Improving Carbon Dioxide Solubility in Ionic Liquids - The Journal of

Jul 4, 2007 - Results show that ionic liquids containing increased fluoroalkyl chains on either the cation or anion do improve CO2 solubility when com...
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J. Phys. Chem. B 2007, 111, 9001-9009

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Improving Carbon Dioxide Solubility in Ionic Liquids Mark J. Muldoon,† Sudhir N. V. K. Aki,‡ Jessica L. Anderson, JaNeille K. Dixon, and Joan F. Brennecke* Department of Chemical and Biomolecular Engineering, UniVersity of Notre Dame, Notre Dame, Indiana 46556 ReceiVed: March 8, 2007; In Final Form: May 4, 2007

Previously we showed that CO2 could be used to extract organic molecules from ionic liquids without contamination of the ionic liquid. Consequently a number of other groups demonstrated that ionic liquid/CO2 biphasic systems could be used for homogeneously catalyzed reactions. Large differences in the solubility of various gases in ionic liquids present the possibility of using them for gas separations. More recently we and others have shown that the presence of CO2 increases the solubility of other gases that are poorly soluble in the ionic liquid phase. Therefore, a knowledge and understanding of the phase behavior of these ionic liquid/ CO2 systems is important. With the aim of finding ionic liquids that improve CO2 solubility and gaining more information to help us understand how to design CO2-philic ionic liquids, we present the low- and high-pressure measurements of CO2 solubility in a range of ionic liquids possessing structures likely to increase the solubility of CO2. We examined the CO2 solubility in a number of ionic liquids with systematic increases in fluorination. We also studied nonfluorinated ionic liquids that have structural features known to improve CO2 solubility in other compounds such as polymers, for example, carbonyl groups and long alkyl chains with branching or ether linkages. Results show that ionic liquids containing increased fluoroalkyl chains on either the cation or anion do improve CO2 solubility when compared to less fluorinated ionic liquids previously studied. It was also found that it was possible to obtain similar, high levels of CO2 solubility in nonfluorous ionic liquids. In agreement with our previous results, we found that the anion frequently plays a key role in determining CO2 solubility in ionic liquids.

Introduction The phase behavior of CO2 with ionic liquids (ILs) is important for the development of several potential IL applications. Since we first showed that CO2 was soluble in ILs and could be used to extract organic solutes,1-3 IL/CO2 systems have been demonstrated for various catalytic reactions.4-17 In some cases, CO2 is used to extract the product from the IL, leaving behind the catalyst and the IL for reuse. Others have adopted IL/CO2 continuous flow systems, where CO2 is used to deliver the substrates to the IL/catalyst phase and extract the products. Recently we also confirmed that CO2 could increase the solubility of gases normally not very soluble on their own in ILs.18 We found that the solubility of CH4 and O2 increased in ILs even at low partial pressures of CO2. Others studying the enantioselective hydrogenation of imines using a cationic iridium catalyst in an IL/CO2 biphasic system found that the catalyst performance was increased dramatically in the IL when CO2 pressure was added.19 High-pressure NMR showed that the addition of CO2 increased the amount of H2 dissolved in the IL. The authors believed that the catalyst, which was known to be sensitive to H2 availability, could operate more effectively due to an increase in H2 concentration in the IL along with a decrease in viscosity of the IL with added CO2. Another * Corresponding author. Telephone: (574) 631-5847. Fax: (574) 6318366. E-mail: [email protected]. † Current address: EaStCHEM, School of Chemistry, University of St Andrews, North Haugh, St Andrews, Fife KY16 9ST, U.K. E-mail: [email protected]. ‡ Current address: Intermediates R&D, Invista S.a.r.l, Sabine River Laboratory B568, P. O. Box 1003, Orange, TX 77631-1003.

potential application of ILs that several research groups along with our group have examined is the separation of CO2 from gas mixtures.20-22 To design and optimize processes using ILs, it is essential that we understand the important factors that determine CO2 solubility. It is possible to design many potential ILs; therefore, the systematic study of different cation/anion combinations will allow us to decipher the key properties in making CO2-philic ILs. To date we have studied the solubility of CO2 in a range of commonly used ILs.2,23-29 We found that there were a number of factors that controlled the CO2 solubility in the chosen ILs, both cation and anion related.26 Of the ionic liquids studied, the anion played the biggest role in CO2 solubility, a fact that was supported by a recent X-ray diffraction study by Kanakubo et al.30 Anions that contain fluoroalkyl groups were found to have some of the highest CO2 solubilities, and as the quantity of fluoroalkyl groups increased, the CO2 solubility also increased. The previous studies showed CO2 solubility for 1-butyl3-methylimidazolium ([bmim]+) based ILs at 60 °C increased in the order nitrate ([NO3]-) < tetrafluoroborate ([BF4]-) < dicyanamide ([DCA]-) ∼ hexafluorophosphate ([PF6]-) ∼ trifluoromethanesulfonate ([TfO]-) < bis[(trifluoromethyl)sulfonyl]imide ([Tf2N]-) < tris(trifluoromethylsulfonyl)methide ([methide]-). For the cations, there were two factors that influenced the CO2 solubility. The biggest effect was seen in increasing alkyl chain length on the cation. For a given cation, the CO2 solubility increased with increasing chain length, and these results are consistent with those reported elsewhere.31-35 It was thought that this may be due to entropic reasons, as the

10.1021/jp071897q CCC: $37.00 © 2007 American Chemical Society Published on Web 07/04/2007

