Mechanistic study of the oxidation of l -ascorbic acid by chloranil at the nitrobenzene ∣ water interface

www.elsevier.nl/locate/jelechem Journal of Electroanalytical Chemistry 490 (2000) 85 – 92

Mechanistic study of the oxidation of L-ascorbic acid by chloranil at the nitrobenzene water interface Toshiyuki Osakai a,*, Naoki Akagi a, Hiroki Hotta a, Jie Ding a, Shigeo Sawada b b

a Department of Chemistry, Faculty of Science, Kobe Uni6ersity, Nada-ku Kobe 657 -8501, Japan Department of Chemistry, Faculty of Science and Technology, Kinki Uni6ersity, Higashiosaka, Osaka 577 -8502, Japan

Received 29 March 2000; received in revised form 9 June 2000; accepted 14 July 2000

Abstract The redox reaction between L-ascorbic acid in water and chloranil in nitrobenzene has been studied by means of polarography with an ascending water electrode as well as cyclic voltammetry with a stationary interface. Through accurate measurement of the limiting currents, it has been suggested that the redox reaction should be a two-electron reaction rather than a one-electron reaction described previously. A spectrophotometric technique has also been used to observe that the redox reaction proceeds spontaneously under certain conditions even without electrochemical control. Based on these findings, it has been concluded that the present heterogeneous charge transfer reaction is the ion transfer of chloranil semiquinone radical, which is driven by the homogeneous electron transfer between ascorbic acid and chloranil in the aqueous phase. © 2000 Elsevier Science S.A. All rights reserved. Keywords: Ascorbic acid; Chloranil; Nitrobenzene water interface; Electron transfer; Ascending water electrode

1. Introduction Electron transfer at the liquid liquid interface (or the oil water interface) is fundamentally important for understanding the energy conversion in biomembranes [1]. Samec et al. [2] first employed cyclic voltammetry to observe an electron transfer between hexacyanoferrate redox couple in water and ferrocene in nitrobenzene. However, a possible interfacial transfer of the oxidation product of ferrocene (i.e. ferricenium ion) and its coupling with the electron transfer were discussed [3]. In order to study the kinetics of true heterogeneous electron transfers, some highly hydrophobic complexes including lutetium(III) biphthalocyanin [4], tin(IV) diphthalocyanine [5], and bis(pyridine)meso-tetraphenylporphyrinato iron(II) and ruthenium(III) [6] were chosen as redox couples in the oil phase. These fundamental studies were made with a view to applying the

* Corresponding author. Tel.: +81-78-8035682; fax: + 81-788035682. E-mail address: [email protected] (T. Osakai).

experimental results to Marcus theory [7–10]. On the other hand, Kihara and co-workers have identified various heterogeneous redox reactions for organic compounds of biological significance which include L-ascorbic acid [11], flavin mononucleotide (FMN) [12], and b-nicotinamide adenine dinucleotide (NADH) [13]. In this study we have re-examined the redox reaction between L-ascorbic acid (AH2) in water (W) and chloranil (2,3,5,6-tetrachloro-1,4-benzoquinone; denoted by Q) in nitrobenzene (NB) [11]. The mechanism of the heterogeneous redox reaction has been closely studied by using cyclic voltammetry with a stationary NB W interface and polarography with an ascending water electrode. It has also been observed by a spectrophotometric technique that a redox reaction between AH2 and Q can occur spontaneously under no electrochemical control. Although the previous paper [11] claimed that a one-electron transfer occurred between AH2 (W) and Q (NB), the present study claims that two electrons per AH2 molecule participate in the heterogeneous redox reaction that includes an ion-transfer process due to a reduction product of Q (i.e. the semiquinone radical ion).

