The sulphides of the metal ions of the zinc group are readily precipitated by ammonium or sodium sulphide, but hydrogen sulphide, in the presence of a small excess of a strong acid, such as hydrochloric acid, does not precipitate any of these sulphides (or any of the sulphides of the aluminium, the alkaline earth and the alkali groups). Under the same conditions the sulphides of the metal ions of the silver group, Ag+, Hg+, Pb2+, of the copper group, Hg2+, Pb2+, Bi3+, Cu2+, Cd2+, and also of the arsenic group, As3+, As5+, Sb3+, Sb5+, Sn2+, Sn4+, Pt2+, Pt4+, Au+, and Au3+, are precipitated. Advantage is taken of these relations in the following way, in systematic analysis: after the separation of the silver group, by precipitation of the difficultly soluble chlorides, hydrogen sulphide, in the presence of an excess of acid, is used to precipitate the sulphides of the ions of the copper and the arsenic groups, the two groups being precipitated together. Hydrogen sulphide, under these conditions, does not precipitate any sulphides of the zinc group or those of any of the remaining groups. Hydrogen sulphide is used, in this way, as one of the most valuable reagents in analytical work, enabling the analyst to separate whole groups of metal ions from other groups. There is also no other agent, equally important, which is more likely to be used in a wrong way and to lead to error.
For the first dissociation, HSH ⇄ H+ + HS−, we have
The value of the constant395 for this primary ionization of hydrogen sulphide is 0.91E−7. It is apparent that hydrogen sulphide, even in the primary ionization, is a very weak acid and produces a very small concentration of hydrosulphide-ion. In a solution, saturated at 25°, the total concentration of hydrogen sulphide is approximately 0.1 molar, and the concentration of hydrosulphide-ion, therefore, in the absence of any foreign acid, at most396 0.95E−4. The concentration of the dissolved, nonionized hydrogen sulphide, [H2S], is practically a constant, if solutions saturated with hydrogen sulphide under a given pressure, say under atmospheric pressure, are considered. For such solutions, then, we may put more simply396
The concentration of hydrosulphide-ion is, therefore, inversely proportional to the concentration of hydrogen-ion. It is clear that the addition of a strong acid, readily yielding concentrations of hydrogen-ion very much greater than 0.95E−4 (the value of [H+] in a saturated aqueous solution of hydrogen sulphide) will, as the result of the greatly increased total hydrogen-ion concentration, reduce the concentration of hydrosulphide-ion to correspondingly low values. For instance, the presence of 0.1 molar hydrochloric acid will increase the concentration of hydrogen-ion close to a thousandfold and will reduce the concentration of hydrosulphide-ion to 0.9E−7.
For the secondary ionization (see p. 101) of hydrogen sulphide, HS− ⇄ H+ + S2−, we have
The value of this constant has recently been determined397 and found to be 1.2E−15. Recalling the fact that the concentrations [H+] and [HS−] of hydrogen-ion and hydrosulphide-ion, [p201] respectively, resulting from the primary ionization, are each398 0.95E−4, we have for the concentration of the sulphide-ion, in aqueous solution saturated with hydrogen sulphide at atmospheric pressure and 25°, [S2−] = 1.2E−15.
Combining equations (II) and (III), we have, further:399
which shows, directly, the relation between the concentration of the sulphide-ion and that of the hydrogen-ion, the relation of primary importance in considering the precipitation of metal sulphides in acid solutions. The concentration of the sulphide-ion is, thus, inversely proportional to the square of the concentration of the hydrogen-ion. A thousandfold increase in the concentration of the latter, which is very nearly the effect produced by the presence of 0.1 molar hydrochloric acid ([H+] = 0.091), reduces the concentration of sulphide-ion in the saturated aqueous solution a millionfold: If we call [S2−]Ac. the concentration of the sulphide-ion in the acid solution, [S2−]Ac. = (1.1E−23) / (0.091)2 = 1.3E−21, whereas, in the absence of acid, as found above, [S2−] = 1.2E−15.