9002 J. Phys. Chem. B, Vol. 111, No. 30, 2007 density of ILs decreases with increasing alkyl chain length; therefore, there may be more free volume within the longer chain ILs. For imidazolium cation based ILs, the effect of substitution at the C2 position was also examined. It is known that the hydrogen in this position is acidic and can hydrogen bond. Therefore, the CO2 solubility in [hmim][Tf2N] was compared to its solubility in [hmmim][Tf2N]. The solubility differed only slightly at high pressures, with [hmim][Tf2N] having the marginally higher CO2 solubility. When there are large amounts of CO2 present in the liquid, some molecules may adopt a secondary position near the cation, with the primary position for CO2 being near the anion. Simulations indicate this may be due to the increased charge on the nitrogen at the three position in [hmim] compared to [hmmim], providing a site with which CO2 can interact.27 Therefore, the secondary interactions with the cation are apparent only at high pressures. A number of other groups have also examined CO2 solubility in ILs,20,21,31,33-56 including ILs that have amine functionality attached in order to chemically bind to CO2.44,46,51 In this paper we present our efforts in understanding the interactions between ILs and CO2 and in development of CO2philic ILs. While we have significant interest in ILs that chemically complex with CO2, here we concentrate on developing ILs with high CO2 capacity through physical absorption. Experimental Section The experimental techniques to measure the solubility of CO2 in ILs and synthesis of new ILs are described. The methods used to determine the concentration of impurities are also discussed. CO2 Solubility Measurement. Two different experimental techniques were used in the current study. The solubility of CO2 at low pressures (P e 13 bar) and at high pressures (13 < P < 150 bar) was measured using a gravimetric microbalance and a static high-pressure apparatus, respectively. The low-pressure measurements were carried out in an Intelligent Gravimetric Analyzer (Hiden Analytical Limited, England), and the solubility was determined by measuring the equilibrium mass uptake by the liquid. Further details of the apparatus are given elsewhere.57,58 In a typical experiment, a known mass of the IL (∼75 mg) was loaded in a quartz sample cell and the sample chamber was evacuated to 10-9 bar at 5060 °C (depending on the IL) to remove dissolved gases. The solubility was determined at a given pressure and temperature by measuring the mass uptake (corrected for buoyancy effects but not volume expansivity) at that pressure. The densities of the pure ionic liquids as a function of temperature were measured in our laboratory, as described elsewhere.59 The solubility was not corrected for the volume expansivity due to the high solubility of CO2 in the ILs and the small expansion of the ILs at low pressure, as shown by the high-pressure data, where calculation of a 5% volume expansion had virtually no effect on the solubilities calculated. The procedure was continued at increasing pressures, up to 13 bar, to obtain an absorption isotherm. For a given isotherm, desorption experiments were also carried out to ensure that the reported measurements were equilibrium solubility values and also to confirm that the solubility behavior is reversible. Henry’s law constants were calculated from the slope of the isotherm in the limit of low solute concentration. The uncertainties in the measurement of IL density and in buoyancy calculations were used to estimate the error in Henry’s law constants, partial molar enthalpies of dissolution, and partial

Muldoon et al. molar entropies of dissolution; the details of the calculations and the error analysis are given elsewhere.28 CO2 solubility at high pressures was measured using a static high-pressure apparatus, and further details of this apparatus are given elsewhere.60,61 Briefly, the apparatus consists of a feed system with a metering pump at a constant temperature and pressure and a sample chamber. In a typical experiment, a known amount of IL (∼1.5 g) is loaded in an 8 in. sapphire sample cell and the cell is connected to the sample chamber and brought to thermal equilibrium in a constant-temperature water bath. A known amount of CO2 is metered into the sample cell and the sample is stirred until equilibrium is attained. The pressure is measured using a Heise pressure gauge. The amount of gas dissolved is obtained by calculating the difference in the amount of gas delivered to the sample chamber and the gas present in the vapor phase using the Span-Wagner equation of state.62 The solubility is measured as a function of pressure at a given temperature. Using a cathetometer, we can also obtain estimates of the molar volume of the mixture as a function of pressure with this apparatus. These values are included in the tables in the Supporting Information for completeness. Measurements to 150 bar are possible with this apparatus. The error in the solubility measurements at high pressures was obtained from propagation of error calculations, the details of which are given elsewhere.26 While it is possible to extrapolate the high-pressure data taken with this apparatus to low pressures, we have chosen not to do this. As a result, we do not report any Henry’s law constants, partial molar enthalpies of dissolution, or partial molar entropies of dissolution for ILs studied only in the high-pressure apparatus. Ionic Liquids. The ILs used in this study, along with their structures, abbreviations, and purities, are shown in Figure 1. Commercial IL samples were used as received without any further purification other than drying the sample as described below. The [C8H4F13mim][Tf2N] IL, originally synthesized at Oak Ridge National Laboratory, was generously provided by Professor Ruth Baltus. The [Et3NBH2mim][Tf2N] was provided by Professor Jim Davis. The [N4111][Tf2N] was received from Dr. Pedi Neta at the National Institute of Standards and Technology. These ILs were used without further purification, except drying. Other ILs were synthesized as described previously in the literature,26,63,64 and [b2-Nic][Tf2N] was prepared using such standard procedures, simply using butyl nicotinate as the starting amine. The purity of all ILs was determined to be greater than 99% pure by 1H NMR spectroscopy. One of the ILs synthesized, [hmim][Tf2N], is the IUPAC standard IL.65 In the synthesis of two other ILs there are some points worth noting. 1-Methyl-3-(nonafluorohexyl)imidazolium Bis[(trifluoromethyl)sulfonyl]imide ([C6H4F9mim][Tf2N]). The synthesis followed the two-step literature procedure,66 where the halide precursor was synthesized in the first step followed by the anion exchange reaction. The synthesis of the iodide salt does not go to completion. Remaining 2-(perfluorobutyl)ethyl iodide (Oakwood Chemicals) was removed by heating under reduced pressure. The iodide salt was then anion exchanged with 10 mol % excess LiTf2N (3M). The [C6H4F9mim][Tf2N] product was washed with water to remove any excess LiI, until a negative result was obtained with a silver nitrate test. It was then washed at least two additional times with water. The IL (dissolved in dry dichloromethane) was further purified and decolorized by mixing with activated carbon and passing through a column containing activated acidic and neutral alumina, to produce a clear, colorless liquid, free from any starting materials, as determined by 1H NMR.

Carbon Dioxide Solubility in Ionic Liquids

Figure 1. Part 1 of 2.

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Figure 1. Part 2 of 2. Structures, abbreviations, and purities of ionic liquids used. (a) Used as received from Merck KGaA. (b) Used as received from Solvent Innovation. (c) Sample synthesized at Oak Ridge National Laboratory and used as received. (d) Sample synthesized at the University of South Alabama and used as received. (e) Sample synthesized at the National Institute of Standards and Technology and used as received. (f) Used as received from Covalent Associates.

