2810
J. Phys. Chem. 1988, 92, 2810-2816
Electrocatalytic Reduction of Viclnal Dibromides by Vitamin B,, Thomas F. Connors,la James V. Arena,lb and James F. Rusling* Department of Chemistry (U-60), University of Connecticut, Storrs, Connecticut 06268 (Received: July 20, 1987; In Final Form: December 15, 1987)
Alkyl vicinal dibromides are reduced by cob(1)alamin generated at a glassy carbon cathode in weakly acidic aqueous acetonitrile in a two-electron electrocatalytic cycle with a lowering in overpotentialof 0.84 V compared to direct reduction. Unlike reactions of primary alkyl halides with cob( 1)alamin that yield stable alkylcobalamins, 1,2-dibromobutane and 1,2-dibromoethane are reduced completely to alkenes at potentials somewhat positive of the Eo’ of Co(II)/Co(I). Spectroelectrochemistry showed that the reaction between Co(1) and 1,2-dibromobutane is spontaneous and regenerates Co(I1). Analysis of voltammetric data by digital simulation/nonlinear regression gave a rate constant of 6 X lo6 L mol-’ s-’ for this rate-determining reaction. A previously developed simulation model for second-order kinetics was used, extended to include intermediate rates of heterogeneous charge transfer and a difference in diffusion coefficients for substrate and catalyst. Intermediacy of a (8-bromoalky1)cobalamin is likely but was kinetically invisible on the millisecond time scale of our experiments. At large excess of substrate, trace crossing observed in cyclic voltammograms may suggest formation of a readily reducible bromobutyl radical in an autocatalytic process.
Vitamin BIZ,, or aquocobalamin, is a Co(II1) corrin complex with the structure shown in Figure 1. It is related to the biologically active adenosylcobalamin, in which adenosine replaces water as an axial ligand. In conjunction with the proper enzyme system, the latter compound catalyzes a variety of group-transfer and other unique reactions in living systemsS2 There is an extensive body of literature on the chemistry of vitamin B12,3*4 and there has been recent interest in its catalytic properties for organic reduction^.^ The catalytic utility of vitamin Blzais due to its facile chemical or electrochemical reduction to BIZa,or cob(I)alamin, a powerful nucleophile that is readily alkylated to yield alkylcob(II1)alamins. The ability to form Co-C bonds provides an inner-sphere pathway for organic reductions, requiring cleavage of the Co-C linkage to obtain reduced organic products. Our interest in vitamin BI2 as a catalyst was further stimulated by reports of its possible role in the biodegradation of organo halide pesticides in the environment6 and in metabolism of small carbon-containing molecules in methanogenic bacteria.’ The pathway for electrocatalytic reduction of such molecules by vitamin BI2 may be relevant to the mechanisms of these reactions in biological systems. These catalytic systems are also of interest for analytical and synthetic applications. A detailed picture of the redox behavior of Vitamin B12has emerged from the work of Lexa and Saveants and othem9 Co(I1) (1) Present addresses: (a) Colgate-Palmolive Co., R. & D. Div., Piscataway, N.J.; (b) Miles Pharmaceuticals, West Haven, CT. (2) E,,, Dolphin, D.; Ed.; Wiley: New York, 1982; Vol. 11. (3) (a) Halpern, J. In 4,;Dolphin, D., Ed.; Wiley: New York, 1982; Vol. I, pp 501-541. (b) Hogenkamp, H. P. C. In E,,; Dolphin, D., Ed.; Wiley: New York, 1982; Vol. I, pp 295-323. (c) Golding, B. T. In E,*; Dolphin, D., Ed.; Wiley: New York, 1982; Vol. I, pp 543-582. (d) Pratt, J. M. In E I 2 ; Dolphin, D., Ed.; Wiley: New York, 1982; Vol. I, pp 325-392. (e) Brown, K. L. In EI2;Dolphin, D., Ed.; Wiley: New York, 1982; Vol. I, pp 245-294. (4) (a) Schrauzer, G. N.; Deutsch, E. J. Am. Chem. Soc. 1969, 91, 3341-3350. (b) Grate, J. H.; Schrauzer, G. N. J . Am. Chem. Soc. 1979, 101, 460 1-46 11. ( 5 ) There is a growing literature on this topic, recently reviewed by Scheffold (In Modern Synthetic Methods; Scheffold, R., Ed.; Wiley: New York, 1983; Vol. 3, pp 355-439. (6) Stotter, D. A. J. Inorg. Nucl. Chem. 1979, 39, 721-727. (7) Zeikus, J. G.; Kerby, R.; Krzycki, J. A. Science (Washington, D.C.) 1985, 227, 1167-1 173. (8) (a) Lexa, D.; Saveant, J. M. Acc. Chem. Res. 1983.16, 235-243. (b) Lexa, D.; Saveant, J. M. J. Am. Chem. SOC.1976,98,2652-2658. (c) Lexa, D.; Saveant, J. M.; Zickler, J. J . Am. Chem. Soc. 1979, 99, 2786-2790. (d) Lexa, D.; Saveant, J. M. J. Am. Chem. SOC. 1978, 100, 3220-3222. (e) DeTacconi, N. R.; Lexa, D.; Saveant, J. M. J. Am. Chem. Soc. 1979, 101, 467-473. (0 Lexa, D.; Saveant, J. M.; Zickler, J. J . Am. Chem. SOC.1980, 102, 2655-2663. (g) Lexa, D.; Saveant, J. M.; Zickler, J. J . Am. Chem. SOC. 1980, 102, 4851-4852. (h) Lexa, D.; Faure, D.; Saveant, J. Electroanal. Chem. 1982, 140, 269-284,285-295.297-309, (i) Lexa, D.; Saveant, J. M.; Soufflet, J. P. J. Electroanal. Chem. 1979, 100, 159-172.
