Solubilities of Carbon Dioxide in 1-Ethyl-3-methylimidazolium

Article Cite This: J. Chem. Eng. Data XXXX, XXX, XXX-XXX

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Solubilities of Carbon Dioxide in 1‑Ethyl-3-methylimidazolium Thiocyanate, 1‑Ethyl-3-methylimidazolium Dicyanamide, and 1‑Ethyl-3-methylimidazolium Tricyanomethanide at (298.2 to 373.2) K and (0 to 300.0) kPa Kuan Huang* and Hai-Long Peng* Poyang Lake Key Laboratory of Environment and Resource Utilization (Nanchang University), Ministry of Education; School of Resources Environmental and Chemical Engineering, Nanchang University, Nanchang, Jiangxi 330031, China S Supporting Information *

ABSTRACT: The solubilities of carbon dioxide (CO2) in three low-viscous ionic liquids (ILs)1-ethyl-3-methylimidazolium thiocyanate ([emim][SCN]), 1-ethyl-3-methylimidazolium dicyanamide ([emim][DCA]), and 1-ethyl-3-methylimidazolium tricyanomethanide ([emim][TCM])were determined by volumetric method at the temperature range of (298.0 to 373.2) K and pressure range of (0 to 300.0) kPa, which are relevant to the capture of CO2 from flue gas and natural gas. The Henry’s law constants and partial molar volume at infinite dilution of CO2 in the three ILs were calculated by fitting the solubility data with the Krichevsky−Kasarnovsky (K-K) model. The absorption enthalpies, Gibbs free energy changes, and absorption entropies of CO2 in the three ILs were also calculated based on the dependence of Henry’s constants on temperature. DFT calculations were further performed to better understand the interaction of ILs with CO2 on the molecular level and provide theoretical basis for the different CO2 solubilities in the three ILs.

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INTRODUCTION Removal of carbon dioxide (CO2) from industrial streams (e.g., flue gas, natural gas, and syngas) is an essential supplement to energy utilization processes, to reduce the emission of greenhouse gas and control the global warming.1 Absorption in liquid solvents, which can be classified as chemical absorption and physical absorption according to the reactivity of solvents to CO2, is the most established technology for CO2 capture in the industry. Chemical solvents such as organic amines,2 offer high loading capacity for CO2 at low partial pressures, but suffer from oxidative degradation, severe corrosion, and intensive energy consumption for regeneration. Physical solvents such as methanol3 and dimethyl ethers of poly(ethylene glycol),4 offer considerable loading capacity for CO2 at moderate to high partial pressures, and can be easily regenerated by temperature and/or pressure swing. However, the volatile nature of organic solvents is a significant issue, which may result in solvents loss in temperature and/or pressure swing processes. Ionic liquids (ILs) have attracted widespread attention over the past decades as alternatives to organic solvents for chemical engineering processes because they have many unique properties such as wide liquid range, negligible volatility, and high thermal stability.5 The application of ILs in CO2 capture is of special interest owing to their excellent affinity to CO2.6 Within this regard, the knowledge of CO2 solubilities in ILs is very important for process design, and a great number of works © XXXX American Chemical Society

have been dedicated to the determination of CO2 solubilities in ILs.7 However, a factor that limits the industrial application of ILs is their high viscosities (mostly >50 cP at room temperature8), which disfavors gas diffusion in them and requires extra pump energy to transfer them in pipelines.9 Among various ILs, 1-ethyl-3-methylimidazolium thiocyanate ([emim][SCN]), 1-ethyl-3-methylimidazolium dicyanamide ([emim][DCA]), and 1-ethyl-3-methylimidazolium tricyanomethanide ([emim][TCM]) are three with much lower viscosities (∼20 cP at room temperature10−12) than others. The chemical structures of [emim][SCN], [emim][DCA], and [emim][TCM] are shown in Scheme 1. To the best of our knowledge, the solubilities of CO2 in [emim][SCN], [emim][DCA], and [emim][TCM] are still very limited in the literature. For example, Lim et al.13 determined the solubilities of CO2 in [emim][SCN], [emim][DCA], and [emim][TCM] at (303.15 to 373.15) K and (1300 to 22860) kPa by the bubble point method; Brennecke et al.14 determined the solubilities of CO2 in [emim][SCN] at (299.4 to 313.3) K and (1090 to 5670) kPa by the volumetric method; Li et al.15 measured the solubilities of CO2 in [emim][DCA] at (303.2 to 343.2) K and (148 to 5944) kPa by the thermogravimetric method; Kroon et al.16 measured the Received: May 26, 2017 Accepted: November 14, 2017

