Electrochemistry of Iodide, Iodine, and Iodine ... - ACS Publications

Dec 27, 2015 - Cameron L. Bentley†‡, Alan M. Bond†, Anthony F. Hollenkamp‡, Peter ... David Hausmann , Ralf Köppe , Silke Wolf , Peter W. Roe...
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Electrochemistry of Iodide, Iodine, and Iodine Monochloride in Chloride Containing Nonhaloaluminate Ionic Liquids Cameron L. Bentley,†,‡ Alan M. Bond,† Anthony F. Hollenkamp,*,‡ Peter J. Mahon,§ and Jie Zhang*,† †

School of Chemistry, Monash University, Clayton, VIC 3800, Australia CSIRO Energy, Private Bag 10, Clayton South, VIC 3169, Australia § Faculty of Science, Engineering and Technology, Swinburne University of Technology, Hawthorn, VIC 3122, Australia ‡

S Supporting Information *

ABSTRACT: The electrochemical behavior of iodine remains a contemporary research interest due to the integral role of the I−/I3− couple in dye-sensitized solar cell technology. The neutral (I2) and positive (I+) oxidation states of iodine are known to be strongly electrophilic, and thus the I−/I2/I+ redox processes are sensitive to the presence of nucleophilic chloride or bromide, which are both commonly present as impurities in nonhaloaluminate room temperature ionic liquids (ILs). In this study, the electrochemistry of I−, I2, and ICl has been investigated by cyclic voltammetry at a platinum macrodisk electrode in a binary IL mixture composed of 1-butyl-3methylimidazolium chloride ([C4mim]Cl) and 1-ethyl-3-methylimidazolium bis(trifluoromethanesulfonyl)imide ([C2mim][NTf2]). In the absence of chloride (e.g., in neat [C2mim][NTf2]), I− is oxidized in an overall one electron per iodide ion process to I2 via an I3− intermediate, giving rise to two resolved I−/I3− and I3−/I2 processes, as per previous reports. In the presence of low concentrations of chloride ([Cl−] and [I−] are both 0.9 V vs saturated calomel electrode in aqueous media).1,15,16 The presence of nucleophilic halides (e.g., Cl−), pseudohalides (e.g., CN−), and amines (e.g., pyridine) can stabilize electrophilic I2 and In+, and thus have a profound effect on the I− oxidation process.1 For example, in aqueous HCl, I− can be oxidized to form interhalogen compounds, such as ICl, [I2Cl]−,

I− is a nucleophile (Lewis base) and I2 is an electrophile (Lewis acid), and these species can combine homogeneously to form the polyhalogen complex anion, tri-iodide: (2)

The driving force for the formation of tri-iodide (i.e., Kstab) is quantified by the equilibrium constant of the reaction given in eq 2. Kstab is extremely sensitive to the donor−acceptor properties of the solvent13 and ranges from ca. 103 in water to ca. 107 in acetonitrile.1 Under conditions where the formation of I3− is extremely favorable (large Kstab and/or high concentrations of I− or I2), the overall I−/I2 process shown in eq 1 occurs in two resolved steps under voltammetric conditions: © 2015 American Chemical Society

(3)

I3− ⇌

(1)

I− + I 2 ⇌ I3−

3I− ⇌ I3− + 2e−

Received: November 16, 2015 Accepted: December 26, 2015 Published: December 27, 2015 1915

