Electrochemistry of Room Temperature Protic Ionic Liquids - The

J. Phys. Chem. B 2008, 112, 6923–6936

6923

Electrochemistry of Room Temperature Protic Ionic Liquids Chuan Zhao,†,‡ Geoff Burrell,§ Angel A. J. Torriero,†,‡ Frances Separovic,*,§ Noel F. Dunlop,| Douglas R. MacFarlane,†,‡ and Alan M. Bond*,†,‡ School of Chemistry, Monash UniVersity, Clayton, Victoria 3800, Australia, ARC Special Research Center for Green Chemistry, Monash UniVersity, Clayton, Victoria 3800, Australia, School of Chemistry, Bio21 Institute, UniVersity of Melbourne, ParkVille, Victoria 3010, Australia, and Orica Ltd, Melbourne, Victoria 3000, Australia ReceiVed: December 16, 2007; ReVised Manuscript ReceiVed: February 28, 2008; In Final Form: March 11, 2008

Eighteen protic ionic liquids containing different combinations of cations and anions, hydrophobicity, viscosity, and conductivity have been synthesized and their physicochemical properties determined. In one series, the diethanolammonium cations were combined with acetate, formate, hydrogen sulfate, chloride, sulfamate, and mesylate anions. In the second series, acetate and formate anions were combined with amine bases, triethylamine, diethylamine, triethanolamine, di-n-propylamine, and di-n-butylamine. The electrochemical characteristics of the eight protic ionic liquids that are liquid at room temperature (RTPILs) have been determined using cyclic, microelectrode, and rotating disk electrode voltammetries. Potential windows of the RTPILs have been compared at glassy carbon, platinum, gold, and boron-doped diamond electrodes and generally found to be the largest in the case of glassy carbon. The voltammetry of IUPAC recommended potential scale reference systems, ferrocene/ferrocenium and cobaltocenium/cobaltocene, have been evaluated and found to be ideal in the case of the less viscous RTPILs but involve adsorption in the highly viscous ones. Other properties such as diffusion coefficients, ionic conductivity, and double layer capacitance also have been measured. The influence of water on the potential windows, viscosity, and diffusion has been studied systematically by deliberate addition of water to the dried ionic liquids. The survey highlights the problems with voltammetric studies in highly viscous room temperature protic ionic liquids and also suggests the way forward with respect to their possible industrial use. Introduction Ionic liquids (ILs) are a class of solvents that typically have negligible vapor pressure and good thermal and chemical stability. If they also are environmental friendly, they may provide an attractive substitute for conventional organic solvents.1,2 ILs are generally divided into two categories, aprotic ionic liquids (AILs) and protic ionic liquids (PILs).1–4 Protic ionic liquids are formed by a variety of proton transfer and association equilibria between neat Bronsted acids and bases.4–11 However, protic ionic liquids may contain neutral species produced by ubiquitous proton transfer equilibrium processes and, therefore, are not necessarily comprised entirely of ionic components. Recently, MacFarlane and Seddon proposed a guideline that less than 1% of the neutral species should be present if the ionic liquid is to be considered a “pure ionic liquid”.7 The very first PIL, to the best of our knowledge, was ethanolammonium nitrate reported by Gabriel in 1888,5 showing a melting point of 52-55 °C. In 1914, Walden reported the synthesis of ethylammonium nitrate by transferring a proton from a Bronsted acid to a Bronsted base without the use of any solvent.6 There has been increasing interest in PILs in recent years.4,7–10 Hirao et al. reported the preparation of a series of heterocyclic amine based PILs and revealed correlations between * Corresponding authors. Email: [email protected] (F.S.); Fax: +61 3 9347 8189, Phone: +61 3 8344 2447. E-mail: [email protected] (A.M.B.); Fax: +61 3 9905 4597: Phone: +61 3 9905 1338. † School of Chemistry, Monash University. ‡ ARC Special Research Center for Green Chemistry, Monash University. § University of Melbourne. | Orica Ltd.

the cation structures and the physical properties such as melting point, glass transition temperature, and ionic conductivity.8 More recently, Greaves et al. reported a series of amine cation based PILs and presented phase transition data and reported a range of physicochemical properities.9–11 The effect of structural change has been investigated and provides the possibility to manipulate the physicochemical properties. From an application perspective, some PILs were discovered to mediate hydrocarbonsolvent interactions and promote amphiphile self-assembly.11 The protonic feature of PILs is a fundamental characteristic that has been explored in a number of areas such as biological applications,12–14 organic synthesis,15–17 and chromatography.18 Other applications of PILs include fuel cell,19,20 selfassembly,11,21–24 catalysis,25 and explosives.21,26 In addition, protic ionic liquids must be considered of interest in electrochemical studies and in electrochemical devices. AILs have already been shown to offer advantages in electrochemical studies when used as either the solvent or electrolyte or both.4,27–33 In some cases, wider potential windows (5-6 V) are available than in the case of traditional solvents. Thus, electrodeposition of various metals at potentials that cannot be obtained in aqueous electrolyte solutions are available.34 The intrinsic conductivity of ILs usually means that the addition of supporting electrolyte may be avoided. Specific capacitance, good stability, and wide potential windows also make them suitable electrolytes for electrochemical capacitors.27 PILs have drawn much attention lately as anhydrous proton conductors in hydrogen-oxygen fuel cells.3,19,20 Wu et al. demonstrated that the conductivities of some PILs can match those found for

10.1021/jp711804j CCC: $40.75  2008 American Chemical Society Published on Web 05/20/2008

