate of 
 chromic hydroxide, Cr 2 (HO) 6 J, partially dissolved by excess of 
 ammonia in the cold, giving a lilac-coloured liquid, but completely 
 precipitated on gently boiling. Cr 2 (HO) 6 is readily soluble in acids. 
 
 KHO and NaHO precipitate Cr 2 (HO) G , readily soluble in 
 excess, giving a deep green solution.ll Reprecipitated by neutraliza- 
 tion with HC1, and by boiling with NH 4 C1, as in the case of Al. 
 
 * At first it may confuse students to find that chrotmV anhydride, and 
 chromzV acid with its salts, should not be in the class of chromic compounds, 
 it must be remembered that the classification is not based upon the nomen- 
 clature of the substances. "Chromic" compounds are those containing 
 chromium as the "base," or the positive radical ; while in those compounds 
 derived from CrO 3 the element is in the " acidic " or negative group. They 
 may, therefore, be conveniently distinguished as " chromic acid " compounds. 
 
 t The green colour is said to be due to the formation of a basic salt, by the 
 action of water upon the normal compound. 
 
 % The composition of the precipitate depends on the conditions under 
 which it is formed. There are several hydrated chromic oxides ; compare 
 also Al. 
 
 For the influence exerted by the presence of organic acids, see footnote, 
 
 P- 39- 
 
 || The soluble compounds produced are similar to those given by 
 aluminium ; the potassium salt has the composition Cr 2 O 3 , K 2 O or K 2 Cr 2 O^. 
 
Group III. Division A. 41 
 
 BaCO precipitates a mixture of the hydroxide and basic 
 carbonate. Complete precipitation only after some hours.- 
 
 K 2 00 3 and Na.CO, give a similar precipitate, the composition 
 of which varies with the conditions of precipitation. 
 
 (NH 4 )>S precipitates Cr 2 (HO) 6 . Precipitation complete. [Cr, 
 like Al, is incapable of forming a sulphide in the wet way.] 
 
 Oxidation of Chromic Compounds. By means of suitable 
 oxidising agents, chromic compounds are readily converted into 
 compounds of chromic acid, the mechanism of the change in all 
 cases being the oxidation of the sesquioxide into the trioxide \ 
 thus 
 
 Cr 2 O 3 4- 3O = 2CrQ 3 
 
 One method, namely, by fusion with oxidising agents, has been 
 explained under Dry reactions. The Cr 2 O 3 in that instance is 
 oxidised into the potassium salt of chromic acid. The oxidation 
 may be accomplished in the wet way by the following reactions : 
 
 (1) Boiling chromic hydroxide with potassium hydroxide and 
 manganese dioxide 
 
 Cr 2 (HO) 6 + 4KHO + 6MnO 2 = 2K 2 CrO 4 + 3Mn 2 O 3 + sHoO 
 
 (2) By substituting lead peroxide for manganese dioxide 
 Cr,(HO) G + 4KHO + 3PbO, = 21^0^ + 3PbO 4- 5H 2 O 
 
 In this case secondary reactions take place, for PbO is soluble 
 in KHO, and the solution so formed then reacts upon the K 2 CrO 4 , 
 producing PbCrO 4 , which also dissolves in KHO. 
 
 (3) By the action of hypochlorites (or hypobromites) in the 
 presence of caustic alkalies, either employed as such, or formed in 
 the solution by the use of chlorine or bromine in the presence of 
 the caustic alkali 
 
 Cr 2 (HO) 6 + 4KHO + 3KC1O = 3KC1 4- 2K 2 CrO 4 + sH 2 O 
 
 (4) By the action of sodium peroxide. If a small quantity of 
 Na.,O, be added to chromium hydroxide suspended in water, and 
 the mixture gently warmed, the chromium compound is immediately 
 converted into the yellow sodium chromate ; thus 
 
 Cr 2 (HO) + 3Na 2 O 2 = 2Na 2 CrO 4 + 2NaHO + 2H 2 O 
 
 0. Chromic Acid and Chromates. The acid, H 2 CrO 4 , has 
 never been isolated. The anhydride, CrO 3 , is readily obtained by 
 
 They are known v&chromites. Chrome iron ore is ferrous chromite, Cr 2 O 3 ,FeO. 
 Although analogous to the aluminrtfer, they must not be called chroma/es, as 
 this name is reserved for the salts of chromic acid. 
 
42 Qualitative Analysis. 
 
 adding strong H 2 SO 4 to a cold strong solution of potassium di- 
 chromate', when the oxide is deposited in the form of red silky needles. 
 It forms two classes of salts, viz. the normal chromates, of which 
 K 2 CrO 4 is a type ; and the dichromates,* of which K 2 Cr 2 O 7 is a 
 familiar example. The chromates are mostly yellow or red in 
 colour, and those which are soluble in water (see p. 40) impart a 
 yellow or orange colour to the liquid. The most important of the 
 insoluble chromates made use of in analysis, and which are all 
 precipitated by the addition of potassium chromate to solutions of 
 the metallic salts, are the following : 
 
 Barium chromate, BaCrO 4 (see Ba reactions, p. 31). 
 
 Lead chromate, PbCrO 4 (see Pb reactions, p. 81). PbCrO 4 
 melts without decomposition, and solidifies on cooling to a brown 
 crystalline mass. At higher temperatures it gives off oxygen 
 
 2PbCrO 4 = Cr 2 O 3 + 2PbO + 30 
 
 PbCrO 4 (known as chrome yelloiv], when digested with NaHO, or 
 with K 2 CrO 4 , is converted into a red basic lead chromate (known 
 as chrome red] 
 
 2PbCrO 4 + 2NaHO = Na2CrO 4 4- H 2 O 4- Pb 2 CrO 5 (or PbCrO 4 ,PbO) 
 2PbCrO 4 + K 2 CrO 4 - PbCrO 4 ,PbO + K 2 CrO 4 ,CrO 3 
 
 Silver chromate, Ag 2 CrO 4 . A dark chocolate-red precipitate, 
 soluble in ammonia and nitric acid. 
 
 Mercurous chromate (basic), Hg 2 CrO 4 ,Hg 2 O. A brick- 
 red precipitate, which, when dried, and heated in a tube, gives a 
 mercury sublimate, evolves oxygen, and leaves a residue of Cr 2 O 3 . 
 
 Oxidation of Chromic Acid. Although CrO 3 is such a 
 highly oxygenated compound, it appears to be capable of still 
 further oxidation by hydrogen peroxide, giving rise to a compound 
 which is believed by some to be perchromic acid, HCrO 4 , or 
 2CrO 3 ,H 2 O 2 , and by others to be a compound of CrO 3 and H 2 O 2 
 in undetermined proportions. The interest of the compound lies 
 in the fact that it has an intense azure-blue colour, and its formation 
 affords an extremely delicate test for either chromic acid or 
 hydrogen peroxide. A few drops of H 2 O 2 (or a few particles of 
 
 * The constitution of the dichromates (sometimes wrongly called 
 mates) may be expressed thus, K 2 CrO 4 ,CrO 3 . They are strictly analogous 
 to the pyrosulphates, K 2 S 2 O 7 , or K 2 SO 4 ,SO 3 , and on this account should 
 consistently be named pyrochromates. By the action of strong acids, the 
 normal potassium chromate is converted into the dichromate ; thus, 2K 2 CrO 4 
 -f H 2 SO 4 = K 2 SO 4 + H 2 O + K 2 CrO 4 ,CrO 3 . And the dichromate is re-con- 
 verted into the normal salt by the action of potash- 
 
 K 2 CrO 4 ,CrO 3 + aKHO = 2K 8 CrO 4 + H 2 O 
 
Group II L Division A. 43 
 
 Na 2 O 2 ) are added to half a test-tube of water, and the mixture 
 acidified with one or two drops of HC1. A single drop of potassium 
 dichromate solution added to this produces an intense blue colour. 
 [The compound is very unstable in aqueous solution, but less so in 
 ether ; therefore, in testing for very minute quantities, ether should 
 be added before the dichromate ; and on shaking the mixture, 
 the ethereal layer which rises to the surface will be coloured blue.] 
 Reduction of Chromic Acid. CrO 3 is a powerful oxidising 
 agent, giving up oxygen to oxidisable substances, and being itself 
 reduced to Cr 2 O 3 ; that is, to the condition of a "chromic" com- 
 pound. Thus, by sulphur dioxide it is reduced to chromium 
 sulphate 
 
 2Cr0 3 + 3 S0 2 = Cr 2 (S0 4 ) 3 
 
 The same action takes place in an acidified solution of potassium 
 dichromate 
 
 K 2 Cr,0, + H 2 S0 4 + 3H 2 S0 3 = Cr,(SO 4 ) 3 + K 2 SO 4 + 4H 2 O 
 
 Similarly, chromic acid and chromates are, reduced by HC1, oxidis- 
 ing the hydrogen of the acid, and liberating chlorine, after the 
 manner of peroxides ; thus 
 
 CrO, + 6HC1 = CrCl s + 3H 2 O + 3 C1 
 K 2 Cr 2 O 7 + UHC1 = 2CrCl s + 2KC1 + ?H 2 O + 3C1 2 
 
 On account of this reaction, a mixture of potassium dichromate 
 and hydrochloric acid is capable of " oxidising " FeCl 2 into FeCl ;{ ; 
 SnCl 2 into SnCl 4 ; As 2 O 3 into As 2 O. v In all cases of oxidation by- 
 chromic acid, the reduction of the chromic acid compound to the 
 state of a " chromic " compound is evidenced by the change of 
 colour from the yellow or orange of the former, to the green colour 
 of the latter. This reduction and change of colour is at once seen 
 by passing sulphuretted hydrogen through acidified potassium 
 dichromate 
 
 K 2 Cr 2 O 7 + ;,H 2 S + 8HC1 = 2CrCl s + 2KC1 + 7H 2 O + 38 
 
 Many organic substances also reduce chromic acid, such as oxalic 
 acid, and alcohol. Thus, one molecule of oxalic acid, C 2 H 2 O 4 , re- 
 quires one atom of O to convert it into CO 2 and H 2 O 
 
 C 2 H 2 4 + O = 2C0 2 + H 2 
 
 Potassium dichromate, in being reduced, has three available atoms 
 of oxygen to give up ; thus, K 2 Cr,O 7 = Cr 2 O 3 ,K 2 O,3O (in the 
 presence of dilute acids the Cr 2 O 3 and K 2 O form salts). Therefore 
 
44 Qualitative Analysis. 
 
 one molecule of K 2 Cr 2 O 7 can oxidise three molecules of oxalic acid, 
 resulting in the evolution of six molecules of CO 2 ; thus 
 
 K 2 Cr. 2 7 + 4H,S0 4 + 3C 2 H 2 4 = K 2 SO 4 + Cr 2 (SO 4 ) 3 + ;H 2 O 
 
 + 6C0 2 
 
 If alcohol be added to a mixture of potassium dichromate and 
 sulphuric acid, the alcohol (C 2 H 6 O) is oxidised first to aldehyde, 
 C 2 H 4 O, and then to acetic acuf (C 2 H 4 O 2 ), and the colour of the 
 mixture changes from orange-red to green. 
 
 Iron, Fe. 
 
 DRY REACTIONS. Iron compounds impart to a borax bead 
 heated in the outer flame, a colour which appears chocolate when 
 hot, and yellow when cold. After heating in the reducing flame, 
 the colour changes to a bottle-green (the green colour of common 
 bottle glass is caused by the presence of iron). When heated on 
 charcoal with Na 2 CO 3 in the inner blowpipe flame, iron com- 
 pounds become reduced, and a dark grey magnetic mass is obtained. 
 If this be washed with water in a small mortar, and the end of 
 a magnet applied, it will be attracted after the manner of iron 
 filings. 
 
 WET REACTIONS. The salts of iron are derived from the two 
 oxides FeO and Fe 2 O 3 .* They are both basic oxides, and give 
 rise to two classes of salts, namely, ferrous and ferric respectively. 
 Ferrous salts readily take up oxygen, and become converted into 
 ferric compounds ; while the latter, under the influence of suitable 
 reducing agents, easily pass back again to the ferrous condition. 
 
 (a) Ferric Compounds. The common ferric salts that are 
 soluble in water are the chloride, FeCl 3 ; nitrate, Fe 2 (NO 3 ) 6 , and 
 sulphate, Fe 2 (SO 4 ) 3 . These all give yellowish-brown solutions. 
 
 NH 4 HO, KHO, and NaHO throw down a brown voluminous 
 precipitate of ferric hydroxide,! Fe 2 (HO) , insoluble in excess, 
 or in NH 4 CU 
 
 K 2 CO t , Na 2 CO 3 , and BaCO 3 give the same precipitate, CO 2 
 being liberated 
 
 2FeCl s + 3Na 2 C0 3 + 3H 2 O = Fe 2 (HO) fi + 6NaCl + 3CO, 
 
 * The oxide known as magnetic oxide of iron, or ferroso-ferric oxide, 
 Fe 3 O 4 or Fe 2 O 3 ,FeO, yields a mixture of ferric and ferrous salts. 
 
 t Several hydrated ferric oxides are known, e.g. Fe 2 O 3 ,3H 2 O; Fe 2 O 3 ,2H 2 O ; 
 Fe 2 O 3 ,H 2 O. Common rust of iron has the composition 2Fe 2 O 3> 3H 2 O. The 
 composition of the precipitate produced by alkalies depends upon the con- 
 ditions of precipitation. 
 
 \ See footnote on p. 39, as to the influence of organic compounds. 
 
Group II L Division A. 45 
 
 The precipitate is soluble in a concentrated solution of K.,CO 3 , 
 giving a deep reddish solution of unknown composition. On the 
 addition of water the hydroxide is reprecipitated. (With BaCO 3 
 basic carbonates are also precipitated.) 
 
 (NH 4 ) 2 S produces a black precipitate of ferrous sulphide. 
 The action may be considered as taking place in two stages : 
 (i) the reduction of the iron to the ferrous state, and (2) the 
 formation of the ferrous sulphide ; thus, using a dissected formula 
 for ammonium sulphide 
 
 (i) :NH 3 NH 3 : H,S + 2FeCl 3 = 2FeCl 2 + 2HC1 + S 
 
 (2) (NH 4 ) 2 S + FeCl 2 = FeS 4- 2NH 4 C1 
 
 In the first' equation the hydrochloric acid formed unites with 
 the ammonia, producing 2NH 4 C1. 
 
 Sulphuretted hydrogen, H,S, brings about the first stage 
 in the above action, reducing the iron from the ferric to the ferrous 
 state with precipitation of sulphur, but in the presence of the free 
 acid which is developed by the action, ferrous sulphide cannot be 
 formed. [Ferric sulphide cannot be produced in the wet way.'] 
 
 Potassium ferrocyanide, K 4 Fe(CN) 6 , or K 4 FeCy 6 ,* pro- 
 duces with ferric salts a dark blue precipitate {Prussian blue] 
 
 3 K 4 (FeCy 6 ) + 4 FeCl 3 = I2KC1 + Fe 4 (FeCy 6 ) 3 
 
 This test is extremely delicate, but where the amount of iron 
 is very small, a blue or greenish coloration only is produced. 
 " Prussian blue " is insoluble in hydrochloric acid, but readily 
 dissolves in oxalic acid. It is decomposed by NaHO or KHO,. 
 with precipitation of ferric hydroxide 
 
 Fe 4 (FeCy 6 ) 3 + I2KHO = 2Fe 2 (HO) 6 + 3 K 4 (FeCy 6 ) 
 
 Potassium ferricyanide, K 3 (FeCy 6 ), gives no precipitate with 
 ferric salts. 
 
 Potassium thiocyanate, K(CN)S, produces with ferric 
 salts a rich wine-red coloration, owing to the formation of ferric 
 thiocyanate, Fe(CNS) 3 , which is soluble in water. The colour of 
 
 * "Cy" is a recognised and convenient symbol for the radical (CN) ; 
 cyanogen. The use of this reagent as a test for iron is unique, as being the 
 only case in which the reagent is itself a compound containing the very metal 
 it is employed to detect. The ferrocyanides and the ferri cyanides, however, 
 although compounds of iron, contain the metal in such a form of combination 
 that it is not detected by the ordinary reagents. Before the iron in these 
 t:ompounds will give any of the ordinary reactions, its union with the cyanogen 
 radical must be first destroyed (see Cyanides, p. 163). 
 
46 Qualitative Analysis. 
 
 this compound is very intense, hence the reaction may be employed 
 to detect very small quantities of iron.* 
 
 Reduction of Ferric to Ferrous Compounds. FernV 
 
 compounds are readily reduced to the ferrous state ; they are 
 therefore oxidising agents of some importance. The action of 
 (NH 4 ) 2 S and of H 2 S has been already mentioned. Nascent 
 hydrogen reduces them in the same way ; therefore, when metallic 
 iron is dissolved in HC1 or H 2 SO 4 , the salts produced are ferrous 
 chloride and sulphate respectively. Nitric acid, on the other- 
 hand, converts the iron into the " ferric " state. 
 
 A ferric salt already in solution is reduced by nascent hydrogen, 
 generated by introducing zinc into the acidified liquid. 
 
 In passing from FeCl 3 to Fed.,, one atom of chlorine is 
 available for oxidising purposes, and is capable of bringing about 
 such actions as the following 
 
 The "oxidation" of stannous chloride, SnCL, to stannic 
 chloride, SnCl 4 . 
 
 The oxidation of sulphurous acid or thiosulphuric acid into 
 sulphuric acid ; thus 
 
 2FeCl 3 4- H 2 SO 3 4- H 2 O = H 2 SO 4 + 2HC1 4- 2FeCL 
 2FeCl 3 4- Na 2 SS0 3 4- H 2 O = Na,SO 4 4- 2HC1 + 2FeCl, 4- S 
 
 (If) Ferrous Compounds. Ferrous salts are usually pale 
 green when crystallised, and white when anhydrous. Of the 
 common salts the chloride and sulphate are soluble. The latter 
 readily forms double salts with the sulphates of the alkalies (such 
 as ferrous ammonium sulphate, FeSO 4 ,(NH 4 ) 2 SO 4 ,6H 2 O), which 
 are also soluble in water, and are less readily oxidised on exposure 
 to the air than ferrous sulphate, which they otherwise closely 
 resemble in appearance. 
 
 NH 4 HO, KHO, and NaHO produce a precipitate of ierrous 
 hydroxide, Fe(HO) 2 , which is at first a dirty white colour, but which 
 rapidly turns first pale greenish-grey, then a dirty grey, and finally 
 brown, owing to its oxidation by atmospheric oxygen. The presence 
 of ammonium salts renders the precipitation incomplete. The 
 
 * This is a "reversible" reaction, and therefore, when equilibrium is 
 established, there will be present in the liquid both ferric chloride and potas- 
 sium thiocyanate, thus, aKCNS + FeCl 3 ; aKCl + Fe(CNS) 3 . That this is 
 so may be proved by the following experiment : Add to a little moderately 
 dilute ferric chloride a small quantity of potassium thiocyanate; then dilute 
 the liquid with water so that the intensity of the red colour is greatly reduced, 
 and divide it into two portions. To one add more ferric chloride, and to the 
 other add more potassium thiocyanate. In each case the liquid becomes a 
 deeper red colour. 
 
Group III. Division A. 47 
 
 precipitate is not soluble in excess of the reagents ; boiling with 
 KHO turns it black, converting it into Fe 3 O 4 . 
 
 K CO > and Na.,CO 3 give a white precipitate of ferrous carbonate, 
 FeCO 3 , which on exposure to the air quickly absorbs oxygen. 
 
 (NH 4 )oS precipitates black ferrous sulphide, FeS. Readily 
 soluble in acids, with evolution of sulphuretted hydrogen ; insoluble 
 in alkalies. The precipitate in the moist state is oxidised on exposure 
 to the air into ferrous and basic ferric sulphate. 
 
 X 4 (FeCy 6 ) precipitates potassium ferrous ferrocyanide, 
 FeK 2 (FeCy 6 ), thus- 
 
 K 4 (FeCy rt ) + Fed, = 2KC1 + FeK,(FeCy 6 ) 
 
 When the solutions are mixed in test-tubes in the ordinary way, 
 the precipitate has a greenish-blue colour ; but when the reaction 
 is made in an atmosphere free from oxygen, and the solutions are 
 previously boiled so as to entirely expel all dissolved oxygen, the 
 precipitate is perfectly white. It rapidly absorbs oxygen and 
 becomes blue, and is also easily oxidised to " Prussian " blue by 
 nitric acid or chlorine ; thus 
 
 4 FeK 2 (FeCy 6 ) + 2C1, = Fe 4 (FeCy 6 ) 5 + K 4 FeCy 6 + 4 KC1 
 
 Potassium ferricyanide, K 3 (FeCy 6 ), gives, with ferrous 
 salts, a precipitate of ferrous ferricyanide, Fe 3 (FeCy G )., (known as 
 TurnbulCs blue), which is indistinguishable by its appearance from 
 Prussian blue 
 
 2K 3 ,FeCy 6 ) + 3 FeCl 2 = Fe 3 (FeCy 6 ), 4- 6KC1 
 
 The precipitate is insoluble in hydrochloric acid, but is decom- 
 posed by caustic alkalies, with the precipitation of ferrous hydroxide ; 
 thus 
 
 Fe 3 (FeCy 6 ) 2 + 6KHO = 2K 3 (FeCy 6 ) + 3F*(HO) 2 
 
 Oxidation of Perrons to Perric Compounds. The ferric 
 salts being the more stable, the ferrous compounds undergo oxida- 
 tion even more readily than the ferric salts become reduced. Mere 
 exposure to the air in many cases causes the change. In analysis 
 the oxidation is usually accomplished either by chlorine (or bromine) 
 or by nitric acid. 
 
 The chlorine may be employed in the form of its aqueous 
 solution (chlorine water), or more conveniently by generating the 
 gas in contact with the ferrous compound by means of hydrochloric 
 acid and potassium chlorate. The solution of the ferrous salt is 
 acidified with concentrated HC1, and heated. A few particles of 
 
48 Qualitative Analysis. 
 
 potassium chlorate arc then dropped into the mixture, and the 
 heating continued for a short time. 
 
 A mixture of HC1 and KC1O 3 evolves both chlorine and chlorine 
 peroxide ; thus 
 
 >x4KClO, + I2HC1 = 6K 2 O + 4KC1 + 3C1O, + gC\ 
 
 Both tfac free chlorine and the chlorine of the chlorine peroxide 
 are available for oxidising the ferrous compound ; hence the equation 
 may be simplified as follows : 
 
 KC10 3 -f 6HC1 + 6FeCl, = 6FeCl 3 + KC1 + 3H,O 
 
 vVhen the oxidation is accomplished with nitric acid, the strong 
 ^Mfcifeaclded, a few drops at a time, to the hot acidulated solution 
 of the ferrous salt. The solution becomes dark in colour, and nitric 
 oxide is disengaged ; thus 
 
 6FeS0 4 + 3H 2 SO 4 + 2HNO 3 = 3Fe 2 (SO 4 ) 3 -I- 4H 2 O + 2NO 
 3FeCl, + 3HC1 + HNO 3 = 3FeCl 3 + 2H 2 O + NO 
 
 Unless the solution of the ferrous salt is acidified, a portion of 
 the iron is converted into Fe 2 O 3 , which is taken up, in the case 
 of the sulphate, by the ferric sulphate, forming insoluble basic ferric 
 sulphates, Fe 2 (S0 4 ) 3 , *Fe,O 3 . 
 
 SEPARATION OF THE METALS OF GROUP II I A. 
 
 The separation of the metals of this subdivision from the other 
 metals of Group III., and also from those of Groups IV. and V., is 
 based upon the fact that their hydrated sesquioxides are precipi- 
 tated by ammonia in the presence of ammonium chloride.* 
 
 The separation of the three metals of this group from each 
 other is based upon 
 
 1. The oxidation of chromic oxide to chromic acid ; and 
 
 2. The solubility of aluminium hydroxide in caustic alkalies. 
 To the solution add NH 4 C1 in considerable quantity ; heat the 
 
 mixture to boiling, and add NH 4 HO carefully until precipitation is 
 complete. Bring the liquid once more " to the boil," when, if suffi- 
 cient ammonia has been added, the steam will smell of it. Filter 
 the mixture while hot.f 
 
 * The separation of Group I1I.\. from Group lllB. by means of NH 4 HO is 
 not sharp and complete in all cases (see Manganese reactions). 
 
 t In the regular course of a complete analysis, the filtrate obtained here 
 will contain the metals of Group IIlB. , IV., and V. 
 
Separation of the Metals of Grotip fllA 49 
 
 The precipitate consists of A1 2 (HO) 6 , Cr 2 (HO) 6 , and Fe 2 (HO) 6 . Wash 
 the precipitate, and transfer it (or a portion of it) to a test-tube with 
 a small quantity of water. Add to the mixture a little sodium per- 
 oxide, and boil for a moment, until the temporary effervescence ceases. 
 The chromium is oxidised to chromate, and the A1 2 (HO) 8 dissolves 
 in the NaHO, which is formed by the action of the sodium peroxide 
 upon the water. Filter. 
 
 The filtrate contains sodium chromate,Na 2 CrO 4 , 
 and sodium aluminate, Al 2 O 3 ,3Na 2 O. The 
 former shows itself by the yellow colour. 
 Divide into two portions 
 
 (1) Acidify with acetic acid, and confirm 
 chromium by special reactions, e.g. lead 
 acetate. 
 
 (2) Acidify with dilute nitric acid, and add 
 NH 4 HO. A white precipitate of A1 2 (HO; 6 
 confirms aluminium. 
 
 The residue consists of 
 Fe 2 (HO) 6 . Dissolve 
 in a little hot dilute 
 HC1, and confirm 
 iron by special re- 
 actions, e.g. K 4 FeCy 6 
 or KCNS.* 
 
 The following alternative methods of separation may also be 
 used. 
 
 (a) The precipitated hydroxides are washed and dried. The 
 residue is then mixed with at least six times its weight si fusion 
 mixture, ,f and fused in a platinum capsule. In this way the 
 chromium is converted into alkaline chromate ; a variable pro- 
 portion of the aluminium into aluminates. 
 
 The fused mass is then dissolved in water, and filtered. The 
 filtrate is tested for aluminium and chromium, while the residue is 
 dissolved and tested for iron, as in the foregoing scheme. 
 
 (b) The precipitated hydroxides are dissolved in a little warm 
 dilute HC1, and pure NaHO J added in quantity considerably 
 more than sufficient to produce precipitation. The mixture is then 
 boiled for a few minutes, and filtered. 
 
 * At this stage in the process, the iron will be in the "ferric" condition. 
 To ascertain whether it was originally present as a "ferrous" or "ferric" 
 compound, separate tests must be made in the solution before it has been 
 subjected to the action of either reducing or oxidising agents. 
 
 f Fusion mixture is a mixture of Na 2 CO 3 and K 2 CO 3 in equivalent propor- 
 tions (or about 10 parts Na 2 CO 3 to 13 of K 2 CO 3 ). It is used in preference to 
 Xa 2 CO 3 alone, because it has the property of melting more easily than either 
 carbonate separately. 
 
 The commercial caustic soda usually employed in the laboratory always 
 contains more or less sodium aluminate. The student should test a sample of 
 the reagent by neutralising it with HC1, and then adding NH 4 HO. In the 
 method of separation given above, this difficulty is avoided, as the sodium 
 peroxide of commerce is usually free from this impurity. 
 
 
 
Qualitative A nalysis. 
 
 The filtrate contains sodium aluminate, Al.,O 3 ,3Na. 2 O. Add 
 dilute HC1 until just acid, and reprecipitate Al,(HO),j with 
 ammonia. 
 
 The precipitate contains Cr 2 (HO) (! and Fe 2 (HO) . This is 
 dried, and fused with fusion mixture. The fused mass is dissolved 
 in water and filtered. The solution contains sodium chromate, 
 while the Fe 2 O 3 remains on the filter. These are confirmed as 
 in the above methods. 
 
 
 14* 
 
 
CHAPTER VI. 
 REACTIONS OF THE METALS OF GROUP III. DIVISION B. 
 
 Manganese, Mn. 
 
 DRY REACTIONS. Manganese compounds, when heated in a 
 borax bead in the oxidising flame, impart to the bead a violet or 
 lilac colour. When heated in the reducing flame, the bead again 
 becomes colourless. 
 
 A more characteristic reaction is based upon the oxidation of 
 manganese to manganic acid. When a manganese compound is 
 fused with KHO, or with Na 2 CO 3 and a little KNO 3 or KC1O 3 
 upon a platinum capsule, the manganese undergoes oxidation, and 
 a deep green-coloured mass is obtained, consisting of manganates 
 of the alkali metals 
 
 MnO, + Na 2 CO 3 + O (from KC1O 3 or KNO 3 ) - Na 2 MnO 4 + CO, 
 MnO 2 + 2KHO + O = K 2 MnO 4 + H 2 O 
 
 The green mass (especially when obtained by fusion with KHO) 
 dissolves in a small quantity of cold water to a deep green solution. 
 When this is either acidified, or warmed, or even largely diluted 
 with water, its colour changes from green to pink, owing to the 
 conversion of the manganate into permanganate ; thus 
 
 3K 3 MnO 4 -f 2H,O = 2KMnO 4 + 4.KHO + MnO 2 
 
 WET REACTIONS. Of the oxides of manganese two are basic : 
 the monoxide, MnO, giving rise to the ma.nga.nous salts ; and the 
 sesquioxide, Mn 2 O 3 , to the manganzV salts. 
 
 [The oxide, Mn 3 O 4 , yields both manganous and manganic salts ; 
 while MnO 2 gives manganous salts with the elimination of available 
 oxygen.] 
 
 The manganic salts are extremely unstable in solution, and are 
 incapable of existing as such under the ordinary conditions of 
 analysis. Thus the chloride, believed to have the composition 
 Mn 2 Cl c , passes at ordinary temperatures into MnCl 2 and chlorine. 
 
52 Qualitative Analysis. 
 
 Of the common manganous salts, the chloride, MnCl 2 , 4H,O, 
 and the sulphate, MnSO4,5H 2 O, are soluble in water. In the 
 crystallised state they have all a pink colour. Manganous salts 
 which are soluble in water do not undergo atmospheric oxidation. 
 
 NH,HO, KHO, and NaHO produce a white precipitate of 
 manganous hydroxide, Mn(HO) 2 . Insoluble in excess of the re- 
 agent, the precipitate quickly absorbs oxygen, and is converted 
 into hydrated manganic oxide, Mn 2 O 3 ,H 2 O, having a brown colour. 
 
 Freshly precipitated Mn(HO) 2 , while still white, is soluble in 
 NH 4 C1, forming the soluble double salt MnCl 2 ,2NH 4 Cl,H 2 O ; there- 
 fore in the presence of NH 4 C1, manganous hydroxide is not precipi- 
 tated by NH 4 HO, and only incompletely by KHO or NaHO. 
 
 The ammoniacal solution of the double chloride, however, is 
 capable of absorbing oxygen just as the precipitated Mn(HO 2 ) does, 
 and the liquid quickly becomes muddy, owing to the precipitation 
 from it of the brown hydrated manganic oxide.* 
 
 K,CO ( , NajjCOg, or (NH 4 ) 2 CO 3 give a white precipitate of 
 manganous carbonate, MnCO 3 ; insoluble in excess. The precipi- 
 tation is not complete in the presence of NH 4 C1. Manganous 
 carbonate absorbs atmospheric oxygen, and is slowly changed into 
 the brown hydrated oxide. 
 
 (NTI 4 ) 2 S precipitates manganous sulphide, MnS, as a pale 
 pinkish-white compound, easily soluble in dilute acids (distinction 
 from Ni and Co), soluble also in acetic acid (distinction from Zn). 
 Precipitation with (NH 4 ) 2 S is only complete in the presence of 
 NH 4 C1. Owing to the ready solubility of MnS in acids, H 2 S is 
 incapable of precipitating manganous sulphide from neutral solu- 
 tions ; for, by double decomposition, the acid of the manganous 
 salt would be set free, and would immediately redissolve the 
 sulphide. 
 
 Manganese Compounds as Oxidising Agents. When 
 manganese dioxide is acted upon by acids, a manganous salt is 
 
 * It is this characteristic property of the double manganous ammonium 
 chloride that renders the complete separation of this metal from those of 
 Group IIlA. extremely difficult to accomplish. The hydrated oxides of Al, 
 Cr, and Fe are sure to carry down more or less of the manganese along with 
 them, and a small quantity of manganese, in the presence of a large proportion 
 of iron, might in this way be altogether overlooked. If the precipitation with 
 NH 4 HO of Group IIlA., be made as quickly as possible in a hot solution, 
 and the excess of ammonia at once boiled off, and the liquid filtered imme- 
 diately, the risk of precipitating the manganese maybe reduced to a minimum. 
 The student will do well to practise the separation of manganese from the 
 metals of Group IIlA. by using solutions containing known small propor- 
 tions of a manganous salt, mixed with large quantities of iron or chromium 
 or aluminium, 
 
Group ILL Division B. 53 
 
 formed, and available oxygen is eliminated, which either appears 
 as free oxygen gas, or as the product of the oxidation of the acid ; 
 thus 
 
 MnO 2 + H 2 SO 4 = MnSO 4 + H O + O 
 MmO a + 4HC1 = MnCL + 2H 2 O + C1 2 
 
 The other oxides of manganese, higher than the monoxide, may 
 be regarded as compounds of MnO 2 with MnO in different propor- 
 tions, and when acted upon by hydrochloric acid they also yield 
 manganous chloride and chlorine, the amount of chlorine being the 
 measure of the amount of MnO 2 in the compound 
 
 Mn 3 4 or 2MnO,MnO, + 8HC1 = 3MnCl 2 + 4H 2 O + CL 
 Mn,O 3 or MnO,MnO, + 6HC1 = 2MnCL + 3H 2 O + CL 
 
 Nitric acid (free from nitrous add} has no action upon man- 
 ganese dioxide ; when, therefore, one of these other oxides is acted 
 upon by nitric acid, the monoxide is converted into the nitrate, and 
 the dioxide is left ; thus 
 
 Mn,0 3 or MnO,MnO 2 + 2HNO 3 = Mn(NO 3 ), + H 2 O + MnO 3 
 
 The manganates and permanganates are still more powerful 
 oxidising agents. When these are acted upon by sulphuric acid, we 
 may suppose that the hypothetical manganic and permanganic 
 acids are first liberated, which, being incapable of existence, break 
 up into the unstable oxides MnO 3 and Mn 2 O 7 ; and that these in 
 contact with sulphuric acid form manganous sulphate, with the 
 evolution of oxygen ; thus 
 
 rK 2 MnO 4 + H.,SO 4 = K 2 SO 4 + H 2 MnO 4 ( = H 2 O + MnO 3 ) 
 I MnO 3 + H 2 "SO 4 = MnSO 4 +*H 2 O + O 2 
 2KMnO 4 *+ H 2 SO 4 = K 9 SO 4 + 2HMnO 4 (= H,O + MnoO.) 
 Mn 2 O 7 + 2H 2 SO 4 = 2MnSO 4 + 2H 2 O + 50 
 
 From these equations, it will be seen that from one molecule of 
 the manganate two atoms of oxygen are given off, while two mole- 
 cules of the permanganate evolve five atoms of oxygen. When 
 acted upon by hydrochloric acid, the equivalent quantity of chlorine 
 is evolved 
 
 KMnO 4 + 8HC1 = KC1 + MnCL, + 4H 2 O + 5C1 
 
 Potassium permanganate is therefore a most important oxidising 
 material. In the presence of hydrochloric or sulphuric acid, it is 
 
 * Some chemists prefer to use the double formula, K 2 Mn 2 Os, for potassium 
 permanganate. 
 
54 Qualitative Analysis. 
 
 capable of oxidising almost every oxidisable compound. The 
 following may be taken as typical examples : 
 
 (1) Sulphurous acid to sulphuric acid 
 
 2KMnO 4 + 5H 2 SO 3 = K 2 SO 4 + 2MnSO 4 + 3H,O + 2H. 2 SO 4 
 
 (2) Oxalic acid to carbon dioxide and water 
 
 H 2 C,O 4 + O = H 2 O + 2CO, 
 
 One atom of oxygen is required for the oxidation of one mole- 
 cule of oxalic acid ; therefore the five atoms of oxygen derivable 
 from two molecules of permanganate will oxidise five molecules of 
 oxalic acid. 
 
 (3) Ferrous sulphate to ferric sulphate, in the presence of free 
 sulphuric acid 
 
 2 FeS0 4 + H 2 S0 4 + O = H 2 O + Fe 2 (SO 4 ) 3 
 
 One molecule of KMnO 4 is therefore capable of oxidising five 
 molecules of FeSO 4 ; or, in the presence of HC1, of oxidising five 
 molecules of FeCL into FeCl 3 .* 
 
 In all these cases of oxidation the change is accompanied by 
 the destruction of the violet colour of the permanganate ; hence it 
 is perfectly easy to watch the progress of the action, and to ascer- 
 tain the exact moment when the process is complete. 
 
 Zinc, Zn. 
 
 DRY REACTIONS. Zinc compounds give no characteristic 
 borax bead. 
 
 When heated on charcoal with sodium carbonate in the reducing 
 
 * Two powerful oxidising substances are often able to oxidise each other in 
 such a way that at first sight it would seem that they were playing the part of 
 reducing agents. Thus, when potassium permanganate and hydrogen peroxide 
 are brought together, both compounds are reduced. The available oxygen 
 from the permanganate oxidises the available oxygen of the hydrogen peroxide, 
 with the result that a number of oxidised atoms' of oxygen (i.e. complete mole- 
 cules) are set free. ,In the presence of H 2 SO 4 , the five oxygen atoms available 
 in two molecules of permanganate oxidise five available atoms of oxygen con- 
 tained in as many molecules of hydrogen peroxide ; thus 
 
 2 KMn0 4 + sH 2 SO 4 + 5H 2 O 2 = sH,O + 3 H 2 O + K 2 SO 4 + 2\InSO 4 + sO 2 
 
 In the absence of the free acid, hydrated sesquioxide of manganese is pre- 
 cipitated, which then becomes a catalytic agent, being alternately oxidised and 
 reduced again 
 
 2KMnO 4 + 4H 2 O 2 = 2KHO + Mn 2 O 3 ,H 2 O + 2H 2 O + 4O 2 
 
 I MnoO 3 + HO = H.O 4- aMnO. 
 
 'I aMn0 2 * H" 2 0-= H 2 O f O 2 + Mn 2 Q 3 
 
Group III. Division B. 55 
 
 flame, zinc compounds are reduced, but the metal is too volatile to 
 appear in the form of globules. As it is reduced, it volatilises ; and 
 the vapour burns as it passes through the outer flame, which thereby 
 becomes tinged a bluish-white colour. The zinc oxide which is 
 thus produced, deposits as an incrustation upon the charcoal, which 
 is canary-yellow while hot, becoming white on cooling. Zinc oxide 
 is not volatile, and therefore the incrustation does not disappear 
 when the oxidising flame is made to play upon it. If the zinc oxide 
 be moistened with a drop of cobalt nitrate, and again heated in the 
 oxidising flame, it assumes a green colour. 
 
 WET REACTIONS. Zinc forms only one series of salts, derived 
 from the only oxide, ZnO. 
 
 Of the common salts the chloride, sulphate, and nitrate are 
 soluble in water. 
 
 NH.HO, KHO, or NaHO throws down a white precipitate of 
 Zn(HO) 2 , readily soluble in excess of the reagent, forming double 
 salts (sometimes called zincates], such as ZnNa 2 O 2 .* Moderately 
 strong solutions of these zincates may be boiled without undergoing 
 any change, but from dilute solutions Zn(HO) 2 is reprecipitated. 
 Zn(HO) 2 is soluble in NH 4 C1, owing to its readiness to form soluble 
 double salts with the alkaline chlorides, having the general formula 
 ZnCl 2 ,2RCl 
 
 Zn(HO) 2 + 4NH 4 Ci = ZnCLj^NH.Cl + 2NH 4 HO 
 
 Hence in the presence of much ammonium chloride, NH 4 HO gives 
 no precipitate, and the precipitation with KHO or NaHO is in- 
 complete. 
 
 K 2 CO 3 , Na 2 CO 3 , or (NH 4 ) 2 CO 3 produces a white precipitate of 
 basic carbonate, the composition of which varies with conditions 
 of precipitation, ,rZnCO 3 , jZn(HO) 2 , H 2 O.f The precipitate is 
 soluble in excess of (NH 4 ) 2 CO 3 . The presence of much NH 4 C1 
 partially or entirely prevents the precipitation. 
 
 * Upon this property is based the separation of zinc from Mn, Ni, and Co, 
 but the solubility of Zn(HO) 2 in caustic alkalies is rendered less easy by the 
 presence of the hydroxides of Mn, Ni and Co, owing to the tendency of these 
 to unite with the zinc oxide to form compounds which are difficultly decom- 
 posed by the alkali. 
 
 + 3ZnCl 2 + 3 Na 2 CO 3 + 5H 2 O = ZnCO 3 ,2Zn(HO),HoO 6NaCl + aCO-, 
 5ZnCl 2 + 5Na 2 CO 3 + 8H,O = 2ZnCO 3 ,3Zn(HO) 2 ,5H 2 O + xoNaCl + 3CO 2 
 
 The former of these basic carbonates is known in pharmacy as sinci carbonas. 
 The normal zinc carbonate is precipitated by hydrogen sodium carbonate. 
 We may suppose that the acid which is in this case liberated, prevents the 
 formation of the hydroxide 
 
 ZnCl- + HNaCOj - ZnCO 3 + HC1 + NuCl 
 
56 Qualitative Analysis. 
 
 The precipitate is soluble also in a concentrated solution of 
 K 2 CO 3 , but is reprecipitated on dilution with water (compare 
 Fe, Ni, Co). 
 
 (NH 4 ) 2 S throws down a white precipitate of zinc sulphide, 
 ZnS. In the presence of NH 4 C1 the precipitation is complete even 
 from dilute solutions. ZnS is soluble in dilute mineral acids, hence 
 H 2 S is incapable of completely precipitating this sulphide from 
 neutral solutions of the zinc salts of such acids 
 
 ZnCl 2 4- H 2 S ZnS +2HC1 
 
 ZnS is insoluble in acetic acid (contrast MnS\ therefore from the 
 acetate, or other sine salts in presence of an alkaline acetate, ZnS 
 is completely precipitated by H 2 S ; thus 
 
 Zn(C 2 H 3 O 2 ) 2 + H 2 S = ZnS 4- 2H(C 2 H 3 O 2 ) 
 ZnCl 2 4 2Na(C 2 H s O 2 ) 4- H,S = ZnS 4 -NaCl + 2H(C>H S (X 
 
 Nickel, Ni. 
 
 DRY REACTIONS. Nickel compounds impart a dark red-brown 
 colour to the borax bead when heated in the oxidising flame, the 
 colour becoming brownish-yellow on cooling. In the reducing 
 flame the borax bead becomes opaque and grey. In a bead of 
 microcosmic salt, the red-brown colour persists in both flames. 
 
 The presence of other colour-producing oxides renders this test 
 uncertain, while even traces of cobalt entirely mask it. Heated 
 on charcoal with Na 2 CO 3 , metallic nickel is obtained as a grey 
 feebly magnetic mass. 
 
 WET REACTIONS. Only one of the oxides of nickel is basic, 
 namely NiO, hence only one series of salts exists. The sesquioxide, 
 Ni 2 O 3 , behaves like a peroxide ; thus 
 
 Ni 2 3 4- 2H 2 SO 4 = 2NiSO 4 + 2H 2 O -f O 
 
 In the crystalline or hydrated condition the nickel salts have a 
 green colour, and dissolve to green solutions. The anhydrous salts 
 are pale yellow. Of the common salts, the chloride, nitrate, and 
 sulphate are soluble in water. Nickel salts form a number of 
 soluble double salts ; those with ammonium salts being important. 
 KHO or NaHO gives a pale bluish-green precipitate of nickelous 
 hydroxide, Ni(HO) 2 ,* insoluble in excess of either reagent; soluble 
 
 * The composition of the precipitate is more exactly expressed by the 
 formula 4Ni(HO) 2 ,H 2 O. When dried and strongly heated, it is converted 
 into NiO. 
 
Group III. Division B. 57 
 
 in ammonium salts. Ni(HO) 2 is not oxidised on exposure to air, 
 but it is converted into black hydrated sesquioxide, Ni. 2 O 3 ,3H 2 O, 
 by hypochlorites, or by the action of chlorine in the presence of 
 caustic alkalies ; thus 
 
 2Ni(HO), + NaCIO + H 2 O = NaCl + Ni 2 O 3? 3H 2 O 
 2Ni(HO) 2 + 2NaHO + CL, = -NaCl + Ni 2 O 3 ,3H,O* 
 
 NH 4 HO forms a number of readily soluble double compounds 
 with nickel salts. When added to an acid solution of a nickel salt, 
 e.g. NiSO 4 , no precipitate is produced, owing to the ready solubility 
 of Ni(HO) 2 in ammonium salts. With neutral solutions partial 
 precipitation takes place, the precipitate quickly dissolving in excess 
 of ammonia to a greenish-blue solution. This blue liquid contains 
 in solution the salt NiSO 4 ,4NH 3 . The nickel in this solution is 
 not oxidised by hypochlorites, but it is completely precipitated as 
 Ni\HO) 2 by KHO. 
 
 K CO or Na 2 CO 3 produces a pale-green precipitate of basic 
 carbonate, ^NiCO 3? >'Ni(HO) 2 . 
 
 The precipitate is soluble to a pale-green solution in a con- 
 centrated solution of K 2 CO 3 , but is reprecipitated on dilution with 
 water. 
 
 (NH 4 ) 2 CO 3 gives no precipitate in acid solutions, but from 
 neutral solutions a similar compound is produced, which dissolves 
 in excess of the reagent to a bluish solution. 
 
 (NH 4 ) 2 S, or H 2 S in presence of ammonia, produces a black pre- 
 cipitate of NiS, soluble to a slight extent in excess; more readily 
 soluble if ammonia or polysulphides of ammonia are present, 
 yielding a brown solution. From this solution the dissolved NiS 
 is reprecipitated slowly on boiling, more quickly after acidifying 
 with acetic acid, or the addition of ammonium acetate. 
 
 NiS is only difficultly soluble in strong HC1, and almost insoluble 
 in the dilute acid ; also in acetic acid. Readily soluble in aqua 
 
 * This higher oxide undergoes rapid alternate reduction and oxidation in 
 the presence of the hypochlorite. As soon as it is formed it is acted upon by 
 the hypochlorite, with evolution of oxygen ; thus, formulating the oxides for 
 simplicity 
 
 Ni 2 O 3 + NaCIO = NaCl + 2NiO + O 2 
 
 The compound therefore becomes a catalytic agent, causing the evolution of 
 oxygen from a relatively infinite quantity of a hypochlorite. Cobalt oxide 
 behaves in the same way (see Methods of obtaining oxygen, Newth's " Inorganic 
 Chemistry "). The hydrated sesquioxide behaves with acids in the same wax- 
 as the anhydrous oxide ; thus with HC1 it yields chlorine 
 
 Ni 2 3 , 3 H 2 + 6HC1 = 2\iCl 2 + 6H 2 O+ Cl- 
 
58 Qualitative Analysis. 
 
 regia, or in HC1 and a crystal of KC1O 3 , yielding NiCl 2 ; soluble 
 also in HNO 3 . 
 
 H 2 S only produces complete precipitation of NiS from a warm 
 solution of the acetate, or from other nickel salts in the presence of 
 an alkaline acetate. In the case of neutral solutions of nickel salts 
 with mineral acids, the precipitation is only partial, while in acid 
 solutions it does not take place at all. 
 
 Cobalt, Co. 
 
 DRY REACTIONS. Cobalt compounds impart to a borax bead 
 a rich blue colour, when heated either in the oxidising or reducing 
 flame. The test is characteristic and delicate. Many metallic 
 oxides which give coloured borax beads become either colourless or 
 lighter in colour when heated in the reducing flame, hence when 
 mixed with these, the blue colour of the cobalt becomes more 
 visible after heating the bead in the inner flame. 
 
 WET REACTIONS. Cobalt forms a number of oxides, two only 
 of which are basic. The cobaltous salts are derived from CoO, 
 while the feebly basic sesquioxide Co.,6 3 forms the very unstable 
 cobaltic salts. 
 
 Of the common cobaltous salts, the sulphate, nitrate, and haloid 
 salts are soluble in water. In the hydrated condition they have a 
 pink colour, dissolving to pink solutions. In the anhydrous state 
 they are blue. A strong aqueous solution (pink) will, however, turn 
 blue when boiled, and return to its original pink colour when again, 
 cooled. In alcohol, cobalt chloride dissolves to a deep blue solu- 
 tion, which turns pink on the addition of water. 
 
 KHO or NaHO gives a greenish-blue precipitate of a basic salt. 
 On boiling with excess of the alkali, the precipitate is converted 
 into the pink hydroxide, Co(HO) 2 , which, however, is coloured 
 more or less brown by the oxidation of a portion of it (by atmo- 
 spheric oxygen) into the hydrated cobaltic oxide, Co.,O 3 ,3H 2 O, or 
 Co 2 (HO) 6 . 
 
 Co(HO) 2 is oxidised by hypochlorites, in the same manner as 
 the corresponding nickel compound. 
 
 NHjHO produces no precipitate in acid solutions. In neutral 
 solutions it causes partial precipitation of a basic salt, which dis- 
 solves easily in excess. The solution, which has a brownish colour, 
 absorbs oxygen, and becomes darker in colour.* 
 
 * Cobalt (in common with platinum and a few other metals) possesses the 
 property of forming a large number of compounds with ammonia. These 
 salts, which are many of them extremely complex in their constitution, and 
 
Group II L Division B. 59 
 
 K CO or Na 2 CO 3 precipitates a lilac-coloured basic carbonate, 
 .rCoC0 3 , yCo(HO) 2 . 
 
 The precipitate is soluble to a violet solution in a concentrated 
 solution of K 2 CO 3 , but is reprecipitated on dilution with water 
 (compare Co, Zn, Cu). In all these cases the solution is believed 
 to contain a double carbonate, which is unable to exist in the 
 presence of much water. 
 
 HKCO gives a pinkish precipitate of normal cobalt carbonate, 
 CoCO 3 . The precipitate so obtained is soluble in hydrogen per- 
 oxide, yielding a deep green solution.* The production of this 
 green solution constitutes a very delicate test for cobalt. It may be 
 carried out in the following way ; to a strong solution of potassium 
 vor sodium) bicarbonate, about an equal bulk of hydrogen peroxide 
 is added. On the addition of a single drop of cobalt chloride, the 
 green colour will appear.t 
 
 (NH 4 ) 2 CO 3 gives no precipitate in acid solutions. In neutral 
 solutions partial precipitation of basic carbonate takes place ; the 
 precipitate readily dissolves in excess to a reddish solution. 
 
 (NH 4 )oS gives a black precipitate of cobaltous sulphide, CoS. 
 The precipitation is complete in the presence of NH 4 C1 (contrast 
 NiS). CoS is soluble in HNO 3 , in "aqua regia," and in HC1 with 
 the addition of a crystal of KC1O 3 ; difficultly soluble in strong 
 HC1 ; practically insoluble in dilute HC1. 
 
 H 2 S precipitates CoS under the same conditions as apply in the 
 case of NiS. 
 
 difficult to classify, are known as cobaltamines. The animonio-cobaltous com- 
 pounds easily absorb oxygen, and pass into the more complex and more 
 numerous ammonio-cobaltic salts. Nickel does not form compounds corre- 
 sponding to the ammonio-cobaltic salts. 
 
 * Durrant. 
 
 f The exact nature of this curious reaction has not yet been established. 
 There is little doubt but that the cobalt is present in a higher state of oxidation 
 than in Co 2 O 3 ; probably the compound is a cobaltate corresponding to the 
 ferrates and manganates. Carbon dioxide is evolved when it is produced, 
 which may be explained by the equation 
 
 CoCO 3 + 2H 2 O 2 = H 2 CoO 4 + CO 2 + H 2 O 
 
 At the same time, the compound is able to act as a catalytic agent towards 
 hydrogen peroxide, causing the evolution of oxygen to take place, the com- 
 pound undergoing rapid alternate reduction and oxidation, in the same way 
 as silver oxide behaves with hydrogen peroxide. The green colour of the 
 solution, curiously enough, is indistinguishable by the eye from the green of 
 nickel salts. Nickel is incapable of giving this reaction, hence it constitutes a 
 ready means of detecting even traces of cobalt in the presence of nickel. 
 
60 Qualitative Analysis. 
 
 The Cyanides. 
 
 The metals of Group 1 1 IB. all form compounds with cyanogen, 
 some of which are of importance in analysis. The addition of 
 potassium cyanide to solutions of either a manganous, zinc, nickel, 
 or cobaltous salt results in the precipitation of the cyanide, of the 
 general formula RCy 2 . In each case, also, the precipitated cyanide 
 readily dissolves in excess of the KCy, producing double cyanides. 
 In the case of Mn, Zn, and Ni, these have the general formula 
 2KCy,RCy 2 , or K 2 RCy 4 , while the cobalt double cyanide has the 
 formula 4KCy,CoCy 2 (or K 4 CoCy 6 , corresponding to potassium 
 ferrocyanide, K 4 FeCy 6 ). 
 
 From the solutions 6f these double cyanides, (NH 4 ) 2 S is incapable 
 of precipitating the metals as sulphides. The addition of dilute 
 HC1 or H 2 SO 4 to these solutions causes the reprecipitation of the 
 cyanide 
 
 2KCy,NiCy 2 + 2HC1 = 2KC1 + 2HCy 4- NiCy, 
 
 In the case of the Mn, Zn, and Ni compounds, boiling with dilute 
 acid completely decomposes the metallic cyanide ; thus 
 
 2KCy,NiCy 2 -f 4HC1 = 2KC1 -I- 4HCy + NiCL. 
 
 The chief interest attaching to these compounds in this connection, 
 is due to the difference between the behaviour of cobalt and of 
 nickel towards cyanogen. The double cyanide of potassium and 
 cobalt, 4KCy,CoCy 2 or K 4 CoCy 6 (like the corresponding iron 
 compound) exhibits a great readiness to undergo oxidation, and 
 to pass from the condition of cobaltocyanide to cobalticyanide ; 
 thus 
 
 2K 4 CoCy 6 + O + H 2 O = -KHO + 2K 3 CoCy 6 
 
 Nickel forms no compounds corresponding to the cobalticyanides. 
 When the nickel double cyanide is oxidised (e.g. by chlorine, 
 bromine, or hypochlorites), the nickel is -converted into the black 
 hydrated sesquioxide and precipitated ; thus 
 
 2K 2 NiCy 4 + NaCIO + 5H 2 O = Ni.,O 3 ,3H.,O + NaCl + 4KCy 
 
 + 4HCy 
 
 While under the same oxidising treatment, the cobalt double 
 cyanide is converted into potassium cobalticyanide, which remains 
 in solution, and can therefore be separated by filtration ; thus 
 
 2K 4 CoCy (! 4- NaCIO 4- H 2 O = NaCl + 2KHO + 2K 3 CoCy 6 
 
Separation of the Metals of Group I I IB. 61 
 
 The cobaltocyanides are so readily oxidised, that by merely boiling 
 the aqueous solution the change is effected ; thus 
 
 K.CoCy, + H 2 = KHO + K 8 CoCy G + H 
 
 Owing to this reaction, the double cobalt cyanide does not behave, 
 when boiled with dilute acid, in the same manner as the nickel 
 double cyanide (see above). We may suppose that the cobalt 
 cyanide first formed, in the presence of free HCy (liberated partly 
 from the double cyanide and partly from the excess of KCy 
 present) and water, is converted into cobalticyanic acid in the 
 following manner : 
 
 (i) 4KCy,CoCy 2 + 4HC1 = 4KC1 + 4HCy + CoCy 2 
 (2) CoCy, + 4 HCy + H 2 O - H 3 CoCy 6 + H 
 
 It follows from this that the addition of dilute HC1 to a solution of 
 a cobalticyanide does not result in the precipitation of a cyanide, 
 but only in the liberation of cobalticyanic acid. 
 
 The metal cobalt in cobalticyanides, like iron in ferricyanides, 
 is not precipitated by the ordinary reagents for these metals. 
 
 SEPARATION OF THE METALS OF GROUP 1 1 IB. 
 
 The separation of the four metals of this group is based upon 
 
 1. The solubility of Zn(HO) 2 in caustic alkalies. 
 
 2. The solubility of MnS in acetic acid. 
 
 3. The different behaviour of the double cyanides of Ni and Co, 
 when subjected to oxidising agents. 
 
 In the course of a systematic analysis, the metals of Group 1 1 IK. 
 are looked for in the solution after Group II I A. has been separated 
 by means of NH 4 HO in the presence of NH 4 C1. H 2 S is passed 
 through this ammoniacal solution, whereby the sulphides of Mn, 
 Zn, Ni, and Co are precipitated. The mixture is gently warmed 
 and filtered.* 
 
 * Even when the exercise is confined only to Group IIlB., the student is 
 advised to go through the step of precipitating the group as sulphides in the 
 ammoniacal solution prepared as described on p. 48. 
 
62 
 
 Qualitative A nalysis. 
 
 The precipitate, consisting of MnS, ZnS, NiS, and CoS, is dissolved in 
 hot dilute HC1, with the aid of a few particles of KC1O 3 .* The solution 
 is boiled until it no longer smells of chlorine, and NaHO added in 
 excess. It is again boiled (to ensure the solution of all the Zn(HO) 2 )j 
 and filtered after cooling (hot NaHO will attack the filter-paper). 
 
 The solution 
 
 contains so- 
 dium zincate, 
 ZnNa 2 O 2 . 
 
 Pass H 2 S 
 through the 
 liquid (or add 
 H 2 S water), 
 when a white 
 precipitate of 
 ZnS is pro- 
 duced. 
 
 The precipitate, consisting of Mn(HO) 2 , Ni(HO) 2 , 
 and Co(HO) 2 , is washed to remove the soluble zinc- 
 salt still adhering to it, and then dissolved in the 
 smallest quantity of warm HC1. The solution is 
 nearly neutralised with NH 4 HO, a considerable 
 quantity of ammonium acetate added, and H 2 S passed 
 through the mixture until precipitation is complete. 
 
 The solution The precipitate (NiS and CoS). Test 
 contains; a small portion with a borax bead, 
 manganous ( Blue colour indicates cobalt. Dis- 
 acetate, solve the precipitate in the least 
 Mn(C 2 H 3 O 2 ) 2 . quantity of aqua regia ; boil off the 
 excess of add, and nearly neutralise 
 Precipitate with Na,CO 3 (avoid dilution). Add 
 MnCO 3 by add- KCy (freshly made solution) until the 
 ing Na 2 CO 3 . precipitated cyanides are just redis- 
 Mlter ; wash so i ve d, add NaHO in considerable 
 dissolve 'in quantity, and then bromine waterf until 
 HC^andcon" tne co ^ our f tne bromine persists, 
 firm' Mn by Filter - 
 precipitation 
 
 of the pinkish 
 MnS, with The solution con- The precipitate 
 
 (NH 4 ) 2 S after tains sodium consists of the 
 NH 4 C1 and cobalticyanide, black hydratedses- 
 NH 4 HO. Na 3 CoCy 8 . quioxide of nickel, 
 
 NuO 3 ,3H 2 O. 
 The formation The cobak may be 
 
 confirmed by the ; Confirm (i)by borax 
 m a n ga n at e borax bead( if the bead . ^ wash pre . 
 
 may also test was not made cipitate, dissolve in 
 used as a con- or was not satis . | ^ and obtain 
 firmatory test. f actO r y , with the ! the characteristic- 
 sulphides, reaction witli 
 (XH 4 ) 2 S. 
 
 * Should the precipitate not be black, only MnS and ZnS can be presarik; 
 and as these are easily soluble in HC1, the KC1O 3 should in this case noe 
 added. 
 
 f Bromine is more soluble in water containing KBr in solution than in water 
 alone. By the use of such a solution, therefore, unnecessary dilution is avoided. 
 
 
"V^ '$K* 
 
 _ 
 
 *1 f * 
 
 , a~A. &. 
 
 
 
 
 THE PHOSPHATES. 
 
 THE salts of phosphoric acid, H 3 PO 4 , are all of them insoluble in 
 water except those of the alkali metals and ammonium. Lithium 
 phosphate is, however, only difficultly soluble. The addition, there- 
 fore, of a solution of one of these soluble phosphates to a solution 
 of a metallic salt results in the precipitation of the phosphate of the 
 metal contained in the salt ; thus, using hydrogen disodium 
 phosphate as being the most convenient reagent 
 
 HNa.,PO 4 + 3AgNO s = Ag 3 PO 4 + 2NaNO s + HNO ;! 
 
 HNa,P0 4 + BaClo = HBaPO 4 * + 2NaCl 
 2HNa2PO 4 -f 3CaCL - Ca 3 (PO 4 ) 2 + 4NaCl + 2HC1 
 
 The most delicate reaction for the detection of a phosphate is 
 the formation of a canary-yellow crystalline precipitate o>i ammonium 
 phosphomolybdate, upon the addition of a nitric acid solution of 
 ammonium molybdate, (NH 4 ) 2 MoO 4 , to a solution of a phosphate. A 
 few drops of the solution of a phosphate are taken, and a very 
 large excess of the nitric acid solution of ammonium molybdate is 
 added, and the mixture slightly warmed when the yellow precipitate 
 crystallises out. It is insoluble in water or acid provided a con- 
 siderable quantity of ammonium molybdate is present. The 
 precipitate is believed to have the composition i2MoO 3 ,(NH 4 ) 3 PO 4 . 
 It is decomposed by ammonia into (NH 4 ) 2 MoO 4 and (NH 4 ) 3 PO 4 . 
 
 When the phosphates, insoluble in water, are acted upon by 
 acids, they either dissolve, or, if the acid employed is capable of 
 forming an insoluble compound with the metal, they are converted 
 into this compound, and phosphoric acid is set free. Thus, calcium 
 phosphate is decomposed by sulphuric acid, insoluble calcium 
 sulphate being formed, while the liberated phosphoric acid passes 
 into solution 
 
 Ca s (P0 4 ) 2 + 3H. 2 S0 4 = 3 CaS0 4 + 2 H 3 PO 4 
 
 * It depends upon the conditions of precipitation whether the so-called acid 
 phosphate, HBaPO 4 , or the normal salt, Ba 3 (PO 4 ) 2 , is produced. 
 
64 Qualitative Analysis. 
 
 In those cases where the phosphate is entirely dissolved by the 
 acid, it is reprecipitated unaltered by the addition of ammonia, 
 ammonium sulphide,* caustic alkalies, alkaline carbonates, or in 
 some cases by barium carbonate. For example, the phosphates 
 of Ni, Co, Mn, Zn, Ba, Sr, Ca, and Mg are all soluble in HC1 ; and 
 if ammonia be added to the hydrochloric acid solution of any one 
 of these, it at once causes the reprecipitation of the phosphate. 
 The significance of this fact, and its bearing upon the method of 
 separation of the metals of Group III., will be obvious. When 
 NH 4 C1 and NH 4 HO are used to separate Group I HA. from 
 Group 1 1 IB. by precipitating the hydroxides of Al, Cr, and Fe, 
 they will at the same time throw down the phosphates of all or 
 any of the above-named metals, if they happen to be present as 
 phosphates in the solution under examination. Hence it is 
 necessary to adopt some method for withdrawing the phosphoric 
 acid which is in combination with these metals, and of converting 
 them all into such salts as chlorides before it is possible to attempt 
 their separation on the basis of the plans laid down on pp. 49 
 and 62. 
 
 It will be well, therefore, at this pointf^ enter somewhat fully 
 upon a study of such reactions with phosfBRes as are made use of 
 in analysis. 
 
 The Solution of Phosphates by Acids. I . By Hydrochloric 
 Acid. The phosphates of the metals of Groups III., IV., and V., 
 which are insoluble in water, are dissolved by hydrochloric acid. 
 When such a phosphate is dissolved in this acid, the addition to 
 the solution of any alkali causes the reprecipitation of the phosphate 
 unchanged. This fact is sometimes explained by supposing that 
 the act of solution in these cases, is of the same simple order as the 
 solution of salt in water ; that the phosphate dissolves unchanged, 
 being present in the liquid as a phosphate; and that, therefore, its 
 reprecipitation by an alkali is simply owing to the removal, by 
 neutralisation, of the solvent acid. 
 
 As a matter of fact, however, the process is more complex, for 
 in reality double decomposition between the acid and the phosphate 
 takes place exactly in the same manner as between the calcium 
 phosphate and sulphuric acid given above, except that the metallic 
 salt formed by the union of the metal with the solvent acid is 
 soluble in water. Thus, in the case of the solution of barium 
 phosphate in hydrochloric acid^ 
 
 HBaPO 4 + 2HC1 =. BaCl 2 + H 3 PO 4 
 
 barium chloride and phosphoric acid are formed ; and both being 
 * Except those of Fe, Mn, Ni, Co, Zn. See p. 65. 
 
The Phosphates. 65 
 
 soluble in water, complete solution results from the action.* 
 BaCL and H 3 PO 4 are incapable of reacting upon each other, and 
 therefore exist together in the solution.! 
 
 Now, when an alkali is added to this solution, it first unites with 
 the phosphoric acid, forming a soluble phosphate of the alkali ; 
 this in its turn, by double decomposition with the barium chloride, 
 forms barium phosphate and a chloride of the alkali. The two 
 stages in the reaction may be thus expressed : 
 
 :BaCl 4- H PO 4- sN^HO - 2 2 4 2 
 
 2 . . 3 4 - : ' -\(2)2NaCl + HBaP0 4 + 2 H 2 O 
 
 In order, therefore, to formulate the various reactions which phos- 
 phates in acid solutions undergo, the solution of the phosphate 
 will be expressed in the equations as a mixture of phosphoric acid 
 and the metallic salt of the solvent acid. In this way the mechanism 
 of the reactions will be rendered easy of understanding. Two or 
 three examples of the precipitation of phosphates from hydrochloric 
 acid solutions by various alkaline reagents may be given. 
 
 (a) Precipitation of calcium phosphate from HC1 solution by 
 NH 4 HO, giving only the final result 
 
 i 3 CaCl,+ 2H 8 P0 4 : + 6NH 4 HO = 6NH 4 C1 + Ca 3 (PO 4 ) 2 + 6H 2 O 
 
 () Precipitation of aluminium phosphate by (NH 4 ) 2 S 
 J2A1C1 3 + 2H 3 P0 4 j + 3(NH 4 ) 2 S = 6NH 4 C1 + 2 A1PO 4 | + 3 H 2 S 
 
 In the case of the metals Fe, Mn, Zn, Ni, and Co (metals which are 
 capable of forming insoluble sulphides), the phosphates are not 
 reprecipitated by (NH 4 ) 2 S ; but the metals are thrown down as 
 sulphides, while the phosphoric acid unites with the ammonia, 
 
 * It is quite easy to prove, in this particular case, that when barium phosphate 
 is dissolved in hydrochloric acid, the solution contains barium chloride and 
 phosphoric acid. This may be done by gently evaporating the solution down 
 upon a steam-bath until crystals begin to separate out. If these are drained, 
 and washed a few times with strong hydrochloric acid (in which BaCL is ne?rrly 
 insoluble), they can then be dissolved in* water. The solution so obtained will 
 be found to give no precipitate with ammonia, therefore showing that it 
 contains no barium phosphate ; but it will give all the ordinary reactions for 
 barium. 
 
 f Hence it follows that the addition of phosphoric acid to such a metallic 
 solution as barium chloride gives no precipitate of barium phosphate. 
 
 % The simpler formulae for the phosphates of Al, Cr, and Fe will be used, 
 instead of A1 2 (PO 4 ) 2 , etc., etc. 
 
 F 
 
66 Qualitative Analysis. 
 
 and therefore remains in the solution: thus, taking iron as an 
 example 
 
 :2FeCl 3 + 2H 3 PO 4 : 
 
 : ..................... : + 2H 2 S + S 
 
 (c) Precipitation of ferric phosphate by BaCO 3 
 
 2FeCl 3 + 2H 3 PO 4 : + 3BaCO 3 = 3BaCl 2 + 2FePO 4 + 3 CO., 
 
 : ...................... : +3H 2 
 
 In the case of solutions of calcium or magnesium phosphates, 
 the addition of BaCO 3 results in the precipitation, not of the calcium 
 or magnesium phosphate, but of barium phosphate ; thus 
 
 (d) !3MgCl 2 +2H 3 P0 4 : + 3BaC0 3 = sMgCU + Ba 3 (PO 4 ) 2 
 
 II. By Acetic Acid, H(C 2 H 3 O 2 ). Of the phosphates of the 
 metals of Groups III., IV., and V., those of aluminium and iron 
 are insoluble in acetic acid, while chromium phosphate is only 
 dissolved with difficulty. The phosphates of the remaining metals 
 of these groups are readily soluble in this acid. 
 
 When a phosphate is dissolved in acetic acid, we may, for the 
 same reasons as apply in the case of hydrochloric acid, consider 
 the action as being strictly a chemical one, and may regard the 
 solution as containing the metallic acetate and phosphoric acid ; 
 thus, in the case of calcium phosphate 
 
 Ca 3 (P0 4 ) 2 + 6H(C 2 H 3 2 ) = 3 Ca(C 2 H 3 O 2 ) 2 + 2 H 3 PO 4 * 
 Phosphoric acid, therefore, is incapable of giving a precipitate 
 with a solution of calcium acetate (or the acetate of any metal 
 whose phosphate is soluble in acetic acid). It can, however, pre- 
 cipitate the phosphate from the acetate of a metal whose phosphate 
 is insoluble in acetic acid. Thus, if phosphoric acid is added to a 
 solution of aluminium acetate, aluminium phosphate is thrown 
 down, and acetic acid formed 
 
 A1(C 2 H 3 2 ) 8 + H 3 P0 4 = A1P0 4 + 3 H(C 2 H 3 2 ) 
 If to a solution of a phosphate (e.g. calcium phosphate) in acetic 
 acid, ferric chloride (or aluminium chloride) be added, by double 
 decomposition acetic acid is formed in the solution, and ferric phos- 
 phate is therefore precipitated ; thus 
 
 l3Ca(C 2 H 3 2 ) 2 +2H 3 PO 4 ': + 2FeCl 3 = 2FePO 4 + 3CaCl 2 
 
 + 6H(C 2 H 3 2 ) 
 
 * In this solution the presence of the phosphoric acid does not prevent the 
 precipitation of the calcium as calcium oxalate by the addition of ammonium 
 oxalate. 
 
The Phosphates. 67 
 
 Now, although phosphoric acid is unable to precipitate ferric 
 phosphate when added to ferric chloride (ferric phosphate being 
 soluble in hydrochloric acid), the presence of a soluble acetate, such 
 as sodium acetate, would bring about the same conditions as exist 
 in the above reaction, and therefore again the phosphate of iron will 
 be thrown down 
 
 3Na(C 2 H 3 2 ) + iH 3 P0 4 + FeCl 3 ! = FePO 4 + 3NaCl 
 : : + 3 H(C 2 H 3 2 ) 
 
 The Removal of Phosphoric Acid. The phosphoric acid 
 may be separated from a phosphate in two ways : namely, either 
 by causing the metal to form some insoluble compound with a 
 reagent, or by making the phosphoric acid unite with a reagent to 
 produce an insoluble compound. For example, the phosphoric 
 acid may be withdrawn from the phosphates of iron, manganese, 
 zinc, cobalt, and nickel, by means of ammonium sulphide, as ex- 
 plained on p. 66. In these cases, the metals, in the form of 
 insoluble sulphides, are precipitated, while the phosphoric acid, in 
 combination with ammonia, remains in solution. The disadvantage 
 of this method lies in the fact that the phosphate of ammonia left 
 in the liquid, is able to undergo double decomposition with any 
 other metallic salts which might be present, and precipitate them 
 as phosphates ; therefore the final result would only be that of sub- 
 stituting one insoluble phosphate for another. The following 
 examples will make this clear : 
 
 Suppose a mixture under examination to consist of alumina and 
 phosphate of iron ; on dissolving in hydrochloric acid, we should 
 get the following reaction : 
 
 A1 2 3 + 2FePO 4 + I2HC1 = 2A1C1 3 + 2FeCl 3 + 2H 3 PO 4 + sH 2 O 
 
 On adding ammonium sulphide to the solution, the iron is precipi- 
 tated as FeS and the aluminium as phosphate ; thus 
 
 2A1C1 3 + i2FeCl 3 +2H 3 PO 4 i + 6(NH 4 ) 2 S = 2FeS + 2A1PO 4 
 
 : : + i2NH 4 Cl + sH 2 S + S 
 
 Or again, if salts of Ba, Sr, Ca, or Mg (other than phosphates) are 
 present with a phosphate which is decomposed by ammonium 
 sulphide, then the ammonium phosphate which is formed would 
 react upon such salts and precipitate the metals as phosphates. 
 Thus, suppose a mixture of zinc phosphate and calcium carbonate 
 is being examined ; as before, it is dissolved in hydrochloric acid, 
 and then, on the addition of ammonium sulphide, a precipitate is 
 
68 Qualitative Analysis. 
 
 obtained consisting of a mixture of zinc sulphide and calcium 
 phosphate 
 
 Zn 3 fPO 4 ) 2 + 3CaCO 3 + I2HC1 = sZnCL, + 2H 3 PO 4 +3CaCL> 
 
 + 3 CO 2 + 3 H 2 
 
 ; 3 ZnCl 2 + 2H 3 P0 4 ; + 3CaCl 2 + 6(NH 4 ) 2 S = 3ZnS + Ca 3 (PO 4 ) 2 
 
 + J2NH 4 C1 + 3H 2 S + S 
 
 In the ordinary course of analysis, on the addition of NH 4 C1 
 and NH 4 HO for the precipitation of Group IIlA., if phosphates are 
 present they are also precipitated along with the hydroxides of Al, 
 Cr, and Fe. If, therefore, ammonium sulphide be added, any 
 phosphates of Fe, Zn, Mn, Ni, and Co which are present will be 
 converted into sulphides, and the ammonium phosphate which is 
 formed will react upon any compounds of Ba, Sr, Ca, and Mg which 
 might be present, not as phosphates, and convert them into phos- 
 phates. In analysis, therefore, the second method for removing 
 phosphoric acid is had recourse to, namely, that of employing a 
 reagent which will form an insoluble compound with the add, and 
 so throw it out of solution. 
 
 The reagent employed is ferric chloride, and the separation is 
 based upon the fact already explained, viz. that when this salt is 
 added to an acetic acid solution of a phosphate, it precipitates the 
 phosphoric acid as ferric phosphate, leaving the metal, which was 
 originally united to the phosphoric acid, in the solution as a 
 chloride. The separation is made in the following manner : 
 
 SEPARATION OF PHOSPHORIC ACID. 
 
 The precipitate produced by NH 4 HO in presence of NH 4 C1 
 (which may contain the hydroxides of Al, Cr, and Fe, as well as all 
 the phosphates), is dissolved in a little warm dilute HC1, and the 
 solution made nearly neutral by the addition of Na 2 CO 3 .* A mixture 
 of sodium acetate and acetic acid \ is then added, and the solution 
 boiled and filtered.:}: 
 
 * This is in order to remove the excess of HC1 used in dissolving the pre- 
 cipitate. This acid would, in any case, be neutralised by double decomposi- 
 tion with the sodium acetate, when that reagent is added ; but it is better to 
 get rid of it previously by means of sodium carbonate, so that the sodium 
 aceta|e may be utilised entirely in bringing about the reactions explained on 
 p. 66, and also in the footnote below. 
 
 t See list of reagents (Appendix). 
 
 % If the whole of the phosphoric acid contained in the substance under- 
 going analysis was originally present in combination with Al, Cr, or Fe, then 
 by this reaction the whole of the phosphoric acid will have been thrown out of 
 
Separation of Phosphoric Acid. 
 
 The precipitate con- 
 tains phosphates of 
 Al, Cr, or Fe. 
 
 This precipitate may be 
 treated with Na 2 O 2 
 exactly as the hy- 
 droxides, p. 49. A1PO 4 
 dissolves in the NaHO, 
 and CrPO 4 is oxidised 
 intoNa 2 CrO 4 . On filter- 
 ing, the iron is left on 
 the filter. The Al may 
 be detected by neutra- 
 lisation and reprecipita- 
 tion with NH 4 HO. The 
 chromium by BaCU 
 in acetic acid solution. 
 (Barium phosphate is 
 not precipitated in 
 presence of acetic acid. ) 
 
 The filtrate, FeCl 3 is added drop by drop * until 
 precipitation is complete, at which point the 
 liquid will begin to assume a distinct reddish 
 colour (due to the formation of soluble ferric 
 acetate).f The mixture is gently boiled for a 
 few minutes (whereby the ferric acetate is con- 
 verted into an insoluble basic acetate), and then 
 filtered. 
 
 The precipitate, containing the whole of the 
 phosphoric acid as ferric phosphate and 
 more or less basic ferric acetate, is thrown 
 away. The solution now contains the metals 
 of Groups IIlB. and IV., as well as Mg, in 
 the form of chlorides. Their separation is 
 made by the usual methods. 
 
 solution. Moreover, even if some of the phosphoric acid had been originally 
 present in combination with other metals, it might happen that there was also 
 present enough aluminium or iron, not as phosphate, to combine with all the 
 phosphoric acid of these other phosphates, and so to cause the complete 
 removal of this acid. The following example will make this clear : Suppose 
 a mixture of aluminium phosphate, ferric oxide, and calcium phosphate is to be 
 examined. When this has been dissolved in HC1, we may represent the action 
 of sodium acetate in the following manner : 
 
 i A1C1 3 + H 3 P0 4 | + 
 
 2FeCl 3 + [ 3 CaCl 2 + 2H 3 PO 4 
 
 : + 9 Na(C 2 H 3 2 ) 
 
 f Vcl + 
 
 j --:;;: 
 
 2 FeP0 4 V 
 
 } precipitated 
 
 3 NaCl 
 3 H(C 2 H 3 2 ) 
 
 6NaCl -" , 
 6H(C 2 H 3 2 H 
 
 I in solution 
 
 If more than enough ferric chloride is present than is required to take the 
 whole of the phosphoric acid from the calcium phosphate (or other similar 
 phosphates), then all the phosphoric acid is thrown down, and the excess of 
 the FeCl 3 passes into the solution. But if, on the other hand, there is an 
 excess of calcium phosphate, it passes into solution owing to the free acetic 
 acid present. If no compounds of Al, Cr, or Fe are present, either as phos- 
 phates or otherwise, no precipitate will be produced by the sodium acetate. 
 
 * If the first drops of FeCl 3 produce no precipitate, it proves that all the 
 phosphoric acid has already been thrown down (see previous note). The 
 metals are all present, therefore, as chlorides, and their separation may be 
 made by the ordinary methods. 
 
 t Ferric acetate is formed by the action of the excess of ferric chloride upon 
 the sodium acetate present ; thus 
 
 FeCl 3 + 3 Na(C 2 H 3 2 ) = Fe(C 2 H 3 O 2 ) 3 + 3 NaCl 
 
 It is very important to avoid the use of any unnecessary excess of FeCl 3 , for the 
 precipitated ferric phosphate is soluble both in ferric chloride and in ferric 
 
Qualitative A nalysis. 
 
 APPENDIX TO CHAPTER VII. 
 THE RARE METALS OF GROUP III.* 
 
 These elements may for convenience be arranged in five groups, 
 based upon the composition of the precipitates which are thrown 
 down by the group-reagents : 
 
 (I) 
 
 (2) 
 
 (3) 
 
 (4) 
 
 (5) 
 
 Beryllium 
 
 Scandium 
 
 Zirconium 
 
 Titanium 
 
 Uranium 
 
 
 Yttrium 
 
 Thorium 
 
 Tantalum 
 
 Indium 
 
 
 Ytterbium 
 
 
 Niobium 
 
 (Thallium) t 
 
 
 Cerium 
 
 
 
 Vanadium \ 
 
 
 Lanthanum 
 
 
 
 
 Those Of the first three divisions are precipitated in the form of 
 hydroxides, which are basic in their character. Those of the fourth 
 division are thrown down in the form of hydrated oxides of a more 
 or less acidic character, while the remaining metals come down as 
 sulphides. The composition of the various compounds is as 
 follows : 
 
 (I) 
 
 Be(HO) 2 
 
 (2) 
 
 Sc(HO) 3 
 Y(HO) 3 
 Yb(HO) 3 
 Ce(HO) 3 
 La(HO) 3 
 
 (3) 
 Zr(HO) 4 
 Th(HO)< 
 
 (4) 
 H 2 TiO 3 
 H 3 TaO 4 
 H 3 NbO, 
 
 (5) 
 (U0 2 )S 
 InS 
 (T1 2 S) 2 
 
 Beryllium. Named from beryl, the chief source of the ele- 
 ment. The sulphate, BeSO 4 , is the most readily obtained salt. 
 
 The reactions of this element resemble very closely those of 
 aluminium so much so, that at one time the two elements were 
 
 acetate. The transformation of soluble ferric acetate into the insoluble basic 
 acetate by boiling, is expressed as follows : 
 
 3 Fe(C 2 H 3 2 ) 3 + 3 H 2 = Fe(C 2 H 3 O 2 ) 3 ,Fe 2 O 3 + 6H(C 2 H 3 O 2 ) 
 
 In cases where very small quantities of a phosphate and a great excess of iron 
 are present in the substance under examination, it is necessary to reduce the 
 iron to the " ferrous " state by means of sulphurous acid. Ferrous salts do not 
 dissolve ferric phosphate. 
 
 * A systematic study of the reactions of the rare metals lies entirely outside 
 the scope of this book. In many cases it is almost impossible to purchase the 
 compounds in anything approaching to a condition of purity, while with most 
 of the metals the cost of the salts practically prohibits such a course (see note, 
 p. 26). These elements, however, are not all equally " rare," and therefore in 
 this section some of the characteristic reactions of a few of the most commonly 
 occurring of these metals are given. 
 
 f Thallium also appears in Group I. (see p. 16), since the chloride is 
 precipitated by hydrochloric acid from thalloits solutions; and, like lead 
 chloride, is slightly soluble in water. 
 
 % See footnote on p. 18. 
 
The Rare Metals of Group III 71 
 
 regarded as belonging to the same natural family. Thus, ammonia, 
 potash, and soda produce, with solutions of either metal, a white 
 flocculent precipitate of the respective hydroxides, A1,(HO) C and 
 Be(HO) 2 . 
 
 Both precipitates are soluble in excess of the fixed alkalies, but 
 the beryllium compound is reprecipitated if the diluted solution is 
 boiled. 
 
 Beryllium hydroxide is decomposed by boiling with ammonium 
 chloride, ammonia being evolved and beryllium chloride passing 
 into solution ; thus 
 
 Be(HO) 2 + 2NH 4 C1 = BeCl 2 + 2NH 3 + 2H 2 O 
 
 Aluminium hydroxide undergoes no change when similarly treated, 
 hence the metals may be separated by this reaction. 
 
 The carbonates of the alkalies give precipitates of beryllium 
 carbonate, or basic carbonate ; readily soluble in excess of ammo- 
 nium carbonate, giving a double carbonate. 
 
 The precipitate is soluble, but far less readily, in the carbonates 
 of the fixed alkalies. Beryllium is therefore readily separated from 
 aluminium by adding ammonium carbonate and warming the 
 solution. Aluminium hydroxide is precipitated, while the beryllium 
 goes into solution as the double carbonate of ammonium and 
 beryllium. On filtering off the aluminium hydroxide, the beryllium 
 hydroxide may be precipitated by ammonia, after first neutralising 
 with hydrochloric acid. 
 
 The salts of beryllium possess a characteristic sweet taste, 
 hence the name ghtcinum, which was formerly applied to the 
 element. 
 
 Zirconium. The oxide of this element, along with others of 
 the so-called " rare earths " (but more especially Zirconia), has the 
 property of remaining unchanged for a long time when heated to 
 incandescence, and of emitting a bright white light when so heated. 
 It is on this account used, with others of the rare earths, in the 
 construction of the " mantles " of the incandescent gas-burners 
 now so common. 
 
 Alkaline hydroxides, as well as ammonium sulphide (group- 
 reagent), give a white precipitate of zirconium hydroxide, Zr(HO) 4 , 
 resembling aluminium hydroxide in appearance. It is distinguished 
 from the latter, in that it is insoluble in excess of potassium or 
 sodium hydroxide. 
 
 Alkaline carbonates give a white precipitate consisting of a 
 basic carbonate, which is soluble in excess, especially of ammonium 
 carbonate. 
 
 Potassium sulphate (a concentrated boiling solution) gives a 
 precipitate of a double sulphate, which, when thrown down from 
 the hot solution, is scarcely soluble in hydrochloric acid (thorium 
 gives a similar precipitate under the same conditions, but the 
 thorium compound is soluble in hydrochloric acid). In dilute 
 solutions the precipitate only appears after standing for some hours. 
 Sodium thiosulphate gives a precipitate of zirconium thiosulphate, 
 
72 Qualitative Analysis. 
 
 which on boiling is complete even in very dilute solutions (thorium 
 behaves similarly). Oxalic acid gives a white precipitate of 
 zirconium oxalate, soluble in ammonium oxalate. (Thorium oxalate, 
 similarly precipitated, is insoluble in ammonium oxalate.) 
 
 Hydrogen peroxide, added to a slightly acid solution of a 
 zirconium salt, gives a white precipitate, believed to be either ZrO 3 
 or Zr 2 O 5 . (Niobium and titanium do not give a precipitate.) 
 
 By means of these three reactions zirconium can be separated 
 from the other " rare earths " of this group. 
 
 The hydrated oxides, which are precipitated by ammonia, are 
 first dissolved in hydrochloric acid, and the solution nearly 
 neutralised with sodium carbonate. Sodium thiosulphate is then 
 added, and the mixture boiled. The precipitate may consist of the 
 thiosulphates of zirconium and thorium, together with titanic acid. 
 The precipitate is treated with boiling hydrochloric acid, which 
 dissolves the two thiosulphates, and possibly a little titanic acid. 
 Excess of ammonium oxalate is added to the solution, which at 
 first precipitates oxalates of zirconium and thorium, but redissolves 
 the zirconium oxalate. The insoluble thorium oxalate is removed 
 by filtration, and any titanium in the solution is precipitated by 
 means of ammonium carbonate, added in excess in order to re- 
 dissolve the zirconium basic salt, which is first thrown down. 
 Any permanent precipitate is filtered off, and the filtrate con- 
 centrated by gentle evaporation. A boiling strong solution of 
 potassium sulphate is then added, which precipitates the double 
 sulphate of potassium and zirconium.* 
 
 Titanium. This element is met with in small quantities in 
 many specimens of iron ores, clays, and igneous rocks. The 
 minerals special to the element, rutile^ anatase (TiO 2 ) ; sphetie 
 (silicate and titanate of calcium), are rare substances titaniferous 
 iron (ferrous titanate) is less rare. 
 
 Titanium oxide, TiO 2 , being insoluble in hydrochloric and 
 nitric acid, is found in the insoluble residue after treatment with 
 acids ; titanates, on the other hand, are dissolved by hydrochloric 
 acid, but on boiling the solution white titanic acid, H 2 TiO. 3 , is 
 precipitated. 
 
 Titanium oxide (in the absence of other metals which colour a 
 bead of microcosmic salt), when heated in a bead of microcosmic 
 salt in the inner blowpipe flame, imparts to the bead a colour 
 which is yellowish when hot, but which becomes violet as the bead 
 
 * This separation, and the reactions for zirconium, may be made by dis- 
 solving up a couple of the incandescent gas "mantles ; " either new ones, or 
 those which have become worn out by use. In the former case they should 
 first be set fire to, in order to burn off the organic matter present. The 
 mantles are boiled in a test-tube with a little strong sulphuric acid, and the 
 liquid when cold diluted with water. The insoluble residue (consisting of 
 the main portion of the material) is filtered off, and the hydrated oxides of the 
 rare earths, along with any alumina, are precipitated by the addition of 
 ammonia. The precipitate is then washed, and dissolved in a small quantity 
 of hydrochloric acid. 
 
The Rare Metals of Group III. 73 
 
 cools. The colour is more readily obtained by the aid of some 
 additional reducing agent besides the reducing flame ; thus if the 
 salt be heated on charcoal instead of a platinum wire, or if a trace 
 of zinc be added, the result is more quickly obtained. In the 
 presence of small quantities of iron the bead appears brown-red. 
 
 Titanium dioxide is separated from silicon dioxide (or silicates) 
 by the action of hydrofluoric acid (sulphuric acid being present to 
 prevent the volatilisation of titanium, as fluoride). Titanium is 
 obtained in solution (and separated from silica, and also from 
 compounds of tantalum and niobium) by fusion with hydrogen 
 potassium sulphate. The " melt " is extracted with cold water, 
 when the titanium passes into solution as a sulphate. 
 
 When titanium oxide is fused with potassium carbonate, 
 potassium titanate, K 2 TiO 3 , is formed, which, being insoluble in 
 water, may be separated from any alkaline silicate by extracting 
 the melt in cold water. 
 
 If the residue of potassium titanate be then treated with cold 
 dilute hydrochloric acid, it dissolves, the solution containing titanic 
 acid, Ti(HO) 4 or TiO 2 ,2H 2 O. If this solution be heated, the 
 titanic acid is rendered insoluble, and is therefore precipitated, 
 the precipitate being the hydrated oxide H 2 TiO 3 , or TiO(HO) 2 , 
 or TiO 2 ,H 2 O. 
 
 If to the cold solution of titanic acid there be added ammonia, 
 or the hydroxide or carbonate of the alkalies, or ammonium 
 sulphide, a white precipitate is produced, consisting of the hydrated 
 compound Ti(HO) 4 , or TiO 2 ,2H 2 O, which redissolves readily in 
 dilute hydrochloric acid or sulphuric acid. Moderate rise of 
 temperature at once converts this soluble compound into the 
 insoluble titanic acid, H 2 TiO 3 , or TiO 2 ,H 2 O. 
 
 Uranium. Like the element chromium, uranium forms 
 uran07/.r and uram^r salts, as well as urinates. 
 
 The uranous salts are derived from uranous oxide, UO 2 , in 
 which the element is tetravalent. Thus, uranous sulphate is 
 represented by the formula U(SO 4 )2. 
 
 These salts readily pass by oxidation into uranic compounds, 
 and are therefore powerful reducing agents. They are for the 
 most part green in colour. The uranic or uranyl salts are derived 
 from the oxide, UO 3 , or (UCX>)O, in which the element is hexa- 
 valent. They are regarded as containing the divalent radical 
 uranyl, (UO 2 ), which takes the place of a divalent metal; thus, 
 uranyl sulphate and nitrate are expressed by the formulas 
 (UO 2 )SO 4 ,3H 2 O and (UO 2 )(NO 3 ) 2 ,6H 2 O respectively. These 
 salts are mostly yellow in colour, and soluble in water. 
 
 The uranates are constituted like the dichromates (uranates 
 corresponding to normal chromates are not known). Thus, sodium 
 uranate, Na 2 U 2 O 7 , analogous to sodium dichromate, Na 2 Cr 2 O 7 . 
 
 Uranyl nitrate and acetate are the salts most commonly met 
 with, being used in the volumetric determination of phosphoric 
 acid. 
 
 Alkaline carbonates give with uranyl salts a yellow precipitate, 
 
74 Qualitative Analysis. 
 
 consisting of a double carbonate of uranium and the alkali. The 
 precipitate is soluble in excess of the reagent. 
 
 The uranium in this solution is not precipitated by ammonium 
 sulphide. If the solution be neutralised with acid, a uranate of 
 the alkali is thrown down. 
 
 With uranous salts, these reagents give a green precipitate, 
 also soluble in excess. 
 
 Caustic alkalies give with uranyl salts a yellow precipitate of 
 the uranate of the alkali, insoluble in excess of the reagent. 
 
 With uranous salts, these reagents give a chocolate-coloured 
 precipitate of uranous hydroxide, U(HO) 4 . 
 
 Ammonium sulphide gives with uranyl salts a brown precipitate 
 of uranyl sulphide, (UO 2 )S, insoluble in excess of the reagent, but 
 soluble in normal ammonium carbonate. 
 
 In acid solutions sulphuretted hydrogen reduces uranyl to 
 uranous compounds. 
 
 Potassium ferrocyanide gives with iiranyl salts a brown pre- 
 cipitate. This reagent is employed as an indicator in the volu- 
 metric estimation of phosphoric acid by means of uranyl salts. 
 
CHAPTER VIII. 
 
 THE METALS OF GROUP II. 
 
 THE metals of this group (as well as those of Group I.) are 
 characterised by their common property of forming sulphides in an 
 acid solution ; that is to say, their sulphides are insoluble in dilute 
 acids. By this property they are all .sharply separated from the 
 metals of Group III. The acid which has been found to be the 
 most convenient to have present, is hydrochloric acid ; and since, 
 on acidifying with this acid preparatory to the precipitation of the 
 sulphides, the chlorides of silver, lead, and mercicrous mercury * are 
 
 * Mercury forms two classes of soils,' mercuric and mercurous ; and the 
 relation in which they stand to each other is very interesting. In the com- 
 pounds of the first type, the metal is playing the part of an ordinary divalent 
 element, replacing two atoms of hydrogen in acids and giving such salts as 
 HgCl 2 , Hg(NO 3 ) 2 , HgSO 4 , HgS, etc. In the mercurous compounds the pro- 
 portion of mercury to the negative radical is twice as great as in the mercuric 
 salts, hence the composition of, say, mercurous chloride may be expressed 
 either by the formula Hg 2 Cl 2 or HgCl. Some chemists adopt the latter 
 formula, and, regarding the mercury in mercurous compounds as acting the 
 part of a monovalent element, express the various salts by such formulae as 
 HgXO 3 , Hg 2 SO 4 , etc. Others prefer to consider the mercurous salts as com- 
 pounds, in which the divalent radical, or double atom, (Hg 2 ), is substituted for 
 the single divalent atom, (Hg), and therefore express the compounds by such 
 formulas as Hg 2 Cl 2 , Hg 2 (NO 3 ) 2 , Hg 2 SO 4 . The density of the vapour yielded 
 by heating mercurous chloride is 117 '59, which, being half that demanded by 
 the formula Hg 2 Cl 2 , gave support to the view that HgCl was the correct 
 formula. But it has since been shown that the compound dissociates on 
 heating, into mercuric chloride, HgCl 2 , and Hg (mercury giving monatomic 
 molecules). In the absence of any conclusive evidence in favour of either view, 
 the formulae adopted for mercurous compounds in this book will be those which 
 represent them as containing the double atom (Hg 2 ). As the two classes of 
 compounds present great differences in their chemical reactions, behaving, 
 indeed, more like compounds of two different metals, one of which belongs to 
 Group I. and the other to Group II., there is at least some advantage in 
 employing a symbol for mercurous mercury (Hg 2 ), which at a glance dis- 
 tinguishes it from that used to denote mercuric mercury, Hg. All the salts of 
 mercury will therefore be formulated as salts of a dibasic metallic radical, 
 which in the mercuric compounds has the symbol Hg, and in the mercurous 
 salts (Hg 2 ) (with or without the bracket). Thus 
 
 Chlorides, HgCl 2 (Hg,)Cl 2 
 
 Sulphates, HgS0 4 (Hg 2 )SO 4 
 
76 Qualitative Analysis. 
 
 thrown down, these three metals are treated separately, and con- 
 stitute Group I. 
 
 The precipitation of these three metals as chlorides, however, 
 is only complete in the case of Ag and (Hg 2 ). Lead chloride is 
 soluble to some extent even in cold water, hence a portion of the 
 lead passes through into Group II. 
 
 The metals of Group II. are divided into two sections, namely 
 
 (1) Metals whose sulphides are insoluble in ammonium 
 sulphide 
 
 Mercury (mercuric), lead, bismuth, cadmium, copper. 
 
 (2) Metals whose sulphides dissolve in ammonium sulphide 
 
 Arsenic, antimony, tin, gold, platinum. 
 
 REACTIONS OF THE METALS OF GROUP II. DIVISION i. 
 Mercury, Hg. 
 
 DRY REACTiONS.-^When heated alone in a tube, many mercury 
 compounds (those with the halogens, for example) volatilise un- 
 changed, giving sublimates of the same compound. The iodide 
 (red) when heated forms a sublimate, consisting chiefly of the yellow 
 allotropic form of HgI 2 , which when cold changes to red if scratched 
 or rubbed. Some mercury compounds when heated decompose, 
 and metallic mercury volatilises and sublimes in the tube. 
 
 If a mercury salt be mixed with several times its weight of 
 sodium carbonate (both being as dry as possible), and the mixture 
 be strongly heated in a dry narrow test-tube, a sublimate of metallic 
 mercury will be obtained. The sublimed mercury will present the 
 appearance of a bright metallic mirror, but if examined by means 
 of a lens, or if rubbed with a glass rod, distinct globules of liquid 
 metal will be visible. 
 
 WET REACTIONS. (a) Mercuric Compounds. Of the com- 
 mon salts, the nitrate, sulphate, chloride, and bromide (but not the 
 iodide) are soluble in water, but the solubility is not very great. 
 
 KHO or NaHO gives with mercuric compounds a yellow 
 precipitate * of mercuric oxide, HgO 
 
 HgCl 2 + 2KHO = HgO + H 2 O + 2KC1 
 
 Nitrates, Hg(NO 3 ) (Hg 2 )(NO 3 ).> 
 
 Basic nitrates, Hg(N0 3 ) 2 ,2HgO,H 2 ... iHg,)(NO,),,{Hg a )O,H t O 
 Double ammonium compounds, NH 2 HgCl NH 2 (Hg 2 )Cl 
 
 * On the first addition of the reagent, the precipitate appears a brownish 
 colour (probably due to the momentary formation of the hydroxide, which is 
 incapable of existing), but almost immediately it becomes yellow. Why the 
 oxide obtained by precipitation should be yellow, while that prepared in the 
 dry way is brick-red, is not known. Compare also the sulphide. 
 
Group II. Division I. 77 
 
 The precipitate is insoluble in excess of the reagent. 
 
 NH 4 HO produces a white precipitate of an ammoniacal mercuric 
 compound, where two atoms of hydrogen from the ammonium 
 radical are replaced by the divalent atom Hg ; thus 
 
 HgCl 2 + 2 NH 4 HO = NH.HgCl + NH 4 C1 + 2H 2 O 
 
 i " ^^^^ -^ 
 
 Or with mercuric nitrate 
 
 Hg(N0 3 ) 2 + 2NH 4 HO = (NH 2 Hg)N0 3 + NH 4 NO 3 + 2H 2 O 
 
 H 2 S produces a black * precipitate of HgS. The precipitation 
 is only complete after some time, and when the solution is con- 
 siderably dilute. The compound is insoluble in HC1, and in 
 HNO 3 even when boiling. (The prolonged action of boiling HNO 3 
 partially converts it into the white compound Hg(NO 3 ) 2 , 2HgS.) 
 Mercuric sulphide dissolves in aqua regia, forming mercuric 
 chloride. In the presence of caustic alkalies it dissolves in sodium 
 or potassium sulphide (not in ammonium sulphide], forming the 
 double sulphides, HgS,Na 2 S and HgS,K 2 S. 
 
 (NH 4 ),S gives the same precipitate. The same result is also 
 obtained by the addition of sodium thiosulphate, Na 2 S 2 O 3 , to a warm 
 solution acidified with HC1. 
 
 II precipitates HgI 2 as a rich scarlet compound, soluble in 
 excess of either solution. When first precipitated it appears 
 yellow, but quickly turns salmon-red and then scarlet. The com- 
 pound is dimorphous, and can be obtained either in the red 
 (quadratic crystals) or the yellow (rhombic prisms'] variety. 
 
 Reduction of Mercuric Compounds. By reducing agents 
 mercuric compounds may be converted into mercurous salts, or the 
 reduction may go a stage further and result in the precipitation of 
 mercury in the metallic state. Thus, on the addition of stannous 
 
 * The progress of this precipitation is accompanied by very characteristic 
 changes of colour. The first action of the H 2 S is to give a -white precipitate, 
 which then passes through various shades of colour, from yellow to yellowish- 
 red, to brown, and lastly, black. The white substance is a compound of HgS 
 with the mercuric salt in solution, HgCl 2 ,2HgS, or Hg(NO ? )2 ,2HgS, and the 
 changes in colour are ascribed to the gradual conversion of this into black, HgS. 
 It does not appear quite obvious, however, how a gradual alteration in the pro- 
 portions of the white double compound and the black sulphide can give the 
 orange and reddish tints which are seen. Mercuric sulphide prepared by other 
 processes is red (the pigment known as vermilion}. Why the compounds pre- 
 pared in different ways should be so very different in colour is not known ; 
 probably it is a case of dimorphism similar to that exhibited by HgI 2 , and it 
 may be that to some extent the red HgS is precipitated, but is not stable 
 under the conditions which are present, t and so passes into the black 
 modification. 
 
78 Qualitative Analysis. 
 
 chloride, SnCl 2 , a white precipitate of mercurous chloride is 
 produced 
 
 2 HgCl 2 + SnCl 2 = (Hg 2 )Cl 2 + SnCl 4 
 
 On gently warming with an excess of stannous chloride, the pre- 
 cipitated mercurous chloride changes to a grey deposit of mercury 
 in a condition of fine powder 
 
 (Hg) 2 Cl 2 + SnCl 2 = SnCl 4 + 2Hg 
 
 Many metals are capable of displacing mercury from its solutions, 
 the mercury being deposited upon the metal. Thus, if a strip of 
 clean copper be immersed in a neutral or slightly acid solution of a 
 mercury salt, it becomes coated with a white silvery deposit of an 
 amalgam of copper and mercury, from which the mercury can be 
 readily volatilised and obtained as a metallic sublimate by heating 
 the copper in a dry test-tube. In the case of mercuric salts, the 
 action may be regarded as taking place in two stages ; thus 
 
 2Hg(N0 3 ) 2 + Cu = (Hg 2 )(N0 3 ) 2 + Cu(N0 3 ) 2 
 (Hg 2 )(N0 3 ) 2 + Cu = 2Hg + Cu(N0 3 ) 2 
 
 (fi) Mercurous Compounds. Of the common salts mer- 
 curous nitrate is the only one which is readily soluble, and this only 
 so long as the water is acid with nitric acid. The addition of much 
 water results in the precipitation of a basic nitrate. Mercurous 
 sulphate is soluble with difficulty. 
 
 KHO or NaHO throws down a black precipitate of (Hg 2 )O. 
 Mercurous oxide is very unstable. When gently warmed, or even 
 upon exposure to light, it is converted into HgO and Hg. 
 
 NH HO precipitates an ammoniacal mercurous compound, 
 which is black. Its composition is exactly analogous to the 
 corresponding mercuric compound 
 
 Hg 2 (N0 3 ) 2 + 2NH 4 HO = NH 2 (Hg 2 )N0 3 + NH 4 NO 3 
 H 2 S produces a black precipitate, which is a mixture of HgS 
 and Hg. (Hg 2 S is not known to exist.) This precipitate, therefore, 
 behaves, on treatment with nitric acid, in the same way as that 
 obtained from a mercuric solution. If we imagine the free atom 
 of mercury as first dissolving in the acid, the mercuric nitrate so 
 formed unites with HgS, giving the white insoluble compound 
 Hg(N0 3 ) 2 ,2HgS. 
 
 (NH 4 ) 2 S gives the same precipitate, but in this case the free 
 mercury will be also converted into HgS in proportion as the 
 ammonium sulphide contains more or less polysulphide, for alkaline 
 poly sulphides convert metallic mercury into HgS. 
 

 Group II. Division i. 79 
 
 HC1 and soluble chlorides precipitate white mercurous chloride, 
 Hg, 2 Cl 2 . Insoluble in dilute acids ; soluble in boiling HNO 3 , being 
 converted into HgCl 2 and Hg, and the mercury then dissolves to 
 mercuric nitrate, with evolution of oxides of nitrogen. 
 
 Long boiling with concentrated HC1 decomposes Hg 2 Cl 2 into 
 HgCl 2 (which dissolves) and Hg, which separates (mercury being 
 insoluble in HC1). Chlorine water converts it into mercuric 
 chloride ; thus 
 
 Hg 2 Cl 2 + C1 2 = 2HgCl 2 
 
 Ammonia converts it into black mercurous ammonium chloride, 
 NH 2 (Hg 2 )Cl. (This constitutes one of the most characteristic 
 reactions for mercurous compounds.) 
 
 Mercurous salts are reduced to metallic mercury by the reducing 
 agents which reduce the mercuric compounds ; thus, with stannous 
 chloride a grey precipitate of mercury is at once produced 
 
 Hg 2 (NO 3 ) 2 + SnCl 2 + 2HC1 = SnCl 4 = 2HNO 3 + 2 Hg 
 
 Lead, Pb. 
 
 DRY REACTIONS. Lead compounds are very readily reduced 
 when heated upon charcoal before the blowpipe flame, either alone 
 or mixed with sodium carbonate or potassium cyanide. Globules 
 of metallic lead are thus obtained, and at the same time a yellowish 
 incrustation is formed, consisting of the oxide, PbO (litharge). 
 When cold, one of the globules can be removed and the properties 
 of the metal examined. Lead may be recognised by its malleability 
 and softness, the latter property enabling it to leave a black mark 
 when rubbed upon paper. It is insoluble in cold HC1 or H 2 SO 4 , 
 but readily dissolves in HNO 3 , forming Pb(NO 3 ) 2 , which, being 
 insoluble in nitric acid, remains as a white deposit, but which 
 dissolves on dilution with water. 
 
 WET REACTIONS. The only salts of lead which are met with 
 in analysis are derived from plumbic oxide, PbO, in which the 
 metal is divalent.* Of the common salts, the nitrate and acetate 
 are readily soluble in water ; the chloride, bromide, and iodide are 
 sparingly soluble. 
 
 KHO, NaHO, or NH 4 HO gives a white precipitate of either 
 lead hydroxide or a basic compound, depending upon whether the 
 lead solution or the alkali is in excess all the time. For example, 
 if the potash be added to the lead solution, the precipitation 
 
 * Salts are known in which lead is tetravalent, e.g. lead tetracetate, 
 Pb(GtHiO*)4. The compound PbCl 4 has also b:en obtained. 
 
82 Qualitative Analysis. 
 
 In the case of bismuth chloride, the oxychloride is thrown 
 down 
 
 BiCl 3 + H 2 O = BiOCl + 2HC1 
 
 This compound is not so easily dissolved by HC1 as the basic 
 nitrate is by HNO 3 , therefore the reaction is complete, the whole 
 of the bismuth being precipitated if the solution is dilute. 
 
 KHO, NaHO, or NH 4 HO precipitates the white hydroxide 
 Bi(HO) 3 , or Bi 2 O 3 ,3H 2 O.* Insoluble in excess of the precipitants. 
 From boiling solutions, or on heating to boiling, the monohydrated 
 oxide is formed, Bi 2 O 3 ,H 2 O. 
 
 K 2 CO 3 , Na 2 CO 3 , or (NH 4 ) 2 CO 3 throws down a white basic 
 carbonate, (BiO) 2 CO 3 .f Insoluble in excess of the reagents. 
 
 H 2 S or (NH 4 ) 2 S precipitates bismuthous sulphide, Bi 2 S 3 , as a 
 dark brown, almost black, compound. Soluble in HNO 3 ; insoluble 
 in alkaline sulphides.^ 
 
 Sulphuric acid produces no precipitate with a bismuth salt. 
 Potassium dichromate, K 2 Cr 2 O 7 , precipitates basic bismuth dichro- 
 mate, (BiO) 2 Cr 2 O 7 . Insoluble in KHO. (These two reactions 
 distinguish between Pb and Bi.) 
 
 TheNmost characteristic reaction for bismuth is the formation 
 of the insoluble oxychloride, BiOCl, on the addition of water to an 
 acid solution of BiCl 3 . As explained above, the reaction with the 
 chloride is more delicate than with the nitrate, hence, if the nitrate 
 is used, it should be converted into the oxychloride. This can be 
 accomplished by the addition of sodium chloride, whereby both 
 the precipitated basic nitrate, as well as the normal nitrate re- 
 maining dissolved, are converted by double decomposition into 
 oxychloride; thus 
 
 Bi(NO 3 ) 3 ,2Bi(HO) 3 + 3NaCl = 
 
 or, Bi(HO) 2 NO 3 + NaCl = BiOCl + H 2 O + NaNO 3 
 Bi(NO 3 ) 3 + sNaCl + H 2 O = BiOCl + 3NaNO 3 + 2HC1 
 
 as a molecule of bismuth nitrate, in which two of the NO 3 groups are re- 
 placed by (HO) ; thus, Bi(HO) 2 NO 3 . The composition in both cases is the 
 same, the simpler formula being merely the other divided by three. 
 
 * Three hydrated oxides are known, BioO^HoO, BioO 3 ,2HoO, and 
 Bi 2 O 3 ,H 2 O. 
 
 f Or this may be regarded as normal bismuth carbonate and oxide, 
 Bi 2 (C0 3 ) 3>2 Bi 2 3 . 
 
 J In this respect bismuth differs from the elements Sb and As, with which 
 it is associated in the " natural " or periodic classification. These two elements 
 form soluble thioantimonates and thioarsenrttes, no such compounds of bismuth 
 being known. 
 
Group II. Division i. 83 
 
 Cadmium, Ccl. 
 
 DRY REACTIONS. Cadmium compounds, heated on charcoal 
 with sodium carbonate, are easily reduced ; but, owing to the ready 
 volatility of the metal, the latter is converted into the oxide, which 
 is deposited as a brown incrustation upon the charcoal.* 
 
 WET REACTIONS. The metal strongly resembles zinc (with 
 which it is associated in the natural classification) in its behaviour 
 towards acids, dissolving readily in dilute acids with evolution of 
 hydrogen, or, in the case of nitric acid, of oxides of nitrogen. Only 
 one series of salts is known, derived from the only oxide, CdO. Of 
 the common salts, the nitrate, sulphate, chloride (bromide, iodide, 
 and acetate) are soluble in water. 
 
 KHO, NaHO, or NH 4 HO precipitates the white hydroxide, 
 Cd(HO) 2 . Insoluble in excess of KHO or NaHO, but soluble in 
 NH 4 'HO. 
 
 K,CO 3 , Na 2 CO 3 , or (NH 4 ) 2 CO 3 gives a white precipitate of 
 CdCO 3 . The presence of NH 4 HO prevents the precipitation. The 
 precipitate is not soluble in excess of the reagent as ordinarily 
 used ; but it dissolves in concentrated solutions of K 2 CO 3 and 
 Na 2 CO 3 . From this solution it is reprecipitated on dilution. 
 
 H 2 S or (NH 4 ) 2 S precipitates cadmium sulphide, CdS, distin- 
 guished from the sulphides of all the other metals of this division 
 by its pure yellow colour. It is more easily soluble in acids than 
 the other sulphides of the group, and therefore the acid solution 
 from which it is precipitated must be dilute, and not too strongly 
 acid, to ensure complete precipitation. 
 
 CdS is insoluble in potassium cyanide, therefore it is capable of 
 being precipitated in the presence of this salt. And for the same 
 reason, H 2 S will throw down a precipitate of CdS from a solution 
 of CdCy 2 in excess of KCy, which contains the double cyanide 
 CdCy 2 ,2KCy (see Method of Separation from Copper). CdS is 
 insoluble in alkaline sulphides, which distinguishes it from arsenious 
 sulphide, which is the only other yellow sulphide (see footnote, 
 P- 77)- 
 
 Copper, Cu. 
 
 DRY REACTIONS. Copper compounds are reduced to metallic 
 copper when strongly heated upon charcoal along with sodium 
 
 * With the exception of the rare mineral greenockife, CdS, cadmium is 
 always found in nature closely associated with zinc ores, and only in small 
 quantities. Being much more volatile, however, than zinc, it is often possible 
 to obtain the brown incrustation of CdO before enough of the zinc has been 
 vaporised to mask the reaction. 
 
 
84 Qualitative Analysis. 
 
 carbonate in the reducing flame. Reddish scales, or even globules, 
 of metal will be found. Heated in a borax bead, copper salts 
 impart a colour which is green while the bead is hot, but bluish 
 When cold. A more delicate dry test for copper is made by heating 
 the compound upon a platinum loop in a Bunsen flame, ancj 
 supplying the flame at the same time with a little hydrochloric acid 
 gas, which is admitted by one of the air-holes of the lamp. A little 
 strong hydrochloric acid is heated in a test-tube having a delivery 
 tube leading into one of the air-holes, in the manner shown in 
 Fig. ii. Under these 'circumstances, the copper compound, which 
 
 FIG. 
 
 may otherwise impart no colour to the flame, gives a brilliant blue 
 colour, which instantly changes to green when the supply of acid 
 gas is momentarily stopped by partially withdrawing the test-tube. 
 WET REACTIONS. Copper is not acted upon by dilute HC1 or 
 dilute H 2 SO 4 . Boiling strong HC1 slowly dissolves it, giving 
 cuprous chloride and hydrogen. Hot concentrated H 2 SO 4 gives 
 SO 2 and copper sulphate. Nitric acid readily dissolves the metal, 
 forming Cu(NO 3 ) 2 and oxides of nitrogen. Copper forms two 
 series of salts, cuprous and cuprtc, derived from the two oxides 
 Cu 2 O and CuO. The former readily pass by oxidation into cupric 
 compounds. 
 
Group II. Division I. 85 
 
 (a] Cupric Salts. Of the common cupric salts, the sulphate, 
 nitrate, chloride (bromide and acetate) are readily soluble in water. 
 In the crystallised or hydrated condition they are blue or green, 
 but in the anhydrous state either white or pale yellow. 
 
 EHO or NaHO produces a pale blue precipitate of cupric 
 hydroxide or hydrated oxide Cu(HO) 2 or CuO,H 2 O. Insoluble in 
 excess.* On warming the mixture (or if the precipitation is con- 
 ducted with hot solutions), this loses a portion of its water of 
 hydration and is changed to a nearly black hydrated oxide 
 
 3CuO,H 2 O or Cu(HO) 2 ,2CuO 
 
 NH,HO or (NH 4 ) 2 CO 3 precipitates a light blue basic com- 
 pound, readily soluble in excess to a deep blue solution (charac- 
 teristic of copper compounds). If the blue solution be allowed to 
 evaporate, it deposits dark blue crystals, having the composition 
 CuSO 4 ,4NH 3 ,H 2 O f in the caseof the sulphate; andCuCl 2 ,4NH 3 ,H 2 O 
 with the chloride. 
 
 K 2 CO 3 or Na 2 CO 3 gives a greenish precipitate of the basic 
 carbonate, Cu(CO 3 ),Cu(HO) 2 4 Insoluble in excess of the reagent 
 under ordinary circumstances. If, however, the reagent be a highly 
 concentrated solution of the alkaline carbonate, the precipitate dis- 
 solves to a deep blue solution, believed to contain a double carbonate 
 of copper and the alkali. Dilution with water decomposes this 
 compound, and reprecipitates the basic carbonate (see also Cd, Ni, 
 Co, Fe). On boiling the mixture, tht precipitate is converted into 
 the black hydrated oxide ; thus 
 
 3[CuC0 3 ,Cu(HO) 2 ] - 2[Cu(HO) 2 ,2CuO] + 3CO 2 + H 2 O 
 
 H 2 S or (NH 4 ) 2 S produces a nearly black precipitate of cupric 
 sulphide, CuS, which, when exposed to the. air in a moist condition, 
 absorbs oxygen and is converted into the sulphate. The precipitate 
 is slightly soluble in ammonium sulphide. It readily dissolves in 
 potassium cyanide, therefore H 2 S fails to give a precipitate of 
 copper sulphide from a solution of Cu 2 Cy 2 1| in KCy (which 
 contains cuprous potassium cyanide, Cu 2 Cy 2 ,6KCy) (compare 
 cadmium). 
 
 * In the presence of tartaric acid (or alkaline tartrates), Cu(HO) 2 is dissolved 
 by excess of KHO (Feliling's solution. See below). 
 
 t Copper sulphate crystals themselves have the composition CuSO 4 ,5H 2 O. 
 
 This compound occurs in nature as the mineral malachite. 
 
 The composition of precipitated cupric sulphide is said to be represented 
 by the formula Cu 4 S 3 (Thomson). 
 
 || See cuprous cyanide. 
 
86 Qualitative Analysis. 
 
 Reduction of Cupric Salts. (i) Many organic substances 
 reduce cupric salts in alkaline solutions, with precipitation of 
 cuprous oxide, Cu 2 O. Thus, if KHO be added to a solution of 
 CuSO 4 in the presence of grape sugar, the Cu(HO) 2 first precipi- 
 tated dissolves in excess of KHO, giving a blue solution ; on gently 
 warming the liquid, a bright red precipitate of Cu 2 O is thrown 
 down.* The reaction may be regarded as an abstraction of one 
 atom of oxygen, by the sugar, from two molecules of cupric oxide 
 
 2CuO = Cu 2 O + O 
 
 Cuprous oxide is soluble in hydrochloric acid, with the formation 
 of cuprous chloride. 
 
 (2) Cupric salts may be reduced by nascent hydrogen. Thus, 
 when metallic copper is placed in a solution of cupric chloride in 
 strong hydrochloric acid, and the mixture boiled, the green cupric 
 chloride becomes colourless, owing to its conversion into cuprous 
 chloride by the hydrogen disengaged by the action of the HC1 
 upon the copper 
 
 2CuCl 2 + H 2 = Cu 2 Cl 2 f + 2HC1 
 
 If the colourless solution be poured into water, the cuprous chloride 
 is thrown down as a white precipitate ; insoluble in water, soluble 
 in HC1, in NH 4 HO, and in NH 4 C1. 
 
 (b] Cuprous Salts. The common cuprous salts are all 
 insoluble in water. For the reactions, a solution of cuprous 
 chloride in hydrochloric acid may be used. 
 
 KHO or NaHO gives a yellow precipitate of cuprous hydroxide, 
 Cu 2 (HO) 2 or Cu 2 O,H 2 O. If the mixture be heated, the precipitate 
 is converted into the red cuprous oxide. 
 
 NH 4 HO gives no precipitate, but forms a soluble compound 
 having the composition Cu 2 Cl 2 ,2NH 3 . The solution is colourless, 
 but has the property (also possessed by the solution of Cu 2 Cl 2 in 
 HC1) of absorbing oxygen from the air, first becoming brown, and 
 finally depositing a greenish precipitate of cupric oxychloride, 
 CuCl 2 ,3CuO,4H 2 04 
 
 * This reaction is utilised as a test for sugar. The reaction is more delicate 
 when an alkaline solution of cupric tartrate (Feeling's solution) is employed. 
 This is prepared by neutralising a solution of tartaric acid with excess of 
 KHO, and adding to the alkaline liquid a snp.ll quantity of copper sulphate 
 solution. A light blue solution is obtained. 
 
 f The cuprous and cupric compounds are related to each other in the same 
 way as the mercurous and mercuric salts. Some chemists regard the cuprous 
 salts as containing monovalent copper. It is preferable, however, to represent 
 them as compounds of the divalent double atom Cu 2 . 
 
 This ammoniacal solution of cuprous chloride also absorbs carbon mon- 
 oxide, forming a compound believed to have the composition COCu 2 Cl2,2H 2 O. 
 
Separation of the Metals of Group II. Division i. 87 
 
 KI produces a yellowish-white precipitate of cuprous iodide, 
 CuJ 2 
 
 Cu 2 Cl 2 + 2KI = Cu 2 I., 4- 2KC1 
 
 Cupric iodide is unknown. When KI is added to a solution of a 
 cupric salt, cuprous iodide is thrown down, and iodine is liberated, 
 which colours the liquid brown 
 
 2CuCl 2 4- 4KI = Cu 2 I 2 4- 4KCL + I 2 
 
 In the presence of a reducing agent, such as a ferrous salt, or 
 sulphurous acid, the cupric salt is first reduced to the cuprous 
 state, which with the KI then gives cuprous iodide without separa- 
 tion of iodine 
 
 2CuCl 2 + 2FeCl 2 = 2FeCl 3 4- Cu 2 Cl 2 
 2CuCl 2 + H 2 SO 3 + H 2 O = H 2 SO 4 + 2 HC1 4- Cu 2 Cl 2 
 
 KCy gives a white precipitate of Cu 2 Cy 2 ; soluble in excess of 
 KCy, giving a double cyanide, Cu 2 Cy 2 ,6KCy. 
 
 \Cupric cyanide, although known, is very unstable, hence when 
 KCy is added to a cupric salt, the cupric cyanide, even if formed, 
 quickly decomposes, and a mixture or compound of cuprous cyanide 
 and cupric cyanide is produced. As with the iodide, if reducing 
 agents are present, cuprous cyanide alone is produced.] 
 
 SEPARATION OF THE METALS OF GROUP II. DIVISION i. 
 
 The separation of these metals is based upon 
 
 1. The difference in the behaviour of their sulphides towards 
 nitric acid ; sulphides of Cu, Cd, and Bi being dissolved, while 
 those of Hg and Pb are either not dissolved or are changed into 
 other insoluble compounds. 
 
 2. The solubility of the hydroxides of Cd and Cu in NH 4 HO, 
 and the insolubility of Bi(HO) 3 . 
 
 3. The solubility of copper sulphide in KCy, and the insolubility 
 of CdS. 
 
 The solution containing salts of these metals is acidified by the 
 addition of a few drops of dilute HC1,* and a stream of H 2 S is 
 
 It likewise absorbs acetylene, C 2 H 2 , producing a red precipitate of cuprous 
 acetylide ; thus 
 
 Cu 2 Cl 2 ,2NH 3 4- H 2 O 4- C 2 H 2 = Cu 2 C 2( H 2 O + 2NH 4 C1 
 * In the ordinary course of analysis, the metals of Group I. are thereby 
 precipitated as chlorides ; nearly the whole of the lead will therefore be 
 removed before Group II. is examined. When the exercise is confined to this 
 group, any precipitate of PbCl 2 obtained may be removed by filtration and 
 examined separately. 
 
88 
 
 Qualitative A nalysis. 
 
 allowed to bubble moderately slowly through the solution until 
 precipitation is complete (see Precipitation, p. 7). During the 
 process the liquid should be frequently stirred with the tube 
 delivering the gas. After filtration, the solution should be diluted 
 with water, and H 2 S should be again passed through the liquid for 
 a moment or two, to make sure that the precipitation has been 
 complete. 
 
 The precipitate contains the sulphides of all the metals of the division. It 
 should be well washed,* and then transferred to a small porcelain dish 
 with the least possible quantity of water. About an equal volume of 
 strong HNO 3 is then added, and the mixture boiled until no further 
 dissolving action can be detected. The mixture is then diluted, and a 
 few drops of dilute H 2 SO 4 added. Before filtering, the mixture should 
 be cooled. 
 
 The residue may contain HgS 
 (black), Hg(NO 3 ) 2 ,2HgS (white), 
 and PbSO 4 (white). Boil with a 
 solution of ammonium acetate. 
 
 The filtrate contains the nitrates 
 of Bi, Cd, and Cu. Add excess 
 ofNH 4 HO, and boil. 
 
 
 The precipi- 
 
 The solution con- 
 
 
 
 The residue. 
 
 The solution. 
 
 tate is white 
 
 tains ammonia- 
 
 Dissolve in the 
 
 Add K 2 CrO 4 . 
 
 Bi 2 3 ,H 2 0.t 
 
 cal Cd and Cu 
 
 least quantity 
 
 A yellow pre- 
 
 
 
 compounds. The 
 
 of aqua regia. 
 Boil to expel 
 
 cipitate of 
 PbCr0 4 . 
 
 Confirm by dis- 
 solving in a few 
 
 presence of Cu is 
 seen by the blue 
 
 chlorine, and 
 
 
 drops of HC1, 
 
 colour. If blue, 
 
 neutralise with 
 
 
 and adding the 
 
 add KCy until 
 
 NaHO.J Acidify 
 
 
 solution drop 
 
 colourless, and 
 
 with II Cl, and 
 
 
 by drop into a 
 
 passH 2 S. Yellow 
 
 introduce a strip 
 
 
 test-tube nearly 
 
 CdS precipitated. 
 
 of clean copper, 
 
 
 filled with 
 
 
 
 which will be- 
 
 
 water. White 
 
 If copper is absent, 
 
 come coated 
 
 
 BiOCl is pre- 
 
 H 2 S may be 
 
 with a silvery 
 
 
 cipitated. 
 
 added at once. 
 
 deposit of mer- 
 
 
 
 
 cury. 
 
 
 
 
 * Unless HC1 (and soluble chlorides) be washed out of this precipitate, the 
 addition of HNO 3 will result in the formation of a little aqua regia, and this 
 will dissolve a portion of the mercuric sulphide. 
 
 t If any lead sulphate escaped precipitation in the previous step, or if any 
 of the mercuric sulphide was dissolved (see previous note), compounds of these 
 metals will be here thrown down by NH 4 HO. Hence it is necessary to con- 
 firm Bi as indicated. 
 
 This step is necessary, in order to remove HNO 3 , the presence of which 
 might prevent the deposition of Hg. 
 
CHAPTER IX. 
 REACTIONS OF THE METALS OF GROUP II. DIVISION 2. 
 
 Arsenic, Antimony, Tin [Gold, Platinum]. 
 
 THE two elements arsenic and antimony belong to the same 
 natural family. Tin, on the other hand, is more nearly related to 
 lead. The three elements are associated together in the same 
 analytical group, for the reason that they possess in common the 
 property of forming " thio " acids, whose alkaline salts are soluble in 
 water ; namely, thio-arsenites, thio-antimonites and thio-stannates. 
 In other words, the sulphides of these metals are soluble in alkaline 
 sulphides. 
 
 Arsenic, As. 
 
 DRY REACTIONS. Compounds of arsenic are easily reduced 
 and the element obtained in the " metallic " state, by heating them 
 with suitable reducing agents. 
 
 Thus, when heated upon charcoal with Na 2 CO 3 and KCy, 
 arsenical compounds are reduced ; but the metal, being extremely 
 volatile and readily combustible, is for the most part burnt to 
 arsenious oxide, As 4 O 6 , which passes off as a white fume. At the 
 same time some of the vapour of the element itself is carried away 
 with the fumes of the oxide, and is readily recognised by its strong, 
 unpleasant, and characteristic garlic-like odour. Arsenic cannot be 
 melted by heat after the manner of most metals, but passes from 
 the solid to the vaporous states without liquefying ; hence it never 
 yields metallic globules upon the charcoal. 
 
 The reduction may be made by heating the arsenic compound 
 in a glass tube with KCy, or a mixture of Na 2 CO 3 and KCy. The 
 reaction is conveniently studied by using arsenious oxide. A small 
 fragment (about the size of a pin's head) is placed in a narrow test- 
 tube * and covered by adding a mixture of Na 2 CO 3 and KCy (equal 
 
 * For such experiments small test-tubes 4 X t> inches answer admirably. 
 Bulb tubes are neither necessary nor desirable. 
 
QO Qualitative Analysis. 
 
 parts), the materials being as dry as possible. The total quantity 
 of material in the tube should not occupy more space than is shown 
 in Fig. 12. On the application of a gentle heat, the first effect will 
 be the expulsion of moisture from the imperfectly dried materials, 
 
 FIG. 12. 
 
 which condenses upon the sides of the tube. This may be driven 
 up the tube by gently warming it, and finally removed by intro- 
 ducing a "spill" of blotting-paper. When no more moisture 
 collects, the mixture may be steadily heated in the tip of a small 
 Bunsen flame. .The arsenic under these circumstances is vaporised 
 without undergoing combustion, and sublimes upon the tube as a 
 metallic mirror. Sufficient of the vapour escapes condensation to 
 enable the strong garlic odour to be detected. 
 
 The reaction which takes place may be expressed by the 
 equation 
 
 As 4 O 6 4- 6KCy = As 4 + 6KOCy 
 
 The oxide of arsenic may be reduced by being heated with charcoal 
 alone in a glass tube, when a metallic sublimate is also obtained. 
 In this case it is necessary to cover the arsenic compound with a 
 layer of charcoal, which should be strongly heated before the oxide 
 of arsenic becomes hot, so that the vapour shall pass through the 
 heated carbon ; otherwise the arsenious oxide might be entirely 
 volatilised before the carbon had become hot enough to reduce the 
 compound. 
 