DOI: 10.1021/acs.analchem.5b04332 Anal. Chem. 2016, 88, 1915−1921

Article

Analytical Chemistry and [ICl2]−.15,17 Similar observations have been made when iodide is oxidized in the presence of Cl− in acetonitrile7 and nitromethane.18 Clearly, such interactions cannot be avoided when, for example, Cl− is an intrinsic part of the electrolyte media as in molten chloride salts19,20 and haloaluminate room temperature ionic liquids.21,22 The interhalogens are themselves an interesting class of compounds, which are used in synthetic chemistry as strong oxidizing and halogenating agents.23 Interhalogen compounds are usually prepared directly from their constituent halogens, and a number of them have been characterized electrochemically (e.g., [IClx]1−x) in aqueous17 and nonaqueous24 media. Over the past 2 decades, nonhaloaluminate room temperature ionic liquids (ILs) have attracted significant interest in academia and industry as replacements for volatile molecular solvents in a range of applications.25,26 A promising application of ILs is as the electrolyte in DSSCs,3,27 and for this reason, the electrochemistry of the I−/I2 couple has been the subject of many investigations in these neoteric solvents.28−33 In a “typical” aprotic IL, such as 1-ethyl-3-methylimidazolium bis(trifluoromethanesulfonyl)imide ([C2mim][NTf2]), the I−/ I2 redox process is analogous to that in acetonitrile,7−12 occurring in two resolved steps (i.e., eqs 3 and 4), via an I3− intermediate. It should be noted that although the most commonly employed IL anions (e.g., [PF6 ]− , [BF 4] −, [CF3SO3]−, and [NTf2]−) are weak Lewis bases34 (i.e., “noncoordinating”), ILs containing nucleophilic Cl− or Br− as the constituent anion can be prepared.35 In addition, ILs are often prepared through the metathesis of a halide salt,25 and thus may contain residual Cl− or Br−. To date, studies on the electrochemistry of I−/I2 couple have yet to be performed in Cl− containing nonhaloaluminate ILs, and so the electrochemical formation of In+ species has yet to be reported in this type of media. Thus, in this work, we investigate the electrochemistry of iodide, iodine, and iodine monochloride at a platinum macrodisk electrode in a chloride containing nonhaloaluminate ionic liquid using voltammetric methodology.

minimize noise in all microelectrode experiments. Positive feedback iRu compensation (Ru = uncompensated resistance) was employed in macroelectrode experiments (Ru was estimated by electrochemical impedance spectroscopy). All voltammetric experiments were carried out using a standard three-electrode arrangement with a working and reference electrode as described below and a Pt wire auxiliary electrode. An Ag wire which had been immersed in the IL under investigation and sealed in a fritted (Vycor glass) glass tube served as the pseudo reference electrode. The pseudo reference electrode potential was calibrated against the formal potential of the IUPAC recommended Fc/Fc+ process36 in the electrolyte of interest, taking into careful consideration the difference in the diffusion coefficients of Fc and Fc+.37,38 The Pt macrodisk working electrode with a nominal diameter of 1.6 mm was purchased from BASi (Bioanalytical Systems, West Lafayette, IN, USA) and the Pt microdisk with a nominal diameter of 20 μm was purchased from Metrohm (Herisau, Switzerland). The Pt macrodisk electrode was activated by polishing with 1 μm and then 0.3 μm aqueous alumina slurries (KEMET, Kent, U.K.) on a clean polishing cloth (Buehler, Lake Bluff, IL, USA). Adherent alumina was removed by sonication in deionized water. The Pt microdisk electrode was activated by polishing with an aqueous slurry of 0.3 μm alumina and rinsed thoroughly with deionized water. Prior to experimentation, the relevant electrodes were preconditioned in 0.1 M sulfuric acid by scanning between the oxygen and hydrogen evolution reactions39 with subsequent rinsing in deionized water and acetone. The active electrode area (A) of each of the electrodes was calibrated with convolution voltammetry,30,40,41 using the oxidation of a 2.0 mM Fc solution in acetonitrile (0.10 M [NBu4][PF6]) and adopting a diffusion coefficient of 2.4 × 10−5 cm2 s−1, as published under these conditions.42 Data Analysis. The algorithm used to calculate the semiintegrated or convolved currents has been reported previously.41,43 The diffusion coefficient (D) of iodide and the stoichiometric number of electrons (n) were calculated simultaneously using chronoamperometry at a microdisk electrode (“Shoup-Szabo method”), as reported by Compton and co-workers.32,37 Background correction was achieved by measuring the voltammetric response in the absence of electroactive species and subtracting the residual current from the total current.

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EXPERIMENTAL SECTION Reagents. 1-Ethyl-3-methylimidazolium bis(trifluoromethanesulfonyl)imide ([C2mim][NTf2], Io-li-tec, Heilbronn, Germany) and 1-butyl-3-methylimidazolium chloride ([C4mim]Cl, Sigma-Aldrich, St. Louis) were commercial samples. Before use, each IL was dried under high vacuum (≤10−2 mbar) at 45 °C for at least 48 h. 1-Ethyl-3methylimidazolium iodide ([C2mim]I, Io-li-tec) was recrystallized twice from a 2:1 mixture of ethyl acetate (EMSURE, Merck, Darmstadt, Germany) and isopropanol (EMSURE, Merck) and then dried under high vacuum prior to use. Care was taken during handling and storage of [C2mim]I to avoid exposure to light. Ferrocene (Fc, Sigma-Aldrich), iodine (I2, Sigma-Aldrich), and iodine monochloride (ICl, Merck) were used as supplied by the manufacturer. All water/oxygensensitive reagents were stored and handled under a dry argon atmosphere in a glovebox. Electrochemistry. Unless otherwise stated, all voltammetric experiments were carried out under benchtop conditions at ambient temperature (24 ± 1 °C) with a Gamry Reference 600 Potentiostat/Galvanostat/ZRA (Gamry Instruments, Warminster, PA, USA). All solvents were initially degassed with N2 and then maintained under a blanket of N2 during the course of the voltammetric experiments. A Faraday cage was employed to

