Electrochemical and Spectroscopic Investigation of the Reduction of

Voltammograms (polarograms) obtained from solutions of cobalt and nickel ...... time (telec), the charge needed for complete reduction was calculated ...
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Anal. Chem. 1998, 70, 1312-1323

Electrochemical and Spectroscopic Investigation of the Reduction of Dimethylglyoxime at Mercury Electrodes in the Presence of Cobalt and Nickel Lesley A. M. Baxter,† Andrzej Bobrowski,*,‡ Alan M. Bond,*,§ Graham A. Heath,† Rowena L. Paul,† Robert Mrzljak,⊥ and Jerzy Zarebski‡

Research School of Chemistry, Australian National University, GPO Box 4, Canberra ACT 2601, Australia, Department of Materials Science and Ceramics, University of Mining and Metallurgy, 30-059 Krakow, al. Mickiewicza 30, Poland, Department of Chemistry, Monash University, Clayton, Victoria 3168, Australia, and School of Chemistry, La Trobe University, Bundoora, Victoria 3083, Australia

Voltammograms (polarograms) obtained from solutions of cobalt and nickel containing dimethylglyoxime (dmgH2) are widely used for the trace determination of these metals. Detailed electrochemical and spectroscopic studies on the reduction process observed in the analytically important ammonia buffer media at mercury dropping, hanging, and pool electrodes are all consistent with an overall 10-electron reduction process, in which both the dmgH2 ligand and cobalt ions are reduced in the adsorbed state: Co(II) + 2dmgH2 h (solution) [CoII(dmgH)2] + 2H+; [CoII(dmgH)2] + Hg h (electrode) [CoII(dmgH)2]adsHg; and [CoII(dmgH)2]adsHg + 10e- + 10H+ f Co(Hg) + 2[2,3-bis(hydroxylamino)butane]. The limited solubility of the nickel complex in aqueous media restricts the range of studies that can be undertaken with this system, but an analogous mechanism is believed to occur. Lowtemperature voltammetric studies in dichloromethane at a frozen hanging mercury drop electrode and in situ electron spin resonance electrochemical measurements on more soluble analogues of the dimethylglyoxime complexes are consistent with an initial one-electron reduction step being available in the absence of water. Deliberate addition of water to acetone solutions enables the influence of the aqueous environment on voltammograms and polarograms to be examined. The results of the present study are compared with the wide range of mechanisms proposed in other studies. Fifty years ago, Stromberg and Zelynskaya1 observed an increase in the polarographic wave height for the reduction of cobalt ions in an ammonia buffer solution when dimethylglyoxime was added to the solution. Since this initial report, the enhancement of the current observed for cobalt and nickel in the presence of dimethylglyoxime has been widely utilized for the determination of trace concentrations of these elements by both polarographic

Chart 1. Structural Representation of (a) Cobalt and Nickel Dimethylglyoxime Complexes, M(dmgH)2, and (b) Cobalt and Nickel Dimethylglyoximate Analogues, M(C8doH)2, where M is Co(II) or Ni(II)

and adsorptive stripping voltammetric methods.2 However, considerable controversy still exists concerning the nature of the reduction mechanism. For convenience, the symbols CoII(dmgH)2 and NiII(dmgH)2 will be used in this paper to denote the complexes (structure a, Chart 1) formed by coordination of two deprotonated molecules of dimethylglyoxime (dmgH2). According to the exclusively adsorptive mechanism, the reaction scheme may be summarized as follows: solution

Co(II) + 2dmgH2 y\z [Co(dmgH)2] + 2H+ electrode

[Co(dmgH)2] y\z [Co(dmgH)2]adsorbed

(1) (2)

[Co(dmgH)2]adsorbed + 2e- + 2H+ f Co(O) + [dmgH2]desorbed (3)

†

Australian National University. University of Mining and Metallurgy. § Monash University. ⊥ La Trobe University. Present address: ICI Botany Operations (Probe Analytical), 16-20 Beauchamp Rd., Matraville, NSW, 2036, Australia. (1) Stromberg, A. G.; Zelyanskaya, A. I. Zh. Obshch. Khim. 1945, 15, 303.

Thus, after the mercury electrode reaches the required potential, the cobalt or nickel dimethylglyoxime complex is reduced in the

1312 Analytical Chemistry, Vol. 70, No. 7, April 1, 1998

S0003-2700(97)00361-2 CCC: $15.00

‡

© 1998 American Chemical Society Published on Web 03/04/1998

adsorbed state, and dmgH2 is released after reduction of the absorbed complex. Jin and coauthors have proposed a slightly different mechanism, in which reactive cobalt(I)3,4 or nickel(I)5 intermediates are generated. electrode

[CoII(dmgH)2] y\z [CoII(dmgH)2]adsorbed

(4)

[CoII(dmgH)2]adsorbed + e- h [CoI(dmgH)2]-adsorbed

(5)

[CoI(dmgH)2]-adsorbed + H+ + e- f V X Co(O) + [dmgH2]desorbed + [dmgH]-desorbed (6)

Compound X (identity unknown) is produced via a chemical decomposition reaction. A similar mechanism was suggested earlier by Pihlar et al.6 for the reduction of the nickel complex, although the formation of product X was not reported. Support for these surface-confined, overall irreversible two-electron processes has been provided by bulk electrolysis experiments carried out on small concentrations of the cobalt dimethylglyoxime complex at a mercury cup microelectrode, which suggested that only two electrons were transferred.4,5 Thus, according to Jin et al.3-5 and Pihlar et al.,6 the metal-based reduction of cobalt and nickel dimethylglyoxime complexes occurs from the adsorbed state and is not associated with catalytic hydrogen evolution, as proposed in the original studies.1 In other noncatalytic mechanisms proposed, reduction of both the central Co(II) or Ni(II) ion and the ligand [dmgH]- occurs as in eq 7:

[M(dmgH)2]adsorbed + xe- + yH+ f M(0) + [dmg]red

(7)

where M represents cobalt or nickel, x ) 10-18 electrons, and [dmg]red is the product(s) of reduction of dimethylglyoxime. According to this mechanism, the coordinated ligand may be partially reduced to 2,3-bis(hydroxylamino)butane,

or totally to 2,3-diaminobutane,

(9)

The concept of reduction of the ligand (eq 7) was first proposed by Weinzierl and Umland7 on the basis of the observation that the reduction wave of dmgH2 (E1/2 ) - 1.7 V) in the ammonia buffer supporting electrolyte disappeared after the addition of Co(II) ions. The mechanism given by eq 7 also was suggested earlier by some of the present authors.8 Recently, Vukomanovic et al.,9 in a study which focused on the Ni(dmgH)2 system, reported predominantly on the basis of voltammetric evidence that the reduction process involved the transfer of 16 or, possibly, 18 electrons. In another recent study,10 Jagner et al. concluded, on the basis of coulometric stripping potentiometric technique, that reductions of the Co(dmgH)2 and Ni(dmgH)2 complexes are 10rather than 16-electron processes. A range of catalytic schemes have been proposed. The cyclic catalytic reduction mechanism may be summarized as follows: solution

Co(II) + 2dmgH2 y\z [Co(dmgH)2] + 2H+ (10) electrode

[CoII(dmgH)2] y\z [CoII(dmgH)2]adsorbed [CoII(dmgH)2]adsorbed + Xe- + yH+ f

Co(II) + 2(dmg)reduced (12)

