Investigation of the Solubilities of Carbon Dioxide in Some Low

Article pubs.acs.org/jced

Investigation of the Solubilities of Carbon Dioxide in Some Low Volatile Solvents and Their Thermodynamic Properties Xiaofeng Li,*,† Yaotai Jiang,‡ Guoqiang Han,‡ and Dongshun Deng*,‡ †

Department of Chemistry, Zhejiang University, Hangzhou 310027, China College of Chemical Engineering, Zhejiang University of Technology, Hangzhou 310014, China

‡

ABSTRACT: Solubilities of CO2 have been measured in five low volatile organic solvents at temperatures ranging from 293.15 to 323.15 K with 10K intervals under a pressure of 0 to 600.0 kPa using an isochoric saturation method. These organic solvents were selected from γ-butyrolactone (GBL), butyl lactate (BL), 1,1,3,3-tetramethylurea (TMU), 1,3-dimethyl-2-imidazolidinone (DMI), and 1,3dimethyl-3,4,5,6-tetrahydro-2(1H)-pyrimidinone (DMPU). Henry’s constants and thermodynamic properties such as Gibbs free energy, enthalpy, and entropy of dissolution were derived from the corresponding solubility data. The gravimetric solubilities of CO2 in these solvents followed the sequence of TMU > DMI ≈ BL ≈ GBL > DMPU. The enthalpies of dissolution were all exothermal at each condition. Henry’s constants of CO2 in these solvents were further compared with those in ionic liquids and other ordinary absorbents, and it was found that they were similar to that of CO2 in polyethylene glycol dimethyl ether (NHD), which is widely used in the present industry.

1. INTRODUCTION Since about the end of the last century, we have had to face a serious problem of balance between sustainable development of economy and good protection of the environment. CO2 is one of recognized greenhouse gases which cause the greenhouse effect. Thus, reduction of carbon emission becomes a focus topic of global interest and attention.1−3 How to reduce the emission of CO2 at its source is a great challenge for scientists. Therefore, capture and sequestration (CCS) of CO2 has become a very important project over the past decades.4 In the present industry, aqueous alkanolmamine solutions are widely used chemicals for absorption of CO2 because of their high absorption performance.5 Meanwhile, it brings with it problems such as secondary pollution, energy consumption, equipment corrosion, and degradation.6,7 In the past 20 years, ionic liquid (IL) has been proposed as a class of green and potential medium in the field of gas separation.8,9 ILs have such attractive properties as negligible vapor pressure, high thermal stability, strong dissolving capacity, and structure diversity. As for CO2 capture, ILs offer more choices for the researchers to design a suitable absorbent,10 and many ILs were reported as new CO2 absorbents in the literature.11−15 However, some facts are still barriers to their large-scale application such as the high cost and viscosity and the relatively low gravimetric absorption capacity because of ILs’ relatively high molecular weights. Conversely, physical absorption technologies to capture CO2 using organic solvents with relatively low molecular weights have such advantages as low energy consumptions to recover absorbents and high gravimetric absorption capacity.16 Any of the abovementioned suitable physical absorbents should possess those © XXXX American Chemical Society

properties such as low viscosity and volatility, thermal stability, and chemical inertness under the operation conditions. Actually, many absorbents have been reported in the literature, including methanol,17 N-methylpyrrolidinone (NMP),16 propylene carbonate (PC),16 dimethyl sulfoxide (DMSO),18 dimethylformamide (DMF),18 poly(ethylene glycol) (PEG),19 surfactants,20 and polyethylene glycol dimethyl ether (NHD).21 However, the basic efforts to investigate structure−property relationships for gas absorbents are still needed.22,23 The solubilities of CO2 in various organic absorbents depend on the functional groups in their molecular structures. Carbonyl, ester, and substituted amino groups are regarded as CO2-philic groups because of their good dissolution capacity for CO2.21,24,25 On the basis of these literature, γ-butyrolactone (GBL), butyl lactate (BL), 1,1,3,3-tetramethylurea (TMU), 1,3-dimethyl-2imidazolidinone (DMI), and 1,3-dimethyl-3,4,5,6-tetrahydro2(1H)-pyrimidinone (DMPU) were selected in this study as CO2 absorbents, and the dissolution behaviors were investigated. The selected five absorbents have low volatility (35.99, 20.79, 154.6, 15.2, and 5.15 Pa, respectively, according to SciFinder), which means that energy to recover absorbents and loss of absorbents in the regeneration of absorbent are both saved. Second, these absorbents are thermally and chemically stable, noncorrosive, and of low toxicity. The abundant CO2-philic groups in the molecules of these absorbents make them more effective for CO2 capture. TMU, DMI, and DMPU are aprotic Received: October 21, 2015 Accepted: January 27, 2016