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RESULTS AND DISCUSSION Iodide Electro-oxidation in the Presence of Millimolar Chloride Concentrations. As noted at the beginning of this article, chloride, bromide, and iodide are common impurities in nonhaloaluminate room temperature ILs. These halide byproducts derived from the starting materials used in IL synthesis are known to significantly influence voltammetric data. For instance, Silvester and Compton44 have reported “halide-induced gold stripping” when voltammetry is carried out in chloride-contaminated ILs at a gold electrode. To investigate the influence that halide impurities (i.e., Cl− or Br−) can have on iodine electrochemistry, the voltammetric iodide oxidation response was initially investigated in halide-free [C2mim][NTf2], and then after controlled quantities of Cl− (from [C4mim]Cl) had been introduced. Cyclic voltammograms obtained from this series of experiments are shown in Figure 1a. In the absence of Cl−, the oxidation of I− is consistent with previous reports,28−33 occurring in two steps, 1916

DOI: 10.1021/acs.analchem.5b04332 Anal. Chem. 2016, 88, 1915−1921

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Analytical Chemistry I3− + xCl− ⇌

3 [I 2Cl 0.67x]0.67x − + e− 2

(5)

The second is that Cl− promotes the further oxidation of I− to I+ (via I3−): I3− + yCl− ⇌ 3[ICl 0.33y]1 − 0.33y + 4e−

(6)

If process (iii)/(iii′) occurs as in eq 5, the total charge (Q) passed during the forward voltammetric sweep should remain constant since it is still an overall one electron per iodide ion process (Q ∝ n). However, if (iii)/(iii′) is due to eq 6, the total charge should increase since, overall, an additional electron is lost per iodide ion oxidized. To distinguish between these two possibilities, the forward sweep of each of the cyclic voltammograms shown in Figure 1a was semi-integrated, as is shown in Figure 1b. The semi-integrals of the voltammograms superficially resemble steady-state curves,43 with two well-defined plateaux evident in the absence of Cl−, corresponding to the I−/I3− (see eq 3) and I3−/I2 processes (see eq 4) at lower and higher potentials, respectively. Upon addition of Cl−, the following occurs: process (i) remains unchanged, qualitatively indicating that the viscosity of the IL is unaltered (see below); process (ii) decreases in magnitude; process (iii) increases in magnitude. The magnitude of the limiting semi-integral current plateau, ML, is given by41,43

ML = nFAD1/2C*

(7)

where n is the stoichiometric number of electrons, F is Faraday’s constant, A is the active electrode area, D is the diffusion coefficient, and C* is the bulk concentration. In Figure 1b, since the addition of millimolar chloride does not appear to change the viscosity of the IL (see above), then A and D can be assumed to be constant. In addition, since the voltammograms have been corrected for changes in [I−] (i.e., C*), the increase in ML with [Cl−] must be due to an increase in n (i.e., the apparent number of electrons lost per iodide ion increases with increasing chloride concentration). This observation implies that process (iii)/(iii′) corresponds to the reaction given in eq 6 and that the presence Cl− leads to a transition to an overall two-electron process, through stabilization of I+ by reaction with Cl−:

Figure 1. Concentration-normalized (a) cyclic and (b) semi-integral voltammograms obtained in [C2mim][NTf2] at a 1.6 mm diameter Pt macrodisk electrode with a scan rate of 50 mV s−1 from the electrooxidation of iodide (approximately 25 mM, from [C2mim]I) in the presence of chloride (from [C4mim]Cl). The increase in [Cl−] progresses from 0, 0.81, 2.7, 4.6, 6.3, 9.9, 12, 16, 19, and 24 mM, respectively. Thus, the [Cl−]:[I−] ratio increases from approximately 0, 0.031, 0.11, 0.18, 0.25, 0.41, 0.51, 0.66, 0.78, and 1.0, respectively.

