Why Is CO 2 So Soluble in Imidazolium-Based Ionic Liquids? Cesar Cadena, Jennifer L. Anthony, Jindal K. Shah, Timothy I. Morrow, Joan F. Brennecke, and Edward J. Maginn* Contribution from the Department of Chemical and Biomolecular Engineering, UniVersity of Notre Dame, Notre Dame, Indiana 46556 Received November 15, 2003; E-mail: ed@nd.edu Abstract: Experimental and molecular modeling studies are conducted to investigate the underlying mechanisms for the high solubility of CO2 in imidazolium-based ionic liquids. CO2 absorption isotherms at 10, 25, and 50 °C are reported for six different ionic liquids formed by pairing three different anions with two cations that differ only in the nature of the “acidic” site at the 2-position on the imidazolium ring. Molecular dynamics simulations of these two cations paired with hexafluorophosphate in the pure state and mixed with CO2 are also described. Both the experimental and the simulation results indicate that the anion has the greatest impact on the solubility of CO2. Experimentally, it is found that the bis(trifluoromethylsulfonyl)- imide anion has the greatest affinity for CO2, while there is little difference in CO2 solubility between ionic liquids having the tetrafluoroborate or hexafluorophosphate anion. The simulations show strong organization of CO 2 about hexafluorophosphate anions, but only small differences in CO2 structure about the different cations. This is consistent with the experimental finding that, for a given anion, there are only small differences in CO 2 solubility for the two cations. Computed and measured densities, partial molar volumes, and thermal expansion coefficients are also reported. 1. Introduction A number of investigations have shown that CO 2 is remark- ably soluble in imidazolium-based ionic liquids. 1-5 We first investigated this phenomenon in connection with the possibility of extracting solutes from ionic liquids (ILs) with supercritical CO 2 . 1,6 More detailed investigations showed that the anion and substituents on the cation did affect the CO 2 solubility some- what, but the solubilities were quite high in all of the different ILs studied. 2 For instance, at just 50 bar of CO 2 pressure, the CO 2 solubility is on the order of 50 mol %. 2 Subsequent investigations at both low and high pressures have confirmed this behavior. 3-5 As mentioned above, supercritical CO 2 can be used as an environmentally benign solvent to extract organic products or contaminants from ILs. This is a particularly attractive technique because the solubility of the IL in CO 2 is immeasurably low. 1 As a result, a number of researchers have adopted this methodology. 7-10 Obviously, understanding the phase behavior of CO 2 with ILs is important for that application. We have also shown that CO 2 can be used to separate organic liquids and water from ILs by inducing a liquid-liquid phase split. 11,12 The amount of CO 2 dissolved in the IL/water or IL/organic mixture is key to determining how much CO 2 pressure is necessary to achieve the phase split. Finally, we have set forth the idea of using ILs for the separation of CO 2 from a wide variety of other gases. 13 Once again, the solubility of CO 2 in the various ILs is vital for the design and evaluation of gas separation processes using ILs. Therefore, we believe that it is important to gain a fundamental understanding of the mechanism for the high solubility of CO 2 in imidazolium-based ILs. Kazarian and co-workers investigated mixtures of CO 2 with 1-n-butyl-3-methylimidazolium hexafluorophosphate ([bmim]- [PF 6 ]) and 1-n-butyl-3-methylimidazolium tetrafluoroborate ([bmim][BF 4 ]) using ATR-IR spectroscopy. 14 They found evidence of a weak Lewis acid-base interaction between the CO 2 and the anions of the ILs. The same group 15 used ATR-IR spectroscopy to probe IL/water solutions, which further con- firmed the importance of interaction, in this case via hydrogen- bonding, of solutes with the anion. This was true for, among (1) Blanchard, L. A.; Hancu, D.; Beckman, E. J.; Brennecke, J. F. Nature 1999, 399, 28-29. (2) Blanchard, L. A.; Gu, Z. Y.; Brennecke, J. F. J. Phys. Chem. B 2001, 105, 2437-2444. (3) Anthony, J. L.; Maginn, E. J.; Brennecke, J. F. J. Phys. Chem. B 2002, 106, 7315-7320. (4) Kamps, A. P.-S.; Tuma, D.; Xia, J.; Maurer, G. J. Chem. Eng. Data 2003, 48, 746-749. (5) Husson-Borg, P.; Majer, V.; Costa Gomes, M. F. J. Chem. Eng. Data 2003, 48, 480-485. 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SOC. 2004, 126, 5300-5308 10.1021/ja039615x CCC: $27.50 © 2004 American Chemical Society