Roles of Ion Pairing on Electroreduction of Dioxygen in Imidazolium

Jan 4, 2008 - The ion-pairing phenomenon during the dioxygen (O2) reduction reaction (ORR) in a room-temperature ionic liquid (RTIL) was recognized an...
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J. Phys. Chem. C 2008, 112, 1269-1275

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Roles of Ion Pairing on Electroreduction of Dioxygen in Imidazolium-Cation-Based Room-Temperature Ionic Liquid Md. Mominul Islam and Takeo Ohsaka* Department of Electronic Chemistry, Interdisciplinary Graduate School of Science and Engineering, Tokyo Institute of Technology, Mail Box G1-5, 4259 Nagatsuta, Midori-ku, Yokohama 226-8502, Japan ReceiVed: October 2, 2007; In Final Form: October 26, 2007

The ion-pairing phenomenon during the dioxygen (O2) reduction reaction (ORR) in a room-temperature ionic liquid (RTIL) was recognized and extensively studied using a voltammetric technique. Cyclic voltammetric two one-electron ORRs at a gold electrode in dimethyl sulfoxide (DMSO) solution containing 0.1 M tetraethylammonium perchlorate and in 1-ethyl-3-methylimidazolium tetrafluoroborate (EMIBF4) RTIL were carried out, and the obtained results were compared. Two cathodic peaks at -0.71 and -1.35 V versus Ag|AgCl|NaCl (sat.) and a well-defined anodic peak at -0.71 V and two ill-defined anodic peaks at potentials more positive than -0.2 V were observed in EMIBF4. The first and second cathodic peak potentials obtained in EMIBF4 were by 0.07 and 0.65 V more positive than those observed for the reductions of O2 to superoxide (O2•-) and O2•- to peroxide (O22-) in DMSO solution, respectively. Such a potential shift in the ORR was considered to be due to the ion-paring phenomenon as observed in the redox reaction of 1,4-benzoquinone in DMSO in the presence of EMIBF4. A 57 mV positive shift of the midpoint potential of the O2/O2•- redox couple with a 10-fold increase of EMIBF4 in DMSO solution was obtained. The association constant and free-energy change (∆G) of the ion-pairing were determined to be 13.5 M-1 and -6.4 kJ mol-1, respectively. The value of ∆G for the comproportionation reaction of O22- and O2 to form O2•- in EMIBF4 solution (-61.7 kJ mol-1) was one-half of that in DMSO solution (-117.7 kJ mol-1), suggesting a favorable stability of O22in EMIBF4.

Introduction Recently, the dioxygen (O2) reduction reaction (ORR) in room-temperature ionic liquids (RTILs) as well as in conventional protic and aprotic solutions has received a great deal of attention in view of its practical applications in fuel cells and the electrosynthesis of reactive oxygen species such as the superoxide ion (O2•-) and also from a viewpoint of fundamental electrode reactions (especially as a one-electron process).1-27 Several research groups have studied the ORR in various RTILs.15-27 AlNashef et al. have reported a quasi-reversible oneelectron ORR (i.e., O2/O2•- redox couple) at a glassy carbon (GC) electrode in 1-n-butyl-3-methylimidazolium hexafluorophosphate RTIL,16,17 while the ORR in imidazolium chloridealuminum chloride molten salts was first observed to be an irreversible two-electron process due to the presence of protonic impurities in the used RTILs.23 Our group has also studied the ORR in imidazolium-cation-based RTILs (ImRTILs) at different electrodes [e.g., GC, platinum (Pt), and gold (Au) electrodes and a hanging mercury drop electrode (HMDE)].15,18 We have reported for the first time that, in contrast to GC, Pt, and Au electrodes, HMDE assists a two-electron quasi-reversible redox reaction in ImRTILs.15 Besides ImRTILs,15-18 the redox reaction of the O2/O2•- couple has also been extensively investigated in quaternary ammonium and pyrrolidinium-cation-based RTILs.19-26 The O2/O2•- redox couple has been extensively studied in different RTILs using a Au microelectrode to examine the diffusion of O2 and O2•- species in such a highly viscous and * To whom correspondence should be addressed. Telephone: +81-45924-5404. Fax: +81-45-924-5489. E-mail: [email protected].

ionic solution (i.e., RTILs) as well as the kinetic analysis of thereactionbetweenelectrogeneratedO2•- andcarbondioxide.19-21,24 Unlike the conventional aprotic solutions, cyclic voltammetric redox peaks of the O2/O2•- couple in RTILs have been generally observed to be unsymmetrical in shape [i.e., the anodic peak corresponding to reoxidation of O2•- to O2 is obtained with slightly distorted shape at the conventional-sized electrode15-18,23,24 (e.g., diameter: 1.6 mm)], and interestingly, the anodic peak has been observed under the usual steady-state cyclic voltammetric condition at the microelectrode19-21 (e.g., diameter: 10 µm) as obtained with a static, conventional-sized electrode. The observation of such an anodic “peak” at the microelectrode has been considered to be due to a specific interaction (e.g., ionpairing) of O2•- with the RTIL cation. In another work,23 it has been claimed that the second step of the ORR in ImRTILs is not a simple reduction of O2•- to O22-, which has been commonly observed in aprotic solutions, but the reduction of the O2•- complex that may be formed by the chemical reaction of electrogenerated O2•- species and the cation of ImRTILs.10-14 Moreover, the diffusion coefficient (DO2•-) of O2•- in RTILs has been observed to be much smaller than that of O2,21 in contrast with the case in aprotic solutions [e.g., (dimethyl sulfoxide (DMSO)12,13] in which DO2•- and DO2 (diffusion coefficient of O2) are almost equal or in which DO2•- is slightly larger than DO2. In the present paper, the effect of ion pairing on the ORR in RTILs will be presented. Ion-pairing effects on the redox reactions in RTILs have been often observed and discussed based on the formation of an ion pair between the cation of RTILs and the negatively charged electrogenerated species.27-37

