Solubility of CO2 and H2S in the ionic liquid 1-ethyl-3-methylimidazolium tris(pentafluoroethyl)trifluorophosphate

Solubility of CO2 and H2S in the ionic liquid 1-ethyl-3-methylimidazolium tris(pentafluoroethyl)trifluorophosphate

J. Chem. Thermodynamics 67 (2013) 55–62 Contents lists available at ScienceDirect J. Chem. Thermodynamics journal homepage: www.elsevier.com/locate/...

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J. Chem. Thermodynamics 67 (2013) 55–62

Contents lists available at ScienceDirect

J. Chem. Thermodynamics journal homepage: www.elsevier.com/locate/jct

Solubility of CO2 and H2S in the ionic liquid 1-ethyl-3methylimidazolium tris(pentafluoroethyl)trifluorophosphate Amir Hossein Jalili a,⇑, Mohammad Shokouhi a, Gerd Maurer b, Masih Hosseini-Jenab a a b

Gas Research Division, Research Institute of Petroleum Industry (RIPI), P.O. Box: 14665-137, Tehran, Iran Department of Mechanical and Process Engineering, University of Kaiserslautern, P.O. Box: 3049, D-67653 Kaiserslautern, Germany

a r t i c l e

i n f o

Article history: Received 25 June 2013 Received in revised form 24 July 2013 Accepted 26 July 2013 Available online 2 August 2013 Keywords: Henry’s law constant Hydrogen sulfide Carbon dioxide Ionic liquid Gas separation Natural gas sweetening

a b s t r a c t The solubility of two single gases carbon dioxide and hydrogen sulfide in the ionic liquid 1-ethyl-3-methylimidazolium tris(pentafluoroethyl)trifluorophosphate ([C2mim][eFAP]) was experimentally determined at temperatures from (303 to 353) K and pressures up to about 2.0 MPa. Results show that hydrogen sulfide is more soluble in that particular ionic liquid than carbon dioxide. At fixed temperature and pressure, the amount of dissolved H2S is more than twice the amount of CO2. The new experimental data were used to determine the Henry’s law constants, which again were used to derive some thermodynamic functions of the gas/solvent systems, such as, for example, the change of the partial molar Gibbs free energy of the gases upon solution in the ionic liquid. Two models were used to correlate the new experimental data: (1) a model comprised of the extended Henry’s law and Pitzer’s virial expansion for the excess Gibbs free energy, and (2) a generic Redlich–Kwong (RK) cubic equation of state recently proposed for gas-ionic liquid systems. Both models are equally suited to correlate the experimental results. The (CO2 + H2S) selectivity of [C2mim][eFAP] was calculated from the RK EoS at various temperatures, pressures and CO2/H2S feed ratios and compared with the recently reported results for the selectivity of other ionic liquids. Ó 2013 Elsevier Ltd. All rights reserved.

1. Introduction Carbon dioxide is one of the most abundant acid gas impurities in many fossil fuel energy resources such as, for example sour natural gas and associated petroleum gas. Carbon dioxide is also produced in large scale in combustion processes that use coal and hydrocarbon fuel for the production of heat and electricity in power plants. Hydrogen sulfide is a highly toxic acid gas, which exists naturally along with methane, light hydrocarbons and CO2 in many oil and gas reservoirs. The presence of CO2 reduces the heating value of hydrocarbon fuel streams and, such as H2S, causes corrosion in transmission pipelines and process equipments. Physical and/or chemical absorption is one of the most versatile processes to remove CO2 and H2S acid gases from industrial gas streams especially from natural gas [1]. In the absorption processes, CO2 and H2S are either physically dissolved in a solvent such as methanol, 1-methyl-2-pyrrolidone, etc., or chemically dissolved in an aqueous solution of a chemical reacting additive such as, for example, 2-aminoethanol (monoethanolamine), 2,20 -iminodiethanol (diethanolamine), and 2,20 -(methylimino)diethanol (N-methyldiethanolamine) [1]. The absorbed acid gases are then stripped

⇑ 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 2013 Elsevier Ltd. All rights reserved. http://dx.doi.org/10.1016/j.jct.2013.07.022

through reverse reactions in the regenerator to yield (CO2 + H2S) sour gas to be treated in the next step (for example, in a Claus process) and the lean solvent is recycled to the absorption stage of the process. There are some disadvantages of a chemical absorption process, for example, the loss of alkanolamine during desorption by vapourization and/or degradation to corrosive and toxic by products, and the high power consumption [2]. The possibility of replacement of conventional alkanolamine solutions by task specific ionic liquids (TSILs) [3] for the removal of CO2 and H2S acid gases in gas sweetening processes is an active research area nowadays. Ionic liquids also known as room-temperature molten salts are liquid over a wide temperature range including ambient [4]. Their negligibly small vapour pressure (meaning that ionic liquids are essentially non-volatile) and the non-flammability are some of their most remarkable properties. They also are assumed to have high thermal and electrochemical stabilities. A class of ionic liquids containing the tris(1,1,2,2,2-pentafluoroethyl)trifluorophosphate anion (commonly referred to tris(pentafluoroethyl)trifluorophosphate and abbreviated as [eFAP]) paired with imidazolium or phosphonium or pyrrolidium cations was first synthesized in 2005 by Ignat’ev et al. [5]. Some experimental data for the solubility of single gases, especially for CO2, in [eFAP]anion based ionic liquids are already available. Muldoon et al. [6] reported experimental data for the solubility of carbon dioxide in 1-hexyl-3methyl-1H-imidazol-3-ium tris(1,1,2,2,2-pentafluoroethyl) trifluo-