In this mechanism, the Co(II) ions released after the reduction process may then react with bulk dmgH2. The coordinationreduction-coordination process will then proceed in cycles until all of the dmgH2 becomes reduced. The reduction of cobalt and nickel as well as copper dimethylglyoxime complexes has been suggested by Skobelkina and Prokhorova11 to involve catalytic hydrogen evolution, according to the following scheme:

[CoII(dmgH)2]adsorbed + 2e- f Co(0) + 2(dmgH)(8)

(11)

(13)

electrode

Co(0) + [dmgH2]adsorbed y\z [Co0(dmgH)]-adsorbed + H+ (14) (2) See, for example: (a) Komarek, K Collect. Czech. Chem. Commun. 1947, 12. (b) Nagniot, P. J. Electroanal. Chem. 1967, 14, 19. (c) Meyer, A.; Neeb, R. Z. Anal. Chem. 1983, 315, 118. (d) Torance, K.; Gatford, C. Talanta 1985, 32, 273. (e) Adeloju, S. B.; Bond, A. M.; Briggs, M. H. Anal. Chem. 1985, 57, 1386. (f) Bobrowski, A. Anal. Chem. 1989, 61, 2178. (g) Brett, C. M. A.; Oliveira Brett, A. M. C. F.; Pereira, J. L. C. Electroanalysis 1991, 3, 683 and references therein. (3) Jin, W.; Lin, K. Acta Chim. Sinica (Chinese Ed.) 1985, 43, 923. (4) Jin, W.; Lin, K. J. Electroanal. Chem. 1987, 216, 181. (5) Jin, W.; Jiang, W.; Yang, Q. Acta Phys.-Chim. Sin. 1988, 4, 445. (6) Pihlar, B.; Valenta, P.; Nu ¨ rnberg, H. W. J. Electroanal. Chem. 1986, 214, 157.

(7) Wienzierl, J.; Umland, F. Z. Anal. Chem. 1982, 312, 608. (8) (a) Bobrowski, A.; Zarebski, J. Report of University of Mining and Metallurgy, December, 1988. (b) Bobrowski, A. D.Sc. Habilitation Thesis, University of Mining and Metallurgy, Krakow, Poland, 1995. (c) Mrzljak, R. Ph.D. Thesis, La Trobe University, Victoria, Australia, 1994. (9) Vukomanovic, D. V.; Page, J. A.; van Luon, G. W. Anal. Chem. 1996, 68, 829 and references therein. (10) Ma, F.; Jagner, D.; Renman, L. Anal. Chem. 1997, 69, 1782. (11) Skobelkina, E. W.; Prokhorova, G. V. Vestn. Moskov. Univ. Khim, 1977, 18, 197 and refs. therein.

Analytical Chemistry, Vol. 70, No. 7, April 1, 1998

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[Co0(dmgH)]-adsorbed + BH+ f {[Co0(dmgH)]-adsorbedH+} + B (15) {[Co0(dmgH)]-adsorbedH+} + e- f /2H2 + [Co0(dmgH)]-adsorbed (16)

1

where BH+ is the source of the donor proton, e.g., NH4+. The catalytic hydrogen evolution mechanism is favored by the majority of the Russian investigators.1,12-14 Burger et al. also have supported this general class of mechanism.15 The principal evidence supporting the thesis of the catalytic hydrogen reduction has been based on the reported formation of gas bubbles during reductive electrolysis on the HMDE surface by Prokhorova et al.16,17 and the presence of the maxima on dc polarographic waves, which is consistent with catalytic reduction of hydrogen ions.1,16-18 The mechanism described by eqs 13-16 has been modified several times by Russian investigators but is always based on the assumption of the catalytic reduction of hydrogen ions and no change in the redox state of dimethylglyoxime. However, the concept of a catalytic evolution of hydrogen has been widely criticized.4,6,7,19-24 For example, Fini and Todeschini,21 using the radio polarographic method and 60Co as a radio tracer, concluded that the cobalt dimethylglyoxime reduction process was not of a catalytic nature, although, in contrast, it should be noted that Li et al.,22 using 3H2 as a radioactive tracer, found that H2 was liberated on a protracted controlled-potential electrolysis time scale (50 h), in accord with possible catalytic reduction of hydrogen ions. In other studies, Chinese investigators23,24 confirmed the appearance of a new compound during the reduction of cobalt and nickel in the presence of dimethylglyoxime and concluded that the electrode process involves the simultaneous reduction of Co(II) to Co(0) or of Ni(II) to Ni(0) and dimethylglyoxime to dmgred rather than catalytic hydrogen evolution. Emons et al.25 also reported that reduction of the cobalt system produced an insoluble new product, while exhaustive reductive electrolysis experiments of cobalt dimethylglyoxime complexes by Nambiar and Subbaraman26 indicated that the ligand is reduced. In acid media, they showed that the ligand is reduced to 2,3-diaminobutane with eight electrons transferred per ligand, with the cobalt(II) central atom being reduced to the metal. This concept is (12) Vinogradova, E. N.; Prokhorova, G. V. Zh. Anal. Khim. 1968, 23, 711. (13) Prokhorova, G. V.; Henrion, G.; Schmidt, R. Z. Chem. 1982, 22, 28. (14) Budnikov, G. K.; Miedianceva, E. P. Zh. Anal. Khim. 1973, 28, 301. (15) Burger, K.; Syrek, G.; Farsong, G. Acta Chim. Acad. Sci. Hung. 1966, 49, 113. (16) Mairanovskii, S. G.; Prokhorova, G. V.; Osipova, E. A. J. Electroanal. Chem. 1989, 266, 205 and references therein. (17) Prokhorova, G. V.; Spigun, L. K.; Vinogradova, E. N. Zh. Anal. Khim. 1972, 27, 780. (18) Prohorova, G. V.; Vinogradova, E. N.; Pronina, N. V. Zh. Anal. Khim. 1970, 25, 2073. (19) Flora, C. J.; Nieboer, E. Anal. Chem. 1980, 52, 1013. (20) Pihlar, B.; Valenta, P.; Nu ¨ rnberg, H. W. Z. Anal. Chem. 1981, 307, 337. (21) Fini, G.; Todeschini, P. Ric. Sci. 1968, 38, 787. (22) Li, K.; Qin, R.; Wang, X.; Wu, X.; Li, H. Yuan Zi Neng Ke Xue Ji Shu 1985, 1, 101. (23) Ni, Y.; Li, L.; Zhou, C.; Gao, X. Acta Chim. Sin. 1987, 47, 242. (24) Li, L.; Ni, Y. M.; Gao, X. Acta Chim. Sin. 1988, 46, 1031. (25) Emons, H.; Schmidt, T.; Werner, G. Anal. Chim. Acta 1990, 228, 55. (26) Nambiar, O. G. B.; Subbaraman, P. R. Aust. J. Chem. 1971, 24, 2089.

