J. Phys. Chem. B 2000, 104, 1025-1032
1025
Electron Transfer Reactions of Hydrophobic Metallocenes with Aqueous Redox Couples at Liquid-Liquid Interfaces. 1. Solvent, Electrolyte, Partitioning, and Thermodynamic Issues Heather O. Shafer, Torri L. Derback, and Carl A. Koval* Department of Chemistry and Biochemistry, UniVersity of Colorado at Boulder, Boulder, Colorado 80309-0215 ReceiVed: August 9, 1999; In Final Form: NoVember 21, 1999
Electron transfer reactions of the form, ML63+(w) + Rn-FER(o) f ML62+(w) + RnFER+(o), where ML63+,2+(w) are hydrophilic aquo or amine transition metal based redox couples and RnFER+,0(o) are hydrophobic alkylferrocence redox couples, can occur at the interface between aqueous (w) and immiscible organic (o) phases. The hydroxymethylferrocene+,0 couple was used as an internal standard in order to compare the formal reduction potentials for a variety of aqueous couples with reduction potentials for the ferrocene couples in organic solvents with dielectric constants ranging from 9 to 35. The ability of various electrolytes to provide adequate conductivity in aqueous/organic phases, without inducing partitioning of alkylferricenium cations into the aqueous phase, was examined. The single electrolyte tetraethylammonium tetrafluoroborate and the electrolyte consisting of potential determining ions tetrapropylammonium bromide/tetrapropylammonium tetraphenylborate were found to be generally suitable in these respects. Decamethylferrocenium ion was found to partition from organic to aqueous phases under certain conditions. In contrast, 1,1′,3,3′-tetrakis(2-methyl-2-hexyl)ferrocenium remained in the organic phases. The rate of partitioning of alkylferricenium ions from benzyl cyanide thin films immobilized on carbon electrode surfaces was found to depend on the hydrophobicity of the cation and on the electrolyte ions. Estimates of the interfacial potential difference, ∆wo φ, induced by the two favored electrolytes for several aqueous/organic solvent interfaces were determined. These values of ∆wo φ combined with relative values of formal reduction potentials for aqueous and organic soluble redox couples can be used to estimate the driving force for a wide variety of electron transfer reactions at liquid/ liquid interfaces. When the tetrapropylammonium ion, a potential determining ion, was used as the electrolyte, the value of ∆wo φ that was established conformed to the Nernst equation. The single electrolyte tetraethylammonium tetrafluoroborate established a value of ∆wo φ that was independent of salt concentration below 0.1 M.
Introduction There is emerging interest in electron transfer reactions (ETRs) that take place across the interfaces between immiscible liquids. Understanding this type of redox process has obvious relevance to biological systems where electron transfer takes place between centers located in media of different polarity. ETRs at liquid-liquid interfaces (LLIs) are also important in microheterogeneous photoconversion schemes that employ redox couples embedded in polymers or miscelles.1,2 There is also practical interest in separation processes that utilize liquid redox extraction.3,4 In 1988, Schiffrin and Girault presented a theoretical framework for ETRs at LLIs.5 In the following years, these researchers reported carefully designed electrochemical cells involving static liquid-liquid junctions to measure interfacial electron transfer rate constants between discrete molecular species.6-10 Shortly thereafter, Marcus adapted his homogeneous and electrochemical electron transfer theories to include liquid-liquid reactions.11-13 Although Schiffrin and co-workers have published a number of papers aimed at examining several redox systems and at improving the interpretation of the experiments, their approach has some severe limitations. Specifically, their approach is limited to ETR’s at LLI’s that are polarizable, i.e., systems in which significant potential drops can be imposed across the interface without inducing ion transfer reactions (ITRs). In practice, these polarization windows are 300 mV) and induce undesirable partitioning of metallocenium cations into the aqueous phase (see below). 5. Tetraphenylborate salts are soluble in many organic solvents but not in water. They are useful in potential determining ion systems as discussed below. In previous studies of ETRs at LLIs, electrolytes containing PDIs involving perchlorate ions have been employed.16,22,23 Since precipitation of multiply-charged, cationic aquo and amine complexes like those in reaction 1 is likely to be problematic with perchlorate salts, attempts were made to identify a suitable electrolyte system composed of a PDI in which the partitioning