6924 J. Phys. Chem. B, Vol. 112, No. 23, 2008 aqueous electrolyte solutions, which are generally assumed to provide superior electrolytic conductivity because of the unique dielectric and fluid properties of water.3 The high conductivities of PILs were attributed to the high fluidity and ionicity rather than existence of the Grotthus mechanism. The combination of high ionicity and proton exchange kinetics with relatively low vapor pressure also makes these PILs excellent candidates for use as fuel cell electrolytes, although concerns regarding evaporative loss of the ionic liquid remain. Watanabe and coworkers describe the performance of a hydrogen electrode utilizing, as the electrolyte, the salt formed by proton transfer from the acid form of bis-trifluoromethanesulfonyl amide to imidazole.19 Belieres et al. also have demonstrated that fuel cells with performance superior to that of the commercially feasible phosphoric acid high temperature fuel cell can be achieved with some protic ionic liquids.20 More recently, Tsunashima et al. have reported phosphonium based ILs as novel electrolytes. Some of them showed quite low viscosity compared with corresponding ammonium ILs, good electrochemical window, and thermal stabilities.35 Despite their well recognized fundamental and practical significance, systematic accounts of the electrochemical properties of PILs, akin to those widely available for AILs, are yet to become available. In this paper, we, therefore, report the outcome of a systematic electrochemical study of room temperature PILs. Eighteen PILs were synthesized and characterized for this purpose. We then deliberately selected and characterized the electrochemical characteristics of the eight that are room temperature protic ionic liquids (RTPILs) using a range of voltammetric techniques. Electrochemical properties of the RTPILs surveyed include their potential windows, the availability of internal potential reference scales based on use of oxidation of ferrocene and reduction of the cobaltocenium cation, and some practical issues such as the influence of water and sometimes very high viscosity. To the best of our knowledge, this is the first systematic study of the electrochemical properties of protic ionic liquids. It is our hope this provision of fundamentally useful information may provide a platform for extension of electrochemical studies on this class of ionic liquids, which has to date been shown to be technically important but still not extensively used in industrial applications. Experimental Reagents. The protic ionic liquids were made from the following chemicals: di-n-butylamine (>99%), methylsulfonic acid (99%), sulfamic acid (>99%), and phosphoric acid (85% in water) from BDH (Poole, England); methanol, toluene, formic (98%), and glacial acetic (99.8%) acids from Ajax Chemicals (Auckland, New Zealand); hydrochloric (35%), nitric (69%), and sulfuric (>96%) acids from Merck (Darmstadt, Germany); diethylamine (99.5%), di-n-propylamine (>99%), triethylamine, (99%), diethanolamine (>99%), and triethanolamine (99%) from Sigma-Aldrich (St. Louis, MO). All of the chemicals were used as received from the manufacturer. Other chemicals used were acetonitrile (CH3CN, Merck), acetone (Merck), cobaltocenium hexafluorophosphate (Co(C5H5)2PF6 or CcPF6, Strem, MA), and ferrocene (Fe(C5H5)2 or Fc, BDH, Australia). These also were used as supplied by the manufacturer. Tetrabutylammonium hexafluorophosphate (Bu4NPF6) was purchased from GFS and recrystallized twice from ethanol before use. PIL Sample Preparation. With the exception of the preparation of sulfamate PILs, the selected base was initially diluted with an equal volume of solvent (water or methanol). An

Zhao et al. TABLE 1: Thermal and Other Properties of the Protic Ionic Liquids compound [DEA][Cl] [DEA][OSA] [DEA][H2PO4] [DEA][HSO4] [DEA][NO3] [DEA][Of] [DEA][Ac] [DEA][Ms] [TEA][Of] [TEA][Ac] [DEtA][Of] [DEtA][Ac] [DPA][Of] [DPA][Ac] [DBA][Of] [DBA][Ac] [TEtA][Of] [TEtA][Ac]

Tga (°C)

Tpb (°C)

Tmc (°C)

Tdd (°C)

∆pKae

H2O %f (w/w)

89 221 206 190 238 114 57 280 88 74 86 53 57 35 63 59 54 48

16.9 7.7 6.9 11.9 10.3 4.1 5.1 10.8 3.0 4.0 7.2 6.2 7.3 6.2 7.5 6.5 7.0 6.0

0.62 0.94 0.46 0.66 0.24 0.54 0.20 0.12 0.91 0.08 1.79 0.14 2.06 0.18 0.81 0.14 0.51 0.42

-86 -62 -35 -65

g

-26

g

g

g

64

g

g

-87 -61 -66

-27 g

g

g

-89

g

69 g

-59

-20, 23

55 48 65 47

g

g

g

-92

-18 -28 13 -1.3

g

-31 -43, 35 -6.6 -15, 11 -18

g

g

g

g

-93

-31

-18

g g g

g

a

Glass transition temperature. b Unidentified phase transition between Tg and Tm. c Melting point. d Thermal decomposition of 5% weight loss temperature. e ∆pKa is defined as the difference in pKa of the acid and base in water.46 f Water content determined by Karl Fischer titration. g No phase transition observed by DSC.

equimolar amount of acid was then added dropwise to the diluted base while stirring rapidly. For reactions with sulfamic acid, the solid acid was mixed with methanol and an equimolar amount of base added dropwise. Since the reaction of bases with strong inorganic acids is extremely aggressive, care should be taken during the course of their synthesis. Thus, the temperature during the reaction was maintained below 35 °C in most cases. Bulk solvent was removed by rotary evaporation, and then water was removed azeotropically from the sample by addition of toluene and subsequent removal in vacuo. Further drying was achieved under vacuum at 0.03 mbar for 24 h. Typically, the final product was hygroscopic and was again dried prior to use. Physicochemical Characterization Procedures. Solution 1H and 13C NMR measurements were conducted with a Varian (Palo Alto, CA) INOVA 400 MHz or Varian INOVA 500 MHz NMR spectrometer. Samples that were liquid at room temperature were also examined by NMR (1H, 13C,15N) as neat liquids and referenced to 1% sodium 3-(trimethylsilyl)propionic acidd4 acid in D2O in a sealed glass capillary (Kontes, Vineland, NJ) for 1H and 13C and externally referenced to TMS as 0 ppm. In the case of 15N, neat PIL samples were externally referenced to NH4Cl. These data confirmed that no unexpected molecular species were formed from acid and base mixing. The water content of the dried PIL was determined by Karl Fischer titration using an 831 Metrohm (Herisau, Switzerland) Karl Fischer coulometer. A 25% w/w solution of the salt in dried methanol (distilled over Mg and stored over 3 Å molecular sieves) was injected for analysis. The average water content is given in Table 1. Thermo-gravimetric analysis (TGA) was carried out using a Mettler Toledo (Columbus, OH) STAR 810 system under a nitrogen/oxygen atmosphere (80:20). TGA samples were run in an open aluminum pan and heated from room temperature to 500 at 5 °C min-1. The thermal decomposition temperature (Td) was taken to be the value when a 5% mass loss was observed. Differential scanning calorimetry (DSC) was performed using a Perkin-Elmer (Waltham, MA) PYRIS Sapphire

Electrochemistry of Ionic Liquids

J. Phys. Chem. B, Vol. 112, No. 23, 2008 6925

system under nitrogen atmosphere from temperatures of -100 °C to within 10 °C of the Td. Samples were run in a sealed aluminum pan using three temperature cycles. Density values were obtained with an accuracy of (1% by measuring the weight of the sample in a 5 mL volumetric flask at room temperature (20 ( 2 °C). The kinematic viscosity of the liquid samples at 25 °C was measured with a Schott (Mainz, Germany) micro-Ubbelohde capillary viscometer in the range of 0.3-3600 cSt. Sample flow times are corrected according to instrument values. Density measurements were then used to convert from kinematic to dynamic viscosity. The ionic conductivity of the liquid samples was obtained at 25 °C by measuring the complex impedance at a frequency between 0.1 Hz and 1 MHz with a Solartron (Farnborough, UK) 1260 immediate response analyzer. A locally designed conductivity cell made from two platinum wires was used for the measurements. The cell constant was determined by calibration after each sample measurement using an aqueous 0.01 M KCl solution. Electrochemical Instruments and Procedures. All voltammetric experiments at stationary macrodisk working electrodes were undertaken with a BAS 100B/W electrochemical workstation (Bioanalytical System, West Lafayette, IN). The uncompensated resistance values (Ru), typically from 1000-10 000 Ohm were measured using the BAS system by applying a potential step (∆E ) 50 mV) in a potential region where no Faradaic reaction occurs.36 The glassy carbon (GC), Au, and Pt working disk electrodes were from Cypress (Cypress Systems, Inc., Lawrence, KS). The boron-doped diamond (BDD) disk (boron carrier concentration of 1020 atoms cm-3) electrode was purchased from Windsor Scientific (Windsor Scientific Ltd, Slough Berkshire, UK). Prior to each experiment, the working electrodes were polished with 0.3 µm alumina (Buehler, Lake Bluff, IL) on a clean polishing cloth (Buehler), sequentially rinsed with distilled water and acetone, and then dried with lint free tissue paper. Effective electrode areas of 0.0299 and 0.0094 cm2 for glassy carbon electrodes (GC), 0.0137 cm2 for the Au electrode, 0.0134 cm2 for the Pt electrode, and 0.0845 cm2 for the BDD were determined from cyclic voltammograms using the peak current from the oxidation of a 1 mM ferrocene solution in CH3CN (0.1 M Bu4NPF6) degassed with N2 and use of the Randles-Sevcik relationship (1)