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DOI: 10.1021/acs.jced.7b00476 J. Chem. Eng. Data XXXX, XXX, XXX−XXX

Journal of Chemical & Engineering Data

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EXPERIMENTAL SECTION Materials. CO2 was supplied by Nanjing Special Gas Co. Ltd., China. [emim][SCN], [emim][DCA], and [emim][TCM] were purchased from Shanghai Chengjie Co. Ltd., China. CO2 was used as received, while [emim][SCN], [emim][DCA], and [emim][TCM] were used after drying at 353.2 K and 0.1 kPa for 48 h. The water content in the ILs were measured by Karl Fischer titration on a Metrohm 787 KF Titrino system, and results indicate that the residual water content in the ILs are lower than 2000 ppm. The basic information on chemicals used in this work are presented in Table 1. Measurement of Gas Solubilities. The apparatus for measuring CO2 solubilities is similar to that introduced in our previous work, and the accuracy has been validated previously.20 Scheme 2 shows the diagram of apparatus for

Scheme 1. Chemical Structures of ILs Investigated in This Work

Scheme 2. Diagram of Apparatus for Gas Solubility Measurements

gas solubility measurements. Basically, the whole system consists of two 316 L stainless steel chambers, the volumes of which are 120.80 cm3 (V1) and 47.37 cm3 (V2), respectively. Volumes V1 and V2 include the test chambers’ volumes and the volumes of connecting parts. They were measured by using helium as the probing gas, with uncertainties of 0.012 and 0.004 cm3, respectively. The bigger chamber is used as a gas reservoir to isolate gas before it contacts with the IL loaded in the smaller chamber. The smaller chamber is used as an equilibrium cell and equipped with a magnetic stirrer. The temperatures (T) of the two chambers are controlled by a water bath with an uncertainty of 0.1 K. The pressures in two chambers are monitored by two Wideplus-8 pressure transducers with an uncertainty of 0.3 kPa. The pressure transducers are connected to a Wideplus-80 digital displayer to record the pressure change online. In a typical run, about 5 g of IL was loaded into the equilibrium cell. The mass of IL loaded (w) was measured by an analytical balance with an uncertainty of 0.0001 g. The whole system was evacuated, and the residual pressure in the equilibrium cell was determined to be P0 (with the values of 0−

solubilities of CO2 in [emim][TCM] at (298.15 to 353.25) K and (0 to 2000) kPa by the gravimetric method; Camper et al.17,18 and Scovazzo et al.19 reported the Henry’s constants of CO2 in [emim][DCA] at 303 and 313 K, but no solubility data were given. The solubilities of CO2 in [emim][SCN], [emim][DCA], and [emim][TCM] at low to moderate pressures (0−500 kPa), which are relevant to the capture of CO2 from flue gas and natural gas are missing. To fill this gap, we measured the solubilities of CO2 in [emim][SCN], [emim][DCA], and [emim][TCM] at (298.2 to 373.2) K and (0 to 300.0) kPa by volumetric method in this work. The experimental temperature range of 298.2−373.2 K was chosen, because it matches the temperature of natural gas and most flue gas. The temperature of some flue gas and syngas may be higher. Therefore, the experimental temperature range in this work is applicable to most cases. Table 1. Basic Information of Chemicals Used in This Work chemical

