DEPENDENCE OF THE AMMONIUM AMALGAM EFFECT ON THE pH OF THE SOLUTION

A. V. KONKOVA

When applying the ammonium amalgam effect (AAE) in order to increase the sensitivity of analytical determination of microimpurities by the method of amalgam polarography with accumulation, it is necessary to know what factors influence the AAE. The aim of the present work was to study the influence of the solution pH on the AAE.

The ammonium amalgam effect is possible when some ammonium salt (or its hydroxide) is present in the solution; therefore, it is necessary to take into account both the dissociation and the hydrolysis of ammonium salts. As is known, in aqueous solutions of ammonium salts there exists the equilibrium

\[ \mathrm{NH}_4^+ + \mathrm{H}_2\mathrm{O} \rightleftarrows \mathrm{NH}_4\mathrm{OH} + \mathrm{H}^+ . \tag{1} \]

During electrolysis of such a solution under conditions of formation of ammonium amalgam, both types of ions are reduced at the mercury cathode—the ammonium ion and the hydrogen ion. In this process hydrogen is removed into the atmosphere in the form of gas, since it is practically insoluble in mercury [1, 2] and only very slightly soluble in water [3]. Ammonium, in the form of a radical, partly penetrates into the mercury with formation of an amalgam and partly interacts with water with formation of \(\mathrm{NH}_4\mathrm{OH}\) and hydrogen. In turn, \(\mathrm{NH}_4\mathrm{OH}\) dissociates with formation of the \(\mathrm{NH}_4^+\) ion. Thus ammonium (more precisely, the greater part of it) as it were undergoes a cyclic process in the solution, whereas hydrogen is continuously removed from the solution. In this connection, the equilibrium according to equation (1) should be shifted to the right, toward the formation of \(\mathrm{NH}_4\mathrm{OH}\). Both processes considered (a decrease in the concentration of hydrogen ions in solution and an increase in the concentration of base) cause an increase in the alkalinity of the solution.

Let us consider equilibrium (1) in a 0.1 N solution of \(\mathrm{NH}_4\mathrm{Cl}\), taking into account hydrolysis of the salt and dissociation of the base. The degree of hydrolysis in a 0.1 N solution of \(\mathrm{NH}_4\mathrm{Cl}\) at \(25^\circ\mathrm{C}\) is 0.01% [4], i.e., hydrolysis in a solution of this concentration is so insignificant that it can practically be neglected. As is known, the concentration of the ammonium salt \((C_{\mathrm{NH}_4\mathrm{Cl}})\) and of ammonium hydroxide \((C_{\mathrm{NH}_4\mathrm{OH}})\) is related to the pH of the solution and to the dissociation constant \((K_{\mathrm{NH}_4\mathrm{OH}})\) by the following relation:

\[ \lg C_{\mathrm{NH}_4\mathrm{OH}} = \lg C_{\mathrm{NH}_4\mathrm{Cl}} + \mathrm{pH} - 14 - \lg K_{\mathrm{NH}_4\mathrm{OH}} . \tag{2} \]

According to this equation, the value of \(C_{\mathrm{NH}_4\mathrm{OH}}\) in a 0.1 N solution of \(\mathrm{NH}_4\mathrm{Cl}\) at \(\mathrm{pH} = 5\) is \(5.6 \cdot 10^{-6}\) g-mol/l, i.e., very small in comparison with \(C_{\mathrm{NH}_4\mathrm{Cl}}\). Such a mixture, of course, does not possess buffer properties, ...

possesses. However, in the course of electrolysis, as was shown above, the difference between the concentration of ammonium ions and the concentration of molecules of its hydroxide in one and the same solution decreases. In connection with this, the buffer capacity of the ammoniacal mixture should increase and, with sufficiently prolonged electrolysis, become sufficient to maintain a constant pH.

In view of the comparatively small volumes of solution in which the AAE was studied (10–15 ml), and also in view of the vigorous stirring of the solution by a stream of gaseous nitrogen, one may apparently speak of a change in pH not only in the cathode layer but also throughout the entire volume of the solution.