I− → I+ + 2e−

(8)

I+ + zCl− → [ICl z]1 − z

(9)

If the complexation reaction shown in eq 9 goes to completion on the voltammetric time scale, eq 6 would be limited by the [Cl−] when a substoichiometric amount of chloride is present. This must be the case in Figure 1b, as the voltammograms depend on the concentration of chloride (explored below). Furthermore, since processes (ii) and (iii) are both observed at a 1:1 [Cl−]:[I−] ratio, the value of “z” in eq 9 must be greater than 1, as only process (iii) would be expected in the presence of a stoichiometric amount of Cl−. Since I− is oxidized in two parallel pathways with different electron stoichiometries (see eqs 1 and 8) under the conditions outlined in Figure 1, the “apparent” stoichiometric number of electrons, napp, calculated using eq 7 from the ML plateaux in Figure 1b, depends on the proportion of iodide entering each pathway. If it is assumed that the two-electron pathway (e.g., eq 8) is limited by the amount of chloride present (see above), the following relationship can be derived:

where the I−/I3− and I3−/I2 processes (see eqs 3 and 4) are labeled (i)/(i′) and (ii)/(ii′), respectively. Introducing Cl− into the IL causes an additional, chemically reversible process [labeled (iii)/(iii′) in Figure 1a] to appear at potentials between processes (i)/(i′) and (ii)/(ii′). The direct oxidation of Cl− (to Cl2) occurs at more positive potentials,45 so the additional process is postulated to arise from interaction between Cl− and the oxidation products of I−. Since the presence of deliberately added Cl− does not affect process (i)/ (i′), it is likely that I3− remains the product of the first oxidation (see below). There are two plausible explanations for process (iii)/(iii′). The first is that Cl− stabilizes I2 by complexation, allowing the oxidation of I3− to occur at a less positive potential: 1917

DOI: 10.1021/acs.analchem.5b04332 Anal. Chem. 2016, 88, 1915−1921

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Analytical Chemistry napp =

[Cl−] +1 z[I−]

(10)

where “z” is the stoichiometric coefficient of Cl−, as shown in eq 9. Substituting eq 10 into eq 7 and rearranging gives ML FAD1/2[Cl−] = + FAD1/2 [I−] z[I −]

(11)

Thus, as shown in Figure S1 of the Supporting Information, a plot of ML/[I−] vs [Cl−]/[I−] is a straight line (r2 = 0.998) with a slope of FAD1/2/z and an intercept of FAD1/2. The value of FAD1/2 was calculated to be 1.1 × 10−3 C cm2 mol−1 s−1 from the intercept. Substitution into the estimated value of the slope (from Figure S1) gives a value for “z” of 2. Therefore, the overall reaction (e.g., see eqs 8 and 9) can be written as follows: I− + 2Cl− → [ICl 2]− + 2e−

(12)

and eq 6 can be rewritten as I3− + 6Cl− ⇌ 3[ICl 2]− + 4e− −

(13) 23

[ICl2] is a well-known interhalogen anion that is formed during the oxidation of iodide in the presence of excess chloride in both aqueous17 and nonaqueous22,24 media. Iodine Electrochemistry in the Presence of Molar Chloride Concentrations. In the above section, the concentrations of Cl− and I− are comparable. This means that the concentration of Cl− in the diffusion layer adjacent to the electrode surface varies during the voltammetric sweep, and that the reaction shown in eq 12 is limited by the amount of Cl− present. In order to simplify the situation, so [Cl−]surf is effectively constant (and equal to [Cl−]bulk) during the voltammetric sweep, the oxidation of I− can be undertaken in the presence of a large excess of Cl−. As noted at the beginning of this article, room temperature ILs containing Cl− as the constituent anion are available.35 However, in practice, these ILs are highly viscous, with even the most fluid chloride containing IL, 1-hexyl-3-methylimidazolium chloride, [C6mim]Cl, having a viscosity of 6416 mPa·s at 30 °C.35 For this reason, the chloride IL, [C4mim]Cl, was mixed in a 2.5:1 mole ratio with [C2mim][NTf2], to give a binary mixture which is relatively fluid (compared to [C6mim]Cl) at room temperature. On the basis of the density of the binary mixture at room temperature of 1.25 g cm−3, the concentration of chloride is calculated to be 3.8 M, which is slightly lower than that present in neat [C6mim]Cl, 5.1 M. However, the behavior of the I−/I2/ I+ redox system in the [C4mim]Cl/[C2mim][NTf2] mixture (2.5:1 mole ratio) should closely resemble that in neat [Cxmim]Cl ionic liquid. Cyclic and semi-integral voltammograms obtained from the oxidation of iodide in the binary [C4mim]Cl/[C2mim][NTf2] mixture are displayed in Figure 2. Five peaks/processes, labeled (I−V), are identified on the cyclic voltammogram shown in Figure 2a. Results from additional sweeps over a range of scan rates are included in the Supporting Information (see Figure S2). Following analysis of data by the Shoup-Szabo chronoamperometric method outlined in the Experimental Section, the overall number of electrons and diffusivity of I− were calculated to be 2 and 2 × 10−8 cm2 s−1, respectively (see Figure S3). Thus, the overall redox process under investigation can be written as −