1314 Analytical Chemistry, Vol. 70, No. 7, April 1, 1998

supported by the work of Spritzer and Meites,27 who demonstrated that, in the pH range from 1 to 3, dimethylglyoxime could be electrochemically reduced to 2,3-diaminobutane. In slightly basic media, the ligand is reduced only to 2,3-bis(hydroxylamino)butane, where four electrons are transferred per molecule of dimethylglyoxime. Interestingly, an unexpected byproduct of the synthesis of TcCl(dmgH)3BEt was TcCl(dmgH)2(BDI)BEt,28 where BDI is butane-2,3-dione imine oxime, a reduced form of dimethylglyoxime. In summary, the existence of discrepancies and controversies induced the present authors to further study the cobalt and nickel dimethylglyoxime reduction processes via a wide range of techniques at different forms of mercury electrodes. Furthermore, the number of electrons involved in the reduction step has been established at both dropping and stationary mercury electrodes by exhaustive electrolysis coupled with coulometry. Additionally, products from electrolysis have been examined by infrared spectroscopy, mass spectrometry, and voltammetry. The insoluble nature of the Co(dmgH)2 and Ni(dmgH)2 compounds in most solvents restricts the range of spectroelectrochemical studies that can be undertaken. The reduction of nickel and cobalt dimethylglyoxime analogues, Co(C8doH)2 and Ni(C8doH)2 (structure b, Chart 1), which are soluble in organic solvents, therefore were also investigated to aid the understanding of the mechanism of the reduction process at mercury electrodes. EXPERIMENTAL SECTION Reagents and Compounds. All reagents were of analytical grade purity unless otherwise stated. Co(dmgH)2 and Ni(dmgH)2 complexes were usually generated in situ in aqueous media by adding the metal ion into an ammoniacal solution at pH 9 containing excess dmgH2. Standard solutions of dmgH2 contained the minimum amount of ethanol required to achieve complete dissolution of the solid. Solid Co(dmgH)2‚2H2O was synthesized by adding cobalt(II) nitrate into an aqueous acetate solution at pH 5 containing dimethylglyoxime and extracting the cobalt complex with chloroform. The chloroform extract was evaporated to dryness, leaving the solid Co(dmgH)2‚2H2O complex. Cobalt and nickel dimethylglyoximate analogues, Co(C8doH)2‚2H2O and Ni(C8doH)2, were synthesized as follows. Co(C8doH)2‚2H2O. Co(OAc)2‚4H2O (0.44 g, 1.75 mmol) was dissolved in 10 mL of degassed H2O, and C8doH2 (0.601 g, 3.5 mmol) dissolved in degassed MeOH (5 mL) was added. On mixing, a brown/orange precipitate formed. The reaction mixture was stirred under N2 for 1 h. The solid was filtered under N2 and washed well with degassed H2O. The brown-orange solid was dried in vacuo (yield 80%). Anal. Calcd for Co(C8doH)2‚ 2H2O ()C16H30CoN4O6): C, 44.34; H, 6.93; N, 12.93. Found: C, 44.30; H, 7.23; N, 12.77. IR (cm-1): 955, 1040 ν(C-H); 3300 ν(O-H) (due to H2O). Ni(C8doH)2. The synthetic route to Ni(C8doH)2 was similar to that for the analogous cobalt complex. Ni(OAc)2‚4H2O was used as the starting material, and a yellow-orange solid was isolated and washed with H2O and EtOH (yield 85%). Anal. Calcd for Ni(C8doH)2 ()C16H26N4O4Ni): C, 48.40; H, 6.55; N, 14.12. (27) Spitzer, M.; Meites, L. Anal. Chim. Acta 1962, 26, 58. (28) Lindner, K. E.; Nowotnik, D. F.; Malley, M. F.; Gougoutas, J. Z.; Nunn, A. D. Inorg. Chim. Acta 1991, 190, 249.

Found: C, 48.35; H, 6.6; N, 14.05. IR (cm-1): 1036, 1239 ν(NO); 1558 ν(CdN); 1780 δ(OHO); 2850, 2920 ν(C-H). Solvents used were water (NANOpure, Barnstead, Dubuque, IA), acetone (Mallinckrodt, Paris, KY, Chrome HPLC grade), dichloromethane (Mallinckrodt, HPLC grade) purified by passing through an active neutral alumina column, and acetonitrile (Mallinckrodt, HPLC grade) purified as described in ref 29. Supporting electrolytes comprised 0.1 M ammonium chloride (BDH, Poole, England) and 0.1 M ammonia solution (May & Barker, Pronalys) for aqueous solution studies and tetrabutylammonium hexaflurophosphate, synthesized by reaction of tetrabutylammonium bromide (Aldrich, Milwaukee, WI) and potassium hexaflurophosphate (Aldrich),30 for the organic solvent studies. To remove oxygen, solutions were purged with high-purity nitrogen or argon for at least 10 min prior to undertaking electrochemical experiments. Electrochemical Instrumentation. Polarographic experiments were conducted with a Metrohm 646 VA processor and 647 VA stand, utilizing the multimode electrode in the dropping mercury mode or with an Unitra-Telpod PP-04 pulse polarograph and a dropping mercury electrode (DME). A saturated calomel (SCE) or Ag/AgCl (3 M KCl) reference electrode and a platinum wire auxiliary electrode were used in these studies. Voltammograms were acquired using a Cypress System CYSY-1190 computercontrolled electroanalysis system in the cyclic staircase mode. A single-compartment cell with a conventional three-electrode arrangement was used with platinum or glassy carbon stationary or rotated macrodisk (1-3-mm diameter), 10-µm-diameter microdisk, or hanging mercury drop (Metrohm EA 410) working electrodes, platinum wire counter electrodes, and either a Ag/ AgCl (3 M KCl) reference electrode for aqueous work or a silver wire quasi-reference electrode for organic solvent studies. The potential of the silver wire quasi-reference electrode was calibrated against the potential of the Fc/Fc+ couple via measurement of the reversible potential found from oxidation of 1.0 mM solution of ferrocene (Fc) in the organic solvent of interest. Potentials in organic solvents are reported versus that found for the Fc/Fc+ couple. For microelectrolysis experiments at a DME, a polarographic cell was designed which enabled a small volume of the solution (1-2 mL) to be exhaustively reduced over a 10-h period of time with a dropping mercury electrode having a drop life of 0.5 s. The three-electrode system containing a DME, reference electrode, and auxiliary electrode was connected to a dc polarograph, which was used to record the polarographic waves after specified periods of microelectrolysis and also provided the voltage to the cell during microelectrolysis. The microelectrolysis was performed in the following manner: 2 mL of a Co(II) solution with addition of various concentration excesses of dmgH2 in 0.1 M ammonia buffer was transferred to the microcoulometric vessel. Argon was passed for 20 min, and a dc polarogram was recorded in the potential range from -0.9 to -1.4 V vs SCE. The potential of the plateau region was reached over the range from -1.2 to -1.4 V vs SCE, and the electrolysis was carried out at constant potential in this region. During the course of electrolysis, dc and differential pulse (DPP) polarograms were recorded each hour.

Exhaustive bulk electrolysis experiments at a mercury pool working electrode were carried out using a Bioanalytical Systems model 100A electrochemical analyzer. Ag/AgCl (3 M NaCl) was used as the reference electrode, and a counter platinum mesh electrode was placed in a separate compartment and connected to the test solution by a salt bridge. An overhead rotating stirrer was used during the electrolysis to provide convection. The cell used 3 mL of mercury and was applied with solution volumes ranging from 4 to 10 mL. The small volumes guaranteed that the electrolysis time was kept to a minimum to reduce errors associated with long time scale electrolysis experiments.31 The counter platinum basket electrode was isolated from the cell, and the circuit was completed with a salt bridge containing 0.1 M ammonia buffer. In both microelectrolysis and bulk electrolysis experiments, the solutions were continuously purged with nitrogen or argon for the duration of the electrolysis period. Spectroscopic Instrumentation. A Perkin-Elmer FT-IR spectrometer (1720X) was used for the identification of electrolysis products in attenuated total reflectance mode using a zinc selenide cell. To obtain a FT-IR spectrum, five drops of sample solution was added onto the zinc selenide cell, and a steel plate was clamped down onto the sample, dispersing the sample over the entire cell and ensuring a thin film. The thickness of the film was controlled by a micrometer controlling the plate. The mass spectra (GC/MS mode) were recorded on a JEOL-DX300 spectrometer. In situ ESR electrochemical measurements were undertaken with a Bruker ECS 106 spectrometer in conjunction with gas-cooled temperature controller and the electrochemical cell described in ref 32 .

