Organic Solvent Mixtures

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J. Phys. Chem. B 2009, 113, 2019–2023

2019

Diffusion of Molecules in Ionic Liquids/Organic Solvent Mixtures. Example of the Reversible Reduction of O2 to Superoxide Dodzi Zigah, Aifang Wang, Corinne Lagrost, and Philippe Hapiot* UniVersite´ de Rennes 1, Sciences Chimiques de Rennes (Equipe MaCSE), CNRS, UMR N° 6226, Campus de Beaulieu, Bat 10C, 35042 Rennes Cedex, France ReceiVed: October 28, 2008; ReVised Manuscript ReceiVed: December 7, 2008

Transport properties of molecules dissolved in room-temperature ionic liquids are highly sensitive to the charge carried by the molecule because of complex ion-ion interactions that could be tuned by addition of a cosolvent. In this connection, the one-electron reduction of oxygen was used as a probe system for studying the effects of the addition of a cosolvent such as dimethylformamide (DMF) into a pure ionic liquid (triethylbutylammonium bis(trifluoromethylsulfonyl)imide) ([Et3BuN][NTf2]) on the diffusion of charged species versus neutral species. Experimental data about the diffusion coefficients of O2 (DO2) and O2•- (DO2•-) and their ratios (γ ) DO2•-/DO2) were extracted using scanning electrochemical microscopy (SECM) in transient mode as a function of the DMF concentration. The ratio γ and both of the diffusion coefficients DO2•- and DO2 were found to increase exponentially with the DMF volume fractions following the same general tendency described for the viscosity. However, DO2•- varies on a much larger range than DO2 (around 1000 times more), and O2•- retains an almost “pure ionic” behavior for higher DMF fractions. All of these results support the occurrence of a sharp transformation in the bonding character of the RTIL cation upon addition of a molecular solvent, as predicted in recent theoretical simulations. 1. Introduction Due to their interesting properties such as chemical and thermal stabilities, nonflammability, nonvolatility, intrinsic conductivity, and so forth, room-temperature ionic liquids (RTIL) are the subject of intense research in many fields of chemistry.1 However, the high viscosity and the resulting slow mass transport of substrates could be a problem for practical applications. Either heating the solutions or adding certain amounts of conventional solvents such as dimethylformamide (DMF), ethanol, or H2O or, even better, by combining these two means have been proposed to overcome these difficulties and to tune the properties of the media.2,3 Several studies have been performed concerning the physical chemistry properties, especially about the evolutions of viscosity, ionic conductivity, and diffusion coefficients as a function of the quantity of added cosolvents in the RTIL.3 In this connection, electrochemical techniques are convenient methods allowing detailed analyses of electron transfer, masstransport processes, or, more generally, the chemical reactivities of the electrogenerated intermediates in RTILs, and the redox behavior of many species dissolved has been examined in such media.4 Specific interactions between substrates and the RTIL’s ions have been detected5 through their electrochemical properties, especially concerning mass transports.4 Indeed, besides the lower diffusion coefficients in RTILs than those in molecular solvents, these values could vary considerably between a neutral molecule and the corresponding species produced upon the electron transfer. This phenomenon, characterized by the ratio γ ) Dred/Dox, leads to γ values that could be very different from unity.6,7 The reversible reduction of oxygen to the superoxide O2•- that has been thoroughly studied in RTILs,6-8 and in organic solvents,9 is an impressive example of such effects. * To whom correspondence should be addressed. E-mail: philippe.hapiot@ univ-rennes1.fr.

