Article pubs.acs.org/jced
Solubilities of Carbon Dioxide in Eutectic Mixtures of Choline Chloride and Dihydric Alcohols Yanfei Chen, Ning Ai,* Guihua Li, Haifang Shan, Yanhong Cui, and Dongshun Deng* Zhejiang Province Key Laboratory of Biofuel, College of Chemical Engineering, Zhejiang University of Technology, Hangzhou 310014, China ABSTRACT: The solubilities of CO2 in eutectic mixtures containing choline chloride and dihydric alcohols (including 1,4-butanediol, 2,3-butanediol, and 1,2propanediol) with a molar ratio of choline chloride to dihydric alcohol of 1:3 and 1:4 were measured at 293.15 K, 303.15 K, 313.15 K, and 323.15 K under pressures up to 600.0 kPa using an isochoric saturation method. Henry’s constant and the dissolution Gibbs free energy, enthalpy, and entropy changes of CO2 solvation were obtained by correlating the experimental data. The solubility of CO2 in the mixtures increased linearly with the increasing pressure or the decreasing temperature. The enthalpies of solution were negative at all conditions.
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CC) as well as its derivatives with metal halide or urea16,17 and were subsequently extended to other hydrogen acceptors in terms of quaternary ammonium or phosphonium salt and hydrogen donors like alcohols, carboxylic acids, amino acids, different sugars, amides, and so on.18−22 Just like RTILs, DESs possess many unusual advantages such as negligibly small vapor pressure, high thermal and chemical stabilities, and high capacity to dissolve CO2. However, unlike RTILs, DESs are easy to be formed by simple blending of cheap and high-purity raw materials. Therefore, the manufacturing cost of DESs is much lower than that of RTILs.23 Properties of DESs can be regulated by selecting various combination and molar ratio of the hydrogen bond pairs. Furthermore, they are environmentally benign and economic media and can be made from biodegradable components.24 Recently, DESs have been attracted as new emerging mediums for CO2 capture because, besides the above-mentioned features, they are similar to ILs and composed of predominantly ionic species, and thus also have interesting properties for high CO2 dissolution. Li et al.25 compared the effect of composition on CO2 solubility for CC− urea mixtures at 313.15 K, 323.15 K, and 333.15 K and pressures until 13 MPa. The results showed that a 1:2 (molar ratio) mixture achieved the highest solubility for CO2. Leron et al.26 further measured the solubilities of CO2 in 1:2 CC−urea DES at more wider temperature scope of T = (303.15 to 343.15) K. Recently, Leron et al.27,28 also presented the solubility of CO2 in DESs containing CC and glycerol (1:2 molar ratio) or ethylene glycol using the thermogravimetric microbalance method with satisfactory results. Francisco et al.22
INTRODUCTION CO2, identified as a chief greenhouse gas, is a major contributor to global warming and climate change.1,2 The continuous rising of the atmospheric concentration of CO2 has raised public concern worldwide. Hence, CO2 capture and sequestration (CCS) have been the focus of many researches. The aminebased absorbents for CO2 are highly effective for CO2 removal, but they are not environmentally benign solvents because their toxicity, degradation sensitivity, volatility, and the intensive energy consumption to enable solvent regeneration.3,4 Consequently, the search for promising alternatives to the traditional solvents which can efficiently absorb CO2 was desirable. In the last two decades, room temperature ionic liquids (RTILs) have been a special subject with remarkable attention because of their intrinsic qualities such as negligible volatility, high thermal stability, and extensive tunability. They are