CHAPTER XIII

«THE ARSENIC GROUP. SULPHO-ACIDS AND SULPHO-SALTS»

The analytical groups, which we have heretofore discussed, contain elements, whose oxides are preëminently ‹base-forming›. The methods of separation of these groups, from each other, involve, primarily, physical[492] differences between the groups—in the matter of the relative insolubility of analogous salts. Thus, barium, strontium and calcium carbonates are precipitated, and separated from the alkalies, by means of ammonium carbonate, not because the alkalies do not form carbonates when their salts, in solution, are treated with ammonium carbonate, but wholly because barium, strontium and calcium carbonates are very difficultly, the alkali carbonates easily, soluble in water. The hydroxides of the aluminium group and the sulphides of the zinc group are less soluble than the hydroxides and sulphides of the alkaline earths and alkalies. The sulphides of the copper and the arsenic groups, again, are still less soluble than the sulphides of the zinc group, and thus the former may be precipitated by hydrogen sulphide, even when its precipitating power is reduced by the suppression of its sulphide (and hydrosulphide) ions by the addition of a strong acid.

On the other hand, the separation of the arsenic group (arsenic, antimony, tin, gold and platinum) from the copper group, with which it is precipitated by hydrogen sulphide from acid solutions, depends, essentially, on a ‹chemical› difference between the groups. The oxides, especially the ‹higher› oxides, of the arsenic group, are preëminently ‹acid›-forming; the higher oxides form such acids as arsenic acid, H_{3}AsO_{4}, antimonic acid, H_{3}SbO_{4}, stannic acid, H_{2}SnO_{3}, platinic acid, H_{2}PtO_{3}, and auric acid, HAuO_{2}. These [p243] hydroxides are, however, all more or less ‹weakly basic› in character ‹as well›. The hydroxides of the lower oxides of the metals are, as one must expect, much more strongly basic, but most of them—arsenious, antimonous and stannous hydroxides—still show sufficient acid character to be distinctly amphoteric in behavior. But, with the exception of arsenious acid, the basic ionization of the hydroxides of the lower oxides is more pronounced than their acid ionization.

The basic ionization of the hydroxides of their lower and higher oxides brings these elements into the plan of analysis for the metal or positive ions in systematic analysis. In the presence of hydrochloric acid they form chlorides, which yield positive ions in sufficient quantity[493] to allow their extremely insoluble sulphides to be precipitated by hydrogen sulphide in acid solution, together with the, likewise, very insoluble sulphides of the copper group.

The acid-forming properties of the oxides of the arsenic group are maintained in their sulphides. Again, this is especially evident in the ‹higher› sulphides. The element sulphur is substituted for the closely related element oxygen without any profound change in the chemical behavior of the compounds. Advantage is taken of this acid-forming power to separate the sulphides of the arsenic group from the sulphides of the copper group, which either are not acid-forming at all, or exhibit this property only to a very slight degree.[494]

«Sulpho-Salts.»—The similarity in the behavior of oxygen and sulphur derivatives, in this respect, is general and is not restricted to the metal sulphides we are discussing. For instance, the acid-forming power of carbon dioxide is shown also by carbon disulphide, the corresponding sulphur derivative of carbon. Just as the former combines with potassium hydroxide to form a carbonate, so carbon disulphide dissolves in solutions of potassium hydrosulphide and potassium sulphide to form potassium sulpho-carbonate:

CO_{2} + 2 KOH ⇄ K_{2}CO_{3} + H_{2}O. (1)

CS_{2} + 2 KSH ⇄ K_{2}CS_{3} + H_{2}S. (2)

CS_{2} + K_{2}S ⇄ K_{2}CS_{3}. (3)

[p244]

The higher sulphides of the arsenic group, and arsenious and antimonious sulphides among the lower ones, combine with the alkali sulphides to form soluble alkali ‹salts› of sulpho-acids, in the same way as carbon bisulphide does. In the case of arsenious sulphide, for instance, we have the action

As_{2}S_{3} + (NH_{4})_{2}S ⇄ 2 NH_{4}AsS_{2}. (4)

The salt, ammonium sulpharsenite, is ionized as follows:

NH_{4}AsS_{2} ⇄ NH_{4}^{+} + AsS_{2}^{−}.