(29) Kiesele, H. Anal. Chem. 1980, 52, 2230. (30) Kissinger, P. T.; Heineman, W. R. Laboratory Techniques in Electroanalytical Chemistry; Dekker: New York, 1984; p 378.

(31) Nemeth, M.; Mocak, J. Collect. Czech. Chem. Commun. 1986, 51, 636. (32) Fiedler, D. A.; Koppenol, M.; Bond, A. M. J. Electrochem. Soc. 1995, 142, 862.

RESULTS AND DISCUSSION Voltammetric Reduction of the Cobalt and Nickel Systems at Mercury Electrodes in Aqueous Media. Initial investigations were undertaken on noncoordinated dimethylglyoxime (added as dmgH2) and are used as a reference for the studies on the cobalt and nickel complexes. Dc polarograms of a solution containing 1 × 10-3 M free ligand in an aqueous 0.1 M ammonia buffer medium showed one reduction wave with an E1/2 value of -1.55 V vs Ag/AgCl when the drop time was 0.6 s. Although Spitzer and Meites27 reported that the response could not be resolved from the solvent reduction limit in neutral or alkaline media, in this study it was found that the process is very well defined when 0.1 M ammonia buffer is used as the electrolyte. Vukomanovic et al.9 have also noted a well-defined reduction wave in alkaline media. Figure 1a shows a polarogram in 0.1 M ammonia for reduction of dmgH2 compared to that of the well-documented twoelectron cadmium reduction, Cd(II) + 2e- h Cd(Hg), at the same concentration of 1 × 10-3 M and with all other experimentally controlled parameters being kept constant. From the data, it is evident from the ratio of the diffusion-controlled limiting currents that the number of electrons transferred during the course of reduction of dmgH2 is significantly greater than the known value of 2 for the cadmium ion, assuming that the diffusion coefficients of the cadmium ion and dmgH2 are similar. Cyclic voltammetric experiments at a hanging mercury drop electrode show that the reduction wave of the free ligand is

Analytical Chemistry, Vol. 70, No. 7, April 1, 1998

1315

Figure 2. Dc polarograms in 0.1 M ammonia buffer at 20 °C of (a) 1 × 10-3 M cobalt(II), (b) with addition of 2 × 10-4 M dmgH2, (c) with addition of 5 × 10-4 M dmgH2, and (d) with addition of 1 × 10-3 M dmgH2.

Figure 1. (a) Dc polarogram for the reduction of 1 × 10-3 M dmgH2 and 1 × 10-3 M cadmium with a drop time of 0.6 s. (b) Cyclic voltammograms of 1 × 10-3 M dmgH2 at a hanging mercury drop electrode: (i) first scan and (ii) second scan using a scan rate of 200 mV s-1. The temperature was 20 °C, and 0.1 M ammonia buffer was the electrolyte.

situated on the shoulder of the process giving rise to the solvent limit (Figure 1b). However, on the reverse scan, a new oxidation wave at about -0.50 V vs Ag/AgCl was observed. This wave is only present after the free ligand has been reduced and, therefore, represents an oxidation process associated with the generation of a reduced form of the ligand. In the second and subsequent cycles, a new reduction wave is evident at about -0.60 V vs Ag/ AgCl, as shown in Figure 1b(ii). This reduction process is present only after the oxidation process at -0.50 V vs Ag/AgCl, and this chemically reversible redox couple is, therefore, the result of an initial reduction of the free ligand. The reduction of dmgH2 is consistent with the formation of 2,3-bis(hydroxylamino)butane (DHAB), as proposed by previous workers,26 and the reversible redox couple observed on the second and subsequent cycles is likely to be the result of the formation of a mercury complex with DHAB at the electrode surface. The following scheme is consistent with the data:

dmgH2 + 4e- f DHAB (-1.70 V) 2DHAB + Hg h Hg(DHAB)2 + 2e- (-0.55 V) 1316 Analytical Chemistry, Vol. 70, No. 7, April 1, 1998

(17) (18)

Importantly, the reversible redox couple observed on the second and subsequent cycles can be used as an indicator to determine whether the reduction of the ligand is involved in the reduction process observed in the presence of cobalt and nickel. Electrochemical reduction of the cobalt system was initially investigated in 0.1 M ammonia buffer by dc polarography. The very low solubility of the Ni(dmgH)2 complex made dc polarographic studies in this system very difficult. Figure 2 shows dc polarograms obtained for free Co(II) ions (Figure 2a) and then with increasing additions of dmgH2 to the solution (Figure 2bd). The polarographic reduction of Co(II) in the absence of dmgH2 has a limiting current value which is the result of a diffusion-controlled two-electron reduction to Co(0) metal under the conditions of Figure 2. The initial addition of dmgH2 to the solution results in an increase in the limiting current as well as the appearance of a polarographic maximum. With a considerable excess of dmgH2, the limiting current attains a constant value which is about 4 times the limiting current for the reduction of noncomplexed Co(II), while the peak current attains a value which is about 6 times greater. This result again indicates that a multielectron reduction process occurs at negative potentials, where the number of electrons partaking in the reduction reaction in the diffusion-controlled region is about eight, assuming equal diffusion coefficients for complexed and noncomplexed forms of cobalt(II). This increase in the polarographic signal is similar to that reported by Stromberg almost 50 years ago.1 Data obtained under a range of conditions are summarized in Table 1, and it can be noted that the maxima are favored by high concentrations of dmgH2 and slow mercury flow rates. Thus, at short drop times and fast flow rates, maxima can be avoided, and diffusioncontrolled reduction can be achieved over a wide potential range.

Table 1. Influence of dmgH2 Concentration on the dc and DPP Polarographic Current Fraction, ICo-dmgH2/Io, in 0.1 M Ammonia Buffer as a Function1 of Drop Time (td) at 20 °C ICo-dmgH2/Io slow DME DPP technique CdmgH2 (M) td ) 2 s td ) 0.5 s 0 5 × 10-6 1 × 10-5 2 × 10-5 4 × 10-5

1 1.72 2.50 3.78 5.66

1 1.35 2.18 3.06 4.35

6 × 10-5

7.69

5.59

8 × 10-5

9.5

6.59

1 × 10-4

11.4

7.53

5 × 10-4

26.3

17.2

1 × 10-3

32.2

18.8

fast DME, dc technique dc technique td ) 2 s td ) 0.5 s td ) 0.5 s 1 1.59 2.29 3.18 3.12 (3.65)b 3.29 (4.12)b 3.29 (4.8)b 3.65 (5.53)b 4.47 (11.6)b

1 1.6 2.13 2.39 2.8

1 1.71 2.48 3.34 3.83

2.93

4.23

3.2

4.0

3.2 (4.4)b 3.47 (5.73)b

4.17 4.17 4.34

a I , limiting or peak current for reduction of 2 × 10-5 M Co(II) in o 0.1 M ammonia buffer. ICo-dmgH2, limiting or peak current after addition of dmgH2. b Value of dc polarographic maximum current.

