J. Phys. Chem. B 2007, 111, 131-138
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Solvatochromic Studies of Ionic Liquid/Organic Mixtures Berlyn R. Mellein, Sudhir N. V. K. Aki,† Rebecca L. Ladewski, and Joan F. Brennecke* Department of Chemical and Biomolecular Engineering, UniVersity of Notre Dame, Notre Dame, Indiana 46556 ReceiVed: August 17, 2006; In Final Form: October 4, 2006
Room-temperature ionic liquids (ILs) have potential for many different applications, including catalysis and synthesis. Organics are often present during IL applications; therefore, a more fundamental understanding of the interactions between IL and organics is necessary. A systematic study of the effects of organic cosolvents, cations, and anions on the solvent strength of IL/organic mixtures will allow for a greater understanding and potential for tuning of ILs for specific purposes. Solvent strength is commonly quantified using spectroscopic probes. We report the solvent strength of IL/organic mixtures using Reichardt’s dyes 30 and 33, KamletTaft parameters, and phenol blue. The results show that the polarity of ILs is largely unaffected by the organic cosolvent; that is, the probes are preferentially solvated by the ILs. However, more specific solvation forces, such as hydrogen bonding, can be influenced indirectly by the strength of the anion/cation interaction, giving counterintuitive results.
Introduction Room-temperature ionic liquids (ILs) belong to a class of potentially benign solvents. They are molten salts that are liquid around ambient conditions, yet have immeasurable vapor pressure at normal operating conditions. The most common ILs pair N,N′-dialkylimidazolium, alkylammonium, alkylphosphonium, or N-alkylpyridinium cations with various anions. ILs have been successfully used for catalysis and synthesis1-3 and extraction.4 The choice of anion and cation allows dexterity in choosing an IL for a specific purpose. However, little is understood concerning how the anions, cations, substituents, and cosolvents interact and affect the solvent strength of the ILs. Organic solvents are often present in IL applications; thus, a way to quantify the solvent strength of ILs, organics, and their mixtures would be useful in choosing or tailoring ILs.1 A simple method for characterizing solvent strength is polarity. Polarity is a general term that refers to all the interaction forces between molecules, both specific, such as hydrogen bonding, and nonspecific, such as inductive. This can be correlated with chemical reaction rates, optical and nuclear spectra, and macroscopic properties like conductivity.5 The more consistent methods for describing polarity reflect molecular level interactions rather than macroscopic bulk properties. Solvatochromism describes the prominent change in the UVvis absorption upon change in polarity of the medium. Empirical parameters of solvent polarity using indicator dyes have been developed on the basis of the idea that the electron excitation energy of an indicator dye changes with the polarity of the solvation shell. Each indicator dye measures different aspects of polarity to a different degree. Multiparameter solvatochromic correlations take advantage of this, classifying the solvent strength into different properties. The Kamlet-Taft linear solvation energy relationship breaks down solvent strength into * Corresponding author. Tel.: +1 574 631 5874. Fax: +1 574 631 8366. E-mail address:
[email protected]. † Present address: Intermediates R&D, Invista S.a.r.l., Sabine River Laboratory B568, P.O. Box 1003, Orange, Texas 77631-1003.
dipolarity and polarizability (π*), hydrogen bond donating acidity (R), and hydrogen bond accepting basicity (β). ILs have complex interaction forces between the anion and the cation; thus, a multiparameter solvatochromic correlation can be beneficial for understanding the solvent strength. The most common measure of general polarity is the ET(30) scale, the molar electron transition energy of Reichardt’s dye 30, pyridinium N-phenolate betaine.6 The polarity, as indicated by ET(30) values, of imidazolium-based ILs is similar to those of short-chain alcohols7-12 and pyridinium-based ILs, as well.12 The ET(30) scale is dependent on the functional group of the cation for imidazolium-based ILs and follows the polarity of the functional group.8,10,13 In addition, the polarity decreases with increasing temperature.9,12 Using 4-aminophthalimide and 4-N,N-dimethylaminophthalimide, we have previously shown that imidazolium and pyridinium-based ILs have polarity similar to primary alcohols.14 Nile Red in imidazolium-based ILs also indicates a polarity similar to that of short-chain alcohols.15 However, this probe shows dependence on both the anion and the cation. Dzyuba and Bartsch showed similar trends for polarity dependence on the substituent as with ET(30) values, but their numbers were slightly lower using Nile Red.13 Another group found the polarity of ILs using Nile Red emission (rather than absorption) to be similar to that of water16 but provided no further explanation for the differences. Pyrene fluorescence spectra suggest that the polarity of imidazolium-based ILs is similar to that of short-chain alcohols17 or acetonitrile (ACN) and dimethyl sulfoxide (DMSO).16 1-Pyrenecarboxaldehyde shows emissions ranging from a low dielectric solvent like hexane for 1-ethyl-3-methylimidazolium bis(trifluoromethylsulfonyl)amide 17 to a relatively high dielectric solvent like ACN or DMSO for 1-butyl-3-methylimidazolium hexafluorophosphate ([bmim][PF6]).16 Although the results are dependent on the probe used, all of the probes indicate that ILs are highly polar. Kamlet-Taft parameters have also been measured for many pure ILs. The π* parameter is relatively high,9,10 depends on both the cation and the anion, and decreases with increasing anion delocalization.10 Baker et al. determined that R is similar