CAS number

M.W. (g/mol)

initial puritya

densitya (g/cm3)

viscosityb (mPa·s)

water contentc (ppm)

CO2 [emim][SCN] [emim][DCA] [emim][TCM]

124-38-9 331717-63-6 370865-89-7 666823-18-3

44.01 169.25 177.21 201.23

99.99 mol % 99.0 wt % 99.0 wt % 99.0 wt %

1.0610 1.1168411 1.0814612

2110 22.1511 14.0012

1517 1268 1344

a

As received by the authors and stated by the supplier. bData at 298.2 K. cDetermined by Karl Fischer titration, and the standard uncertainty is u = 10 ppm. B



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Table 2. Solubilities of CO2 in [emim][SCN]a T (K)

PCO2 (kPa)

298.2 298.2 298.2 298.2 298.2 298.2 313.2 313.2 313.2 313.2 313.2 313.2 333.2 333.2 333.2 333.2 333.2 333.2

61.1 102.2 153.8 202.8 260.1 303.8 52.8 106.1 151.9 204.7 275.1 308.0 59.6 105.6 154.5 203.5 255.0 304.2

xCO2

T (K)

PCO2 (kPa)

± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ±

353.2 353.2 353.2 353.2 353.2 353.2 373.2 373.2 373.2 373.2 373.2 373.2

53.2 100.9 168.4 205.5 251.7 303.8 56.3 103.7 157.6 204.5 258.3 302.4

0.0032 0.0058 0.0094 0.0139 0.0167 0.0189 0.0023 0.0045 0.0067 0.0098 0.0135 0.0151 0.0014 0.0027 0.0043 0.0067 0.0091 0.0102

0.0004 0.0004 0.0004 0.0004 0.0004 0.0004 0.0004 0.0004 0.0003 0.0003 0.0003 0.0003 0.0003 0.0003 0.0003 0.0003 0.0003 0.0002

xCO2 0.0012 0.0014 0.0032 0.0038 0.0054 0.0072 0.0008 0.0010 0.0021 0.0030 0.0040 0.0047

± ± ± ± ± ± ± ± ± ± ± ±

0.0002 0.0002 0.0002 0.0002 0.0002 0.0002 0.0001 0.0001 0.0001 0.0001 0.0001 0.0001

a

The solubility data are reported as the molar fraction of CO2; standard uncertainties u are u(T) = 0.1 K and u(PCO2) = 0.3 kPa; u(xCO2) are reported following the ± sign.

Table 3. Solubilities of CO2 in [emim][DCA]a T (K)

PCO2 (kPa)

298.2 298.2 298.2 298.2 298.2 298.2 313.2 313.2 313.2 313.2 313.2 313.2 333.2 333.2 333.2 333.2 333.2 333.2

56.5 105.8 150.3 199.3 251.5 303.0 52.8 102.6 152.5 204.7 254.1 305.1 52.8 107.2 154.3 204.8 252.5 302.2

xCO2

T (K)

PCO2 (kPa)

± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ±

353.2 353.2 353.2 353.2 353.2 353.2 373.2 373.2 373.2 373.2 373.2 373.2

58.3 104.9 152.3 202.8 252.7 303.1 52.1 99.2 155.3 198.5 254.9 301.2

0.0058 0.0105 0.0147 0.0196 0.0248 0.0298 0.0040 0.0071 0.0112 0.0161 0.0206 0.0249 0.0029 0.0056 0.0079 0.0112 0.0139 0.0169

0.0005 0.0005 0.0005 0.0005 0.0004 0.0004 0.0004 0.0004 0.0004 0.0004 0.0004 0.0004 0.0003 0.0003 0.0003 0.0003 0.0003 0.0003

xCO2 0.0022 0.0038 0.0061 0.0085 0.0110 0.0129 0.0017 0.0034 0.0053 0.0068 0.0086 0.0105

