Review on the Use of Ionic Liquids (ILs) as Alternative Fluids for CO2

Nov 2, 2010 - Alternative materials in technologies for Biogas upgrading via CO 2 capture. Kui Zhou , Somboon Chaemchuen , Francis Verpoort. Renewable...
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Energy Fuels 2010, 24, 5817–5828 Published on Web 11/02/2010

: DOI:10.1021/ef1011337

Review on the Use of Ionic Liquids (ILs) as Alternative Fluids for CO2 Capture and Natural Gas Sweetening Ferdi Karadas,† Mert Atilhan,*,† and Santiago Aparicio‡ †

Department of Chemical Engineering, Qatar University, 2713 Doha, Qatar, and ‡Department of Chemistry, University of Burgos, 09001 Burgos, Spain Received August 22, 2010. Revised Manuscript Received October 18, 2010

The capture of CO2 from flue gases derived from fossil-fueled power plants and the absorption of CO2/H2S for natural gas sweetening purposes are two relevant industrial problems closely related to very important environmental, economical, and technological problems that need to be solved. Amine-based technologies are widely used in the industry for these purposes, but they lead to several problems that have led many researchers to pose new alternatives. Ionic liquids (ILs) have emerged in the last few years as promising new acid gas absorbents, and thus, this remarkable interest, in both industry and academia, has led to a large collection of experimental and theoretical studies in which the most important aspects of the absorption process are analyzed. In this review, we show the most relevant conclusions obtained from the analysis of the literature, analyzing the state-of-the-art results, trying to infer the viability of ILs as an alternative to the available amine-based absorption processes, and showing the possible future directions of research.

problems during gas transportation and commercialization. Natural gas is frequently commercialized through gas-to-liquid (GTL) or liquefied natural gas (LNG) technologies,8,9 for which raw gas conditioning close to the natural gas reservoirs is required. The production of LNG requires the drastic reduction of acid gases, gas sweetening, mainly for the elimination of sulfur content (in the form of H2S) and CO2 (that may crystallize, leading to operational problems).10,11 Likewise, acid gases must be removed from the raw gas stream to meet pipeline sales contract specifications.12,13 Moreover, the connection of CO2 with the global warming problem and the undesirable atmospheric emission of sulfur gases show the importance of gas sweetening for natural gas commercialization. Hence, the increasingly stringent regulations require 99þ% recovery solutions for the sour/acid gases for new gas projects. Furthermore, CO2 emissions from fossil-fueled power plants is another remarkable technological and economical challenge.14,15 Global climate change effects rise from anthropogenic emissions

1. Introduction World natural gas consumption has grown significantly over the years, becoming one of the most important energy sources for the future.1 Natural gas is considered as the most environmentally friendly fossil fuel, because natural gas burning leads to negligible SO2 emissions, low nitrous oxide levels, and less than half of the CO2 emissions when burned in comparison to coal or oil.2 Nevertheless, a remarkable challenge in natural gas processing is that many raw natural gases are contaminated with undesired components, such as hydro gen sulfide and carbon dioxide, resulting in a classification of natural gases as sour (high content of H2S), acid (high content of CO2 or CO2 and H2S), and sweet (low content of H2S and CO2).3 The increase in gas demand required the use of sour gas sources as well as the sweet ones,4 even though purification of sour gases brings extra cost to gas processing because convenient methods are needed to remove CO2 and H2S efficiently from the gas source.5-7 Removal of acid gases from raw natural gases has to be performed not only because of environmental restrictions but also considering technological

(8) Fleisch, T. H.; Sills, R. A.; Briscoe, M. D. Emergence of the gas-toliquids industry: A review of global GTL developments. J. Nat. Gas Chem. 2002, 11, 1–14. (9) Thomas, S.; Dawe, R. A. Review of ways to transport natural gas energy from countries which do not need the gas for domestic use. Energy 2003, 28, 1461. (10) Osman, K.; Vasagam, M. Gas sweetening process. Problems and remedial measures. Proceedings of the Abu Dhabi International Petroleum Exhibition and Conference; Abu Dhabi, United Arab Emirates, 2002. (11) Al-Mohannadi, A. Sour gas treating for LNG production. Proceedings of the International Petroleum Conference; Doha, Qatar, 2005. (12) Gene, G. B. Processing sour natural gas to meet pipeline quality. Proceedings of the SPE Symposium on Sour Gas and Crude; Tyler, TX, 1977. (13) Hubbard, R. The role of gas processing in the natural gas value chain. J. Pet. Technol. 2009, 61, 65–71. (14) Wolsky, A. M.; Daniels, E. J.; Jody, B. J. CO2 capture from the flue gas of conventional fossil-fuel-fired power plants. Environ. Prog. 1994, 13, 214–219. (15) Figueroa, J. D.; Fout, T.; Plasynski, S.; McIlvried, H.; Srivastava, R. D. Advances in CO2 capture technology;The U.S. Department of Energy’s carbon sequestration program. Int. J. Greenhouse Gas Control 2008, 2, 9–20.

*To whom correspondence should be addressed. E-mail: mert.atilhan@ qu.edu.qa. (1) Energy Information Administration. International Energy Outlook; Energy Information Administration: Washington, D.C., 2008; http:// www.eia.doe.gov/oiaf/ieo.index.htm. (2) Energy Information Administration. Natural Gas 1998: Issues and Trends; Energy Information Administration: Washington, D.C., 1998; DOE/ EIA-0560(98). (3) Mokhatab, S.; Poe, W. A.; Speight, J. G. Natural Gas Transmission and Processing; Gulf Professional Publishing: Burlington, MA, 2006. (4) Bates, E. D.; Mayton, R. D.; Ntai, I.; Davis, J. H. CO2 capture by a task-specific ionic liquid. J. Am. Chem. Soc. 2002, 124, 926–927. (5) Stewart, C.; Hessami, M. A study of methods of carbon dioxide capture and sequestration;The sustainability of a photosynthetic bioreactor approach. Energy Convers. Manage. 2005, 46, 403–420. (6) O’Neill, B. C.; Oppenheimer, M. Climate change: Dangerous climate impacts and the Kyoto Protocol. Science 2002, 296, 1971–1972. (7) Intergovernmental Panel on Climate Change (IPCC). IPCC Special Report on Carbon Dioxide Capture and Storage (CCS); IPCC: Geneva, Switzerland, 2009; http://www.ipcc.ch. r 2010 American Chemical Society

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Figure 1. World GHG emissions indicating (a) gases, (b) CO2 emissions origin, (c) CO2 energy-related emissions, and (d) world electricity generation by fuel. Data from refs 1 and 17.

of greenhouse gases (GHGs) into the atmosphere,16 and one of the most important GHGs is carbon dioxide (Figure 1a). For this gas, the most important source is energy-related emissions (Figure 1b). Moreover, CO2 emissions from fossil fuel combustion (coal, oil, and natural gas) for energy production are also the main source of energy-related emissions (panels c and d of Figure 1),1,17 because as a result of their low cost, availability, existing reliable technology for energy production, and energy density, fossil fuels currently supply over 67% of the electricity used worldwide (Figure 1d).1 Hence, approximately 25% of the total CO2 emissions in the world are produced from combustion and nonfuel uses of fossil fuels for electricity generation.1 Moreover, world CO2 energy-related emissions are expected to increase at a rate of 2.1% per year,1 which is in agreement with the forecasted consumption of fossil fuels for electricity generation (natural gas will increase from 19.4% in 2005 to 25% in 2030, and coal will increase from 41.4% in 2005 to 46% in 2030).1 Thus, fossil fuels will remain the major source for electricity generation and supply, because renewable sources would be insufficient to supplant them in the next decades; hence, to control CO2 emissions from these sources should be extremely important to minimize the emissions of GHGs that contribute to climate change without stifling economical and technological development. Although the characteristics of CO2 capture from natural gas, for sweetening purposes, and from flue gas of conventional fossil-fuel-fired power plants (in which low CO2 partial pressures, less than 0.15 atm, are common and with CO2 concentrations of typically 3-13 vol %), are very different, it is clear that both problems are closely related, and thus, similar approaches are frequently used for both cases. Nevertheless, it should be remarked that the economics of CO2 capture are strongly influenced by the partial pressure in the feed gas, with low partial pressures making their use as physical solvents difficult.15 A number of CO2 capture technologies have already been practiced on laboratory or industrial scale, for both natural gas sweetening and flue gases, that require various processes involving physisorption/chemisorption, membrane separation or molecular sieves, carbamation, amine physical absorption,

