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Micropolarity and aggregation behavior in ionic liquid + organic solvent solutions
Fluid Phase Equilibria 248 (2006) 211–216
Micropolarity and aggregation behavior in ionic liquid + organic solvent solutions Wenjing Li, Zhaofu Zhang, Jianling Zhang, Buxing Han ∗ , Bo Wang, Minqiang Hou, Ye Xie Beijing National Laboratory for Molecular Sciences, Institute of Chemistry, Chinese Academy of Sciences, Beijing 100080, China Received 8 December 2005; received in revised form 10 August 2006; accepted 10 August 2006 Available online 18 August 2006
Abstract UV–vis spectroscopy and conductivity measurement techniques were used to study the physicochemical and structural properties of the binary or ternary mixtures of 1-n-butyl-3-methylimidazolium hexafluorophosphate ([bmim][PF6 ]) + organic solvent and 1-n-butyl-3-methylimidazolium tetrafluoroborate ([bmim][BF4 ]) + organic solvent systems. The solvents involved were acetonitrile, water, ethanol, ethyl acetate, and tetrahydrofuran. It was indicated that the micropolarity and the aggregation behavior of the mixtures depend strongly on the dielectric constants of the solvents and the composition of the mixtures. © 2006 Elsevier B.V. All rights reserved. Keywords: Ionic liquids; Micropolarity; Aggregation
1. Introduction Ionic liquids (ILs), organic salts, which are liquids at or close to room temperature, have received considerable attention recently as potential replacements for conventional volatile solvents. ILs have negligible vapor pressure, excellent thermal stability, high electrical conductivity, wide electrochemical window, strong ability to dissolve many chemicals, and can be easily recycled for reused. All of the unusual properties make ILs attractive solvents in electrochemistry [1,2], separations [3,4], chemical reactions [5–7], and material synthesis [8]. It was demonstrated that mixed solvents could be used to adjust the reaction rate and selectivity [9,10]. Moreover, in practical applications, many components exist in the ILs systems, such as reactants, catalysts, products or extractants. Thus, understanding of the structural and physicochemical properties of the mixed systems is desirable. Studies on IL + organic solvent mixtures become more and more attractive in recent years. This drives a rapid increase in the mapping of the physical properties of IL-based mixtures. Brennecke and coworkers investigated the phase behavior of imidazolium-based
∗
Corresponding author. Tel.: +86 10 62562821; fax: +86 10 62559373. E-mail address: [email protected] (B. Han).
0378-3812/$ – see front matter © 2006 Elsevier B.V. All rights reserved. doi:10.1016/j.fluid.2006.08.013
ILs with water and alcohols, respectively [11,12]. Seddon et al. studied the effect of chloride, water and organic solvents on viscosity, density and 1 H NMR chemical shift of ILs [13]. Some other properties, such as activity coefficients, gas solubility, and transport properties, have also been reported [14,15]. The behavior of solvatochromic probe in [bmim][PF6 ] + water, [bmim][PF6 ] + ethanol, and [bmim][PF6 ] + ethanol + water systems were examined [16–18], and the authors observed significant deviations from predicted values. More recently, some reports about the structure properties in ILs and IL-contained solutions have appeared in the literature. The unusual optical behaviors of fluorescent probe in ILs [19,20] and ILs themselves [21] were studied, which suggests the existence of special structure in ILs. On the other hand, several simulation studies have been made to provide information on the structure of the neat ILs [22,23]. The molecular state of water was investigated in some ILs. It is shown that the water molecules are separated from each other at low content, while the self-associated structures are formed at higher concentration [24,25]. Several researchers found that short-chain ILs could form aggregates in water [26] and CHCl3 [27]. However, compared with the intensive study on the pure ILs, the research on the IL + organic solvent mixture is just at early stage and the fundamental information about the structural and physicochemical properties remains to be investigated further. It is no doubt
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that study on this is of great importance from both academic and technological points of view. In this work, UV–vis technique and conductivity measurement were used to study the physicochemical properties of the binary or ternary mixtures containing [bmim][BF4 ] and [bmim][PF6 ]. Different solvents were involved, including acetonitrile, water, ethanol, ethyl acetate, and tetrahydrofuran. 2. Experimental 2.1. Materials
Table 1 π* values for ILs and some conventional solvents and the dielectric constants of the solvents used Solvent
v (cm−1 )
π*
ε (20 ◦ C)
[bmim][BF4 ] [bmim][PF6 ] Water Dimethyl sulfoxide Ethanol + watera Acetonitrile Ethanol Ethyl acetate Cyclohexane
24.53 24.69 23.70 24.56 25.07 25.30 25.79 26.28 28.14
1.008 0.964 1.240 1.000 0.858 0.788 0.656 0.520 0.000
80.37 [33] 47.24 [34] 34 [33] 36.64 [34] 25.3 [34] 6.081 [34] 2.02b [34]
a
The probe N,N-dimethyl-4-nitroaniline (Acros Organics, 99% purity) was used as received. All organic solvents (A.R. grade) were purchased from Beijing Chemical Reagent Factory and dried using appropriate drying agents. The doubly distilled deionized water was used. [bmim][BF4 ] and [bmim][PF6 ] were synthesized and characterized following procedures reported in the literature [28]. The ILs were dried under high vacuum at 70 ◦ C for at least 12 h prior to use. 2.2. Spectroscopic studies The UV–vis absorption spectra were recorded on a Model TU-1201 spectrophotometer with a resolution of 0.1 nm. Quartz cuvettes of 0.5 cm path length with plugs were used. The experiments were conducted at 25 ◦ C. In a typical experiment, known amount of solution of the probe in ethanol was transferred to a quartz cuvette. The cuvette was flushed with nitrogen to remove the ethanol. Suitable amount of IL or IL + organic solvent mixture was charged into the cuvette, which was sealed immediately with the plug. The measurement was done after complete dissolution of the probe, which was confirmed by the fact that the UV–vis absorption spectrum was not changed with time. A DT100 balance with a resolution of ±0.1 mg (Shanghai Science Instrument Company) was used to determine the masses of the chemicals used, and at least 0.5 g chemical was weighed each time. The concentrations of N,N-dimethyl-4-nitroaniline in all samples were lower than 5 × 10−5 M, and absorbance of all samples measured was in the range of 0.3–0.6. Experiments showed that the wavelength at the largest absorbance (λmax ) was independent of the concentration of the probe in the concentration range studied in this work. At each condition, the value of λmax was the average of four scans, and the background spectrum of the solvent has been subtracted. 2.3. Conductivity measurement The conductivity was determined by a conductivity meter which was produced by Shanghai Precision Scientific Instrument Co., Ltd. (DDS-307). The cell constant was calibrated with KCl aqueous solutions of different concentrations. The repeatability and the accuracy of the measurements were 0.3% and 1%, respectively.