On the other hand, the addition of alkali to hydrogen sulphide, by neutralizing and suppressing the hydrogen-ion, and by forming the salts MeSH and Me2S, will very greatly increase the concentrations of the hydrosulphide-ion and of the sulphide-ion. Since the constant for the secondary ionization of hydrogen sulphide shows that HS− is an exceedingly weak acid, its salts, Me2S, are very largely hydrolyzed, the constant for water being somewhat greater400 than its own. According to Knox, in a 0.1 molar solution of Na2S about 99% of the sulphide is hydrolyzed: Na2S + H2O ⥂ NaSH + NaOH. In spite of this almost complete hydrolysis, sufficient sodium sulphide remains in a solution of this [p202] substance, to yield a concentration of the sulphide-ion that is far greater than that obtained from a solution of hydrogen sulphide. In a 0.1 molar solution of sodium sulphide the concentration of the sulphide-ion is, approximately, [S2−]alk. = 0.9E−3, as compared with [S2−] = 1.2E−15 in a saturated solution of hydrogen sulphide (25°, 760 mm.), and with 1.3E−21 in the same solution in the presence of 0.1 molar hydrochloric acid.
Ammonium sulphide (NH4)2S is the salt of an extremely weak acid with a much weaker base than sodium hydroxide, and it is correspondingly more completely decomposed by water. In a 0.1 molar solution of the sulphide (NH4)2S, we find the approximate concentration401 of the sulphide-ion [S2−]am. = 1.8E−6, as compared with 0.9E−3 in a similar solution of Na2S. But the concentration of sulphide-ion is still enormously greater than its concentrations in hydrogen sulphide in the absence and in the presence of acids (see above).
The following table402 contains a summary of the concentrations of sulphide-ion in the various solutions discussed, as well as its concentration in the presence of 0.2 molar hydrochloric acid. In the separation of the copper and arsenic groups from the zinc and aluminium groups, a concentration of hydrogen-ion corresponding to the presence of 0.15 to 0.25 molar hydrochloric acid is satisfactory for an accurate separation for ordinary purposes.
| Solution. | [S2−] |
|---|---|
| 0.1 molar Na2S: | 0.9E−3 |
| 0.1 molar NaSH: | 0.8E−5 |
| 0.1 molar (NH4)2S: | 1.8E−6 |
| 0.1 molar (NH4)SH: | 1.4E−8 |
| 0.1 molar H2S, sat. aq. sol., 25°: | 1.2E−15 |
| 0.1 molar H2S, 0.1 molar HCl:403 | 1.3E−21 |
| 0.1 molar H2S, 0.2 molar HCl: | 3.5E−22 |
The absolute values of the solubilities of the various sulphides, which are involved in the discussion, are known, with any degree of accuracy, only in a few cases. The aim of the discussion will be, therefore, to develop, rather, the relations in the values involved, which may be readily determined. Wherever absolute quantities can be given, they also will be referred to.
If the solution of zinc sulphate is saturated with hydrogen sulphide, under the same conditions as were used with the ferrous salt solution (exp.), or if we add the zinc sulphate solution to the mixture of ferrous sulphate and hydrogen sulphide (exp.), we immediately obtain heavy white precipitates of zinc sulphide. We would decide, therefore, on the basis of the principle of the solubility-product, that in this case [Zn2+] × [S2−] > KZnS. Since we have used hydrogen sulphide under practically the same conditions, we may consider that [S2−], in this experiment,407 is the same as in the [p205] test with ferrous sulphate, and, by the conditions of the experiment, we have also made [Zn2+] = [Fe2+]. The two factors of the product are, therefore, the same, for the first moment, and we may put [Fe2+] × [S2−] = [Zn2+] × [S2−] = P.
Since P is smaller than KFeS and larger than KZnS, it is clear that the solubility-product constant for zinc sulphide must be smaller than that for ferrous sulphide. The solubility-product constants, for similar salts, are a measure of their solubilities in water. We may obtain their values by determining the solubilities of salts in pure water, whenever the solubility is not affected by other chemical changes. In the present instance, the quantitative measurements, that have been made in this way, are open to question, owing to the considerable hydrolysis which sulphides, as salts of a very weak acid, undergo in solutions of such extreme dilution.408 Until such relations have been taken into account quantitatively, it is better to limit ourselves for the present to the more accessible question of relative solubility.