amine dry scrubbing, mineral carbonation, etc.15,18-25 Mostly alkanolamines, such as monoethanolamine (MEA)-, diethanolamine (DEA)-, and methyldiethanolamine (MDEA)-based methods, are being used for CO2 capture through carbamate/ carbonate formation. Amine-based processes are also used for H2S capture together with CO2, for natural gas sweetening purposes,26 and thus, chemical absorption with amines is used for over 95% of all gas sweetening in the United States.27-30 There are some disadvantages in commercial use of these amine solutions,28 including the loss of amine reagents and transfer of water into the gas stream during the desorption stage, degradation of amine reagents to form corrosive byproducts,31,32 and high energy consumption during regeneration, as well as insufficient carbon dioxide/hydrogen sulfide capture capacity. The regeneration step may be 70% of the total operating costs (18) Zelenak, V.; Badanicova, M.; Halamova, D. Amine-modified ordered mesoporous silica: Effect of pore size on carbon dioxide capture. Chem. Eng. J. 2008, 144, 336–342. (19) Ebner, A. D.; Ritter, J. A. State-of-the-art adsorption and membrane separation processes for carbon dioxide production from carbon dioxide emitting industries. Sep. Sci. Technol. 2009, 44, 1273–1421. (20) Chaffee, A. L.; Knowles, G. P.; Liang, Z. CO2 adsorption by adsorption: Materials and process development. Int. J. Greenhouse Control 2007, 1, 11–18. (21) Song, C. S. Global challenges and strategies for control, conversion, and utilization of CO2 for sustainable development involving energy, catalysis, adsorption, and chemical processing. Catal. Today 2006, 115, 2–32. (22) Steeneveldt, R.; Berger, B.; Torp, T. A. CO2 capture and storage closing the knowing-doing gap. Trans. IChemE, Part A 2006, 84, 739–763. (23) Li, J. L.; Chen, B. H. Review of CO2 absorption using chemical solvents in hollow fiber membrane contactors. Sep. Purif. Technol. 2005, 41, 109–122. (24) Bailey, D. W.; Feron, P. H. M. Post-combustion decarbonation processes. Oil Gas Sci. Technol. 2005, 60, 461–474. (25) Astarita, G.; Savage, D. W.; Bisio, A. Gas Treating with Chemical Solvents; John Wiley: New York, 1983. (26) Parrish, W. R.; Kidnay, A. J. Fundamentals of Natural Gas Processing; CRC Press: Boca Raton, FL, 2006. (27) Abu-Zahra, M. R. M.; Schneiders, L. H. J.; Niederer, J. P. M.; Feron, P. H. M.; Versteeg, G. F. CO2 capture from power plants: Part I. A parametric study of the technical performance based on monoethanolamine. Int. J. Greenhouse Gas Control 2007, 1, 37–46. (28) Yang, H.; Xu, Z.; Fan, M.; Gupta, R.; Slimane, R. B.; Bland, A. E.; Wright, I. Progress in carbon dioxide separation and capture: A review. J. Environ. Sci. 2008, 20, 14–27. (29) Feron, P. H. M.; Hendricks, C. A. CO2 capture process principles and costs. Oil Gas Sci. Technol. 2005, 60, 451–459. (30) Plaza, M. S.; Pevida, C.; Arenillas, A.; Pis, J. J. CO2 capture by adsorption with nitrogen enriched carbons. Fuel 2007, 86, 2204–2212. (31) DuPart, M. S.; Bacon, T. R.; Edwards, D. J. Understanding corrosion in alkanolamine gas treating plants. Hydrocarbon Process. 1993, 72, 89–94. (32) Kittel, J.; Idem, R.; Gelowitz, D.; Tontiwachwuthikul, P.; Parrain, G.; Bonneau, A. Corrosion in MEA units for CO2 capture: Pilot plant studies. Energy Procedia 2009, 1, 791–797.

(16) Solomon, S.; Qin, D.; Manning, M.; Chen, Z.; Marquis, M.; Averty, K. B.; Tignor, M.; Miller, H. L. In Climate Change 2007: The Physical Science Basis. Contribution of Working Group I to the Fourth Assessment Report of the Intergovernmental Panel on Climate Change; Solomon, S., Qin, D., Manning, M., Eds.; Cambridge University Press: Cambridge, U.K., 2007. (17) Baumert, K. A.; Herzog, T.; Pershing, J. Navigating the Numbers. Greenhouse Gas Data and International Climate Policy; World Resources Institute: Washington, D.C., 2005; www.wri.org/publication/navigatingthe-numbers.

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of the capture process (4 GJ/ton of CO2).33-35 All of these discrepancies have to be overcome and addressed on a wider scale for more efficiency with less cost,35-37 which requires the emergence of alternative methods for CO2 and H2S capture.15,28,38 A very promising alternative for CO2/H2S capture, for both flue gases and gas sweetening purposes, is the use of ionic liquids (ILs) as absorbents.4,39-41 Because the physical and chemical properties of room-temperature ILs could be enhanced and modified by both their cationic and anionic moieties, they serve a broad range of applications, such as solvents, sensors, solid-state photocells, thermal and hydraulic fluids, lubricants, and several analytical techniques, including mass spectrometry, separation techniques, and electrochemistry.42-45 Therefore, considering that IL properties can be tailor-designed to satisfy the specific application requirements, the use of ILs for CO2 capture and in minor extension for H2S has received interest in the last few years because of their unique characteristics and high solubilities of CO2/H2S (in terms of mole fractions) in some of the studied ILs.46,47 Nevertheless, the available solubility data for CO2/H2S in the ILs studied in the literature should be handled with caution for industrial purposes if these ILs are proposed as possible alternative candidates to aminebased processes. These solubility data should be analyzed considering the moles of gases absorbed per volume of solvent (ILs), especially for ILs with high molar masses for which Henry’s law constants (mole fraction basis) are not very informative (Table 1). ILs have a wide liquid range, a negligible vapor pressure, and an adjustable physicochemical character. ILs can dissolve

Table 1. List of Henry’s Law Constant, H, of CO2 in Common Organic Solvents (at 298.15 K)41,61 solvent

H (bar)

ethanol acetone benzene cyclohexane heptane DMF CHCl3 DMSO

159.2 54.7 104.1 133.3 84.3 71.5 78.8 110

CO2 and H2S and are stable at temperatures up to several hundred degrees centigrade, which make them suitable candidates for CO2 capture/gas sweetening studies.48,49 Because most of the ILs interact with CO2 because of its quadruple moment and dispersion forces (in the case of physiosorption), with the interaction being developed between CO2 and IL anion through a weak Lewis acid-base interaction, the regeneration process could be achieved with little heat.15 It could be argued that that low energy requirements for regeneration (usually performed through stripping) is a property not only for ILs but also for the absorption in any physical solvent with high boiling point in comparison to chemical absorbents. Nevertheless, recent studies reported by Wappel et al.46 and Shiflett et al.50 have shown a slightly better performance for ILs in comparison to amine-based absorption processes at the selected constant process parameters.46 Moreover, IL performance could be improved through the suitable selection of anion/cation combinations that would lead to better absorption ability. The main goal of this review is to summarize the work focusing on the design and measurement of ILs for capturing gaseous carbon dioxide/hydrogen sulfide from fossil-fuel power plants and natural gas sweetening purposes.

(33) Romeo, L. M.; Bolea, I.; Escosa, J. M. Integration of powert plant and amine scrubbing to reduce CO2 capture costs. Appl. Therm. Eng. 2008, 28, 1039–1046. (34) Yeh, J. T.; Resnik, K. P.; Rygle, K.; Pennline, H. W. Semi-batch absorption and regeneration studies for CO2 capture by aqueous ammonia. Fuel Process. Technol. 2005, 86, 1533–1546. (35) Rao, A. B.; Rubin, E. S. A technical, economic, and environmental assessment of amine-based CO2 capture technology for power plant greenhouse gas control. Environ. Sci. Tecnol. 2002, 36, 4467–4475. (36) Rao, A. B.; Rubin, E. S.; Keith, D. W.; Morgan, M. G. Evaluation of potential cost reductions from improved amine-based CO2 capture systems. Energy Policy 2006, 34, 3765–3772. (37) Puxty, G.; Rowland, R.; Allport, A.; Yang, Q.; Bown, M.; Burns, R.; Maeder, M.; Attalla, M. Carbon dioxide postcombustion capture: A novel screening study of the carbon dioxide absorption performance of 76 amines. Environ. Sci. Technol. 2009, 43, 6427–6433. (38) Davison, J.; Thambimuthu, K. An overview of technologies and costs of carbon dioxide capture in power generation. Proc. Inst. Mech. Eng., Part A 2009, 223, 201–212. (39) Anthony, J. L.; Aki, S. N. V. K.; Maginn, E. J.; Brennecke, J. F. Feasibility of using ILs for carbon dioxide capture. Int. J. Environ. Technol. Manage. 2004, 4, 105–115. (40) Pennline, H. W.; Luebke, D. R.; Jones, K. L.; Myers, C. R.; Morsi, B. L.; Heintz, Y. J.; Ilconich, J. B. Progress in carbon dioxide capture and separation research for gasification-based power generation point sources. Fuel Process. Technol. 2008, 89, 897–907. (41) Bara, J. E.; Carlisle, T. K.; Gabriel, C. J.; Camper, D.; Finotello, A.; Gin, D. L.; Noble, R. D. Guide to CO2 separations in imidazoliumbased room-temperature ILs. Ind. Eng. Chem. Res. 2009, 48, 2739–2751. (42) Han, X.; Armstrong, D. W. ILs in separations. Acc. Chem. Res. 2007, 40, 1079–1086.  (43) Berthod, A.; Ruiz-Angel, M. J.; Carda-Broch, S. ILs in separation techniques. J. Chromatogr., A 2008, 1184, 6–18. (44) Rogers, R. D. Reflections on ILs. Nature 2007, 447, 917–918. (45) Brennecke, J. F.; Maginn, E. J. ILs: Innovative fluids for chemical processing. AIchE J. 2001, 47, 2384–2389. (46) Wappel, D.; Gronald, G.; Kalb, R.; Draxler, J. ILs for postcombustion CO2 absorption. Int. J. Greenhouse Gas Control 2010, 4, 486–494. (47) Shokouhi, M.; Adibi, M.; Jalili, A. H.; Hosseini, M.; Mehdizadeh, A. Solubility and diffusion of H2S and CO2 in the ionic liquid 1-(2hydroxyethyl)-3-methylimidazolium tetrafluoroborate. J. Chem. Eng. Data 2010, 55, 1663–1668.