1-Butyl-3-methylimidazolium Pentadecafluorooctanoate ([bmim][C7F15CO2]). This IL was prepared by a standard procedure, anion exchange of 1-butyl-3-methylimidazolium bromide with ammonium pentadecafluorooctanoate (Fluka) in dichloromethane. The synthesis of 1-butyl-3-methylimidazolium bromide has been described elsewhere.67 The solid ammonium bromide precipitate was removed by filtration and further traces of ammonium bromide were removed by water washing of the dichloromethane solution, until a bromide content of the IL of less than 10 ppm was achieved. It should be noted that this compound has surfactant-like properties and water washing to remove halide impurities results in significant loss of IL (into the water phase and also through foaming). The IL is a solid at room temperature, with mp 24.7 °C when completely dry and mp 39 °C when traces of water are present. Prior to study all ILs were dried under vacuum at 70 °C for 48 h and the water content, as measured using Karl Fischer titration, was less than 1500 ppm for all ILs. The [Tf2N]-, [eFAP]-, [pFAP]-, and [bFAP]- ILs typically had lower water contents, on the order of 200 ppm. Previously, we have shown that small amounts of water do not affect CO2 solubility significantly.26 Water was determined using a EM Science Aquastar V-200 Karl Fischer titrator, bromide, chloride, and ammonium contents were determined using a Oakton Ion 510 Series pH/mV/Ion/°C meter with Cole-Parmer specific probes (27502-05 for Br-, 27502-13 for Cl-, and 27502-02 for NH4+), and silver content was determined using inductively coupled

plasma optical emission spectroscopy (Optima 3300 XL). Melting points were determined using a Mettler-Toledo differential scanning calorimeter, Model DSC822.68 A Varian UnityPlus 300 NMR was used to verify structure and verify that the ILs are free of any measurable organic starting materials or by-products. Results and Discussion Although CO2 is an ideal candidate to replace volatile organic solvents, many catalysts, polymers, and molecules have limited solubility in CO2. Therefore, a great deal of work has been carried out developing CO2-philic surfactants, polymers, and catalyst ligands to increase the solubility of polar molecules in CO2. The important factors in designing CO2-philic molecules were highlighted in a recent review.69 The advantages of using nonfluorous methods to increase the CO2-philicity, rather than the more common strategies involving fluorination, were discussed. For example, the addition of carbonyl groups to interact with CO2, along with increasing the free volume by adding ether functionality or branching the alkyl chain, are more environmentally friendly methods. In a bid to improve the CO2 solubility in ILs, we decided to study ILs that possessed CO2philic functionalities. In some cases we systematically changed the IL to allow direct comparisons to be made. Previously we found that the order of CO2 solubility in ILs with a [bmim] cation was [CF3SO3] < [(CF3SO2)2N] < [(CF3-

Carbon Dioxide Solubility in Ionic Liquids

Figure 2. Low pressure CO2 solubility at 25 °C. (9) [hmim][Tf2N]; (0) [hmpy][Tf2N]; (b) [C6H4F9mim][Tf2N]; (O) [C8H4F13mim][Tf2N]; (2) [hmim][eFAP]; (+) [hmim][pFAP]; (4) [p5mim][bFAP]; (1) [Et3NBH2mim][Tf2N].

SO2)3C].26 We believed that this may have been due to increasing fluorination of the anion. Further support of this hypothesis was provided by Pringle et al., who compared the CO2 solubility in 1-ethyl-3-methylimidazolium bis[(trifluoromethyl)sulfonyl]imide ([emim][Tf2N]) with its solubility in an IL with the nonfluorinated version of the anion, 1-ethyl-3methylimidazolium bis(methanesulfonyl)imide ([emim][Nmes2]).47 Their observations confirmed that fluorination results in higher CO2 solubility in ILs: [emim][Tf2N] had a CO2 Henry’s law constant of 47 ( 6 atm, while [emim][Nmes2] was 76 ( 8 atm at 20 °C. It may be that the good CO2 solubility in [Tf2N] and [methide] anion based ILs could be due to a combination of fluorination and the presence of SdO groups. Ab initio calculations suggest that the SdO group could be used to increase CO2-philicity of molecules due to Lewis base-Lewis acid interactions with the carbon atom of CO2.70 Nonetheless, in the design of CO2-philic molecules, fluorination is a proven method of increasing the CO2-philicity, although the exact reasons for this have been the cause of some debate in the literature.71-80 In addition, carbonyl groups in general are known CO2-philes, with the free electrons on the oxygen interacting with the Lewis acidic carbon of the CO2.69,81 The addition of acetate groups has been found to be an effective method of increasing CO2 solubility of polymers,82,83 surfactants,84,85 cyclodextrins,86 and ligands.87 Studies have shown that the oxygen atoms of CO2 can interact with the acidic hydrogens on the methyl group of the acetate, further stabilizing their interactions.70 In trifluoroacetate, the electronegative fluorine atoms reduce the Lewis bascity of the carbonyl oxygen atoms.76 Therefore, experiments were carried out to examine the effect of fluorination and the introduction of other CO2-philic groups on improving the solubility of CO2 in ILs, as described in the following sections. Fluorination of the Anion. The solubility of CO2 in [hmim][eFAP], [hmim][pFAP], and [p5mim][bFAP] was measured to pressures of 13 bar at 25 and 60 °C, and in addition, the solubility in [hmim][eFAP] and [p5mim][bFAP] was measured to pressures of 90 bar at 25, 40 and 60 °C. These data are shown in Figures 2 and 3, and listed in the Supporting Information. The positive effect of fluorinating the anion can be seen by comparing the solubility of CO2 in [hmim][PF6] (Figure 3), as measured by Shariati and Peters,37 to that in [hmim][eFAP]. The FAP-type anion is analogous to the [PF6] anion, where replacement of three fluorine atoms with fluoroethyl groups

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Figure 3. CO2 solubility in ILs at 60 °C. (`) [hmim][PF6];32 ([) [choline][Tf2N]; (9) [hmim][Tf2N];26 (]) [N4111][Tf2N]; (b) [C6H4F9mim][Tf2N]; (2) [hmim][eFAP]; (4) [p5mim][bFAP].

TABLE 1: Henry’s Law Constants for CO2 Dissolution in ILs Henry’s law constant, bar ionic liquid

10 °C

25

[bmim][PF6] [bmim][Tf2N]28 [hmim][Tf2N] [hmpy][Tf2N] [C6H4F9mim][Tf2N] [C8H4F13mim][Tf2N] [hmim][eFAP] [hmim][pFAP] [p5mim][bFAP] [hmim][ACE] [hmim][SAC] [Et3NBH2mim][Tf2N] a

24.2 ( 0.1 25.4 ( 0.1

25 °C

60 °C

53.4 ( 0.3 33.0 ( 0.3 31.6 ( 0.2 32.8 ( 0.2 28.4 ( 0.1 27.3 ( 0.2 25.2 ( 0.1 21.6 ( 0.1 20.2 ( 0.1

81.3 ( 0.8a 48.7 ( 0.9a 45.6 ( 0.3a 46.2 ( 0.3a 48.5 ( 0.4 44.7 ( 0.5b 42.0 ( 0.1 36.0 ( 0.3 32.9 ( 0.2 113.1 ( 16.9 132.2 ( 19.7