ip ) 0.4463nF(nF ⁄ RT)1/2AD1/2V1/2C

(1)

where ip is the peak current (A), n ) 1 is the number of electrons in the charge transfer step, A is the electrode area (cm2), D is the diffusion coefficient of Fc (taken to be 2.3 × 10-5 cm2 s-1),37 C is the concentration (mol cm-3), V is the scan rate (V s-1), and other symbols have their usual meanings. A standard threeelectrode arrangement was used with a Pt wire counter electrode and a Ag/Ag+ (CH3CN, 0.01 M AgNO3) reference electrode. Rotating disk voltammetric experiments were performed with a BAS Epsilon electrochemical analyzer in conjunction with a BAS MF-2066 glassy carbon rotating disk working electrode (area ) 0.707 cm2). Voltammetric experiments in PILs employed the same cell configuration, as described above, except that a silver wire was employed as the quasi-reference electrode. The PILs were dried under vacuum for 12 h prior to measurement, and the electrochemical experiments were conducted inside a homemade nitrogen-filled drybox. Simulations of voltammograms were achieved with Digisim software.38 Cobaltocenium hexaflurophosphate and ferrocene are poorly soluble in the more viscous ILs. For lower viscosity PILs,

TABLE 2: Physical Properties of the Room Temperature Protic Ionic Liquids (RTPIL) from Table 1 RTPIL [DEA][Ac] [DEA][HSO4] [DPA][Of] [TEtA][Ac] [DEA][Of] [TEtA][Of] [DEA][OSA] [DEA][Cl]

base/acid F (g cm - 3) η (mPa · s) σ (mS cm-1) ratio (mol %)a 1.49:1 b

1.00:1 1.17:1 1.37:1 1.32:1 b b

1.22 1.21 0.97 0.96 1.13 1.04 1.45 1.24

336 c

19 11 28 10 720 305

0.14 0.04 1.19 1.27 5.83 0.36 14.21 0.86

a Determined via integration of neat liquid NMR. Actual values refer to sample dried under vacuum in extended periods of time to reduce water concentration (see text for details). b Acid not visible by NMR integration. c Viscosity exceeded 4356 mPa.s.

typically a solution concentration of approximately 1 mM was prepared for voltammetric studies with cobaltocenium hexaflourophosphate and ferrocene. A ME36S microbalance (Sartorius, Australia) was employed to weigh microgram amounts of Fc and CcPF6 for preparation of these solutions. An ultrasonic bath (Unisonics Pty Ltd, Australia) was employed to assist with the dissolution of these compounds. All voltammetric experiments were carried out at ambient temperature (20 ( 2 °C). Results and Discussion 1. Characterization of PILs. In this investigation, two series of protic substituted ammonium salts have been produced from secondary or tertiary amine bases with either organic or inorganic acids. Water content analysis and thermal properties, including glass transition temperature (Tg), melting point (Tm), thermal decomposition (Td), and other thermal solid-solid phase transitions (Tp) are summarized in Table 1. For those samples which were liquid at room temperature, acid-base component ratios of the mixture and physical properties such as density, viscosity, and ionic conductivity are reported in Table 2. In the first series, diethanolammine ([DEA]) was combined with one equivalent of acetic ([Ac]), formic ([Of]), hydrochloric ([Cl]), sulfamic ([OSA]), methanesulfonic ([Ms]), sulfuric ([HSO4]), or phosphoric ([H2PO4]) acids. Initially, [DEA] was employed as a cation, because introducing alkoxy groups into quaternary ammonium cations is known to reduce both melting point and viscosity of the ionic liquid.39,40 The acids were used in their monoprotic forms to examine the effects of pKa and the numbers of hydrogen bond donors and acceptors in the binary mixtures. As shown in Table 2, the formic and acetic acid compounds had comparatively low viscosity and also low Tm and Td values. As these are desirable physical properties for reactive ionic liquids, the formic and acetic acids were then paired with other bases to determine which physical properties were due to the anion. In the second series, [Ac] and [Of] acids were combined with the different secondary and tertiary amine bases: triethylamine ([TEtA]), triethanolamine ([TEA]), diethylamine ([DEA]), di-n-propylamine ([DPA]), and di-n-butylamine ([DBA]). After the compounds were dried, a small portion was dissolved in D2O and analyzed by 1H and 13C NMR spectroscopy to ensure that no side reactions had occurred or unexpected molecular species formed. As the chemical shifts of the ionized species were dependent on concentration, NMR was only used as a rapid, qualitative form of analysis. Also, not all of the acids in the PILs were visible by 1H and 13C NMR, to wit, [H2PO4], [HSO4], [HCl], and [OSA]. Due to the change in properties with

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Zhao et al.

Figure 1. Representative protic ionic liquid (a) 1H and (b) 13C NMR spectra of [DPA][Ac] with alkyl chain ratio imbalance from proton exchange relaxation. δH (400 MHz, neat, TSP) 11.46 (NH, AcOH) 3.03 (4 H, t, J ) 7.4, NCH2) 2.03 (3 H, s, AcOH), 1.86 (4H, dt, J ) 7.2 and 7.4, CH2CH2CH3), 1.08 (6 H, t, J ) 7.2, CH2CH2CH3); δC (100 MHz, neat, TSP) 175.9 (AcOH), 49.3 (NCH2), 23.2 (CH2CH2CH3), 19.8 (CH2CH3), 11.3 (AcOH).