2. Absorption of CO2 Using ILs 2.1. Experimental Studies. Realization of the fact that CO2 is soluble in many ILs prompted many researchers to explore the rich synthetic landscape provided by ILs (Figure 2). Preliminary studies in this field mainly focused on the ILs with imidazolium-based cations because of their observed affinity toward CO2.51-62 [bmim]PF6 and several other ILs have been measured and tested for CO2 solubility to develop phase behavior for CO2-IL pairs as well as to develop methods to obtain higher solubilities. Baltus et al. tested the solubility of carbon dioxide in derivatives of imidazoliumbased ILs, in which the length of the alkyl chain attached to (48) Kumelan, J.; Kamps, A. P.-S.; Tuma, D.; Maurer, G. Solubility of CO2 in the ILs [bmim][CH3SO4] and [bmim][PF6]. J. Chem. Eng. Data 2006, 51, 1802–1807. (49) Kim, Y. S.; Jang, J. H.; Lim, B. D.; Kang, J. W.; Lee, C. S. Solubility of mixed gases containing carbon dioxide in ILs: Measurements and predictions. Fluid Phase Equilib. 2007, 256, 70–74. (50) Shiflett, M. B.; Drew, D. W.; Cantini, R. A.; Yokozeki, A. Carbon dioxide capture using ionic liquid 1-butyl-3-methylimidazolium acetate. Energy Fuels 2010, 24, 5781–5789. (51) Blanchard, L. A.; Gu, Z.; Brennecke, J. F. High-pressure phase behavior of ionic liquid/CO2 systems. J. Phys. Chem. B 2001, 105, 2437– 2444. (52) Kim, Y. S.; Choi, W. Y.; Jang, J. H.; Yoo, K.-P.; Lee, C. S. Solubility measurement and prediction of carbon dioxide in ILs. Fluid Phase Equilib. 2005, 228-229, 439–445. (53) Lee, B.-C.; Outcalt, S. L. Solubilities of gases in the ionic liquid 1-nbutyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide. J. Chem. Eng. Data 2006, 51, 892–897. (54) Shiflett, M. B.; Yokozeki, A. Solubility of CO2 in room temperature ionic liquid [hmim][Tf2N]. J. Phys. Chem. B 2007, 111, 2070–2074.

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: DOI:10.1021/ef1011337 approximately 3.16 bar below 1 atm),64 which requires the investigation of other ILs for CO2 absorption. Shariati et al. reported that CO2 is more soluble in PF6- salt of the [bmim] compound than that of the [emim] derivative, and Peters et al. showed that [hmim] salt has higher solubility than the [emim] salt.65,66 Even though these results show that solubility could be enhanced by changing the substituents in the cation, the change is often subtle and not sufficient to explain the dramatic change in Henry constant values of different ILs. Imidazolium-based ILs have been widely studied not only to measure their CO2 absorption capacity but also to investigate the origin of interaction between CO2 and imidazolium groups. The most acidic proton of the imidazolium group located at the C2 atom is the likely candidate responsible for this interaction. For this reason, Cadena and co-workers studied a series of ILs consisting of methyl groups replaced with the mentioned hydrogen atom and used these derivatives to compare their Henry’s law constants to their hydrogen counterparts.67 Such a modification of the structure leads to a small decrease (1-3 kJ/mol) in the enthalpy of absorption, while the change of anion has a more dramatic effect in CO2 solubility, which suggests that the nature of the anion has the key role in CO2-IL interactions. Furthermore, Scovazzo et al. tested CO2 solubility of emim-based ILs with different anions to investigate the effect of the nature of the anion.68,69 Their study shows that the solubility decreases in the following order: Tf2N-, dca-, OTf-, and Cl-.68 In addition to theoretical studies focusing on understanding the interactions between the anion and CO2, experimental work also suggests the presence of such an interaction. Kazarian et al. used attenuated total reflection infrared (ATR-IR) spectroscopy to confirm the presence of CO2 bending because of the anion-CO2 interaction.70 Similarly, Kanakubo et al. confirmed experimentally that CO2 molecules are preferentially solvated to the PF6- anion.71 The efforts in this field have resulted in not only the measurement of solubility of CO2 in the previously reported ILs but also the design and preparation of new ILs with convenient functional groups that have affinities toward the CO2 molecule. Fluorination of the IL is one of the common ways to increase CO2 solubility because the affinity of CO2 to fluoroalkyl groups

Figure 2. Anions and cations that form ILs discussed in this review.

the nitrogen atom of the imidazolium ring varies (C3mim, C4mim, C6mim, and C8mim) and the octyl derivative is fluorinated (C8F13mim).63 They observed that Henry’s constant for the studied ILs decrease gradually from 39 ( 1 to 30 ( 1 bar when the length of the alkyl side chain in the cation decreases. It should be noted here that the CO2 solubilities in these ILs are too low to compete with the alkanolamine process (Henry’s law constant for the MEA-CO2 system is (55) Urukova, I.; Vorholz, J.; Maurer, G. Solubility of CO2, CO, and H2 in the ionic liquid [bmim][PF6] from Monte Carlo simulations. J. Phys. Chem. B 2005, 109, 12154–12159. (56) Shah, J. K.; Maginn, E. J. Monte Carlo simulations of gas solubility in the ionic liquid 1-n-butyl-3-methylimidazolium hexafluorophosphate. J. Phys. Chem. B 2005, 109, 10395–10405. (57) Shiflett, M. B.; Yokozeki, A. Solubilities and diffusivities of carbon dioxide in ILs: [bmim][PF6] and [bmim][BF4]. Ind. Eng. Chem. Res. 2005, 44, 4453–4464. (58) Camper, D.; Becker, C.; Koval, C.; Noble, R. Low pressure hydrocarbon solubility in room temperature ILs containing imidazolium rings interpreted using regular solution theory. Ind. Eng. Chem. Res. 2005, 44, 1928–1933. (59) Scovazzo, P.; Camper, D.; Kieft, J.; Poshusta, J.; Koval, C.; Noble, R. Regular solution theory and CO2 gas solubility in roomtemperature ILs. Ind. Eng. Chem. Res. 2004, 43, 6855–6860. (60) Kamps, A. P.-S.; Tuma, D.; Xia, J.; Maurer, G. Solubility of CO2 in the ionic liquid [bmim][PF6]. J. Chem. Eng. Data 2003, 48, 746–749. (61) Anthony, J. L.; Maginn, E. J.; Brennecke, J. F. Solubilities and thermodynamic properties of gases in the ionic liquid 1-n-butyl-3methylimidazolium hexafluorophosphate. J. Phys. Chem. B 2002, 106, 7315–7320. (62) Anthony, J. L.; Anderson, J. L.; Maginn, E. J.; Brennecke, J. F. Anion effects on gas solubility in ILs. J. Phys. Chem. B 2005, 109, 6366– 6374. (63) Baltus, R. E.; Culbertson, B. H.; Dai, S.; Luo, H.; DePaoli, D. W. Low-pressure solubility of carbon dioxide in room-temperature ILs measured with a quartz crystal microbalance. J. Phys. Chem. B 2004, 108, 721–727.

(64) Akanksha, K. K. P.; Srivastava, V. K. Mass transport correlation for CO2 absorption in aqueous monoethanolamine in a continuous film contactor. Chem. Eng. Process. 2008, 47, 920–928. (65) Shariati, A.; Peters, C. J. High-pressure phase behavior of systems with ILs: Part III. The binary system carbon dioxide þ 1-hexyl-3-methylimidazolium hexafluorophosphate. J. Supercrit. Fluids 2004, 30, 139–144. (66) Shariati, A.; Peters, C. J. High-pressure phase behavior of systems with ILs: Part II. The binary system carbon dioxide þ 1-hexyl-3-methylimidazolium hexafluorophosphate. J. Supercrit. Fluids 2004, 29, 43–48. (67) Cadena, C.; Anthony, J. L.; Shah, J. K.; Morrow, T. I.; Brennecke, J. F.; Maginn, E. J. Why is CO2 so soluble in imidazolium-based ILs? J. Am. Chem. Soc. 2004, 126, 5300–5308. (68) Scovazzo, P.; Kieft, J.; Finan, D. A.; Koval, C.; DuBois, D.; Noble, R. Gas separation using non-hexafluorophosphate [PF6-] anion supported ionic liquid membranes. J. Membr. Sci. 2004, 238, 57–63. (69) Camper, D.; Scovazzo, P.; Koval, C.; Noble, R. Gas solubilities in room-temperature ILs. Ind. Eng. Chem. Res. 2004, 43, 3049–3054. (70) Kazarian, S. G.; Briscoe, B. J.; Welton, T. Combining ILs and supercritical fluids: In situ ATR-IR study of CO2 dissolved in two ILs at high pressures. Chem. Commun. 2000, 2047–2048. (71) Kanakubo, M.; Umecky, T.; Hiejima, Y.; Aizawa, T.; Nanjo, H.; Kameda, Y. Solution structures of 1-butyl-3-methylimidazolium hexafluorophosphate ionic liquid saturated with CO2: Experimental evidence of specific anion-CO2 interaction. J. Phys. Chem. B 2005, 109, 13847– 13850.

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: DOI:10.1021/ef1011337 Table 2. List of Henry’s Law Constant, H, of CO2 in Selected ILs (at 298.15 K) IL

abbreviation

1-ethyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide 1-butyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide 1-hexyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide 1-butyl-3-methylimidazolium hexafluorophosphate 1-ethyl-3-methylimidazolium tetrafluoroborate 1-butyl-3-methylimidazolium tetrafluoroborate 1-methyl-3-(3,3,4,4,5,5,6,6,6-nonafluorohexyl)imidazolium bis(trifluoromethylsulfonyl)imide 1-methyl-3-(3,3,4,4,5,5,6,6,7,7,8,8,8-tridecafluorooctyl)imidazolium bis(trifluoromethylsulfonyl)imide 1-hexyl-3-methylpyrinium bis(trifluoromethylsulfonyl)imide 1-hexyl-3-methylimidazolium tris(pentafluoroethyl)trifluorophosphate 1-hexyl-3-methylimidazolium tris(heptafluoropropyl)trifluorophosphate 1-pentyl-3-methylimidazolium tris(nonafluorobutyl)trifluorophosphate 1-ethyl-3-methylimidazolium trifluoromethanesulfonate 1,1,3,3-tetramethylguanidium lactate lactic acid 1-octyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide 1-propyl-2,3-methylimidazolium bis(trifluoromethylsulfonyl)imide 1-butyl-3-methylpyrinium bis(trifluoromethylsulfonyl)imide 1-butyl-2,3-methylimidazolium hexafluorophosphate 1-butyl-2,3-methylimidazolium tetrafluoroborate 1-ethyl-2,3-methylimidazolium bis(trifluoromethylsulfonyl)imide 1-ethyl-3-methylimidazolium dicyanamide 1-ethyl-3-methylimidazolium trifluoromethanesulfonate

[emim][Tf2N] [bmim][Tf2N] [hmim][Tf2N] [bmim][PF6] [emim][BF4] [bmim][BF4] [C6H4F9mim][Tf2N]