A

DOI: 10.1021/acs.jced.5b00893 J. Chem. Eng. Data XXXX, XXX, XXX−XXX

Journal of Chemical & Engineering Data

Article

Table 1. Description of Chemicals Used in the Present Work chemicals

abbreviation

CAS no.

source Jingong Special Gas Co., Ltd. Jinan Chemical Elements Co., Ltd. Aladdin Industrial Co., Ltd. Aladdin Industrial Co., Ltd. Aladdin Industrial Co., Ltd. Aladdin Industrial Co., Ltd.

carbon dioxide γ-butyrolactone

CO2 GBL

124-38-9 96-48-0

butyl 2-hydroxypropanoate 1,1,3,3-tetramethylurea 1,3-dimethyl-2-imidazolidinone 1,3-dimethyl-3,4,5,6-tetrahydro-2 (1H)pyrimidinone

BL TMU DMI DMPU

138-22-7 632-22-4 80-73-9 7226-23-5

a

purification method

mass fraction purity

analysis method

none

0.999 0.994

GCa

none none none none

0.993 0.995 0.996 0.993

GCa GCa GCa GCa

Gas−liquid chromatography.

Table 2. Basic Properties of Five Selected Solvents under 101.3 kPaa

Standard uncertainties u are u(T) = 0.05 K, and relative standard uncertainty of density, pressure, and viscosity are ur(ρ) = 0.001, ur(p) ≈ 1%, and ur(η) = 0.001, respectively. The relative standard uncertainty of water content is 0.005. bThe data were obtained from SciFinder. cAt 5.87 kPa. d− g The data were taken from refs 26−29, respectively. hThe data were taken from Chemical Toxicity Database. iIntraperitoneal for mouse. a

Karl Fischer analysis (SF-3 Karl Fischer Titration, Zibo Zifen Instrument Co. Ltd.) and listed in Table 2, together with some basic properties from the literature. An electronic balance (Mettler-Toledo AL204) with the standard uncertainty of 0.0002 g was used to determine the mass of the above-mentioned solvents. Density data of the solvents were determined at T = 293.15, 303.15, 313.15, and 323.15 K under 101.3 kPa using a DMA4500 combined system (Anton Paar) with the standard uncertainty of 0.0001 g·cm−3. The thermal gravimetric analysis of the solvents was investigated by thermogravimetric analysis (Q50) according to a literature method.30 The experiment in N2 (60 mL/min) was started at 313 K, and the temperature was heated to 480 K at the heating rate of 5.0 K/min. 2.2. Measurement of the Solubilities of CO2. The solubilities of CO2 in the solvents were determined using an isochoric saturation method on a modified apparatus on the base of our previous works31 and illustrated in Figure 1. It was composed mainly of a gas equilibrium cell (4, EC) and a gas reservior (5, GR), with the volumes of 141.61 cm3 and 370.99 cm3, respectively. The temperature of solutions in the EC and GR was maintained at a certain value using thermostatic water bath with the standard uncertainty of 0.05 K. Two pressure transmitters (Fujian WIDEPLUS Precision Instruments Co., Ltd., WIDEPLUS-8) with the standard uncertainty of 0.6 kPa were used to record the change of pressure during absorption

dipolar solvents with similar structure. They are known as “universal solvents” in the laboratory and industry owing to their high dielectric constants and dipole moments, with satisfactory dissolution capacity for many materials. In this study, solubilities of CO2 in the selected five organic solvents were measured at the temperature ranging from 293.15 to 323.15 K with 10 K intervals under a pressure of 0 to 600.0 kPa. Henry’s constants were calculated out from the slope of the linear relationship between solubility and pressure. The thermodynamic parameters of Gibbs free energy, enthalpy, and entropy of dissolution were also deduced from the relationship between Henry’s constant and temperature. Furthermore, the comparisons of Henry’s constant in the selected five solvents with those in several other common solvents as well as some ILs were made systematically.