± ± ± ± ± ± ± ± ± ± ± ±

0.0003 0.0003 0.0003 0.0002 0.0002 0.0002 0.0002 0.0002 0.0002 0.0002 0.0002 0.0002

a


0.3 kPa) approaching the detecting limit of pressure transducers. The residual pressure P0 was recorded for 24 h before CO2 solubility measurements started. It was found that the increase in P0 was less than 0.1 kPa, validating the good vacuum tightness of the whole system. The gas from cylinder was then fed into the gas reservoir to a pressure of P1. The needle valve between two chambers was turned on to let the gas be introduced into the equilibrium cell. Absorption equilibrium was considered to be reached when the pressures of two chambers remained constant for at least 2 h. The equilibrium pressures were denoted as P2 for equilibrium cell and P′1 for the gas reservoir. The gas partial pressure in the equilibrium cell was PCO2 = P2 − P0. The gas uptaken(PCO2)was calculated using eq 1:

n(PCO2) = ρg (P1 , T )V1 − ρg (P1′, T )V1 − ρg (PCO2 , T ) (V2 − w/ρL )

(1)

where ρg(Pi,T) is the density of gas in mol/cm at Pi (i = 1, CO2) and T; ρL is the density of IL in g/cm3 at T. Continual measurements of gas solubilities at elevated pressures were performed by introducing more gas into the equilibrium cell to reach a new equilibrium. The gas solubilities were defined as the molar fraction of gas in the IL phase. After the completion of measurements, the gas remaining in two chambers was swept into a tail-gas absorber containing aqueous solutions of NaOH to prevent the gas from leaking into the atmosphere. Duplicate experiments were performed for each IL to give averaged values of gas solubility. The standard uncertainties of gas solubility data were estimated from the uncertainties of pressure by error 3

C



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Table 4. Solubilities of CO2 in [emim][TCM]a T (K)

PCO2 (kPa)

298.2 298.2 298.2 298.2 298.2 298.2 313.2 313.2 313.2 313.2 313.2 313.2 333.2 333.2 333.2 333.2 333.2 333.2

51.6 99.4 150.6 200.4 249.2 306.3 49.7 99.6 150.1 199.9 249.2 299.5 51.5 100.1 151 200.8 250 299.1

xCO2

T (K)

PCO2 (kPa)

± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ±

353.2 353.2 353.2 353.2 353.2 353.2 373.2 373.2 373.2 373.2 373.2 373.2

49.4 100.5 150.4 198.8 249.6 299.3 50.2 101.6 157.2 199.3 254.6 301.6

0.0104 0.0194 0.0288 0.0377 0.0463 0.0563 0.0082 0.0150 0.0219 0.0291 0.0354 0.0420 0.0052 0.0103 0.0155 0.0206 0.0256 0.0304

0.0008 0.0008 0.0008 0.0008 0.0008 0.0008 0.0008 0.0008 0.0007 0.0007 0.0007 0.0007 0.0007 0.0007 0.0007 0.0007 0.0006 0.0006

xCO2 0.0041 0.0079 0.0122 0.0159 0.0199 0.0239 0.0032 0.0064 0.0095 0.0121 0.0155 0.0183

± ± ± ± ± ± ± ± ± ± ± ±

0.0006 0.0006 0.0006 0.0006 0.0006 0.0006 0.0006 0.0006 0.0006 0.0005 0.0005 0.0005

a


propagation, as the contributions of the uncertainties of temperature, volume, mass, and density to the total uncertainties of gas solubility data are very small and can thus be neglected.20 Theoretical Calculations. Gas phase density functional theory (DFT) calculations were performed at standard conditions (298.2 K, 1 bar) with a Gaussian 09 program.21 All the geometries were fully optimized at the B3LYP/6-31G+ +(d,p) level of theory. No restrictions on symmetries were imposed on the initial structures, therefore, the geometry optimization for the saddle points occurred with all degrees of freedom. Vibration frequencies were calculated to confirm the true minima. For each anion−CO2 complex, structures with different binding modes were screened to locate the most stable structure. The binding energies were calculated by subtracting the thermal enthalpies of single anion and CO2 from that of anion−CO2 complex.