0022-3654/88 /2092-2810$01.50/0
[cob(II)alamin or BI2J and Co(1) oxidation states are electrochemically accessible at potentials positive of -1 V versus SCE. In aqueous media, the Co(II1) and Co(I1) complexes can exist in “base-on” (Figure 1) or “base-off“ forms, the latter having protonated benzimidazole side chains. Electron transfer from an electrode is significantly faster to the base-off form of cob(I1)alamin, the predominant Co(I1) species in aqueous media below pH 2.9. Reduction of this species yields a four-coordinate cob(I)alamin,5~s which reacts with primary alkyl monohalides to yield stable alkylcob(III)alamins.3-5~*i The latter can be decomposed by light, heat, or electrochemical potentials about 0.7 V more negative than the Eo’ of Co(II)/Co(I). Alkylcobalamins with a good leaving group such as OH in the alkyl ligand’s & x ~ i t i o n ~ - ~ are more easily decomposed than their unsubstituted analogues. Many toxic organo halides in the environment contain halogen on adjacent carbons. An example is the suspected carcinogeni0 1,2-dibromoethane, or ethylene dibromide (EDB). Direct electrochemical reduction of vicinal dibromides such as EDB requires potentials of about -1.6 V versus SCE.’I However, in solutions of cob( III)alamin, vicinal dibromides may form (bromoalky1)cobalamins stable at the potential of the Co(II)/Co(I) redox Alternatively, couple, about -0.75 V versus SCE in acidic the dibromide could react spontaneously with Co(1) in a process leading to hydrocarbon products at the Co(II)/Co(I) potential. Our goals in studying electrocatalytic reduction of vicinal dibromides with vitamin B12 were to explore the pathways for possible formation of hydrocarbon products. In this article, we report the results of electrochemical and spectroelectrochemical studies on the reduction of 1,2-dibromobutane and 1,2-dibromoethane by electrochemically generated vitamin BIZ,.
Experimental Section Chemicals. Vitamin Blzawas obtained as hydroxocobalamin hydrochloride (crystalline, 99%) from Sigma Chemical Co. and used as received. Purity was confirmed by UV-vis spectroscopy. 1,2-Dibromoethane and 1,2-dibromobutane were obtained from Aldrich Chemical Co., and they were used as received. Spectrograde acetonitrile was obtained from Baker Chemical Co. and used as received. n-Butene standard calibrating gas was from Ideal (9) (a) Jaselskis, B.; Diehl, H. J. Am. Chem. SOC.1954, 76, 4345-4348. (b) Tackett, S . L.; Ide, J. W. J. Electroanal. Chem. 1974, 30, 510-514. (c) Swetik, P. G.; Brown, D. G. J. Electroanal. Chem. 1974, 51, 433-439. (d) Birke, R. L.; Brydon, G. A,; Boyle, M. F. J. Electroanal. Chem. 1974, 52, 237-249. (e) Birke, R. L.; Venkatesan, S. J . Electrochem. SOC.1981, 128, 984-991. (f) Birke, R. L.; Gu, R.-A.; Yau, J.-M.; Kim, M.-H. Anal Chem. 1984, 56, 1716-1722. (10) Van Bladeren, P. J. J . Am. Coll. Toxicol. 1983, 2, 73-78. (11) (a) Casanova, J.; Rogers, H. R. J . Org. Chem. 1974, 39, 2408-2410. (b) Tokoro, R.; Bilewicz, R.; Osteryoung, J. Anal. Chem. 1986, 58, 1964-1969.
0 1988 American Chemical Society
Reduction of Vicinal Dibromides by Vitamin B12
OH2
I*"%
Figure 1. Structure of vitamin BIzain the base-on form.