water concentration, component ratios of PILs were only recorded with neat liquid samples. Proton exchange between the acid, base, water, and hydroxyl groups of [DEA] and [TEA] was observed to cause rapid relaxation of the exchanging sites and some NMR signal intensity was lost in the “dead” time between application of the rf pulse and spectral acquisition. Consequently, the R-C and attached protons adjacent to a site of proton exchange relaxed faster than the β-C and β-protons and, thus, nonequivalent integration was observed for alkyl chains (Figure 1). Hence, integration of proton and carbon resonances furthest from an exchanging site was used for comparison of ratios, i.e., the β-C ([DEtA], [TEtA], and [Ac]), γ-C ([DPA]), and δ-C ([DBA]). However, in the case of [Ms] the R carbon that is β to the exchange site was used. In the case of [Of], [DEA], and [TEA], where all carbons are adjacent to an exchange site, the NMR integrations were expected to have a ratio imbalance between the acid and base mixture. Therefore, the reported ratios of acid molecules compared to base molecules (Table 2) are a qualitative rather than quantitative indication of stoichiometric imbalance only. The PILs produced via the synthetic method employed were a three-component mixture of acid, base, and water. Consequently, it was difficult to exactly reproduce the ratios of components especially after vacuum had been applied to remove water. As expected, asymmetric stoichiometries were obtained from 1H NMR integrals, with the less volatile species in excess,

and additional exposure of the sample to vacuum tended to bias the ratio further away from stoichiometry.1 The inability to exactly reproduce the combination of components led to disparity in physical properties for replicate acid-base mixtures. However, due to the necessity to remove at least a substantial fraction of solvent and trace water in vacuo, a 24 h drying period at constant pressure of 0.3 mbar was adhered to as a standard. Compounds were then stored in a desiccator or glovebox under an inert atmosphere, and physical measurements were then undertaken on the same sample wherever possible to provide a constant proportion of components. The thermal decompositions of the PILs were determined by heating the sample and assigning the decomposition temperature (Td) when a 5% mass loss was obtained. Samples with inorganic acids produced an unidentified black char that remained beyond the heating limit of 500 °C. Several of the room temperature PIL samples with formic and acetic acids started to lose mass at temperatures above 25 °C. This is ascribed to evaporation of free acid or base, which implies that an equilibrium exists between the free acid and base and the ionized cation and anion and that PILs containing weak organic acids are not fully ionized. The TGA instrument used in the present studies required sample pans that were open to the atmosphere. The compounds [TEA][Ac] and [DEA][Ms] were very hygroscopic and visibly deliquesced, hence thermal decomposition data reported in Table 1 was determined by DSC. Samples of [TEA][Ac] and [DEA]-



Figure 2. Examples of data obtained from DSC heating experiment.

[Ms] were handled and sealed in a DSC sample pan in a glovebox under nitrogen atmosphere. Comparison of thermal decomposition data derived from [DEA][NO3] and [DEA][Ac] showed less than (1 °C variance between DSC and TGA analysis. The solid-liquid phase transitions of all compounds were investigated by DSC. The data for Tg, Tp, and Tm of the 18 PILs are reported in Table 1. DSC traces for typical types of phase transitions observed are exemplified in Figure 2. One type of behavior (Figure 2, a) was characterized by PILs exhibiting only a single melting transition (e.g., [DEA][NO3]). A second form or behavior (Figure 2, b) arises from samples such as [DEA][Of] where only a glass transition was observed without a clearly defined Tm. The third type of behavior (Figure 2, c) was from samples with multiple solid-solid transitions before Tm. The fourth type of behavior (Figure 2, d) was observed when upon heating from low temperature (-100 °C), a supercooled liquid was formed which then crystallized at a higher temperature before a more typical Tm. This supercooled state has been observed in other ionic compounds.40,41 Compounds that did not display a clear melting point by DSC did indeed have visually observable phase transitions in the range examined. However, the rate of cooling/heating and small amount of sample analyzed by DSC can result in supercooling such that no transitions are observed. 2. Ionic Conductivity. The ionic conductivity of PILs is an important parameter that is relevant to their possible applications in polymer membrane fuel cells. Ionic conductivity is related to the degree of dissociation, charge number, and ionic mobility of the ions, but a mobile proton transported via a Grotthus mechanism may make a non-negligible contribution to the total conductivity in protic ionic liquids.42,43 The generally low conductivity of many PILs is mainly attributed to the poor mobility of the ions, which is associated with the high viscosity. Deliberate addition of water increased the conductivity significantly as well as decreasing the viscosity. It, therefore, follows that the presence of adventitious water may play a role in the reported conductivity and viscosity. Xu et al. and others have used the Walden plot based upon viscosity and ionic conductivity to classify ionic liquids as superionic (alternative ionic mechanism) or subionic (associated,

Figure 3. Walden plot of compounds from Table 2 assuming full ionization. The straight line plot is generated from data obtained in aqueous 0.01 M KCl solution. The dashed line lies 1 order of magnitude below the aqueous KCl line, illustrating the situation of a liquid that is only 10% ionized.

ion-paired, etc.).3,44 When compared to a solution which is known to be fully dissociated and to have ions of equal mobility, e.g., dilute aqueous solutions of KCl, all PILs examined in this study were subionic and have relatively poor ionic conductivity (Figure 3), except for the sulfonamide which lies well above the ideal line, suggestive of a proton conduction mechanism operating in this case. Given the additional protons and proton sites available on the sulfonamide nitrogen, this possibility seems entirely feasible and is worthy of further investigation in the context of fuel cell electrolytes. It was hypothesized that the ∆pKa of the acid and base species in water would have some relationship to the physical properties of the PILs. Specifically, a sufficiently large ∆pKa is thought to lead to complete proton exchange and ionization between two components and hence different physical properties compared to a weakly ionized pair. Although ∆pKa > 4 would be an estimate to be sufficient in aqueous solution to produce complete proton transfer, Xu et al.3 and others45 have suggested that in fact ∆pKa > 10 is necessary to produce highly ionized


Zhao et al.

TABLE 3: Cyclic Voltammetric Dataa for Ferrocene and Cobaltocinium Hexaflurophosphate in Protic Ionic Liquids Cc+/0 PIL

scan rate (mV/s)

∆Ep (mV)

ip,a/ip,c

10 100 200 10 100 200 20 200 3000 20 100 200 100 300 600 20 100 600 20 100 600 20 100 600

53 67 83 45 245 191 53 53 83 56 53 53 70 76 88 79 75 78 167 96 153 76 72 76

0.98 0.99 0.99 4.30 3.51 1.45 0.99 1.03 1.05 0.97 0.98 0.98 1.02 0.98 1.03 1.00 1.02 1.00 0.95 2.25 2.44 0.06 0.10 0.24

[DEA][Ac] [DEA][HSO4] [DPA][Of] [TEtA][Ac] [DEA][Of] [TEtA][Of] [DEA][OSA] [DEA][Cl]

Fc0/+ D

(cm2

/s)

8.7 × 10-9 7.7 × 10-9 7.7 × 10-9 b b b

1.5 × 10-7 1.4 × 10-7 1.3 × 10-7 6.1 × 10-8 5.7 × 10-8 5.4 × 10-8 1.8 × 10-7 1.8 × 10-7 1.8 × 10-7 3.4 × 10-7 3.4 × 10-7 3.3 × 10-7

D (cm2/s)