35.6 ( 1.4 33.0 ( 0.3 31.6 ( 0.2 53.4 ( 0.3 80 ( 4 59.0 ( 2.6 28.4 ( 0.1

67 67 67 62 58 62 67

[C8H4F13mim][Tf2N]

27.3 ( 0.2

67

[hmpy][Tf2N] [hmim][eFAP] [hmim][pFAP] [p5mim][bFAP] [emim][OTf] TMGL [C8mim][Tf2N] [pmmim][Tf2N] [bmpy][Tf2N] [bmmim][PF6] [bmmim][BF4] [emmim][Tf2N] [emim][dca] [emim][OTf]

32.8 ( 0.2 25.2 ( 0.1 21.6 ( 0.1 20.2 ( 0.1 73 ( 1.0 at 303 K 14.6 at 308 K 30 ( 1 38.5 ( 0.9 33 ( 1 61.8 ( 2.1 61.0 ( 1.6 39.6 ( 1.4 78 ( 3.0 73 ( 1.0

67 67 67 67 67 76 67 67 67 67 67 67 59 59

is well-known. As part of a continuing effort to use CO2 as a solvent for polymers, DeSimone and co-workers found that polyfluoroalkyl acrylate is miscible with CO2 below 150 bar, while nonfluoro polymers were reported to be insoluble at pressures around 200 bar.72 This method has been successfully employed to this field. A summary and comparison of Henry’s law constant values for nonfluoro ILs and their corresponding fluorinated substituents are tabulated in Table 2. For example, Henry’s law constant for [C8H4F13mim][Tf2N] is reported to be 6 ( 1 bar at 25 °C by Baltus et al., and later a value of 27.3 ( 0.2 bar at the same temperature is reported by Brennecke et al. for the same compound, while the value for [C8mim][Tf2N] is 30 ( 1 bar.73 Similarly, Henry’s law constant for [C6H4F9mim][Tf2N] (28.4 ( 0.1 bar at 25 °C) is smaller than that of [hmim][Tf2N] (31.6 ( 0.2 bar at 25 °C). It should be noted that insertion of the fluorine atom to the cationic part of IL leads to a small increase in CO2 solubility. Considering that the anion-CO2 interaction plays a primary role in CO2 solubility, one expects a higher decrease in Henry’s law constant and, thus, a higher CO2 solubility when fluorinated derivatives of the anionic part are used. For this purpose, Brennecke et al. used derivatives of the PF6- anion, eFAP-, pFAP-, and bFAP-, where three fluorine atoms are replaced with pentafluoroethyl, heptafluoropropyl, and nonafluorobutyl groups (Table 2).73 Even though the data clearly indicate an increase in CO2 solubility as the length of the fluoroalkyl group increases, Bara et al. found that the contribution of fluorination increases the CO2 solubility (mol/volume) not more than 10% for the mentioned compounds.74 The origin of obtaining higher solubility with fluorinated substituents could be attributed to the following factors: (i) fluorinated complexes having weaker

H (bar)

reference

self-interactions, leading to higher miscibility with CO2, (ii) a Lewis acid-base interaction between electronegative fluorine atoms and the electron-poor carbon atom of CO2, and (iii) a hydrogen-bonding interaction between oxygen atoms of CO2 and relatively more acidic protons of IL because of the presence of neighboring fluorine atoms. The mentioned fluorinated compounds, however, could be less environmentally friendly because of their high toxicity values. Such characteristics undermine their potential use as green solvents. The use of fluorine contributes significantly to the cost of CO2 capture because it increases the experimental cost and leads to ILs with a high viscosity value, which renders this procedure less favorable. Even though a high viscosity problem could be avoided by mixing these ILs with convenient organic compounds, this alternative method leads to disadvantages, such as volatility, that are common in organic solvents. Therefore, other methods should also be considered in this field to avoid such limitations. ILs with oxygen-containing functional groups have received much attention because the electron-poor carbon atom of CO2 has lability toward electronegative atoms.75 For example, the solubility of CO2 in ether-containing ILs, PEG-5 cocomonium methylsulfate (Ecoeng 500) and 1-butyl-3-methylimidazolium 2-(2-methoxyethoxy)ethylsulfate (Ecoeng 41M), is very similar to that in [hmim][Tf2N], which is claimed to be due to the electronegative oxygen atom as well as the free space created by the flexible ether group.74 Similarly, [b2-Nic][Tf2N], an IL with an ester group, exhibits similar CO2 solubility behavior of CO2 at 60 °C to that in [hmim][Tf2N].75 Zhang et al. investigated the CO2 absorption capacity of an IL consisting of the 1,1,3,3tetramethyl-guanidium lactate (TMGL) anion because of its success for the absorption of CO2.76 The solubility of CO2 in the mentioned IL was found to be slightly larger than that in [bmim][PF6]. The use of alkylguanidinium lactate ILs for

(72) DeSimone, J. M.; Guan, Z.; Elsbernd, C. S. Synthesis of fluoropolymers in superciritical carbon dioxide. Science 1992, 257, 945– 946. (73) Muldoon, M. J.; Aki, S. N. V. K.; Anderson, J. L.; Dixon, J. K.; Brennecke, J. F. Improving carbon dioxide solubility In ILs. J. Phys. Chem. B 2007, 111, 9001–9009. (74) Bara, J. E.; Gabriel, C. J.; Carlisle, T. K.; Camper, D. E.; Finotello, A.; Gin, D. L.; Noble, R. D. Gas separations in fluoroalkylfunctionalized room-temperature ionic liquids using supported liquid membranes. Chem. Eng. J. 2009, 147, 43–50.

(75) Zhang, X.; Liu, Z.; Wang, W. Screening of ILs to capture CO2 by COSMO-RS and experiments. AIChE J. 2008, 54, 2717–2728. (76) Zhang, S.; Yuan, X.; Chen, Y.; Zhang, X. Solubilities of CO2 in 1-butyl-3-methylimidazolium hexafluorophosphate and 1,1,3,3-tetramethylguanidium lactate at elevated temperatures. J. Chem. Eng. Data 2005, 50, 1582–1585.

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Figure 3. Stoichiometric reaction of TSILs with CO2.

at 1666 cm-1. It should also be noted that the TSIL could be regenerated by heating the carbamate product to 80-100 °C at vacuum for several hours. Given the mentioned success, other ILs with amine functional groups have been prepared, some of which are amino-acid-based. Amino-acid-based compounds could be of significant use in this field because they possess both carboxyl and amine functional groups. This approach was successfully used by Gurkan et al. to design ILs that react with CO2 in a 1:1 stoichiometric ratio.82 [P66614][Met] and [P66614][Pro] could be used to absorb CO2 via the chemical complexation mechanism, as shown in Figure 3. The proposed reaction and 1:1 absorption is in good agreement with the results obtained from IR spectroscopy and CO2 absorption studies. Zhang et al.83 developed (3-aminopropyl)tributylphosphonium amino acid ILs for CO2 capture, leading to 1 mol of CO2 captured/1 mol of IL, and the ILs can be repeatedly recycled for CO2 uptake. Nevertheless, the use of amine-functionalized ILs for CO2 uptake shows several problems, with the main one being their high viscosities at ambient temperature, which are even larger upon complexation with CO2;84 this fact would hinder the CO2 diffusion and uptake rate. Likewise, the synthesis of these ILs would require several synthetic steps,4 thus limiting their commercial viability. Some authors have proposed the use of mixed IL-amine solutions for CO2 capture. Camper et al.85 used 50 mol % of ILs with MEA and DEA, leading to rapid and reversible captures of 1 mol of CO2/1 mol of MEA. These mixtures are considered as analogous to the water-amine solutions commonly used for CO2/gas sweetening purposes but with increasing efficiency arising from the properties of used ILs and retaining the capturing ability of used amines. For natural gas sweetening purposes, not only is the CO2/ H2S solubility important but also the selectivity of these

natural gas sweetening purposes has also been discussed.77 The lactate anion, consisting of both carboxyl and hydroxyl groups, as well as alkylguanidinium complexes with three alkylamine groups could both be produced in large scale, which makes this class of materials convenient for industrial purposes. Furthermore, because these ILs could have toxicity in an acceptable range for pharmaceutical purposes, they have been used as model compounds to study their interaction with CO2. Similar to the guanidium-lactate system, one common approach to obtain a low Henry’s law constant is to obtain ILs with more than one specific functional group to increase the number of potential sites for the CO2 interaction. A group of hydroxyl ammonium ILs has been studied by Yuan and co-workers for this purpose. They obtained similar solubility values to those for imidazolium-based ILs.78 Another approach to increasing the CO2 solubility is to incorporate amine functional groups to ILs to mimic the use of aminebased processes for CO2 absorption. This is a clear application of the design of task-specific ILs (TSILs) because they are specifically designed to increase the CO2-IL interaction.79-81 This method has first been applied by Bates et al. to prepare an imidazolium-based IL with an amine functional group attached to one of the alkyl chains (Figure 2).4 For this family of ILs, the absorption of CO2 is carried out through chemisorption, the proposed reaction, in which the studied IL reacting with CO2 in a 2:1 stoichiometric ratio is in good agreement with experimental results, showing that the mole ratio of CO2/TSIL approaches a maximum of 0.5 over the course of 3 h. The authors also use Fourier transform infrared (FTIR) spectroscopy to observe the formation of a carbamate product, which possesses a new absorption band (77) Aparicio, S.; Atilhan, M. Computational study of hexamethylguanidinium lactate ionic liquid: A candidate for natural gas sweetening. Energy Fuels 2010, 24, 4989–5001. (78) Yuan, X.; Zhang, S.; Liu, J.; Lu, X. Solubilities of CO2 in hydroxyl ammonium ILs at elevated temperatures. Fluid Phase Equilib. 2007, 257, 195–200. (79) Wasserscheid, P.; Keim, W. ILs;New “solutions” for transition metal catalysis. Angew. Chem., Int. Ed. 2000, 39, 3772–3789. (80) Cole, A. C.; Jensen, J. L.; Ntai, I.; Tran, K. L. T.; Weaver, K. J.; Forbes, D. C.; Davis, J. H., Jr. Novel Brønsted acidic ILs and their use as dual solvent-catalysts. J. Am. Chem. Soc. 2002, 124, 5962–5963. (81) Yao, Q.; Zhang, Y. Olefin metathesis in the ionic liquid 1-butyl3-methylimidazolium hexafluorophosphate using a recyclable Ru catalyst: Remarkable effect of a designer ionic tag. Angew. Chem., Int. Ed. 2003, 42, 3395–3398.

