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Solubility of CO2 in 1-(2-hydroxyethyl)-3-methylimidazolium ionic liquids with different anions
J. Chem. Thermodynamics 42 (2010) 787–791
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Solubility of CO2 in 1-(2-hydroxyethyl)-3-methylimidazolium ionic liquids with different anions Amir Hossein Jalili a,*, Ali Mehdizadeh a, Mohammad Shokouhi a, Hossein Sakhaeinia b, Vahid Taghikhani c a
Gas Science Department, Gas Research Division, Research Institute of Petroleum Industry (RIPI), National Iranian Oil Company (NIOC), P.O. Box 14665-137, West Blvd., Azadi Sport Complex, Tehran, Iran b Department of Chemical Engineering, Faculty of Engineering, Islamic Azad University, Semnan Branch, Semnan, Iran c Department of Chemical and Petroleum Engineering, Sharif University of Technology, Azadi Ave., Tehran, Iran
a r t i c l e
i n f o
Article history: Received 4 January 2010 Accepted 1 February 2010 Available online 6 February 2010 Keywords: Gas solubility Henry’s law constant Carbon dioxide Ionic liquid Gibbs free energy of solution
a b s t r a c t The solubility of carbon dioxide in a series of 1-(2-hydroxyethyl)-3-methylimidazolium ([hemim]+) based ionic liquids (ILs) with different anions, viz. hexafluorophosphate ([PF6]), trifluoromethanesulfonate ([OTf]), and bis-(trifluoromethyl)sulfonylimide ([Tf2N]) at temperatures ranging from 303.15 K to 353.15 K and pressures up to 1.3 MPa were determined. The solubility data were correlated using the Krichevsky–Kasarnovsky equation and Henry’s law constants were obtained at different temperatures. Using the solubility data, the partial molar thermodynamic functions of solution such as Gibbs free energy, enthalpy, and entropy were calculated. Comparison showed that the solubility of CO2 in the ILs studied follows the same behaviour as the corresponding conventional 1-ethyl-3-methylimidazolium ([emim]+) based ILs with the same anions, i.e. [hemim][NTf2] > [hemim][OTf] > [hemim][PF6] > [hemim][BF4]. Ó 2010 Elsevier Ltd. All rights reserved.
1. Introduction Natural gas, as a source of clean fuel and as a starting material for synthesis of various petrochemical products, is usually contaminated with a variety of undesirable materials, especially CO2 and H2S acid gases. The presence of CO2 lowers the heating value of natural gas while H2S gives it high toxicity and corrosiveness as a clean fuel source. Alkanolamines, especially monoethanolamine, diethanolamine, and methyldiethanolamine are the main constituents of aqueous solutions used in industrial natural gas treating plants [1]. There are some disadvantages in commercial use of these alkanolamine solutions, including loss of alkanolamine and transfer of water into the gas stream during desorption stage and degradation of alkanolamines to form corrosive by products, which make the scrubbing process economically expensive [2]. Room temperature ionic liquids (RTILs) are molten salts that are liquid over a wide temperature range including ambient [3]. Due to the existence of Coulombic attraction between the ions of ILs, they exhibit negligibly small vapour pressure, meaning that ionic liquids are essentially non-volatile and could facilitate the sequestration of gases without loss of the capture agent into the gas stream. They also have high thermal and electrochemical stability. Nowadays, one of the active research areas is to explore task specific * Corresponding author. Tel.: +98 21 48252466; fax: +98 21 44739716. E-mail addresses: [email protected], [email protected] (A.H. Jalili). 0021-9614/$ - see front matter Ó 2010 Elsevier Ltd. All rights reserved. doi:10.1016/j.jct.2010.02.002
ionic liquids to replace conventional alkanolamine solutions for removal of acid gases (CO2 and H2S) in gas sweetening processes [4]. Another research area is concerned with fixation and sequestration of CO2, which is regarded as the most important greenhouse gas, by employing ionic liquids [5]. An important feature in the evaluation of ILs for potential use in industrial natural gas treating processes is the knowledge of the solubility of gases at various temperatures and pressures. In the past few years, a growing number of measurements reporting solubility of CO2 in various conventional ILs have become available. Blanchard et al. experimentally explored the high-pressure phase behaviour of CO2 with six imidazolium-based ILs including 1-ethyl-3-methylimidazolium ethylsulfate ([emim][EtSO4]) by systematically changing the anionic and cationic components of the IL [6]. Camper et al. measured the low-pressure solubility of CO2 and C2H4 in different 1-ethyl-3-methylimidazolium-based ILs with different anions [PF6], [Tf2N] , [OTf], and dicyanamide ([dca]) as well as 1-butyl-3-methylimidazolium hexafluorophosphate ([bmim][PF6]) at T = 303 K, and then showed that the regular solution theory can be used to model the solubility results at low pressures [7]. In a subsequent work, they concluded that CO2/IL complexations are not the sole controlling factor in relative CO2 solubility and the energies of vapourisation and molar volumes appear to be determining factors in relative CO2 solubility in ILs [8]. Later they reported the solubility in the form of Henry’s law constants and diffusion of CO2 at five discrete temperatures from
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303 K to 343 K and those of ethylene, ethane, propane, and propylene in [emim][Tf2N] at 303 K [9]. Shariati and Peters have reported the experimental high-pressure phase behaviour of the binary system CO2/[emim][PF6] within temperature and pressure ranges of 308.14 K to 366.03 K and 1.49 MPa to 97.10 MPa, respectively [10]. In a subsequent work, they reported the experimental highpressure phase behaviour of CO2/ILs at pressures up to 97 MPa and T = 303.15 K [11]. The ILs studied were comprised of imidazolium-based cations from [emim]+ to [bmim]+ to [hmim]+ and [PF6] anion. In their investigation, they showed that the solubility of CO2 in the ILs increased by increasing the alkyl group chain length. Later Schilderman et al. reported experimental values for the solubility of CO2 in [emim][Tf2N] within temperature and pressure ranges of 310 K to 450 K and 0 MPa to 15 MPa respectively, and concluded that the solubility of CO2 in [emim][Tf2N], due to the presence of fluoroalkyl groups in the anion, is much higher than that of [emim][PF6] [12]. Kim et al. have reported the solubility of CO2 in some ILs including [emim][BF4] and [emim][Tf2N] at T = 298.15 K and pressures up to 1 MPa [13]. They used the data to obtain the parameters of a group contribution non-random lattice-fluid equation of state. Shin et al. have reported the solubility values of CO2 in a series of bis(trifluoromethyl) sulfonylimide ILs with 1-alkyl-3methylimidazolium ([Cn-mim]+, n = 2, 4, 6, 8) cations at pressures up to 45 MPa and temperatures from 298.15 K to 343.15 K and used the Peng–Robinson equation of state to correlate the data [14]. The solubility of CO2 in [bmim][PF6] and [emim][BF4] for temperatures ranging from 303.2 K to 343.2 K and pressures up to 5 MPa [15] and in 1-ethyl-3-methylimidazolium 2-(2-methoxyethoxy) ethylsulfate at pressures up to 6.7 MPa and temperatures from 303.2 K to 343.2 K [16] using a thermogravimetric microbalance apparatus is reported by Soriano and his coworkers. Subsequently they reported the solubility of CO2 in [emim][OTf] for temperatures ranging from 303.2 K to 343.2 K and pressures up to 5.9 MPa [17]. The high-pressure phase behaviour of carbon dioxide in 1-methyl-3-pentylimidazolium bis(trifluoromethylsulfonyl)imide ([pmim] [Tf2N]) as well as in [emim][Tf2N] at temperatures up to 363 K and pressures up to 50 MPa is investigated by Carvalho et al. [18]. Finally, Shokouhi et al. have reported the solubility and diffusion of H2S and CO2 in 1-(2-hydroxyethyl)imidazolium tetrafluoroborate ([hemim][BF4]) at temperatures ranging from 303.15 K to 353.15 K and pressures up to 1.1 MPa [19]. They have used the data obtained to estimate Henry’s law constants and diffusion coefficients at different temperatures. In this work, the equilibrium solubility of CO2 in the ionic liquids [hemim][PF6], [hemim][OTf], and [hemim][Tf2N] at six temperatures from 303.15 K to 353.15 K was measured. The values of solubility determined were used to estimate Henry’s law constants and partial molar thermodynamic functions of solution of CO2 at different temperatures. A correlation equation for the solubility obtained and partial molar volume at infinite dilution data with temperature is presented here. The results obtained are compared with [hemim][BF4] found in the previous work [19] and other [emim] based ILs, which are reported in the literature.