10.1021/jp0653353 CCC: $37.00 © 2007 American Chemical Society Published on Web 12/10/2006
132 J. Phys. Chem. B, Vol. 111, No. 1, 2007 to short-chain alcohols for [bmim][PF6],9 while Crowhurst et al. showed that for a variety of ILs, R is moderately high and determined by the cation.10 The β parameter is moderate for [bmim][PF6]9 as well as other imidazolium, pyridinium, and ammonium-based ILs, but depends on the anion.10 Anderson et al. have similarly reported that the hydrogen bond basicity is largely dependent on the anion, while hydrogen bond acidity is dependent on the cation and anion.18 Another measure of the solvent basicity, which correlates with the hydrogen bond accepting ability, is the shift of the maximum absorbance wavelength (λmax) of acetylacetonato-(N,N,N′,N′-tetramethylethylenediamine) copper(II) tetraphenylborate ([Cu(acac)(tmen)][BPh4]), which gives results consistent with β.10 [Cu(acac)(tmen)][BPh4] and ET(30) results show that ILs are highly polar with low nucleophilicities, a combination thought to be responsible for some of the unusual properties of ILs.8 The solvent strengths of some IL mixtures have been studied. Preliminary data on ET(30) values and the basicity, as indicated by [Cu(acac)(tmen)][ClO4], were reported for mixtures of [bmim][PF6] with acetone, ACN, methanol (MeOH), and propylenecarbonate.19 Strong preferential solvation of the probe by the IL was observed, indicating nonideal behavior, consistent with our results shown herein. The effect of water, in terms of “wet” compared to “dry” IL, has been determined on [bmim][PF6] using the ET(30) scale, showing a difference at approximately 2% by volume of water.9,11 The solvent strength of [bmim][PF6] with ethanol up to saturation has also been measured.20 Pyrene shows preferential solvation by [bmim][PF6], while 1-pyrenecarboxaldehyde and 1,3-bis-1-(pyrenyl)propane show preferential solvation by ethanol. The ET(30) scale shows a decrease in the mixture that is lower than either pure solvent. This can be explained with the Kamlet-Taft parameters because the polarity of the pure IL is higher than that of ethanol, but the hydrogen bond donating ability is lower. For the ET(30) probe, 68% of the shift is due to hydrogen bonding with the phenoxide oxygen, and they determined that this explains the unusual behavior. For mixtures with [bmim][PF6], ethanol, and water, results are similar.21 1,3-Bis-1-(pyrenyl)propane and pyrenecarboxaldehyde show preferential solvation by the IL, but the ET(30) scale shows a solvent shell rich in water. A more complete study of IL binary mixtures has been done for 1-butyl3-methylimidazolium tetrafluoroborate ([bmim][BF4]) with MeOH, ethanol, and water using Reichardt’s dye and KamletTaft parameters.22 MeOH and ethanol have similar behavior; there is preferential solvation of the probe by the IL for π* and β and no preferential solvation for R. In addition, certain compositions for R were observed to be higher than either the IL or the alcohol. It was suggested that the IL and alcohol interact to form a hydrogen bond complex that is more polar and a better hydrogen bond donor than either component alone. For [bmim][BF4]/water mixtures, there is preferential solvation by the IL for all parameters. The pure R value reported by these authors22 for MeOH and ethanol is much higher than the established literature values,5 strongly suggesting that their alcohols were contaminated with water. Unfortunately, the water content, which can substantially alter the results, was not reported. Therefore, it is difficult to draw comparisons between their results and ours reported below. Thus, although some IL binary systems have been studied, a systematic study of binary mixtures to examine the different aspects of solvent strength of the ILs is lacking. We propose a study to systematically investigate the solvent effects of ILs and IL/organic mixtures using both phenol blue and KamletTaft probes. Careful choice of organic and IL will highlight
Mellein et al. specific interactions. Phenol blue is a polarity indicator that is responsive to hydrogen bonding and is a good overall polarity indicator. Kamlet-Taft parameters reflect more specific interactions in the ILs. For the phenol blue studies, [bmim][PF6], one of the original ILs to receive intense investigation, was studied with MeOH, dichloromethane (DCM), and carbon tetrachloride. MeOH is a weaker hydrogen bond donor than the IL, while DCM and carbon tetrachloride are nonpolar. For the Kamlet-Taft studies, the ILs and organics were chosen to highlight specific interactions. 1-Hexyl-3-methylimidazolium bis(trifluoromethylsulfonyl)amide ([hmim][Tf2N]) is the IUPAC standard IL and was studied with three different organics. ACN is a polar aprotic solvent, 2-butanone (MEK) is a hydrogen bond acceptor, and 2,2,2-trifluoroethanol (TFE) is a strong hydrogen bond donor. 