ON THE MIGRATION OF GOLD IN AQUEOUS MEDIA IN THIOSULFATE FORM

L. L. SHABYNIN

(Presented by Prof. P. A. Udodov)

From the chemistry of gold [23] it is known that gold dissolves only under strictly definite conditions, the principal ones being the following:

1. The presence in solutions of a complex-forming agent capable of lowering the concentration of free gold ions as a result of binding them into a sufficiently stable soluble complex.

2. The presence in solutions of an oxidizing agent which, under these conditions, can accept electrons from the dissolving gold.

Of the sulfur compounds formed in nature during the oxidation of sulfides (hydrogen sulfide, hydrosulfide, diatomic sulfur, thiosulfate, polythionates, sulfate), most are capable of forming complex ions with gold. The sulfate ion is not among the latter, and on this basis a conclusion has been drawn regarding the impossibility of dissolving gold by sulfuric acid and its salts [20].

Of the indicated sulfur compounds, the thiosulfate ion is of greatest interest; it is stable over a considerable range of pH fluctuations: from 4.6 to 10.0 [18]. Owing to its very strong reducing capacity, it forms stable complex compounds with metals in a lower valence state. The tendency of this ion toward complex formation exceeds that of \(\mathrm{SCN}^{-}\), \(\mathrm{NO}_{2}^{-}\), \(\mathrm{J}^{-}\) [12]. At present, only a small group of thiosulfate complexes with metals of the platinum group—copper, silver, gold, selenium, and tellurium—has been well studied [12, 13, 14].

The dissolution of gold and silver by aqueous solutions of sodium thiosulfate under ordinary conditions showed that in 4 hours, 2.48 mg of gold and 101.3 mg of silver passed into the dissolved state. Gold and silver passed into solution with the formation of the complexes \([\mathrm{Au}(\mathrm{S}_{2}\mathrm{O}_{3})_{2}]^{3-}\) and \([\mathrm{Ag}(\mathrm{S}_{2}\mathrm{O}_{3})_{2}]^{3-}\), according to the scheme

\[ \mathrm{Au} + 2\mathrm{S}_{2}\mathrm{O}_{3}^{-} \rightarrow [\mathrm{Au}(\mathrm{S}_{2}\mathrm{O}_{3})_{2}]^{3-} + e^{-} \]

[18].

From the standpoint of thiosulfate migration, the difference in the behavior of gold and silver in oxidation zones, noted in the literature [10, 15], is well explained. The point is that decomposition of the thiosulfate complex of gold proceeds according to the scheme [19]

\[ 2[\mathrm{Au}(\mathrm{S}_{2}\mathrm{O}_{3})_{2}]^{3-} + 4\mathrm{H}^{+} \rightarrow 2\mathrm{Au} + 4\mathrm{S} + \mathrm{SO}_{4}^{2-} + 3\mathrm{SO}_{2} + 2\mathrm{H}_{2}\mathrm{O}, \]

and of silver according to the scheme

\[ 2[\mathrm{Ag}(\mathrm{S}_{2}\mathrm{O}_{3})]^{3-} + 4\mathrm{H}^{+} \rightarrow \mathrm{Ag}_{2}\mathrm{S} + 3\mathrm{S} + \mathrm{SO}_{4}^{2-} + 3\mathrm{SO}_{2} + 2\mathrm{H}_{2}\mathrm{O}. \]

Silver sulfide, having slight solubility, can nevertheless be carried by waters farther than the place where gold is precipitated.

It is interesting to note the fact that the aggregate state of gold particles does not affect their dissolution by sulfur compounds with the formation of complex ions [7].

Thus, experimental [12, 13, 14, 16, 17, 18, 19, 20] and calculated data show the full possibility of gold dissolution under natural conditions.

Unfortunately, in the literature available to the author, no information has been found on the determination of the thiosulfate ion in the waters of oxidation zones. There are only fragmentary data on the determination in the waters of a number of thermal springs of subsequent products of thiosulfate oxidation—polythionates [8], as well as data on the determination directly of thiosulfates in the waters of the Black Sea [2]. Therefore, let us try by an indirect route to demonstrate the possibility of the formation and existence of thiosulfate in the oxidation zones of sulfide deposits.

I. I. Ginzburg [4, 5], on the basis of experimental studies, showed that during the oxidation of pyrrhotite, especially in distilled water and in a millinormal alkali solution, during the first two to three months a large excess of sulfur over iron was observed in the solution. The smaller content of iron in the solution compared with sulfur was explained not only by its precipitation from the solution, but also by the presence of thiosulfate and polythionates. During the oxidation of pyrite the opposite picture was observed: the predominance in solution of iron over sulfur. However, in this case as well there was a difference in the solution between the content of total sulfur and sulfate sulfur, due to the presence of thiosulfate and polythionates. The existence of the latter was also noted during the oxidation of sphalerite.

A good illustration of the formation of significant quantities of incompletely oxidized forms of sulfur in the first moments of sulfide oxidation is provided by the results of experimental work by T. I. Drobasheva and S. A. Durov [6]. According to their data, during the oxidation of natural sulfides in an aqueous medium under experimental conditions, the difference between total sulfur and sulfate sulfur reaches considerable values. Thus, during the oxidation of a mixture of pyrrhotite and sphalerite, this difference amounts to more than 37 mg-eq when recalculated to the sulfate ion.