We carried out experiments to study the change in pH during electrolysis of \( \mathrm{NH_4Cl} + \mathrm{NH_4OH} \) solutions with different initial pH values. The pH value of the solution was measured at definite intervals during electrolysis (after additional stirring of the solution) with glass and calomel electrodes on an LP-58 pH meter. The results of these experiments are presented in Fig. 1 and testify to the correctness of the considerations set forth. As was to be expected, in the less acidic solution (curve 1) pH constancy is attained after a shorter electrolysis time; for the more acidic solution (curve 2), more time is required for this.

Fig. 1. Dependence of the pH of an \( \mathrm{NH_4Cl} + \mathrm{NH_4OH} \) solution on electrolysis time under AAE conditions.

Table 1

Radius of an ammonium amalgam droplet at different moments of electrolysis time at different pH values of the solution. Experimental conditions:
\( C_{\mathrm{NH_4^+}} = 0.1 \,\mathrm{N}; \quad V = 10 \,\mathrm{ml}; \quad t = 19^\circ \mathrm{C}; \quad \varphi_k = -1.998 \,\mathrm{V} \)

\( r \), mm

The ammonium ion concentration in the solution (\( C_{\mathrm{NH_4^+}} \)) has a large influence on the ammonium amalgam effect; therefore, in our experiments the dependence of the AAE on pH (in the interval from 6 to 10.5) was studied at a strictly constant \( C_{\mathrm{NH_4^+}} \). For this purpose, the solutions under investigation were prepared according to calculations using formula (2). The results of the experiments\(^1\) on the influence of pH on the growth kinetics of an ammonium amalgam droplet are given in Table 1.

From a comparison of the experimental data it is easy to see that, at constant concentration of the ammonium ion in the solution (and other experimental conditions)

\(^1\) The experimental setup and working procedure are described in works [5, 6].

the kinetics of droplet growth is practically independent of pH. For comparison, Table 1 also gives data characterizing the kinetics of droplet growth in a concentrated solution of \(\mathrm{NH_4OH}\) (4 N). In this case the formation of ammonium amalgam proceeds according to the following equation:

\[ \mathrm{NH_4OH}+\bar{e}+x\mathrm{Hg}\rightleftarrows \mathrm{NH_4(Hg)}_x+\mathrm{OH}^{-}. \tag{3} \]

The experimental data presented indicate that, in the reduction of the \(\mathrm{NH_4^+}\) ion, ammonium amalgam is formed more readily than in the reduction of the \(\mathrm{NH_4OH}\) molecule. This conclusion is in good agreement with the literature data [7].

It should be noted that in sufficiently acidic solutions, at \(\mathrm{pH}<4\text{--}5\), ammonium amalgam is not formed. This interesting fact is also noted in works [8] and [9]. This is apparently explained by the fact that in acidic solutions at a mercury cathode it is predominantly hydrogen ions, and not ammonium ions, that are reduced.

The studies we have carried out make it possible to draw the following conclusion: in the practical use of AAA to increase the sensitivity of the analytical determination of elements by the APA method, the pH of the solution will affect the height of the anodic peak only in the case where the given element is subject to hydrolysis. The pH does not affect the kinetics of growth of the ammonium amalgam droplet or its limiting radius.

REFERENCES

1. A. E. Vol. Structure and Properties of Binary Metallic Systems. Fizmatgiz, vol. 1, 1960; vol. 2, 1962.
2. M. Sh. Khashkhozhev. “Metallurg,” No. 12, p. 19, 1939.
3. Brief Handbook of the Chemist. P. 34, 1954.
4. B. V. Nekrasov. Course of General Chemistry. P. 191, 1962.
5. A. V. Konkova, A. G. Stromberg. Izv. TPI, 164, 147, 1968.
6. A. V. Konkova. ZhAKh, vol. 23, issue 3, 451, 1968.
7. C. I. Nyman, J. L. Ragle, F. P. Linde. Anal. Chem., 32, 352, 1960.