«Sulpho-Acids.»—The free ‹sulpho-acids› are liberated, from their salts, by any stronger acid, such as hydrochloric acid; the acids are extremely unstable and revert rapidly to the sulphides, from which their salts were originally obtained. We have, for instance,

2 HAsS_{2} ⥂ As_{2} S_{3} ↓ + H_{2}S ↑.

This instability of the sulpho-acids is entirely analogous to the instability of the metal hydrosulphides (p. 203) and to the instability of certain oxygen acids, notably of carbonic acid. Sulpho-carbonic acid, H_{2}CS_{3}, is the best-known free acid of this type. It may be precipitated, undecomposed, as an oil, and its gradual decomposition into hydrogen sulphide and carbon disulphide may be observed.

«Sulpho-Bases.»—Since one meets, in this group, ‹sulpho-acids› and ‹sulpho-salts›, corresponding to the oxygen acids and their salts, one is naturally led to inquire,[495] whether the third great class of oxygen compounds is not found duplicated among sulphur compounds, whether ‹sulpho-bases›, as well as sulpho-acids and salts, are known. One would look for such bases, with the most pronounced basic character, among the compounds obtained by the substitution of sulphur for oxygen in the strongest oxygen bases. To a certain extent potassium and sodium hydrosulphides and sulphides show, in fact, properties, which are akin to those fundamentally characteristic of ordinary bases. They combine with acid sulphides to form sulpho-salts, as the oxygen bases combine with acid oxides. To a very considerable extent they neutralize all but the very weakest acids; hydrogen sulphide, itself a very weak acid, is driven out of its salts by all stronger acids, and the latter are almost completely neutralized. For instance, we have: [p245] KOH + HCl ⥂ KCl + H_{2}O, and KSH + HCl ⥂ KCl + H_{2}S and K_{2}S + 2 HCl ⥂ 2 KCl + H_{2}S. The concentration of the hydrogen-ion is reduced most decidedly in each of these reactions.

EXP.—A solution of potassium hydrosulphide is saturated with hydrogen sulphide, in order to prevent hydrolysis and the formation of potassium hydroxide[496] (p. 180), as far as possible, and the solution is added to an acid (hydrochloric) solution of methyl orange;[497] the acid color is changed to orange, as a result of the almost complete neutralization of the acid. The potassium hydrosulphide (the hydrosulphide-ion HS^{−}) neutralizes the hydrogen-ion (of hydrochloric acid), that converts methyl orange into its pink salt, and hydrogen sulphide is formed, which is too weak an acid to affect the color of the indicator (p. 79).

The objection that potassium sulphide and hydrosulphide are salts, the salts of hydrogen sulphide, might be raised against the conception of their possessing a certain measure of basic functions; but the common oxygen base, potassium hydroxide, is also a salt, the salt of a still weaker acid, water. Indeed, the characteristic properties of ordinary bases are due essentially to the fact, that they are the more or less readily ionizable salts of an extremely weak acid, water, and these properties may well be duplicated by salts of other ‹weak› acids, duplicated in a ‹very much weaker› way, in proportion as the acids are stronger than water. The difference is, then, really one of ‹degree› and not of kind.[498]

Owing to the fact that hydrogen sulphide is a much stronger acid than water, the action of potassium hydrosulphide on an acid sulphide, like carbon disulphide (equation (2), p. 243), is reversed to a correspondingly greater degree than the action of potassium hydroxide on carbon dioxide[499] (equation (1), p. 243). The dissociation constant for the secondary ionization of hydrogen sulphide (HS^{−} ⇄ H^{+} + S^{2−}) is very much smaller than the constant for the primary ionization (HS^{−} is a much weaker acid than HSH), and so we find that a sulphide like K_{2}S exhibits very much stronger basic functions than do the hydrosulphides, as, for instance, in forming salts with acid-forming sulphides [p246] (equation (3), p. 243) and in neutralizing acids. There can be no question that, if we could have an aqueous solution of potassium oxide, K_{2}O, it would show, similarly, the characteristic actions of strong bases ‹even more powerfully› than the hydroxide, KOH; for instance, in acting on acid-forming oxides (equation (1), p. 243), in neutralizing acids, in saponifying esters (p. 81), and so forth. It is, in fact, on account of this property, that potassium oxide is decomposed by water. It is a salt involving the ‹secondary ionization› of water, (HO^{−} ⇄ H^{+} + O^{2−}), which has a much smaller dissociation constant even than the primary ionization (H_{2}O ⇄ H^{+} + HO^{−}). The oxide, K_{2}O, is decomposed by ‹neutralizing hydrogen ions formed by the primary ionization of water›. We have 2 K^{+} + O^{2−} + H^{+} + HO^{−} ⥂ 2 K^{+} + 2 HO^{−}, which is entirely analogous, in principle, to K^{+} + HO^{−} + H^{+} + Cl^{−} ⥂ K^{+} + Cl^{−} + HOH.