10.1021/jp7096185 CCC: $40.75 © 2008 American Chemical Society Published on Web 01/04/2008

1270 J. Phys. Chem. C, Vol. 112, No. 4, 2008 A strong ion pairing (complexation) between metal cations and electrogenerated species has been well-known since earlier times.9-11,14,33-38 The ion-pairing effects of metal cations on the ORR9-11,14 and the electrooxidation of the hexachloroiridate(III) anion33 and bis(fulvalene)dinickel35 have been reported; for example, the two consecutive one-electron oxidation steps of bis(fulvalene)dinickel have been found to merge together in the presence of sodium ion, resulting in the two-electron single response.35 Fry28 has observed a similar effect in the redox reactions of 1,n-dinitrobenzenes (DNBs) in ImRTILs, where the two consecutive one-electron cyclic voltammetric redox responses that are usually observed in conventional aprotic solutions coalesce into one two-electron redox response due to a strong ion-pairing phenomenon. The ion-pairing effect has been also reported for the redox reactions of benzaldehydes30 and nitrobenzenes29 in ImRTILs. In the present study, we examined the cyclic voltammetric two-step ORR, that is, two one-electron reductions of O2 to O2•and O2•- to peroxide (O22-) at a Au electrode in DMSO solution and 1-ethyl-3-methylimidazolium tetrafluoroborate (EMIBF4). For a clear understanding of these reductions, the redox reaction of 1,4-benzoquinone (BQ) was also studied in both solutions. Furthermore, the effect of EMIBF4 on the redox reactions of BQ and ORR in DMSO solution was examined. Various thermodynamic parameters concerning the so-called ion pairing in the ORR were estimated. Experimental Section Reagents. 1-Ethyl-3-methylimidazolium tetrafluoroborate (EMIBF4) with a purity of more than 99% and less than 30 ppm (ca. 2.1 mM) H2O was obtained from Stella Chemifa Co. (Japan) and was used as received. Dimethyl sulfoxide (DMSO) of reagent grade was purchased from Kanto Chemical Co. Inc. and dried over molecular sieves (4A 1/16, Wako Pure Chemicals Industries). Tetraethylammonium perchlorate (TEAP) (Kanto Chemical Co. Inc.) was used as a supporting electrolyte. 1,4Benzoquinone (BQ) was available from Kanto Chemical Co. Inc. The H2O used in this study was deionized H2O purified with a Milli-Q Millipore system. Apparatus and Procedures. The electrochemical cell was a conventional two-compartment Pyrex glass container with a gold (Au; φ ) 1.6 mm or 10 µm) or glassy carbon (GC, φ ) 1.0 mm) working electrode, a spiral Pt-wire counter electrode, and a silver (Ag)|silver chloride (AgCl)|sodium chloride saturated [NaCl (sat.)] or a Ag wire (quasi-) reference electrode. The working electrodes of Au and GC were polished with alumina powder and then washed with Milli-Q water by sonication for 10 min. Furthermore, the Au electrode was electrochemically pretreated in Ar-saturated 0.05 M H2SO4 solution by repeating the potential scan in the range of -0.2 to 1.5 V until the voltammogram characteristic for a clean Au electrode was obtained.1 Prior to use, the electrode was washed well with Milli-Q water and dried by blowing air. In order to minimize the IR drop yielded across the cell and the contamination of H2O into the cell solution, the reference electrode was placed in a glass beaker containing a NaCl-saturated aqueous solution and connected to the cell solution via a salt bridge filled with the solution of interest (EMIBF4 or DMSO containing TEAP).4 All of the measurements were carried out with an IR drop compensation using a positive feedback circuit. Steadystate voltammograms were measured at a Au microdisk electrode (φ ) 10 µm). Voltammetric measurements were carried out using computer-controlled electrochemical systems (models 50W, Bioanalytical Systems, Inc., and ALS/CHI 832

Islam and Ohsaka

Figure 1. CVs obtained at a Au electrode (φ ) 1.6 mm) in O2-saturated DMSO solution containing 0.1 M TEAP. Potential scan rate: 0.1 V s-1.