potential alternatives to conventional solvents for CO2 capture.5 Numerous works have been devoted to increasing the absorption capacity of RTILs for CO2.6−10 However, the complex preparation process and relatively high cost make RTILs less attractive when compared to traditional solvents.11−13 To overcome such drawbacks, deep eutectic solvents (DESs) have been received more and more attention as alternatives to RTILs. DESs are generally considered to be eutectic mixtures composed of two or three safe and cheap compounds which have the ability to associate with each other through hydrogen bond interactions.14 The charge delocalization along with the hydrogen bond prevents the crystallization of the individual components, thus resulting in the decrease in the freezing point of the mixture relative to the melting points of the starting materials.15 DESs were first obtained by mixing choline chloride (2-hydroxyethyltrimethylammonium chloride, © 2014 American Chemical Society
Received: October 4, 2013 Accepted: March 10, 2014 Published: March 21, 2014 1247
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controlled by temperature controllers with a precision of ± 0.05 K. The pressure was monitored using a pressure transmitter (Fujian WIDEPLUS Precision Instruments Co., Ltd., WIDEPLUS-8, (0 to 600.0) kPa, with an accuracy of 0.1 % full scale). The solubilities of CO2 in DESs were determined using an isochoric saturation method.32 In a typical experiment, the temperature of 2 was always set at 303.15 K. After the desired amount of DES was added into the EC, the temperature of 3 was raised to be 340 K, with valve 10 closed and 9, 11 opened; the whole system was degassed by vacuum pumping while stirring for (1 to 2) h. Then valve 11 was closed, and the temperature of 3 was slowly cooled to experimental temperature, for example, 293.15 K. The pressure of rudimental air was recorded as p1. With valve 9 closed and 10 opened, GR was charged with CO2 from gas cylinder through the pressure reducer until the pressure reached scheduled value, recorded as pressure p2. Subsequently, with valve 10 closed and 9 opened slowly, an amount of CO2 entered into EC from the GR and was absorbed by DESs. The absorption was accelerated by a magnetic stirrer. The system was considered to have reached equilibrium if the pressures had been unchanged within 4 h. Then the pressures of GR and EC were recorded as p3 and p4, respectively. Therefore, the quantity of CO2 absorbed could be obtained from the difference between the CO2 charged into the GR and remaining CO2 in GR as well as gaseous CO2 in EC. p4 was regarded as the equilibrium pressure of gaseous CO2. The next measurement at the same equilibrium temperature was carried out by introducing further amount CO2 into the EC from GR with the similar procedure. The measurement was repeated until the pressures between GR and EC were equal.
reported the CO2 capture using DES composed of CC and natural lactic acid (1:2 molar ratio). Considering that the combination of green DESs with CO2 capture has a potential application for various chemical processes, research on dissolution of CO2 in DESs possess prime importance. In this work, we presented the solubility of CO2 in the DESs composed of CC and dihydric alcohols (including 1,4-butanediol,29 2,3-butanediol,30 and 1,2-propanediol20) with molar ratios of 1:3 and 1:4 at the temperature range of (293.15 to 323.15) K and pressure up to 600.0 kPa. The experimental solubility data were then correlated with Henry’s law model. The related dissolution Gibbs free energy, enthalpy, and entropy changes of CO2 solvation in DESs were also calculated according to the relationship between Henry’s law constants and experimental temperatures.