It is a comparatively easy matter to determine the relative solubility of zinc and ferrous sulphides. If equal quantities of the equivalent solutions are mixed and a precipitant, ammonium sulphide, which would precipitate either sulphide, if its salt were present alone, is carefully and gradually added to the mixture, it will precipitate first the less soluble one (see p. 163); and that one alone can be present permanently (i.e. in equilibrium) in contact with the solution containing the two salts. As a matter of fact,409 [p206] we find that zinc sulphide is precipitated first, under conditions permitting the precipitation of ferrous sulphide if no zinc sulphate were present, and the precipitate of zinc sulphide remains unchanged in the presence of a mixture of the composition indicated (exp.).
It is clear, therefore, that the prediction, based on the conclusion drawn from the application of the principle of the solubility-product, is verified by experiment.
Now, closer examination of the solution of zinc sulphate, from which zinc sulphide has been precipitated by the action of hydrogen sulphide, shows, after the hydrogen sulphide has precipitated as much sulphide as it can and the solution has been passed through a filter, that a very considerable proportion of zinc salt is still present in the filtrate, and we must ask why hydrogen sulphide fails to precipitate the zinc completely. The concentration of the zinc-ion has grown somewhat smaller, but that is not the cause of the nonprecipitation of zinc sulphide under the new conditions, since hydrogen sulphide will precipitate the sulphide, if it is passed into a solution of zinc sulphate which contains even a smaller concentration of zinc-ion than the filtrate, in which it fails to give any further precipitate. If the filtrate is examined, it is found to be strongly acid, since sulphuric acid has been liberated by the action of hydrogen sulphide on zinc sulphate: ZnSO4 + H2S → ZnS ↓ + H2SO4. Sulphuric acid is a strong acid, which is very highly ionized, much more so than the exceedingly weak acid hydrogen sulphide, and consequently, as the precipitation of zinc sulphide proceeds, the concentration of hydrogen-ion in the solution rapidly grows larger and larger. But, the greater the concentration of the hydrogen-ion, the smaller is that of the sulphide-ion, since the product [H+]2 × [S2−] is a constant [equation (IV), p. 201] for a solution kept saturated with hydrogen sulphide. The sulphide-ion is reduced in concentration very much more rapidly than is the zinc-ion.410 As [S2−] is a factor in the [p207] solubility-product of zinc sulphide, it is clear that the value of this product must grow rapidly smaller during the precipitation of zinc sulphide from the solution, and that it may well, eventually, grow too small to surpass the value of the solubility-product constant KZnS. Precipitation of zinc sulphide will then cease. Obviously, the suppression of the sulphide-ion may be accomplished by the addition of hydrochloric, sulphuric or any other strong acid to the zinc sulphate solution in the first place, and then hydrogen sulphide should fail to precipitate any zinc sulphide at all. In fact, if to 50 c.c. of the 0.1 molar zinc sulphate solution 2 c.c. of hexamolar hydrochloric acid is added,411 hydrogen sulphide does not precipitate even a trace of zinc sulphide (exp.).
We have found, then, that 0.1 molar zinc sulphate solution, acidified with a small excess of hydrochloric acid, fails to produce a precipitate of zinc sulphide, when it is saturated with hydrogen sulphide. We must conclude that, under these circumstances, the product of the ion concentrations is smaller than the solubility-product constant for zinc sulphide: ([Zn2+] × [S2−] / x) < KZnS, the new concentration of the sulphide-ion being represented by the symbol [S2−] / x.
It would follow, from these considerations, that the action of hydrochloric or sulphuric acid, in preventing the precipitation of zinc sulphide, depends on their producing a sufficiently high concentration of the hydrogen-ion, to keep the concentration of the sulphide-ion, in a mixture of zinc sulphate and hydrogen sulphide, below the point where the solution could become supersaturated with zinc sulphide. [p208] For exactly similar reasons, none of the sulphides of the zinc group is precipitated by hydrogen sulphide in (sufficiently) acid solutions.
It is evident, further, that, if a solution of zinc acetate (without the addition of any acid) is substituted for the zinc sulphate solution and is treated with hydrogen sulphide, an entirely different result, quantitatively considered, must be obtained. By the action of hydrogen sulphide on the acetate, acetic acid is liberated, according to Zn(CH3CO2)2 + H2S ⥂ ZnS ↓ + 2 CH3COOH. As a weak acid, acetic acid produces much less hydrogen-ion than is formed in equivalent solutions of sulphuric acid. Consequently, a much slighter suppression of the sulphide-ion and a much more complete precipitation of zinc sulphide from the acetate, than from the sulphate solution, must result. Such is the case. Zinc sulphide is, indeed, precipitated quantitatively by hydrogen sulphide from the acetate solution.
