Electroreduction of Diphenyl Disulfide on a Self-Assembled Lipid

Different mechanisms based on Ph2S2 adsorption either directly on a mercury drop or on phospholipids heads are proposed. The charge associated with th...
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Langmuir 2002, 18, 9377-9382

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Electroreduction of Diphenyl Disulfide on a Self-Assembled Lipid Monolayer on Mercury P. Lodeiro, R. Herrero, and M. E. Sastre de Vicente* Departamento de Quı´mica Fı´sica e Ingenierı´a Quı´mica I, University of A Corun˜ a, Alejandro de la Sota 1, 15071 A Corun˜ a, Spain Received May 22, 2002. In Final Form: September 17, 2002 In the present work, a voltammetric study of permeability and redox behavior of diphenyl disulfide (Ph2S2) through a self-assembled monolayer of dioleoylphosphatidylcholine adsorbed on mercury is carried out; the results are compared with those obtained on a bare electrode; in this case, the reduction of Ph2S2 proceeds in a reversible way with the formation of a mercurial compound. The presence of the monolayer of phospholipid provokes an increase in the process irreversibility. Different mechanisms based on Ph2S2 adsorption either directly on a mercury drop or on phospholipids heads are proposed. The charge associated with the adsorbed Ph2S2 electroreduction was employed to calculate its solubility, and the method is compared with semiempirical expressions proposed in the literature for obtaining approximate values of solubility.

Introduction The importance of electrochemical studies of thiol and disulfide groups stems from their presence in many biological molecules of interest. In fact, their redox reactions play an important role in the biological activity of molecules that contain them. Although diphenyl disulfide (Ph2S2) is not present in biological systems, it can constitute a model to interpret the aromatic ring activity, without protonatable groups, associated with thiol/disulfide groups. A voltammetric study of the permeability and redox behavior of Ph2S2 through a biomimetic membrane model has been carried out; different biomimetic models have been recently developed which may provide useful information on the electrochemistry of biomolecules in biological processes.1,2 The biomimetic model is obtained by deposition of a self-assembled phospholipid monolayer on a mercury electrode. The phospholipid coating was obtained by spreading a solution of phospholipid in pentane on the interface of an aqueous electrolyte, allowing the solvent to evaporate, and immersing a hanging mercury drop electrode in the electrolyte. This procedure gives rise to half a bilayer with the hydrocarbon tails directed toward the hydrophobic mercury surface and the polar heads directed toward the solution. The use of a lipid coated mercury electrode as a biomimetic membrane was introduced by Miller3 and subsequently adopted in a modified version by Nelson4 and later by Guidelli.5 The defect-free support to the lipid film provided by liquid mercury and the complete absence of pentane in the film impart to the film high mechanical stability, resistance to electric field and reproducibility that are not shared by planar black lipid membranes. This biomembrane model has been used to measured the intrinsic protonation constants of different phospholipids6,7 as well as their surface dipole potentials8 and (1) Guidelli, R.; Aloisi, G.; Becucci, L.; Dolfi, A.; Moncelli, M. R.; TadiniBuoninsegni, F. J. Electroanal. Chem. 2001, 504, 1-28. (2) Plant, A. L. Langmuir 1999, 15, 5128-5135. (3) Miller, I. R. In Topics in Bioelectrochemistry and Bioenergetics; Milazzo, G., Ed.; John Wiley: Bristol, U.K., 1981; Vol. 4, pp 161-224. (4) Nelson, A.; Benton, A. J. Electroanal. Chem. 1986, 202, 253-270. (5) Moncelli, M. R.; Guidelli, R. J. Electroanal. Chem. 1992, 326, 331-338.