0022-0728/00/$ - see front matter © 2000 Elsevier Science S.A. All rights reserved. PII: S 0 0 2 2 - 0 7 2 8 ( 0 0 ) 0 0 2 4 7 - 3

T. Osakai et al. / Journal of Electroanalytical Chemistry 490 (2000) 85–92

86

2. Experimental

2.1. Reagents L-Ascorbic acid (Wako Pure Chemical Industries, Ltd.; analytical grade) and chloranil (Tokyo Kasei Industry Co., Ltd.) were used as received. The tetra-n-pentylammonium salt of 12-tungstophosphate, (TPnA)3[PW12O40], which can be used as the supporting electrolyte in NB, was prepared by the equimolar addition of an aqueous solution of 12-tungstophosphoric acid n-hydrate (Wako; analytical grade) and an aqueous solution of n-tetrapentylammonium chloride (TPnACl; Wako practical grade); the precipitate was washed five times with distilled water and recrystallized from acetone. An aqueous solution of TPnACl, which was used for phase I in Cell (A), was treated with silver chloride to remove trace amounts of iodide ion; the concentration of TPnACl was determined by potentiometric titration with a standard silver nitrate solution. Analytical grade nitrobenzene (Wako) was treated before use with activated alumina for column chromatography (Wako; 200 mesh). All other reagents were of the highest grades available and were used as received.

Fig. 1. Three-electrode type electrolytic cell with an ascending water electrode: (a) NB solution for phase II in cell (A); (b) aqueous solution for phase III; (c) Ag AgCl rod electrode for potential control; (d) Ag AgCl spiral electrode for potential control and current detection; (e) spiral platinum electrode for current detection; (f) Teflon tip; (g) Teflon heat-shrink tubing; (h) glass tube; (i) glass tube with a glass sinter; (j) Luggin capillary; (k) glass capillary connected to an aspirator to suck off excess solution; (l) glassware with a water jacket; (m) elastic silicone + Teflon rubber distributed by Nippon Rikagaku Kikai Co., Ltd.; (n, o, p) silicone rubber; (q) infusion pump; (r) 10 ml syringe; (s) Teflon tubing (0.5 mm i.d.).

In cyclic voltammetry and polarography electrochemical Cell (A) was used: I Ag AgCl

II

0.02 M TPnACl 0.1 M MgSO4 (W)

III

20 mM Q 0.035 M (TPnA)3[PW12O40]

0.1–0.5 mM AH2 0.1 M NaCl 0.5 M Na2SO4 (W)

(NB)

–––––––––

with a 10-ml syringe (Hamilton 1010 model) was used to send an aqueous solution to the cell at a constant rate. The aqueous solution in the syringe, as well as the NB solution in the cell, was deaerated in advance.

2.2. Electrochemical measurements

IV

0.1 M NaCl 0.5 M Na2SO4 (W)

AgCl Ag

Cell A The NB W interface denoted by is to be tested. The pH (5.7–7.9) of the aqueous phase (III) was adjusted with 0.05 M phosphate buffer. Phases III and IV were separated by means of a glass sinter. For cyclic voltammetry with a stationary NB W interface, the previous three-electrode electrolytic cell [14] was used, but partly modified so as to deaerate the NB and W phases with N2 gas. The surface area of the NB W interface to be tested was 0.049 cm2. For polarography with an ascending water electrode, a modified three-electrode cell [15] was used. As shown in Fig. 1, a Teflon tip with a small orifice of 0.3 mm-f, which was supported on a glass tube with Teflon heatshrink tubing, was incorporated into glassware similar to a Liebig condenser (85 mm length, 45 mm o.d., 20 mm (upper) and 25 mm (lower) i.d.). In place of the solution reservoir with a glass capillary [15], an infusion pump (World Precision Instruments, SP100i) equipped

All electrochemical measurements were performed using a laboratory-constructed computer assisted system. The temperature of the electrolytic cells was always set at 2590.1°C.

2.3. Spectrophotometric measurements As shown in Fig. 2, an interface between a 2.0 ml aqueous solution of 0.5 mM AH2 (containing 0.05 M phosphate buffer) and a 0.2 ml NB solution of 20 mM Q was formed in a 10 mm quartz cell. While stirring the aqueous phase with a direct-mixing motor (SpectroCell, Inc. MTR-11D), the absorbance of the aqueous phase was monitored by means of a photodiode array spectrophotometer (Shimadzu MultiSpec1500).