This behavior of zinc acetate412—and zinc salts of other weak acids show, of course, the same behavior—represents one of the pitfalls, into which the unwary analytical chemist is liable to fall, when he uses hydrogen sulphide. The separation of groups by hydrogen sulphide depends, as stated, on the fact, that, in the presence of a certain concentration of hydrogen-ion, hydrogen sulphide will not precipitate zinc sulphide and the remaining sulphides of the zinc group. To secure this concentration of hydrogen-ion, some hydrochloric acid is added to solutions, from which hydrogen sulphide is expected to precipitate none but sulphides of the copper and arsenic groups—and, as a rule, the purpose is accomplished, as desired. It is evident, however, that if a solution contains an acetate, say sodium acetate, or the salt of any other weak acid, e.g. a borate or a phosphate, the addition of hydrochloric acid will result, at least at first, in the liberation of the weaker acid and will not produce the excess of hydrogen-ion, required for the analysis. Zinc sulphide, and possibly nickel and cobalt sulphides,412 may, under such conditions, be precipitated with the sulphides of the groups mentioned. Unless provision is made, therefore, to insure a certain excess of hydrogen-ion (p. 213), or unless we are on our guard and look for zinc, nickel and cobalt in [p209] the analysis of the precipitate formed by hydrogen sulphide,413 serious errors obviously could result. To add an inordinately large excess of hydrochloric acid to mixtures, in order to avoid this pitfall, will, as we shall presently see, only throw us more certainly into still another error, to which we are exposed in the use of this important reagent, hydrogen sulphide.
By the conditions of the experiment we started with a concentration of the cadmium-ion, [Cd2+], equal to the concentration, [Zn2+], of the zinc-ion in the zinc sulphate solution, and with the same concentration,415 [S2−] / x, of sulphide-ion as was used when hydrogen sulphide failed to precipitate zinc sulphide. The corresponding factors of the products of the ion concentrations are equal, at the beginning of the two experiments, and we may put ([Cd2+] × [S2−] / x) = ([Zn2+] × [S2−] / x) = P′. We recall the fact, that we have already concluded, on the basis of the principle of the solubility-product, that [Zn2+] × [S2−] / x, or P′, is smaller than KZnS (p. 207), and that [Cd2+] × [S2−] / x, or P′, is greater than KCdS.
P′ being smaller than KZnS and larger than KCdS, it is clear that [p210] KCdS is smaller than KZnS and that cadmium sulphide must be the less soluble of the two sulphides. As a matter of fact, if ammonium sulphide is carefully added to a mixture of equal quantities of the two salt solutions, cadmium sulphide is precipitated first, and when practically all of the cadmium is precipitated, a final precipitate of white zinc sulphide is obtained (exp.; see note, p. 205). Or, if zinc sulphide is first precipitated by the addition of a little ammonium sulphide to 25 c.c. of the 0.1 molar zinc sulphate solution, care being taken to have zinc sulphate in excess, and if 25 c.c. of the 0.1 molar cadmium sulphate solution is then added to the mixture, the white zinc sulphide immediately gives way to the less soluble yellow cadmium sulphide (exp.; see p. 165). Cadmium sulphide is thus proved to be the less soluble of the two sulphides, a result which confirms the prediction made above with the aid of the principle of the solubility-product, and we may indeed conclude that the solubility-product constant KCdS of cadmium sulphide must be smaller than the constant KZnS of zinc sulphide (see pp. 163–168, on fractional precipitation). We are, therefore, also justified in deciding that CdS may well be precipitated from acidulated solutions by hydrogen sulphide, when ZnS is not thus precipitated, simply because KCdS is sufficiently small (CdS is sufficiently insoluble) to make the product of the ion concentrations [Cd2+] × [S2−] / x, in spite of the extremely small value of [S2−] / x, greater than the constant KCdS, whereas the same small value of [S2−] / x makes it impossible for the product [Zn2+] × [S2−] / x to reach the value of the larger constant KZnS, required for the precipitation of ZnS. Since cadmium sulphide may be precipitated quantitatively under the conditions given, it is also evident that it may be precipitated even when the concentration of the cadmium-ion also has a rather small value. The relations, in regard to this point, will be discussed presently.