b
Containing 60 mol% ethanol. 25 ◦ C.
3. Results and discussion 3.1. Spectroscopic study The polarity parameter π* measures the non-specific part of van der Waals interactions between solvents and solutes, such as dispersive, inductive and electrostatic interactions. That is, it indicates the dipolarity and polarizability of a solvent [29–32]. π* is related to the electronic transition energy of the probe which is solvent-dependent and can be calculated from the position of maximum absorption of the probe, as shown in Eq. (1): π∗ (S) =
ν(S) − ν(c-C6 H12 ) ν(DMSO) − ν(c-C6 H12 )
(1)
where ν is the wave number (cm−1 ) of the probe in solvent S at maximum absorption and in two reference solvents, cyclohexane (π* = 0.00) and dimethyl sulfoxide (DMSO) (π* = 1.00). Table 1 presents the π* values of the two ILs measured in this work using N,N-dimethyl-4-nitroaniline as the probe, which agree with the literature values [31]. Clearly, the ILs have high π* values and the polarity/polarizability is in the order [bmim][BF4 ] > [bmim][PF6 ]. The reasonable explanation is that the charge on the anions becomes more delocalized and the strength of the interaction with probe decreases with increasing anion size [35,36]. π* values and dielectric constants of other solvents involved in this work are also listed in Table 1. For an binary solvent mixture where the local composition around the probe is the same as the bulk composition, the probe response (π* for the probe used here) is given by [37–39] π∗ (mixture) = Φ1 π1∗ + Φ2 π2∗
(2)
where Φ1 and Φ2 are, respectively, the volume fractions of the organic solvent and IL in the mixture which were calculated simply from the densities and the masses of the chemicals used, and π1∗ and π2∗ are their π* values. From the equation above, it is clear that π* (mixture) should be linear function of Φ2 if the local composition around the probe is the same as the bulk composition. Fig. 1 shows the dependence of π* of [bmim][PF6 ] + organic solvent mixtures on the volume fraction of the IL (Φ2 ). Rogers and coworker [40] reported that the relative hydropho-
W. Li et al. / Fluid Phase Equilibria 248 (2006) 211–216
213
Fig. 1. Dependence of π* of [bmim][PF6 ] + organic solvent mixtures on the volume fraction of [bmim][PF6 ] (Φ2 ). Solid points are experimental values; dotted lines are calculated results from Eq. (2).
Fig. 2. Dependence of π* of [bmim][BF4 ] + organic solvent mixtures on volume fraction of [bmim][BF4 ](Φ2 ); solid points are experimental values; dotted lines are calculated results from Eq. (2).
bic [bmim][PF6 ] is totally miscible with aqueous ethanol as the mole fraction of ethanol is in the range of 0.5–0.9, though the solubility of the IL in either water or ethanol is very limited. In the present work, the aqueous ethanol with 0.6 mol fraction of ethanol is used. The π* value of [bmim][PF6 ] + actetonitrile and [bmim] [PF6 ] + ethyl acetate mixtures increases as the concentration of the IL increases in the whole concentration range. For the ternary system, the π* value first increases with the addition of the IL at Φ2 < 0.7 and then decreases at Φ2 > 0.7. In the certain concentration range, π* value can be higher than those both in the IL and the ethanol + water mixture, and is close to the value in water. The high response of probe in [bmim][PF6 ] + ethanol + water mixture can be contributed to the preferential solvation of the probe by water [18], or there may exist water-rich domains in the system in this concentration range. Moreover, we can known from Fig. 1 that the π* values for [bmim][PF6 ] + ethanol + water and [bmim][PF6 ] + ethyl acetate systems deviate obviously from linearity with changing of Φ2 , while the plot of π* versus Φ2 for [bmim][PF6 ] + acetonitrile mixture is nearly linear. Generally, the deviation from the linearity is partly due to the so-called preferential solvation induced by probe [38,41]. When a probe is added to a mixed solvent, it may interact differently with different components in the mixture, which results in the difference of the compositions of the solvation shell and that of the bulk solution. On the other hand, the non-linear behavior also depends on the change in the microscopic structure of the solvent. Specially, a component may self-associate or form aggregates in the mixture in the absence of the probe [42–44]. Thus, the π* value may derivate from the linear behavior due to the micro-heterogeneity. By careful examination of Fig. 1 we can observe that the drastic change in the solvent polarity occur in the [bmim][PF6 ] + ethanol + water and [bmim][PF6 ] + ethyl acetate systems in the dilute region. From association point of view, we may assume that the two components in [bmim][PF6 ] + acetonitrile mixture are miscible on the molecu-
lar level, i.e., there is no obvious aggregation in the solution. For the other two systems, the reasonable explanation for the obvious non-linear behavior is that the aggregation of [bmim][PF6 ] occurs. The π* values of some mixtures containing [bmim][BF4 ] were also determined, and the solvents were aqueous ethanol (60 mol% ethanol), acetonitrile, and ethanol. The [bmim][BF4 ] + ethyl acetate mixture was not studied because of the poor miscibility. The π* values of the systems at various concentrations are shown in Fig. 2. Clear deviations of π* from linearity is observed when [bmim][BF4 ] is mixed with water and aqueous ethanol, respectively, and the break point in the π* –Φ2 curves appears earlier in the former mixture. However, the degree of non-linearity in [bmim][BF4 ] + acetonitrile mixture is very small in the entire composition range. As discussed above, this indicates that aggregation in [bmim][BF4 ] + acetonitrile system is not considerable. In contrast, the microscopic structure of the other two systems changes as the composition of the mixtures is changed. The dramatic change of π* value for the [bmim][BF4 ] + ethanol mixture in very dilute region hints that the aggregates of the IL or IL-rich domains exist even its concentration is very low. 3.2. Conductivity study As discussed above, aggregation of the molecules in the mixtures may results from the preferential solvation induced by probe. In order to clarify the formation of IL aggregates in the systems in the absence of the probe, we further measured the conductivity of the different IL + organic solvent mixtures. The conductivity measurement has been widely used to evaluate the critical micelle concentration (cmc) of surfactants. According to Debye–H¨uckel–Onsager equation, the molar conductivity (Λm ) of a salt in a diluted solution should be linear with the square root of its concentration: Λm = Λ0m − AC1/2