(82) Gurkan, B. E.; de la Fuente, J. C.; Mindrup, E. M.; Ficke, L. E.; Goodrich, B. F.; Price, E. A.; Schneider, W. F.; Brennecke, J. F. Equimolar CO2 absorption by anion-functionalized ILs. J. Am. Chem. Soc. 2010, 132, 2116–2117. (83) Zhang, Y.; Zhang, S.; Lu, X.; Zhou, Q.; Fan, W.; Zhang, X. Dual amino-functionalised phosphonium ILs for CO2 capture. Chem.;Eur. J. 2009, 15, 3003–3011. (84) Gutowski, K. E.; Maginn, E. J. Amine-functionalized task-specific ILs: A mechanistic explanation for the dramatic increase in viscosity upon complexation with CO2 from molecular simulation. J. Am. Chem. Soc. 2008, 130, 14690–14704. (85) Camper, D.; Bara, J. E.; Gin, D. L.; Noble, R. D. Roomtemperature ionic liquid-amine solutions: Tunable solvents for efficient and reversible capture of CO2. Ind. Eng. Chem. Res. 2008, 47, 8496–8498.

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gases in relation to CH4 solubility. Although experimental studies on methane solubility in ILs are scarce, recent results have shown that functionalization of ILs to increase CO2 solubility also leads to a decrease in CO2/CH4 selectivity, for both fluor- and oxygen-containing ions.74 Therefore, selectivity should also be analyzed in detail for future IL candidates, and thus, through the selection of adequate functional groups, CO2 solubility may be increased without decreasing CO2/CH4 selectivity.74 2.2. Theoretical Analysis. The analysis of the available literature shows that despite the promising properties of many ILs for CO2 capture purposes, they are not competitive with the current amine-based technologies, especially for low CO2 partial pressure conditions. The main reason is the low carrying capacity of studied ILs in comparison to MEA or other amines currently used in the industry. Nevertheless, this should not discard the use of ILs for CO2 uptake purposes; the possibility of developing TSILs should be the most useful approach. The number of possible IL candidates for CO2 capturing is extremely large, and keeping in mind that performing experimental measurements of CO2 solubility is extremely time- and resource-consuming, computational methods for IL screening are very useful. Zhang et al.74 and Maitie86 used the COSMO-RS approach; this method allows for the analysis of CO2 uptake based on the prediction of Henry’s law constant without any specific parameter adjustment at very low computational cost. Zhang et al.73 analyzed 408 kinds of ILs, leading to the conclusion of [FEP] anion (tris(pentafluoroethyl)trifluorophosphate)-based ILs as the most promising fluids. They also reported that a value of 0.2 MPa seems to be the lowest limit of Henry’s law constant of CO2 in ILs at 298 K in physical absorption. Atomistic simulations have also been used to analyze, from a molecular viewpoint, and predict the CO2 absorption in ILs. Maginn et al. developed a continuous fractional Monte Carlo method that was applied successfully to predict gas solubilities in ILs.87,88 The proposed method applied for 1-n-hexyl-3methylimidazolium bis(trifluoromethylsulfonyl)imide predicted CO2 absorption isotherms and Henry’s law constants that agree quantitatively with experimental data. Moreover, the simulation results show how the absorption of CO2 molecules does not disrupt the structure of the IL, and the absorbed molecules occupy available free volume up to high concentrations, in which subtle rearrangements of alkyl chains in cations allow further CO2 absorption.88 This method was also applied to 1-n-hexyl-3-methylimidazolium [FEP] IL, which was considered a suitable candidate for gas absorption from COSMORS results,73 and the predicted absorption isotherms and Henry’s law constants are in fair agreement with highly accurate experimental results obtained by the same authors.89 Atomistic simulations for 1-n-hexyl-3-methylimidazolium [FEP] show that large quantities of absorbed CO2 have little effect on the overall IL structure. The authors show how the

balance of van der Waals/Coulombic interactions of CO2 molecules with cation/anion determine the solubility of the gas in the studied IL.89 Huang et al.90 analyzed using molecular dynamics simulations the reasons underlying the low partial molar volume of CO2 when absorbed in ILs, the authors show that most of the space occupied by CO2 in the IL phase consists of very localized cavities of large size that are formed by small angular rearrangements of the anions. Bhargava et al. developed molecular dynamics simulations of CO2-[bmim][PF6] mixtures using a refined atomistic potential model for the IL. The authors analyzed the volume expansion upon CO2 absorption, which is in good agreement with experimental data, and suggested the development of anion-CO2 interactions as the main factor in CO2 absorption.91 Wang et al.92 developed an all-atom force field for the analysis of CO2 and SO2, where absorption in 1,1,3,3-tetramethylguanidium lactate IL showed relatively weak organization of CO2 around both the cation and anion of the IL. Aparicio and Atilhan77 studied the CO2 absorption in hexamethylguanidinium lactate IL, showing that cavities in pure ILs are too small to fit absorbed molecules, and thus, the absorption of these molecules should lead to a rearrangement of these cavities, leading to larger ones to fit solute molecules without large volume expansion. Moreover, absorption leads to a stronger effect on structuring around the anion than around the cation, although it leads to a weakening of all of the ion-ion interactions. The underlying mechanism of the CO2 interaction with anion/cation has also been studied using quantum chemistry methods. Yu et al.93 analyzed the frontier molecular orbital interactions between CO2 and -NH2 in 1,1,3,3-tetramethylguanidinium lactate and 1-n-propylamine-3-butylimidazolium tetrafluoroborate ILs; it was shown how the addition of electron-donating groups attached to the amino groups and the intramolecular hydrogen bond arising from the hydrogen in the amino group may rise the energy of the frontier occupied molecular orbital of the amino group, therefore improving the CO2/amino interactions and thus leading to a higher solubility of CO2 in the IL. According to this approach, two new TSILs were designed, tetrabutylphosphonium alanine and tetrabutylphosphonium glycine, whose larger absorption capacity was confirmed experimentally.94 Bhargava et al.95 studied anion/CO2 interactions for commonly used anions using density-functional-theory-based calculations for isolated clusters. A Lewis acid-base type of interaction between CO2 and the anions was analyzed, whose strength was found to be directly proportional to the basicity of the anion, showing that the interaction disrupts the CO2 molecule from (90) Huang, X.; Margulis, C. J.; Li, Y.; Berne, B. J. Why is the partial molar volume of CO2 so small when dissolved in a room temperature ionic liquid? Structure and dynamics of CO2 dissolved in [bmimþ] [PF6-]. J. Am. Chem. Soc. 2005, 127, 17842–17851. (91) Bhargava, B. L.; Krishna, A. C.; Balasubramanian, S. Molecular dynamics simulation studies of CO2-[bmim][PF6] solutions: Effect of CO2 concentration. AIChE J. 2008, 54, 2971–2978. (92) Wang, Y.; Pan, H.; Li, H.; Wang, C. Force field of the TMGL ionic liquid and the solubility of SO2 and CO2 in the TMGL from molecular dynamics simulation. J. Phys. Chem. B 2007, 111, 10461– 10467. (93) Yu, G.; Zhang, S.; Yao, X.; Zhang, J.; Dong, K.; Mori, R. Design of task-specific ILs for capturing CO2: A molecular orbital study. Ind. Eng. Chem. Res. 2006, 45, 2875–2880. (94) Zhang, J.; Zhang, S.; Dong, K.; Zhang, Y.; Lu, X. Supported absorption of CO2 by tetrabutylphosphonium amino acid ILs. Chem.; Eur. J. 2006, 12, 4021–4026. (95) Bhargava, B. L.; Balasubramanian, S. Probing anion-carbon dioxide interactions in room temperature ILs: Gas phase cluster calculations. Chem. Phys. Lett. 2007, 444, 242–246.

(86) Maiti, A. Theoretical screening of ionic liquid solvents for carbon capture. ChemSusChem 2009, 2, 628. (87) Shi, W.; Maginn, E. J. Continuous fractional component Monte Carlo: An adaptive biasing method for open system atomistic simulations. J. Chem. Theory Comput. 2007, 3, 1451–1463. (88) Shi, W.; Maginn, E. J. Atomistic simulation of the absorption of carbon dioxide and water in the ionic liquid 1-n-hexyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide ([hmim][Tf2N]. J. Phys. Chem. B 2008, 112, 2045–2055. (89) Zhang, X.; Huo, F.; Liu, Z.; Wang, W.; Shi, W.; Maginn, E. J. Absorption of CO2 in the ionic liquid 1-n-hexyl-3-methylimidazolium tris(pentafluoroethyl)trifluorophosphate ([hmim][FEP]): A molecular view by computer simulations. J. Phys. Chem. B 2009, 113, 7591–7598.