Solubilities vary from salt to salt, and we have already found that, in the zinc group, zinc sulphide is less soluble than the sulphide of a second member of the group, ferrous sulphide, and that the difference is revealed in a somewhat different behavior of their salts toward hydrogen sulphide, when the action is studied in some detail. Similar differences must be expected to exist among the sulphides of the groups that hydrogen sulphide precipitates even in the presence of an excess of hydrochloric acid. As these differences are the sources of some of the most common and most serious errors which analysts are liable to commit, the detailed study of the action of hydrogen sulphide must be continued a little further.
In fact, if a large excess417 of hydrochloric acid is added to 50 c.c. [p212] of the 0.1 molar solution of cadmium sulphate, hydrogen sulphide fails to precipitate any of the sulphide (exp.).
But, if a few cubic centimeters of a 0.1 molar solution of cupric sulphate (25.0 grams of CuSO4, 5 H2O, per liter) are added to the solution from which hydrogen sulphide fails to precipitate cadmium sulphide, cupric sulphide is at once precipitated. And, if 15 c.c. of concentrated hydrochloric acid are added to 50 c.c. of the 0.1 molar cupric sulphate solution, there results a mixture corresponding to the cadmium sulphate solution from which hydrogen sulphide fails to precipitate CdS; we find that hydrogen sulphide will precipitate the sulphide of copper very readily, even under these adverse conditions (exp.). Cupric sulphide must be even less soluble in water than cadmium sulphide,418 and there is no difficulty in showing that such is the case. If ammonium sulphide, or hydrogen sulphide, is gradually introduced into a mixture of 25 c.c. each of the 0.1 molar sulphate solutions, cupric sulphide is precipitated first, and, if the precipitate is collected in fractions, [p213] pure yellow cadmium sulphide is precipitated last.419 Or, if 25 c.c. of 0.1 molar cupric sulphate is added to the mixture in which a precipitate of cadmium sulphide displaced the more soluble zinc sulphide (p. 210), the yellow sulphide will, in turn, give way to the less soluble black sulphide of copper (exp.).
We find thus that the precipitation of cadmium sulphide, by hydrogen sulphide in acid solution, can be prevented by the presence of an excess of hydrochloric acid, which does not prevent the precipitation of the less soluble cupric sulphide.420 The fact, then, that, in an analysis of some unknown mixture, hydrogen sulphide produces a precipitate in acid solution, must not be considered as evidence that the conditions are such as to insure the precipitation of all the sulphides of the groups, which we intend to precipitate. To avoid error, conditions must be such as to insure the complete precipitation of the more soluble as well as the less soluble sulphides. The sulphides of cadmium and lead, in particular, and, to a lesser degree, the sulphides of antimony and tin, are most liable to remain unprecipitated and thus escape detection in systematic analysis. This is a matter of special importance, also, in detecting traces of the ions of these metals, especially of lead, which is a slow cumulative poison, even when absorbed in minute amounts, and which analysts must therefore be able to detect, even in traces, with absolute certainty. It is clear, from a consideration of the product of the ion concentrations, as affecting the precipitation or nonprecipitation of such a sulphide, that a much smaller excess of acid will prevent the precipitation of the last traces of lead sulphide, and, therefore, of all of it, if only traces are present, than will interfere with the precipitation of the sulphide in bulk.
If an analyst aims to find even smaller quantities of a particular metal ion, e.g. traces of lead, the ordinary method of analysis [p215] may be modified, the source of error in the precipitation of traces of lead sulphide being kept in mind.424
Besides the complications mentioned, and provided against in the way discussed, there is still one more complication in the use of hydrogen sulphide: this is in the matter of the precipitation of arsenic sulphide from solutions containing arsenic in the pentavalent condition. Since the interpretation of this complication and the explanation of the methods for avoiding the errors, which may arise therefrom, are necessarily intimately connected with the chemical behavior of arsenic acid, this subject will be considered in the discussion of the arsenic group (Chapter XIII).
[395] Auerbach, Z. phys. Chem., 49, 220 (1904).
[397] Knox (in Abegg's laboratory), Trans. Faraday Soc., 4, 44 (1908).
[398] The concentration of the hydrogen-ion is really a little greater than that of the hydrosulphide-ion, as a result of the ionization of the latter, but the amount of hydrogen-ion formed in this way (about 1E−15) is so minute, compared with that formed by the primary ionization, that it is negligible.
