METHOD OF AMALGAM POLAROGRAPHY WITH ACCUMULATION. II. ON THE QUESTION OF THE NATURE OF DOUBLE ANODIC PEAKS OF GALLIUM

O. S. STEPANOVA, M. S. ZAKHAROV

For the analysis of substances of a high degree of purity, in recent years the method of amalgam polarography with accumulation has become increasingly widespread. However, in a number of cases, evaluation of the quality and quantity of the impurity being determined is made difficult by the appearance of anomalous double or triple peaks of a single element. The literature contains reports on obtaining double peaks of lead [1, 2] and copper [3, 4]. In [4] the influence of various factors on the shape of the anodic copper peak in alkaline and carbonate supporting electrolytes was studied, with the aim of determining the causes of the anomalous behavior of copper amalgams during anodic dissolution in the indicated supporting electrolytes.

The present work is devoted to elucidating the causes of the appearance of double anodic peaks of gallium when recording voltammetric curves by the method of amalgam polarography with accumulation. Study of the causes and conditions for the occurrence of a double anodic peak makes it possible to prevent the formation of this peak, or else to avoid erroneous conclusions about the qualitative composition of the solution or amalgam under investigation.

Experimental Part

A self-recording Hungarian polarograph 7-77-4/B was used in the work. The polarographic cell and the electrodes did not differ from those described earlier [5]. Standard gallium solutions were prepared from a weighed portion of metallic gallium dissolved in hydrochloric acid and subsequently diluted with tridistillate to a final concentration of \(1 \cdot 10^{-5}\) g-ion/l. To prepare the supporting-electrolyte solutions, reagents of chemically pure, analytically pure, or extra-pure grade were used, containing less than \(1 \cdot 10^{-8}\) g-ion/l of gallium in the solution studied.

During one day the work was carried out with the same mercury drop (\(r = 0.04\) cm); when the height of the gallium peak deviated by more than 5%, a new mercury drop was deposited electrolytically. The electrolysis potential was \(-1.4\) V (sat. cal. electrode). At this value of the potential, the limiting diffusion current of gallium is reached in all the supporting electrolytes studied. Removal of dissolved oxygen from the test solutions and stirring of the electrolyte during electrolysis were carried out by bubbling with nitrogen purified of traces of oxygen.

The anodic peaks of gallium were obtained in supporting electrolytes of sodium salicylate acidified with hydrochloric acid to pH 3–5, acetate buffer pH 4.6, sodium nitrate and sodium chloride acidified to pH

4–5 [7]. In all supporting electrolytes (pH 3–5), when recording the voltammetric curves of gallium by the method of amalgam polarography with accumulation, in the presence of only gallium at a mercury stationary electrode not one peak is formed, corresponding to the oxidation of gallium amalgam, but two (Fig. 1). It may be assumed that the formation of the double gallium peak may be due to one of the following causes:

1) stepwise oxidation of gallium amalgam, i.e., oxidation of gallium first to ions of lower valence, which remain

Fig. 1. Investigation of the nature of the double gallium peak in a background electrolyte of 0.1 N KCl + 0.1 N NaSal. Experimental conditions: \(C_{\mathrm{Ga}} = 1 \cdot 10^{-5}\) g-ion/l; \(\tau = 5\) min. 1—double gallium peak; 2—double gallium peak. The potential was stopped after reaching the peak potential of the more negative peak; 3—anodic gallium peak after repeated recording.

in the near-electrode layer, and then the latter are oxidized to the trivalent state;

2) formation on the surface of the stationary mercury electrode of a solid film of a gallium compound.

The validity of the first assumption can be checked using coulometric measurements [4]. Studies have shown that this technique can be successfully applied only in solving analogous problems for elements when there are no other electrode processes except reduction of the metal [6]. In the cathodic reduction of gallium, however, it is practically impossible to obtain a “pure” diffusion current of the latter, since it catalyzes cathodic hydrogen evolution [6, 7].

In addition to coulometric measurements, the problem posed can also be solved qualitatively in the following way. Suppose that at the potential of the first (more negative) peak, oxidation of gallium to ions of lower valence occurs. These ions, in simple form or in the form of complex ions, in the absence of stirring of the solution may remain for some time at the electrode surface (the diffusion coefficient of gallium ions in solution is \(\sim 1 \cdot 10^{-6}\) cm\(^2\)/sec), as a result of which the concentration of gallium ions of lower valence at the surface of the mercury