33.1 ( 1.2

At 50 °C. Estimated density from 25 °C. b

increases the CO2 solubility considerably. Also shown in both graphs is the solubility of CO2 in [hmim][Tf2N], the IUPAC standard IL,18,26,29,40 with the low-pressure data given in the Supporting Information. The new data given appear to be consistent with the data obtained by Maurer and co-workers and Shiflett and Yokozeki for the same IL.40,55 The solubility of CO2 in [p5mim][bFAP] is also shown in Figure 3, and the results indicate that there is a further increase in CO2 solubility by increasing the length of the fluoroalkyl chain on the anion. Previously, we have reported that an increase in the alkyl chain length on the cation increases the CO2 solubility.26 On the basis of these results, we could expect the solubility to be slightly greater in [hmim][bFAP] than in [p5mim][bFAP]. Comparison of the CO2 solubility in the three FAP compounds (Figure 2) shows that solubility increases with increasing fluoroalkyl chain length on the anion. From the low-pressure measurements we calculated Henry’s law constants for these ILs, and the results are shown in Table 1. Smaller values of the Henry’s law constant correspond to higher solubility. The solubility of CO2 in [p5mim][bFAP] is the highest we have observed for any IL when the dissolution is by physical absorption. We measured CO2 solubility in two ILs with fluorinated carboxylate anions ([bmim][TFA] and [bmim][C7F15CO2]). Results for these compounds are shown in Figure 4 at 60 °C, in comparison to [bmim][Tf2N] and [bmim][methide].26 These data and values for [bmim][TFA] at 25 and 40 °C are shown in the Supporting Information. The solubility of CO2 is lower in [bmim][TFA] than in [bmim][C7F15CO2]. The highly fluorinated [bmim][C7F15CO2] has a profile similar to that of [bmim]-

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Figure 4. CO2 solubility at 60 °C. (`) [bmim][TFA]; (0) [bmim][Tf2N];26 ([) [bmim][methide];26 (O) [bmim][C7F15CO2].

[methide]. As mentioned earlier, the good solubility of CO2 in [Tf2N] and [methide] could be due to fluorination, as well as the presence of SdO groups. The relatively poor solubility of CO2 in [bmim][TFA] is not unexpected since the electronegative fluorine atoms reduce the effectiveness of the carbonyl in interacting with CO2.76 Extension to the carbonyl with the fluorinated octyl chain results in better solubility, due to the effect of the fluoroalkyl chain, rather than any effect of the carboxylate functionality. Obviously, one would wish to measure the solubility in [bmim][acetate], which we have done. However, this IL chemically complexes with CO2 and a more detailed discussion will be given elsewhere.88 The results confirm that an increase in the chain length of the fluoroalkyl chain length increases the CO2 solubility. It should be noted that continually increasing the amount of fluorination is unlikely to lead to proportionate increase in CO2philicity. Others studying the CO2 solubility of fluorocarbons using computational methods have stated that there will be an optimum number of fluorine atoms for maximum CO2-philicity and then continued increases in fluorine may lead to reduced CO2-philicity.76 Fluorination of the Cation. We also examined ILs having partially fluorinated alkyl chains on the cation. [hmim][Tf2N]18,26,29,39,40 can be directly compared to [C6H4F9mim][Tf2N], and as seen in Figures 2, 3, and 5, fluorinating the last four carbons of the alkyl chain does increase the CO2 solubility. The Henry’s law constant at 25 °C (Table 1) for [C6H4F9mim][Tf2N] is 28.4 ( 0.1 bar compared to 31.6 ( 0.2 bar for [hmim][Tf2N]. However, this increase in the solubility of CO2 was less than expected, based on the results of a previous report. Earlier, Baltus et al. reported a Henry’s law constant in an analogous IL, [C8H4F13mim][Tf2N], to be 4.5 ( 1 bar at 25 °C38 and later revised this value to 6 ( 1 bar.52 Despite this IL having two more fluorinated carbons on the chain, we did not expect the Henry’s law constants to be so dramatically different from the compound we tested. We measured the CO2 solubility in [C6H4F9mim][Tf2N] at 25, 40, and 60 °C using the high-pressure (∼12-100 bar) stoichiometric phase equilibrium apparatus, and low-pressure measurements (0-13 bar) were done in the Intelligent Gravimetric Analyzer (IGA) at 25 and 60 °C. As can be seen from Figure 5, the results from the low- and highpressure measurements using the two different methods are in good agreement. We are therefore confident of our results. In order to investigate the dissimilarity between the Henry’s law constants of [C6H4F9mim][Tf2N] and [C8H4F13mim][Tf2N], we were kindly given a sample of the original [C8H4F13mim][Tf2N]

Muldoon et al.

Figure 5. Low and high pressure CO2 solubility at 25 °C. (9) [hmim][Tf2N]; (b) [C6H4F9mim][Tf2N]; (O) [C8H4F13mim][Tf2N]; (2) [hmim][eFAP]; (+) [hmim][pFAP]; (4) [p5mim][bFAP].

used by Prof. Baltus’s group and we measured CO2 solubility in this IL at 25 °C in our laboratory using the two different experimental techniques and additionally at 60 °C using the IGA. The solubility of CO2 in both [C6H4F9mim][Tf2N] and [C8H4F13mim][Tf2N] can be seen in Figures 2 and 5. The Henry’s law constant for [C8H4F13mim][Tf2N] is 27.3 ( 0.2 bar at 25 °C, and despite the viscous nature and limited sample size of the IL increasing the error in the high-pressure static method, the results from both methods are in good agreement. As might be expected, the CO2 solubility was higher in [C8H4F13mim][Tf2N] than in [C6H4F9mim][Tf2N] but not as high as previously reported. In the previous report the gas solubility was measured using a quartz crystal microbalance and it would appear that, using this method, the Henry’s law constant for CO2 dissolution in [C8H4F13mim][Tf2N] was substantially underestimated. The experimental data for the high- and low-pressure measurements are given in the Supporting Information. Comparison of the Cation. Figure 2 compares CO2 solubility in imidazolium and pyridinium Tf2N compounds. As we reported previously,29 the low-pressure solubility of CO2 in [hmim][Tf2N] and that in [hmpy][Tf2N] are very similar, with Henry’s law constants of 31.6 ( 0.2 and 32.8 ( 0.2 bar at 25 °C, respectively (see Table 1). Therefore, it can be concluded that at low pressures there is very little difference between imidazolium and pyridinium cations. We have also considered tetraalkylammonium cations, and the solubility of CO2 in [choline][Tf2N] and [N4111][Tf2N] at 60 °C is shown in Figure 3. The [choline] cation lowered CO2 solubility compared to the [hmim] cation, whereas the solubility of CO2 in [N4111][Tf2N] is similar to that in [hmim][Tf2N]. Hydrogen bonding of the [Tf2N] anion with the [choline] cation may make the anion less available for interaction with CO2. In these ILs the primary interaction of the CO2 appears to be with the anion, as we have observed previously.26 Clearly, one can design ILs, especially those developed to chemically complex with the CO2,51 where the CO2 solubility is directly affected by the cation. Enhancing CO2 Solubility without Additional Fluorination. In the field of designing CO2-philic polymers, surfactants, and ligands, researchers have explored the use of nonfluorinated molecules due to the cost and potential environmental implications of their fluorinated counterparts.69,89 Such environmental concerns should also be a factor for the design and use of ILs;