+

I ⇌ I + 2e

−

Figure 2. Background-subtracted (a) cyclic and (b) semi-integral voltammogram obtained at a 1.6 mm diameter Pt macrodisk electrode with scan rates of 25 and 10 mV s−1 in (a) and (b), respectively, for the electro-oxidation of 10.8 mM I− (from [C2mim]I) in a mixture of [C4mim]Cl and [C2mim][NTf2] (2.5:1 mole ratio).

I+ is strongly electrophilic, and hence would be expected to be complexed by nucleophilic Cl−. From the positive potential sweep of the cyclic voltammogram (Figure 2a) and the semiintegral voltammogram (Figure 2b), it is clear that the oxidation of I− to I+ (see eq 11) occurs in two steps, (I) and (II), via an intermediate. As shown in Figure 2b, the ML plateaux associated with processes (I) and (II) are equal in magnitude; there is an inflection point halfway up the wave at −30 mV vs Fc/Fc+ consistent with an electron stoichiometry of 1:1. This observation suggests that the intermediate species has an oxidation state of zero (as in I2), and therefore (I) is assigned to the oxidation of I− to I2, followed by complexation with Cl− via an EC mechanism to form an interhalide species: 2I− → I 2 + 2e− −

(15) −

I 2 + Cl ⇌ [I 2Cl]

(16)

giving the overall process: 2I− + Cl− → [I 2Cl]− + 2e−

(14) 1918

(17) DOI: 10.1021/acs.analchem.5b04332 Anal. Chem. 2016, 88, 1915−1921

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Analytical Chemistry Tri-iodide (I3−) is extremely stable in neat [C2mim][NTf2] (Kstab ≈ 106.4)31 and is expected to be more stable than [I2Cl]− in the [C4mim]Cl/[C2mim][NTf2] mixture.22 It is therefore plausible that I3− is initially formed in process (I) [as per eq 3]. However, as I− becomes depleted in the diffusion layer, the formation of the I3− polyhalide would be compromised by the large concentration of excess chloride present: I3− + Cl− ⇌ I 2Cl− + I−

(18)

The proposed reaction (see eq 17) is consistent with the 1:1 electron stoichiometry observed for processes (I) and (II) in Figure 2b; if I3− were the major species present, the ratio would be closer to 1:2. It follows from eq 17 that process (II) can be assigned to the oxidation of the intermediate produced in (I) to I+, coupled by complexation with Cl− in an EC mechanism, to form a different interhalide species: [I 2Cl]− → 2I+ + Cl− + 2e−

(19)

2I+ + 4Cl− → 2[ICl 2]−

(20)

giving the overall process: [I 2Cl]− + 3Cl− → 2[ICl 2]− + 2e− −

(21) −

In the above equations, [I2Cl] and [ICl2] are assumed to be the intermediate and final product, respectively. Both are known interhalogen compounds23 and have been postulated in the oxidation of iodide in both aqueous HCl solutions15,17 and basic chloroaluminate ionic liquid melts.22 As noted above, [ICl2]− is the stable I+ species formed when the concentrations of I− and Cl− are comparable (see Figure 1). Equations 17 and 21 have an electron stoichiometry of 1:1, in agreement with the magnitude of the ML plateaux in Figure 2b, and in combination give the overall reaction shown in eq 12. Proceeding from I− to the final product, [ICl2]−, via the mechanism proposed represents an ECEC process. To assign reduction processes (III) to (V), the electroreduction of I+ (from iodine monochloride, ICl) was investigated in the binary [C4mim]Cl/[C2mim][NTf2] IL mixture; a representative cyclic voltammogram is shown in Figure 3a. Upon dissolution in this medium, ICl should be complexed by Cl− to form [ICl2]−, as in eq 20. [ICl2]− gives rise to a single, broad reduction process, coincidental with process (V) in Figure 2a, and which is assigned to the reaction: 2[ICl 2]− + 4e− → 2I− + 4Cl−

Figure 3. Background-subtracted cyclic voltammograms obtained at a 1.6 mm diameter Pt macrodisk electrode from (a) 16.0 mM ICl and (b) 8.1 mM I2 in a mixture of [C4mim]Cl and [C2mim][NTf2] (2.5:1 mole ratio). The scan rates in (a) and (b) are 25 and 10 mV s−1, respectively. In (b), individual sweeps were performed toward negative (---) and positive potentials (), starting from the open-circuit potential (−0.124 V vs Fc/Fc+).