2. Experimental 2.1. Materials Carbon dioxide (c.p. grade, mass fraction purity >0.9995) was obtained from the Roham Gas Company. The 1-(2-hydroxyethyl)3-methylimidazolium chloride ([hemim][Cl]) was synthesized by modification of a known literature procedure [20]. A mixture of 1-methylimidazole (5 g, 61 mmol) and 2-chloroethanol (4.9 g, 61 mmol) was placed in a TFM vessel of MICCROSYNTH laboratory microwave oven (Milestone srl, Italy) and subjected to microwave
irradiation at T = 383 K (200 W, 10 bars) for half an hour. After cooling to ambient temperature, the reaction mixture was washed with 20 cm3 of ethyl acetate to remove non-reacted starting materials. The solid product obtained was then heated at T = 343 K for 7 h at reduced pressure to remove the residual water and obtain 6.6 g of the ionic liquid. The yield was 66.6%. 1H NMR (500 MHz, D2O, T = 209 K: d 106 = 3.87 (3H, S, NCH3), 3.89 (2H, t, NCH2CH2OH), 4.28 (2H, t, NCH2CH2OH), 7.42 (1H, d, H-4), 7.47 (1H, d, H-5), 8.71 (1H, S, H-2). The ionic liquids [hemim][PF6], [hemim][OTf], [hemim][Tf2N], were prepared according to the synthesis method described by Yeon et al. [21] and Dubreuil and Bazureau [22], which are briefly described below. 2.1.1. Synthesis of [hemim][PF6] 1-(2-Hydroxyethyl)-3-methylimidazolium chloride (5.0 g, 31 mmol) and potassium hexaflurophosphate (5.66 g, 31 mmol) in deionised water (10 cm3) and acetonitrile (10 cm3) were mixed and the resulting two-phase mixture was magnetically stirred at room temperature for 24 h. After completion of the reaction, the solvents were removed at reduced pressure and the ionic liquid was removed from KCl by extraction in dry methanol (17 cm3) and chloroform (3 cm3). The suspension was then filtered and the filtrate was washed with chloroform (3 7 cm3) and diethyl ether (3 7 cm3) to remove impurities and non-reacted starting materials. The solvents were removed at reduced pressure. The product obtained was then heated at T = 343 K for 7 h at reduced pressure to remove the residual water and obtain 3.56 g of the ionic liquid. The yield was 42.6%. 1H NMR (500 MHz, D2O, T = 343 K): d 106 = 3.87 (3H, S, NCH3), 3.90 (2H, t, NCH2CH2OH), 4.28 (2H, t, NCH2CH2OH), 7.41 (1H, d, H-4), 7.46(1H, d, H-5), 8.68 (1H, S, H-2). 2.1.2. Synthesis of [hemim][OTf] 1-(2-Hydroxyethyl)-3-methylimidazolium chloride (5.0 g, 31 mmol) and ammonium trifluoromethanesulfonate (5.14 g, 31 mmol) in deionised water (10 cm3) and acetonitrile (10 cm3) were mixed and the resulting two-phase mixture was magnetically stirred at room temperature for 24 h. After completion of the reaction, the solvents were removed at reduced pressure and the ionic liquid was removed from NH4Cl by extraction in dry methanol (17 cm3) and chloroform (3 cm3). The suspension was then filtered and the filtrate was washed with chloroform (3 7 cm3) and diethyl ether (3 7 mL) to remove impurities and non-reacted starting materials. The solvents were removed at reduced pressure. The product obtained was then heated at T = 343 K for 7 h at reduced pressure to remove the residual water and obtain 5.21 g of the ionic liquid. The yield was 61.3%. 1H NMR (300 MHz, D2O, 25 °C): d 106 = 3.84 (3H, S, NCH3), 3.87 (2H, t, NCH2CH2OH), 4.25 (2H, t, NCH2CH2OH), 7.39 (1H, d, H-4), 7.44 (1H, d, H-5), 8.68 (1H, S, H-2). 2.1.3. Synthesis of [hemim][Tf2N] 1-(2-Hydroxyethyl)-3-methylimidazolium chloride (5.0 g, 31 mmol) and lithium bis (trifluoromethanesulfonimide) (8.83 g, 31 mmol) in deionised water (10 cm3) and acetonitrile (10 cm3) were mixed and the resulting two-phase mixture was magnetically stirred at room temperature for 24 h. After completion of the reaction, the solvents were removed at reduced pressure and the ionic liquid was removed from LiCl by extraction in dry methanol (17 cm3) and chloroform (3 cm3). The suspension was then filtered and the filtrate was washed with chloroform (3 7 cm3) and diethyl ether (3 7 cm3) to remove impurities and non-reacted starting materials. The solvents were removed at reduced pressure. The product obtained was then heated at T = 343 K for 7 h at reduced pressure to remove the residual water and obtain 7.84 g of the ionic liquid. The yield was 62.6%. 1H NMR (500 MHz, D2O,
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25 °C): d 106 = 3.87 (3H, S, NCH3), 3.90 (2H, t, NCH2CH2OH), 4.28 (2H, t, NCH2CH2OH), 7.42 (1H, d, H-4), 7.47 (1H, d, H-5), 8.69 (1H, S, H-2).
of the method of measurement have been checked in our previous work [23]. The Krichevsky–Kasarnovsky equation, [26] modified for solvents of negligible vapour pressure [23,24], was used to model the experimental results obtained in this work:
ln
2.2. Apparatus and procedure The experimental method for the measurement of gas solubility has previously been presented in detail [23,24] and only a short description will be provided here. In this technique, known quantities of gaseous solute and degassed solvent are contacted at a constant temperature inside an equilibrium cell of known volume. On reaching thermodynamic equilibrium, the pressure above the liquid solution is constant and directly related to the solubility of the gas in the liquid. The quantity of solute present in the liquid solution, nlsolute , is calculated by the difference between two (P, V, T) measurements: first when the gas is introduced from the gas container of known volume into the equilibrium cell containing the IL and secondly after reaching thermodynamic equilibrium (i.e. autoclave pressure remains fixed and no longer changes with time):
nlsolute ¼ ntotal ngsolute ;
V gc Pi P f ; RT gc Z i Z f
ð2Þ
where V gc denotes the volume of the gas container, Z i and Z f are the compressibility factors corresponding to the initial and final pressures Pi and Pf , respectively, in the gas container before and after transferring the gas, and T gc is the temperature of the gas container. The most accurate (P, V, T) data presented by NIST for pure compounds were used to calculate compressibility factors [25]. The ngsolute in equation (1) is the number of moles of gas solute left in the gas phase and was determined from the following equation:
ngsolute ¼
V gP ; ZRT
ð3Þ