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rophosphate (the 1-alkyl-3-methyl-1H-imidazol-3-ium-based ionic liquids commonly referred in the literature as 1-alkyl-3-methylimidazolium-based ILs are abbreviated hereafter as [Cnmim], in which n represents number of carbon atoms in the alkyl substituent). Almantariotis et al. [7] measured the solubility of carbon dioxide, nitrogen, ethane and nitrous oxide in [C2mim][eFAP], [C4mim][eFAP] and [C6mim][eFAP] and reported their densities and viscosities as a function of temperature and pressure. Blath et al. [8] studied the solubility of CO2, N2, CH4 and CO in [C6mim][eFAP]. Zhang et al. [9] investigated the solubility of CO2 in a wide range of ionic liquids with different cations and anions and showed that ILs with the [eFAP]anion are better solvents for CO2 capture than other ILs. By molecular simulation studies, they provided an explanation why the [eFAP]anion enhances the solubility of CO2 in [eFAP]-based ILs [10]. However, experimental data for the solubility of hydrogen sulfide in ILs, especially [eFAP] anion based ionic liquids are scarce. Jou and Mather [11] reported experimental results for the solubility of H2S in 1-butyl-3-methyl-1H-imidazol-3-ium hexafluorophosphate ([C4mim][PF6]) at temperatures between 298 K and 403 K and pressures up to 10 MPa. Subsequently, Pomelli et al. [12] reported experimental data for the solubility of H2S in several [C4mim]-based ILs with different anions and in a series of 1,1,1-trifluoro-N-[(trifluoromethyl)sulfonyl]methanesulfonamide (commonly referred to bis(trifluoromethanesulfonyl)imide and abbreviated as [Tf2N]) ILs with different cations at 298 K and 1400 kPa. Heintz et al. [13] measured the solubility of CO2 and a mixture of (N2 + H2S) in a polymeric ammonium polyether-based IL with chloride anion in the temperature range from (300 to 500) K and pressures up to 0.23 MPa for H2S and 3.0 MPa for CO2, respectively. Shiflett and Yokozeki [14] reported experimental results for the vapour-liquid-liquid equilibrium (VLLE) measurements of single gases H2S and CO2 as well as their binary mixtures with [C4mim][PF6] at temperatures from (273 to 342) K and pressures up to 6.5 MPa. Our Laboratory at RIPI investigated the solubility of CO2 and/or H2S in a variety of ionic liquids at temperatures from 303 K to 353 K and pressures up to about 2.0 MPa. These publications dealt with the solubility of H2S in [C4mim][PF6], [C4mim][BF4], and [C4mim][Tf2N] [15], the solubility of H2S in [C6mim][PF6], [C6mim][BF4], and [C6mim][Tf2N] [16], the solubility of H2S and CO2 in 1-(2-hydroxyethyl)-3-methyl-1H-imidazol-3-ium ionic liquids ([HOC2mim]) with [PF6], [OTf] (trifluoromethanesulfonate), and [Tf2N] anions [17,18], the solubility and diffusion of H2S and CO2 in 1-(2-hydroxyethyl)-3-methyl-1H-imidazol-3-ium tetrafluoroborate ([HOC2mim][BF4]) [19] and 1-ethyl-3-methyl1H-imidazol-3-ium ethylsulfate ([C2mim][EtSO4]) [20], the solubility of H2S in [C2mim][PF6] and [C2mim][Tf2N] [21] and most recently the solubility of single gases H2S and CO2 as well as their binary mixtures in 1-methyl-3-octyl-3H-imidazol-1-ium 1,1,1-tri-

fluoro-N-[(trifluoromethyl)sulfonyl]methanesulfonamide ([C8mim][Tf2N]) [22] and 1-methyl-3-octyl-3H-imidazol-1-ium hexafluorophosphate ([C8mim][PF6]) [23]. The present publication focuses on the solubility of CO2 and H2S in the ionic liquid [C2mim][eFAP]. All measurements are carried out at temperatures ranging from 303 K to 353 K and pressures up to 2.0 MPa and the new experimental solubility results are compared to literature data where other ionic liquids were employed. The new solubility data are modelled using two distinct correlations related to two theoretical approaches: the generic Redlich– Kwong (RK) cubic equation of state recently proposed by Shiflett and Yokozeki for gas-ionic liquid systems [24–27], and the extended Henry’s law combined with a modification of Pitzer’s activity coefficient model for electrolytes [28,29] on one side and an equation of state to describe the non-ideality of the vapour phase on the other side. Some characteristic partial molar thermodynamic properties of gases dissolved at infinite dilution in that particular ionic liquids were calculated from the solubility data. 2. Experimental 2.1. Materials The specifications and sources of the chemicals used in this work are summarized in table 1. The molecular formula of the cation and anion of [C2mim][eFAP] IL is also included in this table for the convenience of the reader. The IL was pre-treated in vacuo (below 1.0 kPa) for about 24 h at a temperature of 343 K to remove trace amounts of moisture and volatile impurities before use. It was found that the water content of the IL was below (10 ± 1)  105 on mass fraction basis (using a model DL-37 Karl– Fischer Mettler volumetric titrator). 2.2. Apparatus and procedure The experimental apparatus described by Jalili et al. [22] was used in the experimental part of this investigation and only a short description is given here. In this technique, known quantities of the gaseous solute and the pre-treated degassed solvent are brought into contact at a constant temperature inside an equilibrium cell of known volume. In thermodynamic equilibrium, the pressure above the liquid solution is measured by a pressure transmitter. The amount of gas present in the liquid solution, is calculated by the difference between two pVT measurements:

nlsolute ¼ ntotal  ngsolute ;

where ntotal is the total number of CO2 or H2S moles injected from the gas container into the autoclave and calculated using the following equation:

ntotal ¼

TABLE 1 Specifications and sources of chemicals used in this work. Mass fraction purity

Chemical name

Molecular formula

CAS registry number

Hydrogen sulfide

H2S

[778306-4]

0.9995

Carbon dioxide

CO2

[124-389]

0.995

1-Ethyl-3methylimidazolium tris(pentafluoroethyl) trifluorophosphate ([C2mim][eFAP])

C12H11F18N2P

[37773943-0]

>0.99

Source

Roham Gas Company Roham Gas Company Merck Chemical Company

ð1Þ

  V gc pi pf ;  RT gc Z i Z f

ð2Þ

where Vgc denotes the volume of the gas container, Zi 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, R is the universal gas constant, and Tgc is the temperature of the gas container. ngsolute in equation (1) is the number of moles of gas solute left in the gas phase, which was determined in the same manner as described previously [15–23]. The temperature of the liquid jacket equilibrium cell, which was connected to a water recirculation bath (LAUDA, model RE415) with temperature stability within ±0.02 K 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 and the gas container pressures were measured using KELLER model