1028 J. Phys. Chem. B, Vol. 104, No. 5, 2000 ion was a monocation. The counterions associated with the PDI in the organic and aqueous phases must be very hydrophobic or hydrophilic, respectively, so their ability to partition into the opposing phase can be neglected. For this reason, tetraalkylammonium salts containing the hydrophobic anion tetraphenylborate (TPB) and hydrophilic halide anions were investigated. Attempts to identify a single electrolyte system containing a PDI that was compatible with a variety of aqueous/organic solvent pairs met with a number of difficulties. Salts containing the tetraphenylborate were not soluble in octanol, precluding octanol’s further use for LLIs involving this type of electrolyte. Tetramethyl- and tetraethylammonium TPB are not sufficiently soluble in most of the organic solvents tested. Salts derived from tetrabutyl- or tetrapentylammonium ions caused undesirable partitioning of decamethylferricenium ions into water, as discussed below. Tetrapropylammonium tetraphenylborate, TPrATPB, together with TPrA+halide salts proved to be an acceptable electrolyte system composed of a PDI for all of the organic solvents examined except octanol. TPrABr is quite soluble in water and TPrATPB is soluble in 2-nonanone to 0.02 M and is at least 0.05 M soluble in nitrobenzene, butyronitrile, benzonitrile, and benzyl cyanide. Electrolyte-Induced Partitioning of DcMFER+ and HEP+. In order for ETRs like those in reaction 1 to occur at a LLI rather than homogeneously, it is advantageous to minimize the ability of the various redox species to partition into the opposing phases. Problems associated with undesired partitioning and precipitation of redox couples have recently been discussed by Quinn et al.36 For the organic solvents and electrolytes used in this study, partitioning of M(III,II) aquo and amine couples was never observed to occur to a large extent. This lack of partitioning was expected since partitioning of multiply-charged metal cations usually requires the presence of hydrophobic complexing or ion-pairing reagents.42 Partitioning of the highlycolored Ru(NH3)5(pyridine)2+ ion, which is the most hydrophobic of the redox couples proposed in Figure 1, from aqueous solutions into the organic phases used in this study could not be detected spectrophotometrically. Appreciable partitioning of metallicenium ions (and even neutral metallocenes in some cases) is more problematic. In principle, partitioning of charged redox species can be predicted from experimentally determined standard transfer potentials compared with values of ∆wo φ, the interfacial potential difference produced by a given electrolyte/solvent combination.9 The application of this procedure, as well as some of the difficulties and limitations, has been described for simple ferrocene derivatives (including DcMFER+) for the aqueous/1,2-dichloroethane interface.8 Partitioning of DcMFER+ into water from nitrobenzene, benzonitrile, and butyronitrile solutions was not detected spectrophotometrically. However, DcMFER+ readily partitioned into water from organic solvents with smaller dielectric constants, e.g., octanol and 2-nonanone. The cationic forms of less hydrophobic metallocenes (e.g., ferrocene and dimethylferrocene) partition extensively into water from all organic solvents. The presence of electrolytes had a significant effect on partitioning of DcMFER+ into aqueous solutions. For example, the use of TPeACl as a single electrolyte system causes DcMFER+ to partition into the aqueous phase from all of the organic solvents investigated. Single electrolytes containing less hydrophobic cations, such as TEtATFB, did not induce extensive partitioning of DcMFER+, provided a stoichiometric quantity of TPB- anions was present in the organic phase. For electrolyte systems composed of tetraalkylammonium PDIs, salts containing
Shafer et al. TMeA+ and TEtA+ ions did not induce DcMFER+ partitioning, but these systems were not as widely applicable due to limited solubility of TMeATPB and TEtATPB in low dielectric organic solvents. Electrolytes containing the PDI TBuA+, as well as TBuATFB, were also frequently observed to induce extensive partitioning of DcMFER+ into the aqueous phase. Electrolytes containing the PDI TPrA+ did not induce detectable partitioning of DcMFER+ ions. This result coupled with the solubility data above suggests that electrolyte systems containing the PDI TPrA+ have the widest applicability for studying ETRs like reaction 1. Partitioning of organic phase redox couples into the aqueous phase can also be minimized by addition of hydrophobic substituents. In the case of metallocenes, the derivatives reported by the Strauss group are excellent in this regard. The presence of four 2-methyl-2-hexyl groups in HEP+ greatly limits its ability to partition into aqueous solutions. Partitioning of either HEP or HEP+ into water was never observed. Recently, Anson and Shi have proposed the use of thin organic films that spontaneously coat the surfaces of carbon electrodes for use in studying ETRs at LLIs.21,22 Since the volume ratio of organic phase to aqueous phase in these experiments is often less than 1:104, partitioning of the organic soluble redox couple into the aqueous phase can be especially problematic. Fortunately, the thin film technique provides a convenient method for monitoring the loss of