the change in surface dipole potential with a change in the charge density of the polar heads of the phospholipids9 and upon adsorption of antitumor drugs.10 This model has also been used to study the redox behavior of molecules with important biological relevance as retinal,11 ubiquinone,12,13 vitamin K114 or gluthatione.15 Further developments have led to the obtaining of supported lipid bilayers based on alkanethiol-tethering chemistry, which are becoming important biomimetic materials.2,16,17 The important role of these metal-supported biomembrane models in the recent developments of bioelectrochemistry has been extensively reviewed in the excellent paper of Guidelli and co-workers.1 The work was then completed by performing an electrochemical study of Ph2S2 directly on a bare mercury electrode in aqueous solution to fill the gap observed in the literature owing probably to the low solubility of this compound. The estimation of the charge involved in the Ph2S2 reduction was employed to determine its solubility, comparing it with that of aldrithiol-2 for solubility calculations; the value obtained in this way was compared with that calculated from equations suggested in the literature.18 (6) Moncelli, M. R.; Becucci, L.; Guidelli, R. Biophys. J. 1994, 66, 1969-1980. (7) Moncelli, M. R.; Becucci, L. J. Electroanal. Chem. 1995, 385, 183189. (8) Becucci, L.; Moncelli, M. R.; Herrero, R.; Guidelli, R. Langmuir 2000, 16, 7694-7700. (9) Moncelli, M. R.; Becucci, L.; Buoninsegni, F. T.; Guidelli, R. Biophys. J. 1998, 74, 2388-2397. (10) Herrero, R.; Moncelli, M. R.; Guidelli, R.; Carla´, M.; Arcangeli, A.; Olivotto, M. Biochim. Biophys. Acta 2000, 1466, 278-288. (11) Nelson, A. J. Electroanal. Chem. 1992, 335, 327-343. (12) Moncelli, M. R.; Becucci, L.; Nelson, A.; Guidelli, R. Biophys. J. 1996, 70, 2716-2726. (13) Moncelli, M. R.; Herrero, R.; Becucci, L.; Guidelli, R. Biochim. Biophys. Acta 1998, 1364, 373-384. (14) Herrero, R.; Buoninsegni, F. T.; Becucci, L.; Moncelli, M. R. J. Electroanal. Chem. 1998, 445, 71-80. (15) Herrero, R.; Barriada, J. L.; Lo´pez-Fonseca, J. M.; Moncelli, M. R.; Sastre de Vicente, M. E. Langmuir 2000, 16, 5148-5153. (16) Krysinski, P.; Zebrowska, A.; Michota, A.; Bukowska, J.; Becucci, L.; Moncelli, M. R. Langmuir 2001, 17, 3852-3857. (17) Tadini-Buoninsegni, F.; Becucci, L.; Moncelli, M. R.; Guidelli, R. J. Electroanal. Chem. 2001, 500, 395-407. (18) Hansch, C.; Quinlan, J. E.; Lawrece, G. L. J. Org. Chem. 1968, 33, 347-50.

10.1021/la025977r CCC: $22.00 © 2002 American Chemical Society Published on Web 10/26/2002

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Ph2S2 shows different electrochemical behavior in the presence or absence of a phospholipid monolayer deposited on a mercury electrode. Different mechanisms based on Ph2S2 adsorption directly, either on a mercury drop or on phospholipid heads, are proposed. Experimental Section The chemicals used in this work were mercury (Merck Suprapur), KCl (Merck p.a., heated at 500 °C in a muffle furnace to remove organic impurities), KH2PO4, Na2HPO4, CH3COOH, and CH3COONa (Merck pro-analysis), HCl (Merck Suprapur), and aldrithiol-2 and Ph2S2, both from Aldrich. Dioleoylphosphatidylcholine (DOPC) (purity grade 1) was supplied by Lipid Products (South Nutfield, Surrey, U.K.) as solutions containing 100 mg of phospholipid in 5 mL of chloroform plus methanol. The working solutions for spreading at the air-water interface were prepared daily by dilution of 50 µL of the stock in 1 mL of pentane and were stored at -20 °C. All measurements were carried out at 25 ( 0.1 °C in aqueous solutions containing 0.1 M KCl to keep the ionic strength constant. The pH was controlled with HCl (range from 2 to 4), CH3COOH/ CH3COO- (range from 4 to 5.5), and H2PO4-/HPO42- (range from 5.5 to 8.3) buffers. The overall concentration of the acidic and basic components of the buffers is referred to as the buffer concentration. The homemade hanging mercury drop electrode employed in the measurements, the cell, and the detailed procedure to produce self-assembled lipid monolayers deposited on mercury are described elsewhere;6,7 the electrode area was 13.6 × 10-3 cm2. All potentials were measured vs a saturated calomel electrode (SCE). Differential capacitance measurements were made by means of a Metrohm Herisau Polarecord E 506 polarograph using phase sensitive ac voltammetry. The baseline potential was overlapped with a sinusoidal signal of 10 mV amplitude and a 75 Hz frequency, using a lag angle of 90°. The differential capacity of the bare and coated mercury electrodes was frequently measured against the applied potential to check the absence of impurities, good behavior of the electrode, and the stability and reproducibility of the lipid film. The experiments using lipid monolayer were conducted on a fully covered mercury drop where the minimum capacity was 1.7 ( 0.1 µF‚cm-2. Cyclic voltammetry measurements were made with a Metrohm E 506 polarograph connected to a Metrohm VA-Scanner E 612 triangular signal generator. Voltammograms were recorded on a Linseis LX 1600 recorder. The pH was measured before and after each experiment by means of a Crison micro pH 2000 pH-meter fitted to a radiometer combined glass electrode with an Ag|AgCl reference electrode. The electrode performance was periodically checked with solutions of accurately known pH.