1-Hexyl-2,3-dimethylimidazolium bis(trifluoromethylsulfonyl)amide ([hmmim][Tf2N]) was measured with MEK. The hydrogen on the C2 position carbon of the imidazolium ring is the most acidic10 and is important for IL/organic interactions. The IL [hmmim][Tf2N] has a methyl substituent instead of an acidic hydrogen on the C2 carbon of the imidazolium ring, decreasing the hydrogen bond donating ability of the IL. The interactions with MEK should reflect this. [TfO] (in 1-hexyl-3-methylimidazolium triflate ([hmim][TfO])) is more coordinated to the cation than [Tf2N] and is a stronger hydrogen bond acceptor. This should be reflected in the interactions with TFE. In addition, the Kamlet-Taft parameters were measured for several pyridinium-based ILs for comparison with other pure ILs. The structures of all the ILs used in this study are shown in Table 1. Experimental Section Materials. The indicator probes 4-nitroaniline (Aldrich, 99+%), N,N-diethyl-4-nitroaniline (Oakwood Products, Inc.), Reichardt’s dye 33s2,6-dichloro-4-(2,4,6-triphenyl-1-pyridino)phenolate (Chemika, 99+% HPLC grade), and phenol blue (Sigma-Aldrich, 97%) were used as received. The organic cosolvents ACN (Fisher Scientific, 99.9% or optima 99.9+%), MEK (Aldrich, 99.5+% HPLC grade), TFE (Sigma-Aldrich, 99.5+% NMR grade), MeOH (Fisher Scientific, 99.9%), DCM (Sigma-Aldrich, 99.9+% PRA grade), dimethylformamide (Sigma-Aldrich, 99.9+% HPLC grade), valeronitrile (SigmaAldrich, 99.5%), and carbon tetrachloride (Sigma-Aldrich, 99.9+% HPLC grade) were used as received with the exception of ACN, MEK, and TFE. These solvents were stored over molecular sieves after the bottle was opened so that water contents were always below 100 ppm. The solvents were filtered prior to use if the molecular sieve particles were interfering with the UV absorption. IL Synthesis. The synthesis of ILs used in this study is described below. The water content was determined by Karl Fischer titration. Halide impurity was measured using an Oakton Ion 510 meter with Cole-Parmer ion specific probes. Silver content was measured using an Optima 3300 KL ICP-OES. The imidazolium-based ILs [hmim][Tf2N], 1-octyl-3-methylimidazolium bis(trifluoromethylsulfonyl)amide ([omim][Tf2N]), [hmmim][Tf2N], and [bmim][PF6] were synthesized in our laboratory, as described elsewhere.23,24 The IL [hmim][TfO] was synthesized with an anion exchange of [hmim][Br] with [Ag][TfO]. The trace amounts of silver in the sample affected the measurements with Reichardt’s dye but did not interact with either N,N-diethyl-4-nitroaniline or 4-nitroaniline. The silver caused λmax of Reichardt’s dye to shift to lower wavelengths, signifying an extremely strong hydrogen bond donor. The
Ionic Liquid/Organic Mixture Solvatochromic Study TABLE 1: List of ILs Used in the Current Study with Nomenclature and Purity
J. Phys. Chem. B, Vol. 111, No. 1, 2007 133 determine R), the [hmim][TfO] synthesized with [Ag][TfO] was used for the N,N-diethyl-4-nitroaniline and 4-nitroaniline dyes (λmax in the 300 to 350 nm range and used to determine β and π* values), and the [hmim][TfO] synthesized with [Li][TfO] was used with Reichart’s dye (to determine R). The pyridinium-based ILs [hmpy][Tf2N], [ompy][Tf2N], and [hmDMApy][Tf2N] were synthesized as described previously.25 The IL [bmpy]Tf2N] was obtained from Strem Chemicals and used as received. Table 1 provides the list of ILs used in the current study along with the nomenclature and abbreviations used. Also shown in Table 1 are the halide and water concentrations of the IL samples, as used. Sample Preparation. A stock solution of DCM with phenol blue and each of the Kamlet-Taft probes was prepared. The stock solution was added dropwise to pure IL until the absorbance was ∼1. The sample was exposed to house vacuum at 40-60 °C for 1-2 days to remove the DCM. The IL was then added to a cuvette (with a stir bar) in a glovebox under nitrogen and sealed with a septum (Sigma-Aldrich, white rubber septa). The mass was measured before and after adding the IL to obtain the mass of the sample. Organic Additions. A Cary 300 UV-visible spectrophotometer was used for measuring the maximum absorption wavelength. After the maximum absorption wavelength of the sample was measured, a syringe was used to add a small amount of organic (∼10 mol %) to the sample through the septum. The additions were done in the glovebox. The sample was stirred using a magnetic stir plate for ∼5 min or until the organic was dissolved throughout the sample. The maximum absorption wavelength was measured. The process was repeated at increments of 10 mol % organic each time unless a large difference in the maximum absorption wavelength (more than 5 nm) was seen. At this point, the composition increments were reduced to 2-5 mol % until a maximum of 99 mol % organic was reached. All measurements were made at room temperature, which is 23 ( 1 °C. Solvatochromic Probe Analysis. The π* parameter is determined from the spectroscopic shift of N,N-diethyl-4nitroaniline (probe 1, Figure 1) using eq 1.26
ν(1)max ) 27.52 - 3.182π*
(1)
The β parameter is determined using the spectroscopic shift of 4-nitroaniline (probe 2) with respect to N,N-diethyl-4nitroaniline and eq 2.27
ν(2)max ) 1.035ν(1)max - 2.80β + 2.64
(2)
The ET(30) parameter is obtained from the spectroscopic shift of Reichardt’s dye 30 (probe 3) and is simply the electron transition energy of the dissolved dye, as defined by eq 3.5 hydrogen bond donating ability of [hmim][TfO] was expected to be approximately the same as [hmim][Tf2N] since the hydrogen bond donating ability depends mainly on the cation.10 To verify that the silver was causing this shift, [Ag][TfO] was added to pure [hmim][Tf2N] with each of the three dyes. The β and π* showed no change in λmax, but R, which is determined from Reichardt’s dye, shifted to lower wavelengths when [Ag][TfO] was added. Another batch of [hmim][TfO] was synthesized from an anion exchange of [hmim][Br] and [Li][TfO]. It was more difficult to remove all of the color from the [hmim][TfO] synthesized with [Li][TfO]. It had an unacceptably high absorbance between 300 and 350 nm but gave a reasonable R value because Reichardt’s dye absorbs at higher wavelengths. Since the silver content only affects Reichardt’s dye (used to
ET(30) ) hcν˜ maxNA ) 2.8591ν(3)max
(3)
In this study, Reichardt’s dye 33 (probe 4) was sometimes used instead of Reichardt’s dye 30 because it is more stable under acidic conditions and has been used previously for studying ILs.9,28 The ET(30) parameter is determined from Reichardt’s dye 33 using the following correlation, shown in eq 4,28 where ET(33) is the electronic transition energy of Reichardt’s dye 33.