Figure 2. Solubilities of CO2 in [emim][DCA].

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RESULTS AND DISCUSSION The solubilities of CO2 in [emim][SCN], [emim][DCA], and [emim][TCM] are presented in Tables 2−4, and plotted in Figures 1−3. As can be seen, the solubility of CO2 increases

Figure 3. Solubilities of CO2 in [emim][TCM].

almost linearly with the increase of CO2 pressure, indicating the typically physical absorption of CO2 in the three ILs, which is in accordance with the literature.13 In addition, the solubility of CO2 decreases with the increase of experimental temperature, which is within expectation as the absorption of CO2 is

Figure 1. Solubilities of CO2 in [emim][SCN]. D



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Figure 7. Comparison of CO2 solubilities in [emim][SCN], [emim][DCA], and [emim][TCM] at 298.2 K.

Figure 4. Comparison of CO2 solubilities in [emim][DCA] from our work with those from Li et al.15

Figure 5. Comparison of CO2 solubilities in three ILs at 313.2 K from our work with those from Lim et al.13

Figure 8. Comparison of calculated CO2 solubilities with experimental CO2 solubilities.

and 152.5 kPa, and 0.0079 at 333.2 K and 154.4 kPa. Obviously, the solubility data provided by Li et al. are higher than ours. For the first point, the relative deviation is as large as 24%; but for the second point, the relative deviation is only 8%. According to Kroon et al.,16 the solubilities of CO2 in [emim][TCM] are measured to be 0.0161 at 298.15 K and 96 kPa, 0.0334 at 298.15 K and 194 kPa, 0.0509 at 298.15 K and 294 kPa, 0.0072 at 353.25 K and 95 kPa, 0.0155 at 353.25 K and 194 kPa, and 0.0227 at 353.25 K and 294 kPa. Our values at similar conditions are 0.0194 at 298.2 K and 99.4 kPa, 0.0377 at 298.2 K and 200.4 kPa, 0.0563 at 298.2 K and 306.3 kPa, 0.0079 at 353.2 K and 100.5 kPa, 0.0159 at 353.2 K and 198.8 kPa, and 0.0239 at 353.2 K and 299.3 kPa. Obviously, the solubility data provided by Kroon et al. agree well with ours, with the average relative deviation being smaller than 10%. If the CO2 solubility data from our work and those from Li et al.15 and Lim et al.13 are put in one graph for comparison (see Figures 4−6), the solubility data determined at the same temperature are along one continuous path as the pressure increases. It is acceptable as long as the whole trend matches well, although one or two certain data points do not match very well. Figure 7 shows the comparison of CO2 solubilities in [emim][SCN], [emim][DCA], and [emim][TCM] at 298.2 K. As can be seen, the CO2 solubilities follow the sequence of [emim][TCM] > [emim][DCA] > [emim][SCN]. Lim et al.13

Figure 6. Comparison of CO2 solubilities in [emim][TCM] from our work with those from Kroon et al.16

normally an exothermic process. As mentioned above, the solubilities of CO2 in [emim][SCN], [emim][DCA], and [emim][TCM], at low to moderate pressures (0−500 kPa) are very scarce in the literature, thus we can identify only several data points for direct comparison with ours. According to Li et al.,15 the solubilities of CO2 in [emim][DCA] are measured to be 0.0148 at 313.2 K and 155 kPa, and 0.0086 at 333.2 K and 149 kPa. Our values at similar conditions are 0.0112 at 313.2 K E



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Table 5. Thermodynamic Parameters of CO2 Absorption in [emim][SCN]a

a

T (K)