A). Before measurements, O2 or N2 gas was bubbled directly into the cell to obtain its saturated solutions, and during the measurements, O2 or N2 gas was flushed over the cell solution. All of the measurements were carried out at room temperature (25 ( 2 °C). Results and Discussion ORR in DMSO Solution. Figure 1 shows the typical cyclic voltammograms (CVs) obtained at a Au electrode in O2saturated DMSO solution containing 0.1 M TEAP. A welldefined cathodic peak at -0.78 V versus Ag|AgCl|NaCl (sat.) for the reduction of O2 to O2•- and an anodic peak at -0.64 V for the reoxidation of O2•- to O2 were observed (Figure 1a), that is, this CV response corresponds to a quasi-reversible redox reaction of the O2/O2•- couple.3-14 When the CV was measured in the potential range of 0.3 f -2.2 f 0.3 V, a second cathodic peak at -2.0 V was observed for further one-electron reduction of O2•- to O22-. In this case, the anodic peak current at -0.64 V was found to be smaller than that obtained when the potential was scanned in the range of 0.3 f -1.2 f 0.3 V, and at the same time, a new anodic peak was observed at approximately -0.15 V. These cyclic voltammetric observations agree very well with those reported previously,3-6,12,13 and the obtained anodic peak at -0.15 V was considered to correspond to the reoxidation of peroxide species to O2.7,9,10,14,24,39 ORR in EMIBF4 Solution. In Figure 2 are shown the CVs obtained at a Au electrode in N2- and O2-saturated EMIBF4 solutions. No remarkable peak was found in the CV measured in N2-saturated EMIBF4, indicating that EMIBF4 is electrochemically stable within the examined potential range of 0.5 to -2.1 V. In the O2-saturated EMIBF4 solution, when the CV was measured in the potential range of 0.4 f -1.1 f 0.4 V, a couple of anodic and cathodic peaks were observed at -0.47 and -0.71 V, respectively, while the shape of the CV (Figure 2b) was rather distorted compared to that observed in DMSO solution (Figure 1a). Similar distortion of the redox peak of ORR has been previously observed at the conventional electrodes15-18,23 and microelectrodes19-21 in ImRTILs. However, when the CV was measured in the potential range of 0.6 f -1.6 f 0.6 V, in addition to the cathodic peak at -0.71 V, an irreversible cathodic peak was found at -1.35 V (Figure 2c). The ratio of the obtained first and second cathodic peak currents was found be almost unity (ca. 1.1) (i.e., the number of electrons involved in the first and second steps of the ORR may be assumed to be equal). On the other hand, the anodic

Roles of Ion Pairing on Electroreduction of Dioxygen

Figure 2. CVs obtained at a Au electrode (φ ) 1.6 mm) in (a) N2saturated and (b and c) O2-saturated EMIBF4 solutions. Potential scan rate: 0.1 V s-1.

peak current at -0.47 V decreased when the electrode potential was scanned in the range of 0.6 f -1.6 f 0.6 V, and two ill-defined anodic peaks were found at potentials more positive than -0.2 V (discussed later). From the analogy to the redox reaction of the O2/O2•- couple in DMSO solution, the observed first redox peaks in ImRTILs were ascribed to the O2/O2•- redox couple.15-26 By comparing the CVs shown in Figures 1 and 2, we can see that the potentials of the first and second cathodic peaks in EMIBF4 are by 0.08 and 0.65 V more positive than those obtained in DMSO solution. This fact may suggest that the ORR observed in EMIBF4 and DMSO solutions would not be entirely comparable with each other. Redox Reactions of BQ in DMSO and EMIBF4 Solutions. For a clear understanding of the difference in the ORR observed in DMSO and EMIBF4 solutions, the cyclic voltammetric redox behavior of BQ was examined in both solutions, and the obtained results are shown in Figure 3. In a DMSO solution containing 0.1 M TEAP, two redox responses of the BQ/BQ•(anion radical of BQ) and BQ•-/BQ2- (dianion of BQ) couples were observed with midpoint potentials (E0′ value; the average of cathodic and anodic peak potentials) of -0.32 (E0′BQ/BQ•-) and -1.12 V (E0′BQ•-/BQ2-) (Figure 3a).3,40-43 It is noted that the magnitude of the peak current of the BQ•-/BQ2- redox couple is slightly smaller than that of the BQ/BQ•- redox couple, as generally observed in aprotic solutions.3,41-43 The decrease in the diffusion coefficient of BQ•- (compared with that of BQ) due to solvation or hydrogen bonding and dimerization of BQ•and the complexation between BQ2- and BQ to form [BQBQ]2have been proposed as the cause for the smaller peak current of the BQ•-/BQ2- redox couple,39-41 but the latter has been proven to be the most rational.41 Similarly to the DMSO solution, the CV measured in the EMIBF4 solution was found to possess two couples of redox peaks with E0′ of -0.37 and -0.71 V (Figure 3b) in which the peak current of the second redox couple is essentially smaller than that of the first redox couple. The differences between anodic and cathodic peak potentials (∆Ep) for the BQ/BQ•- and BQ•-/BQ2- redox couples in EMIBF4 are almost the same [about 70 mV, close to the theoretical value (59 mV) for a oneelectron reversible redox reaction40] and comparable with those obtained for these couples in DMSO solution. In analogy with the redox behavior of BQ in DMSO solution (Figure 3a), the observed two couples of redox peaks in EMIBF4 may be attributed to the BQ/BQ•- and BQ•-/BQ2- redox couples