where V g is the volume of the gas phase above IL phase, T is the equilibrium temperature of the cell, and Z is the compressibility factor of gas solute at P and T. In all experiments, the IL was dried in vacuo (below 1.0 kPa) before solubility measurements for about 48 h at a temperature of 343 K to remove trace amounts of moisture and volatile impurities. Using a Mettler model DL-37 Karl-Fischer volumetric titrator, the water content of ILs was found to be below (100 ± 10)106. The temperature of the double-wall equilibrium cell, which was connected to a water recirculation bath (PMT Tamson model T 2500), was measured with a Lutron model TM-917 digital thermometer with a 0.01 K resolution using a Pt-100 sensor inserted into the cell. The equilibrium cell pressure was measured using a KELLER model PA-33X pressure transmitter sensor in the range of 0 MPa to 2 MPa, which was accurate to within 0.01% of full scale and that of the gas container was measured using a Druck model PTX 1400 pressure transmitter sensor in the range of 0 MPa to 4 MPa, which was accurate to within 0.1% of full scale. The calibration of pressure sensors was carried out against a dead-weight gauge. 3. Results and discussion The results of the measurement of the solubility of carbon dioxide in the ionic liquids [hemim][PF6], [hemim][OTf], and [hemim][Tf2N], at temperatures of (303.15, 313.15, 323.15, 333.15, 343.15, 353.15) K and pressures up to about 1.3 MPa are summarized in tables 1 to 3, respectively. The reliability and accuracy
ð4Þ
where f2 is the fugacity of pure solute (carbon dioxide) in the gas phase, x2 is the mole fraction of solute in the solvent 1, H21 is Henry’s law constant of gas solute 2 1 in solvent 1 at the pressure, P, v 2 is the partial molar volume of gas solute 2 at TABLE 1 Solubility of carbon dioxide gas in [hemim][PF6]: x2, CO2 mole fraction; P, equilibrium pressure; T, equilibrium temperature. P/kPa
x2
T = 303.15 K 133.3 0.013 ± 0.006 310.6 0.035 ± 0.008 579.4 0.064 ± 0.009 667.0 0.075 ± 0.008 745.2 0.084 ± 0.009 906.3 0.104 ± 0.011
ð1Þ
where ntotal is the total number of CO2 moles injected from the gas container into the autoclave and calculated using the following equation:
ntotal ¼
f2 v1P ¼ ln H21 þ 2 x2 RT
152.2 355.9 502.3 663.2 762.1 853.2 935.8
T = 333.15 K 0.009 ± 0.006 0.026 ± 0.008 0.037 ± 0.007 0.049 ± 0.009 0.059 ± 0.009 0.066 ± 0.009 0.070 ± 0.010
P/kPa
x2
T = 313.15 K 139.0 0.012 ± 0.006 324.5 0.033 ± 0.008 459.5 0.045 ± 0.007 606.3 0.060 ± 0.009 699.9 0.068 ± 0.008 783.7 0.076 ± 0.009 854.2 0.085 ± 0.010 1001 0.101 ± 0.011 160.5 370.4 523.4 687.2 794.1 889.3 973.0
T = 343.15 K 0.006 ± 0.006 0.023 ± 0.008 0.033 ± 0.007 0.048 ± 0.009 0.054 ± 0.009 0.061 ± 0.009 0.066 ± 0.010
P/kPa
x2
T = 323.15 K 145.2 0.010 ± 0.006 341.3 0.028 ± 0.008 482.4 0.040 ± 0.007 634.2 0.055 ± 0.009 733.1 0.062 ± 0.009 819.2 0.070 ± 0.009 897.0 0.076 ± 0.010 986.2 0.086 ± 0.011 169.0 386.0 542.0 717.0 827.1 926.2 1008 1127
T = 353.15 K 0.009 ± 0.002 0.020 ± 0.004 0.032 ± 0.006 0.042 ± 0.008 0.049 ± 0.009 0.055 ± 0.010 0.064 ± 0.011 0.070 ± 0.013
TABLE 2 Solubility of carbon dioxide gas in [hemim][OTf]: x2, CO2 mole fraction; P, equilibrium pressure; T, equilibrium temperature. P/kPa
x2
T = 303.15 K 101.2 0.017 ± 0.005 251.6 0.044 ± 0.008 432.0 0.073 ± 0.010 652.8 0.107 ± 0.013 895.6 0.144 ± 0.016 1157 0.188 ± 0.019 115.2 286.8 492.0 745.4 1021
T = 333.15 K 0.014 ± 0.005 0.036 ± 0.008 0.061 ± 0.011 0.089 ± 0.014 0.123 ± 0.016
P/kPa
x2
T = 313.15 K 105.6 0.016 ± 0.005 263.8 0.040 ± 0.008 452.0 0.068 ± 0.010 684.8 0.100 ± 0.014 939.0 0.135 ± 0.016 1216 0.176 ± 0.019 119.2 297.2 510.8 772.8 1063
T = 343.15 K 0.013 ± 0.005 0.034 ± 0.008 0.058 ± 0.011 0.087 ± 0.014 0.118 ± 0.017
P/kPa
x2
T =323.15 K 110.4 0.015 ± 0.005 275.2 0.038 ± 0.008 472.6 0.064 ± 0.011 715.6 0.094 ± 0.014 981.6 0.128 ± 0.016 1280 0.162 ± 0.019 123.6 307.2 558.4 801.2 1105
T = 353.15 K 0.013 ± 0.005 0.034 ± 0.008 0.058 ± 0.011 0.083 ± 0.014 0.112 ± 0.017
TABLE 3 Solubility of carbon dioxide gas in [hemim][Tf2N]: x2, CO2 mole fraction; P, equilibrium pressure; T, equilibrium temperature. P/kPa
x2
T = 303.15 K 97.3 0.022 ± 0.005 243.7 0.054 ± 0.008 443.6 0.093 ± 0.011 564.1 0.119 ± 0.009 654.0 0.139 ± 0.009 795.4 0.164 ± 0.011 954.1 0.191 ± 0.013 T = 333.15 K 112.5 0.018 ± 0.005 281.9 0.043 ± 0.008 514.9 0.075 ± 0.011 656.8 0.096 ± 0.009 764.2 0.111 ± 0.009 930.4 0.132 ± 0.012 1114.3 0.158 ± 0.014
P/kPa
x2
T = 313.15 K 102.7 0.021 ± 0.006 256.7 0.050 ± 0.010 468.4 0.086 ± 0.014 597.4 0.109 ± 0.012 693.5 0.127 ± 0.012 843.3 0.151 ± 0.016 1011 0.177 ± 0.018 T = 343.15 K 117.1 0.017 ± 0.005 293.3 0.041 ± 0.008 535.6 0.072 ± 0.011 683.7 0.092 ± 0.009 796.8 0.106 ± 0.010 968.8 0.127 ± 0.012
P/kPa
x2
T = 323.15 K 107.5 0.019 ± 0.005 269.3 0.046 ± 0.008 491.8 0.080 ± 0.011 627.5 0.102 ± 0.009 730.0 0.118 ± 0.009 888.1 0.141 ± 0.012 1064 0.166 ± 0.014 T = 353.15 K 121.7 0.016 ± 0.005 305.0 0.039 ± 0.008 556.7 0.069 ± 0.011 710.9 0.088 ± 0.009 828.7 0.102 ± 0.010 1007.5 0.122 ± 0.012
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TABLE 4 1 Thermodynamic properties of CO2 in ionic liquids: T, temperature; H21, Henry’s law constant; v 1 2 , partial molar volume at infinite dilution; DsolG , Gibbs free energy of solution; DsolH1, enthalpy of solution; DsolS1, entropy of solution. T/K
H21/MPa
3 1 m1 2 /(cm mol )
303.15 313.15 323.15 333.15 343.15 353.15
8.94 ± 0.08 10.1 ± 0.09 11.7 ± 0.08 13.2 ± 0.05 14.8 ± 0.03 16.5 ± 0.02
118 ± 4 109 ± 4 99.3 ± 5 94.1 ± 3 85.5 ± 5 79.2 ± 5
303.15 313.15 323.15 333.15 343.15 353.15
6.14 ± 0.01 6.89 ± 0.01 7.74 ± 0.05 8.28 ± 0.04 8.93 ± 0.07 9.73 ± 0.06
303.15 313.15 323.15 333.15 343.15 353.15
4.42 ± 0.01 5.11 ± 0.02 5.75 ± 0.05 6.50 ± 0.04 7.02 ± 0.07 7.69 ± 0.01
DsolG1/(kJ mol1)
DsolH1/(kJ mol1)
DsolS1/(J mol1 K1)
[hemim][PF6] 11.3 12.0 12.8 13.5 14.3 15.0
11.07 11.03 10.99 10.96 10.93 10.90
73.87 73.62 73.64 73.51 73.40 73.30
121 ± 3 112 ± 3 102 ± 5 96.9 ± 3 88.4 ± 5 82.2 ± 5
[hemim][OTf] 10.4 11.0 11.7 12.2 12.8 13.4
9.35 8.75 8.19 7.66 7.16 6.70
65.1 63.1 61.5 59.7 58.2 57.0
184 ± 2 173 ± 2 161 ± 2 148 ± 2 133 ± 3 118 ± 3
[hemim][Tf2N] 9.55 10.2 10.9 11.6 12.1 12.8
11.8 10.9 10.1 9.29 8.55 7.85
70.5 67.6 64.9 62.6 60.3 58.3
TABLE 5 Numerical values of the parameters of equations (5) and (6). IL
A0
A1
A2
B0/(cm3 mol1)
B1/(cm3 mol1 K1)
B2/(cm3 mol1 K2)
[hemim][PF6] [hemim][OTf] [hemim][Tf2N]
6.33876 1.80997 0.622176
1186.22 1128.77 1944.24
22 000.8 341 511 510 054
724.52 724.81 1.0193
3.0489 3.0424 2.2464
0.00346 0.00346 0.00542
infinite dilution, R is the universal gas constant, and T is the absolute temperature. The fugacity of carbon dioxide was calculated using the most accurate corresponding states used by NIST for pure compounds [25]. Values of ln H21 at each temperature T were obtained from the intercept of Krichevsky–Kasarnovsky plots (that is plots of lnðf2 =x2 Þ vs P) at the specified temperature. The Henry’s law constants are given in table 4 for the solubility of CO2 in the ILs studied in this work together with their standard deviations. The Henry’s law constants were fit by the equation, 2 X i¼0