ET(30) ) 0.9986ET(33) - 8.6878
(4)
The R parameter can then be calculated using ET(30) with respect to π* from eq 5.29
134 J. Phys. Chem. B, Vol. 111, No. 1, 2007
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Figure 1. Structures of solvatochromic probes.
R ) 0.0649ET(30) - 2.03 - 0.72π*
(5)
Slightly different π*, R, and β values can be obtained depending on the exact correlation and solvatochromic probes used. This explains some minor discrepancies in Table 2 (discussed herein) for pure organics. The general polarity indicators phenol blue (probe 5) and Reichardt’s dye 30 and 33 were also analyzed according to normalized shift, shown in eq 6, where λ2 refers to the compound that is being added dropwise to the first compound.
λmix - λ2 λ1 - λ2
normalized shift, λn )
(6)
Error was calculated using the standard deviation of five or more measurements of λmax. For the following mixtures with phenol blue, less than five measurements were taken, so the error was calculated using a generous estimate of the uncertainty of (1 nm: DCM/TFE, ACN/carbon tetrachloride, MeOH/carbon tetrachloride, TFE/carbon tetrachloride, valeronitrile/carbon tetrachloride, and dimethylformamide/carbon tetrachloride. Results and Discussion Solvatochromic behavior of Reichardt’s dye 30 and 33 and phenol blue in pure ILs and their mixtures with organics was measured. We have also measured Kamlet-Taft parameters for a series of ILs and their mixtures with organics. The results from these measurements are discussed in the following sections. ET(30) ValuessPure ILs. We used Reichardt’s dye 30 and Reichardt’s dye 33 to measure ET(30) and ET(33) values, respectively, for a series of ILs and their mixtures with organics. ET(30) and ET(33) values for pure ILs are given in Table 2 along with those for pure organics and for other ILs reported in the
literature. Here we measured the ET(30) value for [bmim][PF6] and the ET(33) values for [hmim][Tf2N], [hmmim][Tf2N], and [hmim][TfO]. Comparison of our ET(30) value for [bmim][PF6] with the literature value indicates a very good agreement. Replacing the protic hydrogen with CH3 at the C2 position for [hmim][Tf2N] lowered the ET(30) value from 51.8 to 49.3 kcal/ mol. This reduction in the ET(30) value is consistent with the trends reported in the literature for [bmim][Tf2N] and [bmmim][Tf2N] ILs, and this is due to the fact that one of the hydrogen bonding sites has been eliminated. Recall that ET(30) is a measure of both nonspecific (dipolarity and polarizability) and specific (hydrogen bond donating ability) interactions. The ET(30) value for [hmim][TfO] is 52.5 kcal/mol, which is slightly higher than 51.8 kcal/mol for [hmim][Tf2N]. This behavior is similar to the [bmim][Tf2N] (51.5 kcal/mol) and [bmim][TfO] (52.3 kcal/mol) analogues, which are comparable because hydrogen bonding depends primarily on the choice of cation for these ILs. Additionally, an increase in the alkyl chain length decreased the ET(30) value marginally; compare the values for [hmim][Tf2N] and [omim][Tf2N]. A marginal decrease in the ET(30) values was observed for pyridinium-based ILs compared to imidazolium-based ILs. As noted previously,7-12 the overall polarity of ILs based on Reichardt’s dye is similar to those of short-chain alcohols. ET(30) ValuessIL/Organic Mixtures. The effect on the ET(30) values of the addition of organic cosolvents to the ILs is shown in Figure 2. The results are presented in terms of normalized shift; therefore, the probe response varied between 0 and 1 as the concentration of IL was increased from 0 to 1 mole fraction. The response of the Reichardt’s dye to the addition of ACN, DCM, and MEK to the various ILs was found to be extremely similar (all of the data on the upper half of the diagram). The choice of IL did not have much effect on the qualitative behavior. The ET(30) values rapidly increased with the addition of small amounts of IL to the organic, indicating that polar solutes will be preferentially solvated by ILs in mixtures of ILs with aprotic solvents, even if those solvents are highly polar (e.g., ACN). By contrast, the IL (whether it be [hmim][Tf2N] or [hmim][TfO]) does not dominate the interactions with the probe in mixtures with TFE. The hydrogen bond donating ability of this solvent is so great that it contributes substantially to the ET(30) value. This could happen even if the probe is preferentially solvated by the IL since hydrogen bonding is a specific interaction that is saturated by a single alcohol hydrogen bonding with the probe. In simple terms, the ET(30) values for the ILs are greater than those of the organics considered, except for TFE, so the observed behavior is not surprising. However, the strength of interaction between the probe and the IL was found to be dependent on the choice of both the IL and the organic. For example, consider the response of the Reichardt’s dye to the addition of [bmim][PF6] to ACN and DCM, where