Hx (MPa)

298.2 313.2 333.2 353.2 373.2

16.2 24.3 39.8 75.9 90.9

± ± ± ± ±

1.4 0.8 4.8 8.0 9.7

V∞ (L/mol)

ΔH (kJ/mol)

ΔG (kJ/mol)

ΔS (J/mol·K)

−0.29 −1.59 −3.01 −5.57 −3.78

−22.4 ± 1.5

41.2 ± 0.1

−213 ± 6

± ± ± ± ±

8.61 0.30 1.28 1.17 1.30

Standard uncertainties u are u(T) = 0.1 K.

Table 6. Thermodynamic Parameters of CO2 Absorption in [emim][DCA]a

a

T (K)

Hx (MPa)

298.2 313.2 333.2 353.2 373.2

10.1 14.6 20.0 26.7 30.1

± ± ± ± ±

0.1 0.5 0.5 1.3 0.7

V∞ (L/mol)

ΔH (kJ/mol)

ΔG (kJ/mol)

ΔS (J/mol·K)

−0.10 −1.57 −1.01 −1.38 −0.41

−13.6 ± 1.1

40.0 ± 0.1

−180 ± 3

± ± ± ± ±

0.10 0.37 0.26 0.54 0.27


Table 7. Thermodynamic Parameters of CO2 Absorption in [emim][TCM]a

a

T (K)

Hx (MPa)

V∞ (L/mol)

ΔH (kJ/mol)

ΔG (kJ/mol)

ΔS (J/mol·K)

298.2 313.2 333.2 353.2 373.2

5.0 ± 0.1 6.4 ± 0.1 9.7 ± 0.1 12.5 ± 0.2 16.1 ± 0.2

0.69 1.00 0.15 0.07 0.25

± ± ± ± ±

−14.6 ± 0.5

38.2 ± 0.1

−177 ± 3

0.05 0.17 0.06 0.15 0.18


The K-K equation is simple in expression, and the Henry’s constants can be obtained directly by fitting the solubility data. However, many other models, for example the equation of state (EoS), are complex in expression, and the Henry’s constants can not be obtained directly by fitting the solubility data. Although some empirical equations (e.g., polynimials) are also simple in expression, the parameters in them do not have any physical meaning. It should be noted here that, since the experimental pressures are not very high (0−300 kPa), the fugacities of CO2 are almost the same as the pressures of CO2. We calculated the fugacities of CO2 using the PR equation of state, and calculated results are presented in Table S1 in the Supporting Information. It can be seen that the relative deviations between the fugacities and pressures of CO2 are less than 2%, therefore, it is reasonable to use the pressures of CO2 while not the fugacities of CO2 for the K−K equation. The fitting curves are shown in Figures 1−3, and Figure 8 shows the comparison of calculated CO2 solubilities with experimental CO2 solubilities. As can be seen, the K-K model can correlate the solubilities of CO2 in [emim][SCN], [emim][DCA], and [emim][TCM] very well, and the calculated CO2 solubilities agree very well with experimental CO2 solubilities, with an average relative deviation (ARD) of 3.5%. Tables 5−7 present the fitted values of Hx and V∞ for CO2 absorption in [emim][SCN], [emim][DCA], and [emim][TCM]. The Henry’s constants of CO2 in [emim][SCN] and [emim][DCA] are 16.2−90.9 MPa and 10.1−30.1 MPa, respectively, being higher than those in [emim][TCM] (5.0− 16.1 MPa) if compared at the same temperature. This result is consistent with the much lower solubilities of CO2 in [emim][SCN] and [emim][DCA] than those in [emim][TCM], which has been illustrated in Figure 7. Furthermore, the partial molar volumes of CO2 at infinite dilution have

Figure 9. Linear fit of ln Hx and 1/T for calculation of absorption enthalpies.