 WET REACTIONS. The element arsenic has little or no claim to 
 be ranked as a metal in the strict acceptation of the term. Being 
 like a metal in appearance and in some of its physical properties, 
 it is usually called a metalloid. In the natural family to which it 
 belongs, it stands between phosphorus and nitrogen on the one 
 
Group II. Division 2. 91 
 
 hand, and antimony and bismuth on the other, showing in its 
 chemical habits a strong similarity to phosphorus. 
 
 It forms two oxides, both of which are acidic in their character ; 
 and all the salts of arsenic are such as contain this element in the 
 acidic or negative portion of the molecule, united with other metals 
 as the base ; such, for example, as the arsenites and arsenates of 
 various metals. No oxysalts of arsenic are known in which the 
 element plays the part of a base, such as nitrate, sulphate, carbonate, 
 etc., of arsenic. 
 
 (a) Arsenious compounds, derived from arsenious oxide, 
 As 4 O 6 .* The arsenites of sodium, potassium, and ammonium alone 
 are soluble in water. For the following reactions, potassium arsenite, 
 K 3 AsO 3 , or a solution of As 4 O 6 in dilute HC1, may be employed. 
 
 H 2 S or (NH 4 ) 2 S precipitates from slightly acid solutions yellow 
 arsenious sulphide, As 2 S 3 , soluble in excess of ammonium sulphide, 
 giving ammonium thio-arsenite ; thus 
 
 As 2 S 3 + 3(NH 4 ) 2 S = 2(NH 4 ) 3 AsS 3 
 
 With yellow ammonium sulphide (polysulphide) the arsenious 
 sulphide is dissolved, with the formation of ammonium thio-ar senate, 
 the As 2 S 3 first uniting with the sulphur of the polysulphide to form 
 As 2 S 5 . From this solution the higher sulphide is precipitated on 
 the addition of an acid. If H 2 S be passed through an aqueous 
 solution of As 4 O,; no precipitate is produced, but the liquid becomes 
 yellow owing to the presence of arsenious sulphide in solution in 
 the colloidal state. The addition of HC1 causes the precipitation 
 of the yellow sulphide. From neutral or alkaline solutions, H 2 S 
 (or (NH 4 ) 2 S) gives no precipitate, as under these circumstances the 
 soluble thio salt is produced. The addition of an acid decomposes 
 the thio salt, and arsenious sulphide is thrown down. The reactions 
 may be expressed thus 
 
 K 3 As0 3 + 3H 2 S = K 3 AsS 3 + 3H 2 O 
 2K 3 AsS 3 + 6HC1 = 6KC1 + 3H 2 S + As 2 S 3 t 
 
 Arsenious sulphide is soluble in caustic alkalies and ammonia, also 
 
 * The acid corresponding to this oxide is not known. Three series of salts 
 are known, which may be considered as being derived from the three hypo- 
 thetical acids ortho-arsenious acid, H 3 AsO 3 ; pyro-arsenious acid, H 4 As 2 O 5 ; 
 and metarsenious acid, HAsO 2 . 
 
 t The hypothetical thio-arsenious acid, HsAsS 3 , may be supposed to be 
 first formetl, and to at once split up into the thio-anhydride As 2 S 3 and into 
 H 2 S, just as on the addition of an acid to a carbonate the unstable carbonic- 
 acid H 2 CO 3 breaks up into the anhydride CO 2 and water. 
 
92 Qualitative Analysis. 
 
 in ammonium carbonate, forming in each case a mixture of arsenite 
 and thio-arsenite of the alkali ; thus 
 
 As,S, + 6KHO = K 3 AsO 3 + K 3 AsS s + 3H 9 O 
 As 2 S 3 + 3 (NH 4 ) 2 C0 3 - (NH 4 ) 3 As0 3 + (NH 4 ) 3 AsS 3 + 3 CO 2 
 
 On the addition of acid, the mixed arsenite and thio-arsenite are 
 decomposed, and arsenious sulphide is reprecipitated 
 
 K 3 As0 3 + K 3 AsS 3 + 6HC1 = 6KC1 + 3H 2 O + As 2 S 3 
 
 Arsenious sulphide is practically insoluble in HC1 (contrast Sb 2 S 3 ), 
 but readily dissolves in HNO 3 , or in HC1 with addition of a crystal 
 of KC1O 3 ; being oxidised in each case into arsenic acid.* 
 
 CuSO, produces, in a solution of potassium arsenite, a green 
 precipitate of hydrogen cupric arsenite, HCuAsO 3 ;t soluble in 
 ammonia and caustic alkalies. If the solution be boiled, the arsenite 
 is oxidised to arsenate at the expense of the copper, which is thereby 
 reduced to cuprous oxide, and precipitated in this form ; thus 
 
 AgNO : . precipitates, from a solution of potassium arsenite, pale- 
 yellow silver arsenite, Ag 3 AsO 3 , soluble both in NH 4 HO and in 
 HNO 3 . If the precipitate be dissolved in either of these solvents, 
 it can only be reprecipitated by neutralisation with the other, when 
 the utmost care has been taken to avoid excess of the first, because 
 the silver arsenite is soluble also in ammonium nitrate. When the 
 ammoniacal solution is boiled for some time, metallic silver is 
 precipitated, and the arsenite is oxidised to arsenate. The solution 
 of silver arsenite in ammonia, and the decomposition on boiling, 
 may be expressed by the following equations : 
 
 Ag 3 As0 3 + 2NH 4 HO = (NH 4 ) 2 AgAs0 3 + Ag 2 O J + H.O 
 (NH 4 ) 2 AgAs0 3 + Ag 2 = (NH 4 ) 2 AgAs0 4 + Ag 2 
 
 Precipitation by Copper (Reinsch's test). If a strip of clean 
 copper foil be introduced into a solution of arsenious oxide in HC1, 
 or an arsenite acidified with the same acid, and the mixture be 
 
 * The modus operandi of this oxidation of the sulphide by HC1 and KC1O 3 
 is that the chlorine first converts the sulphide into arsenic trichloride, which, in 
 presence of water, forms arsenious acid, and that the chlorine oxidises this into 
 arsenic acid. The changes may be thus expressed 
 
 U) As 2 S 3 + 3 C! 2 - 2 AsCl 3 + 3 S 
 
 (2) 2AsCl 3 + 6H 2 O = 6HC1 + 2H 3 AsO 3 
 
 (3) H 3 AsO 3 + H 2 O + C1 2 = 2HC1 + H 3 AsO 4 
 
 t Used as a pigment, under the name of Scheele s green. 
 t Ag 2 O dissolves in ammonia. 
 
Group II. Division 2. 93 
 
 warmed, metallic arsenic is deposited upon the copper, at the same 
 time uniting with it, forming copper arsenide, Cu 5 As 2 . If the 
 copper be then dried, and gently heated in a dry test-tube, the 
 arsenic will be volatilised, and at the same time oxidised, giving, 
 therefore, a white crystalline sublimate of As 4 O 6 (contrast antimony}. 
 
 (b] Arsen/c compounds, derived from arsenic pentoxide, 
 As 2 O 5 .* Arsenates of sodium, potassium, and ammonium are 
 soluble in water. 
 
 H 2 S. From neutral or alkaline solutions of an arsenate no 
 precipitate is produced, the soluble thio-arsenate being formed. 
 The addition of HC1 to this solution throws down the pentasulphide 
 in the form of a yellow precipitate 
 
 K 3 As0 4 + 4H 2 S = 4H 2 4- K 3 AsS t 
 2K 3 AsS 4 + 6HC1 = 6KC1 + As 2 S 5 4- 3H 2 S 
 
 From acidified solutions of an arsenate, H 2 S gives a precipitate 
 after a short time, which is either As 2 S 5 or a mixture of As 2 S 3 and 
 S, depending upon conditions. 
 
 If the solution is strongly acid, and the gas is passed rapidly, 
 the precipitate which slowly comes down is the pentasulphide. 
 On the other hand, if the solution is less strongly acid, and the 
 H 2 S is passed slowly, the arsenic acid is first reduced to arsenious 
 acid, with deposition of sulphur, and the arsenious acid as it forms 
 is converted into arsenious sulphide ; thus 
 
 (1) H 3 AsO 4 + H 2 S = H 2 O + S + H 3 AsO 3 
 
 (2) 2H 3 AsO 3 + sH 2 S = As 2 S 3 +6H 2 O 
 
 This reducing action of H 2 S is very slow, therefore complete pre- 
 cipitatio7i requires considerable time. Warming the liquid hastens 
 the action. The addition of a more powerful reducing agent, such 
 as sulphurous acid, produces the effect at once. 
 
 As 2 S 5 dissolves in alkaline sulphides, forming thio-arsenates, 
 similar to the thio-arsenites. 
 
 CuSO 4 gives a pale bluish precipitate of hydrogen cupric 
 arsenate, HCuAsO 4 . Soluble, like the corresponding arsenite, in 
 ammonia ; but the copper is not reduced on heating the solution, 
 for the reason that arsenates are incapable of any further oxidation 
 in other words, they do not act as reducing agents (contrast 
 arsenites]. 
 
 * Three arseruV acids derived from this oxide are known, namely, ortho- 
 arsenic acid, H 3 AsO4 1 pyro-arsenic acid, H 3 As2O 7 ; and metarsenic acid, 
 HAsO 3 . The two latter, when dissolved in water, are converted into the ortho- 
 acid ; it is, therefore, only possible to have, an aqueous solution of the ortho-acid. 
 
 
94 Qualitative Analysis. 
 
 AgNO, produces a chocolate-coloured precipitate of silver 
 arsenate, Ag 3 AsO 4 , which, like the arsenite, dissolves in NH 4 HO, 
 in HNO 8 , and in NH 4 NO 3 . If a mixture of silver arsenate and 
 arsenite be carefully dissolved in HNO 3 , avoiding any excess, and 
 then ammonia added drop by drop, the silver arsenate is pre- 
 cipitated first (recognised by its chocolate colour), and afterwards 
 the yellow arsenite. 
 
 The ammoniacal solution of silver arsenate is not reduced on 
 boiling, for the same reason that the copper salt is not reduced 
 (contrast arsenites}. 
 
 MgSO,, in presence of NH 4 C1 and NH 4 HO, gives a white 
 crystalline precipitate of ammonium magnesium arsenate, 
 (NH 4 )MgAsO 4 * ; practically insoluble in water. (The correspond- 
 ing arsenite is known, but, being readily soluble in water, it is not 
 produced by precipitation ; hence this reaction serves to distinguish 
 an arsenate from an arsenite?) 
 
 Marsh's Test. 
 
 In the presence of nascent hydrogen, both arsenic *x\& arseniotis 
 compounds are reduced, and arsenuretted hydrogen is evolved. Thus, 
 if a solution of arsenious or arsenic oxide be subjected to electrolysis, 
 or if such solutions are introduced into a mixture from which hydro- 
 gen is being generated (e.g. zinc or magnesum with dilute acid), this 
 compound of arsenic and hydrogen is produced. The action may 
 be regarded as taking place in two stages, first the reduction of the 
 arsenical compound to metallic arsenic, and then the further action 
 of the nascent hydrogen upon this ; thus 
 
 As 4 O 6 + 6H 2 = As 4 + 6H 2 O 
 As 4 + 6H 2 = 4AsH 3 
 
 The properties of arsenuretted hydrogen which are made use of in 
 analysis are the following : 
 
 (1) The deposition of metallic arsenic f from the flame of the 
 burning gas when a cold object is depressed upon the flame. 
 
 (2) The decomposition of the compound on passing through a 
 heated tube, with deposition of an arsenical mirror. 
 
 * This compound closely resembles the corresponding phosphate, 
 (NH 4 )MgPO 4 , and is precipitated under the same conditions. It may, how- 
 ever, be at once distinguished from the phosphate by dissolving it in HC1, and 
 adding H 2 S, when a yellow precipitate of As 2 S 5 is produced. 
 
 f Recent experiments seem to prove that this deposit is in reality a solid 
 hydride of arsenic AsH (Retgers, Zeifs, / anorg. Chetn., iv. 739). 
 
Group II. Division 
 
 95 
 
 (3) The action of the gas upon a solution of silver nitrate, 
 resulting in the precipitation of metallic silver. 
 
 The reaction is made in a small hydrogen generating apparatus, 
 preferably a Woulffs bottle, of about 200 cub. cms. capacity. In this 
 hydrogen is slowly gene- 
 rated from zinc and dilute 
 sulphuric acid, both mate* 
 rials being free from arse- 
 nic. To the exit-tube is 
 attached a tube, drawn 
 out of a piece of combus- 
 tion tube in such a manner 
 as to present one or two 
 constricted places in its 
 length, as shown in Fig. 13. 
 As soon as the air is all 
 expelled from the appa- 
 ratus, the issuing hydrogen 
 is inflamed.* 
 
 A small quantity of the 
 arsenical solution is now 
 introduced through the 
 thistle-tube. The first effect of this is to cause the precipitation 
 of arsenic upon the zinc ; which, constituting a voltaic couple, at 
 once gives rise to a greatly increased rate of evolution of hydrogen. f 
 The colour of the hydrogen flame will be seen to change, and 
 to assume a lilac tint (resembling the colour given to a flame by 
 potassium compounds), and at the same time white fumes of As 4 O ( . 
 escape from the tip of the flame. If now a porcelain dish be de- 
 pressed upon the flame, a rich brown-black metallic-looking stain 
 will be deposited. The deposit being volatile, and the flame very 
 hot, the stain will again disappear if the flame be allowed to 
 impinge for more than a moment or two on the same spot. 
 
 If the drawn-out tube be heated near one of the constrictions, 
 the arsenuretted hydrogen will be decomposed as it passes the hot 
 spot, and an arsenic mirror will be deposited in the tube. 
 
 * A small test-tube should be filled by upward displacement, and tested by 
 a flame before igniting the gas at the exit-tube of the apparatus. As an 
 additional precaution, it is well to throw a duster lightly over the Woulffs 
 bottle before applying a light, so that, should an explosion happen, the broken 
 glass will be prevented from flying about. 
 
 f On this account, it is necessary that the generation of hydrogen before 
 adding the arsenic solution should be quite slow ; and also that the quantity of 
 the arsenic solution added at a time should be small. 
 
 FIG. 13. 
 
96 Qualitative Analysis. 
 
 It will be noticed that the deposition takes place entirely on 
 that part of the tube which is on the side of the flame farthest from 
 the generating vessel * (antimony is deposited from its hydride on 
 both sides of the heated spot). 
 
 Since antimony also forms a gaseous compound with hydrogen 
 which gives similar stains,! it is necessary to employ further 
 confirmatory tests. 
 
 1. The arsenic stains are readily dissolved by a solution of a 
 hypochlorite. If, therefore, a solution of bleaching powder be 
 poured over such stains they immediately disappear 
 
 5Ca(OCl) 2 + 6H 2 + As, = 5CaCl 2 + 4H 3 AsO 4 
 Or 
 
 3Ca(OCl) 2 + 2H 2 O + 2AsH = 3CaCl 2 + 2H 3 AsO 4 
 
 Antimony stains do not dissolve in hypochlorite solutions. 
 
 2. When a stream of H 2 S is passed through the tube containing 
 the deposit of either arsenic or antimony, slightly warmed, in each 
 case the sulphide is formed. Yellow arsenious sulphide volatilises 
 along from the warm region and condenses on the cold distant part 
 of the tube ; antimonious sulphide, reddish or nearly black, remains 
 unmoved, being non-volatile. (If present together, they can in this 
 way be separated.) 
 
 If a stream of gaseous HC1 be now passed through the tube, 
 antimonious sulphide is converted into antimonious chloride, which 
 passes on with the HC1, and may be led into water and again 
 precipitated as the red sulphide with H 2 S. The yellow arsenious 
 sulphide remains in the tube, being unattacked by HC1. 
 
 3. Arsenuretted hydrogen can also be distinguished from the 
 antimony compound, by the difference in the behaviour of the two 
 gases towards silver nitrate. When passed into the silver solution, 
 each gas produces a black precipitate. In the case of arsenic this 
 consists of metallic silver, while with antimony it consists of 
 antimonide of silver ; thus 
 
 6AgN0 3 + 3H 3 O + AsH 3 = 3Ag 2 + 6HNO 3 + H 3 AsO 3 
 3AgN0 3 + SbH 3 = SbAg 3 + 3HN0 3 
 
 If both antimony and arsenic were originally present, these two 
 
 * When the quantity of arsenic present is very minute, it will not be visible 
 in the flame, neither may it be possible to obtain a stain on cold porcelain. 
 But the formation of the mirror in the heated tube is a method by which 
 extremely small traces of arsenic can be detected. 
 
 f The stains given by antimoniuretted hydrogen have a rather more velvety 
 or sooty appearance, when deposited on porcelain from the flame, than those 
 of arsenic. 
 
Group II. Division 2. 97 
 
 precipitates will be produced together. On filtering, the arsenic 
 (now as arsenious acid) goes into the nitrate along with the excess 
 of silver nitrate used. Its presence may be detected by the cautious 
 addition of ammonia, which causes the precipitation of yellow silver 
 arsenite. 
 
 The antimony (as silver antinionide) remains on the filter ; after 
 being washed, it is boiled with a solution of tartaric acid.* The 
 liquid thus obtained after nitration is acidulated with HC1, and 
 antimonious sulphide precipitated by H 2 S. 
 
 Fleitmamts Test. 
 
 When an arsenite, or a solution of arsenious oxide, is warmed 
 in a test-tube with a solution of sodium hydroxide and metallic zinc, 
 arsenuretted hydrogen is evolved, which can be detected by means 
 of a piece of filter-paper moistened with silver nitrate held to the 
 mouth of the tube. A black stain of precipitated silver is pro- 
 duced. Antimdhiuretted hydrogen is not produced from antimony 
 compounds under similar conditions. 
 
 Antimony, Sb. 
 
 DRY REACTIONS. Antimony compounds may be reduced to 
 metallic antimony by heating them with Na 2 CO 3 and KCy upon 
 charcoal. Globules of the metal are thus obtained, which burn in 
 the blowpipe flame, producing white fumes of antimonious oxide, 
 Sb 4 O 6 , the combustion being continued for a short time after 
 removal from the flame. The charcoal at the same time receives a 
 white incrustation. The bead of metal will be found to be very 
 brittle, and, when broken, to exhibit a highly crystalline appearance. 
 Antimony is unacted upon by dilute HC1 or H 2 SO 4 . Nitric acid 
 oxidises it into antimonic acid, or antimonious oxide, depending 
 upon conditions of concentration. 
 
 WET REACTIONS. Both in its physical properties and chemical 
 relations antimony approaches more nearly to the true metals than 
 is the case with arsenic. It forms two series of compounds, " anti- 
 monious " and " antimonic/' which may be regarded as being 
 derived respectively from the two oxides, antimonious oxide, Sb 4 O 6 , 
 and antimony pentoxide, Sb 2 O 5 . 
 
 * Although ordinary metallic antimony is not soluble in tartaric acid, and is 
 only slowly attacked by strong hydrochloric acid even when the metal is in the 
 form of powder, nevertheless, when combined with silver as it is in the precipitate 
 of silver antimonide, the antimony is comparatively easily dissolved by HC1, 
 and dissolves when boiled in a strong solution of tartaric acid. In the latter 
 case antimony tartrate, (SbO)2C 4 H 4 p 6 , goes into solution, leaving metallic 
 silver. With HC1 antimonious chloride is formed, and silver chloride remains. 
 
 H 
 
96 Qualitative Analysis. 
 
 It will be noticed that the deposition takes place entirely on 
 that part of the tube which is on the side of the flame farthest from 
 the generating vessel * (antimony is deposited from its hydride on 
 both sides of the heated spot). 
 
 Since antimony also forms a gaseous compound with hydrogen 
 which gives similar stains,f it is necessary to employ further 
 confirmatory tests. 
 
 1. The arsenic stains are readily dissolved by a solution of a 
 hypochlorite. If, therefore, a solution of bleaching powder be 
 poured over such stains they immediately disappear 
 
 5Ca(OCl) 2 + 6H 2 + As, = 5CaCl 2 + 4H 3 AsO 4 
 Or 
 
 3Ca(OCl) 2 + 2H 2 O + 2AsH = 3CaCl 2 + 2H 3 AsO 4 
 
 Antimony stains do not dissolve in hypochlorite solutions. 
 
 2. When a stream of H 2 S is passed through the tube containing 
 the deposit of either arsenic or antimony, slightly warmed, in each 
 case the sulphide is formed. Yellow arsenious sulphide volatilises 
 along from the warm region and condenses on the cold distant part 
 of the tube ; antimonious sulphide, reddish or nearly black, remains 
 unmoved, being non-volatile. (If present together, they can in this 
 way be separated.) 
 
 If a stream of gaseous HC1 be now passed through the tube, 
 antimonious sulphide is converted into antimonious chloride, which 
 passes on with the HC1, and may be led into water and again 
 precipitated as the red sulphide with H 2 S. The yellow arsenious 
 sulphide remains in the tube, being unattacked by HC1. 
 
 3. Arsenuretted hydrogen can also be distinguished from the 
 antimony compound, by the difference in the behaviour of the two 
 gases towards silver nitrate. When passed into the silver solution, 
 each gas produces a black precipitate. In the case of arsenic this 
 consists of metallic silver, while with antimony it consists of 
 antimonide of silver ; thus 
 
 6AgN0 3 + 3 H 3 + AsH 3 = 3 Ag 2 + 6HNO 3 + H 3 AsO 3 
 3AgN0 3 + SbH 3 = SbAg 3 + 3HN0 3 
 
 If both antimony and arsenic were originally present, these two 
 
 * When the quantity of arsenic present is very minute, it will not be visible 
 in the flame, neither may it be possible to obtain a stain on cold porcelain. 
 But the formation of the mirror in the heated tube is a method by which 
 extremely small traces of arsenic can be detected. 
 
 f The stains given by antimoniuretted hydrogen have a rather more velvety 
 or sooty appearance, when deposited on porcelain from the flame, than those 
 of arsenic. 
 
Group II. Division 2. 97 
 
 precipitates will be produced together. On filtering, the arsenic 
 (now as arsenious acid) goes into the nitrate along with the excess 
 of silver nitrate used. Its presence may be detected by the cautious 
 addition of ammonia, which causes the precipitation of yellow silver 
 arsenite. 
 
 The antimony (as silver antimonide) remains on the filter ; after 
 being washed, it is boiled with a solution of tartaric acid.* The 
 liquid thus obtained after filtration is acidulated with HC1, and 
 antimonious sulphide precipitated by H 2 S. 
 
 Fleitmanris Test. 
 
 When an arsenite, or a solution of arsenious oxide, is warmed 
 in a test-tube with a solution of sodium hydroxide and metallic zinc, 
 arsenuretted hydrogen is evolved, which can be detected by means 
 of a piece of filter-paper moistened with silver nitrate held to the 
 mouth of the tube. A black stain of precipitated silver is pro- 
 duced. Antimdhiuretted hydrogen is not produced from antimony 
 compounds under similar conditions. 
 
 Antimony, Sb. 
 
 DRY REACTIONS. Antimony compounds may be reduced to 
 metallic antimony by heating them with Na 2 CO 3 and KCy upon 
 charcoal. Globules of the metal are thus obtained, which burn in 
 the blowpipe flame, producing white fumes of antimonious oxide, 
 Sb 4 O 6 , the combustion being continued for a short time after 
 removal from the flame. The charcoal at the same time receives a 
 white incrustation. The bead of metal will be found to be very 
 brittle, and, when broken, to exhibit a highly crystalline appearance. 
 Antimony is unacted upon by dilute HC1 or H. 2 SO 4 . Nitric acid 
 oxidises it into antimonic acid, or antimonious oxide, depending 
 upon conditions of concentration. 
 
 WET REACTIONS. Both in its physical properties and chemical 
 relations antimony approaches more nearly to the true metals than 
 is the case with arsenic. It forms two series of compounds, " anti- 
 monious " and " antimonic," which may be regarded as being 
 derived respectively from the two oxides, antimonious oxide, Sb 4 O 6 , 
 and antimony pentoxide, Sb 2 O 5 . 
 
 * Although ordinary metallic antimony is not soluble in tartaric acid, and is 
 only slowly attacked by strong hydrochloric acid even when the metal is in the 
 form of powder, nevertheless, when combined with silver as it is in the precipitate 
 of silver antimonide, the antimony is comparatively easily dissolved by HC1, 
 and dissolves when boiled in a strong solution of tartaric acid. In the latter 
 case antimony tartrate, (SbO) 2 C4H 4 p 6 , goes into solution, leaving metallic 
 silver. With HC1 antimonious chloride is formed, and silver chloride remains. 
 
 H 
 
9^ Qualitative Analysis. 
 
 (a) Antimonious Compounds. In antimonious oxide we 
 see the gradual fading away, so to speak, of the acidic properties 
 exhibited by the corresponding oxides of arsenic and phosphorus 
 (elements with which antimony is associated in the natural classi- 
 fication), and the beginnings of basic qualities. Thus, no acids 
 corresponding; to this oxide are known, and only a few salts (derived 
 from the hypothetical metantimonious acid) have been obtained. 
 Of these the best known is sodium metantimonite, obtained by 
 dissolving the oxide in sodium hydroxide ; thus 
 
 Sb 4 O 6 + 4NaHO = 4NaSbO 2 + 2H 2 O 
 
 On the other hand, this oxide unites with certain acids forming 
 salts, in which the antimony in combination with oxygen as the 
 monovalent radical antimony '/(SbO) takes the place of the positive 
 or basilous element. Of these salts the tartrate, (SbO) 2 (C 4 H 4 O 6 ), 
 and the double potassium tartrate (tartar emetic}, (SbO)K(C 4 H 4 O 6 ), 
 are the most familiar 
 
 Sb 4 6 + 4HK(C 4 H 4 6 ) = 4(SbO)K(C 4 H 4 6 ) + 2H 2 O 
 
 For the following reactions * an acid (HC1) solution of antimonious 
 chloride may be employed. 
 
 KHO, NaliO, NH 4 HO, as well as alkaline carbonates, pre- 
 cipitate antimonious oxide, Sb 4 O 6 ; thus 
 
 4 SbCl 3 + I2KHO = Sb 4 6 + I2KC1 + 6H,0 
 4SbCl 3 + 6K 2 CO 3 = Sb 4 O 6 + I2KC1 + 6CO 2 
 
 The precipitate redissolves in excess of either potassium or 
 sodium hydroxide, forming the respective metantimonites (equation 
 above). 
 
 HO, added in considerable quantity to the acid solution of 
 SbCl 3 , gives a white precipitate of an oxychloride, SbOCl 
 
 SbCl 3 + H 2 - 2HC1 + SbOCl 
 
 This compound is at once distinguished from the similarly pro- 
 duced BiOCl, by the fact that the antimony oxychloride readily 
 dissolves in tartaric acid, giving antimonyl tartrate ; thus 
 
 2SbOCl + H 2 (C 4 H 4 C ) = (SbO) 2 (C 4 H 4 G ) + 2 HC1 
 
 On boiling the precipitated oxychloride for some time with water, 
 the whole of its chlorine is given up, and antimonious oxide is 
 formed 
 
 4SbOCl -1- 2H 2 O = Sb 4 O 6 + 4HC1 
 
 * Except that with AgNO 3 , in which obviously the presence of chlorine 
 would interfere. 
 
Group If. Division 2. 99 
 
 H 2 S and (NH 4 ) 2 S give a red or orange-red precipitate of anti- 
 monious sulphide, Sb 2 S 3 . Soluble in excess of ammonium sulphide. 
 If yellow ammonium sulphide be employed, the thio salt of Sb 2 S 5 is 
 formed, and on acidifying the liquid, Sb. 2 S 5 is precipitated from it. 
 Antimonious sulphide is soluble also in caustic alkalies, from which 
 solution the trisulphide is again thrown on acidifying 
 
 2Sb 2 S 3 + 4KHO = 3KSbS 2 + KSbO 2 + 2H,O 
 3KSbS 2 + KSbO 2 + 4HC1 = 2Sb 2 S 3 + 4KC1 + 2H,O~ 
 
 Antimonious sulphide is not dissolved by ammonia or by ammonium 
 carbonate (contrast arsenic). 
 
 Antimonious sulphide is decomposed by hot hydrochloric 
 acid 
 
 Sb 2 S 3 + 6HC1 = SbCl 3 + 3H 2 S 
 
 This reaction is the reverse of that by which the sulphide is 
 formed. In this case, however, one of the products, namely 
 the H 2 S, is driven from the sphere of action as fast as it is gene- 
 rated. 
 
 AgNO , added to an alkaline solution of antimonious oxide, 
 gives a black precipitate, consisting of silver oxide and metallic 
 silver. The action takes place in two stages. The silver nitrate 
 interacts with the caustic alkali present, giving silver oxide, and this 
 in its turn gives up oxygen to the antimonz'/i, converting it into 
 antimon#/ ; thus 
 
 2AgNO 3 + 2KHO = Ag 2 O + 2KNO 3 + H 2 O 
 KSb0 2 + Ag 2 = Ag 2 + KSb0 3 
 
 (V) Antimonic Compounds. These are derived from the 
 pentoxide, Sb 2 O 5 . This oxide is, however, so feebly acidic that it 
 is not converted into antimonic acid by the action of water. Only 
 two classes of antimonates are known (compare arsenates and 
 phosphates), namely, pyro-antimonates (e.g. K 4 Sb 2 O 7 ) and met- 
 antimonates (e.g. KSbO 3 ). 
 
 Potassium pyro-antimonate is readily soluble in water, while 
 the sodium salt is difficultly soluble, hence the potassium compound 
 is used as a reagent for sodium (p. 22). 
 
 KHO, NaHO, NH 4 HO, as well as alkaline carbonates, give 
 with an acid solution of antimony pentachloride a white precipitate 
 consisting of metantimonic acid 
 
 b/i's'tfHO" = HS'jO, i ;ICC1 + 2K.;O 
 
ioo Qualitative Analysis. 
 
 H 2 O, added in considerable quantity, produces the same pre- 
 cipitate * 
 
 SbCl fi + 3H 2 O = HSbO 3 + 5HC1 
 
 From aqueous solutions of antimonates, the same precipitate is 
 produced by acids 
 
 KSbO 3 + HNO 3 = KNO 3 + HSbO 3 
 
 In the case of pyro-antimonates, the " pyro " acid is first pro- 
 duced, which then passes into the "meta" compound (footnote). 
 
 H 2 S and (NH 4 ) 2 S produce with acidified solutions of anti- 
 monic compounds an orange-red precipitate consisting of Sb 2 S 5 , 
 Sb 2 S 3 , and S in varying proportions. 
 
 The reducing action of the H 2 S converts a portion of the anti- 
 monic compound to the antimonious state, from which by further 
 action the antimonious sulphide is thrown down. The three follow- 
 ing reactions, therefore, proceed simultaneously 
 
 2KSbO 3 + 5H 2 S = Sb 2 S 5 + 5H 2 O + K 2 O 
 
 KSbO 3 + H 2 S = S + KSbO 2 + H ? O 
 2KSbO 2 + 3H,S = Sb 2 S 3 + 3H 2 O + K 2 O 
 
 The action being carried on in acid solution, the K 2 O is at once 
 neutralised. 
 
 The precipitate behaves towards alkaline sulphides and caustic 
 alkalies in a manner similar to that of Sb 2 S 3 ; thus, with KHO 
 potassium thio-antimonate (ortho) and antimonate (meta) are pro- 
 duced, from which an acid reprecipitates the Sb 2 S 5 
 
 4Sb 2 S a + iSKHO = 5K 3 SbS 4 + 3KSbO 3 + 9H,O 
 5K 8 SbS 4 + 3KSbO s + iSHCl = 4Sb 2 S 5 + iSKCl + 9H 2 O 
 
 AgNO :; gives a white precipitate with aqueous solutions of anti- 
 monates, consisting of silver antimonate. The precipitate dissolves 
 in ammonia (distinction between " antimonic " and " antimonious " 
 compounds). 
 
 KI. When antimonic compounds, in presence of HC1, are 
 boiled with a solution of KI, the antimonic compound is reduced 
 to the " antimonious " state, and iodine is liberated. The free 
 iodine, if in quantity, colours the solution brown ; if in small 
 quantity, it may be recognised by the formation of the blue colour 
 
 * This action takes place in two stages, pyro-antimonic acid being first 
 formed, which loses a molecule of water and passes into the more stable 
 metantimonic acid ; thus 
 
 ^ iOAG'1 
 - H,0 -= Sb 2 5 ,H 2 = 2 HSbO 3 
 
Group II. Division 2. 101 
 
 with starch (distinction between " antimonic ' ; and " antimonious " 
 compounds} 
 
 KSb0 3 + 2HC1 + 2KI = KSbO 2 + 2KC1 + H 2 O + I 2 
 Or, in the case of the oxide or hydrated oxide 
 Sb 2 O 5 ,H 2 + 4HC1 + 4KI = Sb 2 O 3 ,H 2 O + 4KC1 + 2H 2 O + 2l 2 
 
 Precipitation of Metallic Antimony. Antimony com- 
 pounds (of either state of oxidation), when in acid (HC1) solution, 
 readily deposit metallic antimony by galvanic action. The test is 
 applied in the following way : one or two drops of the acid solution 
 are placed upon a piece of clean platinum foil, and a small frag- 
 ment of zinc immersed in the liquid. Immediately a black stain is 
 produced upon the platinum by the deposition of metallic antimony. 
 Even very dilute solutions give the stain, hence the test is a delicate 
 one. HC1 has no action upon the deposit, but warm HNO 3 instantly 
 attacks it, giving antimonic acid (hence the absence of nitric acid 
 from the solution before applying the test is desirable). 
 
 Antimoniuretted hydrogen is evolved when antimony 
 compounds are acted upon by nascent hydrogen. The compound 
 undergoes reactions similar to those of the corresponding arsenic 
 compound. The methods for distinguishing between them are 
 described under arsenic. 
 
 Tin, Sn. 
 
 DRY REACTIONS. Compounds of tin are reduced to the metallic 
 state by being heated on charcoal with Na 2 CO 3 and KCy in the 
 reducing flame. A portion of the metal is oxidised by the flame, 
 and produces a white incrustation of SnO 2 upon the charcoal ; 
 this, on being moistened with cobalt nitrate and reheated, assumes 
 a greenish appearance. 
 
 The beads of reduced metal are malleable (therefore easily dis- 
 tinguished from Bi or Sb), but are not soft enough to mark paper 
 in the manner of lead. 
 
 Tin dissolves in hot strong HC1, forming stannous chloride, 
 SnCl 2 , hydrogen being evolved. Hot H 2 SO 4 converts it into 
 stannous sulphate, SnSO 4 , and gives off SO 2 . 
 
 Cold dilute HNO 3 dissolves tin, forming stannous nitrate, 
 Sn(NO 3 ) 2 , and ammonia ; ordinary strong acid (sp. gr. 1-24) attacks 
 it violently, converting it into white metastannic acid, H 10 Sn 5 O 15 . 
 
 WET REACTIONS. Tin forms two classes of compounds, dis- 
 tinguished as " stannous " and " stannic," derived respectively from 
 stannous oxide, SnO, and stannic oxide, SnO 2 . 
 
IO2 Qualitative Analysis. 
 
 (a) Stannous Compounds. Stannous oxide is basic in its 
 character, giving with acids the Stannous salts, in which the metal 
 is divalent. Of these the chloride, sulphate, and nitrate are soluble 
 in water. The chloride is most common. 
 
 KHO, NaHO, NH 4 HO, as well as alkaline carbonates, give 
 with Stannous chloride a white precipitate of hydrated Stannous 
 oxide (basic hydroxide) ; thus 
 
 2 SnCl 2 + 4KHO = 2SnO,H 2 O + 4KC1 + H 2 O 
 
 The precipitate is soluble in excess of KHO or NaHO (but not in 
 the other precipitants), forming alkaline stannites ; thus 
 
 2SnO,H 2 O + 4KHO = 2K 2 SnO 2 + 3H 2 O 
 
 H 2 S or (NH 4 ) 2 S gives, with dilute solutions of Stannous chloride, 
 a deep brown precipitate of Stannous sulphide. Soluble in yellow 
 ammonium sulphide, with the formation, not of a thio-stannitc, but 
 thio-stannate. (Colourless ammonium sulphide is almost without 
 action upon it.) Stannic sulphide, SnS 2 , is thrown down upon the 
 addition of acids. 
 
 Stannous sulphide dissolves in caustic alkalies, forming a 
 stannite and thio-stannite 
 
 2SnS + 4KHO = K 2 SnO 2 + K 2 SnS 2 + 2H 2 O 
 
 From this, on addition of acid, Stannous sulphide is again pre- 
 cipitated 
 
 K 2 SnO 2 + K 2 SnS 2 + 4HC1 = 4KC1 + 2H 2 O + 2SnS 
 
 Stannous sulphide is insoluble in ammonium carbonate. Boil- 
 ing HC1 converts it into SnCl 2 ; while aqua regia oxidises it to 
 stannic chloride, SnCl 4 . 
 
 Oxidation of Stannous Compounds. These substances 
 readily pass, by oxidation, into stannic compounds ; they therefore 
 act the part of powerful reducing agents in a number of reactions, 
 of which the following are important : 
 
 (i) Mercuric chloride, in the presence of small quantities of 
 Stannous chloride, is reduced to mercurous chloride, which is 
 thrown down as a white precipitate. With excess of the Stannous salt, 
 and on warming, the precipitate is further reduced, and becomes 
 grey through the separation of metallic mercury 
 
 2HgCl 2 + SnCl 2 = Hg.,Cl 2 + SnCl 4 
 Hg 2 Cl 2 + SnCl 2 = 2Hg + SnCl 4 
 
 ^2) Perric salts are reduced to the <; ferrous " state ; thus, 
 
Group II. Division 2. 103 
 
 ferric sulphate, in the presence of hydrochloric acid, gives ferrous 
 sulphate and stannic chloride 
 
 Fe,(SO^ 3 + SnCl, + 2HC1 = 2FeSO 4 + H 2 SO 4 + SnCl 4 
 
 (3) Gold chloride, in the presence of acid, is reduced to 
 metallic gold 
 
 2AuCl a + 3SnCL, = 3SnCl 4 4- 2Au 
 
 In neutral solutions, a reddish or purple precipitate (or colora- 
 tion, if very dilute) is produced, known as purple of Cassius. Its 
 formation is promoted by the presence of a little stannic chloride 
 in the stannous compound. Its composition is believed to be ex- 
 pressed by the formula Au 2 ,3SnO 2 . Acids convert it into metallic 
 gold and a stannic salt. 
 
 (4) Cupric salts, chromates, and permanganates are 
 reduced respectively to cuprous, chromic, and manganous salts ; 
 thus 
 
 2KMnO 4 + i6HCl + 5SnCl 2 = 2KC1 + 2MnCl 2 + 8H 2 O + 5SnCl 4 
 
 (5) Bismuth salts are reduced in alkaline solutions, with the 
 precipitation of black bismuthous oxide, Bi 2 O 2 or BiO, and the 
 oxidation of the stannite to stannate. Thus, if bismuth chloride or 
 nitrate (salts derived from Bi 2 O 3 ) are added to an alkaline solution 
 of stannous oxide (i.e. to a solution of potassium stannite), the 
 following reaction takes place : 
 
 K 2 SnO 2 + 6KHO + 2Bi(NO 3 ) 3 = K 2 SnO s + 6KNO 3 + 3H 2 O + 
 
 2BiO 
 
 The reaction will be simpler if we have regard only to the oxide 
 of bismuth, from which the nitrate is derived 
 
 K 2 SnO 2 + Bi 2 O 3 = 2BiO + K 2 SnO 3 
 
 (6) An aqueous solution of stannous chloride gradually absorbs 
 atmospheric oxygen, being converted partly into stannic chloride, 
 and a white insoluble basic chloride ; thus 
 
 3 SnCl 2 4- H 2 O + O = SnCl 4 + SnCl 2 ,SnO,H 2 O 
 
 (b) Stannic Compounds. Stannic oxide, SnO 2 , may be 
 regarded as being both a " basic " and an acidic oxide, for, although 
 itself insoluble in either acids or alkalies, we may consider both 
 the stannic salts and the stannates as being derived from this 
 oxide. The most important stannic salt is the chloride, SnCl 4 . A 
 solution of this compound in hydrochloric acid may be used for 
 the following reactions : 
 
IO4 Qualitative Analysis. 
 
 KHO, NaHO, NH 4 HO, as well as alkaline carbonates, give 
 a white precipitate of hydrated stannic oxide, or stannic acid, 
 SnO,,H 2 O, or H 2 SnO 3 * ; thus 
 
 SnCl 4 + 4NaHO - H 2 SnO 3 + 4NaCl -f H 2 O 
 
 Soluble in HNO 3 and in HC1. Soluble in KHO and NaHO, with 
 formation of the respective stannates, K 2 SnO 3 , and Na 2 SnO 3 ; thus 
 
 H 2 SnO 3 + 2NaHO = Na 2 SnO 3 + 2H 2 O 
 
 H 2 S or (NH 4 ) 2 S precipitates yellow stannic sulphide, SnS 2 
 (with H 2 S the precipitate appears nearly white at first, and is only 
 complete in dilute solutions). The precipitate is soluble in caustic 
 alkalies, in ammonium sulphide, and sulphides of the alkalies, 
 forming thio-stannates ; thus 
 
 3SnS 2 + 6KHO = 2K 2 SnS 3 + K 2 SnO 3 + 3H 2 O 
 SnS 2 + (NH 4 ) 2 S = (NH 4 ) 2 SnS 3 
 
 From these solutions the yellow stannic sulphide is reprecipitated 
 on the addition of HCl 
 
 (NH 4 ) 2 SnS 3 + 2HC1 = 2NH 4 C1 + SnS 2 + H 2 S 
 
 Stannic sulphide is insoluble in ammonium carbonate, but dis- 
 solves in hot strong HCl. Stannic oxysalts, such as the nitrate or 
 sulphate, are unstable in aqueous solution, and are decomposed 
 into stannic or metastannic acid. Therefore, when by double 
 decomposition such oxysalts might be expected to form, the result 
 is the precipitation of one or both of these stannic acids. Thus, 
 when sulphuric acid is added to stannic chloride and the solution 
 diluted with water, the action may be represented by the two 
 equations 
 
 (1) SnCl 4 + 2H 2 SO 4 = 4HC1 + Sn(SO 4 ) 2 
 
 (2) [Sn(SO 4 ) 2 + 3H 2 O = 2H 2 SO 4 + H 2 SnO 3 ] x 5 to re- 
 
 present metastannic acid. 
 
 A similar precipitation takes place when neutral salts of the 
 alkalies are employed, such as sodium sulphate or ammonium i 
 nitrate ; thus 
 
 * Meta-stannic acid, the white compound obtained by the action of nitric 
 acid upon tin, is expressed by the same formula multiplied by five, 5(H 2 SnO 3 ), 
 or H 10 Sn 5 Oi 5 . It forms salts by the replacement of two hydrogen atoms 
 only, as is the case with stannic acid ; their composition may therefore be ex- 
 pressed by the formula (taking potassium metastannate as an example), 
 KjSnOa,4SnOj,4HjO. Boiling (or fusing) with caustic alkalies converts meta- 
 stannates into stannates ; thus 
 
 5K 2 SnO 3 + 8H 2 O 
 
Group II. Division 2. 105 
 
 (1) SnCl 4 + 4(NH 4 )NO 3 = 4NH 4 C1 + Sn(NO 3 ), 
 
 (2) [Sn(N0 3 ) 4 + 3H 2 - 4 HNO 3 + H 2 SnO 3 ] x 5 
 
 Precipitation of Metallic Tin. When zinc is immersed in 
 an acid solution of stannous or stannic chloride, the tin is displaced 
 by the zinc, and precipitated as a grey-black deposit upon the 
 surface of the zinc ; or, if the whole of the zinc becomes dissolved, 
 the tin is left as a scaly powder. It may be collected, and, after 
 being washed, dissolved in hot hydrochloric acid. Tin produces 
 no stain upon platinum as is the case with antimony ; and, again, 
 differs from this metal in dissolving in HC1. 
 
 Gold and Platinum. 
 
 Although these two metals do not belong to the class of rare 
 elements, nevertheless it is very rarely that the student is called 
 upon to analyse mixtures containing either of them. In actual 
 practice these metals are met with only in the analysis of alloys, 
 where their presence is probably more than suspected at the outset. 
 Neither gold nor platinum are soluble in either sulphuric, nitric, 
 or hydrochloric acid ; but they readily dissolve in a mixture of 
 nitric and hydrochloric acids, yielding the chlorides AuCl 3 and 
 PtCl 4 respectively. Comparatively few simple salts containing 
 gold or platinum are known, and their compounds generally are 
 characterised by the extreme readiness with which they are reduced 
 to the metallic state. 
 
 H 2 S or (NH 4 ) 2 S gives, with solutions of AuCl 3 and PtCl 4 , pre- 
 cipitates of the sulphides. In the case of platinum the precipita- 
 tion is slow. Auro-auric sulphide, Au 2 S,Au 2 S 3 , or AuS (black), 
 and platinic sulphide, PtS 2 (also black) 
 
 8AuCl 3 + 9H 2 S + 4H 2 O = 24HCl + H 2 SO 4 + 2(Au.,S,Au 2 S 3 ), orSAuS 
 PtCl 4 + 2H 2 S = 4HC1 + PtS 2 
 
 If the gold solution be boiling, the reduction goes further, and 
 metallic gold is thrown down 
 
 8AuCl 3 + 3H 2 S + I2H 2 O - 24HC1 + 3H 2 SO 4 + 8Au 
 
 The sulphides of both metals are insoluble in HC1 and HNO 3 , 
 but dissolve in aqua regia. 
 
 AuS dissolves in ammonium sulphide and in sulphides of the 
 alkalies, forming thio-aurates. PtS 2 , unmixed with other sulphides, 
 does not dissolve in ammonium sulphide ; but in the presence of 
 other sulphides of the group, it is partially dissolved.* From the 
 
 * Owing to the fact that PtS 2 is only partially dissolved by ammonium 
 sulphide, if platinum happened to be present in a mixture which was under- 
 going systematic analysis, a portion of the sulphide would pass into Group II. 
 Division 2, along with As, Sb, Sn, Au ; and a part of it would remain in 
 Division i, along with Hg, Bi, Cu, Cd, being ultimately found with the HgS, 
 insoluble in HNO 3 . For this reason it is more advantageous to remove 
 platinum (and also gold) before the separation of Group II. is commenced, 
 as explained on p. 107. 
 
106 Qualitative Analysis. 
 
 solution in both cases, the sulphides are reprecipitated on the 
 addition of HC1. 
 
 KHO. Neither the hydroxides nor carbonates of the alkalies 
 give any precipitate with moderately dilute solutions of AuCl 3 or 
 PtCl 4 . 
 
 From a concentrated solution of AuCl 3 , NH 4 HO gives an 
 orange-red precipitate of fulminating gold, (NH 3 ) 2 ,Au L> O a ; while 
 KHO produces a brown precipitate of hydrated auric oxide, 
 Au 2 O 3 ,3H 2 O, or Au(HO) 3 . This precipitate is soluble in excess of 
 potash, yielding potassium aurate, K 2 O,Au 2 O 3 , or KAuO 2 . 
 
 Both AuCl 8 and PtCl 4 form double salts with alkaline chlorides 
 (rhloro-aurates and chloro-platinates). Ammonium chloro-aurate, 
 NH 4 Cl,AuQ 3 , or NH 4 AuCl 4 , is soluble in water ; while ammonium 
 chloro-platinate, 2NH 4 Cl,PtCl 4 , or (NH 4 ) 2 PtCl 6 , is moderately 
 insoluble, and is produced by precipitation when ammonium 
 chloride is added to a moderately strong solution of PtCl 4 . 
 
 Gold and platinum are readily reduced from their compounds 
 (more especially gold) and precipitated in the metallic State, 
 and their most characteristic reactions are based upon this fact. 
 Thus 
 
 1. Ferrous sulphate gives a brown precipitate of metallic gold ; 
 in weak solutions a bluish coloration 
 
 AuCl 3 + 3FeSO 4 = Fe,(SO 4 ) 3 + FeCl 3 + Au 
 
 With platinum the action only takes place on prolonged boiling 
 3PtCl 4 + !2FeSO 4 = 4Fe 2 (SO 4 ) 3 + 4FeCl 3 + sPt 
 
 2. Oxalic acid, on being gently warmed with AuCl 3 , causes the 
 deposition of the metal either as a scaly precipitate or as a 
 coherent gold film upon the glass, according to the conditions of 
 the experiment. Platinum is not reduced by oxalic acid 
 
 2AuCl 3 + 3C 2 H 2 O 4 = 6CO 2 + 6HC1 + 2Au 
 
 3. Potassium nitrite reduces AuCl 3 , being itself converted into 
 nitrate 
 
 2AuCl 3 + 3H 2 O + 3KNO 2 = 3KNO 3 + 6HC1 + 2Au 
 
 With platinum no precipitate forms at first, but on standing, yellow 
 crystals are deposited of a double nitrite of potassium and platinum 
 (one of the few oxysalts of platinum known) containing platinum in 
 \bzplatinous condition, 2KNO 2 ,Pt(NO 2 ) 2 , or K 2 Pt(NO 2 ) 4 * 
 
 PtCl 4 + 5KN0 2 + H 2 = K 2 Pt(N0 2 ) 4 + 2KC1 + 2HC1 + KNO S 
 
 4. Stannous chloride gives with AuCl 3 a precipitate or colora- 
 tion (depending upon concentration) varying in colour from 
 reddish-brown to purple. The compound is known as purple of 
 
 * The platino-nitrites are remarkable in that the platinum they contain 
 does not answer to the ordinary tests for that metal, just as the iron in ferro- 
 cyanides is not detected by the ordinary reactions for iron. 
 
Separation of the Metals of Group II. ~ Division 2. 107 
 
 Cassius, and its composition is not known with certainty. The 
 presence of a small quantity of stannic chloride (such as is always 
 present in a solution of stannous chloride except when quite 
 freshly made) facilitates the production of the purple. 
 
 With PtCl 4 a brown colour is produced, by the reduction of the 
 platinum chloride to platinous chloride, PtCl 2 . 
 
 In analysis, when gold and platinum are present, it is prefer- 
 able to remove them before the precipitation of Group II. by 
 sulphuretted hydrogen. The gold is precipitated in the metallic 
 state by oxalic acid, and the solution is evaporated down with 
 ammonium chloride, which causes the precipitation of the platinum 
 as ammonium platinum chloride, 2NH 4 Cl ; PtCl 4 . 
 
 SEPARATION OF THE METALS OF GROUP II. DIVISION 2. 
 
 The separation of Group II. from Groups III., IV., and V. 
 depends upon the precipitation of their sulphides from acid 
 solutions. 
 
 The separation of Division I from Division 2 is based upon 
 the solubility of the sulphides of the latter in ammonium sulphide 
 or in caustic alkalies. 
 
 The separation of the metals As, Sb t and Sn from each other is 
 based upon 
 
 (1) The insolubility of arsenic sulphide in hydrochloric acid .or 
 its solubility in ammonium carbonate), whereby arsenic is separated 
 from Sb and Sn. 
 
 (2) The solubility of metallic tin, and insolubility of antimony, 
 in hydrochloric acid. 
 
 The solution, if neutral or alkaline, is acidified with HC1,* and 
 sulphuretted hydrogen passed through (as described on p. 88) until 
 the precipitation of the metals of Group II. is complete. The pre- 
 cipitate is thoroughly washed (in order to make the separation from 
 Groups III., IV., and V. complete), and is then transferred to a 
 small beaker, and gently warmed with yellow ammonium sulphide 
 for a few minutes. f The liquid is then filtered. The residue con- 
 sists of the undissolved sulphides of the metals of Group II., 
 Division i. 
 
 * If the solution under examination is alkaline, it may contain thio salts of 
 As, Sb, or Sn ; the addition of HC1 will result in the precipitation of the sul- 
 phides of these metals. If it is neutral, basic salts of antimony might be pre- 
 cipitated at first, but redissolve on warming with a slight excess of the acid. 
 
 t Ammonium sulphide dissolves CuS to a slight extent (see Reactions, p. 85), 
 hence, if this element is present, a small quantity of it will find its way into the 
 solution along with As, Sb, and Sn. 
 
 , 
 
io8 Qualitative Analysis. 
 
 * 
 
 The solution contains the thio salts of As, Sb, and Sn. It should be some- 
 what diluted, and hydrochloric acid added drop by drop until the sul- 
 phides are completely reprecipitated ; then filtered and washed. The 
 precipitate is then transferred to a boiuig-tube with a small quantity 
 of HC1, and boiled for a few momentj until H 2 S is no longer given 
 off. It is then diluted and filtered. 
 
 i 
 
 The residue consists of The solution. Pour a few drops upon a piece 
 
 arsenic sulphide and of platinum foil, and add a fragment of zinc. 
 
 sulphur. A black stain indicates Sb. If antimony is 
 
 present, place a strip of zinc along with the 
 
 Confirm by dissolving in platinum foil in the remainder of the solution, 
 
 HC1 with a crystal of un til all the antimony and tin are thrown 
 
 KC1O 3 , and applying down. Collect the deposit and boil it with 
 
 ; special reactions for | st Ha and filter< Test the fihrate for 
 
 mTn'so'r Reinsch^ 
 
 APPENDIX TO CHAPTER IX. 
 
 THE RARE METALS OF GROUP II. 
 
 Four of these metals belong to Division I of this group, their 
 sulphides being insoluble in ammonium sulphide ; these are 
 
 Ruthenium, Rhodium, Palladium, and Osmium. 
 
 The remaining members form sulphides which are soluble in 
 ammonium sulphide, and they therefore belong to the antimony, 
 arsenic, and tin subdivision, namely 
 
 Iridium, Tellurium, Selenium, Molybdenum. 
 
 The composition of the precipitates which are thrown down by 
 the group-reagent, is the following : 
 Ru 2 S 3 | 
 
 PdS 3 [l ns l u bl e m ammonium sulphide. 
 OsS } 
 
 T r i 3 1 
 
 6 2 > Soluble in ammonium sulphide. 
 
 MoS 3 J 
 
 * In order to ascertain the condition of oxidation in which the arsenic 
 originally existed in the substance under analysis, special tests must be applied 
 to the solution before it has been exposed either to reducing or oxidising 
 influences, as in the case of iron. 
 
The Rare Metals of Group II. 109 
 
 The metals of the first division, together with iridium in the 
 second section, belong to the natural family of elements known as 
 \hzplatiniim metals, because they all occur associated together in 
 platinum ore. Of these, ruthenium and rhodium are the most 
 rare.* 
 
 Palladium. This metal, in the elemental state, is readily dis- 
 tinguished from all the others of the platinum group by its ready 
 solubility in warm nitric acid. The other platinum metals are 
 unacted upon by any ordinary acid* Aqua regia is without action 
 upon rhodium and iridium; it acts with slowness upon ruthenium, 
 and readily dissolves platinum, forming the chloride, while it con- 
 verts osmium into the tetroxide. 
 
 When palladium is dissolved in nitric acid, the compound formed 
 is palladious nitrate, Pd(NO 3 ) 2 . If the solution be diluted with 
 water, especially if the amount of free acid present 'is only small, a 
 brown-coloured precipitate is produced, consisting of a basic 
 nitrate. Palladious sulphide, PdS, produced by the group reagent, 
 is black. 
 
 Mercuric cyanide, HgCy.,, gives a yellowish precipitate of palla- 
 dious cyanide, PdCy 2 ; slightly soluble in hydrochloric acid, readily 
 soluble in ammonia and in potassium cyanide. The precipitate is 
 distinguished from other metallic cyanides by the reaction common 
 to all palladium salts, namely, that when heated they decompose, 
 leaving spongy metallic palladium. 
 
 Potassium iodide, KI, gives a characteristic black precipitate of 
 palladious iodide, PdI 2 . Palladious salts are readily reduced to the 
 metallic state either by heat or by the action of reducing agents. 
 
 Osmium. The compound of this rare element which is most 
 commonly met with is the so-called osmic add, which is employed 
 in the preparation of microscopic sections of animal tissues. This 
 compound is the tetroxide, OsO 4 , or osmic anhydride. 
 
 It is characterised by its extremely low melting-point (about 40) 
 and boiling-point (100), and by the peculiar and irritating vapour 
 which it gives. The vapour exerts a most injurious effect upon the 
 eyes, and is extremely poisonous. 
 
 This vapour is given off when any osmium compound is heated 
 with nitric acid, and serves as a characteristic test for the element. 
 
 The tetroxide is soluble in water, giving a neutral solution which 
 has powerful oxidising properties j it bleaches indigo, liberates 
 iodine from potassium iodide, and oxidises ferrous sulphate and 
 alcohol, the osmium compound being reduced to the state of 
 hydrated dioxide, OsO 2 ,2H 2 O (or Os(HO) 4 ), which is thrown down 
 as a black precipitate. 
 
 Sulphurous acid, or a sulphite added to the solution, produces a 
 series of colour-changes from yellow to green and lastly blue, the 
 colour of the osmious sulphite, OsSO 3 , which then separates out. 
 
 Iridium. This metal differs from platinum in not being dis- 
 solved by aqua regia. When heated with a fused mixture of sodium 
 
 * See footnote on p. 70. 
 
1 1 o Qualitative A nalysis. 
 
 nitrate and hydroxide in a silver vessel, the metal is oxidised to the 
 trioxide, Ir 2 O 3 ; and on treating the residue with aqua regia, a 
 dark-coloured solution is obtained of the double chloride, 2NaCl, 
 IrCl 4 . 
 
 This solution may be used for the following reactions : 
 
 When sulphuretted hydrogen is passed into the solution, the 
 brown colour disappears owing to the reduction of the IrCl 4 to 
 IrCl 3 (or Ir 2 Cl 6 ) and simultaneous precipitation of sulphur. The 
 further passage of the gas throws down the trisulphide, Ir 2 S 3 , as a 
 dark brownish precipitate. 
 
 The double chlorides, 2NH 4 Cl,IrCl 4 and 2KCl,IrCl 4 , are pre- 
 cipitated by the addition of ammonium chloride and potassium 
 chloride respectively. They are both dark brownish-red precipitates, 
 insoluble in strong solutions of the precipitants. 
 
 By the action of reducing agents (e-g. ferrous or stannous salts, 
 nitrites, etc.), these double chlorides are reduced, giving similar 
 compounds containing the lower chloride of iridium, and having 
 the composition expressed by the formulae 3NH 4 Cl,IrCl 3 and 
 3KCl,IrCl 3 respectively. The solution is at the same time decolor- 
 ised, and the double chloride gradually deposits. When caustic alkali 
 is added to a solution of iridic chloride (or the double sodium salt), 
 and the mixture heated, the solution assumes a deep blue colour 
 owing to the precipitation of iridic hydroxide, Ir(HO) 4 , which when 
 separated appears as an indigo-blue powder. This reaction serves 
 to distinguish iridium from platinum. 
 
 Tellurium and Selenium. These two elements belong to 
 the same natural family as sulphur, which in some respects they 
 closely resemble. They both lie on the borderland between the 
 non-metals and the true metals. While selenium forms no stable 
 compounds in which it forms the positive constituent, tellurium 
 exhibits feeble basic properties ; the oxide, TeO 2 , being both an 
 acid-forming and a salt-forming oxide. 
 
 Both elements form hydrogen compounds corresponding to 
 sulphuretted hydrogen, and closely resembling it in properties. 
 Thus, when compounds of either element are heated on charcoal 
 with sodium carbonate, sodium telluride, Na 2 Te, or selenide, 
 Na 2 Se, is formed. These are decomposed by acids, with forma- 
 tion of the respective hydrogen compounds, whose odour is even 
 more offensive than that of sulphuretted hydrogen. Or if the 
 sodium compounds are moistened with water upon a silver coin, 
 a black stain is produced of silver telluride or selenide. 
 
 Sulphuretted hydrogen gives, with tellurous compounds, a brown 
 precipitate of tellurous sulphide, TeS 2 , but with selenious compounds 
 the precipitate consists of selenium and sulphur, selenious sulphide 
 being too unstable to exist. Both precipitates, however, are soluble 
 in ammonium sulphide. With selemV compounds (selenates) sul- 
 phuretted hydrogen gives no precipitate until the selenate is 
 reduced to selenite. 
 
 Barium chloride, BaCl 2 , gives, with solutions of selenates, a 
 white precipitate of barium selenate, BaSeO 4 . This precipitate is 
 
The Rare Metals of Group //. 1 1 1 
 
 distinguished from barium sulphate in that, when boiled with hydro- 
 chloric acid, it is converted into barium selen/te, which is soluble. 
 
 Sulphurous acid reduces both tellurous and selenious com- 
 pounds, with precipitation of the element ; tellurium being thrown 
 down as a black powder, while selenium is precipitated in the form 
 of the brick-red amorphous variety. 
 
 Both elements burn in the .air or in oxygen, with a blue 
 flame, giving rise to the dioxide ; in the case of selenium, the 
 combustion is accompanied by a smell of putrid horseradish. 
 
 Molybdenum. The compound of this element most commonly 
 met with is ammonium molybdate, (NH 4 ) 2 MoO 4 . 
 
 From a strong aqueous solution of this salt, hydrochloric or 
 nitric acid gives a white precipitate of molybdic acid, H 2 MoO 4 , 
 soluble in excess of acid. 
 
 When sulphuretted hydrogen is passed into an acidulated solu- 
 tion, the solution first assumes a blue colour, which turns green as 
 the dark-brown sulphide is precipitated. The precipitate produced 
 is the trisulphide, MoS 3 , analogous to the trioxide. It is soluble 
 in ammonium sulphide (and alkaline sulphides generally), forming 
 thio-molybdates. These, like the corresponding antimony and 
 arsenic compounds, are decomposed by hydrochloric acid, with the 
 reprecipitation of the tri-sulphide. 
 
 The reaction by which molybdic acid is most readily identified, 
 is the formation of the yellow precipitate of ammonium phospho- 
 molybdate, by adding sodium phosphate to a nitric acid solution 
 of molybdic acid. 
 
CHAPTER X. 
 REACTIONS OF THE METALS OF GROUP I. 
 
 Silver, Lead, Mercury (as Mercurous Mercury). 
 
 THE classification of these three elements into a group is owing to 
 common properties possessed by their chloricfes. In most other 
 respects they are greatly dissimilar, and in the "natural" classifi- 
 cation of the elements they take their places in three different 
 families. 
 
 Silver, Ag. 
 
 DRY REACTIONS. Compounds of silver, when heated on char- 
 coal with sodium carbonate in the reducing flame, yield metallic 
 silver ; which, being non-oxidisable, is not accompanied by any 
 oxide incrustation upon the charcoal. The metal, however, is 
 slightly volatile in the blowpipe flame, and sometimes a faint red- 
 brown incrustation is thus obtained. 
 
 The reduced metal may be removed to a watch-glass, dissolved 
 in nitric acid, and precipitated as chloride. 
 
 WET REACTIONS. The salts of silver are derived from the 
 monoxide Ag 2 O. Of the common salts, the nitrate is readily 
 soluble, the acetate and sulphate sparingly soluble, in water. (The 
 chlorate, nitrite, and fluoride are also soluble. The solubility of 
 the fluoride is noteworthy in view of the insolubility of the other 
 halogen salts.) 
 
 A characteristic property exhibited by a number of the silver 
 salts is their readiness to form soluble compounds with ammonia, 
 which are either double salts of silver and ammonium, such as 
 NH 4 NO 3 ,AgNO 3 , or belong to the class of salts known as metall- 
 amnionitun compounds, or metallo-amines. Thus, silver nitrate 
 absorbs ammonia gas and yields the compound AgNO 3 ,2NH 3 . 
 The same substance is obtained on adding ammonia solution to 
 silver nitrate (a neutral salt) until the precipitate first formed is 
 dissolved. 
 
Group I. U3 
 
 HC1, and soluble chlorides, give a white amorphous precipitate 
 of silver chloride, AgCl, which, on being warmed or stirred, becomes 
 granulated in appearance, and very quickly settles. On exposure 
 to light, the white compound assumes a slate colour or drab tint, 
 which gradually deepens to a violet, and finally appears brown or 
 black. 
 
 Silver chloride is quite insoluble in water, but soluble to a 
 slight extent in strong HC1 ; dilution causes the complete precipi- 
 tation. It readily dissolves in ammonia, forming the compound 
 2AgCl,3NH 3 . Nitric acid decomposes this compound, causing 
 the reprecipitation of AgCl, which is practically insoluble in that 
 acid 
 
 2AgCl, 3 NH 3 + 3 HN0 3 = 3NH 4 NO 3 + 2AgCl 
 
 Silver chloride is soluble also in KCy, being first converted into 
 silver cyanide, which dissolves in excess of KCy, forming the double 
 cyanide KCy,AgCy. It also dissolves in sodium thiosulphate, with 
 the formation of a double thiosulphate ; thus 
 
 AgCl + Na 2 S,0 3 = NaCl + NaAgS,O 3 
 
 When boiled with potassium hydroxide, silver chloride is con- 
 verted into silver oxide (black) 
 
 2AgCl + 2KHO = 2KC1 + H.,0 + Ag 2 O 
 
 Silver chloride melts without decomposition at 451, and re- 
 solidifies to a horny mass (horn silver). 
 
 [Reactions with bromides, iodides, and cyanides are described 
 under the respective acids.] 
 
 KHO, NaHO, or NH 4 HO gives a greyish-black precipitate of 
 silver oxide, Ag.,0. Insoluble in excess of the caustic alkalies, but 
 readily soluble in ammonia. If the silver solution is acid, ammonia 
 gives no precipitate, but forms a soluble double salt (see above). 
 
 K CO or Na 2 CO s gives a white precipitate of silver carbonate, 
 Ag 2 CO 3 . Insoluble in excess of the precipitant ; soluble in am- 
 monium carbonate, ammonia, and nitric acid. 
 
 H. 2 S or (NH 4 ) 2 S produces a black precipitate of silver sulphide, 
 Ag 2 S. Insoluble in dilute acids, except boiling dilute nitric acid, 
 which converts it into nitrate. The H 2 S, which by double de- 
 composition is set free, is acted upon by the nitric acid, with the 
 precipitation of sulphur and evolution of nitric oxide (see Lead 
 reactions, p. 80). 
 
 Silver sulphide is insoluble in ammonia, ammonium sulphide, or 
 potassium .sitlphide. 
 
H4 Qualitative Analysis. 
 
 Reduction of Silver Salts to the Metallic State. 
 
 Silver compounds (more especially the ammoniacal solution of 
 silver oxide) are readily reduced with precipitation of metallic silver 
 (which often deposits as a coherent mirror) by certain organic 
 substances, as sugar, tartrates, aldehydes, etc. Many inorganic 
 salts also, which act as reducing agents, precipitate metallic silver 
 from solutions of its salts, e.g. ferrous sulphate 
 
 3 AgN0 3 + 3FeS0 4 = Fe(NO 8 ) 3 + Fe 2 (SO 4 ) 3 + 3 Ag 
 
 Many metals are capable of reducing silver compounds. Men- 
 tion may be made of the three metals, zinc, iron, and mercury. 
 
 If a strip of zinc be immersed in a solution of silver nitrate, 
 crystals of metallic silver are seen to grow out from the surface of 
 the zinc in the manner of the familiar " lead tree." The reducing 
 action of zinc is often employed in the laboratory to convert pre- 
 cipitated silver chloride into metallic silver. The same action of 
 iron is used in one of the metallurgical processes for the extraction 
 of silver ; the reaction in both cases is similar 
 2AgCl + Zn = ZnCl 2 + 2Ag 
 2AgCl + Fe = FeCl 2 + 2Ag 
 
 Mercury in contact with silver chloride similarly precipitates 
 metallic silver, with formation of mercurous chloride 
 2AgCl + 2Hg = Hg 2 Cl 2 + 2 Ag 
 
 None of these reductions are made use of in qualitative analysis ; 
 the latter reaction, however, with mercury has an important bearing 
 upon the separation of silver from mercurous mercury in the usual 
 course of analysis. The separation of these metals is based upon 
 the solubility of silver chloride in ammonia, mercurous chloride 
 being at the same time converted into the black mercurous 
 ammonium chloride 
 
 (i) Hg 2 Cl 2 + 2NH 4 HO = NH 2 (Hg 2 )Cl + NH 4 C1 + 2 H,O 
 
 This mercurous compound, however, in the presence of silver 
 chloride and excess of ammonia, passes into the corresponding and 
 more stable mercuric salt, NH 2 HgCl, while the mercury it thus 
 loses, reduces a portion of the silver chloride ; thus 
 
 (2) NH 2 (Hg 2 )Cl + 2AgCl = NH 2 HgCl + HgCl 2 + 2Ag * 
 
 (3) HgCl 2 + 2NH 4 HO = NH 2 HgCl + NH 4 C1 + 2H 2 O 
 
 * It is possible, but not probable, that the following reaction goes on 
 simultaneously : 
 
 2 NH 2 (Hg 2 )Cl + 2 AgCl = 2 NH 2 HgCl + Hg 2 Cl 2 + 2 Ag 
 
 The Hg 2 Cl 2 , being reconverted by the excess of ammonia into mercurous 
 ammonium chloride, would then be in a position to react upo%ia fresh pro- 
 portion of silver chloride, so that the action would be continued U^iil the whole 
 of the silver became reduced ; this, however, does not appea; to be the case. 
 
Separation of the Metals of Group L 115 
 
 The 'entire change, therefore, may be summed up in the single 
 equation following : 
 
 In the practical separation, therefore, a small quantity of silver, 
 in presence of a large proportion of mercury, might escape detection 
 by being entirely precipitated and left upon the filter along with the 
 mercury compound. It will be evident that if the action of the 
 ammonia be allowed to continue sufficiently long, it would be 
 possible to precipitate the whole of the silver if the quantity of 
 mercury present only slightly exceeded the equivalent proportion. 
 
 Lead, Mercury. 
 
 The reactions of these metals have already been considered in 
 connection with the metals of Group II., Division i, pp. 76, 79. 
 
 SEPARATION OF THE METALS OF GROUP I. 
 
 The separation of these metals from the other groups, depends 
 upon the insolubility of their chlorides in cold water, and con- 
 sequently their precipitation by hydrochloric acid. The separation 
 of the metals from each other is based upon 
 
 (1) The solubility of lead chloride in hot water. 
 
 (2) The solubility of silver chloride in ammonia. 
 
 To the solution add moderately dilute hydrochloric acid drop 
 by drop, until a slight excess beyond what is required for complete 
 precipitation has been added. Gently warm the mixture,* and 
 after again cooling it, filter. The filtrate contains Groups II., III., 
 IV., and V. 
 
 The precipitate, consisting of PbCl 2 , AgCl, and Hg 2 Cl 2 , is thoroughly 
 washed in cold water, and then boiled with water (or washed while 
 the filter with boiling: water) and filtered. 
 
 The filtrate contains 
 PbCl 2 , which deposits 
 in white needle-shaped 
 crystals on cooling. 
 
 Confirm by special test, e.g. 
 the formation of PbCrO 4 . 
 
 The residue is treated, while still upon the 
 filter, with a small quantity of ammonia, 
 which dissolves the AgCl, and converts the 
 white Hg 2 Cl 2 into black NH 2 (Hg 2 )Cl.t 
 
 The solution yields a white precipitate of 
 AgCl upon being acidified with HNO 3 . 
 
 The black residue may be dissolved in a little 
 aqua regia, and (after being nearly neutralised) 
 the mercury precipitated upon metallic copper. 
 
 * See footnote on p. 107. 
 
 f For conditions under which this separation is incomplete, see p. 114. 
 
 (JL 
 
1 1 6 Qualitative A nalysis. 
 
 SYSTEMATIC SEPARATION OF THE GROUPS.* 
 
 (1) Separation of Group I. By precipitation with hydro- 
 chloric acid (p. 115). 
 
 The precipitate is thoroughly washed with cold water, and 
 examined by the method on p. 115. 
 
 (2) Separation of Group II. The filtrate from Group I. is 
 diluted with water (any precipitation of basic bismuth compounds 
 is to be disregarded, as the subsequent action of SH 2 will convert 
 them into the sulphide), and a moderately slow stream of sulphu- 
 retted hydrogen passed through. The changes which occur during 
 this precipitation should be carefully watched and noted (see 
 Lead, Mercury, Arsenic reactions). Complete precipitation must 
 be ensured (p. 88. See also Arsenic, p. 93). The precipitate is 
 then treated as described on p. 107. 
 
 (3) Separation of Groups IIlA. and IIlB. The filtrate 
 from Group 1 1. is boiled until the sulphuretted hydrogen is entirely 
 expelled. Two or three drops of strong nitric acid are added, and 
 the liquid again boiled for a few minutes, in order to oxidise com- 
 pounds (e.g. iron or chromium) which might have become reduced 
 by the action of sulphuretted hydrogen. The liquid is then carefully 
 evaporated to dryness f in a porcelain dish, and if it shows signs 
 
 * It may not be out of place at this point to impress once more upon 
 the student the supreme importance of cleanliness, neatness, and method in 
 analysis. Slipshod and careless work inevitably ends in disappointment and 
 failure. All the utensils (test-tubes, beakers, etc. ) must be scrupulously clean, 
 being always rinsed once or twice with distilled water after having been washed. 
 Too much care cannot be paid to securing complete precipitation in every 
 separation, or to the thorough washing of precipitates. The time spent in 
 securing these results is never lost, whereas the slovenly neglect of these points 
 may, and frequently does, involve a repetition of the analysis, and a corre- 
 sponding sacrifice of time. A little experience will enable the student to keep 
 two or three operations going on at the same time. Thus, while one precipitate 
 is being washed, another precipitation can be made, and the filtration of this 
 second separation can be carried on simultaneously with the other. While 
 these two are filtering, other tests or special reactions can be made. To do 
 this successfully, however, it must be done methodically, and (especially in 
 cases where the analysis has to be interrupted for a time) the various precipi- 
 tates and solutions should be labelled. Before making a group separation, it 
 is well to make a preliminary test in a smull portion of the solution, and if the 
 group-reagent gives no precipitate, the operation may then be omitted with 
 the bulk of the solution. 
 
 t " Carefully evaporate to dryness" does not mean that the student may 
 place a lamp under the dish and go away and leave it. The operation must 
 be watched, and as the liquid becomes more and more concentrated the lamp 
 flame must be lowered. As the residue begins to dry round the edges, he 
 should notice if it shows signs of charring, for if there is no organic compound 
 present it is better to avoid strongly heating the dried residue, as the operation 
 
Systematic Separation of the Groups. i \ 7 
 
 of charring (owing to the presence of citrates, tartrates, etc.), the 
 residue is cautiously heated until such organic matter is completely 
 decomposed. The residue is moistened with a few drops of strong 
 hydrochloric acid, water is then added, and the mixture boiled. 
 The remaining residue, consisting of silica and carbon, is filtered 
 off.* 
 
 A small portion of the solution is next tested for phosphoric 
 acid by means of ammonium molybdate (p. 63). To the main 
 portion of the solution ammonium chloride is added in consider- 
 able quantity, and the liquid heated to boiling. Ammonia is 
 carefully added until precipitation is complete. 
 
 If phosphoric acid is absent, this precipitate is examined for 
 the metals of Group II I A. by the method on p. 49. 
 
 If phosphoric acid is present,\hz precipitate is treated as explained 
 on p. 68. The filtrate is then saturated with sulphuretted hydro- 
 gen,! gently warmed, and filtered. The precipitate is examined for 
 the metals of Group 1 1 IB. according to the plan on p. 62. 
 
 (4) Separation of Group IV. The filtrate from Group 1 1 IB. 
 is boiled briskly, with the addition of a little hydrochloric acid, 
 to decompose ammonium sulphide and expel all the sulphuretted 
 hydrogen. Any precipitated sulphur is removed by filtration. The 
 solution is then rendered alkaline by addition of ammonia, and 
 ammonium carbonate added until precipitation is complete. The 
 solution may be warmed, but must not be boiled (see p. 34). The 
 precipitate is examined for the metals of Group IV., as indicated 
 on p. 35. \ 
 
 (5) Detection of Metals of Group V. The filtrate from 
 Group IV. is tested for magnesium, potassium, and sodium, J as 
 described Jon p!\25. 
 
 is likely to render certain oxides (e.g. Fe 2 O 3 , Cr 2 O 3 , A1 2 O 3 ) very difficult of 
 solution in hydrochloric acid** 
 
 * If by a preliminary test upon a small portion of the solution (by evapo- 
 rating it upon a platinurtf capsule) it is found that there is no silicious or 
 carbonaceous residue, the mail bulk of the liquid need not be evaporated 
 down. 
 
 f Or ammonium sulphide may be used for the precipitation (see, in this 
 case, the action of ammonium sulphide on nickel sulphide). 
 
 I Ammonium obviously canndtjpe tested for in this solution, since ammo- 
 nium salts have been frequently introduced during the course of analysis. The 
 test for this "metal" must therefore be made in a portion of the original 
 solution. 
 
Qualitative A nalysis. 
 
 , 
 
 APPENDIX TO CHAPTER X. 
 
 THE RARE METALS OF GROUP I. 
 7 
 
 The rare metals of this group are thallium and tungsten. 
 
 The compounds that are precipitated by the group-reagent 
 being thallous chloride, T1C1, and tungstic acid, H 2 WO 4 , in a 
 >? hydrated condition. 
 
 f+ Thallium. Two classes of compounds of thallium are known, 
 
 namely, thallous salts (derived from thallous oxide, T1 2 O) and 
 thallic salts (derived from thallic oxide, T1 2 O 3 ). In the thallous 
 compounds (which are the more stable of the two classes), the 
 element shows a resemblance to the alkali metals on the one hand, 
 and to lead on the other. For example, thallous hydroxide, T1HO, 
 is soluble in water, giving a strongly alkaline solution ; hence 
 KHO, NaHO, and NH 4 HO give no precipitate with thallous 
 solutions. Thallous carbonate, T1 2 CO 3 is moderately soluble in 
 water, and therefore is only precipitated by carbonates of the alka- 
 lies from concentrated solutions (compare Thallic salts). Again, 
 platinic chloride gives with thallous solutions an orange-yellow 
 precipitate of thallous platinic chloride, 2TlCl,PtCl 4 . On the 
 other hand, its resemblance to lead is specially seen in the chloride, 
 iodide, and chromate. 
 
 Hydrochloric acid (or soluble chlorides) gives a white curdy 
 precipitate of T1C1. Like lead chloride it is slightly soluble in 
 cold water, and dissolves more readily in hot water (100 parts 
 of water at 16 dissolve 0-265 parts, while at 100 1*427 parts 
 are dissolved). 
 
 It is at once distinguished from lead chloride by the fact 
 that it is readily soluble in strong sulphuric acid, forming soluble 
 thallous sulphate. It is distinguished from silver chloride in not 
 being soluble in ammonia, and in not changing colour on exposure 
 to light. Potassium iodide gives a bright golden-yellow precipitate 
 of thallous iodide, Til ; while potassium chromate throws down 
 thallous chromate, Tl 2 CrO 4 , also (like the lead compound) yellow 
 in colour. 
 
 Thallous sulphide (black) is readily soluble in mineral acids ; 
 hence sulphuretted hydrogen only gives complete precipitation in 
 acetic acid solutions (or in alkaline solutions). With ammonium 
 sulphide precipitation is complete, but the precipitate rapidly under- 
 goes atmospheric oxidation into the sulphate, which passes into 
 solution. Thallic salts are readily distinguished from thallous 
 compounds. Thus the chloride, T1C1 3 , is soluble in water ; hence 
 hydrochloric acid gives no precipitate. Potassium chromate, simi- 
 larly, gives no precipitate. With potassium iodide, a mixture of 
 thallous iodide and iodine is precipitated, while with sulphuretted 
 hydrogen the compound is reduced to the thallous state, with 
 precipitation of sulphur. 
 
The Rare Metals of Group I. 119 
 
 The caustic alkalies, and also the alkaline carbonates, give a 
 brown precipitate with* thallic solutions, consisting of an oxy- 
 hydroxide, TIO(HO). 
 
 Thallium, like lead, is easily reduced from solutions of its salts 
 by metallic zinc, the thallium being deposited upon the zinc in a 
 spongy metallic state. 
 
 When introduced into a Bunsen flame upon a platinum wire, 
 thallium salts impart a characteristic brilliant green colour, which, 
 when viewed through a spectroscope, is seen to consist of one 
 bright green line. 
 
 Tungsten. The compound of this element most commonly 
 met with in commerce is sodium tungstate. The formula repre- 
 senting the composition of the normal s&\\. is Na 2 WO 4 2H 2 O.* The 
 tungstates of the alkalies are alone soluble.f Hydrochloric acid 
 gives a white precipitate of hydrated tungstic acid. When air-dried, 
 this precipitate has the composition WO(HO) 4 , or H 2 WO 4 ,H 2 O. 
 When dried over sulphuric acid, it loses H 2 O and is converted 
 into the normal acid. The precipitate is insoluble in excess of 
 acid. 
 
 The most characteristic reaction for compounds of tungsten is 
 the blue colour given when either metallic zinc or tin or stannous 
 chloride is added to a solution of a tungstate, and the solution 
 strongly acidified with hydrochloric acid. On the addition of 
 stannous chloride a pale yellow precipitate is produced, which, on 
 the addition of hydrochloric acid and gently warming, turns deep 
 blue.* 
 
 Sulphuretted hydrogen gives no precipitate in acid solutions, 
 but causes the solution to assume a blue colour.^ 
 
 Tungsten trisulphide, WS 3 , is a thio-anhydride, forming soluble 
 thio-tungstates with alkaline sulphides ; hence, ammonium sulphide 
 gives no precipitate with solutions of alkaline tungstates. But on 
 acidulating the mixture with hydrochloric acid, the sulphide is 
 thrown down as a dark brown precipitate. 
 
 * There is quite a number of sodium tungstates, which may all be regarded 
 as compounds of the normal salt with varying quantities of the trioxide. Of 
 these the so-called metatungstate and paratungstate are articles of commerce 
 
 Sodium metatungstate, Na 2 W4Oi3,ioH 2 O ; or Na 2 WO 4l 3WO3,ioH 2 O 
 Sodium paratungstate, Nai W 12 O 41 ,2i(or 28)H 2 O ; 
 
 or 5Na 2 WO 4 ,7WO 3 ,2i(or 28)H 2 O 
 
 The former of these two salts is the common compound sometimes used for 
 rendering fabrics uninflammable. 
 
 f The oxides of tungsten are all acidic in their characters, and no salts of 
 this element are known in which it functions as the positive or basic constituent. 
 
 J The blue colour is believed to be due to the formation of a lower hydrated 
 oxide. 
 
CHAPTER XI. 
 THE NON-METALS AND THEIR ACIDS. 
 
 ONE of the chief chemical distinctions between metals and non- 
 metals is that, while the oxides of the former exhibit, generally 
 speaking, basic properties, the latter elements give oxides which 
 are acidic in their character. Two exceptions to this generalisation 
 respecting the non-metals are seen in the case of hydrogen (whose 
 oxides are not acidic), and the element oxygen itself. Although 
 the number of non-metals is comparatively very small, the number 
 of acids derived from them is very considerable, owing to the fact 
 that many of these elements give rise to several acids. 
 
 All the elements belonging to this section form compounds with 
 hydrogen ; but only in the case of fluorine, chlorine, bromine, 
 iodine, and sulphur are the hydrogen compounds included among 
 the acids in analytical classification. 
 
 The classification of the acids (or acid radicals) is based, as in 
 the case of the metals, upon the solubility or insolubility of certain 
 of their salts produced by interaction with certain specified reagents. 
 They are in this way divided into certain arbitrary groups, but the 
 mode of treatment differs entirely from that adopted in the case of 
 the metals. The reagents referred to are not employed as " group- 
 reagents " to separate one group of acid radicals from another ; 
 but each acid is separately detected by a special test. The use of 
 the general reagent is in order to ascertain by a single operation 
 the absence, or otherwise, of an entire group, whereby the necessity 
 for applying a number of separate tests may be obviated. 
 
 The study of the special reactions for the non-metals and their 
 acids may therefore be made irrespective of their analytical 
 classification, and it will be more advantageous to postpone the 
 consideration of the classification until a knowledge of the special 
 reactions has been gained, and the student is prepared to undertake 
 the systematic detection of the acids. 
 
The Non-metals and their Acids. 121 
 
 The non-metals and their acids which will be included in this 
 section are the following : 
 
 Chlorine : hydrochloric acid, hypochlorous acid, chloric acid, 
 perchloric acid. 
 
 Bromine : hydrobromic acid, bromic acid. 
 
 Iodine : hydriodic acid, lodic acid. 
 
 Fluorine : hydrofluoric acid, hydrofluo-silicic acid. 
 
 Sulphur : sulphuretted hydrogen, sulphuric acid, sulphurous 
 acid, thio-sulphuric acid. 
 
 Nitrogen : nitric acid, nitrous acid. 
 
 Phosphorus : phosphoric acid, phosphorous acid, hypophos- 
 phorous acid. 
 
 Carbon : carbonic acid, formic acid, oxalic acid, acetic acid, 
 tartaric acid, citric acid, cyanogen, hydrocyanic acid, ferrocyanic 
 acid, ferricyanic acid, cyanic acid, thiocyanic acid. 
 
 Silicon : silicic acid. 
 
 Boron : boric acid. 
 
 In some instances the properties of the elements themselves, in 
 the elemental state, are made use of in analysis ; such, for instance, 
 as in the cases of chlorine, bromine, and iodine. When this is the 
 case, those properties of the elements by means of which they are 
 most readily identified will be studied. With others, such as 
 fluorine, silicon, boron, the properties of the isolated elements 
 have nb analytical bearing, and therefore all description of such 
 elements will be omitted.* 
 
 Certain acids, such as arsenious, arsenic, chromic, have already 
 been discussed under their respective metals. 
 
 The Halogens. 
 
 These elements not only form oxy-acids, but they also combine 
 with hydrogen and yield acids. The salts of these hydrogen acids, 
 e.g. chlorides, bromides, and iodides are known as the haloid salts, 
 or sometimes as the halides. 
 
 Chlorine. 
 
 The properties by which this element is identified are the 
 following : 
 
 It is a pale greenish-yellow gas, having a characteristic suffocat- 
 ing smell, irritating and rapidly attacking the mucous membrane 
 of the nose and throat. It dissolves in water, imparting its own 
 
 * It is presupposed that the student of analysis has already become familiar 
 with the common properties of the non-metallic elements ; and if he has not, 
 they must be sought in text-books of general chemistry. 
 
122 Qualitative Analysis. 
 
 colour to the solution. It combines readily with many metals ; it 
 is a powerful oxidising agent ; it liberates bromine and iodine from 
 bromides and iodides respectively ; it possesses powerful bleaching 
 properties. 
 
 The Liberation of Chlorine from its Compounds. The 
 method usually resorted to in analysis for the liberation of chlorine, 
 depends upon the action of peroxides (manganese dioxide being 
 employed) upon the hydrogen acid, or upon a chloride in the 
 presence of sulphuric acid 
 
 2NaCl + 2H 2 SO 4 + MnO 2 = Na 2 SO 4 + MnSO 4 + 2H 2 O + C1 2 
 
 For this purpose the mixture of the chloride with sulphuric acid 
 and manganese dioxide is gently warmed fn a small flask or test- 
 tube, fitted with a delivery tube, and the volved gas is allowed to 
 pass into water contained in a second test-tube. As the solution 
 of the chlorine by water is not complete, the characteristic smell of 
 that which escapes solution may be noted, and its bleaching action 
 can be seen by introducing a strip of litmus paper into the mouth 
 of the tube. 
 
 The presence of the free chlorine in the water may be detected 
 by the following more delicate tests : 
 
 (a) On adding a few drops of the chlorine water to a solution 
 of potassium iodide, to which a little dilute starch paste has been 
 added, a deep blue coloration results, owing to the liberat^plodine 
 (set free by the chlorine) uniting with the starch. 
 
 (b) A crystal of ferrous ammonium sulphate* [FeSO 4 ,(NH 4 ) 2 - 
 SO 4 ,6H 2 O] is dissolved in water, and a few drops of ammonium 
 thiocyanate added. To this colourless mixture a drop or two of the 
 chlorine water is added. This at once oxidises the ferrous salt to 
 the ferric state, which, in the presence of the thiocyanate, im- 
 mediately gives rise to the wine-red coloration due to ferric 
 thiocyanate. 
 
 Hydrochloric Acid and Chlorides. 
 
 Hydrochloric acid is a colourless gas having a sharp choking 
 smell. It fumes in contact with moist air, is strongly acid, but has 
 no bleaching properties. It is extremely soluble in water, the 
 solution constituting the ordinary reagent. 
 
 Hydrochloric acid is liberated in the gaseous state when 
 
 * This double salt is used instead of ferrous sulphate, as it is less easily 
 oxidised by the air, and therefore is more easily obtained free from ferric 
 compounds. 
 
Hydrochloric Acid and Chlorides. 123 
 
 chlorides (except those of tin, lead, mercury, and silver) are gently 
 heated with strong sulphuric acid 
 
 KC1 + H 2 SO 4 = HKSO 4 4- HC1 
 
 The presence of free hydrochloric acid in a solution containing 
 a soluble chloride may be detected by gently warming the liquid 
 with manganese dioxide (without the addition of sulphuric acid) ; 
 chlorine is evolved, which may be detected by the methods already 
 described 
 
 4HC1 + MnO 2 = MnCl 2 + 2H 2 O + C1 2 
 
 Chlorides are all soluble in water, except those of the metals of 
 Group I. (PbClj being soluble; ia-Jiot water), and certain others 
 which are decomposed by ''water. The chief of these latter salts 
 are the chlorides of (phosphorus, arsenic), antimony, bismuth, and 
 tin. In contact with wd%r these either an oxide or an 
 
 oxychloride of the metal, with Formatior of free hydrochloric acid 
 (see Reactions of these meta 
 
 Silver nitrate, AgNO 3 , giv^s, in solution of chlorides or 
 hydrochloric acid, a white precipitate of silver chloride. Insoluble 
 in nitric acid. Readily soluble in ammonia, even dilute (for 
 further properties, see Silver reactions, p. 113)^ 
 
 AgCl is distinguished from either / Agl by the fact j 
 
 that chlorine water is without action upoi Bromides am 
 
 iodide^^ It may also be distinguished in the following way : 
 
 Reduction "by Zinc. If the washed precipitate of AgCl be 
 mixed with a little very dilute sulphuric acid, and a strip of zinc 
 placed in the mixture, the silver chloride turns grey, owing to its 
 reduction to metallic silver, while zinc chloride passes into solution. 
 This, on treatment with manganese dioxide and sulphuric acid, will 
 yield chlorine. 
 
 Fusion with sodium carbonate converts AgCl into metallic 
 silver and sodium chloride. On treatment with water, chlorine can 
 be liberated from the solution, as in the foregoing. 
 
 Formation of Chromyl Chloride. When a chloride is mixed 
 with potassium dichromate (the two salts being powdered together), 
 and the mixture gently warmed with strong sulphuric acid, a red- 
 brown vapour is disengaged (resembling bromine in colour, but very 
 different in smell) consisting of chromyl chloride, CrO 2 Cl 2 
 4KC1 + K 2 Cr 2 O 7 + 3H 2 SO 4 = 3K 2 SO 4 +3H 2 O + 2CrO 2 Cl 2 * 
 
 * We may regard this action as taking place between chromium trioxide 
 (formed by the action of the sulphuric acid upon the dichromate) and hydro-, 
 chloric acid (from the interaction of the chloride and sulphuric acid) ; thus 
 CrO 3 + 2HC1 = H 2 O + CrO 2 Cl 2 
 
1 2 4 Qualitative A nalysis . 
 
 The gas is decomposed by water or alkaline hydroxides, form- 
 ing in the latter case an alkaline chromate and chloride ; thus 
 
 Cr0 2 CL + 4NH 4 HO = (NH 4 ). 2 CrO 4 + 2NH 4 C1 + 2 H 2 O 
 
 If the reaction be made in a test-tube or small flask fitted with a 
 delivery tube, and the vapour of the chromyl chloride (which is a 
 deep red fuming liquid at ordinary temperatures) be passed into a 
 second test-tube containing ammonia, the above decomposition 
 takes place. The presence of ammonium chromate is seen by the 
 yellow colour which the liquid assumes, and the presence of the 
 chromate is proof of the presence of ji chloride in th* 3 first test-tube 
 or flask. [The presence of th*^ ^^Hksyrnny be Confirmed 
 
 by applying a special test, s "vdrogen 
 
 peroxide.] 
 
 By means of this , chloride in the 
 
 Presence ^/either a broflp$B| ' neither bromine nor iodine 
 
 form similar chromyl jp Pa bromide or iodide is 
 
 treated in this way, bSJ3 ^H^ r esca P es alone. 
 
 Bromine. 
 
 BromiiM vn-red, volatile liquid, which passes into 
 
 a brown-njH -.Unary temperatures. It has a powerful, 
 
 ^irritating sij stnerTTUicous membrane of the nose and 
 
 throat, and caus ^ne eyes to smart. It bleaches, but less easily 
 than chlorine, reproduces a yellow colour with starch. j'^t dis- 
 solves in water, giving a reddish solution (bromine wafer), and is 
 soluble in IprfKef and in carbon disulphide, giving reddish-brown 
 solutions. 
 
 Liberation from its Compounds. Bromine is liberated 
 from bromides by the action of manganese dioxide and sulphuric 
 acid (compare -chlorine). The solution of the bromide and the 
 manganese dioxide are placed in a small beaker, and a little strong 
 
 It is important to bear in mind that with an excess of hydrochloric acid (which 
 would result if the proportion of potassium dichromate to the chloride in the 
 mixture was small) the reaction takes a different course, and only chlorine is 
 evolved, the chromium being completely reduced to chromic chloride ; thus 
 
 Cr0 3 + 6HC1 = sH 2 + CrCl 3 + 3d 
 Or, to give the complete equation 
 
 6KC1 + K 2 Cr 2 : + 7H 2 SO 4 = 4K 2 SO 4 -f- Cr 2 (SO' 4 } 3 + 7H 2 O + 3C1 2 
 Hence, in applying this test, it is necessary to employ an excess of the dichromate. 
 * The former tests (p. 123), which enable one to distinguish between a 
 chloride, bromide, and iodide, will not be confounded with a test such as the 
 above, which permits of the detection of one class of salts in the presence of 
 others. 
 
Hydrob^l ^md and Bromides. 125 
 
 sulphuric acid is addeij Ire beaker is then covered over with a 
 piece of moistened V^ottii^Bper upon which a little starch flour 
 has been dusted, jj Rted bromine produces a yellow colour 
 
 with the starch.J|NP 
 
 Bromine mjp ties ^^Herated by means of chlorine 
 /starch wBT 
 
 A small quat^L , 22bn disulphide is added to the solution of 
 
 the bromide irrow test-tube ; a few drops of chlorine water 
 
 idded, and the mixture shaken. The liberated bromine is 
 
 ed by the carbon disulphide, giving a red or brownish 
 
 cv,. ...red liquid (according to the amount of bromine), which settles 
 
 to the bottom of the test-tube. The test must be made with a little 
 
 care, an excess of chlorine being avoided, as otherwise chloride of 
 
 bromine is formed, which, being colourless, destroys the test. 
 
 Hydrobromic Acid and Bromides. 
 
 Gaseous hydrobromic acid closely resembles hydrochloric acid. 
 The propertigfc of the gas are not used in analysis. 
 
 The detection of free hydrobropiic acid in solution in presence 
 of bromides may be accomplished by gently warming the liquid 
 with manganese dioxide, without the addition of sulphuric acid. 
 Bromine is liberated from tM^Bfee acid (not from the bromide), 
 and may be detected by the ^^1 ction given above. 
 
 All bromides are soluble in water, except mercurous bromide and 
 silver bromide ; lead bromide disifilves in boiling water less easily 
 than the chloride. 
 
 Bromides (except Hg 2 Br 2 and AgBr), when acted upon by 
 strong sulphuric acid, evolve hydrobromic acid, bromine, and 
 sulphur dioxide (contrast chlorides, which under these circum- 
 stances give only HC1), the hydrobromic acid first formed being 
 immediately acted upon by the sulphuric acid ; thus 
 
 KBr + H 2 SO 4 = HBr + HKSO 4 
 2HBr + H 2 SO 4 = Br 2 + SO 2 + 2H 2 O 
 
 Silver Nitrate, AgNO 3 , precipitates from solutions of 
 bromides or hydrobromic acid, pale-yellow silver bromide, AgBr 
 (the colour is indistinguishable from white by gaslight). It is 
 insoluble in nitric acid, and difficultly soluble in ammonia (scarcely 
 soluble in dilute ammonia. Contrast AgCl). 
 
 AgBr may be distinguished from AgCl by shaking up a little 
 of the washed precipitate with a few drops of carbon disulphide 
 and chlorine water. 
 
126 Qualitative Analysis. 
 
 Silver bromide is decomposed by metallic zinc in the presence 
 of dilute* sulphuric acid, in the same manner as the chloride. Zinc 
 bromide goes into solution, from which the 'bromine can be sepa- 
 rated by either of the methods given under b^viinc. 
 
 Prolonged boiling with a strong soluf sma ^ "sodium carbonate 
 (or, better, heating the dry substances /^^ chlo n a glass tube ^ 
 decomposes silver bromide. On filteri ratures bt $xtraction with 
 water in the case of the dry reaction), tSv a ove J solution con- 
 taining sodium bromide may be tested as abov8. at 
 
 A bromide may be detected in the presence of a chloride by 
 means of chlorine water, which liberates bromine from the bromide, 
 but obviously is without action upon the chloride. 
 
 Iodine. 
 
 Iodine is a steel-black, shining, crystalline solid. When gently 
 heated it melts and passes into vapour, which has a characteristic 
 deep violet colour. Iodine is very slightly soluble in water, but 
 readily dissolves in water holding hydriodic acid or alkaline iodides 
 in solution, giving a brown liquid. It dissolves in carbon disulphide 
 (also in chloroform), yielding a violet solution. In contact with 
 starch it produces an intense indigo-blue colour, which constitutes 
 one of its most delicate tests. 
 
 Liberation of Iodine Qtfte its Compounds. Iodine is 
 more easily set free from corm|H|ion than either bromine or 
 chlorine, and the methods which are applicable for the liberation 
 of these apply also in the case of iodine. Thus, manganese dioxide 
 and dilute sulphuric acid decompose iodides in a manner precisely 
 similar to that explained on p. 122 for chlorides. Strong acids, as 
 nitric and sulphuric, also expel iodine from iodides, with evolution 
 of oxide of nitrogen, or sulphur dioxide ; e.g. 
 
 2K1 + 2H 2 SO 4 = K 2 SO, + 2H 2 O + SO 2 + I, 
 
 The comparative ease with which iodine is liberated from com- 
 bination, affords the basis of most of the tests by which this element 
 is detected. The following are the reactions most used in analysis: 
 
 i. Chlorine water, when added to a solution of an iodide, 
 expels the iodine. The test may be applied as described under 
 bromine, the carbon disulphide in this case being coloured violet. 
 
 The presence of the liberated iodine may also be recognised by 
 means of starch. A small quantity of starch paste * is mixed with 
 
 * Starch paste is made by mixing a little starch flour into a thin cream with 
 the least quantity of cold water, and then pouring boiling water upon it until the 
 dead-white appearance of the starch is changed to a translucent appearance. 
 
Iodine. 127 
 
 the solution of the iodide, and one or two drops of chlorine water 
 added, when the deep indigo-blue compound of iodine with starch 
 is produced. On the addition of an excess of chlorine, the colour 
 is destroyed. Boiling the liquid also destroys the compound, hence, 
 when small quantities of iodine are being tested, it is necessary to 
 avoid using the starch while hot. 
 
 Detection of Bromides and Iodides in Solution to- 
 gether. Although chlorine water liberates both bromine and 
 iodine, the reaction may be employed to detect both halogens in the 
 same solution, owing to the fact that the chlorine exerts a selective 
 action, expelling first the iodine, and afterwards the bromine. The 
 test is applied in the following way : Carbon disulphide is added to 
 the mixed solution of iodide and bromide in a test-tube, and chlorine 
 water added in small quantities at a time, with agitation. If this be 
 done carefully, it is not difficult to see when the further addition of 
 a drop of chlorine water produces no further precipitation of iodine. 
 At this point the carbon disulphide (coloured deep violet with 
 dissolved iodine) is removed from the aqueous liquor either by 
 decanting the latter, or, better, by withdrawing a portion of it by 
 means of a small pipette and transferring it to another test-tube. 
 A fresh quantity of carbon disulphide is now added to this, and a 
 few drops of chlorine water. If the whole of the iodine had been 
 liberated in the first tube, the bromine now begins to be expelled, 
 and the carbon disulphide becomes brown. If a small quantity of 
 iodine were still left, the first drop of chlorine water causes its 
 liberation, and, on snaking, the disulphide will show a pale violet 
 colour. A few more drops of chlorine water, however, will destroy 
 this, and afterwards liberate the bromine.* 
 
 2. Nitrous Acid. When a solution of an iodide is acidified 
 (preferably with dilute sulphuric acid), and a few drops of a solution 
 of sodium nitrite added, the nitrous acid generated (by the action 
 of the acid upon the nitrite) decomposes the iodide, with the 
 liberation of iodine. The action is in reality between the nitrous 
 acid and hydriodic acid ; thus 
 
 HNO 2 + HI - H,O + NO + I 
 or, to state the complete change 
 NaNO 2 + 2H 2 SO 4 + KI = HNaSO 4 + HKSO 4 + H 2 O + NO +1 
 
 * When the quantity of iodide present is very small in proportion to the 
 bromide, it is not necessary to change the carbon disulphide as here described. 
 The chlorine, after liberating the iodine, next begins to combine with it, form- 
 ing iodine trichloride (therefore destroying the violet colour of the carbon 
 disulphide solution), and afterwards to liberate the bromine. But, if the pro- 
 portion of iodine present is at all considerable, the yellow colour which the 
 iodine chloride itself imparts to the carbon disulphide if present in any 
 quantity renders it practically impossible to distinguish the brownish colour 
 which would be given by moderate quantities of bromine. 
 
128 Qualitative Analysis. 
 
 Neither bromine nor chlorine is liberated by nitrous acid, hence 
 this reaction allows of the separation of iodine from bromides and 
 chlorides, and therefore the subsequent detection of bromine or 
 chlorine in mixtures of the three salts. 
 
 Detection of Iodides, Bromides, and Chlorides in Solu- 
 tion together. The solution containing the three salts, to which 
 a little carbon disulphide has been added, is acidified with two 
 or three drops of dilute sulphuric acid, and a dilute solution of 
 sodium nitrite added drop by drop until the whole of the iodine has 
 been expelled. On shaking the mixture this will be dissolved by the 
 carbon disulphide, giving the violet solution. The aqueous liquid 
 is then withdrawn with a pipette and divided into two portions. 
 The first is neutralised by the cautious addition of ammonia drop 
 by drop. It is then shaken with chlorine water and carbon disul- 
 phide. The bromine is thereby liberated, and imparts its brownish 
 colour to the disulphide. The second portion is evaporated down, 
 mixed with potassium dichromate and sulphuric acid, and the 
 chromyl-chloride test made as described on p. 123. 
 
 3. Ferric Chloride. AVhen ferric chloride, acid with hydro- 
 chloric acid, is added to a solution of an iodide and the mixture 
 heated, the iron salt is reduced to the ferrous condition, and iodine 
 escapes as a violet vapour 
 
 Fed, + KI = KC1 -f FeCl 2 + I 
 
 Hydriodic Acid and Iodides. 
 
 Gaseous hydriodic acid resembles hydrobromic and hydro- 
 chloric acids, except that it is more easily decomposed, being 
 dissociated into iodine and hydrogen by very moderate heat. The 
 properties of the gas are not used in analysis. 
 
 All iodides are soluble in water, except those of silver, Agl ; 
 mercury, Hg 2 I 2 and HgI 2 ; copper, Cu 2 I 2 ; gold, AuI 3 ; platinum, 
 PtI 4 ; palladium, PdI 2 ; (BiI 2 and PbI 2 sparingly soluble). 
 
 The formation of some of these insoluble salts is utilised in 
 analysis. 
 
 Silver nitrate, AgN0 3 , when added to a solution of hydriodic 
 acid or an iodide, gives a pale-yellow precipitate of silver iodide, 
 Agl, insoluble in nitric acid, and more difficult of solution in 
 ammonia than silver bromide. [In very strong solutions the pre- 
 cipitation is incomplete, owing to the partial solubility of Agl in 
 concentrated solutions of HI or KI. Dilution with water repre- 
 cipitates the dissolved compound.] 
 
 Silver iodide may be distinguished from the bromide or chloride 
 
Hypochlorons Acid. 129 
 
 by shaking up the precipitate with a little carbon disulphide and 
 chlorine water. 
 
 Silver iodide is, reduced by metallic zinc in the same way as the 
 bromide and chloride, or it may be decomposed (in common with 
 all the other insoluble iodides) by fusion with sodium carbonate. 
 
 Copper sulphate, CuSO 4 , added to a solution of an iodide, 
 gives a dirty brown precipitate, consisting of cuprous iodide, Cu 2 I 2 , 
 and free iodine (see Copper reactions, p. 87) 
 
 2CuSO 4 + 4KI = 2K 2 SO 4 + Cu 2 I 2 + I 2 
 Cuprous iodide is a white crystalline compound, insoluble in 
 
 water and in dilute acids ; soluble in ammonia and in sodium thio- 
 
 sulphate. 
 
 In the presence of suitable reducing agents, the whole of the 
 
 iodine may be precipitated as cuprous iodide ; and as bromine and 
 
 chlorine are not precipitated under these conditions, this reaction 
 
 is made use of as a means of separation. 
 
 Separation of Iodine as Cuprous Iodide. To the solution 
 containing an iodide, bromide, and chloride, copper sulphate is added 
 in excess (i.e. in quantity more than sufficient to combine with the 
 whole of the iodine). Sodium thiosulphate is added drop by drop 
 from a pipette until the brown colour (due to the free iodine) 
 disappears, and the precipitate is white. When this is filtered, 
 the liquid will be coloured blue with the excess of copper sulphate, 
 and will contain the bromide and chloride, -for which it may be 
 examined by methods already described. 
 
 The first action of the sodium thiosulphate is to convert the 
 free iodine into sodium iodide, and to form sodium tetra-thionate 
 
 2Na 2 S 2 O 3 + I 2 = 2NaI + Na 2 S 4 O c 
 
 and the sodium iodide, then reacting upon copper sulphate in the 
 presence of the thiosulphate, has the whole of its iodine precipi- 
 tated as Cu 2 I 2 . 
 
 Falladious nitrate, Pd(NO 3 ) 2 , when added to a solution 
 of an iodide, gives a black precipitate of palladious iodide, PdI 2 . 
 The reaction is delicate, notwithstanding that the precipitate is 
 soluble to some extent in an excess of potassium iodide. It serves 
 to detect iodides in the presence of chlorides and bromides. 
 
 THE HALOGEN OXYACIDS. 
 Hypochlorous Acid and Hypochlorites. 
 
 Hypocklorous acid, HC1O, is only known in solution in 
 water, as it readily undergoes decomposition. The solution has a 
 pale yellow colour, and a faint smell (familiar in ordinary chloride 
 of lime]. It also has bleaching properties. 
 
 K 
 
130 Qualitative Analysis. 
 
 A solution of the free acid may be distinguished from chlorine 
 water (besides from its odour) by shaking it up with a little mer- 
 cury. With chlorine water white mercurous chloride, Hg 2 Cl 2 , is 
 formed, while with hypochlorous acid a brownish-coloured oxy- 
 chloride is produced 
 
 2Hg + 2HC1O = HgO,HgCl 2 + H 2 
 
 All hypochlorites are soluble in water, therefore no precipitation 
 reactions for their detection are possible. The tests employed for 
 their identification are based upon their oxidising action. They 
 are decomposed by even feeble acids (carbonic), with evolution of 
 chlorine, e.g. 
 
 NaCIO + 2HC1 = NaCl + H 2 O + Cl, 
 Ca(C10) 2 + C0 2 = CaC0 3 + C1 2 + O 
 
 If a few drops of a solution of a hypochlorite be poured upon 
 litmus paper, and the paper exposed to the action of carbon dioxide 
 (by breathing upon it), it will become bleached where it was mois- 
 tened, by the evolved chlorine. 
 
 The evolution of chlorine by the action of dilute acids distin- 
 guishes hypochlorites from chlorides and chlorates. The addition 
 of silver nitrate to a hypochlorite results in the precipitation of 
 silver chloride, one-third of the silver going into solution as silver 
 chlorate ; thus 
 
 (1) NaCIO -I- AgNO 3 = AgCIO + NaNO s 
 
 (2) 3AgC10 = AgC10 3 + 2AgCl 
 
 The commonest salt of hypochlorous acid is bleaching powder, 
 which is a double chloride and hypochlorite of calcium, having the 
 composition Ca(OCl)CL When treated with water it splits up 
 into calcium chloride and calcium hypochlorite, and a clear solu- 
 tion made in this way contains the salt Ca(ClO) 2 
 2Ca(ClO)Cl = CaCL, + Ca(ClO) 2 
 
 Chlorates, Bromates, and lodates. 
 
 The three acids, chloric, HC1O 3 , bromic, HBrO 3 , and iodic, 
 HIO 3 , differ in stability in the opposite order to that usually 
 exhibited by compounds of the halogens ; thus, iodic acid is a 
 comparatively stable solid, while chloric acid can only exist in 
 aqueous solutions containing not more than 20 per cent, of the 
 acid. Beyond this strength, or when heated, it decomposes into 
 oxygen, chlorine, and perchloric acid 
 
 3HC1O 3 = 2O 2 + C1 2 + HC1O 4 + H 2 O 
 
Chlorates, Bromates, and lodates. 131 
 
 Bromic acid under the same conditions breaks up into bromine, 
 oxygen, and water ; thus 
 
 4HBrO 3 = 50, + 2Br> + 2H 2 O 
 
 lodic acid does not give a blue colour with starch, but it readily 
 parts with its oxygen to such reducing agents as sulphuretted 
 hydrogen or sulphur dioxide, with liberation of iodine ; thus 
 
 3HI0 3 + 5H 2 S = 58 + 6H 2 + I a * 
 
 Similarly, hydriodic acid and iodic acid undergo interaction, 
 the whole of the iodine being thrown down 
 
 This reaction affords the means of detecting the presence of an 
 iodate when mixed with an iodide. A little starch is added, and a 
 small quantity of acetic acid (this, while liberating the respective 
 iodine acids from their salts, is not able to decompose the iodic 
 acid), when the blue starch coloration is obtained. 
 
 The salts of these acids are produced (mixed with the corre- 
 sponding halides) by the direct action of the halogens upon caustic 
 alkalies ; thus, in the case of iodine (bromine and chlorine being 
 similar) 
 
 6KHO + 3l 2 = 5KI + KIO 3 + 3H 2 O 
 
 Chlorates are all soluble in water. 
 
 Bromates. The silver and mercurous salts are difficult of 
 solution, and may be obtained by precipitation. 
 
 lodates. Only the iodates of the alkali metals are soluble in 
 water. 
 
 The chlorates, bromates, and iodates are all decomposed by 
 heat, yielding either oxygen and the haloid salt (or oxygen, metallic 
 oxide, and free halogen) ; or, in some cases, leaving both an oxide 
 and halide in the residue for example 
 
 (1) AgC10 3 = AgCl + 3 
 
 (2) Mg(Br0 3 ) 2 = MgO + Br 2 + 50 
 
 (3) 2Pb(BrO 3 ) 2 = PbO + PbBr 2 + Br 2 + iiO 
 
 When heated with oxidisable substances (e.g. charcoal), deflagration 
 of the mixture results. 
 
 Hydrochloric acid decomposes chlorates, with the evolution 
 of chlorine and chlorine peroxide 
 
 4KC1O., + I2HC1 = 4KC1 + 6H 2 O + QC1 + 3C1O 2 
 
 * Owing to this property, if iodic acid is present in a mixture undergoing 
 analysis, iodine will be precipitated in the process of separating Group II. by 
 means of H 2 S. By the continued action of the sulphuretted hydrogen, however, 
 the iodine is converted into hydriodic acid, H 2 S + I 2 S + 2HI. 
 
132 Qualitative Analysis. 
 
 (The use of this mixture as an oxidising agent has frequently been 
 referred to.) 
 
 Bromates and iodates under the same treatment give bromine 
 and iodine respectively. In the latter case, chloride of iodine is 
 also formed 
 
 2KBrO 3 + 2HC1 = 2KC1 + H 2 O + Br 2 + 5O 
 KIO 3 + 6HC1 = KC1 + IC1 3 + C1 2 + 3H 2 O 
 
 Sulphuric acid decomposes chlorates, with the evolution of 
 chlorine peroxide (a deep-yellow unpleasant-smelling gas), which 
 on very slight elevation of temperature, explodes with violence 
 
 3 KC10 3 + 2H 2 S0 4 = KC10 4 + 2HKS0 4 + H 2 O + 2 C1O 2 
 
 A minute crystal of potassium chlorate, with about three or 
 four drops of strong sulphuric acid, may be heated in a test-tube ; 
 sharp detonations will result, characteristic of chlorates. From bro- 
 mates sulphuric acid liberates bromine. Dilute acid first liberates 
 bromic acid, which gradually decomposes into HBr and O. As 
 soon as this decomposition sets in, the HBr reacts upon the HBrO 3 , 
 with liberation of bromine 
 
 HBr0 3 + 5HBr = 3 H 2 O + 3Br 2 
 
 Dilute sulphuric acid, therefore, liberates bromine immediately 
 from a mixture of a bromide and bromate ; but only gradually , after 
 an interval, from a bromate alone. 
 
 Silver nitrate gives with bromates a white precipitate of 
 silver bromate, AgBrO 3 ; soluble in ammonia. It is decomposed by 
 hydrochloric acid, with liberation of bromine (distinction from all 
 silver halides). With iodates, silver nitrate precipitates white silver 
 iodate, AgIO 3 . Difficultly soluble in nitric acid, readily dissolved 
 by ammonia (distinction from Agl). With chlorates silver nitrate 
 gives no precipitate.* 
 
 Hydrogen sodium sulphite added to a solution of an iodate, 
 liberates the whole of the iodine ; thus 
 
 2NaIO 3 + 5HNaSO s = 3HNaSO 4 + 2Na. 2 SO 4 + H 2 O + I 2 
 
 Similarly, if sulphurous acid be added to the ammoniacal solution 
 of silver iodate, the iodate is entirely reduced, and the iodine 
 thrown out as silver iodide 
 
 AgI0 3 + 3H(NH 4 )S0 3 = 3 H(NH 4 )S0 4 + Agl 
 
 * For method of detecting a chlorate and nitrate in solution together, see 
 Nitric acid (p. 147). 
 
Hydrofluoric Acid and Fluorides. 133 
 
 Barium chloride precipitates from soluble iodates, white 
 barium iodate, Ba(IO 3 ) 2 . Slightly soluble in water, insoluble on 
 addition of alcohol ; slightly soluble in nitric acid, easily soluble 
 in hydrochloric acid. (This reaction affords a means of separating 
 iodic from hydriodic acid, as well as from all other halogen acids.) 
 
 A chloride may be separated from a chlorate by adding to the 
 solution silver sulphate. Silver chloride is precipitated, and removed 
 by filtration. Sodium carbonate is then added to remove the 
 excess of silver, and any metals other than the alkalies which 
 might be present ; and the solution evaporated to dryness, and 
 heated. The chlorate is thereby converted into chloride. 
 
 Per chlorates. 
 
 The perchlorates are all soluble in water, and are more stable 
 than the chlorates. Thus, when potassium chlorate is heated, the 
 final products of the decomposition are KC1 and O. But as an 
 intermediate product potassium perchlorate is formed ; thus 
 
 2KC10 3 = KC10 4 -f KC1 4- (X 
 
 and on applying a higher temperature, the potassium perchlorate is 
 finally resolved into potassium chloride and oxygen. (Bromates, 
 when heated, pass at once into bromides, with evolution of oxygen.) 
 Sulphuric acid liberates perchloric acid, HC1O 4 , a colourless, 
 fuming, corrosive, volatile liquid. The mixture of the perchlorate 
 and acid does not become yellow ; and on warming, the acid 
 distils without decomposition (distinction from chlorates, also from 
 bromates]. 
 
 Hydrochloric acid has no action on perchlorates (distinction 
 from chlorates, bromates, iodates}. 
 
 Fluorine, Hydrofluoric Acid, and Fluorides. 
 
 The properties of the element are not made use of in analysis, 
 as fluorine is not liberated by any of the ordinary analytical re- 
 actions. 
 
 Hydrofluoric acid is a colourless, fuming, highly corrosive liquid, 
 boiling at 19-5 C, therefore extremely volatile. It dissolves in 
 water, and even a moderately dilute solution attacks the skin 
 violently. The most characteristic property of the acid, either 
 gaseous or in solution, is its power of dissolving silica (therefore 
 glass) : its application is described below. 
 
 Liberation of Hydrofluoric Acid. Fluorides are decom- 
 posed by strong sulphuric acid, with evolution of the acid in the 
 
134 Qualitative Analysis. 
 
 gaseous condition. Thus with the mineral cryolite, and the still 
 commoner fluor spar 
 
 2Na 3 AlF,. + 6H 2 SO 4 = 3Na 2 SO 4 + A1.,(SO 4 ) S + I2HF 
 CaF 2 + H 2 S0 4 = CaS0 4 + 2HF 
 
 Hydrofluoric acid is also evolved when a fluoride is heated with 
 powdered hydrogen potassium sulphate 
 
 CaF 2 + 2HKSO 4 = CaSO 4 + K,SO 4 + 2HF 
 Certain acid fluorides, when heated alone, decompose with 
 evolution of hydrofluoric acid ; thus 
 
 HF,KF = KF + HF 
 
 The fluorides of the non-metals are all volatile without decom- 
 position. 
 
 The fluorides of the alkali metals, and of silver, mercury (iron, 
 aluminium, tin), are soluble in water. Those of the alkaline earths, 
 and of lead (copper, zinc, manganese), are insoluble. 
 
 Calcium chloride gives a transparent gelatinous precipitate 
 of calcium fluoride, CaF 2 , partially soluble in hydrochloric acid. 
 
 Barium chloride throws down a white precipitate of barium 
 fluoride, BaF 2 , partially soluble in HC1. 
 
 Silver nitrate gives no precipitate, as silver fluoride is soluble 
 in water (distinction between a fluoride and the other halides). 
 Formation of Silicon Fluoride. When hydrofluoric acid 
 comes in contact with silica, or silicates (such as glass), the silicon 
 dioxide is dissolved, and gaseous silicon fluoride is formed 
 
 SiO 2 + 4HF = SiF 4 + 2H 2 O 
 
 The test may be applied in the following ways : 
 (a) Etching Glass. The powdered fluoride is mixed with 
 strong sulphuric acid in a small dish or tray, made of lead (or a 
 platinum capsule). It is covered with a small piece of sheet glass 
 which has been coated on one side with wax,* and some marks or 
 words scratched upon the wax. In a few minutes the exposed 
 parts of the glass will have become eaten into or dissolved away 
 by the acid gas ; so that, on removing the wax with a little hot 
 water, the marks or letters will be found to be etched into the glass. 
 
 () The Decomposition of Silicon Fluoride by Water. 
 If the fluoride (natural fluoride) contains much silica, or if it be 
 
 * A little paraffin wax is melted in a test-tube, and poured upon the 
 previously warmed glass plate. The liquid wax is made to flow all over the 
 surface of the plate, which is then stood up on edge to drain and cool. In this 
 way a thin and uniform film of wax will be obtained. 
 
Hydrofluoric Acid. 135 
 
 intentionally mixed with silica (sand), the action of sulphuric acid is 
 to liberate silicon fluoride, a gas which is without action upon glass. 
 The hydrofluoric acid (from the fluoride), being generated in the 
 immediate presence of silica, at once combines with it ; thus 
 2CaF 2 + 2H 2 S0 4 + Si0 2 = 2CaSO 4 + 2H 2 O* + SiF 4 
 
 This gas is at once decomposed by water, with the formation of 
 hydrofluosilicic acid (soluble in water) and the precipitation of 
 gelatinous silicic acid, H 2 SiO 3 
 
 3 SiF 4 + 3 H 2 = 2H 2 SiF 6 4- H 2 Si0 3 
 
 The mixture of the fluoride and sand is gently warmed in a 
 test-tube with a little strong sulphuric acid, and a glass rod with a 
 drop of water upon the end is lowered into the mouth of the tube. 
 The gas, on coming in contact with the water, is decomposed, and a 
 white deposit of silicic acid is formed upon the rod. [The experi- 
 ment is rendered more delicate by using a piece of narrow glass 
 tube with a drop of water at the end. By means of a piece of 
 rubber pipe attached to the other end of the tube, the drop of water 
 can be slowly sucked up the tube, and the gas which is drawn up 
 after it will then deposit a white film of silica upon the wet walls of 
 the tube. The gas must not be drawn into the lungs.] (See also 
 Hydrofluosilicic acid.) 
 
 The Formation of Boron Fluoride. When a fluoride (finely 
 powdered) is mixed with powdered borax, and the mixture moistened 
 with strong sulphuric acid, gaseous boron fluoride, BF 3 , is evolved. 
 The action takes place between the hydrofluoric acid and boric 
 acid, which are disengaged by the action of the sulphuric acid upon 
 the respective compounds 
 
 B(HO) 3 + 3HF = 3 H 2 + BF 3 
 
 If the mixture be introduced into the edge of a Bunsen flame 
 upon a loop of platinum wire, the flame is tinged a grass-green colour 
 by the escaping boron fluoride. 
 
 [The test may be modified by employing hydrogen potassium 
 sulphate instead of sulphuric acid, and heating the mixture of the 
 three salts upon the platinum wire.] 
 
 Hydrofluosilicic Acid and Silicofluorides. 
 This acid is obtained when silicon fluoride is decomposed by 
 water. The gas is made to bubble into water, and the silicic acid 
 separated by filtration. 
 
 * The two molecules of H 2 O here produced are not liberated as water, but 
 are retained by the calcium sulphate in the form of water of hydration ; the 
 fully hydrated compound having the composition CaSO 4 ,2HoO. 
 
136 Qualitative Analysis. 
 
 The acid and its salts are decomposed by heat 
 
 H 2 SiF 6 = SiF 4 + 2HF 
 BaSiF 6 = SiF 4 + BaF 2 
 
 The silicofluorides of potassium, sodium, and barium are pre- 
 cipitable salts, being only difficultly soluble in water. The addition 
 of alcohol renders the precipitation complete (see Reactions of 
 metals of Group IV. and V.). Barium silicofluoride is only 
 sparingly soluble in dilute hydrochloric acid, and is therefore pre- 
 cipitated by barium chloride in a hydrochloric acid solution. 
 Barium silicofluoride, however, is distinguished from barium sulphate 
 by dissolving in strong hydrochloric acid (the solution being after- 
 wards diluted to dissolve any barium chloride which may be thrown 
 down by the strong acid) ; or by the fact that when dried and 
 heated it evolves silicon fluoride, as shown by the above reaction. 
 
 The addition of ammonia to hydrofluosilicic acid, or a soluble 
 silicofluoride, gives a white precipitate consisting of gelatinous 
 silicic acid 
 
 K,SiF 6 + 4NH,HO = 2KF + 4NH 4 F + H 2 O + H 2 SiO 3 
 
CHAPTER XII. 
 SULPHUR. 
 
 THE chief properties made use of in analysis are the following : 
 Sulphur is a pale-yellow, brittle solid. Insoluble in water ; easily 
 soluble in carbon disulphide, from which it is deposited on evapora- 
 tion in amber-coloured rhombic octahedrons. 
 
 When heated in a test-tube, it melts and boils off as a brownish- 
 yellow vapour, which condenses to brown drops on the moderately 
 hot parts of the tube (turning yellow when cold), and as a yellow 
 sublimate on the upper and cooler part of the tube (flowers of 
 sulphur}. 
 
 When heated with access of air, either on a platinum capsule or 
 in a glass tube open at both ends, sulphur burns with a pale blue 
 flame, producing sulphur dioxide (of characteristic smell), and 
 leaving no residue. Nitric acid oxidises sulphur into sulphuric 
 acid, with evolution of nitrogen peroxide. 
 
 Liberation of sulphur takes place when certain metallic 
 sulphides are heated alone, either in a tube closed at one end, or 
 with partial access of air ; e.g. 
 
 3FeS 2 = Fe 3 S 4 + S 2 
 3FeS 2 + 50, - Fe 3 O 4 + 380, + 3$ 
 
 Sulphuretted Hydrogen* and Sulphides. 
 
 Sulphuretted hydrogen is a colourless gas, easily distinguished 
 from all other gases by its unmistakable odour. It is soluble in 
 water, and imparts its own smell to the liquid. The solution, how- 
 ever, is unstable, undergoing oxidation and depositing sulphur. 
 The gas burns with a flame resembling that of burning sulphur, 
 and yields water and sulphur dioxide. 
 
 Liberation of sulphuretted hydrogen takes place when 
 
 * Sometimes called hydro sulphuric acid ; the solution of the gas in water 
 has a feeble acid reaction. It should be remembered that the gas is poisonous. 
 
138 Qualitative Analysis. 
 
 certain sulphides (see below) are acted upon by acids. The gas may 
 be recognised (i) by its odour ; (2) by its action upon solutions of 
 metallic salts, e.g. lead acetate. The reaction is made in a test- 
 tube, and a piece of paper moistened with lead acetate is held over 
 the mouth of the tube. The sulphuretted hydrogen causes a black 
 stain of lead sulphide. 
 
 Sulphides of the alkalies and alkaline earths are soluble in 
 water ; all other metallic sulphides are insoluble (see Analytical 
 classification of the metals). 
 
 Soluble sulphides are decomposed by dilute acids (HC1 or 
 H 2 SO 4 ), with liberation of sulphuretted hydrogen; in the case of 
 polysulphides, sulphur is also precipitated.* 
 
 K.jS + 2HC1 + Aq = 2KC1 + H 2 S + Aq 
 CaS, + 2HC1 = CaCl 2 + H 2 S +4$ 
 
 Insoluble Sulphides. t The behaviour of these towards 
 acids has already been considered in detail, in studying the separa- 
 tion of the metals. It may be briefly summarised as follows : 
 
 (a) Sulphides decomposed by dilute acids (HC1 or H 2 SO 4 ), with 
 liberation of sulphuretted hydrogen : namely, ZnS, MnS, FeS. 
 
 (b) Sulphides unacted upon by dilute acid, but decomposed by 
 hot strong hydrochloric acid with more or less difficulty : Sb 2 S 3 , 
 PbS, SnS, NiS, CoS. 
 
 (c) Sulphides unacted upon by strong hydrochloric acid, but 
 decomposed by aqua regia, or by a mixture of hydrochloric acid 
 and potassium chlorate ; HgS, As 2 S 3 , PtS 2 , AuS. 
 
 The sulphides of class (b\ when treated with hydrochloric acid 
 in the presence of zinc, or, better, of reduced iron, readily evolve 
 sulphuretted hydrogen. 
 
 Oxidising agents, e.g. nitric acid, convert many of the sul- 
 phides into oxides or sulphates, sulphur being first separated and 
 afterwards oxidised into sulphuric acid. 
 
 When a sulphide is added (in small quantities at a time) to a 
 fused mixture of sodium carbonate and potassium nitrate in a 
 platinum crucible, the sulphide is immediately oxidised. After the 
 mass has cooled, and been extracted with water, the aqueous liquid 
 may be tested for a sulphate. The sulphides of classes (a) and (b\ 
 when fused upon a piece of platinum foil (or, better, silver) with 
 
 * Under certain conditions hydrogen persulphide is formed, without any 
 evolution of sulphuretted hydrogen. 
 
 t Aluminium sulphide, A1 2 S 3 , is not formed in analysis. It is decomposed 
 by water, with the liberation of sulphuretted hydrogen ; thus 
 A1 2 S 3 + 3HoO = A1 2 3 + 3H 2 S 
 
Sulphuric Acid and Sulphates. 139 
 
 sodium hydroxide, are decomposed, with the formation of sodium 
 sulphide. If a fragment of the fused mass, after cooling, be placed 
 upon a silver coin and moistened with a drop of water, or upon 
 a piece of paper which has been moistened with a solution of lead 
 acetate, in either case a black stain will be produced ; silver sulphide 
 on the coin, and lead sulphide upon the paper. 
 