«Sulphoxy-Salts.»—The close relations between the oxygen and the sulphur series are seen also in the fact that an oxygen base may be combined with an acid sulphide, and ‹vice versa›; arsenious sulphide, for instance, dissolves even in the solution of so weak a base as ammonium hydroxide (‹exp.›). The salts produced by this "crossing" are usually "hybrid" salts, partly sulpho-, partly oxygen-salts. There is, for instance, a series of ‹arseniates›,[500] Me_{3}AsO_{4}, Me_{3}AsSO_{3}, Me_{3}AsS_{2}O_{2}, Me_{3}AsS_{3}O and Me_{3}AsS_{4}. In analytical work the pure types are ordinarily utilized, rather than the mixed types.

«Complex Sulphide Ions.»—The ions of the sulpho-acids, like the ions of oxygen-acids (p. 238), may also be treated as ‹complex ions›—of the positive metal ions and the sulphide-ion, S^{2−}. Ammonium sulphide combines with stannic sulphide, forming ammonium sulphostannate: SnS_{2} + (NH_{4})_{2}S ⇄ (NH_{4})_{2}SnS_{3}. If the action is considered to be the result of interactions of the ions of stannic and ammonium sulphides, we can resolve the equation into the following one: Sn^{4+} + 2 S^{2−} + 2 NH_{4}^{+} + S^{2−} ⇄ 2 NH_{4}^{+} + SnS_{3}^{2−}.

The ammonium-ion, appearing with the same coefficient on both sides of the last equation, evidently takes no direct part in the action and we have more simply: Sn^{4+} + 3 S^{2−} ⇄ SnS_{3}^{2−}.

For the condition of equilibrium between the complex and its components we have:[501][502]

[Sn^{4+}] × [S^{2−}]^3 / [SnS_{3}^{2−}] = K.

[p247]

«Sulphurization of Sulphides.»—Since the solubility of the arsenic group of sulphides in ammonium sulphide solution—the reagent commonly used in analysis—depends on the formation of soluble sulpho-salts, due consideration must be taken of the fact that some of the ‹lower› sulphides—notably stannous, aurous and platinous sulphides—do not possess acid-forming properties in any marked degree; even antimonous sulphide is soluble only in a considerable excess of ammonium sulphide. In order, then, to insure a more complete separation of the arsenic from the copper group, precautions are taken to sulphurize the lower sulphides to higher, stronger acid-forming, sulphides, in the course of the separation. For this purpose so-called "yellow" ammonium sulphide, containing persulphides of ammonium, (NH_{4})_{2}S_{2}, etc., is used in place of a solution of ammonium sulphide and ammonium hydrosulphide. Stannous sulphide, for instance, is dissolved by the reagent as ammonium sulphostannate: SnS + (NH_{4})_{2}S_{2} ⇄ (NH_{4})_{2}SnS_{3}.