of the drop is substantially greater than that in the bulk of the solution. At a more positive potential, oxidation of gallium ions of lower valence to gallium ions of higher valence may occur, and in view of their high concentration at the electrode surface a second peak may be obtained. If, however, after obtaining the first peak the change in potential is stopped and the solution is vigorously stirred in order to carry gallium ions away from the electrode surface into the bulk of the solution, then the second, more positive, peak should not appear. To test what has been stated above, when recording the anodic peaks of gallium we stopped the potential, after the potential of the peak of the first (more negative) peak had been reached, until the current fell to the background line, while simultaneously vigorously stirring the solution in order to remove gallium ions of lower valence from the surface of the mercury drop (Fig. 1). After this, a repeated polarogram was immediately recorded. On this polarogram the more negative peak was not observed, and only the more positive peak was obtained (Fig. 1). Consequently, the supposition that the more positive gallium peak is formed as a result of oxidation of gallium ions of lower valence to gallium ions of higher valence was not confirmed. It remains to suppose that both gallium peaks are due to oxidation of the gallium amalgam to the higher valence. Moreover, during the course of the process at the more negative potential, a film of gallium hydroxide is formed on the surface of the mercury electrode, which hinders further oxidation of the amalgam (the activation energy of the process increases). When the electrode potential is shifted into a more positive region, the gallium atoms acquire additional energy, the process of amalgam oxidation resumes, and a second anodic peak is obtained.

The question arises why, when the polarogram is recorded a second time (Fig. 1), only the more positive anodic gallium peak is observed. In our opinion, the absence of the more negative peak can be explained by the fact that, when the polarogram is recorded again, the gallium hydroxide film on the surface of the mercury drop does not have time to break down, since the preceding polarogram was recorded down to \(-0.6\ \mathrm{V}\) (sat. cal. el.). This sparingly soluble film on the surface of the mercury drop retards the oxidation of the gallium remaining in the amalgam (during one cycle of recording the anodic polarogram without stopping the potential after the peak is reached, up to 80% of the metal is removed from the amalgam), and therefore we do not observe the first anodic gallium peak at the more negative potential; upon further depolarization of the electrode, partial destruction of the film and further oxidation of gallium atoms to the ionic state occur.

The process of formation of a film of a sparingly soluble compound between the corresponding anions and the metal, occurring during anodic polarization of the electrode, can in general be represented by the following equations if the classical mechanism of “precipitation from a supersaturated solution” is used:

\[ \mathrm{M} \to \mathrm{M}^{q+} + qe \]

\[ \frac{n\mathrm{M}^{q+} + m\mathrm{A}^{p-} \to (\mathrm{M})_n\mathrm{A}_m}{n\mathrm{M} + m\mathrm{A}^{p-} \to \mathrm{M}_n\mathrm{A}_m \downarrow + q}. \]

The electrode process in this case consists of several stages: transport of anions from the bulk of the solution to the electrode surface; ionization of atoms of the metal of which the electrode consists (the gallium amalgam); and the chemical reaction of compound formation. A sufficiently high rate of the latter reaction, compared with the rate of removal of metal ions into the bulk of the solution, is a necessary condition for the forma-

...of deposition of the precipitate on the electrode surface. A precipitate forms at the electrode surface if the condition is satisfied:

\[ (C_M^S)^n \cdot (C_A)^m \geq L, \]

where \(C_M^S\) is the concentration of metal ions at the electrode surface;
\(C_A\) is the concentration of anions at the electrode surface;
\(L\) is the solubility product of the compound \(M_nA_m\).

It is quite clear that the smaller the value of the solubility product, the lower the concentration of the ion being determined at which the formation of a film of a sparingly soluble compound occurs, that is, passivation of the electrode. In the present case the insoluble film may be formed by gallium hydroxide, \(\mathrm{Ga(OH)_3}\). It is known [9] that gallium oxide hydrate is formed in an acidic medium already at pH \(3.3 \div 3.5\) according to the reaction

\[ \mathrm{Ga(Hg)} + 3\mathrm{OH}' + 3e \to \mathrm{Ga(OH)_3}\downarrow \]

and may exist up to pH \(9 \div 10\). At pH greater than 10 the following reaction takes place at the electrode:

\[ \mathrm{Ga(Hg)} \to (\mathrm{GaO_2})^{2--}. \]

Let us estimate, with an accuracy of up to 5%, the concentration of gallium ions at the surface of the mercury electrode \((C_{\mathrm{Ga}}^S)\) during recording of the anodic voltammetric curve. To calculate \(C_{\mathrm{Ga}}^S\), we take the pH of the supporting solution to be 4. Under these experimental conditions, double anodic peaks are formed during dissolution of the gallium amalgam.

The solubility product for gallium hydroxide at \(25^\circ\mathrm{C}\) is \(1.4 \cdot 10^{-34}\) g-ion/l [10]:

\[ L = a_{\mathrm{Ga}^{3+}}^{3} \cdot a_{\mathrm{OH}'}^{3}, \]

where \(a_{\mathrm{Ga}^{3+}}^{3}\) and \(a_{\mathrm{OH}'}\) are the activities of gallium and hydroxyl ions in the aqueous solution. At \(\mathrm{pH} = 4\), \(a_{\mathrm{OH}'} = 10^{-10}\). The activity coefficient of gallium ions under the experimental conditions was calculated from the generally accepted expression:

\[ \lg f = -\frac{0.5 Z^2 \sqrt{\mu}}{1 + \sqrt{\mu}}, \]

where \(\mu\) is the ionic strength of the solution;
\(Z\) is the valence of the ion.