Carbon Dioxide Solubility in Ionic Liquids

Figure 6. Comparison of CO2 solubility of nonfluorinated and fluorinated ILs at 60 °C. (0) Ecoeng 41M; (3) Ecoeng 500; (9) [hmim][Tf2N];26 (`) [b2-Nic][Tf2N]; (+) [N4444][docusate]; (2) [hmim][eFAP].

therefore, we examined CO2 solubility in ILs with nonfluorous CO2-philic groups. In Figure 6 we can see the CO2 solubility in a number of ILs that were examined in the static high-pressure apparatus at 60 °C. The experimental data at all the temperatures investigated can be found in the Supporting Information. These ILs possess a number of features that we thought may improve CO2 solubility in ILs. As mentioned previously, carbonyl groups are known to be very good CO2-philes, having a low level of self-interaction but the ability to take part in Lewis acid-Lewis base interactions.69,90 As mentioned, the change from the imidazolium cation to a similar pyridinium analogue resulted in almost identical solubility of CO2. Therefore, we examined CO2 solubility in [b2-Nic][Tf2N] with the assumption that the butyl ester group might dramatically improve the solubility of CO2 in pyridinium based ILs. It was found that the solubility of CO2 at 60 °C is similar to that in [hmim][Tf2N] (Figure 6). It appears that the ester group has little effect and that the [Tf2N] anion may be the dominant factor. The ester functionality may also have less influence than expected since it is directly attached to the pyridinium cation, which will change the electronegativity of the oxygen. At higher pressures [b2-Nic][Tf2N] has slightly higher CO2 solubility compared to [hmim][Tf2N]. The increased capacity for CO2 in [b2-Nic][Tf2N] compared to [hmim]Tf2N] at high pressure could be due to the secondary interactions between CO2 and the cation. This behavior is similar to what we have reported earlier for [hmim][Tf2N] and [hmmim][Tf2N] systems.27 Another method of improving the CO2-philicity in polymers and surfactants is to insert ether functionality into the alkyl chain. The addition of ether groups is believed to improve CO2philicity by increasing the flexibility of alkyl chains, leading to increased free volume.69,91 The ether oxygen has also been shown to interact with the carbon of CO2.92 In order to understand the influence of ether functionality on the solubility behavior of CO2 in ILs, we used two commercially available ILs possessing ether groups, Ecoeng 500 and Ecoeng 41M. The solubility of CO2 in Ecoeng 500 and that in Ecoeng 41M at 60 °C are shown in Figure 6. The solubility of CO2 in Ecoeng 500 is similar to that in [hmim][Tf2N] at all pressures. It is worth noting that the higher viscosity of Ecoeng 500 (viscosity at 60 °C ) 300 cP)68 may result in large measurement uncertainties at low pressures, due to inadequate mixing. On the other hand, as the CO2 dissolves in the liquid, it lowers the viscosity of the

J. Phys. Chem. B, Vol. 111, No. 30, 2007 9007 sample, providing better mixing, faster equilibration times, and less uncertainty at higher pressures. The ammonium cation of this IL has a large alkyl chain and two chains with ether functionality. In comparison, the solubility of CO2 in Ecoeng 41M, an IL possessing an ether functionalized alkyl sulfate anion paired with the [bmim] cation, is not as good. The good performance of Ecoeng 500 is most likely due to large free volume the IL possesses, with the long alkyl chains. Ether functionality is also beneficial for ILs as the resultant flexibility allows preparation of ILs that are liquid at room temperature, where in some cases their alkyl analogues would be solids. A further benefit of using an IL such as Ecoeng 500 is that it possesses some ions of known toxicity,93 and that are not likely to be as environmentally persistent as fluorinated ILs, while maintaining good capacity for CO2. Additionally, we measured the solubility of CO2 in another nonfluorous IL, [N4444][doc], and the results are shown in Figure 6 and in the Supporting Information. Once again due to the IL’s high viscosity (12 100 cP at 25 °C),68 the measurements at low pressures have significant uncertainty, even at 60 °C. [N4444][doc] also contains an anion of known low toxicity94 and has a good affinity for CO2. As seen in Figure 6, at the highest pressures measured the solubility of CO2 was nearly identical to that in [hmim][eFAP], one of the best performing fluorinated ILs. Although this nonfluorinated IL has a very large viscosity (12 100 cP at 25 °C), the docusate anion has several features that lead to good solubility of CO2: carbonyl functionality and long, branched, alkyl chains. Thus, it is possible to design nonfluorinated ILs with good capacity for CO2, although not significantly higher than [hmim][Tf2N] except for [N4444][doc] at the highest pressures. The main drawback for the ILs investigated here is relatively high viscosity. Finally, the solubility of CO2 was measured with the lowpressure apparatus in three other ILs: [hmim][SAC],64 [hmim][ACE],64 and [Et3NBH2mim][Tf2N].95 The [SAC] and [ACE] anions contain sulfonyl groups, like [Tf2N]. They also contain carbonyl groups adjacent to the nitrogen that may be particularly nucleophilic, providing opportunity to interact with the carbon in CO2. However, they do not contain fluoroalkyl groups. From the Henry’s law constants in Table 1, it is clear that the sulfonyl and carbonyl functionality is not sufficient to enhance CO2 solubility, leading one to conclude that it is the fluoroalkyl groups in [Tf2N] that play a key role in dissolution of CO2. The [Et3NBH2mim][Tf2N] was tested to see if it may act as a hydride, reacting with the CO2 to produce formate. Clearly, this did not occur since the CO2 solubility in this compound is virtually the same as in [hmim][Tf2N], indicating that the strongest interactions of CO2 with [Et3NBH2mim][Tf2N] are with the [Tf2N] anion. Additionally, this was supported by 1H NMR, where no formate was observed. The data for all three of these compounds are found in the Supporting Information. Enthalpy and Entropy of Gas Dissolution. The partial molar enthalpy and entropy of dissolution of CO2 in the ILs can be estimated from the temperature dependence of the Henry’s law constants, as described previously.28 A larger negative value for the enthalpy indicates stronger IL/CO2 interactions. The partial molar enthalpies of dissolution of CO2 in [hmim][Tf2N] and [hmpy][Tf2N] are -12.1 and -11.4 kJ/mol, respectively. Although only investigated at two temperatures in the lowpressure apparatus, estimates for [hmim][eFAP], [hmim][pFAP], [p5mim][bFAP], [C6H4F9mim][Tf2N], and [C8H4F13mim][Tf2N] all fall in this range, indicating that CO2 dissolves in all of the ILs investigated here by simple physical absorption, consistent with our previous reports.25,28 The numbers for the various ILs