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A.H. Jalili et al. / J. Chem. Thermodynamics 67 (2013) 55–62 TABLE 2 Solubility of CO2 in [C2mim][eFAP]: p, equilibrium pressure;m, molality of CO2; T, equilibrium temperature.a p/MPa

a

m/mol CO2  kg1IL

0.2092 0.4059 0.4170 0.6486 0.9069 1.1670 1.3059 1.5115 1.5199 1.7376 1.8423

T = 303.15 K 0.1317 ± 0.0044 0.2633 ± 0.0079 0.2709 ± 0.0081 0.4323 ± 0.0126 0.6163 ± 0.0172 0.8218 ± 0.0226 0.9088 ± 0.0249 1.0698 ± 0.0292 1.0856 ± 0.0296 1.2725 ± 0.0345 1.3293 ± 0.0360

0.1544 0.2463 0.7734 1.0818 1.5658 1.8328

T = 333.15 K 0.0644 ± 0.0029 0.1034 ± 0.0044 0.3390 ± 0.0135 0.4718 ± 0.0186 0.6898 ± 0.0269 0.8128 ± 0.0317

m/mol CO2  kg1IL

p/MPa 0.1317 0.2633 0.2709 0.4323 0.6163 0.8218 0.9088 1.0698 1.0856 1.2725 1.3293 0.1317 0.2633

T = 313.15 K 0.0283 ± 0.0015 0.0566 ± 0.0024 0.1200 ± 0.0043 0.2407 ± 0.0080 0.3962 ± 0.0130 0.4173 ± 0.0134 0.5655 ± 0.0179 0.7333 ± 0.0226 0.8011 ± 0.0251 0.8222 ± 0.0257 0.9235 ± 0.0288 0.9740 ± 0.0304 1.0956 ± 0.0341

0.1604 0.2577 0.8107 1.1342 1.6435 1.9243

T = 343.15 K 0.0600 ± 0.0029 0.0971 ± 0.0044 0.3168 ± 0.0137 0.4413 ± 0.0188 0.6446 ± 0.0273 0.7606 ± 0.0322

p/MPa

m/mol CO2  kg1IL

0.1118 0.2339 0.4580 0.7332 1.0241 1.4836 1.7345

T = 323.15 K 0.0530 ± 0.0024 0.1129 ± 0.0045 0.2268 ± 0.0085 0.3687 ± 0.0133 0.5154 ± 0.0184 0.7481 ± 0.0264 0.8839 ± 0.0310

0.1666 0.2682 0.5258 0.8484 1.1861 1.7205 1.9411

T = 353.15 K 0.0568 ± 0.0027 0.0919 ± 0.0043 0.1824 ± 0.0086 0.2953 ± 0.0139 0.4137 ± 0.0194 0.6039 ± 0.0284 0.6819 ± 0.0320

Standard uncertainties u are u(T) = 0.02 K and u(p) = 0.0010 MPa.

PA-33X pressure transmitter sensors in the range of 0 to 3 MPa and 0 to 4 MPa, respectively, which were accurate to within 0.01% of their full scale. The pressure sensors were calibrated against a pneumatic dead-weight gauge (DH. Budenberg model 550). The accurate pVT data presented by NIST for pure compounds [30] were used to calculate the compressibility factors. The average relative uncertainty associated with the measurement of the solubility (molality scale) of CO2 and H2S gases in the IL is u(m)/ m = 0.0375.

3. Results and discussion 3.1. Experimental results The experimental results (together with the uncertainties) for the solubility of the single gases carbon dioxide and hydrogen sulfide in [C2mim][eFAP] at temperatures of (303.15, 313.15, 323.15, 333.15, 343.15, 353.15) K and pressures up to about 2.0 MPa are summarized in tables 2 and 3. The reliability and accuracy of the

TABLE 3 Solubility of H2S in [C2mim][eFAP]: p, equilibrium pressure;m, molility of H2S; T, equilibrium temperature.a p/MPa

a

m/mol H2S  kg1IL

0.0604 0.1613 0.1645 0.2854 0.4025 0.7200 0.8905 1.2038 1.2326 1.3957

T = 303.15 K 0.0720 ± 0.0032 0.1998 ± 0.0054 0.2040 ± 0.0055 0.3714 ± 0.0083 0.5618 ± 0.0116 1.0884 ± 0.0205 1.3905 ± 0.0256 2.0552 ± 0.0369 2.1836 ± 0.0391 2.6158 ± 0.0465

0.0683 0.0801 0.1218 0.2103 0.2137 0.2302 0.4588 0.5185 0.6698 0.8472 0.8736 0.9464 1.1836 1.2994 1.4996 1.5891 1.6742

T = 333.15 K 0.0488 ± 0.0028 0.0573 ± 0.0031 0.0881 ± 0.0041 0.1551 ± 0.0063 0.1576 ± 0.0064 0.1705 ± 0.0069 0.3550 ± 0.0131 0.4174 ± 0.0151 0.5483 ± 0.0196 0.7013 ± 0.0248 0.7351 ± 0.0258 0.8089 ± 0.0284 1.0366 ± 0.0360 1.1917 ± 0.0413 1.3849 ± 0.0477 1.4999 ± 0.0516 1.6086 ± 0.0553

p/MPa

m/mol H2S  kg1IL

0.0582 0.1003 0.1764 0.3351 0.5642 0.6688 0.7546 0.9618 0.9656 1.0663 1.2507 1.3195

T = 313.15 K 0.0577 ± 0.0029 0.1008 ± 0.0039 0.1817 ± 0.0057 0.3633 ± 0.0097 0.6518 ± 0.0160 0.8137 ± 0.0196 0.9141 ± 0.0218 1.2426 ± 0.0291 1.2492 ± 0.0292 1.4185 ± 0.0330 1.7509 ± 0.0403 1.8783 ± 0.0431