the results indicate that the dye molecule is more strongly solvated by the IL in DCM compared to ACN. This may be an indirect, rather than direct, effect. As the dielectric constant of ACN is greater than that of DCM, ACN could be interacting with the IL more strongly compared to DCM, and hence, it could reduce the ability of the IL to solvate the probe. The role of the IL on the solvatochromic behavior of Reichardt’s dye can be explained by comparing the systems [bmim][PF6]/ACN and [hmim][Tf2N]/ACN, as well as [hmim][Tf2N]/MEK and [hmmim][Tf2N]/MEK, as shown in Figure 2. In both cases, [hmim][Tf2N] seems to solvate the probe more strongly. In the case of [bmim][PF6] versus [hmim][Tf2N], this may be due to stronger anion/cation interactions for [bmim]-
Ionic Liquid/Organic Mixture Solvatochromic Study
J. Phys. Chem. B, Vol. 111, No. 1, 2007 135
TABLE 2: ET(30), ET(33), Kamlet-Taft Values, and λmax Values for Phenol Blue for a Series of ILs and Organic Solvents (Literature Values are Given in Parentheses) compound
ET(30), kcal/mol
[bmim][PF6] [bmim][Tf2N] [bmmim][Tf2N] [bmim][TfO] [hmim][Tf2N] [hmmim][Tf2N] [hmim][TfO] [omim][PF6] [omim][Tf2N] [bmpy][Tf2N] [hmpy][Tf2N] [ompy][Tf2N] [hmDMApy][Tf2N] ACN acetone MEK MeOH TFE DCM carbon tetrachloride valeronitrile dimethyl formamide 1-methyl imidazole 1-butyl imidazole cyclohexane
52.4 (52.3)8 (51.5)8 (48.6)8 (52.3)8 51.8a 49.3a 52.5a (51.2)8 51.0a 50.2a 50.4a 48.99a 48.3a 45.6 (45.6)5
a
60.6 58.04 61.24 59.8 58.95 59.20 57.76 57.07
40.9 (41.3)5 55.2 (55.4)5 60.6 (59.8)5 40.7 (40.7)5
R
π*
ET(33), kcal/mol
β
(1.032)10 (0.984)10 (1.010)10 (1.006)10 0.98 0.99 0.98
(0.634)10 (0.617)10 (0.381)10 (0.625)10 0.65 0.45 0.667
(0.207)10 (0.243)10 (0.239)10 (0.464)10 0.25 0.26 0.52
0.97 0.965 0.984 0.965 0.984 0.80 (0.799)10
0.597 0.540 0.543 0.463 0.405 0.35 (0.35)10
0.28 0.276 0.306 0.276 0.254 0.38 (0.37)10
0.67 (0.67)33 (0.60)5 1.20 (0.73)33 (0.791)10
0.14 (0.06)33 (0.98)5 1.04 (1.51)33 (0.042)10
0.58 (0.48)33 (0.66)5 0.01 (0.00)33 (-0.014)10
0.00b
0.00b
0.00b
44.6 42.1
phenol blue λmax, nm 598.0
597.2 593.0
583.0 (584)32 (582.0)32 606.8 (608)32 660.0 (660)32 587.9 563.8 (565)32 583.6 593.3 (595)32 601.0 598.1 (552.0)32
ET(30) values were estimated from ET(33) values according to eq 4. b By definition.
Figure 2. Behavior of Reichardt’s dye 30 and 33 in IL/organic mixtures at 23 °C.
Figure 3. Change in π* values with the addition of organics to imidazolium-based ILs at 23 °C.
[PF6] that limit its interactions with the probe. For [hmmim][Tf2N] versus [hmim][Tf2N], it may be decreased hydrogen bonding of the IL with the MEK that makes the MEK a more effective competitor with the IL in solvating the probe. Even though we observed small differences in the behavior of the probe molecule in various IL/organic mixtures, the important result is that the ILs strongly solvated the probe molecule in all solutions. On the other hand, the addition of TFE to [hmim][Tf2N] and [hmim][TfO] resulted in a different response from all other systems described so far. In this case, small additions of TFE to IL caused a major shift toward the alcohol; that is, the probe’s interactions with the solvent mixture are strongly affected by the TFE. As explained above, this behavior is simply due to the fact that TFE is a strong hydrogen bond donor and the probe can accept the hydrogen bond. As shown in Table 1, the R value for TFE is much greater than that for [hmim][Tf2N], and therefore, TFE interacts with the probe through hydrogen bonding more strongly than the IL.
π*sPure ILs. The π* values for a series of imidazolium and pyridinium cation-based ILs are shown in Table 2. In general, the π* values for pure ILs were higher than those for pure ACN and MEK and lower than that of TFE and were found to be independent of the type of IL. π*sIL/Organic Mixtures. The π* values for binary mixtures of [hmim][Tf2N] with ACN, MEK, and TFE are shown in Figure 3. The addition of any of the three organics to [hmim][Tf2N] does not produce a significant change in dipolarity/polarizability until the mixture is approximately 50 mol % organic, as shown in Figure 3. Note that the π* probe for [hmim][Tf2N]/TFE mixtures is not preferentially solvated by the more polar compound, as one might have assumed. This clearly indicates that the ET(30) response in IL/TFE mixtures discussed above is due to specific hydrogen bonding, not general preferential solvation. Behavior similar to that of the [hmim][Tf2N] mixtures for π* is noted with the [hmmim][Tf2N]/MEK and [hmim][TfO]/TFE mixtures. This implies that the probe N,N-diethyl4-nitroaniline is preferentially solvated by the IL until large amounts of organic are present for all mixtures investigated.
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Figure 4. Change in β values with the addition of organics to imidazolium-based ILs at 23 °C.
Figure 5. Change in R values with the addition of organics to imidazolium-based ILs at 23 °C.