9008 J. Phys. Chem. B, Vol. 111, No. 30, 2007 are sufficiently similar within experimental error that no conclusions can be made about the relative strength of the physical interactions of CO2 with the different ILs. Conclusions We have examined the CO2 solubility in a range of ILs, allowing us to determine the important factors in designing ILs with a high CO2 solubility. As might be expected, ILs that contain a level of fluorination have improved CO2 solubility. However, fluorinated ILs may be less environmentally benign than some of the possible nonfluorinated ILs. The high stability and low reactivity of the fluorinated compounds gives them many excellent properties, but these also lead to them being poorly biodegradable and persistent in the environment. For example, [bmim][C7F15CO2] contains an anion that has been widely used in the industrial processing of fluorinated products; however, problems due to its toxicology and environmental persistence have been highlighted recently.96-98 Nonfluorinated ILs containing ether linkages and flexible alkyl chains to increase free volume can be designed to have a high affinity for CO2. The use of such ILs, however, will be dependent on whether their chemical stability and viscosity are suitable for the given application. Acknowledgment. We acknowledge financial support for this project from the State of Indiana 21st Century Research and Technology Fund (909010455) and the U.S. Department of Energy National Energy Technology Laboratory under Award No. DE-FG26-04NT42122. We thank Dr. Jacob Crosthwaite for the synthesis of [bmim][TFA] and [choline][Tf2N], and Professor Jim Davis, Solvent Innovation, Sachem, Merck KGaA, Oak Ridge National Laboratories, and the National Institute for Standards and Technology for supplying IL samples. Supporting Information Available: Solubility data for all of the compounds investigated in this study are provided in tabular form. We have also included color versions of Figures 2-6. This material is available free of charge via the Internet at http://pubs.acs.org. References and Notes (1) Blanchard, L. A.; Brennecke, J. F. Ind. Eng. Chem. Res. 2001, 40, 2550. (2) Blanchard, L. A.; Gu, Z. Y.; Brennecke, J. F. J. Phys. Chem. B 2001, 105, 2437-2444. (3) Blanchard, L. A.; Hancu, D.; Beckman, E. J.; Brennecke, J. F. Nature 1999, 399, 28-29. (4) Bosmann, A.; Francio, G.; Janssen, E.; Solinas, M.; Leitner, W.; Wasserscheid, P. Angew. Chem., Int. Ed. 2001, 40, 2697-2699. (5) Cole-Hamilton, D. J. Science 2003, 299, 1702-1706. (6) Gao, L. A.; Tao, J. A.; Zhao, G. Y.; Mu, T. C.; Wu, W. Z.; Hou, Z. S.; Han, B. X. J. Supercrit. Fluids 2004, 29, 107-111. (7) Hou, Z. S.; Han, B. X.; Gao, L.; Jiang, T.; Liu, Z. M.; Chang, Y. H.; Zhang, X. G.; He, J. New J. Chem. 2002, 26, 1246-1248. (8) Jessop, P. G. J. Synth. Org. Chem. Jpn. 2003, 61, 484-488. (9) Jessop, P. G.; Stanley, R. R.; Brown, R. A.; Eckert, C. A.; Liotta, C. L.; Ngo, T. T.; Pollet, P. Green Chem. 2003, 5, 123-128. (10) Laszlo, J. A.; Compton, D. L. Chymotrypsin-Catalyzed Transesterification in Ionic Liquids and Ionic Liquid/Supercritical Carbon Dioxide. In Ionic Liquids; Rogers, R. D., Seddon, K. R., Eds.; ACS Symposium Series 818; American Chemical Society: Washington, DC, 2002; pp 387398. (11) Lozano, P.; De Diego, T.; Carrie, D.; Vaultier, M.; Iborra, J. L. Enzymatic Catalysis in Ionic Liquids and Supercritical Carbon Dioxide. In Ionic Liquids as Green SolVents: Progress and Prospects; Rogers, R. D., Seddon, K. R., Eds.; ACS Symposium Series 856; American Chemical Society: Washington, DC, 2003; pp 239-250. (12) Lozano, P.; De Diego, T.; Carrie, D.; Vaultier, M.; Iborra, J. L. J. Mol. Catal. A: Chem. 2004, 214, 113-119. (13) Lu, X. B.; Zhang, Y. J.; Liang, B.; Li, X.; Wang, H. J. Mol. Catal. A: Chem. 2004, 210, 31-34.