Determination of the Aqueous Solubility of Ph2S2 The principal inconvenience that we found in performing the experiments was the solubility determination of this apolar compound in aqueous solutions and the absence of any references to this respect. An experimental adsorption isotherm of aldrithiol-2 was employed to determinate it. The method is based in comparing the charge obtained for the reduction of both adsorbed molecules under identical experimental conditions. Both molecules have a very similar structure, are strongly absorbed on the electrode, and exchange the same number of electrons, so we can suppose that the same charge corresponds to the same concentration. Moreover, at these low concentrations, aldrithiol-2 shows a linear isotherm behavior, discharging the possibility of strong lateral interactions between the adsorbed molecules, and at these low concentrations, the uncertainty in the actual surface area occupied by one Ph2S2 or aldrithiol-2 molecule has an almost negligible effect. Different quantities of aldrithiol-2 were added to a cell containing background electrolyte buffered at pH ) 7, to vary its concentration in a range from 1.5 × 10-7 to 6 × 10-6 mol dm-3. The reduction peaks for each concentration of aldrithiol-2 in cell were registered by cyclic voltammetry and obtained, by

Figure 1. Adsorption isotherm of aldrithiol-2, in 0.1 mol‚dm-3 KCl and phosphate buffer of pH 7, on a coated DOPC electrode. integration, the charge value associated with each one. The adsorption isotherm for aldrithiol-2 was obtained by plotting the charge vs aldrithiol concentration. Deviations from the linear range were observed at concentrations greater than 3 × 10-6 mol dm-3. A saturated aqueous solution of Ph2S2 was employed to make the same measurements under identical conditions; the charge for the reductions peaks registered was obtained and then introduced into the isotherm to obtain the aqueous solubility of Ph2S2: (2.2 ( 0.05) × 10-6 mol dm-3 (Figure 1). On the other hand, we tried to obtain this aqueous solubility value (S) using several empirical equations often described in the literature, on the basis of the existing correlation between different physical properties. Two of the ones applied, involving fusion temperature (Tm) and the distribution coefficient in octanol/ water solution (Kd,oct) are as follows.18

log(1/S) ) 1.214 × log(Kd,oct) - 0.850 + 0.0095(Tm - 25) r2 ) 0.912 (1) The log Kd,oct, value for Ph2S2 is 4.41.19The validity range for eq 1 with respect to log(Kd,oct) is 0.3-4.7. The validity range for eq 1 with respect to S (mol dm-3): 2 × 10-5 - 5. The solubility value obtained is S, where S ) 14 × 10-6 mol dm-3.

log(1/S) ) 1.339 × log(Kd,oct) - 0.978 + 0.0095(Tm - 25) r2 ) 0.874 (2) The validity range for eq 2 with respect to log(Kd,oct) is 0.3-4.7. The validity range for eq 2 with respect to S (mol dm-3) is 6 × 10-6 - 5. The solubility value obtained is S, where S ) 5.3 × 10-6 mol dm-3. Both equations are valid for benzene derivates.18 The values obtained are not coincident, this being evidence of the limitations of the empirical equation in this case.

