Experimental Study on the Solubility of Carbon Dioxide in Nitrate and

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Experimental Study on the Solubility of Carbon Dioxide in Nitrate and Thiocyanate-Based Ionic Liquids Babak Mokhtarani,* Abolfazl Negar Khatun, Morteza Mafi, Ali Sharifi, and Mojtaba Mirzaei Chemistry and Chemical Engineering Research Center of Iran, P.O. Box 14335-186, Tehran, Iran S Supporting Information *

ABSTRACT: New experimental results are reported for the solubility of carbon dioxide (CO2) in nitrate- and thiocyanatebased ionic liquids (ILs) at temperatures ranging from 298.15 to 333.15 K and pressure up to 4.5 MPa. The studied ILs are 1methyl 3-octylimidazolium thiocyanate [Omim][SCN], 1-methyl 3-hexylimidazolium thiocyanate [Hmim][SCN], 1-methyl 3octylimidazolium nitrate [Omim][NO3], and 1-butyl 3-methylimidazolium tetrafluoroborate [Bmim][BF4]. The solubility measurements are performed in a known volume stainless steel equilibrium cell. The experimental data indicated the solubility of CO2 decreases with increasing of temperature. Henry’s constant are calculated from the solubility data. The experimental results for CO2 solubility are correlated with the extended Henry’s law and Pitzer virial expansion using binary parameters. The correlation results of gas solubility are agreed with the experimental data.

1. INTRODUCTION

Most of the experimental data in literature shows the superiority of ILs over various conventional solvents.16 This research is the continuation of our previous work to investigate the suitability of nitrate and thiocyanate based ILs in separation process.17−19 In this work, new experimental data for the solubility of CO 2 in 1-methyl 3-octylimidazolium thiocyanate [Omim][SCN], 1-methyl 3-hexylimidazolium thiocyanate [Hmim][SCN], 1-methyl 3-octylimidazolium nitrate [Omim][NO3] and 1-butyl 3-methylimidazolium tetrafluoroborate [Bmim][BF4] are reported at T = 298.15− 333.15 K and pressure up to 4.5 MPa. The experimental data are reported for the first time. The apparatus with an equilibrium cell is used to measure the solubility of CO2 in ILs. Henry’s constant of CO2 in ILs are calculated. The experimental results are correlated with the extended Henry’s law and Pitzer’s virial expansion model.

The concentration of greenhouse gases (e.g., CO2, CH4, and NOx) in atmosphere is increasing dramatically in recent years. Carbon dioxide (CO2) is one the most important greenhouse gases.1 Industrial activities and the fossil fuels burning are two main sources of this gas. Removal of CO2 from flue gas is a major problem for environment and several efforts have been performed to reduce the level of greenhouse gases. Amine compounds (e.g., monoethanolamine (MEA)) are common solvents for absorption of CO2 from flue gas in industries. These compounds have several disadvantages such as much solvent loss during the regeneration, high energy required for regeneration, high operational cost, as well as corrosive amine compounds.2 Ionic liquids (ILs) may be applied as new solvent for CO2 absorption. These materials are liquid at room temperature, have low vapor pressure, are nonflammable, are chemically stabile at high temperatures, and have excellent solubility for organic and inorganic compounds.3 Application of these solvent for CO2 absorption have been studied in recent years. First, Blanchard et al. studied the solubility of CO2 in different types of ILs in 1999.4,5 They found that the anion of ILs has a great effect on the CO2 solubility. Sudhir Aki et al.6 reported experimental CO2 solubility data in ILs with various kinds of anions at different temperature and pressure. According to their results, the solubility of ILs decreases with increasing temperature and increases with increasing pressure.6 Kumelan et al.7−10 studied the solubility of different single gases in ILs. In recent years, the amount of experimental data in literature for these systems has considerably increased.11−15 © XXXX American Chemical Society

2. EXPERIMENTAL SECTION 2.1. Chemicals. The suppliers and the purities of the chemicals are reported in Table 1. The studied ILs were synthesized and prepared in our laboratory. [Omim][SCN] and [Hmim][SCN] were prepared by an anion exchange reaction using [Omim][Cl], [Hmim][Cl], and KSCN.20,21 [Omim][Cl] was synthesized according to the method described in literature.22 The preparation methods were explained in our previous works.17,23 The Received: October 22, 2015 Accepted: February 1, 2016

