Russian Journal of Electrochemistry, Vol. 40, No. 4, 2004, pp. 466–469. Translated from Elektrokhimiya, Vol. 40, No. 4, 2004, pp. 524–527. Original Russian Text Copyright © 2004 by Korshunov, Vykhodtseva, Safonov.

SHORT COMMUNICATIONS

Reduction of Trivalent Chromium Ions on Stationary Mercury Electrode in Concentrated LiCl Solutions V. N. Korshunov, L. N. Vykhodtseva, and V. A. Safonovz Faculty of Chemistry, Moscow State University, Vorob’evy Gory 1, Moscow, 119992 Russia Received July 22, 2003

Abstract—The effect of supporting-electrolyte (LiCl) concentration on the first stage Cr(III) + e Cr(II) of Cr(III) reduction on a stationary mercury electrode is studied by linear voltammetry. It is found that an increase in the LiCl molar concentration m from 0.4 to ≈7 weakly affects the position of the cathodic current peak potential EP1 corresponding to this process. In the mentioned concentration range, [Cr(H2O)6 – nCln]3 – n (0 ≤ n < 3) ions are preferentially reduced. The further increase in m up to 19.1 results in a substantial positive shift of EP1, which is associated with the formation (upon passing the complete hydration limit) of the [Cr(H2O)3Cl3] (n = 3) complex, which is followed by its reduction. Key words: electroreduction of Cr(III), stationary mercury electrode, concentrated solutions of LiCl

INTRODUCTION The [ër(H2O)6]3+ aquacomplex is known to be highly kinetically inactive (the half-time of exchange of water molecules in the inner hydrate sphere is 3.9 × 109 s, whereas for Cr(II) it is 10–9 s) [1]. This factor (in agreement with the series accounting for the degree of the covalent nature of the Cr(III) bond with ligands: 2– F− > H2O > NH3 > C2 O 4 > OH– > Cl– [2, 3]) determines the low rate of the inner-sphere exchange Cl– in the complex at usual concentrations of ç2é the supporting electrolyte. This fact allows us to separately study the polarographic behavior of each form of the [Cr(H2O)6 – nCln]3 – n(aq) complex up to n = 2 [4–7]. However, the compound which corresponds to n = 3 is already unstable [8] and undergoes fast dissociative– hydrolytic decomposition [ Cr ( H 2 O ) 3 Cl 3 ] ( aq ) + H 2 O +

–

[ Cr ( H 2 O ) 4 Cl 2 ] ( aq ) + Cl ( aq ),

(1)

which prevents carrying out precision voltammetric studies. The mentioned complication can be bypassed by using concentrated LiCl solutions as the supporting electrolyte, because in this case equilibrium (1) turns out to be substantially shifted to the left [9]. The mentioned factor was used in this study for gaining information on certain electrochemical properties of [Cr(H2O)3Cl3] and the mechanism of formation of its inner coordination sphere at n 3. z

Corresponding author, e-mail: [email protected]

EXPERIMENTAL Polarization i vs. E curves were measured at 20°ë in ç2 atmosphere on a stationary mercury cathode of a 0.17-cm2 surface area in the potentiodynamic mode (8 mV/s). The scanning was repeated when necessary, then the electrode was renewed. The initial potential was taken equal to –0.2 V with respect to a Ag/AgCl, LiCl (m = 4) reference electrode (+0.18 V, NHE). The concentration of LiCl was varied from m = 0.4 to 19.1. Under these conditions, the diffusion potential estimated using the Henderson equation does not exceed 20–30 mV. The concentration of Cr(III) was maintained constant (m = 4 × 10–3). A 0.2 M KCr(SO4)2 · 12ç2O solution served as the source of Cr(III) ions. According to [10], at a temperature below 40°ë, [ër(H2O)6]3+ aquacomplexes are the main cationic form in such solutions. RESULTS AND DISCUSSION As it follows from our experimental data, throughout the studied LiCl concentration range up to saturation (m = 19.1), the i vs. E dependences measured on a stationary Hg cathode in Cr(III) solutions (e. g., curves 2 and 3, Fig. 1) are characterized by the presence of two current peaks (P1 and P2) separated by the a – b interval. These peaks correspond to the following stages: Cr ( II ) aq , Cr ( III ) aq + e (2) Cr ( II ) aq + 2e + Hg

Cr ( Hg ),

(3)

respectively. Similar results were obtained earlier by many authors [4, 11, 12] but only in solutions with supporting electrolyte concentration m < 7.

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i, mA/cm2 2 ∆H, kJ/mol

–30

P2 d 2

c

–20

0.6

3

~

EP1/2, V

–0.8

1

a

–0.4

0.4

c

b

0

10

20

30 NH2O

4 Fig. 2. Dependences of (1) half-peak potential of the Cr(III) aq + e Cr(II) aq stage and (2) integral heat (∆H) of LiCl dissolution in water on the quantity N H O = 55/m.

