10976
J. Phys. Chem. C 2008, 112, 10976–10981
Electrode Kinetics and Mechanism of Iodine Reduction in the Room-Temperature Ionic Liquid [C4mim][NTf2] Emma I. Rogers,† Ian Streeter,† Leigh Aldous,‡ Christopher Hardacre,‡ and Richard G. Compton*,† Department of Chemistry, Physical and Theoretical Chemistry Laboratory, Oxford UniVersity, South Parks Road, Oxford OX1 3QZ, United Kingdom and School of Chemistry and Chemical Engineering/QUILL, Queen’s UniVersity Belfast, Belfast, Northern Ireland BT9 5AG, United Kingdom ReceiVed: April 4, 2008; ReVised Manuscript ReceiVed: May 5, 2008
The fast electrochemical reduction of iodine in the RTIL 1-butyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide, [C4mim][NTf2], is reported and the kinetics and mechanism of the process elucidated. Two reduction peaks were observed. The first reduction peak is assigned to the process
3I2 + 2e- h 2I3The second reduction peak is assigned to the process
I3- + 2e- h 3IA diffusion coefficient of 6.6 × 10-11 m2 s-1 (298 K) is inferred for I2 in [C4mim][NTf2] with a solubility of 1.70 mM. A mechanistic study was undertaken using a digital simulation program based on the mechanism ka′
I2 + 2e- h 2Ikb′
kf,hom
I- + I2 y\z I3kb,hom
and simulation of the first reduction wave allowed extraction of various kinetic parameters including the diffusion coefficients for I2, I3 , and I , rate constants for the homogeneous process (kf,hom and kb,hom), and the ′ heterogeneous rate constants ka and k′b, and the associated transfer coefficients. The electrode process was found to be consistent with the following form of Butler-Volmer kinetics F F ∂[I2] ) ka′ e-R RT E[I2] - kb′ eβ RT E[I-] ∂z The mechanistic basis for this rate law is discussed.
DI2
1. Introduction Room-temperature ionic liquids (RTILs) are defined as salts that exist in the liquid phase at temperatures around 298 K.1,2 As ‘liquid salts’, RTILs display negligible volatility (low vapor pressures) and high ionic conductivities (due to an abundance of charge carriers), the former suggesting that RTILs are greener for the environment compared to volatile organic compounds (VOCs) and the latter allowing for their use as solvents in electrochemical studies without supporting electrolytes.1–6 RTILs are comprised of an asymmetric, bulky organic cation and a weakly coordinating organic or inorganic anion, and a large number of possible combinations allows for the ability to ‘fine tune’ the solvent properties for a specific purpose.3 The stability of the ions suggest that ILs frequently possess high thermal stability and large electrochemical windows.4 High viscosity, generally 1-2 orders of magnitude more than traditional * To whom correspondence should be addressed. Phone: +44(0) 1865 275 413. Fax: +44(0) 1865 275 410. E-mail:
[email protected]. † Oxford University. ‡ Queen’s University Belfast.
Figure 1. Structures of the anion and cation used as the RTIL in this study.
solvents, e.g., acetonitrile, results in lower diffusion coefficients of the electroactive species than would be observed in conventional solvents.1,2 In this paper, we report the cyclic voltammetry for the reduction of iodine in the room-temperature ionic liquid 1-butyl3-methylimidazolium bis(trifluoromethylsulfonyl)imide,[C4mim][NTf2]. Many authors have used voltammetric techniques for the study of iodine in various media.7–14 Nakata et al.10 used cyclic voltammetry to investigate the electrochemical reduction of iodine on a platinum electrode in a range of nonaqueous/ aprotic solvents, e.g., dimethylformamide (DMF), acetonitrile (MeCN), dimethyl sulfoxide (DMSO), and observed two
10.1021/jp802934y CCC: $40.75 2008 American Chemical Society Published on Web 06/26/2008
Electrode Kinetics and Mechanism of Iodine Reduction
J. Phys. Chem. C, Vol. 112, No. 29, 2008 10977 diffusion coefficient data and electrode kinetic parameters and mechanism for iodine reduction, extracted from the voltammetry obtained using digital simulation techniques.
Figure 2. Cross section of the glass cell used to conduct electrochemical experiments on 20 µL samples of RTILs under a controlled atmosphere.
reduction waves, which were assigned to formation of triiodide (I3 ) followed by formation of iodide (I ), shown in eqs 1 and 2.