 Most sulphides, when heated in a glass tube open at both ends, 
 and held in a slightly inclined position in order to cause an air- 
 current to pass through the tube, are decomposed, and evolve sul- 
 phur dioxide. 
 
 Sodium nitroprusside, Na.,(NO)FeCy ,* gives with alkaline 
 sulphides (but not with an aqueous solution of sulphuretted hydro- 
 gen} a rich reddish-purple colour. One or two drops of the nitro- 
 prusside solution are added to the alkaline sulphide (e.g. ammonium 
 sulphide) upon a white plate. The colour is destroyed by caustic 
 alkalies, which must therefore be absent when applying the test. 
 (By this test, a soluble sulphide may be detected in the presence of 
 dissolved sulphuretted hydrogen.) 
 
 Sulphuric Acid and Sulphates. 
 
 Sulphuric acid is an oily, highly corrosive acid liquid. It com- 
 bines with water with evolution of heat, and is able to abstract the 
 elements of water from many organic compounds. Thus paper, 
 straw, etc., are blackened or charred by the strong acid. This pro- 
 perty is made use of in testing for the free acid in the presence of 
 soluble sulphates : a piece of paper is moistened here and there 
 with the solution, and then carefully dried, when it becomes charred 
 where it had been wetted. Or the solution may be mixed with a 
 little white sugar, and evaporated down in a porcelain dish upon a 
 steam-bath, when a charred residue will be left. 
 
 Most sulphates are soluble in water. Barium, strontium, cal- 
 cium, and lead sulphates are insoluble, or nearly so (see Reactions 
 of the various metals). All sulphates (except ferric sulphate) are 
 precipitated by alcohol, being insoluble in that liquid (hence the 
 use of alcohol in rendering the precipitation of the sulphates of 
 lead, calcium, or strontium complete). 
 
 Soluble Sulphates. Barium chloride, BaCL, gives with sul- 
 phuric acid or soluble sulphates, a white precipitate of barium 
 
 * This compound is readily obtained by boiling a little solid potassium 
 ferrpcyanide with strong nitric acid, diluting with water, and neutralising with 
 sodium carbonate. The salt easily crystallises in deep red crystals. The exact 
 nature of the action of alkaline sulphides upon it is unknown. 
 
140 Qualitative Analysis. 
 
 sulphate (see Barium reactions], insoluble in hydrochloric acid.* 
 The solutions should be dilute, as barium chloride, being insoluble 
 in strong hydrochloric acid, may otherwise be thrown out of solu- 
 tion ; the addition of water dissolves it. 
 
 ; Insoluble sulphates may be decomposed by fusion with 
 sodium carbonate, sodium sulphate being formed. The residue is 
 extracted with water, and the aqueous solution tested with barium 
 chloride after being acidified. 
 
 When a sulphate is fused with sodium carbonate (which must 
 be free from sulphates as impurities) upon charcoal in the reducing 
 flame, a sulphide of the alkali metal is obtained. If this be placed 
 upon a piece of paper moistened with acetate of lead, and touched 
 with a drop of dilute hydrochloric acid, sulphuretted hydrogen is 
 liberated, and the lead paper stained black. 
 
 [This test is only conclusive evidence of a sulphate when other 
 sulphur compounds are proved to be absent.] 
 
 Sulphurous Acids and Sulphites. 
 
 Sulphurous acid, H 2 SO 3 , is only known in solution, being pro- 
 duced when sulphur dioxide is passed into water, or when this gas 
 is liberated from combination (as from sulphites) in the presence 
 of water ; thus 
 
 (1) In dilute solution, Na 2 SO 3 + 2HCl +Aq = 2NaCl + H,SO 3 4-Aq 
 
 (2) In stronger solution, Na 2 SO 3 + 2HC1 = 2NaCl + H 2 O+ SO. 2 
 
 ftVThe anhydride, SO 2 , is recognised by its characteristic suffocating 
 4 odour (familiar as " the smell of burning sulphur "). 
 
 Reducing Action of Sulphurous Acid. Sulphurous acid 
 easily takes up oxygen, and passes into sulphuric acid, and some 
 of its most important reactions are those in which it thus acts as a 
 reducing agent. Many of these have been referred to under the 
 metals ; thus, potassium permanganate is reduced with formation 
 of manganous sulphate 
 
 2KMnO 4 + 5H 2 SO 3 = K 2 SO 4 + 2MnSO 4 + 2H 2 SO 4 + 3H 2 O 
 
 This reaction affords a delicate test for sulphur dioxide. The gas 
 is cautiously decanted (being much heavier than air) into a test- 
 tube containing water slightly tinted with a minute quantity of 
 potassium permanganate. On shaking the gas and water, the pink 
 colour will be destroyed. 
 
 * The insolubility of barium sulphate in hydrochloric acid does not dis- 
 tinguish this compound from barium silicorluoride. The latter compound, 
 however, when dried and heated, gives off silicon-fluoride (see p. 136). 
 
Sulphurous Acid and Sulphites. 141 
 
 Oxidising Action of Sulphurous Acid. Sulphurous acid is 
 also capable of undergoing reduction, acting therefore towards more 
 powerful reducing agents in the capacity of an oxidising substance. 
 Thus, stannous chloride in presence of hydrochloric acid, is oxidised 
 into stannic chloride, the sulphurous acid being reduced to sulphu- 
 retted hydrogen. This latter then reacts upon the stannic chloride, 
 with precipitation of stannic sulphide 
 
 (1) 3 SnCL + 6HC1 + H 2 SO 3 = 3SnCl 4 + 3H 2 O + H 2 S 
 
 (2) SnCl 4 + 2H 2 S = SnS 2 + 4HC1 
 
 Nascent hydrogen,* obtained by the action of hydrochloric acid 
 upon zinc, also reduces sulphurous acid to sulphuretted hydrogen, 
 
 H 2 S0 3 + 3H 2 = 3H 2 + H 2 S 
 
 The test may be made by adding a minute trace of sulphurous acid 
 (or a solution of a sulphite) to a mixture of zinc and hydrochloric 
 acid in a test-tube, and applying acetate of-lead paper to the mouth 
 of the tube. 
 
 Sulphites. The only sulphites soluble in water are those of 
 the alkali metals. They are all decomposed by dilute acids, with 
 evolution of sulphur dioxide (see above). Oxidising agents convert 
 them into sulphates. 
 
 When heated by themselves, most sulphites are converted into 
 sulphides and sulphates 
 
 4K 2 SO 3 = K 2 S + 3K 2 SO 4 
 
 Those of the alkaline earths leave an oxide, and evolve sulphur 
 dioxide 
 
 BaSOo = BaO + SO. 2 
 
 Barium chloride gives a white precipitate of barium sulphite, 
 BaSO 3 , soluble in dilute HC1 (distinction from BaSO 4 ). 
 
 Lead acetate precipitates white lead sulphite, PbSO 3 . The 
 salt undergoes no change when boiled (contrast Lead thiosulphate). 
 
 Silver nitrate gives a white precipitate of silver sulphite, 
 which, on boiling, is converted into black metallic silver 
 
 Ag 2 SO 3 + >H 2 O = H 2 SO 4 H- Ag 2 
 A sulphate and sulphite in solution may be separated by acidulating 
 
 * Nascent hydrogen, obtained by the action of sulphurous acid upon zinc 
 without any other acid, reduces the sulphurous acid to hyposulphurous (not to 
 be confounded with thiosulphuric} acid, an unstable compound which passes 
 into thiosulphuric acid and water ; thus 
 
 (1) H 2 S0 3 + H 2 = HO + H 2 SO, 
 
 (2) 2 H 2 S0 2 = IL,0 + H 2 S.0 3 
 
142 Qualitative Analysis. 
 
 the dilute solution and adding barium chloride. The precipitated 
 sulphate (insoluble in acid) is removed by nitration. To the solu- 
 tion, which now contains barium chloride and sulphurous acid, an 
 oxidising agent, such as chlorine water, is added, when a precipitate 
 of barium sulphate is again thrown down 
 
 BaCl 2 + H 2 SO 3 + H 2 O + C1 2 = BaSO 4 + 4HC1 
 
 Thiosulphuric Acid and Thiosulphates. 
 
 The acid is unknown in the free state ; when liberated from its 
 salts it at once breaks down into water, sulphur dioxide, and sulphur. 
 
 Thiosulphates of barium, lead, and silver are sufficiently difficult 
 of solution in water, to be precipitated on the addition of a solution 
 of a thiosulphate to solutions of the respective metals. All others 
 are more easily soluble. They are all decomposed by acids, with 
 precipitation of sulphur and evolution of sulphur dioxide (this is 
 characteristic, and distinguishes thiosulphates from sulphites, which 
 only give the dioxide without depositing sulphur). 
 
 When heated alone, thiosulphates are converted into sulphates 
 and polysulphides. 
 
 Barium chloride precipitates from a moderately strong solu- 
 tion white barium thiosulphate, BaS 2 O 3 , decomposed by hydro- 
 chloric acid ; thus 
 
 BaS 2 O 3 + 2HC1 = BaCl 2 -f H 2 O + SO 2 + S 
 
 Silver nitrate gives a white precipitate of silver thiosulphate 
 (easily soluble in excess of sodium thiosulphate, forming the double 
 salt, AgNaS 2 O 3 ), which quickly changes colour owing to the for- 
 mation of black silver sulphide (contrast the behaviour of silver 
 sulphite above) 
 
 Ag 2 S 2 3 + H 2 = H 2 SO, + Ag 2 S 
 
 Lead acetate behaves in a similar manner, giving black lead 
 sulphide (contrast lead sulphite). 
 
 Thiosulphates are oxidised by chlorine, bromine, and other 
 oxidising agents, in the same manner as the sulphites 
 
 NasjS.O, + 4C1 2 + 5H,O = 8HC1 + 2HNaSO 4 
 
 and are also reduced by nascent hydrogen, with formation of sul- 
 phuretted hydrogen 
 
 Na 2 S 2 O 3 .+ 2HC1 + 4H 2 = 2NaCl + 3H 2 O + 2H 2 S 
 Separation of the Sulphur Acids. A solution containing 
 
Separation of the Sulphur Acids. 143 
 
 a sulphide, sulphate, sulphite, and thipsulphate may be examined in 
 the following way : 
 
 (1) The sulphide is first separated as an insoluble metallic 
 sulphide in such a way as to avoid the liberation of any acid, 
 which would decompose the sulphite or thiosulphate. This is 
 done either by shaking up the solution with a little lead carbonate 
 (or cadmium carbonate), or by adding a solution of zinc chloride 
 which has been rendered alkaline with ammonia. The precipitated 
 sulphide is then removed by filtration. . Very small traces of sul- 
 phuretted hydrogen will produce a distinct coloration in the white 
 carbonate of lead. 
 
 (2) Barium chloride is added to the filtrate, which precipitates 
 barium sulphate and sulphite. 
 
 (3) The filtrate from this precipitate is acidified with hydro- 
 chloric acid and warmed. The thiosulphate is thereby decomposed, 
 with precipitation of sulphur and evolution of sulphur dioxide. 
 
 (4) The mixed barium sulphate and barium sulphite is treated 
 with hydrochloric acid. The sulphate is left, and the sulphite 
 dissolves. On filtering and adding chlorine water, the sulphurous 
 acid is oxidised to sulphuric acid, and a precipitate of barium sul- 
 phate is produced. 
 
CHAPTER XIII. 
 NITROGEN AND PHOSPHORUS. 
 
 THE properties of the element nitrogen in its free state are not 
 employed in the analytical examination of its compounds.* 
 
 Nitric Acid and Nitrates. 
 
 Nitric acid is a fuming corrosive liquid, miscible with water. 
 It readily dissolves most metals, converting them into nitrates or 
 oxides, with evolution of oxides of nitrogen, and in some cases with 
 the formation of ammonia ; e.g. 
 
 [Sn + 4HNO 3 = H 2 SnO 3 + H 2 O + 4NO,] x 5 (see Meta- 
 
 stannic acid) 
 
 3 Hg + 8HN0 3 = 3Hg(N0 3 ) 2 + 4 H 2 O + 2 NO 
 4Zn + ioHNO 3 = 4Zn(NO 3 ) 2 + 5H 2 O + N 9 O 
 
 4Zn + 9HNO 3 = 4Zn(NO 3 ) 2 + 3H.O + NH 3 
 
 The reduction of the nitric acid may be regarded as due to the 
 action of nascent hydrogen, which is developed by the first action 
 of the acid upon the metal 
 
 2HNO 3 + 3H 2 = 4H 2 O + 2NO 
 
 The extent to which the reduction of the acid will take place 
 depends upon a number of conditions namely, the particular metal, 
 the strength of acid, the temperature, and the amount of nitrate 
 which is present in the mixture. 
 
 * The presence of nitrogen in organic compounds may readily be detected 
 by heating a small quantity of the solid substance in a test-tube with a fragment 
 of the metal sodium (or potassium). Under these circumstances the nitrogen 
 unites with the carbon of the organic compound, forming cyanogen, which, in 
 the presence of the alkali metal, gives a cyanide of the metal. The still hot 
 tube should be plunged into a small quantity of water in a beaker, which 
 immediately breaks it up and allows the cyanide to dissolve, and the excess of 
 sodium to be quietly acted upon by the water. Two or three drops of both 
 ferrous and ferric salts are added, and the liquid then acidified with hydrochloric 
 acid, when Prussian blue will be formed. 
 
Nitric Acid and Nitrates. 145 
 
 With magnesium, and under special conditions with zinc also, 
 free hydrogen is evolved by the action of nitric acid. 
 
 Nitric acid also oxidises many of the non-metals ; thus sulphur, 
 phosphorus, and iodine, are converted respectively into sulphuric, 
 phosphoric, and iodic acids. 
 
 Nitrates are all soluble in water ; their recognition, therefore, 
 is based upon the oxidising reactions of which they, or the nitric 
 acid which they yield, are capable. 
 
 Reduction by Perrons Salts. When ferrous sulphate is 
 brought into contact with a mixture of a nitrate and strong sulphuric 
 acid, the solution assumes a deep brown colour. Three chemical 
 changes go to make up the reaction : (i) the liberation of nitric 
 acid by the action of sulphuric acid upon the nitrate ; (2) the 
 oxidation of the ferrous to a ferric salt, with elimination of nitric 
 oxide ; and (3) the absorption of the nitric oxide so formed by a 
 further portion of ferrous salt, forming an unstable brown compound 
 having the composition NO,2FeSO 4 
 
 (1) KNO 3 + H 2 SO 4 = HKS0 4 + HNO 3 
 
 (2) 2HNO 3 + 6FeSO 4 4- 3H 2 SO 4 = 3Fe 2 (SO 4 ) 3 + 4H 2 O + 2NO 
 The test is extremely delicate, and is carried out in the following 
 manner. The solution of the nitrate is mixed with about its own 
 volume of strong sulphuric acid in a test-tube, and the mixture 
 cooled. To this a little ferrous sulphate solution is cautiously added, 
 the tube being held in an inclined position, so that the ferrous sul- 
 phate shall float upon the denser liquid already in the tube. Where 
 the two liquids meet, the brown colour will be developed. By a 
 gentle movement of the tube, so as to cause a slight admixture of 
 the liquids at the point where they meet, the brown ring will be still 
 more apparent. The coloured compound is decomposed by heat, 
 with evolution of nitric oxide, hence the necessity for making the 
 test with cold solutions.* 
 
 Reduction by Sulphurous Acid. When copper (or mer- 
 cury) is heated with sulphuric acid in the presence of a nitrate, nitric 
 oxide is evolved, which, in contact with the air, gives red vapours 
 of nitrogen peroxide 
 
 3Cu + 4H 2 SO 4 4- 2KNO 3 = 3CuSO 4 + K 2 SO 4 + 4H 2 O + 2NO 
 The sulphur dioxide (developed by the action of the acid upon the 
 
 * Ferrous chloride in the presence of hydrochloric -acid gives a precisely 
 similar result with a nitrate ; nitric oxide is liberated, which is absorbed by the- 
 ferrous chloride to a brown liquid 
 
 KNO 3 + aFeCl. + 4HCI - 3 FeCl, + KC1 -f 2 H,O -f NO 
 
 L 
 
146 Qualitative Analysis. 
 
 copper) is oxidised by the nitric acid (simultaneously generated by 
 the action of the acid upon the nitrate) to sulphuric acid ; thus 
 
 (1) Cu + 2H 2 SO 4 = CuSO 4 + 2H 2 O + SO 2 
 
 (2) 2HN0 3 + 3S0 2 + 2H 2 = 3H 2 S0 4 4- 2NO 
 
 The nitrate is mixed with a little strong sulphuric acid, and a few 
 fragments of copper foil or turnings are introduced. On boiling 
 the mixture, red fumes of nitrogen peroxide, NO 2 , will appear in the 
 tube, which will be more easily seen by looking down through the 
 mouth of the tube.* 
 
 * A similar reduction takes place, with the liberation of nitric oxide, 
 when sulphur dioxide is generated in presence of a nitrate, from 
 either a sulphite or a thiosulphate, by the action of sulphuric acid. 
 
 A small crystal of sodium thiosulphate is added to the solution 
 of a nitrate, and a few drops of strong sulphuric acid added. On 
 gently warming the mixture, brown fumes are seen in the tube. 
 
 Reduction by Nascent Hydrogen, (i) With Formation of 
 Nitrite. When a nitrate in solution is exposed to the gentle action 
 of nascent hydrogen derived by the action of sodium amalgam, 
 zinc amalgam, or copper-zinc couple the nitrate is reduced to 
 nitrite 
 
 KNO 3 + H 2 = H 2 O + KNO, 
 
 This test is very delicate, and may be carried out as follows : 
 A small piece of zinc foil (or granulated zinc) is placed in the solution 
 of the nitrate in a test-tube, and one drop of copper sulphate added 
 (this causes the deposition of a minute quantity of copper upon the 
 zinc, thus creating the " copper-zinc couple "). The mixture is 
 gently boiled for a minute or two. One drop of the liquid (after 
 cooling) is placed upon a piece of potassium-iodide-and-starch 
 paper,f and then touched with a glass rod moistened with dilute 
 sulphuric acid. The paper will be instantly stained blue by the 
 
 * The mechanism of this reaction is sometimes explained by supposing that 
 the nitric acid, set free from the nitrate, acts upon the copper according to the 
 familiar equation sCu -f 8HNO 3 = 3Cu(NO 3 ) 2 + 4H 2 O + aNO. But, as a 
 matter of fact, copper sulphate, not nitrate, is found in solution, the whole of 
 the nitrogen being converted into nitric oxide. Moreover, charcoal may be 
 substituted for the copper. If dilute nitric acid be boiled with charcoal no 
 brown fumes are formed, but on the addition of a little sulphuric acid they at 
 once appear, owing to the action of the sulphur dioxide which is evolved from 
 carbon and sulphuric acid. 
 
 f Potassium-iodide-and-starch paper is made by dipping ordinary white 
 note paper into moderately thin starch containing a little potassium iodide, and 
 allowing it to dry. It may be preserved indefinitely if kept in a stoppered 
 bottle, and thus obviates the necessity of constantly making a little starch paste. 
 Every student should prepare for himself a stock of the paper. 
 
Nitrous Acid and Nitrites. 147 
 
 liberation of iodine and formation of iodide of starch (test for a 
 nitrite). 
 
 (2) With Formation of Ammonia. By the prolonged action of 
 the copper-zinc couple ; or by means of nascent hydrogen produced 
 by boiling a solution of caustic soda with zinc, nitrates (and 
 nitrites) are reduced to ammonia 
 
 KNO 3 + 4H.j = KHO + 2H,O + NH ;? 
 
 A few drops of the solution of a nitrate are added to caustic soda 
 in a test tube, along with a little granulated zinc. On boiling, 
 ammonia is evolved, which may be detected in the usual way. 
 
 Indigo. Nitric acid oxidises indigo blue (C 8 H 5 NO), converting 
 it into yellow isatin (C 8 H 5 NO 2 ). A little concentrated sulphuric 
 acid is tinted with a minute quantity of indigo, and a few drops of 
 a solution of a nitrate added. On heating the mixture the blue 
 colour disappears. (The yellow of the isatin is scarcely visible in 
 small quantities, hence the solution appears bleached.) 
 
 Nitrates all undergo decomposition when strongly heated. 
 Nitrates of alkali metals and alkaline earths, when gently heated, 
 are reduced to nitrites, with evolution of oxygen. 
 
 Ammonium nitrate passes into water and nitrous oxide. Other 
 nitrates, e.g. lead nitrate, leave an oxide of the metal, and give off 
 oxygen and nitrogen peroxide. 
 
 When heated with oxidisable substances (carbon, sulphur, etc.) 
 the decomposition is propagated with explosive violence. Thus, 
 when nitrates are heated before the blowpipe on charcoal, deflagra- 
 tion of the charcoal takes place. 
 
 Nitrates and chlorates, when present together, are examined by 
 being first converted by heat into nitrites and chlorides. If present 
 as salts of metals other than the alkalies, sodium carbonate is 
 added, and the dry mixture heated until the evolution of oxygen is 
 at an end. The residue is extracted with water, and the solution 
 examined for nitrites and chlorides. 
 
 If chlorides are originally present as well as nitrates and 
 chlorates, they must be first removed by precipitation with silver 
 sulphate, as explained on p. 133. 
 
 Nitrous Acid and Nitrites. 
 
 The acid is not known in the pure state. Even when liberated 
 in dilute solutions, it speedily breaks up into nitric acid, nitric oxide 
 and water. Hence when nitrites are decomposed by acids, nitric 
 
148 Qualitative Analysis. 
 
 oxide is evolved, which, in contact with atmospheric oxygen, passes 
 into the brown gas NO 2 ; thus 
 
 6NaNO, + 3H 2 SO 4 = 3Na 2 SO 4 + 2HNO 3 + 2H 2 O + 4NO 
 
 Nitrites are all soluble in water, but the silver salt is suffi- 
 ciently difficult of solution to be precipitated, on the addition of 
 silver nitrate, to a (not too dilute) solution of a nitrite. 
 
 All nitrites are easily decomposed by dilute acids in the cold, 
 with evolution of nitric oxide, as shown in the above equation. If 
 the action takes place in the presence of a ferrous salt, the same 
 brown-coloured compound is produced as in the case of a nitrate. 
 [Nitrites therefore give a " brown ring," when dilute sulphuric, or 
 even acetic, acid is used (distinction from nitrates).'] 
 
 Oxidation Reactions. Nitrous acid and nitrites part with 
 oxygen, and are converted into nitric oxide 
 
 2HNO 2 = 2ND + H 2 O -4- O 
 
 Thus, when a nitrite is acidified with dilute sulphuric acid in the 
 presence of potassium iodide, the nitrous acid first formed oxidises 
 the potassium iodide (or hydriodic acid), setting free the iodine : 
 the liberated iodine is detected by its action upon starch 
 
 2K1 + H 2 O + O = 2KHO 4- L 
 
 A few drops of a solution of a nitrite are placed upon potassium 
 iodide-and-starch paper, and a single drop of very dilute sulphuric 
 acid added (by means of a glass rod dipped in the acid) when a 
 blue stain at once appears upon the paper. [Or starch emulsion 
 may be added to the solution of the nitrite, then a drop or two of 
 potassium iodide, and lastly a small quantity of acid.] 
 
 Sulphuretted hydrogen is similarly oxidised by a nitrite in 
 presence of an acid, with precipitation of sulphur 
 
 H 2 S + 2HNO 2 = 2H 2 O + 2NO + S 
 
 Nitrites of the alkali metals are decomposed by sulphuretted 
 hydrogen without the addition of an acid, giving alkaline sulphides. 
 
 Reduction Reactions. Nitrous acid, by absorption of 
 oxygen, passes into nitric acid ; it therefore is capable of reducing 
 other compounds, such as chromates, permanganates, mercurous 
 (but not mercuric) salts, and gold compounds. Potassium per- 
 manganate (in presence of acid) is converted into manganous salts, 
 the violet colour of the permanganate being destroyed. Gold and 
 mercurous salts are reduced to the respective metals ; thus 
 
Phosphoric Acid and Phosphates. 149 
 
 HgXlo + H,0 + HN0 2 = HN0 3 + 2HC1 + 2Hg 
 2AuCl 3 + 3H 2 6 + 3HNO 2 = 3HNO 3 + 6HC1 + 2Au 
 
 Detection of Nitrites and Nitrates in the Same Solu- 
 tion. Owing to the ready decomposition of nitrites by dilute 
 acids, they are easily detected in presence of nitrates, either by the 
 liberation of iodine, the oxidation of ferrous salts, or reduction of 
 potassium permanganate. To find a nitrate when nitrites are 
 present is less simple. The dilute solution of the mixed nitrate 
 and nitrite is acidified with three or four drops of dilute sulphuric 
 acid, and a little ferrous sulphate solution (or a small crystal of the 
 salt) is added. The solution at once becomes dark brown (owing 
 to the absorption, by the ferrous salt, of the nitric oxide liberated 
 from the nitrite). It is then heated (but not allowed to boil), with 
 frequent shaking, when nitric oxide is expelled, and the liquid 
 gradually becomes colourless. The mixture is cooled, and one drop 
 more dilute acid added, and a little more ferrous sulphate. (If all 
 the nitrite present has been decomposed, this addition gives no 
 further coloration.) This solution is now poured carefully on to a 
 small quantity of strong sulphuric acid in a test-tube, so as to float 
 upon the acid, and where the liquids meet a " brown ring " will be 
 formed, due to the nitrate present. 
 
 Phosphorus. 
 
 The property which phosphorus possesses of emitting a feeble 
 luminosity when exposed to the air, either in the solid state or 
 when vaporised, is made use of in analysis for the detection of the 
 element in the free state. The substance, mixed with water, is 
 boiled in a flask with a narrow neck in a dark room. The issuing 
 steam will then appear luminous as it escapes from the flask ; and 
 presents the appearance of a lambent pale greenish flame. A piece 
 of paper moistened with silver nitrate, when held over the neck of 
 the flask, will become blackened, owing to the formation of silver 
 phosphide.* 
 
 Phosphoric Acid and Phosphates. 
 
 Three phosphoric acids (each with its series of phosphates) are 
 known, namely, orthophosphoric acid, H 3 PO 4 ; pyrophosphoric 
 acid, H^PoO 7 ; and metaphosphoric acid, HPO ;i . 
 
 * Phosphorus is a violent poison, and it is practically only in cases of 
 toxicologieal examination that free phosphorus is sought for. The allotropic 
 form ("red phosphorus ") is not poisonous. It is changed by heat into the 
 ordinary variety. 
 
150 Qualitative Analysis. 
 
 Orthophosphoric acid is tribasic, and hence produces three 
 classes of orthophosphates, by the replacement of one, two, or 
 three of the hydrogen atoms ; e.g. 
 
 Ammonium magnesium phosphate, (NH 4 )MgPO 4 . 
 
 Hydrogen sodium ammonium phosphate (microcosmic salt\ 
 HNa(NH 4 )PO 4 . 
 
 Dihydrogen sodium phosphate, H 2 NaPO 4 . 
 
 Normal orthophosphates (not containing a volatile base, as 
 NH 4 ) are not decomposed when heated alone ; those containing 
 one or two hydrogen atoms, or ammonium as the base, are con- 
 verted into pyro or meta salts (see below). The only orthophos- 
 phates which are soluble in water are those of the alkali metals. 
 
 Silver nitrate gives a yellow precipitate with soluble 
 phosphates, of silver phosphate, Ag 3 PO 4 , which distinguishes ortho 
 from pyro and meta compounds. 
 
 The chief analytical reactions of the orthophosphates have 
 already been considered in Chap. VII., p. 63. 
 
 Fyrophosphoric Acid and Pyrophosphates. When 
 Orthophosphoric acid, Qr a phosphate containing either one hydro- 
 gen atom or one ammonium radical, is heated, it loses water, and 
 is converted into pyrophosphoric acid or a salt ; thus 
 
 2H 3 P0 4 - H 2 + H 4 P 2 7 
 2HNa 2 PO 4 = H 2 O + Na 4 P 2 O 7 
 2NH 4 MgP0 4 = H 2 + 2NH 3 + Mg 2 P 2 O 7 
 
 Boiling with acids retransforms pyrophosphates into ortho- 
 phosphates. 
 
 Only the pyrophosphates of the alkalies are soluble in water. 
 
 Silver nitrate gives a white precipitate of silver pyrophos- 
 phate, Ag 4 P 2 O 7 . 
 
 Magnesium sulphate precipitates white magnesium pyro- 
 phosphate, Mg 2 P 2 O-, soluble in excess of magnesium sulphate, and 
 not reprecipitated in the cold by ammonia (distinction from ortho- 
 phosphates]. 
 
 Ammonium molybdate gives no precipitate until the " pyro " 
 acid has been changed to " ortho " by the action of the nitric acid 
 present. 
 
 Metaphosphoric Acid and Metaphosphates.f The acid 
 is formed when the " ortho " or " pyro " acids are strongly heated, 
 whereby water is expelled. It is known as glacial phosphoric 
 acid 
 
 H 3 PO 4 = H 2 O + HPO 3 
 
 * Metaphosphoric acid has the curious property of forming a number of 
 salts which may be regarded as polymers of the ordinary salts (see Newth's 
 " Inorganic Chemistry," p. 437). 
 
Hypophosphorous Acid. 151 
 
 The reaction is reversible ; for when metaphosphoric acid is dis- 
 solved in water, it passes back to the ortho acid slowly in the 
 cold, quickly when boiled. 
 
 Silver nitrate gives a white precipitate of silver metaphos- 
 phate, AgPOg. 
 
 Magnesium sulphate, in presence of ammonium chloride, 
 gives no precipitate (distinction from " pyro " and " ortho" acids). 
 
 Albumen (white of egg) is coagulated when shaken up with 
 metaphosphoric acid (or metaphosphates acidified with acetic acid) 
 (distinction from "pyro " and " ortho " acids, which are 'without action 
 upon albumen}. 
 
 Phosphorous Acid and Phosphites. 
 
 Phosphorous acid, H 3 PO 3 , is tribasic, but the most stable salts 
 are those which contain one atom of hydrogen. Thus, trisodium 
 phosphite, Na 3 PO 3 , is decomposed by water into hydrogen disodium 
 phosphite, HNa 2 PO 3 . With the exception of phosphites of the 
 alkalies, the salts of this acid are either insoluble in water or dis- 
 solve only with difficulty, and may therefore be precipitated by 
 double decomposition ; thus, on the addition of -barium chloride 
 to a solution of sodium phosphite, barium phosphite, HBaPO 3 , is 
 thrown down as a white precipitate. Both the acid and its salts 
 are powerful reducing agents ; thus, - with silver nitrate a white 
 precipitate momentarily forms, which quickly becomes black from 
 reduced silver. Similarly, mercuric compounds are converted first 
 to mercurous and then to metallic mercury. When strongly heated, 
 phosphites decompose, and give off phosphoretted hydrogen. 
 
 Hypophosphorous Acid and Hypophosphites. 
 
 Hypophosphorous acid, H 3 PO 2 , is monobasic ; its formula may 
 therefore be written H(H 2 PO 2 ). The sodium salt has the com- 
 position Na(H 2 PO 2 ), while the barium salt is expressed by the 
 formula Ba(H 2 PO 2 ) 2 . All the salts are soluble in water, therefore 
 barium chloride gives no precipitate with solutions of hypophos- 
 phites (distinction from phosphites}. 
 
 Hypophosphorous acid and its salts are still more powerful 
 reducing agents than phosphites. One characteristic reaction of 
 this nature, which distinguishes hypophosphites from phosphites, 
 is the reduction of copper sulphate to copper hydride, Cu 2 H 2 . 
 
 On adding copper sulphate to an acidulated solution of a 
 hypophosphite, and gently warming the mixture, a precipitate is 
 obtained at first yellowish-brown, quickly turning to a dark 
 chocolate-brown colour ; thus 
 
 This precipitate of cuprous hydride is distinguishable from the red 
 cuprous oxide, not only by its much darker colour, but by the 
 
152 Qualitative Analysis. 
 
 fact that when treated with strong hydrochloric acid, it evolves 
 hydrogen ; thus 
 
 Cu 2 H 2 + 2HC1 = Cu 2 Cl 2 + 2H 2 
 
 Under the influence of nascent hydrogen from zinc and hydro- 
 chloric acid, hypophosphites as well as phosphites give phospho- 
 retted hydrogen, which imparts to the hydrogen, when it is inflamed, 
 the characteristic colour of the phosphorus flame. 
 
 Hypophosphorous acid and its salts, when gently heated, give 
 off phosphoretted hydrogen, leaving phosphoric acid 
 
 2 H 3 P0 2 = H 3 P0 4 + PH 3 
 
CHAPTER XIV. 
 CARBON, SILICON, BORON. 
 
 Carbon. 
 
 THE properties of carbon by which it is most readily recognised 
 are (i) its combustibility in air or oxygen, with formation of only 
 carbon dioxide ; and (2) its power to withstand the action of chlorine 
 even at high temperatures, which distinguishes it from all metals 
 or black substances. 
 
 Carbonic Acid and Carbonates. 
 
 Carbonic Acid, H 2 CO 3 , is an unstable compound only capable 
 of existence in dilute aqueous solution. It is formed when carbon 
 dioxide is dissolved in water, and has a feeble acid reaction. It is 
 capable of dissolving the normal carbonates of the alkaline earths 
 and of magnesium, forming the so-called "acid" carbonates or 
 " bicarbonates " (the pharmaceutical preparation known as fluid 
 magnesia is a case in point) 
 
 MgCOa + H 2 C0 3 = H 2 Mg(C0 3 ) 2 
 
 All normal carbonates are insoluble in water, except those of the 
 alkalies. The "acid" carbonates are all soluble, but on boiling 
 their solutions, they are converted into normal salts, which (except 
 in the case of the alkaline carbonates) are then precipitated ; thus 
 
 H 2 Ca(C0 3 ) 2 = CaC0 3 + H 2 O + CO, 
 
 (The formation of boiler incrustations and the " furring " of kettles 
 are due to this decomposition.) 
 
 The normal carbonates of the alkalies are by this reaction 
 readily distinguished from the bicarbonates. Thus, on heating a 
 solution of sodium bicarbonate, effervescence rapidly sets in owing 
 to the escape of carbon dioxide 
 
 2HNaCO 3 = Na 2 CO 3 + H 2 O + CO, 
 These two classes of salts may also be distinguished by the 
 
154 Qualitative Analysis. 
 
 difference in their behaviour towards certain metallic solutions, 
 e.g. magnesium or mercuric salts. Thus 
 
 With magnesium sulphate or chloride, normal sodium 
 carbonate gives a white precipitate of a basic carbonate 
 
 5MgS0 4 + 5Na 2 C0 3 = MgO,4MgCO 3 + 5Na 2 SO 4 + CO, 
 Sodium bicarbonate gives no precipitate with magnesium salts, the 
 reason being that in the presence of the large quantity of carbonic 
 acid set at liberty from the bicarbonate, the magnesium carbonate 
 is redissolved, forming the soluble magnesium bicarbonate. 
 
 With mercuric chloride, the normal carbonate gives a 
 reddish precipitate of a basic oxide, while the bicarbonate gives 
 no precipitate. 
 
 All carbonates are decomposed by dilute hydrochloric acid 
 (and by nearly all acids) with effervescence, due to the rapid escape 
 of carbon dioxide. The gas is identified by its action upon lime- 
 water (or baryta-water). 
 
 The test is made by adding a few drops of acid to the carbonate 
 in a test-tube, and decanting the evolved (heavy) gas into a second 
 test-tube containing a little lime-water, Ca(HO) 2 . On shaking the 
 lime-water with the gas, the liquid becomes milky, owing to the 
 precipitation of calcium carbonate. 
 
 The only other gas which gives a white precipitate with lime- 
 water is sulphur dioxide. This is 
 easily distinguished from carbon 
 dioxide by its smell. If the two 
 gases are present together, the 
 sulphur dioxide (recognised by 
 its odour) may be removed by 
 means of potassium permanga- 
 nate. The two gases (liberated 
 simultaneously by the action of 
 an acid upon a mixture of a car- 
 bonate and sulphite) are passed 
 through a little potassium per- 
 manganate solution in a test-tube, 
 fitted as shown in Fig. 14. The 
 sulphur dioxide is absorbed (being 
 oxidised by the permanganate into 
 sulphuric acid), and the carbon 
 dioxide passes on, and can be de- 
 tected by means of lime-water. 
 If the passage of the gases through 
 FlG - J 4- ' the permanganate be continued 
 
 for. a few minutes, the colour of the solution becomes entirely 
 destroyed ; and the liquid may then be tested for sulphuric acid in 
 the usual way. 
 
Oxalic Acid and Oxalates. 155 
 
 When strongly heated, the normal carbonates of the alkali 
 metals (not ammonium) remain unchanged. Those of the alkaline 
 earths are converted at a high temperature into oxides, with 
 evolution of carbon dioxide (illustrated in the process of lime- 
 burning). All other carbonates are more readily decomposed by 
 heat. 
 
 Formic Acid and Formates. 
 
 Formic acid, H 2 CO 2 or H(COHO) (the acid present in ants 
 and in nettles), is a pungent-smelling colourless liquid. The acid 
 is monobasic. All formates are soluble in water. Both the acid 
 and the salts are powerful reducing agents, as one molecule of the 
 acid is capable of withdrawing one atom of oxygen (or its equivalent 
 in chlorine) from compounds capable of reduction. Thus, mercuric 
 chloride is reduced to mercurous chloride, and carbon dioxide is 
 eliminated-T- 
 
 2HgCl 2 + H 2 CO 2 = CO 2 + 2HC1 + Hg 2 CL 
 
 Formic acid and formates are immediately decomposed by concen- 
 trated sulphuric acid without the application of heat, into water 
 (absorbed by the acid) and carbon monoxide 
 
 H(COHO)=.H 2 + CO 
 
 The escaping carbon monoxide is distinguished by burning with a 
 fine pale-blue flame. 
 
 Oxalic Acid and Oxalates. 
 
 Oxalic acid, H 2 C 2 O 4 , is a white solid, soluble in water, and 
 crystallising from the solution with two molecules of water of 
 crystallisation, H 2 C 2 O 4 ,2H 2 O. The acid is dibasic. 
 
 The oxalates of the alkalies are soluble in water, as well as 
 certain acid oxalates (e.g. acid oxalate of barium). All other 
 oxalates are insoluble or only sparingly soluble. 
 
 Insoluble oxalates are- decomposed by mineral acids, with the 
 formation of salts of the mineral acid with the metal contained in 
 the oxalate, and the liberation of oxalic acid. In the case of strong 
 sulphuric acid, the liberated oxalic acid is itself decomposed on 
 the application of a gentle heat ; carbon dioxide and monoxide 
 being disengaged in equal volumes ; thus 
 
 H 2 C 2 O 4 = H 2 O + CO 2 + CO 
 
 the sulphuric acid in this instance, as in the case of formic acid, 
 taking up the elements of water from the molecule. 
 
 In the presence of reducible compounds, oxalic acid passes into 
 water and carbon dioxide only, one molecule of the acid taking up 
 
156 Qualitative Analysis. 
 
 one atom of oxygen. In this way it is able to reduce potassium 
 permanganate, two molecules of which furnish five atoms of 
 oxygen, and therefore oxidise five molecules of oxalic acid ; or one 
 molecule of permanganate causes the liberation of five molecules 
 of carbon dioxide 
 
 2 KMnO 4 + 3H 2 SO 4 = K 2 SO 4 + 2MnSO 4 + 3H,O + 50 
 50 + 5H 2 C 2 O 4 = sH 2 O + ioCO 2 
 
 Similarly, manganese dioxide, having one atom of available oxygen, 
 when passing to the condition of a manganous salt, is able to 
 oxidise one molecule of oxalic acid ; thus 
 
 MnO 2 + H 2 SO 4 + H 2 C 2 O 4 = MnSO 4 + 2H 2 O + 2CO 2 
 
 Of the insoluble oxalates, the calcium salt, CaC 2 O 4 , is of the most 
 importance analytically. Its formation as a special test for 
 calcium has already been considered (p. 33). The precipitate is 
 only slightly soluble in oxalic acid (whereas barium oxalate is 
 somewhat readily dissolved by oxalic acid, forming a soluble acid 
 oxalate), and is scarcely dissolved by acetic acid. It is, however, 
 readily soluble in hydrochloric acid. 
 
 When heated, oxalates are all decomposed. Those of the 
 alkalies and alkaline earths are converted into carbonates, with 
 evolution of carbon monoxide. The oxalates of metals which 
 either do not form carbonates, or whose carbonates are decomposed 
 by heat, leave a metallic oxide, and give off carbon dioxide and 
 monoxide either together or in two stages ; thus 
 
 (1) CaC 2 O 4 = CaCO 3 + CO 
 
 (2) CaC0 3 = CaO + CO 2 
 
 In the case of metals whose oxides are decomposed by heat, e.g. 
 silver oxalate, the metal is left. 
 
 Acetic Acid and Acetates. 
 
 Acetic acid, C 2 H 4 O 2 or H(C 2 H 3 O 2 ), is a white crystalline 
 solid, which melts at the temperature of a warm room (16 C.). 
 The common reagent is a solution of the acid in water. The acid 
 is monobasic. The acetates are all soluble in water, the silver and 
 mercurous salts being sufficiently difficult of solution to allow of 
 their being precipitated by the addition of moderately strong acetic 
 acid to solutions of the respective metals. Ferric chloride, when 
 added to a solution of an acetate, gives ferric acetate, which is a 
 soluble salt imparting a dark red colour to the solution. When the 
 liquid is boiled, a brownish precipitate separates out, consisting of 
 
Tar tar ic and Citric Acids. 157 
 
 basic ferric acetate (p. 69). The acetates are decomposed by mineral 
 acids, acetic acid being liberated ; and as acetic acid is volatile 
 without undergoing decomposition, it can be distilled off and 
 recognised by its odour. When an acetate is mixed with strong 
 sulphuric acid and a little alcohol, (C 2 H 5 )HO, and the mixture 
 gently heated, a reaction takes place, resulting in the formation of 
 ethyl acetate (acetic ether), a volatile, ethereal, and pleasant-smelling 
 compound.* The mechanism of the reaction is as follows: (i) 
 The sulphuric acid acts upon the alcohol, forming a compound 
 known as sulphethylic acid, H(C 2 H 5 )SO 4 ; and (2) this then reacts 
 with the acetic acid liberated by the sulphuric acid from the 
 acetate ; thus 
 
 (1) H 2 S0 4 + (C 2 H 5 )HO = H(C 2 H 5 ;S0 4 + H 2 O 
 
 (2) H(C 2 H 5 )S0 4 + H(C 2 H 3 2 ) = H 2 S0 4 + C 2 H 5 (C 2 H 3 O,) 
 
 When acetates are heated alone, they blacken slightly, and evolve 
 vapours of a volatile liquid known as acetone. The vapour has a 
 characteristic smell,* and is inflammable. 
 
 Tartaric Acid and Tartrates. 
 
 Tartaric acid, H 2 (C 4 H 4 O 6 ), is a white crystalline solid ; decom- 
 posed by heat, and therefore not volatile. It is soluble in water. 
 The acid is dibasic. 
 
 The normal tartrates of the alkalies are freely soluble in water ; 
 the " acid " tartrates (especially of potassium and ammonium (see 
 pp. 2 1 and 23) are sparingly soluble. Other normal tartrates are either 
 insoluble or difficultly soluble, but most of them dissolve in tartaric 
 acid, forming " acid " salts. The tartrates show a great tendency 
 to form soluble double salts with alkalies ; thus the tartrates of the 
 metals of Group III. are dissolved by alkalies, owing to the forma- 
 tion of double tartrates. Hence, in the presence of tartaric acid, 
 alkalies fail to give precipitations with solutions of those metals. 
 
 Tartrates are decomposed by mineral acids, with elimination of 
 tartaric acid. If heated with strong sulphuric acid, the liberated 
 tartaric acid is broken up into water, carbon dioxide, carbon 
 monoxide, and separation of carbon (the mixture blackens there- 
 fore) \ and from the action of the sulphuric acid upon the carbon, 
 sulphur dioxide is formed 
 
 H 2 (C 4 H 4 6 ) = 3H 2 + C0 2 + CO + 2C 
 C + 2H 2 SO 4 = 2H 2 O 4- CO 2 + 2SO 2 
 
 Calcium hydroxide (lime-water), added to a solution of 
 tartaric acid, gives no precipitate, unless added in sufficient excess 
 
 * No description of odours of this kind can convey an exact idea of them. 
 It is therefore only by preparing and smelling the compound, and so actually 
 learning its characteristic smell, that the substance can afterwards be recog- 
 nised by such a test as this. 
 
158 Qualitative Analysis. 
 
 to ensure the formation of the normal tartrate, which then pre- 
 cipitates out. 
 
 Soluble calcium salts (not the sulphate) give the same precipi- 
 tate when added to solutions of normal tartrates. The precipitate 
 is soluble in tartaric acid, and also in acetic acid (contrast Calcium 
 oxalate). 
 
 Silver nitrate gives, with solutions of normal tartrates, a 
 white precipitate of silver tartrate. The precipitate is soluble in 
 ammonia, and on gently warming the ammoniacal solution, the 
 silver is precipitated as a coherent film or mirror upon the glass 
 vessel. The test is made in the following way : 
 
 A small quantity of a normal tartrate (the double tartrate of 
 sodium and potassium, Rochelle salt, is a suitable salt to employ 
 in order to study the reaction), not much larger than a pin's head, 
 is dissolved in a little water in a carefully cleaned test-tube, and a 
 few drops of silver nitrate added. Dilute ammonia is then added 
 drop by drop, until the precipitated silver tartrate is nearly wholly 
 dissolved. The mixture is then diluted with water, so as to about 
 half fill the test-tube. The tube is then placed into boiling water 
 for a few minutes, when the silver will be deposited as a brilliant 
 mirror upon the glass. ( This reaction distinguishes tartaric acid 
 from all the other acids treated in this chapter?) 
 
 When heated alone, tartaric acid and tartrates are decomposed, 
 evolving vapours which are inflammable, and emitting a smell 
 resembling that of burnt sugar. The residue, in the case of 
 tartrates of the alkalies and alkaline earths, consists of carbon and 
 the carbonate of the alkali ; while other tartrates leave either a 
 metallic oxide or metal mixed with carbon. 
 
 Citric Acid and Citrates. 
 
 Citric acid, H 3 (C 6 H^O 7 ), closely resembles tartaric acid in 
 appearance. It is soluble in water. The acid is tribasic. 
 
 The citrates of the alkalies are readily soluble ; those of the 
 alkaline earths are less easily soluble ; calcium citrate, however, 
 being soluble in cold water, but nearly insoluble in hot water. 
 
 Calcium hydroxide, or calcium chloride, if moderately 
 dilute, gives no precipitate in the cold, when added to solutions 
 of citric acid or normal citrates. (A strong solution of calcium 
 ide mixed with a strong solution of a citrate may give a 
 precipitate even in the cold. With more dilute solutions, the 
 calcium citrate is gradually precipitated ; completely only after a 
 few hours) ; but on heating the mixture, a white precipitate of 
 calcium citrate is produced (distinction from tartaric acid). The 
 precipitate is soluble in acetic acid (distinction from calcium 
 oxalate}. 
 
 By means of the reaction with calcium chloride, tartaric and 
 citric acids may be detected when mixed together. Calcium 
 chloride is added to the cold and not too concentrated solution, 
 whereby calcium tartrate is precipitated ; on filtering and heating 
 the filtrate, a further precipitation takes place of calcium citrate. 
 
Hydrocyanic Acid and Cyanides. 159 
 
 Silver nitrate gives a white precipitate of silver citrate, 
 soluble (like the tartrate) in ammonia ; but the ammoniac al solu- 
 tion does not deposit metallic silver when warmed (compare Silver 
 tartrate). 
 
 Cyanogen. 
 
 Cyanogen, (CN) or Cy, is a negative compound radical, which 
 resembles the halogen elements in many of its chemical habits. 
 It is liberated when the cyanides of many of the heavy metals (e.g. 
 mercuric cyanide) are heated alone. 
 
 Cyanogen is a colourless gas, with an odour which is like the 
 taste of bitter almonds or the kernels of cherry-stones.* It 
 burns with a characteristic flame, having a beautiful mauve-pink 
 colour. The gas is soluble in water, but the solution is unstable. 
 It also dissolves in yellow ammonium sulphide, forming ammonium 
 thiocyanate 
 
 (NH 4 ). 2 S 2 + 2Cy = 2NH 4 CyS 
 
 This reaction forms a ready method for the detection of small 
 quantities of cyanogen. The test is made in the following way. 
 A piece of filter-paper, moistened with a drop of yellow ammonium 
 sulphide, is held over the mouth of the test-tube in which the 
 compound is being heated. If the yellow colour of the sulphide is 
 destroyed, the colourless spot is touched with a glass rod dipped 
 in ferric chloride, when a red stain of ferric thiocyanate is at once 
 obtained. If the quantity of cyanogen is so small that the colour 
 of the ammonium sulphide is not destroyed, a single drop of a 
 solution of zinc chloride or sulphate is placed upon it, using a 
 pipette (this removes the excess of ammonium sulphide, forming 
 white zinc sulphide), after which the spot is touched with ferric 
 chloride, when the red stain will appear. (Ferric chloride cannot 
 be added until the ammonium sulphide is removed, or a black 
 stain of ferrous sulphide will be produced.) 
 
 Hydrocyanic Acid and Cyanides. 
 
 Hydrocyanic acid, HCN or HCy, is the analogue of the 
 hydrogen acids of the halogens. In the pure state it is a very 
 volatile liquid (B.P., 27 C.). It is liberated when certain cyanides 
 (see below) are decomposed by acids, and may be recognised by 
 
 * Cyanogen, and also hydrocyanic acid, are most deadly poisons, and 
 therefore the greatest caution must be exercised in smelling either of these 
 compounds. The odour of the greatly diluted gas seems to be perceived more 
 by the sense of taste than of smell, being recognised as a somewhat bitter but 
 not unpleasant sensation at the back of the throat. 
 
160 Qualitative Analysis. 
 
 its odour (see footnote, p. 159), or by applying the reaction with 
 ammonium sulphide as described for cyanogen 
 
 (NH 4 ) 2 S 2 + HCy = (NH 4 )HS + NH 4 CyS 
 
 Single Cyanides. These are binary compounds, analogous to 
 the chlorides. Of these the cyanides of the alkalies, alkaline earths, 
 and mercuric cyanide are soluble in water (barium cyanide with 
 difficulty) ; all others are insoluble. The aqueous solutions of 
 these cyanides are unstable, undergoing gradual decomposition ; 
 chiefly into a formate and free ammonia 
 
 KCN + 2H 2 O = K(CHO 2 ) + NH, 
 
 The cyanides of the alkali metals are unchanged when strongly 
 heated alone, while those of the heavy metals are mostly converted 
 into cyanogen * and the metal, or in some cases into nitrogen, 
 carbon, and metal, or a metallic carbide. Under the influence of 
 heat, the cyanides of the alkalies readily take up oxygen or sulphur, 
 from compounds capable of yielding these elements, and pass into 
 cyanate or thiocyanate. Hence potassium cyanide is constantly 
 employed in analysis as a reducing agent in blowpipe reactions. 
 
 The alkali cyanides, and some of the insoluble single cyanides 
 (e.g. ZnCy 2 , PbCy 2 ) are readily decomposed by dilute mineral 
 acids ; others are only decomposed with difficulty ; in either case, 
 with liberation of hydrocyanic acid. They are all decomposed 
 by strong sulphuric acid (some requiring the aid of heat), with 
 formation of metallic sulphates (see p. 164). Mercuric cyanide is 
 decomposed by sulphuretted hydrogen. Cyanides in solution may 
 be detected (i) by the precipitation of insoluble single cyanides by 
 double decomposition (not very characteristic) ; (2) by the formation 
 of alkali thiocyanate, and subsequently obtaining the red coloration 
 with ferric chloride ; (3) by the formation of Prussian blue. 
 
 (i) Silver nitrate, added to a solution of a soluble cyanide 
 (other than mercuric cyanide), gives a white precipitate of silver 
 cyanide, AgCy ; soluble in ammonia, insoluble in nitric acid, and 
 readily soluble in potassium cyanide, giving a double cyanide, 
 KCy,AgCy. 
 
 Silver cyanide closely resembles silver chloride ; it is, however, 
 readily distinguished by the fact that, when boiled with hydrochloric 
 acid, it is decomposed, with evolution of hydrocyanic acid ; and 
 also that, when heated alone, it gives off cyanogen (detected as 
 above), and leaves a black residue of silver and paracyanogen. 
 
 * And usually some paracyanogen, a non-volatile black compound, which 
 is a polymeride of cyanogen expressed by the formula (CN).t . 
 
Hydrocyanic Acid and Cyanides. 161 
 
 (2) Formation of Alkali Thiocyanate. Yellow ammonium 
 sulphide is added to the solution of the cyanide, and the mixture 
 gently boiled for a moment or two. In the case of an alkali cyanide, 
 the following change takes place : 
 
 (NH 4 ) 2 S 2 + KCy = KCyS + (NH 4 ) 2 S 
 
 With mercuric cyanide, the ammonium sulphide precipitates black 
 mercuric sulphide ; thus 
 
 HgCy 2 + 2(NH 4 ) 2 S 2 = HgS + (NH 4 ) 2 S + 2NH 4 CyS 
 
 A solution of zinc sulphate is then added (without filtering, in 
 the case of mercuric cyanide), until the excess of ammonium 
 sulphide is removed by precipitation, as white zinc sulphide. The 
 mixture is then filtered, and to the colourless solution a drop of 
 ferric chloride is added, which at once gives the deep red coloration 
 due to ferric thiocyanate. (This test is very delicate.) 
 
 (3) Formation of Prussian Blue. To the solution con- 
 taining the soluble cyanide (or hydrocyanic acid) a small quantity of 
 sodium or potassium hydroxide is added, after which three or four 
 drops both of a ferrous and ferric salt (ferrous sulphate and ferric 
 chloride). The visible result is the production of a dirty brown 
 precipitate of the mixed ferrous and ferric hydroxides. The invisible 
 result is the formation of ferrocyanide in the solution. The ferrous 
 sulphate, interacting with the cyanide present, forms (i) ferrous 
 cyanide, FeCy 2 , which in its turn dissolves in the excess of the 
 cyanide, forming the ferrocyanide ; thus, with potassium cyanide 
 
 (1) FeSO, 4- 2KCy = K 2 SO 4 4- FeCy 2 
 
 (2) FeCy 2 + 4 KCy = K 4 (FeCy 6 ) 
 
 On the addition of hydrochloric acid to the mixture, the pre- 
 cipitated iron hydroxides are dissolved, and the ferric chloride thus 
 formed interacts with the potassium ferrocyanide present, giving 
 rise to the precipitation of ferric ferrocyanide, or, in cases when 
 the amount of cyanide is very small, to a blue or green coloration. 
 
 Double Cyanides. Most of the single cyanides are soluble 
 in alkali cyanides, giving rise to double cyanides. Some of these 
 have an alkaline reaction, are easily decomposed by dilute mineral 
 acids with evolution of hydrocyanic acid, and (like the single 
 cyanides) are poisonous. Others, on the contrary, have a neutral 
 reaction, do not give off hydrocyanic acid when treated with dilute 
 acids, and are not poisonous. The double cyanides are therefore 
 divided into two classes, namely (i) those which are easily decom- 
 posed by acids, and (2) those which are decomposed with difficulty. 
 
1 62 Qualitative Analysis. 
 
 All double cyanides, like all single cyanides, are decomposed by 
 strong sulphuric acid, and also by fusion with potassium nitrate 
 (or a mixture of ammonium nitrate and sulphate}. 
 
 (i) Easily decomposed Double Cyanides. From an 
 analytical point of view, these may all be regarded as mixtures 
 of the single cyanide with the solvent cyanide (usually potassium 
 cyanide). Thus, the double cyanides obtained by dissolving the 
 cyanides of zinc, silver, nickel, in potassium cyanide, are expressed 
 by the formulae 2KCy,ZnCy 2 , KCy,AgCy, 2KCy,NiCy 2 . 
 
 Their chemical behaviour is practically the same as would be 
 shown by mixtures of potassium cyanide and the several simple 
 metallic cyanides. Thus, when such double cyanides are heated, 
 the alkali cyanide (which is stable at high temperature) remains, 
 and the other portion of the molecule behaves exactly as the single 
 cyanide \'g. the silver cyanide evolves cyanogen, leaving a black 
 residue of silver and paracyanogen. 
 
 When treated with acids, the same thing is observed ; the 
 alkali cyanide is decomposed with evolution of hydrocyanic acid 
 (just as it would if alone), while the other cyanide in the molecule 
 behaves exactly as- though it were alone. If it is not one which is 
 decomposed by the acid, it remains ; reprecipitated of course, since 
 the solvent cyanide has been decomposed. Thus, the double 
 potassium nickel cyanide, when acted upon by acids, evolves hydro- 
 cyanic acid, and the single nickel cyanide is reprecipitated 
 
 2KCy,NiCy 2 + 2HC1 = 2KC1 + 2HCy + NiCy 2 
 
 In the case of the double potassium zinc cyanide, not only is 
 the potassium cyanide decomposed, but also the zinc cyanide ; 
 thus 
 
 2KCy,ZnCy 2 + 4HC1 = 2KC1 + ZnCl, + 4HCy 
 
 Some of these double cyanides are decomposed by sulphuretted 
 hydrogen, with precipitation of the sulphide of the heavy metal, 
 e.g. 2KCy,HgCy 2 , 2KCy,CdCy 2 , and 2KCy,ZnCy 2 . In other cases, 
 however, sulphuretted hydrogen fails to precipitate the metals as 
 sulphides, e.g. 2KCy,CuCy 2 or 6KCy,Cu 2 Cy 2 , and 2KCy,NiCy 2 (see 
 also pp. 60 and 83). 
 
 All the double cyanides of this class are decomposed by being 
 boiled with mercuric oxide, with the formation of mercuric cyanide 
 and precipitation of the oxide or hydroxide of the heavy metal- - 
 
 The detection of the cyanogen in these double compounds may 
 
Ferrocyan ides. 163 
 
 be accomplished by the same reactions as those already described 
 for single cyanides. To recognise the heavy metal (as in the case 
 of single cyanides), the compound must be completely decomposed 
 by one of the various methods given, either by boiling with dilute 
 hydrochloric acid, or by strong sulphuric acid, or by precipitation 
 with mercuric oxide, or, where available, by means of sulphuretted 
 hydrogen. 
 
 (2) Difficultly decomposed Double Cyanides. These 
 compounds are not decomposed with liberation of hydrocyanic 
 acid by dilute mineral acids. They may be regarded as containing 
 a complex metallo-cyanogen radical, represented by the general 
 formula RCy, ; ; and in the case of those which are of most im- 
 portance in analysis, R stands for Fe or Co, as in the following : 
 
 Potassium ferrocyanide, K 4 (FeCy 6 ) 
 
 Copper ferrocyanide, Cu 2 (FeCy 6 ) 
 
 Potassium ferrous ferrocyanide, K 2 Fe(FeCy 6 ) 
 
 Ferric ferrocyanide (Prussian blue}, Fe 4 (FeCy 6 ) 3 
 
 Potassium ferricyanide, K 3 (FeCy 6 ) 
 
 Ferrous ferricyanide (TurnbulPs blue\ Fe 3 (FeCy 6 ).> 
 
 Potassium cobalticyanide, K 3 (CoCy 6 ) 
 
 The reactions which potassium ferro and ferri cyanides give 
 with ferric and ferrous salts have already been studied in connection 
 with the special tests for iron. The two salts of iron are therefore 
 used for the detection of soluble ferro and ferri cyanides. 
 
 (a) Ferrocyanides. The alkali ferrocyanides alone are easily 
 soluble in water. By double decomposition, therefore, with potas- 
 sium ferrocyanide and a metallic salt, the insoluble ferrocyanides 
 are precipitated ; eg. 
 
 K 4 (FeCy a )+ 2CuCl 2 = 4KC1 + Cu 2 (FeCy 6 ) 
 When treated with cold dilute acids, the ferrocyanides are 
 decomposed, not with evolution of hydrocyanic acid, but with the 
 formation of hydroferrocyanic acid, the complex cyanogen group 
 remaining intact 
 
 K 4 (FeCy 6 ) + 4 HC1 - 4 KC1 + H 4 (FeCy 6 ) 
 
 When boiled with dilute sulphuric acid, the alkali ferro (and 
 i; ferri") cyanides are partially decomposed, with elimination of 
 a portion of the cyanogen as hydrocyanic acid. (This is the usual 
 method for preparing aqueous hydrocyanic acid) 
 
 2K,(FeCy c ) + 3H 2 SO, = K a SO 4 + K 2 Fe(FeCy 6 ) + 6HCy 
 
164 Qualitative Analysis. 
 
 When heated with strong sulphuric acid (in common with all 
 cyanides) the cyanogen radical is completely decomposed, the 
 nitrogen it contains being converted into ammonia 
 
 Insoluble ferrocyanides are decomposed by treatment with 
 caustic alkalies and alkali carbonates ; thus, with Prussian blue 
 (ferric ferrocyanide) 
 
 Fe 4 (FeCy G ) 3 + i2NaHO = 2 Fe 2 (HO) 6 + 3 Na 4 (FeCy 6 ) 
 
 By removing the precipitated metallic hydroxide (in this case 
 ferric hydroxide) by nitration, the presence of the ferrocyanide in 
 the solution can be detected by the addition of ferric chloride. 
 
 Under the influence of oxidising agents (e.g. chlorine or bromine 
 water) potassium ferrocyanide is converted into ferricyanide 
 
 K 4 (FeCy) + Cl - KC1 + K 3 (FeCy fl ) * 
 
 (b} Perricyanides. The alkali salts only are easily soluble 
 in water. Acids and alkalies act upon the ferricyanides as upon 
 ferrocyanides ; e.g. the insoluble compounds, when boiled with 
 alkaline hydroxides, yield the soluble alkali ferricyanide, pre- 
 cipitating the metal as hydroxide ; thus, with ferrous ferricyanide 
 (Turnbull's blue) 
 
 Fe 3 (FeCy 6 ) 2 + 6NaHO = 2Na 3 (FeCy 6 ) + 3Fe(HO) 2 
 
 Soluble ferricyanides are recognised by the formation of a blue 
 precipitate with ferrous sulphate (see Iron). Insoluble ferricyanides 
 (in common with all other double cyanides) are examined for the 
 metals they contain by first decomposing the compound ; either by 
 boiling with alkali hydroxide as shown above, or by boiling with 
 strong sulphuric acid ; or, lastly, by fusion with a mixture of 
 ammonium nitrate and sulphate, whereby the metals are converted 
 into sulphates. 
 
 Cyanic Acid and Cyanates. 
 
 Cyanic Acid, HCyO. This compound cannot exist in con- 
 tact with water, being broken up into carbon dioxide and ammonia ; 
 
 * Although for many reasons it is convenient to regard these compounds 
 as containing the complex radical (FeCy 6 ), this view of their constitution 
 somewhat obscures some of their reactions. The oxidation of potassium 
 ferrocyanide to ferricyanide is a case in point. If these two compounds be 
 represented as double compounds of "ferrous " and " ferric " cyanides respec- 
 tively, then the oxidising action of chlorine in converting the one into the other 
 is at once apparent 
 
 4 KCy,Fe"Cy 2 + Cl = KC1 + 3KCy,Fe"'Cy 3 
 
Thiocyanic Acid and Thiocyanates. 165 
 
 therefore, when cyanates are acted upon by acids, while the vapour 
 of cyanic acid escapes, and is recognised by its powerful and 
 pungent or acrid smell, the greater part of it is decomposed by the 
 water present ; thus 
 
 2K(CN)O + 2H 2 SO 4 + 2H 2 O = K 2 SO 4 + (NH 4 ), SO 4 + 2CO 2 
 
 Cyanates of the alkalies and alkaline earths are soluble in water. 
 Insoluble cyanates are therefore precipitated when salts of most 
 of the heavy metals are added to a solution of potassium cyanate. 
 
 Ammonium cyanate, NH 4 CNO, is isomeric with urea,(NH 2 ) 2 CO, 
 and when an aqueous solution of the former is heated, it passes 
 into the latter, by simple rearrangement of the atoms already in 
 the molecule without any addition or subtraction. 
 
 Cyanates are at once distinguished from cyanides, by the fact 
 that they do not evolve hydrocyanic acid on treatment with dilute 
 acids. They are identified by the evolution of carbon dioxide, 
 with simultaneous formation of ammonium salt. 
 
 Commercial potassium cyanide usually contains potassium 
 cyanate, which is detected by this reaction. 
 
 Thiocyanic Acid and Thiocyanates. 
 
 Thiocyanates are the sulphur analogues of the cyanates. 
 The acid itself, HCyS (like cyanic acid), is unstable in aqueous 
 solution. The chief insoluble thiocyanates are those of silver, lead, 
 mercury, and copper. On the addition of solutions of these metals 
 to a solution of potassium thiocyanate, the various compounds are 
 precipitated. Silver thiocyanate, AgCyS, is obtained as a white 
 precipitate resembling the chloride in appearance. It dissolves in 
 ammonia, but less easily than either the cyanide or chloride, and is 
 insoluble in dilute nitric acid. 
 
 With ferric salts, potassium or ammonium thiocyanates give a 
 red colour, which is not destroyed by hydrochloric acid. The 
 formation of this colour is employed both as a test for " ferric " iron 
 or for thiocyanates. The addition of mercuric chloride destroys 
 the colour. 
 
 When heated, the thiocyanates are decomposed. The potassium 
 salt may be heated to fusion in the absence of air, but in contact 
 with air it is converted into sulphate and cyanate, with evolution of 
 sulphur dioxide. The mercurous and mercuric salts, when heated 
 in contact with the air, take fire and burn, at the same time 
 swelling up in a most remarkable manner,* leaving a straw-coloured 
 voluminous ash, which is many times as bulky as the original 
 compound. At the same time sulphur dioxide, mercury vapour, 
 nitrogen, and cyanogen are evolved. 
 
 * The thiocyanates of mercury are the materials of which the so-called 
 " Pharaoh's serpents " are made. 
 
1 66 Qualitative Analysis. 
 
 Silicon. 
 
 The properties of the element silicon, in the uncombined state, 
 are not made use of in analysis, as in no analytical reactions is this 
 element liberated. 
 
 Silicic Acid and Silicates. 
 
 Silicic acid, H 2 SiO 3 , is obtained, when soluble silicates are 
 decomposed by acids, as a white gelatinous substance, slightly 
 soluble in water, and still a little more soluble in acids 
 
 Na 2 SiO 3 4- 2HC1 = H 2 SiO 3 + 2NaCl 
 
 If the decomposition of the soluble silicate be made to take place 
 in the presence of an excess of acid that is, if the silicate be added 
 to the acid (instead of adding acid to the silicate) the silicic acid 
 which is produced is not precipitated at once, but remains in solu- 
 tion, and may be separated from the sodium chloride and the excess 
 of hydrochloric acid by dialysis. After a short time, however, the 
 silicic acid separates out as a transparent gelatinous mass.* 
 
 Silicic acid is also precipitated from a solution of an alkali 
 silicate by the addition of ammonium carbonate (ammonia does 
 not cause any precipitate). The action may be regarded as taking 
 place in two stages, ammonium silicate being first formed by 
 double decomposition, and immediately breaking up into silicic 
 acid and ammonia ; thus 
 
 Na 2 Si0 3 + (NH 4 ) 2 C0 3 = Na 2 CO 3 + (NH 4 ) 2 SiO 3 
 (NH 4 ) 2 SiO 3 = 2NH 3 + H,SiO 3 
 
 Silicic acid is a very feeble acid, and when heated to 130 C. 
 it parts with a molecule of water and is converted into silicon 
 dioxide (silica), SiO 2 , a compound which is insoluble in water and 
 in acids.t 
 
 Silica, SiO 2 , occurs in a more or less pure state in nature, in the 
 form of quarts, flint, agate, sand, etc. When prepared artificially 
 by heating the hydrated compound, it is a white amorphous powder. 
 It is extremely stable, and capable of standing a very high tempera- 
 ture. It is insoluble in water and all acids except hydrofluoric 
 acid. It dissolves in caustic alkalies ; and, when in the amorphous 
 
 * The acid which is thus obtained in solution is believed to be the less 
 stable tetrabasic silicic acid, H 4 SiO 4 , or Si(HO) 4 , or SiO 2 ,2H 2 O. This com- 
 pound, by giving up one molecule of water, passes into the dibasic acid, H 2 SiO 3) 
 orSi0 2 ,H 2 0. 
 
 f Silicic acid is sometimes spoken of as soluble silica, and silicon dioxide as 
 insoluble silica. 
 
Silicic Acid and Silicates. 167 
 
 state, in boiling carbonates of the alkalies, yielding in all cases the 
 soluble alkali silicates. 
 
 The only silicates which are soluble in water are the alkali 
 silicates (known as water glass, or soluble glass). The natural 
 .silicates form a numerous and complex class of minerals. They 
 may be regarded as compounds of metallic oxides with silicon 
 dioxide ; \h\is, felspar, Al 2 O 3 ,K 2 O,6SiO 2 . 
 
 The silicates, which are insoluble in water, may be classified for 
 analytical purposes into (i) those which are decomposed by acids 
 (other than hydrofluoric acid) ; and (2) those which are unattacked, 
 which comprises by far the larger class. 
 
 The presence of silica in these compounds is detected by means 
 of their behaviour when heated with microcosmic salt. A clear 
 bead of this salt is made upon a loop of platinum wire, and a few 
 small particles of the powdered mineral are heated in it. The 
 metallic oxides are dissolved by the fused microcosmic salt, but not 
 the silicon dioxide, particles of which (the skeletonic remains of the 
 mineral) remain floating about in the molten bead.* 
 
 Silica may also be detected by the formation of silicon fluoride, 
 when a silicate is acted upon by hydrofluoric acid. The test is 
 made as described under fluorine (p. 134).! In this case, however, 
 the drop of water which is suspended in the gas in order to detect 
 the silicon fluoride, should be held on a loop of platinum wire, as the 
 action of the hydrofluoric acid upon the glass rod might be mis- 
 taken for a deposition of silica. When applying this test for the 
 detection of fluorine, this is obviously no disadvantage. 
 
 (i) Silicates which are decomposed by acid may have 
 their silica removed by treatment with hydrochloric acid, just as 
 in the case of the alkali silicates above mentioned. In this instance, 
 however, the mineral, in as finely powdered a condition as possible, 
 is digested with strong hydrochloric acid at a gentle heat, when gela- 
 tinous silicic acid separates, and the powdered mineral gradually 
 dissolves. In order to render the separation of the silica complete, 
 the mixture is next evaporated to dryness (preferably on a steam- 
 bath), and in order to ensure the entire conversion of the silicic acid 
 into silica, it is afterwards gently heated over a flame for a short time : 
 (excessive heating is to be avoided, as certain metallic oxides are 
 thereby rendered very difficult of subsequent solution in acid). The 
 residue is then treated with a small quantity of strong hydrochloric 
 
 * This preliminary test, however, is not absolutely reliable. Some in- 
 soluble silicates are known which do not leave this residue, and there are 
 certain natural phosphates which are not dissolved by fused microcosmic salt. 
 . but remain floating in the bead, and might be mistaken for silica. 
 
 f Except that a platinum crucible must replace the test-tube, 
 
1 68 Qualitative Analysis. 
 
 acid ; water is added, and the solution containing the metals present 
 as chlorides is separated by nitration from the insoluble residue of 
 silica. 
 
 (2) Silicates which, are unattacked by acids are decom- 
 posed by fusion with alkali carbonates. The finely powdered 
 silicate is mixed with several times its weight of fusion mixture, and 
 the mixture heated in a platinum crucible. As the mass melts, 
 effervescence takes place, due to the escape of carbon dioxide ; 
 thus, using a general formula for a simple silicate * 
 
 *RO,ySiO 2 +jyNa 2 CO 3 = ^-RQ + jNa 2 SiO 3 +_yCO 2 
 
 The fusion is continued until effervescence ceases, the temperature 
 being raised towards the end of the operation. The residue, after 
 cooling, is extracted with water, which dissolves the alkali silicate 
 (and excess of carbonate), leaving the metallic oxide (or carbonate). 
 Hydrochloric acid is then gradually added, which causes the pre- 
 cipitation of silicic acid, and at the same time dissolves the metallic 
 oxides. The mixture is then evaporated, and treated in the manner 
 described above. 
 
 The acid radicals of any other compounds present in the silicate 
 (e.g. sulphates, phosphates, borates, arsenates, chlorides) are now 
 present as their sodium or potassium salts, and may be detected 
 by their several special tests in the ordinary way. 
 
 This method of fusion with alkali carbonates is obviously inad- 
 missible in the case of natural silicates which are suspected of 
 containing the alkali metals, and which are insoluble in hydrochloric 
 acid. In this case, one of the following plans may be employed: 
 
 (a) Heating with Barium Oxide. A small quantity of the 
 finely powdered mineral is intimately mixed with three or four 
 times its weight of barium oxide, and strongly heated in a platinum 
 crucible by the blowpipe, the heating being continued for about 
 twenty minutes. The mass is then dissolved in dilute hydrochloric 
 acid. The solution so obtained is made alkaline with ammonia (a 
 precipitate is produced, which need not be removed), and ammo- 
 nium carbonate is added. The mixture is then filtered, and the 
 filtrate evaporated to dryness and heated, and examined for the 
 alkalies in the usual way (p. 25). 
 
 () Heating with Ammonium Chloride and Calcium 
 Carbonate. A small quantity of the powdered silicate is mixed 
 with about an equal weight of ammonium chloride and about eight 
 
 * The possibility of this reaction depends upon the fact that, although 
 silicic acid is an unstable acid, silicon dioxide and the alkaline silicates are 
 extremely stable even at high temperatures. 
 
Boric Acid and B orates. 169 
 
 times its weight of calcium carbonate (precipitated), and the mix- 
 ture gently heated in a platinum crucible until no more fumes of 
 ammonium chloride are given off. It is then strongly heated with 
 the blowpipe for about fifteen minutes, after which the mass is 
 treated with water. The alkalies, in the form of chlorides, together 
 with a small quantity of calcium chloride, pass into solution, and 
 are separated by filtration. The calcium is removed by pre- 
 cipitation with ammonium carbonate, and the filtrate evaporated 
 and examined for alkalies. 
 
 (V) Decomposing the Silicate with Hydrofluoric Acid. 
 This may be accomplished by treating the powdered mineral with 
 aqueous hydrofluoric acid in a platinum crucible, gently evaporating 
 the liquid (in a draught cupboard) to dryness, adding fresh acid 
 and evaporating again, continuing the operation until the residue is 
 entirely soluble in hydrochloric acid. Or the mineral may be 
 moistened with a little strong ammonia in a platinum crucible, and 
 the whole exposed to the action of gaseous hydrofluoric acid 
 (generated from fluorspar and sulphuric acid) in a leaden pot, which 
 can be covered with a leaden lid. The action is allowed to proceed 
 for at least twenty-four hours. The crucible is then removed, and 
 the ammonium fluoride expelled by gently heating in a draught 
 cupboard. 
 
 Boron. 
 
 The properties of the free element are not employed in analysis. 
 
 Boric Acid and Borates. 
 
 Boric acid, H 3 BO 3 , is a white crystalline solid, sparingly 
 soluble in cold, but more readily soluble in hot, water. It is 
 deposited from its solutions in the form of pearly white scales. 
 
 It is also soluble in alcohol, and when either the aqueous or 
 alcoholic solution is boiled, the acid vaporises along with the solvent. 
 
 When heated, boric acid (ortho), H 3 BO 3 , loses water, passing first 
 into metaboric acid, H 2 B 2 O 4 , and finally into pyroboric acid, H 2 B 4 O 7 . 
 
 The borates of the alkalies are readily soluble in water ; most 
 other borates are insoluble. A few (e.g. magnesium borate) are 
 difficultly soluble. The most familiar salts are those of pyroboric 
 acid, e.g. ordinary borax, Na 2 B 4 O 7 . 
 
 The insoluble borates obtained by precipitation are not charac- 
 teristic, and are not used in analysis for the detection of borates. 
 
 All borates are decomposed by mineral acids, with liberation of 
 orthoboric acid ; thus 
 
 Na 2 B 4 O 7 + 2HC1 + 5H 2 O = 2NaCl + 4H 3 BO 3 
 
170 Qualitative Analysis. 
 
 Reaction with. Turmeric. Boric acid produces upon tur- 
 meric paper a characteristic red-brown stain. This coloration is 
 distinguished from that produced by alkalies (which it closely re- 
 sembles in appearance) by the fact that when touched with an alkali 
 the brown colour is changed to a greenish-black, but is restored to 
 its original tint by dilute acids (HC1 or H 2 SO 4 ). The borate is 
 moistened with hydrochloric or sulphuric acid in order to liberate 
 the boric acid, and a drop or two of the liquid is poured upon the 
 turmeric paper. 
 
 Flame Reactions. Volatile boron compounds impart a 
 characteristic green colour to a non-luminous flame ; thus, wh'en 
 an alcoholic solution of boric acid is boiled, and the alcohol vapour 
 inflamed, the green colour due to the volatilised boric acid* is 
 apparent. The test is made in the following manner \~ 
 
 The borate (borax) is moistened with a little strong sulplpric 
 acid in a test-tube or small flask, and alcohol is added. The test- 
 tube is closed with a cork carrying a short straight glass tube. 
 The contents of the tube are then heatejjl, and as the alcohol boils 
 off it is inflamed at the exit tube, when the gre%n colour of the 
 flame is observed. 
 
 Since ethyl chloride (which is liable to be formed if a metallic 
 chloride is treated in the same manner) gives a flame which also 
 has a green edge or fringe (albeit, if this compound were formed, 
 the alcohol flame would at once be rendered luminous, like an 
 ordinary gas flame), the following test may also be applied as a 
 confirmation : 
 
 A small quantity of the powdered borate (or boric acid) is 
 mixed with about its own weight of powdered calcium fluoride 
 and three or four times its weight of hydrogen potassium sulphate. 
 This mixture is moistened with the least drop of water, and a little 
 of the paste is introduced into a Bunsen flame upon a loop of 
 platinum wire. By the action of the acid sulphate upon the 
 fluoride, hydrofluoric acid is formed ; and this in the presence 
 of the borate gives boron fluoride, BF 3 , which causes a green 
 coloration of the flame : 
 
 (1) CaF, + 2HKS0 4 = CaSO 4 + K.,SO 4 + -HF 
 
 (2) B t Oi + 6HF = 3H 2 + 2BF, 
 
 In many cases it is sufficient to mix the borate with the fluoride, 
 and moisten the mixture with a drop of strong sulphuric acid and 
 bring this into the name in the same way. The presence of copper 
 salts masks the reaction. 
 
CHAPTER XV. 
 SYSTEMATIC DETECTION OF THE ACIDS. 
 
 THE acids are classified into three main groups based upon the 
 solubility of their barium and silver salts. 
 
 Group I. Acids whose barium salts are precipitated from 
 neutral solutions by barium Chloride. 
 
 (a) Whose barium salts are insoluble in dilute hydrochloric acid 
 
 Sulphuric acid, H 2 SO 4 Hydrofluosilic acid, H 2 SiF * ' 
 
 
 
 (b} Whose barium salts are soluble in hydrochloric acid 
 
 Carbonic acid, H 2 CO 8 Boric acid, H 3 BO S 
 
 Sulphurous,, H 2 SO 3 lodic HIO 3 
 
 Thiosulphuric acid, H 2 S 2 O 3 IJydrofluoric aq^JiF 
 
 Silicic HgSiO, Oxalic ., H 2 (C 2 O 4 )t 
 
 Chromic H 2 CrO 4 Tartaric ., H,(C 4 H 4 O c )t 
 
 Phosphoric H 3 PO 4 Citric H 3 (C 6 H 5 O 7 )t 
 
 Group II. Acids whose silver salts are precipitated by silver 
 nitrate from solutions acidified with nitric acid 
 
 Hydrochloric acid, HCl Sulphuretted hydrogen, H 2 S 
 
 Hydrobromic HBr Thiocyanic acid, HCyS 
 
 Hydriodic HI Ferrocyanic H 4 FeCy,, 
 
 Hydrocyanic HCy Ferricyanic H 8 FeCy 
 
 Group III. Acids whose barium salts are soluble in water, 
 and ^vhose silver salts are not precipitated, in a nitric acid solution, 
 namely 
 
 Nitric acid, HNO ;i 
 
 Nitrous HNO 2 Cyanic acid, HCyO 
 
 * Soluble in strong hydrochloric acid. 
 
 f In the case of these three acids, the neutral solution must not contain 
 appreciable quantities of ammpniacal salts, or soluble double compounds are 
 formed on addition of barium chloride, which may entirely prevent the forma- 
 tion of a precipitate. 
 
i/2 Qualitative Analysis. 
 
 Chloric acid, HC1O 3 Formic acid, H(CHO 2 ) 
 
 Perchloric acid, HC1O 4 Acetic H(C 2 H 3 O 2 ) 
 
 As already mentioned (p. 120), these general reagents, barium 
 chloride and silver nitrate, are not employed to effect the separa- 
 tion of groups of acids, but rather as indicators of the presence or 
 absence of entire groups. 
 
 In most cases, each acid is individually tested for in separate 
 portions of specially prepared solutions (see below). 
 
 The examination for acids should be made after the metals 
 have been detected, for two reasons : firstly, because during the 
 course of the^ystematic examination for the metals, the presence 
 or absence of quite a number of acids will be incidentally ascertained ; 
 and secondly, because a knowledge of what metals are present 
 will be a guide to the student in deciding what acids may and what 
 cannot possibly be present. For example, if the material under 
 examination is a solution having an add reaction, and it is found 
 to contain silver as one of the metals, it will obviously be useless 
 to expect to find any of the acids of Group II. Or if the substance 
 is a solid which dissolves easily in water, and calcium is found, it 
 is equally unnecessary to apply tests for sulphuric, phosphoric, 
 carbonic, silicic, hydrofluoric or other acids whose calcium salts 
 are not easily^|pluble in water. Or if a solid, soluble in water, 
 were found to contain the metal lead, the acids most probably 
 present would be nitric or acetic, as the lead salts of these two 
 acids are the only common soluble lead compounds. A student 
 who, in spite of this, goes blindly through a series of tests for acids 
 which cannot be present, shows at once that he has no intelligent 
 appreciation of the work he is doing. 
 
 It will be evident that, in order to make the fullest use of the 
 information he. has gained concerning the metals that are present, 
 the student must have a knowledge of the solubilities of a large 
 number of compounds. The solubility of the common compounds 
 of the metals and acids has been mentioned under the separate 
 elements treated in the foregoing chapters ; but in order to make 
 the information more accessible for reference, a condensed epitome 
 of solubilities of the common compounds of the various metals and 
 acids which have been treated, is given at the end of this chapter 
 
 (P- 177). 
 
 As stated above, several acids will be detected during the 
 examination for metals ; thus on acidifying with hydrochloric 
 arid and gently warming (for separation of Group I.), the presence 
 of carbonates, sulphites, sulphides, cyanides, or nitrites will be 
 
Systematic Detection of the Acids. 173 
 
 indicated by the evolution of carbon dioxide, sulphur dioxide, 
 sulphuretted hydrogen, hydrocyanic acid, or nitric oxide (appearing 
 as brown fumes of the peroxide) respectively.* 
 
 The evolution of sulphur dioxide with simultaneous deposition 
 of sulphur indicates thiosutphates. The escape of chlorine may 
 denote hypochlorites, or may arise from the action of the hydro- 
 chloric acid (especially if not very dilute) upon chlorates, iodates, 
 chromates, nitrates, or (in the case of solid substances) of peroxides. 
 
 When the substance under examination is a solid, these indi- 
 cations obtained by the action of hydrochloric acid should be 
 followed up by a second general test, namely 
 
 Treatment with Strong Sulphuric Acid. On gently warm- 
 ing a little of the solid with a small quantity of concentrated 
 sulphuric acid in a test-tube, a number of important indications 
 may be obtained : thus, Cyanides, chlorides, fluorides, and nitrates 
 evolve their respective acids, in the latter case accompanied by 
 brown fumes. Iodides give violet vapours ; bromides give the 
 dark-brown vapour of bromine, condensing to dark drops on the 
 tube ; chlorates yield the deep yellow explosive gas, chlorine 
 peroxide. Formates evolve carbon monoxide, while oxalates give 
 carbon monoxide and dioxide (the residue in neither case is black- 
 ened). Tartrates and citrates evolve sulphur dioxide in addition 
 to the oxides of carbon, at the same time leaving a black residue. 
 In many cases these indications should be followed up at once by 
 special confirmatory tests, which it may be practicable to apply 
 directly to the substance under investigation; e.g. an indication 
 which leads to the suspicion of nitrates being present, should be 
 confirmed by applying the test with copper, or the ferrous sulphate 
 reaction. If, therefore, this test with sulphuric acid should give 
 a purely negative result, it is obvious that a large number of acids 
 are excluded. The chief compounds which give no indications 
 when subjected to this test are sulphates, phosphates, phosphites, 
 arsenates, silicates, and metallic oxides other than peroxides, or 
 such as behave as peroxides ; these latter evolve oxygen. 
 
 On treatment with sulphuretted hydrogen, in the separation 
 of the metals of Group II., indications will have been obtained 
 of the presence of chromates and iodates : the former by the 
 change of colour from orange to green, with simultaneous pre- 
 cipitation of sulphur ; the latter by the elimination of iodine, which 
 
 * If these indications of the decomposition of salts of these acids were not 
 carefully noted during the examination for metals, the reaction should be 
 repeated at this stage of the analysis. 
 
Qualitative Analysis. 
 
 gives a dark brown colour to the solution, the colour gradually 
 disappearing as the excess of sulphuretted hydrogen converts the 
 iodine into hydriodic acid. 
 
 On the preparation of the solution for separating the metals of 
 Group III., the presence or absence of phosphates and silicates 
 will have been ascertained, and also of such organic acids (citric 
 or tartaric) as leave a charred residue when heated. 
 
 Besides these, indications of the presence of several acids will 
 have been obtained during the performance of the general dry 
 tests for the metals (p. 179). 
 
 Preparation of the Solution for the Detection of Acids. 
 
 Before testing for the acids, it is advisable (in many cases it is 
 necessary} that they should be present in the solution as salts of 
 the alkalies (or alkaline earths) ; that is to say, the metals with 
 which the various acids are united in the substance under analysis 
 should be exchanged, by double decomposition, for an alkali metal ; 
 for the reason that the presence of these other metals would in many 
 cases mask the reactions by which the acids are to be detected.* 
 To accomplish this, the solution is boiled,t and sodium carbonate 
 added in quantity slightly in excess of that required to effect 
 complete precipitation. The mixture is then filtered, and the 
 acids, now present as their sodium salts, are detected in the 
 solution.J 
 
 * Whether or not this would be the case will usually be revealed by the 
 result of the examination for metals. For example, suppose the metals found 
 in a solution were magnesium and copper, then the only acids likely to be 
 present would be hydrochloric, sulphuric, and nitric (these acids forming the 
 common soluble salts of the two metals), and the reactions by which these 
 acids may be detected can be made equally well whatever metal they may be 
 united with. Or again, suppose, in addition to these two metals, silver was 
 found ; this would reduce the acids probably present to nitric and sulphuric, 
 and the test for the latter would then merely require the substitution of barium 
 nitrate for barium chloride. The student should therefore use an intelligent 
 judgment at this point, and not proceed by a blind habit to prepare the 
 solution by the removal of the metals with sodium carbonate, whether such a 
 step be necessary or not. The importance of carefully and systematically 
 taking notes of his work as it is in progress cannot be too strongly urged. 
 Such notes should be recorded as important memoranda for his own guidance 
 and instruction, and not as a mere voucher for his teacher that he has carried 
 out a certain prescribed process, or made a certain stereotyped set of tests. 
 
 f In cases where the substance under analysis is insoluble, the product 
 obtained by fusion with sodium carbonate is extracted with water, and the 
 solution so obtained is employed for the detection of the acids (see p. 184). 
 
 J This method for separating the metals from their acids is not, however, 
 of universal application. Thus, in the case of many of the phosphates held 
 in solution by acids, the action of the sodium carbonate is to cause the pre- 
 cipitation of the phosphates themselves ; in such instances, therefore, the acid 
 is not found in the solution. Since, however, phosphoric acid will have been 
 
Systematic Detection of the Acids. 175 
 
 GENERAL TESTS. (i) A small portion of this solution is care- 
 fully neutralised by first adding dilute nitric acid drop by drop 
 (the mixture being heated to expel carbon dioxide), until the liquid 
 is just add, and then adding a drop or two of dilute ammonia. 
 
 This neutral solution is then tested by the addition of barium 
 chloride. If no precipitate is obtained, the acids of Group I. 
 (p. 171) are absent. If a precipitate is formed which redissolves on 
 the addition of hydrochloric acid, the two acids sulphuric and 
 hydrofluosilicic are excluded. 
 
 (2) A second small portion of the solution is acidified with 
 nitric acid, and tested with silver nitrate. A negative result proves 
 the absence of the acids of Group II. 
 
 If these general tests show that acids of both Groups I. and II. 
 are present, special tests must then be applied in separate portions 
 of the solution, for such acids as, from information already gained, 
 are considered likely to be in the substance under analysis, and 
 which have not been definitely discovered during the course of the 
 examination. The tests may be made in accordance with the 
 following outline scheme, in which most of those acids which must 
 certainly have been detected in the earlier stages are not again 
 mentioned. 
 
 A. In portions of the solution acidulated with hydrochloric acid. 
 
 Sulphuric Acid. Barium chloride precipitates white barium 
 sulphate : not dissolved by the addition of strong hydrochloric 
 acid, and boiling the liquid. 
 
 Hydrofluosilic Acid. Barium chloride gives white barium 
 silicofluoride : soluble in strong hydrochloric acid. 
 
 Ammonia precipitates gelatinous silicic acid. 
 
 Silicic Acid. Ammonium carbonate (but riot ammonia) pre- 
 cipitates silicic acid. 
 
 " Ferro " and " ferri "-cyanic acids * give their respective 
 reactions with iron salts. 
 
 Thiocyanic acid * gives the red coloration with ferric salts. 
 
 discovered during the examination for the metals, this fact is of little 
 consequence. 
 
 On the other band, sodium carbonate fails to separate the metals from 
 cyanides soluble in water (mercuric cyanide) or In potassium cyanide (i.e. 
 soluble double cyanides, such as those of silver, copper, zinc, etc.), and also 
 from double tartrates and citrates, etc. In such cases as these, the metals 
 may be separated by means of sulphuretted hydrogen or ammonium sulphide 
 (see Reactions for cyanides, p. 161). 
 
 It will be obvious that carbonates must be detected at an earlier stage in the 
 analysis. 
 
 * If the general test, No. 2 (see above), gave a negative result, these acids 
 will have been proved to be absent. 
 
176 Qualitative Analysis. 
 
 Arsenic Acid. If arsenic has been found among the metals 
 and its state of oxidation (i.e. whether present as an arsenite or 
 arsenate) has not been determined, the following test may be 
 applied : Ammonium chloride, ammonia, and magnesium sulphate 
 are added a white precipitate may be due to either a phosphate 
 or arsenate. If phosphoric acid has been proved to be absent (by 
 the molybdate reaction), it must be the arsenate. In either case 
 it should be filtered, and after being washed free from ammonium 
 chloride, it is dissolved in a little nitric acid, and silver nitrate 
 added (or a drop or two of silver nitrate may be poured upon the 
 washed precipitate in the funnel). A brown precipitate indicates 
 an arsenate. 
 
 B. In portions acidulated with nitric acid. 
 Hydrochloric, Hydrocyanic, and Thiocyanic Acids. 
 
 Silver nitrate gives a white precipitate. 
 
 Hydrobromic and Hydriodic Acids. Silver nitrate gives 
 a yellowish precipitate. 
 
 (For the methods of discriminating between these, see pp. 127 
 and 128.) 
 
 C. In portions acidulated with acetic acid. 
 
 Oxalic Acid. Calcium sulphate precipitates calcium oxalate : 
 readily soluble in hydrochloric acid. 
 
 Hydrofluoric Acid. Calcium sulphate or chloride precipi- 
 tates calcium fluoride : only slightly soluble in hydrochloric acid. 
 
 (Chromic, phosphoric, and arsenic acids may also be looked 
 for in portions of this solution. These acids, however, will have 
 been detected at an earlier stage of the analysis.) 
 
 D. In portions rendered neutral. 
 
 Tartaric and Citric Acids. Calcium chloride gives a white 
 precipitate of calcium tartrate in the cold, and of calcium citrate 
 on boiling (see Special reactions, p. 157). 
 
 Formic and Acetic Acids. Ferric chloride produces a red- 
 coloured solution, browner in colour than that given by thio- 
 cyanates. The colour is destroyed by hydrochloric acid (see 
 Special reactions). 
 
Table of Solubility. 
 
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CHAPTER XVI. 
 
 PRELIMINARY TESTS AND OPERATIONS IN A SYSTEMATIC 
 ANALYSIS. 
 
 A. When the Substance under Examination is a Liquid. 
 
 Before proceeding to the separation of the groups, the liquid 
 should be carefully tested with litmus paper, in order to ascertain 
 whether it is neutral, acid, or alkaline. 
 
 (1) If neutral^ the liquid might be simply water. To ascertain 
 whether or not it contains anything in solution, a little of it should 
 be carefully evaporated to dryness. This should be done upon a 
 watch-glass, which is heated upon a steam-bath. (The glass may 
 be placed over the mouth of a small beaker, in which a little water 
 is being gently boiled.) If there is no residue left in the watch-glass 
 when the liquid has entirely evaporated, it contained no salts in 
 solution, and was therefore simply water. 
 
 [Unless evaporated gently as described, certain compounds 
 which could be present might volatilise entirely, and in this way be 
 altogether overlooked ; e.g. ammonium nitrate would go off as 
 steam and nitrous oxide, if its aqueous solution were merely boiled 
 to dryness.] 
 
 (2) If acid, the solution may contain either a free acid, or salts 
 possessing an acid reaction. [Either certain normal salts, e.g. copper 
 sulphate, alum, etc., or certain " acid " salts, as hydrogen sodium 
 sulphate. It must be remembered that many of the so-called 
 " acid " salts, that is, salts in which the whole of the replaceable 
 hydrogen has not been exchanged for metal, are not acid in the 
 sense of exhibiting an acid reaction towards litmus ; some are 
 neutral, while others are alkaline ; e.g. hydrogen sodium carbonate, 
 hydrogen disodium phosphate, etc.]. 
 
 (3) If alkaline, the liquid may contain either free alkali, or salts 
 having an alkaline reaction. 
 
 B. When the substance under analysis is a solid, it 
 should be critically examined, in order, if possible, to gain any 
 
Preliminary Tests in a Systematic Analysis. 179 
 
 information from its general physical properties which may help to 
 identify it. If crystalline, the colour, shape, etc., of the crystals 
 should be noted. If powdered, it may be examined with a pocket- 
 lens, in order to discover whether or not it is homogeneous ; i.e. 
 whether it is a single compound, or a mixture of more than one. 
 
 The substance should then be subjected to the following general 
 tests * : 
 
 I. The Flame Reaction. A small quantity of the substance 
 is introduced into the edge of the Bunsen flame (first in the cooler 
 region near the base of the flame, and afterwards in the hotter 
 parts near the top of the interior cone) upon a loop of clean platinum 
 wire. By a close observation of the colour imparted to the flame, 
 compounds of sodium, potassium, lithium, barium, strontium 
 calcium, copper, and boron may be indicated (see Reactions of these 
 metals). Borates and certain thio-cynates (e.g. of mercury) behave 
 in characteristic fashion when heated. 
 
 II. The Borax Bead. By heating a minute particle of the 
 substance in a borax bead in the blowpipe flame, both inner and 
 outer flame, indications may be obtained of the presence of com- 
 pounds of cobalt, nickel, manganese, iron, chromium, copper. 
 [Should this test lead to the suspicion that either manganese or 
 chromium is present, it should be followed up by the fusion of a 
 portion of the substance with sodium carbonate and nitre upon 
 platinum foil. See Manganese and chromium reactions.] 
 
 III. The Blowpipe Reaction upon Charcoal, (a) When 
 heated alone. A considerable amount of information may be 
 obtained by heating a little of the substance by itself upon charcoal. 
 
 (1) If it melts and is absorbed into the charcoal, it points to the 
 substance consisting of salts of the alkalies. [Chlorates and nitrates 
 cause vivid combustion of the charcoal.] 
 
 (2) If a white infusible residue is obtained, the substance may 
 consist of oxides (or salts which yield oxides when heated) of the 
 alkaline earth metals, alumina, zinc (ZnO yellow while hot, white 
 when cold), or of silica. [If the residue, when placed upon a piece 
 
 * Every smallest detail of these preliminary tests should be carefully and 
 systematically noted down, whether the interpretation of the observed phe- 
 nomena is obvious to the student or not. Indeed, it is often just those very 
 points which at the time do not convey any definite information to him which 
 may prove to be most important. 
 
 The examples which are here given of the changes that may take place when 
 substances are submitted to the various general tests, must be regarded merely 
 as examples. It would be obviously impossible to describe and tabulate every 
 result which might be obtained by subjecting the infinite variety of possible 
 mixtures to the tests described, and even if it were possible, it would be quite 
 undesirable to do so. 
 
180 Qualitative Analysis. 
 
 of turmeric or litmus paper and moistened with water, shows an 
 alkaline reaction, it will contain one of the alkaline earth metals. 
 Silica may be specially tested for by the bead of microcosmic salt.] 
 
 (3) If a coloured residue is left, it points to one of the metals 
 already indicated by the borax-bead test. 
 
 (4) If reductio7i takes place, resulting in the formation of fumes 
 and an incrustation upon the charcoal, without indication o*f a 
 metallic bead, it points to the presence of compounds of very volatile 
 metals, as arsenic (white incrustation accompanied by garlic odour), 
 cadmium (red-brown incrustation), zinc (incrustation yellow when 
 hot, and forming very close to the substance). Such a volatile 
 compound as ammonium chloride gives white fumes and an in- 
 crustation (the latter being formed at a considerable distance from 
 the heated spot). In this case the blowpipe flame appears of a 
 yellow-ochre colour as it impinges upon the ammonium salt. 
 
 () When heated with Reducing Agents. When the substance 
 is mixed with sodium carbonate and potassium cyanide, and heated 
 upon charcoal in the reducing flame, compounds of copper, silver, 
 and gold are reduced to the metallic state without giving any in- 
 crustation upon the charcoal ; while compounds of antimony, 
 bismuth, tin, and lead give metallic beads accompanied by 
 incrustations (for the characteristics of the incrustations and the 
 metals, see Special reactions of each). 
 
 IV. The Action of Heat. The behaviour of a substance 
 when heated alone in a dry narrow tube closed at one end (a small 
 test-tube), will generally afford important information respecting its 
 composition. 
 
 (a) IF THE SUBSTANCE SIMPLY MELTS, AND SOLIDIFIES AGAIN 
 
 ON COOLING, without giving off any gases or vapours, a large number 
 of compounds are obviously at once excluded. In this case the 
 substance may consist of salts of the alkalies or the alkaline earth 
 metals ; a few white salts of the heavy metals, e.g. silver chloride ; 
 or a few coloured salts, e.g. lead chromate. 
 
 () IF WATER IS GIVEN OFF, collecting in drops upon the upper 
 part of the tube, it may be due (i) to hygroscopic* moisture (in this 
 case the amount will probably be small), (2) to the decomposition 
 of metallic hydroxides, or (3) to water of crystallisation. [Sub- 
 stances containing much water of crystallisation, when heated, often 
 melt twice first in the water of crystallisation, and as this is 
 
 * That is moisture due to the substance being "damp" mechanically 
 adhering moisture. Almost all powders attract a little moisture in this way 
 from the atmosphere. 
 
Preliminary Tests in a Systematic Analysis. 181 
 
 expelled they resolidify, but on the application of a stronger heat 
 they once more undergo fusion.] 
 
 The water should be carefully tested by introducing a small 
 strip of litmus paper. 
 
 Alkalinity would indicate ammonium compounds (e-g- 
 (NH 4 ) 2 MgPO 4 , HNa(NH 4 'P0 4 , etc.). 
 
 'Acidity might be due to the decomposition of certain acid 
 salts, either alone or by interaction with other salts (e.g. hydrogen 
 potassium sulphate, when heated, is converted into the normal 
 salt and sulphuric acid, and if present along with a salt containing 
 water of crystallisation, the water would become acid ; or if mixed 
 with sodium chloride or nitrate, hydrochloric acid or nitric acid 
 would be evolved by double decomposition). 
 
 (<f) IF THE SUBSTANCE CHANGES COLOUR 
 
 (1) From white to yellow, indicates oxide of zinc, tin, bismuth ; 
 
 (2) From yellow to brown (fusing at a red heat), oxide of lead. 
 Most coloured oxides become much darker when heated (e.g- 
 
 HgO, Fe. 2 3 )- 
 
 (d) IF GASES OR VAPOURS ARE EVOLVED, they must be iden- 
 tified by special tests. 
 
 (1) Oxygen and nitrous oxide (re-ignition of glowing splint of 
 wood). The former from chlorates, iodates, bromates, nitrates, or 
 peroxides ; the latter from ammonium nitrate, or salts which by 
 double decomposition give ammonium nitrate. 
 
 (2) Hydrogen (character of flame), from formates, and from 
 oxalates in presence of caustic alkalies or alkaline earths 
 (K 2 C 2 O 4 + 2KHO = 2K 2 C0 3 + H 2 ). 
 
 (3) Nitrogen (extinguishes flame, and gives no reaction with 
 lime-water), from ammonium nitrite or chromate, or from mixtures 
 which yield these salts by double decomposition. 
 
 (4) Chlorine, bromine, and iodine (recognised by colour, smell, 
 etc.), evolved from certain halogen compounds when heated alone, 
 or in admixture with acid salts and peroxides (e.g. a mixture con- 
 taining such compounds as NaCl, HKSO 4 , and MnO.,, when heated, 
 evolves chlorine). 
 
 (5) Sulphur vapour (sublimate of yellow drops), from metallic 
 persulphides. 
 
 (6) Carbon dioxide (action on lime-water), from most carbonates 
 pther than the normal salts of the alkalies (also, mixed with carbon 
 monoxide, from certain formates and oxalates). 
 
 (7) Sulphur dioxide (characteristic odour), from certain sul- 
 phites (e.g. CaSO 3 = CaO + SO.,), certain sulphates (e.g. 2FeSO 4 
 
1 82 Qualitative Analysis. 
 
 = Fe 2 O 3 + SO 3 + SO 2 ; CuSO 4 = CuO + S0 2 + O). It may also 
 be due to the oxidation of sulphur evolved from sulphides or other 
 sulphur compounds, as thiocyanates. 
 
 (8) Cyanogen (formation of thiocyanate of ammonia), from 
 certain metallic cyanides. 
 
 (9) Ammonia (odour, and action on test-paper), from many 
 ammonium salts ; e.g. sulphate, phosphate, carbonate, etc. (also 
 from the decomposition of cyanates). 
 
 (10) Sulphuretted hydrogen (odour, and action on paper 
 moistened with lead acetate), from the decomposition of hydrated 
 metallic sulphides, of sulphites and thiosulphates, in the presence 
 of compounds which yield water or acid. 
 
 (11) Phosphoretted hydrogen (odour, and character of flame), 
 from hypophosphites. 
 
 (12) Nitrogen peroxide and oxygen, from the decomposition 
 of nitrates of heavy metals, or from mixtures which by double 
 decomposition yield such a compound (e.g. PbSO 4 and KNO 3 \ 
 
 (e) IF A SUBLIMATE is PRODUCED, it indicates such volatile 
 solids as the following : 
 
 (1) Giving a white sublimate ammonium haloid salts, arsenious 
 oxide, mercuric chloride, mercurous chloride (yellowish while hot, 
 white on cooling). 
 
 (2) Giving a coloured sublimate mercuric iodide (red and 
 yellow), arsenious sulphide (yellow). 
 
 (3) Giving a black metallic-looking sublimate iodine (violet 
 vapours), mercuric sulphide (red streaks if rubbed with a glass rod), 
 metallic mercury (runs into liquid globules when rubbed with a 
 glass rod), metallic arsenic (garlic smell). [If this test leaves it 
 doubtful whether arsenical or mercurial compounds are present, 
 a little of the substance should be mixed with sodium carbonate 
 and heated in a dry tube. If mercurial, the sublimate consists of 
 minute globules.] 
 
 Ammoniacal compounds may be at once tested for by heating 
 a portion of the substance with a little sodium hydroxide, when 
 ammonia is evolved. 
 
 (/) IF NO CHANGE TAKES PLACE, it will be evident from the 
 foregoing that the number of substances which can possibly be 
 present is extremely limited ; and if in addition the compound is 
 white, the range of probable substances is still further narrowed 
 down to the oxides of a few of the metals, such as the alkaline 
 earths, alumina, silica, etc. 
 
 To obtain a Solution of the Solid Substance, (i) In 
 
Solution of a Solid Substance. 183 
 
 Water. A small quantity of the substance, in a finely powdered 
 condition, is first treated with cold water in a large test-tube. 
 [This may result in the evolution of gas, owing to the chemical 
 action of water upon such compounds as certain metallic phosphides 
 (e.g. calcium phosphide) or carbides (calcium carbide), which 
 evolve phosphoretted hydrogen and acetylene respectively ; or 
 carbon dioxide might be expelled, owing to the presence of soluble 
 carbonates and acid salts, which would interact upon each other.] 
 The mixture is then boiled for a few moments. 
 
 If the substance does not dissolve, it is either wholly or partly 
 insoluble in water. To ascertain whether any of it has dissolved, 
 the mixture should be allowed to settle, and a few drops of the 
 clear liquid evaporated to dryness upon a watch-glass ; or upon a 
 piece of platinum foil, cautiously heated. If a residue is obtained 
 on evaporation (thus showing that partial solution in water has 
 taken place), the aqueous liquid is decanted off (or filtered if the 
 insoluble part does not readily settle), and the insoluble portion 
 again boiled with water and filtered. 
 
 (2) In Acids. The portion insoluble in water is then treated 
 with dilute hydrochloric acid and boiled. If it does not entirely 
 dissolve, the dilute acid is decanted off into another test-tube, and 
 strong hydrochloric acid substituted,* the mixture being submitted 
 to prolonged boiling, if necessary. If the substance wholly dissolves, 
 the two acid solutions may be mixed together. 
 
 If, on adding a few drops of the aqueous extract to a small 
 portion of the acid solution, no precipitation takes place, the main 
 portions of the two solutions may be mixed together f and examined 
 for metals and acids (except hydrochloric acid) by the methods 
 already described. On the other hand, if the two extracts contain 
 compounds which will interact with the formation of an insoluble 
 precipitate, they must be examined separately. 
 
 If the substance is not entirely dissolved by strong hydrochloric 
 acid, it may consist of one of the few compounds which are only 
 dissolved by aqua regia.\ The residue is therefore treated with a 
 small quantity of mixed nitric and hydrochloric acids, and boiled. 
 The solution should be evaporated down in a porcelain dish until 
 
 * By the behaviour of the substance under this treatment, the presence of 
 carbonates, sulphites, and peroxides will be indicated. Sulphur and silicic 
 acid may be precipitated. 
 
 f See remarks on p. 187 with reference to this point. 
 
 J Such as sulphides of nickel, cobalt, mercury. The presence of these should 
 have been indicated by the preliminary tests. The use of aqua regia should 
 only be resorted to when absolutely necessary. 
 
184 Qualitative Analysis. 
 
 only a small bulk remains, in order to expel as much acid as 
 possible, and then diluted with water ; this solution may then be 
 mixed with the hydrochloric acid extract. 
 
 (3) Treatment of the Residue insoluble in Water or 
 in Acids. The number of compounds insoluble in water and 
 acids is very limited. Of commonly occurring substances, the 
 following may be present 
 
 (a) Sulphates of barium, calcium, strontium, and lead (the three 
 last named being sufficiently soluble to be detected in the aqueous 
 extract). 
 
 (b) Silica, and many natural silicates. 
 
 (c} Fluor spar, and other natural fluorides. 
 
 (if) Silver chloride (such insoluble silver compounds as the 
 bromide, iodide, cyanide, and ferrocyanide are converted into the 
 chloride by boiling with aqua regid). 
 
 (e) Oxides of aluminium and chromium (which have been 
 strongly heated). Native stannic oxide (tinstone.} 
 
 (y) A few arsenates and phosphates. Carbon and sulphur. 
 
 These insoluble substances (except carbon and sulphur, which 
 are readily recognised and require no further treatment) are con- 
 verted into soluble compounds by fusion with alkali carbonates.* 
 The insoluble residue (in the absence of lead and silver) is dried, 
 and mixed with about four times its weight of fusion mixture and a 
 little potassium nitrate, and fused in a platinum crucible until all 
 effervescence is at an end. The crucible is then allowed to cool. 
 [If the crucible be quickly chilled by standing it upon a piece of 
 cold iron, e.g. the foot of a retort stand, or by dipping it on to the 
 surface of cold water in a basin, the solidified contents of the crucible 
 at once crack away from the metal, and can be removed as a solid 
 cake, which can be broken up in a mortar.] 
 
 The fused mass is then boiled with water until nothing further 
 dissolves, and filtered. 
 
 The aqueous solution is examined for such acids as in the 
 
 * If lead sulphate or silver chloride be present, they must either be re- 
 moved, or the fusion must be conducted in a crucible of porcelain instead of 
 platinum (in which case a certain amount of alumina from the crucible will 
 pass into combination with the alkali carbonate). Lead sulphate may be re- 
 moved by boiling the insoluble residue in a strong solution of ammonium 
 acetate (the solution being tested for lead and for sulphuric acid). Silver 
 chloride may be dissolved out by a solution of potassium cyanide ; or it can be 
 reduced to metallic silver by means of zinc (see Silver, reaction, p. 114), and 
 after washing free from zinc chloride, the reduced silver can be dissolved in 
 nitric acid. If sulphur be present in the insoluble residue, it may be expelled 
 by heating the substance in a porcelain crucible. 
 
Treatment of Insoluble Substances. 185 
 
 nature of the case could be present namely, sulphuric, silicic, 
 hydrofluoric, chromic, arsenic, phosphoric, hydrochloric.* 
 
 The residue, after being thoroughly washed with hot water, 
 until the wash-water is no longer alkaline, is dissolved in hydro- 
 chloric acid, and the solution examined for such metals as can be 
 present. 
 
 In a few special cases, fusion with hydrogen potassium sulphate 
 is resorted to instead of the fusion mixture above mentioned. This 
 method is applicable, for example, in the case of minerals contain- 
 ing titanium, such as titaniferous iron ores (see pp. 73, 343). 
 
 * The fusion-mixture must obviously be free fro n\ sulphates, chlorides, etc., 
 or tests for these acids are of no value. 
 
CHAPTER XVII. 
 THE RESULTS OF A QUALITATIVE ANALYSIS. 
 
 AT the end of an analysis, the student will be in possession of the 
 knowledge that the substance he has examined contained com- 
 pounds of certain metals or positive radicals, and certain negative 
 or acid radicals. If only one metal and one acid have been found, 
 the substance must, of course, be the compound of these two ; e.g. if 
 magnesium and sulphuric acid have been detected, the compound 
 must have been magnesium sulphate. If a metal has been found, 
 but no acid radicals, the substance will in all probability be an 
 oxide of that metal. In deciding this point, however, the student 
 must consider the general properties of the substance, for if he 
 finds, for example, the metal magnesium in a substance soluble in 
 water, the compound obviously cannot be the oxide. He must 
 therefore conclude that he has overlooked the acid radical. 
 
 When more than two metals and acid radicals are found in the 
 substance, it becomes a much more difficult matter to discover 
 which acid is in combination with which metal. Indeed, in many 
 cases it is not possible to decide this point by a qualitative analysis 
 alone. For example, suppose a mixture is found on analysis to 
 contain the metals cadmium and sodium, and the acid radicals of 
 sulphuric and nitric acids. Does the mixture consist of cadmium 
 nitrate and sodium sulphate, or of cadmium sulphate and sodium 
 nitrate ? All the four salts are white, and none of them is so 
 markedly different from the others in its solubility in water as to 
 throw any light on the question ; so that, if the two salts are present 
 in fairly equal quantities, it is impossible to decide which pair of 
 salts is present. 
 
 There are a number of points, however, which a student who 
 cultivates his powers of observation will be able to note, which 
 will often give a clue, if they do not afford proof, as to the manner 
 in which the metals and acid radicals are united in a mixture. 
 
 For instance, even in such a mixture as the above, by carefully 
 
The Results of a Qualitative Analysis. 187 
 
 noting the relative quantities of various precipitates obtained during 
 the course of analysis, some light might be thrown on this point. 
 Suppose it was noted that the amount of cadmium sulphide thrown 
 down seemed very small, and that the test for sulphuric acid re- 
 sulted in a most copious precipitation of barium sulphate, it would 
 be fair to conclude that the cadmium was not present as sulphate. 
 
 Again, suppose barium and ammonium are the " metals " found 
 along with the acid radicals of hydrochloric and nitric acids. The 
 four possible salts are in this case also all of them white and easily 
 soluble. If the salts are in fairly equal quantities, is it possible to 
 decide how the acids and metals are distributed ? If the dry re- 
 actions had been carefully observed, they would enable the student 
 to decide the point. A mixture consisting of BaCl 2 and NH 4 NO 3 
 when heated would melt, and give off water and nitrous oxide 
 (decomposition of the ammonium nitrate), whereas if BaNO 3 and 
 XH 4 C1 were present, some of the ammonium chloride would at 
 once sublime up the tube ; and later on, brown fumes would appear, 
 due to the decomposition of barium nitrate. 
 
 Very often two colourless and soluble salts contained in a 
 mixture, undergo double decomposition with formation of a pre- 
 cipitate, as soon as water is added to the mixture with a view to its 
 solution. When the precipitate is coloured (as would be the case, 
 for example, if the mixture consisted of silver nitrate and sodium 
 phosphate), the fact that it has been formed during the process of 
 solution, and was not originally present, will be evident. On the 
 other hand, when the product of interaction is also white, it is not 
 always possible to decide this point. If such a mixture as silver 
 nitrate and barium chloride be treated with water, the two soluble 
 salts interact, with the precipitation of silver chloride. In this case 
 a close inspection of the appearance of the mixture (the milkiness 
 imparted to the water) would indicate that the substances were 
 interacting. On the other hand, with a mixture such as barium 
 chloride and magnesium sulphate, it would be almost impossible to 
 say whether the original mixture contained barium sulphate or not. 
 
 When the substance under analysis is partly soluble in water 
 and partly soluble in acids, it will be obvious that much informa- 
 tion as to the distribution of the metals and acids will be gained, if 
 each solution is separately examined. Thus, in a mixture consist- 
 ing of lead carbonate and barium nitrate, water will remove the 
 latter salt, and it becomes easy to settle that the lead, and not the 
 barium, is in combination with the carbonic acid. 
 
 These are simple illustrations of the kind of evidence which has 
 
1 88 Qualitative Analysis. 
 
 to be made use of in ordinary qualitative analysis, when endeavour- 
 ing to settle the question of the distribution of the metals and acid 
 radicals. It will be obvious that the success with which the analyst 
 will meet in this part of his work will largely depend upon the power 
 which he possesses of seeing (i.e. of observing) all the various 
 indications which are given him during the course of the analysis. 
 The chief aim of the analytical student, as it should be the 
 first object of his instructor, must always be the cultivation and 
 development of these powers of exact observation. Without this, 
 anylysis sinks into a mere dull mechanical routine, utterly devoid 
 of all interest or educational value. 
 
BOOK II. 
 
 Q UA NT1TA Tl VE A NA L YSSS. 
 
 PART I. 
 
 GRAVIMETRIC METHODS. 
 
 SECTION I. 
 PRELIMINARY MANIPULATIONS. 
 
 i. Weighing. In gravimetric methods of analysis, as the name 
 indicates, the quantities which it is the object of the analysis to 
 determine, are estimated by the process of weighing. Not only is 
 the amount of the material employed for analysis a weighed quantity 
 (which may also be the case in volumetric methods), but the final 
 products of the various processes are obtained in a form in which 
 they also can be weighed. For instance, in making a quantitative 
 analysis of, say, a silver coin, a weighed quantity of the metal will 
 be operated upon, and the silver and copper will finally be obtained 
 in the form of silver chloride and copper oxide. The exact quanti- 
 ties of these products will then be ascertained by the process of 
 weighing, and from the weights so obtained the proportions of the 
 two metals are deduced by calculation. 
 
 In order to obtain the accurate weight of a substance, it is 
 necessary to employ a delicate balance, and to perform the opera- 
 tions of weighing with careful manipulation. 
 
 Fig. 15 represents a modern form of chemical balance. It is 
 contained in a glass case with counterpoised sliding sashes, to 
 exclude from it as much as possible, dust, damp, and laboratory 
 fumes ; * and also to prevent disturbance due to air-draughts during 
 the process of weighing. 
 
 * Unless circumstances render it unavoidable, the balance should never be 
 kept in the laboratory. In all properly equipped schools a balance-room is 
 provided, the instruments being placed where they will be subject to the least 
 possible vibration, and never being moved. 
 
190 Quantitative Analysis : Gravimetric Methods. 
 
 The case is mounted upon levelling-screws, by means of which 
 the instrument can be placed in a perfectly horizontal position, 
 which is indicated by the spirit-levels attached to the base of the 
 pillar. 
 
 When the balance is at rest, the beam and the scale-pans are 
 held by the mechanical supports, which can be operated upon by 
 means of the handle or " milled-head " placed outside and in front 
 of the case. 
 
 FIG. 15. 
 
 By a slight turn of the milled-head, these supports are simul- 
 taneously lowered, and the beam and pans are free to swing upon 
 their "knife-edges." 
 
 The extent to which the beam swings is indicated by a pointer 
 and scale ; and if the balance is in proper order and adjustment, it 
 will be seen that as the beam moves, the pointer travels practically 
 an equal number of graduations in each direction.* 
 
 * If the pointer swings further in one direction than in the other, either the 
 balance is not level, or is out of adjustment. The devices for adjusting the 
 instrument differ with different balances. In the example shown in the figure, 
 it is effected by means of the two little " bobs " which travel upon two screws 
 attached to the ends of the graduated scale of the beam. 
 
Prelim inary Man ipulations. 
 
 191 
 
 When not in use, the balance must never be left in this condition, 
 but the supports must always be raised. 
 
 It is desirable to keep inside the balance case a glass jar about 
 half filled with dry calcium chloride,* or fragments of quicklime, 
 with a view to render the atmosphere of the case as dry as possible. 
 This is especially necessary when any of the vital working parts of 
 the instrument are made of steel instead of agate. 
 
 A set of gram weights, suitable for chemical analysis, is seen in 
 Fig. 1 6. The weights from I gram to 50 grams (usually arranged 
 to make up 100 grams in all) are of brass, while the fractions of 
 the gram are flat pieces of platinium or aluminium foil. In 
 practice it is usual to employ 
 only the " deci " and " centi " 
 grams of these small weights 
 (i.e. the first and second places 
 of decimals), and for the milli- 
 grams and fractions of a milli- 
 gram (the third and fourth 
 decimals) to use a sliding 
 weight upon the beam of the 
 balance. For this purpose 
 the beam is graduated (in the 
 figure, into tenths and hun- 
 dredths), and the sliding 
 weight, or rider (consisting of 
 a piece of gilded brass wire, 
 bent so that it can be placed astride the beam), is moved to any 
 desired position on the beam by means of the carrier which passes 
 through the side of the balance case. The weight of the rider is 
 one centigram, and when it is placed on the tenth main gradua- 
 tion, its effect is the same as if it were placed in the scale, i.e. 10 
 milligrams or i centigram. The divisions of the beam, therefore, 
 represent milligrams and sub-divisions of a milligram. 
 
 The weights must on no account be touched with the hand, but 
 should be manipulated entirely by means of the forceps supplied 
 with each set. 
 
 Great care should be taken in touching these parts, and it may be accepted 
 as a general rule that the young student should not himself attempt the adjust- 
 ment of the balance. 
 
 * The vessel should not be filled, especially when calcium chloride is used ; 
 as the salt becomes liquid by the absorption of moisture, and is likely to overflow 
 the vessel. This is liable to happen when the balance is left unused for any 
 lengthy period, as, for example, during a vacation. 
 
 FIG. 16. 
 
1 92 Quantitative Analysis: Gravimetric Methods. 
 
 It is essential that the weights all bear the exact relation to each 
 other which is professed by the values indicated upon them ; that 
 is to say, the i-gram must be exactly half the 2-gram weight, and 
 this must be exactly a fifth of the 10 gram, and so on.* To ascer- 
 tain that this is actually the case, they should be tested against each 
 other first by seeing that the several i-gram weights are equal to 
 each other ; then that two of them exactly counterpoise the 2 -gram ; 
 then that the 2-gram and the three i -grams together equal the 
 weight of the 5-gram, and so on.f When using the balance, the 
 object to be weighed should be placed upon the left scale-pan, 
 as it is more convenient for manipulation to use the right-hand 
 pan for the weights (a left-handed operator will reverse this 
 order). 
 
 Whenever anything (either objects to be weighed, or the weights) 
 is put upon the scale or removed from it, the beam must be at rest 
 on its supports. 
 
 The actual operation is conducted as follows : The article to be 
 weighed, say a porcelain crucible with its lid, is placed upon the 
 left scale-pan, the balance being at rest. A weight, which by a 
 guess is judged to be rather greater than that of the object, is 
 deposited by means of the forceps upon the opposite scale, and 
 the beam gently liberated by turning the milled-head with the 
 left hand. Suppose the 2o-gram weight has been selected, and 
 found too heavy ; it is then removed, returned to its place in the 
 box, and replaced by a lo-gram. If this is too little, the weight 
 of the crucible lies between 10 and 20 grams. The 5-gram is 
 then added, and if still insufficient, the 2-gram weight is added. 
 If this is too much, the 2-gram is removed and a i-gram sub- 
 stituted. If this is still too little, the weight of the object lies 
 between 16 and 17 grams. The subdivisions of the gram are then 
 used in the same systematic order, until the second decimal place 
 is established, that is, until the whole number of centigrams is 
 ascertained. Thus, suppose the weight is found to be between 
 1674 and 1675, tne front sash is then closed, and the system 
 brought into complete equilibrium by means of the rider, which is 
 moved from division to division along the graduated beam, until 
 the oscillation of the pointer is equal in both directions. Suppose 
 that when this is the case, the position of the rider is midway 
 
 * It is not necessary, although it is preferable that it should be so, that the 
 gram weight should be absolutely true to the standard gram, so long as all the 
 weights used bear the correct ratios to each other. 
 
 f It is extremely rarely that the weights of a reliable maker, such, for 
 example, as Oertling, are ever found to be untrue. 
 
Drying and weighing a Filter. 193 
 
 between the third and fourth main divisions, the weight of the 
 crucible will then be represented by the number 167435 grams.* 
 
 In recording a weight, which should be done immediately the 
 weighing operation is finished, the value of the weights should first 
 be read off from the empty spaces in the box, and this should then 
 be checked as the weights are returned one by one, beginning with 
 the highest, to their respective compartments. 
 
 Except in the case of pieces of metal, alloys, etc., substances 
 which have to be weighed must never be placed directly upon the 
 scale-pan, but must be contained in a suitable vessel, such as a 
 weighing-bottle or a watch-glass, which itself must be perfectly 
 clean and dry. 
 
 The correct weight of an object cannot be taken while it is hot, 
 owing to the disturbance introduced by the warm air-currents rising 
 from it. Substances, therefore, which require to be heated previously 
 to being weighed, must be allowed to cool down to the temperature 
 of the air in the balance-room (see Desiccator, p. 196). 
 
 When taking a weighed quantity of a substance for analysis, it 
 is usual to obtain its weight by difference. The 
 prepared powder is contained in a light thin glass 
 stoppered bottle (a convenient form of weighing- 
 bottle is shown in Fig. 17), and the bottle with its 
 contents is weighed. 
 
 A sufficient quantity of the powder is then care- 
 fully tipped out into the beaker or other vessel 
 selected in which to operate upon the compound, 
 and the bottle is again weighed. The difference 
 between the two weighings represents the weight of 
 the substance employed for the analysis. 
 
 Sometimes the clean empty vessel in which the substance is to 
 be afterwards treated is itself weighed, then a quantity of the sub- 
 stance is introduced into it, and the vessel again weighed. This 
 would be the method, for instance, when the substance is to be 
 operated upon in a crucible. 
 
 2. Drying and weighing a Filter. The filter-paper most 
 suitable for use in quantitative analysis is a specially prepared 
 Rhenish paper (Sdileicher and Schiill, No. 589), which has been 
 treated with hydrochloric and hydrofluoric acids in order to remove 
 
 * This method of weighing only gives the absolute weight of any object 
 when the bulk or volume of the article is the same as that of the weights used 
 to counterpoise it. In ordinary analytical operations, the difference between 
 the weight of the air displaced by the objects and that displaced by the weights 
 is too insignificant to he considered. 
 
 O 
 
194 Preliminary Gravimetric Manipulations. 
 
 soluble mineral matters. The paper is obtained already cut into 
 circles, the two most useful sizes for general use being respectively 
 9 and 1 1 centimetres (or, roughly, 3! and 4^ inches) in diameter. 
 
 In certain quantitative estimations, as will be seen later, it is 
 necessary to weigh a precipitate which has been dried upon a filter, 
 and in such cases it is needful to know the weight of the dry paper 
 itself, in order to deduct this from the total weight. The process 
 of making this simple determination will form a useful exercise in 
 a number of different operations which have constantly to be made 
 in quantitative analysis. 
 
 The apparatus in which to weigh the filter may be either a light 
 stoppered tube or weighing-bottle (A, Fig. 18), or a pair of watch- 
 
 FIG. 18. 
 
 glasses with ground edges, held together by means of a bent wire 
 clip (B, Fig. 1 8). A duplicate experiment may be made, using 
 each of these forms of apparatus. 
 
 The bottle and the watch-glasses are first rendered perfectly 
 clean by means of a dry clean glass-cloth, and are then made quite 
 dry by being placed for about half an hour in a steam-oven ; the 
 bottle with its stopper out, and the glasses separated from each other. 
 
 The steam-oven consists of a double-walled copper vessel (Fig. 
 19), the space between the two walls being partly filled with water, 
 which is maintained at the boiling temperature by means of a 
 Bunsen lamp. In order to keep the water at a constant level, the 
 oven may be provided with a feeding arrangement, shown at /in 
 the figure. This consists of a syphon, the recurved end of which 
 is constantly immersed in water, which is continuously flowing in 
 at the bottom of the outer tube, and overflowing by the lateral tube 
 into the sink.* 
 
 * In most well-appointed laboratories, stacks of steam-ovens, heated by 
 steam from a steam-boiler, form a part of the regular fittings. The temperature 
 within a steam-oven never quite reaches 100. 
 
Drying and weighing a Filter. 195 
 
 The weighing-bottle and watch-glasses are placed upon the 
 floor of the oven (which must be quite clean), or upon a piece of 
 clean paper placed first upon the floor. At the expiration of about 
 half an hour the apparatus is removed (the stopper being inserted 
 in the bottle, and the glasses placed within the clip), and deposited 
 in a desiccator, where it is allowed to cool. 
 
 FIG. 19. 
 
 The desiccator is an air-tight glass vessel, in which the enclosed 
 air is kept dry by means of some convenient drying agent, such as 
 lime, calcium chloride, or sulphuric acid. In its most simple form, 
 it may be an ordinary bell-glass with a ground edge, standing upon 
 a thick piece of ground glass, over a glass dish containing the 
 desiccating agent. The ground edge of the bell is greased with 
 resin cerate, and the objects to be kept dry are supported upon a 
 perforated zinc cover upon the inner dish (Fig. 20). 
 
 A handy form of desiccator, and one more easily carried about, 
 is shown in Fig. 21. The desiccating material is placed in the 
 lower part of the vessel, and a disc of perforated zinc, resting upon 
 the shoulders, serves as a support for the objects to be dried. The 
 whole is surmounted with a cover, ground to fit the vessel, which, 
 
196 l^reliniinary Gravimetric Manipulations. 
 
 as in the other example, is greased with resin cerate.* When the 
 apparatus has cooled in the desiccator, it is next carefully weighed, 
 the weight of the empty tube and that of the watch-glasses being 
 separately noted. 
 