30
Ai ðT=KÞ
35
i
:
ð5Þ
The parameters Ai of equation (5) obtained for CO2 are summarized in table 5. The Henry’s law constants are compared with each other in figure 1 as a function of temperature. It can be observed that the solubility of carbon dioxide in the ILs decreases by increasing the temperature. The curves also indicate that the solubility of CO2 in the ILs studied in this work is typical of that of physical solvents [22], therefore obeying the Henry’s law. It can be seen that, the solubility of CO2 is the highest in [hemim][Tf2N] and the lowest in [hemim][PF6]. Also variation with temperature of the Henry’s law constant for the solubility of CO2 in [hemim][BF4] investigated in previous work [19] as well as in the corresponding conventional ILs, [emim][BF4], [emim][OTf], and [emim][Tf2N] are shown in figure 1. It can be observed that the solubility of CO2 in [hemim][BF4] is lower than the ILs studied in this work. This can be explained on the basis of the fact that ILs containing anions with fluoroalkyl groups, i.e. [Tf2N] and [OTf], have high affinities for CO2 than those containing [PF6] and [BF4] anions. Schilderamn et al. [12] and Zhang et al. [27] have mentioned same conclusion by experimental and theoretical (using COSMO-RS method) investigations, respectively. Inspection of figure 1 reveals that the solubility of CO2 in both [hemim][Tf2N] and [emim][Tf2N] is comparable. This observation shows that in this case, the anion, which contains two –CF3 groups, is the predominant species in determining CO2 solubility and the presence of a hydroxyl group, though it is capable of hydrogen bond formation, plays a very negligible role in this respect. However, in the case of other ILs, one can see that the effect of the anion gradually ceases as the number of –CF3 groups decreases from 1 for [hemim][OTf] to zero for [hemim][PF6] and [hemim][[BF4] and thus, the solubility of CO2 decreases too. The higher solubility of CO2 in [hemim] ILs relative to the corresponding [emim]s containing the same anion may be ascribed to the presence of hydroxyl group, which participates in hydrogen bond formation with CO2 solute molecules. We are studying these observations, from molecular point of view, for the interaction and configuration of CO2 and H2S molecules in functionalized ILs and their effect on the bulk liquid structure, by using molecular dynamics (MD) simulation and will present the results in the near future.
H 21 / MPa
lnðH21 ðTÞÞ ¼
40
25
20 15 10
5 0 300
310
320
330
340
350
360
T /K FIGURE 1. Plot of Henry’s law constants as a function of temperature for the solubility of CO2 in ILs: N, [hemim][Tf2N]; d, [hemim][BF4] (reference [19]); j, [hemim][OTf]; , [hemim][PF6]; 4, [emim][Tf2N] (reference [18]); s, [emim][BF4] (reference [15]); h, [emim][OTf] (reference [17]); +, [emim][OTf] (reference [7]).
The slope of Krichevsky–Kasarnovsky plots at each temperature, T, yields v 1 2 at the specified temperature. The obtained experimental data of partial molar volumes 1 of gas solute, CO2, at infinite dilution, v 2 , are presented in table 4. They were correlated with temperature using the following simple quadratic formula
A.H. Jalili et al. / J. Chem. Thermodynamics 42 (2010) 787–791
v 12 ðTÞ ¼
2 X
Bi T i ;
ð6Þ
791
5. Supporting information available
i¼0
3 1 where v 1 ) and T is in K. The obtained parameters Bi of equation (6) 2 is in (cm mol are summarized in table 5. It can be observed from table 4 that the v 1 2 values increase with temperature for CO2 in [hemim][PF6] and [hemim][OTf] but they decrease for CO2 in [hemim][Tf2N]. It can be shown that the Gibbs free energy of solution, corresponding to the change in Gibbs free energy when the solute is transferred at constant temperature from the pure perfect gas at the standard pressure to the standard state of infinite dilution of the solute in the solvent, is given by [28]:
Dsol G1 ¼ RT ln
Kh P
0
;
This information is available free of charge via the Internet at http://pubs.acs.org/. Acknowledgements We are thankful to the research council of the Research Institute of Petroleum Industry (RIPI) and also to the Research and Development of the National Iranian Oil Company (NIOC) for their support of this work.
ð8Þ
References where P0 is the standard state pressure. The partial molar differences in enthalpy and entropy between the two states can be obtained by calculating the corresponding partial derivatives of the Gibbs free energy with respect to temperature
Dsol H1 ¼ T 2
Dsol S1 ¼
@ Dsol G1 @ Kh ¼ RT 2 ; ln T @T @T P0
ðDsol H1 Dsol G1 Þ : T
ð9Þ
ð10Þ
The pressure range considered in this work is not too high to cause Henry’s law constant to be a strong function of pressure and Henry’s law is weakly dependent on pressure under the specified conditions [23,24]. Therefore, it does not give rise to large errors if one ignores this pressure dependency. By means of this approximation and using equations (8)–(10), we estimated the thermodynamic functions of solution at infinite dilution for CO2 in the IL. The values for the Gibbs free energy, enthalpy, and entropy of solution are given in table 4 at six discrete temperatures between 303.15 K and 353.15 K. As it can be observed, the Dsol G1 values are positive and increase with temperature in a similar manner for the solubility of CO2 in ILs. The Dsol H1 values and Dsol S1 values are negative. The variation with temperature of the Dsol H1 values and Dsol S1 values are positive for CO2 in the ILs studied in this work and they increase with temperature. In the case of [hemim][PF6] the rate of variation is very slow while for [hemim][Tf2N] the rate of variation is the highest.
4. Conclusions New experimental results for the solubility, partial molar volume at infinite dilution, and thermodynamic functions of solution of carbon dioxide gas in functionalized ionic liquids 1-(2-hydroxyethyl)-3-imidazolium cations with different anions, not previously reported in the literature, have been measured and presented in this work. The solubility of carbon dioxide in [hemim][Tf2N] is the highest and in [hemim][BF4] is the lowest. The solubility of CO2 in the ILs studied in this work is of a physical nature. The solubility of carbon dioxide in [hemim] ILs is greater than corresponding conventional [emim]s containing the same anions, indicating that these solvents are more efficient for CO2 sequestration than the emims.
