Anal. Chem. 2001, 73, 337-342
Electrochemical Monitoring of Proton Transfer across Liquid/Liquid Interfaces on the Surface of Graphite Electrodes Taek Dong Chung and Fred C. Anson*
Arthur Amos Noyes Laboratories, Division of Chemistry and Chemical Engineering, California Institute of Technology, Pasadena, California 91125
Equilibrium partitioning of HClO4 between aqueous solutions and benzonitrile (BN) or nitrobenzene (NB) was measured and used to evaluate the pKa of the acid in the two organic solvents. The potential drop across the BN/ H2O interface was evaluated from the known potential drop across the NB/H2O interface and the voltammetrically measured formal potential of a ferrocenium/ferrocene redox couple confined within thin layers of the two organic solvents. The voltammetric reduction of tetrachloro-1, 4-benzoquinone in thin layers of BN was used to monitor changes in the concentration of protons in the layer during proton-consuming faradaic reactions. The rate of proton transfer from the aqueous to the nonaqueous phase across the BN/H2O interface was shown to be adequate to sustain proton-consuming reactions at the electrode/BN interface. Thin layers of immiscible organic solvents interposed between electrode surfaces and aqueous solutions have been shown to offer attractive advantages compared with conventional electrochemical cell configurations in a variety of applications.1-3 In a recent study, this approach was exploited to examine the electroreduction of O2 as catalyzed by cobalt porphyrins dissolved in thin layers of benzonitrile (BN) on the surface of graphite electrodes immersed in aqueous, acidic solutions.4 The rate of the catalyzed reduction of O2 is sensitive to the quantity of protons that partition into the thin layer of BN from the aqueous solution with which it is equilibrated.4 The protons consumed during the reduction of O2 are replaced by protons that cross the BN/H2O interface in order for sustained electroreduction of O2 to occur. This proton transfer from the aqueous to the nonaqueous phase is a special case of the general process of cross-phase ion transfer5,6 that must occur simultaneously with the electron transfer at the electrode surface in order for faradaic current to flow through the cell (Figure 1). To gain insight into cross-phase proton-transfer processes at the BN/H2O interface we examined the well-understood7 reduc* Corresponding author: (e-mail)
[email protected]; (telephone) (626) 3956000; (fax) (626) 577-4088. (1) Shi, C. N.; Anson, F. C. Anal. Chem. 1998, 70, 3114. (2) Shi, C. N.; Anson, F. C. J. Phys. Chem. B 1998, 102, 9850. (3) Shi, C. N.; Anson, F. C. J. Phys. Chem. B 1999, 103, 6283. (4) Steiger, B.; Anson, F. C. Inorg. Chem. 2000, 39, 4579. (5) Volkov, A. G.; Deamer, D. W.; Tanelian, D. L.; Markin, V. S. Liquid Interfaces in Chemistry and Biology; John Wiley & Sons: New York, 1998. 10.1021/ac0009447 CCC: $20.00 Published on Web 12/14/2000
© 2001 American Chemical Society
Figure 1. Schematic depiction of the cell configuration used in electrochemical experiments with thin layers of benzonitrile interposed between an electrode surface and an aqueous supporting electrolyte solution.
tion of a hydrophobic benzoquinone derivative in thin layers of BN equilibrated with aqueous solutions containing HClO4. A convenient method was devised for evaluating both the equilibrium partitioning of HClO4 from H2O into BN and the pKa of HClO4 in BN. With these data in hand, the dynamics of cross-phase proton transfer were examined by noting how the electroreduction of the quinone responded to changes in the concentration of HClO4 in the aqueous solutions and in the thin layers of BN that were equilibrated with the aqueous solutions. The results are summarized in this report. EXPERIMENTAL SECTION Materials. Benzonitrile (HPLC grade, Aldrich) was passed through activated molecular sieves (MX 1583 D-1, type 3A, 8-12 mesh, EM Industries) and distilled twice under reduced pressure. Tetrachloro-1,4-benzoquinone (Aldrich), tetrabutylammonium perchlorate (electrochemical grade, Fluka), perchloric acid (Mallinck(6) Girault, H. H. J.; Schiffrin, D. J. In Electroanalytical Chemistry; Bard, A. J., Ed.; Marcel Dekker: New York, 1989; Vol. 15. (7) Chambers, J. Q. In The Chemistry of Quinonoid Compounds; Patai, S., Rappoport, Z., Eds.; Wiley: New York, 1988; Vol. II.