Gas Products (f2% accuracy, reference no. 119-57). Distilled water was further purified with a Barnstead Nanopure system and had a specific resistance 2 1 5 MR cm. Apparatus. A Bioanalytical Systems BAS- 100 electrochemical analyzer was used for cyclic voltammetry (CV) and spectroelectrochemistry. A three-electrode, amber glass cell, with a glassy carbon working electrode, platinum wire counter electrode, and an aqueous saturated calomel reference electrode (SCE) were used for CV. The reference electrode was connected to the sample compartment by a salt bridge containing supporting electrolyte. The planar glassy carbon electrodes ( A = 0.088 or 0.071 cm2) were constructed and polished as described previously.12 For spectroelectrochemical experiments, the cell consisted of a 1-cm quartz cuvette with Pt wire counter and Ag/AgCl reference electrodes connected via agar/KCl salt bridges in Pasteur pipet tips extending through holes drilled in the cuvette top. An additional hole in the cell top afforded insertion of a nitrogen inlet tube. The working electrode was a 0.9 X 1.3 X 0.15 cm piece of 60 pore/in. reticulated vitreous carbon (Chemotronics Int.) placed directly in the light path. This cell transmitted about 70% of the incident light. A matching cuvette containing electrolyte solution and a piece of reticulated carbon was placed in the reference beam to balance the spectrophotometer. A Princeton Applied Research Corp. (PARC) Model 363 potentiostat/galvanostat was used for controlled-potential electrolyses (CPE). A gas-tight, three-electrode cell was used, with 2 X 10 X 0.5 cm WDF carbon felt (Union Carbide) as the working electrode. A spectroscopic carbon rod (Ultracarbon Co.) in a compartment separated from the reaction chamber by a medium-porosity glass frit was the counterelectrode. The reference was an SCE. The cell top contained ground-glass openings for electrodes, the nitrogen inlet, and the gaseous outlet to a collection trap. The top fit snugly onto the cell by a ground-glass connection. Gas chromatography/mass spectrometry (GC/MS) was done with a Hewlett-Packard 5985B GC/MS system. Gaseous mix(12) Kamau, G. N.; Willis, W. S.; Rusling, J. F. Anal. Chem. 1985, 57, 545-551.
The Journal of Physical Chemistry, Vol. 92, No. IO. 1988 2811 tures were separated with an Analabs Spherocarb (80/100 mesh), 1/8-in.o.d., 3-ft stainless steel column, with conditions previously described for separation of light hydrocarbon~.]~Ether extracts of electrolyzed solutions were analyzed with an S E 54 Chromosorb W P H 8100, 1/8-in.o.d., 6-ft stainless steel column. UV-vis spectrophotometry was done with Varian Cary 17D or Perkin-Elmer Lambda 3B spectrophotometers. Wavelength was calibrated with holmium oxide.14 A Corning glass electrode and a Corning Model 130 digital pH meter were used to measure apparent pH. Procedures. Stock solutions of vitamin BI2& were prepared with ultrapure water in amber glass volumetric flasks and were stable for up to 2 months at ambient temperature as shown by CV and UV-vis spectroscopy. Buffer solutions were prepared on the day of the experiment by combining the appropriate amounts of stock 1.3 M H3P04/NaH2P04buffer solution (pH 2.3 or 2.9), water, and acetonitrile. The final buffer concentration was 0.1 M in water/acetonitrile (1:1, v/v). The desired concentrations of vitamin BI2*and dibromide were obtained by adding appropriate amounts of stock solutions to the cell. Solutions were thoroughly purged with purified nitrogenI5 before use. Glassy carbon electrodes were polished successively on the day of use with silicon carbide particles, diamond paste, and alumina on a Buehler metallographic polishing wheel with procedures described previously.12 Before each scan, the electrode was polished for 2 min with 0.3-pm alumina, ultrasonicated for 2 min in purified water, polished with 0.05-pm alumina, and finally ultrasonicated again in pure water. Reproducibility on the reversible oxidation of 1.7 mM ferrocene in acetonitrile16 was f 3 mV for peak potentials and f5% for peak currents and was similar for vitamin BI2. With the apparatus used, the anodic-cathodic peak separation for ferrocene was 63 f 1 mV, and this value was used as a standard for heterogeneous rate constant estimates from peak separations (see Results). Ohmic drop of the electrochemical cell was fully compensated by the BAS-100 in all voltammetric experiments. Unless otherwise stated, all experiments were done at ambient temperature (23 f 2 "C). Reticulated carbon working electrodes for spectroelectrochemistry were pretreated with 6 M nitric acid, followed by extensive washing with water. All solutions were thoroughly purged with purified nitrogen before each experiment and during electrolyses. Solutions for CPE at carbon felt electrodes were purged in a similar way and stirred during electrolyses. For collection of products, the outlet from the gas-tight cell was directed into a U-tube containing glass beads immersed in an acetone/dry ice bath (-77 "C). The condensate collected in the U-tube was dissolved in 1 mL of cold hexane and analyzed by GC/MS. Ether extracts of the initial and final electrolysis solutions were also analyzed by GC/MS. Conditions for the GC/MS analysis of gaseous products were as follows: helium flow rate, 40 mL/min; temperature program, -20 "C hold for 3 min, followed by a 10 "C/min ramp. A known mixture of butene and acetonitrile in hexane served as the standard. Standard and sample solutions at -77 "C were purged into the gas sample port of the G C with helium at 50 mL/min. Retention times relative to hexane were 0.47 for acetonitrile and 0.53 for butene. Mass spectra (1.2 s/spectrum) were taken four times during elution of each G C peak. Column conditions for analyzing ether extracts were a helium flow rate of 30 mL/min and a temperature program beginning at 40 "C and increasing 15 OC/min. Ether extracts and vicinal dibromide standards were injected directly into the GC. Computations. Expanding space-grid digital simulations and nonlinear regressions of voltammetric data onto simulation models were done with programs developed previously." A FORTRAN (13) Robillard, M. V.; Siggia, S.; Uden, P. C. Anal. Chem. 1979, 51, 435-439. (14) The Encyclopedia of Chemical Elements; Hampel, C . A,, Ed.; Reinhold: New York, 1968; p 227. (15) Rusling, J . F. J . Elecrroanal. Chem. 1981, 125, 447-458. (16) Kuwana, T.; Bublitz, D. E.; Hoh, G. J . Am. Chem. SOC.1969, 82, 581 1-5817.