Ph2S2 on a Bare Mercury Electrode Results. Ph2S2 is a symmetrical apolar compound. Its low solubility seems to have hindered electrochemical studies in aqueous solution; in fact, no single report on this compound refers to its electrochemical behavior directly on a bare mercury electrode in aqueous media. However, a few studies in water/organic solvent mixtures can be found.20-23 (19) Hansch, C.; Leo, A. Substituent Constants for Correlation Analysis in Chemistry and Biology; John Wiley & Sons: Claremont, CA, 1979; Vol. 1. (20) Kolthoff, I. M.; Stricks, W.; Tanaka, N. J. Am. Chem. Soc. 1955, 63, 520. (21) Magno, F.; Bontempelli, G.; Pilloni, G. J. Electroanal. Chem. 1971, 30, 375-383. (22) Nygard, B. Acta Chem. Scand. 1966, 20, 1710-1732.

Electroreduction of Diphenyl Disulfide

Figure 2. Pseudocapacitance (C) vs potential plot of 0.1 mol‚dm-3 KCl and HCl buffer of pH ) 3.93 on a bare mercury electrode, scan rate 7.5 mV‚s-1: (a) in the absence of Ph2S2 and (b) in the presence of 2.2 × 10-6 mol‚dm-3 Ph2S2.

Figure 3. Cyclic voltammetry of 2.2 × 10-6 mol‚dm-3 Ph2S2, in 0.1 mol‚dm-3 KCl and HCl buffer of pH ) 3.93 on a bare mercury electrode, scan rate 0.2 V‚s-1, stirring time 1 min.

The presence of Ph2S2 in the bulk solution is reflected in the pseudocapacitance curves due to its adsorption on the mercury electrode, producing a typical curve for these kinds of substances (Figure 2).24 The quadrature component of the electrode admittance Y in the absence of faradaic processes is equal to ωC, where ω is the angular frequency and C the differential capacity. In the presence of a faradaic process, Y is normally affected by such a process, and hence, the quantity Y/ω will be referred to as a pseudocapacity and should not be confused with the differential capacity. Cyclic voltammograms were recorded by scanning the potential from -0.1 to -1.1 V and then in the reverse direction. The influence of such experimental factors as scan rate (from 0.1 to 0.6 V s-1), stirring time (varied from 1 to 4 min to obtain quantitative peaks), and pH (from 2.5 to 8.3) was analyzed over the concentrations 4.7 × 10-7, 1.1 × 10-6, and 2.2 × 10-6 mol dm-3. The voltammograms obtained show different shapes depending on the experimental conditions. The most simple shape was obtained at acid pH (Figure 3); it presents a first reduction peak that is very wide and not entirely reproducible, followed by a second reduction bellshaped peak, at more negative potential. We have devoted our voltammetric studies to it. In the positive scan, no oxidation current is normally observed; only very concrete conditions like high scan rates, cause the appearance, and the gradual increase, of (23) Persson, B.; Nygard, B. Electroanal. Chem. Interface Electrochem. 1974, 56, 373-383. (24) Bockris, J. O. M.; Khan, S. U. M. Surface electrochemistry: a molecular level approach; Plenum: New York, 1993.

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Figure 4. Cyclic voltammetry of 1.1 × 10-6 mol‚dm-3 Ph2S2, in 0.1 mol‚dm-3 KCl and phosphate buffer of pH ) 6.87 on a bare mercury electrode, scan rate 0.4 V‚s-1, stirring time 4 min.

an oxidation peak at the same potential as that of the cathodic peak (Figure 4). An increase in stirring time for a Ph2S2-saturated solution involves a shift in the negative direction of the first cathodic peak that overlaps with the main peak, making difficult any quantitative analysis (Figure 4). An analogous behavior could be observed with more dilute Ph2S2 solutions at greater stirring times. An increase in pH to a value greater than 6 involves more complex voltammograms; several overlapping peaks are obtained with different intensities and potentials depending on experimental conditions (Figure 4). For these reasons, we have selected the conditions that allowed us to obtain simple voltammograms; then we performed a systematic study of the effects on the potential and intensity of the main peak at different scan rates, pH, and concentrations of Ph2S2. Discussion. The depressed zone presented in the pseudocapacitance curves (Figure 2) in the presence of Ph2S2 is due to its adsorption; normally this zone appears limited by two capacitance peaks caused by the adsorption and desorption processes taking place at the electrode.24 In this case, the peak at more negative potential is a pseudocapacitance peak due to Ph2S2 electroreduction across the lipid film to the corresponding thiolic form.15 This fact was confirmed by the coincidence of the potential peak, even at different pH values, with that obtained in cyclic voltammetry. Despite this finding, it is significant that the pseudocapacity values at more negative potentials than the pseudocapacity peak of Ph2S2 are the same as those obtained for the background electrolyte; these results suggest that the reduction form of Ph2S2 does not remain adsorbed on electrode. In cyclic voltammetric measurements, no Faradaic signal is observed when stirring is not applied prior to the registration; on the other hand, a bell-shaped symmetric peak was obtained in the negative scan after stirring. This peak shape is characteristic of an electroreduction of an adsorbed species in coincidence with the pseudocapacitance behavior. An additional proof was obtained by plotting log Ipeak vs log v for the reduction peaks; this plot is roughly linear and exhibits a practically unit slope for all curves. Unit slopes were obtained at all pH values and Ph2S2 concentrations investigated. These characteristics are clearly indicative of an electroreduction of an adsorbed species.