2. EXPERIMENTAL SECTION 2.1. Materials. 99.9 wt % CO2 was supplied by Jingong Special Gas Co., Ltd. (Hangzhou, China). 1,3-Dimethyl-3,4,5,6tetrahydro-2 (1H)-pyrimidinone (7226−23−5, 99.3 wt %), 1,1,3,3-tetramethylurea (632−22−4, 99.5 wt %), 1,3-dimethyl-2imidazolidinone (80−73−9, 99.6 wt %), γ-butyrolactone (96− 48−0, 99.4 wt %), and butyl 2-hydroxypropanoate (138−22−7, 99.3 wt %) were obtained from Aladdin Industrial Corporation (Shanghai, China). Table 1 provides a summary of the chemicals used. The moisture content of each solvent was determined by B



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By combining eqs 2 and 3, the quality of CO2 absorbed by the solvents is then given as follows: nCO2 = [(ρg (p2 , T ) − ρg (p3 , T ))VGR − (VEC − Vsolvent)]/(1 + ρg (pe , T )V 2∞)

where Vsolvent is the volume of the pure solvent, which can be calculated from the mass and density values at the equilibrium temperature for each solvent. V∞ 2 is the partial molar volume of CO2 in the solvent at the infinite dilution, which can be estimated according to the literature method.35 Accordingly, the mole fraction (xCO2) and molality (mCO2) of CO2 in the liquid phase were calculated according to the two equations below:

Figure 1. Schematic diagram of the detector of CO2 solubility. 1, CO2 gas cylinder; 2,3, thermostatic water bath and magnetic stirrer; 4, CO2 gas equilibrium cell (EC); 5, CO2 gas reservoir(GR); 6,7, pressure transmitters; 8, digital indicator; 9,10,11, valves.

process. After feeding a known amount of solvents (w) into the EC, the whole system was evacuated to pressure p1 with valve 10 closed and valves 9 and 11 opened. With valve 10 opened and valves 9 and 11 closed, GR was charged with CO2 from the cylinder to a pressure p2. Then, the needle valve 9 was opened, valve 10 was closed, and CO2 was brought into EC to contact with the solvents under stirring. Gas−liquid equilibrium is assumed to be achieved if the pressure of EC keeps stable at least 2 h. The final pressure was recorded as p3 for GR and p4 for EC. Then, the equilibrium partial pressure of CO2 in the EC was denoted as following: pe = p4 − p1 (1)

(5)

mCO2 = nCO2 /w

(6)

3. RESULTS AND DISCUSSION 3.1. Solubilities of CO2 in the Solvents. The solubilities of CO2 in five selected solvents were determined at T = 293.15 to 323.15 K with 10 K intervals and under pressures of 0−600.0 kPa. The experimental data are listed in Tables 3−7, including equilibrium pressure (p), molar fraction (x1), and molality (m1) of CO2 in the liquid phase. Figure 2 demonstrates the typical dependence of the solubility (expressed as xCO2 or mCO2) on temperature and pressure for TMU. Figure 3 illustrates the solubility profile of CO2 in five selected solvents at 303.15 K. As seen from Figures 2 and 3, the solubility of CO2 changed linearly with temperature or pressure, meaning that the dissolution of CO2 in the present solvents follows a physical mode. 3.2. Henry’s Constant. As discussed above, the solubility of CO2 changed linearly with temperature or pressure. Therefore, we can use Henry’s law36 to evaluate quantitatively the absorption capacity for these solvents. Thus, Henry’s constant on the basis of molar fraction (Hx) can be expressed as follows:

nCO2 = ρg (p2 , T )VGR −ρg (p3 , T )VGR −ρg (pe , T )(VEC − Vliquid) (2) −3

where ρg (pi, T) is the density of CO2 in g·cm under pressure pi (i = 2, 3, e) and temperature T, and it was found in the database.32 VGR and VEC are the volumes of GR and EC, respectively. Vliquid is the volume occupied by the liquid solution, which can be calculated from the mass and density at equilibrium condition for each solvent. However, this value needs to be corrected due to the dissolution of CO2 into the solvent.33,34 Vliquid = Vsolvent +

xCO2 = nCO2 /(nCO2 + nsol)

where nsol is the mole amount of the solvent used, which is the quotient of of the mass (w) divided by molecular weight.

The amount of absorbed CO2 (nCO2) was calculated by the following equation:

nCO2V 2∞

(4)

Hx(p , T ) ≡ lim

(3)

f 2liq (p , T , x 2)

x2 → 0

x2

(7)

Table 3. Experimental CO2 Mole Fraction (x1) and Molality (m1) in GBL at Temperature (T) under Equilibrium Pressure (p)a

a

T (K)

p (kPa)

m1 (mol·kg−1)

x1

T (K)

p (kPa)

m1 (mol·kg−1)

x1

293.15 293.15 293.15 293.15 293.15 293.15 293.15 303.15 303.15 303.15 303.15 303.15 303.15 303.15

73.3 121.1 189.8 292.4 392.0 490.6 552.4 70.4 128.6 210.3 312.3 401.1 500.2 553.7

0.0947 0.1589 0.2509 0.3917 0.5317 0.6719 0.7610 0.0746 0.1371 0.2262 0.3396 0.4392 0.5527 0.6138

0.0081 0.0135 0.0211 0.0326 0.0438 0.0547 0.0615 0.0064 0.0117 0.0191 0.0284 0.0364 0.0454 0.0502

313.15 313.15 313.15 313.15 313.15 313.15 323.15 323.15 323.15 323.15 323.15 323.15 323.15

65.7 125.4 206.5 304.4 399.1 510.1 62.2 116.3 196.6 291.1 385.8 488.2 556.7

0.0587 0.1139 0.1884 0.2808 0.3707 0.4765 0.0497 0.0912 0.1540 0.2278 0.3030 0.3849 0.4409

0.0050 0.0097 0.0160 0.0236 0.0309 0.0394 0.0043 0.0078 0.0131 0.0192 0.0253 0.0321 0.0366

Standard uncertainties u are u(T) = 0.05 K and u(p) = 0.6 kPa; and relative standard uncertainties of solubility are ur (x) = 0.02 and ur (m) = 0.02. C



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Table 4. Experimental CO2 Mole Fraction (x1) and Molality (m1) in BL at Temperature (T) and Equilibrium Pressure (p)a

a

T (K)

p (kPa)

m1 (mol·kg−1)

x1

T (K)

p (kPa)

m1 (mol·kg−1)

x1

293.15 293.15 293.15 293.15 293.15 293.15 293.15 303.15 303.15 303.15 303.15 303.15 303.15 303.15

69.7 127.7 212.2 307.2 398.5 503.2 556.5 71.3 127.2 215.9 308.5 407.4 512.1 570.4

0.0845 0.1601 0.2687 0.3906 0.5120 0.6523 0.7284 0.0768 0.1360 0.2310 0.3337 0.4422 0.5618 0.6303

0.0122 0.0229 0.0378 0.0540 0.0696 0.0871 0.0962 0.0111 0.0195 0.0327 0.0465 0.0607 0.0759 0.0844

313.15 313.15 313.15 313.15 313.15 313.15 313.15 323.15 323.15 323.15 323.15 323.15 323.15 323.15

75.9 129.9 223.3 324.2 431.4 515.6 577.6 98.4 152.1 227.0 320.5 424.3 521.4 584.1

0.0694 0.1196 0.2029 0.2960 0.3989 0.4783 0.5436 0.0741 0.1154 0.1710 0.2448 0.3252 0.4039 0.4554

0.0100 0.0172 0.0288 0.0415 0.0551 0.0654 0.0736 0.0107 0.0166 0.0244 0.0345 0.0454 0.0558 0.0624

Standard uncertainties u are u(T) = 0.05 K and u(p) = 0.6 kPa, and relative standard uncertainties of solubility are ur (x) = 0.02 and ur (m) = 0.02.