[399] We can obtain the relation, directly, from H2S ⇄ 2 H+ + S2− and [H+]2 × [S2−] / [H2S] = K = 1.1E−22. For a given pressure of the hydrogen sulphide, [H2S], expressing its solubility (about 0.1 molar at 25°), is constant, and therefore [H+]2 × [S2−] = a constant, as given in equation (IV). Putting [H2S] = 0.1, we have [H+]2 × [S2−] = 0.1 × 1.1E−22 = 1.1E−23.
[400] On account of the great mass of water, we compare (see equation, p. 176) [H+] × [HO−] = 1.2E−14 (at 25°) with [H+] × [S2−] / [HS−] = 1.2E−15.
[401] The calculation was made by the method used by Knox (loc. cit.) for a molar solution. The degree of ionization of the salt was not considered and the correct ionization constant for ammonium hydroxide was used, 1.8E−5 in place of 2.3E−5. The latter, evidently, was used by Knox as the result of overlooking a correction, which Bredig made in his (Bredig's) first calculations of the constant; cf. Bredig, Z. phys. Chem., 13, 293, footnote. The same erroneous constant is found in Kohlrausch and Holborn, loc. cit., p. 194.
[402] For further values and for the method of calculation, see Knox, loc. cit.
[404] Knox's work leads to that conclusion.
[405] The precipitation of sulphides, from a solution containing much more of the hydrosulphide-ion than of the sulphide-ion, is comparable with the precipitation of mercuric oxide, HgO, and of silver oxide, Ag2O, by sodium or potassium hydroxide.
[406] On account of the presence of a small, unknown amount of sulphuric acid in the original solution, resulting from the hydrolysis of ferrous sulphate, the exact value of [S2−] in the first solution cannot be calculated without further examination; but, according to the values given in the table on page 202, the value of x, indicating the growth in the concentration of S2−, is at least 1012, if 2 equivalents of NaOH, and 109, if 2 equivalents of NH4OH are used to convert the 0.1 molar hydrogen sulphide into the corresponding sulphide Me2S, of 0.1 molar concentration.
[407] [S2−] is exactly the same in the two products, when equal volumes of the zinc and ferrous sulphate solutions are mixed and the mixture is saturated with hydrogen sulphide; zinc sulphide is precipitated.
[408] The difference in the values obtained, when hydrolysis is considered or neglected, is very considerable. Vide Bodländer, on the solubility of calcium carbonate, Z. phys. Chem., 35, 23 (1900), and Stieglitz, Carnegie Institution Publications, No. 107, 249 (1909).
[409] In carrying out this fractional precipitation a very dilute solution of ammonium sulphide is used, so as to prevent the mechanical enclosure of black ferrous sulphide, which would discolor the white sulphide. The ammonium sulphide solution should be saturated with hydrogen sulphide, to prevent the precipitation of green ferrous-ferric oxide by an excess of free ammonia. It is best to prepare a set of the precipitates and to preserve them in well-stoppered vessels, and not to try to take the time and care necessary to effect a perfect fractionation as a lecture experiment. The presence of the ferrous sulphate, in the supernatant liquid above the first precipitate of zinc sulphide, may be readily demonstrated by pouring off some of the solution and adding an excess of ammonium sulphide to it. Of course, it is also perfectly legitimate, and easier, to precipitate first zinc sulphide from a pure zinc sulphate solution and then to add ferrous sulphate solution to the mixture and to preserve the mixture. If the zinc sulphide were not the less soluble, it would be rapidly converted into the black ferrous sulphide. (See p. 165, and see below, pp. 210, 213, where similar transformations are carried out as lecture experiments.)
[410] When 10% of the zinc in a 0.1 molar solution has been precipitated, 0.01 molar sulphuric acid has been formed. For the sake of a rough approximation, the acid may be considered completely ionized and then [H+] = 0.02, which is 200 times the value of [H+] in a saturated H2S solution (p. 200); if the presence of a little sulphuric acid in the original zinc sulphate solution, resulting from a slight hydrolysis of the salt, is ignored, the concentration of the sulphide-ion is decreased roughly (200)2 or forty thousandfold, while the concentration of zinc-ion falls 10%. The corrections, that have been indicated, would change the quantities involved, but they would not modify the character of the result.