P1 0.2

2

2 3

a

b

5

4 6 0

–0.6

–1.0

–1.4

1 E, V

Fig. 1. Voltammograms measured on a nonrenewable stationary mercury cathode in solutions of (1) LiCl supporting electrolyte with m = 19.1 and (2–6) Cr(III) (m = 4 × 10–3) + LiCl (m = 19.1). Curves 2 and 3 are measured one after another without delay; curves 4–6 are preceded by a 1-min potential delay at –1.5 V.

Inasmuch as the current peak P2 falls within the range of sufficiently negative potentials where, in addition to reaction (3), the discharge of different proton donors (Çç+) can occur +

BH aq + e

1/2H2(g) + Çaq

(4)

(B = OH–, H2O, etc.), the polarization curve is substantially deformed at the potentials more negative than −1 V, particularly, the iê2/iê1 height ratio which should be close to 2 increases to 2.5–3 and higher. Moreover, the current reproducibility in the vicinity of P2 (in contrast to P1) is impaired in the repeated potential scan. Still more dramatic deformation effects are observed as a result of a short potential delay in the vicinity of the next current increase c – d which follows P2 (curves 4, 5, Fig. 1). In this case, the height of voltammograms gradually decreases throughout the whole potential range studied, and they approach the curve in the supporting electrolyte (curve 1). The described effects are probaRUSSIAN JOURNAL OF ELECTROCHEMISTRY

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bly associated with the fact that the hydrolysis-prone [ër(H2O)6]3+ ions (the hydrolysis constant in the first stage Kh1 = 1.3 × 10–4 [1]) take part in reaction (4) in the role of BH+, get polymerized upon their discharge [13], and the resulting polymers block the electrode, complicating the chromium deposition [14, 15]. The latter process involves stages of diffusion and formation of clusters of Cr atoms in mercury bulk, coarsening and conglomeration of clusters, after which the electrically inactive microcrystalline chromium phase well wettable by mercury is formed [16]. The mentioned facts suggest that when studying the effect of the supporting electrolyte concentration on the structure of discharging [Cr(H2O)6 – nCln]3 – n complex, the potential scanning parameters should be strictly controlled in order to ensure that the system studied stays in the potential range of P1. When these conditions are fulfilled, it was shown that the P1 height increases in proportion with Cr(III) concentration in solution and square root of scan rate. In [17–20], the effect of the supporting electrolyte type and concentration on the electrochemical characteristics of stage (2) was studied by focusing attention on the consequences of changing the corresponding rate constant (ks ~ 10–5 cm/s). In our work, bearing in mind that exchange, hydrolysis, and isomerization processes occurring in Cr(III)-containing electrolytes are very complicated, we chose the half-peak potential EP1/2 of reaction (2) as the main criterion parameter that responds to the structural changes of the discharging [Cr(H2O)6 – nCln]3 – n complexes. As follows from the experiment (Fig. 2), the sought dependence of EP1/2 on the amount of water moles N H2 O accounted for a LiCl mole ( N H2 O = 55/m) is almost horizontal for N H2 O > 7 No. 4

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KORSHUNOV et al.

Ionic radii and thermodynamic characteristics of water structure changes in the zone of inner-sphere hydration of ions [26] Ion Li+ Na+ K+ Cl–

r, nm

∆Gh, kJ/mol

0.07 0.10 0.14 0.18

–521 –421 –347 –308

∆∆Gh, kJ/mol –176 –8 34 15

and, along this segment of the curve, the electrolyte color changes from violet to different shades of green with an increase in N H2 O . According to [21], such color transformations point to successive substitution of chlorine ions Cl– for water molecules in the inner coordination sphere of the complex (n increases from 0 to ~2). However, this brings about a question why the optically revealed formation of new complexes only weakly affects the position of b – c segment of curve 1 in Fig. 2. Apparently, this is associated with the fact that, for moderately high m of the supporting electrolyte, this region corresponds to preferential discharge of [Cr(H2O)5Cl]2+aq ions formed in the reaction

+

+

[ Cr ( H 2 O ) 4 Cl 2 ] + Li …Cl  [ Cr ( H O ) Cl ] + 2 2 4    + –  Li …Cl

–

     +

[ Cr ( H 2 O ) 3 Cl 3 ] + Li ( H 2 O ) .