Table 5. Experimental CO2 Mole Fraction (x1) and Molality (m1) in TMU at Temperature (T) and Equilibrium Pressure (p)a

a

T (K)

p (kPa)

m1 (mol·kg−1)

x1

T (K)

p (kPa)

m1 (mol·kg−1)

x1

293.15 293.15 293.15 293.15 293.15 293.15 293.15 303.15 303.15 303.15 303.15 303.15 303.15

72.6 128.9 226.4 311.5 406.2 497.9 576.0 113.9 225.0 332.9 426.9 528.0 581.3

0.1270 0.2280 0.4050 0.5648 0.7466 0.9259 1.0806 0.1635 0.3279 0.4916 0.6334 0.7914 0.8768

0.0145 0.0258 0.0449 0.0616 0.0798 0.0971 0.1115 0.0186 0.0367 0.0540 0.0685 0.0842 0.0924

313.15 313.15 313.15 313.15 313.15 313.15 313.15 323.15 323.15 323.15 323.15 323.15 323.15 323.15

78.1 132.9 218.5 320.7 418.0 524.4 590.4 88.4 139.0 235.8 332.6 434.4 532.6 589.4

0.0903 0.1554 0.2587 0.3854 0.5078 0.6435 0.7229 0.0887 0.1396 0.2377 0.3392 0.4478 0.5529 0.6153

0.0104 0.0177 0.0292 0.0428 0.0557 0.0695 0.0781 0.0102 0.0160 0.0269 0.0379 0.0494 0.0603 0.0667


Table 6. Experimental CO2 Mole Fraction (x1) and Molality (m1) in DMI at Temperature (T) and Equilibrium Pressure (p)a

a

T (K)

p (kPa)

m1 (mol·kg−1)

x1

T (K)

p (kPa)

m1 (mol·kg−1)

x1

293.15 293.15 293.15 293.15 293.15 293.15 293.15 303.15 303.15 303.15 303.15 303.15 303.15 303.15

47.2 96.9 176.3 260.5 379.1 467.2 528.4 52.5 105.2 172.5 274.8 376.6 479.6 545.1

0.0657 0.1359 0.2489 0.3723 0.5521 0.6872 0.7843 0.0578 0.1183 0.1946 0.3111 0.4341 0.5568 0.6361

0.0074 0.0153 0.0276 0.0408 0.0593 0.0727 0.0822 0.0065 0.0133 0.0217 0.0343 0.0472 0.0598 0.0677

313.15 313.15 313.15 313.15 313.15 313.15 313.15 323.15 323.15 323.15 323.15 323.15 323.15 323.15

52.6 106.8 194.2 287.4 367.5 474.4 544.0 65.3 126.2 211.1 305.8 391.7 498.3 555.8

0.0505 0.1025 0.1885 0.2807 0.3608 0.4685 0.5404 0.0520 0.1011 0.1705 0.2482 0.3204 0.4097 0.4589

0.0057 0.0116 0.0211 0.0310 0.0396 0.0508 0.0581 0.0059 0.0114 0.0191 0.0275 0.0353 0.0447 0.0498


f 2liq (p , T , x 2) = f 2vap (p , T , y2 ) = y2 pϕ2(p , T , y2 )

where f 2 (p, T, x2), p, and x2 represent the fugacity, equilibrium pressure, and molar fraction of CO2 in the liquid phase, respectively. When CO2 reaches the equilibrium of dissolution, the fugacity of CO2 in the liquid phase must be equal that in the gas phase. Then,