[1] A.L. Kohl, R.B. Nielsen, Gas Purification, fifth ed., Gulf Publishing Company, Texas, 1997. [2] L.M. Galán Sánchez, G.W. Meindersma, A.B. de Haan, Trans. IChemE, Part A, Chem. Eng. Res. Des. 85 (2007) 31–39. [3] J.S. Wilkes, P. Wassercheid, T. Welton (Eds.), Ionic Liquids in Synthesis, Wiley VCH Verlag, 2002. [4] E.D. Bates, R.D. Mayton, I. Ntai, J.H. Davis, J. Am. Chem. Soc. 124 (2002) 926– 927. [5] S. Zhang, Y. Chen, F. Li, X. Lu, W. Dai, R. Mori, Catal. Today 115 (2006) 61–69. [6] L.A. Blanchard, Z. Gu, J.F. Brennecke, J. Phys. Chem. B 105 (2001) 2437–2444. [7] D. Camper, P. Scovazzo, C. Koval, R. Noble, Ind. Eng. Chem. Res. 43 (2004) 3049–3054. [8] P. Scovazzo, D. Camper, J. Kieft, J. Poshusta, C. Koval, R. Noble, Ind. Eng. Chem. Res. 43 (2004) 6855–6860. [9] D. Camper, C. Becker, C. Koval, R. Noble, Ind. Eng. Chem. Res. 45 (2006) 445– 450. [10] A. Shariati, C.J. Peters, J. Supercrit. Fluids 29 (2004) 43–48. [11] A. Shariati, C.J. Peters, J. Supercrit. Fluids 34 (2005) 171–176. [12] A.M. Scilderman, S. Raeissi, C.J. Peters, Fluid Phase Equilib. 260 (2007) 19–22. [13] Y.S. Kim, W.Y. Choi, J.H. Jang, K.-P. Yoo, C.S. Lee, Fluid Phase Equilib. 228–229 (2005) 439–445. [14] E.-K. Shin, B.-C. Lee, J.S. Lim, J. Supercrit. Fluids 45 (2008) 282–292. [15] A.N. Soriano, B.T.J. Doma, M.-H. Li, J. Chem. Eng. Data 53 (2008) 2550–2555. [16] A.N. Soriano, B.T.J. Doma, M.-H. Li, J. Chem. Thermodyn. 40 (2008) 1654–1660. [17] A.N. Soriano, B.T.J. Doma, M.-H. Li, J. Chem. Thermodyn. 41 (2009) 525–529. [18] P.J. Carvalho, V.H. Álvarez, J.J.B. Machado, J. Pauly, J.-L. Daridon, I.M. Marrucho, M. Aznar, J.A.P. Coutinho, J. Supercrit. Fluids 48 (2009) 99–107. [19] M. Shokouhi, M. Adibi, A.H. Jalili, M. Hosseini-Jenab, A. Mehdizadeh, J. Chem. Eng. Data, doi:10.1021/je900716q. [20] J.F. Dubreuil, M.-H. Famelart, J.P. Bazureau, Org. Proc. Res. Dev. 6 (2002) 374– 378. [21] S.H. Yeon, K.S. Kim, S. Choi, H. Lee, Electrochim. Acta 50 (2005) 5399–5407. [22] J.F. Dubreuil, J.P. Bazureau, Tetrahedron 59 (2003) 6121–6130. [23] A.H. Jalili, M. Rahmati-Rostami, C. Ghotbi, M. Hosseini-Jenab, A.N. Ahmadi, J. Chem. Eng. Data 54 (2009) 1844–1849. [24] M. Rahmati-Rostami, C. Ghotbi, M. Hosseini-Jenab, A.N. Ahmadi, A.H. Jalili, J. Chem. Thermodyn. 41 (2009) 1052–1055. [25] NIST Scientific and Technical Databases; Thermophysical Properties of Fluid Systems.
JCT 09–433
Contents lists available at ScienceDirect
J. Chem. Thermodynamics journal homepage: www.elsevier.com/locate/jct
Solubility of CO2 in 1-(2-hydroxyethyl)-3-methylimidazolium ionic liquids with different anions Amir Hossein Jalili a,*, Ali Mehdizadeh a, Mohammad Shokouhi a, Hossein Sakhaeinia b, Vahid Taghikhani c a
Gas Science Department, Gas Research Division, Research Institute of Petroleum Industry (RIPI), National Iranian Oil Company (NIOC), P.O. Box 14665-137, West Blvd., Azadi Sport Complex, Tehran, Iran b Department of Chemical Engineering, Faculty of Engineering, Islamic Azad University, Semnan Branch, Semnan, Iran c Department of Chemical and Petroleum Engineering, Sharif University of Technology, Azadi Ave., Tehran, Iran
a r t i c l e
i n f o
Article history: Received 4 January 2010 Accepted 1 February 2010 Available online 6 February 2010 Keywords: Gas solubility Henry’s law constant Carbon dioxide Ionic liquid Gibbs free energy of solution
a b s t r a c t The solubility of carbon dioxide in a series of 1-(2-hydroxyethyl)-3-methylimidazolium ([hemim]+) based ionic liquids (ILs) with different anions, viz. hexafluorophosphate ([PF6]), trifluoromethanesulfonate ([OTf]), and bis-(trifluoromethyl)sulfonylimide ([Tf2N]) at temperatures ranging from 303.15 K to 353.15 K and pressures up to 1.3 MPa were determined. The solubility data were correlated using the Krichevsky–Kasarnovsky equation and Henry’s law constants were obtained at different temperatures. Using the solubility data, the partial molar thermodynamic functions of solution such as Gibbs free energy, enthalpy, and entropy were calculated. Comparison showed that the solubility of CO2 in the ILs studied follows the same behaviour as the corresponding conventional 1-ethyl-3-methylimidazolium ([emim]+) based ILs with the same anions, i.e. [hemim][NTf2] > [hemim][OTf] > [hemim][PF6] > [hemim][BF4]. Ó 2010 Elsevier Ltd. All rights reserved.
1. Introduction Natural gas, as a source of clean fuel and as a starting material for synthesis of various petrochemical products, is usually contaminated with a variety of undesirable materials, especially CO2 and H2S acid gases. The presence of CO2 lowers the heating value of natural gas while H2S gives it high toxicity and corrosiveness as a clean fuel source. Alkanolamines, especially monoethanolamine, diethanolamine, and methyldiethanolamine are the main constituents of aqueous solutions used in industrial natural gas treating plants [1]. There are some disadvantages in commercial use of these alkanolamine solutions, including loss of alkanolamine and transfer of water into the gas stream during desorption stage and degradation of alkanolamines to form corrosive by products, which make the scrubbing process economically expensive [2]. Room temperature ionic liquids (RTILs) are molten salts that are liquid over a wide temperature range including ambient [3]. Due to the existence of Coulombic attraction between the ions of ILs, they exhibit negligibly small vapour pressure, meaning that ionic liquids are essentially non-volatile and could facilitate the sequestration of gases without loss of the capture agent into the gas stream. They also have high thermal and electrochemical stability. Nowadays, one of the active research areas is to explore task specific * Corresponding author. Tel.: +98 21 48252466; fax: +98 21 44739716. E-mail addresses: [email protected], [email protected] (A.H. Jalili). 0021-9614/$ - see front matter Ó 2010 Elsevier Ltd. All rights reserved. doi:10.1016/j.jct.2010.02.002
ionic liquids to replace conventional alkanolamine solutions for removal of acid gases (CO2 and H2S) in gas sweetening processes [4]. Another research area is concerned with fixation and sequestration of CO2, which is regarded as the most important greenhouse gas, by employing ionic liquids [5]. An important feature in the evaluation of ILs for potential use in industrial natural gas treating processes is the knowledge of the solubility of gases at various temperatures and pressures. In the past few years, a growing number of measurements reporting solubility of CO2 in various conventional ILs have become available. Blanchard et al. experimentally explored the high-pressure phase behaviour of CO2 with six imidazolium-based ILs including 1-ethyl-3-methylimidazolium ethylsulfate ([emim][EtSO4]) by systematically changing the anionic and cationic components of the IL [6]. Camper et al. measured the low-pressure solubility of CO2 and C2H4 in different 1-ethyl-3-methylimidazolium-based ILs with different anions [PF6], [Tf2N] , [OTf], and dicyanamide ([dca]) as well as 1-butyl-3-methylimidazolium hexafluorophosphate ([bmim][PF6]) at T = 303 K, and then showed that the regular solution theory can be used to model the solubility results at low pressures [7]. In a subsequent work, they concluded that CO2/IL complexations are not the sole controlling factor in relative CO2 solubility and the energies of vapourisation and molar volumes appear to be determining factors in relative CO2 solubility in ILs [8]. Later they reported the solubility in the form of Henry’s law constants and diffusion of CO2 at five discrete temperatures from
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303 K to 343 K and those of ethylene, ethane, propane, and propylene in [emim][Tf2N] at 303 K [9]. Shariati and Peters have reported the experimental high-pressure phase behaviour of the binary system CO2/[emim][PF6] within temperature and pressure ranges of 308.14 K to 366.03 K and 1.49 MPa to 97.10 MPa, respectively [10]. In a subsequent work, they reported the experimental highpressure phase behaviour of CO2/ILs at pressures up to 97 MPa and T = 303.15 K [11]. The ILs studied were comprised of imidazolium-based cations from [emim]+ to [bmim]+ to [hmim]+ and [PF6] anion. In their investigation, they showed that the solubility of CO2 in the ILs increased by increasing the alkyl group chain length. Later Schilderman et al. reported experimental values for the solubility of CO2 in [emim][Tf2N] within temperature and pressure ranges of 310 K to 450 K and 0 MPa to 15 MPa respectively, and concluded that the solubility of CO2 in [emim][Tf2N], due to the presence of fluoroalkyl groups in the anion, is much higher than that of [emim][PF6] [12]. Kim et al. have reported the solubility of CO2 in some ILs including [emim][BF4] and [emim][Tf2N] at T = 298.15 K and pressures up to 1 MPa [13]. They used the data to obtain the parameters of a group contribution non-random lattice-fluid equation of state. Shin et al. have reported the solubility values of CO2 in a series of bis(trifluoromethyl) sulfonylimide ILs with 1-alkyl-3methylimidazolium ([Cn-mim]+, n = 2, 4, 6, 8) cations at pressures up to 45 MPa and temperatures from 298.15 K to 343.15 K and used the Peng–Robinson equation of state to correlate the data [14]. The solubility of CO2 in [bmim][PF6] and [emim][BF4] for temperatures ranging from 303.2 K to 343.2 K and pressures up to 5 MPa [15] and in 1-ethyl-3-methylimidazolium 2-(2-methoxyethoxy) ethylsulfate at pressures up to 6.7 MPa and temperatures from 303.2 K to 343.2 K [16] using a thermogravimetric microbalance apparatus is reported by Soriano and his coworkers. Subsequently they reported the solubility of CO2 in [emim][OTf] for temperatures ranging from 303.2 K to 343.2 K and pressures up to 5.9 MPa [17]. The high-pressure phase behaviour of carbon dioxide in 1-methyl-3-pentylimidazolium bis(trifluoromethylsulfonyl)imide ([pmim] [Tf2N]) as well as in [emim][Tf2N] at temperatures up to 363 K and pressures up to 50 MPa is investigated by Carvalho et al. [18]. Finally, Shokouhi et al. have reported the solubility and diffusion of H2S and CO2 in 1-(2-hydroxyethyl)imidazolium tetrafluoroborate ([hemim][BF4]) at temperatures ranging from 303.15 K to 353.15 K and pressures up to 1.1 MPa [19]. They have used the data obtained to estimate Henry’s law constants and diffusion coefficients at different temperatures. In this work, the equilibrium solubility of CO2 in the ionic liquids [hemim][PF6], [hemim][OTf], and [hemim][Tf2N] at six temperatures from 303.15 K to 353.15 K was measured. The values of solubility determined were used to estimate Henry’s law constants and partial molar thermodynamic functions of solution of CO2 at different temperatures. A correlation equation for the solubility obtained and partial molar volume at infinite dilution data with temperature is presented here. The results obtained are compared with [hemim][BF4] found in the previous work [19] and other [emim] based ILs, which are reported in the literature.