 Two of the filter-papers are now folded in the usual manner. 
 One of them is then rolled loosely, and introduced into the weigh- 
 ing-bottle ; the other is carefully folded, without cracking the paper 
 (because in actual work it would afterwards be used for filter- 
 ing), so that it can be placed between- the watch-glasses. Each 
 piece of apparatus with the enclosed filter is then weighed. The 
 former weights deducted from those now obtained, give the weights 
 of the imdried filters. The weighing-bottle and the glasses (opened 
 
 as before) are then replaced in the steam-oven, and heated for half 
 an hour, after which they are returned to the desiccator to cool, 
 and then re-weighed. The difference between the first weights 
 and these last is the weight of the dried filters. 
 
 By comparing the weights of the undried and the dried filters, 
 the actual weight of the moisture which such a piece of filter-paper 
 is capable of absorbing from the atmosphere will be ascertained. 
 
 The tube and watch-glasses should be once more heated for 
 
 * Resin cerate, as supplied by the shops, is often rather too stiff. It is 
 well in this case to melt a quantity of it in a porcelain dish, and add a little 
 vaseline to it. The whole should then be poured into a clean vessel to cool. 
 An ordinary ointment pot with a cover (such as the druggists use) is a con- 
 venient vessel in which to keep it. Dust and dirt should not be allowed to get 
 into it, as any small particles of gritty matter will prevent the lid of the desic- 
 cator from closing quite air-tight. 
 
Estimation of Water of Crystallisation. 197 
 
 a similar period, cooled as before in the desiccator, and again 
 weighed. The weights now obtained should be concordant with 
 the last ; if they are not, the heating must be repeated until two 
 consecutive weighings are found to agree. 
 
 In actual practice we do not want to know the weight of the 
 moisture expelled from the filter, or even the weight of the dry paper 
 alone. It is enough to know the weight of the dry paper and the 
 weighing-bottle. The process is therefore reduced to a single 
 operation. The paper is placed in an unweighed bottle, and after 
 half an hour's drying, it is cooled in the desiccator and weighed. 
 The paper is then used to receive the precipitate which has to be 
 weighed. Then, after proper treatment, the filter with the precipi- 
 tate upon it is returned to the same weighing-bottle and heated in 
 the steam-oven for a suitable time, after which it is allowed to cool 
 in the desiccator and weighed ; the process of heating and weighing 
 being repeated until two concordant weights are obtained. The 
 weight of the bottle plus paper deducted from the weight of the 
 bottle plus paper plus precipitate, gives the weight of the precipi- 
 tate, which it is the object of the series of operations to ascertain. 
 
 Most substances, especially when in the state of fine powders, 
 are more or less hygroscopic. On exposure to the air, they absorb 
 and retain a certain amount of moisture. Some things such, for 
 example, as copper oxide, manganese dioxide, charcoal powder, etc.. 
 are able to absorb in this mechanical way very appreciable quan- 
 tities of moisture. The water so retained is in almost all cases 
 entirely expelled by heating the substance for a suitable time in the 
 steam-oven, after which, in order to preserve it in its dry state, it 
 must be kept in the desiccator. 
 
 When the amount of moisture such substances contain has to 
 be estimated, a convenient quantity (from 2 to 10 grams, depend- 
 ing upon the amount of moisture) is weighed out into a pair of 
 watch-glasses, and the process conducted in the manner above 
 described in the case of the filter-paper. If the substance is one 
 which cannot withstand a temperature of 100 without suffering 
 chemical decomposition, other methods have to be adopted for 
 estimating the moisture it contains, which will be described later. 
 
 3. Estimation of Water of Crystallisation. Most salts 
 containing water of crystallisation readily part with some or all of 
 their water when moderately heated. In those cases where the salt 
 suffers no other change in composition at the temperature necessary 
 for the expulsion of the water of crystallisation, the latter may be 
 estimated by ascertaining the loss of weight which results from 
 heating the substance. If the salt loses its water of crystallisation 
 
198 Preliminary Gravimetric Manipulations. 
 
 at or below the temperature of 100, the operation may be con- 
 ducted in the steam-oven in the manner described above ; but 
 when higher temperatures are necessary, an air-oven is employed. 
 
 The air-oven (Fig. 22) is similar in construction to the steam- 
 oven, except that it contains no water between the walls. (The 
 oven is frequently only a single-walled copper chamber, fitted with 
 a shelf or false bottom). It is heated by a Bunsen flame, and the 
 
 FIG. 22. 
 
 temperature of the air within is ascertained by means of a thermo- 
 meter, which passes through the top by means of a cork. The 
 door of the oven is usually furnished with a ventilator, which can 
 be opened or closed in order to regulate the current of air through 
 the oven ; the escaping air passes out through a second opening in 
 the top. 
 
 The objects to be heated must be placed upon the perforated 
 shelf (or upon any other convenient support), which is at some 
 
Estimation of Water of Crystallisation. 199 
 
 considerable distance from the bottom of the oven, and the bulb 
 of the thermometer should be as near to the object as possible. 
 Articles to be heated must not be placed directly upon the floor of 
 the oven, as obviously the metal of that part of the oven will be 
 hotter than the temperature indicated by the thermometer. This 
 is more especially the case when the oven is a single-walled 
 apparatus. 
 
 Within reasonable limits, it is easy to maintain the temperature 
 of such an oven at any desired point by regulating the Bunsen 
 tlame. As the mercury in the thermometer rises nearly to the re- 
 quired point, the lamp is gradually turned down until the upward 
 movement of the mercury ceases. If this point is above the 
 desired temperature, and if after a few minutes it does not fall, the 
 flame should be lowered still more. (Thermostats, or instruments 
 which serve as gas-governors, can be fitted to the air-oven, so that 
 the supply of gas necessary to maintain the oven at a constant 
 temperature is automatically regulated.) 
 
 Crystallised copper sulphate may be conveniently used in order 
 to illustrate the method of determining water of crystallisation. 
 This compound crystallises with five molecules of water. Four of 
 these it gives up slowly at 100, but more rapidly at 1 10, while the 
 remaining molecule is retained until the temperature has risen 
 above 200. In the following exercise, each of these proportions 
 of water will be separately determined. 
 
 A quantity of purified copper sulphate is coarsely powdered in a 
 mortar as quickly as possible, and the powder placed in a stoppered 
 bottle (a slight loss of water will result if the powdering process 
 involves a prolonged exposure of the salt to the air). About two 
 grams of it are then weighed out into a pair of watch-glasses and 
 clip, by first weighing the clean dry glasses and clip, and then 
 placing into them a quantity of the salt which is judged to be about 
 two grams, and weighing again. The glasses are then separated, 
 and placed upon the shelf of the air-oven, the latter being main- 
 tained at a temperature between 110 and 115. 
 
 At the expiration of an hour the glasses are removed from the 
 oven, quickly put together in the clip, and allowed to cool in the 
 desiccator, and then weighed. 
 
 They are then returned to the oven and reheated for about 
 half an hour, after which they are again taken out and cooled 
 and weighed. If this last weight differs appreciably from the 
 previous one, the process must be repeated until the weight is 
 constant. 
 
2OO Preliminary Gravimetric Manipulations. 
 
 From these data the percentage of water lost by heating the 
 salt to i io-i 1 5 is calculated in the manner shown in the following 
 example : 
 
 Grams. 
 Weight of watch-glasses and salt ... ... . . 17*3465 
 
 empty ... ... 15-1890 
 
 Weight of copper sulphate used 2<I 575 
 
 Weight of watch-glasses and salt after first heating 167545 
 
 second ,, 167230 
 
 >, ,, third 167230 
 
 Weight of glasses and original salt 17-3465 
 
 dried ,, 167230 
 
 Weight of water lost 0-6235 
 
 Since 2 f i575 grams of salt lost 0-6235 gram of water 
 
 then 2 2 ~ * = 28-90 = percentage of water lost at iio-ii5 
 
 The calculated percentage for four molecules of water = 28-91 
 therefore the error, as deficit, = o'oi 
 
 The temperature of the air-oven is now raised to between 
 210 and 215, and the opened watch-glasses with the salt are again 
 heated for about an hour to this higher temperature. The glasses 
 are then removed, cooled in the desiccator and weighed. As in 
 the former case, this is repeated until two consecutive weighings 
 agree. 
 
 In the example the following weights were obtained : 
 
 Grams 
 
 Weight of glasses and salt after first heating to 215 16-5685 
 ,, ,, ,, second ,, 16-5680 
 
 Weight of glasses and original salt 17-3465 
 
 ,, ,, dried 16*5680 
 
 Total weight of water lost ... ... ... 07785 
 
 Since 2*1575 grams of salt have lost 07785 gram of water- 
 then 77 2 .^ X I0 = 36*08 = percentage of water lost at 215 
 
 Percentage calculated for five molecules) _ 
 of water of crystallisation \ 
 
 Error, as a deficit, = o - o6 
 
 In determining the water of crystallisation of salts which 
 undergo no other chemical change, except the loss of their water, 
 even when heated to a high temperature, it is more expeditious to 
 
Estimation of Water of Crystallisation. 201 
 
 heat the salt in a crucible by means of a small Bunsen flame. 
 Copper sulphate could not be treated in this way, as at a temperature 
 somewhat over 300 the anhydrous salt 'suffers decomposition. 
 A convenient salt with which to illustrate the method is barium 
 chloride, which crystallises with two molecules of water. 
 
 A crucible (preferably of platinum) with its lid, is first heated 
 upon a pipeclay triangle (as shown in Fig. 23) for a few minutes, 
 and weighed after cooling in the desiccator. 
 
 Into the crucible a quantity of pure, powered barium chloride 
 is introduced, and the weight again taken. 
 
 The weight of salt taken should be between two and three 
 
 FIG. 23. 
 
 grams.* The crucible is at first gently heated by a small Bunsen 
 flame, the lid being upon the crucible. The temperature is 
 gradually raised until the crucible attains a low red heat, at which 
 
 * The quantity of material taken for an estimation must depend upon 
 several considerations. Speaking generally, the smaller the weight employed, 
 the more expeditiously will the analysis be made. At the same time, errors 
 due to experiment will be more magnified, and therefore a greater demand is 
 made upon the manipulative skill of the analyst. The relation between the 
 weight of the substance which is to be estimated, and the formula-weight of 
 the compound to be analysed has also to be considered ; for instance, in the 
 present case, owing to the high atomic weight of barium, and the comparatively 
 small quantity of water present, the weight of the latter constitutes only a 
 small part of the total formula-weight of the crystallised salt. 
 
202 Preliminary Gravimetric Manipulations. 
 
 it is maintained for about ten minutes. It is then placed in the 
 desiccator to cool, and afterwards weighed. 
 
 The operations of heating and weighing are repeated until the 
 weight is constant. The following is an example :-- 
 
 Grams. 
 Weight of platinum crucible and salt ...... 47*6275 
 
 alone ......... 45'3 6 4 
 
 Weight of barium chloride taken ...... 2-2635 
 
 Crucible and salt after first heating ......... 47-2950 
 
 ,, ,, second , .......... 47-2940 
 
 Weight of crucible and original salt ......... 47-6275 
 
 ,, ,, dried ,, ... ... ... 47-2940 
 
 Weight of water lost ......... 0-3335 
 
 Since 2-2635 grams of salt contained o'3335 gram of water 
 
 then -~-~r --- = percentage of water of crystallisation 1473 
 
 Percentage calculated on the formula BaCL,2H 2 O = 1475 
 
 Error, as deficit, = o'o2 
 
 In cases where the substance undergoes other chemical changes 
 at, or near, the temperature at which it parts with its water of 
 crystallisation, or with the moisture which is mechanically retained 
 by it, the weight of the water which it gives up cannot be 
 estimated by loss. The method is therefore so modified that the 
 expelled water is absorbed by calcium chloride, or other suitable 
 desiccating material, which is accurately weighed before and after 
 the experiment. The gain of weight thus observed represents the 
 weight of water absorbed. The apparatus for carrying out the 
 method is shown in Fig. 24. 
 
 A weighed quantity of the substance is heated in the bulb-tube 
 B. By means of an aspirator, a slow stream of air (dried by 
 passing through the calcium-chloride tube A) is drawn through 
 the bulb, and the water-vapour which is being expelled is absorbed 
 in the weighed calcium-chloride tube D. The wash-bottle W, 
 containing sulphuric acid, serves to show the rate at which air 
 is being drawn through the apparatus, and also prevents the 
 drying-tube D from absorbing water-vapour from the aspirating 
 arrangement. 
 
 Fitting iip the Apparatus. The two drying-tubes are first 
 perfectly cleaned and dried, and provided with tight-fitting corks * 
 
 * When ordinary corks are employed in fitting up apparatus generally, 
 those having a fine,' close texture and free from faults should alone bo used. 
 
Estimation of Water of Crystallisation. 203 
 
 (either ordinary or caoutchouc) carrying tubes as shown in the 
 figure. They are then nearly filled with dry granulated calcium 
 chloride, the pieces of which should be about the size of grains of 
 wheat. Powder must not be put into the tubes, or they will be 
 liable to become blocked ; therefore, if the stock of calcium chloride 
 contains much that is too fine for the purpose, it may be passed 
 as quickly as possible through a warm dry wire sieve. A little 
 plug of cotton-wool is loosely pushed down upon the top of the 
 calcium chloride, and the tubes securely closed by the corks. 
 
 FIG. 24. 
 
 In the case of the tube D, a little plug of cotton-wool is first 
 introduced, and drawn round to the top of the bulb by applying 
 suction to the end of the narrow tube, in order to prevent particles of 
 the calcium chloride from getting into the narrow tube. The cork 
 which closes the open limb of this tube should be cut off close to 
 the glass with a sharp knife, and the surface completely covered 
 by shellac or sealing-wax. This prevents the cork from absorbing 
 moisture from the air. 
 
 The cork should be at first a trifle too large for the tube it is intended to fit. 
 It is then carefully pressed in cork-squeezers, or rolled under the foot with a 
 gentle pressure, which will both soften it and make it a little narrower. The 
 hole should be cut with a sharp, clean cork-borer to the exact size required to 
 tightly fit the tube which it is intended to pass through it, without having to 
 use a round file to enlarge it. 
 
2O4 Preliminary Gravimetric Manipulations. 
 
 The ends of the exit tubes are closed by means of little caps, 
 made by plugging one end of a short piece of caoutchouc tube 
 with a piece of glass rod. These caps should remain on the tubes 
 except when they are actually being used, or being weighed. The 
 tube D, which has to be weighed, is furnished with a loop of wire 
 (preferably platinum), by means of which it can be suspended from 
 
 the arm of the balance. 
 The two drying-tubes are 
 attached to a hard glass 
 bulb - tube by means of 
 rubber or ordinary corks 
 in the manner shown in the 
 figure. If ordinary corks 
 are used, they should first 
 be thoroughly dried in the 
 steam-oven. A simple form 
 of aspirator for small pres- 
 sures is shown in Fig. 25. 
 An ordinary T-tube is at- 
 tached to a narrow glass 
 tube about 70 cm. long, 
 which has been bent into 
 a loop near one end. When 
 a gentle stream of water 
 from the tap is allowed to 
 enter through the tube S, 
 air is drawn in through the 
 other limb of the T-tube 
 and carried down the long 
 tube. By means of the 
 clamp D, the rate at which 
 air is drawn through the 
 apparatus can be regu- 
 lated. During the experi- 
 ment it should be such that the bubbles pass through the sulphuric 
 acid in the little wash-bottle W about two per second. 
 
 To carry out the Operation. The drying-tube D is first care- 
 fully weighed, 'without the little caps which close the ends. These 
 are instantly replaced as soon as the weight is ascertained. The 
 bulb-tube, after being carefully wiped and dried, is weighed, and 
 a suitable quantity (2 to 4 grams) of the substance is introduced 
 into the bulb. This may be accomplished by placing the powdered 
 
 FIG. 25. 
 
Estimation of Water of Crystallisation. 205 
 
 substance upon a strip of writing-paper, folded into a gutter 
 sufficiently narrow to pass into the tube (A, Fig. 26), and depositing 
 it in the bulb by twisting the paper (B, Fig. 26). Before with- 
 drawing the paper, it should be turned over again into its original 
 position, so that on its way out it may not leave any traces of the 
 powder in the stem of the bulb. The bulb is now re-weighed, the 
 increase being the weight of the substance taken. 
 
 The two U-tubes are then attached to the bulb, the caps being 
 removed only from the extremities which are fitted to the bulb- 
 
 FIG. 26. 
 
 tube. The aspirator is then set in operation, and when the rate 
 at which it draws air through the acid in bottle W is adjusted to 
 the requirements of the experiment, the caps are removed from the 
 drying-tubes, and the bottle W is attached by a narrow caoutchouc 
 tube to the drying-tube D. Two screens, made of asbestos card- 
 board, are placed in the position shown in the figure to shield the 
 corks and the drying-tubes from the heat of the lamp. 
 
 The bulb is gently heated by means of a small Bunsen flame, 
 the temperature being gradually increased until the whole of the 
 water is expelled. Some of the moisture will be seen to condense 
 
206 Preliminary Gravimetric Manipulations. 
 
 upon the stem of the bulb-tube, but as the heated air passes through 
 the apparatus, this gradually vaporises and passes on into the 
 absorbing-tube D. When the whole of the water has been expelled, 
 tube D is detached and re-weighed. Its gain in weight represents 
 the amount of water contained in the weighed amount of the 
 original compound. 
 
 When it is desirable to heat the substance to some definite 
 temperature, the bulb-tube is replaced by the U-tube (Fig. 27), 
 which can be heated to the required tem- 
 perature in a bath of oil or melted paraffin, 
 the temperature being ascertained by a 
 thermometer suspended in the bath. 
 
 4. Determination of the Weight 
 of the Filter-ash. In the majority of 
 cases, before a precipitate is weighed, it 
 is necessary that it be subjected to a high 
 temperature. It has, in fact, often to be 
 strongly heated in a crucible, in order to 
 convert it into a product suitable for weigh- 
 ing. For example, the precipitate may con- 
 sist of a hydrated oxide (perhaps of indefinite 
 composition) which it is necessary to con- 
 vert into the oxide (having a known and 
 definite composition) by the action of heat. 
 In such a process, obviously, the paper be- 
 comes burnt up, and the ash only is weighed. 
 It is therefore needful to know what is the 
 weight of the ash so produced, that it may 
 be deducted from the total weight. 
 
 For this purpose a number of the papers are incinerated one 
 after the other in the following manner : 
 
 A clean crucible, with its lid (either a platinum or porcelain 
 crucible may be used), is supported upon a pipeclay triangle 
 (Fig. 23), and heated, by means of a Bunsen flame for a few 
 minutes. It is then allowed to cool in the desiccator, and weighed. 
 One of the filters is then folded as shown (a, Fig. 28), and rolled 
 into a compact little bundle (). A piece of moderately stout 
 platinum wire is then wound a few times round it, leaving a 
 sufficient length of wire to serve as a handle,* as shown at c. The 
 
 * If, as is sometimes recommended, the wire is fused into a glass tube for 
 a handle, the greatest care must be taken to prevent any little fragments of 
 glass, which are liable to chip off by the bending of the wire, from falling into 
 the crucible. 
 
Determination of a Filter-ash. 
 
 20; 
 
 crucible is then placed upon a clean sheet of white paper (or upon 
 a square of glass standing on white paper), with the lid by its side. 
 
 FIG. 28. 
 
 The rolled-up paper, being held over the crucible as in Fig. 29, is 
 then ignited (set fire to) by means of a Bunsen flame. After the 
 flame of the burning paper has died out, the charred and shrunken 
 remains still smoulder on for some time, until the carbon has burnt 
 away and a grey mass is left. It may be lightly touched once or 
 twice with the Bunsen flame (it must not be held for any length of 
 time in the gas-flame), and then allowed to drop into the crucible. 
 
 FIG. 29. 
 
 If the ash does not easily fall away from the wire as the latter is 
 inclined over the open vessel, a gentle tap of the wire against the 
 rim of the crucible will usually detach it. 
 
 This process is repeated (two papers may be burnt at a time), 
 until the ash from ten filters has been obtained. If any minute 
 particles have fallen on to the paper or glass, they are carefully 
 swept up into the crucible with a small camel's-hair brush or a 
 feather. The crucible is then covered, heated again for a few 
 minutes as at first, placed in the desiccator to cool, and weighed. 
 The increase in weight is the weight of the ash from ten filters, and 
 by dividing the result by ten, the average weight of the ash of 
 
208 Preliminary Gravimetric Manipulations. 
 
 one is obtained. If the filters mentioned on p. 193 are used, the 
 weight of the ash they yield is indicated on the packets ; the student 
 will therefore be able to compare the result of his determination 
 with the figure there given. It will be obvious that the larger the 
 number of papers incinerated, the more distributed will be the 
 experimental errors, and therefore the more exact will be the result. 
 
 5. Preparation of Pure Salts. In order to test the 
 accuracy of the results of a quantitative analysis, two courses are 
 open. Either duplicate analyses are made, so that two concordant 
 results may be obtained ; or the result of a single analysis is com- 
 pared with the calculated theoretical composition of the compound 
 under examination. It will be obvious that it is only when com- 
 pounds of known composition and of known purity are being 
 analysed that the latter plan is possible ; in all professional work 
 the former method is necessarily adopted. For the student, 
 however, who is beginning the practical study of quantitative 
 methods, it is advantageous that his first exercises in quantitative 
 estimations should be made with salts which are practically pure 
 compounds. Purified salts (i.e. salts which have been so far 
 purified that the amount of impurity present is too small to 
 appreciably affect the analysis) are readily obtained from chemical 
 manufacturers, but it is well for the student himself to prepare 
 some of the compounds he will use, in order that he may gain 
 experience in the general methods employed for this purpose. 
 
 i. Crystallisation. By the process of crystallisation from a 
 suitable solvent (repeated several times, if necessary), a salt may 
 be separated from most soluble substances with which it is 
 admixed. The wider the difference between the solubilities of the 
 substances to be separated, the more easily is purification by this 
 method effected ; e.g. barium chloride is more readily freed from 
 calcium chloride by recrystallisation than from potassium chloride, 
 the solubility of the latter salt at the ordinary temperature being 
 only slightly different from that of barium chloride, while calcium 
 chloride is an extremely soluble salt. 
 
 Certain salts, on the other hand, cannot be separated from each 
 other by the process of crystallisation ; because when the solution 
 containing them is allowed to crystallise, the two are together 
 deposited, not as a mere mixture of crystals of the two individual 
 compounds, but in the form of a double salt, having a definite 
 composition. 
 
 The operation of recrystallising a salt is similar in outline in 
 every case, with certain differences in detail. 
 
Preparation of Pure Salts. 
 
 209 
 
 A hot strong solution of the commercial salt is made in water. 
 For this purpose a quantity of water, say from 200 to 300 c.c., is 
 heated in a beaker, either by means of a rose burner, the beater 
 being supported upon wire gauze as in Fig. 30, or by means of a 
 small Fletcher's burner (Fig. 31), which is a convenient lamp and 
 stand in one. A quantity of the salt to be recrystallised is then 
 added, moderate portions at a time, until a tolerably strong solution 
 is obtained. The quantity of salt 'which will be dissolved will 
 depend obviously upon the solubility of the compound operated 
 
 FIG. 30. 
 
 FIG. 31. 
 
 upon ; but it is sometimes necessary to take other considerations 
 into account in fixing a limit to the strength of the solution which 
 is made. For example, zinc sulphate and magnesium sulphate, 
 when deposited from solutions at a temperature below 40 and 50 
 respectively, form crystals containing seven molecules of water ; 
 but if the crystals are deposited from solutions above these tempe- 
 ratures, they not only have a different form, but contain only six 
 molecules of water of crystallisation. Therefore the solution must 
 either be of such a degree of concentration that crystals do not 
 form until the temperature has fallen to these degrees, or else the 
 temperature of the water in which the salt is dissolved should not 
 be higher than these points. 
 
 Unless the solution is absolutely clear and bright (which, with 
 the ordinary commercial salts, will hardly ever be the case), it 
 
2io Preliminary Gravimetric Manipulations. 
 
 must be freed from suspended impurity by filtration. In order to 
 prevent the solution from crystallising in the funnel, and so block- 
 ing the filter, it is neces- 
 sary to keep the funnel 
 warm during the opera- 
 tion. This may be 
 done by inserting the 
 funnel in a double- 
 walled metal jacket, 
 made in the shape of a 
 truncated cone. The 
 jacket contains water 
 which is heated by 
 means of a Bunsen 
 flame (Fig. 32). The 
 same result may be ac- 
 complished by winding 
 
 FIG. 32. a piece of lead or 
 
 " compo " pipe round 
 the funnel, two or three turns being enough, and blowing steam 
 
 through the pipe. The steam may conveniently be generated in 
 a common tin can. The arrangement is shown in Fig. 33. 
 
Preparation of Pure Salts. 
 
 211 
 
 A still simpler plan for keeping the funnel warm is shown in 
 Fig. 34. A short piece of " compo " pipe is closed at one end, and 
 bent at that end into a loop or ring. Five or six pinholes are 
 bored in the ring, and the pipe is supported by means of a clamp, 
 so that the ring is at some little distance below the body of the 
 funnel. The pipe is con- 
 nected to the gas-supply, 
 which is regulated so 
 that little beads of flame 
 are burning at the holes 
 in the pipe. 
 
 The beaker contain- 
 ing the filtered solution 
 is then placed in a basin 
 or other convenient 
 vessel containing cold 
 water, in order to cool 
 the liquid as quickly as 
 possible. The solution 
 is also continually stirred 
 with a glass rod. In 
 
 this way the salt is deposited in the form of very small crystals, 
 whereas if allowed to cool slowly and to be undisturbed, the crystals 
 would be larger and better defined. In the former condition the 
 salt is more easily freed from the mother-liquor. 
 
 When the solution is cold, the crystals are allowed to settle, 
 and the mother-liquor decanted off. The residue is then trans- 
 ferred to a drainer, which is a circular disk of porcelain or glass, 
 perforated with a number of small holes, and supported in a glass 
 funnel. In this way the greater part of the mother-liquor is 
 separated from the crystals. The process is accelerated by partially 
 reducing the atmospheric pressure beneath the funnel. For this 
 purpose the stem of the funnel is fitted into the mouth of a stout 
 conical flask by means of a perforated caoutchouc stopper, and the 
 branch tube of the flask is connected to any suitable exhausting 
 apparatus, such as a Geissler's or Wetzel's water-pump. The 
 arrangement is shown in Fig. 35, where a Wetzel's water-pump is 
 represented, which is one of the best forms of its kind. Between 
 the pump and the filtering-flask an empty Woulff's bottle should 
 intervene, with the tubes arranged as shown. The object of this 
 is to intercept the water which is sucked back from the pump 
 towards the flask, when the pump is stopped while there is still a 
 
212 Preliminary Gravimetric Manipulations. 
 
 partial vacuum in the apparatus. The water which is thus driven 
 into the Woulff's bottle, is automatically drawn out again when next 
 the pump is set in action. The rubber tube used for the water- 
 supply must bfc canvas-covered or canvas-lined, in order to with- 
 stand the water-pressure ; that employed for the other connections 
 must be sufficiently thick in the walls to resist the pressure of the 
 atmosphere, without collapsing. 
 
 In cases where the crystals are only moderately soluble in cold 
 water, they may be rinsed while upon the drainer by pouring a 
 little cold water upon them. After they have been allowed to 
 
 drain as completely as possible in the funnel, they are then spread 
 out upon a clean porous plate * or tile to dry. 
 
 If the salt being operated upon is one which does not effloresce 
 or undergo any change on exposure to the air, it may be left until 
 quite dry, care being taken to prevent particles of foreign matter 
 from falling upon it. In the case of compounds which cannot 
 safely be exposed so long, the crystals are finally dried by gently 
 pressing them between folds of filter-paper. When dry, the crystals 
 should be transferred to a dry stoppered bottle. 
 
 After this manner, any of the following salts may for practice 
 
 * Porous plates are most convenient for the purpose, and are cheaper than 
 the tiles. They are ordinary plates which, on account of some imperfection 
 or slight damage, are not worth carrying any further in the process of manu- 
 facture than the ' ' biscuit " or unglazed stage. A porous plate (or tile) must 
 not be used a second time (unless for another crystallisation of the same salt), 
 as it would obviously introduce impurities into the compounds being dried. 
 
Preparation of Double Sa//s. 213 
 
 be recrystallised : CuSO 4 ,5H,O ; MnSO^H.O ; ZnSO 4 ,;H,(3 ; 
 NiSO 4 ,7H 2 O; MgSO 4 ,7H,Oj BaCl 2 ,2H 2 O ; KC1 ; NH 4 C1 ; 
 K 2 CrO 4 ; K 2 Cr 2 O 7 . 
 
 Double salts are salts which contain two metals and only 
 one acid radical.* They are usually formulated as associa- 
 tions of molecules of the two single salts which enter into their 
 composition : thus, potassium aluminium sulphate (potassium 
 alum), K 2 SO 4 ,A1 2 (SO 4 ) 3 ,24H 2 O ; ammonium nickel sulphate, 
 (NH 4 ) 2 SO 4 ,NiSO 4 ,6H 2 O ; potassium magnesium sulphate, 
 K 2 SO 4 ,MgSO 4 ,6H 2 O ; and so on.f The exact nature of the union 
 between the molecules of the two single salts is not known, but it 
 is a union which exhibits shades of stability just as in the case of 
 that which exists between the atoms within the molecules. For 
 instance, the union between the two sulphates in potassium alum 
 is a much closer or more stable union than that between the two 
 sulphates in potassium magnesium sulphate. When alum is dis- 
 
 * Although it is true that " double salts" contain two metals and one acid 
 radical, it does not follow that all salts which contain two metals and one 
 acid radical belong to this class of compounds. For instance, microcosmic salt, 
 HNa(NH 4 )PO 4 , and ammonium magnesium phosphate, NH 4 MgPO 4 , are 
 examples of salts containing two metals (regarding NH 4 as a metal) and one 
 acid radical, which are not classed as " double salts." Strictly speaking, a 
 compound is only a " double salt " when the sum of the effective valencies of 
 the two metals is not greater than the basicity of the single acid radical. Thus, 
 the basicity of the acid radical (PO 4 ) is 3. In microcosmic salt the sum of the 
 valencies of the two metals Na and (XH 4 ), is only 2, while in ammonium 
 magnesium phosphate it is 3. On the other hand, in such a compound as 
 ferrous ammonium sulphate, (NH 4 ) 2 SO 4 , FeSO 4 ,6H2O, which is a true double 
 salt, the sum of the valencies of the metals present is 4 (two monovalent X1I 4 
 groups and one di-valent iron), while the single acid radical (SO 4 ) has a basicity 
 only of 2. 
 
 t Some chemists adopt formulae which are constructed on the model of the 
 formula of a single salt ; thus, alum is represented as A1 2 K2(SO 4 ) 4 ,24H 2 O, ferrous 
 ammonium sulphate as Fe(NH 4 )2(SO 4 )2,6H 2 O ; and so on. The use of formula; 
 of this type, however, is more advantageously reserved for a special class of 
 double salts, which differ from those here mentioned in their behaviour towards 
 reagents in a very marked manner. For example, if barium chloride be added 
 to a solution of either of the double sulphates mentioned above, all the sul- 
 phuric acid radical is precipitated as barium sulphate, while chlorides of both 
 metals are formed ; thus 
 
 K 2 SO 4 , A1 2 (SO 4 ) 3 + 4BaCl 2 = 4BaSO 4 + 2KCI + A1 2 C1 6 
 
 Certain double salts, however such, for instance, as the double chlorides of 
 platinum and the alkali metals do not behave in this way. Thus, if silver 
 nitrate be added to a solution of sodium platinic chloride, 2NaCl,PtCl 4 , the 
 whole of the chlorine is not precipitated as silver chloride ; the silver is only- 
 capable of uniting with 2 out of the 6 atoms of chlorine in the compound, the 
 result being the production of the compound 2AgCl,PtCl 4 and two molecules 
 of XaNO 3 . Such a compound as this may with advantage be expressed by the 
 formula Na 2 PtCl 6 , and the action of silver nitrate may be indicated by the 
 '<]uation 
 
 Xa,PtCl 8 + 2 AgX0 3 = 2 XaN0 3 + Ag 2 PtCl fl 
 
214 Preliminary Gravimetric Manipulations. 
 
 solved in water (without entering upon the much-debated question 
 as to the condition of the salt while in solution), the solution, on 
 evaporation, again deposits the double salt unchanged ; but when 
 the double potassium magnesium sulphate is similarly treated, the 
 salt undergoes a partial separation, and the crystals which are 
 deposited consist partly of the double salt and partly of the least 
 soluble of the two single salts, namely, the potassium sulphate. 
 
 These facts have to be kept in mind in the preparation or puri- 
 fication of double salts, for it is obvious that, while the more stable 
 of this class of compounds (such as alum) can be purified by 
 recrystallisation, the process is inadmissible in the case of those 
 salts which readily separate into the simple salts of which they are 
 composed. The general method for the preparation of double salts 
 is to take the two single salts in the proportion of their formula- 
 weights, make a strong solution of each separately in water, and 
 carefully mix the two liquids. For example, to prepare potas- 
 sium alum, the two salts, aluminium sulphate, Al 2 (SO 4 ) 3 ,i8H 2 O 
 (formula- weight = 612), and potassium sulphate, K 2 SO 4 (formula 
 weight = 174), are taken in the proportions indicated by these 
 weights. (A suitable quantity in this case would be one-third 
 of these numbers in grams, i.e. 204 grams of aluminium sulphate 
 and 58 grams of potassium sulphate.) If the salts are both pure 
 (so called) compounds, the weights employed may be as near to 
 the theoretical proportions as will be obtained by weighing the 
 salts on a rough balance. If, however, ordinary commercial salts 
 are used, a slight excess of aluminium sulphate should be taken, as 
 this is likely to be the less pure of the two compounds. The two 
 salts, in separate beakers, are dissolved in hot water, and the solu- 
 tions filtered if they contain any suspended impurities. A slight 
 loss of the salts may occur in this operation, but by using 
 moderately small filters, and exercising care, the loss will be 
 practically equal in both cases. The two clear solutions are then 
 thoroughly mixed by pouring them from one beaker to the other 
 once or twice, carefully avoiding any loss of either liquid. The 
 solution is then cooled, and the crystals drained and dried in the 
 manner already described. 
 
 In those cases where the union between the two component 
 salts in the double compound is less stable, as already mentioned, 
 the more insoluble of the single salts is liable to crystallise out in 
 greater or less quantity along with the double compound. To 
 prevent this, a considerable excess of the more soluble ingredient 
 must be employed ; usually an excess equal to one-fourth or one- 
 
Preparation of Pure Salts, 2 1 5 
 
 third of the formula-weight is employed. Thus, in the case of the 
 double potassium magnesium sulphate above cited, instead of 246 
 (the formula-weight of MgSO 4 ,7H 2 O) and 174 (formula-weight of 
 K 2 SO 4 ), the proportions used should be 246 + 80 to 174. 
 
 In special cases such, for instance, as the double salts contain- 
 ing ferrous sulphate as one of the components the separate solu- 
 tions must be made at temperatures lower than the boiling-point, 
 in order to guard against the precipitation of basic salts. Thus, in 
 preparing ferrous ammonium sulphate, FeSO 4 ,(NH 4 ) 2 SO 4 ,6H 2 O, 
 the temperature of the solutions should not be higher than about 
 40 C. 
 
 The two salts are taken in the proportion of their formula- 
 weights (FeSO 4 ,7H 2 O = 278, and (NH 4 ) 2 SO 4 = 132), and dis- 
 solved in water so as to obtain solutions as strong as possible at a 
 temperature of about 40. The liquids are then mixed, and the 
 mixture acidified by the addition of two or three drops of dilute 
 sulphuric acid. The solution is then cooled and crystallised as usual. 
 
 For practice, such double salts as the following may be 
 prepared : NiSO 4 ,(NH 4 ) 2 SO 4 ,6H 2 O ; MnSO 4 ,K,SO 4 ,6H,O : 
 2(NH 4 )Cl,CuCl 2 ,2H 2 O. 
 
 Precipitation by Change of the Solvent, Salts may be 
 thrown out of their aqueous solution by the addition to the liquid 
 of some substance which is miscible with, or soluble in, the water, 
 but which, when so dissolved or mixed, constitutes a solution in 
 which the salt previously dissolved in the water is insoluble. For 
 instance, barium chloride is soluble in water but is almost in- 
 soluble in a solution of hydrochloric acid in water. If, therefore, 
 hydrochloric acid be added to such an aqueous solution, the barium 
 chloride is at once thrown out of solution as a fine crystalline pre- 
 cipitate. 
 
 Again, ferrous sulphate is soluble in water, but practically 
 insoluble in alcohol. By adding alcohol, therefore, to an aqueous 
 solution of this salt, a precipitate of ferrous sulphate in fine powder 
 is produced. 
 
 As illustrations of this method, the following exercises may be 
 made : 
 
 (i) Precipitation of Sodium Chloride by Hydrochloric Acid. 
 About half a litre of water is saturated with common salt, and the 
 brine filtered.* Into the clear solution, contained in a beaker, a 
 
 * As salt is almost equally as soluble in cold as in hot water, the liquid need 
 not be heated ; and for the same reason, should occasion arise where a hot 
 solution of salt has to be filtered, there is no necessity to use the steam-jacket 
 for the funnel. 
 
216 Preliminary Gravimetric Manipulations. 
 
 stream of gaseous hydrochloric acid is passed. The gas is gene- 
 rated by acting upon common salt with sulphuric acid (previously 
 mixed with water in the proportion of 1 1 volumes of strong acid 
 to 8 of water, and the mixture cooled) in the flask H, Fig. 36. In 
 order to arrest any sodium sulphate which might be carried as fine 
 spray along with the stream of gas, the latter is made to bubble 
 through a strong solution of hydrochloric acid contained in a three- 
 necked bottle, B, before being passed into the brine. The tube 
 which delivers the gas into the salt solution should be sufficiently 
 wide not to become stopped up with the deposited salt, and need 
 
 FIG. 36. 
 
 not dip very far into the solution. An ordinary thistle funnel, sup- 
 ported as shown in the figure, is convenient. 
 
 The generating flask should be provided with a safety funnel, 
 and the wash-bottle B with a tube which just dips beneath the 
 surface of the liquid. In this way, should there be any interruption 
 in the evolution of gas, there is no possibility of the brine being 
 sucked back into B, as, directly a reduction of pressure takes place 
 within the apparatus, air is drawn in through the safety tubes. 
 
 As the gas begins to bubble into the brine, the salt almost 
 immediately commences to separate out as a crystalline powder. 
 
Preparation of Pure Salts. 217 
 
 When a sufficient quantity of the salt has been precipitated, the 
 crystals are allowed to settle, and the bulk of the liquid decanted 
 off. The salt is then thrown upon a drainer (which may be con- 
 nected to a water-pump, as on p. 212), and after as much as possible 
 of the mother-liquor has drained away, the salt is rinsed two or 
 three times with a little strong hydrochloric acid (pure). It is then 
 removed to a porous plate to dry, after which it is heated in a 
 porcelain evaporating-dish, by means of a " rose " burner, until all 
 traces of water and hydrochloric acid are expelled. 
 
 (2) Precipitation of ferrous sulphate, FeSO 4 ,7H 2 O, by alcohol. 
 A convenient quantity of ferrous sulphate (say 100 grams) is 
 dissolved in water at a temperature of about 40 C. The salt should 
 be placed in a beaker, and a quantity of water having a temppra- 
 ture of about 40 added to it. The temperature of the mixture at 
 once falls, perhaps as much as 10, owing to the absorption of heat 
 due to solution. The beaker is then placed in a larger beaker or 
 other vessel containing water at 40, and the solution constantly 
 stirred until the whole of the salt has dissolved and a tolerably 
 saturated solution obtained at that temperature. The solution is 
 acidified with two or three drops of dilute sulphuric acid, and then 
 filtered. To prevent oxidation by undue exposure to the air, the 
 filtering operation should be accelerated by means of the filter- 
 pump * (p. 2 1 2). The perfectly clear greenish solution thus obtained 
 
 * As ordinary filter-paper, when wet, is not strong enough, when folded and 
 used in the ordinary way, to stand the pressure to which it would be exposed by 
 the use of the filter-pump, special devices must be employed for its support. 
 
 One method most suitable for such occasions as that under consideration, 
 where great rapidity is an advantage, is to place in the funnel a porcelain 
 drainer, and to lay upon it a small circular disc of ordinary filter-paper cut a 
 trifle larger (about 2 mm. all round) than the drainer. The paper is moistened 
 and carefully adjusted so as to entirely cover the drainer. [In some cases 
 where a very high degree of exhaustion is necessary, a circular piece of muslin 
 or thin calico should be laid under the paper.] 
 
 Where the filter-pump is used for precipitates that are afterwards to be 
 dried and weighed, as in the ^^___________ _ 
 
 usual quantitative determinations, 
 one of the two following plans is 
 usually adopted : Either a spe- 
 cially prepared toughened paper 
 is used (Schleicher and Schiill, 
 No. 575), or the ordinary paper 
 is supported at the apex of the 
 cone by means of a small muslin 
 or a platinum cone. When muslin 
 
 is used, a small circular piece FIG. 37. 
 
 (conveniently cut to the size of a 
 
 penny piece) is folded as a filter is folded, and placed in the funnel. The 
 
 folded filter-paper is then introduced, being carefully pushed into its position 
 
 with its apex within the muslin cone. The filter-paper must fit the funnel 
 
218 Preliminary Gravimetric Manipulations. 
 
 is poured into a quantity of alcohol (methylated spirit) about equal 
 to the volume of the ferrous sulphate solution. The first effect will 
 be somewhat unexpected, for, instead of any precipitation of crystals, 
 a green liquid separates out at the bottom of the vessel. On stir- 
 ring the mixture, however, for a moment or two, this liquid disap- 
 pears, and the ferrous sulphate is precipitated as a fine crystalline 
 powder. The liquor is then decanted off, the crystals drained upon 
 a drainer and rinsed with alcohol. It must then be dried as 
 quickly as possible upon a porous plate, or by gentle pressure 
 between folds of blotting-paper. 
 
 Precipitation by Double Decomposition. The process 
 of preparing a pure compound by this method is practically identical 
 with that of conducting any quantitative precipitation. It will be 
 obvious that, in order to obtain a pure product, there must 'be nothing 
 present which will give a precipitate with the reagent to be used, 
 other than the compound intended. For instance, suppose it is 
 desired to prepare pure barium carbonate by precipitation from a 
 solution of barium chloride by means of ammonium carbonate ; the 
 barium chloride must itself be a purified salt, and must be free 
 from any metallic compounds which would give a precipitate with 
 ammonium carbonate. The precipitated compound is washed 
 with warm water until perfectly free from any soluble salts, and is 
 finally dried in a steam-oven if it can stand that temperature without 
 decomposition, otherwise it is spread upon a porous plate. 
 
 When it is desired to crystallise a salt with the object of obtain- 
 ing large and well-defined crystals, it is necessary to modify the 
 methods so as to cause the process to take place very slowly, and, 
 as nearly as possible, with uniform regularity. A nearly saturated 
 solution of the salt is made at a temperature about 30 C, usin^ 
 a considerable volume of water. The solution is placed in a beaker, 
 which is then stood in a larger beaker of warm water in order to 
 render the cooling operation slower, and covered with a clock- 
 glass. When the solution is cold, there will be a small crop of 
 
 exactly, and not leave any air-spaces down the sides. If the sides of the funnel 
 do not enclose an angle of exactly 60, the paper must be folded so as to make 
 it exactly fit the funnel, or the funnel must be rejected for one of the right 
 shape. A platinum cone is made by cutting from a piece of foil a segment of 
 a circle whose radius is 15 mm. (or 3 inch). The segment should be nearly 
 three-quarters of the entire circle, as shown in Fig. 37, so that when it is bent 
 together so as to form the cone, there will be an overlap extending almost one- 
 half of the way round. A few fine holes may be pierced in the metal near to 
 the apex of the cone, by laying the flat piece of foil upon a cork, and pricking 
 it with a fine needle. The cone is then put into the funnel with the paper just 
 as the muslin cone. Sometimes a cone of the toughened paper above men- 
 tioned is used ; this is less advantageous, as it practically introduces a double 
 filtering medium, and retards the process considerably. 
 
Preparation of Pure Salts. 219 
 
 crystals on the bottom of the beaker. The liquid is then poured 
 into another beaker, and a few of the best formed of the crystals 
 are picked out. (If this first crystallisation has resulted in the 
 deposition of a solid cake or crust, and no isolated crystals, the 
 liquid must be returned to the beaker, and the whole warmed up 
 again, with the addition of a little more water, and the solution set 
 to cool as before.) These are dropped into the cold mother-liquor 
 so as to lie isolated from each other, and the vessel placed where 
 its temperature will remain as nearly constant as possible. The 
 rate of evaporation of the liquid may be checked, if the salt deposits 
 too quickly, by partially covering the beaker. If circumstances 
 make it necessary that the operation be carried on in a room liable 
 to considerable changes of temperature, it is a good plan to keep 
 the beaker in a larger vessel of water all the time the process is 
 going on. In order that the crystals may become equally developed 
 in every direction, so as to approach as nearly as possible the 
 perfect form for the particular compound being crystallised, they 
 must be turned at regular intervals, say every day, or every second 
 day, according to the rate of growth, so as to rest upon a different 
 face. The more slowly the crystals can be grown, other things 
 being equal, the more perfect in form will they be. 
 
 In cases where the solution cannot be left exposed to the air for 
 so long without undergoing some chemical change (for instance, 
 ferrous sulphate would become oxidised), the salt may be slowly 
 crystallised, if it is one which is insoluble in alcohol, by allowing 
 that liquid to very gradually diffuse into the aqueous solution. To 
 take the case of ferrous sulphate as an example. A strong cold 
 aqueous solution of the salt is placed in a large wide-mouth 
 stoppered bottle, which is less than half filled by it. A layer 
 of water about 2i c.c. (i inch) deep is floated upon the top of 
 the solution, and "upon the top of the water a quantity of alcohol 
 is placed, equal to the volume of liquid already present. The 
 stopper is inserted, and the bottle placed where it can remain for 
 some weeks undisturbed. By slow degrees the alcohol diffuses 
 into the aqueous solution, and thereby causes the gradual deposition 
 of crystals of the salt. 
 
 The object of the shallow layer of water dividing the alcohol 
 from the saline solution is to prevent the precipitation of the salt 
 at the surface of contact, which would be the case if the alcohol 
 floated immediately upon the solution. 
 
 A good plan for floating one liquid upon another without causing 
 admixture, is the following : 
 
 The heaviest solution is first poured into the vessel by means of 
 a funnel with a long stem, so as not to splash the sides of the 
 vessel. A thin circular disc of cork, as large as will pass the 
 mouth of the bottle, is then lowered on to the surface of the solution 
 by means of a fine thread attached to the centre of the cork with a 
 little wire hook. 
 
 By means of a stoppered funnel, attached to a long piece of 
 
220 
 
 Preliminary Gravimetric Manipulations. 
 
 glass tube drawn to a narrow jet at the lower end, the lighter liquid 
 (in the above instance, the ivater) is allowed slowly to stream on to 
 the cork, the funnel being gradually raised as the cork ascends, so 
 as to keep the jet almost close to the latter, as shown in Fig. 38. 
 
 When a sufficient layer of the second liquid has been floated 
 upon the first, the still less dense third liquid (the alcohol in the 
 
 example) is introduced in the same way. As the layer of the added 
 liquid gradually gets deeper, and the cork float is further removed 
 from the line of contact, the liquid may be run in more quickly, 
 without fear of disturbing the solution below. When all has been 
 added, the cork is withdrawn by means of the thread, and the bottle 
 stoppered. (The cork should be put into hot water to soak, if it is 
 intended to use it again for a similar purpose.) 
 
SECTION II. 
 TYPICAL GRAVIMETRIC ESTIMATIONS OF METALS. 
 
 IN order to make a gravimetric estimation of a metal, the latter is 
 either obtained from the compound in the elemental form, and the 
 metal itself weighed ; or it is caused to enter into some convenient 
 combination, yielding a compound of known and definite com- 
 position, and from the weight of the compound so produced the 
 quantity of metal is ascertained by calculation. This latter mode 
 is of more general application than the first, as many of the metals 
 are not easily reduced from their compounds. Of late years, how- 
 ever, electrolytic methods of analysis have been greatly developed 
 and perfected, so that a large number of metals, which formerly 
 could only be determined in combination, are now deposited in the 
 form of metal, and so weighed. Typical examples of this process 
 of analysis will be given in a separate section. 
 
 The forms of combination which are most frequently had recourse 
 to, and which are applicable to various metals, are practically three 
 in number, namely, the oxide, sulphide, and sulphate. 
 
 When weighed in the form of an oxide, the metal is first pre- 
 cipitated either as a hydroxide (as, for example, in the case of 
 Al, Fe, Cr, Ni, Cu, etc.) or as a carbonate (as in the case of Bi, Mn, 
 Zn, etc.). In either case the precipitated compound is converted 
 by heat into the oxide, which is then weighed. 
 
 When estimated as a sulphide, the metal is precipitated in this 
 form of combination either by sulphuretted hydrogen or ammonium 
 sulphide. Examples of metals which may be estimated in this way 
 are seen in mercury, cadmium, antimony, and arsenic. Metals 
 which are to be weighed as sulphates, are either precipitated 
 from solution in this form, as in the case of lead, barium, and 
 strontium, or the sulphate is obtained by evaporating a suitable 
 compound of the metal (e.g. the chloride) with sulphuric acid, 
 whereby it becomes converted into the sulphate. This method is 
 sometimes employed for the estimation of the alkali metals. 
 
222 Typical Gravimetric Estimations. 
 
 Besides these methods, each of which is applicable for the 
 estimation of several metals, there are special forms of combination 
 which are available for special cases. Thus, for example, silver is 
 precipitated and weighed as chloride (or bromide) ; magnesium 
 is precipitated as ammonium magnesium phosphate, which is then 
 converted into magnesium pyrophosphate by the action of heat, 
 and weighed in this form ; and so on. 
 
 Descriptions will be given in this section of a selected number of 
 typical examples, illustrating the most important methods of gravi- 
 metric estimations, and forming a suitable course for practice in 
 the various analytical operations. 
 
 Estimation of Aluminium. 
 
 Epitome of Process. The metal is precipitated from the 
 solution of its compounds as aluminium hydroxide by means of 
 ammonia.* The hydroxide is afterwards converted by heat into 
 the oxide, A1 2 O 3 , in which form the metal is weighed. 
 
 Employ potassium alum, K 2 SO 4 ,Alo(SO 4 ) 3 ,24H 2 O, taking about 
 2 grams.f 
 
 A stoppered weighing-bottle containing powdered alum, which 
 has been purified by re-crystallisation, is first weighed ; after which 
 about 2 grams of the salt are carefully transferred without loss to 
 a beaker, % conveniently of about 550 c.c. (nearly 20 ozs.) capacity. 
 (The stopper should be removed from the bottle while the latter 
 is held immediately over the beaker. A sufficient quantity of the 
 salt is then tipped out and the stopper again replaced, care being 
 taken that no particles are left adhering to the lip of the bottle, 
 by giving the latter a few taps with the finger.) The beaker is 
 
 * The precipitation of metallic hydroxides by alkalies is not complete in the 
 presence of certain organic substances, such as tartaric or citric acids, sugar, 
 etc. In cases where such non-volatile organic matter is present, these must 
 first be decomposed by adding to the solution sodium carbonate and potassium 
 nitrate, and evaporating to rdryness in a platinum dish, upon a steam-bath. 
 The residue is then fused, and afterwards extracted with dilute hydrochloric 
 acid. The solution is then filtered, and the precipitation conducted as above 
 described. 
 
 t See note on p. 201. In this case, the formula weight of alum being 948, 
 while that of A1 2 O 3 is only 102, the final product that is to be weighed amounts 
 to less than one-ninth of the weight of the original salt. Hence a fair quantity 
 must be operated upon. Two grams of the salt will only furnish about o'2i 
 grams of Al 2 Os. When estimating the (804) in alum the case is different, for 
 the weight of the BaSO 4 yielded is nearly equal to the original weight of the 
 alum itself. 
 
 J In special cases, where exceptional accuracy is required, a vessel of 
 platinum, nickel, or porcelain is substituted for the beaker, as alkalies exert ;i 
 slight solvent action upon the glass. With Jena glass this action is extremely 
 slight. 
 
Estimation of Aluminium. 
 
 223 
 
 immediately covered with a clock-glass. The bottle is then weighed 
 again, and the difference gives the weight of alum taken. 
 
 The salt is then dissolved in such a quantity of water that the 
 beaker is about one-fourth part filled with liquid,* and the process 
 of solution is aided by gently warming the mixture. 
 
 For this purpose the beaker may be supported and heated by 
 either of the methods shown on p. 209, Figs. 30 and 31 ; or it 
 may be placed upon an iron plate heated by a Bunsen, as in Fig. 39. 
 
 FIG. 39. 
 
 This is a specially convenient arrangement when several operations 
 are being conducted simultaneously. The heat can be moderated 
 by placing the vessel nearer to, or farther from the flame. A piece 
 of asbestos cloth laid upon the top of the plate prevents any risk of 
 the fracture of glass vessels by placing them suddenly upon the 
 rough iron. It is important that the various pieces of apparatus 
 used should bear a suitable relation to each other as regards their 
 size. For example, the clock-glass employed as a cover should not 
 be much larger than the mouth of the beaker ; and when a flame 
 is used directly under the vessel, care must be taken that it does 
 not extend beyond the bottom of the beaker. Such an arrangement 
 
 * In all the following examples, the size of the beaker, and the quantity of 
 water employed for the solution of the salt, may be the same as those here 
 described, unless special directions to the contrary are given. It ought not to 
 be necessary to repeat at this stage, what was stated on pp. 3 and 116, that 
 in all analytical operations distilled water must be exclusively used. 
 
224' 
 
 Typical Gravimetric Estimations. 
 
 FIG. 40. 
 
 of apparatus as shown in Fig. 40, and which is not unfrequently 
 seen in the laboratory, is almost sure to end in disaster. The heat 
 from the lamp either fractures the clock-glass, fragments of which 
 fall into the solution, or the beaker itself 
 becomes unduly heated upon its sides, so 
 that the first movement of the liquid 
 causes it to crack. The beaker should 
 be kept covered by the clock-glass as 
 much as possible throughout the entire 
 operation. 
 
 A quantity of ammonium chloride 
 solution is added,* equal to about a fourth 
 of the volume of the liquid already present, 
 and then ammonia in the least possible 
 excess beyond that which is required for 
 complete precipitation. 
 
 The mixture is stirred with a glass rod 
 with rounded ends, the rod being of such 
 a length that it will just stand in the 
 beaker while the cover is upon it. A short 
 piece of clean narrow caoutchouc tube 
 should be previously fitted upon the end 
 of the rod. This not only prevents the beaker from becoming 
 scratched by the rod, but it is useful in the subsequent operation of 
 removing particles of the precipitate which may adhere to the sides 
 of the beaker. 
 
 The liquid is then boiled until the greater part of the ammonia 
 is expelled, and the issuing steam possesses only a slight smell of 
 the reagent. 
 
 The precipitate is then allowed to settle ; the clock-glass is re- 
 moved, and the under side rinsed into the beaker by means of a 
 fine jet of hot water from a wash-bottle, in case any particles may 
 have been thrown up during the operation of boiling. The clear 
 liquid is then decanted off through a filter,t the weight of whose ash 
 is known (see p. 206). 
 
 When as much of the liquid as possible has been poured off 
 
 * In the presence of ammonium chloride, aluminium hydroxide is insoluble 
 in ammonia. 
 
 t The operation of filtering must be carried out with all the precautions 
 against loss described in the introduction to qualitative analysis, p. 2. It 
 sometimes happens, owing to the particular shape of the lip of a beaker, that, 
 even when the liquid is carefully poured down against a glass rod, a little of it 
 creeps back under the edge. This may be prevented by slightly greasing the 
 under side of the lip by means of a little touch of vaseline or resin cerate upon 
 the finger. 
 
Estimation of Aluminium. 
 
 225 
 
 without disturbing the precipitate, the beaker is half filled up with 
 boiling water, the mixture well stirred, and, after being allowed to 
 settle, is decanted as before through the same filter. This process 
 is repeated a second time, after which the precipitate itself is trans- 
 ferred to the filter. Every trace is removed from the sides of the 
 beaker by the aid of the glass rod, and the beaker is finally rinsed 
 into the funnel by means of a jet of hot water from the wash-bottle 
 in the manner shown in Fig. 41, the movable nozzle of the bottle 
 being directed round the beaker by the fore-finger. 
 
 The final washing of the precipitate is now made upon the filter 
 by means of the jet of hot water from the wash-bottle, washing 
 downwards from the 
 upper edge of the 
 paper, but never quite 
 filling the paper cone 
 with water, and allow- 
 ing each washing to 
 drain completely 
 through before adding 
 fresh water.* The 
 operation is continued 
 until the precipitate is 
 entirely freed from all 
 soluble salts, which in 
 this case consist of 
 sulphates of potassium 
 and ammonium. To 
 ascertain when this is 
 the case, a few drops 
 of the filtrate are col- 
 lected in a test-tube, 
 
 and a little barium chloride added. If no precipitate forms in a few 
 minutes after warming the mixture, the washing may be considered 
 complete. The last washing should aim at collecting the precipitate 
 together into the cone of the filter. 
 
 The filter and precipitate are next dried in a steam- oven. For 
 this purpose, the funnel containing the filter is covered with a piece 
 
 * When a filter-pump is being used, however, this rule does not apply. In 
 this case it is important not to let the filter run empty, but to constantly add 
 liquid before the former quantity has completely run through, otherwise 
 channels or fissures are liable to be formed in the solid material upon the filter 
 through which subsequent wash-waters run without properly washing the 
 precipitate. 
 
 Q 
 
 FIG. 41. 
 
226 
 
 Typical Gravimetric Estimations. 
 
 of filter-paper, and either stood up in the corner of the oven, or 
 placed in a conical tin support. In order to securely cover the 
 funnel, a common circular filter-paper, somewhat larger in diameter 
 than the top of the funnel, is held upon the latter with the left hand, 
 
 while the projecting 
 paper is folded down 
 all round with the other 
 hand in the manner 
 shown in Fig. 42. 
 
 While the precipi- 
 tate is drying, a plati- 
 num crucible is heated 
 to redness for a few 
 moments, cooled in 
 the desiccator, and 
 then weighed. The 
 dried filter containing 
 the aluminium hydrox- 
 ide is then removed 
 from the funnel, care- 
 fully folded up, and placed in the weighed crucible. The crucible, 
 supported upon a clean pipeclay triangle, is first very gently 
 heated, the lid being placed slightly on one side in order to allow 
 the gaseous products of the combustion of the paper to escape, and 
 also to admit the necessary supply of air for the combustion. When 
 the active combustion of the paper has subsided, the temperature 
 is gradually raised, and when it has reached a red heat, the lid of 
 the crucible may be removed.* The vessel is maintained at a bright 
 * It is very necessary that the tips of the crucible-tongs should be perfectly 
 
 FIG. 42. 
 
 
 / ' Wgp 
 
 clean, and after being in contact with the hot crucible or lid, the tongs should 
 never be placed on the table -with the tips downward. The first impulse of almost 
 
Estimation of Aluminium. 
 
 227 
 
 red heat by means of a good Bunsen flame or a blowpipe for 
 about ten minutes. The lid is then replaced, and the crucible 
 
 every one is to hold and use crucible-tongs in the manner shown in Fig. 43, 
 with the almost inevitable result that when putting them down, the hot points 
 
 
 FIG. 44 . 
 
 FIG. 45. 
 
 come into contact with the 
 wooden table (Fig. 44), and 
 are liable to pick up matter 
 which afterwards may be trans- 
 ferred to the crucible. Tongs 
 should always be held in ex- 
 actly the reverse way, with 
 the tips pointing in the direc- 
 tion of the back of the hand, 
 as in Fig. 45. They are then 
 naturally placed upon the table 
 with the tips pointing upwards 
 (Fig. 46). The habit of mani- 
 pulating tongs in this way 
 should be cultivated from the 
 very beginning. 
 
 FIG. 46. 
 
228 
 
 Typical Gravimetric Estimations. 
 
 transferred to the desiccator to cool, after which it is weighed. 
 It is reheated for a similar time and again weighed, the operation 
 being repeated until the weight is practically constant. 
 
 [When a filter-pump is employed for accelerating the filtration 
 of the aluminium hydroxide, the precipitate may be partially dried 
 by allowing the pump to continue in operation for fifteen minutes 
 or so after all the liquid has passed through. The filter may then 
 be folded up, and at once transferred to the crucible, in which it is 
 completely dried by the cautious application of heat. To avoid 
 spitting, the crucible should be supported rather obliquely in the 
 triangle, and a small flame applied to the extreme upper part, so 
 that the heat may be conducted very gently to the wet materials.] 
 
 From the data obtained by the analysis, the percentage of 
 aluminium is calculated in the following manner : 
 
 Estimation of Aluminium in Potash Alum* Iv.SO,,- 
 A1 2 (S0 4 ) 3 ,24H 2 0.- 
 
 Weighing-bottle and salt (first weight) ... 
 (second weight) 
 
 Weight of salt taken 
 
 Weight of crucible + filter-ash + AUO 3 ... 
 
 Weight of crucible alone 
 
 Weight of ash* 
 
 16-9485 grams. 
 14-8870 ,. 
 
 2-0615 
 
 25-6857 
 
 25 '4635 
 
 O'OOO2 
 25-4637 
 
 Weight of A1 2 O 3 
 
 | Formula- weight of A1 2 O 3 = 102 
 
 Ali = 54 
 therefore weight of aluminium \ 54 x 0*2220 
 
 in o'2220gram Al 2 O s t 
 
 102 
 
 = 0-11753 
 
 0*1 1753 x 100 
 hence the percentage of aluminium found = ^ = 570 
 
 * If the filter employed is that mentioned on p. 193, having an ash equal to 
 0-00017, it may be ignored altogether, as it does not affect the second place of 
 decimals in the final result. 
 
 f The ratio 102 : 54 is equal to i : 0-5294, therefore the weight of Al in any 
 given weight of AlgOa is at once obtained by multiplying the latter by the 
 factor 0-5294. Thus, in the above example, 0-2220x0-5294 0-117526. At 
 the commencement of each estimation in this section, the factor will be found, 
 which, when multiplied by the weight of the compound that is actually weighed, 
 will give the weight of metal it contains. The use oi factors, especially by the 
 student, is always attended with the danger of obscuring the rationale of the 
 calculation, and reducing it to a mechanical or "rule of thumb" operation. 
 To prevent this as far as possible, the ratio of the formula-weights of the com- 
 pound weighed, and the metal to be estimated are given, so as to keep present 
 to the mind the true significance of the factor. 
 
Estimation of Chromium. 229 
 
 r Formula- weight of K 2 SO 4 ,A1 2 (SO 4 ) S ,24H 2 O = 948 
 Formula-weight of A1 2 = 54 
 hence the theoretical percentage of \ 54 x 100 _ 
 aluminium in alum /"" 948 
 
 Chromium. 
 
 Epitome of Process. The chromium is precipitated from 
 solutions of chromic compounds in the form of chromic hydroxide 
 by means of ammonia. The hydroxide is converted by heat into 
 chromic oxide, Cr 2 O 3 , in which form the element is weighed. 
 
 Chromates are first reduced to the " chromic " state by means 
 of sulphur dioxide. Factor 
 
 (Cr,O 3 ) 152 : (Cr.,) 104 = i : 0-68421 
 
 Estimation of Chromium in Chrome Alum, K 2 SO 4 ,- 
 Cr 2 (SO 4 ) 3 ,24H 2 O. Employ about 1*5 gram. The process is con- 
 ducted exactly as in the case of aluminium, except that the addition 
 of ammonium chloride may be omitted. The filter may either be 
 burnt in the crucible along with the precipitate, or it may be in- 
 cinerated separately, as described below in the case of iron. 
 Chromium oxide is not reduced during the process. 
 
 Formula weight of chrome alum = 998. Theoretical percentage 
 of Cr = 10-42. 
 
 Estimation of Chromium in Potassium Bichromate, 
 K 2 Cr 2 O 7 . Take about 0-75 gram. The purified salt is weighed out 
 into a beaker (see notes, pp. 222, 223), and dissolved in water without 
 being warmed. A gentle stream of sulphur dioxide * is passed 
 through the liquid until the chromium is wholly reduced to the 
 " chromic " state. The reduction is complete when the solution, 
 on being gently moved against the sides of the beaker, shows no 
 trace of the yellow colour, but has a clear green tint. The delivery 
 tube is rinsed into the beaker with water from the wash-bottle, and 
 removed. The' solution is then warmed to expel the excess of 
 sulphurous acid, and is afterwards precipitated with ammonia as 
 in the above example. 
 
 Formula-weight of potassium dichromate = 294. Theoretical 
 percentage of Cr = 35 -37. 
 
 * The sulphur dioxide is most conveniently derived from a "syphon" of 
 the liquefied gas. In the absence of such a supply, the gas must be generated 
 from sulphuric acid and copper. In this case the sulphur dioxide must be 
 'washed by being passed through water before being delivered into the solution 
 of the chromate. 
 
230 Typical Gravimetric Estimations. 
 
 Iron. 
 
 Epitome of Process. The iron is precipitated from solutions 
 of ferric compounds as ferric hydroxide by means of ammonia. 
 The hydroxide is converted by heat into ferric oxide, Fe 2 O 3 , in 
 which form the metal is weighed. The filter is burnt apart from 
 the precipitate. The iron \\\ ferrous compounds is first oxidised 
 into the " ferric " state. Factor 
 
 (Fe 2 O 3 ) 160 : (Fe 2 ) 112 = i : 0700 
 
 Estimation of Iron in Ferrous Ammonium Sulphate, 
 (NH 4 ) 2 SO 4 ,FeSO 4 ,6H 2 O. Employ about 1*5 gram. The purified 
 salt is weighed out into a beaker (notes, pp. 222, 223), and dissolved 
 in water with the addition of one or two drops of dilute sulphuric 
 acid. The solution is warmed, and a little strong nitric acid 
 added, sufficient to oxidise the whole of the iron. The mixture is 
 heated nearly to the boiling-point for a short time in a covered 
 beaker. Each drop of nitric acid as it is added produces a brown 
 coloration, owing to the absorption of the disengaged nitric oxide 
 by the ferrous salt ; but when the oxidation is complete, the addi- 
 tion of a drop of the acid produces no visible effect. By adding, 
 therefore, a few drops of nitric acid after its addition has ceased to 
 cause any coloration, sufficient will have been introduced to com- 
 plete the oxidation. As a confirmation, the smallest drop of the 
 solution is placed upon a piece of white porcelain (such as a 
 crucible lid) by means of a glass rod, and the drop is touched with 
 another glass rod which has been dipped in a solution of potassium 
 ferricyanide. If the ferrous iron has been completely oxidised, no 
 blue coloration will result. [The ferricyanide solution must be 
 freshly made, by dissolving a small crystal of the salt in water.] 
 
 The ferric hydroxide is then precipitated by the addition of a 
 slight excess of ammonia, and the mixture boiled until the steam 
 scarcely smells of ammonia. The precipitate is washed with hot 
 water by decantation two or three times, and finally washed upon 
 the filter until the wash-water is entirely free from sulphates (com- 
 pare Aluminium, p. 225). 
 
 The precipitate is then thoroughly dried in the steam-oven. 
 
 When perfectly dry, the filter is withdrawn from the funnel, and 
 as much of the precipitate as possible is detached from the paper 
 by gently squeezing the cone together, and transferred to a platinum 
 crucible, which is placed upon a sheet of glazed paper (see p. 207). 
 By flattening the paper cone and gently rubbing one side against 
 the other, the remaining adhering particles may be detached. The 
 paper is then folded in the manner described on p. 207, bound up 
 
Estimation of Calcium. 23! 
 
 in a platinum wire, and incinerated as completely as possible. The 
 ash is then shaken off the wire into the crucible, and any particles 
 which may have fallen upon the glazed paper are also carefully 
 swept into the crucible by means of a feather. The crucible is 
 then heated for about ten minutes to a bright red heat by means of 
 a Bunsen flame, after which it is placed to cool in the desiccator 
 and weighed ; the heating and weighing being repeated until no 
 further loss of weight results. 
 
 The minute quantity of ferric oxide which is left upon the filter, 
 and which becomes reduced to metallic iron during incineration, is 
 again oxidised during the process of heating ; and, although it 
 may not pass into the same oxide, Fe 2 O 3 , but probably into Fe 3 O 4 , 
 the total quantity is so small, and the ratio Fe 3 O 4 : (Fe 3 ) is so 
 nearly equal to Fe 2 O 3 : (Fe 2 ), that the difference is quite outside the 
 experimental error. 
 
 Formula weight of (NH 4 ) 2 SO 4 ,FeSO 4) 6H< f O = 392. Theoreti- 
 cal percentage of Fe = 14*28. 
 
 Calcium. 
 
 Epitome of Process. The calcium is precipitated in the form 
 of calcium oxalate, CaC 2 O 4 , by means of ammonium oxalate in the 
 presence of ammonia. The calcium oxalate is then converted by 
 a gentle heat into calcium carbonate, CaCO 3 , in which form the 
 metal is weighed ; or it is strongly heated until it is entirely 
 changed into calcium oxide, and weighed in this form. 
 
 As an exercise in manipulation, the oxalate may be converted 
 first into the carbonate and weighed, and afterwards, by the appli- 
 cation of a stronger heat, be changed into the oxide and weighed 
 again. Factors 
 
 (CaCO 3 ) 100 : (Ca) 40 = i : 0-4000 
 (CaO) 56 : (Ca) 40 = i : 071428 
 
 Estimation of Calcium in Calcium Carbonate, CaCO 3 . 
 Employ about 0-5 to 075 gram. Pure precipitated calcium carbonate 
 is weighed out into a beaker, and dissolved in a very little dilute 
 hydrochloric acid. The clock-glass cover should be pushed a little 
 to one side, and the diluted acid poured gradually down the side of 
 the beaker. When the carbonate is completely dissolved, the cover 
 should be rinsed into the beaker by means of the wash-bottle, and 
 water added until about the usual volume is present. Ammonia is 
 then added until the solution smells distinctly of the reagent, and 
 the liquid is heated to boiling. The calcium oxalate is then pre- 
 cipitated by the addition of a slight excess of a warm strong solution 
 of ammonium oxalate, to which a little ammonia has been added. 
 
2 3 2 
 
 Typical Gravimetric Estimations. 
 
 The mixture is boiled for a few minutes, and allowed to settle.* 
 The clear liquid is then decanted off through a filter, without dis- 
 turbing the precipitate. It is washed three or four times by decan- 
 tation with boiling water, allowing it to settle thoroughly each time. 
 The small quantity of the precipitate which is inevitably conveyed 
 to the filter with the washings, will so far fill up the pores of the 
 paper that, when, after the third or fourth wash, the precipitate itself 
 is transferred to the filter, the filtrate will be perfectly clear. The 
 precipitate is washed with warm water until the wash-water is free 
 from ammonium chloride, as indicated by the absence of any 
 milkiness on the addition of silver nitrate after acidifying with 
 nitric acid. 
 
 When washing this precipitate, a fine jet of water should be 
 employed, and it should be made to impinge first upon the glass 
 funnel above the edge of the paper, and then gradually directed 
 down towards the precipitate. 
 
 After being dried in the steam-oven, the precipitate is trans- 
 ferred as completely as possible to a platinum crucible. The filter 
 
 is incinerated in a plati- 
 num wire coil, and the ash 
 deposited in the crucible. 
 The calcium oxalate is 
 then converted by a gentle 
 heat into the carbonate, 
 the temperature being 
 carefully regulated so that 
 it never reaches a visible 
 redness even at the bottom 
 of the crucible. A con- 
 venient plan, in order to 
 avoid unduly heating any 
 spot of the crucible, is to 
 place it upon a triangle 
 arranged above a piece of 
 wire gauze upon a tripod, as shown in Fig. 47. The gauze is 
 heated to redness by means of a rose burner, the heat being thus 
 radiated to the crucible, which must be heated in this manner for 
 
 * Calcium oxalate is a precipitate requiring some special care in its mani- 
 pulation. Under ordinary circumstances it is somewhat slow to settle, showing 
 an inclination to creep up the wet sides of the beaker, and it also passes through 
 the filter very easily. By conducting the precipitation in the manner here 
 described, the compound is obtained in a more coherent form, and with a little 
 care will be perfectly retained by the filter. 
 
 FIG. 47- 
 
Estimation of Calcium. 233 
 
 about fifteen minutes ; it is then cooled in the desiccator and 
 weighed. 
 
 In case any of the calcium carbonate has become converted 
 into the oxide in spite of care in the heating, the residue is 
 moistened with a few drops of a strong solution of ammonium 
 carbonate, which is then evaporated to dryness by placing the 
 crucible in the steam-oven. When dry, it is reheated for a few 
 minutes over the heated gauze, and again weighed. 
 
 If this process results in any increase in the weight, it shows 
 that in the first heating some of the carbonate had been decom- 
 posed ; the series of operations should, therefore, be repeated until 
 the weight remains constant. 
 
 In this particular instance the form of combination in which 
 the metal to be estimated is weighed, happens to be the same as 
 that in which it exists in the original compound, namely, the 
 carbonate; hence the weight of the product obtained will be the 
 s*ame as that of the substance taken for analysis. 
 
 The calcium carbonate is now converted entirely into calcium 
 oxide by heating the crucible to a strong red heat by means of a 
 Bunsen flame for about ten minutes, and finishing off with a blow- 
 pipe flame. The crucible is placed in the desiccator to cool, and 
 is then weighed. It is afterwards reheated with the blowpipe for 
 a few minutes and weighed again, the process being repeated until 
 no further loss of weight results. 
 
 Magnesium. 
 
 Epitome of Process. The magnesium is precipitated by 
 sodium phosphate, in the presence of ammonium chloride and 
 ammonia, as ammonium magnesium phosphate, NH 4 MgPO 4 ,6H 2 O. 
 This is afterwards converted by heat into magnesium pyrophosphate, 
 Mg 2 P 2 O 7 , in which form the metal is weighed. Factor 
 
 (Mg 2 P 2 O 7 ) 222 : (Mg 2 ) 48 = i : 0-2162 
 Estimation of Magnesium in Magnesium Sulphate, 
 
 MgSO 4 ,7H 2 O. Take about 75 gram. The recrystallised salt is 
 weighed out into a beaker, and dissolved in a rather smaller bulk 
 of water than usual (about 60 to 80 c.c.). A quantity of ammonium 
 chloride solution is added, equal to about one-half the volume of 
 the solution of magnesium sulphate, and then ammonia in mode- 
 rate excess.* To this is added a solution of hydrogen disodium 
 
 * Should the addition of ammonia cause precipitation, more ammonium 
 chloride must be introduced until this precipitate is redissolved. Excess of 
 ammonia, over and above that which is required for the formation of the 
 ammonium magnesium phosphate, is necessary to ensure complete precipitation, 
 as the phosphate is slightly soluble in water, but insoluble in the presence of 
 free ammonia. 
 
234 Typical Gravimetric Estimations. 
 
 phosphate, in quantity in excess of that required for complete precipi- 
 tation of the ammonium magnesium phosphate. The mixture must 
 be well stirred with a rubber-tipped glass rod. The precipitate is 
 a crystalline compound, which requires several hours for its com- 
 plete separation. Its formation is accelerated by stirring or shaking, 
 but care is necessary to avoid pressing or rubbing the glass of the 
 stirring-rod against the sides of the beaker, as this would cause the 
 deposition of minute crystals upon the surface of the vessel, which 
 are very difficult to detach. The covered beaker should be put 
 aside for a few hours.* 
 
 The liquid is filtered, and the precipitate washed with dilute 
 ammonia (water 3 parts by volume, strong ammonia sp. gr. "880 
 i part) until the filtrate is perfectly free from chloride, as indicated 
 by the absence of any milkiness when tested with silver nitrate 
 after acidification with nitric acid. 
 
 The precipitate is then dried in the steam-oven. 
 
 The dry precipitate is transferred to a platinum crucible, and 
 the filter incinerated in a platinum coil. Care should be taken not 
 to overheat the filter during this operation, whereby the adhering pre- 
 cipitate becomes fused, and thereby renders the complete combustion 
 of the paper extremely difficult.! After the ash has been placed in 
 the crucible, the latter is gently heated by means of a small Bunsen 
 flame. The crucible must be covered during this operation, and 
 the heat must be applied cautiously, as, during the conversion of 
 the compound into the pyrophosphate, ammonia and water are 
 eliminated, and, in escaping too rapidly, these would cause loss of 
 the precipitate. When the smell of ammonia can no longer be 
 perceived, the temperature is gradually raised to a bright red heat, 
 finishing off with a blowpipe flame. 
 
 The crucible is then removed to the desiccator to cool, and 
 afterwards weighed, the heating process being repeated until the 
 weight is constant. 
 
 Formula weight of MgSO 4 ,7H 2 O = 246. Theoretical percentage 
 of Mg = 9756. 
 
 * If the entire operation be carried out in a 500 c.c. wide-mouth stoppered 
 bottle, and the mixture be briskly shaken for a few minutes, the precipitation 
 will be complete in a very short time. The disadvantage of this plan is the 
 greater difficulty of transferring the precipitate to the filter. When the stopper 
 is removed, it must be carefully rinsed with water. 
 
 f The incineration may be made by the alternative method described for 
 magnesium arsenate, p. 254. 
 
Estimation of Copper, 235 
 
 Copper. 
 
 I. As COPPER OXIDE. 
 
 Epitome of Process.* The copper is precipitated from solu- 
 tions of its salts in the form of copper hydroxide by means of potas- 
 sium hydroxide. The copper hydroxide is subsequently converted 
 by heat into copper oxide, in which form the metal is weighed. 
 The filter is burnt apart from the oxide, and treated in a special 
 manner. Factor 
 
 (CuO) 79'3 : (Cu) 63-3 = i : 079823 
 
 Estimation of Copper in Copper Sulphate, CuSO 4 ,5H 2 O. 
 Employ about i gram. The recrystallised salt is weighed out 
 into a beaker,f and dissolved in water with the aid of heat. The 
 solution is raised to the boiling-point, and potassium hydroxide 
 added in small quantities with constant stirring, until precipitation 
 is complete. When this is the case, the liquid will appear colour- 
 less as the precipitate settles, and a drop of the clear liquid placed 
 on turmeric paper by means of a glass rod, will show that the 
 solution is alkaline. 
 
 The mixture is gently boiled for a few minutes, and then allowed 
 to settle. 
 
 The precipitate is washed two or three times by decantation 
 with boiling water, and then transferred to the filter, when the 
 washing is continued with hot water from the wash-bottle until the 
 filtrate is perfectly free from potassium sulphate, as indicated by 
 barium chloride. The funnel is then covered, and placed to dry in 
 the steam-oven. 
 
 The dried precipitate is transferred to a platinum crucible, and 
 as much as possible of that which adheres to the paper is detached 
 by the process of gently rubbing the sides of the filter together. 
 
 The paper is folded up for incineration, so that that portion 
 which is soiled with the precipitate shall be in the interior of the 
 roll, and shall not come in contact with the platinum wire. 
 
 By giving the crucible a gentle tap, the precipitate is made to 
 take up a position rather to one side, and so leave a clear space, 
 upon which the ash may be deposited. 
 
 After the ash has been dropped into the crucible, a single drop^ 
 
 * See note *, p. 222, with reference to the presence of organic matter. In 
 the case of copper hydroxide, precipitation is also incomplete if much alkali 
 nitrate is present. In such cases the copper may be estimated as sulphide 
 (p. 236). 
 
 f See note J, p. 222, which applies still more in the case of caustic potash 
 or soda. 
 
236 Typical Gravimetric Estimations. 
 
 of strong nitric acid is allowed to fall upon it from a pipette. Any 
 particles of reduced copper are thus converted into the nitrate, 
 which, in the subsequent heating operation is decomposed, leaving 
 copper oxide. 
 
 The crucible is first gently warmed to expel the excess of nitric 
 acid, after which it is raised to a red heat, and maintained at this 
 temperature for about ten minutes. It is then allowed to cool in 
 the desiccator, and weighed. The heating is repeated until the 
 weight is practically constant. 
 
 Formula weight of CuSO 4 .5H 2 O = 249-3. Theoretical per- 
 centage of Cu = 2 5 '3 5. 
 
 II. As COPPER SULPHIDE. 
 
 Epitome of Process. The copper is precipitated from 
 neutral or acid solutions as cupric sulphide, CuS, by means of 
 sulphuretted hydrogen. The cupric sulphide is then converted 
 by heat into cuprous sulphide, Cu 2 S, in which form the metal is 
 weighed. Factor 
 
 (Cu 2 S) 158-6 : (Cu 2 ) 126-6 = i : 079823 
 
 Estimation of Copper in Copper Sulphate, CuS o ,,51-1,0. 
 Take about i gram. The weighed-out salt is dissolved in hot 
 water, in quantity sufficient to about half fill the beaker, and a few 
 cubic centimetres of strong hydrochloric acid are added. A stream 
 of sulphuretted hydrogen is then passed through the solution, until 
 precipitation is complete. 
 
 The precipitate is at once filtered, the funnel being covered with 
 a clock-glass to prevent atmospheric oxidation of the copper 
 sulphide into sulphate, which would pass into solution and be lost.* 
 The filter should be continuously replenished, and not allowed to 
 run dry.' The precipitate is washed down into the apex of the 
 filter with warm sulphuretted hydrogen water, and as in this 
 instance the solution contains no soluble salts, two or three rinses 
 with the warm sulphuretted hydrogen water will be sufficient to 
 wash out the acid present.f The filter is then put to dry in the 
 steam-oven. . 
 
 The dry precipitate is detached from the filter and transferred to 
 a Rose's crucible J (previously heated and weighed). The paper is 
 
 * The filtrate may be tested to see if any copper is thus escaping into the 
 solution by passing sulphuretted hydrogen through a portion of it. 
 
 f By double decomposition sulphuric acid will have been formed, and if 
 this is not washed out the paper will become charred when it is dried in the 
 steam-oven. 
 
 t This is a porcelain crucible provided with a perforated lid, through which 
 a porcelain tube can be passed, by means of which the contents of the crucible 
 
Estimation of Copper. 
 
 237 
 
 then incinerated in the usual way, and added to the precipitate. 
 A little powdered sulphur (purified by redistillation) is introduced 
 into the crucible, which is then covered with its perforated lid. 
 A gentle stream of dry 
 hydrogen is allowed to flow 
 through the bent porcelain 
 tube (Fig. 48), which is 
 then inserted into the hole 
 in the lid until the flange 
 upon the tube rests upon 
 the lid. The crucible is first 
 gently heated until the sul- 
 phur is nearly all burnt off, 
 when the temperature is 
 raised to a bright-red heat 
 and maintained at this point 
 for about ten minutes. The 
 apparatus is then allowed 
 to cool, with the hydrogen 
 still passing through it, until 
 it is nearly cold, when it is 
 removed to the desiccator, 
 and after becoming quite 
 cold is weighed. 
 
 The operation is repeated, with the addition of a little more 
 sulphur, until the weight is practically constant. 
 
 Cuprous thiocyanate, Cu 2 (CyS) 2 . Copper may also be pre- 
 cipitated as cuprous thiocyanate, which can be afterwards converted 
 into cuprous sulphide, Cu 2 S, by the action of heat in an atmosphere 
 of hydrogen in a Rose's crucible ; or it may be weighed on the 
 filter as cuprous thiocyanate. 
 
 The solution of the copper salt, which should be slightly acidified 
 with hydrochloric acid, is first saturated with sulphur dioxide* (see 
 
 can be heated in an atmosphere of hydrogen. In the absence of such an 
 apparatus, an ordinary porcelain crucible may be used ; and a common clay 
 tobacco-pipe (of such a size that the mouth of the inverted bowl will just pass 
 into the crucible) may be substituted for the perforated flange and tube. 
 Hydrogen from a cylinder of compressed gas is most convenient. If the gas 
 has to be generated from sulphuric acid and zinc, it must be dried before 
 entering the crucible, by being passed through strong sulphuric acid. By 
 arranging the flow of gas through the tube before introducing it into the 
 crucible, there is no risk of projecting the contents out by too suddenly 
 admitting the gas. Although not so advantageous, coal-gas may be employed 
 instead of hydrogen. 
 
 * Or a solution containing equal weights of hydrogen ammonium sulphite 
 and ammonium thiocyanate may be used to precipitate the copper salt. 
 
 FIG. 48. 
 
238 Typical Gravimetric Estimations. 
 
 note on p. 229). A solution of ammonium thiocyanate is then added 
 until precipitation is complete, and the mixture heated nearly to 
 the boiling-point for a few minutes. The precipitate, which is 
 coloured when it is first formed, quickly becomes white, and settles 
 very readily. It is washed with warm water by decantation two or 
 three times, and then transferred to the filter, where it is washed 
 until the wash-water is free from the acid of the original copper 
 salt (e.g. sulphuric acid if copper sulphate is employed for the 
 estimation). The precipitate is then dried in the steam-oven. The 
 dry precipitate is then treated as described above in the case of 
 copper sulphide, and finally weighed as cuprous sulphide. 
 
 As an alternative method, the precipitated cuprous thiocyanate 
 may be filtered upon a weighed filter-paper, and dried in a hot-air 
 oven at a temperature of 110 to 115, and finally weighed as 
 cuprous thiocyanate, the heating being repeated until the weight 
 is constant. The compound retains water very persistently. 
 Factor 
 
 (Cu 2 (CyS) 2 ) 242-6 : (Cu 2 ) 126-6 = i : o 52184 
 
 Silver. 
 
 Epitome of Process. Silver is precipitated from solutions 
 of its salts in the form of silver chloride, AgCl, by means of 
 hydrochloric acid, and weighed in this form.* The filter is inciner- 
 ated apart from the precipitate, with special precautions. Factor 
 
 (AgCl) i43'i6 : (Ag) 107*66 = i : 075202 
 
 Estimation of Silver in Silver Nitrate, AgNO 3 . Use 
 about o P 5 gram. The purified salt is weighed out into a beaker, and 
 dissolved in cold water, and the solution acidified with a little 
 nitric acid. Dilute hydrochloric acid is then gradually added, with 
 continued stirring, until precipitation is complete. This point is 
 easily determined, as the precipitated silver chloride coagulates and 
 settles very quickly, so that it is easy to see when the addition of 
 acid produces no further precipitation. The mixture is heated 
 nearly to the boiling-point, and then allowed to settle. The pre- 
 cipitate is washed two or three times by decantation with boiling 
 water acidified with nitric acid, after which it is transferred to the 
 filter, and washed with hot water until the filtrate is free from 
 hydrochloric acid, as indicated by the addition of silver nitrate. 
 It is then placed in the steam-oven to dry. 
 
 * Silver is sometimes precipitated and weighed as silver bromide. The 
 precipitation is effected by means of a solution of ammonium bromide, as 
 hydrobromic acid is not a usual laboratory reagent. For the same reason the 
 filter ash is treated as described in the alternative method, No. 2, for the 
 chloride, given above. Owing to the greater weight of bromide, the experi- 
 mental error is more equally divided between the silver and the halogen than 
 in the case of the chloride. 
 
Estimation of Silver. 239 
 
 The dry precipitate is detached from the paper as thoroughly as 
 possible and transferred to a porcelain crucible, which with its lid 
 has been previously heated and weighed. The paper is then folded 
 up with the soiled parts innermost, and incinerated in a platinum 
 coil. The greatest care must be taken to arrange the folding of the 
 paper and the binding of the wire round it in such a way that 
 neither the silver chloride nor the reduced silver comes in contact 
 with the heated wire ; otherwise the platinum will become alloyed 
 with silver, and the analysis will be vitiated. 
 
 The filter ash is deposited in the inverted lid of the crucible, 
 where it is moistened with a single drop of strong nitric acid, 
 allowed to fall upon it by means of a pipette. The reduced silver is 
 thereby converted into the nitrate. With a similar pipette a single 
 drop of hydrochloric acid is added, which reproduces silver chloride. 
 The crucible lid is then cautiously heated upon a pipeclay triangle 
 by means of a small flame placed at a considerable distance below 
 it, until the acids are completely evaporated. 
 
 The crucible is then gently heated until the precipitate just 
 begins to melt, when it is removed along with the lid to the 
 desiccator, and weighed when cold.* 
 
 Formula weight of silver nitrate - i69'66 theoretical percentage 
 of Ag = 63-45. 
 
 Alternative Methods of Treating the Filter, f i. After 
 the precipitate has been transferred to the crucible, the latter is 
 heated until the silver chloride just begins to fuse. The filter is 
 then rolled up and incinerated as described above, and the ash 
 allowed to fall into the crucible. It is then moistened with the 
 acids as already explained, and the crucible again heated ; at first 
 very gently, to expel the excess of acid, afterwards to the point at 
 which the silver chloride begins to melt. The crucible is then 
 cooled and weighed. 
 
 2. The precipitate is transferred to the crucible, which is then 
 heated until the chloride begins to fuse. The crucible is cooled in 
 the desiccator and weighed. 
 
 The filter is then burnt as before, and the ash deposited in the 
 crucible, which is once more gently heated, cooled, and weighed. 
 The increase in weight is made up of the weight of the ash plus the 
 
 * The fused silver chloride may afterwards be removed from the crucible by 
 adding a little dilute sulphuric acid, and introducing a fragment or two of 
 granulated zinc. The chloride is thereby reduced to metallic silver, and the 
 mass becomes detached from the porcelain. 
 
 f A third alternative process is similar to that given for lead, Method A, 
 p. 241. 
 
240 Typical Gravimetric Estimations. 
 
 weight of a small quantity of metallic silver. This weight of silver 
 is added to the amount which, by calculation, is found to be 
 present in the silver chloride already weighed. The following 
 example will make clear the process of calculating the result : 
 
 Weight of silver nitrate taken 0*292 gram 
 
 Crucible + AgCl 217265 grams 
 
 Crucible alone 21-4850 
 
 Weight of AgCl 0-2415 
 
 Multiplying this result by the factor, the weight of metallic silver 
 it contains is obtained 
 
 0-2415 x 075202 = 0-181612 gram of Ag 
 
 Crucible + AgCl + ash -f reduced silver 2173067 
 
 Crucible -f AgCl 217265 
 
 Ash ... ... ... ... ... 0*00017 
 
 2172667 
 
 Weight of reduced silver 0-004 
 
 Therefore total silver = o f 18161 2 + 0-004 = 0-185612 gram 
 
 , 0-185612 x loo 
 and *-- = 63-5 per cent, of silver 
 
 Lead. 
 
 I. As LEAD OXIDE. 
 
 Epitome of Process. The lead is precipitated from solutions 
 of its salts in the form of basic carbonate, by means of ammonium 
 carbonate. The basic carbonate is subsequently converted into 
 plumbic oxide, PbO, by heat, and the metal is weighed in this 
 form. Factor 
 
 (PbO) 223 : (Pb) 207 = i : 0-92825 
 
 Estimation of Lead in Lead Acetate, Pb(C 2 H 3 O 2 ) 2 ,3H 2 O. 
 Employ about 0*75 gram. The recrystallised salt is weighed out 
 into a beaker, and dissolved in water with the addition of a few drops 
 of acetic acid. The mixture is slightly warmed, and a solution of 
 ammonium carbonate containing a little ammonia is added, until 
 precipitation is complete. Excess of the reagent is detrimental, as 
 the precipitated lead carbonate is somewhat soluble in ammoniacal 
 solutions. The solution should be allowed to become cold before 
 being filtered. The precipitate is washed upon the filter with cold 
 water until a drop of the filtrate gives no alkaline reaction with 
 turmeric paper, and afterwards dried in the steam-oven. 
 
 The further treatment of the dry precipitate is carried out by 
 either of the two following methods : 
 
Estimation of Lead. 
 
 241 
 
 O) The precipitate is detached from the filter and carefully 
 deposited upon a small clock-glass, and covered with an inverted 
 beaker or funnel for protection. The filter is then folded up and 
 placed in a porcelain crucible, which is heated until the paper is 
 completely burnt. When it has again cooled, the ash is moistened 
 with a drop or two of strong nitric acid from a pipette, after which 
 it is gently warmed to expel the excess of acid. (In this way the 
 reduced lead is converted into lead nitrate, which in the subsequent 
 heating is decomposed into the oxide.) The precipitate upon the 
 clock-glass is then carefully transferred to the crucible, by means of 
 a small camel's-hair brush or a feather, in the manner shown in 
 Fig. 49 ; after which the crucible is heated by means of a Bunsen 
 flame to a dull-red heat for about ten minutes. It is then placed in 
 
 FIG. 49. 
 
 the desiccator to cool, and weighed. The heating is repeated until 
 the weight is constant, care being taken not to raise the temperature 
 sufficiently high to fuse the lead oxide. 
 
 () The precipitate is transferred at once to the porcelain 
 crucible. The filter is cut into about three pieces, cutting from the 
 apex of the cone to the upper edge with a clean sharp pair of 
 scissors. Each piece is then separately burnt by holding it with a 
 pair of tongs over the inverted crucible-lid, supported on a pipe-clay 
 triangle, in the manner shown in Fig. 50. The ash is then 
 moistened upon the crucible-lid with a drop of nitric acid, and the 
 lid is gently warmed to expel the excess of acid. The crucible and 
 lid are together heated until the precipitate is completely converted 
 into the oxide as in the previous description. 
 
 R 
 
242 
 
 Typical Gravimetric Estimations. 
 
 Formula-weight of Pb(C,H 3 O 2 ) 2 ,3H 2 O = 379. Theoretical 
 percentage of Pb = 54*617. 
 
 FIG. 50. 
 
 II. As LEAD SULPHATE. 
 
 Epitome of Process. The lead is precipitated as lead 
 sulphate, PbSO 4 , by means of sulphuric acid, in the presence of a 
 considerable volume of alcohol. The precipitate is dried and 
 weighed. Factor 
 
 (PbSO 4 ) 303 : (Pb) 207 -' i : 0-68316 
 
 Estimation of Lead in Lead Acetate, Pb(C 2 H 3 O 2 ) 2 ,3H 2 O. 
 Employ from 0*5 to 075 gram. The salt is dissolved in 
 rather less water than the usual quantity, in order to keep the 
 solution moderately strong (about 80 c.c. will be found convenient). 
 Dilute sulphuric acid is added until no further precipitate is 
 produced. A moderate excess of the reagent is advantageous, 
 since lead sulphate is less soluble in dilute sulphuric acid than 
 in water. The precipitation is made complete by the addition of 
 a quantity of alcohol (methylated spirit may be used) equal to about 
 twice the volume of the aqueous solution (about 200 c.c.). After 
 the precipitate has settled, it is transferred to a filter, and washed 
 with methylated spirit (in which lead sulphate is quite insoluble) 
 
Estimation of Zinc. 243 
 
 until the filtrate is free from sulphuric acid, as shown by tire 
 addition of barium chloride to a few drops of the liquid. The 
 precipitate is then dried in the steam-oven. 
 
 The filter is incinerated according to either of the two methods 
 described in the foregoing estimation. The ash, however, after 
 being moistened with nitric acid, is treated with a single drop of 
 sulphuric acid in order to convert the lead nitrate first formed into 
 lead sulphate. After the excess of the acids has been evaporated, 
 the crucible is heated to dull redness, cooled in the desiccator, 
 and weighed ; the heating being repeated until the weight is 
 constant. 
 
 Zinc. 
 
 I. As ZINC OXIDE. 
 
 Epitome of Process. The zinc is precipitated from solutions 
 of its salts in the form of basic zinc carbonate * by means of sodium 
 carbonate. The zinc carbonate is afterwards converted by heat 
 into zinc oxide, ZnO, in which form the metal is weighed. Factor 
 
 (ZnO)8i : (Zn)6s = I : 0*80247 
 
 Estimation of Zinc in Zinc Sulphate, ZnSO 4 ,7H 2 O. 
 Employ from 0*75 to i gram. The recrystallised salt is weighed out 
 and dissolved in a beaker t in the usual manner. The solution is 
 heated to boiling, and a solution of sodium carbonate is gradually 
 added until precipitation is complete. The reagent should be care- 
 fully added, so as to avoid the loss of any of the solution by the 
 undue rapidity of the effervescence due to the escape of the carbon 
 dioxide. The mixture is boiled in the covered beaker for a few 
 minutes, after which the precipitate is allowed to settle. It is 
 washed with boiling water by decantation two or three times, and 
 is then transferred to the filter. The washing is continued, with 
 boiling water, until a few drops of the filtrate give no precipitate 
 when gently warmed with a little barium chloride. 
 
 With the final washings the precipitate is collected together as 
 much as possible into the apex of the filter. 
 
 The dried precipitate is transferred from the filter to a porcelain 
 crucible (previously heated and weighed), special care being taken 
 by rubbing the sides of the filter one against the other that the 
 least possible quantity of the precipitate is left adhering to the paper. 
 During the process of incineration, the zinc oxide is liable to become 
 
 * For the composition of this precipitate, see p. 55. 
 f See note, p. 223. 
 
244 Typical Gravimetric Estimations. 
 
 reduced to the metallic state, and the metal so formed to burn 
 away and be lost. In order, as far as possible, to prevent this loss 
 by facilitating the combustion of the paper, the filter (after the 
 removal of the precipitate) may be moistened with a little solution 
 of ammonium nitrate, and again placed for a short time in the steam- 
 oven to dry.* It is then rolled up in platinum wire and burnt, the 
 lamp-flame being only momentarily applied to it, so as to avoid 
 unnecessary contact with the reducing gases of the flame. The 
 ash is allowed to fall into the crucible, which is then heated to 
 redness for ten minutes, with the lid upon it, over a Bunsen flame.t 
 After cooling in the desiccator it is weighed, and the heating 
 repeated until the weight is constant. 
 
 Formula- weight of ZnSO,7H 2 O = 287. Theoretical percentage 
 of Zn = 22-648. 
 
 II. As ZINC SULPHIDE. 
 
 Epitome of Process.- The zinc is precipitated from its solu- 
 tions by means of ammonium sulphide as zinc sulphide, ZnS. The 
 precipitate is dried, and heated in a stream of hydrogen in a Rose's 
 crucible, and the metal is weighed as ZnS. Factor 
 
 (ZnS) 97 : (Zn) 65 = i : 0-6701 
 
 Estimation of Zinc in Zinc Sulphate, ZnSO.,,7H 2 O. 
 Take about 075 gram. The salt is dissolved in warm water in a 
 beaker, and a moderate quantity of ammonium chloride is added.J 
 Colourless ammonium sulphide is then added until precipitation is 
 complete. The mixture is boiled for a short time in order to cause 
 it to coagulate, and therefore to settle more easily. It is washed twice 
 by decantation with hot water containing a little ammonium sulphide, 
 and then transferred to the filter, where the washing is continued 
 with the same liquid until the filtrate is free from sulphates. The 
 final washing should collect the precipitate as much as possible in the 
 apex of the paper. The precipitate is dried in the steam-oven, after 
 which it is transferred to a Rose's crucible. The filter is incinerated, 
 with the precautions given above, and the ash added to the pre- 
 cipitate. A little powdered sulphur (redistilled) is sprinkled into 
 
 * If the precipitate has been detached cleanly from the paper, this operation 
 may be omitted. 
 
 f When heating a crucible containing compounds of easily reduced metals, 
 the lamp should be placed at such a distance beneath the crucible that the 
 flame does not envelop it and lick over the top ; otherwise reduction of the 
 compound by the heated gases from the flame is liable to result ; and in such a 
 case as zinc, portions of the metal itself would be volatilised, and therefore lost. 
 
 J In general analyses, if the solution containing the zinc salt is acid, it is first 
 rendered slightly alkaline by the addition of ammonia. 
 
Estimation of Manganese. 245 
 
 the crucible, the lid is put on, and a gentle stream of hydrogen 
 is passed into the apparatus (see note, p. 237). The crucible is 
 first gently heated, the temperature being gradually increased to a 
 bright-red heat, using a blowpipe if necessary. It is then allowed 
 to cool with the gas still passing, until nearly cold, when it is 
 removed to the desiccator, and finally weighed. It is reheated with 
 a little more sulphur in the same manner until the weight is uniform. 
 
 Manganese. 
 
 Epitome of Process. The manganese is precipitated as 
 manganous carbonate by means of sodium carbonate. The car- 
 bonate is afterwards converted by heat into the tetroxide, Mn 3 O 4 , 
 in which form the metal is weighed. Factor 
 
 (Mn 3 O 4 ) 229 : (3Mn) 165 = i : 07205 
 
 Estimation of Manganese in Potassium Perman- 
 ganate, KMnO 4 . Employ about 075 gram. The powdered 
 recrystallised salt is weighed out into a beaker and dissolved in 
 water. A stream of sulphur dioxide is passed through the solution 
 until the liquid is perfectly colourless and clear.* The manganese 
 is then precipitated and washed exactly as described for zinc (p. 243). 
 In this case, however, two special sources of error arise. First, 
 the precipitation of the manganese is never quite complete, causing, 
 therefore, a loss; and, second, the precipitate always carries down 
 with it a certain quantity of the alkali, thus giving rise to gain. 
 Although by chance these two errors might neutralise each other, 
 an exact result can only be certainly obtained by adopting the 
 special methods for their elimination given below. 
 
 The filtrate, together with the whole of the wash-water, is 
 evaporated to dryness in a platinum, nickel, or porcelain dish. 
 The evaporation may be conducted at first over a r,ose burner, 
 provided the flame be so adjusted that the liquid never actually 
 boils. But when the liquid reaches such a state of concentration 
 that salts begin to deposit, the process must be finished upon a 
 steam-bath. During the entire process care must be taken that no 
 particles of foreign matter fall into the dish ; it is therefore desirable 
 
 * This roundabout method of obtaining a solution of manganous sulphate 
 is preferred to the direct use of the crystallised sulphate, for the reason that the 
 latter salt is not easy to obtain in a definite state of hydration. According to 
 the conditions of crystallisation, manganous sulphate forms various hydrated 
 salts. Thus, if deposited below 6, the crystals contain 7H 2 O ; between 7 and 
 20 the salt contains 5H 2 O ; while between 20 and 30 it contains 4H 2 O. When 
 manganous sulphate is used, the latter salt is the best to employ. It is pre- 
 pared by evaporating the solution at a temperature about 25. 
 
246 
 
 Typical Gravimetric Estimations. 
 
 to screen the vessel, either by supporting over it a sheet of filter- 
 paper by means of a glass tube or rod bent into a triangle and 
 held by a clamp, as shown in Fig. 51; or by supporting a large 
 inverted funnel over the dish, taking care that the water which 
 condenses upon the funnel shall drip outside the dish. 
 
 The residue is then treated with a little hot water, which dis- 
 solves the salt, leaving any 
 manganese present in the 
 form of the hydrated oxide. 
 This is filtered through a 
 separate filter, washed with 
 hot water until the wash- 
 water is free from sulphates, 
 and then placed in the steam- 
 oven to dry. 
 
 The contents of the two 
 filters are next transferred to 
 a platinum crucible, and the 
 two filter-papers incinerated 
 in a coil of platinum wire, 
 one after the other, and the 
 two ashes deposited in the 
 crucible. The crucible is 
 then raised to a bright-red 
 heat for about ten minutes 
 (see note, p. 244) ; then 
 cooled in the desiccator and 
 weighed.* 
 
 FIG. 51. 
 
 The residue in the crucible is treated with a little boiling water, 
 the particles being reduced to powder by means of a short piece of 
 stout glass rod with smooth ends. The water is poured off through 
 a filter, and the residue washed in this way by decantation three or 
 four times. It is then entirely transferred to the filter, and thoroughly 
 washed with hot water until the washings cease to give a pronounced 
 yellow coloration to the flame when heated upon a loop of clean 
 platinum wire. The residue is then dried in the steam -oven, and 
 afterwards transferred to the crucible ; the filter is incinerated, and 
 the crucible again strongly heated, allowed to cool in the desiccator, 
 
 * Weighing at this stage is not necessary. By doing so, however, the 
 student will ascertain the extent of the error which is due to the presence of 
 alkali in the precipitate, and will therefore be able to judge of the value of the 
 next operation. 
 
Estimation of Potassium. 247 
 
 and weighed. The heating is repeated until the weight remains 
 constant. 
 
 In calculating the result, it must be remembered that the ash 
 from three filters has to be deducted from the weight. 
 
 Formula-weight of KMnO 4 =158. Theoretical percentage of 
 Mn = 34-81. 
 
 Nickel. 
 
 Epitome of Process. The nickel is precipitated as nickelous 
 hydroxide, Ni(HO) 2 , by means of potassium hydroxide. The dried 
 precipitate is converted by heat into nickelous oxide, NiO, in 
 which form the metal is weighed. Factor 
 
 (NiO) 75 :(Ni)59= i : 078666 
 
 The precipitation is conducted as in the case of copper (p. 235 ; 
 see also note, p. 222). The filter is incinerated in a platinum coil, 
 and the precipitate is strongly heated in a platinum crucible 
 (avoiding the entrance of reducing gases from the flame, see p. 244) 
 until the weight is constant. 
 
 Nickelous hydroxide retains traces of the alkali with some per- 
 sistence ; therefore, when special exactness is desired, the precipitate 
 is afterwards treated as in the case of manganese. 
 
 A convenient salt in which to make an estimation of nickel is 
 the double ammonium nickel sulphate, (NH 4 ) 2 SO 4 ,NiSO 4 ,6H 2 O. 
 Formula-weight = 395. Theoretical percentage of Ni = i4'93. 
 
 Cobalt. 
 
 Epitome of Process. The cobalt is precipitated as hydroxide, 
 Co(HO) 2 , by means of potassium hydroxide. The dried precipitate 
 is reduced in a stream of hydrogen in a Rose's crucible, and 
 weighed in the metallic state. 
 
 The precipitation is carried out as in the case of nickel. The 
 filter is incinerated in a porcelain crucible along with the precipitate, 
 after which a gentle stream of dry hydrogen is delivered through 
 the Rose's apparatus (p. 237), and the precipitate heated to a low red 
 heat in the gas for about ten minutes. It is allowed to cool in the 
 hydrogen atmosphere, and then weighed. 
 
 Like the hydroxides of nickel and of manganese, cobalt hydroxide 
 retains traces of alkali. The contents of the crucible must there- 
 fore be treated with hot water until the washings are free from 
 alkali. The residue is again dried, and once more heated in a 
 stream of hydrogen. 
 
 Potassium. 
 
 I. As POTASSIUM SULPHATE. 
 
 Epitome of Process. The potassium salt is converted into 
 potassium sulphate, K 2 SO 4 , by treatment with strong sulphuric 
 
248 Typical Gravimetric Estimations. 
 
 acid. The mixture is evaporated to dryness and strongly heated, 
 and the dry residue of potassium sulphate weighed. The process 
 is only applicable in the case of compounds of potassium with 
 volatile acids which will be completely expelled by the sulphuric 
 acid, leaving nothing but the potassium sulphate. Such compounds, 
 for example, as the nitrate, chloride, bromide, etc. Factor 
 
 (K 2 S0 4 ) 174 : ( 2 K) 78 - i : 0-44827 
 
 Estimation of Potassium in Potassium Nitrate, KNO 3 , 
 Take about 0-5 gram. The nitre (purified by recrystallisation) is 
 weighed out into a platinum crucible, and a few drops of strong 
 sulphuric acid are added.* The crucible is then supported in an 
 inclined position upon the pipe-clay triangle, the lid being placed 
 
 so that the crucible is 
 very nearly but not en- 
 tirely closed, as shown in 
 Fig. 52. A gentle heat is 
 applied to the crucible, 
 the flame being placed 
 under the edge of the 
 vessel. After the nitric 
 acid has been expelled, 
 the excess of sulphuric 
 acid begins to volatilise, 
 as seen by the escape of 
 white fumes. 
 
 F 2 As the operation pro- 
 
 ceeds, the hydrogen po- 
 tassium sulphate, which is the first product of the action, and which 
 is a moderately easily fusible salt, is converted into the normal sul- 
 phate with elimination of sulphuric acid. The normal salt being 
 much more difficult of fusion, the contents of the crucible gradually 
 become pasty, and finally solid. During this stage, care must be 
 taken that no loss by spirting takes place. 
 
 The completion of this change requires prolonged exposure to 
 a bright-red heat, but in an atmosphere charged with ammonia 
 it takes place much more quickly. Hence, while the mass is heated 
 to a dull red, a few small fragments of ammonium carbonate are 
 introduced from time to time, the lid being replaced immediately. 
 The heat is continued for a few minutes after the addition of the 
 
 * In the case of potassium salts containing acids which are expelled with 
 effervescence, the salt should be first moistened with water, and the sulphuric 
 acid added drop by drop so long as effervescence takes place, the cover 
 being replaced immediately after each addition of acid. 
 
Estimation of Potassium. 249 
 
 ammonium carbonate, and the crucible cooled in the desiccator 
 and weighed. It is afterwards re-heated with a little more of the 
 carbonate and again weighed, the operation being repeated until 
 the weight is constant. 
 
 Formula-weight of KNO 3 = 101. Theoretical percentage of 
 K = 38-61. 
 
 II. As POTASSIUM PLATINIC CHLORIDE, 2KCl,PtCl 4 . 
 
 Epitome of Process. The potassium is precipitated in the 
 form of the double chloride of potassium and platinum (potassium 
 chloroplatinate), and the precipitate is weighed without alteration 
 upon a tared filter. 
 
 This process is only applicable in the case of such potassium 
 compounds as are wholly converted into the chloride by means of 
 hydrochloric acid. Factor 
 
 (2KCl,PtCl 4 ) 486 : (2K) 78 = 1 : 0-1605 
 
 Estimation of Potassium in Potassium Chloride, KC1. 
 Employ about 0-25 gram. The recrystallised salt is weighed 
 out into a small porcelain dish, and dissolved in about 10 c. c. of 
 water, to which two or three drops of hydrochloric acid have been 
 added. A moderately strong solution of platinum chloride is added 
 in slight excess,* and the mixture evaporated on a steam-bath until 
 it becomes semi- solid. 
 
 The crystalline residue is rinsed with alcohol, and, after being 
 allowed to settle, the supernatant liquid is poured off through a 
 previously weighed filter (see p. 194). The precipitate is washed 
 by decantation with alcohol two or three times, and finally trans- 
 ferred to the filter, where it is further washed with alcohol (con- 
 tained in a small wash-bottle) until the filtrate is free from any 
 tinge of yellow colour. The filter is then withdrawn from the 
 funnel, carefully folded, and placed between a pair of watch-glasses 
 to dry in the steam-oven. It is then heated and weighed until the 
 weight is constant.f 
 
 Formula-weight of potassium chloride, KC1 = 74*5. Theoretical 
 percentage of K = 52'349- 
 
 * It is desirable that the strength of the solution of platinum chloride should 
 be approximately known, in which case a quantity of it will be added which 
 contains a weight of PtCl 4 equal to about four times the weight of potassium 
 chloride taken for the analysis. The solution should be strong, so as to avoid 
 unnecessary loss of time in evaporation. If the strength of the platinum 
 chloride is unknown, it is only possible to ascertain whether excess has been 
 added by observing whether the supernatant liquid, as it becomes concentrated 
 by evaporation, has the yellow colour of the reagent. 
 
 t For an alternative method of treating the precipitate, see Estimation 
 of ammonium below. 
 
250 
 
 Typical Gravimetric Estimations. 
 
 Ammonium. 
 
 Epitome of Process. The ammonium salt is first boiled 
 with sodium hydroxide, and the ammonia which is thus expelled 
 is absorbed by dilute hydrochloric acid. The ammonium, in the 
 ammonium chloride so obtained, is then precipitated as the double 
 chloride of ammonium and platinum, 2NH 4 Cl,PtCl 4 , by means of 
 platinic chloride, and is finally weighed in this form. Factor 
 
 (2NH 4 Cl,PtCl 4 ) 644 : (2NH 4 ) 36 = i : 0-08108 
 
 Estimation of Ammonium in Ammonium Sulphate, 
 
 (NH 4 ) 2 SO 4 . Employ about 0^25 gram. The purified salt is weighed 
 out into a flask (F, Fig. 53), and dissolved by the addition of a 
 
 small quantity (30 to 40 c.c.) 
 of water. The rubber cork, 
 carrying a dropping -funnel 
 and leading-tube, is then in- 
 troduced ; the drawn-out end 
 of the funnel tube should 
 nearly, but not quite, touch 
 the surface of the liquid. 
 
 To the other end of the 
 twice-bent leading-tube * the 
 empty flask G is attached by 
 means of a second rubber 
 cork, through which passes 
 the " scrubber" tube S, which 
 is packed with small broken 
 glass, previously thoroughly 
 washed. The leading tube 
 should be moderately wide 
 in the bore, and at both ends 
 it should be cut obliquely, as 
 shown in the figure, f 
 
 When the apparatus is 
 thus arranged, a quantity of dilute hydrochloric acid is introduced 
 into the absorption flask G through, the scrubber tube, in order 
 to thoroughly wet the broken glass with the acid. While this is 
 being done the tap of the funnel must be open, to allow the air 
 
 * It is preferable that the leading-tube should be in one piece, as shown in 
 the figure, as every unnecessary joint is liable to introduce error through 
 leakage. 
 
 t This is done in a few minutes by grinding the tube down upon a sheet of 
 moderately coarse emery cloth, which is freely wetted with turpentine. 
 
 S 
 
 FIG. 53. 
 
Estimation of Ammonia. 251 
 
 to pass out (which it will do if the funnel does not touch the liquid, 
 as directed). When the acid is in, the obliquely cut end of the 
 leading-tube should just dip into the liquid, so as to about half 
 close the orifice. The tap is then closed, and the funnel about 
 three-parts filled with a moderately strong solution of sodium 
 hydroxide (i part of soda to 5 parts of water). 
 
 The solution in flask F is heated nearly to the boiling-point, 
 and the caustic soda allowed to enter from the funnel drop by drop, 
 the liquid being gently boiled the whole time.* When the whole 
 of the soda has been introduced, and the ammonia has been all 
 expelled by continuing the boiling for a few minutes, the tap of the 
 funnel is opened (the end now being beneath the surface of the 
 liquid), and the gaseous contents of the flask are very gently 
 drawn through into the absorption vessel by means of a piece of 
 rubber tube attached to the scrubber tube. The apparatus is then 
 disconnected, and the contents of the absorption flask are trans- 
 ferred to a porcelain dish. The end of the leading-tube, the 
 scrubber, and the flask are thoroughly washed with water, and the 
 washings added to the solution in the dish. 
 
 The requisite quantity of platinum chloride (see note, p. 249) is 
 added, and the solution evaporated upon a steam-bath until it 
 becomes semi-solid. The crystalline precipitate of 2NH 4 Cl,PtCl 4 
 is then washed with alcohol, and afterwards treated in one of the 
 two following alternative ways : 
 
 (a) It is transferred to a tared filter, dried, and weighed as 
 described for the similar potassium compound. 
 
 (b] The precipitate is washed with alcohol by decantation, the 
 washings being poured off through an unweighed filter. The 
 process is continued until the decanted 'liquid is free from any tinge 
 of yellow colour. The filter, which contains small quantities of 
 the precipitate, is placed in the steam-oven to dry. The main 
 portion of the precipitate is rinsed from the dish into a platinum 
 crucible by means of a fine alcohol jet, and the crucible placed in 
 the steam-oven to dry. 
 
 * Care must be taken throughout the whole of the boiling process that 
 none of the alkaline liquid is carried over either as spray or by frothing. This 
 precaution is specially necessary when the ammonia which is absorbed in the 
 Mask G is afterwards to be estimated by volumetric methods (g.v.). If the 
 leading-tube is of moderately large bore, and is cut off diagonally as directed, 
 the spray which is thrown up will run back. Should the mixture show any 
 tendency to froth (which will not be the case with the particular example here 
 chosen, but which is liable to happen with some ammoniacal liquids), a few- 
 fragments of paraffin wax should be added before commencing to boil. 
 
 In order to keep the absorption flask cool, it may with advantage be stood 
 in a dish of water during the operation. 
 
252 Typical Gravimetric Estimations. 
 
 The crucible with its dry contents is then weighed, after which 
 the filter is incinerated in a platinum coil, and the ash added. The 
 crucible is then weighed again. The increase in weight is due to 
 the ash plus a small quantity of metallic platinum derived from 
 that portion of the precipitate which was carried on to the filter 
 during the process of washing. This quantity must therefore be 
 added to the weight of platinum which is found by calculation to 
 be present in the precipitate in the crucible * (as explained in the 
 case of silver, p. 240). 
 
 The result of the analysis may be checked by converting the 
 whole of the precipitate into metallic platinum, and weighing it. 
 The crucible is first gently heated by means of a small flame, and 
 the temperature gradually raised to a bright-red heat, after which 
 the crucible is cooled in the desiccator and weighed. 
 
 Formula-weight of (NH 4 ) 2 SO 4 = 132. Theoretical percentage 
 of (NH 4 ) = 27-27. 
 
 Tin. 
 
 Epitome of Process. The tin is precipitated as sulphide 
 from either " stannous " or " stannic " solutions by means of 
 sulphuretted hydrogen. The dried precipitate is afterwards con- 
 verted by heating in contact with air, into tin dioxide, SnO 2 , in 
 which form the metal is weighed. Factor 
 
 (SnO 2 ) 150 : (Sn) 118 = i : 078666 
 
 Estimation of Tin in Stannous Chloride, SnCl 2 ,2H 2 O. 
 Employ about 0*6 gram. The purified salt is weighed into a 
 beaker, and dissolved in a small quantity of dilute hydrochloric 
 acid. The solution is diluted by the addition of about 250 c.c. 
 of moderately warm water, and sulphuretted hydrogen passed 
 through until the liquid is saturated with gas. The beaker is then 
 allowed to stand for an hour or two in a warm place, or it may be 
 gently heated for a short time to a temperature not higher than 
 can be comfortably borne by the hand. The precipitate has a great 
 tendency to pass through the filter, especially during the process 
 of washing. This is overcome to a great extent by employing a 
 double filter, or by employing a dilute solution of ammonium 
 acetate acidified with a little acetic acid, with which to wash the 
 
 * This method may be adopted as an alternative in the estimation of 
 potassium. In that case, however, it is very important that the least possible 
 quantity of the precipitate is carried on to the filter during the process of 
 washing, for the reason that, unlike the ammonium compound, the potassium 
 compound, when heated, leaves a residue of potassium chloride as well as 
 metallic platinum. 
 
Estimation of Arsenic. 253 
 
 precipitate. The dried precipitate is transferred to a porcelain 
 crucible, and the filter incinerated in a platinum coil. The ash is 
 moistened with a single drop of strong nitric acid, as in the case 
 of copper (p. 235), and the excess of acid expelled by a gentle heat. 
 The crucible is then gradually heated, first with its lid on, and 
 afterwards with it removed, until the sulphide is completely con- 
 verted into oxide ; the temperature being finally raised to a red 
 heat by means of a blowpipe. The last traces of sulphur are 
 expelled by allowing the crucible to partially cool, introducing a 
 few fragments of ammonium carbonate, and again strongly heating 
 until the weight is constant. 
 
 Formula-weight of SnCl 2 ,2H 2 O = 225. Theoretical percentage 
 of Sn = 52-44. 
 
 Arsenic. 
 
 I. As ARSENIOUS SULPHIDE, As 2 S 3 . 
 
 Epitome of Process. The arsenic is precipitated from a 
 solution of an " arsenious " compound in the form of arsenious 
 sulphide, As ? S 3 , by means of sulphuretted hydrogen. The pre- 
 cipitate is weighed upon a tared filter. It is afterwards subjected to 
 treatment with carbon disulphide to remove free sulphur. Factor 
 
 (As 2 S 3 ) 246 : (2As) 150 = i : 0-60975 
 
 Estimation of Arsenic in Arsenious Oxide, As 4 O 6 . 
 Employ about 0*5 gram. The arsenious oxide (purified by re- 
 sublimation) is weighed out into a conical flask capable of holding 
 about half a litre, and fitted with a caoutchouc cork carrying two 
 tubes for the transmission of a stream of gas through the solution, 
 the longer tube reaching nearly to the bottom of the flask. 
 
 The arsenious oxide is dissolved in a small quantity (30 to 
 40 c.c.) of dilute hydrochloric acid, the operation being aided 
 by gently warming the mixture upon a steam-bath. The solution 
 must not be boiled. When the oxide has entirely dissolved, the 
 liquid is largely diluted by the addition of about 200 c.c. of warm 
 water, and a gentle stream of sulphuretted hydrogen passed through, 
 until precipitation is complete. 
 
 Precipitated in this manner from a warm solution, the arsenious 
 sulphide settles quickly, and is easily filtered.* The excess of 
 
 * When conditions exist which cause the precipitation of sulphur, such as 
 the presence of arsenic acid, chromates, ferric salts, etc. , the solution should 
 be cold, or only slightly warm, as from a hot solution the sulphur is deposited 
 in coagulated particles, and in this form it is much more difficult to dissolve in 
 carbon disulphide. In this case the filter containing the dried precipitate must 
 be digested with carbon disulphide in a dish upon a water-bath. The carbon 
 disulphide must itself be free from dissolved sulphur, and should leave no 
 residue when a small quantity of it is evaporated upon a watch-glass. 
 
254 Typical Gravimetric Estimations. 
 
 sulphuretted hydrogen is expelled from the solution by passing a 
 brisk stream of carbon dioxide through the still warm solution, 
 after which the precipitate is transferred to a weighed filter, and 
 washed with warm water containing a little sulphuretted hydro- 
 gen. If any of the precipitate has deposited itself upon the glass 
 tube delivering the sulphuretted hydrogen, it may be dissolved off 
 by means of ammonia in a small test-tube, and reprecipitated by 
 the addition of hydrochloric acid, and added to the main portion. 
 The precipitate, after being washed until the washings are free 
 from hydrochloric acid, is dried in the steam-oven, and weighed 
 in a stoppered tube (p. 194). 
 
 The filter is then carefully replaced in the funnel, and the 
 precipitate rinsed with small quantities of carbon disulphide, until 
 a few drops of the liquid, when evaporated upon a watch-glass, 
 leave no residue. It is then once more dried and weighed. 
 
 Formula-weight of As 4 O 6 = 396. Theoretical percentage of 
 As = 7575- 
 
 II. As MAGNESIUM PYRO-ARSENATE, Mg 2 As 2 O 7 . 
 
 Epitome of Process. The arsenic is precipitated from 
 solutions of " arsenic " * compounds in the form of ammonium 
 magnesium arsenate, (NH 4 )MgAsO 4 ,6H 2 O, by means of magnesium 
 sulphate in the presence of ammonia and ammonium chloride. 
 The precipitate is converted by heat into magnesium pyro-arsenate, 
 Mg 2 As 2 O 7 , in which form the arsenic is weighed. Factor 
 
 (Mg 2 As 2 O 7 ) 310 : (2As) 150 = i : 0-4838 
 
 Estimation of Arsenic in Hydrogen Disodium Arse- 
 nate, HNa 2 AsO 4 ,i2H 2 O. Employ about 075 gram. The re- 
 crystallised salt is weighed out into a wide-mouth stoppered bottle, 
 about 500 c.c. capacity, and dissolved in cold water. A moderate 
 quantity of ammonia is added, and then " magnesia mixture " f in 
 excess. The mixture is then briskly shaken in the stoppered bottle 
 for about five minutes, after which it is filtered, washed, and dried, 
 exactly as directed for the treatment of magnesium phosphate 
 (P- 233). 
 
 The dried precipitate is detached from the filter and deposited 
 upon a clock-glass, or a piece of black glazed paper, and covered 
 
 * "Arsenious" compounds must be first oxidised to the "arsenic" con- 
 dition by gently warming the solution with strong hydrochloric acid, and 
 adding a few crystals of potassium chlorate. The liquid is kept at a gentle 
 heat until it no longer smells of chlorine, and then rendered alkaline by the 
 addition of ammonia. 
 
 t See Reagents, in the Appendix. 
 
Estimation of Antimony. 255 
 
 over with a clean beaker. The filter is then moistened with a few 
 drops of solution of ammonium nitrate, and returned to the steam- 
 oven. As soon as it is dry,* it is incinerated in a porcelain crucible, 
 the temperature being gradually raised to a bright-red heat. The 
 crucible is then allowed to cool, and the precipitate transferred to it 
 (see p. 241), after which it is heated in a gentle stream of oxygen by 
 means of Rose's apparatus. The heat, which is applied very gradu- 
 ally at first, is finally increased to bright redness. After being cooled 
 and weighed, it is reheated in oxygen until the weight is constant. 
 Formula- weight of HNa 2 As0 4 , 1 2H 2 O = 402. Theoretical per- 
 centage of As = 18-65. 
 
 Antimony. 
 
 Epitome of Process. The antimony is precipitated as 
 antimonious sulphide, SboS a , by means of sulphuretted hydrogen. 
 The precipitate, which retains water after drying in the steam-oven, 
 and which also usually contains free sulphur, is treated by one of 
 the following methods : 
 
 (a) It is freed from water and sulphur by being heated in a 
 porcelain boat in a stream of carbon dioxide, and weighed as 
 Sb 2 S 3 ; or 
 
 (b) It is converted into the tetroxide, Sb 2 O 4 , by treatment with 
 strong nitric acid, and weighed in this form.f Factors 
 
 (Sb 2 S 3 )336 : (2Sb) 240 = i : 071428 
 (Sb 2 O 4 ) 304 : (2Sb) 240 = i : 078947 
 
 Estimation of Antimony in Tartar Emetic, (SbO)K- 
 (C 4 H 4 O, ; ). Employ about 075 gram. The dried salt is weighed 
 into a conical flask, and dissolved in a little water. Three or four 
 grams of tartaric acid, dissolved in a little water, are added, and 
 then a few drops of hydrochloric acid. The solution is then 
 diluted by the addition of about 200 c.c. of slightly warm water 
 (if sufficient tartaric acid has been introduced, this dilution will 
 not cause any precipitation), and a slow stream of sulphuretted 
 hydrogen passed through, as described in the case of arsenic (p. 253). 
 When the solution is saturated with the gas, the exit and delivery 
 tubes may be joined by means of a piece of rubber tube, in order 
 to prevent the entrance of air, and the flask is allowed to stand in 
 a warm place for an hour or two. 
 
 The excess of sulphuretted hydrogen is then expelled by passing 
 
 * If left too long in the oven, the ammonium nitrate is expelled, and the 
 object for which it is added, namely, to aid the combustion of the paper, will 
 be defeated. 
 
 f This latter method is preferable in cases where the quantity of antimony 
 is small. 
 
256 Typical Gravimetric Estimations. 
 
 a brisk stream of carbon dioxide through the mixture, which may 
 at the same time be gradually heated to the boiling-point.* The 
 precipitate is transferred to a weighed filter, and washed as quickly 
 as possible (without allowing the filter to run dry) with hot water 
 containing a little sulphuretted hydrogen. It is then dried in the 
 steam-oven until the weight is constant. 
 
 The precipitate, consisting of Sb 2 S 3 plus an unknown quantity 
 of water and free sulphur, is treated in one of the following ways : 
 
 (a) A certain proportion of it (as much as convenient) is trans- 
 ferred to a porcelain boat which has previously been heated and 
 weighed, and the boat with its charge again weighed. The boat 
 is then pushed, by means of a glass rod, into a wide glass tube 
 about 25 c.c. (10 inches) long, supported in a horizontal position, 
 and a gentle current of dry carbon dioxide passed over it. The 
 boat is gently heated by means of a Bunsen flame in the stream 
 of gas, whereby the water and the sulphur are expelled from the 
 precipitate. The tube is allowed to cool with the gas still passing 
 through, after which the boat containing its residue of black 
 anhydrous Sb 2 S 3 is withdrawn and weighed.t 
 
 From the loss of weight the contents of the boat have suffered 
 by this treatment, the quantity of pure Sb 2 S 3 which was present 
 in the whole of the original precipitate can be deduced, and from 
 this the amount of antimony can be calculated. 
 