Indeed, the γ ratio for O2•- and O2 was found to be around 1/30 in common RTILs, as in quaternary ammonium bis((trifluoromethyl)sulfonyl)imide-based ionic liquids.6,7a Just to illustrate the meaning of such unusual γ values, by considering a simple Stokes-Einstein relationship, this would mean that the “solvated” O2•- in the RTIL is more than 10000 times heavier than O2 in the same RTIL.7a However, the phenomenon is not limited to the oxygen reduction; several redox couples display similar effects, but to a lower extent.4 Recently, the variations of γ were investigated for the reduction of a series of nitro derivatives. It was found that γ is directly related to the localization and accessibility of the charge on the radical ion.7b All of these effects were rationalized in terms of ions interactions between the radical anion and the RTIL cations, and γ was found to be a convenient and rapid way for appreciating the strength of ion interactions in RTILs.7b Concerning the transport properties of substrates in morecomplicated systems of RTIL/molecular solvent mixtures, most of the studies have focused on the variations of the viscosity, diffusion coefficient, or conductivity upon the addition of the molecular solvent (see, for example, ref 3). However, the variations of γ in these mixtures remain almost unknown, despite the strong influence of the molecular charge on the transport properties and the interest of these media for applications as mentioned above.10 In this work, we have investigated the influence of the presence of an organic cosolvent on the diffusion coefficients of the species in the solution. Using an organic solvent provides a simple way to vary the composition of the mixture, without the complexity of RTIL/water solutions, because several RTILs are highly soluble in organic solvents.11 More specifically, we focused our study on the variations of the γ ratio that could be seen as a consequence of the ionic interactions considering that the effects of the viscosity should remain similar for the charged and uncharged species. The redox couple O2/O2•- was chosen

10.1021/jp8095314 CCC: $40.75  2009 American Chemical Society Published on Web 01/27/2009

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SCHEME 1: Ions Used for Preparing the Room-Temperature Ionic Liquids, [Et3BuN] and [NTf2]

as a test system because of its intrinsic interest in many applications but also because the small size of the molecule makes this couple an efficient probe for observing the ions’ interactions.6d The [Et3BuN][NTf2] (see Scheme 1 for its structure) was used as a typical and common RTIL, and dimethylformamide (DMF) was used as an example of common polar organic aprotic solvent.3a Moreover, numerous data concerning the reduction of O2 in the two pure media are available in the literature, allowing direct comparisons with published data.8,9,12 Concerning the measurements, the determination of γ was based on the use of the scanning electrochemical method (SECM)13 in transient feedback mode, which was applied to the measurement of γ and of the corresponding diffusion coefficients.7a,14 Interest of the method is that the measurements could be performed for different diffusion lengths, allowing a self-consistent determination of the data. Indeed, influences of other transport mechanisms, such as “natural convection”, are sometimes difficult to detect in steady-state methods but are more likely to be a problem in RTILs15 because of the low diffusion rates.16 In order to facilitate the experiments, the working temperature was settled as 36 °C, so that the RTIL was less viscous and thus electrochemical currents were higher.3,17 By doing this, the feasibility of moderate heating and (or) addition of a cosolvent as the means for obtaining conditions well adapted to electrochemical studies in ionic liquids could also be tested. 2. Experimental Section Chemicals. The RTIL, [Et3BuN][NTf2], was prepared from aqueous [Li][NTf2] (Solvionic) and [Et3BuN]Br according to standard procedures.18 The sample was purified by repeated washing with H2O and filtered over neutral alumina and silica. Prior to each experiment, vacuum pumping carefully dried RTILs overnight, and the amount of residual water was measured with Karl Fischer titration (Karl Fischer 652 Metrohm). The amount of water measured in our samples ranged from 100 to 200 ppm. Dimethyformamide (DMF) was from Fluka, assay g99.5% (GC) quality, containing less than 0.005% of water. Electrodes. The working electrode was a gold disk ultramicroelectrode (UME) with a radius of 5 µm. UMEs were fabricated by sealing gold wires (Goodfellow) into glass capillaries, the ends of which were subsequently polished to expose a disk-shaped surface with an outer radius of ∼50 µm (typically RG ≈ 10). Prior to use, the UMEs were polished using decreasing sizes of diamond and alumina pastes. The UMEs were characterized by cyclic voltammetry and by typical approach curves recorded on conducting and insulating surfaces. The reference electrode was a quasi-reference electrode made with a Pt wire covered by semioxidized polypyrrole and prepared according to the published procedure.19 A platinum wire (0.25 mm radius) was used as the counter electrode. SECM Measurements. Approach curves and current-time experiments were performed using a commercial SECM instrument with closed-loop piezoelectric motors, CHI 900B (CH Instruments, Austin, TX). The SECM instrument was placed on a vibration isolation platform in a Faraday cage, inside of