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The simplest voltammograms showed another peak before the main one or even a series of overlapping peaks could be observed in a reduced range of potential (Figure 4). Because of low solution concentration, it is difficult to suppose that these peaks are due to direct reduction of Ph2S2 that arrives by diffusion to the mercury surface. In fact, no voltammetric signal was obtained in the absence of stirring prior to the measurement. Consequently, this peak can be ascribed to the presence of energetically different states of Ph2S2 adsorbed on the mercury drop. The main bell-shaped reduction peak seems to indicate the reversible nature of the electrode process relative to the main peak. The reversibility of the process was also confirmed by the independence of the main peak potential on scan rate at the different pHs studied. The absence of anodic peak is justified because the reduction product does not stay adsorbed on the electrode: it diffuses to the bulk solution as soon as it is formed. This hypothesis was confirmed by two experimental findings: the pseudocapacitance curves for saturated Ph2S2 solutions coincided with those of the background electrolyte at more negative potentials than the Faradaic peak potential value; moreover, the presence of oxidation peaks with diffusion shape at the greatest scan rates studied confirms the hypothesis since the time that is allowed to the species to diffuse away from the electrode is decreased, producing a small signal due to its oxidation. The slope of the Epeak vs pH plot is about equal to 60 mV. This behavior is expected for the reversible reduction of Ph2S2 to thiol involving the same number of protons and electrons. For an ideal Nernstian reaction under Langmuirian conditions, the total width at half-height of either reduction or oxidation peak is given by 90.6/n mV at 25 °C. Width at half-height of about 100-140 mV was obtained for reduction peaks. This value accounts for the transfer of 1 electron and suggests the presence of lateral interaction energies.25 On the basis of these studies, the following reaction was proposed for the process associated to the main reduction peak: 2Hg

2e- + 2H+

(Ph-S-S-Ph)ads7982(PhSHg)ads798 2Ph-SH + 2Hg pH < 5 The protonated thiol formation was proposed because of the acidic media, lower than thiol pK, at which the studies could be performed. At greater pH values, no reliable measurements of the cathodic peak could be carried out due to the high complexity of the voltammograms obtained. Ph2S2 on a DOPC-Coated Mercury Electrode Results. The first part of the electrochemical behavior of Ph2S2 on a DOPC-coated mercury electrode study consists of the description of the formation and characteristics of a phospholipid self-assembled monolayer on this electrode. Before the recording of cyclic voltammograms, the differential capacity of the lipid monolayer was always measured against the applied potential to check the stability and reproducibility of the film.26 The pseudocapacitance curves were recorded by scanning the potential from -0.2 to -1.5 V. Figure 5 shows (25) Angerstein-Kozlowska, H.; Conway, B. E. J. Electroanal. Chem. 1979, 95, 1-28. (26) Nelson, A. J. Electroanal. Chem. 1988, 244, 99-113.

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Figure 5. Pseudocapacitance (C) vs potential plot of 2.2 × 10-6 mol‚dm-3 Ph2S2 in 0.1 mol‚dm-3 KCl and phosphate buffer of pH ) 6.85 on a DOPC-coated electrode, scan rate 7.5 mV‚s-1.