also observed the same trend in a previous work. They explained that the CO2 solubilities are affected by the number of cyano groups contained in ILs, and those containing more cyano groups have higher CO2 solubilities. However, they did not give any details about the role of cyano groups in CO2 absorption. The Krichevsky−Kasarnovsky (K-K) eq 2 was then used to correlate the solubility data: ln

P V ∞P = ln Hx + x RT

(2)

where P is the pressure of CO2, x is the solubility of CO2, Hx is the Henry’s law constant of CO2 on molar fraction scale, V∞ is the partial molar volume of CO2 at infinite dilution, R is the universal gas constant, and T is the experimental temperature. F



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Figure 10. Optimized structures of [SCN]-CO2, [DCA]-CO2, and [TCM]-CO2 complexes.

negative values in [emim][SCN] (−0.29 to −5.57 L/mol) and [emim][DCA] (−0.10 to −1.38 L/mol), implying that CO2 tends to be accommodated in the cavity (free volume) of solvents, which is thermodynamically unstable. However, the partial molar volumes of CO2 at infinite dilution are positive in [emim][TCM] (0.07−1.00 L/mol), indicating that CO2 tends to penetrate into the bulky phase of the solvent. Therefore, CO2 is relatively more stable in [emim][TCM] than in [emim][SCN] and [emim][DCA]. This result also agrees with the fact that the solubilities of CO2 in [emim][TCM] are much higher than those in [emim][SCN] and [emim][DCA]. The absorption enthalpies (ΔH) of CO2 in [emim][SCN], [emim][DCA], and [emim][TCM] were calculated from the van’t Hoff eq 3 by drawing a linear fit between ln Hx and 1/T (see Figure 9): ∂ ln Hx ΔH = ∂T RT 2

reflects the chemical potential that is minimized when the solvent−solute system reaches equilibrium. The Gibbs free energy changes of CO2 in [emim][SCN] (41.2 kJ/mol) and [emim][DCA] (40.0 kJ/mol) are higher than that in [emim][TCM] (38.2 kJ/mol), indicating the higher chemical potentials of CO2 in [emim][SCN] and [emim][DCA] than that in [emim][TCM]. That is to say, CO2 is relatively less stable in [emim][SCN] and [emim][DCA] than in [emim][TCM], which agrees with the conclusion drawn from the comparison of partial molar volumes at infinite dilution. Absorption entropy reflects the ordering degree of the solvent−solute system. The absorption entropy of CO2 in [emim][SCN] (−213 J/mol·K) is lower than those in [emim][DCA] (−180 J/mol·K) and [emim][TCM] (−177 J/mol·K), indicating that CO2 is more ordered in [emim][SCN] than those in [emim][DCA] and [emim][TCM]; thus the absorption of CO2 in [emim][SCN] is a less spontaneous process than those in [emim][DCA] and [emim][TCM]. To better understand the interaction of [emim][SCN], [emim][DCA], and [emim][TCM] with CO2 on the molecular level, and provide a theoretical basis for the different CO2 solubilities in three ILs, DFT calculations were carried out to optimize the structures of [SCN]-CO2, [DCA]-CO2, and [TCM]-CO2 complexes, as shown in Figure 10. It has been previously pointed out that anions of ILs play the primary role in interacting with CO2;22 therefore, the contribution of [emim] cation to the interaction of [emim][SCN], [emim][DCA], and [emim][TCM] with CO2 was ignored in this work to save computation time. This way of dealing with CO2 absorption in ILs has been widely accepted by researchers.23 It can be seen that CO2 is preferably binded to the cyano group in [SCN], [DCA], and [TCM]. For example, the binding energy of the cyano group in [SCN] with CO2 is −18.7 kJ/mol, being more negative than that of sulfur with CO2 (−14.1 kJ/mol); the binding energy of the cyano group in [DCA] with CO2 is

(3)