Polarographic maxima are attributed to differences in interfacial tension around the mercury drop causing solution streaming33 and have been investigated in detail for the copper dimethylglyoxime system.34 The enhancement of the peak current with increasing dmgH2 concentration is extremely pronounced in DPP. Thus, the peak current of the Co-dmg complex can be enhanced by a factor of 30 or more (Table 1) relative to the peak current obtained for reduction of free Co(II), although the enhancement factor strongly depends on the drop life of the DME. The considerable enhancement of the peak current in DPP is, of course, analytically desirable. Cyclic voltammetry of a solution containing Ni(II) and dmgH2 at a HMDE revealed a large peak attributable to reduction of the nickel complex at about -1.05 V vs Ag/AgCl (Figure 3a). The low solubility in bulk solution precludes solution-phase studies, and the peak current was found to be linearly proportional to scan rate (50-1000 mV s-1), which is characteristic of the reduction of an adsorbed species.6 The shape of the reduction step also is consistent with reduction of a surface confined species. At more negative potentials, the reduction of the free ligand is encountered due to a large excess being present in the solution. Reduction of this excess ligand gives rise to the observation of a chemically reversible process on the second cycle (Figure 3b), as expected when the ligand is reduced (see Figure 1b). However, when the switching potential was changed to -1.20 V vs Ag/AgCl, thus eliminating the possibility of reducing the free ligand, the same chemically reversible redox couple is still evident (Figure 3c). This redox couple, present on second and subsequent cycles, results from the formation of reduced dmgH2; therefore, reduction of the Ni-(dmgH)2 complex also must involve ligand reduction. Analo(33) Milner, G. W. C. The Principles and Application of Polarography; Longmans: London, 1962; p 72.

Figure 3. Cyclic voltammograms at a scan rate of 100 mV s-1 and 20 °C for the Ni(dmgH)2 complex generated in situ with excess dimethylglyoxime at a hanging mercury drop electrode in 0.1 M ammonia buffer: (a) first scan and (b) second scan; (c) second scan, but after switching the potential at -1.2 V.

gous cyclic voltammetric results were observed for the cobalt system. The above data imply that the electrochemical reduction of the cobalt and nickel complexes involves strong adsorption at a HMDE and that reduction of the complex also represents a multielectron system at both the DME and HMDE, where not only the central metal ion can be reduced but also the ligands. Importantly, the same overall process appears to occur irrespective of whether the complex is adsorbed or not. Microcoulometric and Exhaustive Coulometric Study of the Cobalt Complex at a DME. Nambiar and Subbaraman26 performed exhaustive coulometry in acid solutions and found that eight electrons are transferred per molecule of ligand in the cobalt system. Furthermore, Spitzer and Meites27 found, from conventional exhaustive coulometry of the free ligand in acid solutions, that eight electrons were transferred per molecule of dmgH2. However, no coulometric data are available at higher pH, which is the analytically important pH regime. Therefore, in this work, the microelectrolysis and exhaustive electrolysis (coulometry) of the cobalt complex was performed at pH 9 using a DME. The conditions of coulometric microelectrolysis of a small volume of the solution at DME represent the conditions of polarographic experiments. The potential of electrolysis was held at -1.25 V vs SCE, which corresponds to the limiting current region of dc polarograms and the region where adsorption was found to be Analytical Chemistry, Vol. 70, No. 7, April 1, 1998

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Table 2. Number of Electrons (ne) Determined via Coulometry (Bulk Electrolysis at a Mercury Pool Electrode at 20 °C) for the Reduction Process in 0.1 M Ammonia Buffer for Different Concentration Ratios of dmgH2 to Co(II)a calcd no. of electrons, ne

composition of electrolyzed solution

concn ratio [dmgH2]/[Co(II)]

3.4 × 10-5 M Co(II)

0

2.3

3.4 × 10-5 M Co(II) 4 × 10-5 M dmgH2

1.18

5.18

5.0 × 10-5 M Co(II) 1.0 × 10-4 M dmgH2

2

3.4 × 10-5 M Co(II) 1.0 × 10-4 M dmgH2

2.94

11.6 (1.5) 9.43

5.0 × 10-5 M Co(II) 5.0 × 10-4 M dmgH2

10

10.2 (1.7)

3.4 × 10-5 M Co(II) 1.0 × 10-3 M dmgH2

29.41

10.39

5.0 × 10-5 M Co(II) 5.0 × 10-3 M dmgH2

100

9.0 (2.3)

a Values in parentheses are standard deviations from triplicate experiments.

minimal. Mass transport of electroactive species to the electrode was, therefore, governed predominantly by diffusion and convection introduced by stirring of the solution. Unfortunately, electrolysis experiments of this kind could not be performed on the nickel complex because of inadequate solubility. From the plot of the limiting current (Ilim) vs microelectrolysis time (telec), the charge needed for complete reduction was calculated by extrapolating to the time predicted for completion of the electrolysis. The number of electrons taking part in the reduction process at the DME was subsequently calculated from the derived value of the total charge for exhaustive electrolysis. The determination of the number of the electrons transferred as described above gave an n-value of 10.2 ( 0.9 for 11 experiments. The value was independent of the ligand concentration when the dmgH2-to-Co(II) concentration ratio was in excess of 2:1. From this microcoulometric study, it follows that the reduction process at a DME occurs according to eq 7. Moreover, the independence of the determined number of electrons of the dmgH2 excess over its stoichiometric concentration provides strong evidence that no cyclic catalytic electrode reaction occurs, at least at the DME. To confirm that the microcoulometric results at a DME are valid at stationary mercury electrodes, where reduction occurs in the adsorbed state, and also to investigate the nature of the electrolysis products, larger scale bulk, controlled-potential electrolysis experiments were performed at a mercury pool electrode. The results for the exhaustive electrolysis (coulometry) experiments at a mercury pool electrode for solutions with different ratios of dmgH2 and Co(II) are summarized in Table 2. The Co(II) concentration in these experiments was constant at either 3.4 × 10-5 or 5 × 10-5 M, while the dmgH2 concentration varied, and it was found that the number of electrons transferred (n) increases with increasing concentration of dmgH2 from n ≈ 2 for Co(II) reduction in the absence of dmgH2, reaching n ≈ 10 for dmgH2 concentrations higher than the stoichiometric 2:1 concen1318 Analytical Chemistry, Vol. 70, No. 7, April 1, 1998

tration ratio. These results imply that reduction of the Co(dmgH)2 complex, even when adsorbed onto a mercury surface, still represents a multielectron reduction process, with n ) 10 ( 2 electrons/molecule of Co(dmgH)2 complex. During the reductive electrolysis at a mercury pool electrode, the solution changed from a deep brown color to clear, indicating complete loss of complex. Very little Co(dmgH)2 could be voltammetrically detected in the colorless solution. Voltammetric monitoring at very negative potentials also revealed that the excess dmgH2 is not consumed during the reduction. This result indicates that free Co(II) ions are not a product of electrolysis, as they would immediately react with the excess of dmgH2 in the solution and the brown color would persist. That is, the process does not involve, in the overall sense, pure ligand reduction, Co(dmgH)2 f Co(II) + reduced ligand, giving rise to a cyclic catalytic reduction process of the kind given in eqs 10-12. Evidence to support reduction to cobalt metal was obtained via analysis of the mercury pool cathode. After exhaustive electrolysis experiments, the mercury pool electrode was dissolved in nitric acid, and the resulting solution was examined by the dimethylglyoxime adsorptive stripping method. An appreciable amount of cobalt was detected, which is to be expected when cobalt metal present in or on the mercury surface is oxidized to Co(II) by nitric acid. Additionally, a dark red solid, believed to be cobalt oxide, was observed on the mercury surface upon exposing the mercury pool to air immediately after the reduction period. Finally, a controlled-potential electrolysis experiment was performed at the mercury pool electrode on a solution made from a sample of synthesized Co(dmgH)2. Cyclic voltammograms before and after electrolysis are identical to those obtained from complexes generated in situ, and the number of electrons transferred in the reduction was calculated to be n ) 9.5 ( 0.5. This confirms that the complex generated in solution from mixture of Co(II) and dmgH2 is, indeed, Co(dmgH)2. All evidence, therefore, suggests that both the Co(II) metal center and the ligand are reduced. Examination of Electrolysis Products by Spectroscopic Techniques. Infrared Spectroscopic Study. Examination of products formed after reductive electrolysis is difficult with conventional infrared spectroscopic techniques, due to the presence of strong absorption from the aqueous solvent. However, it is possible to examine these solutions by Fourier transform attenuated total reflectance infrared (ATR) spectroscopy using a waterinsoluble zinc selenide crystal cell for the measurements. The ATR infrared spectrum of the ammonia buffer solution containing dmgH2 and Co(II) in a concentration ratio of 2:1 is dominated by the contribution from the background, which consists of a high concentration of ammonia buffer and a small concentration of ethanol associated with the dmgH2 addition. To obtain any information relevant to the electrolysis products, background subtraction was required. Thus, negative bands are associated with consumption of the starting material, and positive bands correspond to formation of new products. Negative bands at 1050 and 1100 cm-1 are attributable to a diminution of ethanol concentration27 during electrolysis as a result of slow evaporation and/or to the presence of N-O groups.35 The major new bands