A

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Table 1. Purities and Suppliers of the Chemicals chemical name

supplier

carbon dioxide [Omim][SCN] [Hmim][SCN] [Bmim][BF4] [Omim][NO3]

Roham gas, Iran synthesized in lab synthesized in lab

mass fraction purity 0.9995 >0.98 >0.98 >0.98 >0.98

mass fraction of water 0 7 1 2 1

× × × ×

molecular weight (g/mol) 44 253.4 225.4 226.0 257.3

10−4 10−3 10−3 10−3

purification method NMR, Karl Fischer titration, and potentiometric titration

Figure 1. Schematic diagram of the experimental setup: (1) water bath, (2) equilibrium cell, (3) circulator, (4) Vacuum pump, (5) gas container, (6) control panel, (7) monitor, (8) pressure transducer, (9) temperature transducer, and (10) mixer.

range of 273.15−373.15 K with the standard uncertainty of 0.1 K. The equilibrium cell is degassed using a vacuum pump and CO2 is injected into it. CO2 gas is injected from the main capsule and the mass of CO2 is measured by a laboratory balance with the standard uncertainty of 0.01 g. The pressure of equilibrium cell is determined by a pressure transducer (Keller PPA-33X, Switzerland) connected to a personal computer. The standard uncertainty of the pressure transducer is 0.01% of total pressure. The variation of pressure in the cell is monitored and at the equilibrium condition the pressure change of the cell is zero. The volume of equilibrium cell is precisely measured (169.6 ± 1 cm3) and subtracted from the volume of IL to obtain the volume of gas phase. Increasing the liquid phase volume due to the solubility of CO2 is small and can be neglected (1). The density of IL is measured using an Anton Paar DMA-5000 digital densitometer with standard uncertainty of 0.002 g·cm−3. At the equilibrium pressure, gas phase contains pure CO2 because no IL is vaporized while the liquid phase consists of IL and CO2.The mass of CO2 in gas phase is calculated by the Span−Wagner equation of state.28 From the

structures of the ILs were checked with nuclear magnetic resonance (NMR) spectroscopy and the NMR spectrum of the ILs are shown in the Supporting Information. The concentration of chloride anion as impurities are measured for studied ILs because they are prepared from 1-chloro hexane or 1-chloro octane. The mass fraction of chloride ion is determined by the potentiometric titration method with silver nitrate (AgNO3) and is less than 1 × 10−4. The purification methods for synthesized ILs are given in Table 1. The ILs were dried and degassed for 24 h at 343.15 K under vacuum and were kept under argon gas. The water content of the ILs was determined by a 684 Karl Fischer coulometer and reported in Table 1. 2.2. Apparatus and Methods. The experimental apparatus for gas solubility measurement is schematically represented in Figure 1. The apparatus consists of a stainless steel equilibrium cell with known volume. A definite mass of IL is injected into equilibrium cell and the cell is inserted into a water bath for temperature control (Julabo model FP50, Germany). The water bath is controlled the temperature in the B

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total mass of CO2 injected into equilibrium cell and the mass of CO2 in gas phase, the mass of CO2 in IL phase is calculated. From the mass of IL and the mass of CO2 in IL phase, the molality of CO2 is calculated. The standard uncertainty for measurement of the CO2 solubility by this method is less than 0.05 mol·kg−1. 2.3. Experimental Results. The densities of studied ILs at different temperatures are reported in Table 2 and compared

Table 3. Experimental Molality (mCO2) and Fugacity (f CO2) of CO2 in ILsa T/K 298.15

Table 2. Density (g·cm−3) of Studied ILs and Comparison with Literature Data at Different Temperature and 88 kPaa IL

298.15 K

[Omim][SCN]17 [Hmim][SCN]

1.00827 1.03621 [1.05692]24 1.20271 [1.20074]25 [1.20129]26 [1.2011]27 1.06386

[Bmim][BF4]