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linearity. Shuda et al.96 analyzed the interaction of several common anions with CO2 using ab initio calculations at the MP2 level. The authors reported deviations from linearity in CO2 molecules upon the interaction with anions and showed that the binding energy of anion-CO2 complexes is inversely proportional to the solubility in the corresponding IL. The results mentioned above briefly summarize the general methods to increase the CO2 solubility of ILs. However, the CO2/hydrocarbon selectivity should also be considered to remove CO2 from natural gas resources, which will be discussed in section 4.41 2.3. Process Design. In addition to the study of CO2 solubilities in several families of ILs, several process options to carry out the gas absorption from flue gases or natural gas sources have been considered in the literature. The use of ILs and membranes for CO2 absorptive purposes has been analyzed in detail by Bara et al.,97 showing the significant possibilities of this approach. Scovazzo et al.98 analyzed gas separation processes using supported liquid membranes (SLMs), with ILs non-containing the [PF6] anion showing highly selective membranes. These authors reported lately that IL membranes showed adequate CO2/CH4 selectivity and some of them may be economically competitive.99 Park et al.100 used IL-based SLMs for the removal of CO2 and H2S from crude natural gas; the reported results showed good performance for the studied membranes (even at high pressure), leading to high selectivities for both CO2/CH4 and H2S/CH4. Baltus et al.101 used SLMs to separate CO2 and reported an economical analysis of the process, showing that this separation process may be competitive with amine scrubbing. Baltus et al.101 also showed that, if the same ILs are used in absorption towers, the economic analysis is not so favorable. Hanioka et al.102 also used SLMs for CO2 separation using 1-butyl-3methylimidazolium bis(trifluoromethylsulfonyl)imide [C4mim][Tf2N], N-aminopropyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide [C3NH2mim][Tf2N], and N-aminopropyl3-methylimidazolium trifluoromethanesulfone [C3NH2mim][CF3SO3] ILs. The authors used the SLMs for the separation of CO2 from CO2-CH4 mixtures, with the partial pressure of CO2 being about 2-50 kPa. The results showed high CO2 selectivity and high stability of SLMs for long periods of time. The use of SLMs has several shortcomings, as reported by Bara et al.;41 the SLMs actually used are too thick for industrial purposes, and these SLMs could not withstand the high pressures required

for natural gas sweetening. Nevertheless, they could be used for CO2 absorption from flue gases, which is performed at atmospheric pressure conditions. To overcome some of these drawbacks, Radosz et al.103-105 have proposed the use of membrane materials for CO2 separation made of poly(IL)s having absorption capacities larger than those of roomtemperature ILs with reversible and fast sorption and desorption. Barghi et al.106 studied CO2/CH4 selectivity using [bmim][PF6] IL supported on an alumina membrane, showing that permeability and solubility of CO2 is much higher than that of CH4; the authors state than the studied membrane setup can provide higher selectivity for CO2/CH4 separation in comparison to traditional polymer membranes at ambient conditions, although economical analysis of the cost is not reported by the authors. Myers et al.107 performed a wide analysis on the use of IL-based membranes for CO2 separation, analyzing multi-component gas interactions with supported IL membranes and showing the complex effects arising from these interactions and the effects on membrane selectivity. Hudiono et al.108 used a complex membrane formed by three components of polymerizable liquids and zeolite materials for CO2 separation, showing adequate CO2/ CH4 selectivity. Yoo et al.109 proposed the use of IL membranes using a Nafion matrix, leading to a good CO2/CH4 separation factor, which is higher than that of free IL. The analysis of SLM based on unconventional ILs was reported by Cserjesi et al.,110 showing that these systems lead to better selection properties, even for the CO2/CH4 pair, than the commonly used, industrial polymer membranes. Scovazzo et al.111 analyzed the available information on the use of ILbased SLMs, recommending the use of ILs with smaller molar volumes for CO2 separation and the study of functionalized ILs that may lead to an increase in capturing efficiency. The typical CO2 capture and recovery setup used in most commercial plants uses an absorber and stripper design, as reported in Figure 4. The mentioned setup using MEA as an absorber could also be implemented for IL-based absorbents. Anthony et al.39 used a standard absorber setup but in a static configuration, with only the gas mixture flowing, (103) Tang, J.; Tang, H.; Sun, W.; Plancher, H.; Radosz, M.; Shen, Y. Poly(ionic liquid)s: A new material with enhanced and fast CO2 absorption. Chem. Commun. 2005, 3325–3327. (104) Tang, J.; Sun, W.; Tang, H.; Radosz, M.; Shen, Y. Enhanced CO2 absorption of poly(ionic liquid)s. Macromolecules 2005, 38, 2037– 2039. (105) Tang, J.; Tabng, H.; Sun, W.; Radosz, M.; Shen, Y. Low pressure CO2 sorption in ammonium-based poly(ionic liquid)s. Polymer 2005, 46, 12460–12464. (106) Barghi, S. H.; Adibi, M.; Rashtchian, D. An experimental study on permeability, diffusivity, and selectivity of CO2 and CH4 through [bmim][PF6] ionic liquid supported on an alumina membrane: Investigation of temperature fluctuations effects. J. Membr. Sci. 2010, 362, 346–352. (107) Myers, C.; Luebke, D. R.; Pennline, H. W.; Ilconich, J. B. Ionic liquid membranes for carbon dioxide separation. In Membrane Gas Separation; Yampolskii, Y., Freeman, B., Eds.; John Wiley and Sons: New York, 2010. (108) Hudiono, Y.; Carlisle, T. K.; Bara, J. E.; Zhang, Y.; Gin, D. L.; Noble, R. D. A three-component mixed-matrix membrane with enhanced CO2 separation properties based on zeolites and ionic liquid materials. J. Membr. Sci. 2010, 350, 117–123. (109) Yoo, S.; Won, J.; Kang, S. W.; Kang, Y. S.; Nagase, S. CO2 separation membranes using ionic liquids in a nafion matrix. J. Membr. Sci. 2010, 363, 72–79. (110) Cserjesi, P.; Nemest othy, N.; Belafi-Bak o, K. Gas separation properties of supported liquid membranes prepared with unconventional ionic liquids. J. Membr. Sci. 2010, 349, 6–11. (111) Scovazzo, P. Determination of the upper limits, benchmarks, and critical properties for gas separations using stabilized room temperature ionic liquid membranes (SILMs) for the purpose of guiding future research. J. Membr. Sci. 2009, 343, 199–211.

(96) Shuda, S. Y.; Khana, A. Evaluation of interactions of CO2-ionic liquid systems through molecular modelling. World Acad. Sci. Eng. Technol. 2009, 57, 539–542. (97) Bara, J. E.; Camper, D. E.; Gin, D. L.; Noble, R. D. Roomtemperature ILs and composite materials: Platform technologies for CO2 capture. Acc. Chem. Res. 2010, 43, 152–159. (98) Scovazzo, P.; Kieft, J.; Finan, D. A.; Koval, C.; DuBois, D.; Noble, R. Gas separations using non-hexafluorophosphate [PF6]-anion supported ionic liquid membranes. J. Membr. Sci. 2004, 238, 57–63. (99) Scovazzo, P.; Havard, D.; McShea, M.; Mixon, S.; Morgan, D. Long-term, continuous mixed-gas dry fed CO2/CH4 and CO2/N2 separation performance and selectivities for room temperature ionic liquid membranes. J. Membr. Sci. 2009, 327, 41–48. (100) Park, Y.; Kim, B.; Byum, Y. Preparation of supported ionic liquid membranes (SILMs) for the removal of acidic gases from crude natural gas. Desalination 2009, 236, 342–348. (101) Baltus, R. E.; Counce, R. M.; Culbertson, B. H.; Luo, H. L.; DePaoli, D. W.; Dai, S.; Duckworth, D. C. Examination of the potential of ILs for gas separations. Sep. Sci. Technol. 2005, 40, 525–541. (102) Hanioka, S.; Maruyama, T.; Sotani, T.; Teramoto, M.; Matsuyama, H.; Nakashima, K.; Hanaki, M.; Kubota, F.; Goto, M. CO2 separation facilitated by task-specific ILs using a supported liquid membrane. J. Membr. Sci. 2008, 314, 1–4.

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: DOI:10.1021/ef1011337 Table 3. List of Henry’s Law Constant, H, of H2S in Selected ILs IL [bmim][PF6] [hmim][PF6] [bmim][BF4] [hmim][BF4] [1-(2-hydroxyethyl)3-methylimidazolium] [BF4] [bmim][Tf2N] [hmim][Tf2N] [emim][ethylsulfate]

T (K)

H (bar)

references

298.15 303.15 303.15 303.15 303.15 303.15

14.3 ( 0.1 18.6 ( 0.2 17.9 ( 0.2 15.5 ( 0.1 12.5 ( 0.1 31.3 ( 0.1

113 117 116 117 116 47

303.15 303.15 303.15

13.7 ( 0.1 17.4 ( 0.1 60.7 ( 0.2

117 107 116

Nevertheless, concentrations of both compounds in raw gases used to be very different, and requirements for reduction levels for both gases are clearly different.13 Therefore, several scenarios are possible for the removal of these acid gases that lead to very complex process design and several problems arising from the use of amine-based absorbents, such as the inability to selectively remove H2S in the presence of CO2.13,25,26 Therefore, the use of ILs to remove H2S from natural gas resources is a recently growing field that has been considered by some authors as a possible alternative to amine-based processes. Even though considerable work is being performed regarding the physical, thermodynamic, and even engineering aspects of use of ILs for CO2 removal, comparatively, much less work has focused on H2S removal. Henry’s law constant for the solubility of H2S in a common IL, [bmim][PF6], has been reported as 14.3 bar at 25 °C by Jou et al.113 They also concluded that the studied IL is not convenient for complete removal of H2S from natural gas because it cannot reduce the H2S concentration to the required range. Moreover, more research upon measuring the H2S solubility in several imidazolium-based ILs (Table 3) was performed by Pomelli et al.,114 Shiflett et al.,115 Rahmati-Rostami et al.,116 Jalili et al.,117,118 and Shokokui et al.47 Nevertheless, the solubility behavior of these imidazolium-based ILs is not competitive with the solubility in amine compounds. Heintz et al.119 studied H2S and CO2 removal from dry fuel gas streams using an IL-containing chloride anion and an ammonium-based cation, whose composition is not fully clarified because it is proprietary of the manufacturing company. Reported results show H2S solubilities in the IL greater than those of CO2, which is in agreement with the behavior reported in the literature for imidazolium-based ILs.114-117

Figure 4. Typical absorber/stripper setup used for CO2 capture. The pressure and temperature operating conditions are approximate values for common absorption units.

with IL coated in glass beads in the absorbent. These authors showed that it is technically feasible to perform CO2 absorption using [bmim][PF6] IL with this absorbent configuration. Arshad112 performed a process evaluation analysis of CO2 absorption using a common absorption column-stripping column setup, with the absorber in a counter current configuration, by ILs [emim][Tf2N] and [bmim][Tf2N] and compared it to a conventional MEA-based technology. The analysis showed that the energy required to capture CO2 from a gas source at 0.1 bar CO2 partial pressure is 76 and 71 times larger using [emim][Tf2N] and [bmim][Tf2N] ILs, respectively, compared to using MEA solutions. The large difference rises from the poor absorption capacity of the proposed ILs that make necessary larger quantities of absorbent per kilogram of absorbed CO2. Nevertheless, increasing the CO2 partial pressure decreases the energy requirements for IL absorbents, whereas it has no effect on the MEA process. It should also be noted here that aqueous MDEA and piperazine solutions, which are also used in the natural gas industry, have lower energy requirements than the MEA process while maintaining higher CO2 absorption capacities than ILs. Arshad112 proposed a theoretical IL that could be competitive with MEA absorption technologies, showing that the main requirement is to increase the carrying capacity of the IL and, thus, proposing a theoretical IL with Henry’s law constant lower than 1 bar at 25 °C (or up to 16 bar if exposed to higher CO2 partial pressures). The development of such an IL seems to be extremely challenging, when considering the thermodynamic principles and the critical properties of CO2. Wappel et al.46 performed an economical analysis of the CO2 absorption process using a mixture of water/IL to decrease viscosity as an absorbent.