(3)
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W. Li et al. / Fluid Phase Equilibria 248 (2006) 211–216
Fig. 3. The molar conductivity of [bmim][PF6 ] in different solutions at 25 ◦ C. (a) [bmim][PF6 ] + ethanol + water; (b) [bmim][PF6 ] + acetonitrile; (c) [bmim][PF6 ] + ethyl acetate; (d) [bmim][PF6 ] + tetrahydrofuran; the unit of C is mmol L−1 .
where Λ0m is the limiting molar conductivity of the solute, and A is an constant. Therefore, cmc can be easily determined by the break point in the curve of Λm plotted against C1/2 . Researchers have also employed the method to investigate the aggregation behavior of ILs in aqueous solutions [26]. Fig. 3 presents the dependence of the molar conductivity of [bmim][PF6 ] on the concentration of the IL in different solutions. In the case of [bmim][PF6 ] + ethanol + water systems (Fig. 3a), Λm decreases as the IL concentration increases in the concentration range studied due to the increased ion-ion interaction strength. An obvious change in the slope of the curve can be observed at C = 505 mmol L−1 , indicating the change of the electrical charged species. The reason may be that [bmim][PF6 ] exists mainly as free ions at concentrations before the break point. With increasing the concentration of [bmim][PF6 ] aggregates of the IL begin to form in the solution. This is consistent with the argument obtained from UV spectroscopy. The associ-
ation of the ILs results mainly from the hydrogen-bond interactions between the hydrogen atoms of the cations and the anions [45–48], the probable – staking interaction between the imidazolium rings [49,50], the aliphatic interaction [22,23], as well as the Coulombic interaction. The molar conductivity of IL in [bmim][PF6 ] + acetonitrile solution demonstrates a smooth curve with no distinct break point. The mixture shows almost linear Λm –C1/2 behavior in the relatively low concentration region, as can be seen from Fig. 3b. This further confirms that [bmim][PF6 ] exist mainly as separate free ions in the [bmim][PF6 ] + acetonitrile mixtures. Fig. 3c shows a very different electrical conducting behavior of [bmim][PF6 ] in [bmim][PF6 ] + ethyl acetate solution. In fact, the molar conductivity was found to increase with IL concentration in the dilute region. After reaching a maximum at C = 1101 mmol L−1 , the molar conductivity of the IL solution decreases up to that of the pure IL. This is similar to the conduc-
W. Li et al. / Fluid Phase Equilibria 248 (2006) 211–216
tivity of some 1,3-dialkyl-imidazolium salts in CHCl3 –CHCN mixed solvents [27]. This kind of conducting behavior has also been observed for electrolyte solutions of low dielectric constant solvents [51–53], owing to the formation of ion pairs and larger neutral aggregates. The maximum of molar conductivity in Fig. 3c can be interpreted qualitatively in the following. The conductivity is related with the number of charge carriers and ion mobility [54]. The degree of aggregation of the ions and viscosity of the mixture change as the composition of the solution is varied. It is reasonable to suppose that equilibria between ions, ion pairs, and larger aggregates of the IL exist in the system, and increasing the concentration of the IL is favorable to forming large aggregates. NMR spectroscopy has shown that there is significant cation–anion interaction in the solution of imidazolium salts in CDCl3 and ion pairs most probably exist [49,55]. The charges in the ion pairs are almost neutralized and the contribution to the conductivity is small. With increasing IL concentration, larger IL domains are formed and the degree of ionization of the IL is enhanced, which is favorable to increasing the conductivity. On the other hand, the viscosity of the mixture increases with increasing concentration of the IL, which is not favorable to enhancing the conductivity. The two factors affect the conductivity in opposite ways. In the lower concentration range, the first fact is dominant and the conductivity increases with increasing concentration. In the higher concentration range, the second factor becomes dominant and the conductivity is reduced as the concentration increases. We also investigated the effect of concentration on the conductivity of [bmim][PF6 ] in tetrahydrofuran, which is another low dielectric solvent. The results are presented in Fig. 3d. The profile of the curve is similar to that of Fig. 3c, suggesting that [bmim][PF6 ] + ethyl acetate and [bmim]PF6 + THF have similar aggregation behavior. The data in Fig. 3 indicate that the Λm of [bmim][PF6 ] at the same concentration is highest in acetonitrile, indicating that [bmim][PF6 ] is ionized more efficiently in acetonitrile. [bmim][PF6 ] has medium conductivity in aqueous ethanol, while the values of Λm in ethyl acetate and THF are much lower, which further support the argument about the formation of ion pairs in the solvents with lower dielectric constants. The conductivity data for [bmim][BF4 ] + organic solvent systems determined in this work are presented in Fig. 4. The same organic solvents were used as in the polarity measurement. The dielectric constants of these solvents are in the order ethanol (ε = 25) < aqueous ethanol (60 mol% ethanol) (ε ≈ 34) < acetonitrile (ε = 36). As can be seen from the figure that the break point for [bmim][BF4 ] + ethanol system appears at C = 45 mmol L−1 , and for [bmim][BF4 ] + ethanol + water system it occurs at higher concentration (C = 350 mmol L−1 ). This suggests that the IL in [bmim][BF4 ] + ethanol begin to aggregate at lower concentration than that in [bmim][BF4 ] + ethanol + water mixture. Fig. 4c shows that [bmim][BF4 ] + acetonitrile mixture has no any distinct break point, suggesting that aggregates are formed in solution. From the conductivity study we can deduce that the higher the dielectric constant of the organic solvent, the higher the critical aggregation concentration of the IL.