[Omim][NO3] a

313.15 K

333.15 K

0.99970 1.02755

0.98841 1.01612

1.19229 [1.19035]25 [1.19068]26 [1.1901]27 1.05455

313.15

333.15

1.17873 [1.17710]25 [1.17670]26 [1.1753]27 1.04242

Standard uncertainty, density =0.002 g·cm−3, T = 0.1 K, P = 2 kPa

298.15

with literature data. According to the Table 2, the comparison of density data for [Hmim][SCN] at 298.15 K shows a big discrepancy. This difference may be due to the water content and other impurities that existed in IL. The solubility of CO2 in [Hmim][SCN], [Omim][SCN], and [Omim][NO3] are reported at a temperature range from 298.15 to 333.15 K and pressure up to 4.5 MPa in Table 3. Each experiment was repeated three times and the average is reported. The fugacities of CO2 are calculated by the Span−Wagner equation of state.28 The experimental results indicated the solubility of CO2 in [Omim][SCN] is more than [Hmim][SCN]. This implies that increasing the alkyl chain length of ILs leads to the increasing of the CO2 solubility. These results are consistent with the previous studies on CO2 solubility in different types of ILs.5,6,29,30 In addition, the solubility of CO2 in [Omim][NO3] are higher than [Omim][SCN] and [Hmim][SCN]. This behavior indicates the ILs with nitrate anion is more favorable than ILs with thiocyanate. The experimental standard uncertainties for CO2 molality are also reported in Table 3. The experimental standard uncertainty for CO2 solubility was performed by calculating the error of the measurements. The errors of the measurements of CO2 solubility contains the errors in the measurement of the mass of CO2 (δwCO2) and IL (δwIL), the pressure transducer error (δp), the error of the measurement of the equilibrium cell volume (δV), the temperature error (δT), and the error of the measurement of the IL density (δpIL). The relation for measurement the CO2 solubility may be written as follows mCO2 =

nCO2 w ( 1000 ) IL

=

⎛ wCO2 ⎞ ⎜ ⎟ − ⎝ MWCO2 ⎠ w ( 1000 )

313.15

333.15

298.15

313.15

PVg

333.15

(ZRT )

IL

(1)

Where mCO2 is the CO2 molality, nCO2 and wIL are the number of moles of CO2 in the IL phase and mass of IL, wCO2 and MWCO2 are the total mass of CO2 injected to the equilibrium cell and molecular weight of CO2, and P, Z, R, and T are pressure, compressibility factor of CO2 gas, gas constant, and temperature, respectively. Vg is the volume of gas phase in the

mCO2/(mol·kg−1) 0.2338 0.5352 0.9353 1.3346 2.3912 0.2541 0.4790 0.7721 1.0874 1.4123 0.2701 0.4452 0.6711 0.9079 1.2678

± ± ± ± ± ± ± ± ± ± ± ± ± ± ±

0.0223 0.0243 0.0280 0.0339 0.0589 0.0220 0.0242 0.0285 0.0352 0.0436 0.0213 0.0232 0.0269 0.0327 0.0420

0.5200 0.8046 1.2333 1.7209 2.1070 0.3667 0.6302 0.9931 1.3329 1.7356 0.2906 0.4759 0.7122 0.9712 1.1821 1.4877

± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ±

0.0239 0.0264 0.0310 0.0387 0.0463 0.0226 0.0253 0.0309 0.0373 0.0465 0.0215 0.0236 0.0275 0.0328 0.0390 0.0474

0.2082 0.4327 0.7259 1.0430 1.4210 1.8818 2.3572 2.7097 0.2418 0.5516 0.8597 1.2219 1.6627 2.1616 0.1938 0.4510 0.6884 0.9756 1.2589 1.6111

± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ±

0.0211 0.0219 0.0234 0.0255 0.0287 0.0337 0.0394 0.0448 0.0205 0.0221 0.0247 0.0285 0.0341 0.0418 0.0195 0.0209 0.0232 0.0267 0.0312 0.0373