(22)

oxidized to [ICl2]− at the diffusion-controlled rate (see eq 12), giving rise to depletion of I− in the vicinity of the electrode surface. In the I− depleted zone, [ICl2]− is stable. However, as it diffuses away from the electrode surface, it encounters I− and is able to undergo comproportionation:

Two oxidation processes also are observed on the reverse sweep (toward positive potentials), which correspond to (I) and (II), adding confidence to the assignment of these processes in eqs 17 and 21, respectively. The fact that processes (III) and (IV) are not present on the initial sweep toward negative potentials in Figure 3a implies that the process shown in eq 21 is not reversible under the experimental conditions; that is, [ICl2]− cannot be reduced via the [I2Cl]− intermediate. This implies that the chemical step shown in eq 20 is essentially irreversible on the voltammetric time scale. Referring back to the oxidation of I− in Figure 2, it may be proposed that during the course of the voltammetric sweep the intermediate species [I2Cl]− formed electrochemically in process (I) would be further oxidized to [ICl2]− in process (II). Therefore, the electroactive intermediates giving rise to processes (III) and (IV) in Figure 2 must have been formed as a consequence of follow-up chemical reactions. For example, at potentials positive of (II) (e.g., >70 mV vs Fc/Fc+), I− is

[ICl 2]− + I− ⇌ [I 2Cl]− + Cl− −

−

(23)

I3−

[I2Cl] can react further with I , forming as per eq 18. Both [I2Cl]− and I3− can subsequently diffuse to the electrode surface and undergo reduction, giving rise to processes (III) and (IV), respectively: [I 2Cl]− + 2e− → 2I− + Cl−

(24)

I3− + 2e− → 3I−

(25)

Equation 25 was assigned to process (IV), because of its coincidence with the tri-iodide reduction process observed in neat [C2mim][NTf2] (see Figure 1). The observations in this 1919

DOI: 10.1021/acs.analchem.5b04332 Anal. Chem. 2016, 88, 1915−1921

Analytical Chemistry

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IL study are consistent with those of Karpinski and Osteryoung22 in basic chloroaluminate ionic liquid melts. To complete this study, the electrochemical behavior of I2 was investigated in the [C4mim]Cl/[C2mim][NTf2] mixture. The initial cyclic voltammetric scans are shown in Figure 3b, and the results for multiple cycles are displayed in Figures S4 and S5. Upon immersion in the I2/[C4mim]Cl/[C2mim][NTf2] mixture, the Pt working electrode adopted an open circuit (equilibrium) potential of −0.124 V vs Fc/Fc+. Starting from this potential, processes (I)/(II)/(III) and (IV)/(V) are evident on both the positive and negative potential direction sweeps, respectively, indicating that I−, I3−, and [ICl2]− are all present. It is likely that, in the [C4mim]Cl/[C2mim][NTf2] mixture, I2 (initially complexed by Cl−, see eq 16) disproportionates, forming I− and [ICl2]−, with the former species reacting with I2 to form I3−. The ratio of [I2Cl]−, I−, I3−, and [ICl2]− should be composition-dependent, and easily varied by changing the [C4mim]Cl:[C2mim][NTf2] ratio in the mixture. In summary, under voltammetric conditions in the presence of large excess concentrations of Cl−, I− is oxidized at a Pt electrode in two single-electron steps (see eqs 17 and 21) to [ICl2]−, predominantly via an [I2Cl]− intermediate. On reversal of the potential scan direction, [ICl2]− is reduced directly to I− in a single, two-electron process (see eq 22). [ICl2]−, on diffusing away from the electrode surface, undergoes a comproportionation reaction with I− in bulk solution, forming [I2Cl]− (eq 23) and I3− (eq 18), which diffuse in from bulk solution and are reduced to I− at potentials more positive than [ICl2]−, as shown in eqs 24 and 25, respectively. Clearly, the electrochemical behavior of iodine in the presence of chloride involves the complicated interplay of multiple electron transfer (e.g., eqs 17 and 21) and various coupled chemical steps (e.g., eqs 16, 18, 20, and 23) which may or may not be at equilibrium on the voltammetric time scale.