0.0779 0.1383 0.2149 0.2627 0.2820 0.3855 0.5137 0.6040 0.7175 0.9384 0.9505 1.0850 1.3286 1.5882 1.8173

T = 343.15 K 0.0472 ± 0.0029 0.0847 ± 0.0044 0.1335 ± 0.0063 0.1645 ± 0.0076 0.1772 ± 0.0081 0.2498 ± 0.0110 0.3364 ± 0.0145 0.4014 ± 0.0171 0.4951 ± 0.0209 0.6630 ± 0.0277 0.6647 ± 0.0278 0.7695 ± 0.0320 0.9593 ± 0.0397 1.1952 ± 0.0419 1.4522 ± 0.0596

Standard uncertainties u are u(T) = 0.02 K and u(p) = 0.0010 MPa.

p/MPa

m/mol H2S  kg1IL

0.0657 0.1162 0.2016 0.2615 0.3654 0.4827 0.5526 0.6503 0.8498 0.9397 1.0328 1.3253

T = 323.15 K 0.0560 ± 0.0030 0.1004 ± 0.0042 0.1783 ± 0.0063 0.2348 ± 0.0079 0.3398 ± 0.0108 0.4571 ± 0.0140 0.5390 ± 0.0162 0.6356 ± 0.0189 0.8810 ± 0.0256 0.9889 ± 0.0286 1.1086 ± 0.0319 1.5096 ± 0.0429

0.0781 0.1376 0.2395 0.2587 0.2763 0.4630 0.5320 0.6414 0.8073 1.0074 1.1431 1.4042 1.9415

T = 353.15 K 0.0404 ± 0.0027 0.0722 ± 0.0042 0.1277 ± 0.0069 0.1382 ± 0.0074 0.1481 ± 0.0078 0.2562 ± 0.0130 0.3008 ± 0.0150 0.3672 ± 0.0183 0.4616 ± 0.0227 0.6057 ± 0.0296 0.7013 ± 0.0341 0.8993 ± 0.0436 1.3432 ± 0.0647

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2.0

increasing temperature. H2S is more soluble in that particular ionic liquid than CO2. For example, to dissolve 0.5 mole CO2 in one kilogram of [C2mim][eFAP] requires at T = (303 and 353) K pressures of 0.74 MPa and1.43 MPa, respectively, whereas dissolving the same amount of H2S requires only 0.37 MPa at 303 K and 0.97 MPa at 353 K.

1.8 1.6 1.4

3.2. Evaluation of Henry’s constant

p / MPa

1.2

0.4

The new experimental gas solubility data were evaluated to determine the corresponding Henry’s constant for the solubility of gas G in the ionic liquid IL at zero pressure on the molality scale ð0Þ ð0Þ kH;m . The fugacities f(T, p), needed for evaluation of kH;m values were calculated by applying the software package Thermofluids [31]. For CO2 that software is based on the equation of state by Span and Wagner [32] and for H2S that due to Lemmon and Span [33] is employed. The isothermal experimental results were correlated by a linear equation:

0.2

f ðT; pÞ m ¼ AðTÞðNÞ þ BðTÞðNÞ 0 : m=m0 m

1.0 0.8 0.6

0.0 0.0

0.2

0.4

0.6

0.8

1.0

1.2

1.4

mCO2/ mol CO2 · Kg-1 IL FIGURE 1. Experimental results for the solubility of CO2 in [C2mim][eFAP]: , T = 303.15 K; N, 313.15 K; j, 323.15 K; d, 333.15 K; +, 343.15 K; , 353.15 K; continuous lines, correlations from Pitzer’s model.

2.0 1.8 1.6 1.4

p / MPa

1.2 1.0 0.8 0.6 0.4 0.2

0.0 0.0

0.5

1.0

1.5

2.0

2.5

3.0

mH2S / mol H2S · Kg-1 IL FIGURE 2. Experimental results for the solubility of H2S in [C2mim][eFAP]: , T = 303.15 K; N, 313.15 K; j, 323.15 K; d, 333.15 K; +, 343.15 K; , 353.15 K; continuous lines, correlations from Pitzer’s model.

method of measurement have been checked in previous investigations [15–20]. The experimental results are also shown (as isothermal plots of solubility pressure versus molality of the dissolved gas, i.e., the amount of gas dissolved in one kilogram of ionic liquid) in figure 1 (CO2) and figure 2 (H2S). As expected the amount of dissolved gas increases with increasing pressure and decreases with

ð3Þ

Superscript (N) indicates the number of experimental data points N used for the evaluation. N was varied from Nmin = 5 (i.e., the 5 data points with the lowest pressures were used) to Nmax = Nexp in steps of one (where Nexp is the number of experimental data points on the isotherm). The results for A(T)(N) were averaged to ð0Þ give kH;m and the standard deviation of A(T)(N) was used to estimate ð0Þ the uncertainty DkH;m of the results for Henry’s constant. The resulting numbers are given in table 4. The influence of temperature on Henry’s constant of CO2 and H2S in [C2mim][eFAP] was described by equations (4a) and (4b) with a standard deviation of 0.6% and with maximum deviations of 1.3% for CO2 and 1.8% for H2S.