J. Phys. Chem. C, Vol. 112, No. 4, 2008 1271 (discussed below). Meanwhile, the smaller peak currents observed in EMIBF4 than those in DMSO solution can be reasonably ascribed to the restricted diffusion of the redox species since the viscosity of EMIBF4 is larger (37.10 cP) than that (1.99 cP) of DMSO. Although the actual reason is not clear yet, the observed smaller peak current of the BQ•-/BQ2- redox couple than that of the BQ/BQ•- redox couple in EMIBF4 may suggest the involvement of any chemical step in the overall redox process as observed in the case of DMSO (described above). In the electrochemical measurements, the polarity of the solvent, the nature of the supporting electrolyte, and the presence of acidic additives significantly change the E0′ value of a redox couple, reflecting nonspecific solvation energies, ion pairing, and protonation equilibria, respectively.41 To understand the reason for the relatively positive value of E0′BQ•-/BQ2- in EMIBF4 compared with that in DMSO solution, the redox reaction of BQ in DMSO solutions was studied by the addition of EMIBF4, and the obtained results are shown in Figure 3. Although E0′BQ/BQ•- remained almost constant in the absence and presence of EMIBF4, a positive shift of E0′BQ•-/BQ2- was found with increasing of the concentration of EMIBF4 (inset in Figure 3), that is, E0′BQ•-/BQ2- becomes closer to that obtained in the pure EMIBF4, as its concentration is increased in DMSO solution. Here, it should be mentioned that, due to complexation between the electrogenerated anion radical and lithium ion, a positive shift of the first redox peak potential of the redox reaction of 9,10-anthraquinone has been found in acetonitrile (ACN) solution.34 As the polarity of the solvent [or the solvent acceptor number (AN)] is increased, the E0′ value of a redox couple generally shifts to positive potential.12 For instance, the E0′ values of the two redox couples of p-DNB (p-DNB/p-DNB•- and p-DNB•-/ p-DNB2-) have been observed at -0.49 and -0.76 V in DMSO solution and at -0.53 and -0.84 V in DMF solution (data are not shown). These observations may be reasonable because the AN value of DMSO is greater (19.3) than that (16.0) of DMF.2 On the contrary, EMIBF4 is ionic in character, and its AN value is larger (26.3)44 than those of DMSO and DMF. Moreover, only E0′BQ•-/BQ2- in EMIBF4 shifts to the positive potential in the presence of EMIBF4 (inset in Figure 3), while E0′BQ/BQ•remains unchanged, or only E0′BQ•-/BQ2- in pure EMIBF4 is more positive compared to that in DMSO solution (Figure 3). Thus, these results cannot be explained by considering simply the AN (polarity) effect because the higher AN value of EMIBF4 should influence both E0′BQ/BQ•- and E0′BQ•-/BQ2-. Gupta et al.41 have extensively studied the effect of protonic additives (e.g., acetic acid derivatives, ethanol, H2O, and t-butanol) having different pKa values on E0′BQ•-/BQ2- in benzonitrile solution, and they have observed a similar positive shift of E0′BQ•-/BQ2-. It has been reported that the additives with low and high pKa values undergo protonation equilibria and hydrogen bonding with electrogenerated BQ2-, respectively. In the present case, the value of pKa of EMIBF4 (pKa value of imidazolium salt is 24 in DMSO solution)45 would be higher than that (12 in benzonitrile solution) of BQ2-.41,44 Hence, the electrogenerated BQ2- should not favorably undergo the protonation reaction by the EMI+ cation. Recently, similarly to the case of BQ (inset in Figure 3), Fry28 has observed a regular positive shift of the E0′ of the DNB•-/ DNB2- redox couple by the addition of BMIBF4 in ACN solution, although the value of E0′ of DNB/DNB•- remains almost constant. It is also known that the species resulting from a second electron-transfer step (e.g., BQ2-) is more prone to form an ion pair than the first one (e.g., BQ•-).31 Therefore,

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Islam and Ohsaka

Figure 3. CVs obtained for the redox reaction of 1 mM BQ at a GC electrode (φ ) 1 mm) in N2-saturated DMSO solution containing 0.1 M TEAP (a) and EMIBF4 solution (b). The inset shows the CVs obtained for 1 mM BQ at the Au electrode (φ ) 1.6 mm) in N2-saturated DMSO solutions containing 0.1 M TEAP in the absence (a) and presence of (b) 0.07 and (c) 0.95 M EMIBF4. Potential scan rate: 0.1 V s-1.