Figure 1. Cyclic voltammetry of saturated O2 in (a) pure [Et3BuN][NTf2] and (b) DMF with 0.1 mol · L-1 [Et3BuN][NTf2] as the supporting electrolyte on a 5 µm radius gold microelectrode. Scan rates: ν ) 0.1 V/s. T ) 36 °C.

which the temperature was maintained at 36 ( 1 °C. The electrochemical cell was kept inside of this environment for at least 30 min before the experiments to allow the equilibration of the temperature in the solution (T ) 36 ( 1 °C). A fourelectrode setup was employed with a gold UME as the tip and a Pt or glassy carbon large surface as the substrate. Data were recorded in O2-saturated solutions containing various amounts of DMF in the ionic liquid [Et3BuN][NTf2] by holding the potential of the UME at the reduction plateau of oxygen (typically 200-300 mV more negative than the half-wave potential) and the substrate potential much more positive than the half-wave potential of O2/O2•-. In such conditions, electron transfer could be considered as infinitely fast both at the UME and at the substrate surfaces. In such conditions, the electrochemical reduction of O2 and the electrochemical oxidation of the superoxide are only limited by the diffusion of the species and not by the kinetics of the electron transfers occurring at the tip and substrate electrodes7,14 that could be associated with the formation or the breaking of ions pairs. Pure oxygen was bubbled through the electrochemical cell for at least 15 min before starting the measurements. In a typical determination of γ, the tip was first located above the substrate at a certain distance that was adjusted by recording the approach curve.7a At this position, a curve of current-time was performed, the working potentials of the tip and the substrate being the same as those for recording the approach curve. Calculations. Simulations of the transient responses were obtained using the commercially available program, Comsol Multiphysics 3.4, which allows the resolution of the diffusion differential equations based on finite elements.20 We followed the same general procedure as depicted in ref 21 and similar geometry representation as previously published.7a,14 The size of the box was chosen to be at least 2-5 times larger than the size of the electrode, with a total number of points around 100000. The stability of the calculations was tested by varying the mesh sizes, especially for the case where γ differed considerably from unity. 3. Results and Discussions Cyclic Voltammetry Studies. Before starting the SECM measurements, O2 solutions were rapidly examined by cyclic voltammetry, allowing us to adjust the values of the applied potentials at the UME and the Pt substrate in such a way that electron-transfer kinetics could be considered as infinitely fast. Typical cyclic voltammograms of the electrochemical reduction of oxygen in pure [Et3BuN][NTf2] and in pure DMF are shown in Figure 1. In the RTIL, the presence of a large anodic peak, which is related to the oxidation of electrogenerated superoxide to produce oxygen, provides evidence that the generated superoxide is chemically stable in our different experimental