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EXPERIMENTAL SECTION Materials. Anhydrous CC with a mass fraction of more than 0.985 was supplied by Jinan Hualing Pharmaceutical Co., Ltd. (Shandong, China). 2,3-Propanol, 1,4-butanediol, and 2,3butanediol with mass fractions of more than 0.990 were supplied by Sinopharm Chemical Reagent Co., Ltd. (Shanghai, China). All of the chemicals were directly used as received. CO2 gas with a mass fraction of more than 0.999 was supplied by Jingong Special Gas Co., Ltd. (Hangzhou, China). The DESs were obtained by mixing CC with the corresponding dihydric alcohols at a suitable temperature and drying under vacuum at 353 K for 48 h before use. The water content of each DES was determined by Karl Fischer analysis (SF-3 Karl Fischer Titration, Zibo Zifen Instrument Co. Ltd.) with a mass fraction of less than 8·10−4 in all cases. Densities of DESs were carefully determined at T = (293.15, 303.15, 313.15, and 323.15) K under atmospheric pressure using a 5.567 ± 0.004 cm3 pycnometer which was previously calibrated using doubledistilled water at 303.15 K. An electronic balance (MettlerToledo AL204, with an uncertainty of ± 2·10−4 g) was used to measure the mass of the DESs. Apparatus and Procedure. The stainless apparatus was an upgrade version on the basis of our glass one reported previously.31 It was composed mainly of a CO2 cylinder (1), two constant temperature water baths (2, 3), a CO2 gas equilibrium cell (4, EC) with a magnetic stirrer, a gas reservoir (5, GR), pressure transmitter (6, 7), and digital indicator (8) as illustrated in Figure 1. The volumes of the EC and GR were determined using the previous reported method31 with the results of 141.61 cm3 and 370.99 cm3, respectively. The temperatures of the temperature-constant water baths were
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RESULTS AND DISCUSSION Solubility Data of CO2 in DESs. Table 1 lists the measured densities of six DESs at several temperatures under normal Table 1. Densities of DESs at Different Temperaturesa ρ/g·cm−3 DESs
293.15 K
303.15 K
313.15 K
323.15 K
nCC:n1,4‑butanediol = 1:3 nCC:n1,4‑butanediol = 1:4 nCC:n2,3‑butanediol = 1:3 nCC:n2,3‑butanediol = 1:4 nCC:n1,2‑propanediol = 1:3 nCC:n1,2‑propanediol = 1:4
1.0524 1.0445 1.0468 1.0350 1.0666 1.0663
1.0490 1.0411 1.0390 1.0309 1.0599 1.0602
1.0432 1.0354 1.0357 1.0265 1.0533 1.0541
1.0384 1.0306 1.0300 1.0197 1.0467 1.0480
a
The standard uncertainty u is u(T) = 0.05 K, and the combined expanded uncertainty Uc is Uc(ρ) = 0.0008 g·cm−3 (0.95 level of confidence).
pressure. The densities of all DESs varied with the kinds of dihydric alcohols and molar ratios as well as temperatures. The densities of the mixtures of CC and 1,2-propanediol were higher than those of CC and butanediols. Because of the negligibly small vapor pressure of DESs, the gas phase composition was regarded to be pure CO2. The amount of CO2 absorbed in DESs can be obtained according to the following equation, nCO2 = n − n1 − n2 (1) Figure 1. Schematic diagram of the CO2 solubility apparatus. 1, CO2 gas cylinder; 2, 3, thermostatic water bath and magnetic stirrer; 4, CO2 gas equilibrium cell; 5, CO2 gas reservoir; 6, 7, pressure transmitter; 8, digital indicator; 9, 10, 11, valves.