P/MPa [Hmim][SCN] 0.450 ± 0.027 1.027 ± 0.036 1.714 ± 0.046 2.479 ± 0.058 4.521 ± 0.089 0.708 ± 0.031 1.288 ± 0.040 2.034 ± 0.051 2.890 ± 0.064 3.737 ± 0.077 0.856 ± 0.033 1.388 ± 0.041 2.089 ± 0.052 2.925 ± 0.064 3.986 ± 0.081 [Omim][SCN] 0.994 ± 0.032 1.479 ± 0.038 2.156 ± 0.047 2.998 ± 0.057 3.658 ± 0.066 0.863 ± 0.031 1.471 ± 0.038 2.335 ± 0.049 3.087 ± 0.058 3.962 ± 0.069 0.915 ± 0.031 1.450 ± 0.038 2.156 ± 0.047 2.907 ± 0.056 3.633 ± 0.065 4.459 ± 0.075 [Omim][NO3] 0.296 ± 0.023 0.636 ± 0.026 1.053 ± 0.031 1.469 ± 0.035 1.983 ± 0.040 2.608 ± 0.046 3.201 ± 0.053 3.669 ± 0.057 0.436 ± 0.024 0.989 ± 0.030 1.562 ± 0.036 2.178 ± 0.042 2.910 ± 0.050 3.717 ± 0.058 0.446 ± 0.025 0.998 ± 0.030 1.573 ± 0.036 2.219 ± 0.043 2.869 ± 0.049 3.631 ± 0.057

( f CO2/mCO2)/(MPa/(mol· kg−1) 1.884 1.822 1.678 1.633 1.476 2.703 2.545 2.412 2.340 2.354 3.075 2.971 2.893 2.794 2.678

± ± ± ± ± ± ± ± ± ± ± ± ± ± ±

0.182 0.087 0.056 0.048 0.043 0.221 0.134 0.096 0.084 0.081 0.246 0.161 0.124 0.109 0.098

1.819 1.705 1.564 1.488 1.436 2.270 2.192 2.124 2.021 1.911 3.050 2.896 2.803 2.701 2.689 2.552

± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ±

0.087 0.060 0.044 0.038 0.036 0.143 0.092 0.071 0.062 0.056 0.229 0.148 0.114 0.097 0.095 0.087

1.400 1.424 1.376 1.307 1.260 1.210 1.147 1.113 1.771 1.719 1.699 1.621 1.539 1.457 2.867 2.440 2.252 2.104 2.060 1.981

± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ±

0.142 0.073 0.046 0.035 0.028 0.025 0.022 0.022 0.151 0.071 0.052 0.041 0.035 0.032 0.229 0.102 0.076 0.061 0.055 0.050

a Standard uncertainty, u(T) = 0.1 K, u(P) = 0.1 MPa, u(mCO2) = 0.05 mol.kg−1, u( f CO2/mCO2) = 0.25 MPa/(mol/kg−1).

equilibrium cell that is calculated from the total volume and the volume of IL as follows C

DOI: 10.1021/acs.jced.5b00894 J. Chem. Eng. Data XXXX, XXX, XXX−XXX

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Article

ρIL wIL

(2)

where V is the volume of the equilibrium cell and ρIL is the IL density. According to the eq 1, the experimental standard uncertainty was calculated based on Moffat relation31 δmCO2

2 ⎛⎡ ⎡ ∂mCO ⎤2 ∂mCO2 ⎤ 2 ⎜ = ⎢ δp⎥ + ⎢ δwCO2 ⎥ ⎜⎣ ∂p ⎢⎣ ∂wCO2 ⎥⎦ ⎦ ⎝

⎡ ∂mCO ⎤2 ⎡ ∂mCO ⎤2 2 2 +⎢ δT ⎥ + ⎢ δwIL ⎥ + ⎣ ∂T ⎦ ⎣ ∂wIL ⎦

⎤2 ⎞ ⎡ ∂mCO ⎤2 ⎡ ∂mCO 2 2 δV ⎥ + ⎢ δρIL ⎥ ⎟ ⎢ ⎥⎦ ⎟ ⎣ ∂V ⎦ ⎣⎢ ∂ρIL ⎠

0.5

(3)