are located at 1630 cm-1, and a group of peaks appear around 3300 cm-1. (i) Reduction of the Carbon-Nitrogen Double Bonds of the Ligand According to Eq 8. A total of four electrons per ligand and a total of eight electrons for the two ligands in the complex are expected according to this interpretation. IR evidence for 2,3bis(hydroxylamino)butane being a product of electrolysis is provided by the occurrence of IR bands attributable to NH groups at 3200 and about 1630 cm-1. When considering the shape of the bands in the 3200-3500-cm-1 range, it is possible that hydrogen intramolecular bands, such as N-H‚‚‚O or N-H‚‚‚N, are present in the spectrum. A negative IR peak would be expected for the consumption of the double bond at approximately 1600 cm-1. Unfortunately, this is difficult to detect in the IR spectrum because it is masked by the large positive peak at about 1630 cm-1, which is assigned to the formation of an NH group as expected according to eq 8. The negative band located in the 1400-cm-1 region may be assigned to the loss of the symmetric CdN stretching band.36 (ii) Reduction of the Oxime Group on the Ligand:

(19)

This seems unlikely to be the major pathway because there was no evidence of hydroxyl reduction in the IR spectrum. (iii) Complete Reduction of Only One of the Ligands in the Complex To Form the Amine (Eq 9). In this scheme, one ligand and the metal center are completely reduced, leaving one unreduced oxime group, but again no IR evidence to support this proposal was found. If the oxime groups were to be reduced to amine or imine groups, the decay of their bands appearing above 3000 cm-1 would have been expected. The great similarity between the IR spectra of the products resulting from the reduction of Co(dmgH)2 (electrolysis potential, - 1.25 V vs Ag/AgCl) and reduction of free dmgH2 ligand (electrolysis potential, - 1.8 V vs Ag/AgCl) is evidence that the ligand undergoes reduction in the complex and that, in both cases, the same product(s) is (are) obtained. Furthermore, results from electrochemistry and IR analysis suggest that neither reduction of the ligand in the cobalt complex to 2,3-diaminobutane (eq 9) nor reduction of only one dimethylglyoxime is likely to be a major pathway since, following the electrolysis with the ratio [dmgH2]: [Co(II)] ) 2:1, neither voltammetrically free dmgH2 nor the decay of OH IR bands was observed. After eliminating 2,3-diiminobutane and 2,3-diaminobutane as major products, it can be proposed that 2,3-bis(hydroxylamino)butane is a major product of reductive electrolysis of the Co(dmgH)2 complex. Mass Spectrometry Study. Solid dmgH2, Co(dmgH)2, and the precipitate formed from evaporation of solutions after bulk electrolysis were analyzed by the GC/MS technique. In the mass spectra of both dmgH2 and Co(dmgH)2, the parent peak corresponding to dmgH2 (m/z ) 116, relative intensity 100) and peaks corresponding to fragments of the decomposition of dmgH2 (m/z ) 99, relative intensity 69, C4H7N2O; m/z ) 84,

relative intensity 9, C4H6NO; m/z ) 68, relative intensity 21, C4H6N; m/z ) 58, relative intensity 24, C2H4NO; m/z ) 42, relative intensity 28, CNO) were identified. In the mass spectra of the electrolysis product(s), the five most intensive peaks have the following characteristics: m/z ) 87, relative intensity 27, C3H7N2O; m/z ) 85, relative intensity 6, C3H5N2O; m/z ) 73, relative intensity 81, C3H7NO or C2H5ON2; m/z ) 60, relative intensity 100, C2H6NO; m/z ) 45, relative intensity 13, CH3NO. Thus, they contain in their composition the hydroxylamine (-CH-NHOH) groups. The peak (m/z ) 60, relative intensity 100%, C2H6NO), which is consistent with half of the 2,3-bis(hydroxylamino)butane molecule, provided the highest intensity in the mass spectrum. The fact that the peak is not observed in the dmgH2 nor in the Co(dmgH)2 mass spectra again allows us to infer that a major product of bulk electrolysis is 2,3bis(hydroxylamino)butane. The absence of a parent peak for 2,3bis(hydroxylamino)butane (m/z ) 120, C4H12N2O2) in the spectrum of the electrolysis product is attributed to the ready and symmetrical splitting of the parent molecule under the conditions of mass spectrometry. It can also be noted that the mass spectrum of the electrolysis product does not fit that expected for other hypothetical products of electrolysis, such as oxygen-free 2,3diiminobutane or 2,3-diaminobutane, which might have been formed as a result of eqs 9 and 19. The results of the GC/MS study of solid dmgH2, Co(dmgH)2, and the product of electrolysis, therefore, also support the conclusion that a major reduction product of the is 2,3-bis(hydroxylamino)butane. Electrochemistry in Organic Solvents. To overcome solubility problems in organic solvents, a dimethylglyoxime analogue was examined in which a six-carbon backbone chain replaced the two methyl groups. The basic structure of the cobalt and nickel complexes of 1,2-cyclooctanedione dioxime is shown in structure b (Chart 1). These derivatives may be referred to as carbocyclic bis-dioxime complexes and are abbreviated as Co(C8doH)2 and Ni(C8doH)2. Co(dmgH)2‚2H2O. The partial solubility of Co(dmgH)2‚2H2O in dichloromethane allows a comparative electrochemical investigation in this solvent to be undertaken in order to verify that it is valid to extrapolate data obtained from the Co(C8doH)2‚2H2O analogue. The presence of waters of solvation, while omitted in further discussion, means that studies have actually been undertaken in the presence of traces of water. The initial reduction of the Co(dmgH)2 complex at a HMDE in dichloromethane, which occurs at about -2.3 V vs Fc/Fc+, is barely resolved from the solvent limit at 20 °C. However, when the temperature is lowered to -50 °C, when the mercury electrode is solid rather than liquid, two redox processes having some degree of chemical reversibility are observed (Figure 4a), although the second reduction process is barely resolved from the solvent limit. At -70 °C, both reduction process are extremely well defined (Figure 4b). The reversible E1/2 values are -2.25 V for the first process and -2.40 V vs Fc/Fc+ for the second process. Similar results were obtained when acetone was used as the solvent. When glassy carbon and platinum electrodes were used, the waves were masked by the solvent reduction process at all temperatures down to -70 °C. Voltammetric studies of the reduction of Co(dmgH)2 at mercury electrodes in aprotic solvents, therefore, indicate that two Analytical Chemistry, Vol. 70, No. 7, April 1, 1998