where nCO2 is the amount of soluble CO2 in the DESs. n and n1 represent the initial and residual amount of CO2 in the GR, respectively. n2 is the amount of insoluble CO2 in the vapor 1248
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Table 2. Experimental CO2 Mole Fraction (x2) and Molality (m2) in DESs at Temperature (T) and Equilibrium Pressure (p)a 293.15 K p
a
303.15 K m2
kPa
x2
110.9 212.5 312.5 412.0 509.7
0.0034 0.0069 0.0100 0.0131 0.0164
109.5 214.5 316.4 411.6 513.4
−1
m2
kPa
x2
0.0330 0.0681 0.0985 0.1292 0.1624
113.1 217.4 307.6 406.3 497.5
0.0031 0.0060 0.0087 0.0117 0.0142
0.0031 0.0062 0.0093 0.0123 0.0154
0.0306 0.0624 0.0938 0.1244 0.1560
105.8 205.8 304.7 403.3 503.6
0.0027 0.0052 0.0082 0.0110 0.0135
114.3 214.0 316.6 407.9 511.3
0.0031 0.0061 0.0091 0.0121 0.0152
0.0308 0.0601 0.0900 0.1190 0.1501
114.0 214.6 315.2 412.3 513.4
0.0028 0.0058 0.0082 0.0112 0.0139
107.1 207.2 309.7 403.0 508.5
0.0038 0.0071 0.0109 0.0148 0.0188
0.0382 0.0713 0.1105 0.1504 0.1915
117.6 218.7 320.1 409.2 511.2
0.0035 0.0068 0.0102 0.0134 0.0167
108.5 204.7 312.2 408.8 514.5
0.0033 0.0063 0.0098 0.0131 0.0165
0.0365 0.0690 0.1074 0.1441 0.1827
117.5 222.1 317.0 422.0 517.0
0.0028 0.0060 0.0084 0.0113 0.0140
104.4 209.6 307.6 401.1 501.5
0.0031 0.0064 0.0094 0.0127 0.0165
0.0355 0.0723 0.1073 0.1452 0.1884
117.2 220.1 321.7 415.6 517.5
0.0027 0.0050 0.0075 0.0101 0.0129
mol·kg
313.15 K
p
mol·kg
p −1
323.15 K m2
kPa
nCC:n1,4‑butanediol = 1:3 0.0302 121.4 0.0589 233.1 0.0857 321.6 0.1156 410.2 0.1410 502.8 nCC:n1,4‑butanediol = 1:4 0.0272 119.5 0.0526 223.2 0.0828 312.5 0.1113 410.0 0.1364 507.2 nCC:n2,3‑butanediol = 1:3 0.0275 119.3 0.0565 222.4 0.0804 325.5 0.1107 424.9 0.1377 528.8 nCC:n2,3‑butanediol = 1:4 0.0353 110.9 0.0689 217.2 0.1026 318.0 0.1361 414.9 0.1695 511.0 nCC:n1,2‑propanediol = 1:3 0.0304 125.4 0.0654 217.8 0.0919 317.6 0.1238 420.5 0.1543 515.4 nCC:n1,2‑propanediol = 1:4 0.0300 121.6 0.0569 225.2 0.0856 325.1 0.1154 421.4 0.1474 525.6
x2
mol·kg
p −1
m2
kPa
x2
mol·kg−1
0.0028 0.0052 0.0073 0.0097 0.0121
0.0269 0.0506 0.0715 0.0954 0.1196
122.8 218.3 325.2 425.9 525.9
0.0025 0.0042 0.0064 0.0085 0.0103
0.0247 0.0415 0.0625 0.0836 0.1014
0.0026 0.0049 0.0071 0.0096 0.0119
0.0262 0.0497 0.0713 0.0969 0.1208
123.0 221.6 332.6 425.3 519.0
0.0022 0.0039 0.0062 0.0081 0.0102
0.0217 0.0395 0.0622 0.0816 0.1031
0.0029 0.0054 0.0079 0.0102 0.0127
0.0284 0.0529 0.0780 0.1010 0.1251
124.0 217.0 322.0 421.5 514.1
0.0026 0.0047 0.0071 0.0092 0.0116
0.0251 0.0458 0.0696 0.0907 0.1140
0.0030 0.0061 0.0087 0.0118 0.0145
0.0302 0.0612 0.0882 0.1192 0.1469
116.4 214.1 315.9 425.7 512.5
0.0030 0.0052 0.0078 0.0106 0.0130
0.0305 0.0524 0.0788 0.1068 0.1320
0.0024 0.0044 0.0067 0.0093 0.0115
0.0266 0.0478 0.0737 0.1016 0.1264
123.0 236.0 334.3 426.3 524.7
0.0019 0.0039 0.0059 0.0078 0.0094
0.0206 0.0428 0.0650 0.0852 0.1031
0.0021 0.0043 0.0060 0.0081 0.0106
0.0241 0.0482 0.0679 0.0916 0.1208
122.5 226.6 326.5 425.0 519.8
0.0020 0.0036 0.0052 0.0071 0.0087
0.0221 0.0406 0.0591 0.0802 0.0989
Standard uncertainties u are u(T) = 0.05 K, u(p) = 0.6 kPa, ur(x) = 0.02, and ur(m) = 0.02.