(113) Jou, F. Y.; Mather, A. E. Solubility of hydrogen sulfide in [bmim][PF6]. Int. J. Thermophys. 2007, 28, 490–495. (114) Pomelli, C. S.; Chiappe, C.; Vidis, A.; Laurenczy, G.; Dyson, P. J. Influence of the interaction between hydrogen sulfide and ionic liquids on solubility: Experimental and theoretical investigation. J. Phys. Chem. B 2007, 111, 13014. (115) Shiflett, M. B.; Yokozeki, A. Separation of CO2 and H2S using room-temperature ionic liquid [bmim][PF6]. Fluid Phase Equilib. 2010, 294, 105–113. (116) Rahmati-Rostami, M.; Ghotbi, C.; Hosseini-Jenab, M.; Ahmadi, A. N. Solubility of H2S in ILs [hmim][PF6], [hmim][BF4], and [hmim][Tf2N]. J. Chem. Thermodyn. 2009, 41, 1052–1055. (117) Jalili, A. H.; Rahmati-Rostami, M.; Ghotbi, C.; Hosseini-Jenab, M.; Ahmadi, A. N. Solubility of H2S in ILs [bmim][PF6], [bmim][BF4], and [bmim][Tf2N]. J. Chem. Eng. Data 2009, 54, 1844–1849. (118) Jalili, A.; Mehdizadeh, A.; Shokokui, M.; Ahmadi, A. N.; Hosseini-Jenab, M.; Faterminassab, F. Solubility and diffusion of CO2 and H2S in the ionic liquid 1-ethyl-3-methylimidazolium ethylsulfate. J. Chem. Thermodyn. 2010, 42, 1298. (119) Heintz, Y. J.; Sehabiague, L.; Morsi, B. I.; Jones, K. L.; Luebke, D. R.; Pennline, H. W. Hydrogen sulfide and carbon dioxide removal from dry fuel gas streams using an ionic liquid as a physical solvent. Energy Fuels 2009, 23, 4822–4830.

3. Absorption of H2S Using ILs H2S is a highly toxic and corrosive gas. Thus, its concentration has to be reduced up to very low levels, therefore, for natural gas transportation through pipelines, or for commercial use, the H2S concentration should be lower than 6 mg/m3.25 Elimination of H2S from natural gas sources is commonly performed using amine-based processes.26 CO2 and H2S are frequently simultaneously present in many natural gas reservoirs. (112) Arshad, M. W. CO2 capture using ILs. Master’s Thesis, Center for Phase Equilibria and Separation Processes, Technical University of Denmark, Lyngby, Denmark, 2009.

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Shiflett et al.120 studied [bmim]MeSO4 IL for the absorption and separation of CO2 and H2S. The reported results showed a high H2S/CO2 selectivity for this IL (13.5) when a feed gas with a higher CO2 concentration is used, which is larger than that for the [bmim]PF6 IL (3.7) previously studied by the same authors.115 The main reason for the considerable solubility/ selectivity of H2S in many ILs is the development of strong interactions with the anions, whose strength is in the range of common hydrogen bonding, as reported by Pomelli et al.114 To the best of our knowledge, no additional results on the absorption of H2S using ILs have been reported in the literature; therefore, despite the promising results reported in the available studies, more thorough and systematic analysis, however, is needed to foresee the use of ILs for H2S removal.

Table 4. List of Henry’s Law Constant, H, of n-Alkanes in Selected ILs

4. Absorption of Hydrocarbons in ILs The solubility of hydrocarbons in natural gases, such as methane and ethane, should also be considered because the ultimate purpose of sweetening is to separate CO2 and H2S from the main components of natural gases.41 These gases have been studied together with CO2 for this purpose. Several studies have shown that the solubility of carbon dioxide in common ILs, such as [bmim][PF6] and [hmim][Tf2N], is much higher than those of nonpolar gases, such as N2, CH4, and C2H6, because of strong interactions between CO2 and the anion.121-124 Henry’s law constants for some relevant hydrocarbons are reported in Table 4, and these values are considerably larger than those reported in previous sections for CO2 and H2S. Nevertheless, solubilities of ethane are higher than for methane, with Henry’s law constant values for ethane in some ILs close to those reported in Table 2 for CO2 for certain ILs. Likewise, as the alkyl chain of the hydrocarbon increases, the solubility in a fixed IL also increases, as reported in Table 4, where Henry’s law constants for ethane, propane, and butane are compared. Therefore, a relevant quantity of ethane and longer hydrocarbons could be absorbed during sweetening if certain ILs were used, and thus, this should be considered during process design. Even though the solubility of hydrocarbons in many ILs is low, which is very favorable for gas sweetening purposes, this is not true for all of the ILs. For example, ILs containing long alkyl chains in their ions would lead to higher solubilities of hydrocarbons. Kou et al.125 studied methane absorption in [NR1R2R3R4][Tf2N] (ammonium-based cations with different alkyl chain lengths, Ri) IL, but for CH4 storage purposes, high solubilities were required. The authors reported a high solubility, up to 27 mol %, for this type of ILs, and thus, these

hydrocarbon

IL

T (K)

H (bar)

references

methane methane methane methane ethane

[bmim][PF6] [hmpy][Tf2N] [hmim][Tf2N] [bmim][CH3SO4] [bmim][PF6]

ethane ethane ethane ethane ethane ethane ethane propane propane propane propane propane butane butane butane butane butane

[bmim][Tf2N] [bmim][PF6] [hmpy][Tf2N] [bmim][BF4] [emim][Tf2N] [emim][dca] [emim][CF3SO3] [bmim][PF6] [bmim][BF4] [emim][Tf2N] [emim][dca] [emim][CF3SO3] [bmim][PF6] [bmim][BF4] [emim][Tf2N] [emim][dca] [emim][CF3SO3]

293.27 298 298 293.15 298 313.15 298 293.28 298 313.15 313.15 313.15 313.15 313.15 313.15 313.15 313.15 313.15 313.15 313.15 313.15 313.15 313.15

693 300 ( 30 329 ( 23 345 ( 5 336 ( 28 357 ( 46 97 ( 10 221 72 ( 2 416 ( 64 169 ( 14 675 ( 154 357 ( 46 190 ( 24 245 ( 22 92 ( 6 291 ( 34 190 ( 24 84 ( 4 112 ( 5 39 ( 1 168 ( 11 98 ( 3

121 122 122 123 62 124 62 121 122 124 124 124 124 124 124 124 124 124 124 124 124 124 124

compounds could be used for reversible storage of methane but obviously not for gas sweetening purposes. Moreover, Hert et al.126 reported that the simultaneous presence of CO2 and CH4 in the raw gas increases the CH4 solubility in [hmim][Tf2N] IL in comparison to absorption from pure methane, thus making gas separation more challenging. Kilaru et al.127 analyzed the solubility of CO2 and several hydrocarbons on a large group of ILs pertaining to very different families. Their results show the strong effect of interactions between equal types of molecules, that is to say, solute-solute and solventsolvent, thus showing how the claimed prevailing role of solute (CO2)-anion interactions on the gas solubility is a simplified viewpoint of the molecular origin of solubility. Therefore, understanding the molecular origins that govern the solubility of different gases in ILs is among the main challenges to obtain truly competitive absorption processes. 5. Drawbacks of the Use of ILs for Natural Gas Sweetening/ CO2 Capture The extremely low vapor pressure of most ILs and their possible use as task-specific fluids are the key features that render them useful in green chemistry. They do not contaminate the environment because of volatility, which is the common problem for organic solvents. The scarce knowledge about their toxicity, however, is the main barrier to fully recognize them as green solvents. Even though some studies claim that ILs with fluorine atoms and long alkyl chains exhibit high toxicity values, a thorough evaluation has not been carried out.128-130 [bmim][PF6] and [bmim][BF4] are also

(120) Shifflett, M. B.; Niehaus, A. M.; Yokozeki, A. Y. Separation of CO2 and H2S using room-temperature ionic liquid [bmim][MeSO4]. J. Chem. Eng. Data 2010, DOI: 10.1021/je1004005. (121) Jacquemin, J.; Husson, P.; Majer, V.; Gomes, M. F. C. Lowpressure solubilities and thermodynamics of solvation of eight gases in 1-butyl-3-methylimidazolium hexafluorophosphate. Fluid Phase Equilib. 2006, 240, 87–95. (122) Anderson, J. L.; Dixon, J. K.; Brennecke, J. F. Solubility of CO2, CH4, C2H6, C2H4, O2, and N2 in 1-hexyl-3-methylpyrinidium bis(trifluoromethylsulfonyl)imide: Comparison to other ILs. Acc. Chem. Res. 2007, 40, 1208–1216. (123) Kumelan, J.; Perez-Salado, A.; Tuma, D.; Maurer, G. Solubility of the single gases methane and xenon in the ionic liquid [bmim][CH3SO4]. J. Chem. Eng. Data 2007, 52, 2319–2324. (124) Camper, D.; Becker, C.; Koval, C.; Noble, R. Low pressure hydrocarbon solubility in room temperature ILs containing imidazolium rings interpreted using regular solution theory. Ind. Eng. Chem. Res. 2005, 44, 1928–1933. (125) Kou, Y.; Xiong, W.; Tao, G.; Liu, H.; Wang, T. Absorption and capture of methane into ionic liquid. J. Nat. Gas Chem. 2006, 15, 282–286.