3I2 + 2e- h 2I3
(1)
I3 + 2e h 3I
(2)
This electrode process is analogous to that obtained for the reduction of iodine in acetonitrile,14 acetic anhydride,7,8 perchloric acid,11 and aluminum chloride-butylpyridinium chloride ionic liquid.9 A recent paper15 studied the cyclic voltammetry for the oxidation of iodide in the ionic liquid [C4mim][NTf2], which involved the reversible oxidation of iodide to iodine via formation of a triiodide intermediate species following the mechanisms denoted above. Iodine reduction is an important two-electron reduction process. The mechanistic and kinetic aspects of this process have been studied in both aqueous16–18 and aprotic solvents,19 and the reduction was found to follow the mechanism given below, where iodine is first reduced to iodide (eq 3; a summary of eqs 1 and 2) followed by a chemical step in which triiodide is formed from the reaction between iodine and iodide (eq 4) ka′
I2 + 2e- y\z 2I-
(3)
k′b
kf,hom
I- + I2 y\z I3
(4)
kb,hom
However, the electrode kinetics and mechanism of such processes are completely unexplored in RTILs; the only studies of electrode kinetics and mechanisms that have been reported in a quantitative manner are essentially simple one-electron reductions of the type A ( e- h B;9 multielectron processes are hitherto unadressed. The current paper reports, in addition to the voltammetry for the reduction of iodine in [C4mim][NTf2], solubility and
2. Experimental Section 2.1. Chemical Reagents. Iodine (BDH Chemicals Ltd., 99%), acetonitrile (Fischer Scientific, dried and distilled, >99%), ferrocene (Aldrich, 98%), and tetra-n-butylammonium perchlorate (TBAP, Fluka, Puriss electrochemical grade, >99%) were used as received without further purification. 1-Butyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide, [C4mim][NTf2] (see Figure 1 for structure) was prepared by standard literature procedures.20,21 2.2. Instrumentation. Electrochemical experiments were performed using a computer-controlled µ-Autolab potentiostat (Eco-Chemie, The Netherlands). A conventional two-electrode arrangement was employed to study the voltammetry with a platinum electrode (10 µm diameter) as the working electrode. A 0.5 mm diameter silver wire (Goodfellow Cambridge Ltd.) was used as a quasi-reference electrode. The working electrode was modified with a section of disposable micropipette tip to create a small cavity above the electrode surface into which a drop (ca. 20 µL) of ionic liquid solvent was placed. The electrodes were housed in a specially designed glass cell22 for investigating microsamples of ionic liquids under a controlled atmosphere (see Figure 2). All experiments were carried out in a Faraday cage, thermostatted at 298 ( 1 K, which also served to minimize background noise. A saturated solution of iodine was made directly in 60 µL of [C4mim][NTf2], which had previously been purged for 2 h under vacuum (0.05 Torr). The solution was then stirred in a foilcovered, sealed vial (to protect from light and atmospheric moisture), and 20 µL of this stock solution was pipetted into the cell. For voltammetry versus a redox couple, in this case cobaltocenium|cobaltocene, Cc+|Cc, a 10 mM solution of cobaltocenium hexafluorophosphate (CcPF6) was made up in acetonitrile and 10 µL of this was pipetted into the T-cell along with the iodine solution to give an overall Cc+ concentration of 5 mM. The solution was purged under vacuum for at least 90 min prior to any measurements being taken and during the course of the experiment, which also served to remove any impurities and trace atmospheric moisture present in the RTIL. The microdisk electrode was polished on soft lapping pads (Kemet Ltd., Kent, U.K.) using 1.0 and 0.3 µm alumina (Buehler, Lake Bluff, Illinois), respectively. The electrode was calibrated electrochemically by analyzing the steady-state voltammetry of a 2 mM solution of ferrocene in acetonitrile containing 0.1 M TBAP using a value for the diffusion coefficient of 2.30 × 10-9 m2 s-1 at 298 K,23 allowing calculation of the electrode radius. 2.3. Potential Step Chronoamperometric Experiments. Potential step chronoamperometric transients for the reduction of iodine were achieved using a sample time of 0.01 s. The solution was pretreated by holding the potential at a point corresponding to zero faradaic current for 20 s, after which the experimental transients were obtained by stepping the potential to a position after the first reduction peak, and the current was
TABLE 1: Boundary Conditions boundary semi-infinite boundary insulating surface symmetry axis electrode surface