the observed positive shift of E0′BQ•-/BQ2- would be ascribed to the ion pairing between the EMI+ cation and the BQ2-, that is, the ion-pairing phenomenon in EMIBF4 is more predominant than the solvation and protonation processes. Here, the mode of ion pairing would be assumed to be accomplished by the π-π stacking (complexation) of the EMI+ cation and the aromatic π system of BQ2- since the π system of aromatic anions can approach closer to the planer moiety of the EMI+ cation.28 In addition, the relative stability of the BQ2- species in EMIBF4 compared to that in DMSO solution would be understood in terms of the free-energy change for the comproportionation reaction (-∆Gcomp) between BQ and BQ2-. In this case, BQ2- was found to be more stable in EMIBF4 than in DMSO solution since the -∆Gcomp value in EMIBF4 is smaller (32.8 kJ mol-1) than that (77.2 kJ mol-1) in DMSO solution, as determined by using the obtained CV results and the equations similar to eqs 6-8 (see below). Effects of EMIBF4 on the ORR in DMSO Solutions. Figure 4 represents the CVs obtained at a Au electrode in O2-saturated DMSO solutions containing 0.1 M TEAP in the absence and presence of EMIBF4. Remarkably, in the presence of EMIBF4, a new cathodic peak was found at -1.4 V, and no positive shift of this peak with increasing of the concentration of EMIBF4 in DMSO solution was observed. Moreover, with increasing of the concentration of EMIBF4, (i) the cathodic peak observed at -0.78 V regularly shifted to more positive potential, (ii) the peak current at -0.78 V initially increased (Figure 4b) and gradually decreased (Figure 4c,d), (iii) the cathodic peak current at -1.4 V regularly increased46 by the cost of the cathodic peak at approximately -1.8 V (Figure 4b-d), (iv) the anodic peak current at -0.65 V decreased and slightly shifted to the positive potential, and (v) the shape of the anodic wave at approximately -0.15 V deformed, and the peak potential gradually shifted to a more positive potential. A similar new cathodic peak in the potential range between the first and second cathodic peaks

Figure 4. CVs obtained at a Au electrode in O2-saturated DMSO solutions containing 0.1 M TEAP in the absence (a) and presence of 0.09 (b), 0.34 (c), and 0.67 M (d) EMIBF4. The inset shows the CVs measured in the same solutions in the potential range for the redox reaction of the O2/O2•- couple. Potential scan rate: 0.1 V s-1.

of the ORR in aprotic solutions (e.g., DMSO, dimethylacetamide, hexamethylphosphoramide, and propylene carbonate solutions containing tetra-n-alkylammonium salt as the supporting electrolyte) by the addition of metal ions (e.g., sodium, potassium, rubidium, and cesium ions) has been observed.5,6,10,11,14 However, in analogy with the case of the BQ•-/BQ2- redox couple, the observed new cathodic peak (at -1.4 V) that does not shift to the positive potential with increasing of the ionpairing agent (EMIBF4) in DMSO solution may not be simply attributed to the ion-pairing process. To clarify the reason of such an observation, the steady-state voltammetric measurement was carried out.

Roles of Ion Pairing on Electroreduction of Dioxygen

J. Phys. Chem. C, Vol. 112, No. 4, 2008 1273 Ka

O2•- + EMI+ y\z O2•-...EMI+ E0′(O2/O2•-)ip ) E0′(O2/O2•-)sol + RT/F ln(1 + Ka[EMI+]) ∆G0′a ) -RT ln Ka

(1) (2) (3)

Now, by rearranging eq 2, eq 4 can be derived

e∆E ′F/RT - 1 ) Ka[EMI +] 0

Figure 5. Steady-state voltammograms obtained for the reduction of O2 at the Au microelectrode (φ ) 10 µm) in DMSO solutions containing 0.1 M TEAP in the absence (a) and presence of (b) 0.007, (c) 0.032, (d) 0.096, and (e and e′) 0.13 M EMIBF4. The voltammograms (a-e) and (e′) were measured in O2- and N2-saturated solutions, respectively. Potential scan rate: 0.01 V s-1.

Figure 5 shows the typical steady-state voltammograms measured for the ORR at a Au microdisk electrode in DMSO solutions in the absence and presence of EMIBF4. Irrespective of the absence and presence of EMIBF4, two-step ORR necessarily took place, wherein the ratio of the limiting current was estimated to be almost unity. In the presence of EMIBF4, unlike the CV measured at the conventional-sized electrode (Figure 4), no new (third) cathodic peak was registered, and the positive shift of half-wave potential (E1/2) of both waves with increasing of the concentration of EMIBF4 was observed, for example, for a 10-fold increase of the concentration of EMIBF4, a positive shift of E1/2 of the second cathodic wave was 160 mV, that is, about three times that (57 mV, see below) of the midpoint potential of the O2/O2•- redox couple (E0′(O2/O2•)). From the cyclic and steady-state voltammetric observations, the new cathodic peak at -1.4 V (Figure 4) may be ascribed to the reduction of the ion-paired species (i.e., EMI+...O2•-, discussed below). Katayama et al.23 have ascribed the second cathodic peak observed during the ORR in ImRTILs to the reduction of some unknown complex of O2•- and the cation of RTIL. Previously, we have reported that a quasi-reversible twoelectron ORR occurs at HMDE in EMIBF4 via an abstraction of the proton at 2 position of the EMI+ ring.15 However, the observation of the positive shift of the second reduction wave of the ORR at the microelectrode is comparable with that of the BQ•-/BQ2- redox couple resulting from the ion-pairing phenomenon. On the other hand, since the magnitude of positive shift of E1/2 of the second cathodic wave is greater by three times than that of E0′(O2/O2•-), the extent of ion pairing of EMI+ and O22- is considered to be, as expected,31 significantly stronger than that of EMI+ and O2•-. In the present case, the ion pairing between O2•- (and O22-) with the EMI+ cation may possibly take place via both the electrostatic interaction between the positively charged EMI+ and the negatively charged O2•- (and O22-) and the hydrogen bonding of O2•- (and O22-) with the (acidic) hydrogen atom22,47,48 at the 2 position of the EMI+ ring. Extent of Ion Pairing of the EMI+ Cation with Electrogenerated O2•- and O22-. The extent of ion pairing between the EMI+ cation and electrogenerated O2•- in DMSO solution in the presence of EMIBF4 could be evaluated by considering the association constant (Ka) and the free-energy change (∆G0′a) for eq 1 as follows31,36,38,40,49