Article

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.analchem.5b04332. Linear plot of ML/[I−] vs [Cl−]/[I−] (Figure S1), cyclic voltammograms obtained from the electro-oxidation of iodide in [C4mim]Cl/[C2mim][NTf2] (Figure S2), experimental and Shoup-Szabo theoretical chronoamperograms obtained from the diffusion-controlled oxidation of iodide in [C4mim]Cl/[C2mim][NTf2] (Figure S3), and cyclic voltammograms obtained from I2 in [C4mim]Cl/[C2mim][NTf2] (Figures S4 and S5) (PDF)

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AUTHOR INFORMATION

Corresponding Authors

*E-mail: [email protected] (A.F.H). *E-mail: [email protected] (J.Z.). Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS C.L.B. acknowledges financial support received from a Monash University Faculty of Science Postgraduate Publication Award. REFERENCES

(1) Bard, A. J.; Lund, H. Encyclopedia of Electrochemistry of the Elements; Marcel Dekker: New York, 1973. (2) Boschloo, G.; Hagfeldt, A. Acc. Chem. Res. 2009, 42, 1819−1826. (3) Hagfeldt, A.; Boschloo, G.; Sun, L. C.; Kloo, L.; Pettersson, H. Chem. Rev. 2010, 110, 6595−6663. (4) Dane, L. M.; Janssen, L. J. J.; Hoogland, J. G. Electrochim. Acta 1968, 13, 507−518. (5) Swathirajan, S.; Bruckenstein, S. J. Electroanal. Chem. Interfacial Electrochem. 1981, 125, 63−71. (6) Swathirajan, S.; Bruckenstein, S. J. Electroanal. Chem. Interfacial Electrochem. 1983, 143, 167−178. (7) Popov, A. I.; Geske, D. H. J. Am. Chem. Soc. 1958, 80, 1340− 1352. (8) Nelson, I. V.; Iwamoto, R. T. J. Electroanal. Chem. 1964, 7, 218− 221. (9) Dryhurst, G.; Elving, P. J. Anal. Chem. 1967, 39, 606−615. (10) Macagno, V. A.; Giordano, M. C.; Arvia, A. J. Electrochim. Acta 1969, 14, 335−357. (11) Sereno, L.; Macagno, V. A.; Giordano, M. C. Electrochim. Acta 1972, 17, 561−575. (12) Magno, F.; Mazzocchin, G.-A.; Bontempelli, G. J. Electroanal. Chem. Interfacial Electrochem. 1973, 47, 461−468. (13) Iwamoto, R. T. Anal. Chem. 1959, 31, 955−955. (14) Pourbaix, M. Atlas d'equilibres electrochimiques; Gauthier-Villars et Cie: Paris, 1963. (15) Dryhurst, G.; Elving, P. J. Anal. Chem. 1967, 39, 606−615. (16) Miller, F. J.; Zittel, H. E. J. Electroanal. Chem. 1966, 11, 85−93. (17) Kolthoff, I. M.; Jordan, J. J. Am. Chem. Soc. 1953, 75, 1571− 1575. (18) Voorhies, J. D.; Schurdak, E. J. Anal. Chem. 1962, 34, 939−943. (19) Marassi, R.; Mamantov, G.; Chambers, J. Q. Inorg. Nucl. Chem. Lett. 1975, 11, 245−252. (20) Marassi, R.; Chambers, J. Q.; Mamantov, G. J. Electroanal. Chem. Interfacial Electrochem. 1976, 69, 345−359. (21) Karpinski, Z. J.; Osteryoung, R. A. J. Electroanal. Chem. Interfacial Electrochem. 1984, 164, 281−298. (22) Karpinski, Z. J.; Osteryoung, R. A. J. Electroanal. Chem. Interfacial Electrochem. 1984, 178, 281−294.

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CONCLUSIONS The electrochemistry of I−, I2, and ICl has been investigated by voltammetric methodology at a platinum macrodisk electrode in a binary mixture of [C4mim]Cl and [C2mim][NTf2]. In neat [C2mim][NTf2], as per previous reports, I− is oxidized in an overall one electron per iodide ion process to I2 via an I3− intermediate, giving rise to two resolved I−/I3− and I3−/I2 processes at less and more positive potentials, respectively. In the presence of millimolar concentrations of chloride and with [Cl−] comparable to that of [I−], a new oxidation process appears at potentials less positive than that for the I3−/I2 process and corresponds to the oxidation of I3− to the interhalide complex anion, [ICl2]−, in an overall two electron per iodide ion process. In the presence of a 300-fold or greater concentration excess of Cl−, I− was found to be oxidized in an overall two electron per iodide ion process to [ICl2]− via an [I2Cl]− intermediate, as confirmed by investigating the voltammetric response of ICl and I2. Overall, the electrochemical behavior of iodine in chloride containing nonhaloaluminate ILs involves a complicated interplay between multiple electron transfer pathways and homogeneous chemical reactions which may or may not be at equilibrium on the voltammetric time scale. The interhalides identified in this study, ICl, [I2Cl]−, and [ICl2]−, are well-known species and have been shown to form during the electro-oxidation of iodide in chloride containing aqueous, nonaqueous, molten salt, and haloaluminate IL media. 1920