T. Osakai et al. / Journal of Electroanalytical Chemistry 490 (2000) 85–92

87

3. Results

3.1. Cyclic 6oltammetry

Fig. 2. Spectrophotometric cell for monitoring a heterogeneous redox reaction at the NB W interface: (a) 10 mm quartz cell; (b) 2 ml aqueous solution; (c) 0.2 ml NB solution; (d) direct-mixing motor; (e) Teflon propeller; (f) light pass through the aqueous phase.

Fig. 3. Cyclic voltammograms of the stationary NB W interface in the absence (A) and presence (B) of 0.5 mM AH2 in W (pH 7.3). The NB phase contained 20 mM Q. For supporting electrolytes, see Cell (A). Scan rate: 50 mV s − 1.

Fig. 3 shows cyclic voltammograms for stationary NB W interfaces. Curve A represents the voltammogram which was recorded in the presence of 0.5 mM AH2 in W. Because no peaks were observed in the absence of AH2 (curve B), it was suggested that the negative current peak should be due to the oxidation of AH2 by Q. The peak separation (DEp) increased with the voltage scan rate (6); under the same conditions as in Fig. 3, DEp = 65, 70, 95, and 110 mV at 6=10, 20, 50, and 100 mV s − 1, respectively. This may be a consequence of the comparatively slow homogeneous charge-transfer process (Eq. (3)) as described later. In this way, the present system was shown to be a quasireversible system. However, the midpoint potential (Emid) between the negative and positive peaks was practically independent of 6 in the range of 10–100 mV s − 1. As shown in Fig. 4, a plot of Emid against pH (5.7–7.9) showed a straight line with a slope of 59 mV/pH. Except for a reversible wave due to a simple charge transfer, it is generally difficult to determine the number (n) of electrons involved in the charge transfer from the voltammetric peak current. In the voltammetric measurements with the stationary NB W interface, the reproducibility of the negative peak current was not very good (R.S.D.= 9 6%). This may suggest that a certain reaction proceeded at the interface even without potential control and changed the initial conditions before the voltammetric potential sweep was started. We then employed a polarographic technique with an ascending water electrode to determine the value of n from the limiting current.

3.2. Polarography Fig. 4. Plots of Emid ( ) and E1/2 ( ) against pH for the heterogeneous redox reaction between 0.5 mM AH2 in W and 20 mM Q in NB. The solid and dashed lines show regression lines with the slopes of 59 and 58 mV/pH, respectively, for Emid and E1/2.

Fig. 5 shows the polarograms which were recorded in (A) the presence and (B) absence of AH2 in W. Since the concentration of AH2 in W was much lower than that of Q in NB, the limiting current (Id) was determined by the diffusion of AH2 in W, depending linearly on the AH2 concentration in the range of 0.1–0.5 mM. It was also found that Id was practically independent of the Q concentration in the range of 5–20 mM. Therefore, Id (mA) should be given by the Ilkovicˇ equation [16]: Id = 3850n(DAH2)1/2m 2/3t 1/6c*

(1) 2

Fig. 5. Polarograms recorded with an ascending water electrode in the presence (A) and absence (B) of 0.5 mM AH2 in W (pH 7.3). The NB contained 20 mM Q. m =1.55 mg s − 1.

−1

where DAH2 is the diffusion coefficient (cm s ) of AH2 in W, m is the flow rate (mg s − 1), t is the drop time (s), and c* is the AH2 concentration (mol dm − 3). The coefficient of the r.h.s. of Eq. (1) was obtained using the measured density (1.068 g cm − 3) of the

88

T. Osakai et al. / Journal of Electroanalytical Chemistry 490 (2000) 85–92

aqueous phase (i.e. an aqueous solution containing 0.1 M NaCl and 0.5 M Na2SO4). The previous study [15] showed that the Ilkovicˇ equation could be employed to evaluate the limiting current for transfer of tetramethylammonium ion at the NB W interface. On the basis of Table 1 The values of Id and n at various pH values pH

Id a/mA

n

6.12 6.32 6.69 7.03 7.17 7.41 7.62

−18.02 −14.14 −15.03 −15.91 −15.25 −15.80 −15.36

2.1 1.7 1.7 1.8 1.8 1.8 1.8

a Obtained for 0.5 mM AH2 in W and 20 mM Q in NB; m=1.55 mg s−1.