(8)

where y2 denotes the molar fraction of CO2 in the gas phase. Because the vapor pressure of each solvent is too small and was neglected at the experimental temperature, the composition of the gas phase was regarded as pure CO2, and y2 is simplified as D



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Table 7. Experimental CO2 Mole Fraction (x1) and Molality (m1) in DMPU at Temperature (T) and Equilibrium Pressure (p)a

a

T (K)

p (kPa)

m1 (mol·kg−1)

x1

T (K)

p (kPa)

m1 (mol·kg−1)

x1

293.15 293.15 293.15 293.15 293.15 293.15 293.15 303.15 303.15 303.15 303.15 303.15 303.15 303.15

53.8 105.1 185.1 276.6 379.9 475.5 534.3 65.3 120.7 201.2 292.6 393.2 490.8 546.6

0.0604 0.1184 0.2092 0.3159 0.4394 0.5552 0.6291 0.0606 0.1116 0.1861 0.2727 0.3689 0.4645 0.5201

0.0077 0.0149 0.0261 0.0389 0.0533 0.0664 0.0746 0.0077 0.0141 0.0233 0.0338 0.0451 0.0562 0.0625

313.15 313.15 313.15 313.15 313.15 313.15 313.15 323.15 323.15 323.15 323.15 323.15 323.15 323.15

56.0 110.1 191.9 283.1 382.4 476.6 543.9 63.9 118.3 207.1 297.1 384.1 486.7 562.9

0.0442 0.0863 0.1517 0.2249 0.3065 0.3832 0.4396 0.0425 0.0784 0.1378 0.1991 0.2589 0.3304 0.3839

0.0056 0.0109 0.0191 0.0280 0.0378 0.0468 0.0533 0.0054 0.0099 0.0174 0.0249 0.0321 0.0406 0.0469


Figure 2. Solubilities of CO2 in 1,1,3,3-tetramethylurea. (a) Expressed as molality; (b) Expressed as molar fraction. □, 293.15 K; △, 303.15 K; ○, 313.15 K; ▽, 323.15k; , linear fit.

Figure 3. Solubilities of CO2 in five solvents at 303.15 K. (a) Expressed as molality; (b) Expressed as molar fraction. ○, GBL; ▽, BL; △, TMU; ◊, DMI; ★, DMPU.

unity. ϕ2 is the fugacity coefficient of CO2, which was calculated using the Virial equation within two-terms because the pressure of CO2 is low enough. Thus, at the very dilute region of CO2 in the liquid phase, Henry’s constant could be calculated by combining eqs 5 and 6 as follows: Hx(p , T ) = lim

x2 → 0

≅

f2 (p , T , x 2) x2

= lim

x2 → 0

On the basis of molality, Henry’s constant (Hm) could also be calculated out as follows: ⎡ f (p , T ) ⎤ pϕ2(p , T ) ⎥≅ Hm(p , T ) ≡ lim ⎢ 2 0 m2 → 0⎣ (m 2 / m ) ⎦ (m 2 / m 0 )

f2 (p , T )

where m2 is the molality of CO2 in the liquid phase and m is 1 mol·kg−1. In the present work, Hx or Hm for each mixture was obtained as the slope of the curve by fitting fugacity with molar fraction or molality of CO2, with the results listed in Table 8. The Hx for five solvents at each experimental temperature was

x2

pϕ2(p , T ) x2

(10) 0

(9) E



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Table 8. Henry’s Constants (Hm, Based on Molality; and Hx, Based on Molar Fraction) of CO2 in Selected Solvents at Different Temperatures Hm (MPa)