where δmCO2 is the uncertainty for CO2 molality. The experimental standard uncertainty for pressure was measured by calculating the pressure transducer uncertainty (0.002 MPa) and statistics uncertainties. The statistics uncertainties are determined from VLE model that described in Thermodynamic Modeling. The average of pressure uncertainty amounts to 0.053 MPa. The contributions of the main parameters involving in the evaluation of standard uncertainty for CO2 solubility according to eq 3 are given in Table 4. Table 4. Standard Uncertainty and the Average Contribution of Parameters for Estimation of CO2 Solubility parameter pressure (P) temperature (T) volume (V) mass of gas (wCO2) mass of IL (wIL) density of IL(ρIL)

standard uncertainty of the measurement

Figure 2. Total pressure above solution of CO2 + IL (a) [Hmim][SCN]; (b) [Omim][SCN]; (c) [Omim][NO3]); ◆, T = 298.15 K; ■, T = 313.15 K; ▲, T = 333.15 K; solid lines indicate correlation results.

contribution in the standard uncertainty of CO2 solubility

0.002 MPa 0.1 K

44.23% 1.32%

1 mL 0.01 g

34.49% 19.65%

0.01 g

0.10%

0.002 g·cm−3

0.20%

experimental data and literature at temperature 313.15 and 333.15 K, however the experimental data are somewhat different at a temperature of 298.15 K. This difference may be due to the impurities exist in IL. [BF4] anion is unstable even at room temperature and in the presence of water.6 The anion degrades rapidly to form fluoride-based impurities. Becausde of the existence of water in [Bmim][BF4], the IL may be degraded. The experimental data shows good agreement with Lie et al.34 at 313.15 and 333.15 K but discrepancy is observed at 298.15 K. This may be due to the use of different types of equipment for gas solubility. The solubility of CO2 is reduced with increasing the temperature, so the magnitude number of solubility is higher at lower temperature and the role of different type of gas solubility measurement shows its effect at lower temperature. The experimental data are compared with the solubility CO2 data in the other types of ILs at temperature 313.15 K and NMP (normal methyl pyrrolone) as an industrial solvent in Figure 4. As shown in the figure, the solubility of CO2 in NMP is more than the ILs and [Bmim][NTf2] has the greatest capacity for CO2 absorption in this comparison. Among the studied ILs, CO2 solubility in [Omim][NO3] is the highest.

This analysis showed that the pressure measurement has the greatest effect on the evaluation of standard uncertainty for CO2 solubility. The effect of temperature, IL density, and the mass of IL are negligible for estimation of standard uncertainty. The total pressure above the solution of CO2 and IL against the solubility of CO2 in IL at different temperatures is plotted in Figure 2. As Figure 2 shows, the pressure above the solution of CO2 + IL increases linearly with an increase in CO2 molality. This behavior is observed for the solubility of CO2 in other types of ILs.6−10 Furthermore, the solubility of CO2 in studied ILs decreases with an increase in temperature. In order to check the reliability of the experimental data, the solubility of CO2 in [Bmim][BF4] is measured and compared with the experimental data reported in literature.6,32−35 The experimental results are reported in Table 5 and compared with literature data at different temperature in Figure 3. According to Figure 3, the solubility pressure linearly increases with CO2 molality at constant temperature. As the figure shows, there is a good agreement existing between the

3. THERMODYNAMIC MODELING In the studied ILs, the experimental results for the gas solubility are correlated with the extended Henry’s law. The Henry’s D

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Table 5. Experimental molality (mCO2) and fugacity ( f CO2) of CO2 in [Bmim][BF4]a T/K 298.15

313.15

333.15

mCO2/(mol·kg−1) 0.3141 0.7181 1.1925 1.8067 2.4320 3.0703 3.6598 0.2274 0.4993 0.8118 1.1548 1.4822 1.8052 2.0370 0.1788 0.3442 0.5304 0.7187 0.8996 1.0572

± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ±

0.0179 0.0189 0.0207 0.0237 0.0272 0.0312 0.0351 0.0172 0.0184 0.0205 0.0240 0.0282 0.0329 0.0373 0.0165 0.0175 0.0193 0.0217 0.0245 0.0270

P/MPa 0.358 0.799 1.287 1.875 2.426 2.941 3.376 0.416 0.930 1.500 2.178 2.820 3.419 3.924 0.480 0.941 1.484 2.025 2.555 2.959

± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ±

0.029 0.039 0.051 0.065 0.078 0.090 0.101 0.030 0.042 0.056 0.072 0.087 0.092 0.097 0.031 0.042 0.055 0.068 0.081 0.091

( f CO2/mCO2)/(MPa/(mol· kg−1)) 1.121 1.069 1.011 0.943 0.879 1.821 0.771 1.798 1.789 1.732 1.716 1.681 1.627 1.615 2.640 2.648 2.658 2.625 2.596 2.521

± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ±

0.069 0.038 0.030 0.026 0.023 0.021 0.020 0.143 0.079 0.060 0.054 0.051 0.049 0.049 0.252 0.149 0.116 0.101 0.094 0.088

Figure 4. Comparison of experimental and literature data for solubility of CO2 in ILs at temperature 313.15 K and different pressure: ▲, NMP34 (T = 298.15 K); ■, [Bmim][CH3SO4];7 ●, [Bmim][NTf2];1 ▲, [Bmim][BF4] this work; *, [Bmim][NO3];5 ◆, [Bmim][PF6];36 ▲, [Omim][NO3] this work; ×, [C8mim][PF6];5 ■, [Hmim][SCN] this work; ◆, [Omim][SCN] this work.

constant for CO2 in ILs are calculated similar to the procedure that was described by Kumelan et al.7−10 and Perez-Salado et al.36 As Perez-Salado et al. have pointed out,36 the vapor−liquid equilibrium condition for CO2 + ILs systems results in the extended Henry’s law

a Standard Uncertainty, u(T) = 0.1 K, u(P) = 0.1 MPa, u(mCO2) = 0.05 mol.kg−1, u(f CO2/mCO2) = 0.25 MPa/(mol/kg−1).

KH ′ CO2(T ′P)aCO2(T ′mCO2) = fCO (T ′P)

(4)

2

where K0H′CO2(T′P) is Henry’s constant of CO2 in IL based on the molality scale, aCO2(T′mCO2) is the activity of CO2 in the IL, and f CO2(T′P) is the fugacity of CO2 in the vapor phase. The Henry’s constant of CO2 in IL is expressed as ⎛ V m∞′ CO P ⎞ 2 ⎟⎟ KH ′ CO2(T ′P) = KH0 ′ CO2(T )exp⎜⎜ ⎝ RT ⎠

(5)

where K0H′CO2(T) is the Henry’s constant of CO2 in IL at zero pressure, and V∞ m′CO2 is the partial molar volume of CO2 in IL at infinite dilution. K0H′CO2(T) is estimated from the following 7 relation

⎡ f (T , P ) ⎤ 0 ⎢ CO2 ⎥ KH,CO ( T ) = lim mCO2 P → 0⎢ 2 ⎥ ⎣ ⎦ m0 −1

(6)

K0H,CO2(T),

where m° = 1 mol·kg . In order to estimate the plot of f CO2/mCO2 versus pressure are performed from the experimental data and extrapolated to zero pressure. The results are shown in Figure 5 and the values are reported in Table 6. In order to calculate V∞ CO2, the Krichevsky−Kasaranovsky equation is used37 ln Figure 3. Comparison of experimental and literature data for solubility of CO2 in [Bmim][BF4] at different pressure, (a) T = 298.15 K. (b) T = 313.15 K. (c) T = 333.15 K. ▲, this work; ■, Sudhir Aki et al.;6 ●, Shiflett and Yokozeki;32 ◆, Kroon et al.;33 ×, Tian et al.;34 +, Lei et al.;35 solid lines indicate correlation results.

f xCO2

S

(PIL) = ln HCO + 2 ,Il

∞ S vCO ̅ 2(P − P1 )

RT

(7)

Because of the negligible vapor pressure of ILs, eq 7 can be rewritten as follows ln E

f xCO2

S

(PIL) = ln HCO + 2 ,Il

∞ vCO ̅ 2*P

RT

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Table 7. Correlation for Binary Parameters (β0CO2−CO2) and the Absolute Relative Deviations (AARD%) for CO2 Solubility in IL β0CO2−CO2

system CO2 CO2 CO2 CO2

[Hmim][SCN]