. 1272:0 ð0Þ ln kH;m MPa ¼ 4:6914  ðT=KÞ

CO2 ;

ð4aÞ

. 1777:6 ð0Þ ln kH;m MPa ¼ 5:6928  ðT=KÞ

H2 S:

ð4bÞ

The correlation equations (4a) and (4b) were used to estimate the changes of the molar Gibbs free energy Dsol G1 m , of the molar 1 enthalpy Dsol H1 m and of the molar entropy Dsol Sm of gas G when it is transferred from the ideal gas state at temperature T and standard pressure p = p0 = 0.1 MPa to its reference state in the liquid solvent (i.e., a one molal solution of the gas in the particular ionic liquid at temperature T where the dissolved gas experiences the same interactions as in infinite dilution). These properties are also given in table 4. As expected, Dsol G1 m is positive and increases 1 with temperature and Dsol H1 m and Dsol Sm are negative. The accuracy of the experimental data does not allow determining an influence of temperature on those properties. The Dsol H1 m is more negative for H2S than for CO2. That result is in accordance with the explanations presented in the previous paragraph, the attractive interaction via hydrogen bonding of H2S with [eFAP] anion results in more negative values for Dsol H1 m when compared to the weaker attractive interaction between CO2 and [eFAP] as CO2 is a weaker Lewis base than H2S. The entropy changes reveal that more ordering occurs with H2S than does with CO2 in [C2mim][eFAP]. 3.3. Comparison with literature data Almantariotis et al. [7] reported experimental results for the solubility of CO2 in [C2mim][eFAP] (as Henry’s constant on the mole ð0Þ fraction scale kH;x ) at temperatures between 303 K and 343 K. By

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TABLE 4 ð0Þ Thermodynamic properties for the solubility of CO2 and H2S in [C2mim][eFAP]: T, temperature; kH;m , Henry’s constant of gas G at zero pressure on the molality scale; Dsol G1 m, 1 changes of the molar Gibbs free energy; Dsol H1 m , changes of the molar enthalpy; Dsol Sm , changes of the molar entropy of gas G. T/K

ðkH;m  DkH;m Þ/MPa

1 Dsol G1 m =kJ  mol

303.15 313.15 323.15 333.15 343.15 353.15

1.623 ± 0.011 1.891 ± 0.003 2.130 ± 0.014 2.426 ± 0.006 2.679 ± 0.020 2.941 ± 0.017

7.052 7.634 8.215 8.797 9.378 9.960

303.15 313.15 323.15 333.15 343.15 353.15

0.850 ± 0.012 1.022 ± 0.004 1.190 ± 0.005 1.420 ± 0.007 1.671 ± 0.009 1.952 ± 0.010

5.372 6.037 6.702 7.366 8.031 8.695

ð0Þ

ð0Þ

1 Dsol H1 m =kJ  mol

1 Dsol S1  K1 m =J  mol

10.6

58.1

14.8

66.5

CO2

H2S

ð0Þ

ð0Þ

using kH;m ¼ kH;x  M IL =1000 relation, with MIL = 0.55617 kg  mol1, the results of both investigations can be compared. That comparison is shown in figure 3a. At around T = 303 K, the results for the Henry’s constant of CO2 in [C2mim][eFAP] from both investigations agree within 1%, however, with increasing temperature the difference increases and at around 343 K the number reported by Almantariotis et al. [7] for Henry’s constant of CO2 in [C2mim][eFAP] is about 20% higher than the results of the present 1 investigation. As a consequence the numbers for Dsol H1 m ¼ Dsol H x 1 1 and Dsol Sm ¼ Dsol Sx þ R lnðM IL =1000Þ from both investigations reveal rather large differences (Almantariotis et al. [7]: 1 1 Dsol H1 , Dsol S1 m ðCO2 Þ = (70 + 2) J  K m ðCO2 Þ = (14 + 2) kJ  mol 1 1 1 1  mol ); here: Dsol Hm ðCO2 Þ = 10.6 kJ  mol , Dsol Sm ðCO2 Þ = 58.2 J  K1  mol1). Also, comparison shows that the Henry’s constant of CO2 in [C2mim][eFAP] reported by Zhang et al. [9] (estimated by COSMO-RS theoretical approach) is about 44% higher than the value obtained by extrapolating equation (8a) of this work to T = 298.15 K. Figure 3a also shows the Henry’s law constants for CO2 in other [C2mim]+-based ILs with different anions as a function of temperature. The solubility of CO2 in the ILs follows the order: [C2mim][BF4] < [C2mim][EtSO4] < [C2mim][MDEGSO4] < [C2mim] [OTf] < [C2mim][Tf2N] < [C2mim][eFAP], where [MDEGSO4] stands for 2-(2-methoxyethoxy) ethylsulfate, i.e., the solubility increases with increasing number of –CF3 groups in the anion of the IL. This observation was already made by Baltus et al. [38]. Figure 3b shows a similar comparison for Henry’s constant of H2S in various [C2mim]+-based ILs. The solubility of H2S increases in the order: [C2mim][EtSO4] < [C2mim][PF6] < [C2mim][Tf2N]  [C2mim][eFAP], i.e., it follows the same order as that of CO2. Obviously the solubility of CO2 and H2S gases in ILs is governed by similar mechanism. With increasing number of –CF3 groups in the anion the molar density of the IL decreases (the void volume increases), which results in a higher solubility of the gas [22]. The higher affinity of ILs towards H2S compared to CO2 can be ascribed to a stronger intermolecular attraction through hydrogen bonding between H2S acting as hydrogen-bond donor and the anion of the ILs acting as hydrogen-bond acceptor [12,22]. This type of intermolecular attraction is replaced by the weaker Lewis acid-base interactions in case of CO2, with the carbon atom of CO2 as the electron acceptor (acid) and the anion as the donor (base) [39].

4.1. Extended Pitzer’s model When the saturation pressure of the ionic liquid can be neglected the solubility of a single gas in a pure ionic liquid can be described by the extended Henry’s law ð0Þ

kH;m ðTÞ exp

Two models are used to correlate the experimental results for the solubility of a single gas (CO2 or H2S) in [C2mim][eFAP].



v 1 p RT

m c ¼ f ðT; pÞ; m0 m

ð5Þ

where v 1 and cm are the partial molar volume of gas G at infinite dilution in the particular ionic liquid and the activity coefficient of that gas on the molality scale. As the maximum pressure of the new experimental solubility data is below 2 MPa and no information on v 1 is available the exponential term in equation (5), i.e., the Krichevsky–Kasarnovsky correction [40] to Henry’s constant, is neglected. The activity coefficient cm is calculated from a simplified version of Pitzer’s virial expansion for the excess Gibbs free energy on the molality scale [28,29]:

ln cm ¼ 2 

 m 2 m bþ3  l; 0 m m0

ð6Þ

where b and l are binary and ternary parameters, respectively describing interactions between gas molecules in the ionic liquid. These parameters depend on temperature and the particular ionic liquid that is the solvent for the gas. The influence of temperature on such an interaction parameter is approximated by