including liquid phase molality (m2) and mole fraction (x2) of CO2 and gas phase equilibrium pressure (p). Figures 2 and 3 demonstrated the typical dependence of the solubility (m2 and x2) on temperature and pressure for DES composed of CC and 1,4-butanediol with the molar ratio of 1:3. Figure 4 showed the isothermal CO2 solubilities in DESs at 303.15 K. For better visualization, the CO2 solubility in the DES of CC and 1,4butanediol was omitted in Figure 3 because of the overlap. It was evident that the solubility of CO2 in the mixtures increased linearly with the increasing pressure and the decreasing temperature. Henry’s Constant. As shown in Figure 2, the isothermal CO2 solubility at 303.15 K increased with increasing CO2 pressure. Such purely physical solubility can be described using Henry’s law.33 The solubility of CO2 in DESs can be expressed with Henry’s constant based on mole fraction (Hx) as follows,34
phase of the EC at equilibrium. All of the amounts of CO2 at different places can be calculated according to the Soave− Redlich−Kwong (SRK) equation of state from experimental PVT data at different conditions, the volumes of GR and EC, the mass and density of DESs at different temperatures. The volume increase of the DESs in the EC as the result of CO2 solubility is very small and neglected. The mole quantity of absorbent (nliq) consisted of the mole quantity of CC (nCC) and the dihydric alcohols (nda) in the mixtures. The mole fraction and molality of CO2 in DESs were obtained by the following equation: xCO2 = nCO2 /(nCO2 + nliq ) = nCO2 /(nCO2 + nCC + nda) (2)
mCO2 = nCO2 /mliq
(3)
The solubility of CO2 in the DESs was measured at T = (293.15 to 323.15) K with 10 K intervals and pressures up to 600.0 kPa. The molar ratios of CC to dihydric alcohols were 1:3 and 1:4. The measured values were given in Table 2
Hx(p , T ) ≡ lim
x2 → 0
1249
f 2liq (p , T , x 2) x2
(4)
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Figure 4. CO2 solubility, expressed as CO2 molality (mCO2), as a function of CO2 equilibrium pressure (pCO2) at T = 303.15 K in DESs.
Figure 2. Solubilities of CO2 (molality) in DES of nCC:n1,4‑butanediol = 1:3. ■, 293.15 K; ●, 303.15 K; ▲, 313.15 K; ▼, 323.15 K; , linear fit.
■ , n CC :n 2,3‑butanediol = 1:3 ; □ , n CC :n 2,3‑butanediol = 1:4; nCC:n1,2‑propanediol = 1:3; △, nCC:n1,2‑propanediol = 1:4; , linear fit.
▲,
With the similar procedure, Henry’s constant based on molality (Hm) can also be expressed as the following, ⎡ f (p , T ) ⎤ pϕ2(p , T ) ⎥≅ Hm(p , T ) ≡ lim ⎢ 2 0 m2 → 0⎣ (m 2 / m ) ⎦ (m 2 / m 0 ) 0
where m = 1 mol·kg , and m2 represents molality of CO2 in the liquid phase. In present work, Henry’s constant was determined from the slope of an isotherm linear fit of fugacity versus mole fraction or molality of CO2 with the results illustrated in Table 3. By comparing Figure 3 and the Henry’s constants in Table 3, it can be found that the DES of CC and 2,3-butanediol with the molar ratio of 1:4 behaved higher CO2 solubility than that with the molar ratio of 1:3 and DESs containing 1,4-butanediol or 1,2-propanediol at the same temperature. However, the solubility behaviors of CO2 in the DESs composed of CC with 1,4-butanediol or 1,2-propanediol showed complex relationships, which were affected by both molar ratio and the temperature. Henry’s constants Hm can be easily used to compare the dissolution capacity for CO2 in different RTILs and other DESs at the same mass basis. As shown in Table 4, the solubilities of CO2 in present DESs are similar with those in the ammonium-based ILs,35 such as 2-hydroxy-N-(2-hydroxyethyl)-N-methylethanaminium lactate ([hhemel]), bis(2hydroxyethyl)ammonium lactate ([bheal]), and 2-hydroxy-N(2-hydroxyethyl)-N-methylethanaminium acetate ([hhemea]), but are slightly lower than those in the imidazolium36−38 and phosphonium-based39 ILs, such as 1,3-dimethylimidazolium dimethylphosphate ([dmim][Me2PO4]), 1-butyl-3-methylimidazolium dibutylphosphate ([bmim][Bu2PO4]), 1-ethyl-3methylimidazolium diethylphosphate ([emim][Et2PO4]), 1hexyl-3-methylimidazolium tetrafluoroborate ([hmim][BF4]), 1-butyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide ([bmim][Tf2N]) and trihexyl (tetradecyl) phosphoniumdodecylbenzenesulfonates ([P6,6,6,14][C12H25PhSO3]). To our knowledge, compared with imidazolium and phosphoniumbased ILs, DESs are more attractive in the price, toxicity, and biodegradability. Furthermore, the solubilities of CO2 in present DESs were found to be low compared with
Figure 3. Solubilities of CO 2 (mole fraction) in DES of nCC:n1,4‑butanediol = 1:3. ■, 293.15 K; ●, 303.15 K; ▲, 313.15 K; ▼, 323.15 K; , linear fit.