«Behavior of Arsenic Acid toward Hydrogen Sulphide.»—In conclusion, special consideration must still be given to the behavior of arsenic acid, H_{3}AsO_{4}. As its name indicates, it is an acid, and, in fact, a rather strong acid, of the order of strength[503] of phosphoric acid, H_{3}PO_{4}, which it resembles in composition and in many of its properties. As a strong acid, arsenic acid, when it is ionized, yields chiefly negative arseniate ions, H_{2}AsO_{4}^{−}, HAsO_{4}^{2−} and AsO_{4}^{3−}. Any basic properties, which it may and, most likely, does possess, must be extremely weak. It is, therefore, not surprising to find that unusual difficulties are experienced in precipitating arsenic sulphide, by hydrogen sulphide, from arseniate [p248] solutions, in as much as hydrogen sulphide is an agent for the precipitation of the sulphides of ‹cations›. Arsenious acid, the hydroxide of the lower oxide, on the other hand, is a much weaker acid and shows more pronounced basic (amphoteric) properties, and arsenic trisulphide is precipitated, without difficulty, from solutions of arsenious acid in hydrochloric acid. We have: 2 As^{3+} + 3 S^{2−} ⥂ As_{2}S_{3} ↓. When arsenic acid is reduced to arsenious acid by sulphurous acid, by iodides (see Chap. XVI) or by hydrogen sulphide (see below), no further difficulty in precipitating a sulphide (As_{2}S_{3}) is experienced.

When a solution of arsenic acid, containing the usual small amount of hydrochloric acid (0.3 molar), is treated with hydrogen sulphide at ordinary temperatures, the following three reactions take place, but ‹exceedingly slowly›:

2 H_{3}AsO_{4} + 5 H_{2}S ⥂ As_{2}S_{5} ↓ + 8 H_{2}O (1)[504]

H_{3}AsO_{4} + H_{2}S ⥂ H_{3}AsO_{3} + S ↓ + H_{2}O (2)[505]

2 H_{3}AsO_{3} + 6 HCl + 3 H_{2}S ⇄ 2 AsCl_{3} + 3 H_{2}O + 3 H_{2}S ⥂ As_{2}S_{3} ↓ + 6 HCl + 3 H_{2}O (3)

Even in the presence of a considerable amount of arsenic acid, precipitation, either of the trisulphide or of the pentasulphide, may not occur for some time, and, unless one takes account of that fact, the dangerous element, arsenic, would ‹easily› be ‹overlooked›. ‹Heat accelerates› both the precipitation of the pentasulphide and the reduction of arsenic acid and the subsequent precipitation of arsenic trisulphide.[506]

The interesting observation has also been made that, in the presence of an unusually ‹large excess› of hydrochloric acid and of a rapid stream of hydrogen sulphide, the precipitation of the ‹pentasulphide› (equation (1)) is favored and accelerated.[507] For instance, if 100 c.c. of concentrated hydrochloric acid (sp. gr. 1.2) are added to 50 c.c. of a 0.1 molar solution of potassium arseniate and a rapid stream of hydrogen sulphide is passed through the mixture at the ordinary temperature, a copious precipitate is formed within a minute (‹exp.›). The precipitate formed under these conditions [p249] is the ‹pentasulphide›.[508] On the other hand, a mixture of 5 c.c. of hexanormal hydrochloric acid and 50 c.c. of 0.1 molar potassium arseniate fails, for a long time, to give a precipitate when treated in the same way (‹exp.›).

The acceleration of the precipitation of the pentasulphide by the presence of a large excess of hydrochloric acid forms a problem of peculiar interest and importance, and no complete explanation of it has yet been offered.[509] The following considerations lead to one explanation, that has been suggested. Arsenic acid, by virtue of its close relations to antimonic, stannic and arsenious acids, may be assumed to have extremely weak basic, as well as pronounced acid, properties. For its ionization, we would have 3 H^{+} + AsO_{4}^{3−} ⇄ H_{3}AsO_{4} (+ H_{2}O) ⇄ As(OH)_{5} ⇄ As^{5+} + 5 HO^{−}. Further, the precipitation of As_{2}S_{5} may be assumed to result, ultimately,[510] from the action of the sulphide-ion S^{2−} on the positive ion As^{5+} (2 As^{5+} + 5 S^{2−} ⇄ As_{2}S_{5} ↓). The favorable action of the hydrochloric acid might, consequently, be thought to result from the fact, that it ‹facilitates› the ‹ionization of arsenic acid› as a base and the formation of a salt[511] AsCl_{5}. It could thus greatly increase the concentration of the ion As^{5+} and facilitate its combination with the sulphide-ion.