Muldoon et al. (14) Lozano, P.; De Diego, T.; Carrie, D.; Vaultier, M.; Iborra, J. L. Biotechnol. Prog. 2003, 19, 380-382. (15) Reetz, M. T.; Wiesenhofer, W.; Francio, G.; Leitner, W. AdV. Synth. Catal. 2003, 345, 1221-1228. (16) Tumas, W.; Liu, F. C.; Baker, R. T. Abstr. Pap. Am. Chem. Soc. 2001, 221st, U615-U615. (17) Webb, P. B.; Sellin, M. F.; Kunene, T. E.; Williamson, S.; Slawin, A. M. Z.; Cole-Hamilton, D. J. J. Am. Chem. Soc. 2003, 125, 1557715588. (18) Hert, D. G.; Anderson, J. L.; Aki, S. N. V. K.; Brennecke, J. F. Chem. Commun. 2005, 2603-2605. (19) Solinas, M.; Pfaltz, A.; Cozzi, P. G.; Leitner, W. J. Am. Chem. Soc. 2004, 126, 16142-16147. (20) Scovazzo, P.; Visser, A. E.; Davis, J. H.; Rogers, R. D.; Koval, C. A.; DuBois, D. L.; Noble, R. D. Supported Ionic Liquid Membranes and Facilitated Ionic Liquid Membranes. In Ionic Liquids; Rogers, R. D., Seddon, K. R., Eds.; ACS Symposium Series 818; American Chemical Society: Washington, DC, 2002; pp 69-87. (21) Scovazzo, P.; Kieft, J.; Finan, D. A.; Koval, C.; DuBois, D.; Noble, R. J. Membr. Sci. 2004, 238, 57-63. (22) Anthony, J. L.; Aki, S. N. V. K.; Maginn, E. J.; Brennecke, J. F. Int. J. EnViron. Technol. Manage. 2004, 4, 105-115. (23) Anthony, J. L.; Crosthwaite, J. M.; Hert, D. G.; Aki, S. N. V. K.; Maginn, E. J.; Brennecke, J. F. Phase Equilibria of Gases and Liquids with 1-N-Butyl-3-Methylimidazolium Tetrafluoroborate. In Ionic Liquids as Green SolVents: Progress and Prospects; Rogers, R. D., Seddon, K. R., Eds.; ACS Symposium Series 856; American Chemical Society: Washington, DC, 2003; pp 110-120. (24) Anthony, J. L.; Maginn, E. J.; Brennecke, J. F. Gas Solubilities in 1-N-Butyl-3-Methylimidazolium Hexafluorophosphate. In Ionic Liquids; Rogers, R. D., Seddon, K. R., Eds.; ACS Symposium Series 818; American Chemical Society: Washington, DC, 2002; pp 260-269. (25) Anthony, J. L.; Maginn, E. J.; Brennecke, J. F. J. Phys. Chem. B 2002, 106, 7315-7320. (26) Aki, S. N. V. K.; Mellein, B. R.; Saurer, E. M.; Brennecke, J. F. J. Phys. Chem. B 2004, 108, 20355-20365. (27) Cadena, C.; Anthony, J. L.; Shah, J. K.; Morrow, T. I.; Brennecke, J. F.; Maginn, E. J. J. Am. Chem. Soc. 2004, 126, 5300-5308. (28) Anthony, J. L.; Anderson, J. L.; Maginn, E. J.; Brennecke, J. F. J. Phys. Chem. B 2005, 109, 6366-6374. (29) Anderson, J. L.; Dixon, J. K.; Maginn, E. J.; Brennecke, J. F. J. Phys. Chem. B 2006, 110, 15059-15062. (30) Kanakubo, M.; Umecky, T.; Hiejima, Y.; Aizawa, T.; Nanjo, H.; Kameda, Y. J. Phys. Chem. B 2005, 109, 13847-13850. (31) Shariati, A.; Peters, C. J. J. Supercrit. Fluids 2004, 29, 43-48. (32) Shariati, A.; Peters, C. J. J. Supercrit. Fluids 2004, 30, 139-144. (33) Shariati, A.; Peters, C. J. J. Supercrit. Fluids 2005, 34, 171-176. (34) Shariati, A.; Gutkowski, K.; Peters, C. J. AIChE J. 2005, 51, 15321540. (35) Gutkowski, K. I.; Shariati, A.; Peters, C. J. J. Supercrit. Fluids 2006, 39, 187-191. (36) Shariati, A.; Peters, C. J. J. Supercrit. Fluids 2003, 25, 109-117. (37) Shariati, A.; Peters, C. J. J. Supercrit. Fluids 2004, 30, 139-144. (38) Baltus, R. E.; Culbertson, B. H.; Dai, S.; Luo, H. M.; DePaoli, D. W. J. Phys. Chem. B 2004, 108, 721-727. (39) Kumelan, J.; Kamps, A. P. S.; Tuma, D.; Maurer, G. J. Chem. Eng. Data 2006, 51, 1802-1807. (40) Kumelan, J.; Kamps, I. P. S.; Tuma, D.; Maurer, G. J. Chem. Thermodyn. 2006, 38, 1396-1401. (41) Lee, B. C.; Outcalt, S. L. J. Chem. Eng. Data 2006, 51, 892-897. (42) Jacquemin, J.; Gomes, M. F. C.; Husson, P.; Majer, V. J. Chem. Thermodyn. 2006, 38, 490-502. (43) Camper, D.; Scovazzo, P.; Koval, C.; Noble, R. Ind. Eng. Chem. Res. 2004, 43, 3049-3054. (44) Zhang, S. J.; Yuan, X. L.; Chen, Y. H.; Zhang, X. P. J. Chem. Eng. Data 2005, 50, 1582-1585. (45) Zhang, S. J.; Chen, Y. H.; Ren, R. X. F.; Zhang, Y. Q.; Zhang, J. M.; Zhan, X. P. J. Chem. Eng. Data 2005, 50, 230-233. (46) Zhang, J. M.; Zhang, S. J.; Dong, K.; Zhang, Y. Q.; Shen, Y. Q.; Lv, X. M. Chem.sEur. J. 2006, 12, 4021-4026. (47) Pringle, J. M.; Golding, J.; Baranyai, K.; Forsyth, C. M.; Deacon, G. B.; Scott, J. L.; MacFarlane, D. R. New J. Chem. 2003, 27, 1504-1510. (48) Kamps, A. P. S.; Tuma, D.; Xia, J. Z.; Maurer, G. J. Chem. Eng. Data 2003, 48, 746-749. (49) Liu, Z. M.; Wu, W. Z.; Han, B. X.; Dong, Z. X.; Zhao, G. Y.; Wang, J. Q.; Jiang, T.; Yang, G. Y. Chem.sEur. J. 2003, 9, 3897-3903. (50) Husson-Borg, P.; Majer, V.; Gomes, M. F. C. J. Chem. Eng. Data 2003, 48, 480-485. (51) Bates, E. D.; Mayton, R. D.; Ntai, I.; Davis, J. H. J. Am. Chem. Soc. 2002, 124, 926-927. (52) Baltus, R. E.; Counce, R. M.; Culbertson, B. H.; Luo, H. M.; DePaoli, D. W.; Dai, S.; Duckworth, D. C. Sep. Sci. Technol. 2005, 40, 525-541.