condition for I2 [I2] ) [I2] bulk ∂[I2]/∂z ) 0 ∂[I2]/∂r ) 0 eq 12 or 14
condition for I[I-] ) 0 ∂[I-]/∂z ) 0 ∂[I-]/∂r ) 0 DI-∂[I-]/∂z ) -2DI2∂[I2]/∂z
condition for I3 [I3] ) 0 ∂[I3 ]/∂z ) 0 ∂[I3 ]/∂r ) 0 ∂[I3 ]/∂z ) 0
10978 J. Phys. Chem. C, Vol. 112, No. 29, 2008
Rogers et al. ka′
I2 + 2e- h 2Ik′b
kb, hom
I- + I2 y\z I3
(8) (9)
kf,hom
The homogeneous step in reaction 9 is described by a secondorder forward rate constant, kf, hom, and a first-order backward rate constant, kb,hom. The two-electron-transfer process in eq 8 is itself considered to be a two-step process
I2 + e- h I•(ads) + I-
(10)
I•(ads) + e- h I-
(11)
I•(ads)
Figure 3. CV for the reduction of iodine in [C4mim][NTf2] on a platinum microelectrode (diameter 10 µm) at a range of scan rates (10, 20, 50, 100, 200, 400, 700, and 1000 mV s-1). (a) CV obtained when scanning to a more negative potential to give a two-step reduction of iodine. (b) CV obtained for the first step of the reduction process occurring at the electrode. (Inset to b) Chronoamperometric transient recorded for the reduction of iodine. The potential was stepped from +1.20 to +0.27 V.
measured for 10 s. Diffusion coefficients and solubility data were extracted from the experimental transients by fitting the data using the nonlinear curve-fitting function available in Origin 7.5 (Microcal Software Inc.) following the Shoup and Szabo24 approximation for the time-dependent current response at microdisk electrodes. The equations used in this approximation (eqs 5, 6, and 7) sufficiently describe the current response to within 0.6% accuracy.
I ) -4nFDcrd f(τ) 1 f(τ) ) 0.7854 + 0.8863τ- 2
(5)
(
+ 0.2146 exp
τ )
1 -0.7823τ- 2
4Dt 4r2d
)
(6) (7)
where n is the number of electrons transferred, F is the Faraday constant, D is the diffusion coefficient, c is the initial concentration of parent species, rd is the radius of the disk electrode, and t is the time. Optimization of the experimental data using a fixed value of the radius of the electrode (determined previously from calibration) allows the value for the diffusion coefficient (D) of the species and the product of the number of electrons multiplied by concentration (nc) to be determined.
where is an iodine atom most likely adsorbed to the electrode surface and all other species are in solution. We consider two different kinetic regimes for reaction steps 10 and 11. Under the first regime, the electron transfer in eq 10 is rate limiting and the electron transfer in eq 11 is fast such that I•(ads) and I- are in a dynamic equilibrium at the electrode surface, controlled only by the electrode potential but in which the surface coverage of I•(ads) is proportional to the solutionphase concentration of [I-]. This is a limit of relatively low I• (ads) coverage. The rate equation for the flux of material through the electrode surface is therefore given by eq 12, which makes use of Butler-Volmer kinetics
DI2
F F ∂[I2] ) kae-R RT (E-Ef)[I2] - kbe(2-R) RT (E-Ef)[I-]2 (12) ∂z
where ka and kb are heterogeneous rate constants, Ef and R are the formal electrode potential and charge-transfer coefficient for reaction 10, and E is the electrode potential. Notice that this rate equation is second order with respect to I- and first order with respect to I2. In the second kinetic regime we again consider reaction step 10 is the rate-determining process for the reduction of iodine, but the rate equation at the electrode surface for this mechanism is given by eq 13
DI2
F F ∂[I2] ) kae-R RT (E-Ea)[I2] - kbeβ RT (E-Eb)[I-] ∂z
(13)
where ka, Ea, and R are the rate constant, formal electrode potential, and charge-transfer coefficient for reaction step 10 and kb, Eb, and β are the equivalent parameters for reaction step 11. Notice that this rate equation is first order with respect to both I- and I2. This equation can arise in the limit where the electrode surface coverage of I•(ads) is saturated (high coverage) and therefore independent of [I-]. Equation 13 can be rewritten as eq 14,
3. Mathematical Model and Numerical Simulation The experimental voltammetry reported below is interpreted by comparison with numerical simulations for the reduction of I2 to give I3 at a microdisk electrode. The electrode reduction is modeled as a two-step process in which I2 is first reduced to I- at the electrode surface, which then reacts with I2 in a homogeneous process to give I3
Figure 4. CV for the reduction of iodine in the presence of cobaltocenium hexafluorophosphate in [C4mim][NTf2] on a platinum microelectrode (diameter 10 µm) at 100 mV s-1.