[Omim][SCN]

[Omim][NO3]

[Bmim][BF4]

298.15 313.15 333.15 298.15 313.15 333.15 298.15 313.15 333.15 298.15 313.15 333.15

K0H,CO2

3 −1 V∞ CO2/cm ·mol

± ± ± ± ± ± ± ± ± ± ± ±

81.56 56.76 48.47 64.95 89.05 67.59 165.10 146.07 155.67 130.64 160.13 147.36

1.901 2.714 3.159 1.921 2.370 3.111 1.457 1.824 2.270 1.161 1.828 2.690

0.083 0.123 0.148 0.053 0.085 0.128 0.050 0.064 0.096 0.034 0.069 0.133

100 n

n

∑

f xCO2

0.9962/T 17.16/T 23.70/T 762.86/T

1.04 0.90 0.76 1.72

fcal − fexp fexp

i

ARD% =

fcal − fexp fexp

i

(11)

× 100 (12)

As Table 8 shows, the experimental data are in good agreement with the correlated results. Thermodynamic properties of CO2 solubility in studied ILs are calculated from the Henry’s constant based on the following eq and reported in Table 9 ⎛ K (T , P ) ⎞ Δsol G = RT ln⎜ H 0 ⎟ ⎝ P ⎠

(13)

⎡ KH(T , P) ⎢ ∂ ln P0 Δsol H = R ⎢ 1 ∂ T ⎢⎣

(14)

(

The slope of the graph ln

+ + + +

where f is the fugacity of CO2 in IL, subscripts cal and exp refer to the calculated and the experimental fugacity, respectively, and n is the number of data points. The comparison of experimental and calculated fugacity with their absolute relative deviation is reported in Table 8. The relative deviation is defined as follows

Table 6. Henry’s Constant of CO2 in ILs on Molality Scale at Zero Pressure (K0H,CO2) and Partial Molar Volume of CO2 in IL at Infinite Dilution (V∞ CO2) at Different Temperature T/K

−0.0909 −0.1544 −0.1755 −0.0338

[Hmim][SCN] [Omim][SCN] [Omim][NO3] [Bmim][BF4]

AARD% =

Figure 5. Plot of the ratio of CO2 fugacity to the molality in gas phase versus total pressure for the system CO2 + [Omim][NO3]; ◆, T = 298.15 K; ■, T = 313.15 K; ▲, T = 333.15 K.

IL

+ + + +

AARD%

()

versus pressure at constant

∞ temperature is V∞ CO2. The numerical values of VCO2 are reported in Table 6. The value of K0H,CO2(T) for [Bmim][BF4] at 298.15 K was reported at 1.219 by Sudhir Aki et al.6 and Hong et al.38 reported this value at 1.193. In order to correlate the gas solubility, the activity of CO2 in IL should be calculated mCO2 aCO2 = γ (9) m0 CO2

Δsol S =

) ⎤⎥P

(Δsol H − Δsol G) T

⎥ ⎥⎦

(15)

where ΔsolG, ΔsolH, and ΔsolS are the Gibbs free energy of solvation, enthalpy, and entropy of solvation, respectively, . At standard temperature and pressure (T0 = 298.15 K, p0 = 0.1 MPa) the thermodynamic properties of CO2 in studied ILs are calculated.

4. CONCLUSION The unique properties of ILs made them the high potential solvent for CO2 absorption. New experimental data for solubility of CO2 in [Hmim][SCN], [Omim][SCN], [Omim][NO3], and [Bmim][BF4] are reported at temperatures 298.15, 313.15, and 333.15 K and pressure up to 4.5 MPa. The solubility of CO2 is reduced with increasing temperature. Henry’s constant are calculated from the solubility data. The structure of ILs has a great impact on the solubility of CO2. The experimental data reveal that the solubility of CO2 is increased with the enlargement of the alkyl chain length of cation of IL. Furthermore, the nitrate-based ILs are more favorable for solubility of CO2 than thiocyanate-based ILs. The experimental data are correlated with the extended Henry’s law and Pitzer’s virial expansion in molality scale. The correlation results have a good agreement with the experimental solubility data.