 (b} A certain proportion of the precipitate is transferred to a 
 weighed porcelain crucible, and the crucible with the precipitate 
 again weighed.t It is moistened by the addition of three or four 
 drops of ordinary strong nitric acid (sp. gr. 1*42), and the crucible 
 covered with an inverted lid. By means of a pipette 3 or 4 c.c. 
 of fuming nitric acid (sp. gr. 1*5) are introduced, the lid being 
 drawn slightly aside and quickly replaced. Rapid oxidation of the 
 antimony sulphide and free sulphur immediately takes place, after 
 which the contents of the crucible are evaporated to dryness upon 
 a steam-bath. 
 
 The crucible is then heated by means of a Bunsen flame, at 
 
 * Since antimonious sulphide is decomposed by boiling hydrochloric acid, 
 this operation would not be safe if much acid were present. 
 
 f A small portion of the original dried precipitate may be tested for free 
 sulphur by heating it in a test-tube with a little strong hydrochloric acid. If 
 it entirely dissolves to a perfectly clear liquid, no sulphur is present, in which 
 case the portion which is treated in the boat will only require to be very gently 
 heated to expel the water. 
 
 J In cases where the total precipitate is very small, the filter with its entire 
 contents is transferred to the crucible. In such a case the filter-ash must be 
 deducted from the final weight. 
 
Estimation of Cadmium. 257 
 
 first very gently, whereby sulphuric acid is expelled, and antimony 
 tetroxide, Sb.^, is left. The crucible is cooled in the desiccator, 
 and weighed ; the heating being repeated until the weight remains 
 constant. From this result the weight of antimony tetroxide which 
 would be yielded by the whole of the precipitated sulphide can be 
 calculated, and therefrom the weight of antimony. 
 
 Formula-weight of (SbO)K(C 4 H 4 O, ; ) = 323. Theoretical per- 
 centage of Sb = 37-15. 
 
 Cadmium. 
 
 Epitome of Process. The cadmium is precipitated as cad- 
 mium sulphide, CdS, by means of sulphuretted hydrogen. The 
 precipitate is dried, and weighed upon a tared filter. Factor 
 
 (CdS) i44:(Cd) 112 = i 107777 
 
 The precipitation is conducted in a beaker, from a solution 
 acidified with hydrochloric acid. The precipitate is first washed 
 with water containing a little sulphuretted hydrogen and a few 
 drops of hydrochloric acid, and afterwards with plain water. 
 
 The dried precipitate, after weighing, is extracted with carbon 
 disulphide to remove traces of free sulphur (see Arsenic, p. 254), and 
 again dried and weighed. 
 
 Mercury. 
 
 Epitome of Process. The mercury is precipitated in 
 the form of mercuric sulphide, HgS, by means of sulphuretted 
 hydrogen. The precipitate is dried, and weighed upon a tared 
 filter. Factor 
 
 (HgS) 232 : (Hg) 200 = i : 0-86206 
 
 The details of the operation are the same as in the case of 
 cadmium. 
 
SECTION III. 
 TYPICAL GRAVIMETRIC ESTIMATIONS OF ACID RADICALS. 
 
 Sulphuric Acid, (SO,). 
 
 Epitome of Process. The sulphuric acid radical is precipi- 
 tated in the form of barium sulphate, BaSO 4 , by means of barium 
 chloride in the presence of hydrochloric acid. The precipitate is 
 dried, and weighed in the same form.* Factor 
 
 (BaSO 4 ) 233 : (SO 4 ) 96 = i : 0*4120 
 
 Estimation of (SO,) in Potassium Alum, K. J So.,.Al._,- 
 (SO 4 ) 3 ,24H 2 O.t Employ about 0*5 gram. The salt is weighed out 
 into a beaker, dissolved in water, and the solution heated to the 
 boiling-point. A small quantity of hydrochloric acid is added, 
 and then a little moderately strong solution of ammonium chloride. 
 While the liquid is still boiling, a solution of barium chloride, 
 heated to boiling in a large test-tube, is added until precipitation 
 is complete. The precipitate is washed three or four times by 
 decantation with hot water, and finally transferred to the filter, 
 where it is washed until the filtrate is free from chlorides. The 
 last washings should be made to collect the precipitate as much as 
 possible into the apex of the filter, after which it is placed in the 
 steam-oven to dry. 
 
 The dry precipitate is detached as completely as possible from 
 the paper and transferred to a platinum crucible, and the filter 
 incinerated in a platinum coil. The precipitate should be shaken 
 slightly to one side in the crucible, so that the ash may be deposited 
 in a clear space upon the bottom. By means of a fine pipette, the 
 ash is moistened with a single drop of hydrochloric acid, in order 
 to convert into barium chloride any sulphide which may have been 
 
 * Barium is estimated as BaSO 4 by precipitation with dilute sulphuric acid, 
 the process being carried out as described above. Strontium is also estimated 
 in the same manner, except that for the complete precipitation of SrSO 4 it is 
 necessary to add a quantity of alcohol (methylated spirit) equal in bulk to the 
 volume of the aqueous liquid. 
 
 f For the treatment of insoluble sulphates, see p. 184 
 
Estimation of Carbon Dioxide. 
 
 259 
 
 formed by the reduction of a portion of the sulphate adhering to 
 the paper. Then a drop of dilute sulphuric acid is similarly added, 
 and the crucible very gently heated to expel the excess of the acids, 
 after which the temperature is raised to a red heat. It is then 
 cooled in the desiccator and weighed.* 
 
 Formula-weight of K 2 SO 4 ,A1 2 (SO 4 ) 3 ,24H 2 O = 948. Theoretical 
 percentage of SO 4 = 40-50. 
 
 Chlorine in Chlorides t 
 
 Epitome of Process. The chlorine is precipitated in the 
 form of silver chloride, AgCl, by means of silver nitrate, and is 
 weighed in this form. The filter is incinerated apart from the 
 precipitate (see Silver, p. 238). Factor 
 
 (AgCl) 142-5 : (Cl) 35-5 = i I 0-2491 
 Estimation of Chlorine in Sodium Chloride, NaCl. 
 
 Employ about 0*25 gram. 
 
 The process is carried out as in the estimation of silver, p. 238. 
 
 Formula-weight of NaCl = 58-5. Theoretical percentage of 
 Cl = 6c-68. 
 
 Carbon Dioxide. 
 
 I. ESTIMATION BY DIFFERENCE. 
 
 Epitome of Process. A known weight of the carbonate 
 is decomposed by a suitable acid in a 
 weighed apparatus, and the carbon dioxide 
 expelled. The loss of weight which re- 
 sults, represents the carbon dioxide in the 
 compound. Various forms of apparatus 
 are used. Fig. 54 represents a simple 
 apparatus which can readily be fitted up 
 by the student himself, and which gives 
 very good results. It consists of a small 
 flask of thin glass fitted with a cork carry- 
 ing a thistle funnel and a bent exit tube. 
 The exit tube is drawn down to a mode- 
 rately fine tube, which passes to the bottom 
 of a small narrow test-tube, t. The portion 
 
 * In the estimation of (SO 4 ) as barium sulphate 
 in general analysis, whether in a sulphate, or in a 
 sulphide after oxidation into (SO 4 ), the following 
 points must be borne in mind : nitric acid, 
 nitrates, and chlorates must be absent ; hydro- 
 chloric acid must not be present in large quantity. 
 
 f The estimation of bromine in bromides is 
 carried out in the same manner, the precipitate 
 being treated by the alternative method, No. 2, 
 given on p. 239. 
 
260 
 
 Typical Gravimetric Estimations. 
 
 of the funnel tube which extends down into the flask is also drawn 
 out. A short piece of small rubber tube is introduced into the 
 thistle funnel at <:, and a piece of drawn-out glass rod, s, pushed 
 into this, serves as a stopper. Dilute acid (for decomposing the 
 carbonate) is introduced into the funnel, and the tube / is half 
 filled with strong sulphuric acid in order to arrest the escape of 
 aqueous vapour along with the issuing carbon dioxide. By slightly 
 lifting the little stopper s, the admission 
 of acid can be regulated with consider- 
 able nicety. 
 
 The apparatus shown in Fig. 55 
 (known as the Schrotter apparatus) 
 possesses the advantage of not con- 
 taining any corks. The carbonate is in- 
 troduced at the stoppered neck s ; the 
 acid is admitted from the stoppered 
 funnel /, while the escaping carbon 
 dioxide is dried by its passage through 
 the strong sulphuric acid contained in d. 
 The gas passes up the central tube 
 within d, as indicated by the arrows, 
 and forces the acid down to the level 
 of the holes near the bottom of the 
 second tube, and then bubbles out 
 through the acid and escapes at the 
 top of the outer tube. 
 
 Estimation of CO in Potas- 
 sium Carbonate, K 2 CO 3 . Employ 
 about i '5 gram. The pure dry salt is 
 weighed into the flask of the apparatus, 
 and the slightly greased stopper in- 
 serted in the neck s. The dropping- 
 FlG - 55- funnel is then nearly filled with dilute 
 
 sulphuric acid, and the drying-tube is about half filled with strong 
 oil of vitriol, and the apparatus carefully weighed. The acid from 
 the funnel is then allowed to very slowly enter the flask, so that the 
 carbon dioxide may escape in single bubbles through the strong 
 acid in the drying-tube. If the bubbles pass in a stream faster than 
 about two to the second (i.e. too quickly to be easily counted), the 
 gas will not be deprived of aqueous vapour by the acid, and the 
 result of the analysis will be vitiated. When the effervescence has 
 ceased, the remainder of the acid is allowed to run in and the tap 
 left open. The flask is then very gently heated over a small flame, 
 in order to expel the whole of the carbon dioxide, until the liquid 
 just begins to boil. During the heating, the gaseous contents of the 
 
Estimation of Carbon Dioxide. 261 
 
 flask are slowly sucked out by means of an aspirator attached by 
 a caoutchouc tube to the exit of the drying-tube. The aspirator 
 shown on p. 204 may be employed, or the gas may be drawn out 
 by the lungs. In the latter case, the peculiar sweet taste of the 
 carbon dioxide makes it easy to tell when the whole of the gas has 
 been withdrawn and its place taken by air. The aspiration must 
 not be unnecessarily prolonged. The apparatus is then allowed 
 to become cold, and afterwards weighed. While the apparatus is 
 cooling, the exit tube of the drying-vessel d should be closed with 
 a cap (consisting of a short piece of caoutchouc tube plugged at 
 one end with a fragment of glass rod) in order to prevent the 
 absorption of atmospheric moisture by the sulphuric acid ; the cap 
 is removed while weighing.* 
 
 Formula-weight of K,CO 3 = 138. Theoretical percentage of 
 CO, = 31-88. 
 
 II. ESTIMATION BY DIRECT WEIGHING. 
 
 Epitome of Process- The carbonate is decomposed by an 
 acid, and the evolved carbon dioxide is absorbed by soda-lime in 
 a weighed apparatus. 
 
 Fig. 56 represents the apparatus which may be employed for 
 the process. The carbonate is decomposed in the flask B by means 
 of acid introduced from the dropping-funnel, the fine drawn-out 
 end of which reaches nearly to the bottom. The carbon dioxide 
 is first dried by passing through the tube E, containing pumice 
 moistened with sulphuric acid. It is then deprived of acid vapours 
 by means of the tube D, filled with pumice impregnated with an- 
 hydrous copper sulphate (note below). The pure dry gas then 
 passes through the weighed absorption-tube A, which is about 
 four-fifths filled with soda-lime, the remaining portion, a, furthest 
 removed from the incoming gas, being filled up with dry calcium 
 chloride. Beyond the absorption-tube is a smaller U-tube contain- 
 ing soda-lime, which is weighed immediately before and after the 
 experiment. This serves as a guard tube, and its weight should 
 remain constant if the absorption- tube A is properly absorbing the 
 whole of the carbon dioxide. After this guard tube is placed a 
 
 * The accuracy of the results obtained by this method depends very largely 
 upon the slowness with which the carbon dioxide is disengaged. The method 
 is more suitable for those carbonates which give up the whole of their carbon 
 dioxide when acted upon by sulphuric acid ; in other words, carbonates of those 
 metals which yield soluble sulphates. When hydrochloric or nitric acid is 
 employed, an additional source of error is introduced by the volatilisation of 
 small quantities of the acid during the process of expelling the carbon dioxide. 
 This may, however, in a measure be prevented by attaching to the drying-tube 
 d another small thin gla:;s tube filled with particles of pumice which have been 
 soaked in strong copper sulphate solution, and afterwards heated in an air-oven 
 to 200. The anhydrous copper sulphate absorbs hydrochloric acid as well 
 as aqueous vapour. 
 
262 
 
 Typical Gravimetric Estimations. 
 
 U-tube containing sufficient strong sulphuric acid to cover the bend. 
 This serves the double purpose of preventing the entrance of 
 moisture to the guard tube, and also as an indicator of the rate at 
 which the gas is passing through the tubes. The tube S, contain- 
 
 FIG. 56. 
 
 ing soda-lime, is for the purpose of removing carbon dioxide from 
 the air which is drawn through the apparatus at the conclusion of 
 the operation by means of the aspirator (see p. 204). This tube 
 and the aspirator are not connected until the final stage of the 
 process. 
 
 Estimation of CO 2 in Calcium Carbonate, CaCO 3 . Em- 
 ploy about i gram. The pure dry carbonate is weighed out into 
 the flask B. The caoutchouc stopper carrying the funnel and 
 exit-tube is inserted, and the funnel filled with dilute hydrochloric 
 acid. The apparatus is then connected up as shown in the figure, 
 the absorption-tube A having previously been weighed.* The acid 
 is allowed slowly to enter the flask, the admission of the acid being 
 regulated by the rate at which bubbles are seen to pass through 
 the acid in the last tube \ they should not escape faster than two 
 
 * It will be obvious that the apparatus must be perfectly free from leakage. 
 Before being used, it may be tested by connecting a long vertical glass tube, 
 dipping into a vessel of water, to the exit of the last U-tube, and then sucking 
 air out of the apparatus, through the dropping-funnel, until a column of water 
 has been drawn up the long tube. On closing the tap of the funnel, the water 
 in the tube should remain stationary. 
 
Estimation of Nitric Acid. 263 
 
 per second. When effervescence is at an end, the remainder of 
 the acid is run in, the tap is left open, and the tube S is attached 
 to the funnel. The aspirator is connected to the other end, and 
 arranged to draw a very slow stream of air through the apparatus. 
 A gentle heat is applied to the flask, and the temperature of the 
 liquid kept for a short time just below the boiling-point. The 
 aspiration should be continued for 1 5 to 20 minutes, to ensure the 
 whole of the carbon dioxide being drawn out of the flask and ab- 
 sorbed in the soda-lime. 
 
 During this absorption the soda-lime becomes perceptibly warm, 
 and loss would result from the escape of moisture from the material. 
 This loss is prevented by the layer of calcium chloride contained in 
 the tube. 
 
 When the process is concluded, the aspirator is stopped, and the 
 absorption-tube A disconnected. The ends of the exit-tubes are at 
 once closed with caps, and the tube carefully wiped and weighed. 
 The caps are removed during the weighing.* 
 
 The increase in weight is the carbon dioxide expelled from the 
 calcium carbonate employed. 
 
 Formula-weight of CaCO 3 = 100 ; theoretical percentage of 
 C0 2 - 44- 
 
 Nitric Acid (NO 3 ). 
 
 Epitome of Process. The nitrate is strongly heated 
 with powdered silicon dioxide, whereby the whole of the N 2 O 5 
 present in the nitrate is expelled, and the oxide of the metal 
 remains associated with the silica.f The loss of weight represents 
 N 2 O f) , from which the proportion of NO 3 present can be calculated. 
 Factor 
 
 (N 2 O 3 ) 108 : (2NO 3 ) 124 = i : 1-148 
 
 Estimation of (NO 3 ) in Potassium Nitrate, KNO 3 . 
 
 Employ o'5 to 075 gram. From 2 to 3 grams of powdered quartz 
 are placed in a platinum crucible ; it is then strongly heated, and 
 weighed after cooling. The dry powdered nitre (previously pre- 
 pared by just melting the salt in a porcelain dish, pouring out 
 the fluid into a clean warm porcelain dish, and afterwards powder- 
 ing it) is introduced into the crucible, and the whole weighed 
 again. The nitrate and the silica are intimately mixed by means 
 
 * At the conclusion of the determination, the various purifying tubes may 
 remain connected together, the extreme exits being closed with caps ; they will 
 then be ready for use in a second experiment. 
 
 f A few metallic nitrates, when heated alone, give up the whole of their 
 N 2 O 3 , and leave only a metallic oxide ; but most nitrates, when thus heated, 
 leave a residue of uncertain composition 
 
264 Typical Gravimetric Estimations. 
 
 of a clean stout piece of wire or a fine glass rod, and the mixture 
 heated in the covered crucible for about 20 minutes to a red heat.* 
 It is then allowed to cool in the desiccator, and weighed ; the 
 heating being repeated until no further loss takes place. 
 
 Formula-weight of KNO 3 = 101. Theoretical percentage of 
 (N0 3 ) = 61-38. 
 
 Phosphoric Acid (FO 4 ). 
 
 Epitome of Process. The phosphoric acid is precipi- 
 tated in the form of ammonium magnesium phosphate by means 
 of "magnesia mixture " (as in the estimation of arsenic, p. 254). 
 The dried precipitate is converted by heat into magnesium pyro- 
 phosphate, in which form it is weighed (see Estimation of mag- 
 nesium, p. 233). Factor 
 
 (Mg 2 P 2 O 7 ) 222 : (2PO 4 ) 190 = i : 0-8558 
 
 Silicic Acid. 
 
 Epitome of Process. The silicic acid is precipitated in the 
 gelatinous state by means of hydrochloric acid. The mixture is 
 evaporated to dryness, whereby the precipitated acid is converted 
 into the anhydride (silicon dioxide), SiO 2 . The silica so obtained 
 is then filtered, washed, and dried, and weighed in that form. 
 
 Silicates which are decomposed by acids, are treated at once 
 with hydrochloric acid ; the majority are converted by fusion with 
 alkaline carbonates into silicates of the alkalies, which are after- 
 wards precipitated by hydrochloric acid (see p. 285). 
 
 * In general analysis, where there is any likelihood of chlorides or sulphates 
 being present, the temperature should not rise above that at which the crucible 
 is just visibly red hot, or loss will result from the decomposition of these com- 
 pounds. In such a case, the heating should be prolonged to at least half an 
 hour. 
 
SECTION IV. 
 EXERCISES IN GRAVIMETRIC ANALYSIS. 
 
 INTRODUCTION. Sampling and Powdering. When the 
 
 substance to be analysed is a metal or an alloy of homogenous 
 composition, the portion taken for analysis may be obtained by 
 boring into the mass with a steel drill in a lathe, carefully catching 
 the borings upon a clean sheet of paper. If this plan is not prac- 
 ticable, the metal may be cut with a clean sharp file, and the filings 
 used for the analysis. In this case, however, there is always risk 
 of particles of iron from the file becoming mixed with the sample. 
 These can be removed (unless the metal happens to be iron, or 
 other metal attracted by the magnet) by means of a magnet. 
 When the bar, ingot, or piece of metal is small, it may be broken 
 or cut into small pieces. In the case of metals, it is not necessary 
 that the portion taken for analysis should be in a state of very fine 
 subdivision ; the fragments obtained by boring or even by break- 
 ing or cutting are sufficiently small, as metals for the most part 
 are dissolved without much difficulty. Should the borings or 
 filings contract any oil or grease during their production, this 
 should be removed by washing them with either benzol or carbon 
 disulphide. 
 
 Solid substances which are capable of being powdered, such as 
 minerals, ores, slags, etc., must be reduced to as fine powder as 
 possible, and a portion selected for analysis which represents a fair 
 sample of the whole. A number of the selected lumps of the sub- 
 stance are first broken down into small fragments : in a wedgwood 
 mortar, if the material is moderately soft ; in a steel mortar, if hard ; 
 or they may be broken up by wrapping them in a clean cloth and 
 striking them with a hammer upon an anvil. These fragments 
 are then reduced to a coarse powder ; with hard substances, the 
 operation being performed in a steel percussion mortar (A, Fig. 57). 
 This consists of three parts : the foot or base, a ; a short steel tube, 
 , which fits into the cavity in the base ; and a steel rod or plunger, 
 
266 Typical Gravimetric Separations. 
 
 c, which will just pass into the tube or cylinder. A larger percussion 
 mortar is shown at B, Fig. 57, in which a massive block of steel 
 takes the place of the parts a and b. The fragments, in small 
 quantities at a time, are placed in the tube, and the plunger 
 or pestle introduced. This is then sharply struck with a fairly 
 heavy hammer several times, which crushes the material into 
 coarse powder. When all the lumps chosen have been crushed in 
 this way, the coarse powder is spread upon a sheet of paper and 
 thoroughly mixed with a spatula or paper-knife. A small portion 
 
 A 
 
 FIG. 57 . 
 
 of this is then withdrawn, and reduced to the finest powder possible 
 in an agate mortar. In carrying out this last operation, small 
 quantities at a time, taken upon the end of a spatula, are put into 
 the mortar, and powdered entirely by a grinding process, and not 
 by using the pestle as a hammer, which would be liable to chip 
 both the pestle and the mortar. The powder should then be passed 
 through a fine sieve, which may be made by tying a piece of fine 
 muslin tightly over the mouth of a broken beaker.* The portion 
 which is retained upon the sieve should be returned to the mortar 
 for further grinding, until the whole of it is reduced to a sufficiently 
 fine powder to pass through. 
 
 * The little piece of muslin should be renewed before using the sieve for 
 another mineral. If metal sieves are employed, the greatest care must be 
 exercised to have them perfectly clean. Any traces of a former operation 
 should be beaten and brushed out, and the sieve rinsed by sifting a small 
 portion of the powder under examination through it, and rejecting the siftings. 
 On no account should a sieve be wetted. 
 
Separation of Silver and Copper. 267 
 
 Analysis of Silver Coinage. 
 
 (Alloy of Silver and Copper?) 
 
 Epitome of Process. The metal is dissolved in nitric acid, 
 and the silver precipitated as silver chloride. This is separated, 
 and weighed as such. The copper is precipitated from the filtrate 
 in the form of copper hydroxide, which is converted by heat into 
 copper oxide, and weighed in this form. 
 
 The silver coin (a new threepenny piece) is carefully weighed, 
 and then dissolved in a little nitric acid in a small conical flask.* 
 The acid, consisting of strong nitric acid, to which about half its 
 volume of water has been added, is introduced by means of a small 
 funnel, which should be allowed to remain in the mouth of the 
 flask during the whole process, as it serves to prevent loss of the 
 liquid by spirting. The solution is started by the application of a 
 gentle heat, after which the action continues briskly until the metal 
 is nearly all dissolved. A few more drops of acid may be added, 
 if necessary, for the complete solution of the coin, but excess of 
 acid is to be avoided.! 
 
 When the metal has entirely dissolved, the excess of acid is 
 evaporated off by gently boiling the solution. The liquid is then 
 diluted with water and transferred to a beaker, the flask and funnel 
 being thoroughly rinsed once or twice with distilled water, so that 
 the whole of the solution is transferred without loss. 
 
 From this solution the silver chloride is precipitated, and the 
 silver estimated in the manner described on p. 238. The filtrate, 
 together with the washings containing the copper, is heated to 
 boiling, and the copper is estimated as described on p. 235. 
 
 A more expeditious way of carrying out the analysis is as follows. 
 The metal is dissolved in a stoppered assay bottle, the operation 
 being carried out exactly as described above, except that the bottle 
 
 * English silver coinage consists of 925 parts of Ag and 75 of Cu ; i 
 gram of the alloy therefore contains only 0*075 Cu. The weight of a threepenny 
 piece is about 1*4 gram, which therefore contains o'iO5 gram Cu, yielding 
 0-133 gram CuO. It is undesirable, therefore, to employ lessih&n this weight of 
 metal for the analysis. If a larger coin is used, it should be flattened out in a 
 rolling machine, and the thin metal cut into fragments with a pair of scissors ; 
 or, in the absence of a roller, the coin may be broken by bending it backward 
 and forward, holding it by means of two pairs of pliers. The requisite quantity 
 is then weighed out. 
 
 f Old silver coins, especially the larger pieces, sometimes contain small 
 quantities of gold, which will remain as a black residue after solution in nitric 
 acid. Minute quantities may be neglected without affecting the result of the 
 analysis. 
 
268 Typical Gravimetric Separations. 
 
 requires to be heated with more care. It should be placed upon 
 a piece of asbestos cloth, upon a hot iron plate (p. 223), or may 
 be heated on a steam-bath. When the excess of acid has been 
 evaporated, the liquid is diluted in the flask, and the silver chloride 
 precipitated by the addition of hydrochloric acid. The stopper is 
 then inserted, and the flask vigorously shaken for a few minutes. 
 The precipitate is thus obtained in a condition which enables it to 
 settle very rapidly, and the perfectly clear liquid is decanted off 
 into a beaker as thoroughly as possible. The flask is then half 
 filled with water, and the contents again thoroughly shaken, and 
 after settling, the liquid is decanted. After two such washings the 
 whole of the copper will have been washed out, and the solution 
 containing the copper is boiled, and the copper estimated as already 
 described. 
 
 The flask containing the silver chloride is then completely filled 
 
 FIG. 58. 
 
 with water, the mouth is closed with the finger, and the vessel 
 inverted into a weighed porcelain crucible filled with water, in the 
 manner shown in Fig. 58. In this way the precipitate is transferred 
 to the crucible, the whole of it being readily caused to fall down by 
 gently tapping the flask. When the whole of the precipitate has 
 settled to the bottom of the crucible, the flask is first gently raised 
 to the level of the edge of the crucible, so as to allow one or two 
 air-bubbles to enter and displace a little water, which will fill up and 
 overflow the crucible, and then it is quickly drawn away sideways 
 
Analysis of Solder. 269 
 
 (being held in the position shown in Fig. 59) and allowed to 
 empty.* The water in the crucible is then carefully decanted as 
 much as possible, and the crucible placed in the steam-oven to 
 dry. When dry, the crucible may be gently heated until the silver 
 
 FIG. 59- 
 
 chloride just begins to melt ; and after cooling in the desiccator, 
 it is weighed. 
 
 Usually unglazed crucibles are employed for this operation. The 
 precipitate dries more quickly in such a crucible, and the chloride 
 is not heated after it is removed from the steam-oven. 
 
 Analysis of Common Solder. 
 
 {Alloy of tin and lead?) 
 
 Epitome of Process. The alloy is acted upon by strong 
 nitric acid, which converts the tin into stannic oxide, SnO.2, in which 
 form the tin is estimated. The lead passes into solution as lead 
 nitrate, from which it is precipitated in the form of lead sulphate, 
 PbSO 4 , by sulphuric acid, and weighed as such. Employ about 
 0*5 gram.f 
 
 * This operation is quite easily accomplished without disturbing the 
 precipitate in the crucible. The student should perform the experiment of 
 filling an empty flask with water, inverting it in a crucible, and then withdrawing 
 the flask as here directed. He will find that there is no difficulty in drawing 
 the flask aside without letting an air-bubble pass up, which would of course 
 cause a disturbance in the crucible, and throw out any precipitate it contained. 
 As a precaution, the crucible may be placed in a shallow glass basin, so that, 
 should any of the precipitate be accidentally thrown out, it will not be lost. 
 
 f Common solder consists of about equal parts of tin and lead, but as there 
 are varieties of solder in commerce containing these metals in different propor- 
 tions, a duplicate estimation of one at least of the constituents should be made 
 as a check. 
 
270 Typical Gravimetric Separations. 
 
 Tin The alloy is weighed out into a porcelain dish, and 
 dissolved in a small quantity of moderately strong nitric acid, the 
 dish being covered with a clock-glass. When all action has ceased, 
 the cover is rinsed into the dish, and the mixture evaporated nearly 
 to dryness upon a steam-bath. Water is added, and the residue of 
 stannic oxide is transferred to a filter and washed. 
 
 After the stannic oxide has been dried in the steam-oven, it is 
 transferred to a porcelain crucible, being detached as completely 
 as possible from the paper. The filter is incinerated in a platinum 
 coil, care being taken to fold the soiled portion of the paper into the 
 interior of the roll. The ash is moistened with a drop of nitric 
 acid, and the crucible with its contents heated, first gently to expel 
 the acid, and afterwards to a red heat. It is then cooled in the 
 desiccator and weighed. 
 
 Lead. The filtrate containing the lead is precipitated with 
 sulphuric acid, and the lead estimated as lead sulphate (see p. 242). 
 
 Analysis of German Silver. 
 
 (Alloy of copper, zinc, and nickel, with traces of tin and iron .) 
 
 Epitome of Process. The alloy is dissolved in nitric acid, 
 and the insoluble residue of stannic oxide separated and weighed. 
 
 The copper is next precipitated as sulphide by means of sul- 
 phuretted hydrogen, and estimated as Cu. 2 S. 
 
 The filtrate containing the zinc, nickel (and iron), is boiled to 
 expel sulphuretted hydrogen, and a few drops of nitric acid added 
 to oxidise the iron, which is then precipitated as hydrated oxide by 
 means of ammonia, and estimated as Fe 2 O 3 . 
 
 The zinc is next precipitated as sulphide by means of hydrogen 
 sodium sulphide, HNaS, in the presence of potassium cyanide, and 
 estimated as ZnS.* 
 
 The filtrate is acidified with nitric acid and boiled to expel the 
 hydrogen cyanide, after which the nickel is precipitated as hydroxide 
 by means of potassium hydroxide, and estimated as NiO. 
 
 Tin. About i gram of German silver wire or borings is dis- 
 solved in strong nitric acid, and the mixture evaporated almost to 
 dryness upon a steam-bath. The residue is dissolved in water, and 
 the solution filtered to remove the traces of stannic oxide. After 
 being thoroughly washed, the filter is neglected, as the estimation 
 of the tin will be made in a separate larger quantity of the alloy, f 
 
 * For alternative methods of separating zinc and nickel, see p. 272. 
 
 f In the case of alloys in which tin forms one of the intended constituents, 
 as in the various bronzes, this residue should be treated as described on p. 272. 
 Approximately, the composition of German silver is Cu, 55 to 60 parts ; Zn, 22 
 to 27 ; and Ni, 15 to 20. In many samples the amount of tin is too small to 
 be determined. A qualitative analysis should be first made, in order to ascertain 
 whether the sample contains any tin and iron. 
 
Analysis of German Silver. 271 
 
 Copper. About 30 c.c. of strong hydrochloric acid are added 
 to the filtrate, and the solution considerably diluted with hot 
 water. It is then heated to boiling, and a stream of sulphuretted 
 hydrogen passed through for about 20 minutes. The liquid is 
 then filtered quickly, and the precipitate washed with warm water 
 containing sulphuretted hydrogen. 
 
 The copper sulphide so precipitated is liable to carry down with 
 it small quantities of the zinc, also as sulphide ; it is necessary, 
 therefore, to dissolve it and re-precipitate the copper. For this 
 purpose the precipitate is transferred to a porcelain dish, and 
 dissolved in strong nitric acid previously mixed with about half 
 its volume of saturated bromine water, and the solution evaporated 
 to dryness upon a steam-bath. 
 
 The residue is dissolved in about 30 c.c. of strong hydro- 
 chloric acid, and the solution rinsed into a beaker and considerably 
 diluted. It is once more boiled, and a stream of sulphuretted 
 hydrogen passed through it. The precipitated copper sulphide is 
 filtered and washed as before, with warm water containing sul- 
 phuretted hydrogen, the washing being continued until the wash- 
 water is free from hydrochloric acid. (Indicated by silver nitrate, 
 after boiling off the sulphuretted hydrogen.) 
 
 The precipitate is dried, and treated as described on p. 237. 
 
 Iron. The filtrates and washings from both precipitations of the 
 copper are mixed together and briskly boiled in a covered beaker, 
 in order to expel the sulphuretted hydrogen, after which the liquid 
 is evaporated until the volume is reduced to about 200 c.c. 
 
 Two or three drops of nitric acid are introduced in order to 
 oxidise the iron, after which a moderate quantity of ammonium 
 chloride is added, and then a slight excess of ammonia. The pre- 
 cipitated ferric hydroxide is filtered and washed with a little hot 
 water. It is then dissolved in a little hot hydrochloric acid, and 
 re-precipitated with ammonia, and again filtered and thoroughly 
 washed. The precipitate is neglected, as the iron, if only present 
 in traces, is estimated in a separate larger quantity of the alloy. 
 
 Zinc. The two filtrates from the iron precipitations are mixed 
 and boiled until the steam no longer smells of ammonia, and the 
 liquid concentrated somewhat by evaporation. A freshly made 
 concentrated solution of potassium cyanide is then added, until 
 the yellowish colour produced by the soluble double cyanide of 
 potassium and nickel becomes no deeper by the further addition 
 of the reagent. Hydrogen sodium sulphide, HNaS, is now added, 
 and the liquid heated to boiling. 
 
272 Typical Gravimetric Separations. 
 
 The precipitation of the zinc sulphide, which begins almost 
 directly the liquid is heated, is complete after the mixture has 
 boiled for a few minutes. The mixture is filtered, and the pre- 
 cipitate washed with hot water, and estimated as zinc sulphide, as 
 described on p. 244. 
 
 Nickel. The solution containing the double cyanide of potas- 
 sium and nickel, with excess of sodium sulphide, is evaporated 
 down to a moderately small bulk,* and strong nitric acid cautiously 
 added drop by drop until there is no further effervescence. Both 
 the cyanides and the sodium sulphide are thereby decomposed ; 
 the former with evolution of hydrocyanic acid (hence the operation 
 should be conducted in a draught-cupboard), while from the latter 
 sulphur is precipitated, which is at once attacked by the acid, with 
 the evolution of brown fumes. The mixture is boiled until the 
 precipitated sulphur is entirely oxidised to sulphuric acid, after 
 which potassium hydroxide is added until the precipitation of the 
 hydrated oxide of nickel is complete. 
 
 The precipitate is then treated as described on p. 247. 
 
 Separate Estimation of Tin and Iron, For this purpose, 
 about 5 grams of the alloy are weighed out and dissolved in nitric 
 acid, and the solution evaporated and treated as described above. 
 The residue of stannic oxide is filtered off, and estimated as on 
 p. 270. 
 
 The filtrate is treated as described above for the separation of 
 copper ; in this case, however, one precipitation only is necessary. 
 The precipitate, after being thoroughly washed, is neglected. The 
 filtrate and washings are together boiled to expel the sulphuretted 
 hydrogen, and the iron precipitated as hydrated oxide by means 
 of ammonia. The precipitate is treated as described on p. 230. 
 
 Alternative methods for the Separation of Zinc and 
 Nickel, (i) Precipitation of Zinc Sulphide from Acetic Acid solu- 
 tion. After the removal of the tin, copper, and iron as described 
 above, the liquid containing the zinc and nickel is acidified with 
 hydrochloric acid, and evaporated down until its volume is reduced 
 to about 150 c.c. Sodium carbonate solution is then added 
 until a slight precipitate persists, after which acetic acid is added 
 until the precipitate is redissolvecl.f Sulphuretted hydrogen is then 
 
 * If the precipitation of the zinc has been made in too dilute a solution, a 
 further slight separation of ZnS may take place during this concentration. In 
 that case the liquid must be filtered, and the precipitate either added to the 
 main portion, or else separately weighed and the result added. 
 
 t Or the precipitate may be redissolved by adding hydrochloric acid drop by 
 drop, and then removing the free acid by the addition of a few drops of a 
 
Separation of Zinc arid Nickel. 273 
 
 passed through the mixture, when the white zinc sulphide is alone 
 precipitated. When precipitation is complete, the liquid may be 
 gently warmed and filtered. The precipitate is then treated as 
 already described. The nitrate is heated to expel sulphuretted 
 hydrogen, and the nickel precipitated as hydrated oxide by means 
 of potassium hydroxide, and the process carried on as above. 
 
 (2) Precipitation of Zinc Sulphide from a Succinic Acid solution. 
 In order to carry out the separation by this method, the iron 
 must previously have been removed, not by means of ammonia, but 
 with sodium acetate.* 
 
 The solution containing the iron, zinc, and nickel (after separa- 
 tion of the copper as described on p. 271) is boiled to expel 
 sulphuretted hydrogen, and oxidised with nitric acid as before. It 
 is then made nearly neutral with sodium carbonate ; sodium acetate 
 is added, and the liquid boiled until the iron is completely precipi- 
 tated as basic acetate (see footnote, p. 70). The precipitate is 
 filtered off and thoroughly washed, after which it may be neglected.f 
 
 The filtrate and washings are together boiled with hydrochloric 
 acid until the acetates are entirely destroyed and the steam no 
 longer smells of acetic acid, after which the zinc and nickel are 
 together precipitated as carbonates by the addition of sodium 
 carbonate. The liquid is filtered and the precipitate washed, after 
 which it is dissolved upon the filter by slowly pouring over it a hot 
 strong solution of succinic acid. The funnel should be kept covered 
 with a clock-glass to avoid loss due to effervescence. When the 
 precipitate is completely dissolved, and the filter thoroughly washed 
 with hot water, sulphuretted hydrogen is passed through the liquid 
 while still hot, until the zinc is completely precipitated. The zinc 
 sulphide thrown down under these circumstances settles very readily. 
 The sulphide is filtered off, and treated as already described. 
 
 The solution is boiled to expel sulphuretted hydrogen, and the 
 nickel precipitated as in the former methods by means of potassium 
 hydroxide. 
 
 solution ot ammonium acetate, which, by double decomposition, forms 
 ammonium chloride and acetic acid. More than quite a small quantity of the 
 acetate must be carefully avoided, or else the nickel will be partially pre- 
 cipitated by the sulphuretted hydrogen. 
 
 * The reason for this is, that the precipitation of zinc as carbonate, by 
 means of sodium carbonate, is not complete in the presence of ammoniacal 
 salts. 
 
 f In cases where iron is present in sufficient quantity to be estimated at 
 this stage, this precipitate of basic acetate must be dissolved in hydrochloric 
 acid ; and, after boiling, the iron is precipitated as hydrated oxide by means 
 of ammonia, and treated as already described. 
 
 T 
 
274 Typical Gravimetric Separations. 
 
 Analysis of Bronze Coinage. 
 
 (Alloy of copper, tin, and zinc , with traces of lead.) 
 
 Epitome of Process. The alloy is dissolved in nitric acid, 
 and the solution evaporated to dryness. The residue is extracted 
 with water, and the insoluble stannic oxide dried and weighed. 
 
 Sulphuric acid is added to the filtrate, which is then evaporated 
 to a small bulk, and the lead sulphate separated and weighed. 
 
 The solution is diluted with water to a definite volume, and a 
 small measured proportion of it taken in which to estimate the 
 copper ; the copper in this separate portion being precipitated as 
 cuprous thiocyanate. 
 
 In the main portion of the solution, the copper is separated from 
 the zinc by means of sulphuretted hydrogen (the precipitated 
 copper sulphide being neglected), and the zinc is estimated in the 
 nitrate. 
 
 Tin. A clean new halfpenny * is weighed, and dissolved in 
 strong nitric acid in a covered beaker. The solution is then trans- 
 ferred to a porcelain dish, and evaporated to dryness on a steam- 
 bath. The residue is treated with water, and the insoluble stannic 
 oxide filtered and thoroughly washed. It is dried, and treated as 
 described on p. 270. 
 
 The stannic oxide, however, is liable to retain traces of the 
 copper ; it should therefore be submitted to the following treatment 
 after it has been weighed : A small quantity of sulphur (about 
 equal to the weight of the stannic oxide) and the same amount of 
 sodium carbonate are mixed together in the crucible, which with 
 its cover is then gently heated until the mass fuses and the excess 
 of sulphur is expelled. When cold, the mass is boiled with water, 
 which dissolves the sodium thiostannate, and if any copper were 
 present, it remains as a black residue of copper sulphide. It is 
 filtered, thoroughly washed, and dissolved in nitric acid. The 
 copper is then precipitated with potassium hydroxide, and estimated 
 as CuO, the weight being deducted from the weight of the stannic 
 oxide. 
 
 Lead.t The filtrate from the stannic oxide is acidulated by 
 the addition of about 10 c.c. of strong sulphuric acid (free from 
 
 * English bronze coinage consists roughly of Cu 95, tin 4, zinc i. A half- 
 penny weighs about 5 grams. This weight, while affording enough material 
 in which to estimate the zinc, is much more than is necessary for the copper 
 determination, hence the latter is estimated in a small portion of the solution. 
 
 t It is seldom that bronze coinage contains more than from 0^05 to o'i per 
 cent, of lead. In 5 grams of the alloy, o'i per cent, of lead would represent 
 0-005 gram of the metal, or o '007 gram of PbSO 4 , a quantity far too small to be 
 estimated, and scarcely a perceptible turbidity will be seen. This step in the 
 analysis may therefore be omitted without any appreciable influence upon the 
 result. It is given in the description because other varieties of bronze contain 
 appreciable quantities of lead, which would be separated as here explained. 
 
Analysis of Bronze Coinage. 275 
 
 lead),* and evaporated until the liquid gives off vapours of sul- 
 phuric acid freely. It is then allowed to cool, diluted with water, 
 and filtered. The precipitated lead sulphate is washed with dilute 
 sulphuric acid (in which it is far less soluble than in water) until 
 the washings are free from copper, when the vessel containing the 
 nitrate is removed, and the filter is washed free from sulphuric 
 acid by means of methylated spirit. (If the acid were not thus 
 washed out, the paper would become charred, and impossible to 
 manipulate, when dried in the oven. The alcoholic washings are 
 thrown away.) 
 
 The lead sulphate is dried, and treated as described on p. 242. 
 
 Copper. The filtrate containing the copper and zinc is diluted 
 with water in a half-litre flask until its volume is 500 c.c. By 
 means of a graduated pipette, 50 c.c. of this solution are with- 
 drawn and transferred to a beaker, and the copper precipitated as 
 cuprous thiocyanate by means of ammonium thiocyanate in the 
 manner described on p. 237. 
 
 The amount of copper present in this 50 c.c. of the solution 
 will obviously be one-tenth of that present in the entire volume, 
 and if any copper was recovered from the stannic oxide in the first 
 separation, this must be added on in order to obtain the total weight 
 of copper in the quantity of alloy taken for the analysis. 
 
 Zinc. The main portion of the solution (450 c.c.) containing 
 the copper and zinc, is acidified by adding about 50 c.c. strong hydro- 
 chloric acid, and diluted by the addition of 300 c.c. of hot water. 
 The mixture is heated to boiling, and the copper completely pre- 
 cipitated by means of sulphuretted hydrogen. 
 
 The precipitated copper sulphide is treated as described on 
 p. 271, the double precipitation being even more necessary in this 
 example, where the proportion of copper to that of zinc is so large. 
 The final precipitate of copper sulphide is here neglected, the copper 
 having been already estimated. 
 
 The solutions containing the zinc are then evaporated down to 
 moderate bulk, and the zinc estimated as on p. 244. 
 
 The weight of the zinc so obtained will be nine-tenths of that 
 present in the amount of alloy employed for the analysis. 
 
 * Ordinary sulphuric acid contains considerable quantities of lead. If pure 
 acid is not available, acid must be used which has been previously largely 
 diluted and allowed to settle, and a proportionately larger volume will be 
 necessary. 
 
276 Typical Gravimetric Separations. 
 
 Analysis of Dolomite. 
 
 {Magnesium and calcium carbonate, containing usually small 
 quantities of iron, aluminium, and insoluble silicious matter}.* 
 
 Epitome of Process. The mineral is dissolved in hydro- 
 chloric acid, and the solution evaporated to dryness. The residue 
 is treated with hydrochloric acid and water, and the insoluble por- 
 tion, consisting of silica and silicates, is filtered off, washed, and dried. 
 
 The iron and alumina are precipitated together as hydroxides, by 
 the addition of ammonium chloride and ammonia to the filtrate. The 
 precipitate is dried, and weighed as oxides of iron and aluminium. 
 
 The filtrate is boiled, and the calcium precipitated as oxalatc 
 by means of ammonium oxalate, the precipitation being repeated 
 to ensure complete separation of the magnesium. The precipitate 
 is dried and converted by heat into calcium oxide, and weighed. 
 
 The filtrate from the calcium is evaporated to dryness, and 
 heated to expel excess of ammonium salts. The residue is dis- 
 solved in hydrochloric acid, diluted with water, and the magnesium 
 precipitated as ammonium magnesium phosphate by the addition 
 of ammonia and sodium phosphate. 
 
 The carbon dioxide is estimated in a separate portion of the 
 mineral by the method described on p. 259. 
 
 Silicious Matter. About 2 grams of the finely powdered 
 mineral (which has been previously dried in an air-oven at a 
 temperature of about 200 t) is weighed out into a porcelain dish 
 covered with a clock-glass, and dissolved by the cautious addition 
 of dilute hydrochloric acid. As soon as action ceases, the cover is 
 rinsed into the dish, a few drops of strong nitric acid added, and 
 the mixture evaporated to complete dryness upon a steam-bath. 
 The residue is then treated with a little strong hydrochloric acid, 
 and gently warmed. Water is then added, and the insoluble residue 
 of silica and insoluble silicates is transferred to a filter, where it is 
 washed free from hydrochloric acid. It is then dried in the 
 steam-oven. 
 
 The dried filter with the residue is folded up and placed in a 
 platinum crucible, where it is strongly heated until the paper is com- 
 pletely incinerated.^ The crucible, after cooling in the desiccator, 
 is weighed. 
 
 * Occasionally manganese is met with in varieties of dolomite, but usually 
 only in traces, which may be neglected. 
 
 f The amount of moisture present in dolomite is usually very small. If 
 desired, it may be estimated by heating about 10 grams of the powder to 200 
 in the air-oven. 
 
 % Usually the amount of silicious matter does not exceed about two per 
 cent. In cases where there is much insoluble residue, it should be detached 
 from the paper, and the latter incinerated apart in a platinum coil, 
 
Analysis of Dolomite. 277 
 
 Aluminium and Iron.* A considerable quantity of ammo- 
 nium chloride is added to the filtrate (see note on p. 224), and the 
 solution gently warmed ; ammonia is then added, the least pos- 
 sible excess being used, and the mixture boiled. The precipitated 
 hydroxides of aluminium and iron are then filtered and washed ; 
 the precipitate is dissolved upon the filter by pouring a little warm 
 dilute hydrochloric acid upon it (the liquid being received in a 
 separate beaker), and the acid solution thoroughly washed out of 
 the paper. 
 
 The aluminium and iron are then re-precipitated from this solu- 
 tion by the addition of a slight excess of ammonia ; the precipitate 
 is filtered and washed until the washings are free from chlorides, 
 and then dried in the steam-oven. The dry precipitate, if small in 
 amount, is heated with the paper in a platinum crucible ; if the 
 amount is considerable, it is detached from the filter, and the latter 
 is separately incinerated in a platinum coil. The precipitate, after 
 being heated, consists of a mixture of A1 2 O 3 and Fe 2 O 3 ; it is 
 weighed and calculated as such.f 
 
 Calcium. The mixed filtrates and washings from the double 
 precipitation of iron and aluminium are heated to boiling, and the 
 calcium precipitated by means of ammonium oxalate, as described 
 on p. 231. Excess of the reagent must in this case be used, since, 
 in the presence of magnesium chloride, the precipitation of cal- 
 cium oxalate is not complete unless excess of ammonium oxalate is 
 present. The reagent may be added in a solid state (instead of a 
 hot strong solution, as directed on p. 231), three or four grams of 
 the powdered salt being introduced into the boiling liquid, and the 
 mixture stirred for a few minutes. 
 
 The precipitated calcium oxalate carries down with it small 
 quantities of magnesium oxalate ; it is therefore necessary to repeat 
 the precipitation. For this purpose the precipitate is allowed to 
 settle, and the clear liquid decanted off through a filter without 
 disturbing the precipitate ; the latter is twice washed by decanta- 
 tion, each time being allowed to thoroughly settle, so that the least 
 possible trace of it only is trr-.nsferred to the filter. It is then 
 dissolved in the beaker by the addition of a little warm dilute 
 hydrochloric acid, and reprecipitated by the addition of ammonia 
 
 * Should the ore contain more than very small quantities of manganese 
 (which, however, is not usually the case), the iron and aluminium must be 
 separated as basic acetates, as described for iron on p. 282. 
 
 f If the quantity of aluminium and iron is considerable, and it is desired 
 to determine the metals separately, the precipitate may be treated as described 
 on p. 278. 
 
278 Typical Gravimetric Separations. 
 
 in slight excess, and a few drops of ammonium oxalate solution.* 
 This precipitate is allowed to settle, and transferred to the same 
 filter employed in the first operation ; it is washed and dried, and 
 treated as described on p. 233. 
 
 Magnesium. The filtrates and washings from the calcium 
 precipitate are evaporated to dryness in a porcelain dish. The 
 evaporation may be conducted at first over a gas-burner, provided 
 the liquid is not permitted to boil ; but as the liquid begins to 
 deposit salts, the further heating must be done upon a steam- 
 bath.f The dish with the dry residue is heated over a ring 
 gas-burner, at first very cautiously, until the ammonium salts are 
 expelled. 
 
 When the dish is cold, the residue is dissolved by the addition 
 of a little strong hydrochloric acid, and gently warming the mix- 
 ture. The solution is diluted with water, and if not perfectly clear 
 it is filtered, the residue in this case being neglected. 
 
 Ammonia in slight excess is added (no precipitate should form, 
 as the ammonium chloride produced by the neutralisation of the 
 solution will be sufficient to prevent the precipitation of the 
 magnesium hydroxide), and the magnesium then precipitated as 
 ammonium magnesium phosphate by the addition of hydrogen 
 disodium phosphate. The precipitation and the treatment of the 
 precipitate are described on p. 233. 
 
 NOTE. Iron and aluminium are associated together in varying 
 proportions in a great number of common minerals which come 
 under analysis, such as iron ores, clays, felspar, and other sili- 
 cates, as well as slags and other artificial products. They may be 
 separated and determined by the following methods : 
 
 Separation of Iron and Aluminium. (i) The precipi- 
 tated hydroxides are dissolved upon the filter in a little warm 
 hydrochloric acid, and the solution poured into a strong solution 
 of sodium or potassium hydroxide (free from alumina) contained 
 in a dish (preferably of either silver, platinum, or nickel, although 
 porcelain may be used in the absence of these). The mixture is 
 then boiled for two or three minutes. The iron is precipitated 
 as ferric hydroxide, while the aluminium passes into solution as 
 sodium aluminate. The precipitate is filtered off, washed with 
 
 * Or the second precipitation may be conducted as in the case of the iron 
 and aluminium ; i.e. the precipitate may be washed upon the filter, and dis- 
 solved by the addition of acid to it while in the funnel. 
 
 f Solutions of ammoniacal salts are very apt, during evaporation, to be 
 drawn up through the crust of deposit which forms round the edge of the 
 liquid, so that the salt is continually deposited higher and higher upon the sides 
 of the dish, until eventually it creeps over the edge. To prevent this, the 
 finger, slightly greased with vaseline, should be drawn round the inside edge 
 of the dish at' the beginning of the operation. 
 
Analysis of Dolomite. 279 
 
 water, then redissolved in hydrochloric acid and reprecipitated 
 with ammonia.* The ferric hydroxide is again filtered, washed 
 and dried in the usual way, and weighed in the form of sesqui- 
 oxide (p. 230). The two filtrates are together acidified with strong 
 hydrochloric acid, and the aluminium precipitated as hydroxide 
 by the addition of ammonia in slight excess. The precipitate is 
 washed and dried, and weighed in the form of sesquioxide, as 
 described on p. 225. 
 
 (2) When pure caustic soda or potash (i.e. free from alumina) 
 is not at hand, the process may be modified in the following way, 
 and the aluminium estimated by difference. The precipitated 
 hydroxides of the two metals are dried and weighed together as 
 sesquioxides. A quantity of powdered sodium peroxide f is then 
 mixed with the precipitate in the crucible (platinum), and the 
 mixture fused. The fused mass when cold is cautiously treated 
 with water in a beaker (as the aluminium is not to be directly 
 determined, glass vessels may be used), when the sodium aluminate 
 will dissolve, leaving the ferric oxide. This is then filtered and 
 washed, the filtrate being rejected. The precipitate is then dis- 
 solved in hydrochloric acid, reprecipitated with ammonia, and 
 determined in the usual way. The weight of the ferric oxide thus 
 obtained, deducted from the weight of the two oxides together, 
 gives the weight of the aluminium sesquioxide. 
 
 (3) The amount of iron may be determined, and the aluminium 
 estimated by difference, by the use of volumetric methods for 
 estimating iron. The mixed oxides are weighed as in the above 
 process ; they are then dissolved in hydrochloric acid, and the 
 iron determined volumetric ally,:}: without separating the aluminium. 
 The amount of sesquioxide of iron corresponding to the iron thus 
 determined, when deducted from the original weight of the mixed 
 oxides, gives the weight of the sesquioxide of aluminium. 
 
 Analysis of Zinc-blende. 
 
 (Zinc sulphide, liable to' contain zinc carbonate ; and lead, 
 copper, cadmium, iron, and manganese, as sulphides er in other 
 forms. A previous qualitative analysis must be made.} 
 
 Epitome of Process. The ore is decomposed by strong 
 hydrochloric and nitric acids. Any lead present is converted into 
 sulphate and weighed along with the silicious matter, or gangue, 
 from which it is afterwards separated by solution in ammonium 
 acetate. 
 
 Copper and cadmium, if present, are precipitated together from 
 the solution as sulphides, which are afterwards separated by 
 dissolving the latter in dilute sulphuric acid. 
 
 * This second precipitation is necessary in order to remove the potash 
 which is retained by the ferric hydroxide. 
 
 t Caustic soda may be substituted, although the peroxide is preferable. 
 
 See Volumetric methods. 
 
280 Typical Gravimetric Separations. 
 
 Iron is next separated in the form of the basic acetate,* the 
 sulphates present having been first converted into chlorides by 
 means of barium chloride. 
 
 The manganese is precipitated as hydrated peroxide by means 
 of chlorine or bromine ; and, lastly, the zinc is determined by 
 precipitation as zinc sulphide. 
 
 The sulphur in the ore is estimated in a separate portion 
 by oxidation into sulphuric acid, and subsequent precipitation as 
 barium sulphate. 
 
 If the ore contains any calamine (zinc carbonate), the carbon 
 dioxide is determined in a separate portion by the method described 
 on p. 259. 
 
 Sulphur. About I gram of the ore (which has been reduced to 
 as fine a powder as possible, and dried in a steam-oven) is weighed 
 into a platinum crucible, and then mixed thoroughly with five or six 
 times its weight of powdered sodium peroxide, the mixing being done 
 by means of a thin glass rod or stout platinum wire. The covered 
 crucible is then cautiously heated by means of a small flame placed 
 some little distance below it. Considerable action takes place, the 
 contents of the crucible undergoing incipient deflagration, and 
 becoming liquid. It is kept in a state of fusion for a few minutes, 
 and allowed to cool. The crucible is then placed upon its side 
 in a beaker, along with the lid, and water cautiously added, the 
 beaker being covered to prevent loss during the effervescence 
 which takes place owing to the escape of oxygen from the excess 
 of peroxide. 
 
 When the solid is detached from the crucible, the latter is 
 lifted out with a clean pair of tongs and rinsed with water ; the 
 lid is removed and washed into the beaker in the same way. The 
 solution is then acidified with hydrochloric acid, filtered if neces- 
 sary, and the sulphuric acid precipitated as barium sulphate by 
 means of barium chloride in the usual manner (p. 258).t From 
 
 * If there is no manganese in the ore, the iron and zinc may be separated 
 by precipitation of the former by means of ammonia, according to the method 
 described on p. 277. 
 
 f Instead of employing sodium peroxide as the oxidising agent, strong 
 nitric acid may be used. For this purpose the powdered ore is digested in a 
 flask at a gentle heat (see Fig. 39) with strong nitric acid, to which, as the 
 operation proceeds, a little hydrochloric acid or a few crystals of potassium 
 chlorate are added; or fuming nitric acid (sp. gr. 1-5) may be employed 
 alone. When the action is complete, and the whole of the sulphur (which 
 at first often separates out and floats about in liquid drops) is oxidised into 
 sulphuric acid, a little pure sodium chloride is added. This converts the 
 sulphuric acid into sodium sulphate, and so prevents its volatilisation during 
 the subsequent process of evaporation. The solution is then evaporated in a 
 porcelain dish upon a steam-bath. When nearly dry, a little strong hydro- 
 chloric acid is added, and the mixture evaporated to dryness. The residue 
 
Analysis of Zinc-blende. 281 
 
 the weight of barium sulphate obtained, the proportion of sulphur 
 is calculated. Factor 
 
 (BaS0 4 )2 33 : (8)32=1:0-1375 
 
 Silicious Matter and Lead. About 2 grams of the dry 
 and finely powdered ore are weighed out into a beaker, and 20 to 
 25 c.c. of strong hydrochloric acid added. The beaker should be 
 immediately covered, since if any carbonate is present there will 
 be effervescence. The mixture is then gently boiled until sul- 
 phuretted hydrogen is no longer given off, when about 20 c.c. of 
 strong nitric acid are added, and the heating continued until the 
 decomposition of the ore is complete. The contents of the beaker 
 are then rinsed into a porcelain dish and evaporated to dryness 
 upon a steam-bath, after which the dish with the dry residue is 
 placed in an air-oven and heated to about 160 to render the silica 
 insoluble. The residue is next moistened with strong hydrochloric 
 acid, and about 5 c.c. of strong sulphuric acid, previously diluted 
 with twice its own volume of water, added. The mixture is then 
 cautiously heated on a sand-bath until fumes of sulphuric acid are 
 given off, by which time all the nitrates and chlorides will have 
 been converted into sulphates, and their acids expelled. The dish 
 is then allowed to cool, water is added, and the mixture filtered. 
 The residue, consisting of silicious matter and lead sulphate, is 
 washed with water acidulated with sulphuric acid, the filtrate and 
 washings being set aside for the estimation of the remaining 
 constituents. 
 
 In order that the residue may be dried without the paper be- 
 coming charred, the dilute acid is first washed out of it by means 
 of alcohol, the washings being neglected. The dried residue is 
 then transferred to a porcelain crucible, and the incineration of the 
 paper conducted as described in the estimation of lead as sulphate 
 (p. 242). The weight obtained is that of the silicious matter and the 
 
 is again treated with hydrochloric acid, and once more taken down to dryness 
 in order to completely expel the nitric acid. It is finally dissolved in water 
 with a little hydrochloric acid, filtered, and diluted to about 100 c.c. The 
 sulphuric acid is then precipitated as barium sulphate in the usual way. 
 The residue may contain lead sulphate along with the gangue. If, therefore, 
 lead is present in the ore, this residue must be treated as described above, 
 and the amount of sulphur which has been thus carried down in combination 
 with the lead must be added to the total weight of sulphur. 
 
 The sulphur in natural sulphides is usually estimated either by this process 
 or by oxidation with sodium peroxide, the latter being the quicker process. 
 Oxidation of the sulphur may, however, be effected by other methods, such 
 as by fusion with a mixture of sodium carbonate and potassium nitrate, or by 
 digesting the ore in a strong solution of caustic potash, and passing a stream 
 of chlorine throxigh the mixture. 
 
282 Typical Gravimetric Separations. 
 
 lead sulphate together. The contents of the crucible are then 
 transferred to a small beaker, and gently boiled with a solution of 
 ammonium acetate and ammonia ; it is allowed to settle, and the 
 liquid decanted through a small filter. It is treated to succes- 
 sive extractions with fresh portions of the ammonium acetate 
 solution until the whole of the lead sulphate is dissolved out, the 
 completion of the process being ascertained by testing a drop of 
 the filtrate with ammonium sulphide. 
 
 The residue is then dried, and the paper incinerated either in 
 the crucible along with the residue, or upon a platinum spiral. The 
 result gives the weight of the silicious matter alone, the difference 
 between this and the former weighing being the proportion of lead 
 sulphate, from which the percentage of lead is deduced by the factor 
 on p. 242. 
 
 Copper and Cadmium. To the solution containing the re- 
 maining metals,* 10 to 15 c.c. of strong hydrochloric acid are 
 added, and sulphuretted hydrogen is passed through until precipita- 
 tion is complete. If cadmium is absent, the precipitate is treated 
 as described on p. 236 (Estimation of copperas sulphide). If, how- 
 ever, the ore contains cadmium, the precipitate of the mixed sulphides, 
 after being thoroughly washed, is boiled with dilute sulphuric acid 
 (strong acid i part, water 5 parts). It is then filtered, and the in- 
 soluble copper sulphide washed and dried, and treated as described 
 on p. 237. The filtrate and washings containing the cadmium are 
 neutralised with ammonia, one or two drops of hydrochloric acid 
 added, and the cadmium precipitated as sulphide by means of 
 sulphuretted hydrogen. The precipitate is treated as described on 
 p. 257. 
 
 Iron.t 1*he filtrate from the first precipitation with sulphuretted 
 hydrogen is boiled until the gas is entirely expelled, and a few 
 drops of nitric acid are added in order to re-oxidise the iron. 
 Barium chloride is then added so long as a precipitate is produced, 
 and the barium sulphate filtered off and washed. The precipitate 
 is neglected. In this way the sulphates present are converted into 
 chlorides. The solution is then nearly neutralised with ammonium 
 carbonate, by adding the reagent cautiously until a slight precipitate 
 persists, and then adding hydrochloric acid drop by drop with 
 
 * This solution should not be less than about 1500.0. in volume. If the 
 washings from the insoluble residue do not bring It up to this, it must be 
 diluted with water. 
 
 t If the ore contains no manganese, the iron may be precipitated as 
 hydroxide by means of ammonia (as described on p. 230) after the sulphuretted 
 hydrogen has been boiled off and the iron re-oxidised with nitric acid. 
 
Analysis of Zinc-blende. 283 
 
 constant stirring, until the precipitate is almost completely re- 
 dissolved, leaving the solution just slightly turbid.* A strong 
 solution of ammonium acetate, acidified with acetic acid, is then 
 added, and the solution, which assumes a red colour, is boiled for a 
 few minutes, when the basic ferric acetate separates as a voluminous 
 reddish precipitate. This is allowed to settle, and the colourless 
 liquid decanted through a filter, to which the precipitate is then 
 transferred and washed with hot water.f The precipitate is then 
 dissolved in hydrochloric acid, the solution again neutralised with 
 ammonium carbonate, and the iron reprecipitated with ammonium 
 acetate, filtered, and washed. The filtrate and washings from this 
 second precipitation (which is necessary for the complete separation 
 of the manganese) are added to those from the first operation. The 
 precipitate is dried, and the filter incinerated with the precipitate, 
 which is then strongly heated in the crucible until it is completely 
 converted into the sesquioxide, Fe 2 O 3 , in which form it is weighed.^ 
 
 Manganese. The filtrates from the double precipitation of 
 the iron are together evaporated down until the volume is reduced 
 to about 100 c.c. Ammonia is then added until the solution is 
 strongly alkaline, and the mixture heated nearly to boiling. Satu- 
 rated bromine water is then added to the boiling liquid until the 
 precipitation of the hydrated peroxide is complete, and the solution 
 is distinctly yellow. The precipitate is then filtered off, and washed 
 and dried. The dried precipitate, after the incineration of the 
 paper, is strongly heated in the crucible until it is completely con- 
 verted into the tetroxide, Mn 3 O 4 , in which form it is weighed.]) 
 
 Zinc. The alkaline filtrate from the manganese separation is 
 saturated with sulphuretted hydrogen (or colourless ammonium 
 sulphide may be added), and the zinc estimated as sulphide, as 
 
 * The least excess of hydrochloric acid must be avoided. 
 
 f This precipitate is difficult to wash, especially if it has been boiled too 
 long, when it is apt to become slimy. In cases where there is a considerable 
 quantity of the precipitate, the funnel should be kept hot by one of the arrange- 
 ments described on p. 210. 
 
 % When sodium acetate is used for the precipitation instead of the ammo- 
 nium salt, the basic ferric acetate must be dissolved in hydrochloric acid, and 
 the iron precipitated as hydroxide by means of ammonia. 
 
 The manganese cannot satisfactorily be precipitated as the carbonate from 
 a solution containing considerable quantities of acetates. 
 
 || If sodium acetate has been employed in the previous separation of the 
 iron instead of ammonium acetate, the process for the estimation of the 
 manganese must be modified, since the hydrated peroxide retains portions of 
 the alkali which are not removed by washing, and being non-volatile they 
 remain behind with the manganese in the crucible. In this case the hydrated 
 peroxide must be dissolved in hydrochloric acid, and the manganese precipitated 
 as carbonate by means of sodium carbonate (p. 245). 
 
284 Typical Gravimetric Separations. 
 
 described on p. 244. Or the precipitated zinc sulphide, after being 
 filtered off, may be dissolved in dilute hydrochloric acid, the solution 
 boiled, and the zinc precipitated as carbonate by means of sodium 
 carbonate, and estimated in the form of zinc oxide according to 
 the method given on p. 243, or by the volumetric method described 
 on p. 371. 
 
 NOTE. The analyses of other natural sulphides, such as copper- 
 pyrites, iron-pyrites, galena, etc., are conducted on the same general 
 lines as are described above, except that certain necessary modifi- 
 cations have to be introduced depending upon the particular metals 
 present in the ore. Thus, copper and iron pyrites are liable to 
 contain arsenic, or arsenic and antimony. These are precipitated 
 as sulphides along with the copper, and afterwards separated from 
 the copper sulphide by solution in sodium sulphide.* From this 
 solution the sulphides of arsenic and antimony are reprecipitated 
 by acidifying with hydrochloric acid and passing sulphuretted 
 hydrogen, and the elements separated by the following method. 
 
 Separation of Arsenic and Antimony. The two sulphides 
 are dissolved in strong hydrochloric acid, with the addition of 
 crystals of potassium chlorate ; tartaric acid and ammonium 
 chloride are added in considerable quantity, then an excess of 
 ammonia. (The addition of ammonia should cause no precipitation ; 
 should a precipitate form, it is due to insufficient tartaric acid and 
 ammonium chloride, and more must be added.) The arsenic is 
 then precipitated as ammonium magnesium arsenate f by means of 
 " magnesia mixture," and weighed as magnesium pyro-arsenate, as 
 described on p. 254. The antimony contained in the solution is 
 precipitated as antimonious sulphide by means of sulphuretted 
 hydrogen, and finally weighed in the form of the tetroxide, as 
 described on p. 256. 
 
 Galena often contains silver besides other metals. The ore in 
 this case is decomposed by strong nitric acid (sp. gr. 1*5), the lead 
 converted into sulphate by means of sulphuric acid, and the lead 
 sulphate and silicious matter separated together. The silver is 
 then precipitated as silver chloride by means of hydrochloric acid, 
 and estimated as directed on p. 238, while the remaining metals, 
 such as copper, antimony, iron, zinc, which are liable to be present, 
 are separated as explained above. 
 
 * Sodium sulphide is prepared by saturating a solution of sodium hydroxide 
 with sulphuretted hydrogen, and then adding fresh sodium hydroxide solution 
 until the mixture ceases to smell of the gas (see " Reagents," Appendix). 
 
 f Precipitated under these conditions, especially if the quantity of arsenic is 
 large, the ammonium magnesium arsenate is liable to contain a little basic 
 magnesium tartrate. It is desirable, therefore, to dissolve the precipitate in 
 hydrochloric acid, and reprecipitate it by the addition of ammonia and the 
 magnesia solution. 
 
Analysis of an Insoluble Silicate. 285 
 
 Analysis of an Insoluble Silicate.* 
 
 (Containing the metals iron, aluminium, calcium, magnesium, 
 potassium, and sodium.} 
 
 Epitome of Process. The finely powdered silicate is fused 
 with alkaline carbonates (fusion mixture). The " melt " is extracted 
 with water, and the silicic acid precipitated with hydrochloric 
 acid, after which the mixture is evaporated to dryness and gently 
 heated. The insoluble silica is filtered off, dried, and weighed. The 
 iron and aluminium are together precipitated as hydroxides, and 
 afterwards separately determined by one of the methods given on 
 p. 278. 
 
 The calcium is next separated as calcium oxalate. The filtrate 
 is evaporated to dryness, and heated to expel ammonium salts ; the 
 residue is dissolved in water, and the magnesium precipitated as 
 ammonium magnesium phosphate. 
 
 The alkalies are estimated in a separate portion by fusion with 
 calcium carbonate and ammonium chloride, whereby they are con- 
 verted into soluble chlorides. The product is extracted with water ; 
 the soluble calcium salts are precipitated as oxalate, and after the 
 removal of ammoniacal salts, the mixed chlorides of potassium and 
 sodium are dried and weighed. The proportion of potassium 
 chloride in this residue is then determined as the double potassium 
 platinum chloride, and the sodium chloride estimated by difference ; 
 and from the results so obtained the percentage of potassium and 
 sodium oxides is calculated. 
 
 Silica. From 1*5 to 3 grams t of the silicate (which has been 
 reduced to the finest possible powder,:}: and dried ), are weighed out 
 into a platinum crucible of fairly large dimensions, and intimately 
 mixed with five or six times its weight of fusion mixture by means 
 
 * A large number of the common natural silicates (e.g. the felspars, micas, 
 hornblende, etc.) consist essentially of varying proportions of silicates of iron 
 and aluminium, calcium, and magnesium, and one or more of the alkali metals. 
 The various kinds of glass are also silicates containing similar constituents, 
 being composed mainly of silicates of sodium (or potassium) and calcium, 
 mixed with varying smaller quantities of aluminium, iron, and manganese. In 
 flint glass the calcium silicate is more or less entirely replaced by lead silicate. 
 
 f In the case of such a silicate as felspar, which may contain less than i per 
 cent, of some bases, a larger quantity must be employed for the analysis than 
 would be necessary in the case of a substance such as glass. 
 
 J It is absolutely necessary, in order to ensure the complete decomposition 
 of the silicate, that it should be extremely finely powdered. There should be 
 no grittiness to the touch when it is rubbed between the thumb and fingers, 
 and the whole sample should pass readily through a sieve of fine muslin. 
 
 The powdered silicate should be ireed from any adherent moisture by 
 being heated in the steam-oven before being weighed out for analysis. Some 
 silicates contain combined water (e.g. serpentine, meerschaum, etc. ). This should 
 be estimated by heating a weighed quantity of the powdered mineral to redness 
 in a platinum crucible until the weight is constant. 
 
286 Typical Gravimetric Separations. 
 
 of a stout platinum wire, or a thin glass rod with carefully rounded 
 ends. The entire mixture should not more than half fill the crucible. 
 The covered crucible is then heated over a Bunsen flame, at 
 first gently in order to expel the moisture present in the mixture, 
 and afterwards more strongly until the mass begins to melt round 
 the edges. It is then heated by means of a blowpipe, until the 
 decomposition is complete and the contents of the crucible are in 
 a state of quiet fusion. Care must be taken, by regulating the heat 
 of the blowpipe, to prevent undue frothing of the mass while the 
 carbon dioxide is being evolved ; the progress of the operation 
 should be watched by momentarily slightly raising the crucible lid 
 from time to time. 
 
 When the operation is complete, the crucible is allowed to cool 
 down, and when cold it is placed upon its side in a beaker with 
 about loo c.c. of water, care being taken that no impurities are 
 conveyed into the solution upon the outside of the crucible. The 
 beaker is heated upon an iron plate or sand-bath, and the water 
 allowed to boil gently until the " melt " is either entirely detached 
 from the crucible or has become honeycombed by the action of the 
 hot water. Hydrochloric acid is then cautiously added in small 
 portions at a time (the clock-glass cover being partially withdrawn 
 for each addition) until effervescence ceases, and no further pre- 
 cipitation of gelatinous silicic acid takes place. The crucible and 
 lid are then withdrawn * and rinsed into the beaker, t 
 
 The mixture is transferred to a dish, and evaporated to dry- 
 ness upon a steam-bath, the gelatinous mass, as it stiffens, being 
 stirred at frequent intervals with a short glass rod, in order to 
 break it up as much as possible and thus expedite the drying. 
 When the mass has become white and pulverulent, the dish is trans- 
 ferred to an air-bath, and heated to about 160 for half an hour. 
 
 The residue is then moistened with a little strong hydrochloric 
 acid, and digested upon the steam-bath for a short time, acid being 
 added as evaporation goes on.J Hot water is added, and the silica 
 
 * Ordinary brass or iron crucible tongs must not be dipped into this acid 
 liquid. In the absence of platinum or bone-tipped tongs, the crucible can be 
 lifted a little way out of the liquid by means of a glass rod, and the projecting 
 part first rinsed with a drop of water from the wash-bottle, after which the 
 crucible can be taken hold of with the tongs. 
 
 f If the fusion has been satisfactorily conducted, there should be no trace 
 of gritty particles at the bottom of the beaker. Should such be detected, it 
 will probably be due to imperfect powdering of the mineral, and in that case 
 the experiment must be set aside, and the operation repeated. 
 
 \ Oxides of iron and aluminium, after being heated, are less easily dissolved 
 by hydrochloric acid, hence the necessity for this step. 
 
Analysis of an Insoluble Silicate. 287 
 
 washed several times by decantation with hot water, after which it 
 is transferred to the filter and washed until the wash-water is free 
 from chlorides. 
 
 The silica is dried in the steam-oven, transferred to a platinum 
 crucible,* and the paper incinerated upon a platinum spiral. The 
 covered crucible is heated, at first very cautiously with a small 
 flame, and afterwards to a red heat, and weighed until constant.f 
 
 Iron and aluminium are precipitated together in the form of 
 hydroxides by the addition of ammonium chloride and ammonia to 
 the filtrate from the siHca. The precipitation is conducted as de- 
 scribed on p. 277, and the iron and aluminium are separately esti- 
 mated by either of the methods given on p. 278. From the results 
 obtained, the percentages of ferric oxide, Fe 2 O 3 , and aluminium 
 sesquioxide, A1 2 O 3 , present in the mineral are calculated. 
 
 Calcium. If the filtrate and washings from the precipitated 
 hydroxides of iron and aluminium are more in volume than 
 about 150 c.c., they should be evaporated down to that bulk, 
 and the calcium then precipitated as oxalate, as on p. 231. From 
 the result obtained, the percentage of calcium oxide, CaO, is 
 calculated. 
 
 Magnesium- In the filtrate from the calcium oxalate the 
 magnesium is precipitated in the form of ammonium magnesium 
 phosphate, as described on p. 233. 
 
 From the weight of magnesium pyrophosphate obtained, the 
 percentage of magnesium oxide, MgO, is calculated. Factor 
 
 (Mg 2 P 2 O 7 ) 222 : (2MgO) 80 = i : 0-3603 
 
 Potassium and Sodium. When the silicate has been 
 "opened up " by fusion with alkaline carbonates as above described, 
 it is obvious that the estimation of the alkalies cannot be made in 
 the solution so obtained ; these constituents must therefore be 
 
 * The silica obtained in this way is an extremely light powder, which is 
 easily blown away ; hence some care is necessary in the manipulation of it, in 
 order to avoid loss from this cause. 
 
 f It is always desirable to test the purity of the silica after it has been 
 weighed. In the case of silicates which are liable to contain small quantities 
 of tin, tungsten, niobium, or tantalum, the oxides of these metals will be ad- 
 mixed with the silica, and in such cases an examination of the silica is necessary. 
 For this purpose the contents of the platinum crucible are digested upon a 
 steam-bath with pure aqueous hydrofluoric acid and a few drops of strong 
 sulphuric acid, the operation being performed in a draught cupboard. The 
 silica is thus converted into silicon fluoride, which is expelled. Fresh hydrofluoric 
 acid is added once or twice as the liquid evaporates, after which the contents 
 of the crucible are evaporated to dryness, strongly heated, and weighed. The 
 residue is again submitted to the same treatment until the weight is constant. 
 The loss of weight represents the silica which was present. 
 
288 
 
 Typical Gravimetric Separations. 
 
 determined in a separate portion of the mineral, which is decomposed 
 by another method. About 1*5 to 2 grams of the powdered silicate 
 are weighed into a platinum crucible, and there intimately mixed 
 with about six times its weight of pure precipitated calcium 
 carbonate,* and about its own weight of pure ammonium chloride. 
 The crucible is then gradually raised to a bright red heat, and 
 maintained at that temperature for an hour. This may be accom- 
 plished by means of a powerful Bunsen, 
 a plumbago crucible with the bottom 
 out being inverted over the platinum 
 crucible, as shown in Fig. 6o.t The 
 crucible is afterwards allowed to cool, 
 and then placed in a covered porcelain 
 dish and digested with water upon a 
 sand-bath. The crucible is withdrawn 
 and rinsed, and the liquid filtered, the 
 residue being thoroughly washed. The 
 solution, which now contains the alkali 
 metals in the form of chlorides, is freed 
 from any lime salts which have dis- 
 solved, by the addition of ammonia, 
 ammonium carbonate, and a small 
 quantity of ammonium oxalate. The 
 precipitate is filtered off and washed, 
 and the filtrate evaporated to dryness 
 in a platinum dish upon a steam-bath. The residue in the dish 
 is heated over a small rose burner without raising the temperature 
 to redness, in order to expel all the ammonium salts. It is then 
 dissolved in a small quantity of water, and the last traces of lime 
 which may have escaped precipitation are thrown down by the 
 addition of one or two drops of ammonia and ammonium oxalate. 
 The solution is passed through a small filter, which must be after- 
 wards thoroughly washed, the filtrate and washings being received 
 in a weighed platinum dish. A few drops of hydrochloric acid are 
 added, and the liquid is then evaporated to dryness, heated to 
 expel ammonia, and weighed. It is again heated and weighed, until 
 the weight of the mixed chlorides of the alkali metals is constant. 
 
 * Prepared by the addition of ammonium carbonate to a solution of barium 
 chloride, and thoroughly washing and drying the precipitate. 
 
 f The blowpipe is not so suitable for the purpose, partly on account of the 
 tediousness of using a blowpipe for so long, but more particularly because of 
 the liability to overheat portions of the mixture, and the subsequent risk of loss 
 of alkaline chlorides by volatilisation. 
 
 FIG. 60. 
 
Estimation of Alkalies. 289 
 
 In order to estimate the relative proportions of the potassium 
 and sodium chlorides, the residue is dissolved in a very small 
 quantity of water, and the potassium precipitated and weighed as 
 the double potassium platinic chloride, as described on p. 249. 
 From the weight of the double salt obtained, the weight of potassium 
 chloride is calculated. Factor 
 
 (2KCl,PtCl 4 ) 486 : (2KC1) 149 - i : 0-3065 
 
 And on deducting the weight of potassium chloride so obtained, 
 from the weight of the mixed chlorides, the proportion of sodium 
 chloride is found. 
 
 Having found the weights of the two separate chlorides, the 
 percentages of potassium and sodium oxides (K 2 O and Na 2 O) 
 which they represent is calculated by means of the factors 
 
 (2KC1) 149 : (K 2 O) 94 = i : 0-6308 
 (2NaCl) 117 : (Na 2 O; 62 = i : 0-530 
 
 NOTE i. Alternative Methods for the Estimation of the 
 Alkalies. (i) In the foregoing method, it will be obvious that, as 
 the potassium alone is determined, the sodium being estimated by 
 difference, the experimental errors fall more heavily upon the latter 
 alkali. This may be obviated by determining the chlorine (instead 
 of the potassium) in the mixed chlorides. The mixed chlorides are 
 dried and weighed as in the process described above ; the residue 
 is then dissolved in water, and the chlorine precipitated as silver 
 chloride by means of silver nitrate. The estimation may be carried 
 out gravimetrically, as described on p. 238, or by the volumetric 
 method given on p. 360. 
 
 From the result, by whichever method obtained, the percentage 
 of chlorine contained in the mixed chlorides is calculated. 
 
 From these data, namely, (W) the weight of the mixed chlorides, 
 and (B) the percentage of chlorine in the mixture, the actual 
 amount of each chloride present can be calculated ; thus 
 
 The theoretical percentage of chlorine in sodium chloride - 60*68 
 potassium chloride = 47*65 
 
 The percentage of chlorine, therefore, in any mixture of these 
 chlorides will obviously lie between these two extremes ; the nearer 
 it approaches to one of them, the less of the other chloride will be 
 present in the mixture. Thus, for example, the percentage of 
 sodium chloride in the mixture will rise from o to 100 as the per- 
 centage of chlorine increases from 47 '65 to 60*68, i.e. passes through 
 a range of 13*03 per cents. (60*68 47*65 = 13*03). 
 
 Hence, if the percentage of chlorine in potassium chloride be 
 deducted from the percentage of chlorine found in the mixed 
 chlorides, the result when multiplied by 100 and divided by 13*03 
 gives the percentage of sodium chloride in the mixture 
 
 U 
 
290 Typical Gravimetric Separations. 
 
 13*03 : B 47*65 : : 100 : :r = percentage of sodium chloride in 
 
 the mixed chlorides 
 
 From this the actual weight of sodium chloride found is calcu- 
 lated by the proportion 
 
 100 : W : : x : y = grams of sodium chloride in the mixed chlorides 
 
 And then 
 
 W y = 2 grams of potassium chloride in the mixed chlorides 
 A concrete example will render this perfectly clear 
 
 The weight of the mixed chlorides (W) was 0*64 gram 
 The percentage of chlorine in the mixture (B) was 50*256 
 
 Deducting from B the theoretical percentage of chlorine in 
 potassium chloride, we get 
 
 Percentage of chlorine found 50*256 
 
 ,, in KC1 47*650 
 
 Difference 2*606 
 
 Then 13*03 : 2*606; \ 100:20*0 = / of NaCl in mixed chlorides 
 and 100 : 0*64 : : 20 : 0*1280 = grams of NaCl in mixed chlorides 
 
 And deducting this from the weight of the mixed chlorides, the 
 weight of the potassium chloride is found ; thus 
 
 Grams. 
 
 Weight of mixed chlorides 0*640 
 
 ,, sodium chloride 0*128 
 
 ,, potassium ,, 0*512 
 
 By multiplying the weights of sodium chloride and of potassium 
 chloride by their respective factors (given above), the weights of the 
 oxides equivalent to the chlorides will be obtained, and from 
 these the actual percentages of Na 2 O and K 2 O in the silicate are 
 calculated ; thus 
 
 0*128 x 0*530 = 0*06784 = grams Na 2 O 
 0-512 x 0*6308 = 0*32297 = grams K 2 O 
 Weight of silicate taken for analysis = 2*0150 grams 
 therefore 2-0150 : 0-06784 : : 100 : 3*31 = / of Na 2 O in the silicate 
 and 2*0150 : 0*32297 : : loo : 15*03 = / K 2 O " 
 
 (2) Another method for the estimation of the alkali metals is an 
 electrolytic process,* by means of which both alkali metals can be 
 separately determined. The mixed chlorides are dried and weighed 
 as before,f and the potassium precipitated as the double potassium 
 platinic chloride. The precipitate is washed perfectly free from 
 
 * See Section V. p. 292. 
 
 t It is not essential to weigh the dry chlorides, although a knowledge 
 of their weight serves as a .useful check upon the results obtained by the 
 electrolysis. 
 
Estimation of Alkalies. 291 
 
 platinum chloride by means of alcohol ; it is then dissolved in 
 water and the solution submitted to electrolysis, using the current 
 from one Daniell cell. The deposited platinum is washed and 
 dried, and from its weight the weight of potassium chloride present 
 is calculated by means of the factor 
 
 (Pt) 195 : (2KC1) 149 = i : 07641 
 
 The nitrate from the precipitated potassium platinic chloride 
 contains the sodium chloride mixed with the excess of platinum 
 chloride. It is electrolysed until the whole of the platinum is 
 deposited, and the liquid, now containing only sodium chloride, is 
 transferred to a dish and evaporated to dryness : after which it is 
 gently heated, and weighed. The weight of sodium chloride thus 
 determined, when added to the weight of potassium chloride cal- 
 culated from the former operation, should equal the weight of the 
 mixed chlorides. 
 
 From the chlorides, the proportion of potash and soda present 
 in the silicate is calculated as in the previous example. 
 
 NOTE 2. Alternative Method for decomposing the Silicate, by 
 means of Hydrofluoric Acid. The finely powdered silicate is treated 
 with strong aqueous hydrofluoric acid (pure) in a capacious 
 platinum crucible, and the mixture digested upon a steam-bath 
 (in a draught chamber) for some time, fresh acid being added once 
 or twice as the liquid evaporates. A few drops of strong sulphuric 
 acid, previously mixed with its own volume of water, are added in 
 order to convert the metals into sulphates, and the mixture 
 evaporated nearly to dryness upon the steam-bath. The crucible 
 is then cautiously heated over a small rose burner (see Fig. 52, p. 248) 
 in order to expel the excess of sulphuric acid. The residue is 
 moistened with a few drops of strong hydrochloric acid and 
 dissolved in water. If the decomposition of the silicate by the 
 hydrofluoric acid has been complete, any undissolved residue will 
 consist of insoluble sulphates (i.e. sulphates of lead or the alkaline 
 earths). 
 
SECTION V. 
 ELECTROLYTIC METHODS. 
 
 THE quantitative estimation of metals by the process of electrolysis 
 is based upon the fact that when solutions of certain metallic salts 
 are submitted to the action of an electric current of suitable 
 strength, the metals are precipitated upon the negative electrode 
 in the form of coherent films or deposits. The operation is, in 
 fact, an electroplating process, where the article to be " plated " 
 is a weighed piece of platinum, and the metal with which it is 
 coated is the metal that is to be estimated. 
 
 Just as in the familiar process of electroplating with silver, it 
 has been found that a solution of the double cyanide of potassium 
 and silver is the most suitable compound for the purpose, so with 
 the other metals it is only certain of their salts which are adapted 
 for the quantitative deposition of the metal they contain. For 
 example, when a solution of zinc sulphate is electrolysed, the first 
 action of the current (as in other cases) is to separate the salt into 
 its ions. The positive ions, or cathions (i.e. the zinc atoms), travel 
 to the cathode, or negative electrode, while the negative ions, or 
 anions (namely, the (SO 4 ) groups), go to the anode, or positive 
 electrode ; thus 
 
 Cathode. Anode. 
 
 ZnSO 4 = Zn + SO 4 
 
 The negative ions, however, at once undergo a secondary 
 decomposition, for in the presence of the water they are converted 
 into sulphuric acid and free oxygen 
 
 2SO 4 + 2H 2 O = 2H 2 SO 4 + O 2 
 
 The sulphuric acid thus generated by this secondary action 
 attacks the deposited zinc, and a condition of equilibrium is 
 established when the metal is dissolved as fast as it is deposited ; 
 hence complete precipitation of zinc under these conditions is 
 impossible. If, instead of the sulphate, the oxalate (or the double 
 oxalate of zinc and ammonium, as being the more soluble com- 
 pound) be employed, no such complication from the action of the 
 products of secondary decompositions will arise. The zinc oxalate 
 
Electrolytic Methods. 
 
 293 
 
 is decomposed into zinc and carbon dioxide, while the ammonium 
 oxalate is separated primarily into ammonium (NH.j) and carbon 
 dioxide, and secondarily into hydrogen and hydrogen ammonium 
 carbonate ; thus 
 
 Cathode. Anode. 
 
 ZnC 2 O 4 - Zn + 2CO., 
 
 (i) (NH 4 ) 2 C 2 O 4 = 2NH 4 ( = H 2 + 2NH 3 ) + 2 CO, 
 (2) 2NH 3 + 2H 2 O + 2CO 2 = 2H(NH 4 )CO 3 
 
 The apparatus required consists essentially of two platinum 
 electrodes. That which is to serve as the cathode, and receive 
 the deposited metal, is usually made in the form either of a cylinder 
 or, better, a cone ; which in either case is 
 riveted or welded to a stout platinum wire, 
 a, Fig. 61. The advantage of the conical 
 shape is that loss from spitting, as the 
 gas-bubbles rise to the surface of the 
 liquid, is reduced to a minimum. The 
 anode is conveniently a thick platinum 
 wire bent in the form shown at b, Fig. 61. 
 
 When only an occasional analysis is 
 to be made, these electrodes may be sup- 
 ported in a beaker by means of ordinary 
 clamps upon two retort-stands ; the beaker 
 being placed upon a wooden block, so that by the withdrawal of 
 the latter the beaker can be lowered and removed. By means of 
 
 FIG. 61." 
 
 FIG. 62 
 
294 Electrolytic Methods. 
 
 binding-screws, the electrodes are connected to the wires leading 
 from the battery. The arrangement is shown in Fig. 62. 
 
 Instead of conducting the electrolysis in a beaker, it may be 
 carried out in a platinum dish. In this case the dish is made the 
 cathode, by being supported either upon the ring of a retort-stand, 
 or upon any other convenient metal support which is connected to 
 
 FIG. 63. 
 
 the negative wire from the battery ; and the anode, held in a clamp, 
 is lowered nearly to the bottom of the dish, as shown in Fig. 63. 
 The metal to be determined is then deposited upon the dish itself. 
 The deposited metal is ultimately removed from the electrode by 
 solution in a suitable acid. 
 
 Typical Examples.* 
 
 Estimation of Copper. About 1-5 gram of recrystallised 
 copper sulphate is weighed out into a beaker and dissolved in 
 200 c.c. of water. About 2 c.c. strong nitric acid are added,f and the 
 mixture gently stirred with the platinum wire which is to be used 
 as the anode ; this is then left standing in the beaker. The platinum 
 
 10 wish for a more extended practice in electrolytic methods, 
 Classen's " Quantitative Analysis by Electrolysis," English 
 
 * Students who 
 should consult Classen' 
 translation (Herrick). 1894. 
 
 t When copper is to be separated from other metals, a larger quantity of 
 nitric acid must be added (see p. 298). 
 
Estimation of Copper. 295 
 
 cone, which has been perfectly cleaned * and carefully weighed, is 
 then lowered into the beaker over the projecting end of the anode. 
 The two electrodes are then supported by the clamps, so that they 
 do not touch each other, and are slightly raised from the bottom 
 of the beaker, the anode reaching slightly below the cone, in the 
 manner shown in Fig. 62. 
 
 The electrodes are then connected to the battery (the platinum 
 cone being attached to the negative wire), and a current of 0*5 to i 
 ampere passed through the solution.t The copper is gradually 
 deposited upon the platinum cone, and the blue colour of the solu- 
 tion becomes fainter and fainter, until at last the liquid appears 
 colourless. The operation takes several hours for its completion, 
 and may be conveniently allowed to go on all night. To ascertain 
 whether the precipitation is finished, one or two drops of the solu- 
 tion are withdrawn by means of a pipette, and a little sulphuretted 
 hydrogen added, which should produce no coloration. The 
 wooden block is then withdrawn, and the beaker lowered away 
 before interrupting the current, the anode at the same time being 
 disconnected and removed with the beaker. The cone is then dis- 
 mounted, and rinsed with water by means of a wash-bottle. It is 
 then dipped once or twice into a beaker of alcohol, and placed in 
 the steam-oven for a few minutes to dry, and then weighed. 
 
 The gain in weight gives the copper in the amount of copper 
 sulphate employed, from which the percentage of copper is calcu- 
 lated. (The result is usually from 0*1 to o'2 per cent, too low.) 
 
 The electrolytic method is one that is much employed in the 
 commercial analysis or valuation of copper ores and of metallic 
 copper, as well as in the estimation of copper in alloys. The 
 analyses are carried out in the following manner : 
 
 I. In copper ores. 
 
 * It is of the greatest importance that the cathodes used in electrolytic 
 analysis should be absolutely clean and free from the slightest trace of grease, 
 even such as would be contracted by touching them with the fingers. 
 
 f When the electric current from a "dynamo" or from storage batteries 
 is employed, it must be reduced to the requisite strength by the introduction of 
 a system of resistance coils. If too strong a current be used, the deposited 
 copper is rough and less coherent. When only an occasional analysis is to be 
 made, two or three Daniell cells may be used. In this case the zinc plate of 
 the battery is the negative, and is connected to the platinum cone. Care must 
 be taken that all the connections are clean, so as to ensure perfect metallic 
 contact. If the electrolysis be conducted in a platinum basin, the latter should 
 be covered, in order to prevent any of the liquid being carried off in the form 
 of fine spray with the gas which escapes from the anode. For this purpose, 
 either a clock-glass, with a small hole bored in the middle to allow the wire to 
 project through, or an inverted funnel of such a size that the mouth will just go 
 inside the dish, may be employed. When the platinum cone is used, the 
 vessel need not be covered. 
 
296 Electrolytic Methods. 
 
 From i to 1 1 gram * of the finely powdered ore is weighed out 
 into a porcelain'dish, and treated with from 10 to 20 c.c. of strong 
 nitric acid. About an equal volume of dilute sulphuric acidf is 
 then added, and the mixture gently evaporated to about half its 
 bulk in the covered dish. Water is then added, and the insoluble 
 residue (the gangue) is filtered and washed.^ The filtrate is diluted 
 up to 200 c.c. with water, 10 c.c. of nitric acid added, and the solu- 
 tion electrolysed as described above, with a current from two Daniell 
 cells. 
 
 II. In commercial copper. 
 
 About 0*5 gram of the metal is weighed out into a beaker, and 
 dissolved in 10 c.c. strong nitric acid, to which an equal volume of 
 water has been previously added. The mixture is boiled in the 
 covered beaker until " nitrous fumes " cease to come off, after 
 which it is diluted up to 200 c.c. with water, and the solution 
 electrolysed. The electrolysis must not be continued longer than 
 is necessary for the complete precipitation of the copper, other- 
 wise traces of antimony and arsenic, which may be present, will 
 also be deposited, which will be evident by the cathode becoming 
 darkened in colour. (For the separation of traces of antimony 
 and arsenic thus deposited, see footnote.) Should the sample of 
 copper contain any silver, this will be deposited and weighed along 
 with the precipitated copper. In this case the silver must be 
 separately estimated (by dissolving 10 to 20 grams of the metal 
 in nitric acid, and precipitating with hydrochloric acid in the usual 
 way, p. 238) and the proportion present deducted from the weight 
 obtained. 
 
 Estimation of Zinc. About 2 grams of crystallised zinc 
 sulphate are weighed out into a beaker, and dissolved in about 
 
 * For ores containing less than 25 per cent, of copper, from i to 5 grams 
 may be taken. 
 
 f If the ore is a sulphide, the addition of sulphuric acid is unnecessary, 
 since the sulphur is oxidised into sulphuric acid. 
 
 \ Copper ores frequently contain organic matter, in which case this residue 
 will be dark coloured, and is liable to retain a small portion of the copper. If 
 the quantity of this bituminous matter is appreciable, it should be destroyed 
 before treatment with nitric acid, by roasting the weighed quantity of powdered 
 ore taken for the analysis in a porcelain crucible, with free access of air. 
 
 If more than small quantities of antimony or arsenic are present in the 
 ore, these elements begin to deposit upon the copper towards the end of the 
 operation. When the amount is quite small, they may be separated by 
 strongly heating the electrode. The arsenic volatilises, while the antimony 
 and copper are oxidised ; the former oxide volatilises, leaving copper oxide. 
 This is then dissolved in nitric acid and redeposited by the current. When the 
 amount of antimony or arsenic is considerable, this plan cannot be adopted, 
 and, moreover, the precipitation of copper in the presence of much arsenic is 
 incomplete. In this case, a weighed quantity of the finely powdered ore is 
 gently heated in a covered porcelain crucible with about four times its weight 
 of ammonium chloride. Antimony and arsenic are thus converted into 
 chlorides and expelled. The residue is dissolved in nitric acid and treated as 
 already described (Classen, " Zts Anal. Ch.," 18, 388). 
 
Analysis of German Silver. 297 
 
 5oc.c. of water. Six or 7 grams of ammonium oxalate, dissolved in 
 a small quantity of warm water, are gradually added, with constant 
 stirring. 
 
 The solution is then diluted to 150 c.c. and electrolysed. The 
 process is complete when a drop of the solution gives no precipitate 
 when warmed with potassium ferrocyanide upon a watch-glass. 
 The platinum electrode containing the deposited zinc is then 
 removed, thoroughly rinsed with water, and finally with absolute 
 alcohol. It is then placed for a few minutes in a steam-oven to 
 dry, and weighed. 
 
 Estimation of Nickel. About 2 grams of ammonium nickel 
 sulphate, (NH 4 ) 2 SO 4 ,NiSO 4 ,6H 2 O, are weighed out into a beaker 
 and dissolved in water, and the solution rendered strongly alkaline 
 by the addition of ammonia. The volume of the liquid is made up 
 to 1 50 c.c. with water, and the solution submitted to electrolysis. 
 The process may be allowed to continue all night : it is complete 
 when a drop of the liquid gives no precipitate with ammonium 
 sulphide. The cathode is then washed and dried, and finally 
 weighed, as in the former examples.* 
 
 Analysis of Silver Coin. About 0*5 gram of the metal is 
 weighed out and dissolved in nitric acid (strong acid diluted with 
 an equal volume of water) in a covered evaporating-dish. The 
 liquid' is evaporated nearly to dryness upon a steam-bath to expel 
 the acid, and the residue dissolved in water. The solution is then 
 transferred to a beaker, and a moderately strong solution of ammo- 
 nium oxalate is added. Oxalates of silver and copper are formed ; 
 the latter, however, dissolves in excess of the reagent, leaving the 
 insoluble silver oxalate as a white precipitate. The liquid is 
 filtered, and the precipitate washed perfectly free from copper, first 
 with a dilute solution of ammonium oxalate, and finally with water. 
 The copper, contained in the filtrate and washings in the form of 
 the double ammonium copper oxalate, is then precipitated by elec- 
 trolysis. The precipitated silver oxalate is dissolved in a solution 
 of potassium cyanide, diluted with water to about 150 to 200 c.c., 
 and the solution electrolysed. The electrode with the deposited 
 silver is washed with water and with alcohol, and dried in the 
 steam-oven. 
 
 Analysis of German Silver. About i gram of the alloy is 
 
 * Nickel may also be precipitated from the solution of the double ammo- 
 nium oxalate, as described for zinc. In this case, the solution should be 
 maintained at a temperature about 40 to 50 throughout the operation, by 
 means of a small flame placed beneath the beaker. Cobalt and iron may be 
 estimated in the same way. 
 
298 Electrolytic Methods. 
 
 weighed out into a beaker, and dissolved in about 15 c.c. of strong 
 nitric acid mixed with the same volume of water. The solution is 
 then diluted up to 150 c.c. with water, and submitted to electro- 
 lysis. Under these circumstances the copper alone is deposited.* 
 When the precipitation is complete,! the beaker containing the 
 solution is lowered away from the electrodes before the current 
 is interrupted. The cathode with its deposit of copper is rinsed 
 with water (the washing being added to the solution in the beaker), 
 and finally dipped into alcohol and dried and weighed. 
 
 The positive electrode is also rinsed into the beaker,J and the 
 solution evaporated to dryness in a dish upon a water-bath, with 
 the addition of a little hydrochloric acid, in order to convert the 
 remaining metals into chlorides. 
 
 The residue is dissolved in water and a few drops of hydro- 
 chloric acid, and the solution transferred to a beaker. Sodium 
 carbonate solution is added until a slight precipitate persists, after 
 which hydrochloric acid is added drop by drop until the preci- 
 pitate just redissolves. The zinc is then precipitated from this 
 solution as zinc sulphide, as described on p. 244. The washed 
 precipitate is dissolved in the smallest possible quantity of strong 
 hydrochloric acid, and the solution evaporated to expel the excess 
 of acid. The residue is dissolved in water, ammonium oxalate is 
 added (as described on p. 297), and the solution diluted with water 
 to 100 c.c. and electrolysed. 
 
 * It must be remembered that nitric acid is "decomposed by the passage 
 of an electric current, nitrogen peroxide and oxygen being evolved at the 
 anode, while hydrogen is liberated at the negative electrode. This nascent 
 hydrogen reacts upon the nitric acid, with the formation of ammonia ; thus 
 
 Hence, if the current is allowed to continue passing through the solution after 
 the copper is all precipitated, the nitric acid is gradually decomposed in this 
 manner, and then the zinc in the solution begins to deposit along with the 
 copper. 
 
 t A drop of the solution, when tested by adding sodium bicarbonate and 
 potassium ferrocyanide, should give no brown coloration of copper ferro- 
 cyanide. 
 
 % In the analysis of alloys in which traces of lead are present as an 
 impurity, the lead will be deposited as the peroxide upon the positive wire, 
 which will appear brownish or black in consequence. By weighing the 
 electrode with the deposited peroxide, after drying in an air-oven at 110, 
 the amount of lead present can be calculated. 
 
 Any traces of iron present in the alloy will be deposited along with the 
 zinc. If the amount is greater than mere traces, it may either be removed 
 before the zinc is precipitated as sulphide (as described on p. 271), or a 
 separate estimation may be made in a larger quantity of the alloy ; and the 
 calculated proportion which will have been deposited with the zinc is deducted 
 from the weight obtained of the electrolytic deposit of that metal. 
 
Analysis of German Silver. 299 
 
 The filtrate and washings from the zinc sulphide contain the 
 nickel, which may be precipitated either from the double ammo- 
 nium sulphate or oxalate. 
 
 (a) From the Double Sulphate. Three or 4 grams of ammo- 
 nium sulphate dissolved in 10 c.c. of water are added to the 
 solution, and then about 20 c.c. of strong ammonia. The mixture 
 is diluted with water to 100 c.c., and electrolysed as on p. 297. 
 
 () From the Double Oxalate. Four or 5 grams of ammonium 
 oxalate, dissolved in about 20 c.c. of warm water, are added to the 
 solution, which is then diluted with water to 100 c.c. and submitted 
 to electrolysis, the liquid being maintained at a temperature about 
 40 to 50 throughout the operation. 
 
 Alloys consisting of copper and nickel only, may be analysed 
 electrolytic ally in the following way : From 0*5 to 075 gram of the 
 alloy is dissolved in the minimum quantity of nitric acid (diluted 
 as above), and the solution evaporated to dryness with the addition 
 of a few drops of sulphuric acid. The residue is taken up with 
 water, and again evaporated to dryness with a few drops of sul- 
 phuric acid to completely expel the nitric acid. The residue is dis- 
 solved in water with a few drops of sulphuric acid, and made up to 
 about 100 c.c., and the solution electrolysed (a current of 0-5 ampere 
 requiring 4 to 6 hours). The electrode is washed and dried and 
 weighed as already described. To the solution 15 c.c. of strong 
 ammonia are added, and the liquid electrolysed with a current of 
 about 0-3 ampere (or three Daniell cells). The nickel will be 
 entirely deposited in about 6 hours. 
 
PART II. 
 
 VOLUMETRIC METHODS. 
 
 SECTION I. 
 
 PRELIMINARY MANIPULATIONS. 
 
 Introductory Remarks. In volumetric methods of analysis,* 
 determinations are made, not by weighing a product obtained by 
 the interaction of a reagent with the substance to be estimated, 
 but by finding the weight of the reagent which it is necessary to 
 employ in order to exactly complete a given chemical interaction 
 with the substance to be determined. For example, in making a 
 gravimetric determination of chlorine in a soluble chloride, the 
 chlorine is precipitated as silver chloride by the addition of silver 
 nitrate solution in quantity a little in excess of that which is required 
 to complete the reaction, and the precipitated product of the inter- 
 action is washed and weighed. The 'volumetric method, on the 
 other hand, consists in finding the weight of the silver nitrate which 
 is required in order to just exactly throw down the whole of the 
 chlorine as silver chloride. Or again, instead of estimating iron by 
 precipitating the hydroxide, and weighing the sesquioxide obtained 
 by heating this hydroxide, the volumetric method consists in find- 
 ing the weight of some oxidising agent, such as potassium per- 
 manganate or potassium dichromate, which is just exactly required 
 to oxidise the iron from the ferrous to \hzferric state. 
 
 The weight of the reagent which is used in a volumetric deter- 
 mination is ascertained by employing a solution of known and 
 definite strength, and measuring the volume of it which is required 
 to carry the chemical change to its completion. Thus, in the above 
 illustration, if the exact strength of the silver nitrate solution be 
 
 * The term ' ' volumetric analysis " is here employed in its more restricted 
 sense, as applying to those methods of analysis where the volumes of solutions 
 only are concerned. Strictly speaking, it also includes those analytical pro- 
 cesses involving the measurement of volumes of gases or vapours, but by general 
 custom these are classed together under the head of " gas analysis," and con- 
 veniently constitute a separate section of analytical practice. 
 
Classification of Volumetric Methods. 301 
 
 known, it is only necessary to carefully measure the volume which 
 must be employed for the exact precipitation of the whole of the 
 chlorine, in order to learn the weight of the silver nitrate used, and 
 from this the weight of chlorine precipitated can be calculated. 
 
 Similarly, if the strength of the potassium permanganate solu- 
 tion be known, then from the number of cubic centimetres employed, 
 the weight of the permanganate which is required to oxidise the 
 iron present can at once be ascertained, and from this the actual 
 quantity of the iron is readily calculated. 
 
 In practice these calculations are made once for all when the 
 reagents of known definite strengths, called standard solutions, are 
 prepared ; so that the volume of such a reagent which is required 
 for the analysis, gives at once the weight of the substance being 
 estimated. Thus, in the above examples, the strength of the 
 standard silver nitrate being known, the weight of chlorine which 
 can be precipitated by I c.c. is calculated, and this weight multiplied 
 by the number of cubic centimetres employed for an analysis gives 
 the weight of chlorine in the compound. Suppose, for example, 
 that the standard silver nitrate contains 0-017 gram of the salt in 
 every cubic centimetre, this would be capable of precipitating 
 0-00355 gram of chlorine ; hence I c.c. of the standard silver 
 nitrate represents, or is equivalent to, 0-00355 gram of chlorine. 
 
 In the same manner, from the known strength of the potassium 
 permanganate solution, the weight of iron which i c.c. of it is 
 capable of oxidising can be calculated ; so that when the reagent 
 is employed for the estimation of iron, the number of cubic centi- 
 metres used, multiplied by the weight of iron which each cubic 
 centimetre is equivalent to, will give the weight of iron in the portion 
 of the substance being analysed. 
 
 The process of carrying out a volumetric estimation by the use 
 of standard solutions, is called titration ; thus, a solution of a 
 chloride is titrated with standard silver nitrate ; a ferrous salt in 
 solution is titrated with potassium permanganate ; and so on. 
 Standard solutions themselves have to be titrated, i.e. their chemical 
 value or power has to be determined in order to ascertain their exact 
 strength. They are, therefore, often spoken of as titrated solutions. 
 
 Classification of Volumetric Methods. Almost all volu- 
 metric processes are based upon one of three principles, and they 
 may, therefore, be conveniently classified in the following manner: * 
 
 * One important process which is not included in this classification, is the 
 estimation of copper by means of potassium cyanide. A description of this 
 method will be found on p. 374. 
 
3Q2 Volumetric Methods. 
 
 I. Methods based upon the neutralisation of alkalies and acids, 
 known also as methods of saturation. These methods are some- 
 times subdivided into two classes, which are included under the 
 terms alkalimetry and acidimetry, the one being the converse of 
 the other. 
 
 II. Methods based upon processes of oxidation or reduction. 
 
 III. Methods based upon precipitation, either alone or in con- 
 junction with one of the other methods. For example, silver may 
 be estimated by precipitation alone by means of standard sodium 
 chloride. As the precipitate readily settles, the point of completion 
 of the reaction may be determined by the failure of a drop of the 
 standard reagent to produce any further turbidity. Barium, on the 
 other hand, may be determined by precipitation as carbonate (from 
 a neutral solution) by the use of an excess of a standard solution 
 of sodium carbonate. The precipitate is filtered off, and the excess 
 of sodium carbonate in the nitrate and washing is estimated by 
 titration with a standard acid. The excess so found, deducted from 
 the total alkali used, gives the amount required for the precipitation 
 of the barium. Or the precipitated barium carbonate, after thorough 
 washing, is dissolved in an excess of standard hydrochloric acid, 
 and the excess of acid estimated by means of standard alkali. The 
 amount of acid required to dissolve the barium carbonate is thus 
 found, and by calculation the weight of barium is ascertained. 
 
 Direct and Indirect Determinations. In a large number 
 of volumetric analyses, substances are determined by indirect pro- 
 cesses ; they are estimated, as it were, by proxy. The difference 
 between direct and indirect processes will be most readily under- 
 stood from the following illustrations. 
 
 Ammonia may be directly determined in an ammoniacal salt 
 by heating a known quantity of the salt along with caustic alkali, 
 and causing the whole of the liberated ammonia to pass into, and 
 be absorbed by, a measured volume of standard sulphuric acid. 
 The exact quantity of this acid, which is saturated or neutralised 
 by the ammonia, is then ascertained by titrating the solution with 
 a standard alkali. From the amount of alkali thus used, the excess 
 of sulphuric acid present is ascertained, and this deducted from the 
 total volume of acid gives the quantity which has been neutralised 
 by the ammonia ,- and as the weight of ammonia to which each 
 cubic centimetre of the standard acid is equivalent is known, it also 
 gives directly the weight of ammonia present in the ammoniacal 
 salt taken for the estimation. 
 
 Indirectly ammonia can be determined by boiling a known 
 
Preliminary Manipulations. 303 
 
 quantity of the ammonium salt with a measured volume of standard 
 caustic alkali, and allowing the ammonia to escape. When the 
 whole of the ammonia has thus been expelled, the excess of caustic 
 alkali is estimated by titration with standard acid. The amount of 
 alkali thus determined, deducted from the original volume employed, 
 represents the quantity of alkali which has been decomposed and 
 converted into a salt of the acid which was originally in combination 
 with the ammonia, e.g. if ammonium chloride be the salt employed, 
 then a certain quantity of the standard caustic soda will be con- 
 verted into sodium chloride, and this gives indirectly the amount 
 of ammonia. 
 
 In the direct process we actually estimate the ammonia, while 
 by the indirect method we, in reality, determine the quantity of the 
 add radical with which the ammonia is combined, and from this 
 find the ammonia by calculation. 
 
 Again, chromium in a chromate may be determined indirectly 
 by boiling the salt with hydrochloric acid, whereby chlorine is 
 liberated in the proportion of 3 equivalents of Cl for every i equiva- 
 lent of chromium trioxide, CrO 3 ; thus 
 
 K 2 CrO 4 + 8HC1 = 2KC1 + CrCl 3 + 4H 2 O + 3d 
 
 The liberated chlorine is made to pass into potassium iodide 
 solution, whereby an equivalent quantity of iodine is liberated, and 
 the amount of iodine so set free is determined by means of a 
 standard solution of sodium thiosulphate (as explained later). 
 
 In this process, iodine is actually determined, and from this 
 indirectly the chlorine, and yet more indirectly the chromium is 
 estimated. 
 
 In order to carry out volumetric analyses, the three following 
 conditions must be fulfilled, namely : (i) the means of accurately 
 measuring the volumes of liquids ; (2) the means of preparing the 
 necessary standard solutions and of testing their accuracy ; and (3) 
 the means of readily ascertaining the exact point when the various 
 chemical reactions involved are complete. 
 
 I. Instruments for Measuring Liquids Graduated 
 glass vessels are employed for measuring the volumes of liquids, 
 four different forms of apparatus being in common use for different 
 purposes, namely, flasks, cylinders, pipettes, and burettes. 
 
 (a) Graduated Flasks. These are flasks of such a size that they 
 shall contain a specified volume of liquid when filled up to a gradu- 
 ation mark upon the neck. They should be provided with a 
 stopper, the neck should be somewhat long and narrow, and the 
 
304 
 
 Volumetric Methods. 
 
 graduation mark should lie fairly low down upon the neck, in order 
 that the contents of the flask, when it is filled to the mark, may be 
 conveniently shaken up (Fig. 64). The three sizes most convenient 
 for ordinary work are the litre, half- litre, and 
 quarter-litre, besides which a loo-c.c. flask is 
 useful. 
 
 If one of these flasks be filled with water up 
 to the graduation mark, it will be obvious that 
 when the liquid is poured out, a certain small 
 proportion of it remains behind adhering to the 
 inside of the flask; in other words, the flask does 
 not deliver absolutely the whole of the liquid it 
 contained, hence, when such a flask is employed 
 to deliver a volume exactly equal to its indicated 
 capacity, it must contain an excess of the liquid 
 equal to the quantity which remains adhering to 
 the glass. Measuring-flasks, therefore, usually 
 contain two graduation marks, the lower one being 
 the mark to which the vessel must be filled in 
 FIG. 6 4 . order to contain the volume represented by the 
 
 denomination of the flask, while the one slightly above it indi- 
 cates the point to which the flask must be filled in order that 
 it shall deliver that volume. This applies more especially to the 
 smaller sizes, as the larger vessels are more exclusively used for 
 making up solutions to a given volume. Measuring-flasks are 
 usually graduated at 15 or 15-5 C. (60 F.). Thus, a litre flask is 
 graduated to contain 1000 c.c. of water measured at a temperature 
 of 15*5, or about the average temperature of the air.* Since, how- 
 ever, liquids undergo an appreciable alteration in volume by change 
 of temperature, their measurements should always be made as 
 nearly as possible at one uniform temperature, otherwise it is 
 
 * This volume of water will obviously weigh slightly less than 1000 grams, 
 since i gram is the weight of i c.c. at its point of maximum density, namely, 
 4. Its actual weight is found by multiplying the number of cubic centimetres 
 by the density of water at 15 '5 (given in the table in the Appendix) ; thus, icoo x 
 0-99911 = 999*11 grams. Similarly, for a ^-litre flask, 25oxo'999ii = 24977 
 grams. 
 
 If TOCO grams of water at a temperature of 15*5 be placed into a flask, and 
 the volume it occupies be called icoo c.c., obviously these cubic centimetres are 
 not the true, or, as they are termed, the absolute cubic centimetres ; they are, 
 of course, slightly greater greater in the proportion of i : 1-00089 (this being 
 the coefficient of expansion of water at 15-5). In this case, the so-called litre 
 flask would in reality have a capacity of 1000-89 absolute cubic centimetres ; 
 and similarly, the ^-litre flask, while containing 150 of the larger arbitrary cubic 
 centimetres, would hold 250 X 1-00089 = 250-22 absolute cubic centimetres. 
 
Calibration of Litre Flasks. 305 
 
 necessary to introduce a correction for temperature. For example, 
 suppose a 250 c.c. flask, which has been graduated at I5'5, is 
 filled to the mark with water having a temperature of 20, then from 
 the table in the Appendix, giving the coefficients of expansion of 
 water (the volume at 4 = i), we get the following proportion : 
 
 1*00169 : 1*00089 : : 250 : 249*8 c.c. 
 
 (vol. at 20) (vol. at i5'5) 
 
 hence the volume of liquid measured at the higher temperature is 
 about 1 of a cubic centimetre short. For all ordinary analytical 
 purposes, however, slight variations of temperature to the extent of 
 one or two degrees on either side of 15*5 may be disregarded. 
 
 The accuracy of the graduations of measuring-flasks should be 
 tested, i.e. the vessels should be calibrated before being used. For 
 this purpose the dry clean flask is first counterpoised upon a balance 
 capable of carrying a heavy load. Weights are then placed upon 
 the scale-pan, along with the tare, equal to the number of grams of 
 water at 15- 5 which the flask is intended to contain.* Water is 
 then poured into the flask (which should be removed from the 
 balance while being filled) until the level of the lowest point of 
 the meniscus (see p. 310) is coincident with the graduation mark. 
 If the flask so filled exactly balances the weights, the graduation 
 may be taken as correct ; if not, water is either withdrawn or added 
 by means of a fine pipette until equilibrium is established, and a 
 fresh graduation is made upon the neck by means of a scratching 
 diamond or a file. 
 
 In order to graduate a flask to deliver a definite volume, exactly 
 the same procedure is carried out, except that the flask, instead of 
 being dry at the beginning, is first filled with water to the mark, 
 and then emptied out and allowed to drain for a few moments (as 
 
 * In graduating or calibrating measuring-flasks, one of two plans may be 
 adopted 
 
 (1) The number of grams of water equal to the denomination of the flask 
 at 15 '5 (or any other temperature which would be more suitable under special 
 climatic conditions) are carefully weighed out, and the volume which they 
 occupy in the flask is indicated by a mark scratched upon the neck. The cubic 
 centimetres in this case have an arbitrary value, as explained in the footnote on 
 p. 304, but if all the measuring vessels are graduated on the same system, so 
 that this arbitrary unit has the same value in them all, no error will arise on 
 this account. 
 
 (2) A quantity of water at 15 '5 is weighed out, which, if cooled to 4, would 
 then occupy a volume equal to the denomination of the flask. The number of 
 grams to be weighed out is ascertained by multiplying the denomination of the 
 flask by the density of water at 15*5, as explained in a previous footnote. 
 In this case, the cubic centimetres represent the true or absolute cubic centi- 
 metres. The first method of graduating is the one most usually adopted. 
 
 X 
 
306 Preliminary Volumetric Manipulations. 
 
 long as it would be allowed to drain when being actually used to 
 deliver). The flask is then counterpoised with the remaining traces 
 of water adhering to the interior surface of the glass. 
 
 (b] Graduated Cylinders. These are sometimes used instead 
 of flasks, when less accurate measurement is all that is necessary. 
 Two forms are in common use, as shown in Fig. 65, the larger 
 stoppered vessel being known as a test-mixer. As shown in the 
 figure, these cylinders are graduated throughout their length into 
 a number of small subdivisions of the total capacity. The correct- 
 ness of these graduations may be tested by introducing successive 
 
 FIG. 65. 
 
 FIG. 66. 
 
 small volumes of water from a pipette or a burette, which has been 
 previously tested. If the graduations upon the cylinder do not 
 coincide with the volumes which are thus introduced, a table must 
 be drawn up giving the actual value of each graduation. 
 
 (c] Pipettes. The pipette is a glass tube, usually with a bulb 
 or enlargement upon it, drawn to a point at one end. Fig. 66 
 shows the most usual forms of a small and a large pipette. The 
 instrument is graduated to deliver its specified volume of liquid. 
 
Calibration of Pipettes. 307 
 
 It is filled by sucking the liquid up until it is somewhat higher than 
 the graduation mark upon the stem, and covering the upper end 
 with the finger, as shown in the figure. Then, by a slight release 
 of the pressure of the finger, the liquid is allowed to flow slowly out, 
 until the level is coincident with the mark upon the stem. 
 
 When the pipette is allowed to empty itself, the last drops 
 which remain in the pointed end may either be blown out, or made 
 to flow out by drawing the point along against the wet sides of the 
 vessel into which the liquid is being delivered. The former plan is 
 the quicker, and is the one which is most instinctively resorted to ; 
 but the latter is the more accurate method. It is not a matter of 
 serious importance which method is adopted, so long as the same 
 plan is uniformly employed. 
 
 Pipettes are calibrated to deliver a definite volume by a method 
 practically the same as in the case of flasks. To carry it out, the 
 pipette is first filled with water by suction, and the end of the tube 
 which was dipped into the water is wiped dry on the outside. The 
 liquid is then allowed to run out, and the last drops removed by 
 one of the above-mentioned plans. The point of the instrument is 
 now closed by placing a minute fragment of wax * upon it, and 
 then just softening the wax by bringing the point near to a small 
 flame. In this way the pipette can be closed without the wax 
 entering the tube. The moist pipette is then placed upon a balance 
 and counterpoised. 
 
 It is next filled with water at a temperature of 15 '5, which can 
 be introduced by means of a glass tube drawn out sufficiently fine 
 to pass down the stem of the pipette. After the first few drops 
 have been put in, the air-bubble which is usually trapped in the 
 bottom of the tube may generally be dislodged by a few shakes ; if 
 this fails, however, it may be removed and the water made to run 
 right into the point by thrusting a long capillary tube down to the 
 bottom. Care must be taken, however, that no fragments of this 
 fine tube become broken off and left in the apparatus. During the 
 filling process the pipette should be supported by a clamp, and not 
 held in the warm hand, otherwise the temperature of the water 
 may be considerably raised during the operation. When the water 
 is coincident with the graduation mark, the instrument is carefully 
 laid upon the scale-pan (the water will not run out of it, although 
 placed in a horizontal position), and gram weights equal to the 
 number of cubic centimetres the pipette is intended to deliver are 
 
 * The black bituminous material known as "bicycle cement" answers 
 admirably for this purpose. 
 
308 Preliminary Volumetric Manipulations. 
 
 placed upon the opposite pan. If this is found to exactly counter- 
 poise the apparatus, the graduation may be taken as correct ; but 
 if not, water must be either withdrawn or added, as the case may 
 be, until the apparatus is exactly equipoised, and a fresh mark 
 scratched upon the stem. 
 
 (d} Burettes. -The burette is a long straight glass tube, one 
 end of which is drawn down and terminated by a glass stop-cock, 
 or connected to a jet by means of a caoutchouc tube which can be 
 
 closed by means of a pinch- 
 cock. Fig. 67 shows the two 
 forms of apparatus. The 
 burette is graduated almost 
 throughout the entire length, 
 the graduations being usually 
 tenths of a cubic centimetre, 
 as shown in Fig. 68. The 
 size most commonly used 
 has a capacity of 50 c.c. For 
 ordinary use the common re- 
 tort-stand and clamp shown 
 to the left in the figure make 
 a very convenient stand for 
 holding burettes. If desired, 
 several can be supported 
 upon the same retort-stand. 
 The instrument is filled by 
 means of a small funnel 
 placed in the top (which 
 should be removed after- 
 wards, lest any adhering 
 drops fall into the burette), 
 until the liquid is consider- 
 ably above the topmost gra- 
 duation. The tap or pinch-cock is then momentarily opened, in 
 order that the liquid may sweep out before it the air which is 
 contained in the tap and jet. This will not be successfully accom- 
 plished if the tap or pinch-cock is only gradtially opened, as the 
 liquid then slips down the walls of the narrowed portions and 
 leaves air-bubbles in the tubes, which it is absolutely necessary 
 to remove before the instrument can be used with exactness. 
 
 The calibration of a burette is a more serious operation than 
 that of a pipette, for since it is obviously out of the question to 
 
 iiiiiiiimiiiiiiiiii 
 
 FIG. 67. 
 
Calibration of Burettes. 309 
 
 regraduate the instrument, the true value of the graduations upon 
 it must be ascertained, and a special table constructed for each 
 instrument. The process is carried out in the following way : The 
 burette is filled with water at a temperature of 15-5 (if possible in 
 a room having that temperature, since the operation occupies a 
 considerable time), the air-bubbles being swept out as described, 
 and the level of the liquid being coincident with the topmost 
 graduation. A clean, dry, stoppered flask, capable of containing 
 as much liquid as the burette, is accurately counterpoised, and the 
 water in the burette is delivered into it in successive small quanti- 
 ties of 3, 4, or 5 of the c.c. graduations at a time. After each small 
 portion of water has been run into it, the flask is weighed until the 
 whole 50 c.c. has been delivered. 
 
 The number of grams which each portion weighs, represents the 
 number of cubic centimetres it actually occupies, and if this is 
 different from the number of c.c. graduations, the error is equally 
 divided between them.* Thus, suppose the water is delivered in por- 
 tions represented by five of the c.c. divisions, and the first portion 
 is found to weigh 5'!$ grams, instead of exactly 5 grams ; then the 
 error, namely 0*15, is divided equally among the five graduations, 
 each one of which will then have the value, not of i c.c., but 
 1*03 c.c. The table, therefore, will run as follows for the first five 
 graduations 
 
 Graduation i = 1*03 c.c. 
 
 2 = 2'06 
 
 3 = 3*09 
 
 4 = 4-12 
 
 5 = 5'i5 ,, 
 
 Suppose, after the second portion has been delivered, the weight 
 
 of the water is 10*17 grams ; then 
 
 1 0*17 5*15 = 5*02 grams = the weight of the second portion of water, 
 
 and ^-^= 1*004 c.c. = the value of each graduation between 5 
 
 and lo 
 hence the table will continue 
 
 * It will be evident that the larger the volume which is delivered at one 
 time, the less accurate is the calibration. Thus, from the above illustration, 
 if instead of delivering the contents of five graduations, that of ten had been 
 weighed at once, the value of each graduation would have been 
 
 "^7 =1-017 
 10 
 
 and the value of the fifth graduation would have come out 
 1-017x5 = 5-085 c.c. instead of 5*15 
 
3io Preliminary Volumetric Manipulations. 
 
 Graduation 6=5-15 + 1*004 = 6*154 c.c. 
 
 7 = + roo8 = 7-158 
 
 8 = + 1-012 = 8-162 
 
 9 = + roi6 = 9*166 
 
 10 = + I '020 = IOT70 
 
 Small burettes holding 10 c.c. are sometimes used after the 
 manner of a pipette. They consist of straight tubes drawn to a 
 point at one end, and graduated into tenths of a cubic centimetre, 
 as in the burette. They are filled by suction, and employed as 
 pipettes for delivering various small measured volumes. 
 
 Beading the Volume of Liquids in Graduated Vessels. 
 Owing to the action of capillarity, the surface of a liquid con- 
 tained in a glass-tube is not plane, but curved, the extent to which 
 it is curved depending (for the same liquid) upon the diameter of 
 the tube. 
 
 This curved surface is called the meniscus. Figs. 68 and 70 
 show the meniscus in the case of water contained in a burette. 
 
 16 -r 
 
 18- 
 
 FIG. 63. 
 
 FIG. 69. 
 
 In reading graduated instruments, it is usual to take the gradu- 
 ation which coincides with the lowest point of the meniscus. Thus, 
 in Figs. 68 and 70 the volume of the liquid would be taken as 16-5. 
 If, as sometimes happens in certain lights, the line of the curve is 
 not very clearly visible, it may usually be made quite distinct by 
 
The Use of Floats. 3 1 1 
 
 holding behind the tube a piece of white paper, inclining it slightly 
 upwards, as in Fig. 68. In order to make a correct reading, it is 
 necessary that the eye of the observer should be in the same hori- 
 zontal plane as the surface of the liquid. The reason for this will 
 be evident from the diagram (Fig. 69), where an error of one gradu- 
 ation would arise by reading from the positions A or A 1 instead of 
 from B. 
 
 Where the graduation mark is continued as a ring right round 
 the vessel, as is usually the case with pipettes and flasks, the eye 
 will be situated in a horizontal plane with the graduation, when the 
 front of the mark exactly overlays the back, so that the ring 
 appears as a line. The lowest point of the meniscus must then 
 just touch this line. A simple device to ensure that the readings 
 of a burette (where the graduations are on one side of the tube 
 only) shall be consistently made from a correct point of observa- 
 tion is shown in Fig. 70. It consists merely of a narrow strip of 
 card folded in the middle, with the two free ends pinned together 
 with a paper-fastener. This is slipped over the burette, and while 
 it allows of being easily slid up and down the tube, it will also hold 
 itself in any position in which it may be put. When taking a read- 
 ing, this little clip is placed so that its upper edge is just a little 
 below the level of the liquid : then, when the eye is in such a 
 position that the back and front edges of the clip just coincide, it 
 will also be practically in the same horizontal plane as the bottom 
 of the meniscus. 
 
 The same object may be gained by the use of floats. These are 
 little weighted glass bulbs or tubes which are placed inside the 
 burette, floating upon the liquid. The graduation which coincides 
 with the horizontal ring-mark upon the float is the one which is 
 read. Two forms of floats are seen in Fig. 71. A represents the 
 most familiar shape. Unless some care is exercised, the use of 
 such a float is very liable to introduce errors. If it is much 
 narrower than the burette, it may sometimes take up positions in 
 which the graduation mark upon it is not perfectly horizontal. On 
 the other hand, if it is too close a fit, it is very prone to lag behind 
 the liquid, or even to stick altogether. B (Fig. 71) represents a 
 newer and better form of float. Owing to its shape it always main- 
 tains a vertical position, and may therefore be sufficiently narrow 
 at its widest point to admit of perfect freedom of movement in the 
 burette. The graduation mark on this float is upon the smaller 
 bulb at the top, which projects above the liquid altogether, hence 
 this form of apparatus may be used equally well whether the liquid 
 
312 Preliminary Volumetric Manipulations. 
 
 in the burette is colourless or even opaque. The reading, as the 
 float stands in the figure, is 14*4. 
 
 It is very necessary, whatever float is used, to keep a close 
 watch upon it as the liquid in the burette gets down towards the 
 bottom graduations. It is well to place an indiarubber ring or 
 other convenient mark upon the burette at the point beyond which 
 it is not admissible to go when the float is being used. 
 
 FIG. 70. 
 
 FIG. 71. 
 