215
Fig. 4. The molar conductivity of [bmim][BF4 ] in different solutions at 25 ◦ C. (a) [bmim][BF4 ] + ethanol; (b) [bmim][BF4 ] + ethanol + water; (c) [bmim][BF4 ] + acetonitrile; the unit of C is mmol L−1 .
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4. Conclusions The properties of some solutions containing [bmim][PF6 ] and [bmim][BF4 ] have been investigated by combination of UV–vis spectroscopy and conductivity measurement. It is indicated that there exist aggregates in [bmim][PF6 ] + ethanol + water, [bmim][PF6 ] + ethyl acetate, [bmim][PF6 ] + tetrahydrofuran, [bmim][BF4 ] + ethanol, and [bmim][BF4 ] + ethanol + water mixtures, and the degree of aggregation depends on the dielectric constants of the solvents. The aggregates of the ILs form at lower concentration in the solvents with smaller dielectric constants, and the probe molecule can resident in the IL droplets. In contrast, there is no obvious aggregation [bmim][PF6 ] + acetonitrile and [bmim][BF4 ] + acetonitrile mixtures, and the mixtures show almost linear π* –Φ2 behaviors. The hydrophobic IL [bmim][PF6 ] have unusual conductivity behavior in weak polar solvents, such as ethyl acetate and tetrahydrofuran, which can be explained based on existence of equilibria between ions, ion pairs, and larger aggregates of the IL exist in the systems. Acknowledgement The authors are grateful to the National Natural Science Foundation of China (20533010). References [1] A. Noda, M. Watanabe, Electrochim. Acta 45 (2000) 1265–1270. [2] S. Forsyth, J. Golding, D.R. MacFarlane, M. Forsyth, Electrochim. Acta 46 (2001) 1753–1757. [3] K.E. Gutowski, G.A. Broker, H.D. Willauer, J.G. Huddleston, R.P. Swatloski, J.D. Holbrey, R.D. Rogers, J. Am. Chem. Soc. 125 (2003) 6632–6633. [4] W.Z. Wu, B.X. Han, H.X. Gao, Z.M. Liu, T. Jiang, J. Huang, Angew. Chem. Int. Ed. 43 (2004) 2415–2417. [5] T. Welton, Chem. Rev. 99 (1999) 2071–2084. [6] R.D. Rogers, K.R. Seddon, Science 302 (2003) 792–793. [7] P. Wasserscheid, W. Keim, Angew. Chem. Int. Ed. 39 (2000) 3772–3789. [8] Y. Zhou, M. Antonietti, Adv. Mater. 15 (2003) 1452–1455. [9] J. Mo, S.F. Liu, J.L. Xiao, Tetrahedron 61 (2005) 9902–9907. [10] D.W. Kim, D.J. Hong, J.W. Seo, H.S. Kim, H.K. Kim, C.E. Song, D.Y. Chi, J. Org. Chem. 69 (2004) 3186–3189. [11] J.L. Anthony, E.J. Maginn, J.F. Brennecke, J. Phys. Chem. B 105 (2001) 10942–10949. [12] J.M. Crosthwaite, S.N.V.K. Aki, E.J. Maginn, J.F. Brennecke, Fluid Phase Equilib. 228/229 (2005) 303–309. [13] K.R. Seddon, A. Stark, M.J. Torres, Pure Appl. Chem. 72 (2000) 2275–2287. [14] A. Heintz, J. Chem. Thermodyn. 37 (2005) 525–535. [15] Z.M. Liu, W.Z. Wu, B.X. Han, Z.X. Dong, G.Y. Zhao, J.Q. Wang, T. Jiang, G.Y. Yang, Chem. Eur. J. 9 (2003) 3897–3903. [16] K.A. Fletcher, S. Pandey, Appl. Spectrosc. 56 (2002) 266–271. [17] K.A. Fletcher, S. Pandey, Appl. Spectrosc. 56 (2002) 1498–1503. [18] K.A. Fletcher, S. Pandey, J. Phys. Chem. B 107 (2003) 13532. [19] P.K. Mandal, M. Sarkar, A. Samanta, J. Phys. Chem. A 108 (2004) 9048–9053.