redox couple from an organic phase. As described by Anson and Shi, examining the cyclic voltammetry (CV) behavior of redox couples contained in the thin films allows the determination of the concentration of the couple by utilizing relatively fast scan rates (e.g., 400 mV/s) and the determination of the total amount of redox species by utilizing relatively slow scan rates (e.g., 5 mV/ s).21 As indicated above, the ability to retain the ionic forms of hydrophobic redox couples in organic phases is affected by the structure of the redox species, the polarity of the organic solvent, and the electrolytes that are present. Thin film CV experiments were used to examine how these factors affected the extent of partitioning of DiMFER+, DcMFER+, and HEP+ ions from benzyl cyanide thin films into aqueous solutions. The choice of this solvent for these experiments was dictated by several factors. As discussed by Anson and Shi,21 the formation of stable and uniform thin films on the roughened carbon electrode surfaces is most readily accomplished using solvents with relatively low volatility. Of the three nitrile solvents investigated, BzCN and BzN readily formed stable organic thin films, but loss of BuCN due to evaporation made its use more problematic. BzCN and BzN both have densities that are similar to water; however, it is much easier to induce phase separation for BzCN/ water mixtures than for BzN/water mixtures. Therefore, it is easier to presaturate water with BzCN than with BzN. Presaturation of the aqueous phase with the appropriate organic solvent leads to much greater stability for thin film experiments, especially if the aqueous solution is convected (e.g., by deaeration). The organic thin films were prepared by placing BzCN solutions containing the neutral forms of the complexes (1 mM) onto a carbon electrode surface. An initial potential sufficiently positive to oxidize the metallocenes to their cations was applied and repeated CVs were performed at 400 mV/s. The observed decrease in the peak currents over time for the waves associated with the +,0 couples was due to partitioning of the cations into the aqueous phase. In these experiments, the aqueous phase was agitated mildly in order to distribute the material partitioning
ETR of Hydrophobic Metallocenes with Redox Couples
J. Phys. Chem. B, Vol. 104, No. 5, 2000 1029
Figure 3. Cyclic voltammetry (400 mV/s) of DcMFER (1 mM) in BzCN a thin film immersed in aqueous solutions (0.05 M TPeACl). Aqueous phase presaturated with BzCN and deaerated with mild Ar bubbling between scans. Times indicate when various scans were initiated after the first scan (time ) 0).
Figure 2. Cyclic voltammetry (400 mV/s) of DiMFER (curve A), DcMFER (curve B), and HEP (curve C) in BzCN thin films immersed in aqueous solutions (0.05 M TEtATFB). Aqueous phase presaturated with BzCN; the initial concentration for each metallocene was approximately 1 mM. Aqueous phase was deaerated with mild Ar bubbling between scans. Times indicate when various scans were initiated after the first scan (time ) 0).
out of the organic layer throughout the aqueous phase. These experiments were performed for two electrolyte systems: the single electrolyte TEtATFB(0.05 M, aq) and the electrolyte containing a PDI TPrATPB(0.05 M, org)/TPrABr(0.05 M, aq). Cyclic voltammograms vs time after immersion for DiMFER, DcMFER, and HEP using TEtATFB as the electrolyte are contained in Figure 2. Since DiMFER+ is readily soluble in water, the wave associated with DiMFER+,0 (curve A) disappears in less than 2 min. In contrast, the waves associated with DcMFER+,0 and HEP+,0 (curves B and C) remain virtually unchanged for times exceeding 20 min. The result for DcMFER+,0 is consistent with the bulk partitioning experiments, i.e., electrolytes containing less hydrophobic cations such as TMeA+ and TEtA+ induced less partitioning of DcMFER+ cations into water than did the more hydrophobic cations TBuA+ and TPeA+. The effect of TPeA+ ion on the partitioning of DcMFER+ from a BzCN thin film is depicted in Figure 3. The hydrophobic TPeA+ electrolyte cation in the aqueous phase rapidly displaces DcMFER+ ion from the organic thin film as shown by the fact that the wave associated with the DcMFER+,0 couple disappears within 45 s. The results for the loss of metallocenes from BzCN thin films where the electrolyte system composed of PDIs was employed are contained in Figure 4. Reduction waves for DiMFER+ disappear from the thin film is less than 2 min. Although the waves for DcMFER+,0 are more persistent, the concentration of DcMFER+ cations decreases by a factor of 2 in about 15 min. As stated earlier, electrolytes derived from the PDI TPrA+ did not induce extensive partitioning of DcMFER+ ions into water. However, due to the small organic:water volume ratio associated with thin film experiments, DcMFER+ is slowly lost from the BzCN thin films. The HEP+ cation is the most hydrophobic couple of the three metallocenes; the HEP+ concentration remains within 95% of its initial value for more than 25 min.