2. Experimental 2.1. Materials Carbon dioxide (c.p. grade, mass fraction purity >0.9995) was obtained from the Roham Gas Company. The 1-(2-hydroxyethyl)3-methylimidazolium chloride ([hemim][Cl]) was synthesized by modification of a known literature procedure [20]. A mixture of 1-methylimidazole (5 g, 61 mmol) and 2-chloroethanol (4.9 g, 61 mmol) was placed in a TFM vessel of MICCROSYNTH laboratory microwave oven (Milestone srl, Italy) and subjected to microwave
irradiation at T = 383 K (200 W, 10 bars) for half an hour. After cooling to ambient temperature, the reaction mixture was washed with 20 cm3 of ethyl acetate to remove non-reacted starting materials. The solid product obtained was then heated at T = 343 K for 7 h at reduced pressure to remove the residual water and obtain 6.6 g of the ionic liquid. The yield was 66.6%. 1H NMR (500 MHz, D2O, T = 209 K: d 106 = 3.87 (3H, S, NCH3), 3.89 (2H, t, NCH2CH2OH), 4.28 (2H, t, NCH2CH2OH), 7.42 (1H, d, H-4), 7.47 (1H, d, H-5), 8.71 (1H, S, H-2). The ionic liquids [hemim][PF6], [hemim][OTf], [hemim][Tf2N], were prepared according to the synthesis method described by Yeon et al. [21] and Dubreuil and Bazureau [22], which are briefly described below. 2.1.1. Synthesis of [hemim][PF6] 1-(2-Hydroxyethyl)-3-methylimidazolium chloride (5.0 g, 31 mmol) and potassium hexaflurophosphate (5.66 g, 31 mmol) in deionised water (10 cm3) and acetonitrile (10 cm3) were mixed and the resulting two-phase mixture was magnetically stirred at room temperature for 24 h. After completion of the reaction, the solvents were removed at reduced pressure and the ionic liquid was removed from KCl by extraction in dry methanol (17 cm3) and chloroform (3 cm3). The suspension was then filtered and the filtrate was washed with chloroform (3 7 cm3) and diethyl ether (3 7 cm3) to remove impurities and non-reacted starting materials. The solvents were removed at reduced pressure. The product obtained was then heated at T = 343 K for 7 h at reduced pressure to remove the residual water and obtain 3.56 g of the ionic liquid. The yield was 42.6%. 1H NMR (500 MHz, D2O, T = 343 K): d 106 = 3.87 (3H, S, NCH3), 3.90 (2H, t, NCH2CH2OH), 4.28 (2H, t, NCH2CH2OH), 7.41 (1H, d, H-4), 7.46(1H, d, H-5), 8.68 (1H, S, H-2). 2.1.2. Synthesis of [hemim][OTf] 1-(2-Hydroxyethyl)-3-methylimidazolium chloride (5.0 g, 31 mmol) and ammonium trifluoromethanesulfonate (5.14 g, 31 mmol) in deionised water (10 cm3) and acetonitrile (10 cm3) were mixed and the resulting two-phase mixture was magnetically stirred at room temperature for 24 h. After completion of the reaction, the solvents were removed at reduced pressure and the ionic liquid was removed from NH4Cl by extraction in dry methanol (17 cm3) and chloroform (3 cm3). The suspension was then filtered and the filtrate was washed with chloroform (3 7 cm3) and diethyl ether (3 7 mL) to remove impurities and non-reacted starting materials. The solvents were removed at reduced pressure. The product obtained was then heated at T = 343 K for 7 h at reduced pressure to remove the residual water and obtain 5.21 g of the ionic liquid. The yield was 61.3%. 1H NMR (300 MHz, D2O, 25 °C): d 106 = 3.84 (3H, S, NCH3), 3.87 (2H, t, NCH2CH2OH), 4.25 (2H, t, NCH2CH2OH), 7.39 (1H, d, H-4), 7.44 (1H, d, H-5), 8.68 (1H, S, H-2). 2.1.3. Synthesis of [hemim][Tf2N] 1-(2-Hydroxyethyl)-3-methylimidazolium chloride (5.0 g, 31 mmol) and lithium bis (trifluoromethanesulfonimide) (8.83 g, 31 mmol) in deionised water (10 cm3) and acetonitrile (10 cm3) were mixed and the resulting two-phase mixture was magnetically stirred at room temperature for 24 h. After completion of the reaction, the solvents were removed at reduced pressure and the ionic liquid was removed from LiCl by extraction in dry methanol (17 cm3) and chloroform (3 cm3). The suspension was then filtered and the filtrate was washed with chloroform (3 7 cm3) and diethyl ether (3 7 cm3) to remove impurities and non-reacted starting materials. The solvents were removed at reduced pressure. The product obtained was then heated at T = 343 K for 7 h at reduced pressure to remove the residual water and obtain 7.84 g of the ionic liquid. The yield was 62.6%. 1H NMR (500 MHz, D2O,
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25 °C): d 106 = 3.87 (3H, S, NCH3), 3.90 (2H, t, NCH2CH2OH), 4.28 (2H, t, NCH2CH2OH), 7.42 (1H, d, H-4), 7.47 (1H, d, H-5), 8.69 (1H, S, H-2).