the preliminary examination, we employed Eq. (1)) to calculate the n value for the present system. The n values obtained at various pH values are shown in Table 1. In this calculation, we used the value of DAH2 (5.0× 10 − 6 cm2 s − 1) which was determined from the diffusion current due to the oxidation of AH2 on a glassy carbon electrode in an aqueous solution of the same composition as phase III in Cell (A); the diffusion current was measured by potential-step chronoamperometry with a conventional three-electrode electrolytic cell (the applied potential was carefully set at a sufficiently positive potential, +0.9 V versus Ag AgCl KClsat, where the current should be limited only by the diffusion of AH2. The values of t (e.g. 11.4 s for Fig. 5) for the calculation of DAH2 were evaluated accurately by recording current–time curves during some drops at the potential where the limiting current was observed. As seen in Table 1, the n values thus determined are about 2 or slightly lower, suggesting that the present system should be a two-electron reaction rather than a one-electron reaction [11]. Fig. 4 also shows the pH dependence of the halfwave potential (E1/2), which was very close to Emid in cyclic voltammetry. In the pH range of 6.3–7.6, E1/2 as well as Emid showed a straight line with a slope of 58 mV/pH.

3.3. Spectrophotometric measurements

Fig. 6. Absorption spectrum recorded at 90 min after the stirring was started. The aqueous phase (pH 7.2) contained 0.5 mM AH2, while the NB phase contained 20 mM Q.

Fig. 7. Time dependence of the absorbances at (A) 456 and (B) 600 nm. The concentrations of AH2, and Q are as in Fig. 6, but the pH of the aqueous phase was (1) 7.8, (2) 7.2, or (3) 6.5. Curve (4) was obtained in the absence of AH2 in the aqueous phase (pH 7.8).

The heterogeneous redox reactions between 0.5 mM AH2 in W and 20 mM Q in NB were observed under no electrochemical control using the spectrophotometric cell shown in Fig. 2. The spectral changes were monitored while stirring the aqueous phase constantly (with pH 6.5, 7.2, or 7.8). Fig. 6 shows the absorption spectrum for pH 7.2 which was recorded 90 min after stirring was started. The absorption peaks at around 450 and 425 nm show that the semiquinone radical anion (Q’−) [17–19] was produced in the aqueous phase from Q initially added to NB. Furthermore, the broad peak around 550 nm suggested the formation of chloranilic acid (o550 = 183 at pH 7.4). Fig. 7 shows the time dependence of the absorbances at (A) 456 nm and (B) 600 nm, which correspond to Q’− and chloranilic acid, respectively; the wavelength (600 nm) for monitoring the formation of chloranilic acid was chosen as fairly long to minimize the interference from Q’− by tailing. As seen in the figure, after no absorbance was observed at either wavelength for several tens of minutes, a sudden increase in the absorbance at 456 nm was observed, followed by a gradual increase in the absorbance at 600 nm. These absorbance changes showed that chloranilic acid was formed through Q’− as the final product. At the present stage, we cannot declare what the trigger of this

T. Osakai et al. / Journal of Electroanalytical Chemistry 490 (2000) 85–92

89

the well-developed positive-current peak observed in cyclic voltammetry (see Fig. 3) could not be elucidated. Thus, we have proposed a new mechanism as shown in Fig. 8. In the pH range 6–8, AH2 dissociates into ascorbate (AH−) in the aqueous phase (for the first dissociation, pKa = 4.10 [22]). In the proposed mechanism, AH− undergoes an irreversible two-electron oxidation by Q being only partly distributed to the aqueous phase: AH − + H+ + Q“DHA+ QH2 (in W)

(3)

In this way, AH− is irreversibly transformed into DHA, whereas Q is reduced to QH2. The resultant QH2 should lead to a proportionation reaction with Q to produce the semiquinone radical anion (Q’−): Fig. 8. Proposed mechanism for the heterogeneous redox reaction between AH2 (“ AH− + H+) added to the aqueous phase (pH 6 – 8) and Q added to the NB phase. For details, see the text.