V. M. Levchenko [11], studying the kinetics of oxidation of sodium sulfide, came to the conclusion that this process proceeds in three stages.

1. Accumulation of intermediate products of sulfur oxidation with slight oxidation of sulfide to sulfate.

2. Oxidation until the complete disappearance from the solution not only of sulfides, but also of intermediate products, with accumulation of elemental sulfur.

3. Oxidation of sulfur to sulfate.

It should be noted that the thiosulfate content under the experimental conditions [11] increased during the first 33 days.

It is also necessary to say a few words about the time of existence in solution of incompletely oxidized forms of sulfur, in particular thiosulfate. Experimental studies have shown [5] that, regardless of the pH of the medium, in the first moments of sulfide oxidation in solutions there exists a greater or lesser amount of incompletely oxidized sulfur. The latter is completely oxidized, in the case of pyrite and pyrrhotite, after prolonged standing of the solution (4–6 thousand hours). During the oxidation of sphalerite, however, these forms exist even after the solution has stood for 12 thousand hours. If this time is compared with the time for dissolution of 2.5 mg of gold over 4 hours in Turin’s experiment [18], then one may with a certain degree of probability assume the dissolution of gold by the thiosulfate ion in the oxidation zones of gold-sulfide deposits.

Table 1

In oxidation zones, the pH of solutions is determined to a significant extent by the mineralogical composition of the sulfides. In character these solutions approximate very dilute sulfur-bearing solutions. Their pH values generally range from 2.0 to 9.0, and the redox potentials from 0.20 to 0.86 V [1]. It should also be noted that the redox potentials of waters in the zone of oxidized ores are considerably higher than these values for the zone of primary ores [21, 22].

A number of experimental works [4, 5, 6] on the oxidation of sulfides made it possible to assume that the pH and redox potentials of solutions under experimental conditions, in addition to the reaction of the reversible oxygen electrode

\[ 2\mathrm{H}_2\mathrm{O} \rightleftarrows 4\mathrm{H}^+ + \mathrm{O}_2 + 4e;\quad E^0 = 1.229\ \mathrm{V}, \]

also depend on the form in which sulfur occurs in solution, the cationic composition of the primary sulfides, the chemical properties of the oxidation products, and mechanical impurities in the sulfides.

V. M. Levchenko [11], in accordance with the three distinguished stages of oxidation of sodium hydrosulfide (see above), showed the change in redox-potential values. Thus, in the first stage of accumulation of incompletely oxidized sulfur, the redox-potential value lay within the limits of \((-0.08)\) V–\((-0.175)\) V. The pH value is not given in this case, but sodium hydrosulfide gives an alkaline medium upon dissolution. In the second stage of sulfide oxidation, the redox-potential value rises to 0.660 V. With further oxidation and accumulation of the sulfate ion in solution, the acidity of the solution increases and the value of \(Eh\) decreases (to 0.440 V).

Fig. 1. Dependence of \(Eh\) on pH at total sulfur activity \([\Sigma S] = 10^{-5}\).

The author has calculated the values of isobaric-isothermal potentials for most sulfur oxidation reactions given by Garrels [3] (Table 1). In considering the table, one should recall the well-

…a rule known in chemistry, namely that the order in which individual reactions proceed in a mixture is determined not by their rates, but by the magnitudes of the free energies of these reactions [9]. Examination of the table shows that the free energy of formation of the sulfate ion in all cases exceeds, by a small amount, the free energy of formation of the thiosulfate ion. It would seem that oxidation of divalent sulfur should proceed directly to the sulfate ion. Here, however, the course of the reactions is affected by the values of the redox potentials and by pH. Table 1 also gives the author’s calculated dependences of the redox potentials on pH for these reactions. The total activity of sulfur (\([\Sigma S]=10^{-5}\) and \([\Sigma S]=10^{-2}\ \mathrm{mole/l}\)) adopted in the calculations corresponds to contents of tens of milligrams and hundreds of milligrams of sulfate ion per liter. The first value was taken because we carried out field investigations at gold–quartz–sulfide deposits whose waters contain a small amount of sulfates. The high value of activity was taken in order to show that an increase in the sulfur-containing components in solution changes the magnitude of the redox potentials only slightly (hundredths, rarely tenths, of a volt). The dependences presented are expressed as equations of straight lines, which are plotted on a diagram in \(Eh\)—pH coordinates (Fig. 1). The diagram shows that incompletely oxidized forms of sulfur of the type \(S_mO\) and \(S_mO_n\) require, for their complete oxidation, higher values of the redox potential than does the direct oxidation of divalent sulfur to the sulfate ion.

Thus, the calculated data confirm the possibility that intermediate forms of sulfur may exist in solutions during the oxidation of sulfides.

In summary, it may be noted that the presence of thiosulfates and other intermediately oxidized forms of sulfur in natural waters, in particular in waters of oxidation zones, is confirmed by isolated field determinations [2, 8], experimental studies [3, 4, 5, 6, 11], and theoretical calculations. On the basis of this material and of experimental work on the dissolution of gold by the thiosulfate ion [18] and on the formation of a stable gold–thiosulfate complex [14], one may assume, among other forms, a thiosulfate form of migration of gold in waters.

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