2812
The Journal of Physical Chemistry, Vol. 92, No. 10, 1988 T tora
A
Connors et al. TABLE I: Experimental and Simulated" Voltammetric Characteristics for 1.7 mM Vitamin BI2JBlb Couple in pH 2.3 Buffer on Glassy Carbon
-Ep,
E IUOLT 1 Figure 2. CV of 1.7 m M vitamin BIzs in pH 2.3 aqueous acetonitrile (1:l) phosphate buffer at glassy carbon electrode at 0.50 V s-l.
77 version of a general program for nonlinear regressionlg based on the Marquardt algorithmL9was employed. This program was modified for our purposes by writing the desired equation or model to be fit to the experimental data in a subroutine called YCALC. (These subroutines are available from us.) Simulation models previously used for two-electron electrocatalysis needed to be modified to apply to planar diffusion and to include different diffusion coefficients for catalyst and substrate. Thus, the simulations employed the same equations as in ref 17a except that (1) the surface area of the outer boundary of the ith space element became cyi = A (electrode area), (2) the volume of the ith space element became 5 = A(xi - x,,), where xi is the distance of the boundary of the ith element from the electrode, and (3) separate equations for computing AC in the ith element were used for each diffusing species. The diffusion coefficient of the catalyst was fixed a t that measured by CV, and that of the substrate chosen to best-fit experimental catalytic currents. Usually, 45 i-E pairs 5 mV apart from the foot of the catalytic peak to a potential past its maximum were used for the nonlinear regressions. The Co(II1) peak was small and far enough away relative to the catalytic peak to consider the base line linear with slope and intercept determined by the regression ana1ysis.l' All computations were done with double precision on a IBM 3081D computer.
Results Initial Electrochemical Studies. Although reduction of vitamin B12has been studied extensively at mercury, platinum, and gold e l e ~ t r o d e s ,little ~ . ~ work has been done at highly polished glassy carbon.I2 We chose the latter electrode for the majority of our voltammetric studies to avoid hydrogen evolution and chemical reactions of alkyl dihalides and their reduction intermediates (e.g., on Hg)20 associated with other electrode materials. Acetonitrile-water (l:l, v/v) was chosen as solvent to accommodate solubility needs of the water-soluble vitamin and nonpolar substrates. The base-off form is required to ensure fast electron transfer to cob(II)alamin, necessitatingga p H values below 2.9. Voltammograms on glassy carbon at pH 1 (0.1 M perchloric acid) reflected catalytic reduction of hydrogen ions at the reduction potential of C O ( I I ) , ~and ~ , ~this ~ medium was not pursued further. Cyclic voltammograms (CV) of vitamin BLhat pH 2.3 on glassy carbon between scan rates of 0.01 and 50 V s-l showed slow Co(III)/Co(II) reduction but relatively fast Co(II)/Co(I) reduction (Figure 2). The slope of the linear variation in peak current with square root of scan rate ( u ) at u € 0.12 V s-] was (17) (a) Arena, J. A.; Rusling, J. F. Anal. Chem. 1987, 58, 1481-1488 (b) Arena, J. A.; Rusling, J. F. J. Phys. Chem. 1987, 91, 3368-3373. (18) NLLSQ Nonlinear Least Squares Program; CET Research Group, Norman, OK, 198 1. (19) Marquardt, D. W. J. SOC.Indust. Appl. Math. 1963, 11, 431-441. (20) (a) Hawley, M . D. In Encyclopedia of Electrochemistry of the Elements: Bard, A. J., Lund, H., Eds.; Dekker: New York, 1980 Vol. XIV. (b) Fry, A. J. Synthetic Organic Electrochemisrry; Harper and Row: New York, 1972; pp 170-187. (c) Hill, H. A. 0.; Pratt, J. M.; ORiordan, M. P.; Williams, F. R.; Williams, R.J. P. J. Chem. SOC.A 1971, 1859-1862. These authors isolated alkyl mercury compounds after controlled-potential electrocatalytic reduction of alkyl halides by vitamin BI2at mercury pool cathodes
V vs S C E
io+,, A
scan rate, v s-I
found
calcd"
found
calcd"
0.025 0.050 0.10 0.50 1.oo 5.12 10.24
0.803 0.801 0.800 0.803 0.819 0.824 0.851
0.796 0.797 0.798 0.804 0.808 0.825 0.836
1.1 1.6 2.3 4.7 6.7 15.2 17.6
1.1 1.5 2.1 4.6 6.3 13.5 18.7
60 62 59 68 108 102 144
ipalipc 1.o 1.o 1.o 1.o 1.1 1.1 0.9
Digital simulation parameters: ko = 0.02 cm s-l, Eo' = 0.766 V, D X cm s-I; A = 0.071 cm2.