∆Ep (mV)

ip,a/ip,c

58 80 85 27 66 71

1.00 1.06 0.98 1.74 1.45 1.24

d

d

d

e

d

d

144 62 64 73 87 69 66 58 55 58

1.06 0.96 0.97 0.98 3.98 2.95 1.26 0.99 0.93 0.96

d d b b b

e

8.5 × 10-8 9.2 × 10-7 8.1 × 10-7 7.9 × 10-7 d d

2.6 × 10-7 8.6 × 10-7 8.7 × 10-7 8.7 × 10-7

b

c

c

c

b

c

c

c

b

c

c

c

b

c

c

c

b

c

c

c

b

c

c

c

a Midpoint potentials (Em) for the Cc+/0 process calculated as the average of the oxidation (Epox) and reduction (Epred) peak potentials, Em ) (Epox + Epred)/2, are always -1.34 ( 0.01 V versus the midpoint potential for the Fc0/+ process. b Surface interaction prevents calculation of D values. c Fc is too insoluble for measurement of parameters. d Not determined. e Chemically irreversible at slow scan rate, so ratio cannot be calculated.

PILs, due to the different solvation environment in the PIL as opposed to an aqueous solution. Generally, the PILs with higher ∆pKa values46 lie closer to the ideal line in Figure 3 and well above the 10% line that very approximately indicates the region in which the system is only 10% ionized.44 The PILs studied reflect a much greater range of acid strengths, from HCl to formic acid, than base strengths and, therefore, the trends in the Walden plots are more strongly related to the anion present. Nonetheless, a cation dependence was also clear, with the weaker, more substituted, bases such as triethylamine producing less ionized PILs. 3. Development of a Reference Scale for Potential. In order to compare voltammetric data obtained in different RTPILs, it is necessary to use either a reference electrode of known potential against a standard reference electrode or reference all data to a process whose reversible potential is assumed to be independent of the ionic liquid. The ferrocene/ ferrocenium (Fc0/+) couple is most commonly used as an internal potential scale standard in voltammetry in traditional organic solvent medium containing an electrolyte.47,48 This process also is used in ionic liquid electrochemistry. In addition, Bond and co-workers also have demonstrated that the Cc+/0 process probably is a broadly useful reference scale in both organic solvent and ionic liquids.49–51 In this study, the two IUPAC recommended potential references, Cc+/0 and Fc0/+, were tested for all the RTPILs using cyclic voltammetry at a GC working electrode. The results are summarized in Table 3. In some ionic liquids, the cobaltocenium cation exhibits two reversible one-electron reduction processes (eqs 2, 3). +

Cc

(ionic

Cc0 (ionic

-

+ e a Cc

0 liquid)

(2)

liquid) + e a Cc (ionic liquid)

(3)

liquid)

(ionic

However, the second reduction process has been reported to be solvent dependent,50 so only this first reduction process is used for reference potential calibration purposes.

Figure 4. A comparison of (a) an experimental cyclic voltammogram obtained for reduction of 3 mM Cc+ in [DEA][Ac] at a GC electrode at a scan rate of 100 mV s-1, and (b) the theoretical response obtained by digital simulation of a reversible process. Simulation parameters: V ) 100 mV s-1, T ) 20 °C, double layer capacitance ) 5.0 × 10-7 F, DCc+ ) 7.7 × 10-9 cm2 s-1, effective electrode area ) 0.03 cm2, E0 ) -900 mV.

Ferrocene gives rise to a reversible one-electron oxidation process (4).

Fc(ionic liquid) a Fc+(ionic liquid) + e-

(4)

The voltammetric behavior of Cc+/0 and Fc0/+ in the PILs of interest in this study varies dramatically. In the case of [DEA][Ac], essentially ideal voltammetric behavior is found for both the Cc+/0 and Fc0/+ processes at a GC electrode. The cyclic voltammogram presented in Figure 4 for reduction of 3



Fc(ionic liquid) f Fc+(ionic

liquid)

(5)

V Product

Figure 5. Cyclic voltammograms obtained at 20 °C for oxidation of 4 mM ferrocene in [DPA][Of] at a GC electrode using scan rates of (a) 20, (b) 200, (c) 300, (d) 600, and (e) 1000 mV s-1.

mM Cc+ at a scan rate of 100 mV s-1 exhibits close to an ideal one-electron reversible behavior. Cyclic voltammograms of the first reduction process were recorded as a function of scan rate. The peak-to-peak potential separation (∆Ep) of 67 mV is close to the theoretical value (56 mV at 20 °C) for a reversible process at slow scan rates (100 mV s-1), ∆Ep increases as expected, due to the enhanced impact of Ohmic drop. A linear relationship was found between the peak current (ip) and the square root of scan rate (V1/2) as expected for a diffusion controlled process. The diffusion coefficient (D) of Cc+ in [DEA][Ac] was calculated from use of the Randles-Sevcik relatioship to be 7.7 × 10-9 cm2 s-1, a value which is much lower than that in CH3CN (1.9 × 10-5 cm2 s-1)50 but close to that in the widely used AIL [BMIM]PF6 (1.0 × 10-8 cm2 s-1).51 Comparisons of experimental and simulated voltammogram further confirm the reversibility of the process. Thus, as shown in Figure 4, the experimental voltammogram nicely matches the voltammogram simulated using known experimental parameters. A similarly close to reversible voltammetric response was observed for the oxidation of Fc in [DEA][Ac]. The separation in reversible potential for the Fc0/+ and Cc+/0 couples in these PILs is 1.34 ( 10 mV, which is the same as reported in organic solvents and many other ionic liquids, e.g., [BMIM]PF6.50,51 On the basis of studies related to those described above, the Cc+/0 process has been found suitable for use as a reference potential scale system in [DEA][Ac], [DPA][Of], [TEtA][Ac], [DEA][Of], and [TEtA][Of]. Analogously, ferrocence also could be used as a reference potential scale for voltammetric studies in [DEA][Ac], [TEtA][Ac], and [TEtA][Of]. In [DPA][Of] and [DEA][Of], chemically irreversible voltammetric responses were observed for oxidation of Fc at low scan rates, but these become chemically reversible at high scan rates. Figure 5 shows the oxidation of Fc in [DEA][Of]. At a low scan rate of 20 mV s-1, a well-defined oxidation peak appeared at 0.76 V vs Ag wire. However, no reduction peak was observed on the reverse scan until the scan rate increased to 0.2 V s-1. In this ionic liquid, a close to ideal reversible response is exhibited at scan rates higher than 0.3 V s-1. The scan rate dependence indicates that the oxidation of Fc to Fc+ in [DPA][Of] and [DEA][Of] follows an EC mechanism (5)