Analytical Chemistry, Vol. 73, No. 2, January 15, 2001 337
rodt), nitrobenzene (>99.5%, Fluka), and dimethyl sulfoxide (>99.9%, EM Science) were used as received. 1,1′,3,3′-Tetrakis(2-methyl-2-hexyl)ferrocene was a gift from Prof. S. Strauss.8 Pyrolytic graphite electrodes (Advanced Ceramics Corp.): with 0.32 cm2 of the edges of the graphite planes exposed were mounted and pretreated as previously described.1 Apparatus and Procedures. Conventional electrochemical cells and instrumentation were employed. The procedure for introducing thin layers of organic solvents on the surface of graphite electrode has been described.1 Typically, 1 µL of BN was spread across the 0.32-cm2 electrode surface to produce a ∼30µm thin layer. The quantities of HClO4 that partitioned into benzonitrile or nitrobenzene from aqueous solutions were determined titrimetrically: 25 mL of the organic solvent was agitated vigorously with an equal volume of the aqueous solution of HClO4. The resulting mixture was transferred to a centrifuge tube and centrifuged for 20 min to facilitate the separation into two phases. After phase separation appeared to be complete, 10.0-mL aliquots of the upper, organic phase were transferred to a flask containing 20 mL of dimethyl sulfoxide, which was used so that only a single phase was present during the subsequent titration. The solutions were titrated with a standard aqueous solution of NaOH (0.01 or 0.1 M) to an indicator (phenolphthalein) end point. Blank corrections were obtained by repeating the equilibration with pure H2O in place of the aqueous HClO4 solution. RESULTS AND DISCUSSION Partitioning of HClO4(aq) into BN. Because of the substantial simplifications that result when ionic equilibration across two immiscible liquid phases involves only a single, uni-univalent electrolyte,5,6 we limited our initial experiments to the evaluation of the partition equilibrium of HClO4 between H2O and BN. Aqueous solutions of HClO4 were shaken with BN, and after separation of the two equilibrated liquid phases, aliquots of the BN phase were titrated with a standard solution of NaOH to determine the total quantity of HClO4 that had entered the BN phase. Relevant Equilibrium Relationships. The HClO4 that enters the BN phase consists of both H+ and ClO4- ions as well as undissociated HClO4 molecules. The titration with NaOH, of course, measures the total of dissociated and undissociated acid. With only a single electrolyte present, the Karpfen-Randles relationship9 applies to the galvani potential difference, ∆wo Φ, between the BN and H2O phases.
∆wo Φ ) 0.5(∆wo Φ°H+ + ∆wo Φ°ClO4-)
(1)
where the ∆wo Φ°i values are the standard ionic-transfer potentials for the transfer of H+ or ClO4- from water to BN and the activities of the ions in both phases are approximated by their concentrations.5,6 The concentration of protons in the BN phase, [H+]BN, that is equilibrated with an aqueous solution with a proton concentration [H+]W, is given by eq 2.5
[H+]BN ) [H+]W{exp[F/RT(∆wo Φ - ∆wo Φ°H+)]} where F/RT ) 39.1 V-1 at 25 °C. 338
Analytical Chemistry, Vol. 73, No. 2, January 15, 2001
(2)
Substitution of eq 1 in eq 2 yields eq 3.
{ [ (
)]}
∆wo Φ°ClO4- - ∆wo Φ°H+
[H+]BN ) [H+]W exp F/RT
2
(3)
The concentration of undissociated HClO4 in the BN, [HClO4]BN, is given by eq 4
[HClO4]BN )
[H+]BN[ClO[H+]2BN 4 ]BN ) Ka Ka
(4)
where Ka is the equilibrium constant for the dissociation of HClO4 in BN. The titration of aliquots of the equilibrated BN phase with NaOH measures S, the sum of [HClO4]BN + [H+]BN.