Figure 4. Plots of the normalized cathodic peak currents (scan rate ) 400 mV/s) vs time after immersion for DiMFER (curve A), DcMFER (curve B) and HEP (curve C) in BzCN thin films. Peak currents were normalized via division by the peak current for the first scan. Aqueous phase presaturated with BzCN and deaerated with mild Ar bubbling between scans. The initial concentration for each metallocene was approximately 1 mM. The electrolyte was TPrATPB (0.05 M, org)/ TPrABr (0.05 M, aq).
These experiments indicate that by using either extremely hydrophobic redox couples such as HEP+,0 or less hydrophobic cations such as TEtA+ as the electrolyte, it is possible to obtain voltammetry in the organic thin films that is stable for times longer that 20 min. However, if the solution is convected, this stability requires that the aqueous solution is thoroughly presaturated with the organic solvent used to create the thin layer. The effect of not presaturating the aqueous phase is demonstrated in Figure 5. When the wave for HEP+,0 is examined in a thin film placed in an aqueous phase that has not been presaturated with BzCN, the waves increase in size and eventually become highly distorted. Presumably, this is caused by loss of BzCN from the thin film which caused the concentration of HEP to increase. At sufficiently long times, much of the BzCN in the thin film is lost and the voltammetry becomes highly distorted. Estimation and Variation of the Interfacial Potential Difference. For ETRs at LLIs like those in reaction 1, the thermodynamic position of equilibrium is described by the equation
∆wo φ + E°′rxn )
RT ln K nF
(2)
where ∆wo φ is the interfacial potential difference, E°′rxn is the difference in reduction potentials for the redox couples in the two phases and K is the equilibrium constant for the biphasic reaction. Detailed discussions describing how the individual terms in eq 2 can be determined have appeared in the literature.9
1030 J. Phys. Chem. B, Vol. 104, No. 5, 2000
Shafer et al.
TABLE 2: Estimated Interfacial Potentials (∆owO) organic solvent
E°′rxn (V)a (Table 1)
nitrobenzene benzonitrile benzyl cyanide 2-nonanone
0.160 0.178 0.252 0.075
TEtATFB
b,c
E°′rxn + ∆wo φ TPrA+, PDIb,c
-0.044 -0.046 -0.025 -0.205
∆wo φ(est) TEtATFB
-0.109 -0.104 -0.096 0.023
c
-0.204 -0.224 -0.277 -0.280
TPrA+, PDIc -0.269 -0.282 -0.248 -0.052
Calculated for the reaction Ru(NH3)63+(w) + DcMFER0(o) ) Ru(NH3)62+(w) + DcMFER+(o) using the data in Table 1. b Values of E°′rxn + measured using the organic thin film technique (see text). c 0.1 M TEtATFB in the aqueous phase 0.05 M TPrABr in the aqueous phase; 0.05 M TPrATPB in the organic phase (except for 2-nonanone where TPrATPB was 0.02 M). a
∆wo φ
Figure 5. Cyclic voltammetry (400 mV/s) of HEP (1 mM) in BzCN thin film immersed in aqueous solution (0.05 M TEtATFB): (top) aqueous phase presaturated with BzCN; (bottom) aqueous phase not presaturated with BzCN. The aqueous phase was deaerated with mild Ar bubbling between scans. Times indicate when various scans were initiated after the first scan (time ) 0).