 II. Standard Solutions. Solutions of known strength which 
 are used in volumetric analysis are called standard solutions. 
 These various solutions are usually made of such strengths that 
 the quantities of the positive or negative constituents, or of what 
 may be called the active constituents, of the compounds in the solu- 
 tions shall bear the same relation to each other as the numbers 
 which express their chemical equivalence that is to say, equal 
 volumes of the different solutions will contain equivalent propor- 
 tions of the effective substances in solution. 
 
 For example, standard solutions of sodium chloride and silver 
 nitrate will be of such strengths that whatever volume of the former 
 contains 35^5 grams of chlorine, the same volume of the other 
 shall contain 108 grams of silver ; or, again, standard solutions of 
 
Standard Solutions. 313 
 
 hydrochloric acid, of sodium hydroxide, and of sodium carbonate 
 will be of such strengths that whatever volume of the first contains 
 i gram of hydrogen, the same volume of the others shall each 
 contain 23 grams of sodium. 
 
 Normal Standard Solutions. When the solutions are of 
 such a strength that the particular volume which contains I gram 
 of hydrogen, 23 grams of sodium, 35-5 grams of chlorine, etc., is 
 one litre, then the solution is known as a normal solution. Thus, a 
 normal solution of hydrochloric acid will contain I gram of hydro- 
 
 H ci 
 
 gen in the litre, therefore it must contain i + 35 *5 = 36*5 grams of 
 HC1 in that volume. Again, normal caustic soda will contain 
 
 Na H O 
 
 23 grams of sodium per litre, therefore it must contain 23+1 + 16 
 = 40 grams of NaHO in the loco c.c. Or, again, normal sodium 
 carbonate must also contain 23 grams of sodium per litre ; there- 
 
 Na 3 C O 3 
 
 fore the volume must contain - L = 53 grams of Na 2 CO 3 . 
 
 Similarly, if normal sulphuric acid is to be of such a strength 
 that i litre shall contain i gram of hydrogen, or, in other words, 
 that the amount of the negative radical SO 4 shall be the chemical 
 equivalent of i gram of. hydrogen, then the litre must contain 
 H 2 s 4 
 2 +32 + 6 4 = 49 grams of H 2 SO 4 . 
 
 In some cases, it is only a. portion, and not the whole, of some 
 particular element or radical present in the solution, which takes 
 an active part in the chemical reactions for which the solution 
 is used. Thus, in the case of potassium dichromate, K 2 Cr 2 O 7 , only 
 three out of the seven oxygen atoms contained in the salt are 
 available for purposes of oxidation. We may regard the compound 
 as breaking up into K 2 O,Cr 2 O 3 ,3O. A normal solution, therefore, 
 of this salt (that is, one which shall contain per litre a weight of 
 available oxygen chemically equivalent to i gram of hydrogen) 
 
 K 2 Cr 2 O, 
 
 must contain - = 49*13 grams of the salt in every 
 
 loco c.c. The formula-weight of the salt, in this instance, being 
 divided by 6 instead of by 3, because the equivalent of oxygen is 8, 
 and not 16. 
 
3 14 Preliminary Volumetric Manipulations. 
 
 Again, the available oxygen in potassium permanganate is five 
 out of eight atoms, 2(KMnO 4 ) = K 2 O,2MnO,5O. Hence a normal 
 solution of this salt, containing in the litre a weight of available 
 oxygen equivalent to i gram of hydrogen (i.e. 8 grams of available 
 oxygen), must contain in that volume of the solution a weight of 
 salt in grams equal to one-tenth of the formula- weight K 2 Mn 2 O 8 , or 
 
 one-fifth of that of KMnO 4 , namely, ^- = 31-6 grams.* 
 
 Solutions of one-half, one-tenth, and one-hundredth of the 
 strength of a normal solution are called respectively semi-normal, 
 fleet-normal, and centi-normal solutions, and are distinguished by 
 
 N N N 
 
 the signatures ' ' and - 
 
 2 10 IOO 
 
 In preparing standard solutions, the weighed quantity of salt is 
 not added to and dissolved in the measured volume of water, but 
 it is first dissolved in a moderate quantity of water, and the solution 
 afterwards diluted up to the requisite volume by the addition of 
 water. At first sight it might be supposed that the two processes 
 would give the same result, but this is not the case, because the act 
 of solution of many substances is attended by a change (usually a 
 contraction) in the volume, and therefore in one case the solution 
 obtained would be different in volume, and therefore in strength, 
 from that prepared in the other way. 
 
 Further details for the preparation of the various standard solu- 
 tions will be given in the sections devoted to the special analytical 
 processes for which the solutions are employed. 
 
 (III.) Methods for ascertaining the Completion of 
 Volumetric Reactions. It is essential to the success of a volu- 
 metric process that the exact point when the chemical reaction is 
 complete should be readily determinable. This is usually accom- 
 plished by means of a third substance, known as an indicator, which 
 is effected in some visible way either by the volumetric reagent or 
 by the solution of the substance being estimated. For example, in 
 order to determine the exact point when an alkali has been neu- 
 tralised by an acid, a small quantity of litmus solution may be added 
 
 * Some chemists, unfortunately, apply the term normal to solutions which 
 contain in one litre the formula-weight in grams (i.e. the gram-molecule) of the 
 compound. According to this method, while a normal solution of caustic soda 
 contains 23 grams of sodium per litre, the normal sodium carbonate will contain 
 46 grams of sodium in the same volume. To prevent confusion, these solutions 
 should be distinguished from the true normal solutions by the name molecular 
 solutions. 
 
The Use of Indicators. 315 
 
 to the solution. Again, the point of complete precipitation of 
 chlorine by means of standard silver nitrate, can be ascertained 
 with exactness by adding a small quantity of a solution of potassium 
 chromate to the solution containing the chloride to be estimated. 
 The silver combines by preference with the chlorine, forming silver 
 chloride ; but when the whole of the chlorine has been precipitated 
 as silver chloride, then the silver attacks the chromate, precipitating 
 the red silver chromate hence the production of a red precipitate 
 which persists, is the sign that the precipitation of the chlorine is 
 complete. The reagent, or standard solution, may itself be the 
 indicator ; this is the case, for example, with potassium perman- 
 ganate, in which the characteristic violet colour of the solution 
 serves to show the precise moment when the oxidising process is 
 complete. Very often an outside indicator is made use of; that is 
 to say, instead of adding the indicator to the solution which is being 
 tested, the progress of the reaction is watched by removing a drop 
 of the liquid upon a glass rod at frequent intervals, and applying it 
 to the indicator. Thus, in estimating the chlorine in bleaching 
 powder by the oxidation of sodium arsenite, the indicator employed 
 is a mixture of potassium iodide and starch, which gives the 
 characteristic blue colour when in contact with chlorine. The 
 iodide, however, cannot be added to the solution of bleaching 
 powder, but by bringing a drop of the latter into contact with 
 potassium iodide and starch paper until it no longer produces the 
 blue coloration, it is easy to hit the exact point when the oxidation 
 of the arsem'ous salt into the arsenzV state is completed. 
 
 Similarly, the point when a ferrous salt is completely oxidised 
 to the ferric condition by means of potassium dichromate, is ascer- 
 tained by withdrawing a drop of the mixture upon a glass rod and 
 applying it to a drop of potassium ferricyanide solution upon a 
 white plate or porcelain tile ; so long as any ferrous salt is present, 
 a blue coloration is produced, but as soon as the iron is entirely 
 oxidised, a drop of the solution so applied to the ferricyanide 
 indicator will produce no coloration. 
 
SECTION II. 
 
 VOLUMETRIC METHODS BASED UPON SATURATION. 
 Alkalimetry and. Acicliinstry. 
 
 Normal Alkalies and Acids. The four following standard 
 solutions may be prepared, namely, normal sodium carbonate, 
 sodium hydroxide, sulphuric acid, and hydrochloric acid. 
 
 In preparing such a series of solutions, it is necessary to select 
 one of them as the foundation or basis for the others ; and the con- 
 ditions which must be fulfilled by the one which is chosen for this 
 purpose are, that it admits of being made up with the utmost possible 
 accuracy. Of these four solutions the alkaline carbonate is the one 
 which most satisfactorily fulfils these requirements, for it is a salt 
 which can readily be obtained in a state of practical purity, and it 
 also admits of being weighed with perfect exactness. 
 
 Having once obtained a normal solution of sodium carbonate 
 of exact strength, it can be employed as a basis for the preparation 
 of normal sulphuric and hydrochloric acids, and by means of either of 
 these the standardisation of the caustic alkali may be effected. 
 The strength of the acids can be checked, if desired, by precipitation 
 (the one with barium chloride, and the other with silver nitrate) 
 and gravimetric determination. 
 
 Indicators. There are a number of substances which may 
 be employed in order to indicate the exact point of neutrality, or 
 the end reaction, as it is termed, in processes of alkalimetry and 
 acidimetry. The most important and useful * are litmus, methyl 
 orange, and phenolphthalein. 
 
 Litnms Solution. This is prepared by crushing about 10 grams 
 of the solid in a clean mortar with hot distilled water. The liquid 
 is poured off, and the residue further extracted by one or two fresh 
 supplies of hot water. The extract is then allowed to settle for 
 some hours, and the clear blue liquid, diluted to about 200 c.c., is 
 
 * For a full account of the various indicators, their relative sensitiveness, 
 and special uses, see Thomson, Chew. News. 
 
Volumetric Analysis. 317 
 
 transferred to a bottle. As this solution loses its colour if closely 
 corked or stoppered up so as to exclude the air, it is a good plan to 
 place in the neck of the bottle a flat- topped stopper which drops in 
 quite loosely, and serves merely as a cover to prevent the entrance 
 of foreign matter. A few drops of chloroform added and shaken 
 up with the solution, will prevent mould from forming in it, which 
 otherwise the liquid is very prone to develop. 
 
 Litmus solution is turned red by acids, and blue by alkalies ; 
 the colour which it should exhibit when perfectly neutral is violet j 
 it is usually necessary, therefore, to add a drop or two of very dilute 
 acid (preferably nitric acid) by dipping a glass rod into the acid 
 and stirring it into the litmus until a purple or violet tint is imparted 
 to the solution.* 
 
 When carbon dioxide is disengaged (as when an alkaline 
 carbonate is being titrated with an acid, or vice versa], litmus 
 cannot be employed as the indicator unless the liquid be boiled. 
 This is owing to the fact that the dissolved carbon dioxide itself 
 exerts a feeble acid reaction upon the litmus, and it is only when 
 the gas is entirely expelled by boiling that the litmus becomes the 
 true indicator of the point of neutrality. It is under these circum- 
 stances that methyl orange is the more convenient indicator. 
 
 Methyl Orange. This substance dissolves in water, giving 
 an orange-coloured solution, which is turned yellow by alkalies 
 and a pink-red colour by acids. The solution for use as an indi- 
 cator is prepared by dissolving 0*5 gram of the solid in 500 c.c. of 
 water ; and one or two drops only should be employed, as its 
 indications are less sensitive if much of it be used. 
 
 Methyl orange is unaffected by carbon dioxide ; it is, therefore, 
 specially adapted for titrations of alkaline carbonates with mineral 
 acids, or vice versa. With this indicator, therefore, these determi- 
 nations can be carried out in the cold. 
 
 Methyl orange is only suitable as an indicator where mineral 
 acids are concerned ; it cannot be used with organic acids. 
 
 Phenolphthalein. A solution of this compound is prepared 
 by dissolving i gram of the solid in 100 c.c. of alcohol. The dilute 
 neutral solution is colourless, but on the addition of an alkali it 
 
 * The aqueous infusion of litmus obtained in this way contains other things 
 besides the blue colouring matter, and the presence of these tends to diminish 
 the sensitiveness of the reagent. The saline substances may be removed by 
 acidifying the extract with hydrochloric acid, and submitting the mixture to 
 the process of dialysis. The purified colouring matter which remains behind 
 on the dialyser is then dissolved in hot water. It is only for very special 
 purposes, however, that it is necessary to prepare such a purified litmus 
 solution. 
 
318 Volumetric Methods of Analysis. 
 
 becomes a deep red colour. This colour is immediately discharged 
 when the liquid is acidified either with mineral or organic acids. 
 
 The chief value of this indicator lies in its applicability to the 
 titration of organic acids. It cannot be employed in cases where 
 carbon dioxide is evolved, since its colour is destroyed by carbonic 
 acid ; but as the acid carbonates (bi-carbonates) do not give the 
 red colour with this compound, it can be made use of to indicate 
 the completion of \hzfirst stage in the neutralisation of a normal 
 carbonate, namely, the conversion of the normal into the acid car- 
 bonate. Phenolphthalein cannot be used in the presence of 
 ammonia. 
 
 Normal Sodium Carbonate. 
 
 (53 grams 0fNa 2 CO 3 per litre?) 
 
 In order to obtain this salt in a state of purity, it is prepared by 
 heating the purest sodium bicarbonate to a dull red heat until no 
 further loss of carbon dioxide and water takes place.* Theoretically, 
 84 grams of bicarbonate should yield 53 of the normal salt ; a 
 slight excess of this proportion, therefore, should be employed. To 
 prepare half a litre of the standard solution, about 43 grams of the 
 pure sodium bicarbonate are heated in a weighed platinum dish to 
 a low red heat for about 10 or 15 minutes. The salt must not be 
 allowed to fuse. It is then cooled in a desiccator and weighed. 
 To ensure that the decomposition has been completed, the dish is 
 again heated for another 10 minutes, and, after cooling in the 
 desiccator, weighed again. The weight of the salt (after deducting 
 the weight of the dish) will be a little over 26*5 grams. By means 
 of a clean spatula or pen-knife, a small quantity is removed, so as 
 to bring the weight to exactly 26*5 grams, the operation being 
 performed without undue exposure of the dry salt. The contents 
 of the dish are then washed out into a beaker, the dish being 
 
 * Owing to its much slighter solubility in water, the bicarbonate is the 
 more easily purified salt ; hence of the two so-called " pure " salts of commerce, 
 the bicarbonate will have a higher degree of purity than the normal salt. The 
 sample employed here, however, must first be tested in the following way : A 
 few grams are dissolved in a small quantity of hot water, which at once reveals 
 the presence or absence of any traces of insoluble impurity. The solution is 
 then acidified with nitric acid (which must be absolutely free from either hydro- 
 chloric or sulphuric acid), and tested for chlorides and sulphates in the usual 
 way. If the salt does not dissolve to a perfectly clear solution, and is not free 
 from chlorides and sulphates, a quantity of it must be purified by recrystal- 
 lisation. A hot saturated solution is made, and filtered by the use of one of 
 the arrangements for warming the funnel shown on p. 210. The filtered liquid 
 is continually stirred as it cools, whereby the salt is deposited in a fine granular 
 condition. The mother-liquor is then decanted, and the crystals drained and 
 dried upon a porous plate. 
 
Acidimetry and Alkalimetry. 319 
 
 thoroughly rinsed with warm water, and the salt completely dissolved 
 by stirring the mixture with a glass rod. The solution is then care- 
 fully poured into a half-litre flask, the beaker being several times 
 rinsed with water. The liquid is then cooled to I5'5 , and water 
 at the same temperature is added, with continual gentle shaking, 
 until the neck is reached. The vessel is then carefully filled to the 
 graduation mark, after which the stopper is inserted, and the 
 contents thoroughly mixed by shaking. 
 
 Instead of bringing the weight of the sodium carbonate to the 
 exact quantity required for half a litre of solution, the whole of the 
 salt in the dish may be employed, and the exact volume of water 
 which will be required in order to make the solution of normal 
 strength is calculated. Thus, suppose the weight of sodium car- 
 bonate in the dish after heating is 26749 grams, instead of 26*5 ; 
 then 
 
 26-5 : 26749 : 500 c.c. : 5047 c.c. 
 
 Hence, after the solution has been made up to the graduation mark 
 in the manner described above, 47 c.c. of water are added from a 
 burette, and the solution then finally shaken up to ensure thorough 
 mixing. 
 
 Normal Sulphuric Acid. 
 (49 grams of H 2 SO per litre.} 
 
 The specific gravity of ordinary oil of vitriol being about r8 
 49 grams will be rather less than 30 c.c. Hence, if this volume be 
 measured out and diluted up to a litre, a solution will be obtained 
 which will have a rough approximation to the required strength. 
 If this solution is then titrated with the standard sodium carbonate, 
 its actual strength can be ascertained ; and by calculation, the 
 volume of water which must be added in order to bring it to the 
 exact normal strength is determined. 
 
 Dilution of the Acid. Thirty cubic centimetres of pure 
 sulphuric acid are gradually poured into about 150 c.c. of water 
 in a flask. The mixture is then cooled by holding the flask under 
 the water-tap, and allowing a stream of water to run over it, at 
 the same time shaking the liquid round within the vessel. When 
 cold, the solution is transferred to a litre flask,* and the volume 
 made up to the 1000 c.c. mark by the addition of water at a 
 temperature of 15 '5. 
 
 * It is not advisable to dilute the strong acid in the litre flask itself, as 
 there is always some risk (although perhaps not much when experience has 
 been gained) of fracturing the vessel ; and it is obviously wiser that a common 
 flask, and not one which has been carefully graduated, should take this risk. 
 
320 Volumetric Methods of Analysis. 
 
 Titratiou of the Acid. A burette is first filled with the 
 dilute acid,* care being taken to remove all air-bubbles from the 
 tap, as explained on p. 308. Twenty-five cubic centimetres of the 
 normal sodium carbonate are then transferred by means of a pipette * 
 to a small beaker, and a single drop of the methyl orange indicator 
 (see p. 317) is added. The beaker is then placed upon a white 
 glazed tile beneath the burette, and the acid gradually run into the 
 alkaline liquid, the solution being gently rotated in order to ensure 
 thorough mixing after each addition of acid. At first 2 or 3 c.c. 
 of the acid may be added at a time ; but as the point of neutrality 
 is approached, smaller and smaller quantities are added at once, 
 until they are reduced to a few drops only, and at last the addition 
 of a single drop produces a permanent red colour in the liquid. 
 
 Correction of the Acid. The exact volume of the acid 
 which has been used in order to neutralise the 25 c.c. of normal 
 sodium carbonate is noted, and from it the volume of water which 
 must be added in order to make the acid exactly normal is calcu- 
 lated. Thus, suppose instead of 25 c.c. of acid, 23'S c.c. were used 
 in neutralising 25 c.c. of the normal alkali ; then 
 
 25 : 23-8 : : 1000 : 952 
 
 That is to say, 952 c.c. of the acid contain as much sulphuric 
 acid as should be contained in 1000 c.c., if it were exactly normal. 
 If, therefore, 952 c.c. of this acid be measured out into a litre-flask, 
 and the volume be then made up to the litre by the addition of 
 water, a correct normal acid will be obtained. 
 
 The solution thus obtained should be once more titrated with 
 the sodium carbonate. 
 
 When the strength of the acid is very nearly, but not exactly, 
 normal, instead of attempting to bring it to the precise normal 
 strength, it may be used as it is, and a correction introduced every 
 time by means of a factor. For example, suppose 25 c.c. of normal 
 sodium carbonate required 24*4 c.c. of the acid instead of 25 c.c. 
 in order to neutralise it, then 
 
 24-6 : 25 :: i : roi6 
 
 That is to say, every cubic centimetre of this acid is equal to 
 i -oi 6 c.c. of an exactly normal acid ; therefore, if the number of 
 
 * Whenever burettes or pipettes are employed for measuring standard 
 solutions, they must either be dry before use, or, if moist, they must be first 
 rinsed out with a small quantity of the solution ; otherwise that portion which 
 is measured out for use would be slightly diluted by the water adhering to the 
 walls of the instrument. 
 
Acidimetry and Alkalimetry. 321 
 
 cubic centimetres of this acid used in any titration be multiplied 
 by the factor roi6, the result will be the number of cubic centi- 
 metres which would have been required if the acid had been strictly 
 normal. 
 
 Or again, if the acid should be a little weaker than the exact 
 normal, a similar correction can be made. Thus, suppose 25 c.c. 
 of normal sodium carbonate required 25-2 c.c. of acid instead of 
 25 c.c. in order to reach the neutral point, then 
 
 25-2: 25 :: i : 0-992 
 
 That is to say, each cubic centimetre of the acid is in reality only 
 equivalent to 0*992 c.c. of normal acid, therefore in this case 0*992 
 is the factor by which the number of cubic centimetres of this acid 
 which might be used, would have to be multiplied in order to con- 
 vert them into cubic centimetres of normal acid. 
 
 Titration of the Acid, using Litmus as Indicator. 
 In the absence of methyl orange, the titration of the acid by means 
 of sodium carbonate may be carried out with litmus as the indi- 
 cator. In this case, however, it is necessary to boil the solution in 
 order to expel the carbon dioxide. The acid is added gradually, 
 until the colour of the litmus changes from blue to purple-red. The 
 solution is then boiled, and as the carbon dioxide is expelled the 
 colour returns to the original blue shade. After boiling for a few 
 minutes, acid is admitted in small quantities, and the liquid boiled 
 up after each addition, until at last the addition of a single drop 
 gives a permanent bright-red colour, which does not change on 
 boiling. 
 
 As a check to the result thus obtained, the following method 
 may be carried out. After noting the volume of acid which has 
 been used in order to reach the point of neutrality, a measured 
 volume of acid in excess is run into the beaker from the burette, 
 and the mixture (which for this purpose is more conveniently 
 contained in a flask than in a beaker) quickly cooled. This excess 
 of acid is then titrated back with normal caustic soda, which 
 (since no carbon dioxide is evolved) can be carried out in the cold. 
 From the amount of caustic alkali used, the actual volume of free 
 acid present is ascertained. If the former titration was correct, 
 this should exactly agree with the measured excess volume which 
 was added ; and if not, it gives the amount of acid which, in the 
 first operation, had been added beyond what was actually necessary 
 for neutrality. 
 
 Control Experiments. The titration of the normal sul- 
 
 Y 
 
322 Volumetric Methods of Analysis. 
 
 phuric acid may be checked by either one or both of the following 
 experiments, which should be made in duplicate 
 
 (1) A small quantity, say 3 or 4 grams, of pure sodium bicarbon- 
 ate is heated in a weighed platinum crucible until the weight is 
 constant, as described above in the preparation of normal sodium car- 
 bonate solution. The contents of the crucible are then dissolved in 
 water, and washed into a beaker. This solution is then titrated in the 
 cold with the standard acid, using methyl orange as indicator ; or, 
 if litmus is employed, the titration is conducted in the boiling solu- 
 tion. From the result, the exact strength of the acid is calculated 
 as follows : Suppose the weight of sodium carbonate obtained after 
 heating the bicarbonate was 2 '66 grams, and that this weight 
 required 50 c.c. of the acid for complete neutralisation, then 
 
 53 : 2*66 ! \ 49 : 2^459 grams = weight of H 2 SO 4 in the 50 c.c. 
 
 Then 49 : 2-459 ; \ 1000 : 50-19 c.c. = volume of strictly normal 
 acid which contains 2*459 grams of H 2 SO 4 , and which therefore 
 would be required to neutralise 2'66 grams of sodium carbonate. 
 Hence the acid used is slightly above the normal strength. Its 
 factor, therefore, will be i '0038 ; thus 
 
 50 : 50-19 : : i : 1-0038 
 
 (2) The strength of the acid may also be determined gravimetri- 
 cally by taking 50 c.c. of the acid, and, after diluting with two or 
 three times its volume of water, precipitating it with barium 
 chloride, and proceeding as described in Part I., p. 258. 
 
 Normal Sodium Hydroxide. 
 
 (40 grams of NaH O per litre. ) 
 
 About 45 grams of pure caustic soda, prepared from the metal,* 
 are dissolved in water which has been recently boiled to expel 
 
 * If pure sodium hydroxide is not to hand, it may be prepared by either of 
 the two following methods, namely, (i) by the action of lime upon sodium 
 carbonate, or (2) by the action of sodium upon water. 
 
 (1) About 70 grams of sodium carbonate are placed in a clean iron saucepan 
 with 700 c.c. of distilled water, and the mixture heated to boiling. Forty grams 
 of good quicklime are made into a paste with water, and gradually added to 
 the boiling solution, and the mixture boiled until the whole of the carbonic acid 
 in the sodium carbonate is precipitated as calcium carbonate. The completion 
 of the reaction is ascertained by withdrawing a small quantity of the mixture 
 and allowing it to settle in a test-tube. If the sodium carbonate has been 
 wholly converted into caustic, there will be no effervescence "when a little of the 
 clear liquid is acidified with a dilute acid. The saucepan is carefully covered, 
 and the contents allowed to settle and become cold. The clear liquid is then 
 decanted off into a stoppered bottle, and a portion of it titrated with standard 
 acid, as described above. 
 
 (2) A clean piece of sodium weighing about 25 grams is cut into small pieces, 
 which are introduced one at a time into recently boiled and cooled distilled 
 water in a platinum or nickel dish. A clock-glass should be placed over tru* 
 cjish after the addition of each piece of sodium, in order to prevent the caustic 
 
A ci dime try and Alkalimetry. 323 
 
 carbon dioxide, and again cooled and the volume of the liquid 
 then made up to one litre. Twenty-five cubic centimetres of this 
 solution are then transferred to a small flask by means of a pipette, 
 and titrated with the standard sulphuric acid in the cold, with 
 methyl orange as indicator. 
 
 If the caustic soda be free from carbonate, the " end reaction " 
 is very sharply defined when litmus is used as the indicator in the 
 cold ; but as the sodium hydroxide, even if free from carbonates 
 to begin with, is always liable to absorb atmospheric carbon 
 dioxide, it is necessary, when employing litmus, to boil off the 
 carbon dioxide during the titration. 
 
 From the results obtained, the volume of water which must be 
 added to the solution of caustic soda in order to bring it to normal 
 strength is calculated as explained in the case of the sulphuric acid 
 (p. 320). When this has been done, and the liquid well shaken up 
 to ensure perfect admixture, a fresh titration should be made ; and 
 if the alkali is still not absolutely normal, \.}\Q factor is found by 
 which the volume used is to be multiplied in order to make the 
 necessary correction. 
 
 The standard sodium hydroxide should be quickly transferred 
 to a store-bottle, the stopper of which is greased with a touch of 
 vaseline ; and it should be exposed as little as possible to the air. 
 
 Normal Hydrochloric Acid. 
 
 (36*5 grams of HC1 per litre."} 
 
 The ordinary pure strong hydrochloric acid contains 28 per cent, 
 of HC1 ; 40 grams, therefore, of the real acid are contained in 143 
 grams of this solution. Since the specific gravity of this aqueous 
 acid is 1*14, 143 grams will be 125 c.c. This volume, therefore, is 
 measured out into a litre flask, and the volume made up to the 
 looo c.c with water at I5'5- 
 
 The actual strength of the acid so obtained is then ascertained 
 by titration with normal sodium hydroxide. Twenty-five cubic 
 
 soda from being thrown out in the event of explosions or deflagrations taking 
 place. The process of thus introducing the sodium in fragments which are 
 sufficiently small to render the operation free from danger, is a somewhat long 
 and tedious operation. It may be considerably hastened, however, and the 
 metal exposed much less to the action of atmospheric carbon dioxide, by placing 
 rather more than the requisite quantity of sodium in a sodium wire-press, and 
 slowly squeezing the metal into wire, which is received in water in a nickel 
 or platinum dish. The fine wire of sodium, as it touches the water, fuses at 
 the end, and is rapidly dissolved in the liquid, and the rate at which the wire 
 is driven out can be exactly regulated so that it shall be continuously fed into 
 the water. 
 
324 Volumetric Methods of Analysis. 
 
 centimetres of the alkali are transferred to a small flask, either litmus 
 or methyl orange being used as the indicator, and the acid added 
 from a burette until the point of saturation is reached. From the 
 result obtained, the amount of dilution necessary to bring the acid 
 to normal strength is calculated. 
 
 When the acid has been thus corrected, it is once more titrated 
 with the caustic alkali. 
 
 As a check upon the result, the acid may be titrated with the 
 normal sodium carbonate, using methyl orange as indicator, and 
 operating in the cold. (Litmus cannot well be used in this case, as 
 the necessary boiling involves loss of hydrochloric acid.) 
 
 The exact strength of the acid may also be controlled gravi- 
 metrically by precipitating 25 c.c. of the liquid with silver nitrate. 
 
 Typical Analyses by means of Standard Acids and 
 Alkalies. 
 
 1. Estimations of the total alkali in samples of caustic 
 alkali, or in soda-ash ; or determinations of the acid in commercial 
 acids or acid liquids, are conducted in practically the same manner 
 as already described for the titration of the alkalies and acids used 
 for standard solutions. In estimating alkalies containing car- 
 bonates, if litmus is employed as indicator instead of methyl orange, 
 the " end reaction " is rendered more certain if a slight excess of 
 the normal acid is run in from the burette, and the mixture boiled 
 to expel the carbon dioxide. The excess of acid is then carefully 
 titrated back with normal caustic soda, which is cautiously added 
 from a burette until one drop restores the blue colour to the litmus. 
 The amount of acid which is represented by the volume of the 
 normal caustic soda thus used, is then deducted from the original 
 total volume of acid employed, which gives the exact amount of 
 acid required to neutralise the alkali under examination. 
 
 2. Estimation of Caustic and Carbonated Alkali in 
 Soda-ash. About 5 grams of the soda-ash are weighed into a flask, 
 and dissolved in water, and the solution made up to 250 c.c. Fifty 
 cubic centimetres of this solution are then transferred by means of 
 a pipette to a small flask, and the total alkali in it determined by 
 titration with normal sulphuric acid, using either litmus or methyl 
 orange as indicator (see above). The amount of alkali (calculated 
 as Na 2 O) contained in 50 c.c. when multiplied by 5 (as only one- 
 fifth of the total was used) gives the amount in the original weight 
 taken, and from this the percentage may be calculated. 
 
 Another 50 c.c. of the solution are next transferred to a loo-c.c. 
 flask, and a solution of barium chloride added so long as any 
 
Acidimetry and Alkalimetry. 3 2 5 
 
 precipitate of barium carbonate is produced.* The liquid is then 
 diluted with water up to 100 c.c. and allowed to settle in the 
 stoppered flask. 
 
 Fifty cubic centimetres of the clear solution (equivalent to one- 
 tenth of the original weight taken for the analysis) are withdrawn 
 with a pipette, transferred to a small beaker and titrated with deci- 
 normal hydrochloric acid in the cold, with either litmus or methyl 
 orange as indicator. (Sulphuric acid cannot be employed on account 
 of the barium present.) 
 
 EXAMPLE. Weight of soda-ash taken = 5*05 grams. 
 First Solution. 50 c.c. = one-fifth total soda-ash. 
 50 c.c. taken and excess of normal sulphuric acid added 
 
 (i c.c. = 0-031 gram Na 2 O). Volume of acid used... 18-50 c.c. 
 Titrated back with normal caustic soda (i c.c. = i c.c. 
 
 normal acid). Volume required ... 2*37 
 
 Hence volume of acid employed in neutralising 50 c.c. 
 
 soda-ash l6>1 3 
 
 Therefore weight of Na 2 O in \ _ , 
 original weight of soda-ash/ 
 
 x 0-031 x 5 = 2-500 grams 
 
 Percentage of Na 2 O,\ 2-50 x 100 _ A0 .-- 
 or total alkali } ~ 5-05 
 
 Second Solution. 50 c.c. = one-tenth total soda-ash. 
 50 c.c. taken, and titrated with deci-normal hydrochloric 
 acid (i c.c. = 0-0031 gram Na 2 O). Volume of acid 
 required ... ... 6-3 c.c. 
 
 Therefore weight of Na 2 O\ 
 (present as NaHO) in I = 6 
 original weight of soda- 1 
 ash J 
 
 Hence the percentage of j _ 0*1953 x 100 _ g.gg 
 Na 2 O present as caustic / ~ 5^5 
 
 Then total Na 2 O in original weight of soda-ash = 2*500 
 Na 2 O present as caustic alkali =0-1953 
 
 Therefore Na 2 O present as carbonated alkali = 2-3047 grams 
 Hence the percentage of Na 2 O present as^ 2 '347X iQQ^g.gg 
 carbonate j 5*05 
 
 * As soon as the whole of the sodium carbonate has been decomposed and 
 the carbonic acid precipitated as barium carbonate, the barium chloride begins 
 to interact with the sodium hydroxide present, forming sodium chloride and 
 barium hydroxide, the latter being equivalent to the sodium hydroxide. The 
 amount of barium chloride used, however, should be as little in excess of that 
 required for the decomposition of the alkaline carbonate as possible, for it is 
 found in practice that when barium hydroxide is produced in any quantity, the 
 results obtained are too low. 
 
326 Volumetric Methods of Analysis. 
 
 The actual percentage of caustic alkali and carbonated alkali 
 may readily be calculated ; thus 
 
 31 : 40 :: 0*1953 = 0*2474 = NaHO equivalent to 0-1953 grams Na 2 O 
 ^474_x_Loo =4 . 8 Q 
 
 55 
 
 Similarly, 31 : 53 :: 2*3047 = 3-938 = Na 2 CO 3 equivalent to 
 
 2*3047 grams Na. 2 O 
 2-3047 x loo = y7 . 98 = 0/olTM!<)3 
 
 55 
 
 Alternative Methods. (a) The soda-ash is dissolved as 
 before, and the total alkali determined in the manner above 
 described. Then a fresh portion (50 c.c.) of the same solution is 
 transferred to a flask, and the whole of the carbonate precipitated 
 by adding barium chloride. Phenolphthalein is then added for the 
 indicator, and the mixture at once titrated with deci-normal hydro- 
 chloric acid until the red colour is discharged. The presence of 
 the precipitated barium carbonate does not affect the indicator, 
 and the moment the point is reached when the caustic alkali is 
 neutralised, and the standard acid begins to act upon the barium 
 carbonate, the liberated carbon dioxide discharges the colour of the 
 indicator. 
 
 (0) In this method the total alkali is not first determined. The 
 process depends upon the fact that when sodium carbonate is 
 acted upon by acids, the neutralisation takes place in two stages, so 
 that when half the acid necessary for complete saturation has been 
 added, the carbonate has been converted into the bi-carbonate, 
 which has no power to give a red colour with phenolphthalein. If, 
 therefore, to a mixture of sodium hydroxide and carbonate coloured 
 red with this indicator, normal acid be added, the phenolphthalein 
 will show the point when the 'whole of the caustic soda and one- 
 half of the carbonate have become saturated. 
 
 Fifty cubic centimetres of the solution prepared as described 
 above is coloured with phenolphthalein, and titrated with normal 
 sulphuric acid. As soon as the colour is discharged, a drop of 
 methyl orange is added, and the titration continued until the yellow 
 liquid becomes pink. 
 
 The volume of acid required in the second stage (i.e. during the 
 second half of the saturation of the carbonate present) will be one- 
 half of that which has altogether been used up in neutralising the 
 alkaline carbonate. Therefore, if this volume of acid be doubled, 
 and the result subtracted from the total acid used in the whole 
 
Acidimetry and Alkalimetry. 327 
 
 process, the difference will represent the volume which was used 
 in saturating the caustic alkali present. 
 
 Thus, in the case of the solution of soda-ash used in the first 
 example (containing 5 '05 grams in 250 c.c.), 50 c.c. were taken, and 
 titrated with normal sulphuric acid until the phenolphthalein lost 
 its colour. 
 
 Volume of acid required ... ... ... 8*7 c.c. 
 
 Methyl orange was then added, and the volume 
 cf acid required for complete neutralisation 7-45 
 
 Total acid used ... 16-15 > 
 = 7'45 x , = ,4-90 c.c. 
 ' = r c.c. 
 
 These, when calculated as percentages, as in the former example, 
 give 
 
 Percentage carbonate = 78' i 
 Percentage caustic = 4/9 
 
 3. Estimation of Ammonia. 
 
 (a) Direct Method. A solution of the ammonium salt in water 
 is boiled with a solution of caustic soda, and the ammonia thus 
 expelled is absorbed in a measured volume of normal sulphuric 
 acid. The excess of acid is then ascertained by titration with 
 normal sodium hydroxide ; and by deducting this excess from the 
 total volume of acid employed, the volume of acid which has been 
 saturated by the ammonia is found, and from this the weight of 
 ammonia is calculated. 
 
 From i to 2 grams of ammonium chloride are weighed out into 
 the flask F (Fig. 53, p. 250), and dissolved by the addition of a 
 small quantity (40 to 50 c.c.) of water. One hundred cubic centi- 
 metres of normal sulphuric acid are run into the absorption flask 
 through the scrubber-tube, and the process of distilling and absorb- 
 ing the ammonia is carried out exactly as described on p. 251. 
 
 After the whole of the ammonia has been expelled, the apparatus 
 is disconnected, and the end of the leading tube and the contents 
 of the scrubber-tube are thoroughly rinsed into the flask, so as to 
 wash out of the latter every trace of the acid with which the broken 
 glass was moistened. A few drops of methyl orange are then 
 added, and the liquid titrated with normal caustic soda. 
 
 The quantity of acid which is equivalent to the volume of the 
 normal soda required for neutralisation, is the excess which has 
 
328 Volumetric Methods of Analysis. 
 
 remained over after the whole of the ammonia has been saturated. 
 On deducting this from the total volume of acid used (namely, 100 
 c.c.), the actual amount of acid neutralised by the ammonia is 
 found, and by calculation the weight of ammonia, NH 3 (or of 
 ammonium, N H 4 ), is ascertained. 
 
 (b) Indirect Method. In this method the weighed quantity of 
 ammonium salt is boiled with a measured volume of normal sodium 
 hydroxide, until the ammonia is all expelled. A certain quantity 
 of the soda is thus decomposed, being converted into the sodium 
 salt of the acid which was in combination with the ammonia. For 
 example, supposing the ammonium salt to have been the sulphate 
 
 (NH 4 ) 2 SO 4 + 2NaHO = Na 2 SO 4 + 2H,O + 2NH S 
 If, when the whole of the ammonia has been expelled, the remain- 
 ing undecomposed caustic soda be determined by means of normal 
 acid, the amount of soda which has been used up in expelling the 
 ammonia will be found ; and from this the quantity of ammonia so 
 displaced can be calculated. 
 
 About 2 grams of pure ammonium sulphate are weighed out 
 into a small flask, and 30 c.c. of normal caustic soda added. The 
 mixture is then boiled until the escaping steam produces no brown 
 stain upon a strip of turmeric paper. Three or four drops of 
 litmus solution are then added, and the liquid titrated with normal 
 sulphuric acid. The volume of normal soda equivalent to the 
 number of cubic centimetres of acid used to produce neutrality, 
 when deducted from the total volume of soda employed (namely, 
 30 c.c.) gives the number of cubic centimetres of soda which were 
 decomposed in the process of expelling the ammonia. 
 
 From the above equation, it is seen that 40 grams of caustic 
 soda are used up in the expulsion of 17 grams of ammonia ; there- 
 fore each cubic centimetre of normal soda (0*040 gram) is equivalent 
 to 0*017 gram of ammonia, NH 3 . 
 
 4. Estimation of the " Hardness " of Water * (Hehner's 
 method). 
 
 (a) Temporary hardness (due to the presence of calcium or 
 magnesium carbonate) is determined by means of standard sul- 
 phuric acid of one-fiftieth normal strength. This is prepared by 
 diluting 20 c.c. of the normal acid up to 1000 c.c. with water. One 
 cubic centimetre of this acid is equivalent to o'ooi gram CaCO 3 . 
 
 One hundred cubic centimetres of the water to be examined are 
 transferred to a beaker, and 3 drops of methyl orange added. 
 
 * For Clark's soap-test for hardness in water, see Precipitation methods, 
 p. 37 1. 
 
Addimetry and Alkalimetry. 329 
 
 The sulphuric acid is then delivered from a burette until the 
 
 50 P 
 solution becomes red. 
 
 Each cubic centimetre of acid represents i milligram of CaCO 3 
 in loo c.c. of water ; therefore it also stands for parts per 100,000. 
 E.g. suppose 13 c.c. of acid were required to exactly neutralise the 
 carbonate ; then 100 c.c. of the water contain 0-013 gram of CaCO 3 ; 
 or its equivalent of MgCO 2 , and 100,000 parts contain 13 parts of 
 CaCO 3 (or the equivalent of MgCO 3 ). 
 
 One part of CaCO 3 per 100,000 of water is termed i degree of 
 hardness ; hence the water in question has a temporary hardness 
 of 13 degrees. 
 
 (b} Permanent hardness (due mainly to the presence of calcium 
 or magnesium sulphate) is estimated by means of standard sodium 
 carbonate of one-fiftieth the normal strength (prepared by diluting 
 20 c.c. of the normal solution to i litre with water). One cubic 
 centimetre of this solution will precipitate o'ooi gram CaCO 3 . 
 Excess of this is added to the water, which is then evaporated to 
 dryness. The residue is treated with water, and after removing 
 the precipitated carbonates by filtration, the excess of sodium 
 
 N 
 carbonate is determined by means of sulphuric acid. 
 
 One hundred cubic centimetres of the water are transferred to 
 
 N 
 a platinum dish, and a measured volume of the sodium carbonate 
 
 considerably in excess of that required for complete precipitation 
 is added. The solution is then evaporated to dryness on a steam 
 bath. The residue is taken up with water and filtered, the dish 
 and filter being thoroughly washed. The sodium carbonate in the 
 
 N 
 liquid is then determined by titration with the sulphuric acid. 
 
 The excess of sodium carbonate thus found, deducted from the 
 total volume added, gives the number of cubic centimetres used in 
 removing the permanent hardness. Each cubic centimetre so 
 employed represents o'ooi gram of calcium carbonate (formed from 
 the sulphate *) precipitated from 100 c.c. of water ; it therefore 
 also represents parts per 100,000. 
 
 * From the following equations, it will be seen that, although sodium 
 carbonate precipitates calcium carbonate from the soluble bicarbonate of lime 
 (temporary hardness) as well as from the sulphate (permanent hardness) that, 
 in other words, it removes both temporary and permanent hardness it is only 
 in decomposing the sulphate that the alkali is destroyed. In the other case 
 the solution contains the equivalent of sodium bicarbonate 
 
 H 8 CO 3 ,CaC0 3 + Na 2 C0 3 = CaCO 3 + 2HNaCO 3 
 CaS0 4 + NajCO, ^ CaCO 3 + Na 2 SO, 
 
330 Volumetric Methods of Analysis. 
 
 N 
 For example, 100 c.c. of the water were taken, and 50 c.c. of - 
 
 sodium carbonate added. After evaporating to dryness and ex- 
 
 N 
 tracting with water, the solution required 41 c.c. of sulphuric 
 
 acid to neutralise it. 50 41 = 9, therefore 9 c.c. of the sodium 
 carbonate were used up in removing the permanent hardness, 
 hence the water contained 9 degrees of permanent hardness. The 
 temporary hardness added to the permanent hardness gives the 
 total hardness of the water. 
 
SECTION III. 
 
 VOLUMETRIC METHODS BASED UPON OXIDATION AND 
 REDUCTION. 
 
 IN the various processes which are included in this section, there 
 is always an oxidising agent (in other words, a reducible substance) 
 and an oxidisable body, i.e. a reducing agent. 
 
 Sometimes the oxidising material is the standard reagent, 
 which, by its own reduction, furnishes the measure of the amount 
 of the oxidisable substance which is present. Thus, the reduction 
 of a known volume of standard potassium permanganate, or potas- 
 sium dichromate, will afford a measure of the quantity of iron 
 which has been oxidised from the ferrous to the ferric condition ; 
 the ferrous salt in this case being the oxidisable substance, or the 
 reducing agent. 
 
 In other cases the standard reagent is the reducing agent, 
 which undergoes oxidation at the expense of some oxidising (i.e. 
 reducible) compound present, and thereby affords the means of 
 estimating the amount of such a substance. Thus, the oxidation 
 of a known volume of standard sodium arsenite into arsenate, may 
 become the measure of the amount of chlorine present in bleach- 
 ing powder ; or of the quantity of free iodine in a solution. The 
 chlorine or the iodine in this case is the oxidising agent, not by 
 directly giving oxygen, as in the case of potassium permanganate, 
 but by indirect oxidation through the intervention of water ; the 
 hydrogen of which is taken by the halogen, while an equivalent 
 of oxygen is liberated. 
 
 The most important of the oxidising standard solutions are 
 potassium permanganate, potassium dichromate, and iodine ; while 
 the reducing compounds in commonest use are ferrous salts, oxalic 
 acid, sodium thiosulphate, sodium arsenite. 
 
332 Volumetric Analysis. 
 
 Deci-normal Potassium Permanganate. 
 
 (yi6 grams ofKMuO^per litre*} 
 
 3' 1 6 grams of the purest recrystallised and dry salt f are weighed 
 out and dissolved in water in a litre flask, and the solution made 
 up to looo c.c. with water at 15. 
 
 Solutions of potassium permanganate slowly undergo decom- 
 position by light ; if, therefore, the solution is to be kept for any 
 length of time, it is preferably transferred, either to a deep-blue 
 bottle, or to one covered with black paper, and kept in a cupboard. 
 The solution must not be allowed to come in contact with any 
 organic matter ; on this account, burettes with glass taps must 
 always be employed for this reagent. 
 
 Titration of Deci-normal Permanganate. (a) By means 
 of Ferrous Sulphate 
 
 2KMn0 4 + ioFeSO 4 + 8H 2 SO 4 = K 2 SO< + 2MnSO 4 + 8H 2 O 
 
 + 5Fe 2 (S0 4 ) 3 
 
 or, expressing only the relation between the available oxygen and 
 the ferrous oxide 
 
 50 + loFeO = 5Fe 2 O 3 
 
 The deci-normal solution being one-tenth of the strength of the 
 normal solution which, in its turn, contains (in the case of per- 
 manganate) only one-tenth of the formula-weight of the salt per 
 litre will contain 0*0008 gram of available oxygen per cubic centi- 
 metre ; or each cubic centimetre will be equivalent to 0-0056 gram 
 of iron in the ferrous state ; i.e. it will be capable of oxidising this 
 weight of iron from the ferrous to the ferric condition. 
 
 For the purposes of this titration, the ferrous sulphate is pre- 
 pared by dissolving 0*5 gram of the purest soft iron wire in 
 sulphuric acid, with exclusion of air, the wire being clean and free 
 from rust.t 
 
 About 80 c.c. of dilute sulphuric acid (i part acid to 5 parts 
 water) are placed in a 25o-c.c. flask fitted with a rubber cork and 
 bent glass tube. The air in the flask is then expelled by removing 
 the cork and introducing two or three crystals of pure sodium 
 carbonate, the flask being in a vertical position. As soon as the 
 carbonate has dissolved, the weighed quantity of iron is dropped 
 
 * See p. 314. 
 
 f If the ordinary "pure" commercial salt be used, about 3-2 to 3-5 grams 
 should be taken. 
 
 J The fine iron binding wire used for flowers is the best for the purpose ; it 
 contains 99-6 per cent, of iron. 
 
Oxidation and Reduction Methods. 
 
 333 
 
 in. The cork is then inserted, and the flask supported in the 
 manner shown in Fig. 72, with the tube dipping into a solution of 
 sodium carbonate in a small beaker. The flask being in this in- 
 clined position, the fine spray thrown up during the solution of the 
 iron strikes against the sides of the flask and falls back into the 
 liquid. The flask is gently 
 heated by means of a small 
 flame until the iron is wholly 
 dissolved, and only a few minute 
 particles of carbon remain. The 
 lamp is then withdrawn, and the 
 flask allowed to cool. As it does 
 so, the solution in the beaker is 
 gradually drawn up the tube, 
 but the first drops which enter 
 the flask at once cause an effer- 
 vescence of carbon dioxide 
 which drives the liquid down 
 again, and at the same time fills 
 the flask with carbon dioxide. 
 When it has partially cooled in 
 this way, the cork is removed, 
 and air-free distilled water (pre- 
 pared by boiling the water, and 
 again quickly cooling it) is 
 added until the solution is within 
 about 20 or 30 c.c. of the gra- 
 duation mark. The flask is then 
 closed with a rubber stopper, and the contents made quite cold by 
 holding the vessel in a stream of cold water. The solution is then 
 made up to 250 c.c. by the further addition of cold air-free water. 
 
 Fifty cubic centimetres of this solution are transferred by means 
 of a pipette to a small flask, and diluted by the addition of about 
 half the volume of air-free distilled water. The flask is placed 
 upon a white tile, and the deci-normal permanganate solution added 
 from a burette until the colour of the reagent ceases to be destroyed, 
 and a faint pink tint is imparted to the solution. 
 
 Four separate experiments should be made, taking 50 c.c. of the 
 iron solution each time, in order to gain practice in judging when 
 the first appearance of the permanent pink colour takes place. 
 After the experience thus gained, in subsequent duplicate titrations 
 the volume of the reagent used should agree to cri of a cubic 
 
334 Volumetric Analysis. 
 
 centimetre. From the results obtained, the exact strength of the 
 permanganate is calculated ; thus 
 
 o'5 gram of iron wire was dissolved in 250 c.c. of liquid. 
 Fifty cubic centimetres of the solution therefore contain o'i gram 
 of iron. But since the iron wire contained 99*60 per cent, of iron, 
 the actual weight of iron present is found by multiplying the 
 weight of the wire by 0-996. Hence 50 c.c. of the solution contain 
 not o - i gram, but 0*0996 grams of iron. Three titrations were 
 made, which required respectively 17*64 c.c., 1774 c.c., and 17*70 
 c.c. of the permanganate, or 17*69 as a mean figure. 
 
 Then 17*69 c.c. : i c.c. : : 0*0996 gram : 0*00563 gram 
 
 Therefore the solution is very slightly stronger than the exact 
 deci-normal, since i c.c. should be equivalent to 0-0056 gram of 
 iron. 
 
 (^) Titration of Permanganate with Ferrous Ammonium Sul- 
 phate, FeSO 4 ,(NH 4 ) 2 SO 4 ,6H 2 O. The formula-weight of this salt 
 (392) is exactly seven times the weight of the iron it contains (56) ; 
 hence in 3*5 grams of the compound there will be 0-5 gram of iron. 
 
 Exactly 3*5 grams of the pure salt are weighed out and dissolved 
 in air-free water, and the solution made up to 250 c.c. Fifty cubic 
 centimetres of this solution, therefore, will contain 0*1 gram of iron. 
 For the titration, 50 c.c. of the solution are transferred to a small 
 flask, and 10 or 15 c.c. of dilute sulphuric acid added. The deci- 
 normal solution is then run in, as in the former case, until the pink 
 colour persists. 
 
 (c) Titration of Permanganate with Oxalic Acid, H 2 C 2 O 4 ,2H 2 O. 
 
 2KMnO 4 + 5H 2 C 2 O 4 4- 3H 2 SO 4 = K 2 SO 4 + 2MnSO 4 + ioC(X 
 
 + 8H 2 
 or 50 + SHoCO, = 5H 2 O + ioCO, 
 
 Hence 8 parts of oxygen (i equivalent) is capable of oxidising 45 
 parts of oxalic acid (calculated as anhydrous), or 63 parts of the 
 crystallised compound. Therefore i c.c. of the deci-normal per- 
 manganate, which is equivalent to 0*0056 gram of iron, will be 
 equivalent also to 0*0063 gram of crystallised oxalic acid. 
 
 6*3 grams of recrystallised and dry * oxalic acid are weighed 
 out, dissolved in water, and the solution diluted to i litre. The 
 solution so obtained is deci-normal oxalic acid ; but, in order to make 
 sure that the acid is pure, and of the exact state of hydration repre- 
 sented by the formula C 2 H 2 O 4 ,2H 2 O, the precise strength of the 
 
 * The term dry signifies freedom from adherent moisture, and does not 
 refer to the anhydrous compourfd. 
 
Oxidation and Reduction Methods. 335 
 
 solution should be determined by titration with a deci-normal caustic 
 alkali, using phenolphthalein as indicator. 
 
 Twenty-five cubic centimetres of the deci-normal oxalic acid are 
 transferred to a small flask, and about half the volume of dilute 
 sulphuric acid added.* The mixture is then gently heated to a 
 temperature from 50 to 60, and the permanganate solution 
 slowly run in from a burette. In this reaction the discharge of 
 the colour of the permanganate is not so instantaneous as in the 
 case of the ferrous compounds, there being an intermediate forma- 
 tion of a brown colour, which then fades away, leaving the liquid 
 colourless. Hence it is necessary to add the permanganate 
 cautiously, especially as the end of the reaction approaches, in 
 order that the first signs of the permanent pink colour may be 
 detected. Two or three titrations should be made (using various 
 quantities of the deci-normal oxalic acid) for practice in observing 
 the end reaction. 
 
 If the oxalic acid be strictly deci-normal (as determined by 
 means of deci-normal alkali), and the permanganate be likewise 
 exact, 25 c.c. of the acid should require 25 c.c. of permanganate.! 
 
 Typical Analyses "by means of Potassium Perman- 
 ganate. 
 
 (i) Estimation of Iron in Iron Ores4 If the iron is 
 present in \heferrous state, the substance is dissolved with careful 
 exclusion of air, and the solution titrated as already described. 
 If the iron is in \heferric condition, it is reduced (after solution) 
 by the method given below. When the metal is present partly as 
 ferrous and partly as ferric iron, the ferrous iron is first estimated 
 in one portion, and a second portion is then subjected to the 
 reducing agent, and the total amount of iron determined. The 
 difference between the two results represents the iron which existed 
 in the ferric state. 
 
 (a) Iron in the Ferrous State. About 0-5 gram of spathic iron 
 ore, in fat finest possible powder, is weighed out and dissolved in 
 
 * The presence of acid (preferably sulphuric acid) is necessary in all titra- 
 tions with potassium permanganate, in order to keep the manganous oxide in 
 solution. Nitric and hydrochloric acids are prejudicial ; with the latter acid 
 (unless present only in small quantities and the liquid is cold), chlorine is 
 liberated by the oxidising action of the permanganate. 
 
 f Dilute solutions of oxalic acid, such as the deci-normal, are liable to 
 undergo decomposition if preserved ; hence, if the solution has been prepared 
 any length of time, it should be retitrated with alkali before being used. 
 
 J Such as magnetic oxide, red or brcnvn hematite, clay ironstone, spathose, 
 black band, 
 
336 Volumetric Analysis. 
 
 20 c.c. of a mixture of equal volumes of strong hydrochloric acid 
 and water, with exclusion of air, in the apparatus described on p. 
 333.* The acid is gently boiled by means of a small flame until 
 the iron is entirely extracted. The solution is then largely diluted 
 by adding 100 c.c. of cold air-free water, and at once titrated by 
 means of deci-normal permanganate.! 
 
 (Jj) Total Iron. A second portion of ore, about o'5 gram, is 
 weighed out and dissolved in a similar quantity of acid. The 
 operation may be conducted in a flask supported as shown in Fig. 
 72, but the cork with the leading tube may be removed. 
 
 When the solution of the ore is complete, the liquid is con- 
 siderably diluted with water, and about 1 5 c.c. of strong sulphuric 
 acid added. A few fragments of pure zinc (i.e. free from iron J) 
 are then introduced into the flask, and the cork with its leading 
 tube is inserted. Hydrogen is evolved by the solution of the zinc 
 in the acid liquid, and the iron existing in the ferric state is thereby 
 reduced to \hzferrous condition. The action is allowed to continue 
 until the zinc is entirely dissolved, the process being aided towards 
 the end by the application of heat. In order to test whether the 
 reduction of the iron is complete, a drop of the liquid is with- 
 drawn upon the end of a fine glass rod, and brought into contact 
 with a drop of a solution of ammonium thiocyanate upon a white 
 tile. 
 
 Any remaining ferric salt will be revealed by the formation of 
 the red colour ; in which case the process must be continued by 
 the addition of more zinc, and, if necessary, of more acid also. 
 When the iron is entirely in the ferrous state, the solution is quickly 
 cooled, and then titrated with the permanganate. 
 
 By deducting the amount of iron present as ferrous iron from 
 the total iron in the ore, the amount of iron which is in the ferric 
 state is ascertained. 
 
 * In the case of the carbonate of iron, the carbon dioxide evolved will serve 
 to displace the air without the addition of sodium carbonate. 
 
 t It is very necessary that the amount of hydrochloric acid present should 
 be quite small, and the solution cold and dilute, otherwise the analysis will be 
 vitiated by the action of the acid upon the permanganate, whereby chlorine is 
 liberated; thus 
 
 KMn0 4 + 8HC1 = KC1 + MnCl 2 + 4 H,O + sCl 
 
 J If pure zinc is not available, the zinc that is employed must be weighed 
 before being introduced, and the amount of iron present must be found by dis- 
 solving a similar quantity of the zinc in acid, and titrating with permanganate. 
 Or magnesium may be used instead of zinc. Other methods by which iron 
 may be reduced, but which are more applicable when the titration is to be made 
 with potassium dichromate, are described on pp. 343, 344. 
 
Analyses by means of Potassium Permanganate. 
 
 
 EXAMPLE. 0-5 gram of spathic iron ore. 
 (a) Ferrous Iron. 
 
 Permanganate required ............ 34*4 c.c. 
 
 (i c.c. = 0-0056 Fe, - 0-0072 FeO, = 0*0080 Fe a O 3 ) 
 
 Then 34-4 x 0-0072 = 0-24768 = weight of FeO 
 
 , 0*24768 x 100 
 and - = 49*53 = per cent, of FeO in the ore 
 
 (6) Ferric Iron. 0-5 gram of ore. 
 Permanganate required ........ ...... 36-6 c.c. 
 
 Deducting the volume used in the first titration ...... 34-4 
 
 Gives the volume required for the iron present in the 
 
 ferric state ... ... ... ... ... ... 2*2 ,, 
 
 But the zinc used contained 0*08 per cent, of iron. Five grams were 
 
 employed ; therefore IO( ^ 5 = 0*004 gram of iron which was 
 
 introduced. 
 
 The volume of permanganate used in oxidising this 
 
 quantity of iron is 0*0056 : 0*0040 :: I : 0*71 c.c. 2" 2 
 
 0-71 
 
 Deducting this, gives the true volume of permanganate 
 
 required for the iron originally present as ferric iron 1-49 c.c. 
 
 Then 1-49 x o'ooSo = 0*01192 = weight of Fe 2 O 3 
 
 and = 2-38 - per cent, of Fe 2 O 3 in the ore 
 
 (2) Estimation of Available Oxygen in Manganese 
 
 Ores. The available oxygen in a manganese ore is that oxygen 
 which can be made use of for oxidising purposes when the ore is 
 decomposed by an acid. Thus, in the sesquioxide one-third, and 
 in the dioxide one-half, of the total oxygen is available for this pur- 
 pose. When acted upon by sulphuric acid, this oxygen is evolved 
 as such ; thus 
 
 Mn Og + 2H 2 SO 4 = 2MnSO 4 + 2H 2 O + O 
 MnO 2 -1- H 2 SO 4 = MnSO 4 + H 2 O + O 
 
 while, when decomposed by hydrochloric acid, chlorine is evolved, 
 in each case in quantity equivalent to the oxygen 
 
 Mn 2 O 3 + 6HC1 = 2MnCl, + 3H 2 O + C1 2 
 MnO 2 + 4HC1 = MnCl 2 + 2H 2 O + C1 2 
 
 Methods for valuing manganese ores are based upon both of 
 
 z 
 
338 Volumetric Analysis. 
 
 these reactions. In the first case the evolved oxygen is used in 
 oxidising either a ferrous salt or oxalic acid ; while in the second, 
 the chlorine is made to liberate its equivalent of iodine, which is 
 then determined by means of sodium thiosulphate. 
 
 In this place, examples based upon the first reaction are given ; 
 the second method is described on p. 356. 
 
 (a) By the Oxidation of Ferrous Sulphate. 
 
 Mn0 2 + 2H 2 S0 4 + 2FeS0 4 = Fe 2 (SO 4 ) 3 + MnSO 4 + 2H,O 
 
 That is to say, I atom or 16 parts of oxygen (which is the 
 representative of i equivalent (or 87 parts) of manganese dioxide), 
 will oxidise 2 atoms or 112 parts of ferrous iron. 
 
 If, therefore, manganese dioxide be dissolved in sulphuric acid, 
 in the presence of a known quantity of ferrous salt (which must be 
 in excess), every 8 parts of oxygen, or 43*5 parts of MnO., will oxidise 
 56 parts of Fe ; and if the amount of ferrous salt which remains 
 unoxidised be then determined by means of deci-normal perman- 
 ganate, the quantity of iron oxidised, and therefore the amount 
 of manganese dioxide, can be calculated. 
 
 About .1 gram of soft iron wire is weighed out and dissolved in 
 about 40 c.c. of dilute sulphuric acid (i part acid to 4 parts water) 
 in the apparatus shown on p. 333. 
 
 As soon as it is completely dissolved, 075 to i gram of finely 
 powdered pyrolusite, previously dried at 100, is introduced into 
 the flask, and the cork with its tube replaced. The mixture is 
 gently heated until the ore has entirely dissolved, or until all visible 
 black particles have disappeared. The solution is cooled in the 
 manner described on p. 333, diluted with air-free water, and the 
 excess of ferrous salt titrated with permanganate. 
 
 The excess of unoxidised iron thus found, deducted from the 
 weight of iron taken, gives the quantity of iron which has been 
 oxidised by the manganese dioxide ; and from this, by the equation 
 given above, the percentage of available oxygen either as oxygen, 
 or in terms of manganese dioxide can be calculated. Thus 
 
 Volume of permanganate \ 
 
 (i c.c. = 0-0056 Fe) required / ~ 
 
 22-4 x o'oo56 =0*125 gram = excess of Fe 
 
 i gram iron wire (99-6 per j 6 = ^ iron taken 
 
 cent, of Fe) / 
 
 therefore 0-996 - 0-125 = 0-871 gram = iron oxidised by 
 075 gram of the ore 
 
Analyses by means of Potassium Permanganate. 339 
 
 Fe Fe MnO a 
 
 Then 56 : 0-871 : : 43-5 : 0-6765 gram MnO, 
 
 and gZ.5 x ioo <== 90 o = percentage of MnO 2 in the 
 
 '75 ore 
 
 (b] By the Oxidation of Oxalic Add 
 
 MnO 2 -f H 2 SO, + C,,H,O 4 = MnSO 4 + 2CO 2 + 2H 2 O 
 
 Hence every 8 parts of oxygen, or 43-5 parts of MnO 2 , will oxidise 
 63 parts of crystallised oxalic acid. About 0*25 gram of the 
 manganese ore is weighed into a flask, and 10 c.c. of normal oxalic 
 acid added. Twenty-five cubic centimetres of dilute sulphuric acid 
 (i part acid to 4 parts water) are added, and the mixture gently 
 warmed until the decomposition is complete and no black particles 
 are visible. The solution is then cooled, and made up to 100 c.c. 
 with cold water. The excess of oxalic acid remaining is titrated 
 with deci-normal permanganate, taking 25 c.c. of the solution. 
 The amount of oxalic acid thus found, deducted from the total 
 employed, gives the quantity which has been oxidised by the 
 manganese ore, from which the percentage of available oxygen 
 or of manganese dioxide can be calculated. For example, after 
 acting upon 0-25 gram of pyrolusite 
 
 25 c.c. of the solution required of permanganate (i c.c. 
 
 = 0-0063 gram oxalic acid) ............ 12*1 c.c. 
 
 Multiplying by 4 gives for the total excess of oxalic acid 48-4 
 
 48*4 x 0-0063 = weight of oxalic acid in excess = 0-3049 gram 
 
 Weight of oxalic acid in 10 c.c. normal solution... 0*63 
 Deducting the excess ............ 0-3049 
 
 Gives the weight of oxalic acid oxidised by the ore 0-325 1 
 Then 63 : 0*3251 : : 43-5 = 0-223 = weight of MnO 2 
 
 and - 22 Q .* I0 = 89-2 = percentage of MnO, 
 
 (3) Estimation of Tin. When stannous chloride is mixed 
 with ferric chloride, or when metallic tin is dissolved in ferric 
 chloride mixed with hydrochloric acid, the stannous chloride is 
 oxidised to stannic chloride at the expense of the ferric chloride. 
 A quantity of ferrous iron, equivalent to the weight of tin, is thus 
 formed in the solution according to the equation 
 
 Fe 2 Cl 6 + SnCl 2 = SnCl 4 + 2FeCl 2 
 
 Every 56 parts of iron converted into the ferrous state, therefore, 
 is equivalent to 59 parts of tin. The amount of ferrous salt so 
 produced is titrated by means of potassium permanganate. 
 
34O Volumetric Analysis. 
 
 About 0-5 gram of granulated tin is weighed into a 250-0.0. 
 flask, and a few cubic centimetres of a moderately strong solution 
 of ferric chloride, together with a little hydrochloric acid, are added. 
 A crystal of sodium carbonate is dropped into the mixture in order 
 to replace the air with carbon dioxide, and the flask loosely corked 
 and gently warmed. When the tin has wholly dissolved, the 
 solution (still yellow with excess of ferric chloride) is diluted up to 
 250 c.c. with cold air-free water. Fifty cubic centimetres of this 
 liquid are then withdrawn by means of a pipette, and titrated with 
 deci-normal permanganate.* 
 
 (4) Estimation of Calcium. 
 
 (a] The calcium is precipitated in the form of calcium oxalatc 
 by means of ammonium oxalate, in the manner described for the 
 gravimetric estimation, p. 231. The precipitate, after being 
 thoroughly washed to remove every trace of ammonium oxalate, 
 is dissolved in a small quantity of warm hydrochloric acid, and the 
 solution diluted with water. Sulphuric acid is then added, and the 
 mixture heated to between 60 and 70, and the free oxalic acid 
 titrated by means of permanganate. Sixty-three parts of oxalic 
 
 N 
 acid found, are equivalent to 20 parts of calcium ; or i c.c. 
 
 permanganate, which is equal to 0-0063 gram oxalic acid, represents 
 o'oo2 gram of calcium. 
 
 About o'5 gram of marble or limestone is weighed out, and dis- 
 solved and precipitated by means of ammonium oxalate in the 
 manner described on p. 231. The perfectly washed precipitate is 
 then washed into a flask by thrusting a glass rod through the apex 
 of the filter, and that which remains adhering to the paper is dis- 
 solved off by treating it with a little warm hydrochloric acid. The 
 filter is thoroughly rinsed with water, and if the whole of the 
 precipitate in the flask is not dissolved, a few more drops of acid 
 are added, avoiding unnecessary excess. A few cubic centimetres 
 of strong sulphuric acid are now added, and the solution made up 
 to 250 c.c. Fifty cubic centimetres of this solution are then with- 
 drawn, and gently warmed to a temperature about 60, and then 
 titrated with deci-normal permanganate.t 
 
 * A method for estimating copper is based upon a similar reaction. The 
 copper is first reduced either to cuprous oxide by means of grape sugar, or to 
 the metallic state by means of pure metallic zinc, and the reduced product is 
 then dissolved in a mixture of ferric chloride and hydrochloric acid 
 
 Cu 2 + Fe 2 Cl 6 + 2HC1 = 2 CuCl 2 + H 2 O + 2FeCl 2 
 Cu + Fe 2 Cl 6 = CuCl 2 + 2FeCl 2 
 
 t Other metals (e.g. lead) which admit of complete precipitation in the 
 form of oxalates or basic oxalates (bismuth) are capable of being estimated on 
 
Oxidation and Reduction Methods. 341 
 
 (b) Instead of determining the oxalic acid which is actually in 
 combination with the calcium, the calcium may be precipitated by 
 means of an excess of a standard oxalic acid solution, and after 
 removing the precipitate by filtration, the excess of oxalic acid is 
 found by means of permanganate. By deducting this excess 
 from the total amount used, the quantity of oxalic acid which was 
 precipitated by the calcium is found, and from this the percentage 
 of calcium is calculated. 
 
 About o'5 gram of calc-spar is weighed out, and dissolved 
 in dilute hydrochloric acid in a flask. Twenty cubic centimetres 
 of normal oxalic acid (63 grams per litre) are delivered into the 
 flask, and a slight excess of ammonia added. The mixture is 
 boiled for a few minutes, and then allowed to cool, after which it 
 is diluted with water up to 250 c.c. The liquid is then filtered 
 through an unwetted filter, and when sufficient has passed through, 
 50 c.c. of the perfectly clear filtrate * are withdrawn with a pipette 
 and transferred to a small flask, in which the solution is acidified 
 with sulphuric acid, heated to about 60, and titrated with deci- 
 normal permanganate. 
 
 Since one-fifth of the total solution was used for the titration, 
 the result multiplied by 5 will give the excess of oxalic acid ; and 
 this deducted from the original quantity taken will give the amount 
 which was precipitated as calcium oxalate. 
 
 Deci-normal Potassium Bichromate. 
 
 (4*913 grams per litre.\} 
 
 4-913 grams of pure potassium dichromate (previously dried by 
 being gently fused in a porcelain dish, and then powdered in a dry 
 mortar) are exactly weighed out and dissolved in water in a litre 
 flask, and the solution made up to 1000 c.c. with water at 15. 
 
 N 
 One litre of this solution contains one-tenth of an equivalent 
 
 of available oxygen in grams, i.e. o'8 gram ; hence I c.c. = o'oooS 
 gram available oxygen, and is equivalent to 0-0056 Fe, or to the 
 one-thousandth of the equivalent weight of any substance capable 
 of being fully oxidised by one equivalent of oxygen. 
 
 Unlike potassium permanganate, the dichromate solution is 
 
 this principle. Thus zinc and cadmium oxalates also are completely preci- 
 pitated by oxalic acid on the addition of a moderate quantity of alcohol. 
 Alkali salts must be absent, as these produce soluble double oxalates. 
 
 * Any calcium oxalate passing through the filter will be subsequently 
 decomposed by the acid, and will be therefore returned as excess oxalic acid, 
 and so vitiate the analysis. 
 
 t See p. 313. 
 
34 2 Volumetric Analysis. 
 
 stable, and as it has no action upon rubber, it may be used in a 
 burette which is closed by a rubber tube and pinchcock. 
 
 Titration of Decinormal Potassium Bichromate. 
 
 (a) By means of Ferrous Sulphate 
 
 K,Cr 2 7 + 6FeS0 4 + ;H 2 SO 4 = K 2 SO 4 + Cr 2 (SO 4 ) 3 + 3Fe 2 (SO 4 ) 3 
 
 + 7H 2 
 
 or, since the dichromate contains 3 atoms of available oxygen, the 
 reaction may be expressed 
 
 3 + 6FeO = 3Fe 2 3 
 
 For the purposes of the titration, the ferrous sulphate is prepared 
 by dissolving pure iron in dilute sulphuric acid, with exclusion of 
 air, precisely as described for permanganate, p. 333. An aliquot 
 part of the solution say 50 c.c. is withdrawn by means of a 
 pipette, and transferred to a small flask, and the dichromate solution 
 gradually added from a burette. 
 
 In this process, the end of the reaction is ascertained by means 
 of a freshly made and dilute solution of potassium ferricyanide, 
 used as an outside indicator. A number of drops of the ferricyanide 
 are placed about upon a white plate or tile, and from time to time, 
 during the addition of the dichromate, a drop of the mixture is 
 withdrawn upon a glass rod and brought into contact with one of 
 the drops of the indicator. At first a strong blue coloration is 
 produced, but as the amount of ferrous salt is gradually diminished 
 by the addition of the dichromate, the blue becomes less and less 
 intense, until at last a drop of the liquid so tested fails to give any 
 coloration. At this point the whole of the ferrous salt has been 
 oxidised, and the reaction is therefore complete.* 
 
 () By means of Ferrous Ammonium Sulphate. The process 
 is carried out as in the case of potassium permanganate, the end 
 reaction being determined by means of ferricyanide, as described 
 above. 
 
 Analyses by means of Potassium Dichromate. 
 
 In a number of instances, potassium dichromate may be sub- 
 stituted for permanganate in volumetric analysis. This is the case, 
 for example, with all estimations that are based upon the oxidation 
 of ferrous to ferric salts ; but not with those in which oxidation of 
 oxalic acid is the foundation of the process. 
 
 * It will be evident that it is absolutely essential to the success of this 
 operation that the ferricyanide should be perfectly free from ferrocyanide, 
 otherwise the oxidised iron will itself give rise to a blue coloration. 
 
Analyses by means of Potassium Dichromate. 343 
 
 The precautions as to the presence of hydrochloric acid, which 
 have to be taken when using potassium permanganate, do not apply 
 in the case of the dichromate ; and the solution to be titrated may 
 be less dilute when the latter reagent is employed. 
 
 The reducing agent to be used, in order to convert the iron into 
 the ferrous state previous to titration with potassium dichromate, is 
 either an alkaline sulphite, or stannous chloride. Reduction by 
 means of metallic zinc (which is the method most suitable when 
 permanganate is to be used, see p. 336) is not so well adapted for 
 solutions which are to be titrated with the dichromate solution, as 
 the presence of the dissolved zinc salt interferes with the reaction 
 with the ferricyanide indicator. 
 
 (i) Estimation of Iron in Iron Ores.* About i gram 
 of finely powdered and dry red haematite is weighed out into a 
 flask and boiled with a small quantity of strong hydrochloric acid, 
 diluted with about half its own volume of water, until the whole of 
 the iron has been extracted,! and the residue is free from dark- 
 coloured particles. 
 
 Reduction by means of Stannous Chloride. The solution is cooled 
 and diluted up to 100 c.c. with cold water. Twenty-five cubic 
 centimetres of this solution are then transferred to a small flask by 
 means of a pipette, and heated to the boiling-point over a small 
 flame. A clear freshly prepared solution of stannous chloride is 
 cautiously added, drop by drop, until the yellow colour of the ferric 
 chloride is just discharged, carefully avoiding unnecessary excess 
 
 Fe,Cl 6 + SnCl 2 = SnCl 4 + 2FeCU 
 
 The slight excess of stannous chloride is removed from the 
 solution (as stannous chloride is oxidised by potassium dichromate, 
 and its presence, therefore, would vitiate the subsequent titration), 
 by adding a few drops of a solution of mercuric chloride, whereby 
 
 * In the example here given, the total iron is alone estimated. When the 
 amoiint of ferrous and ferric oxide present is to be separately determined, the 
 procedure is the same as that described on p. 335 ; the reduction of the iron 
 in the second portion (in which the total iron is estimated) is accomplished 
 either by stannous chloride or alkaline sulphite, and the solutions titrated with 
 dichromate. 
 
 f In cases where the iron is difficult to extract with hydrochloric acid alone, 
 a few small crystals of potassium chlorate may be added, after which the 
 solution must be evaporated to dryness, and the residue dissolved in hydro- 
 chloric acid. Iron ores (such as the titaniferous ores) from which the iron 
 cannot be wholly extracted by acids, are mixed in a state of fine powder with 
 from 6 to 10 times their weight of hydrogen potassium sulphate (bisulphate q/ 
 potash], and gently fused in a platinum dish for from 20 to 30 minutes. The 
 " melt " is extracted with dilute hydrochloric acid. 
 
344 Volumetric Analysis. 
 
 mercurous chloride is precipitated (which is without action upon 
 the dichromate), and stannic chloride produced * 
 SnCl 2 + 2HgCl 2 = Hg 2 Cl 2 + SnCl 4 
 
 The solution is now titrated with the deci-normal dichromate 
 without unnecessary exposure to the air. The volume of the di- 
 chromate used, multiplied by 4 (as 25 c.c. out of the 100 c.c. were 
 used for the titration), gives the volume required to oxidise the iron 
 in the original weight of ore taken. 
 
 Reduction by means of Alkaline Sulphite. The solution of the 
 ore obtained by boiling with hydrochloric acid is slightly diluted, 
 and, if necessary, filtered through a small filter into a flask, the 
 filter and residue being thoroughly washed. Ammonia is added to 
 the liquid until slightly alkaline, as shown by the formation of a 
 faint permanent precipitate of ferric hydroxide. One or two drops 
 of dilute hydrochloric acid are then added, in order tojtisf re- dissolve 
 this precipitate. 
 
 Ten cubic centimetres of a strong solution of hydrogen ammonium 
 sulphite f (ammonium bisulphite), H(NH 4 )SO 3 , are then added, 
 and thoroughly mixed by shaking, and the solution gradually heated 
 to boiling. Ten cubic centimetres of strong sulphuric acid arc 
 mixed with an equal volume of water, and this mixture is added to 
 the heated liquid, and the boiling continued for 15 or 20 minutes 
 to expel the whole of the sulphur dioxide. The solution is then 
 cooled quickly and diluted to 100 c.c. with cold water. Twenty-five 
 cubic centimetres of this are then titrated with the deci-normal 
 dichromate. 
 
 (2) Estimation of Chromium in Chrome Iron Ore. 
 
 By the exact converse of the foregoing process, chromium (in 
 the condition of chromate) is most readily estimated. In the case 
 of chrome iron ores (as also with ferro-chromium), the chromium 
 is converted into sodium chromate by fusion with sodium peroxide. 
 The chromate is then reduced by means of a measured or weighed 
 excess of a ferrous salt, and the excess of the latter then estimated 
 by titration with deci-normal dichromate solution. 
 
 * If no precipitate is produced on the addition of the mercuric chloride, it 
 shows that an insufficient quantity of stannous chloride was added. Too 
 copious a precipitate, due to too great an excess of the tin solution, will 
 interfere with the titration. 
 
 t That is to say, about i c.c. for each o'i gram of ore which was dissolved. 
 The bisulphite may be prepared by passing a gentle stream of sulphur dioxide 
 through a small quantity of strong ammonia solution, diluted with half its 
 volume of water. The ammonia may be contained in a boiling-tube, which is 
 kept cool by being immersed in a beaker of cold water. The gas is bubbled 
 through it until the liquid smells strongly of sulphur dioxide, and until the 
 crystals of the normal sulphite which first form are wholly re-dissolved. 
 
Analyses by means of Potassium Bichromate. 345 
 
 About 0-5 gram of \&t finely powdered ore * is weighed out into 
 a nickle crucible f and intimately mixed with 3 to 4 grams of sodium 
 peroxide. The crucible is then heated by means of a Bunsen flame, 
 and the mixture kept in a state of fusion for 5 to 10 minutes It is 
 then allowed to partially cool, until a crust is formed upon the 
 surface of the fused mass, when I gram of sodium peroxide is added, 
 and the materials again melted and kept in a state of fusion for 
 5 minutes longer. 
 
 After cooling, the crucible is transferred to a porcelain beaker 
 (or, in the absence of this, a porcelain dish), and the melt exhausted 
 with hot water, the beaker or dish being covered with a clock-glass. 
 The solution is then boiled for 10 minutes, in order to decompose 
 completely the whole of the remaining sodium peroxide.^ [Should 
 the solution first obtained by dissolving the melt be purple in colour, 
 due to sodium ferrate, a little more sodium peroxide should be added 
 before boiling the liquid.] 
 
 The iron present in the ore (the manganese, in the case of ferro- 
 chromium), and the nickel derived from the crucible, are left in 
 the form of oxides, which are separated by filtering and carefully 
 washing the residue upon the filter. The yellow solution containing 
 the chromium, as sodium chromate, is acidulated with sulphuric 
 acid, and an excess of ferrous ammonium sulphate added. This is 
 most conveniently done by adding a weighed quantity of the solid 
 salt. A weighing-bottle, filled with the salt, is first weighed, and 
 then small quantities are carefully taken out upon a small clean 
 spatula (holding the bottle over the beaker containing the chromate 
 solution, lest any particles should fall), and added to the liquid 
 until the chromium is entirely reduced (the solution turning green) 
 
 * Too much stress cannot be laid upon the importance of reducing the ore to 
 the finest possible powder, as the success of the " fusion " hangs entirely upon 
 this point. The ore should be crushed in a steel mortar until the powder is 
 fine enough to pass through a linen bag ; it is then ground in an agate mortar 
 with the addition of a little water, until it is reduced to an impalpable powder. 
 Rideal and Rosenblum (the authors of the method), Journal of the Sociefv 
 of Chew. Industry, December 31, 1895. 
 
 f Platinum and silver vessels are attacked by the fused peroxide. Nickel is 
 also attacked to som.e extent, but the dissolved nickel is subsequently removed. 
 
 J It is essential to the success of the process that the whole of the sodium 
 peroxide be decomposed at this stage, otherwise, when the liquid is subsequently 
 acidified, hydrogen peroxide will be generated, and this, reacting upon the 
 chromate in solution, causes the reduction of a portion of the chromium. The 
 reducing action of hydrogen peroxide upon chromic acid is similar in its 
 character to that which it exerts upon other highly oxidised compounds (e.g. 
 peroxides, permanganates, etc.), whereby oxygen is withdrawn from both com- 
 pounds, and is liberated as free gas. In the case of chromium, however, an 
 intermediate stage occurs, in which the blue compound, supposed to be per- 
 chromic acid, is first formed. 
 
346 Volumetric Analysis. 
 
 and the ferrous salt is present in slight excess. That excess of iron 
 has been introduced is ascertained by bringing a drop of the liquid 
 (by means of a glass rod) into contact with a drop of a fresh and 
 dilute solution of potassium ferricyanide upon a white plate, when 
 a blue colour is obtained. The stopper is replaced in the weighing- 
 bottle, and the latter set aside to be re-weighed after the titration 
 has been completed. 
 
 The excess of ferrous salt present is now determined by means 
 of deci-normal potassium dichromate, which is added until a drop 
 of the liquid no longer gives a blue colour with potassium 
 ferricyanide. 
 
 The weighing-bottle is now re-weighed, the difference being 
 the total weight of the ferrous salt used. From this is deducted 
 the weight of ferrous salt equivalent to the volume of deci-normal 
 dichromate employed, i.e. the excess of ferrous salt, which gives 
 the weight of ferrous salt oxidised by the chromate derived from 
 the ore. From the equation 2CrO 3 + 6FeO = Cr 2 O 3 + 3Fe 2 O 3 , 
 it will be seen that I atom of Cr (52-4) is capable of oxidising 3 
 atoms of Fe (56 x 3 = 168). 
 
 The formula-weight of ferrous ammonium sulphate being exactly 
 seven times the weight of the iron it contains (see p. 334), the 
 weight of this salt which is equivalent to 52-4 parts of chromium is 
 1176 parts. From this proportion the percentage of chromium in 
 the ore is readily calculated. 
 
 Iodine. 
 
 Free iodine, in the presence of water and certain oxidisible 
 substances, is able to take the hydrogen from water and liberate 
 the oxygen. The oxygen thus eliminated oxidises the oxidisible 
 compound present, and therefore the iodine acts as an indirect 
 oxidising agent, 127 parts of iodine being equivalent to 8 parts of 
 oxygen. 
 
 Thus, when a solution of iodine is added to dilute sulphurous 
 acid, the latter compound is oxidised to sulphuric acid 
 H 2 S0 3 + I 2 + H 2 = H 2 S0 4 + 2HI 
 
 Similarly, the lower oxides of arsenic and of antimony (or the 
 salts of these oxides) are oxidised by iodine in the presence of 
 water ; thus with sodium arsenite 
 
 Na 3 As0 3 + I 2 + H 2 = Na 3 AsO 4 + 2HI 
 
 Therefore a standard solution of iodine may be used to estimate 
 
Oxidation Reactions by means of Iodine. 347 
 
 sulphur dioxide in solution, and also for the determination of 
 arsenic and antimony. 
 
 When iodine is added to a dilute solution of sulphuretted 
 hydrogen, a definite reaction takes place, which is usually expressed 
 by the equation * 
 
 H 2 S + I 2 = 2HI + S 
 
 Hence a solution of iodine may be employed for the direct 
 determination of sulphuretted hydrogen in solution. Of all the 
 reactions for which a solution of iodine is employed, the most 
 useful and important in volumetric analysis is that with thio- 
 sulphuric acid, or sodium thiosulphate, which becomes converted 
 into the sodium salt of tetrathionic acid, while hydriodic acid (or 
 the sodium salt) i? produced ; thus 
 
 2Na 2 S,O 3 + I 2 = 2NaI + Na 2 S 4 O G t 
 
 In volumetric processes in which iodine is employed, the end 
 reaction is indicated by means of starch. When brought in contact 
 with free iodine, starch produces a deep blue coloration, and the 
 formation of this blue colour constitutes a delicate indication of 
 the smallest excess of iodine over and above that which is required 
 for the oxidising reaction in hand. 
 
 The blue-coloured substance is usually regarded as a compound 
 of starch and iodine, although unstable in character and of un- 
 known composition ; the idea of combination between the substances 
 is also conveyed by the name iodide of starch, by which it is called. 
 Towards sodium thiosulphate, however, in common with many 
 other reagents, it behaves exactly as free iodine : if a solution of 
 
 * In this equation the oxidising character of the process is obscured by 
 the omission from it of the water which is necessary for its progress. At the 
 ordinary temperature, hydrogen sulphide and iodine do not interact in the 
 absence of water. One function of the water is, by dissolving the hydriodic 
 acid, to furnish the necessary heat required for the combination of the hydrogen 
 and iodine ; while another, no doubt, is to supply the intermediate step in the 
 process whereby oxygen is available for the oxidation of the sulphuretted 
 hydrogen, thus 
 
 iI 2 + H 2 ;JO + H 2 JS = 2 HI + H 2 + S 
 
 f This reaction, just as in the case of sulphuretted hydrogen, is in reality 
 one of indirect oxidation through the intervention of water. The mechanism 
 of the reaction will be understood by the following dissected equations 
 
 |I 2 + H 2 JO + aNalo* = 2HI 
 
 2HI + OXa 2 = aNal + H 2 O 
 
348 Volumetric Analysis. 
 
 the thiosulphate be added to it, the iodine at once attacks the 
 thiosulphate according to the above reaction, just as though there 
 were no starch present ; and the instant the whole of the iodine 
 is thus used up, the blue colour of the liquid disappears. 
 
 It is on this account that a solution of sodium thiosulphate used 
 in conjunction with the iodine solution constitutes so important 
 a volumetric reagent. Thus, all substances which can be made 
 to yield chlorine either directly (as from hypochlorites or chlorates) 
 or indirectly (as by the action of hydrochloric acid upon peroxides, 
 chromates, etc.) are capable of being estimated by these reagents. 
 The liberated chlorine is passed into potassium iodide, and there 
 displaces its equivalent of iodine, and the exact amount of iodine 
 can then be determined by means of sodium thiosulphate. 
 
 Deci-normal Iodine Solution. 
 
 (i 27 grams of iodine per litre.} 
 
 127 grams of the purest re-sublimed iodine, in a state of 
 powder, are exactly weighed out into a litre flask. About 20 grams 
 of potassium iodide (free from iodate) dissolved in about 200 c.c. 
 of water are added, and the mixture gently shaken until the iodine 
 is wholly dissolved. The solution is then diluted up to the litre 
 with water at 15, and preserved in a well- fitting stoppered bottle 
 in the dark. 
 
 If strictly deci-normal, i c.c. will contain 0*0127 grams of iodine, 
 equivalent to 0*0008 gram oxygen or to O'oo355 gram of chlorine. 
 
 Deci-normal Sodium Thiosulphate. 
 
 (24-8 grams 0fNa 2 S 2 O 3 ,$H 2 O per litre.} 
 
 24*8 grams of the purified salt, which has been carefully 
 powdered,* are weighed out into a litre flask, and dissolved in 
 water. The solution is then diluted up to 1000 c.c. with water at 
 1 5. If strictly deci-normal, i c.c. of the solution will be exactly 
 oxidised by i c.c. of the iodine solution. 
 
 Titration of Deci-normal Iodine, by means of sodium 
 thiosulphate. 
 
 * When powdering a salt which contains a large number of molecules of 
 water of crystallisation, such as sodium thiosulphate, there is always a risk of 
 a slight loss of water. This is due to the fact that the pressure produced by 
 the blow of the pestle causes the momentary partial liquefaction of portions o'f 
 the salt. This explains the fact that, although perfectly free from extraneous 
 moisture, such a salt appears to become damp, and " clogs" upon the pestle 
 during the process of powdering. While in this condition the salt may lose 
 traces of its water of crystallisation by evaporation. 
 
Iodine and Sodium Thiosulphate. 349 
 
 Twenty-five cubic centimetres of the deci-normal thiosulphate 
 solution are transferred by means of a pipette to a beaker, diluted 
 with a little water, and a few drops of a clear solution of starch 
 added.* The iodine solution is then run in from a stoppered 
 burette, until a single drop gives a permanent blue coloration. 
 
 The titration may equally as well be made in the opposite 
 order. Twenty-five cubic centimetres of the iodine solution are 
 transferred to a beaker, diluted with about an equal volume of 
 water, and the starch solution added. The thiosulphate is then 
 delivered from a burette until the blue colour is just discharged. 
 As a rule, when the titration is made in this way, it is an advantage 
 not to introduce the starch until towards the end of the reaction, 
 especially if the amount of iodine is considerable. The thio- 
 sulphate is added until the brown colour of the iodine solution 
 begins to perceptibly pale, and the liquid assumes a straw colour. 
 One or two drops of starch are then introduced, causing the blue 
 colour, and the thiosulphate delivered drop by drop until the 
 colour is discharged. 
 
 The result of this titration will show the exact relation in which 
 these two solutions stand to each other. If they bear the correct 
 relative values, then 25 c.c. of the one will require 25 c.c. of the 
 other in order to exactly complete the reaction between them. If, 
 on the other hand, one is found to be proportionately a little stronger 
 than the other, the value of the one in terms of the other must be 
 notified, preferably on the labels of the bottles. Thus, suppose the 
 25 c.c. of iodine solution required 24*9 c.c. of the thiosulphate solu- 
 tion, then i c.c. of the latter equals -f - or 1-004 c.c. of the iodine ; 
 
 and i c.c. of the iodine is equivalent to 0-996 c.c. of thiosulphate. 
 
 Although these solutions might be found to bear the exact 
 relative value towards each other, so that equal volumes are 
 chemically equivalent, it is obviously possible that they might still 
 not be strictly deci-normal solutions. It is possible, although per- 
 haps not highly probable, that an error to exactly the same extent 
 might have arisen in the preparation of each ; so that each of them 
 
 * The starch solution is prepared by mixing about i or 2 grams of white 
 potato-starch in powder with 5 or 6 c.c. of cold water in the bottom of a 
 moderately large beaker. A considerable quantity of boiling water is then 
 poured upon it. As the hot water is being added, the opaque white appearance 
 which the mixture presents at first, changes almost suddenly to that of a semi- 
 translucent gelatinous substance. At this point the addition of the boiling 
 water is stopped, and the beaker nearly filled up with cold water. It is 
 allowed to settle, and the clean liquid poured off for use. This starch solution 
 should not be kept for more than one day, or at the outs'de two days. 
 
350 Volumetric Analysis. 
 
 should be either too strong or too dilute to the same degree. 
 It is necessary, therefore, to make an independent titration of one 
 of them to serve as a check. Either of the following methods can 
 be used. 
 
 (a) Titration of Deri-normal Iodine by means of Arsenious 
 Oxide. 
 
 As 4 O a + 4l 2 + 4H 2 O = SHI + 2As 2 O 5 
 
 That is to say, 127 parts of iodine will oxidise 49*5 parts of 
 arsenious oxide. 
 
 4-95 grams of finely powdered arsenious oxide, which has been 
 purified by resublimation, are exactly weighed out into a litre flask, 
 and about 500 c.c. of water added. Thirty grams of pure sodium 
 bicarbonate are then introduced, and the mixture gently warmed 
 upon a steam-bath to a temperature about 60, and continually 
 shaken until the arsenious oxide is wholly dissolved.* The solu- 
 tion is then cooled and made up to a litre with water at 15. This 
 solution, containing one-tenth of the equivalent of arsenious oxide, 
 is of deci-normal strength, I c.c. being equivalent to o - oi27 gram 
 of iodine, and therefore equal to i c.c. of deci-normal iodine solu- 
 tion. 
 
 Twenty-five cubic centimetres of this solution are transferred to 
 a small beaker, a few drops of clear starch solution added, and the 
 deci-normal iodine solution delivered into it from a burette until 
 the blue coloration is permanent. From the volume of the iodine 
 solution required to completely oxidise the arsenious oxide in 25 c.c. 
 of the solution, the exact strength of the iodine solution is obtained. 
 If it is found to be strictly deci-normal, then obviously the thio- 
 sulphate solution which has the same relative strength is also 
 exactly deci-normal. 
 
 (b} Titration of Deci-normal Thiosulphate Solution by means of 
 Potassium DichromateVJten. a solution of potassium dichromate, 
 
 * An excess of the alkali is necessary in order to neutralise the hydriodic 
 acid which is afterwards produced during the titration, the necessary alkali 
 for this purpose being supplied by sodium bicarbonate, instead of either the 
 normal carbonate or caustic alkali. As already stated, the iodine in this blue 
 compound behaves towards many reagents much a.sfree iodine would, and free 
 iodine is dissolved by caustic alkali, and also (but less easily) by the normal 
 carbonates of the alkalies ; thus 
 
 6NaHO + 3X2 = 5NaI + NaIO 3 
 6Na 2 C0 3 + 3l 2 + 6H 2 = 6NaHCO 3 + 5NaI + NaIO 3 
 
 Hence if either of these alkaline reagents be added to the blue starch iodide, 
 the blue colour will be destroyed (see note on p. 355). 
 
Iodine and Sodium Thiosulphate. 351 
 
 acidified with hydrochloric acid, is mixed with potassium iodide, 
 and the mixture gently warmed, iodine is liberated ; thus 
 
 K 2 Cr 2 O 7 + HHCl + 6KI = 3l 2 + Cr 2 Cl 6 + 8KC1 + ;H 2 O 
 
 The amount of iodine so liberated is the equivalent of the available 
 oxygen in the dichromate, three atoms of available oxygen being 
 equivalent to six atoms of iodine, i.e. 8 parts of oxygen liberate 127 
 parts of iodine. 
 
 If a standard solution of the dichromate be used in the reaction, 
 the available oxygen is a known quantity, and therefore the amount 
 of iodine it can liberate is also a known quantity ; so that if this 
 iodine be titrated with sodium thiosulphate, the strength of the 
 latter can be determined- 
 
 Thus i c.c. deci-normal dichromate solution contains 0*0008 gram 
 
 available oxygen ; 
 therefore i c.c. deci-normal dichromate solution liberates 0-0127 gram 
 
 iodine. 
 
 i c.c. deci-normal thiosulphate oxidises 0*0127 gram iodine ; 
 therefore i c.c. deci-normal dichromate is equivalent to i c.c. deci- 
 
 normal thiosulphate. 
 
 Twenty-five cubic centimetres of deci-normal dichromate solution 
 are placed in a stoppered bottle, and about 2 grams of potassium 
 iodide together with 5 or 6 c.c. of strong hydrochloric acid are added. 
 The mixture is thoroughly shaken, and gently heated by standing 
 the closed bottle in warm water for some time. The bottle with 
 its contents is then cooled, the stopper withdrawn and rinsed into 
 the bottle with water, and the liberated iodine titrated with the 
 thiosulphate solution ; the starch solution being added when the 
 reddish colour of the iodine solution begins to fade. The dis- 
 appearance of the blue colour of the iodide of starch leaves the 
 solution not colourless, but pale green (due to the chromic chloride), 
 hence a little care is necessary in judging the end reaction.* If 
 the deci-normal thiosulphate is of exact strength, 25 c.c. will be 
 required for the titration. 
 
 * Deci-normal permanganate, acidified with dilute sulphuric acid, may be 
 substituted for the dichromate and hydrochloric acid, the operation being con- 
 ducted in the same manner. In this case, when the blue colour of the starch 
 iodide is discharged by the thiosulphate, the liquid is left colourless. The 5 
 atoms of available oxygen in permanganate liberate 10 atoms of iodine ; thus 
 
 + 8H 2 SO 4 + loKI = sI 2 + 6K = SO 4 + aMnSO 4 + 8H 2 O 
 
35 2 Volumetric Analysis. 
 
 Estimations by means of Deci-normal Iodine. 
 
 (1) Antimony in Antimonious Oxide. 
 
 Sb 2 3 4- 2H 2 + 2l 2 - 2HI + Sb 2 5 
 
 About i gram of antimonious oxide is weighed out into a |-litre 
 flask. Two grams of tartaric acid dissolved in about 25 c.c. of 
 water are then added, and the mixture gently shaken until the 
 antimonious oxide is dissolved. Sodium carbonate (the normal 
 salt) is added until the solution is just neutral, and the liquid is 
 diluted with water to 250 c.c. Fifty cubic centimetres of this solu- 
 tion are withdrawn with a pipette and transferred to a beaker, and 
 from 20 to 30 c.c. of a cold saturated solution of pure sodium 
 bicarbonate* added, together with a few drops of starch. The 
 deci-normal iodine solution is then run in from a burette until a 
 permanent blue colour is produced. A duplicate titration should 
 be made in a second portion of the solution. 
 
 EXAMPLE. Fifty cubic centimetres of the solution required 
 27'5 c.c. of the iodine solution ; 
 
 therefore 250 c.c. (containing i gram of the original oxide) would 
 
 require 27-5 x 5 = 137-5 
 i c.c. deci-normal iodine = 0-0127 gram I = 0-0060 gram Sb 
 
 = 0*0072 gram Sb 2 O 3 
 therefore 137-5 x 0-0060 = 0-8250 gram Sb in i gram of the oxide 
 
 25 X I0 = 82-50 percentage Sb, or 99 per- 
 centage of Sb 2 O 3 
 
 (2) Arsenic in Arsenious Compounds. The process is 
 carried out as already described on p. 350. 
 
 (3) Tin, in the Condition of Stannous Chloride, may 
 
 also be estimated by means of iodine, but in this case the principle 
 involved is different to that which obtains with either arsenic or 
 antimony. Stannous chloride unites directly with iodine, forming 
 the compound SnCl 2 I 2 
 
 SnCl 2 + I 2 = SnCl 2 I 2 
 
 hence two atoms of iodine become the measure of one atom of tin, 
 or 127 parts of iodine are equivalent to 59 parts of tin. Tin cannot, 
 however, be determined by this reaction in the presence of lead ; 
 because, although lead chloride alone is not acted upon by free 
 iodine, it is acted upon by the stannic chlor-iodide with the pre- 
 cipitation of yellow lead iodide, which completely obscures the end 
 reaction. In tin alloys, therefore, which contain lead, the latter 
 metal must be first removed. 
 
 * See note on p. 350. 
 
Valuation of Bleaching Powder. 353 
 
 Estimations by means of Iodine and Thiosulphate. 
 
 (1) Sulphur Dioxide in Sodium Sulphite, Na 2 SO 3 ,7H 2 O 
 
 Na 2 SO 3 + I 2 + H 2 O = 2HI + Na 2 SO 4 
 
 That is to say, 126 parts of the crystallised salt are oxidised by 
 
 N 
 127 parts of iodine; therefore I c.c. iodine (containing o - oi27 
 
 gram I) is equivalent to 0-0126 gram Na 2 SO 3 ,7H 2 O, or 0^0032 gram 
 S0 2 . 
 
 About i gram of the powdered salt is weighed out into a flask, 
 and loo c.c. of the iodine solution added. This volume ensures an 
 excess of iodine, so long as the weight of salt taken has not exceeded 
 1*25 grams,* and the excess will be manifest by the colour of the 
 liquid. 
 
 The excess of iodine is now titrated by means of the deci-normal 
 thiosulphate solution, added from a burette. As soon as the brown 
 colour of the iodine begins to pale to a yellow, starch is added, and 
 then the addition of the thiosulphate is continued until the blue 
 colour is discharged. The excess of iodine thus found, deducted 
 from the volume originally taken (100 c.c. in this case), gives the 
 amount of iodine which has been used in oxidising the sodium sul- 
 phite in accordance with the above reaction, and from this the 
 percentage of sulphur dioxide can be calculated.! 
 
 (2) Available Chlorine in Bleaching Powder. The 
 active ingredient in bleaching powder is the compound expressed 
 by the formula Ca(OCl)Cl. Bleaching powder, however, besides 
 containing some unacted-upon calcium hydroxide, usually con- 
 tains also more or less chlorinated compounds (such as cal- 
 cium chlorate, for example), which are of no use for the purposes 
 to which the bleaching powder is put. The chlorine which is 
 present as calcium chloro-hypochlorite, Ca(OCl)Cl, is known as the 
 available chlorine. This chlorine is evolved when the bleaching 
 powder is acted upon by dilute acids ; thus 
 
 Ca(OCl)Cl + H 2 S0 4 = H 2 + CaSO 4 + C1 2 
 Ca(OCl)Cl + 2HC1 = H 2 O + CaCl 2 + C1 2 
 
 Hence, if bleaching powder consisted of the pure compound, 
 
 * In cases where the composition of the sulphite is unknown, a preliminary 
 titration must be made, using a considerable quantity of iodine, so as to gain 
 an approximate idea of the amount required. 
 
 f This method may also be applied for the estimation of sulphuretted 
 hydrogen in solution (as in hepatic waters), and also of aqueous solutions of 
 chlorine and bromine. 
 
 2 A 
 
354 Volumetric Analysis. 
 
 having the formula Ca(OCl)Cl, it would contain 55 per cent, of 
 available chlorine ; but in reality this is never the case> and the 
 best samples of commercial bleaching powder seldom exceed about 
 36 per cent., while in others it falls considerably below this. 
 
 About 10 grams of bleaching powder are weighed out into a 
 porcelain mortar, and rubbed down with a small quantity of water 
 until the mixture has the consistency of thin cream. The heavier 
 particles are allowed a moment or two to settle, and the 
 milky liquid is poured off into a litre flask. The residue is then 
 ground down with a little more water, and the process repeated 
 until the last traces have been transferred to the flask. The mix- 
 ture is then diluted with water up to 1000 c.c., and thoroughly 
 shaken. 
 
 Twenty-five cubic centimetres of the milky fluid are transferred 
 by means of a pipette to a small beaker, the contents of the flask 
 being briskly shaken up immediately before the withdrawal of the 
 sample. Excess of potassium iodide is added by introducing 10 
 c.c. of a solution containing 5 grams KI in 100 c.c. water.* The 
 mixture is then acidified with acetic acid.f This liberates the 
 chlorine from the bleaching powder, and that, in the presence of 
 the excess of potassium iodid^ sets free an equivalent quantity of 
 iodine 
 
 C1 2 + 2KI = 2KC1 + I 2 
 
 The amount of the liberated iodine is then estimated by the addition 
 of deci-normal sodium thiosulphate, after the addition of starch. 
 
 N 
 i c.c. thiosulphate = 0*0127 gram I = 0*00355 gram Cl 
 
 Therefore the volume of thiosulphate required to discharge the 
 blue colour of the starch iodide, is a measure of the chlorine which 
 was liberated from the bleaching powder contained in 25 c.c. of the 
 original mixture. From this the percentage of available chlorine is 
 calculated. 
 
 Valuation of Bleaching Powder. A Iter native Method by means 
 of Arsenious Oxide.\ Instead of estimating the available chlorine 
 
 * It is convenient to remember that a suitable excess of potassium iodide 
 will be ensured if twice the weight of potassium iodide is added as there is of 
 bleaching powder present. In the above example, 25 c.c. of the bleach solution 
 contains 0^25 gram of the compound (10 grams to the litre), and the weight of 
 potassium iodide in 10 c.c. of the KI solution is 0-5 gram. 
 
 f Acetic acid does not liberate chlorine from any chlorate which the bleach- 
 ing powder may contain. 
 
 \ Although this method does not involve the use of sodium thiosulphate, it 
 is most conveniently introduced at this point. 
 
Valuation of Bleaching Powder. 355 
 
 in bleaching powder by measuring the amount of iodine which that 
 chlorine is capable of liberating, it may be determined by measuring 
 the amount of arsenious oxide which the oxygen present in the 
 active constituent of the powder is capable of oxidising. This 
 active ingredient, as stated above, is the calcium chloro-hypo- 
 chlorite, Ca(OCl)Cl, which contains I atom of oxygen to every 2 
 atoms of chlorine. Hence the weight of bleaching powder capable 
 of furnishing 35 '5 parts of available chlorine, will also supply 8 parts 
 of oxygen ; this oxygen, therefore, may be taken as the measure of 
 the chlorine 
 
 2Ca(OCl)Cl* + As 2 3 = AsA + 2CaCl 2 
 
 Twenty-five cubic centimetres of the freshly made and well- 
 shaken up bleaching powder solution are transferred to a small 
 beaker by means of a pipette, and deci-normal arsenious oxide 
 solution f is gradually added from a burette. The end of the reaction 
 
 * When bleaching powder is mixed with water, it is -converted into calcium 
 chloride and hypochlorite ; thus 
 
 2 Ca(OCl)Cl = Ca(OCl) 2 + CaCl 2 
 
 The composition of bleaching powder, therefore, may be, and often is, expressed 
 by the formula Ca(OCl) 2 ,CaCl 2) and it will be obvious that the proportion 
 which the oxygen bears to the chlorine is the same in both cases, i.e. 8 parts 
 by weight of O to 35*5 parts of Cl. The compound calcium hypochlorite 
 Ca(OCl) 2 itself contains all the oxygen but only half the chlorine originally 
 present in the bleaching powder. 
 
 Whether or not the decomposition of the bleaching powder into hypochlorite 
 and chloride by the agency of water is complete, is not known with certainty, 
 probably it is not. Neither is the exact modus operandi of the water, in bring- 
 ing about this decomposition, clearly understood. Probably a cycle of changes 
 takes place, which may be explained by the following equations : 
 
 Ca(OCl)Cl HC1 OC1 
 
 -f- H 2 O = + Ca< + CaO 
 Ca(OCl)Cl HC1 X OC1 
 
 HOC1 
 
 = CaCU + + CaO 
 
 HOC1 
 
 ,oci 
 
 = CaCl 2 + Ca< +H 2 
 
 X OC1 
 
 f The deci-normal solution used for estimating the value of bleaching 
 powder must either \>e freshly prepared, as described on p. 3:50, or else it must 
 be made with normal sodium carbonate instead of the bicarbonate, using 20 
 grams in place of the 30 grams of the latter salt. It has been found in practice 
 that the solution made from the bicarbonate undergoes some change on being 
 kept (the precise nature of which is not known), which introduces errors into 
 the analysis greater than are caused by the action of the normal carbonate upon 
 the starch iodide. 
 
356 Volumetric Analysis. 
 
 is determined by means of starch and potassium iodide, used as an 
 outside indicator. After each addition of the arsenious solution, 
 the mixture is stirred, and a drop of it removed upon a glass rod, 
 and brought into contact with a piece of filter-paper, which has been 
 previously impregnated with a mixture of starch and potassium 
 iodide.* So long as undecomposed bleaching powder is present, 
 the liquid will cause a blue stain upon the paper. The arsenious 
 solution is cautiously added, until a drop of the liquid brought into 
 contact with potassium-iodide-and-starch paper gives no blue 
 colour. 
 
 If the exact point is overstepped, a little additional excess of the 
 arsenious solution may be added, together with a few drops of starch, 
 and the excess titrated by means of deci-normal iodine solution, 
 until the blue colour is produced. On deducting this excess from 
 the total volume of the arsenious solution used, the exact volume 
 which was oxidised by the bleaching powder is obtained, and from 
 this the percentage of available chlorine may be calculated. 
 
 (3) Estimation of Manganese Dioxide. When man- 
 ganese dioxide is heated with hydrochloric acid, an amount of 
 chlorine is evolved which is the chemical equivalent of the avail- 
 able oxygen in the manganese dioxide (see p. 337). The chlorine 
 so evolved may be absorbed by means of a solution of potas- 
 sium iodide, whereby an amount of iodine equivalent to the 
 chlorine (and therefore equivalent to the oxygen) is liberated. The 
 iodine thus set free can then be determined by means of sodium 
 thiosulphate. 
 
 About o'5 gram of the manganese dioxide, in fine powder, is 
 weighed out into a small flask (Fig. 73), about 80 to 100 c.c. 
 capacity. The flask is provided with a delivery tube, which is fitted 
 either by means of a good cork which has been previously soaked 
 in melted paraffin wax, or a rubber stopper which has been boiled 
 in caustic soda in order to remove sulphur ; or the tube is made 
 by drawing out a wider tube, the wide portion of which is slightly 
 conical, and has pushed over it a short piece of black rubber tube. 
 This constitutes a perfect fitting stopper and exit tube in one. The 
 delivery tube, bent as shown, passes through a cork into a U-tube 
 placed in a beaker of cold water. The U-tube contains moderately 
 strong potassium iodide solution. (2*5 grams of KI for 0-5 gram 
 
 * This paper is readily prepared by adding a small quantity of potassium 
 iodide to a little clear starch solution, and dipping strips of unsized paper 
 into the mixture. If these strips are then hung up to dry, they may be pre- 
 served indefinitely in a stoppered bottle. 
 
Estimations by means of Iodine and Thiosnlphate. 357 
 
 of MnO 2 will ensure a suitable excess.*) A few fragments of 
 magnesite are introduced into the flask along with the manganese 
 ore, and about 25 to 30 c.c. of strong hydrochloric acid added. 
 The delivery tube is then attached, and a gentle heat applied to 
 the flask. The object of the magnesium carbonate is to furnish a 
 slow stream of carbon dioxide during the entire operation (magnesite 
 being only slowly dissolved by the acid), which not only sweeps 
 out the chlorine, but prevents the liquid in the absorption tube 
 from being sucked back into the flask during the last stages of the 
 
 FIG. 73. 
 
 operation when the acid is being boiled. As the chlorine enters 
 the potassium iodide solution it is wholly absorbed, so long as the 
 evolution of the gas is not allowed to proceed too fast. When the 
 manganese dioxide has all passed into solution, the temperature is 
 raised to the boiling-point, and the liquid allowed to boil for a few 
 minutes. When the distillation is completed, the U-tube is dis- 
 connected, and the contents washed out and diluted to 100 c.c. in 
 
 * A second U-tube is sometimes connected to the first to serve as a guard, 
 lest any chlorine should escape unabsorbed from the first. This, however, will 
 not happen unless the distillation is made far too rapidly. 
 
35$ Volumetric Analysis. 
 
 a graduated flask. Twenty-five cubic centimetres of this solution 
 are withdrawn by means of a pipette, and the liberated iodine is 
 then estimated by titration with deci-normal sodium thiosulphate. 
 As soon as the brown colour of the liquid fades to yellow, a little 
 starch is added, and the addition of the thiosulphate continued 
 until the blue colour is just discharged. 
 
 EXAMPLE. Weight of manganese dioxide taken = 0*495 gram. 
 Iodine solution diluted to 100 c.c. ; 25 c.c. of the liquid 
 
 !N 
 required of thiosulphate (i c.c. = 0-0127 gram I) ... 21 c.c. 
 
 0*0127 x 21 = 0-2667 gram iodine in 25 c.c. 
 and 0*2667 x 5 = 1*3335 gram iodine liberated by 0*495 gram 
 
 of the ore 
 From the equation MnO 2 + 4HC1 = MnCl 2 + 2H 2 O + C1 2 (= I 2 ), 
 
 127 parts of iodine are liberated by 43*5 parts of manganese 
 dioxide 
 
 Therefore 127 : 1*3335 : : 43*5 : 0*45674 = grams of MnO 2 
 
 and 0*45674 x loo . ., _. . 
 
 ** . = 92*27 = percentage of MnO 2 in 
 
 the ore 
 
 A number of other oxygenated compounds are susceptible of 
 being estimated by the same method, namely, by distillation with 
 hydrochloric acid, and causing the evolved chlorine to liberate its 
 equivalent of iodine. Thus, a chromate treated in this manner 
 yields 3 atoms of chlorine for each (CrO 3 ) group according to the 
 equations 
 
 K 2 CrO 4 + 8HC1 = 4H 2 O + 2KC1 + (CrCl s ) + 3d 
 K 2 Cr 2 O 7 + I4HC1 = 7H 2 O + 2KC1 + 2(CrCl 3 ) + 6C1 
 
 Similarly, a chlorate, when treated with hydrochloric acid, yields 
 a mixture of chlorine and chlorine peroxide, which is capable of 
 liberating iodine to the extent of 6 atoms of iodine for each (C1O 3 ) 
 group ; thus 
 
 4KC10 3 + I2HC1 = 4KC1 + 6H 2 O + Qd + 3C1O 2 
 9C1 + 3C1O 2 + 24X1 = I2KC1 + 6K 2 O + 24!^ 
 
 Or the relation of the chlorate to the available oxygen may be more 
 clearly seen by the following equations : 
 
 (1) KC10 3 = KC1 + 30 
 
 (2) 30 + 6HC1 = 3H 2 O + 6C1 
 
 In the case of bromates and iodates, only 4 atoms of chlorine 
 are disengaged for each group (BrO 3 ) and (IO 8 ), and therefore only 
 
Estimations by means of Iodine and Thio snip hate. 359 
 
 4 atoms of iodine are liberated when the evolved gas is passed into 
 potassium iodide. The reason for this difference is due to the fact 
 that one-third of the chlorine is taken up in decomposing the mole- 
 cule of potassium bromide (or iodide), forming potassium chloride 
 and the monochloride of the halogen, which remain behind in the 
 distilling-flask ; thus 
 
 (1) KIO 3 = KI + 30 
 
 (2) 30 + 6HC1 = 6C1 + sH 2 O 
 
 (3) KI + 6C1 = KC1 + IC1 + 4C1 
 
SECTION IV. 
 
 VOLUMETRIC METHODS BASED UPON PRECIPITATION. 
 
 THE number of volumetric processes which fall into this section 
 is extremely large. In the following typical examples some depend 
 upon precipitation alone, while others are based upon precipitation 
 in conjunction with one of the other methods described in the fore- 
 going sections. 
 
 I. PRECIPITATION BY MEANS OF SILVER NITRATE. 
 
 Deci-normal Silver Nitrate Solution. 
 (16*966 grams 0/~AgNO 3 /r litre.) 
 
 4-2415 grams* of pure silver nitrate are exactly weighed out 
 into a 250-c.c. flask, and dissolved in water. The solution is made 
 up to the quarter litre with cold water. 
 
 One cubic centimetre of this solution contains 0*010766 gram 
 Ag, and is equivalent to 0*00355 gram Cl. 
 
 (i) Estimation of Chlorine in Sodium Chloride.- 
 About i gram of pure sodium chloride is weighed out and dis- 
 solved in water, the solution being diluted up to 250 c.c. 
 
 Fifty cubic centimetres of this solution are transferred to a 
 beaker by means of a pipette, and three or four drops of a solution 
 of potassium chromate f added to serve as indicator. The deci- 
 normal silver solution is then gradually run in from a burette. 
 
 As soon as the whole of the chlorine is precipitated, in the form 
 of silver chloride, any additional silver nitrate interacts with the 
 chromate, precipitating red silver chromate. As soon, therefore, 
 as a drop of the silver nitrate produces a permanent reddish tinge 
 in the liquid, the titration is complete. 
 
 * Unless a number of determinations are to be made, it is not necessary to 
 prepare more than a quarter of a litre of the silver solution. 
 
 f The neutral or yellow chromate K 2 CrO 4 . For this titration the solutions 
 must be neutral and cold 
 
Volumetric Precipitation Methods. 361 
 
 The process may be once or twice repeated for the sake of 
 practice ; and in order to gain experience in detecting the first 
 appearance of the red colour, a second beaker, containing about 
 an equal volume of sodium chloride and the chromate indicator, 
 to which a few drops of silver nitrate solution have been added, 
 may be placed alongside the one in which the titration is being 
 made, so as to compare the colours.* 
 
 Alternative Method. The precipitation of chlorine by silver 
 nitrate (or vice versa, of silver by means of a soluble chloride) is 
 one of the very few examples of a precipitation process in which 
 the end reaction can be determined without recourse to a third 
 substance as indicator. This is due to the peculiarities of the 
 precipitate, which allow of its ready subsidence, leaving a clear 
 liquid after being shaken. Moreover, the chromate indicator is 
 inadmissible in the presence of acids (silver chromate being soluble) ; 
 hence, before it can be employed, the solution to be titrated, if acid, 
 must be first carefully neutralised with sodium carbonate. The 
 following method is carried out in an acid solution : 
 
 Fifty cubic centimetres of the solution of sodium chloride pre- 
 pared as above, are transferred to a well-fitted stoppered bottle 
 (narrow mouth), and a few drops of strong nitric acid added. The 
 bottle should be screened from bright daylight during the process. 
 This is conveniently done by rolling a strip of black paper (or 
 common brown paper) into a short cylinder about the same height 
 as the bottle, and into which the bottle will just fit ; or the bottle 
 may be wrapped in a black cloth. The deci-normal silver is then 
 added from a burette, a considerable proportion of the total 
 quantity required being run in at once. The bottle is then briskly 
 shaken without exposure to daylight, and the precipitate allowed to 
 settle. If any floats or adheres to the sides, it can be made to sink 
 by gently tapping the bottle upon the table. As soon as the liquid 
 has clarified, more of the silver is added, and the shaking repeated. 
 This is repeated until the addition of the silver solution causes no 
 further precipitation. As the point is reached when the precipita- 
 tion is nearly complete, the solution clarifies less readily ; hence, 
 until experience is gained, the exact point of completion is liable to 
 be overstepped. It is desirable, therefore, when the precipitation 
 
 * Instead of the chromate indicator being employed, the chlorine may be 
 completely precipitated by the addition of an excess of the deci-normal silver, 
 a few drops of nitric acid being first added. The mixture is then filtered, and 
 the excess of silver present in the solution is determined by titration with 
 ammonium thiocyanate, with a ferric salt as indicator (see p. 365). Bromides 
 and iodides mav be determined in the same manner. 
 
362 Volumetric Analysis. 
 
 is judged to be complete, to add a few drops of deci-normal sodium 
 chloride from a burette (see p. 364), which, if excess of silver has 
 been added, will cause a further turbidity in the liquid. The 
 amount of silver equivalent to the deci-normal sodium chloride 
 thus added, is deducted from that which has been used. 
 
 In cases where the chloride is present only in very small 
 quantities, as in the estimation of chlorides in natural waters, the 
 analysis may be carried out by the first of the methods here given. 
 If the second process be employed, a carefully measured volume 
 (about 20 c.c.) of deci-normal sodium chloride should be added to 
 50 c.c. of the water, and the mixture titrated with the deci-normal 
 silver. The volume of the silver solution equivalent to the amount 
 of added chlorine, deducted from the total silver solution used, 
 gives the amount of silver required to precipitate the chlorine 
 originally present in the water. 
 
 (2) Estimation of Cyanogen in Potassium Cyanide. 
 
 When excess of silver nitrate is added to potassium cyanide, 
 the reaction which takes place is analogous to that between silver 
 nitrate and potassium chloride, silver cyanide being precipitated 
 
 KCN + AgNO 3 = AgCN 4- KNO 3 
 
 N 
 And therefore i c.c. silver nitrate, containing 0^01766 gram Ag, 
 
 is equivalent to 0-0026 gram (CN). Silver cyanide, however, is 
 soluble in potassium cyanide, molecule for molecule, forming the 
 double cyanide (soluble) KCN, AgCN ; therefore, when silver 
 nitrate is gradually added to a solution of potassium cyanide, the 
 first action is the formation of this double salt ; thus 
 
 2KCN + AgN0 3 = KCN,AgCN + KNO 3 
 
 As soon as the whole of the potassium cyanide has been converted 
 into this soluble double salt, the addition of more silver results in 
 its decomposition and the precipitation of silver cyanide ; thus 
 
 KCN,AgCN + AgNO 3 = 2AgCN + KNO 3 
 
 Hence the completion of the first stage can be taken as the end 
 reaction, which will be indicated by the first appearance of a 
 permanent precipitate. 
 
 N 
 For the completion of the first stage, i c.c. silver nitrate will 
 
 be equivalent to 0*0026 x 2 = 0^0052 gram (CN) or 0-0130 gram 
 KCN. 
 
 The first step in the reaction between silver nitrate and potas- 
 sium cyanide is not interfered with by the presence of soluble 
 
Precipitations by means of Silver Nitrate. 363 
 
 halides (chlorides, bromides, iodides). If, therefore, a chloride be 
 added to the solution, no silver chloride is formed until the cyanide 
 has been entirely converted into the double salt. When that stage 
 is complete, then the further addition of silver nitrate results in the 
 precipitation of silver chloride. In fact, the completion of the first 
 reaction is more distinctly seen by introducing a few drops of 
 sodium chloride solution to serve as indicator, than by depending 
 on the decomposition of the double cyanide and the precipitation 
 of silver cyanide. 
 
 About 0*5 gram of potassium cyanide is weighed out and dis- 
 solved in water in a loo-c.c. flask, and the liquid diluted up to the 
 graduation. Twenty-five cubic centimetres of this are transferred 
 to a small beaker by means of a pipette, and deci-normal silver 
 nitrate added from a burette until the first indications of a faint 
 permanent precipitate appear. The beaker should be placed upon 
 a dead black surface. A second portion may be similarly titrated 
 after the addition of one or two drops of a dilute solution of sodium 
 chloride to serve as the indicator. 
 
 (3) Indirect Estimation of Nitric Acid in Nitre.* 
 
 If potassium nitrate is evaporated with strong hydrochloric acid, 
 the radical (NO 3 ) is exchanged for (Cl) ; that is to say, nitric 
 acid is expelled, and potassium chloride is left 
 
 KN0 3 + HC1 = KC1 + HN0 3 
 
 The residue, after being heated for a short time to about 120, is 
 dissolved in water, the solution diluted to a definite volume, and 
 aliquot portions titrated with deci-normal silver, using the chromate 
 indicator. The chlorine which is estimated is thus the measure of 
 the nitric acid, 35*5 parts of Cl displacing 62 parts of (NO 3 ) or 54 
 
 N 
 parts of (N 2 O 5 ). Therefore i c c. silver is equivalent to 0-0062 
 
 gram (NO 3 ) or 0*0054 gram (N 2 O 6 ). 
 
 The nitrate of any metal which is thus capable of being con- 
 verted into a chloride, and whose chloride can be heated without 
 undergoing decomposition or volatilisation, may be determined in 
 this way. 
 
 (4) Indirect Estimation of Calcium in Calcium Car- 
 bonate. The calcium carbonate is dissolved in hydrochloric acid, 
 and the solution evaporated to dryness. The residue is dissolved in 
 water and again evaporated to dryness, and finally heated to about 
 120 to expel the excess of acid. The residue consisting of calcium 
 chloride is dissolved in water, diluted to a definite volume, and the 
 chlorine estimated in an aliquot portion by means of silver. The 
 
 * The four following examples serve to indicate the various ways in which 
 indirect estimations can be made by means of standard silver nitrate. 
 
364 Volumetric Analysis. 
 
 chlorine is thus the measure of the calcium, 35*5 parts of chlorine 
 being equivalent to 20 parts of calcium 
 
 CaCO 3 + 2HC1 = CaCl, + H 2 O + CO, 
 
 N 
 Hence i c.c. silver is equivalent to 0*0020 gram of Ca. 
 
 (5) Indirect Estimation of Carbon Dioxide in Sodium 
 Carbonate. The sodium carbonate is dissolved in water, and pre- 
 cipitated with barium chloride. The barium carbonate is filtered 
 and thoroughly washed, and then dissolved upon the filter in dilute 
 hydrochloric acid. The liquid is thoroughly washed through the 
 filter and evaporated to dryness at least twice (as directed above), 
 and afterwards heated to about 120. It is then dissolved in 
 water, and made up to a definite volume. An aliquot portion is 
 taken and a slight excess of sodium sulphate added, in order to 
 precipitate the barium. The chromate indicator is then added (the 
 precipitate of barium sulphate exerting no disturbing influence upon 
 the titration), and deci-normal silver run in until the red colour 
 appears 
 
 Na 2 CO 3 + BaCl 2 = 2NaCl + BaCO 3 
 BaCO 3 + 2HC1 = BaCl 2 + CO 2 + H 2 O 
 
 The chlorine estimated is the measure of the carbon dioxide, 
 2 atoms of Cl being equivalent to i molecule of CO 2 ; or 35'5 
 
 parts of Cl are equivalent to 22 parts of CO 2 . Hence i c.c. - 
 
 silver is equivalent to 0*0022 gram of CO 2 . 
 
 (6) Indirect Estimation of Cadmium in Cadmium 
 Sulphate. The sulphate is converted into carbonate by precipi- 
 tation with sodium carbonate. The precipitate is filtered, washed, 
 and dissolved in hydrochloric acid. The solution is evaporated to 
 dryness once or twice, and the residue heated to expel the excess of 
 acid. It is then dissolved, and the chlorine in it estimated by deci- 
 normal silver with chromate indicator 
 
 CdSO 4 + Na^COa = CdCO 3 + Na 2 SO 4 
 CdC0 3 + 2HC1 - CdCl 2 + C0 2 + H 2 
 
 Chlorine is thus the measure of cadmium, 35*5 parts of Cl being 
 equivalent to 56 parts of cadmium. Therefore i c.c. silver is 
 equivalent to 0*0056 gram Cd. 
 
 II. PRECIPITATION BY MEANS OF SODIUM CHLORIDE. 
 Deci-normal Sodium Chloride Solution. 
 
 (5-85 grams NaCl per litre.} 
 
 5-85 grams of purified sodium chloride are weighed out and 
 dissolved in water, the solution then being diluted up to I litre. 
 
Ammonium Thiocyanate. 365 
 
 One cubic centimetre contains 0*00585 gram Nad, or 0x30355 
 gram Cl, and is equivalent to 0-010766 gram Ag. 
 
 Estimation of Silver. The details of the process are 
 practically the same as for the estimation of chlorine by means of 
 silver, p. 360. When the chromate indicator is used to determine 
 the end reaction, it should not be added to the silver nitrate ; but 
 a measured volume of the sodium chloride should be delivered 
 into a beaker, the indicator added to this, and the silver solution 
 delivered from a burette until the red colour makes its appearance. 
 
 III. PRECIPITATION BY MEANS OF AMMONIUM THIOCYANATE. 
 Deci-normal Ammonium Thiocyanate Solution. 
 
 (7-6 grams of (NHJCNS per litre.} 
 
 About 8 grams of the crystallised salt (which is too deliquescent 
 to allow of exact weighing) are roughly weighed out and dissolved 
 in a litre of water. The solution is then titrated by means of deci- 
 normal silver nitrate. 
 
 When ammonium thiocyanate is added to silver nitrate, a white 
 precipitate is produced, consisting of silver thiocyanate, insoluble 
 in nitric acid 
 
 AgNO 3 + (NH 4 )CNS = AgCNS + NH 4 NO 3 
 
 This reaction takes place even in the presence of a ferric salt, 
 and it is not until the whole of the silver has been precipitated that 
 the characteristic blood-red ferric thiocyanate makes its appearance. 
 A ferric salt, therefore (but obviously not the chloride], constitutes 
 an extremely delicate indicator for this reaction. The most con- 
 venient ferric salt is the sulphate, a solution of which may be pre- 
 pared by dissolving a few crystals of ferrous sulphate in water in 
 a boiling-tube, adding about half the volume of strong nitric acid 
 (pure), in order to oxidise the iron, and then briskly boiling the 
 mixture for a few minutes, to expel all the nitrous fumes. This 
 solution, which is acid with excess of nitric acid, is then somewhat 
 diluted with water, and a measured volume say 3 or 4 c.c. should 
 be uniformly employed in the titrations.* 
 
 Titration of the Thiocyanate Solution. Twenty-five 
 cubic centimetres of deci-normal silver nitrate are transferred to a 
 small flask, and 4 c.c. of the ferric sulphate solution added.f The 
 
 * When a number of titrations are to be made, the ferric solution should be 
 made up of some definite strength, and the same volume of it used for each 
 operation. 
 
 f The presence of nitric acid is necessary, and if the ferric sulphate solution 
 is prepared as described, it will contain sufficient acid. 
 
366 Volumetric Analysis. 
 
 solution of the ammonium thiocyanate is then run in from a burette. 
 As each drop enters the mixture, the red colour of ferric thiocyanate 
 appears for a moment, but at once disappears on gently shaking 
 the flask. The solution is cautiously added until the first indications 
 of a permanent red coloration are seen. From the volume used, 
 the exact strength of the thiocyanate solution is ascertained, and 
 the amount of dilution which it requires in order to bring it strictly 
 to deci-normal strength is readily calculated. After the solution 
 has been thus adjusted, it should be once more titrated. 
 
 N 
 i c. c. thiocyanate contains 0*0076 gram (NH 4 )CNS, and 
 
 is equivalent to o'o 10766 gram Ag, o'oo355 gram Cl, or 0-0127 
 gram I. 
 