= 2.7
T
1MVFI
A
-0.5
-I.O
E LUOLTI
Figure 3. CV of 1.7 m M vitamin BIzaand 1.7 m M 1,2-dibromobutane (0.50 V s'l). Conditions are as in Figure 2.
usedz1to estimate the diffusion coefficient (0) of vitamin B12 from
the Randles-Sevcik equation. Anodic-cathodic peak separations (Table I) were used with working curves22a,bto obtain apparent standard heterogeneous rate constants ( k o ) ,by assuming electrochemical transfer coefficients of 0.5 (Table 11). Cathodic peak currents and potentials computed by digital simulation" of quasireversible charge transfer using the estimated values of k" and D were in good agreement with experimental results (Table I) at 0.025 d u d I O V s-l. Vitamin B12r/B12s electrochemistry was also investigated by double potential-step chronocoulometry, which allows uncorrelated estimates of diffusion coefficient and surface concentrations of possible adsorbed species. Chronocoulometric charge versus time data gave excellent fits to the modelZZC for purely diffusion-controlled charge transfer with an average D = 2.75 X 10" cm2 s-l. The latter value agrees very well with that found by CV (Table 11). Adsorption of reactant or product was not detectable by chronocoulometry. The above results suggest that the diffusion-controlled charge-transfer model is satisfactory to describe vitamin BlZr/Bl2electrochemistry under our experimental conditions. Effects of catalytic hydrogen evolution or adsorption of the redox species are not detectable. Furthermore, values of D are consistent with those found in water (Table II), which has a slightly larger viscosity than the MeCN/water medium. When 1,2-dibromobutane (DBB) was added in equimolar concentration to the pH 2.3 vitamin Blzasolution, the current of the Co(1I) reduction peak increased a b o u t eightfold at 0.5 V s-l, and the anodic peak for oxidation of Co(1) disappeared (Figure 3). This behavior is characteristic for homogeneous electrocat a l y ~ i s .A~ peak ~ ~ is not observed for Co(1) because it reacts with substrate to recycle Co(I1) at the electrode, thus increasing the cathodic current. The catalytic peak occurred at more positive (21) Connors, T. F. Ph.D. Thesis, University of Connecticut, 1986. (22) (a) Nicholson, R. S. Anal. Chem. 1965,37,1351-1355. (b) Amatore, C.; Saveant, J. M.; Tessier, D.J. Electroanal. Chem. 1983, 146, 37-45. (c) Rusling, J. F.; Brooks, M. Y. Anal. Chem. 1984, 56, 2147-2152. (23) (a) For two-electron redox electrocatalysis with an initial fast (reversible) electron transfer at the electrode, double peaks are just detectable23b at X = 200, or k , / u = 5 X lo6 L mol-l V-I with equal concentrations of catalyst and substrate. (b) Andrieux, C. P.; Blocman, C.; Dumas-Bouchiat, J. M.; M'Halla, F.; Saveant, J. M. J. Elecrroanal. Chem. 1980, 213, 19-40.
The Journal of Physical Chemistry, Vol. 92, No. 10, 1988 2813
Reduction of Vicinal Dibromides by Vitamin B12
TABLE I 1 Thermodvnamic and Kinetic Constants for Electron Transfer to Cobalamins at Classy Carbon Co(III)/Co(II) medium (electrodel pH 2.3, aq MeCN (glassy carbon) 0.1 M HC104 (Hg, ref 8a) pH 4-7 ( W E ) 0.02 M KCI
106D. cmz s-I
ko. cm s-I 0.001 f 0.0007 8 X 10"
2.7 1.4
Co(II)/Co(I) koa
EO'.
V vs SCE 0.0" -0.04
cm s-I 0.022 k 0.01 1 0.10
EO'.
V vs SCE -0.766 -0.74
2.0 1.6b
"Estimated as midpoint between anodic and cathodic CV peaks. b B y capillary method in ref 8e; other determinations by voltammetry.