[
S ) [H+]WA 1 +
]
[H+]WA Ka
(5)
where
(∆ Φ° [0.5F RT
A ) exp
w o
ClO4-
]
- ∆wo Φ°H+)
(6)
Thus, the slope and intercept of a plot of S/[H+]w versus [H+]w allow values of A and Ka to be obtained from which [H+]BN and [HClO4]BN can be calculated. Comparison of the Partitioning of HClO4 from H2O into Nitrobenzene and Benzonitrile. The standard ionic-transfer potentials, ∆wo Φ°ClO4- and ∆wo Φ°H+, have been previously measured for the transfer of these ions between H2O and nitrobenzene (NB).5,6 We therefore carried out equilibration experiments with NB as the second, immiscible phase in order to compare the results with those expected on the basis of eq 5. Shown in Figure 2A is a plot of S/[H+]w versus [H+]w prepared from the titration data obtained when aqueous solutions of HClO4 were equilibrated with NB. The plot is reasonably linear and its intercept on the y-axis corresponds to ANB ) 2.6 ((0.03) × 10-4 and (Ka)NB ) 2.7 ((0.2) × 10-3 M. Using the tabulated values of ∆wo Φ°ClO4- ) -0.083 V and ∆wo Φ°H+ ) 0.337 V,5 the value of A calculated from eq 6 was 2.8 × 10-4. The good agreement between the measured and calculated values of A for the NB/H2O interface encouraged us to apply the same procedure to the BN/H2O interface. Titration data for BN that was equilibrated with varying concentrations of aqueous HClO4 are plotted in Figure 2B. The slope and extrapolated intercept of the linear portion of the plot were used to obtain ABN ) 1.2 ((0.2) × 10-3 and (Ka)BN ) 2.9 ((0.5) × 10-4 M. With lower concentrations of HClO4, the data points in Figure 2B lie above the extrapolated linear plot. We attribute these deviations to small quantities of the aqueous HClO4 that remain in the BN phase even after centrifugation to induce phase separation. With lower concentrations of HClO4, the densities of the aqueous and BN phase become more nearly equal, and (8) Clark, J. F.; Clark, D. L.; Whitener, G. D.; Schroeder, N. C.; Strauss, S. H. Environ. Sci. Technol. 1996, 30, 3124. (9) Karpfen, F. M.; Randles, J. E. B. Trans. Faraday Soc. 1953, 49, 823.
nium/ferrocene couple was the same in BN and NB. Thus,
(Ef)NB - (Ef)BN = ∆wo ΦNB - ∆wo ΦBN
Figure 2. Plots of S/[H+]w versus [H+] w based on eq 5 (see text). (A) Equilibration between H2O and NB phases. (B) Equilibration between H2O and BN phases.
complete phase separation is more difficult to ensure so that microscopic droplets of aqueous acid may remain in the organic phase. Titration of the acid in such droplets would produce positive deviations from the extrapolated lines in Figure 2B. For this reason, the slope and intercept of the linear plot obtained at higher concentrations of HClO4 were used to determine the values of ABN and (Ka)BN. Standard Ion-Transfer Potentials for the BN/H2O Interface. Values of ∆wo Φ°ClO4- and ∆wo Φ°H+ for the BN/H2O interface are not included in the available tabulations. To compare the properties of the BN/H2O interface with the extensively examined NB/H2O interface, the value of ANB (eq 6) evaluated for the BN/ H2O interface was combined with voltammetric data for a reversible redox couple in BN to obtain estimates of ∆wo Φ°ClO4- and ∆wo Φ°H+ using the following procedure: Dilute solutions of a hydrophobic ferrocene, 1,1′,3,3′-tetrakis(2-methyl-2-hexyl)ferrocene (MHFc), were prepared in NB and BN. Very low concentrations of MHFc were employed (80 µM) to minimize changes in the ionic composition of the thin layer caused by the flow of faradaic current.3 Thin layers of the resulting solutions were employed in a cell configuration like that shown in Figure 1 and cyclic voltammograms were recorded to measure the formal potentials, Ef, for the redox couple in each solvent with respect to the reference electrode in the aqueous phase that contained HClO4 at a convenient concentration (1.0 M). The difference in Ef values for the two solvents was assumed to reflect the difference in the potential drop, ∆wo Φ, across the interface in each case. That is, it was assumed that the formal potential of the ferroce-
(7)