 (1) Estimation of Silver. The estimation of silver is con- 
 ducted exactly as in the titration of the standard thiocyanate 
 solution. The presence of other metals, as lead, copper, zinc, 
 exerts no disturbing effect so long as the solution contains free 
 nitric acid. Hence the silver in such an alloy as silver coinage, 
 may be determined directly in the nitric acid solution, by diluting 
 the liquid to a definite volume, and titrating an aliquot portion with 
 deci-normal thiocyanate, and the ferric sulphate indicator. 
 
 (2) Indirect Estimation of Chlorine (Bromine or 
 Iodine). The solution of the chloride is acidified with nitric acid, 
 and precipitated by the addition of a measured excess of deci-normal 
 silver nitrate. The mixture is briskly shaken and filtered, and 
 the precipitate washed free from silver nitrate. The filtrate and 
 washings are then titrated with deci-normal thiocyanate, after the 
 addition of the ferric indicator. The excess of silver nitrate thus 
 ascertained, deducted from the total volume taken, gives the amount 
 which was used in precipitating the chlorine ; and from this the 
 percentage of the halogen is calculated. 
 
 The necessity for filtering the solution lies in the fact that silver 
 chloride is dissolved by alkaline thiocyanate. Silver bromide is 
 less easily affected, while upon silver iodide it exerts practically no 
 solvent action. 
 
 IV. PRECIPITATION BY MEANS OF URANIUM ACETATE. 
 
 When a solution of uranium acetate or nitrate * is added to a 
 solution of an orthophosphate, in the presence of sodium acetate 
 
 * These compounds are often spoken of as uranyl salts, and are represented 
 as containing the divalent group (UO 2 )", analogous to the monovalent groups, 
 bismuthyl, BiO (in basic bismuth nitrate, or bismuthyl nitrate, (BiO)NO 3 ,H.O), 
 and antimonyl, SbO (in tartar emetic, (SbO)K(C 4 H 4 O 6 ). 
 
Precipitation Methods, 367 
 
 and acetic acid, a yellow precipitate of uranium phosphate is pro- 
 duced, according to the equation 
 
 (U(X)"(C 2 H 3 2 ) 2 + HNa 2 P0 4 = 2Na(C 2 H 3 O 2 ) + H(UO 2 )"PO 4 
 
 The completion of the reaction may be determined by means of 
 a solution of potassium ferrocyanide, used as an outside indicator. 
 As soon as the uranium solution is in excess, the ferrocyanide gives 
 a red-brown coloration with a drop of the mixture brought in 
 contact with it upon a white plate. 
 
 Standard Uranium Acetate, (UO 2 )"(C 2 H 3 O 2 ),2H 2 O. This 
 solution is made of such a strength that i c.c. is equivalent to 
 0-005 gram P 2 O 5 . 
 
 From the equation given above, it will be seen that 425-8 parts 
 of the crystallised salt are equivalent to i^ 2 = 71 parts of P 2 O 5 ; 
 hence 
 
 71 : 5 :: 425*8 : 29^986 = weight of uranium acetate required 
 per litre * 
 
 About 31 or 32 grams, therefore, of the crystallised salt are 
 weighed out and dissolved in water in a litre flask, with the 
 addition of 30 to 40 c.c. of glacial acetic acid. The solution is then 
 diluted up to 1000 c.c. with water. 
 
 Titration of the Uranium Solution, (a] By means of 
 Hydrogen Disodium Phosphate. 10*085 grams of pure recrystallised 
 hydrogen disodium phosphate (which has not become effloresced 
 by exposure to the air) are exactly weighed out and dissolved in 
 water, the solution then being made up to one litre. 
 
 One cubic centimetre of this solution will contain 0*002 gram, 
 P 2 O 5 . It is therefore two-fifths of the strength of the standard 
 uranium solution, 20 c.c. of which should be equivalent to 50 c.c. 
 of the sodium phosphate. 
 
 Fifty cubic centimetres of the sodium phosphate solution (con- 
 taining 0*1 gram, P 2 O 5 ) are transferred to a beaker by means of a 
 pipette, and heated until it just begins to boil t (the titration is 
 best made at a temperature about 95). The uranium solution 
 is then run in from a burette, until, on placing a drop of the hot 
 
 * If uranium nitrate, (UO^"(NO 3 }o,6H 2 O, be used instead of the acetate, 
 the exact weight of the salt per litre to give a solution of the same value will 
 be 35 '47 grams, the formula- weight of this salt being 503 '8. About 36 or 37 
 grams, therefore, should be weighed out. 
 
 f If uranium nitrate has been employed, 5 c.c. of a lo-per-cent. solution of 
 sodium acetate must be added. This is not necessary in the case of uranium 
 acetate, as by double decomposition alkaline acetate is produced during the 
 titration. 
 
368 Volumetric Analysis. 
 
 mixture upon a plate, and adding a drop of potassium ferrocyanide 
 to it, a red-brown coloration is seen.* 
 
 Duplicate titrations giving concordant results having been 
 made, the exact value of the uranium solution is ascertained, and 
 if necessary, it may be diluted by the addition of the requisite 
 amount of water to bring it to the correct strength. Thus if 19 c.c. 
 instead of 20 c.c. were required for the titration of 50 c.c. of the 
 phosphate, then every 190 c.c. must be made up to 200, and the 
 solution titrated once more. If the solution is more nearly exact, 
 say I9'5 c.c. being required instead of 20 c.c., it may be employed 
 in this state, and the volume used in an analysis must then be 
 
 multiplied by the necessary factor ; in this case -- 1-02564. 
 
 () By means of Tricaldum Phosphate. In order to accurately 
 determine phosphoric acid in its combinations with lime or magnesia 
 (as in bone-ash, for example), it is necessary that the uranium 
 solution should be titrated upon a solution of calcium phosphate, 
 so that the conditions of the titration of the standard solution may 
 be the same as those obtaining in the actual analysis. 
 
 The solution of tricalcium phosphate, Ca 3 (PO 4 ) 2 , employed is 
 prepared so as to contain 5 grams of the phosphate per litre. This 
 corresponds very nearly to the strength of the sodium phosphate 
 solution above, and contains 0-00229 gram P 2 O 5 , per cubic centimetre. 
 
 A quantity of the purest commercial tricalcium phosphate is 
 weighed out, which contains exactly 5 grams of Ca 3 (PO 4 ) 2 .t This 
 
 * The formation of the red-brown coloration (uranium ferrocyanide) is 
 influenced by the presence of varying quantities of other salts (alkaline acetates, 
 etc. ), which may be present. It is therefore necessary to make the conditions 
 as nearly as possible the same in all cases. The same volumes of solutions of 
 approximately the same strength and at the same temperature being used ; and 
 the same interval of time allowed for the development of the brown colour 
 with the indicator. Again, if the uranium solution is to be employed for 
 estimating phosphoric acid in solutions which contain ammonia (such as in 
 urine, for example), a definite quantity of ammonium acetate (5 c.c. of a 10- 
 per-cent. solution to 50 c.c. of the liquid) should be added to the solution of 
 sodium phosphate when standardising the uranium solution. 
 
 f Even the purest commercial product cannot be trusted to consist wholly of 
 tricalcium phosphate, therefore it is necessary to estimate the phosphoric acid 
 in the particular sample used, by gravimetric methods. For this purpose a 
 weighed quantity, about 0-4 to o'6 gram, is dissolved in nitric acid, and completely 
 precipitated by means of ammonium molybdate. After standing in a warm 
 place some hours (conveniently all night), the precipitate is washed first by 
 decantation, and finally upon a filter with water containing ammonium 
 molybdate and nitric acid. It is then dissolved in ammonia, and the phosphoric 
 acid precipitated as ammonium magnesium phosphate by the addition of 
 "magnesia mixture," and finally weighed as magnesium pyrophosphate (see 
 p. 264). From the phosphoric acid thus found the percentage of Ca 3 (PO 4 ) 2 
 in the commercial calcium phosphate is calculated. 
 
Precipitation Methods. 369 
 
 is dissolved in the smallest quantity of hydrochloric acid, and the 
 solution diluted to 1000 c.c. 
 
 Fifty cubic centimetres of this solution are then transferred to 
 a beaker, 5 c.c. of a lo-per-cent. solution of sodium acetate, together 
 with 3 or 4 drops of acetic acid, are added, and the uranium 
 solution run in from a burette until nearly the whole amount neces- 
 sary has been added.* The mixture is then heated to about 95,t and 
 the titration continued by adding the uranium solution cautiously, 
 until a drop of the solution withdrawn on a rod and tested with potas- 
 sium ferrocyanide gives the first indications of red-brown coloration. 
 
 Estimation of Phosphoric Acid in Bone-ash. About 
 i '5 grams of bone-ash, which has been dried in a steam-oven, are 
 weighed out and dissolved in the least possible quantity of strong 
 hydrochloric acid. The solution is then diluted to 250 c.c. with 
 water. Fifty cubic centimetres of this solution are transferred to 
 a beaker, and 5 c.c. of sodium acetate (lo-per-cent. solution) added, 
 together with 3 or 4 drops of acetic acid. 
 
 Uranium acetate (standardised by means of calcium phosphate) 
 is then added from a burette, until the first indications of colour 
 are perceived when a drop of the mixture is tested on a white plate 
 with a drop of potassium ferrocyanide. The solution is then 
 heated to 9o-95, and again tested. If the coloration again 
 appears, the titration is complete ; if not, the uranium solution is 
 added drop by drop, until the liquid gives a faint colour with the 
 indicator. A second titration is then made, in which nearly the 
 whole volume of uranium (which the first experiment has shown 
 to be required) is added at once before the mixture is heated, and 
 the process finally completed in the hot solution. 
 
 V. PRECIPITATION BY MEANS OF SODIUM SULPHIDE. 
 
 When a solution of sodium sulphide is added to an ammoniacal 
 solution of a zinc salt, zinc sulphide is precipitated, the end of the 
 
 * When standardising the uranium solution, the volume required to be 
 added is known beforehand within a very near limit ; but when estimating 
 phosphoric acid in a phosphate of unknown composition, the first titration will 
 be made with a view to ascertain approximately the volume of uranium 
 required, after which a second experiment is made more exactly. 
 
 f It is necessary that the solution should not be heated until nearly the 
 whole of the phosphoric acid has been precipitated by the uranium solution, 
 otherwise the calcium phosphate itself will be precipitated, and will then no 
 longer interact with the uranium. This difficulty is sometimes evaded by 
 conducting the titration in the reverse way, i.e. by heating a measured volume 
 of the uranium solution in a beaker, and adding the solution of the phosphate 
 from a burette until a drop of the liquid just ceases to give a coloration with 
 ferrocyanide. 
 
 2 B 
 
370 Volumetric Analysis. 
 
 reaction being indicated by means of either an alkaline lead 
 solution or sodium nitroprusside, used as outside indicators. 
 
 Standard Sodium Sulphide Solution. A solution of con- 
 venient strength should contain about 12 grams of Na 2 S per 
 litre, and may be prepared as follows: 12 grams* of sodium 
 hydroxide are weighed out, and dissolved in 50 c.c. of water. One- 
 half of this solution is then saturated with sulphuretted hydrogen, 
 after which the other half is added, and the mixture thoroughly 
 shaken. If the solution still smells of the gas, a small additional 
 quantity of sodium hydroxide is added, and the liquid diluted to 
 one litre. 
 
 Titration of the Sodium Sulphide Solution. The solu- 
 tion of the sodium sulphide is titrated by means of a standard 
 solution of zinc sulphate, which is made of such a strength that i cc. 
 shall contain o'oi gram Zn. Such a solution is prepared by weigh- 
 ing out 44*15 grams of pure crystallised zinc sulphate, ZnSO 4J 7H 2 O, 
 dissolving in water, and diluting up to I litre.t 
 
 Twenty-five cubic centimetres of this solution are transferred to 
 a beaker by means of a pipette, a small quantity of ammonium 
 carbonate added, and then ammonia, until the precipitate is entirely 
 redissolved. The sodium sulphide solution is gradually delivered 
 into this mixture from a burette, with constant stirring. The 
 progress of the operation is watched by withdrawing a drop of the 
 liquid upon the end of a fine glass rod, and bringing it into contact 
 with a drop of sodium nitroprusside upon a white plate. As soon 
 as the slightest excess of sodium sulphide is present, a violet 
 colour is imparted to the nitroprusside solution. The alkaline lead 
 solution, which may be used as indicator in place of the nitro- 
 prusside, is made by adding a solution of lead acetate to a mixture 
 of tartaric acid with an excess of caustic soda, and heating the 
 solution until it is clear. Drops of this solution are placed at 
 intervals upon a piece of filter-paper, and a drop of the solution 
 undergoing titration is placed near to, but not immediately upon, 
 one of these.J As the latter drop spreads and touches the lead 
 
 * The formula-weight of Na 2 S and aNaHO being almost identical, namely, 
 78 and 80. 
 
 f This solution is a little over three times the strength of a deci-normal 
 solution, which would contain 14 '35 grams of the salt per litre. An alternative 
 method of preparing the solution is to dissolve 10 grams of pure zinc in the 
 smallest excess of dilute sulphuric acid, and then make the solution up to i 
 litre with water. 
 
 % The reason for not bringing the drop immediately upon the indicator is, 
 that any of the precipitate of ZnS which is brought out upon the drop of the 
 mixture would be decomposed by the lead salt, with formation of lead sulphide, 
 
Precipitation by means of Sodium Sulpliide. 371 
 
 solution, a black or brown stain of lead sulphide will be shown as 
 soon as the sodium sulphide is in excess. If the sodium sulphide 
 solution contains exactly 12 grams of Na 2 S per litre; then i c.c. 
 (containing 0-012 gram Na 2 S) will be equivalent to 0-04415 gram 
 of ZnSO 4 ,7H 2 O, or to o'oi gram of Zn, as deduced from the 
 following equation : 
 
 ZnSO 4 ,7H 2 O + Na 2 S = ZnS + Na 2 SO 4 + ?H 2 O 
 287 78 
 
 287 : 78 :: 0-04415 : 0-012 
 or 65 : 78 : : o-oi : 0-012 
 
 Hence 25 c.c, of the zinc solution would be exactly precipitated by 
 25 c.c. of the sodium sulphide. 
 
 Estimation of Zinc in Zinc Ores or Alloys. In the 
 
 analysis of zinc-blende (p. 279) or of bronze (p. 274), the zinc may 
 be estimated by the volumetric process here given. The solution 
 is rendered alkaline by the addition of excess of a mixture of 
 ammonium carbonate and ammonia, and the standard sodium 
 sulphide added from a burette, until a drop of the liquid withdrawn 
 upon a glass rod produces a brown coloration with the alkaline 
 lead indicator when applied in the manner described above. 
 
 As the sodium sulphide solution rapidly undergoes change, it 
 should be titrated against the standard zinc solution shortly before 
 being used. 
 
 Other metals, e.g. cadmium and copper, may also be estimated 
 by means of sodium sulphide. 
 
 VI. CLARK'S METHOD FOR ESTIMATING THE HARDNESS OF 
 WATER.* This is a process of precipitation by means of a standard 
 solution of soap. The soap (potassium oleate) interacts with the 
 calcium and magnesium salts present in the water, with the pre- 
 cipitation of calcium and magnesium oleates. In this process, the 
 soap solution itself serves as the indicator, not by any colour-change, 
 but by its property of causing a lather to form when water contain- 
 ing the solution is shaken up. So long as the hardness-producing 
 salts are undecomposed by the soap, no lather is formed on shaking 
 the mixture, but directly the soap solution is in excess a permanent 
 lather is instantly raised on shaking. By this method of estimating 
 hardness, the total hardness is first determined. Another portion 
 is then boiled until the temporary hardness is destroyed by the 
 
 and the titration would appear to be complete when in reality it was not. By 
 depositing the drop at some little distance from the area wetted with the lead 
 solution, the zinc sulphide precipitate is retained where the drop first touches, 
 and only the solution is able to extend, by capillarity, along the paper. 
 * For Hehner's acidimetric method, see p. 328. 
 
3/2 Volumetric Analysis. 
 
 decomposition of the soluble bicarbonates, and the permanent 
 hardness estimated in the sample. The difference between the 
 two estimations represents the temporary hardness. 
 
 Standard Potassium Oleate Solution. A strong solution 
 of this soap is first prepared by rubbing together in a mortar 75 
 grams of the so-called " lead plaster " of the druggist (plicmbi 
 cmplast., Brit. Pharm., which consists practically of lead oleate) 
 and 20 grams of dry potassium carbonate. When the two are 
 thoroughly incorporated, a small quantity of methylated alcohol is 
 added, and the mixture worked up into the consistency of a thin 
 smooth cream. More spirit is then added, and the contents of the 
 mortar rinsed, by means of spirit, into a bottle, and the mixture 
 placed aside to settle. The clear liquid is decanted off through a 
 filter, and the sediment finally washed upon the filter with more 
 spirit. The volume of the liquid may now be made up to 200 or 
 250 c.c. by adding a mixture of equal volumes of spirit and water. 
 
 Titration and Adjustment of the Soap Solution. For 
 this purpose a standard solution of calcium chloride is prepared, 
 containing 0*222 gram of CaCl 2 per litre. This is equivalent to 
 0*200 gram CaCO 3 per litre, and 50 c.c. will therefore be equivalent 
 to 0*01 gram of CaCO 3 . 
 
 0*2 gram of pure calcite, or Iceland spar, is exactly weighed out 
 and dissolved in dilute hydrochloric acid in a covered vessel, pre- 
 ferably a platinum dish. The solution is then evaporated to dryness 
 upon a steam-bath. The residue is treated with water, and the 
 solution again taken down to dryness. This process should be once 
 or twice repeated in order to ensure the removal of all the acid. 
 The residue is finally dissolved, and the solution made up to i litre. 
 
 Fifty cubic centimetres are then transferred to a stoppered 
 bottle of about 250 c.c. capacity by means of a pipette.* The 
 soap solution is then added from a burette I c.c. at a time, 
 and the contents of the bottle briskly shaken after each addition. 
 As the point is approached when a lather is raised, the volume 
 added each time is reduced, and the process continued until the 
 lather remains permanent for two minutes, the bottle being laid 
 upon its side. 
 
 From this preliminary experiment, the amount of dilution which 
 the soap solution will require, in order that exactly 14*25 c.c. shall 
 be necessary to produce the permanent lather with 50 c.c. of the 
 
 * The contents of the pipette must be allowed to flow out "without being 
 blown out by the breath (see p. 307), because it is very important to avoid intro- 
 ducing any carbon dioxide into the bottle, as this gas affects the soap solution. 
 
Precipitation Methods. 373 
 
 calcium chloride solution, can be calculated. The soap should then 
 be diluted by the addition of a mixture of methylated spirit and 
 water in equal volumes, so as to reduce its strength nearly, but not 
 quite, to this strength say to such a strength that about 13*5 c.c. 
 shall give the lather with 50 c.c. of the standard calcium chloride. 
 
 It is then allowed to stand in a stoppered bottle for 24 hours, 
 during which time a slight sediment usually settles out, and the 
 solution slightly loses strength. It is then decanted or filtered, 
 and titrated again against 50 c.c. of the lime solution, and finally 
 adjusted to the exact strength by the addition of more alcohol and 
 water. 
 
 14*25 c.c. of the soap will therefore be equivalent to 0*01 gram 
 CaCO 3 ; and o'oi gram CaCO 3 in 50 c.c. of water equal 2O'oo parts 
 in 100,000, or 20 degrees of hardness. 
 
 (1) Estimation of the Total Hardness in Ordinary Tap- 
 water. Fifty cubic centimetres of the water are introduced into the 
 stoppered bottle as above (see footnote), and briskly shaken in order 
 to expel dissolved carbon dioxide, and the air in the bottle is then 
 either swept out by means of a current of air from a pair of bellows 
 (the blowpipe bellows), or is sucked out by means of a glass tube. 
 The soap solution is then added, i c.c. at a time at first, but 
 diminishing to one or two drops as the operation proceeds, with a 
 brisk shake up after each addition. As soon as the soap is in 
 excess and a permanent lather will remain for two minutes, the 
 bottle being placed upon its side, the titration is complete, and the 
 degree of hardness corresponding to the volume of soap solution 
 employed is ascertained from the table of hardness (see Appendix). 
 If the sample of water is so hard that 50 c.c. require more than 
 16 c.c. of the soap, the water must be diluted by taking, say, half 
 this volume, and diluting it up to 50 c.c. with distilled water which 
 has been recently boiled for 10 minutes to expel carbon dioxide, 
 and quickly cooled. The degree of hardness which this diluted 
 sample contains is doubled in order to express the actual hardness 
 of the water. 
 
 (2) Estimation of the Permanent Hardness in Water. 
 One hundred and fifty cubic centimetres of the water are placed 
 in a half-litre flask and counterpoised upon a rough balance. The 
 water is gently boiled for half an hour in order to completely pre- 
 cipitate the salts which cause the temporary hardness. It is then 
 cooled and replaced upon the balance, and cold, recently boiled 
 distilled water added until the original weight is restored. The 
 water is then filtered (the filter being used without being previously 
 
374 Volumetric Analysis. 
 
 wetted), and 50 c.c. of the clear filtrate titrated with the standard 
 soap solution, as in the previous case. 
 
 The difference between the permanent and total hardness gives 
 the temporary hardness. 
 
 APPENDIX TO SECTION IV. 
 
 Estimation of Copper by means of Potassium Cyanide. 
 When a solution of potassium cyanide is added to an alkaline 
 solution of a copper salt, a soluble double cyanide is formed (see 
 p. 87). As this compound is colourless, while the original copper 
 solution is blue, the end of the reaction is readily seen by the dis- 
 appearance of the blue colour. 
 
 Standard Potassium Cyanide Solution. This solution 
 is made of such a strength that I c.c. shall be equivalent to 0*005 
 gram of copper. As the solution cannot be made of the exact 
 strength by weighing out the actual quantity of cyanide required, 
 an approximate solution is first prepared by dissolving about 25 
 grams of potassium cyanide in a litre of water. The precise copper- 
 value of this solution is then determined by titration against a 
 standard copper solution, from which the exact amount of dilution 
 necessary to reduce it to the required strength is calculated. 
 
 Titratiou and Adjustment of the Cyanide Solution, 
 by means of Standard Copper Solution. Five grams of pure 
 electrolytic copper are weighed out into a litre flask, and dissolved 
 in a mixture consisting of equal volumes of strong nitric acid and 
 water. The flask may be gently warmed upon a steam-bath. When 
 this metal is wholly dissolved, the solution is cooled and diluted up 
 to i litre with cold water. 
 
 One cubic centimetre of this solution will contain 0*005 g ram Cu. 
 
 Twenty-five cubic centimetres of the standard copper solution 
 are transferred to a beaker, and made alkaline by the addition of a 
 slight excess of sodium carbonate (the solution of sodium carbonate 
 being cautiously added, and the beaker covered, to avoid loss from 
 effervescence). 
 
 To this mixture, in which the precipitate is suspended, is added, 
 by means of a pipette, i c.c. of ammonia solution, prepared by 
 mixing i volume of strong ammonia (sp. gr. 0*88) with 2 volumes 
 of water. This dissolves the precipitate, and produces a deep blue 
 solution.* The potassium cyanide solution is then delivered into 
 
 * The reason for neutralising with sodium carbonate before adding ammonia, 
 is to avoid the introduction of ammoniacal salts, which by their presence 
 
Estimation of Copper by means of Potassium Cyanide. 375 
 
 this solution from a burette, until the blue colour is just discharged. 
 The amount of dilution which the cyanide solution requires in order 
 to reduce its strength to the exact standard, i.e. so that 25 c.c. of 
 the copper solution shall require exactly 25 c.c. of the cyanide, is 
 calculated, and the necessary water added. 
 
 Thus, suppose 21 c.c., instead of 25, were required in the first 
 titration, then 210 c.c. must be diluted to 250 c.c., or 840 c.c. will 
 be diluted to i litre. 
 
 Estimation of Copper in Copper Ores.* 
 
 Epitome of Process. The ore is dissolved in hydrochloric 
 and nitric acids, and after filtration and dilution, the copper is 
 precipitated by means of a rod of metallic zinc in contact with a 
 piece of platinum. The spongy copper is washed, and then dissolved 
 in nitric acid. The solution is made alkaline with sodium carbonate, 
 the precipitate dissolved by the addition of ammonia, and the blue 
 liquid is finally titrated with standard potassium cyanide in the cold. 
 
 Exactly 5 grams of the ore, in fine powder, are weighed into a 
 flask, and gently warmed with a mixture of 40 c.c. of strong hydro- 
 chloric acid and 10 c.c. of water. Six cubic centimetres of a mixture 
 of strong nitric acid and water in equal volumes are then added, 
 and the solution kept gently simmering for about half an hour, after 
 which it is briskly boiled for about 10 minutes. 
 
 The insoluble gangue is then separated by filtration, the filter 
 being thoroughly washed, and the copper precipitated from the 
 warm liquid by introducing into it a rod of zinc, upon which a 
 strip of platinum is loosely twisted. It is essential that there shall 
 be an excess of zinc, the rod therefore should weigh about 50 grams, 
 and should be as free from lead as possible. In about half an hour 
 the whole of the copper will be precipitated, after which the zinc 
 rod is withdrawn (leaving the platinum in the vessel, as a portion 
 of the copper is deposited upon it), and rinsed into the beaker. The 
 spongy copper is washed free from zinc chloride and other dissolved 
 salts by carefully decanting the liquid, and once or twice washing 
 the precipitate by decantation. The last wash-water is drained 
 
 influence the decolorising of the copper solution by the cyanide. This is some- 
 times obviated to a considerable extent by evaporating the acid solution of the 
 copper down to dryness, to expel the acid. In the ordinary process of the 
 analysis of copper ores, however, this involves much loss of time, hence the acid 
 solutions are neutralised as here described, and therefore the standardisation 
 of the potassium cyanide solution should be conducted with a copper solution 
 under similar conditions as obtain in the analyses for which the cyanide is 
 employed. 
 
 * The method here given for the estimation of copper is essentially that 
 known as the " Mansfield," or as Steinbeck's (the discoverer) process. 
 
376 Volumetric Analysis. 
 
 away as thoroughly as possible, and the spongy copper (both that 
 which is loose, as well as that which is deposited on the platinum) 
 is dissolved by adding 16 c.c. of nitric acid (i volume strong acid 
 to i volume water). The solution is then diluted to 100 c.c. with 
 water.* 
 
 Fifty cubic centimetres of this solution are transferred to a 
 beaker, made alkaline by the addition of sodium carbonate, i c.c. 
 of ammonia (i volume strong ammonia to 2 volumes of water) 
 added, and the blue solution titrated with the standard cyanide. 
 The number of cubic centimetres of the cyanide used, multiplied 
 by 2, gives the volume which would be required for the whole of 
 the copper solution ; and since each cubic centimetre of the cyanide 
 is equivalent to o - oo5 gram Cu, and since, also, 5 grams of the 
 ore were employed for analysis, the number of cubic centimetres of 
 cyanide, multiplied by o'i, gives at once the percentage of copper 
 in the- ore. For example, suppose the 50 c.c. of the copper solution 
 required 41 c.c. cyanide, then 
 
 41 x 2 = 82 cc. = cyanide required for the whole of 
 
 the copper in the 5 grams of ore 
 i c.c. cyanide = 0-005 S ram Cu 
 therefore 82 x 0-005 = 0*410 gram Cu in 5 grams ore 
 Q-4IQ x loo . 
 
 and 
 
 or 82 x o-i = 8*2 
 
 * If the ore contains less than 6 per cent, of copper, the reduced metal may 
 be dissolved in half this volume of nitric acid, the solution diluted somewhat, 
 and whole solution used for one titration. 
 
SECTION V. 
 GAS ANALYSIS. 
 
 General Principles. The principles upon which the various 
 methods of gas analysis are based are themselves extremely simple, 
 although in some cases the apparatus employed, and the manipu- 
 lative details involved, are often somewhat complicated and intricate. 
 
 This is more especially the case when the analysis is to be 
 carried out with that high degree of refinement and exactness 
 which is often necessary in scientific research. Until comparatively 
 recent times these exact, extremely slow, and somewhat recondite 
 methods of gas analysis were practically the only processes avail- 
 able ; but of late years the demand for simpler and more rapid 
 methods, sufficiently exact for technical purposes, has resulted in 
 the invention of numerous forms of "gas apparatus" which are 
 capable of yielding good results with the minimum expenditure of 
 time and trouble.* 
 
 The methods by which gases are estimated may be classified in 
 the following order : 
 
 (1) By Absorption in a Suitable Reagent, and Subse- 
 quent Titration. For example, the carbon dioxide present in 
 a gaseous mixture, such as the atmosphere, may be estimated by 
 bringing a known volume of the gas into contact with an excess of 
 a standard solution of barium hydroxide. Barium carbonate is 
 thereby precipitated, and the excess of barium hydroxide is deter- 
 mined by titration with standard oxalic acid. 
 
 Or, again, the chlorine present in the exit gases from the 
 " bleach " chambers may be estimated by aspirating a measured 
 volume of the gas through a solution of potassium iodide, and 
 subsequently determining the liberated iodine by titration with 
 sodium thiosulphate. 
 
 (2) By Absorption, and Subsequent Measurement of 
 
 * For a description of the more refined and elaborate methods of gas 
 analysis, the student is referred to Sutton's "Volumetric Analysis." In this 
 book only the simpler processes will be described. 
 
378 Gas Analysis. 
 
 the Residual Gas. For example, the carbon dioxide in a 
 mixture of gases is determined by exposing a measured volume of 
 the mixture to the action of caustic potash, and after the whole 
 of the carbon dioxide has been absorbed, the volume of the residual 
 gas is measured. The difference, or the contraction^ represents 
 the carbon dioxide which was present. 
 
 (3) By the Combustion of the Gas, with the Sub- 
 sequent Measurement of the Contraction, and Estima- 
 tion of the Carbon Dioxide, if any is formed. For 
 example, hydrogen in a gaseous mixture may be estimated 
 by adding to a known volume of the gas a measured volume of 
 oxygen (or air) more than sufficient to combine with all the 
 hydrogen. These are caused to unite (by methods described later), 
 and the "contraction" ascertained by again measuring the gas. 
 Since 2 volumes of hydrogen and i volume of oxygen unite to 
 form water (which then practically occupies no volume), two-thirds 
 of the contraction represents the hydrogen originally present. 
 
 When the gas to be estimated contains carbon and hydrogen 
 (as in marsh gas, ethylene, etc.), after the contraction due to 
 combustion has been measured, the volume of carbon dioxide 
 produced is determined by absorption with caustic potash and 
 measurement of the residue, as in (2) above. 
 
 I. ESTIMATION OF GASES BY ABSORPTION AND SUBSEQUENT 
 TITRATION. 
 
 (a) Carbon Dioxide in the Air (Pettenkofer's method). 
 
 Epitome of Process. The sample of air to be examined is 
 contained in a large glass jar of known capacity, which can be 
 closed with a well-fitting bung. The temperature of the gas and 
 the atmospheric pressure are noted, A measured volume of a 
 solution of barium hydroxide of known strength is introduced into 
 the jar, and thoroughly shaken up with the enclosed air until the 
 carbon dioxide is entirely absorbed. Aliquot portions of the 
 solution are then withdrawn, and titrated with a standard solution 
 of oxalic acid. 
 
 Standard Oxalic Acid. In order to simplify the subsequent 
 calculations, the strength of the oxalic acid solution may be 
 arranged so that I c.c. is equivalent to I c.c. of carbon dioxide 
 measured under standard conditions that is to say, i c.c. of the 
 acid is capable of saturating a quantity of barium hydroxide which 
 would be decomposed by this volume of carbon dioxide, the weight 
 of which is 0-00197 gram. 
 
Estimation of Carbon Dioxide in the Air. 379 
 
 Oxalic acid solution of this strength would contain 5 '6442 grams 
 of the crystallised acid, H 2 C 2 O4,2H 2 O, in one litre ; but since 
 dilute solutions of oxalic acid are unstable, it is better to prepare 
 the solution of ten times this strength, and dilute it when required. 
 
 56*442 grams, therefore, of the pure dry crystallised compound 
 are exactly weighed out, dissolved in cold air-free water,* and the 
 solution made up to one litre. For use, 10 c.c. are transferred to 
 a ico-c.c. flask by means of a pipette, and the solution diluted up 
 to loo c.c. with air-free water. 
 
 Barium Hydroxide Solution (baryta-water). forty to fifty 
 grams of crystallised barium hydroxide, Ba(HO) 2 ,8H 2 O, are 
 roughly powdered, and placed in a large stoppered bottle, and 
 a litre of water added. The mixture is thoroughly shaken from 
 time to time until the water is saturated, after which it is allowed 
 to settle. The clear liquid is decanted off or filtered into a stop- 
 pered bottle, and diluted with an equal volume of air-free water. t 
 The stopper of the bottle should be greased with resin cerate, so 
 as to exclude the air as much as possible. 
 
 A large glass bottle or jar t (whose mouth is sufficiently wide to 
 admit the hand and arm, so that it may be readily wiped dry with 
 a glass-cloth) is fitted with a good bung, preferably of rubber. A 
 hole is bored in the bung, and this hole is closed with a small cork 
 (see Fig. 74). The capacity of the jar is ascertained by filling it 
 with water up to the bung, and then measuring the volume of the 
 water. The capacity, which should not be less than 6 litres, and 
 may with advantage be 8 or 10 litres, should be scratched upon 
 the vessel. 
 
 The jar is filled with the air to be tested, by leading into it, right 
 to the bottom, a piece of rubber tube, which is attached to a pair 
 of bellows (either hand or foot-bellows), and blowing a stream of 
 air for several minutes in order to ensure the complete displacement 
 
 * Distilled water which has been recently boiled to expel carbon dioxide, 
 and quickly cooled. 
 
 f This solution will roughly approximate the same relative strength as the 
 oxalic acid above mentioned. A solution of exactly equivalent strength, of 
 which i c.c. o'ooigy gram CO 2 (i.e. i c.c. CO 2 at N.T.P.), would contain 
 14-11 grams of Ba(HO) 2 ,8H 2 O per litre. Such a solution cannot be obtained 
 by direct weighing, as, the salt being efflorescent, its state of hydration is 
 uncertain ; and it also absorbs atmospheric carbon dioxide. A saturated 
 solution at the ordinary temperature contains about 32 grams per litre ; hence 
 if this be diluted with its own volume of water, the solution, as first prepared, 
 will contain about 16 grams per litre. As the solution constantly undergoes 
 change by the absorption of atmospheric carbon dioxide, its value must be 
 determined every time it is used, by titration against the standard oxalic acid. 
 
 J The glass jars used by confectioners answer the purpose very well. 
 
38o 
 
 Gas Analysis. 
 
 of the air already in the vessel. The bung is then inserted. The 
 temperature of the air in the immediate vicinity of the bottle is 
 noted, and also the height of the barometer at the time. 
 
 One hundred cubic centimetres of the baryta-water are then 
 delivered into the jar by means of a pipette, through the hole in the 
 bung, and the cork quickly replaced. The liquid is then made to 
 wet the interior surface of the glass by slowly revolving the vessel 
 upon its side ; and it is left in contact with the gas, being shaken 
 up at intervals, for about half to three-quarters of an hour, by 
 which time the whole of the carbon dioxide will be absorbed. 
 
 While this absorption is going on, the exact value of the baryta- 
 water is determined by transferring 25 c.c. to a small beaker, and 
 adding the dilute standard oxalic acid from a burette until the 
 liquid ceases to give any indication of a brown colour when a drop 
 of it is placed upon a piece of turmeric paper by means of a glass 
 rod. As the alkali is gradually neutralised by the oxalic acid, the 
 brown colour, which at first is very evident, gradually shows more 
 and more as a faint fringe of colour round the edge of the moistened 
 spot upon the turmeric paper, until finally it 
 disappears. 
 
 Suppose, for example, that 27*5 c.c. of the 
 oxalic acid are used, then 100 c.c. of this 
 baryta solution would require 1 10*0 c.c. of the 
 oxalic acid to exactly neutralise it. 
 
 When the absorption of the carbon dioxide 
 is complete, 25 c.c. of the turbid liquid are 
 withdrawn by means of a pipette, to which a 
 piece of narrow glass tube has been attached 
 in order that it may reach to the bottom of 
 the jar. This is introduced through the hole 
 in the bung, in the manner shown in the figure, 
 and when the pipette is full, the piece of tube 
 is detached, and the liquid allowed to drip 
 from the pipette until it reaches the graduation 
 mark. The measured volume is then trans- 
 ferred to a beaker, and titrated with the oxalic 
 acid as quickly as possible, so as to expose it 
 to the air for the shortest time. The end of 
 the reaction is indicated by means of turmeric 
 paper used as described above.* 
 
 * The oxalic acid has no action upon the precipitated barium carbonate. 
 Some chemists use a dilute standard hydrochloric acid instead of oxalic acid ; 
 
 either a - - solution, or one arranged so that i c.c. is equivalent to o'ooi gram 
 
Estimations by A bsorption and Subsequent Titration. 3 8 1 
 
 A duplicate titration should be made with a second portion of 
 the liquid. 
 
 Since 25 c.c. (out of the 100 c.c. of baryta-water originally 
 placed in the jar) have been used for the titration, the volume of 
 oxalic acid used, multiplied by four, will give the amount of oxalic 
 acid required to neutralise the 100 c.c. of baryta-water after the 
 absorption of the carbon dioxide in the known volume of air. 
 This obviously will be less than that required by 100 c.c. of the 
 original baryta-water by just the volume of carbon dioxide which 
 was contained in the sample of air. 
 
 EXAMPLE. 
 
 Capacity of glass jar = 10*5 litres 
 temperature of sample = 16 
 
 barometer reading = 767 mm. 
 
 then I0 ' 5 * 273 * ?67 = 10-01 litres - volume of the 
 sample reduced to N.T.P. 
 
 Value of the baryta solution 
 
 ico c.c. = no c.c. of the standard oxalic acid 
 
 100 c.c. of the baryta used for absorption ; 25 c.c. taken out for 
 titration. 
 
 25 c.c. required 26-45 c.c. standard oxalic acid 
 /. 100 c.c. would require 26*45 * 4 = 105-80 c.c. oxalic acid. 
 1 10 105*80 = 4-2 c.c. = the volume of oxalic acid which is equi- 
 valent to the carbon dioxide absorbed by the 
 loo c.c. of baryta-water 
 
 But since i c.c. oxalic acid = i c.c. CO a at N.T.P. 
 
 4*2 c.c. = 4-2 c.c. CO 2 present in 10*01 litres (or 
 10,010 c.c.) of air 
 
 therefore ^ - = 0'0419 = percentage of CO 2 by 
 ' volume 
 
 (U) Estimation of Sulphur Dioxide in Furnace Gases. 
 
 Epitome of Process. By means of a pipe inserted into the 
 flue, a measured volume of the furnace gases is aspirated, with a 
 suitable aspirator, through a known volume of a dilute standard 
 solution of iodine. The excess of iodine is then titrated with a 
 standard solution of sodium thiosulphate. 
 
 The standard solutions used for this estimation may be the 
 
 N 
 
 1 - solutions described on p. 348, in which case i c.c. of the iodine 
 
 CO 2 . Hydrochloric acid of such a strength that i c.c. shall be equivalent to 
 0-00197 gram CO 2 (i.e. i c.c. CO 2 ) cannot be used, as this would act upon the 
 precipitated barium carbonate. 
 
382 
 
 Gas Analysis. 
 
 (and therefore indirectly i c.c. of the thiosulphate) is equivalent to 
 0-0032 gram of sulphur dioxide. It is more convenient, however, 
 and simplifies the calculations, to employ solutions of such a 
 strength that I c.c. shall equal I c.c. SO 2 measured at N.T.P., i.e. 
 0-002867 gram SO 2 instead of 0*0032 gram. 
 
 Such solutions will contain 11*379 grams of iodine, and 22*22 
 grams of sodium thiosulphate per litre respectively, and they may 
 be prepared either by weighing out these quantities, or by diluting 
 
 the ordinary solutions ; 100 c.c. of the solution (both the 
 
 iodine and the thiosulphate) being diluted to ur6 c.c.* 
 
 In cases where the percentage of sulphur dioxide in the gas 
 under examination is very small, it is better to employ solutions of 
 one-tenth this strength, in which i c.c. is equivalent to OT c.c. SO 2 . 
 
 FIG. 75. 
 
 One hundred cubic centimetres of the standard iodine solution 
 are placed in a flask fitted with a cork carrying two tubes, one 
 reaching to the bottom, while the other ends just below the cork, 
 as shown in Fig. 75. The former of these tubes is connected to a 
 piece of narrow metal pipe, which is thrust into the flue through 
 
 * Standard solutions of such a strength that i c.c. = i c.c. of the gas to be 
 estimated (at N.P.T. ) are sometimes spoken of as normal solutions, or normal 
 gas solutions. The use of the term normal solution for any other solution than 
 those already described on p. 313, is greatly to be deprecated, as tending to 
 introduce confusion. 
 
Sulphur Dioxide in Furnace Gases. 383 
 
 which the furnace gases are passing.* The other tube is attached 
 to the large bottle A, filled with water, which serves as an aspirator. 
 The water is allowed to flow slowly out of the bottle by means of 
 the screw clamp, and is received in a litre cylinder. The flask 
 containing the iodine is shaken at frequent intervals during the 
 process. 
 
 When the cylinder is full to the litre mark, it is either changed, 
 or the clamp is closed for a moment while the cylinder is emptied, 
 and then reopened ; a record being made of the number of litres 
 thus drawn off. 
 
 During the experiment the temperature of the water is taken 
 (which will be the temperature of the gas), and the height of the 
 barometer is noted. 
 
 When 8 or 10 litres t of gas have in this way been aspirated 
 through the iodine solution, the process is stopped, and 25 c.c. of 
 the iodine are transferred to a beaker and titrated with the standard 
 thiosulphate. As soon as the red-brown colour of the iodine solu- 
 tion changes to a straw colour, one or two drops of dilute starch 
 solution are added, and the thiosulphate admitted drop by drop, 
 until the blue colour is discharged. A duplicate titration is made 
 in a second portion of the solution. 
 
 EXAMPLE. Gas drawn from the flue of a coke furnace. J 
 
 Volume of water drawn from aspirator 8 litres 
 
 Temperature ... ... ... ... ... 17 
 
 Atmospheric pressure 760 mm. 
 
 Therefore volume of gas operated upon = = 7*5 37 litres 
 
 2 9 at N.T.P. 
 
 * This pipe should be bent at a right angle at a short distance from the end, 
 and a number of small holes bored along the bent part, while the extreme end 
 is closed either with a metal plug or by being roughly hammered together. In 
 this way the gas is drawn from various points in the horizontal sectional area 
 of the flue. 
 
 f When the gas is very rich in sulphur dioxide, as, for instance, in the flue 
 gases from the sulphur-burners in vitriol works, a much smaller volume, say i 
 litre, need only be drawn through the apparatus. In such a case, the pipes 
 leading from the absorption flask to the flue must be filled with the flue gases 
 before the tube is attached to the flask. Also, in the final correction of the 
 volume of gas operated upon, the volume of sulphur dioxide which has been 
 absorbed must be added to the volume of water measured out of the aspirator. 
 In the example here given, where the amount of sulphur dioxide is small, 
 these corrections would not affect the third decimal figure. 
 
 J The furnace had been running many hours previous to taking this sample, 
 and at the time the analysis was made, the draught of air was checked by 
 partially closing the inlet at the bottom. These conditions were arranged 
 in order to intentionally send up the percentage of sulphur dioxide some- 
 what. 
 
384 Gas Analysis. 
 
 ioo c.c. of iodine employed in the absorption flask 
 
 (i c.c. = i c.c. thiosulphate = o'l c.c. SO 2 at N.T.P.) 
 
 After absorption, 25 c.c. iodine solution required 17 c.c. thio- 
 sulphate. 
 
 .*. volume of thiosulphate required for ioo c.c. = 17 x 4 = 68 c.c. 
 
 and volume of SO 2 absorbed = = 3*2 c.c. 
 
 Hence 75307 c.c. of the furnace gases contain 3*2 c.c. SO 2 
 or 3 2 - X o IO = 425 percentage SO 2 by volume 
 
 II. ESTIMATION BY ABSORPTION, AND MEASUREMENT OF THE 
 RESIDUAL GAS. 
 
 These processes differ from the foregoing, in that they involve 
 
 (1) the manipulation of comparatively small volumes of gas, and 
 
 (2) the accurate measurement of these volumes. 
 
 For the manipulation of small volumes of gas special appa- 
 ratus is required, and for the accurate measurement of gaseous 
 volumes special precautions are necessary. 
 
 The Measurement of Gases. The gas whose volume is 
 to be measured must be contained in a graduated and accurately 
 calibrated tube,* and confined over a suitable liquid. Since no 
 gas is absolutely insoluble in water, the confining liquid which is 
 used in the most exact processes of gas analysis is mercury ; but, 
 in the simpler forms of apparatus about to be described, the con- 
 fining liquid is water, which, with certain precautions, and in the 
 case of a considerable number of gases, can be employed with 
 fairly accurate results. 
 
 The apparatus in which the measurements are made is shown 
 in Fig. 76. This is known as a gas-burette, and consists of a 
 graduated measuring-tube, M, and a pressure-tube, P. The 
 measuring-tube has an ordinary stop-cock at the top, and is fur- 
 nished at the bottom with a three-way tap, which allows of com- 
 munication being established with the pressure-tube or with the 
 outer air at will. This will be seen more clearly in Fig. 77- The 
 space between the two taps has a capacity of ioo c.c., and is 
 graduated into cubic centimetres, with subdivisions of fifths of a 
 
 * Formerly these tubes were called "eudiometers," from the Greek, 
 signifying to measure the clearness or goodness of air ; and this term is still 
 usually applied to such vessels when mercury is employed to confine the gas. 
 It has become the custom, however, in the more modern forms of gas apparatus 
 to call these measuring-tubes " gas-burettes." 
 
Gas-burettes. 
 
 385 
 
 cubic centimetre,* and the graduations are numbered in both 
 directions. Each of the tubes is fitted into a loaded wooden foot 
 or stand, cut in the manner shown in the figure, so that the tubes 
 may be brought as near together as possible when adjusting the 
 level of the liquid to the same height in each. The foot of the 
 
 FIG. 76. 
 
 FIG. 77. 
 
 measuring-tube is made in the form of a clamp, which is shown 
 open in Fig. 77. 
 
 The two tubes are connected by means of a rubber tube, long 
 enough to allow one of the tubes to be raised to the top of the 
 
 * The apparatus here figured is Hempel's modified Winkler's gas-burette. 
 In the simple " Hempel " burette there are no taps at all, the top being closed 
 with a piece of rubber tube and a pinch-cock. 
 
 2 C 
 
386 
 
 Gas Analysis. 
 
 other. In order to avoid taking this rubber tube off and on the 
 burette when the latter requires cleaning, it is better to introduce 
 a piece of glass tube as a connector, as shown at /. When the 
 measuring-tube and pressure-tube have to be disunited, the separa- 
 tion is made by removing this glass connector. 
 
 The gas in the measuring-tube is confined over water, and its 
 volume is measured at atmospheric pressure, which is secured by 
 placing the pressure-tube in such a position relative to the measuring- 
 tube that the water stands at the same level in each (Fig. 76).- 
 When reading the volume of the gas, the lower line of meniscus is 
 taken, just as when reading the volume 
 of liquid in an ordinary burette, and the 
 aids may be employed to render the 
 meniscus visible, as are described on p. 
 310. Before reading the volume of a 
 gas, a regular interval of time should be 
 allowed for the water to drain down the 
 walls of the burette ; 60 seconds is a 
 suitable time. 
 
 Calibration of the Gas-burette. 
 Just as in the case of the ordinary burette 
 used for liquids, the gas-burette requires 
 to be calibrated before being used for 
 analytical purposes. This is done in the 
 following way. The measuring-tube is 
 disconnected from the pressure-tube at 
 the glass union / (Fig. 76), and (without 
 its foot) is supported in an inverted posi- 
 tion in an ordinary clamp, as shown in 
 Fig. 78. The tap a is opened, and B is 
 turned so as to communicate between the 
 burette and the rubber tube. A beaker 
 of water, at the temperature of the room, 
 is brought under the bottom end of the 
 pipette, and the liquid sucked up until the 
 burette is filled above the tap B, when 
 the bottom tap is closed, and the rubber 
 removed. By cautiously opening a (B 
 being left open all the time), water is 
 allowed to drop out until the liquid is 
 just level with the under side of the tap B, when a is again closed. 
 A definite volume (3, 4, or 5 c.c.) of the water is then run out 
 
 FIG. 78. 
 
Correction of Gaseous Volumes. 3^7 
 
 at a time,* and weighed exactly in the manner described on p. 309 ; 
 and from the weight so obtained, the true capacity of the graduations 
 upon the tube are calculated and tabulated, as in the case of ordinary 
 burettes for liquids. 
 
 Correction of Gaseous Volumes. The volume which a 
 given weight of gas will occupy depends upon three conditions, 
 namely (i) the temperature, (2) the pressure, and (3) the degree 
 of humidity of the gas, at the time the measurement is made. 
 
 In order, therefore, that the various volumes observed during an 
 analysis shall be comparable with each other, it is necessary either 
 that the conditions mentioned should remain constant throughout, 
 or that the volumes measured under different, but observed, con- 
 ditions, should be reduced by calculation to a common standard. 
 
 In exact methods of analysis, the latter plan is invariably 
 adopted ; but in the more rapid and somewhat rougher methods 
 employed for technical purposes, the analysis may usually be carried 
 out without disturbing the uniformity of the conditions to an extent 
 which will introduce any material error in the results. 
 
 The recognised standard to which gaseous volumes are reduced, 
 is the volume which the gas would occupy at o and under a pressure 
 of 760 mm., when in the dry state. 
 
 (a) Temperature Correction. The coefficient of expansion of 
 gases is taken as 0-003665, or ^ ; therefore the volume at o equals 
 the volume at t divided by i + oxx>3665/ ; 
 
 V V x 273 
 
 hence V = 
 
 where V = volume at o, and V = volume at /. 
 
 (b] Pressure Correction. The volume of a gas being inversely 
 as the pressure 
 
 VP 
 V - 7 6o 
 
 where V = the volume at P pressure : or, making both corrections 
 together 
 
 = VP VP x 273 
 
 ~ 760(1 + 0-003665/) 760 x (273 + /) 
 
 (c) Aqueous Vapour. The aqueous vapour present in a gas 
 exerts a pressure in opposition to the barometric pressure ; hence 
 the volume of a gas is increased by the presence of aqueous vapour. 
 
 * At the narrow part of the burette, near the tap a, each graduation should 
 be calibrated. 
 
388 Gas Analysis. 
 
 If the gas is saturated with aqueous vapour, and an excess of water, 
 however minute, be present, then the pressure or tension of the 
 aqueous vapour is independent of change of pressure, varying only 
 with change of temperature. The tension of aqueous vapour has 
 been exactly determined for every degree of temperature within 
 limits, and in the table in the Appendix will be found the tension or 
 pressure, in millimetres of mercury, of the vapour of water between 
 the temperatures of 5 and 25. 
 
 In making the necessary correction for aqueous vapour, there- 
 fore, the number of millimetres of mercury representing the tension 
 of aqueous vapour at the particular temperature at which the gas 
 measurement is made, is deducted from the barometric pressure to 
 which the gas is exposed. For example, suppose a volume of gas 
 is measured at atmospheric pressure when the barometer is at 
 760 mm., and the temperature is 15 ; then, on referring to the 
 table, the tension of aqueous vapour at 15 is seen to be 127 mm. 
 Deducting this from the barometric pressure gives 760 127 
 = 747-3 mm. as the true pressure. 
 
 If p stands for the pressure due to aqueous vapour, then the 
 formula 
 
 /) x 273 
 
 V - or 
 
 760(1 + 0-0036650 760 x (273 + /) 
 
 embraces all the corrections necessary to reduce a volume of a gas, 
 saturated with aqueous vapour, to the standard conditions. 
 
 When gases are confined over water, as in the Hempel-Winkler 
 burette, the condition of complete saturation with aqueous vapour 
 is of course always present.* With this apparatus also, the gas- 
 volumes are always read at the atmospheric pressure ; and as the 
 analytical operations are rapidly performed, changes of barometric 
 pressure sufficient to influence the results need not be anticipated. 
 
 Changes of temperature, however, must be guarded against as 
 far as possible, and with this object in view, the apparatus should 
 be handled entirely by the wooden feet, and not by the glass 
 portions. In order to ascertain the temperature of the gas and see 
 how far it is being maintained uniform throughout, a simple and 
 convenient plan is to suspend a thermometer inside the pressure- 
 tube P (Fig. 76) by means of a thread, so that it reaches nearly 
 to the bottom, and remains there during the whole analysis. As 
 the water is continually being passed backwards and forwards from 
 
 * When mercury is employed as the confining liquid, complete saturation 
 of the gas with aqueous vapour is ensured by introducing a drop of water into 
 the measuring-tube, or eudiometer. 
 
Correction of Gaseous Volumes. 389 
 
 the pressure-tube to the measuring-tube, the temperature of the gas 
 may be taken as the same as that of the water over which it is 
 confined ; and if the temperature of the latter does not materially 
 change, that of the gas may be considered as practically uniform. 
 
 As stated above, when the conditions under which gas measure- 
 ments are made are constant, it is not necessary to reduce the ob- 
 served volumes to standard conditions. This will be rendered more 
 obvious by the following concrete example.* 
 
 The original volume of gaseous mixture in the burette measured 
 loo c.c. at atmospheric pressure (i.e. when the water was at the 
 same level in both tubes, M, P (Fig. 76). One constituent, x, was 
 then removed by absorption, and the gas measured again. Its 
 volume now was 75 c.c. at atmospheric pressure. 
 
 The temperature was 15, and the barometric pressure 755 mm. 
 throughout. 
 
 Then (i) without making reduction to standard conditions, 
 we get 
 
 loo-*- 75 = 25 = percentage of x in the mixture 
 (2) On reducing the two volumes by means of the formula 
 
 - - 
 
 760(1 +o-oo 3 665/)' 
 
 (n\ v IPO x (755 ~ 127) 
 
 W V = ^ x (I . 055) = 92 56 
 
 W Vo = 75 x (755 - = 6 
 
 760 x (1*055) 
 
 therefore 92*56 69*43 = 23*13 = volume of x in 92-56 c.c. original 
 gas 
 
 and - * 3 x ICQ - 25 = percentage of x in the mixture. 
 
 Since the tension of aqueous vapour is independent of pressure, 
 then, in the event of any alteration of the barometric pressure 
 taking place during an analysis, it is only necessary to make a 
 correction for pressure ; not necessarily by reducing all the volumes 
 to the standard, but by reducing all to the same pressure as any 
 one of them. 
 
 Thus, in the above example, suppose that between the two 
 measurements the barometer fell from 755 to 750 mm., the tem- 
 perature remaining constant at 1 5, then the following are the data: 
 
 Original volume = 100 c.c. at 15 and 755 mm. 
 After absorbing x, volume = 75 15 75 
 
 * The student who has not had practice in the reduction of gases to normal 
 conditions, is strongly recommended to carefully work out these calculations 
 for himself. 
 
390 Gas Analysis. 
 
 then / 3 = 74*5 = volume which the residual gas 
 
 "55 would occupy if measured under the 
 
 same conditions as the original 
 
 volume 
 
 hence 100 - 74 5 = 2 5'5 = percentage of. r in this mixture 
 
 If now, from the above data, the two volumes be reduced to 
 standard conditions by means of the formula, as in the last 
 example, it will be found that the same result is obtained, namely, 
 25*5 percentage of jr. 
 
 Again, since the tension of aqueous vapour depends upon the 
 temperature, increasing with rise of temperature, change of tempe- 
 rature obviously will produce an alteration of the pressure, even 
 though the barometric pressure remains constant. For example, 
 suppose in the above illustration the 100 c.c. original volume is 
 measured at 15, and the 75 c.c. residual gas is measured at 20, 
 the barometer standing uniformly at 760 mm. ; then the actual 
 pressure in the first case is 760 12*7 (tension of aqueous vapour 
 at 15), and in the second it is 760 17-4 (the tension at 20). 
 
 Hence, if any change of temperature is observed in the 
 gas during the progress of an analysis, the observed volumes 
 must be reduced to the standard by means of the formula 
 
 v - v(P-/) _ , 
 
 760(1 + 0-003665/0 ' 
 
 Collection of Gas for Analysis. If the gas for analysis 
 is collected in the laboratory, say, a sample of ordinary coal-gas, 
 it may be introduced into the burette by first placing the measuring- 
 tube on a higher level than the pressure-tube, and allowing the 
 former to empty. By means of the three-way tap, communication 
 between the measuring-tube and the outer air is then opened, and 
 a rapid stream of the gas passed through the tube from the top, 
 
 * In the Appendix there will be found a table of factors for reducing volumes 
 of gas to standard temperature and pressure, by the use of which much time 
 and calculation will be saved. In the columns opposite to the pressure and 
 temperature at which a gas has been measured, a figure is found, which when 
 multiplied into the observed volume will at once give the corrected volume. 
 Thus, suppose 60 c.c. of gas saturated with aqueous vapour were measured at 
 12 with the barometer standing at 766 mm. Then 
 
 766 10 5 (tension aq. vap. at 12) = 755*5 mm. = actual pressure 
 
 On referring to the table, the figure in the 12 column which falls halfway 
 between 755 and 756 mm. is 0-9522 ; 
 
 therefore 60X0-9522 = 57*132 = corrected volume 
 
 The student is warned against using such factors before he \s> perfectly familiar 
 with the method of calculating out the correction for himself. Such a table is 
 not intended to save him the trouble of learning how to make the necessary 
 calculations, but only to save the time and trouble of continually repeating the 
 &a.ne rather lengthy calculations when he does know how to make them, 
 
Correction of Samples oj Gas, 
 
 391 
 
 until the air has been entirely swept out. The upper tap is then 
 closed, and the lower one turned so as to re-establish communica- 
 tion with the pressure-tube. 
 
 When the available supply of the gas is comparatively small, 
 it may be collected in a glass tube over water (or, if necessary, over 
 mercury), and afterwards transferred to the burette as described 
 below. The tube may conveniently have the form shown in Fig. 
 79. It is first filled with water by sucking the liquid up and 
 
 t 
 
 a 
 
 FIG. 79. 
 
 FIG. 80. 
 
 closing the rubber with a pinch-cock, and the gas is then passed 
 up from below in the usual manner. The tube is then connected 
 to the gas burette, with the same precautions against enclosing air 
 in the joints as given below. 
 
 When the gas is collected away from the laboratory, it should 
 be taken in glass tubes drawn out to a capillary constriction at 
 each end. These tubes are filled either by aspirating the gas 
 
392 
 
 Gas Analysis. 
 
 through them so as to sweep out the air, and then hermetically 
 sealing them at the constrictions ; or by taking them to the spot 
 in a vacuous and sealed condition, and then breaking open one end 
 in the gas to be collected. After the gas has filled the tube, the 
 end is again hermetically sealed. 
 
 In order to transfer the gas from the sealed tube to the burette, 
 a piece of capillary tube, /, bent twice at right angles, is attached 
 to the latter, as shown in Figs. 79 and 80, the joints being wired 
 round. The pressure-tube is then raised, until water completely 
 fills the measuring-tube and drops from the open end of the bent 
 capillary.* 
 
 Near to each end of the sealed tube, a slight scratch with a file 
 is made. Over one end a short piece of rubber tube is slipped, 
 and the projecting portion of it filled up with water. The bent 
 capillary (already entirely full of water) is then introduced into this 
 tube, and the latter secured with binding wire. In this way all air 
 is excluded from the joint. The lower end of the tube is dipped 
 into a vessel of water. 
 
 The tube is broken at the file-mark within the indiarubber 
 
 joint, and the end beneath the 
 water is broken off by means of a 
 pair of pliers. On lowering the 
 pressure-tube and opening the tap 
 a at the top of the measuring-tube, 
 the gas will be drawn over into 
 the burette. 
 
 Absorption of Gases. 
 Except when the absorbing liquid 
 is water (in which case the absorp- 
 tion is made in the measuring- 
 tube of the burette), the absorption 
 of the gases in a mixture is carried 
 out in a separate piece of apparatus. 
 Fig. 8 1 shows the Hempel gas- 
 pipette, which is employed for 
 this purpose. It consists of two 
 bulbs connected by a bent tube, 
 
 mounted upon a wooden stand. The capacity of the larger bulb 
 should be at least 150 c.c., so that when the gas from the full 
 
 * Whenever the available supply of the gas to be analysed renders such a 
 course possible, the water used in the burette should be first saturated with the 
 gas by shaking a quantity of water with some of the gas in a stoppered bottle 
 for a short time. 
 
 FIG. 81. 
 
Absorption Pipettes. 
 
 393 
 
 burette (i.e. 100 c.c.) is transferred t j it, there will be room for a 
 considerable quantity of the reagent also. The bent tube d has a 
 capillary bore, and a strip of white glass or porcelain is fixed 
 behind it, in order to render the position of the liquid in the tube 
 more visible. Upon the end of this tube a short piece of moderately 
 stout rubber tube is secured with wire, and the tube furnished with 
 a pinch-cock. The pipette is filled by pouring the reagent through 
 the rather wide tube b (for which purpose a thistle-funnel should 
 be used, so as to avoid spilling the reagent over the outside), the 
 clamp c being open. Such a volume of the reagent is introduced, 
 that when it is sucked up into the capillary tube nearly to the top, 
 there shall be just a small quantity of the liquid in the smaller bulb. 
 
 When not in use, the tube b should be closed with a small cork, 
 the rubber tube being closed, not with the pinch-cock (which spoils 
 the tube after a time), but by inserting a short piece of glass rod 
 into the tube. 
 
 A separate pipette is used for each reagent, and they should be 
 distinctly labelled so as to indicate the reagent they contain. 
 
 In cases where it is necessary to guard the reagent from contact 
 with the atmosphere (as, for example, with alkaline pyrogallol, 
 cuprous chloride, etc.), the 
 double pipette shown in Fig. 82 
 is employed. It differs from 
 the simple pipette only by the 
 two extra bulbs b, b', which, by 
 being partially filled with water, 
 serve as a guard or water-seal. 
 
 This simple addition to the 
 pipette, however, makes it 
 somewhat more difficult to fill. 
 It is necessary so to arrange 
 matters that when the absorb- 
 ing reagent fills the bulb , 
 the water shall occupy bulb b, 
 so that when the reagent passes 
 up into a', the water shall be 
 driven into b'. If this condi- 
 tion is not properly secured, as the reagent is made to pass back- 
 wards and forwards between a and a', either air will be drawn in 
 through the water in the water-seal, or else some of the water itself 
 will be drawn over into a'. The following is the best method of 
 filling the apparatus : 
 
 FIG. 82. 
 
394 
 
 Gas Analysis. 
 
 The empty pipette is supported in an inverted position, and an 
 ordinary ic-c.c. pipette is connected to the capillary tube, in the 
 manner shown in Fig. 83. To the free end of the latter a piece of 
 narrow glass tube, /, is attached by means of a rubber union provided 
 
 with a pinch-cock. A short 
 length of rubber tube, also 
 carrying a pinch - cock, is 
 attached to the other end of 
 the gas-pipette. The air within 
 the apparatus is then swept 
 out by passing a stream of 
 some inert gas (i.e. inert to- 
 wards the particular reagent 
 which is destined to fill the 
 pipette), such as nitrogen, or 
 carbon dioxide. The narrow 
 glass tube is then dipped into 
 the bottle containing the re- 
 agent in question, which is 
 drawn up into the apparatus 
 by applying suction through 
 the rubber tube s. As soon 
 as the bulb a is completely 
 filled (care being taken not to 
 draw any of the liquid over 
 into bulb a'}, the clamps c and 
 c f are closed, and the tube / 
 disconnected. The apparatus 
 is then turned over into its 
 normal position, the little 
 pipette P being supported by 
 the hand. 
 
 The burette (which has previously been filled with the same 
 inert gas used for the pipette) is now attached to the rubber union 
 at c by means of the bent capillary tube, as shown in Fig. 84. The 
 reagent should now occupy the entire space between the clamp <:, 
 and a point x in the bent tube, except probably for a small bubble 
 which will have collected at y. 
 
 The rubber tube S is now removed, and 3 or 4 c.c. of water are 
 introduced into the bulb b' by means of a thistle funnel. This 
 water will partly descend into the bent tube uniting b' with b, and 
 is intended to serve as the temporary water-seal. 
 
 FIG. 83. 
 
Filling the Double Pipette. 
 
 395 
 
 The inert gas in the burette is now slowly passed into the 
 pipette by raising the pressure-tube and opening the tap of the 
 measuring-tube and the clamps c and d. As the gas passes in, the 
 reagent is driven from bulb a to a', while the gas which was in a 
 is expelled through the small quantity of water in b' . When the 
 
 FlG. 84. 
 
 whole of the gas from the burette has been transferred to the pipette, 
 clamp c is closed, and water is introduced into the bulb b' until it 
 is nearly filled. 
 
 The clamp d is now closed, and the little pipette removed. 
 The clamp should be opened again just to allow the liquids to sink 
 to their natural level, and the rubber then closed by means of a 
 
396 Gas Analysis. 
 
 plug of glass rod. The gas-pipette is now properly charged, the 
 space between the reagent and the water-seal being occupied by 
 inert gas, while the confining water occupies such a position in the 
 bulbs b and b' that it will neither pass over into a', nor allow air to 
 pass through when the reagent is being transferred backwards and 
 forwards from a to a'. 
 
 The interposition of the ro-c.c. pipette in the filling operation 
 will have secured the introduction of rather more than enough of 
 the reagent to fill the bulb a. When, therefore, in the process of 
 returning a gas from the absorption-pipette to the burette, the re- 
 agent completely fills the bulb a and the capillary tube, there will 
 still remain a few cubic centimetres in bulb a'. 
 
 Since 100 c.c. of gas (the capacity of the burette) may at any 
 time be introduced into bulb a, it will be evident that the capacity 
 of a', b, and b' must not be less than this volume, otherwise they 
 will overflow. Badly blown specimens are sometimes met with 
 which have this defect ; it will be at once discovered during the 
 operation of filling, and the apparatus must be rejected. 
 
 Gases commonly estimated by Direct Absorption. 
 The gases which are most commonly estimated by simple absorp- 
 tion in a gas-pipette, and the reagents that are employed for the 
 purpose of their absorption, are the following : 
 
 (1) Carbon dioxide; absorbed by potassium hydroxide. 
 The reagent is made by dissolving 150 grams of commercial 
 caustic potash in 500 c.c. of water, and is used in the simple two- 
 bulb pipette. 
 
 (2) Carbon monoxide; absorbed by cuprous chloride. 
 Cuprous chloride is soluble both in ammonia and in hydrochloric 
 acid, and a solution in either solvent may be employed for the 
 absorption of carbon monoxide (see footnote, p. 86). The acid 
 solution is preferable, however, except under the circumstances 
 described on p. 408, when the ammoniacal solution must be used. 
 
 (a) Add Solution. Thirty grams of cuprous chloride* are 
 
 * Cuprous chloride is supplied by the dealers in a fairly pure state. Should 
 the student require to prepare it for himself, it may readily be obtained by dis- 
 solving 25 grams of copper oxide in 200 c.c. of strong hydrochloric acid, and 
 adding about 25 grains of copper clippings or wire. The mixture is then gently 
 boiled for about an hour, more acid being added if the solution becomes much 
 reduced in bulk by evaporation. The dark solution is then poured into a large 
 beaker of water, and the white precipitate of cuprous chloride allowed to settle. 
 The clear liquid is decanted off, and the precipitate transferred to a flask or 
 bottle, where it is again allowed to settle and the liquid poured off. Hydro- 
 chloric acid (strong acid 3 parts, water i part) is then added in sufficient 
 quantity just to dissolve the cuprous ohloride. A few strips of copper are 
 introduced, and the flask or bottle closed with a good cork (a glass stopper 
 should not be used). 
 
Gases Estimated by Direct Absorption. 397 
 
 mixed with 50 c.c. of water in a ^-litre flask, and 150 c.c. of strong 
 hydrochloric acid added. If a few copper turnings or thin strips 
 of the rnetal be put into the brownish solution, and the flask corked 
 up and left for a day or so, the liquid will become colourless. 
 
 () Ammoniacal Solution. Twenty grams of cuprous chloride 
 are mixed with 150 c.c. of water in a flask fitted with a cork carry- 
 ing two tubes, one reaching to the bottom, while the other ends 
 just below the cork. The air is swept out of the flask by a stream 
 of indifferent gas (hydrogen or carbon dioxide), after which the 
 exit tube is made to dip beneath water. A stream of ammonia is 
 then passed into the solution (obtained by gently heating a strong 
 solution of ammonia in a separate flask, the latter not being con- 
 nected until the issuing ammonia has expelled the air) until the 
 cuprous chloride has entirely dissolved. Any unnecessary excess 
 of ammonia should be avoided. This reagent is used in the double 
 pipette, and must be exposed to the air as little as possible during 
 the process of filling the pipette. 
 
 (3) Oxygen ; absorbed by alkaline pyrogallol.* The 
 reagent is prepared by dissolving 20 grams of pyrogallol in 200 c.c. 
 of the potassium hydroxide solution described above (No. i). 
 
 This reagent is used in the double pipette. 
 
 Should the sample of gas under analysis contain more than 
 about 25 per cent, of oxygen (which is seldom the case), it should 
 be diluted with a known volume of hydrogen until the proportion 
 of oxygen is reduced to within this amount ; otherwise the traces 
 of carbon monoxide which are liberated during the absorption of 
 oxygen by alkaline pyrogallol will prejudice the result. If the 
 original gas, however, is known to contain no carbon monoxide, 
 this difficulty may be obviated by transferring the gas, after absorp- 
 tion in the pyrogallol, to the pipette containing cuprous chloride. 
 
 (4) Hydrocarbons (olifines) ; absorbed by faming sul- 
 phuric acid. This reagent is used in the single pipette. 
 
 The same hydrocarbons may also be absorbed by bromine 
 Water, prepared by saturating water with bromine. After ex- 
 posure to either of these reagents, the gas is freed from sulphur 
 dioxide or from the vapour of bromine by being transferred to the 
 caustic potash pipette. The tubes of the pipette containing the 
 fuming sulphuric acid must be closed by means of a piece of glass 
 rod and rubber tube when the apparatus is not in use. 
 
 Benzene vapour is absorbed by fuming sulphuric acid, but not 
 by bromine water. It is also absorbed by fuming nitric acid. 
 * Sometimes, but incorrectly, called " pyrogallic acid." 
 
Gas Analysis. 
 
 (5) Nitric oxide; absorbed by ferrous sulphate. The 
 
 solution of ferrous sulphate is prepared by dissolving 70 grams 
 of the salt in 150 c.c. of water. It is used in a double pipette. 
 
 [Potassium permanganate acidified with sulphuric acid may 
 also be used for the absorption of nitric oxide ; in this case the 
 simple pipette is employed.] 
 
 (6) Chlorine, sulphuretted hydrogen, sulphur dioxide, hydro- 
 chloric acid, and acid gases generally, may be absorbed from 
 gaseous mixtures by caustic potash. 
 
 Estimation of Carbon Dioxide, Oxygen, Carbon Mon- 
 oxide and Nitrogen. As a first exercise in manipulating the 
 gas-apparatus, an absorption-analysis maybe made of an artificially 
 prepared mixture of these four gases, most of which are very 
 constantly met with associated together in such gaseous mixtures 
 as furnace gases, generator gases, water-gas, coal-gas, etc. 
 
 The mixture may be prepared by partially filling (say about 
 three-fourths) the collecting-tube shown in Fig. 79, p. 391, with air. 
 A small quantity of crystallised oxalic acid is heated with strong 
 sulphuric acid in a test-tube fitted with a cork and delivery-tube, 
 and the mixed oxides of carbon collected in the tube so as to fill 
 the remaining one-fourth. The tube now contains the four gases 
 in question. 
 
 They will be estimated by absorption in the following order : * 
 
 (1) Carbon dioxide ; absorbed by caustic potash. 
 
 (2) Oxygen ; absorbed by alkaline pyrogallol. 
 
 (3) Carbon monoxide ; absorbed by cuprous chloride. 
 
 (4) Nitrogen ; estimated by difference. 
 
 (i) Saturating the Water for the Burette. A stoppered 
 bottle of about 300 c.c. capacity is filled with water and inverted 
 in a water-trough. About 100 c.c. of the gaseous mixture is 
 bubbled up into the bottle, which is then closed with the stopper, 
 and the gas thoroughly shaken up with the remaining water for 
 a few minutes. 
 
 Some of this water (now saturated with the gas to be analysed) 
 is poured into the pressure-tube, and the tap at the foot of the 
 measuring-tube is turned so as to establish communication between 
 the two tubes. (This tap is not again touched throughout the 
 analysis, the passage of the gas to and from the measuring-tube 
 
 * It will be obvious that the order in which the gases are to be absorbed 
 from a mixture must be carefully considered. Thus, in the above illustration 
 the oxygen must be absorbed before the carbon monoxide, otherwise the 
 reagent used to remove the latter gas would be acted upon by the oxygen 
 present (see p. 86). 
 
Typical Analysis by Absorption. 
 
 399 
 
 and the various pipettes being controlled entirely by the upper 
 tap.) 
 
 The bent capillary connecting-tube is then attached to the top 
 of the measuring-tube, and the latter is completely filled with water 
 by raising the pressure-tube and opening the tap at the top, until 
 the liquid drips from the end of the capillary tube. The tap is 
 then closed. The indiarubber connector upon the collecting-tube 
 containing the gas is rilled up with a drop of water, and then joined 
 to the end of the bent capillary (see Fig. 79, p. 391), and the con- 
 nections secured with wire. The pressure-tube is lowered, and 
 the pinch-cock upon the collecting-tube and the upper tap of the 
 measuring-tube are opened, whereby gas is drawn over into the 
 latter tube. When sufficient gas has thus been transferred, the pinch- 
 cock and tap are closed, and the two tubes disconnected. One 
 minute is allowed to elapse for the water to drain down the walls 
 of the measuring-tube, when the volume of the gas introduced is 
 read off by lowering the pres- 
 sure-tube until the level of the 
 water in it and in the measur- 
 ing-tube is the same. The 
 graduation mark which coin- 
 cides with the bottom of the 
 meniscus represents the volume 
 of gas taken for the analysis.* 
 
 (2) Estimation of Car- 
 bon Dioxide. The burette 
 is now attached to the absorp- 
 tion pipette containing caustic 
 potash in the manner shown in 
 Fig. 85. Before the two pieces 
 of apparatus are thus united, 
 the potash solution is drawn 
 up so as to completely fill the 
 bulb a, and the bent capillary 
 tube to the mark c, which is 
 made upon the white tablet. The pinch-cock keeps it in this 
 position. After the rubber connections have been secured with 
 
 * It is convenient, when possible, to employ 100 c.c. of the gas under 
 analysis, in which case the number of cubic centimetres of the various consti- 
 tuents which are absorbed, represents the percentage of each ingredient in the 
 gas mixture. If, therefore, more than 100 c.c. has been first introduced into 
 the apparatus, the excess may be removed by raising the pressure-tube until 
 the gas is compressed to exactly 100 c.c. ; then, keeping the water in that 
 
 FlG - 
 
4OO Gas Analysis. 
 
 wire, the pinch-cock is opened, when, if the joints are tight, the 
 reagent will not sink from its position at c. The pinch-cock may 
 be kept open by raising it so that it nips the glass tube. 
 
 The pressure-tube is now raised (being handled exclusively by 
 the foot), and the tap d at the top of the measuring-tube is opened. 
 The gas is thus transferred completely to the bulb a of the pipette. 
 Two or three drops of water from the measuring-tube are allowed 
 to follow the gas into the bulb (by so doing the capillary tube of 
 the pipette is washed each time the apparatus is used), after which 
 the tap d is closed. 
 
 The gas is allowed to remain in contact with the potash for 
 about five minutes, during which time the apparatus is gently 
 shaken so as to moisten the sides of the bulb with the reagent. 
 
 The pressure-tube is then lowered, and tap d opened, whereby 
 the gas is returned to the measuring-tube. As soon as the potash 
 reaches the point in the capillary tube opposite the mark c, the 
 tap is closed. 
 
 On no account must the reagent be allowed to reach the rubber 
 connection, or to pass over into the measuring-lube. 
 
 The pressure-tube is then held in such a position that the water 
 stands at the same level as in the measuring-tube, and, after waiting 
 a minute for the water to drain down the walls of the tube, the 
 volume is read off. 
 
 In order to be sure that the absorption of the carbon dioxide 
 has been complete, the gas is transferred once more to the pipette, 
 and gently shaken with the potash as before, after which it is 
 again returned to the measuring-tube, and its volume read off at 
 atmospheric pressure, allowing the same interval for the water to 
 run down. If the two readings agree, the first absorption was 
 complete.* 
 
 Original volume of gas 100 c.c. 
 
 Volume after absorption by potash .,. 88 
 
 Carbon dioxide 12 = 12 per cent. 
 
 position by pressing a finger upon the rubber tube, the tap at the top of the 
 measuring-tube is momentarily opened. This lets the excess of gas escape, 
 leaving 100 c.c. at atmospheric pressure. This is controlled by again lowering 
 the pressure-tube until the water in each tube is at the same level, when the 
 gas should be found to occupy 100 c.c. 
 
 * If due care be taken to avoid handling the apparatus, except by the foot 
 of the pressure-tube, the temperature of the gas will not undergo any material 
 alteration during the course of an analysis. The temperature may be watched 
 by suspending a thermometer in the water in the pressure-tube, as explained 
 on p. 388. 
 
Typical Analysis by Absorption. 401 
 
 (3) Estimation of the Oxygen. The potash pipette is 
 detached from the bent capillary tube at the joint immediately 
 above the pinch-cock, Fig. 85, and replaced by the double pipette 
 containing the alkaline solution of pyrogallol. Before the latter 
 is connected, the reagent is drawn up into the capillary tube to a 
 marked point, which is as nearly as possible the same distance 
 from the pinch-cock as that upon the previous pipette. The gas is 
 transferred to the pipette for absorption exactly as in the former 
 case, and is left in contact with the reagent, with occasional gentle 
 shaking, for ten minutes. It is then returned to the measuring- 
 tube, the same care being taken to bring the reagent exactly to 
 the mark upon the capillary tube. The tap is closed, and, after 
 allowing time for the water to drain off the walls of the measuring- 
 tube, the volume of the residual gas is read off at atmospheric 
 pressure. 
 
 Volume of gas before absorption of oxygen 88 c.c. 
 after ,, 73 
 
 Oxygen 15 = 15 per cent. 
 
 (4) Estimation of the Carbon Monoxide. The pyrogallol 
 pipette is disconnected and replaced by the one containing cuprous 
 chloride in acid solution, the reagent being previously drawn over 
 into the capillary tube to a mark in the same relative position as in 
 the previous cases. 
 
 The gas is transferred to the pipette and allowed to remain ex- 
 posed to the reagent for ten minutes ; after which it is returned to 
 the measuring-tube, and its volume determined with the same 
 precautions as before. 
 Volume before absorption of carbon 
 
 monoxide ... ... ... ... 73*0 c.c. 
 
 Volume of residual gas (nitrogen) ... 60*5 
 
 Carbon monoxide absorbed ... 12*5 - 12 5 per cent. 
 Hence the composition of the mixture under analysis is- 
 
 Carbon dioxide i2'o 
 
 Carbon monoxide 125 
 
 Oxygen 150 
 
 Nitrogen (by difference) 60*5 
 
 ICO'O 
 
 A second analysis of the same mixture should be made. 
 
 2 E 
 
402 Gas Analysts. 
 
 III. ESTIMATION BY COMBUSTION. 
 
 Hydrogen. When hydrogen and oxygen combine, according 
 to the equation 
 
 2H 2 + O 2 = 2H 2 O 
 
 the two volumes of hydrogen and one volume of oxygen practically 
 cease to occupy space, since the volume of the condensed water is 
 inappreciable. By measuring the shrinkage or contraction which 
 takes place under these circumstances, and multiplying this by , 
 the volume of hydrogen which was burnt is ascertained. 
 
 In order, therefore, to estimate hydrogen, a measured volume of 
 the gas is mixed with a measured volume of air rather greater 
 than is required to furnish the necessary amount of oxygen and 
 the mixture of hydrogen and oxygen caused to unite by one of the 
 methods described below. After the combination the residual gas 
 is measured, when two-thirds of the contraction will represent the 
 volume of the hydrogen consumed. 
 
 Two methods are here given for bringing about the combustion 
 of the mixture of oxygen and hydrogen. 
 
 (a) Combustion of Hydrogen by means of Palladium - 
 ised Asbestos. 
 
 Epitome of Process. The hydrogen is mixed with the 
 necessary excess of air in a gas-burette, which is attached to an 
 absorption-pipette charged only with water. The capillary attach- 
 ment between the two pieces of apparatus contains a thread of 
 asbestos, upon which has been deposited a quantity of finely divided 
 palladium. As the gas is transferred slowly from the burette to the 
 pipette, it passes over this palladiumised asbestos (which is gently 
 warmed by a small flame), and the oxygen and hydrogen are thereby 
 caused to unite. The gas is finally returned to the burette and 
 measured. 
 
 As an exercise in the process, a mixture of pure hydrogen and 
 air may be employed. About 20 c.c. of hydrogen, as pure as 
 possible, are introduced into the gas-burette, and the volume 
 measured. The pressure-tube is then lowered, and a quantity of 
 air drawn into the apparatus until the total volume is about 
 80 or 85 c.c. ; that is to say, from 60 to 65 c.c. of air are introduced, 
 and the volume is again exactly measured. 
 
 The gas-burette is then attached to a gas-pipette charged with 
 water, by means of the capillary combustion-tube containing the 
 palladiumised asbestos * instead of the usual connecting-tube. The 
 
 * The little tube is prepared in the following manner : 0^25 gram, of palla- 
 dium foil is dissolved in the minimum quantity of aqua regia in a small porcelain 
 dish, and the solution evaporated to dryness'on a steam-bath. The residue is 
 
Hydrogen by Combustion. 403 
 
 palladium is gently heated by brushing a Bunsen flame along the 
 tube, which must be kept warm during the whole operation, so as 
 to prevent water (the product of the combustion) from condensing 
 in it. The temperature of the tube must not approach a visible 
 redness. 
 
 The gas is allowed to pass slowly over the warm palladium, 
 which will be seen to glow at the end towards the incoming gas. 
 When the whole of the gas has been transferred to the pipette, it is 
 drawn back again into the burette. This process is repeated once 
 or twice (if the palladiumised asbestos is in good order, one 
 repetition of the operation usually completes the combustion), after 
 which the residual gas is measured. It is then passed once more 
 into the pipette and back, and measured again. If the two 
 measurements agree, the process is complete. 
 
 EXAMPLE. Original volume of hydrogen (by electrolysis) 
 20*5 c.c. 
 
 Excess of air added. Total volume of mixture ... 83-8 c.c. 
 Volume after combustion 53-2 
 
 Therefore contraction 30*6 
 
 30-6 x f = 20*4 c.c. = volume of hydrogen found 
 
 moistened with three or four drops of strong hydrochloric acid, added from a 
 dropping-tube, and 20 drops (i c.c.) of water added. The mixture may be 
 gently warmed to complete the solution. To this red-brown solution, when 
 cold, 20 drops of a cold saturated solution of sodium formate are added. 
 
 Into this mixture, which will not exceed 2^5 c.c. in volume, 0^25 gram of 
 asbestos thread is immersed, which will soak up the whole of the liquid. 
 
 [This thread must be sufficiently fine to admit of being pushed into a 
 capillary tube of i mm. bore. It may be obtained by unravelling a piece of 
 asbestos cloth so as to get single strands. It is cleansed from grease by 
 treating it once or twice with a little warm carbon disulphide in a test-tube, and 
 then spreading it out on a clean piece of paper to dry ; after which it should be 
 heated for a few minutes on a piece of platinum foil. A quarter of a gram of 
 this thread will be about 60 cms. in length.] 
 
 A strong solution of sodium carbonate is added by means of a dropping-tube, 
 and gently worked into the soaked thread with a glass rod, until the mixture is 
 alkaline, and the dish placed upon a steam-bath. A gentle heat suffices to 
 reduce the palladium, which is then precipitated throughout the asbestos as a 
 black deposit. When the contents of the dish are dry, they are rinsed three or 
 four times with hot water in order to dissolve out the soluble salts. The thread 
 is then removed, and cut into short pieces about 4 cms. long. One of these 
 pieces is straightened out with a gentle twisting of the fingers, and laid upon a 
 piece of blotting-paper for a few minutes to remove the superfluous water. It 
 is then introduced into a piece of thick-walled capillary tube, i mm. bore 
 and about 15 cms. long. When the thread has been pushed a little way into 
 the tube, it may readily be drawn into the middle by applying a gentle suction 
 to the other end. The thread is then dried by gently warming the tube and 
 slowly drawing air through it, after which the tube is bent at right angles, 
 about 3 cms. from each end. The same piece of palladiumised asbestos may 
 be used for several combustions. 
 
404 
 
 Gas Analysis. 
 
 (&) Combustion of Hydrogen by Explosion with Air. 
 
 For this purpose the mixture of hydrogen with excess of air is 
 transferred to a special " explosion-pipette," in which are sealed two 
 platinum wires, whereby the gaseous mixture may be ignited by a 
 spark from a Ruhmkorf coil. 
 
 Fig. 86 shows a form of explosion-pipette in which the gas is 
 confined over water.* Water which has been acidulated with 
 
 sulphuric acid, and boiled to 
 expel dissolved gases, is intro- 
 duced at a until the bulb b is 
 just full, and the liquid stands 
 level in the other limb. Upon 
 the tubes a and , pieces of thick- 
 walled rubber tube are securely 
 wired, and a pinch-cock is placed 
 on each. 
 
 At d two platinum wires are 
 sealed into the glass, between 
 which the electric spark is passed 
 when the gas is to be fired.f 
 In the lower part of the tube, at 
 c, two platinum electrodes are 
 fused into the glass. These are 
 for the purpose of adding a small 
 quantity of " electrolytic gas " 
 to the mixture, when the pro- 
 portion of combustible gas is 
 so small that no explosion will 
 take place when the electric 
 spark is passed. Before the 
 " electrolytic gas " is generated, 
 the mixture under analysis is 
 transferred to the measuring- 
 tube, which is then detached 
 from the " explosion-pipette." The two wires from the battery (not 
 from the coil) are then connected to the electrodes (c\ and the 
 oxygen and hydrogen which are evolved are allowed to escape, 
 
 * Another form of explosion-pipette, in which mercury is the confining 
 liquid, is shown and described on p. 407. 
 
 f To prevent the loops of these thin wires from being broken off, it is well to- 
 attach to each a short piece of stouter platinum wire, and pass these through 
 holes in the wooden support so that their ends project and can be bent over 
 into loops, as shown in the figure. The same may be done for the wires at c. 
 
 FIG. 86. 
 
Hydrogen by Explosion. 405 
 
 The current is allowed to pass for about fifteen minutes, in order 
 to saturate the water, after which the current is stopped and the 
 liquid driven up to the usual mark upon the capillary. The burette 
 is reconnected, and the gas returned to the pipette. A small quan- 
 tity of electrolytic gas is then generated (the amount depending 
 upon circumstances), and thoroughly mixed with the gas already 
 present, before exploding. It is not necessary to know the volume 
 of the gas thus added, since it entirely disappears when fired. 
 
 As a first exercise, a mixture of pure hydrogen and air may be 
 exploded. About 10 to 15 c.c. of hydrogen are introduced into the 
 gas-burette, and after being exactly measured, about 60 to 70 c.c. 
 of air are added, and the mixed gases again measured. The burette 
 is then attached to the explosion-pipette, the liquid in the latter 
 being previously drawn up to a mark upon the capillary tube e. 
 
 The gas is then passed over into the pipette, and the clamps 
 upon the rubber tubes both closed. The wires from the induction 
 coil are attached to the wires at d, and the electric spark allowed 
 to pass. The explosion, although not at all violent, will cause a 
 momentary expansion within the apparatus, but if sufficient liquid 
 is present, no gas will be driven out of the bulb-tube. The thick 
 rubber tube upon a being closed, the small quantity of air which 
 is enclosed in the bulb b serves as a cushion or " buffer " at the 
 moment of explosion, and thus relieves the other part of the 
 apparatus from undue pressure. 
 
 The moment after passing the spark, the tube upon a is opened ; 
 and then the gas is returned to the burette and measured. The 
 contraction represents the hydrogen and the atmospheric oxygen 
 with which it has combined to form water, and two-thirds of this 
 shrinkage is the volume of hydrogen which was present. 
 
 EXAMPLE. Volume of hydrogen taken = 12*5 c.c. 
 
 Volume of hydrogen and air ... ... 75'8 c.c. 
 
 Volume after explosion ... 57'5 
 
 Contraction 18*3 
 
 18*3 x f = i2'2 c.c. = volume of hydrogen found 
 
 For practice in the use of " electrolytic gas," the following ex- 
 periment may be made : A quantity of air (about 60 c.c.) is intro- 
 duced into the burette, and its volume measured in the usual way. 
 
 The current from four or five Grove's cells (or its equivalent from 
 any other convenient source) is passed through the dilute acid in 
 the explosion-pipette by means of the electrodes at c (Fig. 86) for 
 
406 Gas Analysis. 
 
 about ten minutes, and the liquid shaken up once or twice with 
 the gas, in order that it may become saturated. The current is 
 then interrupted, and the liquid drawn up to the mark upon the 
 capillary. The burette containing the measured volume of air is 
 then attached in the usual manner. About 12 c.c. of electrolytic 
 gas are then generated in the pipette, and drawn over into the 
 burette.* The mixture of air and electrolytic gas is then passed 
 twice backward and forward from the burette to the pipette, in 
 order to ensure their complete admixture, after which the clamps 
 are closed, and the mixture fired. The residual gas is transferred 
 to the burette and measured, when the volume should be the same 
 as that of the air originally introduced.! 
 
 Methane (Marsh-gas). When a mixture of methane and air, 
 or oxygen, is exploded, carbon dioxide and water are formed ; and 
 from the following equation : 
 
 CH 4 + 2O 2 = CO 2 + 2H 2 O 
 
 it will be seen that methane, when burnt, gives its own "volume of 
 carbon dioxide ; and also that three volumes of mixed gases shrink 
 to- one volume (i.e. the i volume of CO 2 ), since the water ceases to 
 occupy space. The contraction, therefore, is two-thirds of the total 
 volume of the reacting gases ; or, in other words, the contraction 
 is equal to twice the carbon dioxide produced, or twice the volume 
 of the marsh-gas burnt. 
 
 Owing to the effect of pressure in increasing the solubility of 
 carbon dioxide in water, it is only possible to obtain accurate 
 results when mercury is used as the confining liquid.! An explosion- 
 
 * The actual amount of electrolytic gas which has been added may be 
 ascertained by measuring the total volume of gas now in the burette. It is not 
 necessary, however, to know this volume exactly, and a very little experience 
 will enable the operator to judge of the volume by the space it occupies in the 
 explosion-pipette as it is generated. 
 
 t Usually after the first experiment the volume of the residual air is not 
 quite identical with that which was originally taken, owing to the imperfect 
 saturation of the liquid with the various gases. If this is the case, a similar 
 quantity of electrolytic gas should be again added, and the mixture fired once 
 more after thorough admixture. The volume of the residual gas after this 
 second explosion should then exactly agree with that which was measured after 
 the first operation. The process may be repeated with varying amounts of 
 electrolytic gas, and the volume of the residue will be found to remain constaat. 
 
 With a view to ascertain the extent of the loss of carbon dioxide by 
 solution in water when the explosion is conducted over that liquid, the following 
 experiments were made with the apparatus figured on p. 404 : 
 
 (i) 67 c.c. of a mixture of air and carbon dioxide, containing 9 per cent, of 
 CO 2 , were introduced into the apparatus, and 12 c.c. of electrolytic 
 gas added. After explosion, the volume remaining occupied 667 c.c. 
 Loss of CO 2 = o'3 c.c. = o'4 per cent. 
 
Marsh- Gas by Explosion. 
 
 407 
 
 FIG. 87. 
 
 pipette in which mercury is employed is shown in Fig. 87. It 
 differs from the ordinary absorption-pipette only in containing two 
 platinum wires, fused into the upper 
 part of bulb a, and in being fur- 
 nished with a stopcock, in order 
 to close the communication between 
 the two bulbs. 
 
 Before the apparatus is con- 
 nected to the burette containing 
 the measured mixture fo! explosion, 
 the mercury is driven over into 
 the capillary tube, to a fixed mark,* 
 by blowing through the rubber tube 
 upon the upper bulb ; and similarly, 
 when the gas is transferred from 
 the burette to the explosion- pipette, 
 it must be drawn over by applying 
 suction to the same rubber tube. 
 The water from the burette should 
 be made to follow the gas so as to just fill the capillary tube, 
 but without allowing any to enter the bulb. Before firing the 
 mixture, the pinch-cock and the tap are closed. 
 
 (2) 78 c.c. of a similar mixture, containing 20 per cent. CO 2 ; 15 c.c. electro- 
 
 lytic gas added. Volume after explosion 77-2 c.c. 
 
 Loss of CO = 0*8 c.c. = 1*0 per cent. 
 
 (3) 50-8 c.c. of similar mixture, containing 40 per cent. CO 2 ; 15 c.c. elec- 
 
 trolytic gas added. Volume after explosion = 49 c.c. 
 
 Loss of CO 2 = r8 c.c. = 3'3 per cent. 
 
 About the same volume of electrolytic gas was added three times, and the 
 mixture exploded and measured after each addition ; the volumes obtained 
 were 47^4 c.c., 46*2 c.c., and 45*2 c.c., showing a fairly regular loss of carbon 
 dioxide. 
 
 (4) 47-2 c.c. of air, containing 36-4 per cent. CO 2 
 
 (a) 20 c.c. electrolytic gas added ; after explosion, volume 45^6 c.c. 
 () 12 ,, ,, ,, exploded again, volume = 45 'o ,, 
 
 (c) 20 ,, ,, ,, ,, ,. volume- 43-6 ,, 
 
 (a) loss =1-6 c.c., or 3-4 per cent. 
 
 (b) loss = o'6 ,, 
 
 (c) loss = i'4 ,, 
 
 Experiments i, 2, and 3 show that, with about the same force of explosion, 
 the loss of CO 2 increases as the percentage present rises ; while Experiment 4 
 shows that, with the same percentage of carbon dioxide, the amount absorbed 
 depends upon the force of the explosion. 
 
 * In the apparatus here figured, the capillary tube is enamelled white on 
 the back, like the stem of a thermometer ; in other specimens the tube is of 
 clear glass, with a white tablet fixed behind, as in Figs. 81, 82. 
 
408 Gas Analysis. 
 
 After the explosion, the gas is transferred to the burette and 
 measured. The explosion-pipette is then disconnected and replaced 
 by the simple absorption-pipette containing caustic potash, and the 
 carbon dioxide absorbed in the usual way. 
 
 Mixtures of Hydrogen, Methane, and Nitrogen. 
 Many gaseous mixtures which constantly come under analysis 
 (such as coal-gas, producer gas, water-gas, blast-furnace gases) 
 contain varying quantities of these three gases along with others. 
 After all the other gases have been estimated by absorption in their 
 respective reagents, the hydrogen and marsh-gas in the residue 
 are determined by one of the two following methods, while the 
 nitrogen is estimated by difference : 
 
 (a) The gas is mixed with an excess of air, and the hydrogen 
 estimated by combustion by means of palladiumised asbestos (p. 
 402). Under these conditions, the marsh-gas does not burn. 
 
 The marsh-gas is then determined by exploding the residual 
 mixture and absorbing the carbon dioxide, as described above. 
 
 The volume of nitrogen is found by deducting from the original 
 volume of gas the hydrogen and marsh-gas thus determined. 
 
 EXAMPLE. One hundred cubic centimetres of coal-gas were 
 exposed to the action of the following reagents 
 
 (a) Caustic potash ; to absorb carbon dioxide. 
 
 () Alkaline pyrogallol ; to absorb oxygen. 
 
 (c} Fuming sulphuric acid ; to absorb defines and benzene 
 
 vapour. 
 
 (d) Ammoniacal cuprous chloride ; to absorb carbon mon- 
 oxide.* 
 
 The residual gas, measuring 86-6 c.c., was returned to the 
 cuprous chloride pipette, while the burette was disconnected and 
 the water in it (previously saturated with coal-gas) was replaced by 
 water saturated with air. 
 
 Twenty cubic centimetres of the gas were transferred to the 
 burette (the rest being reserved for a subsequent experiment), and 
 air added in more than sufficient quantity for the complete com- 
 bustion of the hydrogen. 
 
 Volume of gas = 20*0 c.c. 
 volume of gas + air = 64/4 
 
 This mixture was then passed over the palladiumised asbestos 
 (p. 402). 
 
 * When the absorption of carbon monoxide is to be followed by the com- 
 bustion of hydrogen with palladiumised asbestos, the ammoniacal solution ot" 
 cuprous chloride should be used. 
 
Mixtures of Hydrogen, Methane, and Nitrogen. 409 
 
 Volume after combustion of hydrogen = 48-2 c.c. 
 
 therefore 64/4 48*2 = 16-2 = contraction 
 
 and 16*2 x = 10*8 = volume of hydro- 
 gen in 20'o c.c. of gas 
 
 To the residual gas (consisting of marsh-gas, nitrogen, and a small 
 surplus of oxygen) an excess of oxygen was added. 
 
 Volume of residual gas = 48*2 c.c. 
 volume of residual gas + oxygen = 69-6 
 
 The mixture was then exploded, and the carbon dioxide absorbed. 
 Volume after explosion = 52*8 c.c. 
 
 therefore 69-6 52*8 = 16*8 = contraction 
 volume after absorption of CO 2 = 44*4 
 
 therefore 52*8 44*4 = 8*4 c.c. = volume of CO 2 pro- 
 duced 
 and therefore 84 c.c. = volume of marsh-gas in 20 c.c. 
 
 of gas 
 
 20 - o c.c. (10*8 + 8*4) = o'8 c.c. = volume of nitrogen in 
 20 c.c. 
 
 Since the original volume of gas taken for analysis was 100 c.c. 
 
 io'8 x 86'6 
 then = 46^46 = per cent, of hydrogen 
 
 and - = 36-07 = per cent, of marsh-gas 
 
 ,0-8 x 86-6 f .. 
 
 and = 3*46 = per cent, of nitrogen 
 
 (b} By this method the mixture of hydrogen, marsh- gas, and 
 nitrogen is mixed with air or oxygen sufficient for the complete 
 combustion of both the combustible gases, and the mixture ex- 
 ploded. The contraction is then measured, after which the carbon 
 dioxide is absorbed and the volume again measured. From these 
 data the volumes of the hydrogen and marsh-gas can be calculated. 
 As already explained, p. 406, the contraction due to the combustion 
 of marsh-gas is twice the volume of carbon dioxide ; if, therefore, 
 the volume of carbon dioxide is ascertained (by absorption with 
 potash), and twice this volume be deducted from the contraction 
 on explosion, the product will represent the contraction due to the 
 combustion of the hydrogen. 
 
 Let C = contraction on explosion, and C' = volume of CO 2 pro- 
 duced (i.e. contraction on absorption with potash) ; 
 
 Then C 2C' contraction due to the hydrogen 
 and |(C - 2C') = volume of hydrogen 
 
410 Gas Analysis, 
 
 Again, since the volume of carbon dioxide produced is the same as 
 the volume of marsh-gas burnt 
 
 C = volume of marsh-gas 
 
 EXAMPLE. A portion of the mixture of hydrogen, marsh-gas, 
 and nitrogen employed in the previous example (being the residual 
 gas after the removal of the absorbable constituents from a sample 
 of coal-gas) was measured in the burette, and an excess of air 
 added. 
 
 Volume of gas taken = 14*2 c.c. 
 volume of gas + air = 97^6 
 after explosion, volume = 74*2 
 therefore contraction, C, = 97^6 74^2 = 23*4. 
 After absorption by KHO, volume = 68'2 c.c. 
 
 therefore C' - 74-2 - 68'2 = 6 
 Hence volume of H in 14*2 c.c. of the gas = f (23^4 12) = 7*6 c c. 
 
 and volume of CH 4 in 14^2 c.c. of the gas = 6'O ,, 
 and volume of N in I4'2 c.c. of the gas= 14 2 (7*6 + 6-0) = O'6 
 Calculating the percentage as in the previous example 
 
 7*6 x 86'6 
 
 = 47 'O per cent, of hydrogen 
 14*2 
 
 and - = 36' 5 per cent, of marsh-gas 
 
 14-2 
 
 = 3'6 per cent, of nitrogen 
 
 14-2 
 
 In cases where the gas under analysis contains a relatively large 
 proportion of nitrogen ; as, for example, in the case of producer 
 gas, or blast-furnace gases, the addition of air would dilute the gas 
 to such an extent as to render it non-combustible. Under these 
 circumstances, therefore, either oxygen must be substituted for air, 
 or else, after sufficient air has been added to furnish the requisite 
 amount of oxygen, a few cubic centimetres of electrolytic gas may 
 be added to the mixture. The electrolytic gas for this purpose 
 may be generated in the pipette described on p. 40.], and then 
 transferred to the mercury explosion-pipette. 
 
 The Nitrometer. 
 
 Many simple processes of gas estimations (such as when it is 
 only desired to determine one gas by absorption) are conveniently 
 and quickly performed by means of the Lunge nitrometer.* This 
 
 * So called because originally designed by Lunge, for the estimation of 
 nitrogen oxides in " nitrous vitriol." 
 
The Nitrometer, or Gas-volumeter. 
 
 411 
 
 consists of a calibrated measuring-tube, m (Fig. 88), connected by 
 means of stout rubber tubing to the pressure-tube p. By means 
 of the two-way tap upon the measuring-tube, communication can 
 be established at will either with the bent capil- 
 lary tube / or the little reservoir n. Mercury 
 is usually employed as the confining liquid. 
 
 To introduce the gas, the pressure-tube is 
 raised until the measuring-tube is completely 
 filled with mercury. The capillary tube / is 
 connected to the supply of the gas, and the 
 two-way tap turned so as to open communication 
 between the measuring-tube and the capillary, 
 as seen in the figure, and the gas drawn over 
 by lowering the pressure-tube. By turning the 
 tap one quarter of a revolution, communication 
 with both the exits is cut off. 
 
 The absorbing reagent is poured into the 
 reservoir ;/, and introduced into the measuring- 
 tube by first slightly lowering the pressure-tube, 
 and then gently turning the tap so as to open 
 communication between m and n* 
 
 Besides simple absorption operations, this 
 apparatus is suitable for the estimation of the 
 volume of gas which is evolved in certain defi- 
 nite chemical reactions. These processes are 
 sometimes spoken of as gas-volumetric analysis. 
 The following are typical examples : 
 
 (i) The estimation of nitrates, either in 
 commercial nitre, or in the residue obtained on 
 evaporating water for the determination of the nitrates and nitrites. 
 
 This process depends upon the fact that when a nitrate is 
 decomposed by strong sulphuric acid in the presence of mercury, 
 this metal is acted upon by the liberated nitric acid with the 
 evolution of nitric oxide. The nitric oxide is therefore the measure 
 of the nitric acid or the nitrate present. One cubic centimetre NO 
 measured at N.T.P. represents 000452 gram of KNO 3 or 0-0038 
 gram of NaNO 3 . 
 
 About o'i gram of the nitrate is placed in the reservoir , and 
 dissolved in 2 or 3 c.c. of water. This solution is carefully drawn 
 
 * As a safeguard, lest by accident the tap should be turned the wrong way, 
 a rubber connector should be left upon the end of /, and closed with a pinch- 
 cock. 
 
 FJG. 
 
412 Gas Analysis. 
 
 into the measuring-tube (previously filled with mercury), and the 
 reservoir rinsed with I c.c. of water, which is also allowed to pass 
 into the tube. (With a little experience these operations will be 
 performed without the admission of any air ; but should air be 
 drawn in, it may be expelled by raising the pressure-tube slightly, 
 and cautiously opening the tap.) The reservoir is again rinsed by 
 the introduction of 5 or 6 c.c. of strong sulphuric acid, which is 
 followed by the addition of about 10 c.c. more acid. The measuring- 
 tube is then released from its clamp, and the contents carefully 
 shaken by inclining the tube, and then quickly (almost with a jerk) 
 bringing it into the vertical position. In fifteen to twenty minutes 
 the process is complete. The tube is replaced in its clamp and 
 allowed to cool. Before reading the volume of gas, in order to 
 compensate for the column of acid in the tube, a similar volume 
 of acid of the same dilution is poured into the pressure-tube. 
 
 (2) By means of a short rubber tube, the tube t may be attached 
 to a small flask fitted with a caoutchouc stopper carrying a short 
 glass tube. If in the flask a chemical reaction resulting in the 
 evolution of gas at the ordinary temperature is carried out, the 
 volume of the evolved gas can be measured. 
 
 With such an arrangement, the estimation of carbon dioxide 
 in carbonates may be made ; also the indirect estimation of 
 manganese dioxide, by measuring the carbon dioxide evolved by 
 the action of the dioxide upon oxalic acid in presence of sulphuric 
 acid, according to the equation 
 
 MnO 2 + H 2 SO 4 + C 2 H 2 O 4 = MnSO 4 + 2H 2 O + 2CO 2 
 
 Similarly, by the action of hydrogen peroxide upon manganese 
 dioxide, potassium permanganate, or bleaching powder, the oxygen 
 evolved is a measure of the available oxygen in these compounds ; 
 half the oxygen given off in each case being derived from the 
 compound, while the other half is from the hydrogen peroxide, as 
 seen by the following equations : 
 
 MnO 2 + H 2 O 2 = MnO + H 2 O + O 
 
 * 
 
 2KMnO 4 4- 3H 2 SO 4 + 5HO 2 = K 2 SO 4 + 2 4 
 
 Ca(OCl)Cl + H 2 2 = CaCl a + H 2 O + O 2 
 
 A weighed quantity of the substance to be tested is introduced 
 into the flask. In the case of manganese dioxide, or potassium 
 permanganate, dilute sulphuric acid is added, but with bleaching 
 
 * In the presence of sulphuric acid. 
 
 f The importance of ensuring the presence of sufficient sulphuric acid in 
 this and the preceding reaction will he seen from the footnote on p. 54. 
 
The Nitrometer, or Gas-volumeter. 413 
 
 powder a small quantity of water only. The reagent (in the above 
 three cases, the hydrogen peroxide) is placed in a test-tube, and 
 deposited within the flask without allowing any of it to come in 
 contact with the materials already present. The rubber cork is 
 inserted in the mouth of the flask, and the apparatus connected to 
 the nitrometer. 
 
 To ensure that the air within the flask is under atmospheric 
 pressure, the two-way tap is turned so as to open communication 
 between the flask and the measuring-tube, and the pressure-tube 
 adjusted so that the mercury stands at the same level in both 
 tubes. The tap is then turned so as to connect the reservoir n 
 with the measuring-tube, and the air from the latter entirely 
 expelled, and the tube filled with mercury by raising the pressure- 
 tube. The tap is then closed and the pressure-tube slightly 
 lowered. Communication with the flask is again established, and 
 the contents of the little tube tipped out into the flask so as to 
 bring about the desired reaction. As the gas is evolved, the 
 pressure-tube is gradually lowered so as to avoid the creation of 
 any unnecessary pressure in the apparatus. When the action is 
 completed, the mercury is brought to the same level, and the 
 apparatus allowed to stand for half an hour for the gas to assume 
 the atmospheric temperature, when the volume is read off. The 
 atmospheric temperature and pressure are noted, and the usual 
 corrections made (see p. 387). 
 
PART III. 
 
 SECTION I. 
 
 ESTIMATION OF CAREON, HYDROGEN, NITROGEN, CHLORINE, 
 SULPHUR, AND PHOSPHORUS IN ORGANIC COMPOUNDS. 
 
 I. Carbon and Hydrogen by Combustion. 
 
 Epitome of Process. These two elements are estimated 
 simultaneously by one and the same series of operations, whereby 
 the organic compound is burnt, and the carbon and hydrogen 
 oxidised to carbon dioxide and water respectively. The oxidation, 
 or combustion, of the organic compound is accomplished by heating 
 it in a combustion-tube with copper oxide, in a gentle stream of air 
 or oxygen. The products of the combustion are thus expelled from 
 the tube, and are made to pass first through a weighed calcium 
 chloride tube, and then through weighed potash bulbs (or soda-lime 
 tubes). The increase in weight of the calcium chloride tube is due 
 to the water, one-ninth of which represents the hydrogen in the 
 organic compound ; while the gain in weight suffered by the potash 
 bulbs is caused by the absorbed carbon dioxide, three-elevenths of 
 which is the weight of the carbon in the compound. 
 
 Fitting up the Apparatus, (i) The Combustion-tube 
 A piece of combustion-tube, about 8 cms. longer than the com- 
 bustion furnace (which itself will be 70 to 80 cms. long*), and about 
 15 mm. bore, is rounded at the edges by just fusing the ends in the 
 blowpipe flame. The tube is then thoroughly cleaned inside. For 
 this purpose one end is closed with a cork, and a few cubic centi- 
 metres of strong sulphuric acid poured in at the other end. This is 
 made to flow over the whole interior surface by tipping and rotating 
 the tube. The acid is then poured away, and the tube thoroughly 
 rinsed out with water, the cork being removed and a stream of water 
 run through the tube. The tube is allowed to drain, and finally 
 dried by gently warming it and blowing air through it from the 
 bellows. 
 
 * For rough measures, 2*5 cms. - i inch. 
 
Carbon and Hydrogen Determination. 415 
 
 Into one end of the tube a roll of copper gauze, c, Fig. 89, is 
 thrust to a distance of about 20 cms. This is made by rolling a 
 strip of fine gauze, 2*5 cms. 
 wide, into a cylinder or plug, 
 which will fit moderately 
 tightly into the tube, so as to 
 retain its place. The tube is 
 then filled up to within 12 
 cms. of the other end with 
 granular copper oxide ; * and 
 a second roll of copper gauze, 
 c, the same length as the 
 other, is pushed into the tube 
 so as to confine the copper 
 oxide. 
 
 The substance under ana- 
 lysis iscontained inaplatinum 
 or porcelain boat, b, which is 
 pushed into the tube nearly 
 close up to the copper-gauze 
 roll. 
 
 Behind the boat is intro- 
 duced a loose-fitting plug, p, 
 the object of which is to con- 
 centrate the current of air or 
 oxygen at this point, in order 
 to prevent the backward diffu- 
 sion of vapours or products 
 of combustion of the sub- 
 stance in the boat. This plug 
 may be either a short roll of 
 copper gauze, having a little 
 wire loop by means of which 
 it can easily be withdrawn ; 
 or, better, a short piece of 
 hard glass tube (which will 
 just pass freely into the com- 
 bustion-tube), closed at one 
 end, and narrowed, but left 
 
 * Copper oxide made by heating the nitrate should not be used, as it is 
 liable to evolve oxides of nitrogen when heated. Granular copper oxide is 
 obtained by direct oxidation of the metal. 
 
41 6 Ultimate Organic Analysis. 
 
 open, at the other end. This can readily be withdrawn by means 
 of a piece of wire, bent into a small hook at the end. 
 
 The combustion-tube is closed at this end with a rubber stopper 
 carrying a short piece of narrow glass tube, which projects about 
 3 mm. into the combustion-tube, the other end being connected to 
 the apparatus for purifying and drying the air and oxygen. The 
 space, S, at the other end of the tube is for the reception of a roll 
 of either copper or silver gauze, about 6 cms. long, and sufficiently 
 loose to admit of easy removal. If the organic compound to be 
 analysed contains nitrogen, either the copper or silver roll will 
 serve to decompose the nitrogen peroxide,* which, if allowed to pass 
 from the tube undecomposed, would be absorbed by the potash, and 
 thereby vitiate the carbon determination. When, however, the 
 compound contains a halogen, the silver roll must be employed. 
 
 Before use, such a copper gauze roll should be first superficially 
 oxidised, by heating for a few moments in a flame, and then heated 
 in a stream of hydrogen. A number of rolls may be prepared at 
 one operation, and may be kept bright by being corked in a tube. 
 
 (2) The Combustion Furnace. When the combustion-tube is 
 fitted up, it is laid in the furnace,t and should project about 3^ cms. 
 at each end. A shallow trough or tray of sheet iron runs the entire 
 length of the furnace, and in this a bed for the tube is made by first 
 laying in the trough a strip of asbestos cloth, wide enough to pro- 
 ject over the sides of the iron trough, so that the tube shall rest 
 uniformly upon the asbestos without coming in contact with the 
 iron. 
 
 (3) The Calcium Chloride Tube. This is filled and prepared 
 precisely as described on p. 203.$ When not in use, the ends of 
 the tube are closed by short rubber tubes plugged with pieces of 
 glass rod. These caps are always removed while the tube is being 
 weighed, but immediately replaced until the apparatus is required 
 for the combustion. 
 
 The calcium chloride tube is attached to the combustion-tube 
 by means of a rubber stopper, and is conveniently supported by a 
 retort stand and ring, as shown in the figure. 
 
 * 4Cu + 2NO 2 = 4CuO + N 2 . 
 
 t Erlenmeyer's gas furnace, which consists practically of a row of Bunsen 
 burners, is one of the best forms. 
 
 As samples of calcium chloride are sometimes slightly alkaline, it is 
 desirable, when the tube is freshly filled, to pass a stream of dry carbon dioxide 
 through it for a few minutes, and then to displace the carbon dioxide by a 
 current of dry air, which should be passed through the tube for about half an 
 hour. 
 
Carbon and Hydrogen Determination. 417 
 
 (4) The Potash Bulbs. These may be either the ordinary 
 Liebig's bulbs, or the 
 more modern Geissler's 
 apparatus, shown at B, 
 Fig. 89. They are 
 charged with a strong 
 solution of caustic 
 potash by dipping the 
 tube projecting from 
 the larger bulb into a 
 small beaker or dish 
 containing the potash, 
 and sucking the liquid 
 into the apparatus by 
 means of a piece of 
 rubber tube upon the 
 opposite end. A little 
 guard tube, G, filled 
 with fragments of solid 
 caustic potash, is at- 
 tached to the bulbs in 
 order to prevent the 
 loss of water-vapour, 
 which would otherwise 
 result from the passage 
 through the solution 
 of the dry gases passing 
 from the combustion- 
 tube. Just as with the 
 calcium chloride tube, 
 when not in actual use 
 the ends of the tubes 
 are closed with rubber 
 caps, which are re- 
 moved while the appa- 
 ratus is being weighed. 
 (5) The Apparatus 
 for Purifying the Air 
 and Oxygen. This 
 apparatus must be 
 fitted up in duplicate, 
 
 so that either air or oxygen can be admitted into the combustion- 
 
 2 E 
 
41 8 Ultimate Organic Analysis. 
 
 tube in a purified condition without the loss of time which would 
 result if one gas were made to follow the other through the whole 
 apparatus. Fig. 90 shows a convenient arrangement for the purpose. 
 
 Each gas (contained in gas-holders, or compressed in cylinders) 
 is passed first through a two-necked bottle, B, B', nearly filled with 
 pumice moistened with sulphuric acid. They then pass through 
 the eprouvettes E, E', which are filled with fragments of solid 
 caustic potash. This part of the apparatus, when once charged, may 
 remain undisturbed for several months, and may conveniently be 
 placed upon a shelf out of the way. The gases may be conducted 
 to it and led away from it by means of fine " compo " pipe 3 mm. 
 in the bore, which may be carried down the wall to the side of the 
 draught chamber, in which the combustion furnace is placed. At 
 this point they are joined to two limbs of a bridle-tube, T, by means 
 of short rubber connections, each with its screw-clamp or stop-cock ; 
 these joints should all be securely wired. The third limb of the 
 bridle-tube is attached to a U-tube filled with soda-lime, and this in 
 its turn is connected to a set of Liebig's bulbs containing strong 
 sulphuric acid. By manipulating the taps, either oxygen or air can 
 be admitted into these two tubes, and the rate at which the gas 
 passes is seen by the bubbles travelling through the bulbs containing 
 the sulphuric acid. 
 
 Testing the Apparatus. When the apparatus is fitted up, 
 it is necessary, before proceeding to make an actual combustion, 
 to ensure the complete removal from the combustion-tube of every 
 trace of moisture and of combustible matter. The calcium chloride 
 tube and the potash bulbs are for this purpose removed, and the 
 roll of copper gauze is introduced into the space S (Fig. 89).* The 
 boat and the glass plug p are also removed. 
 
 The combustion-tube is then heated throughout its entire length 
 to a low red heat for about 20 minutes, while a slow stream of the 
 dried and purified air is passed through it. The air-supply should 
 be then stopped, and the air contained in the oxygen purifying 
 apparatus made to pass through the tube by sweeping it out with 
 a stream of oxygen. This may be passed rather more rapidly than 
 at first, so that in about 15 minutes oxygen will be issuing from the 
 extreme end of the tube. 
 
 * It is advisable to place this roll in the tube at this stage, even although 
 the combustions, which are to be first made, are of substances which do not 
 contain nitrogen ; so that afterwards, when the copper may be required, its 
 freedom from organic matter will be assured. Moreover, in case the ordinary 
 copper oxide (prepared from the nitrate) is the only preparation at hand, the 
 presence of this copper roll will serve to decompose the oxides of nitrogen which 
 are derived from this source. 
 
Carbon and Hydrogen Determination. 419 
 
 While this is going on, the calcium chloride tube and the potash 
 bulbs are weighed, the caps being removed for the short time 
 occupied by the weighing. 
 
 The passage of the oxygen is then stopped, and the weighed 
 tubes attached, the ends of the glass tubes of the two pieces 
 which are connected by the rubber union, being pushed up within 
 the rubber so as to touch each other.* 
 
 A slow stream of air is now passed through the apparatus, the 
 rate being that which would be adopted during an actual combustion, 
 namely, about two bubbles per second, as they pass through the 
 bulbs. This is continued for about 15 minutes, after which the 
 stream of air is replaced by a stream of oxygen at about the same 
 rate for another 15 minutes. The potash bulbs and calcium 
 chloride tube are then carefully disconnected, and their ends im- 
 mediately closed with their caps. After a short interval, during 
 which they are allowed to assume the ordinary temperature, they 
 are weighed (as usual, without the caps), and if the operation has 
 been successfully carried out, their weight should have undergone 
 no increase. If, on the other hand, either piece of apparatus has 
 gained more than 3 or 4 milligrams, it shows that the first part of the 
 operation has not been complete, and that either moisture or carbon 
 dioxide is still being evolved. In this case the passage of air through 
 the still heated tube must be continued for a short time, and another 
 " blank " experiment made by replacing the calcium chloride tube 
 and potash bulbs, and allowing first air, and then oxygen, to pass 
 through the entire apparatus for 15 minutes each. The absorption 
 apparatus is then detached and reweighed, when, unless something 
 is seriously wrong, the weights will be found to have undergone no 
 change. 
 
 The apparatus is now in a condition to be used for a series of 
 " combustions," and need not be further tested until it becomes 
 necessary to replace the combustion-tube. If the asbestos bed has 
 been properly arranged, so that the tube rests uniformly all along 
 it, and if care is exercised in heating the tube, so as not to heat too 
 
 * It is of the greatest importance that the stopper or cork through which the 
 calcium chloride tube passes should not become over-heated during a " com- 
 bustion," and yet at the same time the tube must be sufficiently warm right up 
 to the cork to prevent moisture from condensing before it passes into the bulb. 
 If the combustion-tube projects 4 cms. beyond the furnace, as directed on p. 416, 
 the cork is practically quite safe, but the additional precaution may be adopted 
 of placing a small screen, cut out of asbestos cardboard, over the tube just 
 outside the furnace. This screen, however, should be removed towards the 
 end of the process to make sure that no moisture has deposited before passing 
 into the absorption-tube. 
 
420 Ultimate Organic Analysis. 
 
 suddenly at first, and not to raise the temperature unnecessarily 
 high, whereby the glass is softened, a combustion-tube made of 
 the best hard glass will stand being used over and over again for 
 a great many analyses. 
 
 Typical Examples of Combustions. 
 
 (a) Combustion of Non-volatile Solids. One of the most 
 suitable substances to employ for a first experiment is cane sugar, 
 as it is readily obtained in a state of purity, and the percentage of 
 hydrogen it contains is comparatively high. The pure crystallised 
 compound is finely powdered, and dried in a steam-oven. 
 
 A porcelain or platinum boat (previously proved to be of such 
 a size that it will pass easily into the combustion-tube) is supported 
 on a pipeclay triangle, and heated to low redness by means of a 
 Bunsen flame ; if a porcelain boat is employed, it must be heated 
 cautiously. 
 
 It is then placed to cool in the desiccator, and when cold, it is 
 weighed. It must be handled as little as possible by the fingers, 
 and not at all if the hands of the operator are at all moist. 
 
 From 0*2 to o'3 gram of the dry sugar is weighed into the boat, 
 which is then pushed into the combustion-tube, almost close up to 
 the roll of copper gauze. The clean dry glass plug is next intro- 
 duced, and the tube is then connected to the apparatus for supplying 
 the stream of dry and pure air. 
 
 The weighed calcium chloride tube and potash bulbs are then 
 attached, the projecting end of the combustion-tube being firmly 
 held with one hand while the calcium chloride tube is being fitted 
 on, so as not to disturb the boat and its contents. 
 
 The portion of the tube from S (Fig. 89) to the copper gauze 
 roll c, that is, the part containing the whole of the copper oxide, is 
 first heated to redness, after which a slow stream of air (not faster 
 than two bubbles per second) is passed through. The heating is 
 now gradually extended towards the boat by turning on one burner 
 at a time at first with a small flame, and then increasing it. 
 
 The success of the operation hinges upon this part of the 
 process, therefore it must be carefully watched, lest the heat be 
 applied too rapidly. 
 
 The combustion of sugar is usually completely effected with a 
 stream of air, without the admission of pure oxygen ; but if any 
 traces of a black residue are seen to persist in the boat towards the 
 end of the operation, a stream of oxygen may be substituted for 
 
Typical Examples of Combustions. 421 
 
 the air for a few minutes, after which the oxygen is stopped, and the 
 air again passed through for about 10 or 15 minutes. 
 
 In this way the whole of the water-vapour and carbon dioxide 
 is swept out of the tube into the absorption apparatus ; and the 
 portions of copper oxide which during the combustion had become 
 reduced are re- oxidised. 
 
 The potash bulbs and calcium chloride tube are then removed, 
 their ends capped, and, after being allowed to assume the ordinary 
 temperature, they are weighed. 
 
 H 2 O : H, = 9 : i 
 
 18 2~ 
 
 therefore wei ht of H 2Q = weight of hydroge n 
 
 and CO 2 : C = 11:3 
 44~ 12 
 
 , , weight of CO<> X 3 
 therefore & 8 9 = weight of carbon 
 
 From the numbers thus obtained, the percentage of hydrogen 
 and carbon are calculated. The oxygen is estimated by difference.* 
 
 The theoretical composition of sugar, calculated from the 
 formula C 12 H 22 O n , is hydrogen, 6*43 ; carbon, 42*10 ; oxygen, 51*47. 
 
 The experimental errors of a carbon and hydrogen determination 
 are usually in the direction of a deficit of the former, and a slight 
 excess of the latter, element. If, in the analysis of sugar, the carbon 
 falls more than 0-3 below the theoretical value, or the hydrogen 
 more than 0*1 above it, a duplicate experiment should be made. 
 
 (b) Combustion of Volatile Solids. For practice in the 
 combustion of a volatile solid, benzoic acid may be employed. 
 
 Benzoicacid, C 6 H 5 .COOH : carbon, 68*85 2 per cent. ; hydrogen, 
 4*918 per cent. 
 
 From o'4 to o* 5 gram of the compound is weighed out into the 
 boat, but before the latter is introduced into the combustion-tube, 
 the weighed calcium chloride tube and potash bulbs are attached, 
 and the tube is heated to redness to within about 2^ cms. of the 
 copper roll c (Fig. 89). The boat is then carefully pushed in, the 
 glass plug quickly inserted, and connection made with the gas- 
 drying apparatus. A slow stream of air is admitted, and then the 
 heating very gradually and cautiously extended in the direction of 
 
 * In cases where the substance under analysis leaves an incombustible 
 residue or ash, the boat is withdrawn after the tube has cooled, and placed in 
 a desiccator until cold, when it is weighed. Deducting the weight of the boat 
 gives the weight of the ash. 
 
422 Ultimate Organic Analysis. 
 
 the boat. The more volatile the substance is, the more gradually 
 must the heating be extended towards the boat containing it. 
 
 The remainder of the process is conducted as in the former 
 example, the stream of air being replaced by pure oxygen towards 
 the end of the operation. 
 
 (c) Combustion of Liquid Compounds. If the liquid is 
 not appreciably volatile at ordinary temperatures, and at the same 
 time does not absorb atmospheric moisture, it may be weighed out 
 into the boat and treated as in the above example. 
 
 Moderately volatile liquids are enclosed in small glass bulbs, 
 made by drawing out a piece of narrow glass tube, as shown at A 
 (Fig. 91), and then blowing the small tube into a bulb, B. The 
 
 bulb must be of such a size that 
 it will lie in the boat, and pass 
 into the combustion-tube in this 
 position. 
 
 As an example, benzene, 
 
 A C 6 H 6 , may be employed. One 
 
 FlG - 91- of the little bulbs, with a capil- 
 
 lary stem about 3 cms. long, is weighed. It is then gently warmed, 
 and the stem dipped into the liquid in a small beaker. When about 
 o'25 gram has entered the bulb, the tip of the stem is sealed by 
 bringing it into a small gas-flame. The bulb is then weighed. 
 
 As in the case of the volatile solid, the absorption apparatus is 
 attached, and the combustion-tube is raised to a red heat to within 
 about 4 cms. of the copper coil c (Fig. 89) before the boat containing 
 the bulb is introduced. 
 
 A very slight scratch with a fine-cut file is made on the capillary 
 stem of the bulb, near to the closed end,* and the bulb is laid in 
 the boat in such a position that the fine point projects over the end 
 of the boat beyond the file-mark. In this way the liquid is pre- 
 vented from running up into the capillary tube. When all is in 
 readiness, the end of the fine tube is broken off, the fragment is 
 deposited in the boat (in order that any traces of liquid it contains 
 may be included), and the latter immediately pushed into the 
 combustion-tube, the stem of the bulb foremost. The plug is 
 quickly introduced, and the tube connected with the gas-drying 
 apparatus. The process is then continued as in the former example. 
 
 * The liquid should not be more in the capillary tube than can be avoided. 
 It may be driven back into the bulb so as to have a clear place near the end 
 where the stem is afterwards to be broken, by gently warming the tube with a 
 very small flam,;. 
 
Combustion of Volatile Liquids. 
 
 423 
 
 Liquids which are highly volatile are vaporised outside the 
 combustion-tube, and the vapour slowly passed over the heated 
 copper oxide. For this purpose a piece of moderately narrow glass 
 tube is drawn out into thick-walled capillaries in two places, so as 
 to leave a piece of the original tube about i\ cms. (i inch) long. 
 In order that the walls of the capillary tubes shall be thick, the 
 glass tube is heated in the blowpipe over as long a surface as the 
 flame will allow, until the walls of the tube have softened and 
 thickened as much as possible without falling together. The tube 
 is then slowly drawn out the desired 
 length, about 8 cms. (3 inches). It 
 is then bent in the manner shown 
 in Fig. 92. The tube is first 
 weighed, after which about 0*3 to 0-5 
 gram of the liquid is introduced, 
 and the fine drawn-out ends sealed 
 in a flame. The tube is then re- 
 weighed. The liquid must not be 
 allowed to get up into the hori- 
 zontal capillary tube. A cork, 
 previously selected to fit the com- 
 bustion-tube, and dried in an air- 
 oven, is fitted on to the horizontal 
 limb of the little tube, and the 
 apparatus is cooled by being immersed in ice-water, or in a freezing 
 mixture (depending upon the volatility of the liquid), so as to reduce 
 the tension of its vapour to an insignificant pressure. Thus, for 
 such a liquid as ether, ice-water may be used. 
 
 The weighed absorption apparatus having been attached, and 
 the combustion-tube heated to redness, the extreme end of the 
 capillary tube projecting through the cork is cut off, and the cork 
 inserted. The glass plug is not introduced in this case. 
 