Example of the Reversible Reduction of O2 to Superoxide conditions, in agreement with previous reports.8 As explained before, the effect of a difference in the diffusion coefficients of the redox couple could be evidenced by cyclic voltammetry under conditions where a change of diffusion regime occurs. The dissymmetry observed between the forward (plateau) and the backward (peak) sweeps, similar to the results reported in the literature, is indicative of a change in the diffusion regime.6,7 The competition between the linear diffusion, for which peak-shaped voltammograms are obtained (backward scan in Figure 1a), and the spherical diffusion, for which S-shaped voltammograms are obtained (forward scan in Figure 1a), depends on a single parameter p ) (RT/FV)(D/r2) (where R is the gas constant, D the diffusion coefficient, F the Faraday constant, r the radius of the microelectrode, and V the scan rate).6c In our case, this difference of shape is thus indicative of a smaller diffusion coefficient for O2•- than that for O2 or indicates that the ratio γ ) DO2•-/DO2 is considerably smaller than unity. Upon addition of DMF in the ionic liquid, the difference between the forward and backward sweeps diminishes and then almost disappears after the volume fraction of DMF reaches about 60%. For higher DMF concentrations, the voltammograms display an S-shaped cyclic voltammogram, such as the one shown in Figure 1b, for the reduction of oxygen in pure DMF with 0.1 mol · L-1 [Et3BuN][NTf2] as the supporting electrolyte. Qualitatively, this modification indicates that γ tends to unity upon DMF addition and has a negligible effect on the electrochemical response for the largest concentrations of DMF in our experimental conditions (for a given electrode radius a and scan rate V). Transient SECM Experiments. In order to get a better understanding of this phenomenon, measurements of γ of the redox couple were obtained by SECM experiments. The method is based on a change in the diffusion regime that passes from a semilinear diffusion to a steady-state regime as it could be obtained with micrometer size electrodes.22 As other methods based on this effect, the measurements does not require a priori knowledge of the initial concentration of the dissolved O2 in the solution.7,14 This benefit simplifies the measurements as the saturated concentration depends on the RTIL/DMF mixture or other environmental conditions (temperature, air pressure, and so forth). The principle of the determination of γ by transient SECM was proposed several years ago by Unwin et al.14 and was recently extended to the measurement of small γ.7 In short, the current at the tip electrode is measured as a function of the time when the tip is maintained at a chosen distance from a conducting surface. The way that the dimensionless current Imin/ IST (Imin is the minimum value of the current reached after the first several seconds of the measurement) reaches its steadystate value, especially through the passage by a local minimum, provides a unique determination of γ ) (DO2•-)/(DO2) by comparison of the data with working curves; hence, the absolute values of DO2•- and DO2 can be deduced by considering that the radius of the UME is known.14 Typical transient SECM curves, which were recorded in the RTIL solutions containing different volume fractions of DMF over the ionic liquid [Et3BuN][NTf2], are presented in Figure 2a-f. The normalized current I ) IT/Iinf (IT is the current at the microelectrode and Iinf is the steady-state current when the electrode is localized at an infinite distance from the substrate) was plotted as a function of the time at a chosen distance. From Figure 2a-f, it is visible that upon increasing the DMF volume fraction, (i) the steady-state current, IST, increases, (ii) the time needed for the current to reach the steady state largely

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Figure 2. Experimental (black) and theoretical (red) transient SECM curves on a 5 µm radius gold microelectrode in solutions containing different volume fractions of DMF over [Et3BuN][NTf2]: (a) 0% DMF (pure [Et3BuN][ NTf2]), theoretical curve calculated with L ) 1.14, γ ) 0.017; (b) 10% DMF, theoretical curve calculated with L ) 1.00, γ ) 0.03; (c) 40% DMF, theoretical curve calculated with L ) 0.76, γ ) 0.06; (d) 60% DMF, theoretical curve calculated with L ) 0.78, γ ) 0.09; (e) 80% DMF, theoretical curve calculated with L ) 0.95, γ ) 0.14; (f) 100% DMF using [Et3BuN][ NTf2] at 0.1 mol · L-1 as the supporting electrolyte, theoretical curve calculated with L ) 0.78, γ ) 0.30. T ) 36 °C.

TABLE 1: Values of the Diffusion Coefficients of O2 and O2•- and 1/γ in Solutions with Different Volume Fractions of DMF %DMF (v/v) DO2/DO2•- (1/γ) DO2 /105 cm2 · s-1 DO2•- /105 cm2 · s-1 0 10 40 60 80 100

59 33 17 11 7.1 3.3

0.60 0.71 1.1 1.7 2.3 4.5

0. 010 0.022 0.065 0.15 0.32 1.4

decreases, and (iii) the relative amplitude of the minimum, Imin, that reflects the difference between the diffusion coefficients of O2/O2•- becomes less visible. All of these observations fall in line with both an increase of the two diffusion coefficients and γ values being close to unity. Analyses of the curves according to the previously published procedure7 provide quantitative measurements of the diffusion coefficients, as displayed in Table 1. The smallest value of γ (0.017) was obtained in the pure ionic liquid (Figure 2a), in agreement with literature data.23 Meanwhile, for the other extreme situation, the 100% DMF case (Figure 2f), the largest value of γ was found to be 0.3. This value is not sufficiently different from unity for influencing the shape of the reversible voltammogram that appears in our experimental conditions with a typical symmetrical S-shape.10b The value also agrees with literature as similar γ values were reported for the reduction of oxygen to superoxide in organic solvents, for example, in dimethylsulfoxide (DMSO).24 On the

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Figure 4. Experimental values of 1/γ ) DO2/DO2•- as a function of the DMF volume fraction. The inset shows that the value of γRTIL/γ decreases exponentially with the increase of the DMF volume fraction (γRTIL is the value in the pure RTIL prior DMF addition).