(4)

where E0′(O2/O2•-)sol and E0′(O2/O2•-)ip are the midpoint potentials of the O2/O2•- redox couple in the absence and presence of EMIBF4, respectively. ∆E0′ is the difference between E0′(O2/O2•0 + )ip and E ′(O2/O2•-)ip, and [EMI ] is the concentration of EMIBF4 in DMSO solution. According to eq 4, a plot of e(∆E0′F/RT) versus [EMI+] should be a straight line passing through the origin, with a slope of Ka. The obtained plot is shown in Figure 6, in which the values of ∆E0′ were estimated from the CVs measured by varying [EMI+] (inset in Figure 4). The expected straight line (r ) 0.9939) was obtained, though some deviation from this was observed at lower [EMI+]. From the slope, the value of Ka was determined to be 13.5 M-1 (and ∆G0′a ) -6.4 kJ mol-1) at 25 °C. For comparison with the above-mentioned ion pairing between EMI+ and O2•-, the effect of H2O that is known to undergo association with O2•- species on the ORR in DMSO solution was also examined. In this case, the similar positive shift of E0′O2/O2•- with increasing of the concentration of H2O (up to 0.6 M) in DMSO solution was obtained (data are not shown), and the association constant (Kaw) of the complexation between H2O and the electrogenerated O2•- was estimated to be 15.7 M-1. Singh et al.49 have examined the degree of association of O2•- with different alcohols and H2O in ACN solutions. Generally, the association constant between O2•- and alcohol (30 M-1) has been observed to be larger than that (20 M-1) between O2•- and H2O. In the present study, the observation of a smaller Kaw value in DMSO solution compared to that (20 M-1) found in ACN solution49 may be reasonably accepted since the solvation of O2•- is more favorable in DMSO solution than that in ACN solution, resulting in a weaker ionpairing phenomenon. In addition, the almost equal values of Ka for O2•-...EMI+ (13.5 M-1) and O2•-...H2O (15.7 M-1) indicate that the ion pairing of O2•- species with both the EMI+ cation and H2O in DMSO may be comparable. As mentioned above, in DMSO solution containing EMIBF4, the solvated O2•- species (O2•-sol) coexist in equilibrium with the ion-pared species with EMI+ (O2•-ip). According to eq 2, when Ka[EMI+] . 1, eq 2 can be reduced to eq 5, whereas when Ka[EMI+] , 1, E0′(O2/O2•-)ip ≈ E0′(O2/O2•-)sol

E0′(O2/O2•-)ip ) E0′(O2/O2•-)sol + RT/F ln(Ka[EMI+])

(5)

A plot of E0′(O2/O2•-)ip versus ln(Ka[EMI+]) is shown in Figure 7. As expected, the linear plot was obtained when [EMI+] was larger than ∼0.09 M. This concentration of EMIBF4 may be considered as the critical concentration for a complete ion-pair formation of EMI+ with O2•- species in DMSO solution. Thus, all of the O2•- species formed during the ORR in EMIBF4 can be considered to exist as O2•-ip actually because the concentration of EMI+ in EMIBF4 is much higher (ca. 6.4 M) than the estimated critical concentration (0.09 M).18 On the linear part of the plot observed at [EMI+] > 0.09 M (Figure 7), a 57 mV positive shift of E0′(O2/O2•-)sol with a 10-fold increase of [EMI+]

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Islam and Ohsaka solutions, respectively. Thus, the value of -∆Gcomp in DMSO solution is about twice that obtained in EMIBF4 solution, suggesting that the comproportionation reaction (eq 6) is much more favorable in DMSO solution. In other words, the EMI+ cation effectively stabilizes the electrogenerated O22-.31 It is also likely, though yet clarified, that the EMI+ cation reacts with the electrogenerated O22- to form some reactive oxygen species (e.g., HO2-, HO2•).15,50 Thus, the anodic peak(s) observed at potentials more positive than -0.3 V (Figures 1b and 2b) is (are) possibly attributed to the oxidations of such reactive oxygen species. Conclusions

Figure 6. A plot of (e∆E0′F/RT - 1) versus [EMI+]. The data were obtained from the CVs measured in the absence and presence of EMIBF4, as shown in Figure 4.