From a practical point of view, investigations of reactions like those in reaction 1 are aided by the ability to estimate ∆wo φ and E°′rxn for various redox couples and electrolyte systems. An ability to independently vary ∆wo φ and E°′rxn is useful for testing theoretical predictions of how the two contributions to driving force affect the rate constants for these reactions. One relatively simple method for estimating E°′rxn involves assuming that E°′ for a particular redox couple is solvent independent (the ferrocene assumption9,38). Rather than using the ferrocene couple itself, HMFER+,0 was chosen because it is readily soluble in water as well as most organic solvents. Table 1 contains formal reduction potentials for one aqueous and two organic soluble redox couples vs E°′ for HMFER+,0. By making the assumption that E°′ vs HMFER+,0 is solvent independent and by comparing E°′ (V) for the couples in Table 1 to values for other couples in the same solvent, values of E°′rxn ()E°′(w) - E°′(o)) for ETRs at LLIs can be readily estimated. Table 2 contains estimated values of E°′rxn for the reaction
Ru(NH3)63+(w) + DcMFER0(o) ) Ru(NH3)62+(w) + DcMFER+(o) at a number of aqueous/organic solvent interfaces for two electrolyte systems. Values for the sum of ∆wo φ and E°′rxn are readily determined using electrochemical measurements by determining the position of waves for the two couples vs a reference electrode that is placed in the aqueous phase. The working electrode for these
experiments is placed alternately in the two electrolyte phases contained in the cell. This can be accomplished using either a conventional cell containing macroscopic amounts of each phase or the organic thin film technique. In the case where the concentrations of potential determining electrolyte ions far exceed the concentrations of the redox active ions, ∆wo φ is presumed to be controlled by the electrolyte ions. Values of ∆wo φ + E°′rxn were determined using the thin film technique for the two electrolytes that we have found to be widely useful: the single electrolyte TEtATFB (0.1 M aq) and the electrolyte composed of potential determining ions TPrABr (0.2 M aq)/TPrATPB (0.05 M org). Using CV, values of ∆wo φ + E°′rxn for the reaction Ru(NH3)63+(w) + DcMFER0(o) were determined for the organic solvents and electrolytes in Table 2. Using the values of E°′rxn estimated above, values of ∆wo φ can be estimated. These values of ∆wo φ in Table 2 can be utilized to estimate the position of equilibrium for any ETR at these water/organic solvent interfaces where E°′ (V) vs Ru(NH3)63+,2+ for the aqueous couple and E°′ (V) vs DcMFER+, for the organic couple are known. The concept of choosing aqueous and organic soluble redox couples with interesting relative formal reduction potentials is illustrated in Figure 1. For a given organic solvent and electrolyte system, ∆wo φ serves to shift the potential scales causing the interfacial ETR to have a greater or smaller driving force. One advantage of electrolyte systems composed of PDIs is the ability to vary the portion of the overall driving force attributable to ∆wo φ in a predictable way. The ability to achieve this goal using the TPrABr/TPrATPB electrolyte system at water/BzCN interfaces was investigated. Also, due to an interest in using aquo couples in the aqueous phase, it was also necessary to examine whether the values of ∆wo φ established by this electrolyte were sensitive to the pH of the aqueous solution. The interfacial potential difference, ∆wo φ, for this electrolyte system composed of PDIs is governed by the equation at 298 K:41
∆owφ
)
∆owφ°TPrA+
-0.0592 V log
[TPrA+]o [TPrA+]w
(3)
where ∆wo φ° is the standard interfacial potential difference and assuming that concentrations can be substituted for activities. According to eq 3, ∆wo φ varies by 0.0592 V for every decade change in the ratio of TPrA+ between the two phases. This prediction can readily be tested, either for conventional cells or for experiments involving organic thin films, by measuring the sum of ∆wo φ and E°′rxn as described above for a series of experiments in which the electrolyte concentrations in one or both of the phases are varied (and assuming the value of E°′rxn is independent of electrolyte concentration). For NB organic thin films and an electrolyte consisting of tetrahexylammonium perchlorate (o) and sodium perchlorate
ETR of Hydrophobic Metallocenes with Redox Couples
J. Phys. Chem. B, Vol. 104, No. 5, 2000 1031 Acknowledgment. This work was supported by the CU Boulder Deans Small Grants and by the Department of Chemistry and Biochemistry. The authors gratefully acknowledge many useful discussions with F. C. Anson and R. D. Noble. We also thank F. C. Anson and Chunnian Shi for preprints of their manuscripts. List of Abbreviations
Figure 6. Plot of E°′rxn ()E°′(w) - E°′(o)) for the Ru(NH3)63+,2+(w) and HEP+,0(o) couples vs the Nernst ratio of TPrA+ ions for BzCN/ water interfaces. The concentration of TPrATPB in the BzCN phase was 0.02 M and the concentration of TPrABr in the aqueous phase was varied between 0.02 and 0.5 M. All potentials were measured vs a Ag/AgCl reference electrode placed in the aqueous phase.