Figure 6. Cyclic voltammetry of 2.2 × 10-6 mol‚dm-3 Ph2S2, in 0.1 mol‚dm-3 KCl and phosphate buffer of pH ) 6.23 on a DOPC coated mercury electrode, scan rate 0.6 V‚s-1.

the pseudocapacity curve of a dioleoylphosphatidylcholinecoated mercury electrode in a background solution saturated with Ph2S2, in which a pseudocapacitance peak due to Ph2S2 reduction can be observed in the minimum capacity region. It is interesting to emphasize that the presence of this compound does not alter the monolayer characteristics. It does not originate any changes either in the capacitance minimum or in the reorientation cathodic peaks. Changes in pH values did not alter the potentials of the capacitance peaks but did shift the potential of the reduction peak to more negative values until pH ) 6, where it rose to a constant value. Similar to the situation with the bare mercury electrode, a systematic study using cyclic voltammetry was conducted. A KCl background electrolyte saturated with Ph2S2 was employed in all the experiments, varying scan rate, pH (from 2.5 to 8.3), and buffer concentration. The presence of the DOPC monolayer induced great changes in the electrochemical behavior of Ph2S2. Contrary to that observed on a bare electrode, a Ph2S2 reduction peak obtained on a freshly formed drop was visualized in the absence of stirring, prior to the registration. Cyclic voltammograms were recorded by scanning the potential from -0.1 to -1.1 V and then in the reverse direction. Figure 6 shows a cyclic voltammogram for a saturated solution of Ph2S2 with no stirring at pH ) 6.23;

Electroreduction of Diphenyl Disulfide

Figure 7. Cyclic voltammetry of 2.2 × 10-6 mol‚dm-3 Ph2S2, in 0.1 mol‚dm-3 KCl and phosphate buffer of pH ) 6.81 on a DOPC-coated mercury electrode, scan rate 0.4 V‚s-1, stirring time 1 min.

in this situation, a simpler voltammogram than the respective one for Ph2S2 on a bare electrode was obtained. The voltammograms showed only one bell-shaped asymmetrical peak in the cathodic scan and the capacity peaks corresponding to reorientations of the adsorbed phospholipid monolayer at more negative values in cathodic and anodic scans at the same potential. The application of a specific stirring time prior to the registration involved an increase in the complexity of the voltammograms; several overlapping peaks that form a width band in cathodic scan preventing a quantitative analysis are obtained (Figure 7). Sometimes it was possible to register a peak in the positive scan, but in all these cases was a small, badly developed and diffusion shape peak that appeared at more positive potentials than those of the cathodic peak. The voltammetric study was carried out in situations where only one reduction peak was observed. An important difference in relation to the study in the absence of phospholipid was the possibility of performing experiments at basic pH values without additional voltammogram complexity. Discussion. The DOPC-coated electrode shows great affinity to Ph2S2 molecules. The rapid penetration through the polar heads of the phospholipid gives a reduction peak without agitation. It could be observed that the presence of Ph2S2 does not alter the characteristics of the phospholipid monolayer. No changes in differential capacitance minimum, especially in the first reorientation cathodic peak, show that the Ph2S2 does not modify the phospholipid concentration adsorbed on the mercury drop and, hence, neither the characteristics of the monolayer. Thus, the disulfide molecules that penetrate into the monolayer stay adsorbed, probably in the polar head region and not directly on the mercury surface where it arrives because of a rapid translocation from the polar heads of phospholipid molecules.27-30 If Ph2S2 was adsorbed directly on the mercury drop, it would replace the phospholipid molecules; this would (27) Miller, I. R.; Bach, D.; Teuber, M. J. Membr. Biol. 1978, 39, 49-56. (28) Miller, I. R. J. Membr. Biol. 1988, 101, 113-118. (29) Lecompte, M. F.; Miller, I. R.; Elion, J.; Benarous, R. Biochemistry 1980, 19, 3434-3439. (30) Moncelli, M. R.; Herrero, R.; Becucci, L.; Guidelli, R. J. Phys. Chem. 1995, 99, 9940-9951.