Figure 7. A plot of E0′O2/O2•- versus (Ka[EMI+]). The values of Ka and E0′O2/O2•- were obtained from the plots and CVs such as those shown in Figure 6 and the inset in Figure 4, respectively.

in DMSO solution was observed. This slope is in good agreement with the theoretical expectation for a moderate ionpairing process (i.e., 59 mV at 25 °C).31 Thus, the extent of the ion pairing between the O2•- species and the EMI+ cation may be less compared to the complexation of metal cations (e.g., lithium, sodium, and potassium ions) with the O2•- species.10,11,14 The extent of thermodynamic stability of the ion-paired O22with the EMI+ cation was also evaluated by comparing the values of ∆Gcomp and Kcomp for the comproportionation reaction between O2 and O22- (eq 6) in EMIBF4 and DMSO solutions, which were determined using the following expressions (eqs 6-8)36,40 Kcomp

O2 + O22- y\z 2 O2•-

(6)

∆Epc ) Epc1 - Epc2 ) (RT/nF) ln Kcomp

(7)

-∆Gcomp ) n∆EpcF

(8)

where Epc1 and Epc2 are the potentials of the first and second reduction peaks, respectively (in which the cathodic peak potential40 was conveniently used instead of the formal potential because the latter was not determined experimentally). From the CVs shown in Figures 1 and 2, the values of ∆Epc were calculated to be 0.64 and 1.22 V in EMIBF4 and DMSO solutions, respectively. According to eqs 7 and 8, the values of -∆Gcomp were estimated to be 117.7 (Kcomp ) 4.3 × 1020) and 61.7 kJ mol-1 (Kcomp ) 6.7 × 1010) in DMSO and EMIBF4

The effects of ion pairing on the ORR in EMIBF4 RTIL were investigated. Due to the ion pairing of O2•- (and O22-) with the EMI+ cation, the cyclic voltammetric redox peaks of the ORR in EMIBF4 were observed at more positive potentials compared to those obtained in DMSO solution containing 0.1 M TEAP. The second reduction peak of the ORR in EMIBF4 suffers from a greater ion-pairing effect, compared to the first reduction peak, as indicated by a larger positive shift of this peak potential. The association constant and the change of free energy of the ion pairing between the electrogenerated O2•- and EMI+ cation in DMSO solution were estimated. It was observed that in DMSO solution, the degree of association of O2•- with the EMI+ cation is very comparable to that of O2•- with H2O. Although the degree of ion pairing of the EMI+ cation with the electrogenerated O2•- was examined in DMSO solution, the present observation may imply the occurrence of a strong ion pairing in pure RTILs. Thus, it can be concluded that O2•- may often undergo complexation with the imidazolium cation of RTILs, which has been recently clarified with UV-visible and voltammetric measurements and ab initio molecular orbital calculation,26 resulting in a significantly smaller DO2•- than DO2. Acknowledgment. The present work was financially supported by Grant-in-Aids for Scientific Research on Priority Areas (No. 417) and Scientific Research (A) (No. 19206079) to T.O. and from the Ministry of Education, Culture, Sports, Science and Technology, Japan (Monbu-Kagakusho) and the VBL program at Tokyo Tech. M.M.I. gratefully acknowledges the Government of Japan for a Monbu-Kagakusho Scholarship and the Interdisciplinary Graduate School of Science and Engineering of Tokyo Tech for the Postdoctoral Fellowship. References and Notes (1) El-Deab, M. S.; Okajima, T.; Ohsaka, T. J. Electrochem. Soc. 2003, 150, A851. (2) Saha, M. S.; Che, Y.; Okajima, T.; Kiguchi, T.; Nakamura, Y.; Tokuda, K.; Ohsaka, T. J. Electroanal. Chem. 2001, 496, 61. (3) Islam, M. M.; Okajima, T.; Ohsaka, T. J. Phys. Chem. B 2004, 108, 19425. (4) Islam, M. M.; Saha, M. S.; Okajima, T.; Ohsaka. T. J. Electroanal. Chem. 2005, 577, 145. (5) Johnson, E. L.; Pool, K. H.; Hamm, R. E. Anal. Chem. 1966, 38, 183. (6) Johnson, E. L.; Pool, K. H.; Hamm, R. E. Anal. Chem. 1967, 39, 888. (7) Maricle, D. L.; Hodgson, H. G. Anal. Chem. 1965, 37, 1562. (8) Peover, M. E.; White, B. S. Electrochim. Acta 1966, 11, 1061. (9) Goolsby, A. D.; Sawyer, D. T. Anal. Chem. 1968, 40, 83. (10) Sawyer, D. T.; Chlerlcato, G., Jr.; Angells, C. T.; Nannl, E. J., Jr.; Tsuchiya, T. Anal. Chem. 1982, 54, 1720. (11) Sawyer, D. T.; Valentine, J. S. Acc. Chem. Res. 1981, 14, 393. (12) Vasudenvan, D.; Wendt, H. J. Electroanal. Chem. 1995, 392, 69. (13) Ortiz, M. E.; N.-Vergara, L. J.; Squella, J. A. J. Electroanal. Chem. 2003, 549, 157. (14) Fujinaga, T.; Sakura, S. Bull. Chem. Soc. Jpn. 1974, 47, 2781.