∆wo φ
(w), Shi and Anson have observed changes in for this type of experiment with slopes that are close to 0.030 V rather than the 0.059 V predicted by eq 3.23 They have interpreted these lower slopes in terms of a theoretical treatment by Kakiuchi that specifically relates to situations where the volume ratio of the two phases is far from unity.43,44 Typical results for BzCN thin films utilizing the TPrABr/TPrATPB electrolyte are presented in Figure 6. The slope of the plot is very close to the value of 0.0592 V predicted by eq 3. The reason for the different slopes obtained by Shi and Anson and in this study is currently unknown. One possible explanation is that use of eq 3 assumes that the concentration of the PDI (TPrA+) is controlled by the addition of salts to the organic and water phases. If one (or both) of the salts partition extensively into the opposing phase, the prediction of eq 3 will not be observed either due to changes in the anticipated concentration of the PDI or to the fact that more than one ion is influencing the value of ∆wo φ. As shown by Kakiuchi, partitioning of salts has especially large effects when the ratio of the volumes of the organic and aqueous phases becomes extremely large or small. In order to determine whether the interfacial potential was sensitive to the acidity of the aqueous phase, additional measurements were made at a constant ratio of [TPrA+]o/ [TPrA+]w as the pH of the aqueous phase was varied between 1 and 7. The measured value of ∆wo φ remained constant to within a few millivolts. In contrast to electrolytes composed of PDIs, the magnitude of ∆wo φ for single electrolytes should be independent of the concentration of the salt. For TEtATFB, measured values of ∆wo φ using BzCN thin films were independent of TEtATFB concentration in the aqueous phase between 0.02 and 0.10 M. At higher concentrations, significant changes in ∆wo φ were observed, presumably due to activity effects. Conclusions The experiments described above indicate that it should be possible to examine the kinetics for a wide variety of ETRs involving aquo and amine complexes with metallocenes at LLIs involving solvents of different polarities. Two types of electrolytes, a single electrolyte and an electrolyte composed of potential determining ions that are compatible with many organic solvents, have been identified and interfacial potential differences for several LLIs have been estimated. The partitioning of metallocenium cations into opposing aqueous phases can be controlled to some extent through the judicious choice of alkylated derivatives and electrolyte systems.
SolVents nitrobenzene benzonitrile butyronitrile benzyl cyanide octanol 2-nonanone
NB BzN BuN BzCN OcOH NoN
Electrolytes tetramethylammonium tetraethylammonium tetrapropylammonium tetrabutylammonium tetrapentylammonium tetraphenylborate tetrafluorborate
TMeA TEtA TPrA TBuA TPeA TPB TFB
Redox Couples ferrocene hydroxymethylferrocene dimethylferrocene decamethylferrocene 1,1′,3,3′-tetrakis(2-methyl-2-hexyl)ferrocene
FER HMFER DiMFER DcMFER HEP
Miscellaneous electron transfer reaction liquid-liquid interface potential determining ion
ETR LLI PDI