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Figure 4. Cyclic voltammograms for reduction of 1 × 10-3 M Co(dmgH)2 in dichloromethane (0.1 M Bu4NPF6) at a hanging mercury drop electrode and using a scan rate of 1 V s-1 at (a) -50 and (b) -70 °C.

primary one-electron transfers, which could be metal or ligand based, are available:

Co(dmgH)2 + e- h [Co(dmgH)2]-

(20)

[Co(dmgH)2]- + e- h [Co(dmgH)2]2-

(21)

Co(C8doH)2‚2H2O. The reduction of Co(C8doH)2 in dichloromethane, as is the case with Co(dmgH)2, also occurs close to the solvent limit and, again, is only well-defined at low temperatures at a solid frozen mercury electrode. In dichloromethane at a hanging mercury drop electrode at -40 °C, two reduction waves are again observed, with the second not being resolved from the solvent limit. The reversible E1/2 value for the first reduction process is -2.35 V vs Fc/Fc+, which is 100 mV more negative than the value for reduction of Co(dmgH)2. However, clearly, the voltammetry of the two complexes is closely related. An investigation of the effect of water addition on the reduction process was undertaken in acetone at 20 °C. Unlike the case with dichloromethane, water is very soluble in acetone, so the water addition experiments are readily undertaken in this solvent. Voltammograms obtained at a scan rate of 500 mV s-1 for reduction of 1 × 10-3 M Co(C8doH)2 at a hanging mercury drop electrode in acetone in the absence of deliberately added water, with 1% water added and with 2% water added, showed that reduction of the peak potential shifts to less negative potentials 1320 Analytical Chemistry, Vol. 70, No. 7, April 1, 1998

by about 300 mV from -2.45 to -2.15 V vs Fc/Fc+, and the peak height increases by a factor of 3 with addition of 2% water. Furthermore, in the presence of water, a new oxidation peak is generated on the reverse scan (-0.8 V vs Fc/Fc+ in the presence of 1% water), which has an oxidation potential that is dependent on the amount of water added and has characteristics similar to those of the oxidation peak observed after reduction of dimethylglyoxime in water. Similar results were obtained when a pH 9 ammonia buffer solution was added instead of pure aqueous solvent. The voltammogram for reduction of Co(C8doH)2 in a 100% water solution containing 0.1 M ammonia buffer at pH 9 exhibits only one large reduction peak and has most of the features observed for reduction of Co(dmgH)2 in aqueous media. The peak potential for the reduction process is shifted by about 600 mV to more positive values upon changing from pure organic to pure aqueous solvent. Analogous experiments based on addition of 0.2, 0.6, and 1% water to acetone solutions were undertaken under dc polarographic conditions with a drop time of 0.5 s. With increasing water content, not only is there a shift in the half-wave potential toward less negative potentials (-2.40 to -2.20 V vs Fc/Fc+ on addition of 1% water) and an increase in the limiting current (50% increase on addition of 1% water), but the formation of a polarographic maximum is also observed. Polarograms in acetone with addition of small amounts of added water, therefore, have most of the characteristics observed for reduction of Co(dmgH)2 in aqueous media. Clearly, water and acid-base reactions play an important role in the reduction of Co(dmgH)2. Ni(C8doH)2. Ni(C8doH)2 was found to be sufficiently soluble in dichloromethane that electrochemical experiments could be conveniently performed at a concentration of 10-3 M. The voltammetry of Ni(C8doH)2 in dichloromethane at fast scan rates at a HMDE exhibited a reversible reduction wave at -1.75 V vs Fc/Fc+. The use of platinum or glassy carbon electrodes extends the positive potential range so that oxidation processes also could be observed. Figure 5 shows cyclic voltammograms in dichloromethane at a platinum electrode at scan rates of 200 and 1000 mV s-1. With this electrode, an oxidation process at about +0.80 V vs Fc/Fc+ is observed as well as the reduction process at -1.75 V vs Fc/Fc+. However, apparently both the one-electron reduced and oxidized forms of Ni(C8doH)2 have a short half-life at room temperature, since fast scan rates (1000 mV s-1) are required in order to achieve chemical reversibility. When the voltammetry is performed in dichloromethane at a glassy carbon electrode, the reduction and oxidation redox processes each have a peak-to-peak separation of 110 mV and a peak height ratio approaching unity at a scan rate of 1 V s-1. At a rotating platinum disk electrode, the limiting currents for each process (apart from the sign) are almost identical. Experiments at a 10-µm-diameter platinum disk microelectrode under nearsteady-state conditions (scan rate, 10 mV s-1) show the equivalent result. Ferrocene, which is known to give rise to a one-electron oxidation process, has voltammetric characteristics which are very similar in dichloromethane to those found for the oxidation and reduction of Ni(C8doH)2. All these data suggest that the oxidation and reduction of Ni(C8doH)2 are one-electron processes in dichloromethane when short time domain experiments (fast scan

Figure 5. Cyclic voltammograms of 1 × 10-3 M Ni(C8doH)2 in dichloromethane (0.1 M Bu4NPF6) at 20 °C using a platinum disk electrode at scan rates of (a) 200 s-1 and (b) 1000 mV s-1.

rate cyclic rotating disk and steady-state microelectrode voltammetry) are used. The instability of the one-electron reduced species is apparent by noting the process present at -0.9 V vs Fc/Fc+ on second and subsequent cycles of cyclic voltammetric experiments. An ESR spectrum of a solution containing 5 × 10-3 M Ni(C8doH)2 in dichloromethane produced no signal. The electrochemical cell was then switched on and the potential held constant at a value sufficiently positive to oxidize the nickel complex. Since the half-life of the one-electron oxidized species is short, detection by the ESR instrument requires their in situ generation in the ESR cavity. This was achieved by using an electrolysis cell specifically designed to fit inside an ESR tube.32 A solution volume of 200 µL was used for the electrolysis, and all solution transfers were performed under an argon atmosphere to avoid reactions of products with oxygen. Under these conditions, a single line ESR signal at g ) 2.02 was detected (Figure 6a). The lack of nitrogen coupling suggests that a transient nickel(III) species is formed upon electrochemical oxidation. Platinum(III) compounds of this ligand are known,37 but they are dimeric and diamagnetic. The half-life of this oxidized species was estimated to be about 2 s by observing the decay of the ESR signal intensity at a constant magnetic field after switching off the potential. The initial oxidative electron-transfer process is, therefore, believed to be (34) Bobrowski, A. Electroanalysis 1996, 8, 79. (35) Burger, K.; Ruff, I.; Ruff, F. J. Inorg. Nucl. Chem. 1965, 27, 179. (36) Egneus, B. Talanta 1972, 19, 1387. (37) Baxter, A. M. L.; Heath, G. A.; Raftis, R. G.; Willis, A. C. J. Am. Chem. Soc. 1992, 114, 6944.