Figure 3. Experimental values of DO2 (top) and DO2•- (bottom) as a function of the DMF volume fraction. The insets show the linear variations of ln(D/DRTIL) ) VS/a, where DRTIL are the values of DO2 or DO2•- in the pure RTIL prior DMF addition (a ) 55.2 and 21.5, respectively, and VS is in %DMF; see text). They show that both DO2 and DO2•- vary exponentially with the increase of the DMF volume fraction.

other hand, in 100% DMF, the results agree well with those reported in organic solvents,25 confirming the validity of the SECM method to determine the diffusion coefficients of the redox couple when they are different. As a final test, the experimental SECM transient curves were compared with the simulated curves corresponding to the derived diffusion coefficients reported in Table 1. A good agreement was found between the simulated curves (red) and experimental data (black), confirming the quality of the measurements. Variations of the Diffusion Coefficients upon DMF Addition. As a general tendency (see Table 1), both of the diffusion coefficients of O2 and O2•- raise upon DMF addition, as expected from a decrease of the viscosity of the RTIL with the addition of an organic solvent.3a,11 Surprisingly, both DO2 and DO2•- vary exponentially with the increase of the DMF volume fraction (see Figure 3), and we noted that addition of DMF leads to critical changes of DO2 and DO2•- only for the highest DMF volume fractions. By comparison, a larger DMF concentration is required for the same relative variation O2•-, and the variation is more sudden. In this last case, at least 80% of DMF has to be added in the ionic liquid to obtain a substantial increase of the mass transport and a media that behaves almost like the organic solvent. Another way for observing this difference is the plot of the ratio of 1/γ ) DO2/ DO2•- (Figure 4) to the DMF volume fraction. As for the individual diffusion coefficients, the data show a sharp decrease of 1/γ upon increasing of the quantity of DMF that is close to an almost exponential law. From a practical point of view, it means that any values of the diffusion coefficients for O2 and O2•- (or of the γ ratio) could be estimated from an empirical formula D ) DRTIL exp[(VS)/ (a)] (where DRTIL is the diffusion coefficient in the pure ionic liquid, Vs the volume fraction of the added solvent, and a a constant that depends on the media and the molecule, a ) 55.2 and 21.5 for DO2 and DO2•-, respectively, with VS in %).

A key question remains about the origin of all of these exponential variations. In a direct relation with our work, a comparable behavior has been reported for the diffusion coefficient of the neutral ferrocene in a series of comparable ionic liquids when DMF was added.3a,26 Interestingly, the authors noticed that the exponential increase of the diffusion coefficient is accompanied by a parallel exponential decrease of the viscosities of the RTIL/DMF mixture when the DMF fraction increases.27 This similarity allows the authors to conclude that the diffusion of ferrocene in a RTIL/DMF mixture follows the classical Stokes-Einstein equation, D ) (kBT)/(pπηr), where kB is the Boltzmann constant, T the absolute temperature, η the viscosity, r the hydrodynamic radius, and p a constant equal to 4 or 6, depending if slip or stick conditions applied.4 This point was supported by recent investigations of Compton et al. who found a Stokes-Einstein behavior for the ferrocene and other large organic compounds in different classes of RTILs.28 However, for small molecules like O2, the same authors reported that the agreement with this relation is poor.28 In the literature, most of fundamental investigations concern RTIL/water systems. Recent simulations concerning the ion association in an organic/ RTIL mixture ([BMIM][PF6]/naphthalene mixture, [BMIM]) 1-butyl-3-methylimidazolium) support the same general idea about a fairly sharp but continuous transformation in the bonding character of the RTIL cation upon addition of the molecular solvent.11 As discussed previously,6,7 such properties of the RTIL cation play a fundamental role in ion associations with a charged substrate, which are expected to be strong for the superoxide because of its small size.8e Generally, authors divide the range of concentrations in two regions, a salt-rich region and a solventrich region.2 In the first region, the solvent is accommodated in the ionic liquid structure, and in the second region, for high molecular solvent concentration, an appreciable amount of the solvent may be regarded as “free” solvent that will replace the ionic phase with a purely molecular phase. Returning to our experiments, if the exponential variations exist for DO2•- and DO2, the exponential factors, a, are considerably different for O2•- and O2. It reflects that the amplitudes of the changes as a function of the amount of DMF are larger for DO2•- than those for DO2, indicating stronger interactions between the substrate and the media in the first case. As previously discussed, such ion associations between O2•- and one or several cations of the RTIL explain the low value of γ as these interactions are negligible with the neutral species.7a,8e On the basis of simple electrostatic considerations, addition of a polar solvent, like DMF, should increase the apparent dielectric constant of the mixture and thus reduce the strength of ion-ion interactions. This explains why DO2•- varies on a much larger range than