T
T
T
016 A
1 1
t0.m
to.0
4.5
I
T
b
1
8
1
-1 .o
i\ C
b
A, nm 1
t0.m
T
to.0
-0.5
C
SUA
6
8
1
-1 .o
/P
E[VOLTI Figure 4. CVs of 1.7 mM vitamin BIzaand 17 mM 1,2-dibromobutane with conditions as in Figure 2 except scan rate: (a) 1, (b) 5, and (c) 16 v s-'.
potentials than Eo' of Co(II)/Co(I), and a secondary peak was observed at the potential where Co(I1) was reduced in the absence of substrate. Similar results were obtained upon adding 1,2-dibromoethane to the Blzasolutions. With this more volatile substrate, the catalytic current decreased, and the more negative cathodic peak became relatively larger as substrate was removed from solution by purging with nitrogen. Double peaks are pred i ~ t e din~homogeneous ~ redox electrocatalysis for large values of the kinetic parameter X = RTk,C*/(Fu), where R is the gas constant, T is absolute temperature, kl is the chemical rate constant for reaction of reduced catalyst with substrate, C is the bulk concentration of catalyst, and F is Faraday's constant. At very large A, only a small amount of catalyst is needed to drive the reaction, so it occurs at potentials positive of the Eo' for Co(II)/Co(I). Thus, the first peak corresponds to total consumption of substrate in the reaction layer close to the electrode. When the potential of the electrode approaches the Eo' of Co(II)/Co(I), the unreacted Co(I1) is reduced in the second, diffusion-controlled peak. No change in the Co(III)/Co(II) peaks were observed upon addition of substrate. Large increases in cathodic Co(I1) current and disappearance of the anodic Co(1) peak were also observed when a 10- or 20-fold molar excess of EDB or DBB was added to the pH 2.3 or 2.9 solutions of vitamin Blk However, as expected with a large excess of substrate,23bonly single catalytic peaks were found. Furthermore, on the reverse scan the anodic i-E trace crossed over that of the cathodic trace (Figure 4) at u > 0.5 V s-l. Similar features were found for the above systems in voltammograms at hanging-drop mercury electrodes (HDME), but substrate was slowly removed from solution by a reaction with mercury.20 Catalytic reduction of 1,2-dibromoethane was also briefly investigated on an H D M E in N,N-dimethylformamide/propanol,
Figure 5. UV-vis spectroelectrochemistryof 0.025 mM vitamin Biz, in pH 2.3 aqueous acetonitrile buffer: (a) before electrolysis, (b) after 40 min of electrolysis at -1.2 V versus SCE, (c) after 105 min of electrolysis.
where similar voltammetric characteristics were observed. In the latter system, trace crossing with higher currents and at lower scan rates were found. Dissolving vicinal dibromides in dimethylformamide/2-propanolcontaining tetrabutylammonium iodide gave red solutions, suggesting formation of iodine during reduction of the dibromide by iodide in a well-known @-elimination reaction yielding alkene.24a Bulk Electrolysis and Product Identification. Because of the volatility of EDB, electrolyses were done only on DBB. Solutions 1.7 mM in vitamin Blzaand 120 mM in DBB were electrolyzed in the gas-tight cell at a potential on the rising portion of the catalytic wave for 8 h. A Baeyer's test (alkaline permanganate)24b on the gaseous effluent was positive, the decoloration of permanganate suggesting formation of alkene. A control experiment with no applied potential but with other conditions identical gave a negative Baeyer's test. Separate electrolyses were run in which the gaseous effluent was trapped at -11 OC and analyzed by GC/MS. Comparison of retention times and mass spectra with those of standards showed butene to be the sole gaseous product. Both standard and product butene GC peaks showed characteristic M S peaks at 56, 41, 39, and 28 amu. Analysis of ether extracts of solutions before and after electrolyses revealed only one G C peak, corresponding to 1,2-dibromobutane. About 10% of the substrate was lost by evaporation. Integration of G C peaks as well as catalytic CV peaks after electrolysis indicated that approximately 90% of the remaining DBB reacted in an 8-h electrolysis. This corresponds to conversion of 57 mol of substrate/mol of catalyst present. CV after exhaustive electrolyses indicated no decomposition of the catalyst. Large background currents at the carbon felt cathode precluded accurate determination of the charge passed through the cell, but the approximate value was many times greater than 2 F/mol of catalyst. Spectroelectrochemistry. UV-vis absorption spectra25of vitamin BIZcan be used to follow changes in oxidation state8cvi,26 (24) (a) Gould, E. S. Mechanism and Srrucrure in Organic Chemistry; Holt, Reinhart, and Winston: New York, 1959; pp 494-495. (b) Shiner, R. L.; Fuson, R. C.; Curtin, D. Y. The Systematic Identification of Organic Compounds, 5th ed.; Wiley: New York, 1965; p 149. (25) (a) Gionotti, C. In B I Z ;Dolphin, D., Ed.; Wiley: New York, 1982; Vol. I, pp 393-430. (b) Pratt, J. M. Inorganic Chemistry of Vitamin B I 2 ; Academic: New York, 1972.
2814
The Journal of Physical Chemistry, Vol. 92, No. 10, 1988
-r
TABLE 111 Parameters from Nonlinear Regression Analysis of Voltammograms of Vitamin BIzr[Co(II)/Co(I)y 106D, cm2 s-I 2.82 f 0.06 k o , cm s-I 0.016 f 0.001 E O ' , V versus S C E -0.7656 f 0.0005 av RSD,*% 0.34
06A
\
" E O ' from experiments at u d V s-'; ko from data at u < 10.24 V sd. Parameters given with average standard errors. bSD X loo%, three trials.
'--P\
I
1
200
Connors et al.