Electrode Kinetics and Mechanism of Iodine Reduction
J. Phys. Chem. C, Vol. 112, No. 29, 2008 10979
TABLE 2: Summary of the Peak Potentials (vs Ag) for Iodine and Cc+ Redox Couples in [C4mim][NTf2] at scan rates of 10, 20, 50, 100, 200, 400, 700, and 1000 mV s-1 ν/mV s-1
Ep(red) I2/V
Ep(ox) I2/V
E1/2 I2/V
Ep(red) Cc+/V
Ep(ox) Cc+/V
E1/2 Cc+/V
∆Epp/V
10 20 50 100 200 400 700 1000
+0.413 +0.419 +0.424 +0.435 +0.455 +0.475 +0.477 +0.488
+0.684 +0.681 +0.676 +0.670 +0.667 +0.662 +0.656 +0.652
+0.548 +0.550 +0.550 +0.552 +0.561 +0.562 +0.567 +0.570
-1.096 -1.108 -1.106 -1.116 -1.106 -1.096 -1.095 -1.088
-1.028 -1.018 -1.006 -1.006 -0.995 -0.995 -0.994 -0.996
-1.062 -1.063 -1.056 -1.061 -1.051 -1.046 -1.044 -1.042
+1.610 +1.613 +1.606 +1.613 +1.612 +1.614 +1.611 +1.612
which uses the alternative heterogeneous rate constants k′a and k′b and does not contain the formal electrode potentials
DI2
F F ∂[I2] ) k′ae-R RT E[I2] - k′beβ RT E[I-] ∂z
(14)
The diffusion reaction equations for the three species I2, I-, and I3 are given by equations 15, 16, and 17 using the cylindrical polar coordinates r and z
(
∂[I2] ∂2[I2] 1 ∂[I2] ∂2[I2] + ) DI2 + ∂t r ∂r ∂r2 ∂z2
(
)
)
(15)
∂[I-] ∂2[I-] 1 ∂[I-] ∂2[I-] + - kf[I2][I-] + kb[I) DI+ 3] 2 2 ∂t r ∂r ∂r ∂z (16)
(
)
2 ∂[I∂2[I3] 3] 1 ∂[I3 ] ∂ [I3 ] + + kf[I2][I-] - kb[I) DI3+ 3] ∂t r ∂r ∂r2 ∂z2 (17)
The concentration profiles of these three species are found by solving eqs 15, 16, and 17 subject to the boundary conditions in Table 1. The electrode potential is varied over time in a simulated cylic voltammetry experiment. The Faradaic current, I, is then found from the concentration profile through eq 18
I ) 4πFD
∫0r
e
∂ [I2] ∂z
|
rdr
z)0
(18)
Equations 15, 16, and 17 are discretized over a simulation grid using the fully implicit finite difference method. The resulting equations are nonlinear and solved using the NewtonRaphson method in conjunction with a generalized version of the Thomas algorithm. The simulation grid used is that described by Gavaghan,25 which uses a fine mesh density near the singularity at the edge of the disk electrode. The simulation program is tested for both spatial and temporal convergence
Figure 5. Tafel plot for reduction of iodine. A gradient of 112 mV was achieved to give a value of R ≈ 0.5.
such that any further increase in the grid density leads to a negligible change in the simulated current. 4. Results and Discussion [C4mim][NTf2] was the RTIL of choice in this study as it exhibits a wide electrochemical window and when fully purged shows no measureable voltammetric features. The use of this RTIL also allows comparison in voltammetric data, of a qualitative nature, to be made with that published previously for iodide.15 4.1. Electroreduction of Iodine in [C4mim][NTf2]. Figure 3 shows typical voltammetry of iodine (ca. 1.70 mM, as determined by chronoamperometry, see later) in [C4mim][NTf2] on a 10 µm diameter platinum electrode at scan rates from 10 to 1000 mV s-1. Two well-defined chemically reversible redox couples are observed (Figure 3a) at E1/2 ) +0.12 and +0.58 V vs Ag (at 1000 mV s-1) which result from two electrode processes. It has previously been determined9 that the reduction of iodine involves formation of a triiodide (I3-) intermediate, and it is assumed that this is also taking place for the reduction of iodine in this work. The first reduction wave corresponds to the electrochemical reduction of iodine to iodide followed by a chemical step (I2 + I-), which results in formation of this triiodide species. The second reduction wave can be assigned to dissociation of electrogenerated triiodide followed by reduction of iodine to iodide. These processes are summarized in eqs 19 and 20 for peaks 1 and 2, respectively.