Example of the Reversible Reduction of O2 to Superoxide DO2 (see the insets in Figure 3) and falls in line with the higher addition of DMF required for observing the “molecular solvent behavior” for the diffusion of O2•-. Concerning the exponential variation of γ, in the framework of the Stokes-Einstein relation, this corresponds to a sudden increase of the hydrodynamic radius of the superoxide in the salt-rich region. Such an increase could be associated with a fairly sharp transformation in the bonding character of the RTIL cation, as was described in the case of the [BMIM][PF6]/naphthalene mixtures,11 resulting in considerable modifications of the superoxide ion-RTIL cation interactions. 4. Conclusion The reduction of oxygen has been used to probe the effects of the addition of a cosolvent such as DMF into [Et3BuN][NTf2] (a commonly used ionic liquid) on the diffusion of charged species versus neutral species. SECM measurements in the transient mode with simulated working curves provide the needed information about diffusion coefficient ratios and the absolute values of the diffusion coefficients of the redox couple O2/O2•-, which are different because of their different interactions with the ions of the ionic liquid. Experimentally, all of the values of γ, DO2•-, and DO2 were found to vary exponentially with the increase of the DMF volume fractions. However, DO2•- was found to change on a much larger range than DO2 (around 1000 times more), in line with the expected strong ionic interactions between the superoxide and the RTIL cation. All of these data support the occurrence of a sharp transformation in the bonding character of the RTIL cation upon addition of a molecular solvent, as predicted in recent theoretical simulations. References and Notes (1) (a) Ionic Liquids in Synthesis; Wasserscheid, P., Welton, T., Eds.; Wiley-VCH: Weinheim, Germany, 2003. (b) Seddon, K. R. Nat. Mater. 2003, 2, 363. (c) Green Industrial Applications of Ionic Liquids; Roger, R. D., Seddon, K. R., Volkov, S., Eds.; NATO Sciences Series; Kluwer: Dordrecht, The Netherlands, 2002; Vol. 92. (d) Rantwijk, F.; van Sheldon, R. A. Chem. ReV. 2007, 107, 2785. (2) (a) Seddon, K. R.; Stark, A.; Torres, M. J. Pure Appl. Chem. 2000, 72, 2275. (b) Xu, H. T.; Zhao, D. C.; Xu, P.; Liu, F. Q.; Gao, G. J. Chem. Eng. Data 2005, 50, 133. (c) Paul, A.; Samanta, A. J. Phys. Chem. B 2008, 112, 947, and references therein. . (3) (a) Comminges, C.; Barhdadi, R.; Laurent, M.; Troupel, M. J. Chem. Eng. Data 2006, 51, 680. (b) Najdanovic-Visak, V.; Esperanca, J. M. S. S.; Rebelo, L. P. N.; Nunes da Ponte, M.; Guedes, H. J. R.; Seddon, K. R.; de Sousa, H. C.; Szydlowski, J. J. Phys. Chem. B. 2003, 107, 12797. (c) SirieixPlenet, J.; Gaillon, L.; Letellier, P. Talanta 2004, 63, 979. (d) Zhang, J.; Wu, W.; Jiang, T.; Gao, H.; Liu, Z.; He, J.; Han, B. J. Chem. Eng. Data 2003, 48, 1315. (4) (a) Buzzeo, M. C.; Evans, R. G.; Compton, R. G. ChemPhysChem 2004, 5, 1106. (b) Silvester, D. S.; Compton, R. G. Z. Phys. Chem. 2006, 220, 1247. (c) Hapiot, P.; Lagrost, C. Chem. ReV. 2008, 108, 2238. (5) (a) See, for example, ref 5b for an example of the strong ion-pairing effect on the electrochemical reduction of aromatic compounds. (b) Fry, A. J. J. Electroanal. Chem. 2003, 546, 35. (6) (a) Buzzeo, M. C.; Klymenko, O. V.; Wadhawan, J. D.; Hardacre, C.; Seddon, K. R.; Compton, R. G. J. Phys. Chem. A 2003, 107, 8872. (b) Buzzeo, M. C.; Klymenko, O. V.; Wadhawan, J. D.; Hardacre, C.; Seddon, K. R.; Compton, R. G. J. Phys. Chem. A 2004, 108, 3947. (c) Evans, R. G.; Klymenko, O. V.; Saddoughi, S. A.; Hardacre, C.; Compton, R. G. J. Phys. Chem. B 2004, 108, 7878. (d) Barnes, A. S.; Rogers, E. I.; Aldous, L.; Hardacre, C.; Wildgoose, G. G.; Compton, R. G. J. Phys. Chem. C 2008, 112, 13709. (7) (a) Ghilane, J.; Lagrost, C.; Hapiot, P. Anal. Chem. 2007, 79, 7391. (b) Zigah, D.; Ghilane, J.; Lagrost, C.; Hapiot, P. J. Phys. Chem. B 2008, 112, 14952.