∆wo ΦNB was calculated from eq 1 using the known values of ∆wo Φ°ClO4- and ∆wo Φ°H+ for the NB/H2O interface.5 Then, ∆wo ΦBN was calculated from eq 7 and (∆wo Φ°ClO4-)BN and (∆wo Φ°H+)BN were calculated from eqs 1 and 6. The various experimental and calculated parameters are listed in Table 1. The two standard ion-transfer potentials have similar magnitudes for both solvent pairs. However, much more total acid partitions into BN than into NB under the same conditions. The origin of the difference is traceable to the larger pKa of HClO4 in BN (3.5) than in NB (2.5), which is in accord with the smaller dielectric constant of BN (25.2) compared with NB (35.7). Monitoring the Concentration of Acid in BN. To follow the quantity of acid present in BN solutions when faradaic current is flowing, tetrachloro-1,4-benzoquinone (TCQ) was employed as a probe reactant because its electrochemical reduction yields a response that is sensitive to the concentration of acid at the electrode surface. Shown in Figure 3 are the responses obtained during the reduction of TCQ in bulk solutions of water-saturated BN. The voltammogram in Figure 3A was recorded in the absence of acid. The two peaks correspond to the reduction of TCQ to the anion radical, TCQ•-, followed by reduction to the dianion TCQ2-.7 Addition of an equimolar quantity of HClO4 to the solution caused the response to change to that shown in Figure 3B. The peak at the more positive potential corresponds to the two-electron reduction to the hydroquinone, TCQH2.7 The subsequent peaks can be assigned to unprotonated and partially protonated reduction products that are formed when the concentration of acid at the electrode surface is decreased by its consumption during the first reduction step. In Figure 3C is shown the result obtained in the presence of a stoichiometric excess of acid. Now all of the TCQ is reduced to TCQH2 so that only a single reduction peak is observed. The difference in the potentials where TCQ is reduced in the absence (Figure 3A) and presence (Figure 3C) of acid provides a convenient basis for determining whether acid is present at an electrode surface where the reduction of TCQ is proceeding. Cyclic Voltammetry of TCQ in Thin Layers of BN. TCQ was dissolved in the BN that was used to form a thin layer on the surface of a graphite electrode that was subsequently introduced into an electrochemical cell in a configuration like that shown in Figure 1. The hydrophobic TCQ is not soluble in the aqueous solution adjacent to the thin layer of BN, so it remained in the BN phase where its reduction at the graphite electrode was compared with that obtained in bulk solutions of BN (Figure 3). Shown in Figure 4A is the response obtained when the aqueous phase contained only 0.05 M NaClO4 as supporting electrolyte. (The concentration of NaClO4 in the BN layer was determined by the equilibrium partitioning of the salt from the aqueous into the BN layer.) The response in Figure 4A arises from the reduction of TCQ in the absence of protons (other than those available from the H2O in the water-saturated thin layer of BN). There is no anodic counterpart to the prominent cathodic peak during the first potential scan, and almost no response remains during a second potential scan. This behavior, which contrasts Analytical Chemistry, Vol. 73, No. 2, January 15, 2001
339
Table 1. Estimation of Standard Ion-Transfer Potentials interfacea NB/H2O BN/H2O
interfacea NB/H2O BN/H2O
103
Ab
A. Transfer of H+ and ClO4- at the BN/H2O and NB/H2O Interfaces 103 [H+],c M Ef, Vd ∆wo Φ°ClO4-, V
0.26 ((0.003) 1.2 ((0.2)
-0.035 0.007
0.26 1.2
-0.083f -0.002i
B. Transfer of Na+ at the BN/H2O and NB/H2O Interfaces Ef, Vd ∆wo Φ°ClO4-, V ∆wo Φ°Na+, V -0.060 -0.004
-0.083f -0.002j
0.358f 0.389g
∆wo Φ°H+, V
∆wo Φ, Ve
0.377f 0.340i
0.127g 0.169h
∆wo Φ, Ve 0.138g 0.194h
a Thin layers of each solvent were interposed between the surface of a graphite electrode and aqueous HClO (1.0 M) or NaClO (1.0 M). 4 4 From intercepts of plots such as those in Figure 2. c Concentration of protons in the nonaqueous phase as calculated from eq 3. d Formal potential, taken as the average of the anodic and cathodic peak potentials in cyclic voltammograms for 80 µM MHFc in a thin layer of NB or BN, measured with respect to a reference electrode in 1.0 M aqueous HClO4 or NaClO4. e Potential difference across the liquid/liquid interface. f From Table 1 in ref 5. g Calculated from eq 1. h Calculated from eq 7. i Calculated from eqs 1 and 6. j From part A of this table.
b
Figure 3. Cyclic voltammograms for 1 mM TCQ dissolved in bulk solutions of water-saturated BN containing 0.1 M tetrabutylammonium perchlorate. Concentrations of perchloric acid in the solutions: (A) 0, (B) 1.0, and (C) 4.0 mM. Scan rate, 50 mV s-1. Graphite electrode, 0.32 cm2.