Simulation of Voltammetric Results. The model for digitally simulating cyclic voltammograms was based on Scheme I. The SCHEME I
1
600
400 X,nm
Co(I1)
Figure 6. UV-vis spectrum after addition of a 10-fold molar excess of deoxygenated 1,2-dibromobutane to the Co(1) solution of Figure 5c.
Co(1)
+ RBr,
k,
Co(1)
h nm
and to identify alkylcob(III)alamins.26b Electrolysis of vitamin Blzain the spectroelectrochemical cell at a potential past the Co(I1) reduction peak provided spectra of Co(III), Co(II), and Co(1) species in pH 2.3 aqueous acetonitrile buffer. Before electrolysis, the spectrum was that of base-on cob(II1)alamin. After 40 min of electrolysis, the spectrum was identicalZ5with that of base-off cob(I1)alamin (Figure 5). The spectrum of Co(I1) is observed midway through electrolysis because of the reaction of Co( I) formed at the electrode with Co(II1) in the bulk solution (eq l ) , 2Co(II)
-
Co(II1)
+ Co(1)
(1)
for which equilibrium lies far to the left.27 The spectrum of Co(1) appeared only after further electrolysis (Figure 5), when nearly all the Co(II1) in the cell had reacted. After 105 min of electrolysis, the solution contained mostly Co(1) with a little Co(I1) remaining, as evidenced by the shoulder at 473 nm (Figure 5c). Addition of a 10-fold molar excess of 1,2-dibromobutane (oxygen free) to the Co(1) solution caused immediate regeneration of the cob(I1)alamin spectrum (Figure 6 ) , with no evidence for an alkyl cobalt intermediate. This shows that the reaction between Co(1) and vicinal dibromide is an energetically downhill process. Electrolysis of a mixture of vitamin BIzaand a 10-fold excess of 1,2-dibromobutane gave the spectrum of cob(I1)alamin during electrolysis (Figure 7) until all the substrate was gone. Thereafter, the spectrum of cob(1)alamin began to appear. After 120 min of electrolysis, the spectrum indicated mainly Co(1) and a small amount of Co(I1). Spectral intensities after electrolysis indicated no decomposition of the catalyst. As on the carbon felt electrode, high background currents prohibited quantitative coulometry, but a steady-state current significantly larger than background was observed during the catalytic reduction of DDB. (26) Kenyhercz, T. M.; DeAngelis, T. P.; Norris, B. J.; Heineman, W. R.; Mark, H. B. J . Am. Chem. SOC.1976,98, 2469-2477. (b) Robinson, K. A,; Itabashi, E.; Mark, H. B. Inorg. Chem. 1982, 21, 3571-3573. (27) The equilibrium constant for this reaction3ais about between pH 4.7 and 7.8.
-
Co(1)
[BrCo(III)RBr]
+ RBr'
RBr-
Figure 7. UV-vis spectroelectrochemistryduring catalytic reduction of 1,2-dibromobutane (0.25 mM) with vitamin B12(0.025 mM). Conditions are as in Figure 5: (a) before electrolysis, (b) after 50 min of electrolysis, (c) after 65 min of electrolysis, and (d) after 120 min of electrolysis.
+e
-
-
Co(I1)
-
ko
(2)
Co(I1)
+ RBr' + Br(3)
+ RBr-
alkene + Br-
fast
fast
(4) (5)
simulation was a modification of a recently developed modelI7 for two-electron electrocatalytic reduction where the electron-transfer reactions2' in solution (eq 3 and 4) are outer sphere. Results should apply to inner-sphere situations where an intermediate complex between reduced catalyst and substrate (eq 3) decomposes more rapidly than it is formed. Scheme I explains the disappearance of the anodic wave of cob(1)alamin upon addition of substrate, since it is consumed in eq 3 and 4. The cathodic current is larger than that of catalyst alone because of regeneration of Co(I1) in the reaction layer. Electron transfer to the RBr' radical is considered to arise only from C O ( I ) . ~The ~ simulation includes the ability to vary k" (eq 2) and assumes semiinfinite linear diffusion to the planar electrode and unequal diffusion coefficients of substrate and catalyst. This model was coupled to a nonlinear regression program for analysis of individual i-E curves (see Experimental Section). So that values of ko and initial estimates of standard formal potential ( E O ' ) of Co(II)/Co(I) for fitting the catalytic i-E curves could be obtained, data for the Co(II)/Co(I) reduction alone were fit separately with nonlinear regression analysis onto a simulation model17afor heterogeneous electron transfer (eq 2). Data at low scan rates were analyzed by assuming reversible electron transfer to obtain an estimate of EO'; data at higher scan rates were used to obtain ko and D. Values of E O ' , D,and ko were in excellent agreement (Table 111) with those obtained by conventional analysis of CV and chronocoulometric data (Table 11). As shown previously (Table I), simulated voltammograms for the quasireversible (28) (a) Equation 3 reflects spontaneous reaction of cob(1)alamin with substrate. In the simulation, it is separated into eq 3a and 3b: Co(1)
-
+ RBr2 RBr,'.