Hx (MPa)

solvents

293.15 K

303.15 K

313.15 K

323.15 K

293.15 K

303.15 K

313.15 K

323.15 K

GBL BL TMU DMI DMPU

0.74 0.77 0.54 0.68 0.86

0.92 0.91 0.67 0.86 1.06

1.08 1.08 0.82 1.02 1.25

1.28 1.30 0.97 1.22 1.48

8.97 5.75 5.11 6.41 7.14

11.02 6.72 6.23 8.02 8.72

12.93 7.84 7.51 9.33 10.16

15.19 9.34 8.82 11.14 11.98

are structurally similar and related, they illustrate differences in physiochemical properties. As seen from Table 2, with the increase of polarity (μTMU = 3.45D, μDMI = 4.09D, and μDMPU = 4.23D),37 the density and viscosity increase while saturated vapor pressure at 298.15 K decreases. The substituted amino group in TMU still needs slightly weak basicity to bind with CO2; therefore, TMU demonstrates the highest solubility among five solvents. When the two amino groups were cyclized by ethylene (DMI) or the propylene group (DMPU), the binding ability decreases evidently because of steric hindrance. GBL demonstrates the lowest solubility of CO2 among the five solvents because it has the smallest molar volume (76.16 cm3·mol−1 at 298.15 K). Similarly, BL still possesses good solubility because of the highest molar volume (148.80 cm3·mol−1 at 298.15 K). Conversely, when the molecular mass was taken into account, the gravimetric solubility sequence (as shown with Hm) was changed to be TMU > DMI ≈ BL ≈ GBL > DMPU for these solvents. In our opinion, the later sequence reveals the valuable capture ability for CO2 from the viewpoint of atomic economy. When considering solubility capacity, safety, and physiochemical properties, BL and GBL maybe more attractive than the others as an absorbent for CO2. 3.3. Thermodynamic Properties. Thermodynamic parameters are helpful to study the gas dissolution behavior and is essential for the design process in gas separation technology. On the basis of the dependence of Henry’s constants on temperature, three thermodynamic properties of ΔdisG, ΔdisH, and ΔdisS can be deduced from the following equations:38

graphically illustrated in Figure 4. It is obvious that the values of Hx lie in the scope of 5.17 to 15.29 MPa and increase linearly with increasing temperature.

Figure 4. Dependence of Henry’s constants (based on mole fraction) on temperatures. ★, GBL; ○, DMPU; ◊, DMI; △, BL; ▽, TMU.

The values of Hx are based on molar quantity without consideration of the molecular mass. Therefore, they are helpful for recognizing the structure−CO2 dissolution relationship for each solvent. Generally speaking, the gas solubility in a liquid comprehensively reflects the group information and molar volume. As shown in Table 9, the solubility in the present solvents follows the sequence of TMU > BL > DMI > DMPU > GBL at the same temperature. Although the present three amides

Table 9. Comparison of Selected Solvents with Other Absorbents and ILs at 313.15 K

a

solutions

Hm (MPa)

ΔdisG0 (kJ mol−1)

ΔdisH0 (kJ mol−1)

ΔdisS0 (J mol−1 K−1)

GBL+ CO2 BL + CO2 TMU + CO2 DMI + CO2 DMPU + CO2 NHD + CO2 NMP + CO2 PC + CO2 EEA + CO2 PEG-400 + CO2 Tween 80 + CO2 TX-100 + CO2 FAPE-1815 + CO2 [EMIM][Et2PO4] + CO2 [BMIM][PF6] + CO2 [bheaa] + CO2 [bheal] + CO2

1.08 1.08 0.82 1.02 1.25 1.05 0.56a 0.72 0.75 1.91 1.60b 1.68b 1.85 1.80 1.89 2.38 3.75

12.66 11.36 11.24 11.81 12.03 9.78 10.8 12.1 10.31 10.51 0.32 0.74 7.54 11.06 7.64 0.08 0.17

−13.71 −12.66 −14.34 −14.24 −13.42 −10.28 −16.4 −15.9 −12.71 −11.18 −10.20 −20.40 −34.63 −11.14 −16.21 −14.63 −25.32

−84.21 −76.70 −81.69 −83.19 −81.27 −64.03 −86.9 −89.4 −73.51 −69.26 −33.60 −67.50 −134.67 −70.88 −76.16 −46.97 −81.39

The temperature is 298.15 K. bThe temperature is 308.15 K. F



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⎛ H (T , p) ⎞ ⎟ Δdis G = RT ln⎜ 0 ⎠ ⎝ p