βsPure ILs. The dependence of β values, the ability of the solvent to accept a hydrogen bond, on the choice of cation for a series of [Tf2N] anion-based ILs is shown in Table 2. Also shown in this table is the value for [hmim][TfO]. For [Tf2N] anion-based ILs, replacing the imidazolium cation with pyridinium-based cations had little or no effect on β values. Furthermore, the addition of functional groups to the cation also had negligible effect on β values. Replacing the [Tf2N] anion with [TfO] increased the β value for [hmim] cation-based ILs. This finding is consistent with Crowhurst et al., where it was shown that β values are sensitive primarily to the choice of anion.10 βsIL/Organic Mixtures. The effect of composition on the hydrogen bond accepting ability of [hmim][Tf2N]/organic mixtures, Figure 4, is similar to π*, with no effect of the organics until large amounts are present and the domination of interactions with the 4-nitroaniline by the IL. This is expected for TFE, which is a very weak hydrogen bond acceptor. ACN and MEK have higher values for β than [hmim][Tf2N] (Table 2), but the probe is still dominated by interactions with [hmim][Tf2N]. The IL [hmmim][Tf2N] was chosen to examine the importance of the acidic hydrogen in the C2 position on the imidazolium ring. The β values, as shown in Figure 4, express the difference due to the methylation of the C2 carbon. To interpret all of these results, one must realize that the β values do depend on the choice of anion,10 which suggests that there may be hydrogen bond interactions between the 4-nitroaniline (a hydrogen bond donor) and the anion (a hydrogen bond acceptor). Thus, there are four possible hydrogen bonding interactions in these systems: cation/anion, cation/MEK, probe/ MEK, and probe/anion, where the cation and probe act as hydrogen bond donors and the organic and anion act as hydrogen bond acceptors. For both [hmim][Tf2N]/MEK and [hmmim][Tf2N]/MEK mixtures, the π* values indicated that the probe is prefentially solvated by the IL. Yet the β values indicate preferential interaction of the probe with the IL in the [hmim][Tf2N]/MEK case but not in the [hmmim][Tf2N]/MEK mixture. Clearly, this is because β measures specific interactions, not general preferential solvation. For the [hmmim][Tf2N]/MEK mixture, the probe/anion hydrogen bonding mechanism would not be disturbed, but it is reasonable to expect that the hydrogen bonding due to cation/anion and cation/organic interactions is lessened compared to [hmim][Tf2N]/MEK. If the cation/anion interactions are interfered with, then the results should show a stronger interaction between the [hmmim][Tf2N] and the probe
because more anion would be free to interact with the probe. This is not the case. Instead, there is less interaction. Thus, the experimental results point to decreased interactions between the [hmmim] cation and the MEK. This agrees with results from Crosthwaite et al. that show the methyl substitution on the C2 carbon results in less interaction with alcohols.30,31 When MEK is unable to participate in a hydrogen bond with the cation, then it interferes with the hydrogen bonding between the probe and the anion. Thus, there is clear evidence that [hmim][Tf2N] is hydrogen bonding with both the probe 4-nitroaniline and the organic MEK. It is expected that [hmim][TfO] will interact more strongly than [hmim][Tf2N] with TFE because [hmim][TfO] is a stronger hydrogen bond acceptor and TFE is a strong hydrogen bond donor. β values (Figure 4) show different behavior. In this case, there are three hydrogen bonding possibilities: TFE/anion, probe/anion, and cation/anion. The TFE should have a stronger bond with the [TfO] anion compared to the [Tf2N] anion, as should the probe. This is likely why the interactions between [hmim][TfO] and TFE behave closer to ideal (a straight line)s the probe and the TFE exhibit similar hydrogen bond donating abilities so neither saturates the interactions with [hmim][TfO]. It seems counterintuitive that the IL with the weaker hydrogen bond accepting ability would interact strongly with the probe. Yet, this can be explained by the fact that [Tf2N] is less coordinated with [hmim] than [TfO], which allows more interactions between the anion and the probe, as well as the anion and the organic.8 If the anion is less coordinated to the cation, then it will be free to stabilize the probe and TFE with hydrogen bonding and nonspecific forces, since the probe is also susceptible to general polarity effects. RsPure ILs. The R values for a series of imidazolium and pyridinium cation-based ILs were determined, and the values are given in Table 2. For [Tf2N] anion-based ILs, replacing the imidazolium cation with the pyridinium cation lowered the R value. Changing the cation from [hmim] to [hmmim], that is, replacing the protic hydrogen with a CH3 group, lowered the R value, as expected. RsIL/Organic Mixtures. The hydrogen bond donating ability of [hmim][Tf2N] is affected by the addition of TFE, while ACN and MEK have almost no effect on the R value until ∼85 mol % organic is reached, Figure 5. The hydrogen bonding of the probe with the IL/organic solvent mixtures is dominated by the IL in mixtures with ACN and MEK. This is reasonable since the probe is most likely preferentially solvated by the IL (recall