Fig. 9. Cyclic voltammograms for a saturated solution of Q in a 0.1 M NaCl aqueous solution (containing 0.05 M phosphate buffer; pH 7.0) at the PFC electrode (BAS Inc.; surface area = 0.071 cm2). For the preparation of the test solution, see the text. The potential scan was made as: (A) + 0.2 “ −0.2 “ + 0.3“ +0.2 V; (B) 0 “ +0.5 “ − 0.2“ 0 V. Scan rate: 1.0 V s − 1.

heterogeneous redox reaction is. However, we can conclude that if the necessary conditions are provided, a certain redox reaction can occur between AH2 and Q added to the respective phases, even without electrochemical control. This implies the possibility that neutral molecules of Q were partly distributed to the aqueous phase and then reacted with AH2.

4. Discussion It is known that AH2 undergoes an irreversible twoelectron oxidation on a metal electrode to form dehydro-L-ascorbic acid (DHA) [20,21]. If a similar irreversible two-electron reaction had occurred at the NB W interface, i.e. AH2 (W)+Q (NB) “ DHA (W) + QH2 (NB)

(2)

QH2 + Q X 2Q’ − + 2H+ (in W)

(4)

It can be thought that the interfacial transfers of Q’− from W to NB and vice versa can be observed as the ion-transfer current under electrochemical control. In this mechanism, the oxidation of 1 mol of AH2 leads to the interfacial transfer of 2 moles of monovalent Q’−. This is consistent with the fact that the value of n estimated from the polarographic limiting currents was about 2. Although a series of reactions up to the formation of Q’− may occur without electrochemical control, these reactions seem to be difficult to carry out, since Q is too hydrophobic to be abundantly distributed in W. However, once the transfer of Q’− to NB is raised by electrochemical control, a series of reactions is facilitated by the ion transfer. Thus, the reaction of interest can be regarded as a homogeneous electron-transfer reaction facilitated by a heterogeneous ion-transfer reaction. In the proposed mechanism shown in Fig. 8, the oxidation of AH2 by Q should occur, though gradually, in the absence of an externally applied potential. Accordingly, if the oil water interface to be tested is allowed to stand for a long time, the initial solution conditions near the interface should be changed to some extent. This seems to be the origin of the poor reproducibility of the cyclic voltammetric measurements with stationary interfaces, where the phase contact time could not be controlled accurately. For examining the validity of reaction (3), it is essential to know the standard redox potential of Q in the aqueous solution. Despite the low solubility of Q in water, we could obtain voltammetric waves for Q with a plastic formed carbon (PFC) electrode in the aqueous test solution, which was prepared by addition of 0.2 ml of a stock solution (10 mM) of Q in acetonitrile to 20 ml of a 0.1 M NaCl aqueous solution (pH 7.0), followed by mixing for 10 min (Q seemed to be only slightly dissolved, but mostly suspended). As shown in Fig. 9(A), Q gave a well-defined wave at a relatively

90

T. Osakai et al. / Journal of Electroanalytical Chemistry 490 (2000) 85–92

rapid scan rate (1.0 V s − 1); the potential scan was made from + 0.15 (around the rest potential) to − 0.20 V in the negative direction, followed by the positive to + 0.30 V, then the negative back to the initial potential. In this case, the negative and positive peaks were small, and the wave shape resembled an adsorption wave. At scan rates B 1.0 V s − 1, we could obtain only ill-defined waves which were affected by possible adsorption–desorption processes. However, when the potential scan was initiated at 0 V where Q can be reduced to QH2, a well-developed wave could be obtained at 1.0 V s − 1, as shown in Fig. 9(B). This suggests that QH2, being more soluble than Q, was produced electrolytically in good yield at the electrode surface, and electrolyzed reversibly under these conditions, hardly suffering from adsorption–desorption processes. From the midpoint potential of the quasi-reversible wave, we then estimated the apparent redox potential of Q to be + 0.08 V versus Ag AgCl KClsat at pH 7.0. This potential shows a high probability of Q oxidizing AH2 in the aqueous solution. In practice, we could observe the oxidation of AH2 by Q (i.e. reaction (3)), as shown in Fig. 10. The cyclic voltammograms were recorded using the PFC electrode 10 min after the addition of different amounts of Q to a 0.1 M NaCl aqueous solution (pH 7.0) containing 0.5 mM AH2. The positive peak around +0.25 V observed in the absence of Q (curve a) shows the irreversible oxidation of AH2. As seen in Fig. 10, the addition of Q made the positive peak shift negatively to around + 0.13 V corresponding to the oxidation peak of QH2, even if the added concentration was lower than the