where f 2(p, T, x2) represents the fugacity of component 2 (CO2), p is the equilibrium pressure, and x2 is liquid phase mole fraction of CO2. According to the theory of vapor liquid equilibrium, the fugacity in the liquid phase is equal to that in the vapor phase for each component at an equilibrium state. f 2liq (p , T , x 2) = f 2vap (p , T , y2 ) = y2 pϕ2(p , T , y2 )
(5)
where y2 is the mole fraction of component 2 in the gaseous phase. Because DESs exhibit a negligibly small vapor pressures, y2 is simplified to be unity. ϕ2 is the fugacity coefficient of component 2 and can be calculated using the Peng−Robinson equation of state. At a very diluent region of CO2 in the DESs, it holds that: Hx(p , T ) = lim
x2 → 0
≅
f2 (p , T , x 2) x2
= lim
x2 → 0
pϕ2(p , T ) x2
pϕ2(p , T ) x2
(7)
−1
(6) 1250
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Table 3. Henry’s Constants (Hm, Based on Molality, and Hx, Based on Mole Fraction) of CO2 in DESs at Various Temperaturesa Hm/MPa
a
Hx/MPa
DESs
293.15 K
303.15 K
313.15 K
323.15 K
293.15 K
303.15 K
313.15 K
323.15 K
nCC:n1,4‑butanediol = 1:3 nCC:n1,4‑butanediol = 1:4 nCC:n2,3‑butanediol = 1:3 nCC:n2,3‑butanediol = 1:4 nCC:n1,2‑propanediol = 1:3 nCC:n1,2‑propanediol = 1:4
3.09 3.25 3.37 2.64 2.78 2.68
3.48 3.61 3.69 2.98 3.32 3.53
4.23 4.18 4.12 3.44 4.09 4.43
5.07 5.09 4.51 3.87 5.02 5.24
30.49 32.82 33.22 26.84 30.68 30.59
34.38 36.48 36.41 30.23 36.46 40.29
43.66 42.32 40.68 34.79 44.88 50.29
49.81 51.34 44.30 39.19 55.08 59.42
Standard uncertainties u are ur(Hx) = 0.02 and ur(Hm) = 0.02.
liquid, and ΔdisS shows the order degree in the liquid/gas mixture. The obtained values of ΔdisG, ΔdisH, and ΔdisS are listed in Table 5.