It is tempting to postulate that a homogeneous reaction between Fc+ and formate is the cause of nonreversibility. However, chemically reversible voltammetry was observed at all scan rates in [TEtA][Of], which also has the formate anion. In general, we observed that the Cc+/0 process provides a more widely useful potential reference scale for PILs, since Cc+ salts typically have a higher solubility and commonly exhibit reversible behavior in the chemical and electrochemical senses. Mechanisms of operation that are significantly different for the Cc+/0 and Fc0/+ processes were found in other PILs. Figure 6a shows a voltammogram for the reduction of Cc+ at a GC electrode in [DEA][HSO4] at a scan rate of 100 mV s-1. In this PIL, the Cc+ reduction peak was observed at around -1.11 V vs Ag wire when scanning the potential in the negative direction. On reversal of the scan direction, a larger and sharper oxidation peak was observed at -1.06 V. The ∆Ep value of 53 mV, the peak current ratio |ip,c/ip,a| of 0.327, and the oxidation peak shape all differ from characteristics expected for a diffusion controlled process. The symmetric peak shape for oxidation of Cc0 to Cc+ and the characteristics are typical for an electrode process in which product adsorption occurs after the initial electron transfer reaction.52,53 That is, adsorption of Cc occurs after reduction of Cc+, followed by oxidative stripping. This implies that the neutral Cc0 has a lower level of solubility than the charged Cc+ form, which may be expected, taking into account the ionic nature of the PILs. The postulated mechanism for the Cc+/0 process is given in eq 6a-c.

Cc+(ionic liquid) + e- f Cc0(ionic liquid)

(6a)

Fast adsorption

Cc0(ionic liquid) 98 Cc0(ads)

(6b)

Stripping

Cc0(ads) 98 Cc+(ionic liquid) + e-

(6c)

The oxidation of Fc in [DEA][HSO4] (Figure 6b) has analogous characteristics. However, on this occasion, the voltammetric experiment starts with the neutral form, Fc, which is oxidized to Fc+ via adsorbed Fc when the potential is scanned in the positive direction. On the reverse scan, a much smaller reduction peak was observed, corresponding to the reduction of dissolved Fc+ to Fc, to give the suggested mechanism presented in eq 7a-c. Fast adsorption

Fc(ionic liquid) 98 Fc(ads)

(7a)

Stripping

Fc(ads) 98 Fc+(ionic liquid) + e-

(7b)

Fc+(ionic liquid) + e- f Fc(ads)

(7c)

Analogous voltammetric characteristics for Cc+/0 and Fc0/+ were also found when using Au, Pt, or BDD electrodes, implying that adsorption is not a result of specific interaction with a particular electrode surface. Closely related voltammetric responses that involved adsorption of Fc and Cc also were observed in [DEA][OSA] and [DEA][Cl]. Interestingly, all three PILs where this kind of behavior is found exhibit the common properties of very high viscosity and relatively low conductivity (Table 3). Reduction of Cc+ followed by surface adsorption of


Zhao et al.

Figure 6. Cyclic voltammograms obtained for the (a) Cc+/0 and (b) Fc0/+ processes in [DEA][HSO4] at a GC electrode. Scan rate ) 100 mV s-1, T ) 20 °C.

Cc0 has been previously observed in voltammograms in acetonitrile in the absence of supporting electrolyte and in aqueous electrolyte media where Cc and Fc are known to be sparingly soluble.53 The phenomenon in the absence of added supporting electrolytes was concluded to be associated with the modified diffusion layer, which extends much further into the solution in the absence of supporting electrolyte. A multiple double layer model has been proposed for ILs, in which the excess charge on the IL is located several layers deep within multilayers of alternating cations and anions.54 Therefore, the presence of adsorption in the voltammetric behavior in highly viscous PILs may be related to the poor solubility of Cc and Fc in such media and/or the modification of the diffusion layer at the electrode/ ionic liquid interface owing to the poor conductivity. Reversible Fc0/+ and Cc+/0 processes are widely used for internal voltammetric reference potential scale purpose in aprotic ionic liquids,55,56 as is the case in organic solvents. In aprotic ionic liquids, anomalous behavior has been reported for the Fc0/+ couple.55,57 However, the Cc+/0 couple appears to be close to ideal.55 In our work, while potential scales could be established for some protic ionic liquids using either of the Fc0/+ or Cc+/0

couples, or both, adsorption and chemical reactions coupled to the charge transfer process can cause difficulties in several systems. Direct use of a conventional reference electrode in aprotic ionic liquids has also been reported by several groups.58–61 Different type of reference electrodes have been constructed and evaluated in these studies, including reference electrodes of the first kind (Ag/Ag+ couple dissolved in the ionic liquid)58,60 and of the second kind (Ag/AgCl in ionic liquids containing either dissolved Ag+ or Cl-).59,60 However, no data are available based on use of conventional reference electrodes in protic ionic liquids. Clearly, in principle, these types of reference electrodes may be fabricated in protic ionic liquids where Ag salts are soluble or where the anions, for example, are halides or sulfate and could be considered in future studies. 4. Potential Windows in RTPILs. The potential windows for all the RTPILs have been measured by cyclic voltammetry at a scan rate of 100 mV s-1 with GC, Pt, Au, and BDD electrodes as the working electrodes. All these data for the RTPILs were obtained under dry box conditions after drying the RTPIL under vacuum. A silver wire was used as the quasireference electrode. The reduction and oxidation potential limits



TABLE 4: Electrochemical Potential Windows for the PILs Obtained at Different Electrode Materials PIL [DEA][Ac]

[DEA][HSO4]]

[DPA][Of]

[TEtA][Ac]

[DEA][Of]

[TEtA][Of]

[DEA][OSA]

[DEA][Cl]

a

electrode materials GC GC GC Pt Au BDD GC GC GC Pt Au BDD GC GC GC Pt Au BDD GC GC GC Pt Au BDD GC GC GC Pt Au BDD GC GC GC Pt Au BDD GC GC GC Pt Au BDD GC GC GC Pt Au BDD

anodic limit (V)

cathodic limit (V)

potential windows (V)

0.59 1.80 0.90 0.77 0.99 1.15

-1.84 -0.63 -1.53 -0.73 -0.93 -1.24

2.43 2.43 2.43 1.50 1.92 2.39

a

a

a

a

a

a

1.11 1.45 1.44 1.11 0.23 1.43 0.74

-1.05 -0.48 -0.38 -1.79 -2.46 -1.26 -1.95

2.16 1.93 1.82 2.90 2.69 2.69 2.69

b

b

b

-0.93 -1.38 -2.77 -1.49 -2.39 -0.58 -1.17 -1.33 -3.39 -2.05 -2.73

1.14 2.48 3.42 3.42 3.42 1.53 2.06 2.32 3.60 3.60 3.60

-0.21 1.10 0.65 1.93 1.03 0.95 0.89 0.99 0.21 1.55 0.87 b

b

b

0.16 0.93 0.39 1.73 0.98

-1.00 -1.42 -3.04 -1.70 -2.45

1.16 2.35 3.43 3.43 3.43

b

b

b

-1.08 -1.26

1.23 2.13

a

a

a

a

a

a

1.84 0.89 1.00 1.05

-2.41 -0.58 -0.97 -1.98

4.25 1.47 1.97 3.03

a

a

a

a

a

a

1.47 1.24 0.73 1.39

-2.54 0.10 -0.10 -1.11

4.01 1.14 0.83 2.50

-0.15 -0.87

potential reference Fc0/+ Cc+/0 Ag wire Ag wire Ag wire Ag wire Fc0/+ Cc+/0 Ag wire Ag wire Ag wire Ag wire Fc0/+ Cc+/0 Ag wire Ag wire Ag wire Ag wire Fc0/+ Cc+/0 Ag wire Ag wire Ag wire Ag wire Fc0/+ Cc+/0 Ag wire Ag wire Ag wire Ag wire Fc0/+ Cc+/0 Ag wire Ag wire Ag wire Ag wire Fc0/+ Cc+/0 Ag wire Ag wire Ag wire Ag wire Fc0/+ Cc+/0 Ag wire Ag wire Ag wire Ag wire