of the method of measurement have been checked in our previous work [23]. The Krichevsky–Kasarnovsky equation, [26] modified for solvents of negligible vapour pressure [23,24], was used to model the experimental results obtained in this work:
ln
2.2. Apparatus and procedure The experimental method for the measurement of gas solubility has previously been presented in detail [23,24] and only a short description will be provided here. In this technique, known quantities of gaseous solute and degassed solvent are contacted at a constant temperature inside an equilibrium cell of known volume. On reaching thermodynamic equilibrium, the pressure above the liquid solution is constant and directly related to the solubility of the gas in the liquid. The quantity of solute present in the liquid solution, nlsolute , is calculated by the difference between two (P, V, T) measurements: first when the gas is introduced from the gas container of known volume into the equilibrium cell containing the IL and secondly after reaching thermodynamic equilibrium (i.e. autoclave pressure remains fixed and no longer changes with time):
nlsolute ¼ ntotal ngsolute ;
V gc Pi P f ; RT gc Z i Z f
ð2Þ
where V gc denotes the volume of the gas container, Z i and Z f are the compressibility factors corresponding to the initial and final pressures Pi and Pf , respectively, in the gas container before and after transferring the gas, and T gc is the temperature of the gas container. The most accurate (P, V, T) data presented by NIST for pure compounds were used to calculate compressibility factors [25]. The ngsolute in equation (1) is the number of moles of gas solute left in the gas phase and was determined from the following equation:
ngsolute ¼
V gP ; ZRT
ð3Þ
where V g is the volume of the gas phase above IL phase, T is the equilibrium temperature of the cell, and Z is the compressibility factor of gas solute at P and T. In all experiments, the IL was dried in vacuo (below 1.0 kPa) before solubility measurements for about 48 h at a temperature of 343 K to remove trace amounts of moisture and volatile impurities. Using a Mettler model DL-37 Karl-Fischer volumetric titrator, the water content of ILs was found to be below (100 ± 10)106. The temperature of the double-wall equilibrium cell, which was connected to a water recirculation bath (PMT Tamson model T 2500), was measured with a Lutron model TM-917 digital thermometer with a 0.01 K resolution using a Pt-100 sensor inserted into the cell. The equilibrium cell pressure was measured using a KELLER model PA-33X pressure transmitter sensor in the range of 0 MPa to 2 MPa, which was accurate to within 0.01% of full scale and that of the gas container was measured using a Druck model PTX 1400 pressure transmitter sensor in the range of 0 MPa to 4 MPa, which was accurate to within 0.1% of full scale. The calibration of pressure sensors was carried out against a dead-weight gauge. 3. Results and discussion The results of the measurement of the solubility of carbon dioxide in the ionic liquids [hemim][PF6], [hemim][OTf], and [hemim][Tf2N], at temperatures of (303.15, 313.15, 323.15, 333.15, 343.15, 353.15) K and pressures up to about 1.3 MPa are summarized in tables 1 to 3, respectively. The reliability and accuracy
ð4Þ
where f2 is the fugacity of pure solute (carbon dioxide) in the gas phase, x2 is the mole fraction of solute in the solvent 1, H21 is Henry’s law constant of gas solute 2 1 in solvent 1 at the pressure, P, v 2 is the partial molar volume of gas solute 2 at TABLE 1 Solubility of carbon dioxide gas in [hemim][PF6]: x2, CO2 mole fraction; P, equilibrium pressure; T, equilibrium temperature. P/kPa
x2
T = 303.15 K 133.3 0.013 ± 0.006 310.6 0.035 ± 0.008 579.4 0.064 ± 0.009 667.0 0.075 ± 0.008 745.2 0.084 ± 0.009 906.3 0.104 ± 0.011
ð1Þ
where ntotal is the total number of CO2 moles injected from the gas container into the autoclave and calculated using the following equation:
ntotal ¼
f2 v1P ¼ ln H21 þ 2 x2 RT
152.2 355.9 502.3 663.2 762.1 853.2 935.8
T = 333.15 K 0.009 ± 0.006 0.026 ± 0.008 0.037 ± 0.007 0.049 ± 0.009 0.059 ± 0.009 0.066 ± 0.009 0.070 ± 0.010
P/kPa
x2
T = 313.15 K 139.0 0.012 ± 0.006 324.5 0.033 ± 0.008 459.5 0.045 ± 0.007 606.3 0.060 ± 0.009 699.9 0.068 ± 0.008 783.7 0.076 ± 0.009 854.2 0.085 ± 0.010 1001 0.101 ± 0.011 160.5 370.4 523.4 687.2 794.1 889.3 973.0
T = 343.15 K 0.006 ± 0.006 0.023 ± 0.008 0.033 ± 0.007 0.048 ± 0.009 0.054 ± 0.009 0.061 ± 0.009 0.066 ± 0.010
P/kPa
x2
T = 323.15 K 145.2 0.010 ± 0.006 341.3 0.028 ± 0.008 482.4 0.040 ± 0.007 634.2 0.055 ± 0.009 733.1 0.062 ± 0.009 819.2 0.070 ± 0.009 897.0 0.076 ± 0.010 986.2 0.086 ± 0.011 169.0 386.0 542.0 717.0 827.1 926.2 1008 1127
T = 353.15 K 0.009 ± 0.002 0.020 ± 0.004 0.032 ± 0.006 0.042 ± 0.008 0.049 ± 0.009 0.055 ± 0.010 0.064 ± 0.011 0.070 ± 0.013
TABLE 2 Solubility of carbon dioxide gas in [hemim][OTf]: x2, CO2 mole fraction; P, equilibrium pressure; T, equilibrium temperature. P/kPa
x2
T = 303.15 K 101.2 0.017 ± 0.005 251.6 0.044 ± 0.008 432.0 0.073 ± 0.010 652.8 0.107 ± 0.013 895.6 0.144 ± 0.016 1157 0.188 ± 0.019 115.2 286.8 492.0 745.4 1021
T = 333.15 K 0.014 ± 0.005 0.036 ± 0.008 0.061 ± 0.011 0.089 ± 0.014 0.123 ± 0.016
P/kPa
x2
T = 313.15 K 105.6 0.016 ± 0.005 263.8 0.040 ± 0.008 452.0 0.068 ± 0.010 684.8 0.100 ± 0.014 939.0 0.135 ± 0.016 1216 0.176 ± 0.019 119.2 297.2 510.8 772.8 1063
T = 343.15 K 0.013 ± 0.005 0.034 ± 0.008 0.058 ± 0.011 0.087 ± 0.014 0.118 ± 0.017
P/kPa
x2
T =323.15 K 110.4 0.015 ± 0.005 275.2 0.038 ± 0.008 472.6 0.064 ± 0.011 715.6 0.094 ± 0.014 981.6 0.128 ± 0.016 1280 0.162 ± 0.019 123.6 307.2 558.4 801.2 1105
T = 353.15 K 0.013 ± 0.005 0.034 ± 0.008 0.058 ± 0.011 0.083 ± 0.014 0.112 ± 0.017
TABLE 3 Solubility of carbon dioxide gas in [hemim][Tf2N]: x2, CO2 mole fraction; P, equilibrium pressure; T, equilibrium temperature. P/kPa
x2
T = 303.15 K 97.3 0.022 ± 0.005 243.7 0.054 ± 0.008 443.6 0.093 ± 0.011 564.1 0.119 ± 0.009 654.0 0.139 ± 0.009 795.4 0.164 ± 0.011 954.1 0.191 ± 0.013 T = 333.15 K 112.5 0.018 ± 0.005 281.9 0.043 ± 0.008 514.9 0.075 ± 0.011 656.8 0.096 ± 0.009 764.2 0.111 ± 0.009 930.4 0.132 ± 0.012 1114.3 0.158 ± 0.014
P/kPa
x2
T = 313.15 K 102.7 0.021 ± 0.006 256.7 0.050 ± 0.010 468.4 0.086 ± 0.014 597.4 0.109 ± 0.012 693.5 0.127 ± 0.012 843.3 0.151 ± 0.016 1011 0.177 ± 0.018 T = 343.15 K 117.1 0.017 ± 0.005 293.3 0.041 ± 0.008 535.6 0.072 ± 0.011 683.7 0.092 ± 0.009 796.8 0.106 ± 0.010 968.8 0.127 ± 0.012
P/kPa
x2
T = 323.15 K 107.5 0.019 ± 0.005 269.3 0.046 ± 0.008 491.8 0.080 ± 0.011 627.5 0.102 ± 0.009 730.0 0.118 ± 0.009 888.1 0.141 ± 0.012 1064 0.166 ± 0.014 T = 353.15 K 121.7 0.016 ± 0.005 305.0 0.039 ± 0.008 556.7 0.069 ± 0.011 710.9 0.088 ± 0.009 828.7 0.102 ± 0.010 1007.5 0.122 ± 0.012
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TABLE 4 1 Thermodynamic properties of CO2 in ionic liquids: T, temperature; H21, Henry’s law constant; v 1 2 , partial molar volume at infinite dilution; DsolG , Gibbs free energy of solution; DsolH1, enthalpy of solution; DsolS1, entropy of solution. T/K
H21/MPa
3 1 m1 2 /(cm mol )
303.15 313.15 323.15 333.15 343.15 353.15
8.94 ± 0.08 10.1 ± 0.09 11.7 ± 0.08 13.2 ± 0.05 14.8 ± 0.03 16.5 ± 0.02
118 ± 4 109 ± 4 99.3 ± 5 94.1 ± 3 85.5 ± 5 79.2 ± 5
303.15 313.15 323.15 333.15 343.15 353.15
6.14 ± 0.01 6.89 ± 0.01 7.74 ± 0.05 8.28 ± 0.04 8.93 ± 0.07 9.73 ± 0.06
303.15 313.15 323.15 333.15 343.15 353.15
4.42 ± 0.01 5.11 ± 0.02 5.75 ± 0.05 6.50 ± 0.04 7.02 ± 0.07 7.69 ± 0.01
DsolG1/(kJ mol1)
DsolH1/(kJ mol1)
DsolS1/(J mol1 K1)
[hemim][PF6] 11.3 12.0 12.8 13.5 14.3 15.0
11.07 11.03 10.99 10.96 10.93 10.90
73.87 73.62 73.64 73.51 73.40 73.30