X ¼ AX þ

BX ; ðT=KÞ

ð7Þ

where X represents b or l. The experimental results for temperature and pressure were used to calculate the fugacity of a pure gas as described above. Parameters AX and BX were adjusted to minimize the sum of squared differences between the calculated and the experimental results for the fugacity of a gas. For all those calculations, the correlations for Henry’s constants given by equations (4a) and (4b) were used. The results of the correlations (interaction parameters) are given in table 5. The quality of the correlation is discussed here using the average relative deviation ARD% and the maximum relative deviation MRD% in the molality of gas G

ARD % ¼ 4. Modeling



 N  cor   mexp 100 X i ðT; pÞ mi ðT; pÞ exp ; N i¼1  mi ðT; pÞ

  cor  m ðT; pÞ  mexp  i ðT; pÞ MRD % ¼ max  i  100 : exp  m ðT; pÞ i

ð8aÞ

ð8bÞ

60

A.H. Jalili et al. / J. Chem. Thermodynamics 67 (2013) 55–62

30

(a)

[C2mim][EtSO4], ref [20] [C2mim][Tf2N], ref[34]

25

[C2mim][eFAP] this work [C2mim][OTf], ref[35] [C2mim][MDEGSO4], ref[36]

kH,m(0) / MPa

20

[C2mim][BF4], ref[37] [C2mim][eFAP], ref[7]

15

10

5

0

300

320

340

360

T/K 14

(b)

[C2mim][EtSO4], ref[20]

12

[C2mim][Tf2N], ref[21] [C2mim][eFAP], this work

kH,m(0) / MPa

10

[C2mim][PF6], ref[21]

8 6 4 2 0 300

320

340 T/K

360

380

FIGURE 3. Comparison of Henry’s law constants as a function of temperature for [C2mim]+-based ILs with different anions: (a) CO2 in ILs; (b) H2S in ILs; continuous lines, correlations from equations (4a) and (4b) (See above-mentioned references for further information [34-37]).

TABLE 5 BX Interaction parameters for the Pitzer model, X ¼ AX þ ðT=KÞ : T, temperature; X and AX and BX, parameters of equation (7). System

Parameter X

l

b

CO2 + [C2mim][eFAP] H2S + [C2mim][eFAP]

AX

BX

AX

BX

0.0142 0.3170

40.44 45.46

0.2882 0.0887

79.71 21.96

The values of ARD% and MRD% calculated from equations (8a) and (8b) together with the corresponding experimental uncertainties ARD%exp and MRD%exp are given in table 6. Here, ARD%exp and MRD%exp are respectively the average and maximum of the relative experimental uncertainty of the molality of gas G, ðDmexp Þi ¼ 100  ðDexp mðT; pÞ=mðT; pÞexp Þi . For both gases ARD% and MRD% are smaller than ARD%exp and MRD%exp, i.e., the correlation represents the new gas solubility data within experimental uncertainty.

61

A.H. Jalili et al. / J. Chem. Thermodynamics 67 (2013) 55–62 TABLE 6 Comparison between accuracy of the two models to correlate the experimental solubility data: ARD%, average relative deviation of calculated from experimental molality data; MRD%, maximum relative deviation of calculated from experimental molality data. System

CO2 + [C2mim][eFAP] H2S + [C2mim][eFAP]



Exp. results

Pitzer s model

RK model

ARD%

MRD%

ARD%

MRD%

ARD%

MRD%

1.0 2.5

1.9 4.5

0.7 1.2

1.8 4.4

1.8 1.4

5.6 3.8

4.2. The RK cubic EoS Shiflett and Yokozeki [24–26] proposed to apply a generic Redlich–Kwong type of cubic equation of state (RK EoS) to correlate the solubility of gases in an ionic liquid. Here, the mathematical form of a(T) is chosen the same as used by Shiflett and Yokozeki [24–26]:

aðTÞ ¼

63 X k kk  ðT 1 r  TrÞ ;

ð9Þ

TABLE 8 Binary interaction parameters l12, l21, s12, s21, m12, m21 of the RK EoS in equations (11)–(13). System

l12

l21

s12 = s21 (K)

m12 = m21

CO2(1) + [C2mim][eFAP](2) H2S(1) + [C2mim][eFAP](2) CO2 + H2S

0.12796 0.21689 0.04014

0.15663 0.23572 2.7134

5.91625 71.5385 23.37

0.100095 0.252904 0.0

is discussed here using the average relative deviation ARD% and the maximum relative deviation MRD% in the molality of gas G. The results are also shown in table 6. The correlation with the RK EoS for {H2S + [C2mim][eFAP]} (the new gas solubility data) shows about the same quality as the extended Pitzer’s model. However, for {CO2 + [C2mim][eFAP]} the correlation with the RK EoS results in deviations which are somewhat larger than the experimental uncertainties and it is less accurate than the correlation with the extended Pitzer model. 5. CO2/H2S separation

k¼0

where Tr = T/Tc is the reduced temperature, and kk,s are adjustable parameters. The critical properties of the pure compounds were either taken from the NIST database [30] (for CO2 and H2S) or estimated by the modified Lydersen–Joback–Reid method [41], proposed by Valderrama and coworkers [42,43] (for [C2mim][eFAP]). For CO2 and H2S the parameters k0 through k3 were taken from Shiflett and Yokozeki [14] whereas for [C2mim][eFAP] they were considered as adjustable parameters and either set to zero (k2 and k3) or fit (together with the binary parameters of the model – see below) to the new gas solubility data. All pure component parameters are given in table 7. The EoS was extended to mixtures by applying the modified van der Waals–Berthelot mixing rule proposed by Yokozeki [44]:

aðTÞ ¼

N X ðai aj Þ1=2 fij ðTÞð1  kij Þxi xj ;