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Micropolarity and aggregation behavior in ionic liquid + organic solvent solutions Wenjing Li, Zhaofu Zhang, Jianling Zhang, Buxing Han ∗ , Bo Wang, Minqiang Hou, Ye Xie Beijing National Laboratory for Molecular Sciences, Institute of Chemistry, Chinese Academy of Sciences, Beijing 100080, China Received 8 December 2005; received in revised form 10 August 2006; accepted 10 August 2006 Available online 18 August 2006
Abstract UV–vis spectroscopy and conductivity measurement techniques were used to study the physicochemical and structural properties of the binary or ternary mixtures of 1-n-butyl-3-methylimidazolium hexafluorophosphate ([bmim][PF6 ]) + organic solvent and 1-n-butyl-3-methylimidazolium tetrafluoroborate ([bmim][BF4 ]) + organic solvent systems. The solvents involved were acetonitrile, water, ethanol, ethyl acetate, and tetrahydrofuran. It was indicated that the micropolarity and the aggregation behavior of the mixtures depend strongly on the dielectric constants of the solvents and the composition of the mixtures. © 2006 Elsevier B.V. All rights reserved. Keywords: Ionic liquids; Micropolarity; Aggregation
1. Introduction Ionic liquids (ILs), organic salts, which are liquids at or close to room temperature, have received considerable attention recently as potential replacements for conventional volatile solvents. ILs have negligible vapor pressure, excellent thermal stability, high electrical conductivity, wide electrochemical window, strong ability to dissolve many chemicals, and can be easily recycled for reused. All of the unusual properties make ILs attractive solvents in electrochemistry [1,2], separations [3,4], chemical reactions [5–7], and material synthesis [8]. It was demonstrated that mixed solvents could be used to adjust the reaction rate and selectivity [9,10]. Moreover, in practical applications, many components exist in the ILs systems, such as reactants, catalysts, products or extractants. Thus, understanding of the structural and physicochemical properties of the mixed systems is desirable. Studies on IL + organic solvent mixtures become more and more attractive in recent years. This drives a rapid increase in the mapping of the physical properties of IL-based mixtures. Brennecke and coworkers investigated the phase behavior of imidazolium-based
∗
Corresponding author. Tel.: +86 10 62562821; fax: +86 10 62559373. E-mail address: [email protected] (B. Han).
0378-3812/$ – see front matter © 2006 Elsevier B.V. All rights reserved. doi:10.1016/j.fluid.2006.08.013
ILs with water and alcohols, respectively [11,12]. Seddon et al. studied the effect of chloride, water and organic solvents on viscosity, density and 1 H NMR chemical shift of ILs [13]. Some other properties, such as activity coefficients, gas solubility, and transport properties, have also been reported [14,15]. The behavior of solvatochromic probe in [bmim][PF6 ] + water, [bmim][PF6 ] + ethanol, and [bmim][PF6 ] + ethanol + water systems were examined [16–18], and the authors observed significant deviations from predicted values. More recently, some reports about the structure properties in ILs and IL-contained solutions have appeared in the literature. The unusual optical behaviors of fluorescent probe in ILs [19,20] and ILs themselves [21] were studied, which suggests the existence of special structure in ILs. On the other hand, several simulation studies have been made to provide information on the structure of the neat ILs [22,23]. The molecular state of water was investigated in some ILs. It is shown that the water molecules are separated from each other at low content, while the self-associated structures are formed at higher concentration [24,25]. Several researchers found that short-chain ILs could form aggregates in water [26] and CHCl3 [27]. However, compared with the intensive study on the pure ILs, the research on the IL + organic solvent mixture is just at early stage and the fundamental information about the structural and physicochemical properties remains to be investigated further. It is no doubt
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that study on this is of great importance from both academic and technological points of view. In this work, UV–vis technique and conductivity measurement were used to study the physicochemical properties of the binary or ternary mixtures containing [bmim][BF4 ] and [bmim][PF6 ]. Different solvents were involved, including acetonitrile, water, ethanol, ethyl acetate, and tetrahydrofuran. 2. Experimental 2.1. Materials
Table 1 π* values for ILs and some conventional solvents and the dielectric constants of the solvents used Solvent
v (cm−1 )
π*
ε (20 ◦ C)
[bmim][BF4 ] [bmim][PF6 ] Water Dimethyl sulfoxide Ethanol + watera Acetonitrile Ethanol Ethyl acetate Cyclohexane
24.53 24.69 23.70 24.56 25.07 25.30 25.79 26.28 28.14
1.008 0.964 1.240 1.000 0.858 0.788 0.656 0.520 0.000
80.37 [33] 47.24 [34] 34 [33] 36.64 [34] 25.3 [34] 6.081 [34] 2.02b [34]
a
The probe N,N-dimethyl-4-nitroaniline (Acros Organics, 99% purity) was used as received. All organic solvents (A.R. grade) were purchased from Beijing Chemical Reagent Factory and dried using appropriate drying agents. The doubly distilled deionized water was used. [bmim][BF4 ] and [bmim][PF6 ] were synthesized and characterized following procedures reported in the literature [28]. The ILs were dried under high vacuum at 70 ◦ C for at least 12 h prior to use. 2.2. Spectroscopic studies The UV–vis absorption spectra were recorded on a Model TU-1201 spectrophotometer with a resolution of 0.1 nm. Quartz cuvettes of 0.5 cm path length with plugs were used. The experiments were conducted at 25 ◦ C. In a typical experiment, known amount of solution of the probe in ethanol was transferred to a quartz cuvette. The cuvette was flushed with nitrogen to remove the ethanol. Suitable amount of IL or IL + organic solvent mixture was charged into the cuvette, which was sealed immediately with the plug. The measurement was done after complete dissolution of the probe, which was confirmed by the fact that the UV–vis absorption spectrum was not changed with time. A DT100 balance with a resolution of ±0.1 mg (Shanghai Science Instrument Company) was used to determine the masses of the chemicals used, and at least 0.5 g chemical was weighed each time. The concentrations of N,N-dimethyl-4-nitroaniline in all samples were lower than 5 × 10−5 M, and absorbance of all samples measured was in the range of 0.3–0.6. Experiments showed that the wavelength at the largest absorbance (λmax ) was independent of the concentration of the probe in the concentration range studied in this work. At each condition, the value of λmax was the average of four scans, and the background spectrum of the solvent has been subtracted. 2.3. Conductivity measurement The conductivity was determined by a conductivity meter which was produced by Shanghai Precision Scientific Instrument Co., Ltd. (DDS-307). The cell constant was calibrated with KCl aqueous solutions of different concentrations. The repeatability and the accuracy of the measurements were 0.3% and 1%, respectively.