(11)

⎛ ∂ ln(H(T , p)/p0 ) ⎞ ⎟ Δdis H = R ⎜ ∂(1/T ) ⎠p ⎝

(12)

but evidently higher than those in other solvents with higher boiling points and ordinary ILs, such as fatty PEG-400,40 Tween 80,20 TX-100,20 amine polyoxyethylene ether 1815 (FAPE1815), 41 1-ethyl-3-methylimidazolium diethylphosphate ([emim][Et2PO4]),42 and 1-butyl-3-methylimidazolium hexafluorophosphate ([bmim][PF 6 ]), 36 bis(2-hydroxyethyl)ammonium lactate ([bheal]), and bis(2-hydroxyethyl)ammonium acetate ([bheaa]).43 To our knowledge, compared with traditional absorbents and these new emerging ILs, the present solvents possess low toxicity and low vapor pressure. The enthalpies of absorption in the present solvents were similar to those in common organic absorbents and ILs, but slightly less negative than those in TX-100, FAPE-1815, and IL of [bheal].

Δdis H − Δdis G (13) T 0 where p is the pressure at standard station. ΔdisG is the standard Gibbs free energy of dissolution which reaches the minimum at the gas−liquid equilibrium. ΔdisH is the standard enthalpy of dissolution which reflects quantitatively the intermolecular interaction between gas and solvents in the liquid phase. ΔdisS is the standard entropy of dissolution which represents the dispersion degree of the solution. In Table 9 are listed the ΔdisG, ΔdisH, and ΔdisS at 313.15 K and 0.1 MPa. As shown in Table 9, the negative values of ΔdisH indicate that the absorption of CO2 by these solvents is an exothermic process and that there are strong interactions between CO2 and the solvents. The negative values of ΔdisS mean that the solutions become less chaotic than CO2 in the gas form because of lower movements of CO2 molecules at random after dissolution of CO2 on the molecular level. The positive values of ΔdisG at the studied temperature show that the process of dissolution is not spontaneous and that it is mainly pushed by pressure. 3.4. Thermal Gravimetric Analysis of the Solvents. Figure 5 illustrates the thermal gravimetric curves of the solvents. Δdis S =

4. CONCLUSIONS In this study, the data of CO2 solubility in five solvents with high boiling point were presented at 293.15, 303.15, 313.15, and 323.15 K under pressures of 0−600.0 kPa using an isochoric saturation method. The results indicate that the dissolution of CO2 in these solvents follows a physical mode and that TMU (1,1,3,3-tetramethylurea) has the highest gravimetric absorption capacity of CO2 among the five solvents. The process of desorption was advantageous because of the low value of dissolution enthalpy. The positive values of both ΔdisS and ΔdisG show that the process of dissolution is not spontaneous and that it is mainly pushed by pressure.

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AUTHOR INFORMATION

Corresponding Authors

*(X.L.) E-mail: [email protected]. *(D.D.) E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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REFERENCES

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Figure 5. Thermal gravimetric analysis of solvents: red line, TMPU; black line, DMI; blue line, GBL; dark cyan line, BL; magenta line, TMU (all weight percentages).

Here, the onset temperature (Tonset) was used to demonstrate the temperature where the weight loss begins on the TGA analysis. The Tonset values of TMPU, DMI, GBL, BL, and TMU are 418, 400, 393, 377, and 361 K, respectively. Indeed, such relatively low Tonset values suggest that these physical solvents should be used near or below room temperature. 3.5. Comparison with Literature Solvents. For further evaluation of the present five solvents as CO2 absorbents, the dissolution capacity of CO2 in these solvents was compared with that of other physical absorbents in the industry and some new emerging ILs. Hm provided a convenient scale for different absorbents owing to its definition on mass, with the detailed results listed in Table 9. As shown in Table 9, the solubilities of CO2 in selected solvents are lower than those in NMP,16 PC,16 and 2-(2-ethoxyethoxy)ethyl acetate (EEA),39 similar to NHD,21 G



Article

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