3 I + e- h I3 2 2
(19)
I3 + 2e h 3I
(20)
The overall electrode process for both peaks can be written as follows
3 I + e- h 3I2 2
(21)
From the mechanism shown above, the theoretical height of the first reduction peak to the height of the second is 1:2; this ratio was also approximately observed experimentally (1:1.9 (() 0.13). The shape of the peaks suggests that both electrode processes show relatively fast electrode kinetics and peak current increases with scan rate. Although plots of peak current vs square root of scan rate are linear, the line does not typically pass through the origin, which is characteristic of voltammetry at macro- rather than microelectrodes.26 A decrease in peak separation, ∆Epp, with increasing scan rate is also opposite to the behavior observed at macroelectrodes, and a transition from steady-state behavior at slower scan rates to transient behavior at higher scan rates occurs. Slower diffusion in the more viscous RTIL
10980 J. Phys. Chem. C, Vol. 112, No. 29, 2008
Rogers et al.
Figure 6. Comparison of the experimental and simulation cyclic voltammograms for reduction of iodine in [C4mim][NTf2] at scan rates of (a) 10 mV s-1 (ka′ ) 272.56, k′b ) 4.25 × 10-9), (b) 100 mV s-1 (ka′ ) 312.36, k′b ) 5.69 × 10-9), (c) 400 mV s-1 (ka′ ) 312.36, k′b ) 5.69 × 10-9), and (d) 1000 mV s-1 (ka′ ) 434.92, k′b ) 9.26 × 10-9) using values of D and c close to that obtained from chronoamperometric data. Parameters are given in Table 3.
TABLE 3: Parameters for I2 Reduction from the Simulation Fits Shown in Figure 6 (in all cases, r ) β ) 0.5, kf,hom ) 1 × 106 mol-1 cm3 s-1 and kb,hom ) 10 s-1) parameter
average value
DI2/× 10-7 cm2 s-1 ka′ /cm s-1 k′b × 10-9/cm s-1
7.22 ( 0.30 333.05 ( 70.46 6.22 ( 2.14
medium suggests that the voltammetry obtained is intermediate of the micro/macroelectrode region. According to the following inequality26–28
V,
RTD nFr2d
(22)
where R is the universal gas constant (8.314 J K-1 mol-1) and T is the absolute temperature (298 K), a true steady-state response is achieved at a scan rate, V, of 29.9 mV s-1 assuming the diffusion coefficient D ) 6.6 × 10-11 m2 s-1 (from electrochemical reduction of I2, see later), n ) 2, and rd ) 5.3 µm. Figure 3b shows voltammetry obtained when scanned to a potential after the appearance of the first peak at scan rates from 10 to 1000 mV s-1. Potential-step chronoamperometry was undertaken on the first reduction peak, and the data obtained was analyzed to give the diffusion coefficient and solubility data. The potential was stepped from +1.20 to 0.27 V vs Ag, and the resulting transient is shown as an inset to Figure 3b, along with the best theoretical fit to the Shoup and Szabo24 equation. A nc value of 3.40 ( ( 0.10) mM was determined (where n is the number of electrons transferred, in this case n ) 2) to give a solubility of ca. 1.70 mM with a diffusion coefficient of 6.56 ( ( 0.05) × 10-11 m2 s-1 for iodine in [C4mim][NTf2 ] at 298 K. These values are 1-2 orders of magnitude smaller than diffusion coefficients obtained in pure water (9.4 × 10-10 m2 s-1 with a solubility of 1.10 mM)29 and in other aqueous solutions (1.4 × 10-9 m2 s-1 at 298 K in H2 SO413 and 1.0 × 10-9 m2 s-1 in HClO4),11,30,31 which is reasonable considering
the higher viscosity of the RTIL (44 cP)1 compared to water (0.89 cP)32 at 298 K. For simulation purposes (see section 4.2), it is useful to know the potential of the iodine|iodide couple relative to a known redox couple, for example, ferrocene|ferrocenium (Fc|Fc+) or cobaltocenium|cobaltocene (Cc+|Cc), to know if the potential varies over scan rate. Potentials of Fc|Fc+ and Cc+|Cc couples are not expected to change over scan rate, although it is not known if the potential of the I2|I- couple is constant. To investigate this the voltammetry of both I2 and cobaltocenium hexafluorophosphate (CcPF6) was studied at a range of scan rates. Figure 4 shows the voltammetry obtained; two iodine reduction peaks are observed (as reported previously) with a redox couple at -1.06 V vs Ag (at 100 mV s-1) corresponding to reduction and subsequent oxidation of the Cc+|Cc system. The peak potentials and half-wave potentials for the first I2 reduction and the first CcPF6 reduction are reported in Table 2. It is observed that the potential of iodine reduction relative to cobaltocenium reduction does not vary significantly with scan rate (maximum 8 mV). 