121 ± 3 112 ± 3 102 ± 5 96.9 ± 3 88.4 ± 5 82.2 ± 5
[hemim][OTf] 10.4 11.0 11.7 12.2 12.8 13.4
9.35 8.75 8.19 7.66 7.16 6.70
65.1 63.1 61.5 59.7 58.2 57.0
184 ± 2 173 ± 2 161 ± 2 148 ± 2 133 ± 3 118 ± 3
[hemim][Tf2N] 9.55 10.2 10.9 11.6 12.1 12.8
11.8 10.9 10.1 9.29 8.55 7.85
70.5 67.6 64.9 62.6 60.3 58.3
TABLE 5 Numerical values of the parameters of equations (5) and (6). IL
A0
A1
A2
B0/(cm3 mol1)
B1/(cm3 mol1 K1)
B2/(cm3 mol1 K2)
[hemim][PF6] [hemim][OTf] [hemim][Tf2N]
6.33876 1.80997 0.622176
1186.22 1128.77 1944.24
22 000.8 341 511 510 054
724.52 724.81 1.0193
3.0489 3.0424 2.2464
0.00346 0.00346 0.00542
infinite dilution, R is the universal gas constant, and T is the absolute temperature. The fugacity of carbon dioxide was calculated using the most accurate corresponding states used by NIST for pure compounds [25]. Values of ln H21 at each temperature T were obtained from the intercept of Krichevsky–Kasarnovsky plots (that is plots of lnðf2 =x2 Þ vs P) at the specified temperature. The Henry’s law constants are given in table 4 for the solubility of CO2 in the ILs studied in this work together with their standard deviations. The Henry’s law constants were fit by the equation, 2 X i¼0
30
Ai ðT=KÞ
35
i
:
ð5Þ
The parameters Ai of equation (5) obtained for CO2 are summarized in table 5. The Henry’s law constants are compared with each other in figure 1 as a function of temperature. It can be observed that the solubility of carbon dioxide in the ILs decreases by increasing the temperature. The curves also indicate that the solubility of CO2 in the ILs studied in this work is typical of that of physical solvents [22], therefore obeying the Henry’s law. It can be seen that, the solubility of CO2 is the highest in [hemim][Tf2N] and the lowest in [hemim][PF6]. Also variation with temperature of the Henry’s law constant for the solubility of CO2 in [hemim][BF4] investigated in previous work [19] as well as in the corresponding conventional ILs, [emim][BF4], [emim][OTf], and [emim][Tf2N] are shown in figure 1. It can be observed that the solubility of CO2 in [hemim][BF4] is lower than the ILs studied in this work. This can be explained on the basis of the fact that ILs containing anions with fluoroalkyl groups, i.e. [Tf2N] and [OTf], have high affinities for CO2 than those containing [PF6] and [BF4] anions. Schilderamn et al. [12] and Zhang et al. [27] have mentioned same conclusion by experimental and theoretical (using COSMO-RS method) investigations, respectively. Inspection of figure 1 reveals that the solubility of CO2 in both [hemim][Tf2N] and [emim][Tf2N] is comparable. This observation shows that in this case, the anion, which contains two –CF3 groups, is the predominant species in determining CO2 solubility and the presence of a hydroxyl group, though it is capable of hydrogen bond formation, plays a very negligible role in this respect. However, in the case of other ILs, one can see that the effect of the anion gradually ceases as the number of –CF3 groups decreases from 1 for [hemim][OTf] to zero for [hemim][PF6] and [hemim][[BF4] and thus, the solubility of CO2 decreases too. The higher solubility of CO2 in [hemim] ILs relative to the corresponding [emim]s containing the same anion may be ascribed to the presence of hydroxyl group, which participates in hydrogen bond formation with CO2 solute molecules. We are studying these observations, from molecular point of view, for the interaction and configuration of CO2 and H2S molecules in functionalized ILs and their effect on the bulk liquid structure, by using molecular dynamics (MD) simulation and will present the results in the near future.
H 21 / MPa
lnðH21 ðTÞÞ ¼
40
25
20 15 10
5 0 300
310
320
330
340
350
360
T /K FIGURE 1. Plot of Henry’s law constants as a function of temperature for the solubility of CO2 in ILs: N, [hemim][Tf2N]; d, [hemim][BF4] (reference [19]); j, [hemim][OTf]; , [hemim][PF6]; 4, [emim][Tf2N] (reference [18]); s, [emim][BF4] (reference [15]); h, [emim][OTf] (reference [17]); +, [emim][OTf] (reference [7]).
The slope of Krichevsky–Kasarnovsky plots at each temperature, T, yields v 1 2 at the specified temperature. The obtained experimental data of partial molar volumes 1 of gas solute, CO2, at infinite dilution, v 2 , are presented in table 4. They were correlated with temperature using the following simple quadratic formula
A.H. Jalili et al. / J. Chem. Thermodynamics 42 (2010) 787–791
v 12 ðTÞ ¼
2 X
Bi T i ;
ð6Þ
791
5. Supporting information available
i¼0
3 1 where v 1 ) and T is in K. The obtained parameters Bi of equation (6) 2 is in (cm mol are summarized in table 5. It can be observed from table 4 that the v 1 2 values increase with temperature for CO2 in [hemim][PF6] and [hemim][OTf] but they decrease for CO2 in [hemim][Tf2N]. It can be shown that the Gibbs free energy of solution, corresponding to the change in Gibbs free energy when the solute is transferred at constant temperature from the pure perfect gas at the standard pressure to the standard state of infinite dilution of the solute in the solvent, is given by [28]:
Dsol G1 ¼ RT ln
Kh P
0
;
This information is available free of charge via the Internet at http://pubs.acs.org/. Acknowledgements We are thankful to the research council of the Research Institute of Petroleum Industry (RIPI) and also to the Research and Development of the National Iranian Oil Company (NIOC) for their support of this work.
ð8Þ
References where P0 is the standard state pressure. The partial molar differences in enthalpy and entropy between the two states can be obtained by calculating the corresponding partial derivatives of the Gibbs free energy with respect to temperature
Dsol H1 ¼ T 2
Dsol S1 ¼
@ Dsol G1 @ Kh ¼ RT 2 ; ln T @T @T P0
ðDsol H1 Dsol G1 Þ : T
ð9Þ
ð10Þ
The pressure range considered in this work is not too high to cause Henry’s law constant to be a strong function of pressure and Henry’s law is weakly dependent on pressure under the specified conditions [23,24]. Therefore, it does not give rise to large errors if one ignores this pressure dependency. By means of this approximation and using equations (8)–(10), we estimated the thermodynamic functions of solution at infinite dilution for CO2 in the IL. The values for the Gibbs free energy, enthalpy, and entropy of solution are given in table 4 at six discrete temperatures between 303.15 K and 353.15 K. As it can be observed, the Dsol G1 values are positive and increase with temperature in a similar manner for the solubility of CO2 in ILs. The Dsol H1 values and Dsol S1 values are negative. The variation with temperature of the Dsol H1 values and Dsol S1 values are positive for CO2 in the ILs studied in this work and they increase with temperature. In the case of [hemim][PF6] the rate of variation is very slow while for [hemim][Tf2N] the rate of variation is the highest.
4. Conclusions New experimental results for the solubility, partial molar volume at infinite dilution, and thermodynamic functions of solution of carbon dioxide gas in functionalized ionic liquids 1-(2-hydroxyethyl)-3-imidazolium cations with different anions, not previously reported in the literature, have been measured and presented in this work. The solubility of carbon dioxide in [hemim][Tf2N] is the highest and in [hemim][BF4] is the lowest. The solubility of CO2 in the ILs studied in this work is of a physical nature. The solubility of carbon dioxide in [hemim] ILs is greater than corresponding conventional [emim]s containing the same anions, indicating that these solvents are more efficient for CO2 sequestration than the emims.
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JCT 09–433