Figure 6. ESR spectra obtained at 22 °C in dichloromethane (0.1 M Bu4NPF6) after (a) electrochemical oxidation of Ni(C8doH)2 and (b) electrochemical reduction of Ni(C8doH)2: (i) initial spectrum, (ii) spectrum after 20-min reduction, and (iii) spectrum after 40-min reduction.

described by eq 22:

NiII(C8doH)2 h [NiIII(C8doH)2]+ + e-

(22)

where electron transfer occurs predominantly via the metal center and not the ligand. This mechanism is supported by X-ray photoelectron spectroscopic studies.38 To study the reduction of Ni(C8doH)2 in dichloromethane by the electrochemical ESR method, the potential was held constant at a value which was sufficiently negative to reduce the complex. Multiple ESR scans were performed, and the signal was accumulated in order to detect the ESR signal of the reactive species. Initially, a weak nine-line ESR signal was observed with g ) 2.09 and an intensity ratio of 1:2:3:4:5:4:3:2:1 (Figure 6b). The hyperfine coupling is consistent with the interaction of the electron with four equivalent nitrogen atoms, which would be expected for a paramagnetic nickel complex in which four nitrogen atoms surround the central metal atom. As the reduction reaction proceeds, a three-line ESR signal of g ) 2.01, of intensity 1:1:1 Analytical Chemistry, Vol. 70, No. 7, April 1, 1998

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(Figure 6b), is observed and grows at the expense of the nineline spectrum until, finally, the only signal evident in the ESR spectrum is the three-line signal (Figure 6b). This three-line signal is consistent with an organic radical species in which an electron is coupled with one nitrogen atom only. The ESR experiments combined with voltammetric data, therefore, suggest that the initial step in the reduction process has considerable ligand character, while the oxidation process is predominantly metal based. The extent to which the processes are metal or ligand could, conceivably, be determined from ESR data obtained after synthesis and subsequent oxidation and reduction of 61Ni (I ) 3/2)-enriched complexes. In aqueous media, reactions of the one-electron reduced initial product with water and/or protons presumably give rise to the multielectron reduction process observed in aqueous media. CONCLUSIONS The results obtained during the course of this study suggest that the electrochemical reduction of cobalt and nickel dimethylglyoxime complexes in aqueous media involves the overall reduction of both the central metal atom and the surrounding ligands in an overall 10-electron reduction process. In aprotic solvents such as dichloromethane, the nickel and cobalt complexes undergo an initial one-electron reduction which is believed to have significant ligand character. However, in water, electron transfer is coupled with proton transfer, giving rise to a stable electrolysis product, 2,3-bis(hydroxylamino)butane. This is evidenced by the number of electrons obtained during coulometric and microcoulometric experiments and by identification of electrolysis products by IR spectroscopy and mass spectrometry, as well as the voltammetric observation of a new process after reduction. Disruption of the ligand, liberating M2+, must then lead to M(0) at the prevailing potential. Hence, during the determination of cobalt or nickel by adsorptive stripping voltammetry, the predominant overall process that occurs at the mercury electrode surface is as follows: solution

M(II) + 2dmgH2 y\z [MII(dmgH)2] + 2H+ (23) electrode

[MII(dmgH)2] + Hg y\z [MII(dmgH)2]adsHg

(24)

[MII(dmgH)2]adsHg + 10e- + 10H+ f M(Hg) + 2DHAB (25)

where M represents Co or Ni and DHAB is 2,3-bis(hydroxylamino)butane. Additional support for the mechanism is the polarographic behavior of dmgH2 in alkaline medium, coupled with an earlier finding that dmgH2 undergoes a four-electron reduction, with 2,3bis(hydroxylamino)butane being the possible reduction product.39 Our data also are consistent with the prediction of Nambiar and Subbaraman26 that the ligand is reduced in the Co(dmgH)2 complex during electrolysis in neutral media and that the product is 2,3-bis(hydroxylamino)butane. The half-wave potential of -0.43 (38) Young, V. Y.; Chang, F. C.; Cheng, K. I. Appl. Spectrosc. 1986, 41, 636. (39) Bossa, V.; Morpurgo, G.; Morpurgo, L. Ric. Sci. 1967, 37, 402.

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V40 for the oxidation process at a mercury electrode in alkaline solution also is consistent with that found for the reduction product (of dmgH2 at -1.7 V vs Ag/AgCl or of Co(dmgH)2 or Ni(dmgH)2 at about -1.3 V vs Ag/AgCl), being 2,3-bis(hydroxylamino)butane. Quantum mechanical calculations by Ni et al.41 on the Ni(dmgH)2 complex emphasize the proximity of the metal and ligand orbitals. Combinations of metal- and ligand-based reductions are, therefore, consistent with these theoretical results. The results proposed in this work also may be applied to other results related to the electrochemical reduction of Ni(dmgH)26 or Co(dmgH)2.4 In the cited works, the anomalously high maximal surface concentration for the nickel complex, Γ ) 1.1 × 10-9 M cm-2, was calculated from charge measurement under the assumption that two electrons are transferred in the overall electrode process. For the value ne ) 10, determined in our experiments, the recalculated value, Γ ) 2.2 × 10-10 M cm-2, lies within the normally expected range. It is apparent from our results that reduction of Co(dmgH)2 and Ni(dmgH)2 complexes does not occur with catalytic evolution of hydrogen as stated by a number of researchers.11-17,42 That is, there is no evidence of hydrogen formation during exhaustive coulometry, and the reduction products observed are not consistent with the proposed mechanism associated with catalytic reduction of hydrogen. Likewise, hydrogen is not observed during reduction of Co(C8doH)2. Weinzierl and Umland7 predicted that the ligand in the complexes becomes reduced. However, they assumed that the process occurred according to eqs 10 and 11 and that the metal center is released to react with a further 2 mol of dmgH2 and is not reduced. This has been disproved in this study, since cobalt metal is identified in the mercury pool as a product after exhaustive reduction. The mechanism summarized by eqs 4-6, and proposed in ref 6, also appears to be inadequate, because it describes only part of the electrode process and does not include the reduction of the ligand. Also, the mechanism proposed by Jin and Lin4 is not correct, since it is based on their experimental observation that, during microelectrolysis, the number of electrons involved in overall reduction reaction is 2. The recent paper by Vukomanovic et al.,9 which provides extensive data on the nickel dimethylglyoxime system in ammonia buffer media, led to the conclusion that a 16- or 18-electron reduction is involved. This result implies that dimethylglyoxime is reduced in an eight-electron step in alkaline media to give 2,3-diaminobutane. The insolubility of the nickel system in water makes this a more difficult system to study. However, all of our data on the more tractable cobalt system are completely consistent with a 10-electron reduction and with 2,3bis(hydroxylamino)butane being formed. That is, our results strongly favor an overall process combining partial (4e-) reduction of each ligand with exhaustive reduction of the metal ion. This result is completely consistent with the recent findings of Jagner et al.,10 who also report that an overall 10-electron reduction occurs. Our observations in organic solvents also do not support the hypothesis25,26,38 that Co(I), Co(O), or the Ni(I) intermediates (40) Meites, L. Polarographic Techniques; Interscience: New York, 1955. (41) Ni, Y.; Ren, J.; Li, J.; Wang, D.; Liang, W.; Zhu, Z.; Gao, X. Sci. China 1990, 33, 393. (42) Prokhorova, G. V.; Osipova, E. A.; Torocheshnikova, I. I.; Spighun, L. K. Vestn. Moskov. Univ. Khim. 1991, 32, 1991 and references therein.

induce the catalytic reduction of the ligand in the presence of a proton source such as water.

which facilitated the international collaboration required for this project.

ACKNOWLEDGMENT The authors gratefully acknowledge financial support from the Research School of Chemistry (ANU), the Australian Research Council, and the University of Mining and Metallurgy (Poland),

Received for review April 3, 1997. Accepted December 23, 1997. AC9703616

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