DOI: 10.1021/acs.analchem.5b04332 Anal. Chem. 2016, 88, 1915−1921

Article

Analytical Chemistry (23) Surles, T. Interhalogen Compounds. AccessScience. McGraw-Hill Education: Columbus, OH, 2014. http://www.accessscience.com. Accessed: 16 Nov 2015. (24) Popov, A. I.; Geske, D. H. J. Am. Chem. Soc. 1958, 80, 5346− 5349. (25) Welton, T. Chem. Rev. 1999, 99, 2071−2083. (26) Galinski, M.; Lewandowski, A.; Stepniak, I. Electrochim. Acta 2006, 51, 5567−5580. (27) Armand, M.; Endres, F.; MacFarlane, D. R.; Ohno, H.; Scrosati, B. Nat. Mater. 2009, 8, 621−629. (28) Bentley, C. L.; Bond, A. M.; Hollenkamp, A. F.; Mahon, P. J.; Zhang, J. Electrochim. Acta 2013, 109, 554−561. (29) Bentley, C. L.; Bond, A. M.; Hollenkamp, A. F.; Mahon, P. J.; Zhang, J. Anal. Chem. 2013, 85, 11319−11325. (30) Bentley, C. L.; Bond, A. M.; Hollenkamp, A. F.; Mahon, P. J.; Zhang, J. J. Phys. Chem. C 2014, 118, 29663−29673. (31) Bentley, C. L.; Bond, A. M.; Hollenkamp, A. F.; Mahon, P. J.; Zhang, J. J. Phys. Chem. C 2015, 119, 22392−22403. (32) Rogers, E. I.; Silvester, D. S.; Aldous, L.; Hardacre, C.; Compton, R. G. J. Phys. Chem. C 2008, 112, 6551−6557. (33) Zhang, Y.; Zheng, J. B. Electrochim. Acta 2007, 52, 4082−4086. (34) Schmeisser, M.; Illner, P.; Puchta, R.; Zahl, A.; van Eldik, R. Chem. - Eur. J. 2012, 18, 10969−10982. (35) Seddon, K. R.; Stark, A.; Torres, M. J. In Clean Solvents: Alternative Media for Chemical Reactions and Processing; Abraham, M. A., Moens, L., Eds.; American Chemical Society: Washington, DC, 2002; pp 34−49. (36) Gritzner, G.; Kuta, J. Pure Appl. Chem. 1984, 56, 461−466. (37) Rogers, E. I.; Silvester, D. S.; Poole, D. L.; Aldous, L.; Hardacre, C.; Compton, R. G. J. Phys. Chem. C 2008, 112, 2729−2735. (38) Meng, Y.; Aldous, L.; Belding, S. R.; Compton, R. G. Phys. Chem. Chem. Phys. 2012, 14, 5222−5228. (39) Swain, G. M. In Handbook of Electrochemistry; Zoski, C. G., Ed.; Elsevier: Amsterdam, 2007; pp 111−153. (40) Mahon, P. J.; Oldham, K. B. J. Electroanal. Chem. 1999, 464, 1− 13. (41) Bentley, C. L.; Bond, A. M.; Hollenkamp, A. F.; Mahon, P. J.; Zhang, J. Anal. Chem. 2014, 86, 2073−2081. (42) Bard, A. J.; Faulkner, L. R. Electrochemical Methods: Fundamentals and Applications, 2nd ed.; Wiley: New York, 2001; p 833. (43) Bentley, C. L.; Bond, A. M.; Hollenkamp, A. F.; Mahon, P. J.; Zhang, J. Anal. Chem. 2013, 85, 2239−2245. (44) Silvester, D. S.; Compton, R. G. Z. Phys. Chem. 2006, 220, 1247−1274. (45) Huang, X. J.; Silvester, D. S.; Streeter, I.; Aldous, L.; Hardacre, C.; Compton, R. G. J. Phys. Chem. C 2008, 112, 19477−19483.

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DOI: 10.1021/acs.analchem.5b04332 Anal. Chem. 2016, 88, 1915−1921