Carbon Dioxide Solubility in Ionic Liquids (53) Hutchings, J. W.; Fuller, K. L.; Heitz, M. P.; Hoffmann, M. M. Green Chem. 2005, 7, 475-478. (54) Kim, Y. S.; Choi, W. Y.; Jang, J. H.; Yoo, K. P.; Lee, C. S. Fluid Phase Equilib. 2005, 228, 439-445. (55) Shiflett, M. B.; Yokozeki, A. J. Phys. Chem. B 2007, 111, 20702074. (56) Chen, Y. H.; Zhang, S. J.; Yuan, X. L.; Zhang, Y. Q.; Zhang, X. P.; Dai, W. B.; Mori, R. Thermochim. Acta 2006, 441, 42-44. (57) Anthony, J. L.; Maginn, E. J.; Brennecke, J. F. J. Phys. Chem. B 2001, 105, 10942-10949. (58) Anthony, J. L. Gas Solubilities in Ionic Liquids: Experimental Measurements and Applications. Ph.D. Dissertation, University of Notre Dame, 2004. (59) Anderson, J. L.; Dixon, J. K.; Maginn, E. J.; Brennecke, J. F. Manuscript in preparation, 2007. (60) Stradi, B. A.; Kohn, J. P.; Stadtherr, M. A.; Brennecke, J. F. J. Supercrit. Fluids 1998, 12, 109-122. (61) Tarantino, D. E.; Kohn, J. P.; Brennecke, J. F. J. Chem. Eng. Data 1994, 39, 158-160. (62) Span, R.; Wagner, W. J. Phys. Chem. Ref. Data 1996, 25, 15091596. (63) Bonhote, P.; Dias, A. P.; Papageorgiou, N.; Kalyanasundaram, K.; Gratzel, M. Inorg. Chem. 1996, 35, 1168-1178. (64) Carter, E. B.; Culver, S. L.; Fox, P. A.; Goode, R. D.; Ntai, I.; Tickell, M. D.; Traylor, R. K.; Hoffman, N. W.; Davis, J. H. Chem. Commun. 2004, 630-631. (65) Marsh, K. Thermodynamics of Ionic Liquids, Ionic Liquid Mixtures, and the DeVelopment of Standardized Systems; http://www.iupac.org/ publications/ci/2005/2705/pp4_2002-005-1-100.html, 2005. (66) Singh, R. P.; Manandhar, S.; Shreeve, J. M. Tetrahedron Lett. 2002, 43, 9497-9499. (67) 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. (68) Crosthwaite, J. M.; Muldoon, M. J.; Dixon, J. K.; Anderson, J. L.; Brennecke, J. F. J. Chem. Thermodyn. 2005, 37, 559-568. (69) Beckman, E. J. Chem. Commun. 2004, 1885-1888. (70) Raveendran, P.; Wallen, S. L. J. Am. Chem. Soc. 2002, 124, 1259012599. (71) Yee, G. G.; Fulton, J. L.; Smith, R. D. J. Phys. Chem. 1992, 96, 6172-6181. (72) Dardin, A.; DeSimone, J. M.; Samulski, E. T. J. Phys. Chem. B 1998, 102, 1775-1780. (73) Cece, A.; Jureller, S. H.; Kerschner, J. L.; Moschner, K. F. J. Phys. Chem. 1996, 100, 7435-7439. (74) Yonker, C. R.; Palmer, B. J. J. Phys. Chem. A 2001, 105, 308314. (75) Diep, P.; Jordan, K. D.; Johnson, J. K.; Beckman, E. J. J. Phys. Chem. A 1998, 102, 2231-2236.

J. Phys. Chem. B, Vol. 111, No. 30, 2007 9009 (76) Raveendran, P.; Wallen, S. L. J. Phys. Chem. B 2003, 107, 14731477. (77) Fried, J. R.; Hu, N. Polymer 2003, 44, 4363-4372. (78) Rindfleisch, F.; DiNoia, T. P.; McHugh, M. A. J. Phys. Chem. 1996, 100, 15581-15587. (79) McHugh, M. A.; Park, I. H.; Reisinger, J. J.; Ren, Y.; Lodge, T. P.; Hillmyer, M. A. Macromolecules 2002, 35, 4653-4657. (80) Han, Y. K.; Jeong, H. Y. J. Phys. Chem. A 1997, 101, 56045604. (81) Raveendran, P.; Ikushima, Y.; Wallen, S. L. Acc. Chem. Res. 2005, 38, 478-485. (82) Kilic, S.; Michalik, S.; Wang, Y.; Johnson, J. K.; Enick, R. M.; Beckman, E. J. Ind. Eng. Chem. Res. 2003, 42, 6415-6424. (83) Shen, Z.; McHugh, M. A.; Xu, J.; Belardi, J.; Kilic, S.; Mesiano, A.; Bane, S.; Karnikas, C.; Beckman, E.; Enick, R. Polymer 2003, 44, 1491-1498. (84) Shervani, Z.; Liu, J. C.; Ikushima, Y. New J. Chem. 2004, 28, 666668. (85) Fan, X.; Potluri, V. K.; McLeod, M. C.; Wang, Y.; Liu, J. C.; Enick, R. M.; Hamilton, A. D.; Roberts, C. B.; Johnson, J. K.; Beckman, E. J. J. Am. Chem. Soc. 2005, 127, 11754-11762. (86) Potluri, V. K.; Hamilton, A. D.; Karanikas, C. F.; Bane, S. E.; Xu, J. H.; Beckman, E. J.; Enick, R. M. Fluid Phase Equilib. 2003, 211, 211217. (87) Ablan, C. D.; Sheppard, D.; Beckman, E. J.; Olmstead, M. M.; Jessop, P. G. Green Chem. 2005, 7, 590-594. (88) Dixon, J. K.; Anderson, J. L.; Muldoon, M. J.; Aki, S. N. V. K.; Maginn, E. J.; Brennecke, J. F. Manuscript in preparation, 2007. (89) Sarbu, T.; Styranec, T. J.; Beckman, E. J. Ind. Eng. Chem. Res. 2000, 39, 4678-4683. (90) Kazarian, S. G.; Vincent, M. F.; Bright, F. V.; Liotta, C. L.; Eckert, C. A. J. Am. Chem. Soc. 1996, 118, 1729-1736. (91) Sarbu, T.; Styranec, T.; Beckman, E. J. Nature 2000, 405, 165168. (92) Van Ginderen, P.; Herrebout, W. A.; van der Veken, B. J. J. Phys. Chem. A 2003, 107, 5391-5396. (93) Toxcity Data AVailable for Ecoeng 500 Ionic Liquid; http:// www.solvent-innovation.com/Englisch/ecoeng_unten.htm (accessed 2007). (94) Davis, J. H.; Fox, P. A. Chem. Commun. 2003, 1209-1212. (95) Fox, P. A.; Griffin, S. T.; Reichert, W. M.; Salter, E. A.; Smith, A. B.; Tickell, M. D.; Wicker, B. F.; Cioffi, E. A.; Davis, J. H.; Rogers, R. D.; Wierzbicki, A. Chem. Commun. 2005, 3679-3681. (96) Lau, C.; Butenhoff, J. L.; Rogers, J. M. Toxicol. Appl. Pharmacol. 2004, 198, 231-241. (97) Boulanger, B.; Vargo, J.; Schnoor, J. L.; Hornbuckle, K. C. EnViron. Sci. Technol. 2004, 38, 4064-4070. (98) Hekster, F. M.; Laane, R.; de Voogt, P. ReV. EnViron. Contam. Toxicol. 2003, 179, 99-121.