(126) Hert, D.; Anderson, J. L.; Aki, S. N. V. K.; Brennecke, J. F. Enhancement of oxygen and methane solubility in 1-hexyl-3-methylimidazolium bis(trifluoromethylsulfonyl) imide using carbon dioxide. Chem. Commun. 2005, 2603–2605. (127) Kilaru, P.-K.; Conderim, R. A.; Scovazzo, P. Correlations of lowpressure carbon dioxide and hydrocarbon solubilities in imidazolium-, phosphonium-, and ammonium-based room-temperature ionic liquids. Part 1. Using surface tension. Ind. Eng. Chem. Res. 2008, 47, 900–909. (128) Jastorff, B.; St€ ormann, R.; Ranke, J.; M€ olter, K.; Stock, F.; Oberheitmann, B.; Hoffmann, W.; Hoffmann, J.; Nuchter, M.; Ondruschka, B.; Filser, J. How sustainable are ILs? Structure-activity relationships and biological testing as important elements for sustainability evaluation. Green Chem. 2003, 5, 136–142.

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known to decompose in the presence of water, leading to the formation of hydrofluoric and phosphoric acids.131 Therefore, although several studies have analyzed toxicity and biodegradability issues of the most commonly used families of ILs,132-134 it should be addressed with more detailed studies before employing these fluids in large-scale industrial applications, such as those considered in this review. Alternative methods should also be developed to clean ILs once they are contaminated because the cleaning process of ILs involves washing with water or volatile organic compounds, leading to another waste stream. Even though this could be avoided by some of the recently reported methods that use membrane separation and supercritical extraction technologies to remove contaminations from ILs, they require other aqueous waste streams for the regeneration process, resulting in extra expenses.45 Another important drawback of ILs is their high viscosities. For example, the viscosity of [bmim][BF4] (79.5 cP) is reported to be much higher than that of pure monoethanolamine (25 cP) and 30% aqueous MEA solution (2 cP).135 Some authors have shown how the absorption of acid gases in ILs is too slow;104 the rate of absorption is limited by the slow diffusion of the ILs, which is directly related to their high viscosity. Therefore, the high viscosity of the many ILs, including those NH2-functionalized, limits the rate of absorption, and equilibrium can take even 48 h at 303 K. This problem could be avoided by specifically choosing a convenient anion and cation to adjust the viscosity value in an acceptable range, although this approach would require obtaining accurate structureviscosity relationships that may lead to the proposal of anion/cation combinations that are cost-effective. It is also possible to minimize this limitation by mixing ILs with convenient organic compounds, which lead to other disadvantages regarding volatility issues that are common in organic solvents. One of the main problems arising from the use of amine-based processes for gas sweetening/CO2 absorption is corrosionderived issues; if ILs are considered as alternative absorbents, their corrosion properties should also be analyzed. Some studies have shown a certain corrosion of ILs toward some metals and alloys, especially at high temperatures.136-138 Moreover, certain impurities present in IL samples, such as halides, may increase the corrosive ability of these ILs toward metals and alloys.

It should be mentioned that certain ILs have also some inherent drawbacks: (i) Various ILs have been found to be combustible and require careful handling.139 A brief exposure (5-7 s) to a flame torch will ignite some ILs, and some of them can be completely consumed by combustion.139 (ii) The higher viscosity of some ILs compared to conventional solvents would increase the pumping and related operating costs, such as reduced mass-transfer rates and poor heat transfer. (iii) Thermal and chemical stability of ILs140 should be analyzed with caution, especially when long thermal stability is required141 to prevent degradation during the absorption/stripping cycles. (iv) Some ILs are highly hygroscopic,142 which dramatically changes their physicochemical properties,143,144 as well as potentially toxic to aquatic environments.145,146 This aquatic toxicity should not be ignored, because it was reported to be equal to or greater than that of many conventional solvents. (v) The high prices of most ILs hinder their extension to large-scale applications, such as those considered in this work.

(129) MacFarlane, D. R. ILs symposium. Aust. J. Chem. 2004, 57, 111–112. (130) Sheldon, R. A.; Lau, R. M.; Sorgedrager, F.; van Rantwijk, K.; Seddon, R. Biocatalysis in ILs. Green Chem. 2002, 4, 147–151. (131) Cammarata, L.; Kazarian, S. G.; Salter, P. A.; Welton, T. Molecular states of water in room temperature ILs. Phys. Chem. Chem. Phys. 2001, 23, 5192–5200. (132) Romero, A.; Santos, A.; Rodrı´ guez, A. Toxicity and biodegradability of imidazolium ILs. J. Hazard. Mater. 2008, 151, 268–273. (133) Pham, T. P.; Cho, W.; Yun, Y. S. Environmental fate and toxicity of ILs: A review. Water Res. 2010, 44, 352–372. (134) Coleman, D.; Gathergood, N. Biodegradation studies of ILs. Chem Commun. 2010, 39, 600–637. (135) Sanchez, L. M. G.; Meindersma, G. W.; de Haan, A. B. Solvent properties of functionalized ILs for CO2 adsorption. Chem. Eng. Res. Des. 2007, 85, 31–39. (136) Uerdingen, M.; Treber, C.; Schmitt, G.; Werner, C. Corrosion behaviour of ILs. Green. Chem. 2005, 7, 321. (137) Perissi, I.; Bardi, U.; Caporali, S.; Lavachi, A. High temperature corrosion properties of ILs. Corros. Sci. 2006, 48, 2349–2362. (138) Tolstoguzov, A. B.; Bardi, U.; Chenakin, S. P. Study of the corrosion of metal alloys interacting with an ionic liquid. Bull. Russ. Acad. Sci. Phys. 2008, 72, 605–608. (139) Smiglak, M.; Reichert, W. M.; Holbrey, J. D.; Wilkes, J. S.; Sun, L.; Thrasher, J. S.; Kirichenko, K.; Singh, S.; Katritzky, A. R.; Rogers, R. D. Combustible ILs by design: Is laboratory safety another ionic liquid myth? Chem. Commun. 2006, 2554–2556.

(140) Huddleston, J. G.; Visser, A. E.; Reichert, W.; Willauer, H. D.; Broker, G. A.; Rogers, R. D. Characterization and comparison of hydrophilic and hydrophobic room temperature ILs incorporating the imidazolium cation. Green Chem. 2001, 3, 156–164. (141) Kosmulski, M.; Gustafsson, J.; Rosenholm, J. B. Thermal stability of low temperature ILs revisited. Thermochim. Acta 2004, 412, 47–53. (142) Cuadrado, S.; Domı´ nguez, M.; Garcı´ a, S.; Segade, L.; Franjo, C.; Cabeza, O. Experimental measurement of the hygroscopic grade on eight imidazolium based ILs. Fluid Phase Equilib. 2009, 278, 36–40. (143) Liu, W.; Cheng, L.; Zhang, Y.; Wang, H.; Yu, M. The physical properties of aqueous solution of room-temperature ILs based on imidazolium: Database and evaluation. J. Mol. Liq. 2008, 140, 68–72. (144) Vila, J.; Gines, P.; Rilo, E.; Cabeza, O.; Varela, L. M. Great increase of the electrical conductivity of ILs in aqueous solutions. Fluid Phase Equilib. 2006, 247, 32–39. (145) Latala, A.; Nedzi, M.; Stepnowski, P. Toxicity of imidazolium and pyridinium based ILs towards algae. Bacillaria paxillifer (a microphytobenthic diatom) and Geitlerinema amphibium (a microphytobenthic blue green alga). Green Chem. 2009, 11, 1371–1376. (146) Pretti, C.; Chiappe, C.; Baldetti, I.; Brunini, S.; Monni, G.; Intorre, L. Acute toxicity of ILs for three freshwater organisms: Pseudokirchneriella subcapitata, Daphnia magna and Danio rerio. Ecotoxicol. Environ. Saf. 2009, 72, 1170–1176. (147) Wilms, D.; Klos, J.; Kilbinger, A. F. M.; Lowe, A.; Frey, H. Ionic liquids on demand in continuous flow. Org. Process Res. Dev. 2009, 13, 961–964.

6. Conclusions ILs offer promising advantages in natural gas sweetening and flue gas CO2 absorption processes because of their unique characteristics. Their negligible vapor pressure is one of the main factors that make this class of materials superior to the conventional organic solvents. Solubility of CO2/H2S in common ILs is too low to compete with the state-of-the-art amine process. Chemical modification of ILs to prepare TSILs has produced promising results, attracting more research groups to this field. ILs have received much attention to replace volatile organic solvents. Even though the toxicity of the ILs is still in debate because of their incomplete toxicological data, the interest in ILs has further grown. It was clearly shown that selection of ILs consisting of cation and anion groups with appropriate functional groups could improve the interactions of ILs with CO2/H2S, making them as effective as the alkanolamine molecules. Their high thermal stability is also very beneficial, considering that high-pressure conditions are required while the natural gas is purified. Aspects relating to the use of ILs in such harsh conditions are, however, still not clear because very few projects have focused on the

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industrial-scale applications of ILs in high-pressure gas-processing systems, as well as their toxicological properties.

Cost-performance analysis should also be considered prior to using these chemicals in pilot-scale applications. Although the present costs of the ILs do not make them favorable for commercial applications, current synthetic studies could overcome this barrier in the near future, making ILs economically more viable.41,147,148

(148) Davis, J. H.; Gordon, C. M.; Hilgers, C.; Wassercheid, P. Syntheis and purification of ionic liquids. In Ionic Liquids in Synthesis; Wasserscheid, P., Welton, T., Eds.; Wiley-VCH: Weinheim, Germany, 2003.

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