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Figure 8. Plots of Epeak vs pH for reduction peak as obtained from cyclic voltammogram of 2.2 × 10-6 mol‚dm-3 Ph2S2 in 0.1 mol‚dm-3 KCl, scan rate 0.3 V s-1, at a DOPC-coated electrode.

involve a decrease in the coated fraction of the electrode and consequently a decrease in the intensity value of the first cathodic reorientation peak of the phospholipid. The voltammogram confirms the behavior observed in the capacitance curves on the ability of Ph2S2 to penetrate into the phospholipid monolayer. The bell-shaped peak obtained by cyclic voltammetry is characteristic of an adsorption process. Again a plot of log Ipeak vs log v leads to a straight line with unit slope at all pHs investigated, confirming an adsorption process for the Ph2S2 molecule. In contrast to the study in the absence of DOPC, the voltammetric process presented a certain degree of irreversibility, as indicated the following facts: the separation in the potential value for the cathodic and anodic peaks, when the last one appears, the peak potential dependence with scan rate, and the asymmetrical shape of the cathodic peak. The slope of the plot of Epeak vs log v for the reduction peak measures a charge-transfer coefficient (R) that is practically equal to unity. A value of R equal to unity denotes a rate-determining chemical step preceded by the reversible uptake of the first transferring electron.13 These results were confirmed by the width at half-height peak analysis. The total width at half-height of either reduction or oxidation peak in a irreversible process is given by 62.5/R mV at 25 °C. Widths at half-height about 60 mV were obtained for the reductions peaks; this value accounts for the transfer of one electron and suggests the absence of lateral interaction energies,25 in contrast with the behavior observed on a bare electrode. This difference can be justified by taking into account that the Ph2S2 does not stay directly adsorbed on the mercury drop in this case; its presence in the polar heads of the phospholipid, where Ph2S2 interaction takes place with less intensity, was proposed. Figure 8 shows two different regions for the Epeak vs pH plot. This entails modifications in the electrode process as a function of pH. The slope for the first part of the curve, from pH ) 3 to 6.3, is practically equal to 60 mV slope. This is the expected behavior for a reaction involving n protons and n electrons (n ) 1), and then, we can conclude that the rate-determining chemical step is the protonation of the Ph-S- anion. For pH > 6.3 the peak potential is once again pH independent. The intersection of both curves leads to a pK value of 6.30 ( 0.02, within the experimental

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error, in good agreement with a pK ) 6.5 determined by spectrophotometry.31 At these pH values the concentration of proton is very low so we can suppose again a rate-determining protonation step following the reversible uptake of the first transferring electron, in which the proton donor is now represented by water. The Ph2S2 behavior on a DOPC-coated electrode can be justified according to the following overall reactions:

pH < 6.30 2Hg

Ph-S-S-Ph f Ph-S-S-Phads 798 2e-

2H+

2Ph-S-Hg 798 2Ph-S- + 2Hg 798 2Ph-SH + 2Hg pH > 6.30 2Hg

Ph-S-S-Ph f Ph-S-S-Phads 798 2e-

2H2O

2Ph-S-Hg 798 2Ph-S- + 2Hg 798 2Ph-SH + 2Hg The reduction of Ph2S2 on a DOPC monolayer does not depend on the buffer concentration at constant pH and hence does not satisfy the principles of general acid-basic catalysis. This behavior can be explained by assuming that the protonation takes place well inside the polar head region of the DOPC monolayer. This monolayer is practically impermeable to the buffer molecules within the time (31) Danehy, J. P.; Paramewaran, K. N. J. Chem. Eng. Data 1968, 13, 3.

scale of the experiment, while it may be permeated by the proton itself. In this case, the role of the buffer is exclusively that of maintaining the pH just outside the lipid layer constant during Ph2S2 reduction, through a dissociation reaction in quasi-equilibrium.15 Hence, the rate-determining protonation step is affected by a change in pH, but not by a change in the buffer concentration at constant pH. Conclusions The Ph2S2 reduction on a bare mercury electrode in aqueous solution proceeds in a reversible way with the formation of a mercurial compound adsorbed on an electrode like it was reflected in the corresponding mechanism. The presence of a phospholipid monolayer coating the mercury drop modifies the previous mechanism, with an increase in the process irreversibility and facilitating Ph2S2 adsorption into the polar region of the lipid layer. The study of Epeak reduction dependence with pH value allow us to obtain the pK value for benzenethiol: 6.30 ( 0.02. The charge associated with the adsorbed Ph2S2 reduction was employed to calculate its solubility by comparing it with the aldrithiol adsorption isotherm; on the other hand, empirical equations described in the literature were employed to obtain an approximate value for this solubility. Acknowledgment. R.H. wishes to thank Xunta de Galicia for financial support through Project PGIDT99PXI10302A. LA025977R