b
Containing 60 mol% ethanol. 25 ◦ C.
3. Results and discussion 3.1. Spectroscopic study The polarity parameter π* measures the non-specific part of van der Waals interactions between solvents and solutes, such as dispersive, inductive and electrostatic interactions. That is, it indicates the dipolarity and polarizability of a solvent [29–32]. π* is related to the electronic transition energy of the probe which is solvent-dependent and can be calculated from the position of maximum absorption of the probe, as shown in Eq. (1): π∗ (S) =
ν(S) − ν(c-C6 H12 ) ν(DMSO) − ν(c-C6 H12 )
(1)
where ν is the wave number (cm−1 ) of the probe in solvent S at maximum absorption and in two reference solvents, cyclohexane (π* = 0.00) and dimethyl sulfoxide (DMSO) (π* = 1.00). Table 1 presents the π* values of the two ILs measured in this work using N,N-dimethyl-4-nitroaniline as the probe, which agree with the literature values [31]. Clearly, the ILs have high π* values and the polarity/polarizability is in the order [bmim][BF4 ] > [bmim][PF6 ]. The reasonable explanation is that the charge on the anions becomes more delocalized and the strength of the interaction with probe decreases with increasing anion size [35,36]. π* values and dielectric constants of other solvents involved in this work are also listed in Table 1. For an binary solvent mixture where the local composition around the probe is the same as the bulk composition, the probe response (π* for the probe used here) is given by [37–39] π∗ (mixture) = Φ1 π1∗ + Φ2 π2∗
(2)
where Φ1 and Φ2 are, respectively, the volume fractions of the organic solvent and IL in the mixture which were calculated simply from the densities and the masses of the chemicals used, and π1∗ and π2∗ are their π* values. From the equation above, it is clear that π* (mixture) should be linear function of Φ2 if the local composition around the probe is the same as the bulk composition. Fig. 1 shows the dependence of π* of [bmim][PF6 ] + organic solvent mixtures on the volume fraction of the IL (Φ2 ). Rogers and coworker [40] reported that the relative hydropho-
W. Li et al. / Fluid Phase Equilibria 248 (2006) 211–216
213
Fig. 1. Dependence of π* of [bmim][PF6 ] + organic solvent mixtures on the volume fraction of [bmim][PF6 ] (Φ2 ). Solid points are experimental values; dotted lines are calculated results from Eq. (2).
Fig. 2. Dependence of π* of [bmim][BF4 ] + organic solvent mixtures on volume fraction of [bmim][BF4 ](Φ2 ); solid points are experimental values; dotted lines are calculated results from Eq. (2).
bic [bmim][PF6 ] is totally miscible with aqueous ethanol as the mole fraction of ethanol is in the range of 0.5–0.9, though the solubility of the IL in either water or ethanol is very limited. In the present work, the aqueous ethanol with 0.6 mol fraction of ethanol is used. The π* value of [bmim][PF6 ] + actetonitrile and [bmim] [PF6 ] + ethyl acetate mixtures increases as the concentration of the IL increases in the whole concentration range. For the ternary system, the π* value first increases with the addition of the IL at Φ2 < 0.7 and then decreases at Φ2 > 0.7. In the certain concentration range, π* value can be higher than those both in the IL and the ethanol + water mixture, and is close to the value in water. The high response of probe in [bmim][PF6 ] + ethanol + water mixture can be contributed to the preferential solvation of the probe by water [18], or there may exist water-rich domains in the system in this concentration range. Moreover, we can known from Fig. 1 that the π* values for [bmim][PF6 ] + ethanol + water and [bmim][PF6 ] + ethyl acetate systems deviate obviously from linearity with changing of Φ2 , while the plot of π* versus Φ2 for [bmim][PF6 ] + acetonitrile mixture is nearly linear. Generally, the deviation from the linearity is partly due to the so-called preferential solvation induced by probe [38,41]. When a probe is added to a mixed solvent, it may interact differently with different components in the mixture, which results in the difference of the compositions of the solvation shell and that of the bulk solution. On the other hand, the non-linear behavior also depends on the change in the microscopic structure of the solvent. Specially, a component may self-associate or form aggregates in the mixture in the absence of the probe [42–44]. Thus, the π* value may derivate from the linear behavior due to the micro-heterogeneity. By careful examination of Fig. 1 we can observe that the drastic change in the solvent polarity occur in the [bmim][PF6 ] + ethanol + water and [bmim][PF6 ] + ethyl acetate systems in the dilute region. From association point of view, we may assume that the two components in [bmim][PF6 ] + acetonitrile mixture are miscible on the molecu-
lar level, i.e., there is no obvious aggregation in the solution. For the other two systems, the reasonable explanation for the obvious non-linear behavior is that the aggregation of [bmim][PF6 ] occurs. The π* values of some mixtures containing [bmim][BF4 ] were also determined, and the solvents were aqueous ethanol (60 mol% ethanol), acetonitrile, and ethanol. The [bmim][BF4 ] + ethyl acetate mixture was not studied because of the poor miscibility. The π* values of the systems at various concentrations are shown in Fig. 2. Clear deviations of π* from linearity is observed when [bmim][BF4 ] is mixed with water and aqueous ethanol, respectively, and the break point in the π* –Φ2 curves appears earlier in the former mixture. However, the degree of non-linearity in [bmim][BF4 ] + acetonitrile mixture is very small in the entire composition range. As discussed above, this indicates that aggregation in [bmim][BF4 ] + acetonitrile system is not considerable. In contrast, the microscopic structure of the other two systems changes as the composition of the mixtures is changed. The dramatic change of π* value for the [bmim][BF4 ] + ethanol mixture in very dilute region hints that the aggregates of the IL or IL-rich domains exist even its concentration is very low. 3.2. Conductivity study As discussed above, aggregation of the molecules in the mixtures may results from the preferential solvation induced by probe. In order to clarify the formation of IL aggregates in the systems in the absence of the probe, we further measured the conductivity of the different IL + organic solvent mixtures. The conductivity measurement has been widely used to evaluate the critical micelle concentration (cmc) of surfactants. According to Debye–H¨uckel–Onsager equation, the molar conductivity (Λm ) of a salt in a diluted solution should be linear with the square root of its concentration: Λm = Λ0m − AC1/2