Treatment of a solution of arsenic acid with a concentrated acid, yielding a large concentration of hydrogen-ion, would carry the series of actions, represented in the above ionization equation for arsenic acid, decidedly toward the right—suppressing the arseniate-ion AsO_{4}^{3−} and increasing the concentration of the arsenic-ion As^{5+}. Since we cannot apply the equilibrium laws (or the principle of the solubility-product) to solutions as concentrated as the one under discussion, a quantitative theoretical treatment of the subject cannot be given. The following may be suggested: The action of the acid would be favorable to the precipitation of As_{2}S_{5} by suppressing the arseniate-ion AsO_{4}^{3−} and thus increasing the concentration of the hydroxide As(OH)_{5}, available for ionization as a base and for the production of the ion As^{5+}. But the further favorable effect of the hydrochloric acid, in converting the hydroxide into a salt AsCl_{5} and increasing thereby the concentration of As^{5+}, would be ‹very largely offset› by the action of the acid in suppressing the [p250] sulphide-ion ([S^{2−}] = ‹k› / [H^{+}]^2; see p. 201). For systems to which the equilibrium laws could be applied, the concentration of As^{5+} (except for the suppression of the ion AsO_{4}^{3−}) would grow, approximately, with the fifth power[512] of the concentration of the hydrogen-ion, and the concentration of the sulphide-ion would decrease, approximately, proportionally to the square of the concentration of the hydrogen-ion. Further, the precipitation of As_{2}S_{5}, in a system to which the principle of the solubility-product were applicable, would depend on the relation of the product [As^{5+}]^2 × [S^{2−}]^5 to the solubility-product constant; it is evident that the value for [As^{5+}]^2 would ‹increase› proportionally to the tenth power of [H^{+}] and the value of [S^{2−}] ‹decrease› proportionally to the tenth power of the same factor [H^{+}]. The two effects would consequently offset each other under such conditions. However, the equilibrium laws cannot legitimately be applied to such concentrated solutions and the relation has been developed only to indicate opposing factors, which must be taken into account. An experimental study of the problem would be extremely interesting.[513] Since it involves the question of the ‹minute basic ionization of a moderately strong acid› (H_{3}AsO_{4}), which may be open to ‹measurement› (see Chap. XVI), the problem is one of particular interest and importance.

The analytical precautions, taken to insure the precipitation, by hydrogen sulphide, of arsenic sulphide, when arsenic is present in quinquivalent form, are based on the observations described; in quantitative analysis, for the sake of securing a precipitate of ‹uniform composition›, the aim is to precipitate the pure ‹pentasulphide› and a considerable excess of hydrochloric acid is used. In qualitative analysis, where the composition of the precipitate is a matter of indifference and a large excess of acid would seriously interfere with the precipitation of certain sulphides (‹e.g.› CdS, see p. 211), a smaller excess of acid is used and the precipitation of arsenic sulphide is insured by prolonged treatment of a solution with hydrogen sulphide ‹at a high temperature›.

FOOTNOTES:

[492] The ‹weak basic properties› of the hydroxides of the aluminium group, as compared with those of the zinc group, ‹a chemical difference›, and the resulting great instability of the carbonates of the former group, are used in the separation of the aluminium from the zinc group, by barium carbonate; but the physical element of extreme insolubility of the trivalent hydroxides enters also as an important factor (see footnote 3, p. 194).

[493] See p. 247, in regard to the behavior of arsenic acid in this respect.

[494] See p. 246, footnote 3, in regard to the action of sodium sulphide on mercuric and bismuth sulphide.

[495] ‹Vide› Nilson, ‹J. prakt. Ch.›, «14», 150 (1876).

[496] Such a solution does not react alkaline to phenolphthaleïn.

[497] Hydrogen sulphide rapidly destroys the indicator and the experiment is best carried out by preparing 50 c.c. of a saturated aqueous solution of hydrogen sulphide, containing 1 or 2 c.c. of normal hydrochloric acid, and by adding a considerable excess of methyl orange to the solution immediately before the addition of potassium hydrosulphide solution, which has been prepared as described in the text.

[498] See the discussion on p. 177. See also the discussion by Remsen on acidic and basic halides, ‹Am. Chem. J.›, «11», 300 (1889) Stud.