4.2. Modeling the Reduction of Iodine in [C4mim][NTf2]. The next step was to simulate the first iodine reduction peak (voltammetry shown in Figure 3b) using the computer simulation programs based on the theory for both mechanisms detailed in section 3. The programs were used to simulate the experimental data. Both mechanisms denoted by the Butler-Volmer kinetics reported in eqs 12 and 14 were investigated. The kinetics following eq 12, which is second order with respect to iodide, failed to provide any set of parameterssphysically realistic or otherwiseswhich allowed adequate simultaneous fitting of the reduction and oxidation peaks. Simulation of the cyclic voltammograms using only the kinetics following eq 14 gave a set of parameters which allowed successful fitting of the experimental data. The parameters involved in the simulation program include heterogeneous rate constants (k′a and k′b), charge-transfer coefficients (R and β), forward and backward homogeneous rate
Electrode Kinetics and Mechanism of Iodine Reduction constants (kf,hom and kb hom, respectively), diffusion coefficients (D) of I2, I-, and I3 , and the concentration (c) of I2. DI2 was determined by chronoamperometry (see section 4.1) and optimized slightly to give the best fit over the range of scan rates investigated. Values of DI- and DI3- were optimized to be, and optimally fixed as, 2.50 × 10-7 and 1.00 × 10-7 cm2 s-1, respectively. The concentration was determined by chronoamperometry also, and this was fixed in all simulations (1.70 × 10-6 mol cm-3). A value for the transfer coefficient R (for reduction of I2) was determined to be 0.5 from a Tafel plot (shown in Figure 5), which gives a straight line with a gradient of 112 mV when potential is plotted as a function of log [1/i 1/ilim ]. This confirms the fast electron transfer to I2 as rate limiting. β was also chosen to be 0.5 when simulating the voltammetry; simulations in which β was approximated to 1.5 (or thereabouts) were unsuccessful in fitting the data. Figure 6 shows the best theoretical fit to the experimental cyclic voltammograms obtained at scan rates of (a) 10, (b) 100, (c) 400, and (d) 1000 mV s-1. The excellent fits shown in Figure 6 between experimental and fitted data, using the parameters given in Table 3, suggest that the mechanism which is first order with respect to both iodine and iodide (eq 14) is correct. Observation of R and β close to 0.5 fixes the rate-limiting step as being I2 + e- h I•(ads) + I-. The observed first-order kinetics for the overall process of 2I- - 2e- f I2 then suggests that the equation I- - e- h I•(ads) lies to the right at potentials corresponding to the I2|I- voltammetric waves so that the surface coverage of I•(ads) is saturated. Thus, the reaction I•(ads) + I+ e- f I2 is simply first order in [I-]. We note the work of Hubbard and co-workers33 in which the high affinity of platinum surfaces for I• is established. We consider that the kinetics of the process are given by F
rate ⁄ mol cm-2 s-1 ) 333.05e-0.5 RT E[I2] F
6.22 × 10-9e+0.5 RT E[I-] 5. Conclusions A kinetic and mechanistic study has been undertaken on iodine, I2, in the RTIL [C4mim][NTf2]. The fast, two-electron reduction of I2 has been investigated by cyclic voltammetry at a platinum microelectrode. Two reduction peaks were observed and assigned to reduction of iodine to triiodide, I3 , followed by further reduction to iodide, I-. A digital simulation program was developed based on an EC mechanism, which was used to simulate the voltammetric response and successfully identify the electrode reaction mechanism and extract kinetic parameters. Acknowledgment. E.I.R. and I.S. thank the EPSRC, and L.A. thanks the Department of Education and Learning in Northern Ireland and Merck GmBH for financial support.
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