(3)
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Fig. 3. The molar conductivity of [bmim][PF6 ] in different solutions at 25 ◦ C. (a) [bmim][PF6 ] + ethanol + water; (b) [bmim][PF6 ] + acetonitrile; (c) [bmim][PF6 ] + ethyl acetate; (d) [bmim][PF6 ] + tetrahydrofuran; the unit of C is mmol L−1 .
where Λ0m is the limiting molar conductivity of the solute, and A is an constant. Therefore, cmc can be easily determined by the break point in the curve of Λm plotted against C1/2 . Researchers have also employed the method to investigate the aggregation behavior of ILs in aqueous solutions [26]. Fig. 3 presents the dependence of the molar conductivity of [bmim][PF6 ] on the concentration of the IL in different solutions. In the case of [bmim][PF6 ] + ethanol + water systems (Fig. 3a), Λm decreases as the IL concentration increases in the concentration range studied due to the increased ion-ion interaction strength. An obvious change in the slope of the curve can be observed at C = 505 mmol L−1 , indicating the change of the electrical charged species. The reason may be that [bmim][PF6 ] exists mainly as free ions at concentrations before the break point. With increasing the concentration of [bmim][PF6 ] aggregates of the IL begin to form in the solution. This is consistent with the argument obtained from UV spectroscopy. The associ-
ation of the ILs results mainly from the hydrogen-bond interactions between the hydrogen atoms of the cations and the anions [45–48], the probable – staking interaction between the imidazolium rings [49,50], the aliphatic interaction [22,23], as well as the Coulombic interaction. The molar conductivity of IL in [bmim][PF6 ] + acetonitrile solution demonstrates a smooth curve with no distinct break point. The mixture shows almost linear Λm –C1/2 behavior in the relatively low concentration region, as can be seen from Fig. 3b. This further confirms that [bmim][PF6 ] exist mainly as separate free ions in the [bmim][PF6 ] + acetonitrile mixtures. Fig. 3c shows a very different electrical conducting behavior of [bmim][PF6 ] in [bmim][PF6 ] + ethyl acetate solution. In fact, the molar conductivity was found to increase with IL concentration in the dilute region. After reaching a maximum at C = 1101 mmol L−1 , the molar conductivity of the IL solution decreases up to that of the pure IL. This is similar to the conduc-
W. Li et al. / Fluid Phase Equilibria 248 (2006) 211–216
tivity of some 1,3-dialkyl-imidazolium salts in CHCl3 –CHCN mixed solvents [27]. This kind of conducting behavior has also been observed for electrolyte solutions of low dielectric constant solvents [51–53], owing to the formation of ion pairs and larger neutral aggregates. The maximum of molar conductivity in Fig. 3c can be interpreted qualitatively in the following. The conductivity is related with the number of charge carriers and ion mobility [54]. The degree of aggregation of the ions and viscosity of the mixture change as the composition of the solution is varied. It is reasonable to suppose that equilibria between ions, ion pairs, and larger aggregates of the IL exist in the system, and increasing the concentration of the IL is favorable to forming large aggregates. NMR spectroscopy has shown that there is significant cation–anion interaction in the solution of imidazolium salts in CDCl3 and ion pairs most probably exist [49,55]. The charges in the ion pairs are almost neutralized and the contribution to the conductivity is small. With increasing IL concentration, larger IL domains are formed and the degree of ionization of the IL is enhanced, which is favorable to increasing the conductivity. On the other hand, the viscosity of the mixture increases with increasing concentration of the IL, which is not favorable to enhancing the conductivity. The two factors affect the conductivity in opposite ways. In the lower concentration range, the first fact is dominant and the conductivity increases with increasing concentration. In the higher concentration range, the second factor becomes dominant and the conductivity is reduced as the concentration increases. We also investigated the effect of concentration on the conductivity of [bmim][PF6 ] in tetrahydrofuran, which is another low dielectric solvent. The results are presented in Fig. 3d. The profile of the curve is similar to that of Fig. 3c, suggesting that [bmim][PF6 ] + ethyl acetate and [bmim]PF6 + THF have similar aggregation behavior. The data in Fig. 3 indicate that the Λm of [bmim][PF6 ] at the same concentration is highest in acetonitrile, indicating that [bmim][PF6 ] is ionized more efficiently in acetonitrile. [bmim][PF6 ] has medium conductivity in aqueous ethanol, while the values of Λm in ethyl acetate and THF are much lower, which further support the argument about the formation of ion pairs in the solvents with lower dielectric constants. The conductivity data for [bmim][BF4 ] + organic solvent systems determined in this work are presented in Fig. 4. The same organic solvents were used as in the polarity measurement. The dielectric constants of these solvents are in the order ethanol (ε = 25) < aqueous ethanol (60 mol% ethanol) (ε ≈ 34) < acetonitrile (ε = 36). As can be seen from the figure that the break point for [bmim][BF4 ] + ethanol system appears at C = 45 mmol L−1 , and for [bmim][BF4 ] + ethanol + water system it occurs at higher concentration (C = 350 mmol L−1 ). This suggests that the IL in [bmim][BF4 ] + ethanol begin to aggregate at lower concentration than that in [bmim][BF4 ] + ethanol + water mixture. Fig. 4c shows that [bmim][BF4 ] + acetonitrile mixture has no any distinct break point, suggesting that aggregates are formed in solution. From the conductivity study we can deduce that the higher the dielectric constant of the organic solvent, the higher the critical aggregation concentration of the IL.