Potential reference scales were not established using Fc0/+ or Cc+/0. b No well-defined potential window could be obtained.

are also referred, where possible, to the IUPAC recommended potential reference scales (Cc+/0 and Fc0/+). The results are summarized in Table 4. Figure 7 contains cyclic voltammograms from which the potential windows available in [DEA][Of] have been examined for all four electrode materials. At a GC electrode, the reduction current, presumably associated with the reduction of [DEA]+ or the irreversible reduction of H+ to H2, starts to increase dramatically above the baseline value at -2.73 V vs Ag. The positive potential limit, presumably resulting from the oxidation of [Of]-, appears at 0.87 vs Ag. The difference gives rise to an overall potential window of 3.60 V. At a BDD electrode, a welldefined low background current could be observed in the potential range of -1.42 and 0.93 V, amounting to a potential window of 2.35 V, which is significantly larger than the value obtained at an Au electrode of 1.16 V. There is no well-defined potential window measurable at a Pt electrode.

As revealed by evaluation of data in Figure 7 and Table 4, the potential window is strongly dependent on electrode material. However, the largest potential window is normally found with a glassy carbon electrode. BDD electrodes also generally show excellent voltammetric performance, typically providing the second widest potential window, as well as the lowest background current per unit area. Platinum and Au electrodes normally exhibit smaller cathodic potential limits than carbon based electrodes. These noble metal electrodes have a much smaller overpotential for reduction of protons and impurities such as water in conventional solvent (electrolyte) media.52 The same situation appears to be true in aprotic ionic liquids.27 At the GC electrode, the potential window magnitudes follow the ionic liquid sequence of [DEA][OSA] > [DEA][Cl] > [DEA][Of] > [TEtA][Of] > [TEtA][Ac] > [DPA][Of] > [DEA][Ac] > [DEA][HSO4]. At a BDD electrode, the order is slightly modified to give [DEA][OSA] > [DEA][HSO4] >


Zhao et al.

HCOO(ad) f CO2 + H+ + e-

Figure 7. Voltammetric potential windows for [DEA][Of] obtained under dry box conditions at (a) platinum, (b) gold, (c) boron-doped diamond, and (d) glassy carbon electrodes. Scan rate ) 100 mV s-1, T ) 20 °C.

[DEA][Cl] > [DPA][Of] > [DEA][Ac] > [DEA][Of] > [TEtA][Ac] > [TEtA][Of]. At a Au electrode, the sequence follows the order [TEtA][Ac] > [DEA][OSA] > [DEA][Ac] > [DEA][HSO4] > [DEA][Of] > [DPA][Of] > [TEtA][Of] > [DEA][Cl]. At the most active Pt electrode surface, in PILs that give a well defined potential window, the sequence is [TEtA][Ac] > [DEA][Ac] > [DEA][OSA] > [DEA][Cl]. The cathodic and anodic potential limits are indicative of the electrochemical stability of the RTPILs. However, establishing a sequence of electrochemical stability is difficult for all cation and anion combinations because a suitable potential reference scale is not always available. However, when this is not a problem, the cathodic and anodic limiting potentials are found to vary even when the cations or anions present are common. For example, the oxidation potential limit occurs at 0.6 V (vs Fc0/+) in [DPA][Of], 0.21 V (vs Fc0/+) in [DEA][Of], and 0.39 V (vs Fc0/+) in [TEtA][Of] at a GC electrode, and the reduction potential limit for the [DEA]+ series is -1.84 V (vs Fc0/+) in [DEA][Ac], but this limit is far more negative in [DEA][Of] (-3.39 V vs Fc0/+) at a GC electrode. The variation of these potentials implies that association of anions or cations plays an important role in their electrochemical stabilities. This is not surprising since strong interactions in PILs related to hydrogen bonding and ion-ion interactions have been reported to dramatically modify the physicochemical properties in this and other PIL studies.9,10 Table 4 reveals that a large potential window of 3.60 V is observed for [DEA][Of] at a GC electrode. In contrast, no welldefined potential window could be obtained via voltammetry at a Pt electrode for this or any other formate-based RTPIL. Figure 7a illustrates the complexity of cyclic voltammograms obtained in [DEA][Of] at Pt electrode. When the initial potentials were set at negative values, large peaks appeared when the potential was scanned in the positive direction while a major sharp peak and a small peak appeared on the reverse scan or negative potential direction. The production of gas bubbles was detected visually at the electrode surface when the potential was very positive. It has been demonstrated in acidic solution that formate is oxidized at platinum surface to CO2 via adsorption and then oxidation (eq 8).62

(8)