J. Phys. Chem. B, Vol. 113, No. 7, 2009 2023 (8) (a) AlNashef, I. M.; Leonard, M. L.; Kittle, M. C.; Mattews, M. A.; Weidner, J. W. Ind. Eng. Chem. Res. 2002, 482, 87. (b) Noda, A.; Susan, M. A. B. H.; Kudo, K.; Mitsushima, S.; Hayamizu, K.; Watanabe, M. J. Phys. Chem. B 2003, 107, 4024. (c) Zhang, D.; Okajima, T.; Matsumoto, F.; Ohsaka, T. J. Electrochem. Soc. 2004, 151, D31. (d) Katayama, Y.; Sekiguchi, K.; Yamagata, M.; Miura, T. J. Electrochem. Soc. 2005, 152, E247. (e) Islam, M. M.; Ohsaka, T. J. Phys. Chem. C 2008, 112, 1269. (9) Sawyer, D. T.; Valentine, J. S. Acc. Chem. Res. 1981, 14, 393. (10) (a) The consequences of an inequality of diffusion coefficients on electrochemical measurements are complex and often difficult to detect even if g values are considerably different from unity.10b-e Large inequalities could lead to substantial errors in the extracted data. For example, E1/2 ) E° + RT/F log(Dred/Dox)1/2 in cyclic voltammetry or E1/2 ) E° + RT/F log(Dred/Dox) in steady-state voltammetry on an ultramicroelectrode. Concerning the effects of g on the determinations of kinetics data, see for example refs 10c-e and the references therein. (b) Bard, A. J.; Faulkner, L. R. Electrochemical Methods: Fundamentals and Applications, 2nd ed.; John Wiley & Sons: New York, 2001; pp 180 and 230. (c) Andrieux, C. P.; Hapiot, P.; Save´ant, J.-M. J. Electroanal. Chem. 1984, 172, 49. (d) Andrieux, C. P.; Hapiot, P.; Save´ant, J.-M. J. Electroanal. Chem. 1985, 186, 237. (e) Andrieux, C. P.; Hapiot, P.; Save´ant, J.-M. J. Electroanal. Chem. 1985, 189, 121. (11) Del Popolo, M. G.; Mullan, C. L.; Holbrey, J. D.; Hardracre, C.; Ballone, P. J. Am. Chem. Soc. 2008, 130, 7033. (12) Andrieux, C. P.; Hapiot, P.; Save´ant, J.-M. J. Am. Chem. Soc. 1987, 109, 3768. (13) Scanning Electrochemical Microscopy; Bard, A. J., Mirkin, M. V., Eds.; Marcel Dekker: New York, 2001. (14) Martin, R. D.; Unwin, P. R. J. Electroanal. Chem. 1997, 439, 123. (15) Carano, M.; Bond, A. M. Aust. J. Chem. 2007, 60, 29. (16) For more quantitative descriptions of diffusion concentration profiles at an ultramicroelectrode corrupted by natural convection, see: Amatore, C.; Knobloch, K.; Thouin, L. J. Electroanal. Chem. 2007, 106, 17. (17) The reason for using the 36 °C temperature instead of the classical room temperature is to reduce the viscosity and thus to increase the diffusion coefficients. The drawback is also a likely