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Fig. 4. The molar conductivity of [bmim][BF4 ] in different solutions at 25 ◦ C. (a) [bmim][BF4 ] + ethanol; (b) [bmim][BF4 ] + ethanol + water; (c) [bmim][BF4 ] + acetonitrile; the unit of C is mmol L−1 .
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4. Conclusions The properties of some solutions containing [bmim][PF6 ] and [bmim][BF4 ] have been investigated by combination of UV–vis spectroscopy and conductivity measurement. It is indicated that there exist aggregates in [bmim][PF6 ] + ethanol + water, [bmim][PF6 ] + ethyl acetate, [bmim][PF6 ] + tetrahydrofuran, [bmim][BF4 ] + ethanol, and [bmim][BF4 ] + ethanol + water mixtures, and the degree of aggregation depends on the dielectric constants of the solvents. The aggregates of the ILs form at lower concentration in the solvents with smaller dielectric constants, and the probe molecule can resident in the IL droplets. In contrast, there is no obvious aggregation [bmim][PF6 ] + acetonitrile and [bmim][BF4 ] + acetonitrile mixtures, and the mixtures show almost linear π* –Φ2 behaviors. The hydrophobic IL [bmim][PF6 ] have unusual conductivity behavior in weak polar solvents, such as ethyl acetate and tetrahydrofuran, which can be explained based on existence of equilibria between ions, ion pairs, and larger aggregates of the IL exist in the systems. Acknowledgement The authors are grateful to the National Natural Science Foundation of China (20533010). References [1] A. Noda, M. Watanabe, Electrochim. Acta 45 (2000) 1265–1270. [2] S. Forsyth, J. Golding, D.R. MacFarlane, M. Forsyth, Electrochim. Acta 46 (2001) 1753–1757. [3] K.E. Gutowski, G.A. Broker, H.D. Willauer, J.G. Huddleston, R.P. Swatloski, J.D. Holbrey, R.D. Rogers, J. Am. Chem. Soc. 125 (2003) 6632–6633. [4] W.Z. Wu, B.X. Han, H.X. Gao, Z.M. Liu, T. Jiang, J. Huang, Angew. Chem. Int. Ed. 43 (2004) 2415–2417. [5] T. Welton, Chem. Rev. 99 (1999) 2071–2084. [6] R.D. Rogers, K.R. Seddon, Science 302 (2003) 792–793. [7] P. Wasserscheid, W. Keim, Angew. Chem. Int. Ed. 39 (2000) 3772–3789. [8] Y. Zhou, M. Antonietti, Adv. Mater. 15 (2003) 1452–1455. [9] J. Mo, S.F. Liu, J.L. Xiao, Tetrahedron 61 (2005) 9902–9907. [10] D.W. Kim, D.J. Hong, J.W. Seo, H.S. Kim, H.K. Kim, C.E. Song, D.Y. Chi, J. Org. Chem. 69 (2004) 3186–3189. [11] J.L. Anthony, E.J. Maginn, J.F. Brennecke, J. Phys. Chem. B 105 (2001) 10942–10949. [12] J.M. Crosthwaite, S.N.V.K. Aki, E.J. Maginn, J.F. Brennecke, Fluid Phase Equilib. 228/229 (2005) 303–309. [13] K.R. Seddon, A. Stark, M.J. Torres, Pure Appl. Chem. 72 (2000) 2275–2287. [14] A. Heintz, J. Chem. Thermodyn. 37 (2005) 525–535. [15] Z.M. Liu, W.Z. Wu, B.X. Han, Z.X. Dong, G.Y. Zhao, J.Q. Wang, T. Jiang, G.Y. Yang, Chem. Eur. J. 9 (2003) 3897–3903. [16] K.A. Fletcher, S. Pandey, Appl. Spectrosc. 56 (2002) 266–271. [17] K.A. Fletcher, S. Pandey, Appl. Spectrosc. 56 (2002) 1498–1503. [18] K.A. Fletcher, S. Pandey, J. Phys. Chem. B 107 (2003) 13532. [19] P.K. Mandal, M. Sarkar, A. Samanta, J. Phys. Chem. A 108 (2004) 9048–9053.
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