The ionic strength of the solution under study was 0.1, and the activity coefficient of gallium ions was 0.09. Hence

\[ a_{\mathrm{Ga}} = f_{\mathrm{Ga}} \cdot C_{\mathrm{Ga}}^S \]

and

\[ C_{\mathrm{Ga}}^S = \frac{L}{f_{\mathrm{Ga}} \cdot a_{\mathrm{OH}'}^{3}} = \frac{1.4 \cdot 10^{-34}}{0.09 \cdot 10^{-30}} = 1.55 \cdot 10^{-3}\ \text{g-ion/l}, \]

that is, the concentration of gallium ions in the near-electrode layer must be equal to \(1.55 \cdot 10^{-3}\) g-ion/l in order for a film of gallium hydroxide to form on the surface of the mercury electrode.

Since the concentration of the metal in the mercury drop in the method of amalgam polarography with accumulation, under optimal experimental conditions, is approximately two orders of magnitude greater than the concentration of the latter in the solution under study, it becomes clear that, at a concentration of gallium ions in the solution of \(9 \cdot 10^{-5}\) g-ion/l and higher, the gallium hydroxide film on the mercury drop begins to be visible to the naked eye during recording of the anodic peaks of gallium. The drop first becomes cloudy, then turns gray, and upon further depolarization of the mercury electrode again becomes shiny.

If, however, one takes into account that near the electrode, during the cathodic process, the pH may increase even by two units, then the formation of gallium hydroxide on the electrode surface is quite realistic. As a result of the inhibition of the process of anodic dissolution by a film of gallium hydroxide, at a potential of \(-1.0\) V (sat. cal. el.) an anodic peak is obtained, the height of which in the concentration interval \(1 \cdot 10^{-6}\)—\(10^{-4}\) g-ion/l and an electrolysis time of 10 minutes remains practically constant (Figs. 2, 3).

Fig. 2. Change in the magnitude of the anodic current of gallium as a function of the concentration of the latter in a solution of \(0.1N\,KNO_3 + n1N\,HNO_3\):
1—for the more negative peak;
2—for the more positive peak.

The curves obtained are analogous; only at pH 4.6 is the anodic peak at the more negative potential somewhat smaller in height than the same peak at pH 3.2. This is quite understandable, since at pH 4.6 the solubility product of gallium hydroxide is reached at a lower concentration of gallium ions at the electrode surface than at pH 3.2, because at this pH value it is only beginning to form. At gallium concentrations of \(1 \cdot 10^{-6}\) g-ion/l in solution, the more negative peak is larger in magnitude than the more positive peak (Fig. 4). This is explained by the fact that during oxidation of gallium amalgam almost the entire amount of gallium goes to the formation of a hydroxide film, and only the small remaining part of gallium in the amalgam is oxidized to \(Ga^{3+}\) at the more positive potential and passes into solution. At gallium concentrations in solution of \(1 \cdot 10^{-7}\) M and less, we observed only one peak in the process of anodic oxidation of gallium amalgam at the more negative potential, since at these concentrations of gallium in the solution, in the near-electrode layer during the anodic cycle the solubility product is not reached and there are no kinetic hindrances to the oxidation of the latter.

Fig. 3. Change in the magnitude of the anodic current of gallium as a function of the concentration of the latter in an acetate-buffer solution:
1—for the more negative peak;
2—for the more positive peak.

Summarizing the above, the formation of double anodic peaks of gallium may be represented as follows: in the cathodic process the gallium ion is reduced to the metal and an amalgam of the latter is formed. In the anodic cycle, upon oxidation of this amalgam, a peak is formed, caused by a sparingly soluble film of gallium hydroxide on the surface of the mercury drop, which prevents further oxidation of gallium. With further anodic polarization, oxidation of gallium occurs and its ions penetrate from the amalgam through microcracks in the film,

Fig. 4. Change in the shape of the anodic peak of gallium as a function of its concentration in an acetate buffer solution:

$1—C_{\mathrm{Ga}} = 1 \cdot 10^{-6}$ g-ion/l;

$2—C_{\mathrm{Ga}} = 3 \cdot 10^{-6}$ g-ion/l;

$3—C_{\mathrm{Ga}} = 5 \cdot 10^{-6}$ g-ion/l.

Electrolysis time 10 min.

as a result of which a second anodic peak is obtained at a more positive potential.

The constancy in magnitude of the anodic peak at different pH values of the solution, caused by the film, can be explained by the fact that the kinetics of the process of formation of gallium hydroxide is determined by the concentration of hydroxyl ions in the near-electrode layer.

Conclusions

The cause of the formation of double anodic gallium peaks when recording voltammetric curves by the method of amalgam polarography with accumulation has been established. In the quantitative determination of microconcentrations of gallium in solution, it is necessary to take into account the total current of the anodic dissolution of gallium amalgam.

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