equivalent point, i.e. 0.5 mM (curve d). This suggests that QH2 produced by the reaction with AH2 should function as an electron mediator for the oxidation of the unreacted AH2 remaining. It should also be noted that the positive peak current was not very dependent on the added concentration of Q, probably being limited by the initial concentration of AH2. On the other hand, the negative peak current due to the rereduction of Q increased remarkably with an increase in the added concentration of Q, showing that Q certainly reacted with AH2 and dissolved in the solution as QH2. Since the negative peak showed the characteristics of an adsorption wave, we failed to deduce the reaction stoichiometry from the concentration dependence. However, when Q was added at higher concentration than 0.5 mM, the test solution turned turbid owing to unreacted Q, indicating that the mole ratio of AH2 to Q is 1:1. Thus, it has been confirmed that AH2 undergoes two-electron oxidation by Q in the aqueous solution, in accordance with reaction (3). Similar two-electron oxidation reactions of AH2 by stable free radicals were reported [23–25]. Based on the mechanism shown in Fig. 8, the current–potential curve is given by the following equation (for the derivation and the symbols in the equation, see Appendix A): nmAH – 2K 2D RT RT E= Dw − ln − ln ’ − o f° Q F F FAm °Q’− mkKP[Q]2o





 





RT RT I2 pH− ln + DEref (5) F F I− Id By setting I= Id/2 (Id being given by (Eq. (A10)), we can obtain an expression for E1/2: n 2m 2AH – RT RT E1/2 = Dw ln − ln c* o f° Q’− − F m °Q’− F + 2.303

−









K 2D RT RT + 2.303 pH+ DEref ln 2 F F mkKP[Q]o (6)

Fig. 10. Cyclic voltammograms recorded with the PFC electrode 10 min after the addition of different amounts of Q to a 0.1 M NaCl aqueous solution (pH 7.0) containing 0.5 mM AH2. Q was added as the stock solution in acetonitrile so that the concentration became: (a) 0 mM; (b) 0.10 mM; (c) 0.25 mM; (d) 0.5 mM; (e) 0.75 mM (suspension). The potential scan was made as: 0“ + 0.3“ −0.2 “ 0 V. Scan rate: 0.1 V s − 1.

The observed slope of the E1/2 versus pH plot as well as that of the Emid versus pH plot (see Fig. 4) coincided well with the theoretical slope of 59 mV/pH (2.303RT/ F) expected from Eq. (6). On the whole the system of interest could be elucidated by the mechanism shown in Fig. 8. However, it may be necessary to take account of some subreaction(s). For example, the interfacial transfer of QH2 formed by the reaction with ascorbic acid cannot be completely excluded, because QH2 is somewhat hydrophobic. For redox reactions between NADH and quinone derivatives at the 1,2-dichloroethane water interface [13], it has been reported that a neutral species such as QH2 is formed in the DCE phase as a reaction product. Supposing that the interfacial transfer of QH2 occurs in the present system, the faradaic current should be reduced, because the electric charge provided