Ionic Liquid/Organic Mixture Solvatochromic Study the π* results), and ACN and MEK are poorer hydrogen bond donors than the ILs. The behavior is the same for [hmmim][Tf2N]/MEK (Figure 5), although the R values for the pure IL are different. However, even a small amount of TFE affects the R value in the [hmim][Tf2N]/TFE mixtures. In this instance, the R value for pure TFE is essentially reached at ∼65 % organic. This indicates that TFE dominates the hydrogen bonding with the probe, even though the probe is likely preferentially solvated by the IL (recall the π* results). The strong hydrogen bond donating ability of TFE means that it can saturate the interaction with the probe, even if it is present in relatively low concentrations in the cybotactic region around the probe. The IL [hmim][TfO] has a higher β value than [hmim][Tf2N] (Table 2) and thus should interact more with the TFE. Both the [hmim][TfO]/TFE and [hmim][Tf2N]/TFE mixtures show domination of hydrogen bonding with the probe by the TFE at compositions higher that ∼65 mol % (Figure 5). However, at lower compositions, the [hmim][TfO]/TFE mixture shows less domination of hydrogen bonding with the probe by TFE, as is noted by the straighter line. This supports the idea that TFE interacts more strongly with the [TfO] anion than the [Tf2N] anion and thus is less available to hydrogen bond with the probe. Even though [hmim][Tf2N] moderately hydrogen bonds and is moderately polar, it preferentially solvates and dominates the hydrogen bonding with all of the probes except in the presence of an extremely electrophilic organic cosolvent like TFE. This is likely due to the relatively weaker association of the anion and cation in [hmim][Tf2N] compared to ILs with other anions. Phenol Blue. We also used phenol blue as a probe molecule to understand the overall polarity of ILs and their mixtures with organic solvents. Phenol blue is sensitive to both hydrogen bond donating ability as well as accepting ability of solvents. It is also influenced by the nonspecific properties of solvents such as dipole moment, dielectric constant, dipole-dipole interactions, and dipole-induced dipole interactions. It is a very common probe molecule, and a solvent strength scale based on the wavelength at maximum absorbance of the dye molecule in various solvents has been developed. Phenol BluesPure ILs. The λmax values for a series of organic compounds are given in Table 2 with literature values in parentheses. A comparison of our values for organic compounds with literature values indicates a reasonable agreement. As shown in Table 2, phenol blue was used to understand the solvent strength of three different ILs. Comparison with the organic solvents indicates the overall solvent strength of ILs is similar to that of polar organic compounds. Replacing the protic hydrogen at the C2 position in the imidazolium ring with a CH3 group decreased the λmax value by 4 nm. As the replacement of the protic hydrogen with a CH3 group decreases the number of available hydrogen bonding sites, the decrease in the overall solvent strength of [hmmim][Tf2N] compared to [hmim][Tf2N] IL is expected. These results are consistent with those found using Reichardt’s dye 30 and 33. Phenol BluesIL/Organic Mixtures. We further examined the response of phenol blue to the addition of ILs to organic compounds, and these results are shown in Figure 6. Addition of small amounts of IL to organic induced a dramatic shift in the λmax value toward that of the IL, indicating that the probe is preferentially solvated by the IL. A more pronounced shift was observed when DCM was used as the organic solvent compared to ACN (see [bmim][PF6]/DCM and [bmim][PF6]/ACN systems). ACN is a slightly more polar solvent compared to DCM as indicated by the ET(30) values in Table 2. Therefore, the
J. Phys. Chem. B, Vol. 111, No. 1, 2007 137
Figure 6. Behavior of phenol blue in binary mixtures of organics or IL/organics at 23 °C.
Figure 7. Behavior of phenol blue in organic binary mixtures containing CCl4 at 23 °C.
ability of the IL to solvate the probe in ACN is lowered in comparison to its ability in solutions with DCM. On the other hand, the ability of the IL to solvate phenol blue was found to be independent of the nature of the IL in a given organic solvent. Even the replacement of the protic hydrogen with a CH3 group had little effect on the ability of [hmmim][Tf2N] to solvate phenol blue (compare [hmim][Tf2N]/ACN and [hmmim][Tf2N]/ ACN). One would have expected that the removal of protic hydrogen would have some effect on the ability of the IL to solvate the probe as phenol blue is sensitive to the hydrogen bonding ability of the solvent. For example, as shown in Figure 7, the addition of small amounts of MeOH and TFE to carbon tetrachloride, a nonpolar solvent, resulted in the domination of interactions with phenol blue by alcohols. Therefore, the results indicate that the interactions between the IL and phenol blue include interactions other than hydrogen bonding and that ACN plays a significant role in these interactions. ACN being an aprotic solvent, one would expect no specific interactions with phenol blue. Therefore, the addition of either MeOH or TFE, both good hydrogen bond donating solvents, to ACN should lead to preferential solvation of phenol blue by the alcohols. As shown in Figure 6, a linear or close to linear change in the normalized shift was observed upon the addition of alcohol (MeOH or TFE) to ACN. This is in contrast to what we observed when MeOH was added to ACN in the presence of Reichardt’s dye molecule, where the dye was strongly solvated by the alcohol (Figure 6). Even though ACN is an aprotic solvent, it