Subsequently, the Gibbs free energy changes (ΔG) and absorption entropies (ΔS) of CO2 in [emim][SCN], [emim][DCA], and [emim][TCM] were calculated by eqs 4 and 5: ΔG = RT ln Hx

(4)

ΔS = (ΔH − ΔG)/T

(5)

Calculated values of ΔH, ΔG, and ΔS are also presented in Tables 5−7. Absorption enthalpy reflects the strength of interaction between solvent and solute. It is surprising to find that the absorption enthalpy of CO2 in [emim][SCN] (−22.4 kJ/mol) is more negative than those in [emim][DCA] (−13.6 kJ/mol) and [emim][TCM] (−14.6 kJ/mol), indicating the stronger interaction of [emim][SCN] with CO2 than those of [emim][DCA] and [emim][TCM] with CO2, although [emim][SCN] has much lower CO2 solubilities than [emim][DCA] and [emim][TCM]. The Gibbs free energy change G



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−15.1 kJ/mol, being more negative than that of amide nitrogen with CO2 (−13.1 kJ/mol). Furthermore, the binding strength of the cyano group in [SCN] with CO2 is stronger than those of the cyano group in [DCA] and [TCM] with CO2 (−18.7 vs −15.1 vs −10.9 kJ/mol), which accounts for the more negative absorption enthalpy of CO2 in [emim][SCN] than those in [emim][DCA] and [emim][TCM] (−22.4 vs −13.6 vs −14.6 kJ/mol). However, there is only one cyano group in [SCN], while there are two and three cyano groups in [DCA] and [TCM], respectively. The binding energy of the second cyano group in [DCA] with CO2 is −14.4 kJ/mol, being comparable to that of first cyano group in [DCA] with CO2. The binding energy of second and third cyano group in [TCM] with CO2 are −10.4 and −10.0 kJ/mol, respectively, also being comparable to that of first cyano group in [TCM] with CO2. Therefore, the overall loading capacities of CO2 in [emim][DCA] and [emim][TCM] are much higher than those in [emim][SCN]. The DFT calculation results also validate the inference of Lim et al.,13 that the solubilities of CO2 in these ILs are determined by the number of cyano group in ILs.

Notes

The authors declare no competing financial interest.

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CONCLUSIONS The solubilities of CO2 in [emim][SCN], [emim][DCA], and [emim][TCM] at (298.2 to 373.2) K and (0 to 300.0) kPa were reported in this work. The solubility of CO2 increases almost linearly with pressure, implying the physical absorption behavior of CO2 in the two ILs. Thermodynamic parameters including Henry’s law constants, partial molar volumes at infinite dilution, absorption enthalpies, Gibbs free energy changes, and absorption entropies were calculated by fitting the solubility data with the Krichevsky−Kasarnovsky (K-K) model. DFT calculations disclosed that CO2 is preferably binded to the cyano group in [SCN], [DCA], and [TCM], and the binding strength of the cyano group in [SCN] with CO2 is stronger than those of the cyano groups in [DCA] and [TCM] with CO2. However, there are less cyano groups in [emim][SCN] than in [emim][DCA] and [emim][TCM], which accounts for the much lower solubilities but more negative absorption enthalpy of CO2 in [emim][SCN] than those in [emim][DCA] and [emim][TCM].

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ASSOCIATED CONTENT

* Supporting Information S

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jced.7b00476.

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Calculated fugacities of CO2 (PDF)

AUTHOR INFORMATION

Corresponding Authors

*E-mail: [email protected]. *E-mail: [email protected]. ORCID

Kuan Huang: 0000-0003-1905-3017 Hai-Long Peng: 0000-0003-3150-4618 Funding

This work was supported by the Natural Science Foundation of Jiangxi Province (Grant No. 20171BAB203019), and the Natural Science Foundation of China (Grant No. 31660482). The authors also appreciate the sponsorship from Nanchang University. H



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Solubilities of Carbon Dioxide in 1-Ethyl-3-methylimidazolium

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