T. Osakai et al. / Journal of Electroanalytical Chemistry 490 (2000) 85–92

for chloranil by ascorbic acid is transported not only by Q’− but also by neutral QH2 giving no faradaic current. This may be the reason why the n-values determined by polarography are a little smaller than 2 (see Table 1). In this study, however, we have not made a rigorous theoretical analysis in which the interfacial transfer of QH2 is also considered, since it seems quite cumbersome or impossible. According to Eq. (5), the E versus log{I 2/(I − Id)} plot should show a straight line with the slope of − 59 mV (− 2.303RT/F). In practice, the logarithmic analysis based on Eq. (5) showed a straight line (data not shown), but the slope ( − 40 mV) was smaller than the theoretical slope. Furthermore, although the theoretical slopes of the E1/2 versus log c* plot and the E1/2 versus log[Q]o plot are, respectively, −59 and +118 mV (see Eq. (6)), they were rather different from the observed ones (− 30 and +24 mV, respectively). These gaps between theory and experiment seem to indicate certain defects in the present theory based on a semiempirical treatment of the reaction layer. We should recognize that the reaction layer thickness (m) should be not a constant but a function of the rate constant. For a simple electrode reaction (Ox+ne− “Red) with a prek ceding chemical reaction (A X Ox), m is known to be k% given by D/2k% (D being the diffusion coefficient) [26,27]. Unfortunately, this relation cannot be applied directly to the present complicated system. However, it seems that a more rigorous theoretical analysis could be made by employing a digital simulation technique for cyclic voltammograms. Finally, though further theoretical as well as experimental studies will be required to reach a conclusion, the proposed mechanism has been shown to give a better account of the present system than the previous mechanism based on a one-electron transfer at the NB W interface [11]. Acknowledgements The authors wish to thank Dr K. Kano of Kyoto University for his helpful advice on the reaction mechanism. The present study is partially supported by a JSPS research grant for the Future Program. Appendix A Here, Eq. (5) for the current – potential curve for the proposed mechanism (Fig. 8) is derived. In the mechanism, a series of chemical reactions occurring in the aqueous phase is somewhat complicated. Accordingly, it has been assumed that the chemical reactions take place in a definite thickness (m) of the reaction layer [26,27]; only AH− has been considered as the diffusing species in the aqueous phase.

91

Regarding the distribution equilibrium of Q at the NB W interface, we can write KD =

[Q]o [Q]

(A1)

In this equation and the following, [ ] stands for the concentration of a species, and the subscript ‘o’ represents the oil phase and no script, the reaction layer in the aqueous phase. The proportionation reaction between Q and QH2 is assumed to be in equilibrium: [Q’−]2[H+]2 (A2) KP = [Q][QH2] therefore [Q’−]2 =

KP[Q] [QH2] [H+]2

(A3)

Under the experimental conditions employed in this study, the concentrations of Q and H+ can be regarded as constant. Hence, differentiating Eq. (A3) with respect to t yields d[Q’−] KP[Q] d[QH2] 2[Q’−] = +2 (A4) dt [H ] dt The formation rate of QH2 can be given by using the rate constant (k) as d[QH2] = k[AH−][Q][H+] dt

(A5)

Using Eqs. (A1), (A4) and (A5), the formation rate of Q’− can be obtained: d[Q’−] kKP [Q]2o[AH−] = 2 (A6) dt 2K D [Q’−][H+] In the proposed mechanism (Fig. 8), only Q’− is responsible for the faradaic current. Accordingly, the amount of Q’− formed in unit time should be equated with the faradaic current (I) flowing through the interface. Therefore, we can write I d[Q’−] − =m (A7) FA dt where A is the surface area and F is the Faraday constant. Substituting Eq. (A6) in Eq. (A7) yields −

I kKP [Q]2o[AH−] =m FA 2K 2D [Q’−][H+]

(A8)

In this study the added concentration of ascorbic acid in W was chosen as much lower than that of Q in NB. Under these conditions, a necessary and sufficient amount of Q can be supplied by the rapid distribution to W, and the faradaic current will be expressed only by the diffusion of AH− in W: I= − nFAmAH – (c*−[AH – ])

(A9)

T. Osakai et al. / Journal of Electroanalytical Chemistry 490 (2000) 85–92

92

where mAH – is the mass transfer coefficient of AH− in the aqueous phase. Then the limiting current (Id) is given by (A10)

Id = − nFAmAH – c* From Eqs. (A9) and (A10), we obtain [AH – ]=

I −Id nFAmAH –

(A11)

Substituting Eq. (A11) in Eq. (A8) yields an expression of [Q’−]: [Q’−]=m

kKP [Q]2o Id −I 2K nmAH – [H+] I

(A12)