–

3+

accompanied by active salting out of Cl– ions by much stronger hydrated Li+ ions. A much more important fact is the formation of a more compact solvent structure in the inner coordination sphere of Li+, which results in the additional gain in hydration energy ∆∆Gh (see table). According to the magnitude of ∆∆Gh, lithium cation has an exceptionally high dehydration ability, and, when the CHL is exceeded, the LiCl solution resembles a glassy melt containing practically no free water [27]. Such a medium favors the formation of compact network structures with strong electrostatic interaction between fragments, which blocks stage (1) and stabilizes the [Cr(H2O)3Cl3] complex. Its formation (in the context of scheme (1)) in the discussed medium can be assumed to proceed as

[ Cr ( H 2 O ) 6 ] aq + Cl aq 2+

[ Cr ( H 2 O ) 5 Cl ] aq + H 2 O, and, for small m (not exceeding 7), it corresponds to the discharge of the same complexes arising as a result of fast catalytic decomposition of [Cr(H2O)4Cl2]+(aq) complexes. The decomposition of the latter is induced by Cr(II) ions generated according to scheme (2) [22, 23] +

[ Cr ( H 2 O ) 4 Cl 2 ] aq + H 2 O Cr(II)

2+

–

[ Cr ( H 2 O ) 5 Cl ] aq + Cl aq

(the decomposition stops when n = 1 is reached). This situation substantially changes at the verge of N H2 O ~ 7. Upon passing this boundary, the solution becomes yellowish green, which indicates the appearance of [Cr(H2O)3Cl3] complexes which can be synthesized under conditions of severe water deficiency [8, 9, 21]. It is this N H2 O region which corresponds to the most significant positive shift of EP1/2 (Fig. 2, curve 1, segment a – b). The appearance of a break in the EP1/2 vs. N H2 O dependence precisely near N H2 O = 7 can be explained if we compare this dependence with the ∆H vs. N H2 O curve (curve 2, Fig. 2) which reflects the changes of the integral heat of LiCl dissolution (∆H) in the gradually decreasing number of solvent moles. According to [24, 25], the breaks in dependences of such a kind (which involve, e. g., activity coefficients, specific conductivity, etc.) serve as an indicator that the complete hydration limit (CHL) of supporting electrolyte ions is reached. In our case, passing this limit is

CONCLUSIONS The study carried out showed that the existence of neutral [Cr(H2O)3Cl3] complexes in an aqueous chloride medium can take place only in the presence of strongly hydrated lithium cations and supporting electrolyte concentrations exceeding the complete hydration limit. Under these conditions, the mentioned complex is reduced on the mercury cathode near the potential –0.38 V (–0.20 V, NHE). ACKNOWLEDGMENTS The study was supported by the Russian Foundation for Basic Research (project no. 02-03-33221) and the Council for Grants of President of the Russian Federation (project NSh 2089.2003.3). REFERENCES 1. Cotton, F.A. and Wilkinson, G., Advanced Inorganic Chemistry, New York: Interscience, vol. 1. 2. Fialkov, Ya.A. and Nazarenko, Y.P., Zh. Neorg. Khim., 1956, vol. 1, p. 565. 3. Falicheva, A.I., Burdykina, R.I., and Shatalova, V.I., Vopr. Khim. Tekhnol. (Kharkov), 1981, no. 65, p. 30. 4. Chrom, vol. 52 of Gmelin’s Handbook der Anorganische Chemie, Weinheim: Chemie, 1962, part B, p. 942. 5. Hamm, R.E. and Shull, C.M., J. Am. Chem. Soc., 1951, vol. 73, p. 1240. 6. Virgili, J. and Costa, J.M., An. Real. Soc. Esp. Fis. y Quim., 1961, vol. 57, p. 479.

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REDUCTION OF TRIVALENT CHROMIUM IONS 7. Ibarz, J., Virgili, J., and Costa, J.M., An. Real. Soc. Esp. Fis. Quim., 1961, vol. 57, p. 489. 8. Kubota, H. and Costanzo, D.A., Anal. Chem., 1964, vol. 36, p. 2454. 9. Elving, P.J. and Zemel, B., J. Am. Chem. Soc., 1957, vol. 79, p. 1281. 10. Nisihara, K., Kurati, M., Nisii, H., et al., Trans. Min. Metallurg. Alumni Assoc., 1964, vol. 15, p. 203. 11. Demassieux, N. and Heyrovsky, J., J. Chim. Phys. Phys.Chim. Biol., 1929, vol. 26, p. 219. 12. Marczak, S., Wrona, P., and Galus, Z., J. Electroanal. Chem., 1995, vol. 396, p. 419. 13. Kopylovich, M.N., Radion, E.V., and Baev, A.K., Zh. Neorg. Khim., 1995, vol. 40, p. 1037. 14. Edigaryan, A.A. and Polukarov, Yu.M., Zashch. Met., 1998, vol. 34, p. 117. 15. Danilov, F.I. and Protsenko, V.S., Zashch. Met., 2001, vol. 37, p. 251. 16. Ochodzky, P. and Wrona, P., J. Electroanal. Chem., 1990, vol. 277, p. 225. 17. Weaver, M.J. and Satterberg, T.L., J. Phys. Chem., 1977, vol. 81, p. 1772.

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