[499] In both cases acid salts, KHCO_{3} and KHCS_{3}, are also formed.

[500] McCay, ‹Z. anorg. Chem.›, «29», 36 (1901).

[501] On p. 238 the analogous equation for the condition of equilibrium of the anion of an oxygen acid with its components was developed. Applying the result to the ion SnO_{3}^{2−} of stannic acid, H_{2}SnO_{3}, we have:

[Sn^{4+}] × [O^{2−}]^3 / [SnO_{3}^{2−}] = K.

It is evident, from the form of the equation, that for the stronger oxygen acids, ‹which are most stable as acids and ionize as bases at most in traces›, the value of the constant must be extremely small.

[502] Mercuric sulphide is somewhat soluble in potassium and sodium sulphides, forming the salts Me_{2}HgS_{2}, and the complex ion HgS_{2}^{2−}. A liter of 0.1 molar Na_{2}S dissolves, at 25°, 1.9 grams (0.0082 mole) of HgS [Knox, ‹Trans. Faraday Society›, «4», 36 (1908)]. While the oxide (hydroxide) shows no perceptible tendency toward acid ionization, mercuric salts, it will be recalled, show in many cases an abnormally small tendency to form the mercuric-ion (see p. 115), and the latter also shows a particularly great tendency towards forming very stable complex ions of all kinds (‹e.g.› HgI_{4}^{2−}, in K_{2}HgI_{4}, HgCl_{4}^{2−} in K_{2}HgCl_{4}, Hg(CN)_{4}^{2−}, etc.). Knox found for [Hg^{2+}] × [S^{2−}]^2 / [HgS_{2}^{2−}] = ‹k›, the approximate value of ‹k› to be 1 / 10^{53}. Bismuth sulphide is also very sparingly soluble in sodium or potassium sulphide, but not in ammonium sulphide. Solid salts, KBiS_{2} and NaBiS_{2} are known [Knox, ‹J. Chem. Soc.› (London), «95», 1760 (1909)].

[503] See the table, p. 104.

[504] See the discussion of the reaction, given below.

[505] See Chap. XVI, for the interpretation of the reduction as an ionic reaction.

[506] Bunsen, ‹Ann.› (Liebig), «192», 305 (1878). Brauner and Tomicek, ‹J. Chem. Soc.› (London), «53», 145 (1888). Usher and Travers, ‹ibid.›, «87», 1370 (1905).

[507] Neher, ‹Z. anal. Chem.›, «32», 45 (1893).

[508] Neher, ‹loc. cit.›

[509] The theory of the relations favoring the precipitation expressed in equation (1) as against the reduction expressed in equation (2), forms a second interesting problem.

[510] Intermediate derivatives, such as H_{3}AsSO_{3} (p. 246), could be the result of the ionization of As(OH)_{5}, or of AsCl_{5}, in stages (see p. 106). Neher, ‹loc. cit.›, McCay, ‹loc. cit.›

[511] Neher (‹loc. cit.›) suggested that the favorable action of the large excess of hydrochloric acid might well be due to the formation of AsCl_{5}. McCay, ‹J. Am. Chem. Soc.›, «24», 661 (1902), discusses the ionization of arsenic acid as a base, in connection with the precipitation of As_{2}S_{5}.

[512] As(OH)_{5} is considered to be an ‹extremely weak› base and AsCl_{5} to be an ionizable salt.

[513] To a certain extent, the effect of the acid may be to coagulate and precipitate the ‹colloidal› sulphide. Possibly, also, the concentrated acid renders inactive a considerable portion of the water present (forming oxonium salts OH_{3}Cl, etc., see p. 238), which tends, by hydrolysis, to reverse the formation of the chloride As(OH)_{5} + 5 HCl ⇄ AsCl_{5} + 5 H_{2}O. Possibly, the formation of the pentasulphide is not wholly an ‹ionic reaction›, its precipitation being always a more or less ‹slow process›, and there may be intermediate products whose formation could be ‹accelerated› by the presence of acids (see Bredig and Walton, ‹Z. Elektrochem.›, «9», 114 (1903) for the study of a simple inorganic action involving such ‹catalytic› effects of acids).

[p251]