where γCO2 is the activity coefficient of CO2. This parameter is expressed based on pitzer’s virial expansion in molality scale that can be rewritten after simplification as follows7 mCO 0 ln γCO = 2 o 2 βCO 2 2 − CO2 (10) m The parameter β0CO2−CO2 is the binary interactions between CO2 molecules in the IL. The new experimental data for CO2 solubility in IL are correlated with the results for Henry’s constant, partial molar volume of CO2 in IL at infinite dilution, and Pitzer’s virial model with binary interaction parameter β0CO2−CO2. The correlation results for solubility of CO2 in IL are reported in Table 7 and the average absolute relative deviation (AARD%) is defined as F

DOI: 10.1021/acs.jced.5b00894 J. Chem. Eng. Data XXXX, XXX, XXX−XXX

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Table 8. Comparison of the Experimental Fugacities for CO2 Solubility in IL and Correlated and the Absolute Relative Deviations (ARD%) f exp CO2

f cal CO2

ARD%

CO2 + [Omim][SCN] 0.946 0.926 2.08 1.372 1.374 0.16 1.929 1.973 2.25 2.561 2.561 0.00 3.025 2.962 2.09 0.832 0.832 0.03 1.381 1.386 0.29 2.109 2.092 0.81 2.694 2.692 0.05 3.317 3.334 0.51 0.886 0.870 1.81 1.378 1.390 0.83 1.996 2.015 0.97 2.623 2.654 1.19 3.192 3.149 1.37 3.797 3.797 0.00 CO2 + [Omim][NO3] 0.291 0.297 1.95 0.616 0.605 1.82 0.999 0.987 1.23 1.363 1.371 0.60 1.791 1.798 0.42 2.276 2.272 0.18 2.703 2.703 0.00 3.015 2.995 0.65 0.428 0.431 0.58 0.948 0.953 0.49 1.461 1.442 1.27 1.981 1.974 0.34 2.560 2.563 0.13 3.149 3.156 0.22 0.439 0.433 1.36 0.964 0.986 2.18 1.489 1.479 0.68 2.053 2.047 0.30 2.593 2.582 0.41 3.191 3.205 0.43

f exp CO2

f cal CO2

*E-mail: [email protected]. Tel.: +98 21 44787770. Fax: +98 21 44787781.

ARD%

Notes

The authors declare no competing financial interest.

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Table 9. Numerical Values for the Thermodynamic Properties of CO2 in Studied IL ΔsolG (J mol−1)

ΔsolH (J mol−1)

ΔsolS (J·K−1·mol−1)

[Hmim][SCN] [Omim][SCN] [Omim][NO3] [Bmim][BF4]

7310 7340 6660 6090

−11740 −11620 −13220 −19710

−63.88 −63.64 −66.54 −86.55

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REFERENCES

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0.03 0.03 0.00 0.53 1.50 2.89 4.50 0.42 0.38 1.22 0.69 1.29 2.95 2.92 1.69 1.23 0.86 2.08 3.23 5.99

IL

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CO2 + [Hmim][SCN] 0.440 0.433 1.67 0.975 0.958 1.73 1.570 1.597 1.72 2.179 2.179 0.00 3.529 3.470 1.67 0.687 0.683 0.49 1.219 1.211 0.71 1.862 1.882 1.04 2.545 2.550 0.20 3.324 3.271 1.60 0.831 0.827 0.48 1.323 1.334 0.83 1.942 1.956 0.76 2.537 2.577 1.57 3.395 3.442 1.38 CO2 + [Bmim][BF4] 0.352 0.352 0.768 0.768 1.206 1.206 1.703 1.694 2.139 2.107 2.520 2.447 2.822 2.695 0.409 0.411 0.893 0.890 1.406 1.424 1.981 1.995 2.492 2.524 2.937 3.024 3.291 3.388 0.472 0.480 0.912 0.923 1.410 1.422 1.887 1.926 2.336 2.411 2.665 2.825

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S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jced.5b00894. The NMR spectra of studied ILs. (PDF) G

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DOI: 10.1021/acs.jced.5b00894 J. Chem. Eng. Data XXXX, XXX, XXX−XXX