A related reaction is assumed to apply in formate based PILs. The oxidation of formic acid at a Pt electrode is regarded as the simplest system for prototyping the electrocatalytic oxidation pathway for methanol, formaldehyde, and formic acid (HCOOH) at Pt and Pt-based alloys. Their oxidation reactions are of interest in low temperature fuel cells and other industrial applications.63–65 Given the potentially very high concentration of formic acid in the formate RTPILs, applications in these areas may be possible. 5. Double Layer Capacitance. Double layer capacitance at electrode-solution interfaces may be crucial in term of understanding the electrode performance in an ionic liquid.52 In this study, the double layer capacitances per unit area (Cdl) for the RTPILs were obtained by plotting the background or charging currents (ic) at a selected potential range where no Faradic process is evident, versus scan rates (V). Under ideal conditions, the relationship ic ) Cdl V exists and so Cdl may be calculated using the slope of the plot. Results show that the higher capacitance values were obtained at a GC electrode for acetate and formate RTPILs (53 µF cm-2) and [DEA][Of] (27 µF cm-2), which are also the least viscous. Capacitance values for more viscous PILs were 6.4 µF cm-2 for [DEA][Cl] and [DEA][OSA], which are similar to values reported for [C2mim]+ and [C4mim)]+ AILs of 5.2 - 6.9 µF cm-2.66 The accessibility of the electrolyte to the electrode surface is assumed to limit the specific capacitance. The capacitance value of 14 µF cm-2 for [DEA][HSO4] is comparable to data reported at a GC electrode in aqueous solutions with supporting electrolyte, e.g., 15.1 µF cm-2 for aqueous 0.1 M KCl and 14.6 µF cm-2 for aqueous 3 M H2SO4, and also values of 10.6 - 12.4 µF cm-2 for a range of [C2MIM]+ based AILs.67 It should be noted that the capacitance data reported in this study are based on the Stern-diffuse double layer model,52 which is widely used to describe aqueous-electrode-electrolyte interfaces. Nanjundiah et al. have suggested that such a model cannot be used to describe the interface developed in RTILs.68 A multiple layer model, as established for molten salts, may be preferable.27 Nevertheless, PILs are still seen as suitable solventelectrolytes in electrochemical studies since they have adequate capacitance values, good stability and wide potential windows.9 6. Steady-state voltammetry. In principle, steady state measurements based on rotating disk or microelectrode methods would minimize iR drop and background capacitance terms and would be valuable in PIL-based media. However, usually only transient behavior could be observed for the PILs under voltammetric conditions when steady state behavior is found in aqueous or organic solvent electrolyte media. Figure 8a shows rotating disk electrode voltammograms for reduction of 3 mM Cc+ in [DEA][Ac]. A sigmoidal shape curve is expected to be achieved under steady state conditions where the peak shape form is indicative of transient diffusion controlled behavior. Simulations confirm that only transient behavior is expected when the diffusion coefficient is very small (Figure 8b). At 22 °C, the viscosity of [DEA][Ac] is 336 cP, the density is 1.01 g cm-3, and the diffusion coefficient for Cc+ obtained from Randles--Sevcik equation is 7.7 × 10-9 cm2 s-1. This combination of data gives a Schmidt number (Sc) of 3.6 × 108, which is about 10 times larger than that obtained in [BMIM][PF6]69 and more than 105 times larger than values of about 1000 reported in aqueous electrolyte solutions.52 Such unusually high values of Sc give rise to very low sensitivity to convection so that mass transport is governed predominantly by linear diffusion (transient voltammetry).



Figure 8. (a) Rotated disk electrode (RDE) voltammograms obtained for the reduction of 3 mM Cc+ in [DEA][Ac] at a GC electrode using a scan rate of 5 mV s-1 and rotation rate of 500 rpm. (b) Simulated voltammograms for a reversible process with parameters as follows: V ) 5 mV s-1, T ) 20 °C, electrode radius ) 0.15 cm, kinematic viscosity ) 0.41 cm2 s-1, rotation rate ) 500 rpm, E0 ) -1.033 V, DCc+ ) 7.7 × 10-9 cm2 s-1. (c) Simulated RDE voltammograms for reduction of 3 mM Cc+ as a function of Schmit numbers (Sc) with other simulation parameters as in (b) except that scan rate ) 10 mV s-1. A RDE simulation for a typical IL [BMIM][PF6] (kinematic viscosity ) 2.29 cm2 s-1, DCc+ ) 1.0 × 10-9 cm2 s-1)49 is included for comparison. (d) Simulated RDE voltammograms predicted for reduction of 3 mM Cc+ in [DEA][Ac] as a function of rotation rate (ω) of 100, 250, 500, 1000, 1500, 2500, and 3500 rpm. Other simulation parameters are as in (b) except that scan rate ) 10 mV s-1 and kinematic viscosity ) 2.86 cm2 s-1.

Clearly viscosity plays a significant role in PIL voltammetry. It affects the rotating disk electrode voltammetry in two ways. First, high viscosity results in low diffusion coefficients, as predicted by the Stokes-Einstein equation, and second very high Sc numbers are encountered, both factors minimizing the contribution of convection. Figure 8c shows a series of rotating disk voltammograms that simulate the reduction of Cc+ with a variation of kinematic viscosity and diffusion coefficients. The kinematic viscosity mimic values of 2.86 cm2 s-1 found for [TEtA][Ac], 2.29 cm2 s-1 for [BMIM][PF6], 0.11 cm2 s-1 for [TEtA][Ac], and 0.07 cm2 s-1 for [TEtA][Of]. Diffusion coefficients used in simulation mimic reported values of 7.7 × 10-9 cm2 s-1 for [DEA][Ac], 5.7 × 10-8 cm2 s-1 for [TEtA][Ac], and 3.4 × 10-7 cm2 s-1 for [TEtA][Of]. The Sc numbers calculated under these conditions vary from 2 × 105 to 4 × 108. Simulations predict that the sigmoidal steady state response would be achieved when Sc e 2 × 105, about 200 times larger than that encountered in aqueous electrolyte solutions.52 The diffusion coefficient of Cc+ in [TEtA][Of] (3.4 × 10-7cm2 s-1) is about 40 times smaller than that obtained in acetonitrile and 9 times smaller than that in DMSO.53 In principle, a near steady-state response can be achieved by increasing the rotation rate until the Levich limiting current value is achieved. However, access to sufficiently high rotation rates

are not straightforward to achieve in viscous protic ionic liquids. Figure 8d provides simulated RDE voltammograms at rotation rates over the range of 10-10 000 rpm. A steady state limiting current is predicated to be reached when the rotation rate is >5000 rpm. In practice, this rotation rate is difficult to achieve in RTPILs. A practical problem again arises from their high viscosity. At fast electrode rotation rates, air bubbles become quickly trapped at the PIL-electrode interface, producing very noisy data. Figure 9a-c shows a series of photographs taken during the course of a rotating disk voltammetric measurement in a viscous PIL, [DEA][HSO4]. A large number of air bubbles are already accumulated in the bulk solution and at the electrode-PIL interface within 30 s when using a rotation rate of 1000 rpm (Figure 9b). Many of these air bubbles even remained trapped in the viscous ionic liquid 30 min after completion of the experiment (Figure 9c). Ultramicroelectrodes are also often used to achieve steady state measurements in electrochemical, e.g., in scanning electrochemical applications microscopy.70–75 Under microelectrode slow scan rate conditions, radial diffusion is expected to dominate and give rise to a steady state rather than transient response. Since the electrode is used in the stationary mode, the problems with air bubbles produced by electrode rotation and hence stirring can be avoided. Nevertheless, the transient


Zhao et al.

Figure 9. Photographs of electrochemical cell (a) before, (b) during, and (c) after a rotating disk voltammetric experiment in the viscous ionic liquid, [DEA][HSO4]. Image b was taken after rotation for 30 s at 1000 rpm, and image c was taken 30 min after the experiment.

response still dominates because of the small diffusion coefficient. Simulations mimicking the reduction of Cc+ at a 5 µm radius microelectrode at the scan rate of 1 mV s-1 revealed that a near steady state voltammogram could be achieved when DCc+ ) 5.7 × 10-8 cm2 s-1. In the case of a relatively viscous PIL, e.g., [DEA][Ac] where DCc+ ) 7.7 × 10-9 cm2 s-1, simulations predict steady state could only be achieved at very low scan rates of

Electrochemistry of Room Temperature Protic Ionic Liquids - The

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