decrease of the O2 concentration in the ionic liquid. Experimentally, we found that globally passing from 20 to 36°C improves the situation both in terms of signal/noise ratio and diffusion rates. (18) (a) Bonhote, P.; Diaz, A. P.; Papageorgiou, N.; Kalyasundaram, K.; Gratzel, M. Inorg. Chem. 1996, 35, 1168. (b) Sun, J.; Forsyth, G. R.; MacFarlane, D. R. J. Phys. Chem. B 1998, 102, 8858. (19) Ghilane, J.; Hapiot, P.; Bard, A. J. Anal. Chem. 2006, 78, 6868. (20) Comsol Multiphysics 3.4. http: //www.comsol.com. (21) (a) Lefrou, C. J. Electroanal. Chem. 2006, 592, 103. (b) Cornut, R.; Lefrou, C. J. Electroanal. Chem. 2007, 604, 91. (c) Cornut, R.; Lefrou, C. J. Electroanal. Chem. 2007, 608, 59. (22) (a) Amatore, C.; Azzabi, M.; Calas, P.; Jutand, A.; Lefrou, C.; Rollin, Y. J. Electroanal. Chem. 1990, 298, 45. (b) Stulik, K.; Amatore, C.; Holub, K.; Marecek, V.; Kutner, W. Pure Appl. Chem. 2000, 72, 1483. (23) DO2 and DO2•- in the pure ionic liquid [Et3BuN][NTf2] (DO2 ) 3.2 × 10-6 cm2 s-1 and DO2•- ) 1.45 × 10-7 cm2 s-1 at 36 °C) are in the same range as those previously reported in the same RTIL [Et3BuN][NTf2] (DO2 ) 3.9 × 10-6 cm2 s-1 and DO2•- ) 9.2 × 10-8 cm2 s-1 at 25 °C7a) or as those in another quaternary ammonium-based ionic liquid [N6222][NTf2],6c taking into account the higher working temperature of our experiments that also leads to a decrease of solvent viscosity.2,3 (24) Sawyer, D. T.; Roberts, J. L. J. Electroanal. Chem. 1966, 12, 90. (25) (a) In DMSO, DO2 ) 3.23 × 10-5 cm2 s-1 and DO2•- ) 1.08 × 10-5 cm2 s-1 (γ ) 0.33),24 and in DMF, DO2 ) 4.76 ×10-5 cm2 s-1.25b. (b) Tsushima, M.; Tokuda, K.; Ohsaka, T. Anal. Chem. 1994, 66, 4551. (26) The ferrocene/ferrocenium couple is usually chosen as a model of electroactive systems in many studies performed in organic solvents and RTILs.4 (27) (a) Concerning the exponential variation of the viscosity with the molecular solvent molar fraction, such observations have been reported in several instances and more than a decade ago for the variations of the viscosity in pure fused electrolyte/organic solvent mixtures.27b This empirical law was proposed as an effective general formula in RTIL/molecular solvents.2a. (b) Kumar, A. J. Am. Chem. Soc. 1993, 115, 9243. (28) Rogers, E. I.; Silvester, D. S.; Poole, D. L.; Aldous, L.; Hardacre, C.; Compton, R. G. J. Phys. Chem. C 2008, 112, 2729.

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