ð10Þ

i;j¼1



N 1X ðbi þ bj Þð1  mij Þð1  kij Þxi xj ; 2 i;j¼1

fij ðTÞ ¼ 1 þ

kij ¼

sij T

ð11Þ

ð12Þ

;

lij lji ðxi þ xj Þ ; lji xi þ lij xj

ð13Þ

wheresij = sji, sii = 0, mij = mji, mii = 0 and kii = 0. There are four interaction parameters per binary system: lij, lji, mij, and sij, which together with the pure component parameters (k0 and k1) of [C2mim][eFAP] were adjusted to the new gas solubility data. The numerical values of the binary interaction parameters are given in table 8. For preset temperature and pressure the composition of the liquid phase was calculated and the quality of the correlation

As in a previous publication [22], RK EoS was applied to discuss the separation of CO2 and H2S by absorption in an ionic liquid (here in [C2mim][eFAP]). The binary interaction parameters, lCO2 ;H2 S , lH2 S ;CO2 , mCO2 ;H2 S and sCO2 ;H2 S for interactions between CO2 and H2S were adopted from Shiflett and Yokozeki [14] and are given in table 8. The selectivity aCO2/H2S [14,45,46]

aCO2 =H2 S ¼

yCO2 =xCO2 yH2 S =xH2 S

ð14Þ

was calculated. In all investigated cases the selectivity is above one (as [C2mim][eFAP] is a better solvent for H2S than for CO2). The influence of the mole fraction of [C2mim][eFAP] on the selectivity is small as long as its mole fraction is below about 0.8. For example at T = 303.15 K and p = 0.1 MP the selectivity shows a very slight decrease from 1.82 to 1.78 when the CO2 mole fraction in the feed increases from 0.2 to 0.8. The selectivity without the absorbent as a function of pressure at 303.15 K for the CO2 mole fractions xCO2 = 0.8, 0.5, and 0.2 was calculated in our previous study [22]. Comparison with previous study [22] reveals that there is a slight selectivity enhancement due to the addition of [C2mim][eFAP] ionic liquid only for the CO2 mole fraction of 0.8 and at high pressures. When the CO2 mole fraction decreases from 0.8 to 0.5 to 0.2, the improvement in selectivity caused by the ionic liquid also decreases. Also, the selectivity decreases from 1.82 to 1.63 with increasing temperature from (303 to 343) K, i.e., a 40 K increase in temperature causes more than 10% decrease in selectivity. This behaviour exists at all CO2 mole fractions and all pressures from (0.1 to 1.0) MPa. In addition, the selectivity increases slightly by increasing pressure, i.e. the selectivity enhancement is about 6% in going from p = 0.1 MPa (aCO2 =H2 S ¼ 1:82) to p = 0.5 MPa (aCO2 =H2 S ¼ 1:93). This trend can be observed for all CO2 mole fractions. The most important fact to consider is that at higher temperatures, e.g., 343.15 K and high CO2 mole fractions (0.5 and 0.8) without the ionic liquid, no vapour-liquid equilibria exists [14] and the gas separation cannot be accomplished using traditional

TABLE 7 Pure Compound Constants Used for the RK EoS: Tc, critical temperature; Pc, critical pressure; k0, k1, k2 and k3, pure component parameters used for the RK EoS in equation (9). Compound

Molar mass/kg  mol1

Tc/K

Pc/kPa

k0

k1

k2

k3

CO2a H2Sa

0.04401 0.03408 0.55617

304.13 373.10 830.67

7377 9000 10030

1.00049 0.99879 1.00000

0.43866 0.33206 0.69896

0.10498 0.04942 0.00000

0.06250 0.004639 0.00000

[C2mim][eFAP] a

The values were taken from reference [14].

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A.H. Jalili et al. / J. Chem. Thermodynamics 67 (2013) 55–62

distillation methods. However, the ionic liquid allows separating of CO2 from H2S. A comparison with previous results [14,22] shows that [C2mim][eFAP] is much less effective than [C8mim][Tf2N] and [C4mim][PF6] for the separation of CO2 and H2S gases from each other in gaseous streams, but more effective for the removal of these two acid gases from natural gas. 6. Conclusions New experimental results are presented in this work for the solubility of CO2 and H2S in an ionic liquid with an perfluorinated anion, e.g., the [C2mim][eFAP] IL. Results show that [C2mim][eFAP] has more affinity for dissolution of CO2 and H2S compared to other [C2mim]-based ILs, except [C2mim][OAc] (where [OAc] stands for ethanoate or acetate anion), which shows exceptionally ultrahigh dissolving power for CO2 and H2S. The solubility of H2S is higher than that of CO2 in [C2mim][eFAP] and the solubility of both gases in the IL studied in this work is of a physical nature. Also the results obtained show that as the number of –CF3 groups increases in the anion part of the IL the solubility of acid gases CO2 and H2S increases too. The solubility behaviour of CO2 + H2S + [C2mim][eFAP] ternary system were predicted by using the RK EoS and the parameters for CO2 + [C2mim][eFAP] and H2S + [C2mim][eFAP] binary systems to investigate the feasibility of separation of CO2 and H2S from each other especially at high temperatures where the traditional distillation procedure fails to work at these conditions. It was found that [C2mim][eFAP] has much lower capability compared with [C8mim][Tf2N] and [C4mim][PF6] for separation of CO2 and H2S from each other. Acknowledgement We are thankful to the Research Council of the Research Institute of Petroleum Industry (RIPI) for their support of this work. References [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] E.D. Bates, R.D. Mayton, I. Ntai, J.H. Davis, J. Am. Chem. Soc. 124 (2002) 926– 927. [4] J.S. Wilkes, P. Wassercheid, T. Welton (Eds.), Ionic Liquids in Synthesis; Wiley VCH Verlag, 2002. [5] N.V. Ignat’ev, U. Welz-Biermann, A. Kucheryna, G. Bissky, H. Willner, J. Fluorine Chem. 126 (2005) 1150–1159. [6] M.J. Muldoon, S.N.V.K. Aki, J.L. Anderson, J.K. Dixon, J.F. Brennecke, J. Phys. Chem. B 111 (2007) 9001–9009. [7] D. Almantariotis, S. Stevanovic, O. Fandino, A.S. Pensado, A.A.H. Padua, J.-Y. Coxam, M.F. Costa Gomes, J. Phys. Chem. B 116 (2012) 7728–7738.