sharply with that shown in Figure 3A, can be understood by considering the low level of supporting electrolyte present in the BN layer. Using the values of (∆wo Φ°Na+)BN and (∆wo Φ°ClO4-)BN from Table 1, the calculated concentration of Na+(and ClO4-) in BN equilibrated with 0.05 M NaClO4 is only 0.02 mM, which is much smaller than the concentration of TCQ in the BN layer of Figure 4A. Electric field-induced migration of the ions in the BN must occur during current flow. Thus, as TCQ•- anions are generated at the electrode surface, they migrate to and across the BN/H2O interface and are lost from the BN layer. As a result, little TCQ•- remains to be oxidized after the direction of the scan is reversed, and the second scan toward more negative potentials also produces a diminished response. The loss of TCQ from the thin layer would be smaller if Na+ ions from the aqueous phase 340 Analytical Chemistry, Vol. 73, No. 2, January 15, 2001
Figure 4. Cyclic voltammograms for 2.8 mM TCQ dissolved in a thin layer of BN in a cell configuration like that in Figure 1. (A) Supporting electrolyte in the aqueous phase, 0.050 M NaClO4; solid line, first scan; dashed line, second scan; S ) 3 µA. (B) Aqueous supporting electrolyte, 0.050 M HClO4, S ) 5 µA; first and second scans as in (A). Scan rate, 5 mV s-1.
crossed the BN/H2O interface as TCQ was reduced to TCQ•-. That such cross-phase transfer of Na+ is apparently not extensive presumably reflects the high electrostatic barrier that opposes the transfer (Table 1B). In Figure 4B is shown the response obtained when the aqueous solution was changed to 0.05 M HClO4. Now the reduction of TCQ occurs under a single peak with the symmetrical shape expected for voltammetry conducted in thin layers of solution.10 Comparison of the peak position with that in Figure 3C indicates that the TCQ is reduced to TCQH2. In contrast to Figure 4A, very little loss of (10) Hubbard, A. T.; Anson, F. C. In Electroanalytical Chemistry; Bard, A. J., Ed.; Marcel Dekker: New York, 1970; Vol. 4.
TCQ from the thin layer accompanies its reduction as is evident from the comparable magnitude of the cathodic and anodic peaks. In addition, a second scan toward more negative potentials produces an only slightly smaller cathodic peak current. The reason for the difference in the responses obtained with NaClO4 or HClO4 as the aqueous supporting electrolyte reflects the basicity of TCQ•- anions. With protons present in the thin layer, the anions produced by the reduction of TCQ at the electrode are neutralized by protonation. As a result, the only available ionic charge carriers are protons whose field-induced migration into and across the thin layer provides the ionic current required for the electrode reaction to proceed. The initial concentration of protons in the BN (0.06 mM) is somewhat greater than that of Na+ ions in Figure 4A (0.02 mM), but most of the protons consumed during the electrode reaction must be transferred across the BN/H2O interface. The higher efficiency of proton transfer compared with the transfer of Na+ reflects the high affinity of TCQ•- for protons as well as the intrinsically higher mobilities of protons in water-saturated BN. The area under the peak in Figure 4B, 4.4 × 10-4 C, corresponds to the reduction of 2.3 × 10-9 mol of TCQ to TCQH2 with the consumption of 4.6 × 10-9 mol of protons. (A total of 2.8 × 10-9 mol of TCQ was present in the volume of BN used to form the thin layer. The missing 0.5 × 10-9 mol of TCQ is believed to be lost in the portion of the BN that wets the heat-shrinkable tubing or is displaced from the electrode surface when it is immersed in the aqueous solution.) Mixed Supporting Electrolytes. When only a single cation is present in the aqueous phase, there is no ambiguity about the cationic component of the ionic current that passes across the BN/H2O interface. However, with supporting electrolytes containing more than one cation, the situation becomes less simple.5 We explored this case by repeating the experiments of Figure 4 with a supporting electrolyte in the aqueous phase that consisted of a mixture of HClO4 and NaClO4. The results are shown in Figure 5. When 0.10 M NaClO4 is added to the aqueous phase, the cathodic peak becomes broader (Figure 5B) and with 2.0 M NaClO4 the cathodic peak is followed by an extensive “tail” as the reduction of TCQ proceeds at ever more negative potentials (Figure 5C). There is a small second peak near the potential where TCQ is reduced in the absence of acid (Figure 4A). The oxidation of the TCQH2 in Figure 5C begins at the same position as in the absence of NaClO4 (Figure 5A) because the Na+ cations that entered the thin layer during the reduction of TCQ are replaced by protons that enter the thin layer to react with the basic TCQ•anions that are generated with Na+ as their counterion. However, HClO4 is generated within the thin layer during the oxidation of TCQH2 and the concentration of acid eventually exceeds its initial concentration and causes the oxidation of TCQH2 to shift to more positive potentials. This faradaic generation of acid is believed to be responsible for the double anodic peak in Figure 5C. Second scans (not shown), such as those in Figure 4, showed that very little TCQ was lost from the BN layer under the conditions employed in Figure 5. In Figures 4 and 5 the concentrations of supporting electrolyte present in the BN thin layers were small (