Co(I1) + RBr2'-
RBr'
+ Br-
fast
k,
(3a) (3b)
This formalism is necessary because the simulation program cannot handle a step with more than two products. Equations 3a and 3b are not meant to imply existence of RBr,'-. By analogy with alkyl halides, this species should not exist in our systems. See the following: (b) Andrieux, C. P.; Merz, A,; Saveant, J. M. J . Am. Chem. SOC.1985, 107, 6097-6103 and references therein. (29) (a) This assumption is based on standard potentials (E'") of Co(II)/Co(I) [-0.77 V versus SCE] and RBr'/RBr-. The Eo' of the latter can be roughly estimated from that of n-butyl bromide [-0.99 by assuming about a 0.4 increase for addition of a second Br,20aplacing the E'" of DBB at about -0.6 V. Reduction of the vicinal dibromides by I- (ref 24a) suggests an even more positive E O ' . Since alkyl radicals are invariably much more easily reduced than their alkyl halide parents,29EEo' for RBr'/RBr' must be well positive of -0.6 V. This gives a negative standard free energy for reduction of RBr' by Co(I] (eq 4), implying that the reaction has a nearly diffusion-controlled rate. 9d Thus, when Co(1) is present, RBr' produced in the reaction layer will not have time to diffuse back to the electrode to be reduced. (b) Eberson, L. Acra Chem. S c a d . Secr B 1982, 36, 533-543. (c) Bard, A. J.; Merz, A. J . Am. Chem. SOC.1979, 101, 2959-2965. (d) Miller, J. R.; Calcaterra, L. T.; Closs, G.L. J . Am. Chem. SOC.1984, 105, 3047-3048.
The Journal of Physical Chemistry, Vol. 92, No. 10, 1988 2815
Reduction of Vicinal Dibromides by Vitamin B12 TABLE I V Estimates of Eo' and k l for DBB/Co(I) Reaction from Nonlinear Regression Analysis of Voltammetric Data" V, ko, -EO', 10dkl, V s-l cm s-l V vs S C E L mol-'s-I RSD, % x2 25.60 15.75 10.24 10.24b
0.016 0.016 0.016 0.016 mean f sd 25.60 0.05 0.05 15.75 0.05 10.24 0.05 10.24b mean f sd 25.60 0.10 0.10 15.75 0.10 10.24 0.10 10.24b mean f sd
0.6934 f 0.0005 5.53 f 0.18 0.6481 f 0.0004 10.44 0.03 0.6715 f 0.0006 5.62 f 0.27 0.6813 f 0.0004 8.74 f 0.37 7.6 f 2.4 0.673 f 0.019 0.7479 f 0.0003 4.42 f 0.10 0.7031 f 0.0003 8.96 f 0.17 0.7338 f 0.0002 6.84 f 0.17 4.37 f 0.14 0.7294 f 0.0004 6.1 f 2.2 0.728 0.020 2.84 0.16 0.7708 f 0.0007 0.7284 f 0.0002 6.63 f 0.10 0.7441 f 0.0002 2.86 f 0.04 0.7547 f 0.0004 4.32 f 0.17 4.1 f 1.8 0.750 f 0.017
*
*
0.54 0.56 0.88 0.65 0.66 0.36 0.41 0.37 0.55 0.42 0.79 0.41 0.25 0.55 0.50
12.6 20.5 29.5 26.3 22.2 8.3 11.0 7.4 23.3 12.5 20.8 11.0 3.3 23.3 14.6
" d = 10, y = 1; ko is fixed at the value shown; electrode area = 0.071 cm2 unless otherwise stated. E" and kl values are given with standard errors from regression program. Electrode area = 0.088 cm2.
one-electron reduction over the entire range of u showed good agreement with Co(1I) peak currents and potentials.21 Thus, the model, with the assumption that all species are dissolved in solution, was extended to the catalytic case with confidence. In preliminary regression analyses of electrocatalytic data, it was necessary to fix the ratio (d) of substrate to catalyst diffusion coefficients30to obtain a unique value of rate constant k,. Good quality simulation/regression analyses of the experimental i-E curves over a wide range of scan rates were obtained with d = 10. Values of d significantly smaller than 10 gave regression solutions with peak currents much smaller than experimental values, no matter how large the value of kl. This agrees with the finding that d has a major influence on relative catalytic currents at large kl/v for systems with fast heterogeneous electron transfer. 30b Thus, nonlinear regression analysis of i-E data required fixing d and k". The parameters determined in the regressions were EO', kl,and background slope and intercept." Voltammograms at u > 10 V s-, were analyzed since these should contain the most information about k, when it is large. Initial guesses of k, on the order of 106-107 L mol-' s-l were obtained by comparing catalytic efficiencies ( & / ( 2 ~ i ~ ) with ) ~ &working curves for two-electron catalysis with fast heterogeneous electron transfer.30b Nonlinear regressions with k" = 0.016 cm s-l gave good fits to the experimental data with average standard deviation of regression (SD) of 0.66% of the peak current (Table IV). However, tests of several different values of ko showed that k" = 0.05 cm s-l gave better fits in terms of average S D and xz. S D was