References and Notes (1) Lewis, N. S.; Kamat, P.; Spitler, M. In Research Opportunities in Photochemical Sciences; Nozik, A. J., Ed.; U.S. DOE: Estes Park, CO, 1996; p 142. (2) Koval, C.; Sutin, N.; Turner, J. In Research Opportunities in Photochemical Sciences; Nozik, A. J., Ed.; U.S. DOE: Estes Park, CO, 1996; p 219. (3) Clark, J. F., et al. EnViron. Sci. Technol. 1996, 30, 3124. (4) Matsuno, S., et al. Chem. Lett 1981, 1543. (5) Girault, H. H. J.; Shiffrin, D. J. J. Electroanal. Chem. 1988, 244, 15. (6) Cheng, Y.; Shriffrin, D. J. J. Chem. Soc., Faraday Trans. 1993, 89, 199. (7) Cheng, Y.; Shiffrin, D. J. J. Chem. Soc., Faraday Trans. 1994, 90, 2517. (8) Cunnane, V. J.; Geblewicz, G.; Shiffrin, D. J. Electrochim. Acta 1995, 40, 3005. (9) Girault, H. H. J.; Schiffrin, D. J. Electroanal. Chem. 1989, 15, 1. (10) Geblewicz, G.; Shiffrin, D. J. J. Electroanal. Chem. 1988, 244, 27. (11) Marcus, R. A. J. Phys. Chem. 1990, 94, 4152. (12) Marcus, R. A. J. Phys. Chem. 1990, 94, 1050. (13) Marcus, R. A. J. Phys. Chem. 1991, 95, 2010. (14) Bard, A. J.; Fan, F.-R. F.; Mirkin, M. V. Electroaanl. Chem. 1993, 18, 243. (15) Solomon, T.; Bard, A. J. J. Phys. Chem. 1995, 99, 17487. (16) Tsionsky, M.; Bard, A. J.; Mirkin, M. J. Phys. Chem. 1996, 100, 17881. (17) Wei, C.; Bard, A. J.; Mirkin, M. V. J. Phys. Chem. 1995, 99, 16033. (18) Shao, Y.; Mirkin, M. V.; Rusling, J. F. J. Phys. Chem. 1997, 101, 3202. (19) Mirkin, M. V. Anal. Chem. 1996, 68, 177.
1032 J. Phys. Chem. B, Vol. 104, No. 5, 2000 (20) Liu, B.; Mirkin, M. V. J. Am. Chem. Soc. 1999, 121, 8352. (21) Shi, C.; Anson, F. C. Anal. Chem. 1998, 70, 3114. (22) Shi, C.; Anson, F. C. J. Phys. Chem. B 1998, 102, 9850. (23) Shi, C.; Anson, F. C. J. Phys. Chem. B 1999, 103, 6283. (24) Schmickler, W. Interfacial Electrochemistry; Oxford University Press: New York, 1996. (25) Schmickler, W. J. Electroanal. Chem. 1997, 428, 123. (26) Girault, H. H. Mod. Aspects Electrochem. 1993, 25, 1. (27) Bottcher, W.; Brown, G. M.; Sutin, N. Inorg. Chem. 1979, 18, 1447. (28) Brown, G. M.; Sutin, S. J. Am. Chem. Soc. 1979, 101, 883. (29) Chou, M.; Creutz, C.; Sutin, N. J. Am. Chem. Soc. 1977, 99, 5615. (30) Hupp, J. T., et al. J. Electroanal. Chem. 1984, 168, 313. (31) Gennett, T.; Weaver, M. J. Anal. Chem. 1984, 56, 1444. (32) Pladziewicz, J. R.; Espenson, J. H. J. Am. Chem. Soc. 1973, 95, 56. (33) Bond, A. M., et al. Anal. Chem. 1988, 60, 1878. (34) Penner, R. M., et al. Science 1990, 250, 1118.
Shafer et al. (35) Yang, S. H.; Chan, M. S.; Wahl, A. C. J. Phys. Chem. 1980, 84, 3094. (36) Quinn, B., et al. Electrochim. Acta 1998, 44, 47. (37) Mann, C. K. In Electroanalytical Chemistry: A Series of AdVances; Bard, A. J., Ed.; Marcel Dekker: New York, 1969. (38) Gagne´, R. R.; Koval, C. A.; Lisensky, G. C. Inorg. Chem. 1980, 19, 2854. (39) Lim, H. S.; Barclay, D. J.; Anson, F. C. Inorg. Chem. 1972, 11, 1460. (40) Koepp, H. M.; Wendt, H.; Strehlow, H. Z. Electrochem. 1960, 64, 483. (41) Melroy, O. R.; Buck, R. P. J. Electroanal. Chem. 1983, 143, 23. (42) Alegret, S. DeVelopments in SolVent Extraction; Algret/Ellis Horwood Ltd.: Chichester, UK, 1988. (43) Kakiuchi, T. Anal. Chem. 1996, 3658. (44) Kakiuchi, T. In Liquid-Liquid Interfaces; Volkov, A. G., Deamer, D. W., Eds.; CRC Press Inc.: New York, 1996; p 1.