 The temperature of the little tube containing the liquid is then 
 allowed gradually to rise. If it is in ice-water, this maybe done by 
 gently applying heat to the water by bringing a small flame under 
 the beaker, while, if it is in a freezing mixture, the gradual addition 
 of ordinary water will be sufficient to raise the temperature. 
 
 In this way the liquid is slowly vaporised, and the vapour 
 caused to pass over the heated copper oxide. When the liquid has 
 entirely volatilised, a fine scratch is made with a file close to the 
 sealed end of the apparatus, and the rubber tube from the gas- 
 purifying apparatus is slipped on. The end is then snapped off 
 
 FIG. 93. 
 
424 Ultimate Organic Analysis. 
 
 inside the rubber, and a gentle stream of air (afterwards followed 
 by oxygen if the organic compound is rich in carbon) is passed 
 through the tube until the combustion is complete, and the products 
 have been entirely expelled into the absorption apparatus. 
 
 II. Estimation of Nitrogen in Organic Compounds. 
 Nitrogen in organic substances may be estimated either by burning 
 the compound with copper oxide in a combustion-tube (very much 
 as in a carbon and hydrogen determination already described), and 
 collecting and measuring the nitrogen that is evolved ; or it may 
 be indirectly determined by converting it into ammonia, and esti- 
 mating the ammonia by one of the processes described in former 
 sections. The former process (Dumas) is applicable to all nitro- 
 genous organic compounds, while the latter cannot be employed in 
 the case of those substances in which the nitrogen is present in a 
 " nitro " or " azo " group. 
 
 i. Nitrogen by Dumas' Method. 
 
 Epitome of Process. The organic compound * is mixed 
 with copper oxide in a combustion-tube closed at one end, contain- 
 ing at the closed end a quantity of pure sodium bicarbonate. 
 (This furnishes a supply of carbon dioxide with which to sweep 
 out first the air, and finally the products of the combustion, from 
 the tube.) A roll of bright copper gauze is introduced at the open 
 end, to prevent the escape of any nitrogen in the form of oxides of 
 nitrogen. The tube is closed with a cork and delivery tube, and 
 the nitrogen is collected in a measuring-tube over caustic potash, 
 whereby the carbon dioxide is removed. 
 
 A piece of combustion-tube (cleaned in the manner described 
 on p. 414) is drawn off and sealed at one end before the blowpipe. 
 The tube should be about 76 cms. (30 inches f) long, or of such a 
 length that it can be heated to the extreme end when in the furnace. 
 
 A quantity of pure dry sodium bicarbonate % is introduced, so 
 that when the closed end is gently tapped on the table, the salt 
 shall occupy a space of about 20 cms. A layer of about 12 cms. of 
 powdered copper oxide is introduced, the tube being conveniently 
 supported by a clamp in a vertical position. About 0*5 gram of the 
 organic substance is then carefully introduced into the tube (the 
 
 * A suitable compound to employ for practice in the process is urea, 
 CO(NH 2 ) 2 . 
 
 f For rough measures, 2*5 cms. i inch. 
 
 j Sodium bicarbonate of commerce sometimes contains appreciable quanti- 
 ties of ammonia, when manufactured by the ammonia-soda process. If the 
 sample to be used is of unknown origin, a " blank" experiment should first be 
 made, in which sugar or some other pure organic compound containing no 
 nitrogen is used instead of urea. 
 
Nitrogen by Dumas' Method. 425 
 
 weighing-bottle containing the compound is first weighed, and after 
 a suitable quantity has been transferred to the tube, the bottle is 
 reweighed), and by means of a long stout clean wire, bent like a 
 corkscrew at the end, this is thoroughly stirred into the copper 
 oxide below, to a distance of about one-half the depth of the copper 
 oxide. A little more copper oxide (about 2\ cms.) is added (with- 
 out withdrawing the wire), and the top layers of the lower stratum 
 are mixed upwards into this. Once more a similar quantity of 
 copper oxide is added, and the corkscrew thoroughly freed from 
 any traces of the organic substance by being twisted up through 
 this. 
 
 The tube is then filled to within about 12 cms. of the top with 
 granular copper oxide, and lastly, a tolerably tight-fitting roll of 
 bright copper gauze, 7 cms. long, is introduced. The tube is then 
 removed from the clamp, and, while held in a horizontal position 
 by the thumb and forefinger of both hands, it is gently tapped upon 
 the table. In this way the contents are made to settle down, so as 
 to leave a narrow channel or air-space all along the top. The tube 
 is then laid in the furnace. 
 
 A well-fitting cork with a bent delivery tube is inserted, and 
 that portion of the tube containing the granular copper oxide is 
 heated to dull redness. 
 
 As soon as the tube is visibly red hot, the most forward portions 
 of the sodium carbonate are heated so as to sweep out the air from 
 the tube by means of the carbon dioxide so generated. Not more 
 than half the carbonate should be thus used at this stage. 
 
 As soon as the heating of the carbonate is commenced, the 
 delivery tube is attached, by means of a rubber connection, to the 
 lower branch tube of a Schiff's burette, as shown in Fig. 93. At 
 the bottom of the burette there is previously introduced a small 
 quantity of mercury, sufficient to reach about 12 mm. ( inch) 
 above the junction of the branch tube, and a strong solution of 
 caustic potash (i part of solid KHO in 2 parts of water) is poured 
 into the movable reservoir R. The tap t is then opened, and the 
 burette filled with potash solution by raising the reservoir, after 
 which the tap is again closed, and the reservoir placed in about the 
 position shown in the figure. So long as air is still being expelled 
 from the combustion-tube, gas will collect in the burette ; but as it 
 becomes swept out by the carbon dioxide, the ascending bubbles 
 get smaller and smaller, until at last they are practically entirely 
 absorbed by the potash. When this point is reached, the air which 
 has collected is expelled by again cautiously raising the reservoir 
 
426 
 
 Ultimate Organic Analysis. 
 
 and opening the tap, and the heating of the combustion-tube ex- 
 tended first to the roll of copper gauze, and then gradually along the 
 
 tube towards the mixture of 
 copper oxide and the organic 
 substance. 
 
 The nitrogen which is 
 evolved, together with the 
 carbon dioxide now filling the 
 air-space of the tube, pass up 
 into the burette, and, the car- 
 bon dioxide being absorbed 
 by the potash, the nitrogen 
 alone collects. When the 
 evolution of nitrogen is com- 
 pleted, the remainder of the 
 sodium bicarbonate is heated, 
 whereby a fresh supply of 
 carbon dioxide is generated, 
 which drives out the remain- 
 der of the nitrogen now filling 
 the tubes. This is continued 
 until the bubbles, which enter 
 the burette through the mer- 
 cury, are absorbed as before 
 by the potash. 
 
 In order to make sure that the carbon dioxide is completely 
 removed, the cup c is filled up with fresh potash solution, which is 
 slowly admitted into the burette by cautiously opening the tap. 
 When the operation is completed, the delivery tube is withdrawn 
 from the rubber connector, which is then closed with a pinch-cock.* 
 The reservoir is raised until the surface of the liquid it contains is 
 level with that in the burette, in which position the apparatus is 
 left for about a quarter of an hour, to ensure the perfect absorption 
 of any traces of carbon dioxide which may be present. The levels 
 are then exactly adjusted, and the volume of the nitrogen read off 
 upon the graduated tube.f 
 
 * In some forms of Schiff s burette, the branch tube is sufficiently wide to 
 admit of its being closed with a small cork ; but in this case it is less easy to 
 secure an air-tight connection with the delivery tube. 
 
 f Where a Schiff s burette is not at hand, the gas may be collected in a 
 graduated glass tube filled with strong potash solution, and inverted in a trough 
 of mercury. The point at which the air in the combustion-tube has been all 
 swept out by carbon dioxide, must in this case be ascertained by placing the 
 
 FIG. 93. 
 
Nitrogen. 427 
 
 The atmospheric temperature and pressure are noted, and the 
 volume of gas reduced to N.T.P. by the formula on p. 388. As the 
 tension of aqueous vapour exerted by such a strong solution of 
 potash as is here used is considerably less than that of water alone, 
 it is usual to give to p (see p. 388) the value of half the tension 
 of vapour of water, taken from the table in the Appendix. 
 
 Since i c.c. of nitrogen, under normal conditions, weighs 1*254 
 milligrams, the weight in milligrams of nitrogen contained in that 
 quantity of the organic compound employed for the analysis, is 
 obtained by multiplying the corrected volume (in cubic centimetres) 
 by 1-254. 
 
 Thus, suppose the corrected volume of nitrogen = 50 c.c., then 
 the weight of nitrogen would be 50 x 1*254 = 6270 milligrams 
 = 0*0627 gram. 
 
 2. Nitrogen by the Soda-lime Process. 
 
 This method is based upon the fact that many organic substances 
 containing nitrogen (see p. 424), when strongly heated with soda- 
 lime, give up their nitrogen in combination with hydrogen as 
 ammonia. By estimating the ammonia so evolved, the weight of 
 nitrogen can be determined. 
 
 Epitome of Process. The organic compound is mixed with 
 dry granular soda-lime in a combustion-tube closed at one end, and 
 containing a short layer of dry oxalic acid at its closed end (in order 
 to furnish a stream of hydrogen wherewith to sweep out the ammo- 
 nia at the end of the process). The evolved ammonia is absorbed 
 in dilute acid contained in a Will and Varrentrap's bulb-tube, which 
 is attached to the combustion-tube by means of a cork. The 
 ammonia may be estimated either gravimetrically , by precipitation 
 as the double ammonium platinic chloride, in which case the gas is 
 absorbed in dilute hydrochloric acid (i vol. acid to 4 vols. 
 water), and the estimation carried out as described on p. 251 ; or 
 it may be determined volumetrically by absorbing it in an excess 
 of standard sulphuric acid, and then titrating with standard sodium 
 hydroxide. The latter process is the more rapid. 
 
 end of the delivery tube beneath a separate tube (a test-tube, for example) 
 filled with potash and inverted in the same trough. 
 
 At the conclusion of the combustion, the measuring-tube containing the gas 
 is withdrawn from the mercury trough (by slipping a small dish underneath it) 
 and transferred to a tall cylinder of water. The mercury and the remaining 
 potash flow out, and are replaced by water, and after the whole has assumed a 
 uniform temperature, the tube is lowered into the water until the level of the 
 liquid inside and outside are the same, when the volume is read off. The 
 temperature of the water and the barometric pressure are noted. In correcting 
 the volume to N.T.P., the full value of the pressure due to aqueous vapour 
 must be taken (compare above). 
 
428 Ultimate Organic Analysis. 
 
 Into a clean dry combustion-tube, from 40 to 45 cms. long, is in- 
 troduced a quantity of oxalic acid (previously rendered anhydrous by 
 being heated in a steam-oven), sufficient to occupy about 5 cms. 
 Upon this is added about the same quantity of dry granular soda- 
 lime, which has been recently heated moderately strongly in a 
 porcelain dish. 
 
 A quantity of soda-lime (judged to be sufficient to occupy about 
 10 cms. of the tube) is powdered in a dry porcelain mortar, and by 
 means of a small clean spatula a weighed quantity (from 0-3 to 0-5 
 gram) of the organic compound (urea may be used as an exercise) 
 is thoroughly mixed with the soda-lime.* The mixture is then 
 transferred to a sheet of clean writing-paper, and carefully poured 
 into the tube by holding the paper in the palm of one hand and 
 bending it into a gutter. The mortar is then rinsed out with a 
 little more powdered soda-lime, which is similarly transferred to 
 the tube. The tube will now be rather more than half full. Granu- 
 lar soda-lime is added until it reaches to about 5 cms. of the mouth, 
 and a plug of asbestos (previously heated to redness) is inserted to 
 keep the materials in position. The tube is then tapped length- 
 ways upon the table (see p. 425), in order to create a free air-passage 
 along the top of the materials, and it is then laid in the furnace. 
 
 A measured volume of normal sulphuric acid (15 or 20 c.c., 
 depending upon the capacity of the bulbs (Fig. 94), which should 
 
 contain about the 
 volume of liquid re- 
 presented) is intro- 
 duced into the bulbs, 
 which are then con- 
 nected to the tube by 
 a tight - fitting cork. 
 The portion of the 
 tube extending back 
 from the asbestos plug, 
 FlG> 94- containing soda - lime 
 
 only, is first heated to a low redness ; after which the heating is 
 gradually extended along until the whole of the mixture of soda- 
 lime with the organic compound has become heated. As the 
 heat approaches tne end of the tube, a little care must be taken 
 that the oxalic acid is not decomposed before its time. The column 
 of pure soda-lime which separates the mixture should serve to 
 
 * Instead of mixing the materials in the mortar, the process may be carried 
 out within the tube exactly as in the case of the former method (p. 425). 
 
Nitrogen. 429 
 
 protect the acid, but the additional precaution may be taken of 
 inserting in the furnace a small screen, made of asbestos cardboard, 
 before beginning the operation. 
 
 As soon as the evolution of ammonia is complete, the heating 
 is extended to the oxalic acid, which is decomposed in the presence 
 of excess of alkali with evolution of hydrogen 
 
 Na 2 C 2 O 4 + 2NaHO = 2Na 2 CO 3 + H 2 
 
 In this way, the ammonia still filling the tube is driven out into 
 the acid in the bulbs.* The bulbs are then disconnected and the 
 contents transferred to a beaker, the bulbs being thoroughly rinsed 
 out with water. The excess of sulphuric acid present is titrated 
 with normal sodium hydroxide according to the method described 
 on p. 327. 
 
 Each cubic centimetre of normal acid which has been neutral- 
 ised by the ammonia, represents o'ooi7 gram of ammonia, or 
 0*0014 gram of nitrogen. 
 
 3. Nitrogen by Kjeldakl's Method. This process depends 
 upon the fact that many nitrogenous organic compounds, when 
 heated with strong sulphuric acid, have their nitrogen converted 
 into ammonia, which at once unites with the acid, forming am- 
 monium sulphate. From this salt the ammonia is afterwards 
 expelled by means of caustic soda, and estimated by one of the 
 usual methods. 
 
 Epitome of Process. The nitrogenous compound is digested 
 with strong sulphuric acid in a small round-bottom flask with a 
 long neck, potassium sulphate being added towards the end of 
 the operation in order to raise the boiling-point of : the liquid. The 
 mixture, after cooling, is diluted with water, and transferred to the 
 apparatus (Fig. 53, p. 250), where it is treated with an excess of 
 strong sodium hydroxide solution, and the ammonia determined 
 by the volumetric method given on p. 327. 
 
 From 0*5 to i 'o gram of the organic compound is weighed out 
 
 * In this process, and also in the Dumas' method previously described, the 
 combustion may be conducted in a tube open at both ends (as described in 
 the estimation of carbon and hydrogen, p. 415), and a stream of carbon 
 dioxide or of hydrogen, as the case may be, passed through from a reservoir 
 or generating apparatus outside. This modification of the processes, how- 
 ever, is not found to yield better results ; and as in both cases the compound 
 must be mixed with the contents of the tube, and cannot be introduced in a 
 boat, the tube has to be dismounted and recharged for each combustion. In 
 the carbon and hydrogen combustion, the open-tube method enables a series 
 of determinations to be made in rapid succession without dismantling the 
 apparatus, but in a nitrogen determination this is not the case, and it not 
 only involves the use of unnecessary apparatus, but introduces manipulative 
 difficulties in filling the combustion-tube. 
 
43 Ultimate Organic Analysis. 
 
 into the digestion-flask, and 20 c.c. of strong sulphuric acid added 
 to it.* The flask is then supported in an inclined position, and 
 heated by means of a small flame nearly to the boiling-point until 
 the mixture ceases to froth. The temperature is then raised, and 
 the mixture (usually dark-coloured or black) boiled for about 15 
 minutes, carefully guarding against frothing. Ten to 12 grams 
 of potassium sulphate are then added, and the mixture allowed to 
 continue gently boiling until it is nearly colourless ; after which 
 the flame is removed, and the liquid allowed to cool. A little water 
 is added, and the contents of the flask thoroughly rinsed out into 
 the distilling apparatus shown on p. 250. It is there treated with 
 caustic soda, and the ammonia distilled into an excess of standard 
 sulphuric acid as described on p. 327. 
 
 III. Estimation of Chlorine t in Organic Compounds. 
 
 Epitome of Process (Cartus* method)* k weighed quantity 
 of the substance, contained in a thin tube or bulb, is introduced 
 into a piece of stout combustion-tube, sealed at one end, along 
 with a few crystals of silver nitrate and a quantity of strong nitric 
 acid. The open end of the tube is then drawn out and sealed. 
 The thin tube containing the compound is then broken, and the 
 combustion-tube is placed in a hot-air bath, specially constructed 
 for the purpose, and heated to a temperature necessary to ensure 
 complete decomposition of the compound, for two or three hours. 
 In this way the carbon and hydrogen are converted into carbon 
 dioxide and water, while the liberated halogen combines with the 
 silver, present as nitrate, and forms silver chloride. The tube is 
 afterwards opened and the contents washed out, and the silver 
 chloride filtered, dried, and weighed in the usual way. 
 
 A piece of combustion-tube about 60 cms. long, not too wide in 
 the bore, and with tolerably thick walls, is sealed at one end before 
 the blowpipe, care being taken not to allow the end to be thinned 
 in the process. The organic compound is introduced into a thin 
 glass tube, the end of which has been slightly enlarged before 
 the blowpipe into a still thinner bulb (a. Fig. 95). This little 
 bulb-tube is first weighed ; about 0^3 to 0*5 gram of the substance is 
 introduced, and the tube drawn off and sealed by means of a fine 
 pointed blowpipe-flame at a'. The two portions a and b are then 
 together weighed. The difference is the weight of the substance. 
 
 * As the sulphuric acid of commerce is liable to contain ammonium 
 sulphate as an impurity, a blank experiment should be made with the sample 
 of acid which is to be used ; and the weight of nitrogen per cubic centimetre 
 of the acid which is represented by this impurity, must be deducted from the 
 result obtained in subsequent analyses. 
 
 f The process is equally applicable in the cases of bromine and iodine. 
 
Chlorine. 
 
 43* 
 
 FIG. 95. 
 
 The portion b is left upon the scale-pan, while the bulb is carefully 
 introduced into the combustion-tube, into which a few crystals (i 
 or 2 grams) of silver nitrate have previously been dropped. Fuming 
 nitric acid (free from 
 chlorine) is then added 
 in quantity sufficient 
 to occupy about 5 cms. 
 of the tube. The tube 
 is then held in an in- 
 clined position before 
 the blowpipe, and 
 drawn out into a thick 
 capillary. This is done 
 by heating the tube in 
 the flame, revolving it 
 slowly, until the soft- 
 ened walls thicken and 
 almost fall together, 
 assuming the condition shown in C, Fig. 95. The glass is then 
 gently pulled out, when the thick capillary shown at D will be 
 obtained ; the end is then sealed in the flame. When the tube has 
 become quite cold, the little bulb within is broken by shaking the 
 tube in such a way as to cause the bulb to strike against the walls, 
 but without allowing any of the contents to get into the capillary 
 end. The tube is then laid in the air-bath and gradually heated 
 up to a temperature of 1 50 to 200. 
 
 The air-bath consists of an oblong iron chamber or box, with 
 a number of iron tubes closed at one end passing through length- 
 ways ; the tubes are screwed into the ends of the box, and fixed 
 in a slightly inclined position, so that when the glass tubes are 
 pushed in, the liquid contents will be kept at one end. The whole 
 is supported on an iron stand, and heated by a row of small flames 
 placed beneath. A thermometer, passing through a small hole in 
 the top, registers the temperature. As great pressure is often 
 developed in the glass tubes, it is desirable to place an iron screen 
 or a brick opposite the end of the bath, so that, in the event of a 
 tube bursting, the glass shall be prevented from flying about. 
 After about two hours' heating, the gas is extinguished, and the 
 glass tube allowed to become quite cold before being removed. 
 
 It must never be forgotten that the tube is now a somewhat 
 dangerous object, and it must therefore be handled with caution. 
 Such a tube should never be left lying about, as they sometimes 
 
43 2 Ultimate Organic Analysis. 
 
 burst with great violence, without any apparent reason, even many 
 hours after they have become cold. 
 
 It is carefully withdrawn from the bath, and immediately en- 
 veloped in a cloth, with the end of the capillary tube only projecting, 
 and keeping it all the time in such a position that the liquid does 
 not run up to the narrow end. Should any liquid have got into 
 the end, it is carefully driven out by gently brushing a flame along 
 that part of the tube. 
 
 The tube is then opened by cautiously introducing the tip of 
 the capillary into a blowpipe-flame, so as to gradually soften the 
 glass, which is then blown out by the outrushing gases from the tube. 
 
 On no account whatever must the tube be opened by scratching 
 the capillary with a file and breaking off the point^ as this is 
 extremely likely to cause the tube to burst. 
 
 When the gas has escaped, the tube is cut at a point just below 
 the narrowed portion by first making a scratch with a sharp file, 
 and then touching the end of the scratch with the red-hot end of 
 a piece of glass or wire. The contents of the tube are then trans- 
 ferred to a beaker, the tube and the cut-off piece being thoroughly 
 rinsed into the beaker. 
 
 Besides the silver chloride, there will be the remains of the 
 glass bulb-tube. As the latter will probably have been broken at 
 the blown-out and thinner end, it is generally possible to pick out, 
 with a platinum wire, the greater part of the glass in a single large 
 piece. This is rinsed thoroughly, and carefully placed on a piece 
 of filter-paper in a steam-oven to dry. It is then conveyed to the 
 balance, and weighed along with the drawn-out end which has 
 been left there. The difference between this weight and the original 
 weight of the empty bulb-tube, will represent the weight of the 
 fragments of glass still remaining admixed with the silver chloride, 
 and which therefore must afterwards be deducted from the weight 
 finally obtained. 
 
 To the acid liquid sufficient sodium hydroxide (free from 
 chloride) is added to neutralise most of the acid, and the mixture 
 is then filtered, every particle of the broken glass, as well as the 
 silver chloride, being carefully brought on to the filter. The pre- 
 cipitate is washed, dried, and treated in the usual way. 
 
 From the weight of silver chloride (after deducting the weight 
 of the fragments of glass) the percentage of chlorine in the com- 
 pound is calculated. 
 
 I V. Estimation of Sulphur and Phosphorus in Organic 
 Compounds. The oxidation of the organic compound by means 
 
SulpJiur and Phosphorus. 433 
 
 of strong nitric acid in a sealed tube converts sulphur into sulphuric 
 acid, and phosphorus into phosphoric acid. In either case, there- 
 fore, the process is carried out as in the estimation of chlorine, 
 but without the silver nitrate. 
 
 In the case of sulphur, the contents of the tube are transferred 
 to a dish, and the nitric acid expelled by evaporating the mixture 
 nearly down to dry ness. Dilute hydrochloric acid is added, and 
 the sulphuric acid precipitated as barium sulphate in the usual 
 way (p. 258). When phosphorus is being estimated, the acid 
 liquid is neutralised with ammonia, and the phosphoric acid pre- 
 cipitated as ammonium magnesium phosphate (p. 264). 
 
 2 V 
 
SECTION II. 
 
 MISCELLANEOUS PHYSICO-CHEMICAL DETERMINATIONS. 
 
 SPECIFIC gravity of solids and liquids. 
 Boiling-point. 
 Melting-point. 
 Vapour density. 
 
 (i) Determination of the Specific Gravity of a Solid.* 
 
 (a) Solid Substances which are not acted upon by Water. 
 Principle. When a body is weighed suspended in a liquid, its 
 weight is diminished by the weight of 
 the liquid displaced. Hence, therefore, 
 if the body is weighed first in air, and 
 again when immersed in water, the dif- 
 ference between the two weights is the 
 weight of the water displaced by the 
 body ; that is, the volume of water which 
 is equal to the volume of the body. 
 
 EXAMPLE. A fragment of quartz or 
 flint or other mineral, about the size of a 
 walnut, is weighed in the ordinary way. 
 It is then suspended by means of a 
 thread of silk, or a horsehair, from the 
 beam of the balance, in such a manner 
 
 * The specific gravity of a solid or liquid 
 substance is the ratio between the weight of a 
 given volume of the substance and the weight 
 of the same volume of water at its point of 
 maximum density (taken as 4 C.). It is, 
 therefore, obtained by dividing the former 
 weight by the latter. As, however, it is more 
 difficult to maintain a uniform temperature at 
 4 than at the common atmospheric tempera- 
 ture, the operations are usually conducted at 
 FIG. 96. 1 5 '5. 
 
Specific Gravity. 435 
 
 that it will reach to about the middle of a small beaker supported 
 upon a wooden bridge which stands over the pan, as in Fig. 96, 
 leaving the latter free to move. The beaker is nearly filled with 
 water at I5'5, and any air-bubbles adhering to the solid substance 
 are carefully removed by means of a feather or small camel's-hair 
 brush.* In this position the substance is again weighed. t 
 
 If a the weight in air, and b weight in water 
 then a b the weight of water displaced 
 
 therefore . = specific gravity of the substance. 
 
 (b} Solid Substances which are acted upon by Water. When 
 the solid is dissolved, or otherwise acted upon by water, some other 
 liquid of known specific gravity is employed which exerts no action 
 upon the body. 
 
 If x = density of the solid as compared with the liquid, 
 and_y = specific gravity of the liquid 
 
 (sp. gr. water) 
 
 then i : y '. : x \ specific gravity of solid compared with water. 
 
 EXAMPLE. A crystal of copper sulphate (or alum, nitre, sodium 
 thiosulphate, etc.) is first weighed in the usual manner. It is then 
 suspended in a beaker containing benzene (C 6 H 6 , sp. gr. 0-885), an d 
 again weighed, care being taken, as in the former example, to 
 remove all air-bubbles. 
 
 If a = weight in air, and b = weight in benzene 
 
 then = density as compared with benzene = x 
 and x x y = specific gravity compared with water. 
 
 (t) Powdered Substances. When the substance is in a powdered 
 condition, the weight of water which it displaces is ascertained by 
 weighing it in a small bottle which is filled up with water. The 
 weight of the empty bottle is known, and the weight of water it 
 contains when full is also known. 
 
 * A porous substance, like charcoal, for instance, must be first placed in 
 water beneath an air-pump receiver and exhausted. 
 
 t If the substance is lighter than, water, a weight is attached to it in order to 
 sink it. The weight of this sinker \\hen. suspended in water is first ascertained. 
 Then, if a = weight of substance in air, 
 
 b weight of substance and sinker together in water, 
 c weight of sinker in water 
 
 then -_- --. . __ : - specific gravity of the substance 
 
436 Physico-chemical Determinations. 
 
 If ,z = weight of powder in air, 
 b weight of powder and water, 
 c weight of water 
 
 then - -7T ; = specific gravity of the powder 
 
 A small bottle with a perforated glass stopper (known as a 
 specific-gravity bottle) is exactly weighed in a perfectly clean and 
 dry condition. It is then filled with water at I5 0t 5, and the stopper 
 inserted. The small quantity of water which is thereby extruded 
 through the bored stopper is wiped off with a clean cloth, and the 
 bottle quickly weighed. (As sold, these bottles are marked with 
 the weight of water they are supposed to hold at this temperature.) 
 From this the weight of water it contains (c in the above formula) 
 is ascertained once for all. 
 
 Into the dry bottle a quantity of the powder is placed, and the 
 bottle weighed. From this the weight of powder in air (a) is found 
 by deducting the known weight of the empty bottle. The bottle 
 containing the powder is then filled with water at I5'5, being 
 gently tapped to disentangle any air-bubbles. The stopper is in- 
 serted, and after carefully wiping off the overflow the bottle is again 
 weighed. By subtracting the weight of the empty bottle, this gives 
 the weight of powder + water (b}. From these data the specific 
 gravity is calculated by the above formula. 
 
 (2) Determination of the Specific Gravity of a Liquid. 
 This may be ascertained by means of the specific-gravity bottle 
 described above. It is filled with the liquid in question at a 
 temperature of I5*5 and weighed. Deducting the weight of the 
 empty bottle gives the weight of the liquid. The weight of the 
 same volume of water at the same temperature being already known, 
 the specific gravity of the liquid is at once ascertained by dividing 
 the former by the latter. 
 
 Instead of the specific-gravity bottle, a thin glass U-tube, drawn 
 out at each end to a capillary, may be employed. The liquid is drawn 
 into the tube by dipping one extremity into the bottle containing it, 
 and applying a gentle suction to a piece of rubber tube attached to 
 the other end. The liquid is then brought to the temperature of 
 1 5'5 by immersing it in a beaker of water at this temperature, 
 after which it is carefully wiped and weighed. 
 
 (3) Determination of Boiling-point. The liquid is gently 
 heated in a tubulated distillmg-flask (a " Wurtz " flask) fitted with 
 a cork, through which is inserted a thermometer, the bulb of the 
 
Boiling-point. 437 
 
 latter being a little above the surface of the liquid. The flask 
 may be connected to a condenser in order to recover the portion of 
 the liquid which is vaporised during the process. A fragment of 
 scrumpled platinum foil placed in the flask causes the liquid to 
 boil more regularly. 
 
 The point at which the mercury in the thermometer remains 
 constant is noted, and where only moderate exactness is required 
 this temperature may be taken as the boiling-point. 
 
 In more exact determinations, either a correction must be made 
 for the cooling of the mercury in such portions of the column as 
 extend above the cork, or else the operation must be performed in 
 such an apparatus as to ensure the equal heating of the entire thread 
 of mercury. 
 
 In the former case a second thermometer is lashed to the first 
 outside the flask, in such a manner that the bulb lies midway 
 between the cork and the highest point to which the mercury 
 reaches, so as to record the average temperature of that portion of 
 the mercury thread which extends above the cork. The correction 
 which is added to the observed temperature is then made by means 
 of the following formula : 
 
 ;/(/ - /') x 0-000143 
 
 where n number of degrees over which the mercury is not exposed 
 
 to the heated vapour ; 
 / = observed temperature ; 
 t' = mean temperature of the outside mercury (given by 
 
 the second thermometer) ; and 
 0000143 = apparent coefficient of expansion of mercury in glass. 
 
 More exact determinations of boiling-points are made by means 
 of the apparatus shown in Fig. 97. The vapour from the boiling 
 liquid passes up the central tube, and then down through the 
 annular space between the inner and outer tubes before it finally 
 issues through the lateral branch tube. This side tube, which 
 extends for a considerable length, is connected to a condenser. 
 
 In this apparatus, not only is the whole of the mercury thread 
 of the thermometer exposed to the hot vapour, but the tube through 
 which it passes is itself surrounded by a jacket of the same vapour. 
 In this way cooling due to radiation is reduced to a minimum.* 
 
 * It will be obvious that when a boiling-point is to be determined with 
 sufficient exactness to necessitate the use of this apparatus, a thermometer 
 which has been standardised must be employed, and the instrument must be 
 one having a sufficiently open scale to render it possible to read small fractions 
 of degrees. 
 
438 
 
 Physico-chemical Determinations. 
 
 It must be remembered that the temperature at which a 
 liquid boils is influenced by pressure ; thus when we say that the 
 boiling-point of water is 100, it is understood that the liquid boils 
 at this temperature under normal atmospheric 
 pressure. When making a boiling-point 
 determination, therefore, the barometric pres- 
 sure at the time must be noted. If this is 
 greater or less than 760 mm., then either the 
 observed boiling-point is recorded as having 
 been taken under that particular pressure, or 
 a correction for pressure must be introduced. 
 There is, however, no formula of general 
 application by means of which this correction 
 can be made, for the reason that the extent 
 to which the boiling-point is influenced by 
 variations of pressure is different with different 
 liquids. If, therefore, the liquid under exami- 
 nation is an unknown substance, its exact 
 normal boiling-point (i.e. its boiling-point at 
 normal pressure) can only be ascertained by 
 making the determination at that actual 
 pressure. 
 
 The extent to which liquids differ from 
 each other with respect to the effect of pres- 
 sure upon their boiling-points is, however, 
 very slight, so that an approximate correction 
 may be introduced by regarding them as all 
 behaving as water does. 
 
 For each millimetre pressure (in the neighbourhood of 760 
 mm.) the boiling-point of water varies 0-0375 5 or ? * n other words, 
 the boiling-point is lowered o'i for a decrease of pressure amount- 
 ing to 2*68 mm., and raised by the same faction by a similar 
 increase of pressure. 
 
 (4) Determination of Melting-point. Provided a sub- 
 stance melts without undergoing decomposition, the temperature 
 at which it fuses when in a state of purity is constant, and is 
 termed its melting-point. This is determined by introducing the 
 solid substance into a thin glass tube drawn down to a capillary 
 end which is sealed. This little tube is lashed to the bulb of a 
 thermometer either by small rubber rings (cut from ordinary rubber 
 tube) or by wire, depending upon the particular liquid to be 
 employed in the bath in which it is to be heated. Thus if a bath 
 
Vapour -density Determination. 
 
 439 
 
 of strong sulphuric acid is used, platinum wire should be employed. 
 Fig. 98 shows the tube attached to the thermometer. The ther- 
 mometer is then supported in a clamp, and lowered into a tall 
 narrow beaker containing a liquid which may be heated 
 well above the melting-point of the substance, without 
 boiling. A small flame is placed beneath the beaker, 
 and the liquid is kept constantly mixed by means of 
 a stirrer, consisting of a glass rod bent in the form of 
 a ring at one end. This is moved up and down in 
 the beaker, and the moment the solid substance melts, 
 the temperature is read off. 
 
 The thermometer is then raised out of the bath, 
 which is allowed to cool somewhat, and the operation 
 repeated. With some solids, such as fats, etc., the 
 repetition of the determination must be made with a 
 fresh specimen, as each time they are melted they 
 undergo a slight decomposition or change which lowers 
 the melting-point. If the substance is translucent, 
 or is one in which from other causes it is difficult 
 to observe the exact fusion point, the capillary tube 
 may be open at the end, and the substance introduced 
 by dipping the tube into a small quantity of the melted 
 compound. The liquid rises by capillarity into the 
 tube, which is then removed and the substance allowed 
 to re-solidify. When the tube so charged is intro- 
 duced as before into the bath, the moment the sub- 
 stance melts it is driven up the tube by the entrance of 
 the liquid from the bath, and the temperature at which 
 this movement is observed is noted. As in the boiling- 
 point determination, the same correction must be made 
 for the portion of the mercury thread which projects 
 out of the bath. 
 
 (5) Determination of Vapour Density The density of 
 a vapour is the ratio between the weight of a given volume of that 
 vapour, and the weight of the same volume of another gas (under 
 the same conditions of temperature and pressure) which is taken 
 as the standard or unit. The standard usually adopted is hydrogen; 
 therefore the vapour density of a substance is in reality the specific 
 gravity of the vapour of that substance compared with hydrogen. 
 
 Vapour Density "by Victor Meyer's Method. Of the 
 various methods which have been devised for determining vapour 
 densities, that of Victor Meyer is the one most suitable for general 
 
 FIG. 98. 
 
440 
 
 Physico-chem ical Deter m ina tions. 
 
 laboratory practice, where a very high degree of accuracy is not 
 desired.* The apparatus consists of a bulb-tube, b (Fig. 99), with a 
 
 long, rather narrow stem, upon 
 which near the top there is a branch 
 tube of narrow bore. This bulb- 
 tube is placed in a wider tube, in 
 which a suitable liquid may be 
 boiled in order that the bulb-tube 
 may be heated by the vapour. A 
 weighed quantity of the liquid 
 whose vapour density is to be taken 
 is introduced into the hot bulb- 
 tube, which is then immediately 
 closed with a cork. The liquid in 
 vaporising displaces a volume of 
 air equal to the volume of the 
 vapour, and the air so expelled is 
 collected over water in the gra- 
 duated tube, under the ordinary 
 atmospheric conditions of tempera- 
 ture and pressure, which are noted. 
 The volume in cubic centimetres 
 so measured is reduced to N.T.P., 
 and the corrected number of cubic 
 centimetres when multiplied by 
 0-0000896 (the weight in grams of 
 i c.c. of hydrogen) gives the 'weight 
 of an equal volume of hydrogen. 
 
 For practice the vapour den- 
 sity of carbon disulphide may be 
 
 FIG. gg. 
 
 taken.f A little plug of glass wool or asbestos is pushed down to 
 the bottom of the bulb-tube , which is then placed in the outer 
 tube, into which a small quantity of water has been introduced. 
 The tube b must be perfectly dry inside. The top is closed with 
 a good cork, and the water boiled. As the temperature rises and 
 the air in b expands, the excess bubbles out through the water in 
 the small trough. When the temperature is constant, and no more 
 
 * For descriptions of other methods, as Dumas', Hofmann's, etc., the 
 student is referred to text-books on Physics. 
 
 f The specimen should be as pure as possible. It may be distilled, and 
 that portion of the distillate which passes over at 46 collected separately for 
 the determination. 
 
Vapotir- density Determination. 
 
 441 
 
 air is expelled, the graduated measuring-tube filled with water is 
 brought over the end of the bent delivery tube. The apparatus is 
 now ready for the introduction of the liquid.* This is contained 
 in a tiny stoppered bottle, or in a fine glass tube sealed at one end 
 and drawn to a capillary at the other (shown the actual size at a 
 and b, Fig. 100). The bottle or tube is weighed first empty, and 
 again when filled with the liquid (the tube is 
 filled by gently heating in a flame, and then 
 dipping the end into the liquid : and the 
 capillary end is left unsealed). About o'i 
 gram is a suitable weight of liquid to employ. 
 
 The cork is removed from the bulb-tube, 
 and the little bottle (or tube), held in a pair 
 of forceps, is dropped in, and the cork quickly 
 replaced. If the bottle is used, the stopper is 
 just loosened before dropping it in ; if the 
 tube is employed, it is dropped in capillary 
 end downwards. The pad of glass wool or 
 asbestos prevents the bulb from being frac- 
 tured by the falling bottle. The liquid imme- 
 diately vaporises and expels air from the 
 apparatus, which collects in the graduated tube. When no more 
 bubbles escape, the cork is removed. 
 
 The measuring-tube is then transferred to a large beaker filled 
 with water at the temperature of the room, the transference being 
 done by slipping a small porcelain dish under the tube as it stands 
 in the water-trough. The tube is depressed in the water until the 
 
 * Other liquids than water are used in the outer tube or bath, when a 
 higher temperature is required to vaporise the substance to be examined. It 
 is only necessary that the temperature of this bath should be considerably 
 higher than the boiling-point of the liquid to be vaporised, and that it should 
 remain constant throughout the operation. It is not necessary to know what 
 the actual temperature is. This fact often presents a difficulty to the young 
 student, which is expressed by the constantly repeated question, " If the 
 volume of gas or vapour is increased by rise of temperature, then surely the 
 hotter the tube is, the greater the volume of vapour which will be produced 
 when the liquid is introduced, and therefore a correspondingly greater volume 
 of air will be displaced from the apparatus ; how is it, therefore, that it is not 
 necessary to take the temperature of the bath ? " This difficulty vanishes when 
 the fact is taken into account that, whatever is the temperature of the bath 
 when the volatile liquid is introduced, the air within has already been heated to 
 that temperature, and proportionately expanded or rarefied. Consequently 
 at the higher temperature, although the volume occupied by the vapour is 
 greater, and therefore the volume of air expelled is greater, it is only more 
 rarefied air, which when cooled by collection over the water, is reduced to the 
 same volume as the lesser bulk of less expanded air which would be displaced 
 by the less expanded vapour at the lower temperature. 
 
442 Physico-chemical Determinations. 
 
 level of the liquid inside and outside is the same, when the volume 
 is read off. The temperature of the water is taken, and height of 
 the barometer is read. 
 
 The gas being saturated with aqueous vapour, the volume is 
 reduced to N.T.P. by the formula given on p. 388. 
 
 The corrected volume so obtained is the number of cubic 
 centimetres which the vapour of the carbon disulphide employed 
 would occupy if it could exist under N.T.P. 
 
 Multiplying this by 0-0000896 gives the weight in grams of an 
 equal volume of hydrogen at N.T.P. ; and by dividing the weight 
 of carbon disulphide employed by this weight of hydrogen, the 
 vapour density of carbon disulphide is obtained. 
 
 6. Determination of Molecular Weight by the "Freez- 
 ing-point Method." When pure water is cooled under ordinary 
 conditions, it freezes at o C. ; but if the water be rendered impure 
 by the presence of dissolved salts or other substances, it becomes 
 necessary to cool it below o before ice begins to form. Thus, 
 water containing i per cent, of sodium chloride in solution will 
 require to be cooled to o'6 before it begins to freeze. It is a 
 familiar fact that a lower degree of cold is necessary to freeze sea- 
 water than fresh. 
 
 When such aqueous solutions are cooled, the solid which begins 
 to separate out is pure ice, and not the substance which is in solu- 
 tion in the water, provided the solution is sufficiently dilute. 
 
 What is true of water in this respect is true also of other liquids 
 which are capable of being solidified. Thus, the freezing-point of 
 pure benzene is 6, but if a small quantity of benzoic acid (a solid) 
 or aniline (a liquid) be dissolved in the benzene, then it will be 
 found that this solution will require to be cooled below 6 before 
 the benzene begins to freeze ; and, just as in the case of water, it is 
 the solvent alone which freezes out of the solution. 
 
 The discovery that a quantitative relation existed between the 
 lowering of the freezing-point of a liquid and the amount of dissolved 
 substance present was made by Blagden as long ago as 1788, who 
 formulated the law that the depression of the freezing-point of 
 aqueous solutions of the same substance was proportional to the 
 amount of substance dissolved. 
 
 It was not, however, until the year 1871 that the important 
 observation was made by Coppet that, in the case of certain chemi- 
 cally allied substances, if quantities which \\Qreproporttonal to the 
 molecular weights were dissolved in equal volumes of the same 
 solvent, the freezing-point was depressed to the same extent ; or, 
 
Molecular Weight by Freezing-point Method. 443 
 
 in other words, the solutions would have the same solidifying- 
 point. 
 
 Upon these fundamental principles is based the process developed 
 by Raoult, and known as Raoult's method (or the cryoscopic method) 
 for determining molecular weights. 
 
 The extent to which the freezing-point of a liquid would be 
 depressed by dissolving in 100 grams of it a gramme-molecule (i.e. 
 a weight in grams equal to the molecular weight) of any substance, 
 is spoken of as the molecular lowering, or molecular depression, of 
 the freezing-point of that liquid ; and with certain qualifications, 
 this is found, for one and the same solvent, to approximate to a 
 constant.* Thus when water is the solvent, the molecular lowering 
 of the freezing-point is about 18*5. For acetic acid the constant is 
 about 39, while for benzene the value is about 49. 
 
 For example, a gramme-molecule of acetamide dissolved in 100 
 grams of water would depress the freezing-point 17*8, while the 
 same weight dissolved in 100 grams of acetic acid lowers the 
 freezing-point of this solvent 36' i. 
 
 This molecular lowering is calculated from experiments made 
 with more dilute solutions in the following manner. Suppose a 
 solution of cane-sugar (C 12 H 22 O n , mol. wt. = 342) in water, con- 
 taining 2*5 grams in 100 grams of water, is found by experiment to 
 freeze at o'i37. Then 
 
 2'5 : 342 :; 0*137 = 187 = molecular depression 
 or, molecular depression = 
 
 * This does not hold true in the case of substances whose molecules either 
 dissociate or undergo some molecular aggregation in the particular solvent. 
 Thus, in the case of water as solvent, it is found that such things as strong acids 
 and bases and their salts electrolytes, in other words produce a molecular 
 depression much greater than 18-5. This will be seen from the following 
 table (Raoult) : 
 
 Non-electrolytes. 
 Glycerol 
 Mannite 
 Sugar 
 Ethyl alcohol 
 Pyrogallol 
 Acetamide 
 
 In the case of non-electrolytes the phenomenon is concerned only with 
 the molecules of the substance dissolved in water, while with electrolytes a 
 certain proportion of the molecules (depending upon the degree of concentra- 
 tion) are believed to be dissociated into their ions. 
 
 Mol. 
 
 Mol. 
 
 depression. 
 
 Electrolytes. depression. 
 
 
 18-5 
 
 Sulphuric acid 
 
 38-2 
 
 
 i8'6 
 
 Hydrochloric acid 
 Nitric acid ... 
 
 35'8 
 
 
 16-6 
 
 Sodium hydroxide 
 
 36-3 
 
 
 18-0 
 
 Sodium chloride ... 
 
 35' 1 
 
 
 17 '7 i Potassium chloride 
 
 33'8 
 
444 PJiysico- chemical Determinations. 
 
 where m molecular weight of the dissolved substance, 
 
 / = observed depression of the freezing-point of the solvent, 
 
 and 
 
 g = grams of substance in 100 grams of the solvent. 
 Having once established the constant representing the molecular 
 lowering of freezing-point for any particular solvent, we are then in 
 possession of all the data for determining the molecular weight of 
 a substance which will dissolve in that liquid ; for, substituting C 
 for the constant in the above formula, we get 
 
 The determination may be made by means of the apparatus 
 shown in Fig. 101. This consists of a delicate thermometer 
 graduated into T o ths of degrees, which is fitted by means of a cork 
 into the mouth of a large test-tube, A. The same cork also carries 
 a short piece of narrow glass tube, C, through which the wire 
 stirrer passes. The tube A fits into a wider tube (a boiling-tube), 
 B, by means of a cork or a short piece of wide rubber tube. The 
 whole is supported by a clamp and retort stand, and can be lowered 
 into or raised out of the beaker or glass jar, F, containing the 
 freezing mixture.* 
 
 As a convenient exercise in the use of the apparatus, the 
 molecular weight of sugar may be determined by the lowering of 
 the freezing-point of water.f 
 
 Adjusting the Thermometer. The thermometer employed 
 differs from ordinary thermometers, in that it has an adjustable 
 zero. The scale covers a range of usually about 7 or 8 degrees, 
 and in the most recent forms the zero is at the upper end of the 
 scale. The total quantity of mercury in the thermometer is usually 
 greater than is required to bring the column to zero when the bulb 
 is placed in ice ; hence it is necessary to so get rid of some of the 
 mercury that the freezing-point of the solvent to be employed (in 
 the present example, water) shall fall somewhere near the top of 
 the scale not necessarily exactly at zero, but conveniently near it. 
 
 * The Beckmann apparatus usually figured in instrument makers' price 
 lists, contains a number of unnecessary appendages which give it an appear- 
 ance of complication, and considerably add to its cost. It is only necessary to 
 purchase the thermometer ; for the rest, two boiling-tubes and a beaker or 
 glass jar are all that is required. 
 
 f When the solvent employed is hygroscopic (glacial acetic acid, for 
 example), the glass tube c, through which the stirrer passes, should be replaced 
 by a \- tube ; and a slow stream of air, dried by means of calcium chloride or 
 sulphuric acid, allowed to pass in through the side branch and escape at the top. 
 In this way the air which is introduced by the movement of the stirrer is dry air. 
 
Molecular WeigJit by Freezing-point Method. 445 
 
 FIG. 101. 
 
446 Physico-chemical Determinations. 
 
 To accomplish this, the bulb is placed in water 3 or 4 degrees 
 warmer than the freezing-point of the solvent, that is to say, about 
 + 3, since water is to be the solvent.* The effect of this is to 
 send the mercury up until it begins to enter the wider part of the 
 bent stem E. By gently tapping the side of the instrument with 
 the finger-tips, a small piece of the mercury becomes detached, and 
 falls to the bottom of E, so that when the bulb is then placed in 
 ice, the top of the column of mercury will come to rest upon the 
 scale somewhere near the top graduations. 
 
 Into the tube A, perfectly clean and dry, a weighed quantity of 
 water (about 20 to 25 grams) is introduced either by weighing 
 the tube, first empty, and then with the water, or by weighing the 
 liquid out from a small flask. When the cork carrying the ther- 
 mometer and stirrer is inserted, the bulb should be entirely 
 immersed in the water. Tube A is then placed inside tube B, and 
 the whole system lowered into a freezing mixture of ice and salt in 
 the jar F. The object of the outer tube B is to render the cooling 
 more gradual and uniform than would be the case if it were directly 
 dipped into the freezing mixture.! As the water cools, it should be 
 kept gently moving by means of the stirrer, the handle of which 
 should be thrust into a small cork to prevent the warmth of the 
 hand from being conducted to the liquid. The temperature will 
 fall considerably below the freezing-point before solidification 
 actually sets in, but the moment the supercooled liquid begins to 
 freeze, the mercury in the thermometer will run up, and finally 
 come to rest. This point, which may be conveniently read by 
 means of a lens, is taken as the freezing-point of the water. The 
 observation should be repeated two or three times, by removing 
 tube A, and gently warming it with the hand until the ice is re- 
 melted, and then going through the freezing process again. 
 
 When the exact freezing-point of the solvent is established, a 
 weighed quantity of pure cane-sugar is introduced by lifting the 
 cork with the thermometer, and tipping the powdered substance 
 out of a weighing-bottle. \ When the substance has completely 
 
 * If benzene (M.P. 6) were to be used as the solver.t, then about 3 above 
 its freezing-point ; while if acetic acid (M. P. i6'5), then a temperature of about 
 20 would be suitable. 
 
 t It is sometimes convenient, and a saving of time, to cool the water in A 
 down to the freezing-point by first dipping the tube direct into the freezing 
 mixture. Then removing it, wiping it dry, and, when the contents have again 
 entirely remelted, placing it inside tube B, and then lowering the whole into 
 the mixture. 
 
 \ In the form of apparatus usually figured, tube A has a branch tube blown 
 into it near the top, which is intended tor the introduction of the substance ; 
 
Molecular Weight by Freezing-point Method. 447 
 
 dissolved (the process being aided by means of the stirrer), the 
 system is lowered into the freezing mixture, and the freezing-point 
 of the solution determined as before, by two or three repetitions of 
 the operation. 
 
 For example : 
 
 Weight of water taken 20' 1 15 grams 
 
 Freezing-point indicated by thermometer r 1 1 5 
 
 Weight of sugar introduced 1*020 gram 
 
 Freezing-point of solution - r 395 
 
 Percentage strength of solution = - - = 5*07 g 
 
 lowering of freezing-point = 1*395 ~ 1*115 = 0*28 / 
 molecular lowering for water = 18*5 = c 
 
 18*5 x 5-07 
 therefore molecular weight = * - - = 3 34' 9 
 
 O'2o 
 
 When a supercooled solution begins to freeze, and the temperature runs up 
 to a point at which the thermometer remains for a time practically stationary, 
 it will be seen that a certain quantity of the solvent has crystallised out. With 
 the same solution, the amount which so separates depends upon the extent to 
 which it was previously supercooled. This separation of a portion of the 
 solvent obviously renders the remaining liquid a little more concentrated than 
 the original solution, and from this it follows (i) that the observed temperature 
 is slightly Iffiaer than the true freezing-point of a solution having the original 
 strength ; and (2) that, unless the extent to which the liquid is supercooled is the 
 same each time the operation is repeated, the results obtained will not exactly 
 agree. To obviate this, it is necessary to make a correction for the amount of 
 solvent which thus separates. 
 
 Let t= number of degrees to which the solution is supercooled, 
 
 S = specific heat of the solution, 
 
 L = latent heat of the solvent, 
 
 M = mass of the solution, and 
 
 m = mass of solvent which separates out ; 
 
 Then S x t x M = L;/ 
 and therefore (the fraction which separates) = ^- 
 
 For practical purposes, the specific heat of the solution may bs taken as the 
 same as that of the solvent ; that is to say, the specific heat of the small 
 quantity of the dissolved substance may be neglected. Therefore, in the case of 
 
 S/ ' / 
 
 water, S = i and L = 80 ; hence ~j- becomes ---. That is to say, for each 
 
 LJ oo 
 
 degree Centigrade to which the solution is supercooled, uV of the solvent 
 separates out when the freezing-point is indicated, and the solution thereby 
 becomes more concentrated to the same extent. And since the depression of 
 the freezing-point is proportional to the concentration, the observed depression 
 will be too low bv the same fraction. 
 
 but, as a matter of fact, the substance is more conveniently introduced at the 
 top as here described, and with the greater certainty that the whole of it actually 
 falls into the liquid. 
 
44-S Physico-chemical Determinations. 
 
 Applying this correction in the example given above 
 
 The lowest temperature to which the mercury sank -- 2-695 
 observed freezing-point 1*395 
 
 therefore degrees of supercooling 1-3 
 Hence becomes ~ = o '01625 
 
 therefore o'28 x o '01625 = '00455 
 
 and o'28 '00455 = 0^27545 = corrected depression 
 
 Substituting this new value for / in the first calculation gives 
 
 * ^ = 34 J '3 = corrected molecular weight. 
 
 o '27545 
 
 7. Determination of Molecular Weight by the " Boil- 
 ing-point Method." Just as the freezing-point of a liquid is 
 lowered by the presence of dissolved substances, so the boiling- 
 point is raised from the same cause, provided the substance is 
 either non-volatile or exerts no appreciable vapour-pressure at the 
 boiling-point of the solution. 
 
 The elevation of the boiling-point of a liquid caused by the solu- 
 tion in loo grams of it of a gramme-molecule of such a substance 
 is termed the molecular elevation, or the molecular rise of the 
 boiling-point ; and (with the same reservations as to substances 
 which dissociate, as applied in the freezing-point method) for one 
 and the same liquid this value is practically a constant. 
 
 In the following table are given the boiling-points and the mole- 
 cular rise of boiling-point of a few common solvents : 
 
 Mol. rise of 
 Boiling-point. boiling-point. 
 
 Water 1000 5-2 
 
 Benzene 80*5 27*0 
 
 Chloroform 61-2 36-6 
 
 Carbon disulphide 46*2 237 
 
 Ether 34-9 21-1 
 
 The operation may be carried out by means of the " Beck- 
 mann" apparatus, shown in a dissected condition in Fig. 102. 
 
 A weighed quantity of the solvent is introduced into tube A 
 in the case of very volatile liquids, such as ether, by weighing the 
 tube with the ether, the tube being closed by corks. This tube is 
 similar to tube A in the freezing-point apparatus, except that here 
 the branch tube is necessary in order to attach a small reflux con- 
 denser, Ci, to return the liquid which would otherwise escape as 
 vapour during the boiling. Small glass beads, b, are placed in the 
 tube to a depth of about 20 mm. (a short inch), and the quantity of 
 liquid employed must be such that the bulb of the thermometer is 
 
Molecular Weight by Boiling-point Method. 449 
 
 well covered when the bottom of it nearly touches the beads. 
 Sometimes a short piece of platinum is fused through the bottom of 
 tube A, as seen at fi, with a view to aid in the heating. 
 
 The thermometer is similar to that employed in a freezing-point 
 determination, and must be previously adjusted, so that the boiling- 
 
 FlG. 102. 
 
 point of the liquid to be used may be indicated upon the lower 
 parts of the graduated scale. The zero in this case is at the bottom 
 of the scale. 
 
 Tube A is then placed within the jacket-tube B, which is open 
 at both ends, but having a "jacket" blown upon it, in which can 
 be boiled a small quantity of the same liquid as is in tube A. 
 
 2 G 
 
45O Physico-chemical Determinations. 
 
 This jacket is also connected to a reflux condenser, in order to 
 return the vaporised liquid. The object of this tube is to surround 
 tube A with an envelope which shall have a temperature as nearly 
 as possible the same as that within A itself. 
 
 The apparatus is heated upon the asbestos support D. This 
 somewhat resembles the lid of a box, provided with two asbestos 
 chimneys for carrying away the waste heat from the lamps. From 
 the centre hole, over which tube B is placed, there depends a short 
 wide asbestos tube, and the lamp (or lamps) are placed so that their 
 small flames are outside this tube, and therefore not directly below 
 the glass tubes. The heat must be so regulated that a steady and 
 brisk boiling of the liquid shall take place ; any irregularity or 
 "bumping "is quite fatal to the success of the operation. It is 
 desirable to still further shield the apparatus by a screen of asbes- 
 tos paper. When the mercury is stationary (after one or two 
 gentle taps upon the stem of the thermometer with the finger-tips), 
 the reading is taken using a pocket-lens after which the appa- 
 ratus is allowed to cool, and the operation repeated, if necessary, 
 once or twice. 
 
 When the exact boiling-point of the liquid, as registered by the 
 thermometer, has been ascertained, a weighed quantity of the sub- 
 stance whose molecular weight is to be determined is then intro- 
 duced, and after it has entirely dissolved, the boiling-point of the 
 liquid is again taken. The molecular weight is calculated from 
 the formula (similar to that for freezing-points) 
 
 R 
 
 where C = constant (or mol. elevation of boiling-point of solvent), 
 g grams of substance in 100 grams of solvent, and 
 R = observed rise of boiling-point. 
 
 The method is less capable of yielding accurate results than the 
 freezing-point method, partly owing to conditions of experiment, 
 and partly to variations of atmospheric pressure; but since the 
 object is usually merely to decide between two values, one of which 
 is a multiple of the other, a considerable experimental error is not 
 of serious consequence. 
 
 As an exercise for practice in the method, the molecular weight 
 of turpentine (C 10 H 16 = 136) may be determined, using ether as 
 the solvent.* 
 
 * Water is unfortunately not a convenient solvent, partly owing to disso- 
 ciation suffered by so many substances in solution in this liquid, and partly on 
 
Molecular Weight by Boiling-point Method, 451 
 
 EXAMPLE : 
 
 Weight of ether taken ... 18*45 grams 
 
 Boiling-point indicated by the thermometer ... 3*25 
 By means of a bent pipette, a weighed quan- 
 tity of turpentine was introduced at the 
 branch tube ... ... ... ... ... 0*8650 gm. 
 
 Boiling-point of the solution ... ... 3'94 
 
 O'86sO X TOO 
 
 Percentage strength of solution = ^ = 4*688 = g 
 
 rise of boiling-point = 3*94 - 3*35 = 0*69 = R 
 mol. rise of boiling-point for ether 21*1 = C 
 
 Hence mol. wt. of turpentine = 2ri x 4'688 = 143 .g 
 
 0^69 
 
 account of its very low molecular elevation of boiling-point. A volatile liquid 
 like ether, on the other hand, necessitates very careful condensation. The 
 most effective condensers are Cribb's double-surface condensers ; and the water 
 with which they are supplied should be cooled by being passed through a few- 
 turns of " compo " pipe placed in a bowl of ice. The temperature of the water 
 issuing from the condenser may thus easily be maintained at 5 or 6 C. 
 
APPENDIX 
 
 TABLE OF ATOMIC \VEIGHTS 
 
 Atomic weights. 
 
 Atomic weight? 
 
 Name. 
 
 Aluminium 
 Antimony 
 Arsenic ... 
 Barium ... 
 Beryllium 
 Bismuth ... 
 Boron 
 Bromine... 
 Cadmium 
 C&sium ... 
 Calcium ... 
 Carbon ... 
 Cerium ... 
 Chlorine... 
 Chromium 
 
 Cobalt 
 
 Copper (cuprum} 
 Didymium 
 Erbium ... 
 Fluorine... 
 Gallium ... 
 Germanium 
 Gold (aurttfft) 
 Hydrogen 
 Indium ... 
 Iodine 
 Iridium ... 
 Iron (ferrum} .. 
 Lanthanum 
 Lead (plumbum} 
 Lithium ... 
 Magnesium 
 Manganese 
 Mercury (hydrar- 
 gyrum 
 
 W 
 
 PI 
 
 3 
 
 a 2. 
 
 ^ V 
 
 Name. 
 
 1! 
 H 
 
 o 2J 
 
 < > 
 
 u w 
 
 o rt 2 
 
 a9| 
 
 Al 
 
 27 
 
 27-04 
 
 j\Iolybdenum 
 
 Mo 
 
 96 
 
 95 '9 
 
 Sb 
 
 120 
 
 119-6 
 
 Nickel 
 
 Ni 
 
 59 
 
 58-6 
 
 As 
 
 75 
 
 74 '9 
 
 Niobium 
 
 Nb 
 
 937 
 
 
 
 Ba 
 
 137 
 
 136-86 
 
 Nitrogen 
 
 N 
 
 14 
 
 14-01 
 
 Be 
 
 9 
 
 9-08 
 
 Osmium 
 
 Os 
 
 191 
 
 
 
 Bi 
 
 207-5 
 
 
 
 Oxygen ... 
 
 O 
 
 16 
 
 15-96 
 
 B 
 
 11 
 
 10-9 
 
 Palladium 
 
 Pd 
 
 106 
 
 106-2 
 
 Br 
 
 80 
 
 7976 
 
 Phosphorus 
 
 P 
 
 31 
 
 30-96 
 
 Cd 
 
 112 
 
 1117 
 
 Platinum 
 
 Pt 
 
 195 
 
 1 94 '3 
 
 Cs 
 Ca 
 
 133 
 40 
 
 132-7 
 39'9i 
 
 Potassium (ka- \ 
 Hum} ... ) 
 
 K 
 
 39 
 
 39 '03 
 
 C 
 
 12 
 
 11-97 
 
 Rhodium 
 
 Rh 
 
 101 
 
 104-1 
 
 Ce 
 
 141-2 
 
 
 Rubidhtm 
 
 Rb 
 
 85 
 
 85-2 
 
 Cl 
 
 355 
 
 35-37 
 
 Ruthenium 
 
 Ru 
 
 103-5 
 
 
 Cr 
 
 52 
 
 52-45 
 
 Samarium 
 
 Sm 
 
 150 
 
 
 
 Co 
 
 59 
 
 58-6 
 
 Scandium 
 
 Sc 
 
 44 
 
 43'97 
 
 Cu 
 
 63 
 
 63-18 
 
 Selenium 
 
 Se 
 
 79 
 
 78-87 
 
 Di 
 
 145 
 
 
 
 Silicon ... 
 
 Si 
 
 28 
 
 28-3 
 
 Er 
 
 166 
 
 
 
 Silver ... 
 
 Ag 
 
 108 
 
 107*66 
 
 F 
 
 19 
 
 19*06 
 
 Sodium ... 
 
 Na 
 
 23 
 
 22-995 
 
 Ga 
 
 70 
 
 6986 
 
 Strontium 
 
 Sr 
 
 87-3 
 
 
 Ge 
 
 72 
 
 
 
 Sulphur 
 
 s 
 
 32 
 
 31*98 
 
 Au 
 
 197 
 
 196-8 
 
 Tantalum 
 
 Ta 
 
 182 
 
 
 H 
 
 1 
 
 
 
 Tellurium 
 
 Te 
 
 125 
 
 
 
 In 
 
 113 
 
 1 13 '4 
 
 Thallium 
 
 Tl 
 
 203-7 
 
 
 
 I 
 
 127 
 
 126-54 
 
 Thorium 
 
 Th 
 
 232 
 
 
 
 jlr 
 
 192*5 
 
 
 Tin 
 
 Sn 
 
 118 
 
 H7'35 
 
 Fe 
 
 56 
 
 55-88 
 
 Titanium 
 
 Ti 
 
 48 
 
 
 La 
 
 138-5 
 
 
 Tungsten 
 
 W 
 
 184 _ 
 
 Pb 
 
 207 
 
 206-39 
 
 Uranium 
 
 U 
 
 239-8 
 
 
 Li 
 
 7 
 
 7-01 
 
 Vanadium 
 
 V 
 
 51-1 
 
 
 
 Mg 
 
 21 
 
 23 '94 
 
 Ytterbium 
 
 Yb 
 
 173 
 
 __ 
 
 Mn 
 
 55 
 
 54-8 
 
 Yttrium 
 
 Y 
 
 896 
 
 
 
 Hg 
 
 200 
 
 199-8 
 
 Zinc 
 Zirconium 
 
 Zn 
 Zr 
 
 65 
 904 
 
 64-88 
 
454 Reagents. 
 
 REAGENTS. 
 
 It is important that the reagents used in qualitative analysis should 
 be made up of known definite strength, and that the strength should be 
 indicated upon the bottles.* It is a great additional advantage, also, that 
 the reagents should, as far as possible, have either a uniform strength 
 (that is to say, equal volumes should possess equal chemical values, or be 
 chemically equivalent), or, when this is not possible or desirable, the 
 strength should bear some simple ratio to the equivalent strength. 
 
 By this is not meant that the reagents are to be prepared with an 
 exactness in any way comparable to that required in making up the 
 standard solutions used in volumetric analysis. For qualitative analysis, 
 it is enough that the reagents bear a rough approximation to the exact 
 strength indicated. 
 
 Unless some of the reagents were to be inconveniently dilute, it would 
 not be possible to have them all chemically equivalent, owing to differences 
 of solubility. For example, a solution of oxalic acid, which is saturated 
 at 10, is only equivalent to sulphuric acid consisting of I volume of strong 
 acid in about 18 volumes of water. 
 
 In volumetric analysis, standard solutions are designated normal 
 solutions when they are of such a strength that i litre contains a weight of 
 the reagent in grams, equal to the chemical equivalent of that reagent. 
 Thus, normal sulphuric acid contains 49 grams of H 3 SO 4 per litre ; normal 
 hydrochloric acid 36*5 grams of HC1 per litre, and so on (see p. 313). 
 The signature employed to denote ftiese normal or equivalent solutions 
 is the letter N. 
 
 For qualitative analysis a convenient strength for the common acids and 
 alkalies and general reagents is five times this normal strength. These 
 are distinguished by the signature 5\. 
 
 A large number of other reagents, such as are used for special tests, are 
 conveniently prepared of normal strength, and should bear the signature N ; 
 while others, which for special reasons (such as slight solubility, e.g. calcium 
 sulphate ; or the delicacy of their reactions, e.g. ammonium thiocyanate} 
 should be more dilute, are prepared of suitable strengths, given in the 
 following list of reagents, and the signature of the strength in all cases 
 should be indicated upon the labels. 
 
 Thus, in the case of calcium sulphate, a saturated solution is employed ; 
 this contains one-thirtieth of the quantity of calcium sulphate required to pro- 
 
 N 
 
 duce a normal or equivalent solution, and its strength is denoted by . 
 
 3 
 
 * It is greatly to be regretted that in many laboratories the practice still 
 exists of providing the student with reagents concerning whose strength he is 
 left absolutely in the dark. The effect of this upon the analytical work is only 
 too familiar to those who have had much practical experience in teaching. 
 
Reagents of Standard Strength. 455 
 
 Again, ammonium thiocyanate being so delicate a reagent for the 
 detection of ferric salts, a solution of one-fifth of the equivalent strength 
 is conveniently strong. This is indicated upon the label by the 
 
 N* 
 signature . 
 
 Reagents of Five Times Normal Strength. Signature, 5N. 
 
 Sulphuric acid, H.,SO 4 . Equivalent = 49. Ordinary concentrated 
 acid, sp. gr. I '84, has a strength 36 times normal, and may be designated 
 36N. One volume of strong acid diluted with 6 volumes of water = 5N. 
 
 Nitric acid, HNO 3 .| Equivalent = 63. Concentrated acid, sp. gr. 
 i '42 = i6N. Five volumes of strong acid diluted with 1 1 volumes of 
 water = 5N. 
 
 Hydrochloric acid, HCl.f Equivalent = 36*5. Concentrated acid, 
 sp. gr. i'i6 = loN. One volume of strong acid diluted with I volume of 
 water = 5N. 
 
 Acetic acid, C 2 H 4 O 2 . Equivalent = 60. Glacial acid, M.P. 10 C., 
 = lyN. One volume of glacial acid diluted with 2\ volumes of 
 water = 5N. 
 
 Potassium hydroxide, KHO. Equivalent = 56. 
 
 280 grams dissolved in water and diluted to one litre = 5N. 
 
 Sodium hydroxide, NaHO. Equivalent = 40. 
 
 200 grams, dissolved in water and diluted to one litre = 5N. 
 
 Ammonia (solution in water, NH 4 HO). Equivalent = 35. 
 
 The strong solution, sp. gr. O'88o = 2oN. One volume of the strong 
 solution diluted with 3 volumes of water = 5N. 
 
 Ammonium sulphide, (NH 4 ) a S. Equivalent = 34. 
 
 600 c.c. of 5N ammonia are saturated with sulphuretted hydrogen ; 
 this gives hydrogen ammonium sulphide H(NH 4 )S. This is made up to 
 I litre by adding 5N ammonia. (This reagent is slowly decomposed by 
 atmospheric oxygen, ammonia is evolved, and yellow ammonium sulphide 
 is formed (NH 4 ) 2 S 2 .) 
 
 Sodium sulphide, Na 2 S. Equivalent = 39. 
 
 200 grams of sodium hydroxide are dissolved in 800 c.c. of water. 
 
 400 c.c. of this solution are saturated with sulphuretted hydrogen, and 
 the remaining half added, together with water sufficient to make the 
 
 * The system of standard reagents for use in qualitative analysis, upon 
 which the instructions here given are based, was devised by Reddrop, and is 
 published in full detail in the Chemical News, May, 1890. Mr. Reddrop 
 employs the signature E (equivalent] instead of N (normal] in his system ; but 
 as the designation normal is the term almost universally employed for such 
 solutions when used for volumetric analysis, it appears more harmonious, and 
 brings qualitative analysis more into line with quantitative methods, to adopt a 
 uniform nomenclature throughout. 
 
 f Aqua regia is prepared as required by mixing i volume of i6N nitric acid 
 with 3 volumes of loN hydrochloric acid. 
 
456 Reagents of Standard Strength. 
 
 volume up to I litre. When hydrogen sodium sulphide, HNaS, is 
 required, the sodium hydroxide is simply saturated with sulphuretted 
 hydrogen, without the further addition of sodium hydroxide. 
 
 Ammonium chloride, NH 4 C1. Equivalent = 53*5. 
 
 267*5 grams of the salt dissolved in water and diluted to one litre = 5X. 
 
 Ammonium carbonate,* (NH 4 ) 2 CO 3 . Equivalent = 48. 
 
 200 grams of ammonium sesquicarbonate (commercial carbonate) dis- 
 solved in 350 c.c. of 5N ammonia, and diluted with water to one litre 
 = S N. 
 
 Ammonium acetate, (NH 4 )C 2 H 3 O 2 . Equivalent = 77. 
 
 300 c.c. of glacial acetic acid (I7N) neutralised with strong ammonia 
 and diluted to one litre = $N. 
 
 Reagents of Normal Strength. Signature N. 
 
 The following normal reagents are prepared by dissolving the equi- 
 valent weight in grams of the various salts in water, and diluting to one 
 litre. In all cases the nearest whole number to the equivalent weight 
 may be taken as sufficiently exact. 
 
 Ammonium sulphate, (NH4) 2 S'V Equivalent weight ... 66 -o 
 
 Barium chloride, BaCl 2 ,2H 2 O. ,, ,, ... 122-0 
 
 Calcium chloride, CaCl 2 ,6H 2 O. ,, ,, ... 109/5 
 
 Copper sulphate, CuSO 4 ,5H 2 O. ,, ,, .. 12475 
 
 Ferric chloride, FeCl 3 . ,, ,, ... 54-17 
 
 Ferrous sulphate, FeSO 4 ,7H 2 O. ,, ,, ... 139-0 
 
 Hydrogen disodium phosphate, HNa 2 PO 4 , 1 2H 2 O. ,, ,, ... 119*3 
 Lead acetate Pb(C 2 H 3 O 2 ) 2 , 3 H 2 O. ...189-5 
 
 Lead nitrate, Pb(NO 3 ) 2 . ,, ,, ... 165-5 
 
 Magnesium chloride, MgCl 2 ,6H 2 O. ,, ... 101-5 
 
 Magnesium sulphate, MgSO 4 , 7 H 2 O. ,, ,, ... 123-0 
 "Magnesia mixture" (MgCl 2 , NH 4 C1, and NH 4 I1O). 
 
 68 grams of MgCl 2 ,6H 2 O, together with 165 grams of 
 NH 4 C1, are dissolved in 300 c.c. of water ; 300 c.c. of 
 5N ammonia are added, and the solution diluted with water 
 up to one litre. 
 
 Potassium chromate, K 2 CrO 4 . Equivalent weight ... 97-25 
 
 Potassium cyanide, KCy. ,, ,, ... 65^0 
 
 Potassium ferrocyanide, K 4 FeCy a , 3 1 IoO. ,, ,, ... 105-5 
 
 Potassium ferricyanide, K 3 FeCy 6 . ,, ., ... 109-7 
 
 Potassium sulphate, K 2 SO 4 . ,, ,, ... 87-0 
 
 Sodium acetate, NaC 2 H 3 O 2 ,3H 2 O. ,, ,, ... 136*0 
 
 * Hydrogen ammonium carbonate, H(NH 4 )CO 3 , cannot be prepared 
 stronger than 3N (see below). 
 
Reagents. 457 
 
 Sodium acetate and acetic acid. 
 
 136 grams of the salt dissolved in 800 c.c. of water and 200 
 
 c.c. of glacial acetic acid added. 
 
 Sodium carbonate,* Na 2 CO 3 ,ioH 2 O. Equivalent weight ... 143-0 
 
 Sodium thiosulphate, Na 2 S 2 O 3 ,5H 2 O. Equivalent weight ... 124-0 
 
 Stannous chloride, SnCl 2 ,2H 2 O. ,, ,, ... 112-5 
 
 112 grams of the salt are dissolved in 200 c.c. of 5N HCI, 
 and the solution diluted with water to one litre. Fragments 
 of granulated tin should be placed in the solution. 
 
 Reagents of Various Strengths. 
 
 Ammonium oxalate, (NH 4 ) 2 C 2 O 4 ,2H 2 O. Equivalent weight ... So-o 
 
 N 
 40 grams of the salt in one litre of water = solution. 
 
 Ammonium bicarbonate, H(NH 4 )CO 3 . 
 
 Saturation solution = 3N solution. Prepared by passing 
 
 excess of CO 2 into 3^ ammonia. 
 Ammonium thiocyanate, NH 4 CyS. Equivalent weight ... 76-0 
 
 N 
 15 grams of the salt in one litre of water = -- solution. 
 
 Barium hydroxide, Ba(HO) 2 ,8H 2 O. Equivalent weight ... 157-5 
 
 31 grams in one litre of water = solution. 
 Barium nitrate, Ba(NO 3 ) 2 . Equivalent weight ... 130*5 
 
 65 grams dissolved in one litre of water' = solution. 
 
 Bromine water, Br. Equivalent weight ... 80*0 
 
 Saturated solution, obtained by shaking up an excess of 
 
 N 
 bromine with water. Gives a solution of strength. 
 
 Calcium hydroxide (lime-water) Ca(HO) 2 . Equivalent weight ... 37'O 
 A saturated solution, obtained by shaking up an excess of 
 lime with water and allowing the mixture to settle, 
 
 = about solution. 
 20 
 
 Calcium sulphate, CaSO 4 . Equivalent weight ... 66'O 
 
 A saturated solution, obtained by shaking up an excess of 
 
 N 
 calcium sulphate with water, = solution. 
 
 Chlorine water, Cl. Equivalent weight ... 35-5 
 
 A saturated solution = . 
 
 * A strong solution of sodium carbonate, 3N strength, should also be 
 prepared by dissolving 429 grams of the salt in 'the litre. This is practically 
 a saturated solution at ordinary temperatures. 
 
458 Reagents. 
 
 Mercuric chloride, HgCL. Equivalent weight ... 135-5 
 
 27 grams dissolved in one litre of water = solution. 
 
 Mercurous nitrate, Hg 2 (NO 3 ) 2 ,2H 2 O. Equivalent weight ... 280-0 
 
 56 grams dissolved in 40 c.c. of 5N HNO 3 and diluted to one 
 
 N 
 litre = solution. A small quantity of mercury should be 
 
 placed in the bottle. 
 Potassium iodide, KI. Equivalent weight ... 1660 
 
 N 
 33 grams dissolved in one litre of water = solution. 
 
 Potassium metantimoniate, KSbO 3 ,3H 2 O. Equivalent weight ... 291*0 
 A saturated solution, obtained by shaking up an excess of the 
 
 N 
 salt with water, = ^x 
 
 Silver nitrate, AgNO s . Equivalent weight ... 170-0 
 
 N* 
 3*4 grams dissolved in ioo c.c. = solution. 
 
 Fusion Mixture. Pure anhydrous sodium and potassium carbonates 
 are intimately mixed together in the proportion of their equivalent 
 weights. Sodium carbonate 53 parts, and potassium carbonate 69 parts. 
 
Appendix, 
 
 459 
 
 TABLE OF HARDNESS. 
 Degrees of hardness expressed as parts of CaCO 3 per 100,000. 
 
 C.C of 
 
 CaCO 3 
 
 C.C. of 
 
 CaC0 3 
 
 C.C. of CaCOa 
 
 C.C. of 
 
 CaCO 3 
 
 soap 
 solution. 
 
 per 100,000 
 parts. 
 
 soap 
 solution. 
 
 per 100,000 
 parts. 
 
 soap 
 solution. 
 
 per 100,000 
 parts. 
 
 soap 
 solution. 
 
 per 
 100,000 
 parts. 
 
 I'O 
 
 48 
 
 5' 
 
 6 -oo 
 
 9 o I i i 'So 
 
 I3-0 
 
 1 8 -02 
 
 I 
 
 63 
 
 'I 
 
 14 
 
 i -95 
 
 I 
 
 17 
 
 2 
 
 79 
 
 '2 
 
 29 
 
 '2 
 
 I2'II 
 
 '2 
 
 '33 
 
 '3 
 
 '95 
 
 '3 
 
 "43 
 
 3 
 
 26 
 
 '3 
 
 '49 
 
 "4 
 
 I'll 
 
 '4 
 
 '57 
 
 '4 
 
 41 
 
 '4 
 
 65 
 
 6 
 
 27 
 '43 
 
 6 
 
 
 
 6 
 
 56 
 71 
 
 6 
 
 81 
 '97 
 
 7 
 
 56 
 
 7 
 
 7-00 
 
 7 
 
 86 
 
 7 
 
 19-13 
 
 8 
 
 69 
 
 8 
 
 14 
 
 8 
 
 13-01 
 
 8 
 
 29 
 
 "9 
 
 82 
 
 '9 
 
 29 
 
 '9 
 
 16 
 
 '9 
 
 '44 
 
 2'0 
 
 '95 
 
 6-0 
 
 '43 
 
 IO'O 
 
 3i 
 
 14-0 
 
 60 
 
 I 
 
 2-08 
 
 i 
 
 '57 
 
 i 
 
 46 
 
 i 
 
 76 
 
 "2 
 
 21 
 
 '2 
 
 71 
 
 '2 
 
 61 
 
 '2 
 
 92 
 
 '3 
 
 '34 
 
 '3 
 
 86 
 
 '3 
 
 76 
 
 '3 
 
 20-08 
 
 "4 
 
 '47 
 
 '4 
 
 8-00 
 
 '4 
 
 91 
 
 '4 
 
 24 
 
 3 
 
 60 
 
 73 
 
 
 H 
 29 
 
 r 
 
 6 
 
 14-06 
 
 21 
 
 6 
 
 40 
 56 
 
 7 
 
 86 
 
 7 
 
 '43 
 
 7 
 
 '37 
 
 7 
 
 71 
 
 8 
 
 '99 
 
 8 
 
 '57 
 
 8 
 
 52 
 
 8 
 
 87 
 
 '9 
 
 3-12 
 
 '9 
 
 7i 
 
 '9 
 
 68 
 
 '9 
 
 21-03 
 
 3-0 
 
 2 5 
 
 7-0 
 
 86 
 
 I IX) 
 
 84 
 
 15-0 
 
 19 
 
 i 
 
 38 
 
 i 
 
 9-00 
 
 i 
 
 15-00 
 
 i 
 
 '35 
 
 2 
 3 
 
 i; 
 
 2 
 
 '3 
 
 '14 
 
 29 
 
 '2 
 '3 
 
 16 
 32 
 
 "2 
 '3 
 
 :M 
 
 '4 
 
 77 
 
 4 
 
 '43 
 
 '4 
 
 48 
 
 '4 
 
 85 
 
 I 
 
 90 
 4'03 
 
 3 
 
 57 
 
 7i 
 
 "4 
 
 63 
 
 79 
 
 3 
 
 22*02 
 
 18 
 
 7 
 
 16 
 
 7 
 
 86 
 
 7 
 
 '95 
 
 7 
 
 '35 
 
 8 
 
 29 
 
 8 
 
 lO'OO 
 
 8 
 
 16-11 
 
 8 
 
 S 2 
 
 9 
 
 '43 
 
 '9 
 
 15 
 
 '9 
 
 27 
 
 '9 
 
 6 9 
 
 4-0 
 
 57 
 
 8-0 
 
 30 
 
 12*0 
 
 '43 
 
 16-0 
 
 86 
 
 i 
 
 71 
 
 i 
 
 "45 
 
 'I 
 
 "59 
 
 
 
 2 
 
 86 
 
 '2 
 
 60 
 
 2 
 
 75 
 
 
 
 '3 
 
 5-00 
 
 '3 
 
 75 
 
 '3 
 
 90 
 
 
 
 '4 
 
 H 
 
 '4 
 
 90 
 
 "4 
 
 17-06 
 
 
 
 1 
 
 29 
 '43 
 
 1 
 
 11-05 
 20 
 
 3 
 
 22 
 
 38 
 
 
 
 7 
 
 '57 
 
 7 
 
 35 
 
 7 
 
 54 
 
 
 
 8 
 
 71 
 
 8 
 
 So 
 
 8 
 
 70 
 
 
 
 '9 
 
 86 
 
 '9 
 
 65 
 
 '9 
 
 86 
 
 
 
46o 
 
 Appendix. 
 
 EXPANSION OF WATER BETWEEN o AND 25 C. 
 
 Temperature 
 
 Volume. 
 (Vol. at 4 = i.) 
 
 Density. 
 
 (Density at 4 - i.) 
 
 u 
 
 o 
 
 I-OOOI2 
 
 0-99988 
 
 2 
 
 1-00003 
 
 0-99997 
 
 ^T 
 
 I "OOOOO 
 
 I '00000 
 
 6 
 
 I -00003 
 
 0-99997 
 
 8 
 
 I -0001 1 
 
 0-99989 
 
 10 
 
 1-00025 
 
 Q'99975 
 
 12 
 
 1-00044 
 
 0-99956 
 
 13 
 
 1-00055 
 
 Q'99945 
 
 14 
 
 1-00068 
 
 0-99932 
 
 15 
 
 1-00082 
 
 0-99918 
 
 I5'5 
 
 1-00089 
 
 0-99911 
 
 16 
 
 1-00097 
 
 0-99903 
 
 17 
 
 1-00113 
 
 0-99887 
 
 18 
 
 1-00131 
 
 0-99869 
 
 19 
 
 1-00149 
 
 0-99851 
 
 20 
 
 1-00169 
 
 0-99831 
 
 22 
 
 I -0021 2 
 
 0-99787 
 
 2 4 
 
 I-00259 
 
 0-99742 
 
 25 
 
 I 00284 
 
 0-99717 
 
 The coefficients of expansion of water (the values in the middle column) 
 between the temperatures o and 25 may be calculated from the 
 formula 
 
 V = V 0-000061045^ + 0*000007 7 1 83/ 2 0-00000003 7 34/ 3 
 
 where V = the volume at o. In the above table V = roooi2 ; but if 
 the coefficients are required when the volume at o is taken as unity, 
 then V = i. 
 
 For temperatures above 25 a different formula is required. 
 
Appendix. 
 
 461 
 
 TENSION OF AQUEOUS VAPOUR IN MILLIMETRES OF MERCURY. 
 For each fifth of a degree from 5 to 25 C. 
 
 o 
 
 mm. 
 
 
 
 mm. 
 
 o 
 
 mm. 
 
 o 
 
 mm. 
 
 5' 
 
 6-5 
 
 I0'0 
 
 9'2 
 
 15-0 
 
 I2'7 
 
 20'0 
 
 I7-4 
 
 2 
 
 6-6 
 
 2 
 
 9'3 
 
 '2 
 
 I2'9 
 
 '2 
 
 17-6 
 
 '4 
 
 67 
 
 "4 
 
 9'4 
 
 '4 
 
 I 3 -0 
 
 '4 I7'8 
 
 6 
 
 6-8 
 
 6 
 
 9'5 
 
 6 
 
 I3-2 
 
 6 18-0 
 
 8 
 
 69 
 
 8 
 
 97 
 
 8 
 
 I3'4 
 
 8 
 
 18'3 
 
 6-0 
 
 7-0 
 
 iro 
 
 0-8 
 
 i6'o 
 
 I3-5 
 
 2TO 
 
 18-5 
 
 '2 
 
 7*1 
 
 2 9-9 
 
 "2 
 
 r 3 7 
 
 '2 
 
 187 
 
 4 
 
 7-2 
 
 4 io*i 
 
 '4 
 
 I3-9 
 
 '4 
 
 19-0 
 
 6 
 
 7'3 
 
 6 10-2 
 
 6 
 
 14-1 
 
 6 
 
 19-2 
 
 8 
 
 7'4 
 
 8 i 10-3 
 
 8 
 
 14-2 
 
 8 1 19-4 
 
 7-0 
 
 rs 
 
 I2'0 I0'5 
 
 17-0 
 
 14-4 
 
 22'0 197 
 
 2 
 
 7-6 
 
 '2 
 
 io p 6 
 
 '2 
 
 14*6 
 
 '2 
 
 19-9 
 
 '4 
 
 77 
 
 '4 
 
 107 
 
 '4 
 
 14-8 
 
 4 
 
 20' I 
 
 6 
 
 7'8 
 
 6 
 
 10*9 
 
 6 
 
 15-0 
 
 6 
 
 20'4 
 
 8 
 
 7 9 
 
 8 
 
 iro 
 
 8 
 
 15-2 
 
 8 
 
 20'6 
 
 8-0 
 
 8-0 
 
 13-0 
 
 1 1 '2 
 
 18-0 
 
 iS'4 
 
 23-0 
 
 20*9 
 
 2 
 
 8-1 
 
 '2 
 
 11-3 
 
 '2 
 
 15-6 
 
 2 
 
 2TI 
 
 '4 
 
 8-2 
 
 '4 
 
 "*5 
 
 '4 
 
 157 
 
 '4 
 
 21'4 
 
 6 
 
 8'3 
 
 6 
 
 it -6 
 
 6 
 
 I5-9 
 
 6 
 
 217 
 
 8 
 
 8-5 
 
 8 
 
 11-8 
 
 8 
 
 16-1 
 
 8 
 
 21-9 
 
 9-0 
 
 8-6 
 
 14*0 1 1*9 
 
 19-0 
 
 16-3 
 
 24-0 
 
 22'2 
 
 '2 
 
 87 
 
 '2 I2T 
 
 2 
 
 16-6 
 
 '2 
 
 22'5 
 
 4 
 
 8-8 
 
 4 12-2 
 
 '4 
 
 16-8 
 
 '4 
 
 227 
 
 6 
 
 89 
 
 6 
 
 12-4 
 
 6 
 
 17-0 
 
 6 
 
 23-0 
 
 8 
 
 9-0 
 
 I2'5 
 
 8 
 
 17-2 
 
 8 
 
 23*3 
 
 
 
 
 
 
 
 25-0 
 
 23'S 
 
 Incases where the tension rises O'l mm. for a rise of 0*2, the pressure 
 for the intermediate tenth degree may be taken as the same as that given 
 for the temperature immediately preceding it. Thus, for the temperature 
 iO'i, the tension 9^2 mm. will be taken. P'or very accurate work, fuller 
 tables given to the third decimal should be consulted. 
 
462 
 
 Appendix. 
 
 FACTORS FOR REDUCING GASEOUS VOLUMES TO N.T.P. 
 
 The observed volume, when multiplied by the factor corresponding to 
 the temperature and the pressure, will give the volume reduced to o and 
 760 mm. See footnote on p. 390. 
 
 mm. 
 
 9- 
 
 10. 
 
 it . 
 
 12. 
 
 I3- 
 
 14. 
 
 720 
 
 9I7II 
 
 91388 
 
 91066 
 
 90746 
 
 90426 
 
 90113 
 
 
 91839 
 
 91515 -91192 
 
 9o8/2 
 
 90552 
 
 90238 
 
 2 
 
 91966 
 
 91642 
 
 9I3I8 
 
 90998 
 
 90677 
 
 90363 
 
 3 
 
 92093 
 
 91769 
 
 9H45 
 
 9II24 
 
 90803 
 
 90488 
 
 4 
 
 92221 
 
 91896 
 
 9I57I 
 
 91250 
 
 90929 
 
 90614 
 
 5 
 
 92348 
 
 92023 
 
 91698 
 
 91376 
 
 91054 
 
 90739 
 
 6 
 
 92476 
 
 9II50 
 
 91824 
 
 91502 
 
 9Il80 
 
 90864 
 
 7 
 
 92603 
 
 92277 
 
 9I95I 
 
 91628 
 
 91305 
 
 90989 
 
 8 
 
 92730 
 
 92403 
 
 92077 
 
 91754 
 
 '9I43 1 
 
 9III4 
 
 9 
 
 92858 
 
 92530 
 
 92204 
 
 91880 
 
 91556 
 
 91239 
 
 730 
 
 92985 
 
 92657 
 
 92330 
 
 92OO6 
 
 91682 
 
 91365 
 
 i 
 
 93 II2 
 
 92784 
 
 92457 
 
 92132 
 
 91808 
 
 91490 
 
 2 
 
 93240 
 
 929II 
 
 92583 
 
 92258 
 
 91933 
 
 9I6I5 
 
 3 
 
 93367 
 
 93038 
 
 92710 
 
 92384 
 
 92059 
 
 91740 
 
 4 
 
 '93495 
 
 93165 
 
 92836 
 
 92510 
 
 92184 
 
 91865 
 
 
 93622 
 
 93292 
 
 92963 
 
 92636 
 
 92310 
 
 91990 
 
 6 
 
 '93749 
 
 93419 
 
 93089 
 
 92762 
 
 92436 
 
 92II5 
 
 7 
 
 93877 
 
 93546 
 
 93216 
 
 92888 
 
 92561 
 
 92241 
 
 8 
 
 94004 
 
 93673 
 
 93342 
 
 93014 
 
 92687 
 
 92366 
 
 9 
 
 94132 
 
 93800 
 
 93469 
 
 93I4I 
 
 92812 
 
 92491 
 
 740 
 
 94259 
 
 93927 
 
 '93595 
 
 93267 
 
 92938 
 
 92616 
 
 i 
 
 94386 
 
 94054 
 
 93722 
 
 '93393 
 
 93064 
 
 92741 
 
 2 
 
 945H 
 
 94180 
 
 93848 
 
 93519 
 
 93189 
 
 92866 
 
 3 
 
 94641 
 
 94307 
 
 '93975 
 
 93 6 45 
 
 93315 
 
 92992 
 
 4 
 
 94768 
 
 '94434 
 
 '94IOI 
 
 '93771 
 
 93440 
 
 93U7 
 
 5 
 
 94896 
 
 94561 
 
 94228 
 
 93897 
 
 93566 
 
 93242 
 
 6 
 
 95023 
 
 94688 
 
 '94354 
 
 94023 
 
 93692 
 
 93367 
 
 7 
 
 95I5I 
 
 94815 
 
 94480 
 
 94149 
 
 93817 
 
 '93492 
 
 8 
 
 95 2 78 
 
 94942 
 
 94607 
 
 94275 
 
 '93943 
 
 93617 
 
 9 
 
 95405 
 
 95069 
 
 "94733 
 
 94401 
 
 94068 
 
 93742 
 
 750 
 
 '95533 
 
 '95^6 
 
 94860 
 
 94527 
 
 94194 
 
 93868 
 
 i 
 
 95660 
 
 95323 
 
 94986 
 
 94653 
 
 '943 i 9 
 
 '93993 
 
 2 
 
 95788 
 
 95450 
 
 95H3 
 
 '94779 
 
 "94445 
 
 94II8 
 
 3 
 
 95915 
 
 95577 
 
 95239 
 
 94905 
 
 94571 
 
 94243 
 
 4 
 
 96042 
 
 95704 
 
 95366 
 
 95031 
 
 94696 
 
 94368 
 
 5 
 
 96170 
 
 95831 
 
 95492 
 
 95157 
 
 94822 
 
 "94493 
 
 6 
 
 96297 
 
 95958 
 
 95619 
 
 95283 
 
 '94947 
 
 94619 
 
 7 
 
 96424 
 
 96084 
 
 95745 
 
 95409 
 
 95073 
 
 '94744 
 
 8 
 
 96552 
 
 Q62II 
 
 95872 
 
 '95535 
 
 '95 r 99 
 
 94869 
 
 9 
 
 96679 
 
 96338 
 
 95998 
 
 95661 ^5324 
 
 94994 
 
Factors for reducing Gaseous Volumes to N.T.P. 453 
 
 mm. 
 
 9 
 
 10. 
 
 n. 
 
 12. 
 
 13- 
 
 14. 
 
 7 60 
 
 96806 
 
 96465 
 
 96125 
 
 9578/ 
 
 95450 
 
 95H9 
 
 I 
 
 96934 
 
 96592 
 
 96251 
 
 95913 
 
 '95575 
 
 95244 
 
 2 
 
 97061 
 
 96719 
 
 96378 
 
 96039 
 
 95701 
 
 95370 
 
 3 
 
 97189 
 
 96846 
 
 96504 
 
 '(36165 
 
 95827 
 
 '95495 
 
 A 
 
 97316 
 
 96973 
 
 96631 
 
 96291 
 
 95952 
 
 95620 
 
 s 
 
 '97443 
 
 97100 
 
 96757 
 
 96417 
 
 96078 
 
 '95745 
 
 6 
 
 '9757 1 
 
 97227 
 
 96884 
 
 96543 
 
 96203 
 
 95870 
 
 8 
 
 97698 
 97825 
 
 97354 
 97481 
 
 97010 
 97137 
 
 95670 
 96796 
 
 96329 
 96455 
 
 '95995 
 96120 
 
 9 
 
 '97953 
 
 97608 
 
 97263 
 
 9 6922 
 
 96580 
 
 96246 
 
 770 
 i 
 
 98080 
 98208 
 
 '97734 
 97861 
 
 97390 
 97516 
 
 97048 
 97174 
 
 96706 
 96831 
 
 96371 
 96496 
 
 2 
 
 98335 
 
 97988 
 
 97642 
 
 97300 
 
 96957 
 
 96621 
 
 3 
 
 98462 
 
 98115 
 
 97769 
 
 97426 
 
 97083 
 
 96746 
 
 4 
 
 98590 
 
 98242 
 
 97895 
 
 97552 
 
 97208 
 
 96871 
 
 s 
 
 98717 
 
 98369 
 
 98022 
 
 97678 
 
 '97334 
 
 96997 
 
 6 
 
 98844 
 
 98496 
 
 98148 
 
 97804 
 
 '97459 
 
 97122 
 
 7 
 
 98972 
 
 '98623 
 
 98275 
 
 979 3 
 
 97585 
 
 97247 
 
 8 
 
 99099 
 
 98750 
 
 98401 
 
 9 8056 
 
 97710 
 
 973/2 
 
 9 
 
 99227 
 
 98877 
 
 98528 
 
 98l82 
 
 97836 
 
 '97497 
 
 780 
 
 '99354 
 
 99004 
 
 98654 
 
 98308 
 
 97962 
 
 97622 
 
 mm. 
 
 15. 
 
 1 6, 
 
 i?. 
 
 1 
 18. 
 
 19. 
 
 20. 
 
 720 
 
 89800 
 
 '89489 
 
 89180 
 
 88873 
 
 88569 
 
 88266 
 
 I 
 
 89924 
 
 89613 
 
 89304 
 
 88997 
 
 88692 
 
 88389 
 
 2 
 
 3 
 
 90050 
 
 90174 
 
 89737 
 89862 
 
 89428 
 89551 
 
 89120 
 89243 
 
 88815 
 88938 
 
 885II 
 88634 
 
 4 
 
 90299 
 
 89986 
 
 89675 
 
 89367 
 
 89061 
 
 88757 
 
 5 
 
 90423 
 
 90IIO 
 
 89799 
 
 89490 
 
 89184 
 
 88879 
 
 b 
 
 90548 
 
 90234 
 
 89923 
 
 89614 
 
 89307 
 
 89002 
 
 7 
 
 90673 
 
 90359 
 
 90047 
 
 89737 
 
 89430 
 
 89124 
 
 8 
 
 90798 
 
 90483 
 
 90I7I 
 
 89861 
 
 89553 
 
 89247 
 
 9 
 
 90922 
 
 90607 
 
 90295 
 
 89984 
 
 89676 
 
 89369 
 
 730 
 
 91047 
 
 90732 
 
 90418 
 
 90107 
 
 89799 
 
 89492 
 
 i 
 
 9II72 
 
 90856 
 
 90542 
 
 90231 
 
 89022 
 
 89615 
 
 2 
 
 91296 
 
 90980 
 
 90666 
 
 90354 
 
 90045 
 
 89737 
 
 3 
 
 9I42I 
 
 91104 
 
 90790 
 
 90478 
 
 90168 
 
 89860 
 
 4 
 
 91546 
 
 91229 
 
 90914 
 
 90601 
 
 90291 
 
 89982 
 
 5 
 
 91671 
 
 91353 
 
 91038 
 
 90725 
 
 90414 
 
 90105 
 
 o 
 
 91795 
 
 91477 
 
 91162 
 
 90848 
 
 90537 
 
 90228 
 
 7 
 
 91920 
 
 91602 
 
 91285 
 
 90971 
 
 90660 
 
 90350 
 
 8 
 
 92045 
 
 91726 
 
 91409 
 
 91095 
 
 90784 
 
 90473 
 
 9 
 
 92169 
 
 91850 
 
 91533 
 
 9I2I8 
 
 90907 
 
 90594 
 
4 6 4 
 
 Appendix. 
 
 mm. 
 
 15. 
 
 16. 17. 
 
 18. 
 
 19. 
 
 740 
 
 92294 
 
 91975 
 
 91657 
 
 91342 
 
 91030 
 
 I 
 
 92419 
 
 02099 
 
 91781 
 
 9M65 
 
 9H53 
 
 2 
 
 
 92223 
 
 91905 
 
 91589 
 
 91276 
 
 3 
 
 92668 
 
 92347 
 
 92029 
 
 9I7I2 
 
 91399 
 
 4 
 
 92793 
 
 92472 
 
 92152 
 
 91836 
 
 91522 
 
 5 
 
 92918 
 
 92596 
 
 92276 
 
 91959 
 
 91645 
 
 6 
 
 93043 
 
 92720 
 
 92400 
 
 92082 
 
 91768 
 
 7 
 
 93167 
 
 92845 
 
 92524 
 
 92206 
 
 91891 
 
 8 
 
 93292 
 
 92969 
 
 92648 
 
 92329 
 
 92014 
 
 9 
 
 93417 
 
 93093 
 
 92772 
 
 92453 
 
 92137 
 
 750 
 
 93541 
 
 93217 
 
 92896 
 
 92576 
 
 92260 
 
 i 
 
 93666 
 
 93342 
 
 93020 
 
 92700 
 
 92383 
 
 2 
 
 93791 
 
 93466 
 
 93143 
 
 92823 
 
 92506 
 
 3 
 
 93916 
 
 93590 
 
 93267 
 
 92946 
 
 92629 
 
 4 
 
 94040 
 
 93715 
 
 93391 
 
 93070 
 
 92752 
 
 6 
 
 94165 
 94290 
 
 93839 
 93963 
 
 93515 
 93639 
 
 93193 
 93317 
 
 92875 
 2998 
 
 7 
 
 94414 
 
 94087 
 
 93763 
 
 93440 
 
 93I2I 
 
 8 
 
 '94539 
 
 94212 
 
 93887 
 
 93564 
 
 93244 
 
 9 
 
 94664 
 
 94336 
 
 94010 
 
 93687 
 
 93367 
 
 760 
 
 94789 
 
 94460 
 
 94134 
 
 93811 
 
 93489 
 
 i 
 
 94913 
 
 94585 
 
 94258 
 
 '93934 
 
 93613 
 
 2 
 
 95038 
 
 94709 
 
 94382 
 
 94057 
 
 93736 
 
 3 
 
 95163 
 
 94833 
 
 94506 
 
 94181 
 
 93859 
 
 4 
 
 95288 
 
 '94957 
 
 94630 
 
 94304 
 
 93982 
 
 
 95412 
 
 95082 
 
 '94754 
 
 94428 
 
 94105 
 
 6 
 
 '95537 
 
 95206 
 
 94877 
 
 94551 
 
 94228 
 
 7 
 
 95662 
 
 95330 
 
 95001 
 
 94675 
 
 '9435 1 
 
 8 
 
 95786 
 
 '95455 
 
 95125 
 
 94798 
 
 '94474 
 
 9 
 
 959H 
 
 '95579 
 
 '95249 
 
 94921 
 
 '94597 
 
 770 
 
 96036 
 
 95703 
 
 '95373 
 
 95045 
 
 94720 
 
 i 
 
 96161 
 
 95827 
 
 '95497 
 
 95168 
 
 94843 
 
 2 
 
 96285 
 
 95952 
 
 95621 
 
 95292 
 
 94966 
 
 3 
 
 96410 
 
 96076 
 
 '95744 
 
 95415 
 
 95089 
 
 4 
 
 96535 
 
 96200 
 
 95868 
 
 '95539 
 
 95212 
 
 5 
 
 96659 
 
 96325 
 
 95992 
 
 95662 
 
 '95335 
 
 6 
 
 96784 
 
 96449 
 
 96116 
 
 957S5 
 
 95458 
 
 7 
 
 96909 
 
 96573 
 
 96240 
 
 95909 
 
 95581 
 
 . 8 
 
 97034 
 
 96698 
 
 96364 
 
 96032 
 
 95704 
 
 9 
 
 97158 
 
 96822 
 
 96488 
 
 96156 
 
 95827 
 
 7 8o 
 
 97283 
 
 96946 
 
 96611 
 
 96279 
 
 95950 
 
Factors for reducing Gaseous Volumes to N.T.P. 465 
 
 mm. 
 
 21. 
 
 22. 
 
 01. 
 
 mm. 
 
 21. 
 
 22. 
 
 *3. 
 
 720 
 
 87972 
 
 87668 
 
 87371 
 
 751 
 
 91759 
 
 91442 
 
 9H33 
 
 
 88094 
 
 87790 
 
 87493 
 
 2 
 
 91882 
 
 91564 
 
 91254 
 
 2 
 
 88216 
 
 879H 
 
 87614 
 
 3 
 
 92OO4 
 
 91686 
 
 91376 
 
 3 
 
 88339 
 
 88033 
 
 87735 
 
 4 
 
 92126 
 
 91808 
 
 '9H97 
 
 4 
 
 88460 
 
 88155 
 
 87857 
 
 5 
 
 92248 i ^1929 
 
 9l6l8 
 
 
 88583 
 
 882 7 7 
 
 87978 
 
 6 
 
 92370 | ^OSI 
 
 91740 
 
 6 
 
 88705 
 
 88398 
 
 88099 
 
 7 
 
 92492 
 
 92173 
 
 9l86l 
 
 7 
 
 88827 
 
 88520 
 
 88221 
 
 8 
 
 92615 
 
 92295 
 
 91982 
 
 8 
 
 88949 
 
 88642 
 
 88342 
 
 9 
 
 92737 
 
 92416 
 
 92104 
 
 9 
 
 89071 
 
 88764 : -88463 
 
 760 
 
 92859 
 
 92538 
 
 92225 
 
 730 
 
 89193 
 
 88885 -88585 
 
 
 92981 
 
 9266O 
 
 92346 
 
 
 89316 
 
 89007 
 
 88706 
 
 2 
 
 9 3 I03 
 
 92782 
 
 92468 
 
 2 
 
 89438 
 
 89129 -88827 
 
 3 
 
 93226 
 
 92904 
 
 92589 
 
 3 
 
 89560 
 
 89251 
 
 88949 
 
 4 
 
 93348 
 
 93025 
 
 927II 
 
 4 
 
 89682 
 
 89372 
 
 89070 
 
 5 
 
 93470 
 
 93147 
 
 92832 
 
 5 
 
 8980 4 
 
 89494 
 
 89I9I 
 
 6 
 
 93592 
 
 9 3269 
 
 92953 
 
 6 
 
 89927 
 
 80616 
 
 89313 
 
 7 
 
 937H 
 
 93391 
 
 93075 
 
 7 
 
 90049 
 
 89738 
 
 ^9434 
 
 8 
 
 9 3 836 
 
 93512 
 
 93196 
 
 8 
 
 90I7I 
 
 89860 
 
 ^9555 
 
 9 
 
 '93959 
 
 93634 
 
 93317 
 
 9 
 
 9 0293 
 
 89981 
 
 896 7 7 
 
 770 
 
 0408l 
 
 93756 
 
 '93439 
 
 740 
 
 90415 ^OIOS 
 
 8 9 798 
 
 i 
 
 94203 
 
 93878 
 
 93560 
 
 i 
 
 90538 
 
 90225 
 
 89920 
 
 2 
 
 94325 
 
 '93999 
 
 93681 
 
 2 
 
 90660 
 
 90347 
 
 90041 
 
 3 
 
 '94447 
 
 94I2I 
 
 93803 
 
 3 
 
 90782 
 
 90468 
 
 9OI62 
 
 4 
 
 94570 
 
 94243 '93924 
 
 4 
 
 90904 ^0590 
 
 90284 
 
 5 
 
 94692 
 
 94365 
 
 94045 
 
 5 
 
 9IO26 
 
 90712 
 
 90405 
 
 6 
 
 94814 
 
 94486 
 
 94167 
 
 6 
 
 9II48 
 
 90834 
 
 90526 
 
 7 
 
 94936 
 
 94608 
 
 94288 
 
 7 
 
 9I27I 
 
 90955 
 
 9C648 
 
 8 
 
 95059 
 
 94730 
 
 94409 
 
 8 
 
 91393 -91077 
 
 90769 
 
 9 
 
 95180 
 
 94852 
 
 94531 
 
 9 
 
 9i5!5 
 
 91199 
 
 90890 
 
 780 
 
 95303 
 
 '94973 
 
 94652 
 
 75 
 
 91637 
 
 91321 
 
 9IOI2 
 
 
 
 
 
 2 H 
 
INDEX 
 
 ABSORPTION pipettes, gas analysis, 
 
 393 
 
 , mode of filling, 394 
 
 Acid, acetic, 156 
 
 , boric, 169 
 
 , carbonic, 153 
 
 , , gasometric estimation, 
 
 368-396 
 
 , , gravimetric estimation, 259 
 
 , , volumetric estimation, 364 
 
 , citric, 158 
 
 , cyanic, 164 
 
 , formic, 155 
 
 , hydriodic, 128 
 
 , hydrobromic, 125 
 
 , hydrochloric, 122 
 
 , , gravimetric estimation, 259 
 
 , , standard solution, 323 
 
 , hydrocyanic, 159 
 
 , hydrofluoric, 133 
 
 , hydrofluosilicic, 135 
 
 , hypochlorous, 129 
 
 , hypophosphorous, 151 
 
 , metaphosphoric, 150 
 
 , metastannic, 104 
 
 , nitric, 144 
 
 , .gravimetric estimation, 263 
 
 , , volumetric estimation, 363 
 
 , nitrous, 147 
 
 , orthophosphoric, 150 
 
 , osmic, 109 
 
 , oxalic, 155 
 
 , , standard solution, 378 
 
 , phosphoric, 63, 149 
 
 , , gravimetric estimation, 264 
 
 , , volumetric estimation, 369 
 
 , phosphorous, 151 
 
 , pyrophosphoric, 150 
 
 , radicals, 15 
 
 , silicic, 166 
 
 , sulphuric, 139 
 
 , , gravimetric estimation, 258 
 
 , , standard solution, 319 
 
 Acid, tartaric, 157 
 
 , thiocyanic, 165 
 
 , thiosulphuric, 142 
 
 Acetates, 156 
 
 Acidimetry, 316 
 
 Acids, analytical classification of, 171 
 
 Acids and alkalies, 12 
 
 Agate, 166 
 
 Air-oven, 198 
 
 Alkalies, estimation of in silicates, 
 287 
 
 , by electrolytic method, 290 
 
 Alkalimetry, 316 
 
 Alloy, silver and copper, electrolytic 
 analysis, 297 
 
 , , , gravimetric analysis, 
 
 267 
 
 , tin and lead, gravimetric analy- 
 sis, 269 
 
 , zinc, copper, nickel, electrolytic 
 
 analysis, 297 
 
 , , , gravimetric analysis, 
 
 270 
 
 , zinc, copper, tin, 274 
 
 Aluminium, gravimetric estimation in 
 dolomite, 277 
 
 , potash alum, 222 
 
 reactions, 37 
 
 Alums, 37 
 
 Ammonia, volumetric estimation, 327 
 
 Ammonium, gravimetric estimation, 
 250 
 
 , phosphomolybdate, 63 
 
 , reactions, 19 
 
 , thiocyanate, deci-normal solu- 
 tion, 365 
 
 Analytical classification, 13 
 
 Analytical groups, 16 
 
 Antimoniuretted hydrogen, 101 
 
 Antimony, gravimetric estimation, 255 
 
 , separation from, arsenic, 
 
 284 
 
 reactions, 97 
 
468 
 
 Index. 
 
 Antimony, volumetric estimation, 352 
 
 Aqueous vapour, tension of, 461 
 
 Arsenic, gravimetric estimation as sul- 
 phide, 253 
 
 , magnesium pyro- 
 
 arsenate, 254 
 
 reactions, 89 
 
 , volumetric estimation, 352 
 
 Aspirator, 204 
 
 Atomic weights, table of, 453 
 
 BALANCE, 190 
 
 Barium, gravimetric estimation, 258 
 
 Barium hydroxide, standard solution, 
 
 379 
 
 Barium reactions, 30 
 
 , spectrum of, 34 
 
 Baryta-water, 379 
 
 Beryl, 70 
 
 Beryllium, 70 
 
 Bismuth, reactions, 81 
 
 Bleaching powder, 130 
 
 , available chlorine in, 353 
 
 , valuation of by arsenious 
 
 oxide, 354 
 
 , valuation by means of gas- 
 volumeter, 412 
 
 Blowpipe flame, 9 
 
 Blowpipe reactions, exercises, 9 
 
 Boiling-point determination, 436 
 method, 448 
 
 Borates, 169 
 
 Borax, 169 
 
 Borax beads, u 
 
 Boron, 169 
 
 Boron fluoride, 135, 170 
 
 Bromates, 130 
 
 , volumetric estimation, 358 
 
 Bromides, 125 
 
 Bromides, -iodides, and chlorides, 
 detection in solution together, 128 
 
 Bromine, 124 
 
 , gravimetric estimates in bro- 
 mides, 259 
 
 , estimation in organic com- 
 pounds, 430 
 
 Bronze coinage, analysis of, 274 
 
 Burettes, 308 
 
 , calibration of, 309 
 
 , gas, 385 
 
 C ADMIUM, gravimetric estimation , 257 
 , in zinc-blende, 282 
 
 , reactions, 83 
 
 , volumetric estimation, 364 
 Caesium, 26 
 Calcium, gravimetric estimation in 
 
 calcium carbonate, 231 
 
 , dolomite, 277 
 
 , silicates, 287 
 
 Calcium reactions, 32 
 
 , volumetric estimation, 340, 363 
 
 Calibration of burettes, 309 
 
 gas-burettes, 386 
 
 litre flasks, 305 
 
 pipettes, 307 
 
 Carbon, 153 
 
 , estimation in organic com- 
 pounds, 414 
 
 Carbon dioxide, estimation by means 
 of gas-volumetre, 412 
 
 , gravimetric estimation, 259 
 
 , indirect volumetric estima- 
 tion, 364 
 
 in air (Pettenkofer's 
 
 method), 378 
 
 in other gas mixtures, 399 
 
 Carbon monoxide, estimation of, 401 
 
 Carbonates, 153 
 
 Cerium, 16 
 
 Chlorates, 130 
 
 and nitrates, detection together, 
 
 147 
 
 , separation from chlorides, 133 
 
 , volumetric estimation, 358 
 
 Chlorides, 121 
 
 , detection in presence of bromide 
 
 or iodide, 124 
 
 Chlorine, 121 
 
 , gravimetric estimation in chlo- 
 rides, 259 
 
 , in organic com- 
 pounds, 430 
 
 , volumetric estimation by pre- 
 cipitation, 360, 366 
 
 , in bleaching powder, 
 
 353 
 
 Chrome iron ore, 39, 344 
 
 Chromium, gravimetric estimation, 
 229 
 
 reactions, 39 
 
 , volumetric estimation in chrome 
 
 iron ore, 344 
 
 Chromyl chloride, 123 
 
 Citrates, 158 
 
 Clark's method for estimating hard- 
 ness of water, 371 
 
 Coal-gas, analysis of, 408 
 
 Cobalt, gravimetric estimation, 247 
 
 , reactions, 58 
 
 Cobaltamines, 59 
 
 Cobalticyanides, 60 
 
 Cobaltocyanides, 60 
 
 Coefficients of expansion of water, 
 table of, 460 
 
 Collection of gases for analysis, 390 
 
 Combustion analysis, 414 
 
 Combustion tube, 414 
 
 Copper, electrolytic estimation, 294 
 
Index. 
 
 469 
 
 Copper, gravimetric estimation in 
 copper sulphate, (i) as oxide, 235 ; 
 (2) as sulphide, 236 ; (3) as cuprous 
 thio-cyanate, 237 
 
 , in German silver, 271 
 
 , silver coinage, 267 
 , bronze coinage, 275 
 
 , zinc-blende, 282 
 
 - reactions, 83 
 
 , volumetric estimation (Mans- 
 field process), 375 
 
 Correction of gaseous volumes, 387 
 
 , factors for, .,62 
 
 Crucible tongs, manipulation of, 227 
 
 Cryoscopic method, 443 , 
 
 Crystallisation, 208 
 
 Cupric salts, 85 
 
 Cuprous chloride for gas absorption, 
 396 
 
 hydride, 151 
 
 salts, 86 
 
 Cyanates, 164 
 
 Cyanides, double, 161 
 of Group IIlB., 60 
 , reactions, 159 
 , single, 160 
 
 Cyanogen, 159 
 
 , volumetric estimation, 362 
 
 DESICCATOR, 196 
 Dolomite, analysis of, 276 
 Double cyanides, 161 
 
 easily decomposed, 162 
 
 difficultly decomposed, 163 
 Double gas-pipette, mode of filling, 
 
 394 
 
 Double salts, 213 
 Drying and weighing a filter, 193 
 
 ELECTROLYTIC methods of analysis, 
 292 
 
 Etching glass, 134 
 
 Evaporating to dryness, 4 
 
 Evaporation, 4 
 
 Expansion of water, table of coeffi- 
 cients, 460 
 
 Explosion-pipettes, 404, 407 
 
 FACTORS for correction of gaseous 
 
 volumes, 462 
 Feeling's solution, 85 
 Felspar, 167 
 Ferricyanides, 164 
 Ferrocyanides, 163 
 Filter, incineration of, 207 
 Filter ash, estimation of, 206 
 Filter-paper, quantitative, 193 
 Filter-pump, 212 
 Filtration, i 
 
 Flame, oxidising and reducing, 9 
 Fleitmann's test, 97 
 Flint, 166 
 Floats, 311 
 .Fluid magnesia, 153 
 Fluorides, 133 
 Fluorine, 133 
 Formates, 155 
 Freezing-point method, 442 
 Fulminating gold, 106 
 Furnace gases, sulphur dioxide in, 381 
 Fusion, 5 
 
 Fusion mixture, 49 
 Fusion with borax, exercises, 10 
 
 GAS analysis, 377 
 Gas-burettes, 385 
 Gas-volumeter, 411 
 Gases, absorption analysis, 396, 398 
 , correction of volumes for tem- 
 perature and pressure, 387 
 , estimation by absorption and 
 
 titration, 378 
 
 , measurement, 384 
 
 , measurement of, 384 
 
 Geissler's potash bulbs, 417 
 
 General reagents, 17 
 
 German silver, electrolytic analysis of, 
 
 297 
 
 , gravimetric analysis of, 270 
 
 Gold reactions, 105 
 Graduated cylinders, 306 
 Graduated flasks, 303 
 Gravimetric methods of analysis, 189 
 Gravimetric separations 
 
 Antimony and arsenic, 284 
 
 Copper and cadmium, 282 
 
 silver, 267 
 
 nickel, 271 
 
 zinc, 271 
 
 Iron and aluminium, 278 
 manganese, 283 
 
 zinc, 271 
 
 Magnesium and calcium, 277 
 
 Manganese and zinc, 283 
 
 Nickel and zinc, 272 
 
 Tin and copper, 272 
 
 Tin and lead, 269 
 Group separations, 116 
 Group V. , detection of metals of, 25 
 
 IV. , separation of metals of, 35 
 
 IIlA, ,, ,, 49 
 
 IIlB, ,, ,,62 
 
 II., Division i, separation of 
 
 metals of, 88 
 
 II., Division 2, separation of 
 
 metals of, 108 
 
 I., separation of metals of, 115 
 
 Group reagents, 17 
 
 2 H 2 
 
470 
 
 Index. 
 
 HALIDES, 121 
 
 Halogen oxyacids, 129 
 
 Halogens, 121 
 
 Hardness of water, Clark's process, 
 
 37i 
 
 , Hehner's method, 328 
 
 Hardness, table of, 459 
 
 Hempel's gas-burette. 385 
 
 Hot-water funnel, 210 
 
 Hydrogen, estimation in gaseous 
 
 mixtures, 402, 404, 408 
 
 , organic compounds, 414 
 
 Hypochlorites, 129 
 Hypochlorous acid, 129 
 Hypophosphites, 151 
 
 IGNITION, 8 
 Indicators, use of, 315 
 
 for alkalimetry, 317 
 
 Indium, 16 
 
 Insoluble substances, treatment of, 1 84 
 
 Instruments for measuring liquids, 
 
 T 303 
 lodates, 130 
 
 , volumetric estimation of, 358 
 
 Iodides, 129 
 
 , detection in solution with 
 
 bromides, 127 
 Iodine, 126 
 
 , deci-normal solution, 348 
 
 , estimations by means of, 352 
 
 , separation from bromine and 
 
 chlorine, 129 
 
 , volumetric reactions with, 346 
 Ions, 292 
 Iridium, 109 
 Iron, gravimetric estimation in a 
 
 ferrous salt, 230 
 
 , German silver, 272 
 
 , dolomite, 278 
 
 , zinc-blende, 282 
 
 , magnetic oxide, 44 
 
 reactions, 44 
 
 , volumetric estimation of, in ores, 
 
 335. 343 
 
 LEAD, gravimetric estimation in lead 
 acetate, (i) as oxide, 240; (2) as 
 sulphate, 242 
 , bronze, 274 
 
 , solder, 270 
 
 , zinc-blende, 281 
 
 reactions, 79 
 
 Lithium reactions, 26 
 
 , spectrum of, 28 
 
 Litmus, 12, 316 
 
 Litre flasks, calibration of, 305 
 
 MAGNESIA mixture, 456 
 
 Magnesium, gravimetric estimation in 
 magnesium sulphate, 233 
 
 , dolomite, 278 
 
 , pyro-arsenate, 254 
 
 reactions, 23 
 
 Manganese dioxide, valuation of 1^7 
 356 
 
 , gravimetric estimation in potas- 
 sium permanganate, 245 
 
 , silicates, 287 
 
 , zinc-blende, 283 
 
 reactions, 51 
 
 Mansfield process for copper, 375 
 I Marsh-gas, estimation of, 406, 408 
 
 Marsh's test, 94 
 
 Melting-point determination, 438 
 
 Meniscus, 310 
 
 Mercuric compounds, reactions, 76 
 
 Mercurous compounds, 78 
 
 Mercury, gravimetric estimation, 257 
 
 Mercury reactions, 76 
 
 Metals, 15 
 
 Metalloidal elements, 16 
 
 Meta-phosphates, 150 
 
 Methane, estimation of, 406, 408 
 
 Methyl orange, 317 
 
 Miscellaneous physico-chemical deter- 
 minations, 434 
 
 Molecular weight by boiling-point, 448 
 
 Molecular weight by freezing-point, 442 
 
 Molybdenum, in 
 
 NEGATIVE radicals, 15 
 Nessler's solution, 20 
 Neutralisation, exercises, n 
 Nickel, electrolytic estimation of, 297, 
 299 
 
 , gravimetric estimation, 247 
 
 , in German silver, 272 
 
 reactions, 56 
 
 Niobium, 16 
 Nitrates, 144 
 and chlorates, detection together, 
 
 147 
 
 and nitrites, detection together, 149 
 
 , reduction by ferrous salts, 145 
 
 , nascent hydrogen, 146 
 
 , sulphurous acid, 145 
 
 Nitric acid, gravimetric estimation of, 
 
 263 
 
 , volumetric estimation, 363, 411 
 
 Nitrites, 147 
 
 , distinction from nitrates, 148 
 
 Nitrogen, estimation in gaseous 
 
 mixture, 408 
 
 , organic compounds, 424 
 
 , by Dumas' method, 
 
 424 
 
Index. 
 
 471 
 
 Nitrogen, estimation in organic com- 
 pounds by Kjeldahl's method, 429 
 
 , by soda-lime process, 
 
 427 
 
 Nitrometer, 410 
 
 Nitroprusside of sodium, 139 
 
 ORTHOPHOSPHATES, 150 
 
 Osmium, 109 
 
 Oxalates, 155 
 
 Oxalic acid, standard for gas analysis, 
 
 378 
 Oxygen, estimation of, 401 
 
 PALLADIOUS nitrate, 129 
 
 Palladium, 109 
 
 Palladiumised asbestos, 402 
 
 Per chlorates, 133 
 
 Pettenkofer's method for carbon 
 dioxide, 378 
 
 Phenolphthalein, 317 
 
 Phosphites, 151 
 
 Phosphates, Group III., 63 
 , reactions of, 63, 149 
 
 Phosphoric acid, gravimetric estima- 
 tion of, 264 
 
 , removal of, in Group III., 67 
 
 , separation of, 69 
 
 , volumetric estimation, 369 " 
 
 Phosphorus, 149 
 
 , estimation in organic com- 
 pounds, 432 
 
 Physico-chemical determinations, 434 
 
 Pipettes, 306 
 
 , calibration of, 307 
 
 Platinum reactions, 105 
 
 Positive radicals, 15 
 
 Potassioscope, 22, 33 
 
 Potassium cyanide, standard solution 
 for copper, 374 
 
 dichromate, deci-normal, 341 
 
 , analyses by means of, 342 
 
 , gravimetric estimation of, (i)as 
 
 potassium sulphate, 247 ; (2) as 
 potassium platinic chloride, 249 
 
 , in silicates, 287 
 
 Potassium oleate, standard solution, 
 372 
 
 permanganate, deci-normal solu- 
 tion, 332 
 
 , typical analyses by means 
 
 of, 335 
 
 Potassium reactions, 22 
 Powdering and sampling, 265 
 Precipitation, 6 
 
 Preliminary exercises, qualitative, i 
 Preliminary manipulations, quantita- 
 tive, 189 
 
 Preliminary tests in qualitative ana- 
 lysis, 178 
 
 Pure salts, preparation of, 208 
 Purple of Cassius, 103, 107 
 Pyrogallol, 397 
 Pyrophosphates, 150 
 
 QUANTITATIVE analysis, 189 
 Quartz, 166 
 
 R ADULT'S method, 443 
 Rare metals of Group I., 118 
 
 II., 108 
 
 III., 70 
 
 V., 26 
 
 Reactions of the metals of Group I., 
 112 
 
 II., Division i, 76 
 
 II., Division 2, 89 
 
 IIIA.,37 
 
 IIIB.,51 
 
 IV., 30 
 
 Reactions, wet and dry, 13 
 Reagents, 14 
 
 , group or general, 17 
 
 , preparation of, 454 
 Reinsch's test, 92 
 
 Results of a qualitative analysis, 186 
 Rhodium, 16, 108 
 Rose's crucible, 237 
 Rubidium, 26 
 Ruthenium, 16, 108 
 
 SAMPLING and powdering, 265 
 
 Sand, 1 66 
 
 Scandium, 16 
 
 Scheele's green, 92 
 
 Schiff's burette, 426 
 
 Schrotter's apparatus for carbon 
 
 dioxide, 260 
 Selenium, no 
 Separation of phosphoric acid, 69 
 
 the metals of Group I., 115 
 
 II., Division i, 88 
 
 II. , Division 2, 108 
 
 IIlA.,49 
 
 -IIlB., 62 
 
 IV., 35 
 
 Silica, 1 66 
 
 , estimation of, in silicates, 285 
 
 Silicates, 166 
 
 , gravimetric analysis of, 285 
 
 , treatment of insoluble, 168 
 
 Silicic acid, gravimetric estimation, 
 
 264, 285 
 
 Silicofluorides, 135 
 Silicon, 166 
 Silicon fluoride, 134 
 Silver coinage, 267 
 
4/2 
 
 Index. 
 
 Silver coinage, electrolytic analysis of, 
 
 297 
 Silver, gravimetric estimation in silver 
 
 nitrate, 238 
 
 coinage, 267 
 Silver reactions, 112 
 
 nitrate, deci-normal, 360 
 
 volumetric estimation of, 365, 366 
 
 Single cyanides, 160 
 Soap solution, standard, 372 
 Soda-ash, volumetric estimation of, 324 
 Sodium nitroprusside, 139 
 
 - chloride, deci-normal solution, 
 
 364 
 
 , gravimetric estimation in a 
 
 silicate, 287 
 
 , gravimetric separation from 
 
 potassium, 288 
 
 reactions, 21 
 
 sulphide, standard solution, 370 
 
 thiosulphate, deci-normal, 348 
 
 Solder, gravimetric analysis of, 269 
 
 Soluble glass, 167 
 
 Solubilities, table of, 177 
 
 Solvent, 3 
 
 Solution, 3 
 
 Specific-gravity determinations, 434 
 
 Spectra of alkali metals, 28 
 
 alkaline earth metals, 34 
 
 Spectroscope, 27 
 
 Spontaneous evaporation, 4 
 
 Standard solutions, 312 
 
 Stannic compounds, 103 
 
 Stannous compounds, 102 
 
 Steam-oven, 195 
 
 Strontium reactions, 31 
 , spectrum of, 34 
 
 Sulphates, 139 
 
 Sulphides, 137 
 
 Sulphites, 140 
 
 Sulphur, 137 
 
 , estimation in organic com- 
 pounds, 432 
 , estimation in zinc-blende, 280 
 
 dioxide, estimation by means of 
 
 iodine, 353 
 
 in furnace gases, 381 
 
 Sulphuretted hydrogen, 137 
 
 Sulphuric acid, gravimetric estimation 
 of, 258 
 
 , standard solution, 319 
 
 Systematic detection of the acids, 171 
 
 separation of the groups, 116 
 
 TARTAR emetic, 98 
 
 Tartrates, 157 
 
 Tellurium, no 
 
 Tension of aqueous vapour, table of, 
 
 461 
 
 Test-mixer, 306 
 Thallium, 118 
 Thiocyanates, 16^ 
 Thiosulphates, 142 
 Thorium, 16 
 Tin, gravimetric estimation, 252 
 
 , in solder, 270 
 
 , in German silver, 272 
 
 , in bronze coinage, 274 
 
 reactions, 101 
 
 , volumetric estimation of, 339, 
 
 352 
 
 Titanium, 72 
 Tungsten, 119 
 
 URANIUM, 73 
 
 acetate, standard solution, 367 
 
 VANADIUM, 18 
 
 Vapour-density determination, 439 
 
 Volumetric methods of analysis, 300 
 
 based on oxidation, 331 
 
 precipitation, 360 
 
 WATER, hardness of, 328, 371 
 
 - of crystallisation, direct esti- 
 mation, 203 
 
 , estimation by loss, 197 
 Weighing, 189 
 Weighing-bottle, 193 
 Weights, 191 
 
 YTTERBIUM, 16 
 Yttrium, 16 
 
 ZINC, electrolytic estimation of, 296 
 , gravimetric estimation in zinc 
 
 sulphate, (i) as oxide, 243; (2) as 
 
 sulphide, 244 
 
 , German silver, 271 
 
 , bronze, 275 
 
 , zinc-blende, 283 
 
 reactions, 54 
 
 , volumetric estimation, 371 
 
 Zinc-blende, analysis of, 279 
 Zinci carbonas, 55 
 Zirconium, 71 
 
 THE END 
 
 PRINTED BY WILLIAM CLOWES AND SONS, LIMITED, LONDON AND BECCLES. 
 
DATE DUE SLIP 
 
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 STAMPED BELOW 
 
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:^D75 ''th, C-.S 
 
 manual o: 
 
 T chemical 
 
 inalysis, qualitative 
 
 cci. 
 
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