[8] J. Blath, J. Christ, N. Deubler, T. Hirth, T. Schiestel, Chem. Eng. J. 172 (2011) 167–176. [9] X. Zhang, Z. Liu, W. Wang, AIChE J. 54 (2008) 2717–2728. [10] X. Zhang, F. Huo, Z. Lio, W. Wang, W. Shi, E.J. Maginn, J. Phys. Chem. B 113 (2009) 7591–7598. [11] F.-Y. Jou, A.E. Mather, Int. J. Thermophys. 28 (2007) 490–495. [12] C.S. Pomelli, C. Chiappe, A. Vidis, G. Laurenczy, P.J. Dyson, J. Phys. Chem. B 111 (2007) 13014–13019. [13] Y.J. Heintz, L. Sehabiague, B.I. Morsi, K.L. Jones, D.R. Luebke, H.W. Pennline, Energy Fuels 23 (2009) 4822–4830. [14] M.B. Shiflett, A. Yokozeki, Fluid Phase Equilibr. 294 (2010) 105–113. [15] A.H. Jalili, M. Rahmati-Rostami, C. Ghotbi, M. Hosseini-Jenab, A.N. Ahmadi, J. Chem. Eng. Data 54 (2009) 1844–1849. [16] M. Rahmati-Rostami, C. Ghotbi, M. Hosseini-Jenab, A.N. Ahmadi, A.H. Jalili, J. Chem. Thermodyn. 41 (2009) 1052–1055. [17] H. Sakhaeinia, V. Taghikhani, A.H. Jalili, A. Mehdizadeh, A.A. Safekordi, Fluid Phase Equilibr. 298 (2010) 303–309. [18] A.H. Jalili, A. Mehdizadeh, M. Shokouhi, H. Sakhaeinia, V. Taghikhani, J. Chem. Thermodyn. 42 (2010) 787–791. [19] M. Shokouhi, M. Adibi, A.H. Jalili, M. Hosseini-Jenab, A. Mehdizadeh, J. Chem. Eng. Data 55 (2010) 1663–1668. [20] A.H. Jalili, A. Mehdizadeh, M. Shokouhi, A.N. Ahmadi, M. Hosseini-Jenab, F. Fateminassab, J. Chem. Thermodyn. 42 (2010) 1298–1303. [21] H. Sakhaeinia, A.H. Jalili, V. Taghikhani, A.A. Safekordi, J. Chem. Eng. Data 55 (2010) 5839–5845. [22] A.H. Jalili, M. Safavi, C. Ghotbi, A. Mehdizadeh, M. Hosseini-Jenab, V. Taghikhani, J. Phys. Chem. B 116 (2012) 2758–2774. [23] M. Safavi, C. Ghotbi, V. Taghikhani, A.H. Jalili, A. Mehdizadeh, J. Chem. Thermodyn. 65 (2013) 220–232. [24] M.B. Shiflett, A. Yokozeki, J. Chem. Eng. Data 44 (2005) 4453–4464. [25] M.B. Shiflett, A. Yokozeki, J. Phys. Chem. B 111 (2007) 2070–2074. [26] M.B. Shiflett, A. Yokozeki, J. Chem. Eng. Data 51 (2006) 1931–1939. [27] A. Yokozeki, M.B. Shiflett, C.P. Junk, L.M. Grieco, T. Foo, J. Phys. Chem. B 112 (2008) 16654–16663. [28] K.S. Pitzer, Activity Coefficients in Electrolyte Solution, CRC Press, Boca Raton, FL, 1991. [29] K.S. Pitzer, J. Phys. Chem. 77 (1973) 268–277. [30] NIST Scientific and Technical Databases, Thermophysical Properties of Fluid Systems. http://webbook.nist.gov/chemistry/fluid/ (accessed March 2013). [31] W. Wagner, U. Overhoff, ThermoFluids (Version 1.0, Build 1.0.0), Springer: Berlin, Heidelberg (Germany), 2006. [32] R. Span, W. Wagner, J. Phys. Chem. Ref. Data 25 (1996) 1509–1596. [33] E.W. Lemmon, R. Span, J. Chem. Eng. Data 51 (2006) 785–850. [34] D. Camper, C. Becker, C. Koval, R. Noble, Ind. Eng. Chem. Res. 45 (2006) 445– 450. [35] A.N. Soriano, B.T. Doma, M.-H. Li, J. Chem. Thermodyn. 41 (2009) 525–529. [36] A.N. Soriano, B.T. Doma, M.-H. Li, J. Chem. Thermodyn. 40 (2008) 1654–1660. [37] A.N. Soriano, B.T. Doma, M.-H. Li, J. Chem. Eng. Data 53 (2008) 2550–2555. [38] R.E. Baltus, B.H. Culbertson, S. Dai, H. Luo, D.W. DePaoli, J. Phys. Chem. B 108 (2004) 721–727. [39] B.L. Bhargava, S. Balasubramanian, Chem. Phys. Lett. 444 (2007) 242–246. [40] I.R. Krichevsky, J.S. Kasarnovsky, J. Am. Chem. Soc. 57 (1935) 2168–2171. [41] B.E. Poling, J.M. Prausnitz, J.P. O’Connell, The Properties of Gases and Liquids, fifth ed., McGraw-Hill, New York, 2001. [42] J.O. Valderrama, P.A. Robles, Ind. Eng. Chem. Res. 46 (2007) 1338–1344. [43] J.O. Valderrama, W.W. Sanga, J.A. Lazzus, Ind. Eng. Chem. Res. 47 (2008) 1318– 1330. [44] A. Yokozeki, Int. J. Thermophys. 22 (2001) 1057–1071. [45] A. Yokozeki, M.B. Shiflett, Appl. Energy 84 (2007) 351–361. [46] A. Yokozeki, M.B. Shiflett, Ind. Eng. Chem. Res. 47 (2008) 8389–8395.

JCT 13-351