Table 4. Comparison of Hm in Present DESs with Some ILs and Other DESs at 313.15 K
a
solution
Hm/MPa
[bheal] [hhemel] [hhemea] [dmim][Me2PO4] [emim][Et2PO4] [bmim][Bu2PO4] [hmim][BF4] [bmim][Tf2N] [P6,6,6,14][C12H25PhSO3] nCC:n2,3‑butanediol = 1:4 nCC:n1,4‑butanediol = 1:4 nCC:n1,2‑propanediol = 1:4 nCC:nurea = 1:2.5 nCC:nurea = 1:2 nCC:nglycerol = 1:2 nCC:nethyleneglycol = 1:2 nCC:nlacticacid = 1:2
4.35 5.38 4.65 2.35 1.84 1.73 1.81 1.06 2.62a 3.44 4.18 4.43 1.37 1.29 1.70 2.71 4.00
Table 5. Standard Gibbs Free Energy (ΔdisG), Enthalpy (ΔdisH), and Entropy (ΔdisS) of Dissolution of CO2 in DESs at 0.1 MPa and 303.15 K ΔdisG DESs
⎛ ∂ ln(H(T , p)/p0 ) ⎞ ⎟ ΔdisH = R ⎜ ∂(1/T ) ⎠p ⎝
kJ·mol
−12.82 −9.89 −7.55 −9.75 −14.45 −18.82
ΔdisS J·mol−1·K−1 −89.84 −80.80 −73.18 −78.65 −95.28 −109.67
CONCLUSIONS The solubilities of CO2 in different DESs composed of CC and dihydric alcohols were determined at 293.15 K, 303.15 K, 313.15 K, and 323.15 K under pressures up to 600.0 kPa using an isochoric saturation method. The molar ratios of CC to dihydric alcohols were selected at 1:3 and 1:4. Henry’s constants of CO2 in DESs were calculated by fitting the experimental data. The results indicate that the solubilities of CO2 in DESs increase with decreasing temperature and increasing pressure. Moreover, the DES composed of CC and 2,3-butanediol with the molar ratio of 1:4 demonstrated the highest capacity to dissolve CO2 among all of the DESs. The dissolution Gibbs free energy, enthalpy, and entropy changes of CO2 solvation were also obtained.
(8)
the coefficients Bi were optimized using a linear regression of multiple-variables calculation. Several thermodynamic properties of the system such as ΔdisG, ΔdisH, and ΔdisS, which are the dissolution Gibbs free energy, enthalpy, and entropy, respectively, can be obtained by correlating Henry’s constants using the following equations: ⎛ H (T , p ) ⎞ ⎟ ΔdisG = RT ln⎜ 0 ⎝ p ⎠
−1
■
n i=0
14.42 14.60 14.63 14.10 14.43 15.42
ΔdisH
For CO2 solvation in DESs under all conditions, the value of ΔdisH always shows negative, which indicates the absorption process is exothermic and that the dissolution of CO2 in DESs is also favorable from the enthalpic view. As seen from the molecular level, the ΔdisS can largely illustrate the DESs organization around the soluble CO2.7 The more negative value of entropy is, the higher ordering degree means when CO2 dissolves in the DESs. As result, the ΔdisG shows a positive value.
those25−28 composed of CC and urea or glycerols, almost similar with that composed of CC and lactic acid.22 Thermodynamic Properties. The behavior of Henry’s constants can be expressed as a function of temperature using the following empirical equation:28
∑ Bi (T /K)−i
kJ·mol
nCC:n1,4‑butanediol = 1:3 nCC:n1,4‑butanediol = 1:4 nCC:n2,3‑butanediol = 1:3 nCC:n2,3‑butanediol = 1:4 nCC:n1,2‑propanediol = 1:3 nCC:n1,2‑propanediol = 1:4
Henry’s constant Hm was calculated at T = 312.45 K.
⎡ H (T ) ⎤ ln⎢ ⎥= ⎣ 0.1 MPa ⎦
−1
(9)
■
(10)
ΔdisH − ΔdisG (11) T 0 where p refers to 0.1 MPa. ΔdisH reflects the strength of intermolecular interaction between the DESs and CO2 in the ΔdisS =
AUTHOR INFORMATION
Corresponding Authors
*E-mail:
[email protected] (N.A.). *E-mail:
[email protected] (D.D.). 1251
dx.doi.org/10.1021/je400884v | J. Chem. Eng. Data 2014, 59, 1247−1253
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Funding
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Financial support from the Natural Science Foundation of Zhejiang Province (no. Y4100699) and the Natural Science Foundation of China (no. 21006095) is deeply appreciated. Notes
The authors declare no competing financial interest.
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