138 J. Phys. Chem. B, Vol. 111, No. 1, 2007 could interact with potential suitors through electron-pair acceptor-donor (EPA-EPD) complex formation. Therefore, hydrogen bonding interactions between phenol blue and alcohol are being balanced by ACN interactions with phenol blue via EPA-EPD complex formation, leading to the observed linear dependence on composition. The ability of ACN to solvate phenol blue is illustrated in Figure 7, where addition of small amounts of ACN to carbon tetrachloride had a similar effect to that of adding small amounts of alcohol (either MeOH or TFE). In summary, when Reichardt’s dye was used as the probe molecule, both IL and MeOH preferentially solvated the dye molecule in ACN. On the other hand, when phenol blue was used as the probe, addition of IL and MeOH to ACN had different effects on the solvatochromic behavior of phenol blue. Even though ACN was able to balance the ability of alcohols to solvate phenol blue, it was unable to compete with the ability of ILs to solvate phenol blue. Because of their ionic nature, ILs can interact with solutes through electrostatic interactions, and these types of interactions with phenol blue might explain the observed behavior. The normalized shift of phenol blue in IL/organic mixtures gives results similar to Reichardt’s dyes 30 and 33, as well as the Kamlet-Taft polarity parameters, namely, that the IL tends to preferentially solvate the probe regardless of the organic chosen. The general agreement between probes indicates that ILs will preferentially solvate polar molecules. However, these results highlight the unique property of ILs to interact with solutes or cosolvents through a wide array of interactions. Thus, it emphasizes the need to consult multiple spectroscopic probes in order to understand the solvent strength of IL/organic mixtures. A single probe, like phenol blue, that responds to several types of interactions, is likely to provide a limited or even misleading picture of the interactions in IL solution. Conclusions We report the results for a systematic study of the effects of cations, anions, and cosolvents on the solvent strength of IL/ organic mixtures using different spectroscopic probes. The general solvent strength, as indicated by ET(30), showed large changes for small amounts of ILs in organic, indicating that ILs preferentially solvate polar molecules. The strength of interaction depends on the choice of organic and IL. The general polarities, as indicated by π* and phenol blue, show similar results. The hydrogen bonding is shown by the Kamlet-Taft R and β parameters, as well as phenol blue. The results presented here indicate the important competition between cation/anion and organic/IL interactions. In general, the IL tends to preferentially solvate the probes, which suggests that they are also preferentially solvating the organic. This is supported by the comparison between C2-substituted cations, which shows that decreasing hydrogen bonding between the cation and the organic will actually result in a less preferential solvation of the probe by the IL. Acknowledgment. Acknowledgment is made to the donors of the Petroleum Research Fund, administered by the American
Mellein et al. Chemical Society, for partial support of this research. The work described herein was also supported by the State of Indiana 21st Century Research and Technology Fund (no. 909010455) and theDepartmentofEducationGAANNprogram(no.P200A01044801). We would like to thank Dr. Mark J. Muldoon and Dr. JaNeille K. Dixon for IL synthesis and Lindsay J. Miller for pyridinium IL measurements. References and Notes (1) Welton, T. Coord. Chem. ReV. 2004, 248, 2459-2477. (2) Wasserscheid, P.; Keim, W. Angew. Chem., Int. Ed. 2000, 39, 3772-3789. (3) Holbrey, J. D.; Seddon, K. R. Clean Products and Processes 1999, 1, 223-236. (4) Huddleston, J. G.; Willauer, H. D.; Swatloski, R. P.; Visser, A. E.; Rogers, R. D. Chem. Commun. 1998, 1765-1766. (5) Reichardt, C. SolVents and SolVent Effects in Organic Chemistry, 3rd ed.; Wiley-VCH: Weinheim, 2003. (6) Reichardt, C. Chem. ReV. 1994, 94, 2319-2358. (7) Wasserscheid, P.; Gordon, C. M.; Hilgers, C.; Muldoon, M. J.; Dunkin, I. R. Chem. Commun. 2001, 13, 1186-1187. (8) Muldoon, M. J.; Gordon, C. M.; Dunkin, I. R. J. Chem. Soc., Perkin Trans. 2 2001, 4, 433-435. (9) Baker, S. N.; Baker, G. A.; Bright, F. V. Green Chem. 2002, 4, 165-169. (10) Crowhurst, L.; Mawdsley, P. R.; Perez-Arlandis, J. M.; Salter, P. A.; Welton, T. Phys. Chem. Chem. Phys. 2003, 5, 2790-2794. (11) Fletcher, K. A.; Pandey, S. Appl. Spectrosc. 2002, 56, 266-271. (12) Reichardt, C. Green Chem. 2005, 7, 339-351. (13) Dzyuba, S. V.; Bartsch, R. A. Tetrahedron Lett. 2002, 43, 46574659. (14) Aki, S. N. V. K.; Brennecke, J. F.; Samanta, A. Chem. Commun. 2001, 5, 413-414. (15) Carmichael, A. J.; Seddon, K. R. J. Phys. Org. Chem. 2000, 13, 591-595. (16) Fletcher, K. A.; Storey, I. A.; Hendricks, A. E.; Pandey, S. Green Chem. 2001, 3, 210-215. (17) Bonhoˆte, P.; Dias, A.-P.; Papageorgiou, N.; Kalyanasundaram, K.; Gra¨tzel, M. Inorg. Chem. 1996, 35, 1168-1178. (18) Anderson, J. L.; Ding, J.; Welton, T.; Armstrong, D. W. J. Am. Chem. Soc. 2002, 124, 14247-14254. (19) Koel, M. SolVatochromic Studies of Ionic Liquid Solutions in Organic SolVents. Presented at the EUCHEM Conference on Molten Salts and Ionic Liquids, Hammamet, Tunisia, 2005. (20) Fletcher, K. A.; Pandey, S. Appl. Spectrosc. 2002, 56, 1498-1503. (21) Fletcher, K. A.; Baker, S. N.; Baker, G. A.; Pandey, S. New J. Chem. 2003, 27, 1706-1712. (22) Harifi-Mood, A. R.; Habibi-Yangjeh, A.; Gholami, M. R. J. Phys. Chem. B 2006, 100, 7073-7078. (23) Aki, S. N. V. K.; Mellein, B. R.; Saurer, E. M.; Brennecke, J. F. J. Phys. Chem. B 2004, 108, 20355-20365. (24) Fredlake, C. P.; Crosthwaite, J. M.; Hert, D. G.; Aki, S. N. V. K.; Brennecke, J. F. J. Chem. Eng. Data 2004, 49, 954-964. (25) Crosthwaite, J. M.; Muldoon, M. J.; Dixon, J. K.; Anderson, J. L.; Brennecke, J. F. J. Chem. Thermodyn. 2005, 37, 559-568. (26) Kamlet, M. J.; Abboud, J. L.; Taft, R. W. J. Am. Chem. Soc. 1977, 99, 6027-6038. (27) Kamlet, M. J.; Taft, R. W. J. Am. Chem. Soc. 1976, 98, 377-383. (28) Fredlake, C. P.; Muldoon, M. J.; Aki, S. N. V. K.; Welton, T.; Brennecke, J. F. Phys. Chem. Chem. Phys. 2004, 6, 3280-3285. (29) Marcus, Y. Chem. Soc. ReV. 1993, 22, 409-416. (30) Crosthwaite, J. M.; Aki, S. N. V. K.; Maginn, E. J.; Brennecke, J. F. J. Phys. Chem. B 2004, 108, 5113-5119. (31) Crosthwaite, J. M.; Aki, S. N. V. K.; Maginn, E. J.; Brennecke, J. F. Fluid Phase Equilib. 2005, 228-229, 303-309. (32) Figueras, J. J. Am. Chem. Soc. 1971, 93, 3255-3236. (33) Kamlet, M. J.; Abboud, J.-L. M.; Abraham, M. H.; Taft, R. W. J. Org. Chem. 1983, 48, 2877-2887.