2 D

The faradaic current is also determined by the diffusion of Q’− in the oil phase, i.e. (A13) I = −FAm oQ’−[Q’−]o where m oQ’− is the mass transfer coefficient of Q’− in the oil phase. Here, it is assumed that the transfer of Q’− across the NB W interface is so rapid that its interfacial concentrations obey the Nernst equation: RT [Q’−]o w Dw ln ’− (A14) o f =Do f ° Q’− − F [Q ] w o where Dw o f ( f –f ) is the Galvani potential difference of the interface, Dw o f° Q’− is the standard ion-transfer potential of Q’−, and R and T have their usual meanings. By substituting Eqs. (A12) and (A13) in Eq. (A14), we obtain a theoretical expression for the current –potential curve: RT nmAH – RT 2K 2D w − ln − ln Dw ’ − o f =Do f ° Q F F FAm °Q’− mkKP[Q]2o





 





I2 RT RT ln[H+]− ln (A15) F F I −Id The potential (E) applied to the electrolytic cell is related to Dw o f as −

E = Dw o f + DEref

(A16)

where DEref is a constant which is determined only by the reference electrodes employed. Finally, Eq. (A15) can be rewritten as Eq. (5).

References [1] A.G. Volkov, D.W. Deamer, D.L. Tanelian, V.S. Markin, Liquid Interfaces in Chemistry and Biology, Wiley, New York, 1998. [2] Z. Samec, V. Marecˇek, J. Weber, J. Electroanal. Chem. 103 (1979) 11. [3] J. Hanzlı´k, Z. Samec, J. Hovorka, J. Electroanal. Chem. 216 (1987) 303. [4] G. Geblewicz, D.J. Schiffrin, J. Electroanal. Chem. 244 (1988) 27. [5] V.J. Cunnane, D.J. Schiffrin, C. Beltran, G. Geblewicz, T. Solomon, J. Electroanal. Chem. 247 (1988) 203. [6] Y. Cheng, D.J. Schiffrin, J. Electroanal. Chem. 314 (1991) 153. [7] R.A. Marcus, J. Phys. Chem. 94 (1990) 1050. [8] R.A. Marcus, J. Phys. Chem. 94 (1990) 4152. [9] R.A. Marcus, J. Phys. Chem. 95 (1991) 2010. [10] H.H. Girault, J. Electroanal. Chem. 388 (1995) 93. [11] M. Suzuki, S. Umetani, M. Matsui, S. Kihara, J. Electroanal. Chem. 420 (1997) 119. [12] M. Suzuki, M. Matsui, S. Kihara, J. Electroanal. Chem. 438 (1997) 147. [13] H. Ohde, K. Maeda, Y. Yoshida, S. Kihara, Electrochim. Acta 44 (1998) 23. [14] T. Osakai, S. Himeno, A. Saito, Bull. Chem. Soc. Jpn. 64 (1991) 1313. [15] S. Sawada, T. Osakai, M. Senda, Bunseki Kagaku 45 (1996) 1045. [16] A.J. Bard, L.R. Faulkner, Electrochemical Methods, Wiley, New York, 1980, p. 148. [17] J.J. Andre, G. Weill, Mol. Phys. 15 (1968) 97. [18] K. Kawai, Y. Shirota, H. Tsubomura, H. Mikawa, Bull. Chem. Soc. Jpn. 45 (1972) 77. [19] R. Gschwind, E. Haselbach, Helv. Chim. Acta 62 (1979) 941. [20] J.J. Ruiz, A. Aldaz, M. Domı´nguez, Can. J. Chem. 55 (1977) 2799. [21] P. Karabinas, D. Jannakoudakis, J. Electroanal. Chem. 160 (1984) 159. [22] D.R. Lide (Ed.), CRC Handbook of Chemistry and Physics, 80th ed., CRC Press, Boca Raton, FL, 1999. [23] M.S. Blois, Nature 181 (1958) 1199. [24] M. Ohnishi, H. Morishita, H. Iwahashi, S. Toda, Y. Shirataki, M. Kimura, R. Kido, Phytochemistry 36 (1994) 579. [25] M.B. Arnao, A. Cano, J. Herna´ndez-Ruiz, F. Garcı´a-Ca´novas, M. Acosta, Anal. Biochem. 236 (1996) 255. [26] K. Wiesner, Anal. Chem. 27 (1955) 1712. [27] M. Senda, Rev. Polarogr. (Kyoto) 6 (1958) 95.