[411] This proportion of acid, making the concentration of the hydrogen-ion, approximately, [H+] = 0.2, is used, not because it represents the minimum concentration of the hydrogen-ion, which will prevent the precipitation of zinc sulphide in 0.1 molar zinc sulphate solution, but because it represents the practical conditions under which the precipitation of zinc sulphide is avoided, when the copper and arsenic groups are precipitated in qualitative analysis (see p. 213).
[412] Nickel and cobalt sulphides are also precipitated by hydrogen sulphide in the presence of free acetic acid, if sodium or potassium acetate is added, to suppress the hydrogen-ion of the acetic acid (p. 112).
[413] They would be found in the copper group.
[414] The sulphate, of this composition, is obtained by drying the crystallized sulphate in an air bath at 100–105°.
[417] Any immediate precipitation of cadmium sulphide will be prevented by the addition of 10 c.c. of concentrated acid (sp. gr. 1.19) to 50 c.c. of the 0.1 molar solution, and 15 c.c. will completely prevent any precipitation of the sulphide. Of course, a smaller excess would prevent the precipitation of small quantities of the sulphide (e.g. a half milligram of cadmium), which should easily be found in 50 c.c. (see p. 214).
[418] The value of the solubility-product constant for cupric sulphide, at 25°, was determined by Knox (loc. cit.): [Cu2+] × [S2−] = 1.2E−42, corresponding to a concentration of 1.1E−21 of cupric-ion. Mercuric sulphide was found even less soluble: [Hg2+] × [S2−] = 2.8E−54, and its behavior agrees with such a relation (Lab. Manual, p. 50, § 2). The solubility-product constant for lead sulphide, which resembles cadmium sulphide in the fact that a large excess of acid prevents its precipitation, was found to be [Pb2+] × [S2−] = 2.6E−15, the constant being about 1027 times as large as the constant for cupric sulphide. This value for the solubility-product constant for lead sulphide must either be considerably larger than the true value or lead must be easily precipitated as a hydrosulphide, Pb(SH)2, since solutions in which the product of the ion concentrations, [Pb2+] × [S2−], is very much smaller than the constant given, readily precipitate lead sulphide. Thus Noyes and Bray [J. Am. Chem. Soc., 29, 137 (1907)] report it possible to precipitate 1 to 2 milligrams of lead-ion in 100 c.c. of solution (say [Pb2+] = 1E−4) with hydrogen sulphide in the presence of 4 c.c. of hydrochloric acid (sp. gr. 1.12), for which, approximately, [H+] = 0.25. Then (equation (IV), p. 201) [S2−] = (1.1E−23) / (0.25)2 = 1.8E−22, and [Pb2+] × [S2−] = 1E−4 × 1.8E−22 = 1.8E−26, which is a much lower value than that given by Knox, and which still is not claimed to represent the limit of insolubility. Experiments, made in this laboratory, confirm this result and show further, that lead-ion in a concentration of 1E−5 is precipitated in the presence of 0.25 molar hydrochloric acid ([H+] = 0.22). Then [S2−] = 2.3E−22 and [Pb2+] × [S2−] = 2.3E−22 × 10−5 = 2E−27, which does not yet express the limit of insolubility.
[419] The fractions are not prepared in the lecture, but the first fraction is kept suspended in part of the solution of the two sulphates and may be kept so for years. The last fraction is kept in a separate container.
[420] A large excess of acid is liable to interfere with the precipitation of the last traces of cupric sulphide and is avoided in exact work.
[421] Noyes and Bray use, approximately, [H+] = 0.25 [J. Am. Chem. Soc., 29, 137 (1907)]. Tests in this laboratory showed that 1 milligram of cadmium-ion, or of lead-ion, in 100 c.c., is readily precipitated by hydrogen sulphide in the presence of 0.25 molar hydrochloric acid, ([H+] = 0.22).
[422] Kahlbaum's "Krystallviolett," [(CH3)2NC6H4]2C : C6H4N(CH3)2Cl, is referred to.
[423] An indelible ink pencil (violet) may, in most cases, be used in place of the solution. The details for the application of the indicator are given in the instructions for laboratory practice, Lab. Manual, pp. 31, 102, 103.
[424] See Blyth, Poisons, etc., p. 608 (1895), in regard to the detection of traces of lead.