Roles of Ion Pairing on Electroreduction of Dioxygen (15) Islam, M. M.; Ferdousi, B. N.; Okajima, T.; Ohsaka. T. Electrochem. Commun. 2005, 7, 789. (16) AlNashef, I. M.; Leonard, M. L.; Kittle, M. C.; Mattews, M. A.; Weidner, J. W. Electrochem. Solid-State Lett. 2001, 4, D 16. (17) AlNashef, I. M.; Leonard, M. L.; Kittle, M. C.; Mattews, M. A.; Weidner, J. W. Ind. Eng. Chem. Res. 2002, 482, 87. (18) Zhang, D.; Okajima, T.; Matsumoto, F.; Ohsaka, T. J. Electrochem. Soc. 2004, 151, D31. (19) Buzzeo, M. C.; Klymenko, O. V.; Wadhawan, J. D.; Hardacre, C.; Seddon, K. R.; Compton, R. G. J. Phys. Chem. A 2003, 107, 8872. (20) Evans, R. G.; Klymenko, O. V.; Saddoughi, S. A.; Hardacre, C.; Compton, R. G. J. Phys. Chem. B 2004, 108, 7878. (21) Buzzeo, M. C.; Klymenko, O. V.; Wadhawan, J. D.; Hardacre, C.; Seddon, K. R.; Compton, R. G. J. Phys. Chem. B 2004, 108, 3947. (22) Noda, A.; Susan, M. A. B. H.; Kudo, K.; Mitsushima, S.; Hayamizu, K.; Watanabe, M. J. Phys. Chem. B 2003, 107, 4024. (23) Katayama, Y.; Onodera, H.; Yamagata, M.; Miura, T. J. Electrochem. Soc. 2004, 151, A59. (24) Katayama, Y.; Sekiguchi, K.; Yamagata, M.; Miura, T. J. Electrochem. Soc. 2005, 152, E247. (25) Carter, M. T.; Hussey, C. L.; Strubinger, S. K. D.; Osteryoung, R. A. Inorg. Chem. 1991, 30, 1149. (26) Islam, M. M.; Imase, T.; Okajima, T.; Takahashi, M.; Niikura, Y.; Kawashima, N.; Ohsaka, T. J. Phys. Chem. A, submitted. (27) Schoder, U.; Wadhawan, J. D.; Compton, R. G.; Marken, F.; Suarez, P. A. Z.; Consorti, C. S.; De Souza, R. F.; Dupont, J. New J. Chem. 2000, 24, 1009. (28) Fry, A. J. J. Electroanal. Chem. 2003, 546, 35. (29) Lagrost, C.; Preda, L.; Volanschi, E.; Hapiot, P. J. Electroanal. Chem. 2005, 585, 1. (30) Brooks, C. A.; Doherty, A. P. J. Phys. Chem. B 2005, 109, 6276. (31) Saveant, J.-M. J. Phys. Chem. B 2001, 105, 8995. (32) Fry, A. J. Electrochem. Commun. 2005, 7, 602.

J. Phys. Chem. C, Vol. 112, No. 4, 2008 1275 (33) Watkins, J. J.; White, H. S. J. Electroanal. Chem. 2005, 582, 57. (34) Wain, A. J.; Wildgoose, G. G.; Heald, C. G. R.; Jiang, L.; Jones, T. G. J.; Compton, R. G. J. Phys. Chem. B 2005, 109, 3971. (35) Barriere, F.; Camire, N.; Geiger, W. E.; Mueller-Westerhoff, U. T.; Sanders, R. J. Am. Chem. Soc. 2002, 124, 7262. (36) Fry, A. J.; Britton, W. E. Topics in Organic Electrochemistry; Plenum Press: New York, 1986. (37) Fry, A. J. Electroanalysis 2006, 18, 391. (38) Nakanishi, T.; Murakami, H.; Sagara, T.; Nakashima, N. J. Phys. Chem. B 1999, 103, 304. (39) Andrieux, C. P.; Hapiot, P.; Saveant, J.-M. J. Am. Chem. Soc. 1987, 109, 3768. (40) Bard, A. J.; Faulkner, L. R. Electrochemical MethodssFundamentals and Applications, 2nd ed.; Wiley: New York, 2001. (41) Gupta, N.; Linschitz, H. J. Am. Chem. Soc. 1997, 119, 6384. (42) Marcus, M. F.; Hawley, M. D. Biochim. Biophys. Acta 1970, 222, 163. (43) Eggins, B. R. Chem. Commun. 1969, 1267. (44) Kimura, Y.; Fukuda, M.; Fujisawa, T.; Terazima, M. Chem. Lett. 2005, 34, 338. (45) Aggarwal, V. K.; Emme, I.; Mereu, A. Chem. Commun. 2002, 1612. (46) At first glance, the increase of the current of the new cathodic peak at -1.4 V cannot be seen, but this can be actually traced by considering the decrease of the height of the shoulder (e.g., -1.0 V), which is the baseline for measuring the peak current at -1.4 V, of the CVs after the first cathodic peak with respect to the charging current level (i.e., the original base line). (47) Lynden-Bell, R. M.; Atamas, N. A.; Vasilyuk, A.; Hanke, C. G. Mol. Phys. 2002, 100, 3225. (48) Morrow, T. I.; Maginn, E. J. J. Phys. Chem. B 2002, 106, 12807. (49) Singh, P. S.; Evans, D. H. J. Phys. Chem. B 2006, 110, 637. (50) Mery, D.; Aranzaes, J. R.; Astruc, D. J. Am. Chem. Soc. 2006, 128, 5602.